what do molecules look like?
DESCRIPTION
What Do Molecules Look Like?. The Lewis Dot Structure approach provides some insight into molecular structure in terms of bonding, but what about 3D geometry?. Recall that we have two types of electron pairs: bonding and lone. - PowerPoint PPT PresentationTRANSCRIPT
What Do Molecules Look Like
Recall that we have two types of electron pairs bonding and lone
The Lewis Dot Structure approach provides some insight into molecular structure in terms of bonding but what about 3D geometry
Valence-Shell Electron-Pair Repulsion (VSEPR) 3D structure is determined by minimizing repulsion of electron pairs
Electron pairs (both bonding and lone) are distributed around a central atom such that electron-electron repulsions are minimized
2 electron pairs
3 electron pairs
4 electron pairs
5 electron pairs
6 electron pairs
Period 1 2
Period 3 amp beyond
Electron pairs (both bonding and lone) are distributed around a central atom such that electron-electron repulsions are minimized
Arranging Electron Pairs
bull Example CH4 (bonding pairs only)
bull Must consider both bonding and lone pairs when minimizing repulsion
Lewis Structure VSEPR Structure
H CH
HH
Arranging Electron Pairs (cont)
Example NH3 (both bonding and lone pairs)
Lewis Structure VSEPR Structure
Noteldquoelectron pair geometryrdquo vsldquomolecular shaperdquo
H NH
H
VSEPR Structure GuidelinesThe previous examples illustrate the strategy for applying VSEPR to predict molecular structure
1 Construct the Lewis Dot Structure2 Arrange bondinglone electron pairs in space such
that repulsions are minimized (electron pair geometry)
3 Name the molecular shape from the position of the atoms
VSEPR Shorthand
1 Refer to central atom as ldquoArdquo2 Attached atoms are referred to as ldquoXrdquo3 Lone pair are referred to as ldquoErdquo
ExamplesCH4 AX4NH3 AX3EH2O AX2E2BF3 AX3
VSEPR 2 electron pairsLinear (AX2) angle between bonds is 180deg
Example BeF2
180deg
Experiments show that molecules with multiple bonds can also be linear
Multiple bonds are treated as a single effective electron groupF Be F
F Be F
More than one central atom Determine shape around each
Be FF
VSEPR 3 electron pairsTrigonal Planar (AX3) angle between bonds is 120degExample BF3
Multiple bond is treated as a single effective electron group
FBFF
B
FF
F
120deg
VSEPR 4 electron pairs (cont)
Tetrahedral (AX4) angle between bonds is ~1095degExample CH4
1095degH C
HH
H
tetrahedral e- pair geometry AND tetrahedral molecular shape
Bonding vs Lone pairsBond angle in a tetrahedral arrangement of electron pairs may vary from 1095deg due to size differences between bonding and lone pair electron densities
bonding pair is constrained by two nuclear potentials more localized in space
lone pair is constrained by only one nuclear potential less localized (needs more room)
VSEPR 4 electron pairsTrigonal pyramidal (AX3E) Bond angles are lt1095deg and structure is nonplanar due to repulsion of lone pair
Example NH3
107deg
tetrahedral e- pair geometry trigonal pyramidal molecular shape
H NH
H
VSEPR 4 electron pairs (cont)
Classic example of tetrahedral angle shift from 1095deg is water (AX2E2)
ldquobentrdquo1045o
tetrahedral e- pair geometry bent molecular shape
VSEPR 4 electron pairs (cont)
Comparison of CH4 (AX4) NH3 (AX3E) and H2O (AX2E2)
AX2E
AX3E
AX2E2
1 Refer to central atom as ldquoArdquo2 Attached atoms are referred to as ldquoXrdquo3 Lone pair are referred to as ldquoErdquo
HH
H
H
CNOF
Central Atom
Compound Electron-Pair Geometry
Molecular Shape
Carbon C CH4 tetrahedral tetrahedralNitrogen N NH3 tetrahedral trigonal pyramidalOxygen O H2O tetrahedral bentFluorine F HF tetrahedral linear
Molecular vs Electron-Pair Geometry
What is the electron-pair geometry and the molecular shape for HCFS
a) trigonal planar bent
b) trigonal planar trigonal planar
c) tetrahedral trigonal planar
d) tetrahedral tetrahedral
C
S
H
F
VSEPR Beyond the OctetSystems with expanded valence shells will have five or six electron pairs around a central atom
P
Cl
Cl
Cl
Cl
Cl
FF
F
F
FFS
90deg
120deg SF
F
F
F
F
F 90deg
90deg
VSEPR 5 electron pairsbull Consider the structure of SF4 (34 e- AX4E)bull What is the optimum arrangement of electron pairs
around S
SFF
FF
S
FF
FF
SF
FF
F
lone-pair bond-pair
two at 90o two at 120o
Repulsive forces (strongest to weakest)lone-pairlone-pair gt lone-pairbond-pair gt bond-pairbond-pair
bond-pair bond-pair
three at 90o four at 90o one at
120othree at 90o three at 120o
Compare endash pair angles
VSEPR 5 electron pairsThe optimum structure maximizes the
angular separation of the lone pairs
I3- (AX2E3)
AX3E2
AX4E
AX2E3
5-electron-pair geometries
our previous example
Which of these is the more likely structure
Square Planar
VSEPR 6 electron pairs
See-saw
AX5E
AX4E2
6-electron-pair geometries
our previous example
Molecular Dipole Moments
1 Draw Lewis structures to determine 3D arrangement of atoms
Shortcut completely symmetric molecules will not have a dipole regardless of the polarity of the bonds
2 If one ldquosiderdquo of the molecule has more EN atoms than the other the molecule has a net dipole
We can use VSEPR to determine the polarity of a whole molecule
Molecular Dipoles
The C=O bonds have dipoles of equal magnitude but opposite direction so there is no net dipole moment
The O-H bonds have dipoles of equal magnitude that do not cancel each other so water has a net dipole moment
Molecular Dipoles (cont)
symmetric
symmetricasymmetri
c
F
ClF
Cl
Molecular Dipole Examplebull Write the Lewis dot and VESPR
structures for CF2Cl2 Does it have a dipole moment
C
F
FClCl
32 e-
Tetrahedral
Advanced VSEPR Application
Molecules with more than one central atomhellip methanol (CH3OH)
H CH
OH
H
tetrahedral e- pairstetrahedral shape
tetrahedral e- pairsbent shape
The VSEPR Table
2 AX2 BeF2 linear linear
3 AX3 BF3 trigonal planar
trigonal planar
AX2E O3 trigonal planar
bent
4 AX4 CH4 tetrahedral tetrahedral
AX3E NH3 tetrahedral pyramidal
AX2E2 H2O tetrahedral bent
e- pairs e- Geom Molec Geom
The VSEPR Table5 AX5 PF5 trigonal
bipyramidaltrigonalbipyramidal
AX4E SF4 trigonal bipyramidal
see saw
AX3E2 ClF3 trigonal bipyramidal
T-shaped
AX2E3 I3- trigonal bipyramidal
linear
6 AX6 SF6 octahedral octahedral
AX4E2 XeF4 octahedral square planar
e- pairs e- Geom Molec Geom
What is the expected shape of ICl2+
A linear
B bent
C tetrahedral
D square planar
AX2E2ICl Cl+
20 e-
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons occupies the region between the two nuclei
Rule 1 Maximum overlap The bond strength depends on the attraction of nuclei to the shared electrons soThe greater the orbital overlap the stronger the bond
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 2 Spins pair The two electrons in the overlap region occupy the same space and therefore must have opposite spins
There may be no more than 2 electrons in a molecular orbital
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Electron pairs (both bonding and lone) are distributed around a central atom such that electron-electron repulsions are minimized
2 electron pairs
3 electron pairs
4 electron pairs
5 electron pairs
6 electron pairs
Period 1 2
Period 3 amp beyond
Electron pairs (both bonding and lone) are distributed around a central atom such that electron-electron repulsions are minimized
Arranging Electron Pairs
bull Example CH4 (bonding pairs only)
bull Must consider both bonding and lone pairs when minimizing repulsion
Lewis Structure VSEPR Structure
H CH
HH
Arranging Electron Pairs (cont)
Example NH3 (both bonding and lone pairs)
Lewis Structure VSEPR Structure
Noteldquoelectron pair geometryrdquo vsldquomolecular shaperdquo
H NH
H
VSEPR Structure GuidelinesThe previous examples illustrate the strategy for applying VSEPR to predict molecular structure
1 Construct the Lewis Dot Structure2 Arrange bondinglone electron pairs in space such
that repulsions are minimized (electron pair geometry)
3 Name the molecular shape from the position of the atoms
VSEPR Shorthand
1 Refer to central atom as ldquoArdquo2 Attached atoms are referred to as ldquoXrdquo3 Lone pair are referred to as ldquoErdquo
ExamplesCH4 AX4NH3 AX3EH2O AX2E2BF3 AX3
VSEPR 2 electron pairsLinear (AX2) angle between bonds is 180deg
Example BeF2
180deg
Experiments show that molecules with multiple bonds can also be linear
Multiple bonds are treated as a single effective electron groupF Be F
F Be F
More than one central atom Determine shape around each
Be FF
VSEPR 3 electron pairsTrigonal Planar (AX3) angle between bonds is 120degExample BF3
Multiple bond is treated as a single effective electron group
FBFF
B
FF
F
120deg
VSEPR 4 electron pairs (cont)
Tetrahedral (AX4) angle between bonds is ~1095degExample CH4
1095degH C
HH
H
tetrahedral e- pair geometry AND tetrahedral molecular shape
Bonding vs Lone pairsBond angle in a tetrahedral arrangement of electron pairs may vary from 1095deg due to size differences between bonding and lone pair electron densities
bonding pair is constrained by two nuclear potentials more localized in space
lone pair is constrained by only one nuclear potential less localized (needs more room)
VSEPR 4 electron pairsTrigonal pyramidal (AX3E) Bond angles are lt1095deg and structure is nonplanar due to repulsion of lone pair
Example NH3
107deg
tetrahedral e- pair geometry trigonal pyramidal molecular shape
H NH
H
VSEPR 4 electron pairs (cont)
Classic example of tetrahedral angle shift from 1095deg is water (AX2E2)
ldquobentrdquo1045o
tetrahedral e- pair geometry bent molecular shape
VSEPR 4 electron pairs (cont)
Comparison of CH4 (AX4) NH3 (AX3E) and H2O (AX2E2)
AX2E
AX3E
AX2E2
1 Refer to central atom as ldquoArdquo2 Attached atoms are referred to as ldquoXrdquo3 Lone pair are referred to as ldquoErdquo
HH
H
H
CNOF
Central Atom
Compound Electron-Pair Geometry
Molecular Shape
Carbon C CH4 tetrahedral tetrahedralNitrogen N NH3 tetrahedral trigonal pyramidalOxygen O H2O tetrahedral bentFluorine F HF tetrahedral linear
Molecular vs Electron-Pair Geometry
What is the electron-pair geometry and the molecular shape for HCFS
a) trigonal planar bent
b) trigonal planar trigonal planar
c) tetrahedral trigonal planar
d) tetrahedral tetrahedral
C
S
H
F
VSEPR Beyond the OctetSystems with expanded valence shells will have five or six electron pairs around a central atom
P
Cl
Cl
Cl
Cl
Cl
FF
F
F
FFS
90deg
120deg SF
F
F
F
F
F 90deg
90deg
VSEPR 5 electron pairsbull Consider the structure of SF4 (34 e- AX4E)bull What is the optimum arrangement of electron pairs
around S
SFF
FF
S
FF
FF
SF
FF
F
lone-pair bond-pair
two at 90o two at 120o
Repulsive forces (strongest to weakest)lone-pairlone-pair gt lone-pairbond-pair gt bond-pairbond-pair
bond-pair bond-pair
three at 90o four at 90o one at
120othree at 90o three at 120o
Compare endash pair angles
VSEPR 5 electron pairsThe optimum structure maximizes the
angular separation of the lone pairs
I3- (AX2E3)
AX3E2
AX4E
AX2E3
5-electron-pair geometries
our previous example
Which of these is the more likely structure
Square Planar
VSEPR 6 electron pairs
See-saw
AX5E
AX4E2
6-electron-pair geometries
our previous example
Molecular Dipole Moments
1 Draw Lewis structures to determine 3D arrangement of atoms
Shortcut completely symmetric molecules will not have a dipole regardless of the polarity of the bonds
2 If one ldquosiderdquo of the molecule has more EN atoms than the other the molecule has a net dipole
We can use VSEPR to determine the polarity of a whole molecule
Molecular Dipoles
The C=O bonds have dipoles of equal magnitude but opposite direction so there is no net dipole moment
The O-H bonds have dipoles of equal magnitude that do not cancel each other so water has a net dipole moment
Molecular Dipoles (cont)
symmetric
symmetricasymmetri
c
F
ClF
Cl
Molecular Dipole Examplebull Write the Lewis dot and VESPR
structures for CF2Cl2 Does it have a dipole moment
C
F
FClCl
32 e-
Tetrahedral
Advanced VSEPR Application
Molecules with more than one central atomhellip methanol (CH3OH)
H CH
OH
H
tetrahedral e- pairstetrahedral shape
tetrahedral e- pairsbent shape
The VSEPR Table
2 AX2 BeF2 linear linear
3 AX3 BF3 trigonal planar
trigonal planar
AX2E O3 trigonal planar
bent
4 AX4 CH4 tetrahedral tetrahedral
AX3E NH3 tetrahedral pyramidal
AX2E2 H2O tetrahedral bent
e- pairs e- Geom Molec Geom
The VSEPR Table5 AX5 PF5 trigonal
bipyramidaltrigonalbipyramidal
AX4E SF4 trigonal bipyramidal
see saw
AX3E2 ClF3 trigonal bipyramidal
T-shaped
AX2E3 I3- trigonal bipyramidal
linear
6 AX6 SF6 octahedral octahedral
AX4E2 XeF4 octahedral square planar
e- pairs e- Geom Molec Geom
What is the expected shape of ICl2+
A linear
B bent
C tetrahedral
D square planar
AX2E2ICl Cl+
20 e-
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons occupies the region between the two nuclei
Rule 1 Maximum overlap The bond strength depends on the attraction of nuclei to the shared electrons soThe greater the orbital overlap the stronger the bond
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 2 Spins pair The two electrons in the overlap region occupy the same space and therefore must have opposite spins
There may be no more than 2 electrons in a molecular orbital
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
2 electron pairs
3 electron pairs
4 electron pairs
5 electron pairs
6 electron pairs
Period 1 2
Period 3 amp beyond
Electron pairs (both bonding and lone) are distributed around a central atom such that electron-electron repulsions are minimized
Arranging Electron Pairs
bull Example CH4 (bonding pairs only)
bull Must consider both bonding and lone pairs when minimizing repulsion
Lewis Structure VSEPR Structure
H CH
HH
Arranging Electron Pairs (cont)
Example NH3 (both bonding and lone pairs)
Lewis Structure VSEPR Structure
Noteldquoelectron pair geometryrdquo vsldquomolecular shaperdquo
H NH
H
VSEPR Structure GuidelinesThe previous examples illustrate the strategy for applying VSEPR to predict molecular structure
1 Construct the Lewis Dot Structure2 Arrange bondinglone electron pairs in space such
that repulsions are minimized (electron pair geometry)
3 Name the molecular shape from the position of the atoms
VSEPR Shorthand
1 Refer to central atom as ldquoArdquo2 Attached atoms are referred to as ldquoXrdquo3 Lone pair are referred to as ldquoErdquo
ExamplesCH4 AX4NH3 AX3EH2O AX2E2BF3 AX3
VSEPR 2 electron pairsLinear (AX2) angle between bonds is 180deg
Example BeF2
180deg
Experiments show that molecules with multiple bonds can also be linear
Multiple bonds are treated as a single effective electron groupF Be F
F Be F
More than one central atom Determine shape around each
Be FF
VSEPR 3 electron pairsTrigonal Planar (AX3) angle between bonds is 120degExample BF3
Multiple bond is treated as a single effective electron group
FBFF
B
FF
F
120deg
VSEPR 4 electron pairs (cont)
Tetrahedral (AX4) angle between bonds is ~1095degExample CH4
1095degH C
HH
H
tetrahedral e- pair geometry AND tetrahedral molecular shape
Bonding vs Lone pairsBond angle in a tetrahedral arrangement of electron pairs may vary from 1095deg due to size differences between bonding and lone pair electron densities
bonding pair is constrained by two nuclear potentials more localized in space
lone pair is constrained by only one nuclear potential less localized (needs more room)
VSEPR 4 electron pairsTrigonal pyramidal (AX3E) Bond angles are lt1095deg and structure is nonplanar due to repulsion of lone pair
Example NH3
107deg
tetrahedral e- pair geometry trigonal pyramidal molecular shape
H NH
H
VSEPR 4 electron pairs (cont)
Classic example of tetrahedral angle shift from 1095deg is water (AX2E2)
ldquobentrdquo1045o
tetrahedral e- pair geometry bent molecular shape
VSEPR 4 electron pairs (cont)
Comparison of CH4 (AX4) NH3 (AX3E) and H2O (AX2E2)
AX2E
AX3E
AX2E2
1 Refer to central atom as ldquoArdquo2 Attached atoms are referred to as ldquoXrdquo3 Lone pair are referred to as ldquoErdquo
HH
H
H
CNOF
Central Atom
Compound Electron-Pair Geometry
Molecular Shape
Carbon C CH4 tetrahedral tetrahedralNitrogen N NH3 tetrahedral trigonal pyramidalOxygen O H2O tetrahedral bentFluorine F HF tetrahedral linear
Molecular vs Electron-Pair Geometry
What is the electron-pair geometry and the molecular shape for HCFS
a) trigonal planar bent
b) trigonal planar trigonal planar
c) tetrahedral trigonal planar
d) tetrahedral tetrahedral
C
S
H
F
VSEPR Beyond the OctetSystems with expanded valence shells will have five or six electron pairs around a central atom
P
Cl
Cl
Cl
Cl
Cl
FF
F
F
FFS
90deg
120deg SF
F
F
F
F
F 90deg
90deg
VSEPR 5 electron pairsbull Consider the structure of SF4 (34 e- AX4E)bull What is the optimum arrangement of electron pairs
around S
SFF
FF
S
FF
FF
SF
FF
F
lone-pair bond-pair
two at 90o two at 120o
Repulsive forces (strongest to weakest)lone-pairlone-pair gt lone-pairbond-pair gt bond-pairbond-pair
bond-pair bond-pair
three at 90o four at 90o one at
120othree at 90o three at 120o
Compare endash pair angles
VSEPR 5 electron pairsThe optimum structure maximizes the
angular separation of the lone pairs
I3- (AX2E3)
AX3E2
AX4E
AX2E3
5-electron-pair geometries
our previous example
Which of these is the more likely structure
Square Planar
VSEPR 6 electron pairs
See-saw
AX5E
AX4E2
6-electron-pair geometries
our previous example
Molecular Dipole Moments
1 Draw Lewis structures to determine 3D arrangement of atoms
Shortcut completely symmetric molecules will not have a dipole regardless of the polarity of the bonds
2 If one ldquosiderdquo of the molecule has more EN atoms than the other the molecule has a net dipole
We can use VSEPR to determine the polarity of a whole molecule
Molecular Dipoles
The C=O bonds have dipoles of equal magnitude but opposite direction so there is no net dipole moment
The O-H bonds have dipoles of equal magnitude that do not cancel each other so water has a net dipole moment
Molecular Dipoles (cont)
symmetric
symmetricasymmetri
c
F
ClF
Cl
Molecular Dipole Examplebull Write the Lewis dot and VESPR
structures for CF2Cl2 Does it have a dipole moment
C
F
FClCl
32 e-
Tetrahedral
Advanced VSEPR Application
Molecules with more than one central atomhellip methanol (CH3OH)
H CH
OH
H
tetrahedral e- pairstetrahedral shape
tetrahedral e- pairsbent shape
The VSEPR Table
2 AX2 BeF2 linear linear
3 AX3 BF3 trigonal planar
trigonal planar
AX2E O3 trigonal planar
bent
4 AX4 CH4 tetrahedral tetrahedral
AX3E NH3 tetrahedral pyramidal
AX2E2 H2O tetrahedral bent
e- pairs e- Geom Molec Geom
The VSEPR Table5 AX5 PF5 trigonal
bipyramidaltrigonalbipyramidal
AX4E SF4 trigonal bipyramidal
see saw
AX3E2 ClF3 trigonal bipyramidal
T-shaped
AX2E3 I3- trigonal bipyramidal
linear
6 AX6 SF6 octahedral octahedral
AX4E2 XeF4 octahedral square planar
e- pairs e- Geom Molec Geom
What is the expected shape of ICl2+
A linear
B bent
C tetrahedral
D square planar
AX2E2ICl Cl+
20 e-
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons occupies the region between the two nuclei
Rule 1 Maximum overlap The bond strength depends on the attraction of nuclei to the shared electrons soThe greater the orbital overlap the stronger the bond
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 2 Spins pair The two electrons in the overlap region occupy the same space and therefore must have opposite spins
There may be no more than 2 electrons in a molecular orbital
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Arranging Electron Pairs
bull Example CH4 (bonding pairs only)
bull Must consider both bonding and lone pairs when minimizing repulsion
Lewis Structure VSEPR Structure
H CH
HH
Arranging Electron Pairs (cont)
Example NH3 (both bonding and lone pairs)
Lewis Structure VSEPR Structure
Noteldquoelectron pair geometryrdquo vsldquomolecular shaperdquo
H NH
H
VSEPR Structure GuidelinesThe previous examples illustrate the strategy for applying VSEPR to predict molecular structure
1 Construct the Lewis Dot Structure2 Arrange bondinglone electron pairs in space such
that repulsions are minimized (electron pair geometry)
3 Name the molecular shape from the position of the atoms
VSEPR Shorthand
1 Refer to central atom as ldquoArdquo2 Attached atoms are referred to as ldquoXrdquo3 Lone pair are referred to as ldquoErdquo
ExamplesCH4 AX4NH3 AX3EH2O AX2E2BF3 AX3
VSEPR 2 electron pairsLinear (AX2) angle between bonds is 180deg
Example BeF2
180deg
Experiments show that molecules with multiple bonds can also be linear
Multiple bonds are treated as a single effective electron groupF Be F
F Be F
More than one central atom Determine shape around each
Be FF
VSEPR 3 electron pairsTrigonal Planar (AX3) angle between bonds is 120degExample BF3
Multiple bond is treated as a single effective electron group
FBFF
B
FF
F
120deg
VSEPR 4 electron pairs (cont)
Tetrahedral (AX4) angle between bonds is ~1095degExample CH4
1095degH C
HH
H
tetrahedral e- pair geometry AND tetrahedral molecular shape
Bonding vs Lone pairsBond angle in a tetrahedral arrangement of electron pairs may vary from 1095deg due to size differences between bonding and lone pair electron densities
bonding pair is constrained by two nuclear potentials more localized in space
lone pair is constrained by only one nuclear potential less localized (needs more room)
VSEPR 4 electron pairsTrigonal pyramidal (AX3E) Bond angles are lt1095deg and structure is nonplanar due to repulsion of lone pair
Example NH3
107deg
tetrahedral e- pair geometry trigonal pyramidal molecular shape
H NH
H
VSEPR 4 electron pairs (cont)
Classic example of tetrahedral angle shift from 1095deg is water (AX2E2)
ldquobentrdquo1045o
tetrahedral e- pair geometry bent molecular shape
VSEPR 4 electron pairs (cont)
Comparison of CH4 (AX4) NH3 (AX3E) and H2O (AX2E2)
AX2E
AX3E
AX2E2
1 Refer to central atom as ldquoArdquo2 Attached atoms are referred to as ldquoXrdquo3 Lone pair are referred to as ldquoErdquo
HH
H
H
CNOF
Central Atom
Compound Electron-Pair Geometry
Molecular Shape
Carbon C CH4 tetrahedral tetrahedralNitrogen N NH3 tetrahedral trigonal pyramidalOxygen O H2O tetrahedral bentFluorine F HF tetrahedral linear
Molecular vs Electron-Pair Geometry
What is the electron-pair geometry and the molecular shape for HCFS
a) trigonal planar bent
b) trigonal planar trigonal planar
c) tetrahedral trigonal planar
d) tetrahedral tetrahedral
C
S
H
F
VSEPR Beyond the OctetSystems with expanded valence shells will have five or six electron pairs around a central atom
P
Cl
Cl
Cl
Cl
Cl
FF
F
F
FFS
90deg
120deg SF
F
F
F
F
F 90deg
90deg
VSEPR 5 electron pairsbull Consider the structure of SF4 (34 e- AX4E)bull What is the optimum arrangement of electron pairs
around S
SFF
FF
S
FF
FF
SF
FF
F
lone-pair bond-pair
two at 90o two at 120o
Repulsive forces (strongest to weakest)lone-pairlone-pair gt lone-pairbond-pair gt bond-pairbond-pair
bond-pair bond-pair
three at 90o four at 90o one at
120othree at 90o three at 120o
Compare endash pair angles
VSEPR 5 electron pairsThe optimum structure maximizes the
angular separation of the lone pairs
I3- (AX2E3)
AX3E2
AX4E
AX2E3
5-electron-pair geometries
our previous example
Which of these is the more likely structure
Square Planar
VSEPR 6 electron pairs
See-saw
AX5E
AX4E2
6-electron-pair geometries
our previous example
Molecular Dipole Moments
1 Draw Lewis structures to determine 3D arrangement of atoms
Shortcut completely symmetric molecules will not have a dipole regardless of the polarity of the bonds
2 If one ldquosiderdquo of the molecule has more EN atoms than the other the molecule has a net dipole
We can use VSEPR to determine the polarity of a whole molecule
Molecular Dipoles
The C=O bonds have dipoles of equal magnitude but opposite direction so there is no net dipole moment
The O-H bonds have dipoles of equal magnitude that do not cancel each other so water has a net dipole moment
Molecular Dipoles (cont)
symmetric
symmetricasymmetri
c
F
ClF
Cl
Molecular Dipole Examplebull Write the Lewis dot and VESPR
structures for CF2Cl2 Does it have a dipole moment
C
F
FClCl
32 e-
Tetrahedral
Advanced VSEPR Application
Molecules with more than one central atomhellip methanol (CH3OH)
H CH
OH
H
tetrahedral e- pairstetrahedral shape
tetrahedral e- pairsbent shape
The VSEPR Table
2 AX2 BeF2 linear linear
3 AX3 BF3 trigonal planar
trigonal planar
AX2E O3 trigonal planar
bent
4 AX4 CH4 tetrahedral tetrahedral
AX3E NH3 tetrahedral pyramidal
AX2E2 H2O tetrahedral bent
e- pairs e- Geom Molec Geom
The VSEPR Table5 AX5 PF5 trigonal
bipyramidaltrigonalbipyramidal
AX4E SF4 trigonal bipyramidal
see saw
AX3E2 ClF3 trigonal bipyramidal
T-shaped
AX2E3 I3- trigonal bipyramidal
linear
6 AX6 SF6 octahedral octahedral
AX4E2 XeF4 octahedral square planar
e- pairs e- Geom Molec Geom
What is the expected shape of ICl2+
A linear
B bent
C tetrahedral
D square planar
AX2E2ICl Cl+
20 e-
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons occupies the region between the two nuclei
Rule 1 Maximum overlap The bond strength depends on the attraction of nuclei to the shared electrons soThe greater the orbital overlap the stronger the bond
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 2 Spins pair The two electrons in the overlap region occupy the same space and therefore must have opposite spins
There may be no more than 2 electrons in a molecular orbital
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Arranging Electron Pairs (cont)
Example NH3 (both bonding and lone pairs)
Lewis Structure VSEPR Structure
Noteldquoelectron pair geometryrdquo vsldquomolecular shaperdquo
H NH
H
VSEPR Structure GuidelinesThe previous examples illustrate the strategy for applying VSEPR to predict molecular structure
1 Construct the Lewis Dot Structure2 Arrange bondinglone electron pairs in space such
that repulsions are minimized (electron pair geometry)
3 Name the molecular shape from the position of the atoms
VSEPR Shorthand
1 Refer to central atom as ldquoArdquo2 Attached atoms are referred to as ldquoXrdquo3 Lone pair are referred to as ldquoErdquo
ExamplesCH4 AX4NH3 AX3EH2O AX2E2BF3 AX3
VSEPR 2 electron pairsLinear (AX2) angle between bonds is 180deg
Example BeF2
180deg
Experiments show that molecules with multiple bonds can also be linear
Multiple bonds are treated as a single effective electron groupF Be F
F Be F
More than one central atom Determine shape around each
Be FF
VSEPR 3 electron pairsTrigonal Planar (AX3) angle between bonds is 120degExample BF3
Multiple bond is treated as a single effective electron group
FBFF
B
FF
F
120deg
VSEPR 4 electron pairs (cont)
Tetrahedral (AX4) angle between bonds is ~1095degExample CH4
1095degH C
HH
H
tetrahedral e- pair geometry AND tetrahedral molecular shape
Bonding vs Lone pairsBond angle in a tetrahedral arrangement of electron pairs may vary from 1095deg due to size differences between bonding and lone pair electron densities
bonding pair is constrained by two nuclear potentials more localized in space
lone pair is constrained by only one nuclear potential less localized (needs more room)
VSEPR 4 electron pairsTrigonal pyramidal (AX3E) Bond angles are lt1095deg and structure is nonplanar due to repulsion of lone pair
Example NH3
107deg
tetrahedral e- pair geometry trigonal pyramidal molecular shape
H NH
H
VSEPR 4 electron pairs (cont)
Classic example of tetrahedral angle shift from 1095deg is water (AX2E2)
ldquobentrdquo1045o
tetrahedral e- pair geometry bent molecular shape
VSEPR 4 electron pairs (cont)
Comparison of CH4 (AX4) NH3 (AX3E) and H2O (AX2E2)
AX2E
AX3E
AX2E2
1 Refer to central atom as ldquoArdquo2 Attached atoms are referred to as ldquoXrdquo3 Lone pair are referred to as ldquoErdquo
HH
H
H
CNOF
Central Atom
Compound Electron-Pair Geometry
Molecular Shape
Carbon C CH4 tetrahedral tetrahedralNitrogen N NH3 tetrahedral trigonal pyramidalOxygen O H2O tetrahedral bentFluorine F HF tetrahedral linear
Molecular vs Electron-Pair Geometry
What is the electron-pair geometry and the molecular shape for HCFS
a) trigonal planar bent
b) trigonal planar trigonal planar
c) tetrahedral trigonal planar
d) tetrahedral tetrahedral
C
S
H
F
VSEPR Beyond the OctetSystems with expanded valence shells will have five or six electron pairs around a central atom
P
Cl
Cl
Cl
Cl
Cl
FF
F
F
FFS
90deg
120deg SF
F
F
F
F
F 90deg
90deg
VSEPR 5 electron pairsbull Consider the structure of SF4 (34 e- AX4E)bull What is the optimum arrangement of electron pairs
around S
SFF
FF
S
FF
FF
SF
FF
F
lone-pair bond-pair
two at 90o two at 120o
Repulsive forces (strongest to weakest)lone-pairlone-pair gt lone-pairbond-pair gt bond-pairbond-pair
bond-pair bond-pair
three at 90o four at 90o one at
120othree at 90o three at 120o
Compare endash pair angles
VSEPR 5 electron pairsThe optimum structure maximizes the
angular separation of the lone pairs
I3- (AX2E3)
AX3E2
AX4E
AX2E3
5-electron-pair geometries
our previous example
Which of these is the more likely structure
Square Planar
VSEPR 6 electron pairs
See-saw
AX5E
AX4E2
6-electron-pair geometries
our previous example
Molecular Dipole Moments
1 Draw Lewis structures to determine 3D arrangement of atoms
Shortcut completely symmetric molecules will not have a dipole regardless of the polarity of the bonds
2 If one ldquosiderdquo of the molecule has more EN atoms than the other the molecule has a net dipole
We can use VSEPR to determine the polarity of a whole molecule
Molecular Dipoles
The C=O bonds have dipoles of equal magnitude but opposite direction so there is no net dipole moment
The O-H bonds have dipoles of equal magnitude that do not cancel each other so water has a net dipole moment
Molecular Dipoles (cont)
symmetric
symmetricasymmetri
c
F
ClF
Cl
Molecular Dipole Examplebull Write the Lewis dot and VESPR
structures for CF2Cl2 Does it have a dipole moment
C
F
FClCl
32 e-
Tetrahedral
Advanced VSEPR Application
Molecules with more than one central atomhellip methanol (CH3OH)
H CH
OH
H
tetrahedral e- pairstetrahedral shape
tetrahedral e- pairsbent shape
The VSEPR Table
2 AX2 BeF2 linear linear
3 AX3 BF3 trigonal planar
trigonal planar
AX2E O3 trigonal planar
bent
4 AX4 CH4 tetrahedral tetrahedral
AX3E NH3 tetrahedral pyramidal
AX2E2 H2O tetrahedral bent
e- pairs e- Geom Molec Geom
The VSEPR Table5 AX5 PF5 trigonal
bipyramidaltrigonalbipyramidal
AX4E SF4 trigonal bipyramidal
see saw
AX3E2 ClF3 trigonal bipyramidal
T-shaped
AX2E3 I3- trigonal bipyramidal
linear
6 AX6 SF6 octahedral octahedral
AX4E2 XeF4 octahedral square planar
e- pairs e- Geom Molec Geom
What is the expected shape of ICl2+
A linear
B bent
C tetrahedral
D square planar
AX2E2ICl Cl+
20 e-
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons occupies the region between the two nuclei
Rule 1 Maximum overlap The bond strength depends on the attraction of nuclei to the shared electrons soThe greater the orbital overlap the stronger the bond
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 2 Spins pair The two electrons in the overlap region occupy the same space and therefore must have opposite spins
There may be no more than 2 electrons in a molecular orbital
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
VSEPR Structure GuidelinesThe previous examples illustrate the strategy for applying VSEPR to predict molecular structure
1 Construct the Lewis Dot Structure2 Arrange bondinglone electron pairs in space such
that repulsions are minimized (electron pair geometry)
3 Name the molecular shape from the position of the atoms
VSEPR Shorthand
1 Refer to central atom as ldquoArdquo2 Attached atoms are referred to as ldquoXrdquo3 Lone pair are referred to as ldquoErdquo
ExamplesCH4 AX4NH3 AX3EH2O AX2E2BF3 AX3
VSEPR 2 electron pairsLinear (AX2) angle between bonds is 180deg
Example BeF2
180deg
Experiments show that molecules with multiple bonds can also be linear
Multiple bonds are treated as a single effective electron groupF Be F
F Be F
More than one central atom Determine shape around each
Be FF
VSEPR 3 electron pairsTrigonal Planar (AX3) angle between bonds is 120degExample BF3
Multiple bond is treated as a single effective electron group
FBFF
B
FF
F
120deg
VSEPR 4 electron pairs (cont)
Tetrahedral (AX4) angle between bonds is ~1095degExample CH4
1095degH C
HH
H
tetrahedral e- pair geometry AND tetrahedral molecular shape
Bonding vs Lone pairsBond angle in a tetrahedral arrangement of electron pairs may vary from 1095deg due to size differences between bonding and lone pair electron densities
bonding pair is constrained by two nuclear potentials more localized in space
lone pair is constrained by only one nuclear potential less localized (needs more room)
VSEPR 4 electron pairsTrigonal pyramidal (AX3E) Bond angles are lt1095deg and structure is nonplanar due to repulsion of lone pair
Example NH3
107deg
tetrahedral e- pair geometry trigonal pyramidal molecular shape
H NH
H
VSEPR 4 electron pairs (cont)
Classic example of tetrahedral angle shift from 1095deg is water (AX2E2)
ldquobentrdquo1045o
tetrahedral e- pair geometry bent molecular shape
VSEPR 4 electron pairs (cont)
Comparison of CH4 (AX4) NH3 (AX3E) and H2O (AX2E2)
AX2E
AX3E
AX2E2
1 Refer to central atom as ldquoArdquo2 Attached atoms are referred to as ldquoXrdquo3 Lone pair are referred to as ldquoErdquo
HH
H
H
CNOF
Central Atom
Compound Electron-Pair Geometry
Molecular Shape
Carbon C CH4 tetrahedral tetrahedralNitrogen N NH3 tetrahedral trigonal pyramidalOxygen O H2O tetrahedral bentFluorine F HF tetrahedral linear
Molecular vs Electron-Pair Geometry
What is the electron-pair geometry and the molecular shape for HCFS
a) trigonal planar bent
b) trigonal planar trigonal planar
c) tetrahedral trigonal planar
d) tetrahedral tetrahedral
C
S
H
F
VSEPR Beyond the OctetSystems with expanded valence shells will have five or six electron pairs around a central atom
P
Cl
Cl
Cl
Cl
Cl
FF
F
F
FFS
90deg
120deg SF
F
F
F
F
F 90deg
90deg
VSEPR 5 electron pairsbull Consider the structure of SF4 (34 e- AX4E)bull What is the optimum arrangement of electron pairs
around S
SFF
FF
S
FF
FF
SF
FF
F
lone-pair bond-pair
two at 90o two at 120o
Repulsive forces (strongest to weakest)lone-pairlone-pair gt lone-pairbond-pair gt bond-pairbond-pair
bond-pair bond-pair
three at 90o four at 90o one at
120othree at 90o three at 120o
Compare endash pair angles
VSEPR 5 electron pairsThe optimum structure maximizes the
angular separation of the lone pairs
I3- (AX2E3)
AX3E2
AX4E
AX2E3
5-electron-pair geometries
our previous example
Which of these is the more likely structure
Square Planar
VSEPR 6 electron pairs
See-saw
AX5E
AX4E2
6-electron-pair geometries
our previous example
Molecular Dipole Moments
1 Draw Lewis structures to determine 3D arrangement of atoms
Shortcut completely symmetric molecules will not have a dipole regardless of the polarity of the bonds
2 If one ldquosiderdquo of the molecule has more EN atoms than the other the molecule has a net dipole
We can use VSEPR to determine the polarity of a whole molecule
Molecular Dipoles
The C=O bonds have dipoles of equal magnitude but opposite direction so there is no net dipole moment
The O-H bonds have dipoles of equal magnitude that do not cancel each other so water has a net dipole moment
Molecular Dipoles (cont)
symmetric
symmetricasymmetri
c
F
ClF
Cl
Molecular Dipole Examplebull Write the Lewis dot and VESPR
structures for CF2Cl2 Does it have a dipole moment
C
F
FClCl
32 e-
Tetrahedral
Advanced VSEPR Application
Molecules with more than one central atomhellip methanol (CH3OH)
H CH
OH
H
tetrahedral e- pairstetrahedral shape
tetrahedral e- pairsbent shape
The VSEPR Table
2 AX2 BeF2 linear linear
3 AX3 BF3 trigonal planar
trigonal planar
AX2E O3 trigonal planar
bent
4 AX4 CH4 tetrahedral tetrahedral
AX3E NH3 tetrahedral pyramidal
AX2E2 H2O tetrahedral bent
e- pairs e- Geom Molec Geom
The VSEPR Table5 AX5 PF5 trigonal
bipyramidaltrigonalbipyramidal
AX4E SF4 trigonal bipyramidal
see saw
AX3E2 ClF3 trigonal bipyramidal
T-shaped
AX2E3 I3- trigonal bipyramidal
linear
6 AX6 SF6 octahedral octahedral
AX4E2 XeF4 octahedral square planar
e- pairs e- Geom Molec Geom
What is the expected shape of ICl2+
A linear
B bent
C tetrahedral
D square planar
AX2E2ICl Cl+
20 e-
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons occupies the region between the two nuclei
Rule 1 Maximum overlap The bond strength depends on the attraction of nuclei to the shared electrons soThe greater the orbital overlap the stronger the bond
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 2 Spins pair The two electrons in the overlap region occupy the same space and therefore must have opposite spins
There may be no more than 2 electrons in a molecular orbital
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
VSEPR 2 electron pairsLinear (AX2) angle between bonds is 180deg
Example BeF2
180deg
Experiments show that molecules with multiple bonds can also be linear
Multiple bonds are treated as a single effective electron groupF Be F
F Be F
More than one central atom Determine shape around each
Be FF
VSEPR 3 electron pairsTrigonal Planar (AX3) angle between bonds is 120degExample BF3
Multiple bond is treated as a single effective electron group
FBFF
B
FF
F
120deg
VSEPR 4 electron pairs (cont)
Tetrahedral (AX4) angle between bonds is ~1095degExample CH4
1095degH C
HH
H
tetrahedral e- pair geometry AND tetrahedral molecular shape
Bonding vs Lone pairsBond angle in a tetrahedral arrangement of electron pairs may vary from 1095deg due to size differences between bonding and lone pair electron densities
bonding pair is constrained by two nuclear potentials more localized in space
lone pair is constrained by only one nuclear potential less localized (needs more room)
VSEPR 4 electron pairsTrigonal pyramidal (AX3E) Bond angles are lt1095deg and structure is nonplanar due to repulsion of lone pair
Example NH3
107deg
tetrahedral e- pair geometry trigonal pyramidal molecular shape
H NH
H
VSEPR 4 electron pairs (cont)
Classic example of tetrahedral angle shift from 1095deg is water (AX2E2)
ldquobentrdquo1045o
tetrahedral e- pair geometry bent molecular shape
VSEPR 4 electron pairs (cont)
Comparison of CH4 (AX4) NH3 (AX3E) and H2O (AX2E2)
AX2E
AX3E
AX2E2
1 Refer to central atom as ldquoArdquo2 Attached atoms are referred to as ldquoXrdquo3 Lone pair are referred to as ldquoErdquo
HH
H
H
CNOF
Central Atom
Compound Electron-Pair Geometry
Molecular Shape
Carbon C CH4 tetrahedral tetrahedralNitrogen N NH3 tetrahedral trigonal pyramidalOxygen O H2O tetrahedral bentFluorine F HF tetrahedral linear
Molecular vs Electron-Pair Geometry
What is the electron-pair geometry and the molecular shape for HCFS
a) trigonal planar bent
b) trigonal planar trigonal planar
c) tetrahedral trigonal planar
d) tetrahedral tetrahedral
C
S
H
F
VSEPR Beyond the OctetSystems with expanded valence shells will have five or six electron pairs around a central atom
P
Cl
Cl
Cl
Cl
Cl
FF
F
F
FFS
90deg
120deg SF
F
F
F
F
F 90deg
90deg
VSEPR 5 electron pairsbull Consider the structure of SF4 (34 e- AX4E)bull What is the optimum arrangement of electron pairs
around S
SFF
FF
S
FF
FF
SF
FF
F
lone-pair bond-pair
two at 90o two at 120o
Repulsive forces (strongest to weakest)lone-pairlone-pair gt lone-pairbond-pair gt bond-pairbond-pair
bond-pair bond-pair
three at 90o four at 90o one at
120othree at 90o three at 120o
Compare endash pair angles
VSEPR 5 electron pairsThe optimum structure maximizes the
angular separation of the lone pairs
I3- (AX2E3)
AX3E2
AX4E
AX2E3
5-electron-pair geometries
our previous example
Which of these is the more likely structure
Square Planar
VSEPR 6 electron pairs
See-saw
AX5E
AX4E2
6-electron-pair geometries
our previous example
Molecular Dipole Moments
1 Draw Lewis structures to determine 3D arrangement of atoms
Shortcut completely symmetric molecules will not have a dipole regardless of the polarity of the bonds
2 If one ldquosiderdquo of the molecule has more EN atoms than the other the molecule has a net dipole
We can use VSEPR to determine the polarity of a whole molecule
Molecular Dipoles
The C=O bonds have dipoles of equal magnitude but opposite direction so there is no net dipole moment
The O-H bonds have dipoles of equal magnitude that do not cancel each other so water has a net dipole moment
Molecular Dipoles (cont)
symmetric
symmetricasymmetri
c
F
ClF
Cl
Molecular Dipole Examplebull Write the Lewis dot and VESPR
structures for CF2Cl2 Does it have a dipole moment
C
F
FClCl
32 e-
Tetrahedral
Advanced VSEPR Application
Molecules with more than one central atomhellip methanol (CH3OH)
H CH
OH
H
tetrahedral e- pairstetrahedral shape
tetrahedral e- pairsbent shape
The VSEPR Table
2 AX2 BeF2 linear linear
3 AX3 BF3 trigonal planar
trigonal planar
AX2E O3 trigonal planar
bent
4 AX4 CH4 tetrahedral tetrahedral
AX3E NH3 tetrahedral pyramidal
AX2E2 H2O tetrahedral bent
e- pairs e- Geom Molec Geom
The VSEPR Table5 AX5 PF5 trigonal
bipyramidaltrigonalbipyramidal
AX4E SF4 trigonal bipyramidal
see saw
AX3E2 ClF3 trigonal bipyramidal
T-shaped
AX2E3 I3- trigonal bipyramidal
linear
6 AX6 SF6 octahedral octahedral
AX4E2 XeF4 octahedral square planar
e- pairs e- Geom Molec Geom
What is the expected shape of ICl2+
A linear
B bent
C tetrahedral
D square planar
AX2E2ICl Cl+
20 e-
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons occupies the region between the two nuclei
Rule 1 Maximum overlap The bond strength depends on the attraction of nuclei to the shared electrons soThe greater the orbital overlap the stronger the bond
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 2 Spins pair The two electrons in the overlap region occupy the same space and therefore must have opposite spins
There may be no more than 2 electrons in a molecular orbital
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
VSEPR 3 electron pairsTrigonal Planar (AX3) angle between bonds is 120degExample BF3
Multiple bond is treated as a single effective electron group
FBFF
B
FF
F
120deg
VSEPR 4 electron pairs (cont)
Tetrahedral (AX4) angle between bonds is ~1095degExample CH4
1095degH C
HH
H
tetrahedral e- pair geometry AND tetrahedral molecular shape
Bonding vs Lone pairsBond angle in a tetrahedral arrangement of electron pairs may vary from 1095deg due to size differences between bonding and lone pair electron densities
bonding pair is constrained by two nuclear potentials more localized in space
lone pair is constrained by only one nuclear potential less localized (needs more room)
VSEPR 4 electron pairsTrigonal pyramidal (AX3E) Bond angles are lt1095deg and structure is nonplanar due to repulsion of lone pair
Example NH3
107deg
tetrahedral e- pair geometry trigonal pyramidal molecular shape
H NH
H
VSEPR 4 electron pairs (cont)
Classic example of tetrahedral angle shift from 1095deg is water (AX2E2)
ldquobentrdquo1045o
tetrahedral e- pair geometry bent molecular shape
VSEPR 4 electron pairs (cont)
Comparison of CH4 (AX4) NH3 (AX3E) and H2O (AX2E2)
AX2E
AX3E
AX2E2
1 Refer to central atom as ldquoArdquo2 Attached atoms are referred to as ldquoXrdquo3 Lone pair are referred to as ldquoErdquo
HH
H
H
CNOF
Central Atom
Compound Electron-Pair Geometry
Molecular Shape
Carbon C CH4 tetrahedral tetrahedralNitrogen N NH3 tetrahedral trigonal pyramidalOxygen O H2O tetrahedral bentFluorine F HF tetrahedral linear
Molecular vs Electron-Pair Geometry
What is the electron-pair geometry and the molecular shape for HCFS
a) trigonal planar bent
b) trigonal planar trigonal planar
c) tetrahedral trigonal planar
d) tetrahedral tetrahedral
C
S
H
F
VSEPR Beyond the OctetSystems with expanded valence shells will have five or six electron pairs around a central atom
P
Cl
Cl
Cl
Cl
Cl
FF
F
F
FFS
90deg
120deg SF
F
F
F
F
F 90deg
90deg
VSEPR 5 electron pairsbull Consider the structure of SF4 (34 e- AX4E)bull What is the optimum arrangement of electron pairs
around S
SFF
FF
S
FF
FF
SF
FF
F
lone-pair bond-pair
two at 90o two at 120o
Repulsive forces (strongest to weakest)lone-pairlone-pair gt lone-pairbond-pair gt bond-pairbond-pair
bond-pair bond-pair
three at 90o four at 90o one at
120othree at 90o three at 120o
Compare endash pair angles
VSEPR 5 electron pairsThe optimum structure maximizes the
angular separation of the lone pairs
I3- (AX2E3)
AX3E2
AX4E
AX2E3
5-electron-pair geometries
our previous example
Which of these is the more likely structure
Square Planar
VSEPR 6 electron pairs
See-saw
AX5E
AX4E2
6-electron-pair geometries
our previous example
Molecular Dipole Moments
1 Draw Lewis structures to determine 3D arrangement of atoms
Shortcut completely symmetric molecules will not have a dipole regardless of the polarity of the bonds
2 If one ldquosiderdquo of the molecule has more EN atoms than the other the molecule has a net dipole
We can use VSEPR to determine the polarity of a whole molecule
Molecular Dipoles
The C=O bonds have dipoles of equal magnitude but opposite direction so there is no net dipole moment
The O-H bonds have dipoles of equal magnitude that do not cancel each other so water has a net dipole moment
Molecular Dipoles (cont)
symmetric
symmetricasymmetri
c
F
ClF
Cl
Molecular Dipole Examplebull Write the Lewis dot and VESPR
structures for CF2Cl2 Does it have a dipole moment
C
F
FClCl
32 e-
Tetrahedral
Advanced VSEPR Application
Molecules with more than one central atomhellip methanol (CH3OH)
H CH
OH
H
tetrahedral e- pairstetrahedral shape
tetrahedral e- pairsbent shape
The VSEPR Table
2 AX2 BeF2 linear linear
3 AX3 BF3 trigonal planar
trigonal planar
AX2E O3 trigonal planar
bent
4 AX4 CH4 tetrahedral tetrahedral
AX3E NH3 tetrahedral pyramidal
AX2E2 H2O tetrahedral bent
e- pairs e- Geom Molec Geom
The VSEPR Table5 AX5 PF5 trigonal
bipyramidaltrigonalbipyramidal
AX4E SF4 trigonal bipyramidal
see saw
AX3E2 ClF3 trigonal bipyramidal
T-shaped
AX2E3 I3- trigonal bipyramidal
linear
6 AX6 SF6 octahedral octahedral
AX4E2 XeF4 octahedral square planar
e- pairs e- Geom Molec Geom
What is the expected shape of ICl2+
A linear
B bent
C tetrahedral
D square planar
AX2E2ICl Cl+
20 e-
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons occupies the region between the two nuclei
Rule 1 Maximum overlap The bond strength depends on the attraction of nuclei to the shared electrons soThe greater the orbital overlap the stronger the bond
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 2 Spins pair The two electrons in the overlap region occupy the same space and therefore must have opposite spins
There may be no more than 2 electrons in a molecular orbital
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
VSEPR 4 electron pairs (cont)
Tetrahedral (AX4) angle between bonds is ~1095degExample CH4
1095degH C
HH
H
tetrahedral e- pair geometry AND tetrahedral molecular shape
Bonding vs Lone pairsBond angle in a tetrahedral arrangement of electron pairs may vary from 1095deg due to size differences between bonding and lone pair electron densities
bonding pair is constrained by two nuclear potentials more localized in space
lone pair is constrained by only one nuclear potential less localized (needs more room)
VSEPR 4 electron pairsTrigonal pyramidal (AX3E) Bond angles are lt1095deg and structure is nonplanar due to repulsion of lone pair
Example NH3
107deg
tetrahedral e- pair geometry trigonal pyramidal molecular shape
H NH
H
VSEPR 4 electron pairs (cont)
Classic example of tetrahedral angle shift from 1095deg is water (AX2E2)
ldquobentrdquo1045o
tetrahedral e- pair geometry bent molecular shape
VSEPR 4 electron pairs (cont)
Comparison of CH4 (AX4) NH3 (AX3E) and H2O (AX2E2)
AX2E
AX3E
AX2E2
1 Refer to central atom as ldquoArdquo2 Attached atoms are referred to as ldquoXrdquo3 Lone pair are referred to as ldquoErdquo
HH
H
H
CNOF
Central Atom
Compound Electron-Pair Geometry
Molecular Shape
Carbon C CH4 tetrahedral tetrahedralNitrogen N NH3 tetrahedral trigonal pyramidalOxygen O H2O tetrahedral bentFluorine F HF tetrahedral linear
Molecular vs Electron-Pair Geometry
What is the electron-pair geometry and the molecular shape for HCFS
a) trigonal planar bent
b) trigonal planar trigonal planar
c) tetrahedral trigonal planar
d) tetrahedral tetrahedral
C
S
H
F
VSEPR Beyond the OctetSystems with expanded valence shells will have five or six electron pairs around a central atom
P
Cl
Cl
Cl
Cl
Cl
FF
F
F
FFS
90deg
120deg SF
F
F
F
F
F 90deg
90deg
VSEPR 5 electron pairsbull Consider the structure of SF4 (34 e- AX4E)bull What is the optimum arrangement of electron pairs
around S
SFF
FF
S
FF
FF
SF
FF
F
lone-pair bond-pair
two at 90o two at 120o
Repulsive forces (strongest to weakest)lone-pairlone-pair gt lone-pairbond-pair gt bond-pairbond-pair
bond-pair bond-pair
three at 90o four at 90o one at
120othree at 90o three at 120o
Compare endash pair angles
VSEPR 5 electron pairsThe optimum structure maximizes the
angular separation of the lone pairs
I3- (AX2E3)
AX3E2
AX4E
AX2E3
5-electron-pair geometries
our previous example
Which of these is the more likely structure
Square Planar
VSEPR 6 electron pairs
See-saw
AX5E
AX4E2
6-electron-pair geometries
our previous example
Molecular Dipole Moments
1 Draw Lewis structures to determine 3D arrangement of atoms
Shortcut completely symmetric molecules will not have a dipole regardless of the polarity of the bonds
2 If one ldquosiderdquo of the molecule has more EN atoms than the other the molecule has a net dipole
We can use VSEPR to determine the polarity of a whole molecule
Molecular Dipoles
The C=O bonds have dipoles of equal magnitude but opposite direction so there is no net dipole moment
The O-H bonds have dipoles of equal magnitude that do not cancel each other so water has a net dipole moment
Molecular Dipoles (cont)
symmetric
symmetricasymmetri
c
F
ClF
Cl
Molecular Dipole Examplebull Write the Lewis dot and VESPR
structures for CF2Cl2 Does it have a dipole moment
C
F
FClCl
32 e-
Tetrahedral
Advanced VSEPR Application
Molecules with more than one central atomhellip methanol (CH3OH)
H CH
OH
H
tetrahedral e- pairstetrahedral shape
tetrahedral e- pairsbent shape
The VSEPR Table
2 AX2 BeF2 linear linear
3 AX3 BF3 trigonal planar
trigonal planar
AX2E O3 trigonal planar
bent
4 AX4 CH4 tetrahedral tetrahedral
AX3E NH3 tetrahedral pyramidal
AX2E2 H2O tetrahedral bent
e- pairs e- Geom Molec Geom
The VSEPR Table5 AX5 PF5 trigonal
bipyramidaltrigonalbipyramidal
AX4E SF4 trigonal bipyramidal
see saw
AX3E2 ClF3 trigonal bipyramidal
T-shaped
AX2E3 I3- trigonal bipyramidal
linear
6 AX6 SF6 octahedral octahedral
AX4E2 XeF4 octahedral square planar
e- pairs e- Geom Molec Geom
What is the expected shape of ICl2+
A linear
B bent
C tetrahedral
D square planar
AX2E2ICl Cl+
20 e-
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons occupies the region between the two nuclei
Rule 1 Maximum overlap The bond strength depends on the attraction of nuclei to the shared electrons soThe greater the orbital overlap the stronger the bond
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 2 Spins pair The two electrons in the overlap region occupy the same space and therefore must have opposite spins
There may be no more than 2 electrons in a molecular orbital
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Bonding vs Lone pairsBond angle in a tetrahedral arrangement of electron pairs may vary from 1095deg due to size differences between bonding and lone pair electron densities
bonding pair is constrained by two nuclear potentials more localized in space
lone pair is constrained by only one nuclear potential less localized (needs more room)
VSEPR 4 electron pairsTrigonal pyramidal (AX3E) Bond angles are lt1095deg and structure is nonplanar due to repulsion of lone pair
Example NH3
107deg
tetrahedral e- pair geometry trigonal pyramidal molecular shape
H NH
H
VSEPR 4 electron pairs (cont)
Classic example of tetrahedral angle shift from 1095deg is water (AX2E2)
ldquobentrdquo1045o
tetrahedral e- pair geometry bent molecular shape
VSEPR 4 electron pairs (cont)
Comparison of CH4 (AX4) NH3 (AX3E) and H2O (AX2E2)
AX2E
AX3E
AX2E2
1 Refer to central atom as ldquoArdquo2 Attached atoms are referred to as ldquoXrdquo3 Lone pair are referred to as ldquoErdquo
HH
H
H
CNOF
Central Atom
Compound Electron-Pair Geometry
Molecular Shape
Carbon C CH4 tetrahedral tetrahedralNitrogen N NH3 tetrahedral trigonal pyramidalOxygen O H2O tetrahedral bentFluorine F HF tetrahedral linear
Molecular vs Electron-Pair Geometry
What is the electron-pair geometry and the molecular shape for HCFS
a) trigonal planar bent
b) trigonal planar trigonal planar
c) tetrahedral trigonal planar
d) tetrahedral tetrahedral
C
S
H
F
VSEPR Beyond the OctetSystems with expanded valence shells will have five or six electron pairs around a central atom
P
Cl
Cl
Cl
Cl
Cl
FF
F
F
FFS
90deg
120deg SF
F
F
F
F
F 90deg
90deg
VSEPR 5 electron pairsbull Consider the structure of SF4 (34 e- AX4E)bull What is the optimum arrangement of electron pairs
around S
SFF
FF
S
FF
FF
SF
FF
F
lone-pair bond-pair
two at 90o two at 120o
Repulsive forces (strongest to weakest)lone-pairlone-pair gt lone-pairbond-pair gt bond-pairbond-pair
bond-pair bond-pair
three at 90o four at 90o one at
120othree at 90o three at 120o
Compare endash pair angles
VSEPR 5 electron pairsThe optimum structure maximizes the
angular separation of the lone pairs
I3- (AX2E3)
AX3E2
AX4E
AX2E3
5-electron-pair geometries
our previous example
Which of these is the more likely structure
Square Planar
VSEPR 6 electron pairs
See-saw
AX5E
AX4E2
6-electron-pair geometries
our previous example
Molecular Dipole Moments
1 Draw Lewis structures to determine 3D arrangement of atoms
Shortcut completely symmetric molecules will not have a dipole regardless of the polarity of the bonds
2 If one ldquosiderdquo of the molecule has more EN atoms than the other the molecule has a net dipole
We can use VSEPR to determine the polarity of a whole molecule
Molecular Dipoles
The C=O bonds have dipoles of equal magnitude but opposite direction so there is no net dipole moment
The O-H bonds have dipoles of equal magnitude that do not cancel each other so water has a net dipole moment
Molecular Dipoles (cont)
symmetric
symmetricasymmetri
c
F
ClF
Cl
Molecular Dipole Examplebull Write the Lewis dot and VESPR
structures for CF2Cl2 Does it have a dipole moment
C
F
FClCl
32 e-
Tetrahedral
Advanced VSEPR Application
Molecules with more than one central atomhellip methanol (CH3OH)
H CH
OH
H
tetrahedral e- pairstetrahedral shape
tetrahedral e- pairsbent shape
The VSEPR Table
2 AX2 BeF2 linear linear
3 AX3 BF3 trigonal planar
trigonal planar
AX2E O3 trigonal planar
bent
4 AX4 CH4 tetrahedral tetrahedral
AX3E NH3 tetrahedral pyramidal
AX2E2 H2O tetrahedral bent
e- pairs e- Geom Molec Geom
The VSEPR Table5 AX5 PF5 trigonal
bipyramidaltrigonalbipyramidal
AX4E SF4 trigonal bipyramidal
see saw
AX3E2 ClF3 trigonal bipyramidal
T-shaped
AX2E3 I3- trigonal bipyramidal
linear
6 AX6 SF6 octahedral octahedral
AX4E2 XeF4 octahedral square planar
e- pairs e- Geom Molec Geom
What is the expected shape of ICl2+
A linear
B bent
C tetrahedral
D square planar
AX2E2ICl Cl+
20 e-
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons occupies the region between the two nuclei
Rule 1 Maximum overlap The bond strength depends on the attraction of nuclei to the shared electrons soThe greater the orbital overlap the stronger the bond
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 2 Spins pair The two electrons in the overlap region occupy the same space and therefore must have opposite spins
There may be no more than 2 electrons in a molecular orbital
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
VSEPR 4 electron pairsTrigonal pyramidal (AX3E) Bond angles are lt1095deg and structure is nonplanar due to repulsion of lone pair
Example NH3
107deg
tetrahedral e- pair geometry trigonal pyramidal molecular shape
H NH
H
VSEPR 4 electron pairs (cont)
Classic example of tetrahedral angle shift from 1095deg is water (AX2E2)
ldquobentrdquo1045o
tetrahedral e- pair geometry bent molecular shape
VSEPR 4 electron pairs (cont)
Comparison of CH4 (AX4) NH3 (AX3E) and H2O (AX2E2)
AX2E
AX3E
AX2E2
1 Refer to central atom as ldquoArdquo2 Attached atoms are referred to as ldquoXrdquo3 Lone pair are referred to as ldquoErdquo
HH
H
H
CNOF
Central Atom
Compound Electron-Pair Geometry
Molecular Shape
Carbon C CH4 tetrahedral tetrahedralNitrogen N NH3 tetrahedral trigonal pyramidalOxygen O H2O tetrahedral bentFluorine F HF tetrahedral linear
Molecular vs Electron-Pair Geometry
What is the electron-pair geometry and the molecular shape for HCFS
a) trigonal planar bent
b) trigonal planar trigonal planar
c) tetrahedral trigonal planar
d) tetrahedral tetrahedral
C
S
H
F
VSEPR Beyond the OctetSystems with expanded valence shells will have five or six electron pairs around a central atom
P
Cl
Cl
Cl
Cl
Cl
FF
F
F
FFS
90deg
120deg SF
F
F
F
F
F 90deg
90deg
VSEPR 5 electron pairsbull Consider the structure of SF4 (34 e- AX4E)bull What is the optimum arrangement of electron pairs
around S
SFF
FF
S
FF
FF
SF
FF
F
lone-pair bond-pair
two at 90o two at 120o
Repulsive forces (strongest to weakest)lone-pairlone-pair gt lone-pairbond-pair gt bond-pairbond-pair
bond-pair bond-pair
three at 90o four at 90o one at
120othree at 90o three at 120o
Compare endash pair angles
VSEPR 5 electron pairsThe optimum structure maximizes the
angular separation of the lone pairs
I3- (AX2E3)
AX3E2
AX4E
AX2E3
5-electron-pair geometries
our previous example
Which of these is the more likely structure
Square Planar
VSEPR 6 electron pairs
See-saw
AX5E
AX4E2
6-electron-pair geometries
our previous example
Molecular Dipole Moments
1 Draw Lewis structures to determine 3D arrangement of atoms
Shortcut completely symmetric molecules will not have a dipole regardless of the polarity of the bonds
2 If one ldquosiderdquo of the molecule has more EN atoms than the other the molecule has a net dipole
We can use VSEPR to determine the polarity of a whole molecule
Molecular Dipoles
The C=O bonds have dipoles of equal magnitude but opposite direction so there is no net dipole moment
The O-H bonds have dipoles of equal magnitude that do not cancel each other so water has a net dipole moment
Molecular Dipoles (cont)
symmetric
symmetricasymmetri
c
F
ClF
Cl
Molecular Dipole Examplebull Write the Lewis dot and VESPR
structures for CF2Cl2 Does it have a dipole moment
C
F
FClCl
32 e-
Tetrahedral
Advanced VSEPR Application
Molecules with more than one central atomhellip methanol (CH3OH)
H CH
OH
H
tetrahedral e- pairstetrahedral shape
tetrahedral e- pairsbent shape
The VSEPR Table
2 AX2 BeF2 linear linear
3 AX3 BF3 trigonal planar
trigonal planar
AX2E O3 trigonal planar
bent
4 AX4 CH4 tetrahedral tetrahedral
AX3E NH3 tetrahedral pyramidal
AX2E2 H2O tetrahedral bent
e- pairs e- Geom Molec Geom
The VSEPR Table5 AX5 PF5 trigonal
bipyramidaltrigonalbipyramidal
AX4E SF4 trigonal bipyramidal
see saw
AX3E2 ClF3 trigonal bipyramidal
T-shaped
AX2E3 I3- trigonal bipyramidal
linear
6 AX6 SF6 octahedral octahedral
AX4E2 XeF4 octahedral square planar
e- pairs e- Geom Molec Geom
What is the expected shape of ICl2+
A linear
B bent
C tetrahedral
D square planar
AX2E2ICl Cl+
20 e-
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons occupies the region between the two nuclei
Rule 1 Maximum overlap The bond strength depends on the attraction of nuclei to the shared electrons soThe greater the orbital overlap the stronger the bond
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 2 Spins pair The two electrons in the overlap region occupy the same space and therefore must have opposite spins
There may be no more than 2 electrons in a molecular orbital
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
VSEPR 4 electron pairs (cont)
Classic example of tetrahedral angle shift from 1095deg is water (AX2E2)
ldquobentrdquo1045o
tetrahedral e- pair geometry bent molecular shape
VSEPR 4 electron pairs (cont)
Comparison of CH4 (AX4) NH3 (AX3E) and H2O (AX2E2)
AX2E
AX3E
AX2E2
1 Refer to central atom as ldquoArdquo2 Attached atoms are referred to as ldquoXrdquo3 Lone pair are referred to as ldquoErdquo
HH
H
H
CNOF
Central Atom
Compound Electron-Pair Geometry
Molecular Shape
Carbon C CH4 tetrahedral tetrahedralNitrogen N NH3 tetrahedral trigonal pyramidalOxygen O H2O tetrahedral bentFluorine F HF tetrahedral linear
Molecular vs Electron-Pair Geometry
What is the electron-pair geometry and the molecular shape for HCFS
a) trigonal planar bent
b) trigonal planar trigonal planar
c) tetrahedral trigonal planar
d) tetrahedral tetrahedral
C
S
H
F
VSEPR Beyond the OctetSystems with expanded valence shells will have five or six electron pairs around a central atom
P
Cl
Cl
Cl
Cl
Cl
FF
F
F
FFS
90deg
120deg SF
F
F
F
F
F 90deg
90deg
VSEPR 5 electron pairsbull Consider the structure of SF4 (34 e- AX4E)bull What is the optimum arrangement of electron pairs
around S
SFF
FF
S
FF
FF
SF
FF
F
lone-pair bond-pair
two at 90o two at 120o
Repulsive forces (strongest to weakest)lone-pairlone-pair gt lone-pairbond-pair gt bond-pairbond-pair
bond-pair bond-pair
three at 90o four at 90o one at
120othree at 90o three at 120o
Compare endash pair angles
VSEPR 5 electron pairsThe optimum structure maximizes the
angular separation of the lone pairs
I3- (AX2E3)
AX3E2
AX4E
AX2E3
5-electron-pair geometries
our previous example
Which of these is the more likely structure
Square Planar
VSEPR 6 electron pairs
See-saw
AX5E
AX4E2
6-electron-pair geometries
our previous example
Molecular Dipole Moments
1 Draw Lewis structures to determine 3D arrangement of atoms
Shortcut completely symmetric molecules will not have a dipole regardless of the polarity of the bonds
2 If one ldquosiderdquo of the molecule has more EN atoms than the other the molecule has a net dipole
We can use VSEPR to determine the polarity of a whole molecule
Molecular Dipoles
The C=O bonds have dipoles of equal magnitude but opposite direction so there is no net dipole moment
The O-H bonds have dipoles of equal magnitude that do not cancel each other so water has a net dipole moment
Molecular Dipoles (cont)
symmetric
symmetricasymmetri
c
F
ClF
Cl
Molecular Dipole Examplebull Write the Lewis dot and VESPR
structures for CF2Cl2 Does it have a dipole moment
C
F
FClCl
32 e-
Tetrahedral
Advanced VSEPR Application
Molecules with more than one central atomhellip methanol (CH3OH)
H CH
OH
H
tetrahedral e- pairstetrahedral shape
tetrahedral e- pairsbent shape
The VSEPR Table
2 AX2 BeF2 linear linear
3 AX3 BF3 trigonal planar
trigonal planar
AX2E O3 trigonal planar
bent
4 AX4 CH4 tetrahedral tetrahedral
AX3E NH3 tetrahedral pyramidal
AX2E2 H2O tetrahedral bent
e- pairs e- Geom Molec Geom
The VSEPR Table5 AX5 PF5 trigonal
bipyramidaltrigonalbipyramidal
AX4E SF4 trigonal bipyramidal
see saw
AX3E2 ClF3 trigonal bipyramidal
T-shaped
AX2E3 I3- trigonal bipyramidal
linear
6 AX6 SF6 octahedral octahedral
AX4E2 XeF4 octahedral square planar
e- pairs e- Geom Molec Geom
What is the expected shape of ICl2+
A linear
B bent
C tetrahedral
D square planar
AX2E2ICl Cl+
20 e-
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons occupies the region between the two nuclei
Rule 1 Maximum overlap The bond strength depends on the attraction of nuclei to the shared electrons soThe greater the orbital overlap the stronger the bond
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 2 Spins pair The two electrons in the overlap region occupy the same space and therefore must have opposite spins
There may be no more than 2 electrons in a molecular orbital
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
VSEPR 4 electron pairs (cont)
Comparison of CH4 (AX4) NH3 (AX3E) and H2O (AX2E2)
AX2E
AX3E
AX2E2
1 Refer to central atom as ldquoArdquo2 Attached atoms are referred to as ldquoXrdquo3 Lone pair are referred to as ldquoErdquo
HH
H
H
CNOF
Central Atom
Compound Electron-Pair Geometry
Molecular Shape
Carbon C CH4 tetrahedral tetrahedralNitrogen N NH3 tetrahedral trigonal pyramidalOxygen O H2O tetrahedral bentFluorine F HF tetrahedral linear
Molecular vs Electron-Pair Geometry
What is the electron-pair geometry and the molecular shape for HCFS
a) trigonal planar bent
b) trigonal planar trigonal planar
c) tetrahedral trigonal planar
d) tetrahedral tetrahedral
C
S
H
F
VSEPR Beyond the OctetSystems with expanded valence shells will have five or six electron pairs around a central atom
P
Cl
Cl
Cl
Cl
Cl
FF
F
F
FFS
90deg
120deg SF
F
F
F
F
F 90deg
90deg
VSEPR 5 electron pairsbull Consider the structure of SF4 (34 e- AX4E)bull What is the optimum arrangement of electron pairs
around S
SFF
FF
S
FF
FF
SF
FF
F
lone-pair bond-pair
two at 90o two at 120o
Repulsive forces (strongest to weakest)lone-pairlone-pair gt lone-pairbond-pair gt bond-pairbond-pair
bond-pair bond-pair
three at 90o four at 90o one at
120othree at 90o three at 120o
Compare endash pair angles
VSEPR 5 electron pairsThe optimum structure maximizes the
angular separation of the lone pairs
I3- (AX2E3)
AX3E2
AX4E
AX2E3
5-electron-pair geometries
our previous example
Which of these is the more likely structure
Square Planar
VSEPR 6 electron pairs
See-saw
AX5E
AX4E2
6-electron-pair geometries
our previous example
Molecular Dipole Moments
1 Draw Lewis structures to determine 3D arrangement of atoms
Shortcut completely symmetric molecules will not have a dipole regardless of the polarity of the bonds
2 If one ldquosiderdquo of the molecule has more EN atoms than the other the molecule has a net dipole
We can use VSEPR to determine the polarity of a whole molecule
Molecular Dipoles
The C=O bonds have dipoles of equal magnitude but opposite direction so there is no net dipole moment
The O-H bonds have dipoles of equal magnitude that do not cancel each other so water has a net dipole moment
Molecular Dipoles (cont)
symmetric
symmetricasymmetri
c
F
ClF
Cl
Molecular Dipole Examplebull Write the Lewis dot and VESPR
structures for CF2Cl2 Does it have a dipole moment
C
F
FClCl
32 e-
Tetrahedral
Advanced VSEPR Application
Molecules with more than one central atomhellip methanol (CH3OH)
H CH
OH
H
tetrahedral e- pairstetrahedral shape
tetrahedral e- pairsbent shape
The VSEPR Table
2 AX2 BeF2 linear linear
3 AX3 BF3 trigonal planar
trigonal planar
AX2E O3 trigonal planar
bent
4 AX4 CH4 tetrahedral tetrahedral
AX3E NH3 tetrahedral pyramidal
AX2E2 H2O tetrahedral bent
e- pairs e- Geom Molec Geom
The VSEPR Table5 AX5 PF5 trigonal
bipyramidaltrigonalbipyramidal
AX4E SF4 trigonal bipyramidal
see saw
AX3E2 ClF3 trigonal bipyramidal
T-shaped
AX2E3 I3- trigonal bipyramidal
linear
6 AX6 SF6 octahedral octahedral
AX4E2 XeF4 octahedral square planar
e- pairs e- Geom Molec Geom
What is the expected shape of ICl2+
A linear
B bent
C tetrahedral
D square planar
AX2E2ICl Cl+
20 e-
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons occupies the region between the two nuclei
Rule 1 Maximum overlap The bond strength depends on the attraction of nuclei to the shared electrons soThe greater the orbital overlap the stronger the bond
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 2 Spins pair The two electrons in the overlap region occupy the same space and therefore must have opposite spins
There may be no more than 2 electrons in a molecular orbital
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
AX2E
AX3E
AX2E2
1 Refer to central atom as ldquoArdquo2 Attached atoms are referred to as ldquoXrdquo3 Lone pair are referred to as ldquoErdquo
HH
H
H
CNOF
Central Atom
Compound Electron-Pair Geometry
Molecular Shape
Carbon C CH4 tetrahedral tetrahedralNitrogen N NH3 tetrahedral trigonal pyramidalOxygen O H2O tetrahedral bentFluorine F HF tetrahedral linear
Molecular vs Electron-Pair Geometry
What is the electron-pair geometry and the molecular shape for HCFS
a) trigonal planar bent
b) trigonal planar trigonal planar
c) tetrahedral trigonal planar
d) tetrahedral tetrahedral
C
S
H
F
VSEPR Beyond the OctetSystems with expanded valence shells will have five or six electron pairs around a central atom
P
Cl
Cl
Cl
Cl
Cl
FF
F
F
FFS
90deg
120deg SF
F
F
F
F
F 90deg
90deg
VSEPR 5 electron pairsbull Consider the structure of SF4 (34 e- AX4E)bull What is the optimum arrangement of electron pairs
around S
SFF
FF
S
FF
FF
SF
FF
F
lone-pair bond-pair
two at 90o two at 120o
Repulsive forces (strongest to weakest)lone-pairlone-pair gt lone-pairbond-pair gt bond-pairbond-pair
bond-pair bond-pair
three at 90o four at 90o one at
120othree at 90o three at 120o
Compare endash pair angles
VSEPR 5 electron pairsThe optimum structure maximizes the
angular separation of the lone pairs
I3- (AX2E3)
AX3E2
AX4E
AX2E3
5-electron-pair geometries
our previous example
Which of these is the more likely structure
Square Planar
VSEPR 6 electron pairs
See-saw
AX5E
AX4E2
6-electron-pair geometries
our previous example
Molecular Dipole Moments
1 Draw Lewis structures to determine 3D arrangement of atoms
Shortcut completely symmetric molecules will not have a dipole regardless of the polarity of the bonds
2 If one ldquosiderdquo of the molecule has more EN atoms than the other the molecule has a net dipole
We can use VSEPR to determine the polarity of a whole molecule
Molecular Dipoles
The C=O bonds have dipoles of equal magnitude but opposite direction so there is no net dipole moment
The O-H bonds have dipoles of equal magnitude that do not cancel each other so water has a net dipole moment
Molecular Dipoles (cont)
symmetric
symmetricasymmetri
c
F
ClF
Cl
Molecular Dipole Examplebull Write the Lewis dot and VESPR
structures for CF2Cl2 Does it have a dipole moment
C
F
FClCl
32 e-
Tetrahedral
Advanced VSEPR Application
Molecules with more than one central atomhellip methanol (CH3OH)
H CH
OH
H
tetrahedral e- pairstetrahedral shape
tetrahedral e- pairsbent shape
The VSEPR Table
2 AX2 BeF2 linear linear
3 AX3 BF3 trigonal planar
trigonal planar
AX2E O3 trigonal planar
bent
4 AX4 CH4 tetrahedral tetrahedral
AX3E NH3 tetrahedral pyramidal
AX2E2 H2O tetrahedral bent
e- pairs e- Geom Molec Geom
The VSEPR Table5 AX5 PF5 trigonal
bipyramidaltrigonalbipyramidal
AX4E SF4 trigonal bipyramidal
see saw
AX3E2 ClF3 trigonal bipyramidal
T-shaped
AX2E3 I3- trigonal bipyramidal
linear
6 AX6 SF6 octahedral octahedral
AX4E2 XeF4 octahedral square planar
e- pairs e- Geom Molec Geom
What is the expected shape of ICl2+
A linear
B bent
C tetrahedral
D square planar
AX2E2ICl Cl+
20 e-
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons occupies the region between the two nuclei
Rule 1 Maximum overlap The bond strength depends on the attraction of nuclei to the shared electrons soThe greater the orbital overlap the stronger the bond
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 2 Spins pair The two electrons in the overlap region occupy the same space and therefore must have opposite spins
There may be no more than 2 electrons in a molecular orbital
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
HH
H
H
CNOF
Central Atom
Compound Electron-Pair Geometry
Molecular Shape
Carbon C CH4 tetrahedral tetrahedralNitrogen N NH3 tetrahedral trigonal pyramidalOxygen O H2O tetrahedral bentFluorine F HF tetrahedral linear
Molecular vs Electron-Pair Geometry
What is the electron-pair geometry and the molecular shape for HCFS
a) trigonal planar bent
b) trigonal planar trigonal planar
c) tetrahedral trigonal planar
d) tetrahedral tetrahedral
C
S
H
F
VSEPR Beyond the OctetSystems with expanded valence shells will have five or six electron pairs around a central atom
P
Cl
Cl
Cl
Cl
Cl
FF
F
F
FFS
90deg
120deg SF
F
F
F
F
F 90deg
90deg
VSEPR 5 electron pairsbull Consider the structure of SF4 (34 e- AX4E)bull What is the optimum arrangement of electron pairs
around S
SFF
FF
S
FF
FF
SF
FF
F
lone-pair bond-pair
two at 90o two at 120o
Repulsive forces (strongest to weakest)lone-pairlone-pair gt lone-pairbond-pair gt bond-pairbond-pair
bond-pair bond-pair
three at 90o four at 90o one at
120othree at 90o three at 120o
Compare endash pair angles
VSEPR 5 electron pairsThe optimum structure maximizes the
angular separation of the lone pairs
I3- (AX2E3)
AX3E2
AX4E
AX2E3
5-electron-pair geometries
our previous example
Which of these is the more likely structure
Square Planar
VSEPR 6 electron pairs
See-saw
AX5E
AX4E2
6-electron-pair geometries
our previous example
Molecular Dipole Moments
1 Draw Lewis structures to determine 3D arrangement of atoms
Shortcut completely symmetric molecules will not have a dipole regardless of the polarity of the bonds
2 If one ldquosiderdquo of the molecule has more EN atoms than the other the molecule has a net dipole
We can use VSEPR to determine the polarity of a whole molecule
Molecular Dipoles
The C=O bonds have dipoles of equal magnitude but opposite direction so there is no net dipole moment
The O-H bonds have dipoles of equal magnitude that do not cancel each other so water has a net dipole moment
Molecular Dipoles (cont)
symmetric
symmetricasymmetri
c
F
ClF
Cl
Molecular Dipole Examplebull Write the Lewis dot and VESPR
structures for CF2Cl2 Does it have a dipole moment
C
F
FClCl
32 e-
Tetrahedral
Advanced VSEPR Application
Molecules with more than one central atomhellip methanol (CH3OH)
H CH
OH
H
tetrahedral e- pairstetrahedral shape
tetrahedral e- pairsbent shape
The VSEPR Table
2 AX2 BeF2 linear linear
3 AX3 BF3 trigonal planar
trigonal planar
AX2E O3 trigonal planar
bent
4 AX4 CH4 tetrahedral tetrahedral
AX3E NH3 tetrahedral pyramidal
AX2E2 H2O tetrahedral bent
e- pairs e- Geom Molec Geom
The VSEPR Table5 AX5 PF5 trigonal
bipyramidaltrigonalbipyramidal
AX4E SF4 trigonal bipyramidal
see saw
AX3E2 ClF3 trigonal bipyramidal
T-shaped
AX2E3 I3- trigonal bipyramidal
linear
6 AX6 SF6 octahedral octahedral
AX4E2 XeF4 octahedral square planar
e- pairs e- Geom Molec Geom
What is the expected shape of ICl2+
A linear
B bent
C tetrahedral
D square planar
AX2E2ICl Cl+
20 e-
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons occupies the region between the two nuclei
Rule 1 Maximum overlap The bond strength depends on the attraction of nuclei to the shared electrons soThe greater the orbital overlap the stronger the bond
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 2 Spins pair The two electrons in the overlap region occupy the same space and therefore must have opposite spins
There may be no more than 2 electrons in a molecular orbital
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
What is the electron-pair geometry and the molecular shape for HCFS
a) trigonal planar bent
b) trigonal planar trigonal planar
c) tetrahedral trigonal planar
d) tetrahedral tetrahedral
C
S
H
F
VSEPR Beyond the OctetSystems with expanded valence shells will have five or six electron pairs around a central atom
P
Cl
Cl
Cl
Cl
Cl
FF
F
F
FFS
90deg
120deg SF
F
F
F
F
F 90deg
90deg
VSEPR 5 electron pairsbull Consider the structure of SF4 (34 e- AX4E)bull What is the optimum arrangement of electron pairs
around S
SFF
FF
S
FF
FF
SF
FF
F
lone-pair bond-pair
two at 90o two at 120o
Repulsive forces (strongest to weakest)lone-pairlone-pair gt lone-pairbond-pair gt bond-pairbond-pair
bond-pair bond-pair
three at 90o four at 90o one at
120othree at 90o three at 120o
Compare endash pair angles
VSEPR 5 electron pairsThe optimum structure maximizes the
angular separation of the lone pairs
I3- (AX2E3)
AX3E2
AX4E
AX2E3
5-electron-pair geometries
our previous example
Which of these is the more likely structure
Square Planar
VSEPR 6 electron pairs
See-saw
AX5E
AX4E2
6-electron-pair geometries
our previous example
Molecular Dipole Moments
1 Draw Lewis structures to determine 3D arrangement of atoms
Shortcut completely symmetric molecules will not have a dipole regardless of the polarity of the bonds
2 If one ldquosiderdquo of the molecule has more EN atoms than the other the molecule has a net dipole
We can use VSEPR to determine the polarity of a whole molecule
Molecular Dipoles
The C=O bonds have dipoles of equal magnitude but opposite direction so there is no net dipole moment
The O-H bonds have dipoles of equal magnitude that do not cancel each other so water has a net dipole moment
Molecular Dipoles (cont)
symmetric
symmetricasymmetri
c
F
ClF
Cl
Molecular Dipole Examplebull Write the Lewis dot and VESPR
structures for CF2Cl2 Does it have a dipole moment
C
F
FClCl
32 e-
Tetrahedral
Advanced VSEPR Application
Molecules with more than one central atomhellip methanol (CH3OH)
H CH
OH
H
tetrahedral e- pairstetrahedral shape
tetrahedral e- pairsbent shape
The VSEPR Table
2 AX2 BeF2 linear linear
3 AX3 BF3 trigonal planar
trigonal planar
AX2E O3 trigonal planar
bent
4 AX4 CH4 tetrahedral tetrahedral
AX3E NH3 tetrahedral pyramidal
AX2E2 H2O tetrahedral bent
e- pairs e- Geom Molec Geom
The VSEPR Table5 AX5 PF5 trigonal
bipyramidaltrigonalbipyramidal
AX4E SF4 trigonal bipyramidal
see saw
AX3E2 ClF3 trigonal bipyramidal
T-shaped
AX2E3 I3- trigonal bipyramidal
linear
6 AX6 SF6 octahedral octahedral
AX4E2 XeF4 octahedral square planar
e- pairs e- Geom Molec Geom
What is the expected shape of ICl2+
A linear
B bent
C tetrahedral
D square planar
AX2E2ICl Cl+
20 e-
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons occupies the region between the two nuclei
Rule 1 Maximum overlap The bond strength depends on the attraction of nuclei to the shared electrons soThe greater the orbital overlap the stronger the bond
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 2 Spins pair The two electrons in the overlap region occupy the same space and therefore must have opposite spins
There may be no more than 2 electrons in a molecular orbital
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
VSEPR Beyond the OctetSystems with expanded valence shells will have five or six electron pairs around a central atom
P
Cl
Cl
Cl
Cl
Cl
FF
F
F
FFS
90deg
120deg SF
F
F
F
F
F 90deg
90deg
VSEPR 5 electron pairsbull Consider the structure of SF4 (34 e- AX4E)bull What is the optimum arrangement of electron pairs
around S
SFF
FF
S
FF
FF
SF
FF
F
lone-pair bond-pair
two at 90o two at 120o
Repulsive forces (strongest to weakest)lone-pairlone-pair gt lone-pairbond-pair gt bond-pairbond-pair
bond-pair bond-pair
three at 90o four at 90o one at
120othree at 90o three at 120o
Compare endash pair angles
VSEPR 5 electron pairsThe optimum structure maximizes the
angular separation of the lone pairs
I3- (AX2E3)
AX3E2
AX4E
AX2E3
5-electron-pair geometries
our previous example
Which of these is the more likely structure
Square Planar
VSEPR 6 electron pairs
See-saw
AX5E
AX4E2
6-electron-pair geometries
our previous example
Molecular Dipole Moments
1 Draw Lewis structures to determine 3D arrangement of atoms
Shortcut completely symmetric molecules will not have a dipole regardless of the polarity of the bonds
2 If one ldquosiderdquo of the molecule has more EN atoms than the other the molecule has a net dipole
We can use VSEPR to determine the polarity of a whole molecule
Molecular Dipoles
The C=O bonds have dipoles of equal magnitude but opposite direction so there is no net dipole moment
The O-H bonds have dipoles of equal magnitude that do not cancel each other so water has a net dipole moment
Molecular Dipoles (cont)
symmetric
symmetricasymmetri
c
F
ClF
Cl
Molecular Dipole Examplebull Write the Lewis dot and VESPR
structures for CF2Cl2 Does it have a dipole moment
C
F
FClCl
32 e-
Tetrahedral
Advanced VSEPR Application
Molecules with more than one central atomhellip methanol (CH3OH)
H CH
OH
H
tetrahedral e- pairstetrahedral shape
tetrahedral e- pairsbent shape
The VSEPR Table
2 AX2 BeF2 linear linear
3 AX3 BF3 trigonal planar
trigonal planar
AX2E O3 trigonal planar
bent
4 AX4 CH4 tetrahedral tetrahedral
AX3E NH3 tetrahedral pyramidal
AX2E2 H2O tetrahedral bent
e- pairs e- Geom Molec Geom
The VSEPR Table5 AX5 PF5 trigonal
bipyramidaltrigonalbipyramidal
AX4E SF4 trigonal bipyramidal
see saw
AX3E2 ClF3 trigonal bipyramidal
T-shaped
AX2E3 I3- trigonal bipyramidal
linear
6 AX6 SF6 octahedral octahedral
AX4E2 XeF4 octahedral square planar
e- pairs e- Geom Molec Geom
What is the expected shape of ICl2+
A linear
B bent
C tetrahedral
D square planar
AX2E2ICl Cl+
20 e-
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons occupies the region between the two nuclei
Rule 1 Maximum overlap The bond strength depends on the attraction of nuclei to the shared electrons soThe greater the orbital overlap the stronger the bond
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 2 Spins pair The two electrons in the overlap region occupy the same space and therefore must have opposite spins
There may be no more than 2 electrons in a molecular orbital
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
VSEPR 5 electron pairsbull Consider the structure of SF4 (34 e- AX4E)bull What is the optimum arrangement of electron pairs
around S
SFF
FF
S
FF
FF
SF
FF
F
lone-pair bond-pair
two at 90o two at 120o
Repulsive forces (strongest to weakest)lone-pairlone-pair gt lone-pairbond-pair gt bond-pairbond-pair
bond-pair bond-pair
three at 90o four at 90o one at
120othree at 90o three at 120o
Compare endash pair angles
VSEPR 5 electron pairsThe optimum structure maximizes the
angular separation of the lone pairs
I3- (AX2E3)
AX3E2
AX4E
AX2E3
5-electron-pair geometries
our previous example
Which of these is the more likely structure
Square Planar
VSEPR 6 electron pairs
See-saw
AX5E
AX4E2
6-electron-pair geometries
our previous example
Molecular Dipole Moments
1 Draw Lewis structures to determine 3D arrangement of atoms
Shortcut completely symmetric molecules will not have a dipole regardless of the polarity of the bonds
2 If one ldquosiderdquo of the molecule has more EN atoms than the other the molecule has a net dipole
We can use VSEPR to determine the polarity of a whole molecule
Molecular Dipoles
The C=O bonds have dipoles of equal magnitude but opposite direction so there is no net dipole moment
The O-H bonds have dipoles of equal magnitude that do not cancel each other so water has a net dipole moment
Molecular Dipoles (cont)
symmetric
symmetricasymmetri
c
F
ClF
Cl
Molecular Dipole Examplebull Write the Lewis dot and VESPR
structures for CF2Cl2 Does it have a dipole moment
C
F
FClCl
32 e-
Tetrahedral
Advanced VSEPR Application
Molecules with more than one central atomhellip methanol (CH3OH)
H CH
OH
H
tetrahedral e- pairstetrahedral shape
tetrahedral e- pairsbent shape
The VSEPR Table
2 AX2 BeF2 linear linear
3 AX3 BF3 trigonal planar
trigonal planar
AX2E O3 trigonal planar
bent
4 AX4 CH4 tetrahedral tetrahedral
AX3E NH3 tetrahedral pyramidal
AX2E2 H2O tetrahedral bent
e- pairs e- Geom Molec Geom
The VSEPR Table5 AX5 PF5 trigonal
bipyramidaltrigonalbipyramidal
AX4E SF4 trigonal bipyramidal
see saw
AX3E2 ClF3 trigonal bipyramidal
T-shaped
AX2E3 I3- trigonal bipyramidal
linear
6 AX6 SF6 octahedral octahedral
AX4E2 XeF4 octahedral square planar
e- pairs e- Geom Molec Geom
What is the expected shape of ICl2+
A linear
B bent
C tetrahedral
D square planar
AX2E2ICl Cl+
20 e-
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons occupies the region between the two nuclei
Rule 1 Maximum overlap The bond strength depends on the attraction of nuclei to the shared electrons soThe greater the orbital overlap the stronger the bond
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 2 Spins pair The two electrons in the overlap region occupy the same space and therefore must have opposite spins
There may be no more than 2 electrons in a molecular orbital
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
VSEPR 5 electron pairsThe optimum structure maximizes the
angular separation of the lone pairs
I3- (AX2E3)
AX3E2
AX4E
AX2E3
5-electron-pair geometries
our previous example
Which of these is the more likely structure
Square Planar
VSEPR 6 electron pairs
See-saw
AX5E
AX4E2
6-electron-pair geometries
our previous example
Molecular Dipole Moments
1 Draw Lewis structures to determine 3D arrangement of atoms
Shortcut completely symmetric molecules will not have a dipole regardless of the polarity of the bonds
2 If one ldquosiderdquo of the molecule has more EN atoms than the other the molecule has a net dipole
We can use VSEPR to determine the polarity of a whole molecule
Molecular Dipoles
The C=O bonds have dipoles of equal magnitude but opposite direction so there is no net dipole moment
The O-H bonds have dipoles of equal magnitude that do not cancel each other so water has a net dipole moment
Molecular Dipoles (cont)
symmetric
symmetricasymmetri
c
F
ClF
Cl
Molecular Dipole Examplebull Write the Lewis dot and VESPR
structures for CF2Cl2 Does it have a dipole moment
C
F
FClCl
32 e-
Tetrahedral
Advanced VSEPR Application
Molecules with more than one central atomhellip methanol (CH3OH)
H CH
OH
H
tetrahedral e- pairstetrahedral shape
tetrahedral e- pairsbent shape
The VSEPR Table
2 AX2 BeF2 linear linear
3 AX3 BF3 trigonal planar
trigonal planar
AX2E O3 trigonal planar
bent
4 AX4 CH4 tetrahedral tetrahedral
AX3E NH3 tetrahedral pyramidal
AX2E2 H2O tetrahedral bent
e- pairs e- Geom Molec Geom
The VSEPR Table5 AX5 PF5 trigonal
bipyramidaltrigonalbipyramidal
AX4E SF4 trigonal bipyramidal
see saw
AX3E2 ClF3 trigonal bipyramidal
T-shaped
AX2E3 I3- trigonal bipyramidal
linear
6 AX6 SF6 octahedral octahedral
AX4E2 XeF4 octahedral square planar
e- pairs e- Geom Molec Geom
What is the expected shape of ICl2+
A linear
B bent
C tetrahedral
D square planar
AX2E2ICl Cl+
20 e-
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons occupies the region between the two nuclei
Rule 1 Maximum overlap The bond strength depends on the attraction of nuclei to the shared electrons soThe greater the orbital overlap the stronger the bond
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 2 Spins pair The two electrons in the overlap region occupy the same space and therefore must have opposite spins
There may be no more than 2 electrons in a molecular orbital
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
AX3E2
AX4E
AX2E3
5-electron-pair geometries
our previous example
Which of these is the more likely structure
Square Planar
VSEPR 6 electron pairs
See-saw
AX5E
AX4E2
6-electron-pair geometries
our previous example
Molecular Dipole Moments
1 Draw Lewis structures to determine 3D arrangement of atoms
Shortcut completely symmetric molecules will not have a dipole regardless of the polarity of the bonds
2 If one ldquosiderdquo of the molecule has more EN atoms than the other the molecule has a net dipole
We can use VSEPR to determine the polarity of a whole molecule
Molecular Dipoles
The C=O bonds have dipoles of equal magnitude but opposite direction so there is no net dipole moment
The O-H bonds have dipoles of equal magnitude that do not cancel each other so water has a net dipole moment
Molecular Dipoles (cont)
symmetric
symmetricasymmetri
c
F
ClF
Cl
Molecular Dipole Examplebull Write the Lewis dot and VESPR
structures for CF2Cl2 Does it have a dipole moment
C
F
FClCl
32 e-
Tetrahedral
Advanced VSEPR Application
Molecules with more than one central atomhellip methanol (CH3OH)
H CH
OH
H
tetrahedral e- pairstetrahedral shape
tetrahedral e- pairsbent shape
The VSEPR Table
2 AX2 BeF2 linear linear
3 AX3 BF3 trigonal planar
trigonal planar
AX2E O3 trigonal planar
bent
4 AX4 CH4 tetrahedral tetrahedral
AX3E NH3 tetrahedral pyramidal
AX2E2 H2O tetrahedral bent
e- pairs e- Geom Molec Geom
The VSEPR Table5 AX5 PF5 trigonal
bipyramidaltrigonalbipyramidal
AX4E SF4 trigonal bipyramidal
see saw
AX3E2 ClF3 trigonal bipyramidal
T-shaped
AX2E3 I3- trigonal bipyramidal
linear
6 AX6 SF6 octahedral octahedral
AX4E2 XeF4 octahedral square planar
e- pairs e- Geom Molec Geom
What is the expected shape of ICl2+
A linear
B bent
C tetrahedral
D square planar
AX2E2ICl Cl+
20 e-
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons occupies the region between the two nuclei
Rule 1 Maximum overlap The bond strength depends on the attraction of nuclei to the shared electrons soThe greater the orbital overlap the stronger the bond
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 2 Spins pair The two electrons in the overlap region occupy the same space and therefore must have opposite spins
There may be no more than 2 electrons in a molecular orbital
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Which of these is the more likely structure
Square Planar
VSEPR 6 electron pairs
See-saw
AX5E
AX4E2
6-electron-pair geometries
our previous example
Molecular Dipole Moments
1 Draw Lewis structures to determine 3D arrangement of atoms
Shortcut completely symmetric molecules will not have a dipole regardless of the polarity of the bonds
2 If one ldquosiderdquo of the molecule has more EN atoms than the other the molecule has a net dipole
We can use VSEPR to determine the polarity of a whole molecule
Molecular Dipoles
The C=O bonds have dipoles of equal magnitude but opposite direction so there is no net dipole moment
The O-H bonds have dipoles of equal magnitude that do not cancel each other so water has a net dipole moment
Molecular Dipoles (cont)
symmetric
symmetricasymmetri
c
F
ClF
Cl
Molecular Dipole Examplebull Write the Lewis dot and VESPR
structures for CF2Cl2 Does it have a dipole moment
C
F
FClCl
32 e-
Tetrahedral
Advanced VSEPR Application
Molecules with more than one central atomhellip methanol (CH3OH)
H CH
OH
H
tetrahedral e- pairstetrahedral shape
tetrahedral e- pairsbent shape
The VSEPR Table
2 AX2 BeF2 linear linear
3 AX3 BF3 trigonal planar
trigonal planar
AX2E O3 trigonal planar
bent
4 AX4 CH4 tetrahedral tetrahedral
AX3E NH3 tetrahedral pyramidal
AX2E2 H2O tetrahedral bent
e- pairs e- Geom Molec Geom
The VSEPR Table5 AX5 PF5 trigonal
bipyramidaltrigonalbipyramidal
AX4E SF4 trigonal bipyramidal
see saw
AX3E2 ClF3 trigonal bipyramidal
T-shaped
AX2E3 I3- trigonal bipyramidal
linear
6 AX6 SF6 octahedral octahedral
AX4E2 XeF4 octahedral square planar
e- pairs e- Geom Molec Geom
What is the expected shape of ICl2+
A linear
B bent
C tetrahedral
D square planar
AX2E2ICl Cl+
20 e-
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons occupies the region between the two nuclei
Rule 1 Maximum overlap The bond strength depends on the attraction of nuclei to the shared electrons soThe greater the orbital overlap the stronger the bond
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 2 Spins pair The two electrons in the overlap region occupy the same space and therefore must have opposite spins
There may be no more than 2 electrons in a molecular orbital
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
AX5E
AX4E2
6-electron-pair geometries
our previous example
Molecular Dipole Moments
1 Draw Lewis structures to determine 3D arrangement of atoms
Shortcut completely symmetric molecules will not have a dipole regardless of the polarity of the bonds
2 If one ldquosiderdquo of the molecule has more EN atoms than the other the molecule has a net dipole
We can use VSEPR to determine the polarity of a whole molecule
Molecular Dipoles
The C=O bonds have dipoles of equal magnitude but opposite direction so there is no net dipole moment
The O-H bonds have dipoles of equal magnitude that do not cancel each other so water has a net dipole moment
Molecular Dipoles (cont)
symmetric
symmetricasymmetri
c
F
ClF
Cl
Molecular Dipole Examplebull Write the Lewis dot and VESPR
structures for CF2Cl2 Does it have a dipole moment
C
F
FClCl
32 e-
Tetrahedral
Advanced VSEPR Application
Molecules with more than one central atomhellip methanol (CH3OH)
H CH
OH
H
tetrahedral e- pairstetrahedral shape
tetrahedral e- pairsbent shape
The VSEPR Table
2 AX2 BeF2 linear linear
3 AX3 BF3 trigonal planar
trigonal planar
AX2E O3 trigonal planar
bent
4 AX4 CH4 tetrahedral tetrahedral
AX3E NH3 tetrahedral pyramidal
AX2E2 H2O tetrahedral bent
e- pairs e- Geom Molec Geom
The VSEPR Table5 AX5 PF5 trigonal
bipyramidaltrigonalbipyramidal
AX4E SF4 trigonal bipyramidal
see saw
AX3E2 ClF3 trigonal bipyramidal
T-shaped
AX2E3 I3- trigonal bipyramidal
linear
6 AX6 SF6 octahedral octahedral
AX4E2 XeF4 octahedral square planar
e- pairs e- Geom Molec Geom
What is the expected shape of ICl2+
A linear
B bent
C tetrahedral
D square planar
AX2E2ICl Cl+
20 e-
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons occupies the region between the two nuclei
Rule 1 Maximum overlap The bond strength depends on the attraction of nuclei to the shared electrons soThe greater the orbital overlap the stronger the bond
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 2 Spins pair The two electrons in the overlap region occupy the same space and therefore must have opposite spins
There may be no more than 2 electrons in a molecular orbital
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Molecular Dipole Moments
1 Draw Lewis structures to determine 3D arrangement of atoms
Shortcut completely symmetric molecules will not have a dipole regardless of the polarity of the bonds
2 If one ldquosiderdquo of the molecule has more EN atoms than the other the molecule has a net dipole
We can use VSEPR to determine the polarity of a whole molecule
Molecular Dipoles
The C=O bonds have dipoles of equal magnitude but opposite direction so there is no net dipole moment
The O-H bonds have dipoles of equal magnitude that do not cancel each other so water has a net dipole moment
Molecular Dipoles (cont)
symmetric
symmetricasymmetri
c
F
ClF
Cl
Molecular Dipole Examplebull Write the Lewis dot and VESPR
structures for CF2Cl2 Does it have a dipole moment
C
F
FClCl
32 e-
Tetrahedral
Advanced VSEPR Application
Molecules with more than one central atomhellip methanol (CH3OH)
H CH
OH
H
tetrahedral e- pairstetrahedral shape
tetrahedral e- pairsbent shape
The VSEPR Table
2 AX2 BeF2 linear linear
3 AX3 BF3 trigonal planar
trigonal planar
AX2E O3 trigonal planar
bent
4 AX4 CH4 tetrahedral tetrahedral
AX3E NH3 tetrahedral pyramidal
AX2E2 H2O tetrahedral bent
e- pairs e- Geom Molec Geom
The VSEPR Table5 AX5 PF5 trigonal
bipyramidaltrigonalbipyramidal
AX4E SF4 trigonal bipyramidal
see saw
AX3E2 ClF3 trigonal bipyramidal
T-shaped
AX2E3 I3- trigonal bipyramidal
linear
6 AX6 SF6 octahedral octahedral
AX4E2 XeF4 octahedral square planar
e- pairs e- Geom Molec Geom
What is the expected shape of ICl2+
A linear
B bent
C tetrahedral
D square planar
AX2E2ICl Cl+
20 e-
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons occupies the region between the two nuclei
Rule 1 Maximum overlap The bond strength depends on the attraction of nuclei to the shared electrons soThe greater the orbital overlap the stronger the bond
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 2 Spins pair The two electrons in the overlap region occupy the same space and therefore must have opposite spins
There may be no more than 2 electrons in a molecular orbital
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Molecular Dipoles
The C=O bonds have dipoles of equal magnitude but opposite direction so there is no net dipole moment
The O-H bonds have dipoles of equal magnitude that do not cancel each other so water has a net dipole moment
Molecular Dipoles (cont)
symmetric
symmetricasymmetri
c
F
ClF
Cl
Molecular Dipole Examplebull Write the Lewis dot and VESPR
structures for CF2Cl2 Does it have a dipole moment
C
F
FClCl
32 e-
Tetrahedral
Advanced VSEPR Application
Molecules with more than one central atomhellip methanol (CH3OH)
H CH
OH
H
tetrahedral e- pairstetrahedral shape
tetrahedral e- pairsbent shape
The VSEPR Table
2 AX2 BeF2 linear linear
3 AX3 BF3 trigonal planar
trigonal planar
AX2E O3 trigonal planar
bent
4 AX4 CH4 tetrahedral tetrahedral
AX3E NH3 tetrahedral pyramidal
AX2E2 H2O tetrahedral bent
e- pairs e- Geom Molec Geom
The VSEPR Table5 AX5 PF5 trigonal
bipyramidaltrigonalbipyramidal
AX4E SF4 trigonal bipyramidal
see saw
AX3E2 ClF3 trigonal bipyramidal
T-shaped
AX2E3 I3- trigonal bipyramidal
linear
6 AX6 SF6 octahedral octahedral
AX4E2 XeF4 octahedral square planar
e- pairs e- Geom Molec Geom
What is the expected shape of ICl2+
A linear
B bent
C tetrahedral
D square planar
AX2E2ICl Cl+
20 e-
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons occupies the region between the two nuclei
Rule 1 Maximum overlap The bond strength depends on the attraction of nuclei to the shared electrons soThe greater the orbital overlap the stronger the bond
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 2 Spins pair The two electrons in the overlap region occupy the same space and therefore must have opposite spins
There may be no more than 2 electrons in a molecular orbital
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Molecular Dipoles (cont)
symmetric
symmetricasymmetri
c
F
ClF
Cl
Molecular Dipole Examplebull Write the Lewis dot and VESPR
structures for CF2Cl2 Does it have a dipole moment
C
F
FClCl
32 e-
Tetrahedral
Advanced VSEPR Application
Molecules with more than one central atomhellip methanol (CH3OH)
H CH
OH
H
tetrahedral e- pairstetrahedral shape
tetrahedral e- pairsbent shape
The VSEPR Table
2 AX2 BeF2 linear linear
3 AX3 BF3 trigonal planar
trigonal planar
AX2E O3 trigonal planar
bent
4 AX4 CH4 tetrahedral tetrahedral
AX3E NH3 tetrahedral pyramidal
AX2E2 H2O tetrahedral bent
e- pairs e- Geom Molec Geom
The VSEPR Table5 AX5 PF5 trigonal
bipyramidaltrigonalbipyramidal
AX4E SF4 trigonal bipyramidal
see saw
AX3E2 ClF3 trigonal bipyramidal
T-shaped
AX2E3 I3- trigonal bipyramidal
linear
6 AX6 SF6 octahedral octahedral
AX4E2 XeF4 octahedral square planar
e- pairs e- Geom Molec Geom
What is the expected shape of ICl2+
A linear
B bent
C tetrahedral
D square planar
AX2E2ICl Cl+
20 e-
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons occupies the region between the two nuclei
Rule 1 Maximum overlap The bond strength depends on the attraction of nuclei to the shared electrons soThe greater the orbital overlap the stronger the bond
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 2 Spins pair The two electrons in the overlap region occupy the same space and therefore must have opposite spins
There may be no more than 2 electrons in a molecular orbital
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
F
ClF
Cl
Molecular Dipole Examplebull Write the Lewis dot and VESPR
structures for CF2Cl2 Does it have a dipole moment
C
F
FClCl
32 e-
Tetrahedral
Advanced VSEPR Application
Molecules with more than one central atomhellip methanol (CH3OH)
H CH
OH
H
tetrahedral e- pairstetrahedral shape
tetrahedral e- pairsbent shape
The VSEPR Table
2 AX2 BeF2 linear linear
3 AX3 BF3 trigonal planar
trigonal planar
AX2E O3 trigonal planar
bent
4 AX4 CH4 tetrahedral tetrahedral
AX3E NH3 tetrahedral pyramidal
AX2E2 H2O tetrahedral bent
e- pairs e- Geom Molec Geom
The VSEPR Table5 AX5 PF5 trigonal
bipyramidaltrigonalbipyramidal
AX4E SF4 trigonal bipyramidal
see saw
AX3E2 ClF3 trigonal bipyramidal
T-shaped
AX2E3 I3- trigonal bipyramidal
linear
6 AX6 SF6 octahedral octahedral
AX4E2 XeF4 octahedral square planar
e- pairs e- Geom Molec Geom
What is the expected shape of ICl2+
A linear
B bent
C tetrahedral
D square planar
AX2E2ICl Cl+
20 e-
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons occupies the region between the two nuclei
Rule 1 Maximum overlap The bond strength depends on the attraction of nuclei to the shared electrons soThe greater the orbital overlap the stronger the bond
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 2 Spins pair The two electrons in the overlap region occupy the same space and therefore must have opposite spins
There may be no more than 2 electrons in a molecular orbital
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Advanced VSEPR Application
Molecules with more than one central atomhellip methanol (CH3OH)
H CH
OH
H
tetrahedral e- pairstetrahedral shape
tetrahedral e- pairsbent shape
The VSEPR Table
2 AX2 BeF2 linear linear
3 AX3 BF3 trigonal planar
trigonal planar
AX2E O3 trigonal planar
bent
4 AX4 CH4 tetrahedral tetrahedral
AX3E NH3 tetrahedral pyramidal
AX2E2 H2O tetrahedral bent
e- pairs e- Geom Molec Geom
The VSEPR Table5 AX5 PF5 trigonal
bipyramidaltrigonalbipyramidal
AX4E SF4 trigonal bipyramidal
see saw
AX3E2 ClF3 trigonal bipyramidal
T-shaped
AX2E3 I3- trigonal bipyramidal
linear
6 AX6 SF6 octahedral octahedral
AX4E2 XeF4 octahedral square planar
e- pairs e- Geom Molec Geom
What is the expected shape of ICl2+
A linear
B bent
C tetrahedral
D square planar
AX2E2ICl Cl+
20 e-
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons occupies the region between the two nuclei
Rule 1 Maximum overlap The bond strength depends on the attraction of nuclei to the shared electrons soThe greater the orbital overlap the stronger the bond
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 2 Spins pair The two electrons in the overlap region occupy the same space and therefore must have opposite spins
There may be no more than 2 electrons in a molecular orbital
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
The VSEPR Table
2 AX2 BeF2 linear linear
3 AX3 BF3 trigonal planar
trigonal planar
AX2E O3 trigonal planar
bent
4 AX4 CH4 tetrahedral tetrahedral
AX3E NH3 tetrahedral pyramidal
AX2E2 H2O tetrahedral bent
e- pairs e- Geom Molec Geom
The VSEPR Table5 AX5 PF5 trigonal
bipyramidaltrigonalbipyramidal
AX4E SF4 trigonal bipyramidal
see saw
AX3E2 ClF3 trigonal bipyramidal
T-shaped
AX2E3 I3- trigonal bipyramidal
linear
6 AX6 SF6 octahedral octahedral
AX4E2 XeF4 octahedral square planar
e- pairs e- Geom Molec Geom
What is the expected shape of ICl2+
A linear
B bent
C tetrahedral
D square planar
AX2E2ICl Cl+
20 e-
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons occupies the region between the two nuclei
Rule 1 Maximum overlap The bond strength depends on the attraction of nuclei to the shared electrons soThe greater the orbital overlap the stronger the bond
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 2 Spins pair The two electrons in the overlap region occupy the same space and therefore must have opposite spins
There may be no more than 2 electrons in a molecular orbital
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
The VSEPR Table5 AX5 PF5 trigonal
bipyramidaltrigonalbipyramidal
AX4E SF4 trigonal bipyramidal
see saw
AX3E2 ClF3 trigonal bipyramidal
T-shaped
AX2E3 I3- trigonal bipyramidal
linear
6 AX6 SF6 octahedral octahedral
AX4E2 XeF4 octahedral square planar
e- pairs e- Geom Molec Geom
What is the expected shape of ICl2+
A linear
B bent
C tetrahedral
D square planar
AX2E2ICl Cl+
20 e-
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons occupies the region between the two nuclei
Rule 1 Maximum overlap The bond strength depends on the attraction of nuclei to the shared electrons soThe greater the orbital overlap the stronger the bond
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 2 Spins pair The two electrons in the overlap region occupy the same space and therefore must have opposite spins
There may be no more than 2 electrons in a molecular orbital
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
What is the expected shape of ICl2+
A linear
B bent
C tetrahedral
D square planar
AX2E2ICl Cl+
20 e-
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons occupies the region between the two nuclei
Rule 1 Maximum overlap The bond strength depends on the attraction of nuclei to the shared electrons soThe greater the orbital overlap the stronger the bond
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 2 Spins pair The two electrons in the overlap region occupy the same space and therefore must have opposite spins
There may be no more than 2 electrons in a molecular orbital
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from
two atoms overlap and a pair of electrons occupies the region between the two nuclei
Rule 1 Maximum overlap The bond strength depends on the attraction of nuclei to the shared electrons soThe greater the orbital overlap the stronger the bond
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 2 Spins pair The two electrons in the overlap region occupy the same space and therefore must have opposite spins
There may be no more than 2 electrons in a molecular orbital
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 2 Spins pair The two electrons in the overlap region occupy the same space and therefore must have opposite spins
There may be no more than 2 electrons in a molecular orbital
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Valence Bond Theory
Basic Principle of Localized Electron Model A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies
the region between the two nuclei
Rule 3 Hybridization To explain experimental observations Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms We call this concept
Hybridization
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
What is hybridizationbull Atoms adjust to meet the ldquoneedsrdquo of the
moleculebull In a molecule electrons rearrange in an
attempt to give each atom a noble gas configuration and to minimize electron repulsion
bull Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy
bull The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Example Methane
bull 4 equivalent C-H covalent bondsbull VSEPR predicts a tetrahedral geometry
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
How do we explain formation of 4 equivalent C-H bonds
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Hybridization Mixing of Atomic Orbitals to form New Orbitals for Bonding
+ndash
+ndash
ndash+
+ndash +
ndash
+ +ndash
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Other Representations of Hybridization
y1 = 12[(2s) + (2px) + (2py) + (2pz)]y2 = 12[(2s) + (2px) - (2py) - (2pz)]y3 = 12[(2s) - (2px) + (2py) - (2pz)]y4 = 12[(2s) - (2px) - (2py) + (2pz)]
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Methane (CH4)VSEPR AB4 tetrahedral sp3 hybridized
10947 ordmElectron pair geometry determines hybridization not vice versa
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
NH
HH
1081 ordm
Ammonia (NH3) VSEPR AB3E tetrahedral sp3 hybridized
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Hybridization is related to the number ofvalence electron pairs determined from VSEPR
Water (H2O) VSEPR AB2E2 tetrahedral sp3 hybridized
1056 ordm
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
s bonding and p bondingbull Two modes of bonding are important for1st and 2nd row elements s bonding and p bonding
bull These two differ in their relationship to the internuclear axiss bonds have electron density ALONG the axis
p bonds have electron density ABOVE AND BELOW the axis
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Problem Describe the hybridization and bonding of the carbon orbitals in ethylene
(C2H4) VSEPR AB3 trigonal planar sp2 hybridized orbitals for s bonding
sp2 hybridized orbitals used for s bondingremaining p orbital used for p bonding
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Bonding in ethylene (C2H4)
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Problem Describe the hybridization and bonding of the carbon orbitals in Carbon
Dioxide (CO2) VSEPR AB2 linear sp hybridized orbitals for s bonding
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Bonding in Carbon Dioxide (CO2)
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Atoms of the same kind can have different hybridizations
C2 AB4
sp3
sp
C1 AB2
2s2 2px2pysp sp p p
N ABE2s2 2px2py2pz
sp spp plone pair
s
Bonds
s p p
C1C2 NH
H
H
Acetonitrile (important solvent and industrial chemical)CH3 C N
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
What have we learned so far
bull Molecular orbitals are combinations of atomic orbitals
bull Atomic orbitals are ldquohybridizedrdquo to satisfy bonding in molecules
bull Hybridization follows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Hybridizationbull sp3 Hybridization (CH4)
ndash This is the sum of one s and three p orbitals on the carbon atom
ndash We use just the valence orbitals to make bonds
ndash sp3 hybridization gives rise to the tetrahedral nature of the carbon atom
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Hybridizationbull sp2 Hybridization (H2C=CH2)
ndash This is the sum of one s and two p orbitals on the carbon atom
ndash Leaves one p orbital uninvolved ndash this is free to form a p bond (the second bond in a double bond)
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Hybridizationbull sp Hybridization (O=C=O)
ndash This is the sum of one s and one p orbital on the carbon atom
ndash Leaves two p orbitals free to bond with other atoms (such a O in CO2) or with each other as in HCequivCH
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
General Notesbull This is a model and only goes so far but it is
especially helpful in understanding geometry and expanding Lewis dot structures
bull Orbitals are waves Hybridized orbitals are just the sums of waves ndash constructive and destructive interference
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
What is important to know about hybridization
bull You should be able to give the hybridization of an atom in a molecule based on the formula given
bull Example CH3-CH2-CHObull Step 1 Draw the Lewis Dot Structure
C C C O
H
H
H
H
H H
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
What is important to know about hybridization
bull Step 2 What is the electron pair geometry and molecular shape
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
What is important to know about hybridization
bull Step 3 Use the molecular shape to determine the hybridization
C C C O
H
H
H
H
H H
AX4
AX4
AX3
AXE2
Tetrahedral
Tetrahedral
Trigonal Planar
Trigonal Planar
sp3
sp3 sp2
sp2
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
The Localized Electron Model is very powerfulfor explaining geometries and basic features ofbonding in molecules but it is just a modelMajor limitations of the LE model
bull Assumes electrons are highly localized between the nuclei (sometimes requires resonance structures)
bull Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
bull Doesnrsquot provide direct information about bond energies
Example O2 Lewis dot structure O=O
All electrons are paired Contradicts experiment
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
The Molecular Orbital ModelBasic premise When atomic orbitals interact to form a bond the result is the formation of new molecular orbitals
HY = EY
Important features of molecular orbitals1 Atomic Orbitals are solutions of the Schroumldinger equation for atoms Molecular orbitals are the solutions of the same Schroumldinger equation applied to the molecule
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
2 Atomic Orbitals can hold 2 electrons with opposite spins Molecular Orbitals can hold 2 electrons with opposite spins
3 The electron probability for the Atomic Orbital is given by Y2 The electron probability for the Molecular Orbital is given by Y2
4 Orbitals are conserved - in bringing together 2 atomic orbitals we have to end up with 2 molecular orbitals
How does this work
Molecular Orbital Theory
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
+ -
+
Molecular Orbitals are simplyLinear Combinations of Atomic Orbitals
s bonding
s anti-bonding (s)Example H2
Next Question Why does this work
Molecular Orbitals have phases (+ or -)
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Constructive and Destructive Interference
Constructive interference between two overlapping orbitals leads to a bonding orbital
Destructive interference between two orbitals of opposite sign leads to an anti-bonding orbital
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Bonding is driven by stabilization of electronsbull Electrons are negatively
chargedbull Nuclei are positively charged
The bonding combination puts electron density between the two nuclei - stabilizationThe anti-bonding combination moves electron density away from region between the nuclei - destabilization
= = nucleus+
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
bull We can depict the relative energies of molecular orbitals with a molecular orbital diagram
MO Diagrams
The new molecular orbital is lower in energy than the atomic orbitals
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
s MO is raised in energy
s MO is lowered in energyH atom (1s)1 electron configurationH2 molecule (s1s)2 electron configuration
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Same as previous description of bonding
s
s
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Review of Orbital Fillingbull Pauli Exclusion Principle
ndash No more than 2 e- in an orbital spins must be paired (uarrdarr)
bull Aufbau Principle (aka ldquoBuilding-Uprdquo)ndash Fill the lowest energy levels with electrons first
bull 1s 2s 2p 3s 3p 4s 3d 4p hellipbull Hundrsquos Rule
ndash When more than one orbital has the same energy electrons occupy separate orbitals with parallel spins
Yes No No
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing Energy (Aufbau principle)
H2
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
2) An orbital has a maximum capacity of two electrons with opposite spins (Pauli exclusion principle)
He2
Filling Molecular Orbitals with Electrons
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals) are half filled with spins parallel before any is filled completely (Hundrsquos rule)
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Bond Order
Bond Order =
bonding anti-bonding electrons electrons
2
The bond order is an indication of bond strengthGreater bond order Greater bond strength
(Shorter bond length)
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Bond Order Examples
H2
Bond order = (2-0)2 = 1Single bondStable molecule (436 kJmol bond)
He2
Bond order = (2-2)2 = 0No bondUnstable molecule (0 kJmol bond)
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
He2+
H2+
Bond order = (2-1)2 = 12Half of a single bondCan be made but its not very stable (250 kJmol bond)
Bond order = (1-0)2 = 12Half of a single bondCan be made but its not very stable (255 kJmol bond)
Fractional bond orders are okay
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
bull A s bond can be formed a number of ways
ndashs s overlap
ndashs p overlapndashp p overlap
Forming Bonds
Only orbitals of the same phase (+ +) can form bonds
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
bull For every bonding orbital we form we also form an anti-bonding orbital
Anti-bonding Orbitals
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
MO Theory in Bondingbull Homonuclear atoms (H2 O2 F2 N2)
H2
(Only 1s orbitals
available for
bonding)
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
bull Atomic orbitals must overlap in space in order to participate in molecular orbitals
bull Covalent bonding is dominated by the valence orbitals (only valence orbitals are shown in the MO diagrams)
Covalent Bonding in Homonuclear Diatomics
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
+
+
ndash
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Valence configurations of the 2nd row atoms
Li Be B C N O F2s1 2s2 2s22p1 2s22p2 2s22p3 2s22p4
2s22p5
So far we have focused on bonding involving the s orbitals
What happens when we have to consider the p orbitals
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
For diatomic molecules containing atoms with valence electrons in the p orbitals we must consider three
possible bonding interactions
s-typep-type p-type
= nucleus
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
(+)constructivemixing
(ndash)destructivemixing
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
Major limitations of the LE model
2) Doesnrsquot easily deal with unpaired electrons (incorrectly predicts physical properties in some cases)
Example O2 - Lewis dot structure O=O
- All electrons are paired Contradicts experiment
Experiments show O2 is paramagnetic
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
A quick note on magnetismhellip
Paramagnetic The molecule contains unpaired electrons and is attracted to (has a positive susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired electrons and is not attracted to (has a negative susceptibility to) an applied magnetic field
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
____ s2p ___ ___ p2p
___ ___ ___2p ___ ___ ___ 2p___ ___ p2p
Energy
____ s2p
____ s2s___ 2s ___ 2s
____ s2s
Example the O2 Diatomic
Bond Order = (8-4)2 = 2O2 is stable
(498 kJmol bond strength)
Both have degenerate orbitals
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
Oxygen atom has a 2s22p4 valence configuration
O atom O atomMOO2
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
A prediction from the MO diagram of O2
O=O
The unpaired electrons predicted by the MO diagram should behave as small magnets- O2 should be magnetic
The Lewis dot structure predicts O2 should be diamagnetic-all
electrons are paired
N2 Video
O2 Video
N2 Video
O2 Video
What have we learned so far1 Molecular orbitals (MO) are linear combinations of atomic orbitals
2 Both s and p atomic orbitals can be mixed to form MOs
3 Molecular orbitals are bonding and anti-bonding
4 Bonding and anti-bonding MOs lead to the definition of the bond order
5 Bond order is related to the bond strength (bond dissociation energy)
MO Diagram for H2 vs N2
H2
sp
p2p
s2p
p2p
s2s
s2s
N2
Atomic orbital overlap sometimes forms both s and p bonds Examples N2 O2 F2
1s(N) + 1s(N) 1s(N) ndash1s(N)-37875 -37871
-2965
Elec
tron
ener
gy (k
J mol
-1)
Core Core
Valence Valence
s(2s)
s(2s)
-1479
-1240 -1240
p p
s(2p)
-1155
p p
s(2p)
MO Diagram for N2
A ComplicationhellipMO Diagram for O2
(similar for F2 and Xe2)MO Diagram for B2
(similar for C2 and N2)O OO2
A ComplicationhellipMO Diagram for O2
(similar for F2 and Ne2)MO Diagram for B2
(similar for C2 and N2)
s-p mixing
No s-p mixing
Electron repulsion
Why does s-p mixing occur
s2s and s2p both have significant e- probability between the nuclei so e- in s2s
will repel e- in s2pEffect will decrease as you move across the Periodic Table increased nuclear charge pulls the s2s e-
closer making the s2s orbital smaller and decreasing the s2s and s2p interaction
Molecular Orbitals of X2 Molecules
sp orbital mixing (a little hybridization)bull lowers the energy of the s2s orbitals andbull raises the energy of the s2p orbitals
bull As a result E(s2p) gt E(p 2p) for B2 C2 and N2 bull As one moves right in Row 2 2s and 2p get
further apart in energy decreasing sndashp mixing E(s2p) lt E(p2p) for O2 F2 and Ne2 See text pages 680-681
bull Note that sndashp mixing does not affect bond order or magnetism in the common diatomics (N2 O2 and F2) Hence it is not of much practical importance
s-p mixing
No s-p mixing
B C and N all have 12 filled 2p orbitalsWhen does s-p mixing occur
O F and Xe all have gt 12 filled 2p orbitals
bull If 2 electrons are forced to be in the same orbital their energies go up
bull Electrons repel each other because they are negatively charged
bull Having gt 12 filled 2p orbitals raises the energies of these orbitals due to e- - e- repulsion
s-p mixing only occurs when the s and p atomic orbitals are close in energy ( 12 filled 2p orbitals)
Relating the MO Diagrams to Physical Properties
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11
Removing an electron from N2 decreases the bond energy of the resulting ion whereas removing an electron from O2 increases the bond energy of the resulting ion Explain these facts using MO diagrams
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11
Plan We first draw the MO energy levels for the four species recalling that they differ for N2 and O2 Then we determine the bond orders and compare them with the data bond order is related directly to bond energy and inversely to bond length
Sample Problem - ContinuedSolution The MO energy levels are
N2
sp
p2p
N2+
s2p
p2p
s2s
s2s
O2 O2+
sp
s2s
s2s
p2p
p2p
s2p
Bond Orders(8-2)2 = 3 (7-2)2 = 25 (8-4)2 = 2 (8-3)2 = 25
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11Bond Order 3 25 2 25
What have we learned so far1 Molecular orbitals (MO) explain the properties of valence electrons in molecules (Example O2)
2 s and p atomic orbitals can be mixed to form s s p and p molecular orbitals
3 Electrons in p or p molecular orbitals can have the same energies Degenerate orbitals
4 The ordering of s2p and p2p molecular orbitals depends on the electron occupancy s-p mixing
Bonding in Diatomic Molecules
Covalent
Ionic
IonicCovalent
HomonuclearH2
Nonpolar covalent bond(450 kJmol bond)
HeteronuclearHF
Polar covalent bond(565 kJmol bond)
Electronegativity
Electronegativity
Because F (EN = 40) is more electronegative than H (EN = 22) the electrons move closer to F
This gives rise to a polar bond
H F
Figure 1445
Electrons are not equally sharedin heteronuclear bondsHF
s Antibonding (s)Mostly H(1s)
s BondingMostly F(2p)
H F
H F
MOs of a Polar Covalent Bond HF
This approach simplifies model and only considers electrons involved in bond
MOs OF XY MOLECULES
Equal or unequal e sharing between 2 atoms is reflected in the composition of the MOsmiddot When 2 atoms X and Y have the same electronegativity (purely
covalent bond) their overlapping AOs have the same energy and the bonding and anti-bonding MOs are each half X and half Y AO All electrons spend equal time near X and Y Examples N2 O2 F2
middot If EN(Y) gt EN(X) (polar covalent X+Y) the Y AO has lower energy than the X AO The bonding MO is more like the Y AO and the anti-bonding MO more like the X AO Bonding e spend more time near Y than X vice versa for anti-bonding e Example CO
MOs OF XY MOLECULES ____ s ___ ___ p ___ ___ ___2p ___ ___ ___ 2p uarr ____ s
Energy ___ ___ p ____ s ___ 2s ___ 2s ____ s
middot CO Bond Order = 30 (same as N2)middot CO Bond Energy = 1076 kJmol (N2 = 945 kJmol)middot Isoelectronic to CO and N2 CNndash NO+middot NO has 1endash in p bond order = 25 this endash is more on N than O
NO NO+ easyhellip
C Atom (4endash) Cδ+Oδndash (10endash) O Atom (6endash)
___ ___ ___ 2p ___ 2s
Electronegativity
Bonding in NObull Two possible Lewis dot structures
for NObull The simplest structure minimizes
formal charges and places the lone (unpaired) electron on the nitrogen
bull The Lewis structure predicts a bond order of 2 but experimental evidence suggests a bond order between 2 and 3
bull How does MO theory help us understand bonding in NO
N=O
N=O
+1-1
When the electronegativities of the 2 atoms are more similar the bonding
becomes less polar
EN(N) = 30EN(O) = 34
N=O
Electronegativity
2s2s
2p2p
NON OBond order = 25 unpaired electron is in a N-like orbital
NO+
oxidation
NO
NO is easily oxidized to form NO+ Why What changes can we predict in the bonding and magnetism of the molecule
Bond Order = (8-3)2 = 25Paramagnetic
Bond Order = (8-2)2 = 3Diamagnetic
s2s
s2s
p2p p2p
p2p
-3320
-1835
-1444 -1374
s2p -1307
-597 p2p (empty)
MO diagram for NO
Key Points of MO Theory ndash Heteronuclear Molecules
bull The more electronegative atom has orbitals lower in energy than the more positive atom
bull Electrons in bonding orbitals are closer to the more electronegative atom anti-bonding electrons are closer to the more positive atom
bull For most diatomic molecules s-p mixing changes the orbital energy levels but since these orbitals are almost always fully occupied their order is less important to us
Combining the Localized Electron and Molecular Orbital Models (into a convenient working
model)
Figure 1447
Only the p bonding changes between these resonance structures - The MO model describes this p bonding more
effectively
Figure 1451
Atomic Orbitals Molecular Orbitals
Another example Benzene
p bonding
s bonding
p atomic orbitals p molecular orbital
MO Theory Expectationsbull You should be able to
ndash predict which atomic orbitals are higher or lower in energy (based on electronegativity differences)
ndash correctly fill a molecular orbital diagramndash correctly calculate bond orderndash predict molecular magnetic properties
based on orbital occupationndash understand how molecular properties
change upon ionization (oxidation or reduction) of molecules
- What Do Molecules Look Like
- Slide 2
- Slide 3
- Arranging Electron Pairs
- Arranging Electron Pairs (cont)
- VSEPR Structure Guidelines
- VSEPR 2 electron pairs
- VSEPR 3 electron pairs
- VSEPR 4 electron pairs (cont)
- Bonding vs Lone pairs
- VSEPR 4 electron pairs
- VSEPR 4 electron pairs (cont) (2)
- VSEPR 4 electron pairs (cont) (3)
- Slide 14
- Molecular vs Electron-Pair Geometry
- Slide 16
- VSEPR Beyond the Octet
- VSEPR 5 electron pairs
- VSEPR 5 electron pairs (2)
- 5-electron-pair geometries
- VSEPR 6 electron pairs
- Slide 22
- Molecular Dipole Moments
- Slide 24
- Molecular Dipoles (cont)
- Molecular Dipole Example
- Advanced VSEPR Application
- The VSEPR Table
- The VSEPR Table (2)
- Slide 30
- Slide 31
- Slide 32
- Slide 33
- Slide 34
- Slide 35
- Slide 36
- Slide 37
- Slide 38
- Slide 39
- Slide 40
- Slide 41
- Slide 42
- Slide 43
- Slide 44
- Slide 45
- Slide 46
- Slide 47
- Slide 48
- Hybridization
- Hybridization (2)
- Hybridization (3)
- General Notes
- What is important to know about hybridization
- What is important to know about hybridization (2)
- What is important to know about hybridization (3)
- Slide 56
- Slide 57
- Molecular Orbital Theory
- Slide 59
- Constructive and Destructive Interference
- Slide 61
- MO Diagrams
- Slide 63
- Slide 64
- Review of Orbital Filling
- Slide 66
- Slide 67
- Filling Molecular Orbitals with Electrons
- Slide 69
- Slide 70
- Slide 71
- Forming Bonds
- Anti-bonding Orbitals
- MO Theory in Bonding
- Slide 75
- Slide 76
- Slide 77
- Slide 78
- Slide 79
- Slide 80
- Slide 81
- Slide 82
- Slide 83
- Slide 84
- Slide 85
- MO Diagram for H2 vs N2
- Slide 87
- Slide 88
- Slide 89
- Slide 90
- Molecular Orbitals of X2 Molecules
- Slide 92
- Slide 93
- Slide 94
- Slide 95
- Slide 96
- Slide 97
- Slide 98
- Slide 99
- Slide 100
- Slide 101
- Slide 102
- Slide 103
- Slide 104
- Slide 105
- Bonding in NO
- Slide 107
- Slide 108
- Slide 109
- Key Points of MO Theory ndash Heteronuclear Molecules
- Slide 111
- Slide 112
- Slide 113
- MO Theory Expectations
-
MO Diagram for H2 vs N2
H2
sp
p2p
s2p
p2p
s2s
s2s
N2
Atomic orbital overlap sometimes forms both s and p bonds Examples N2 O2 F2
1s(N) + 1s(N) 1s(N) ndash1s(N)-37875 -37871
-2965
Elec
tron
ener
gy (k
J mol
-1)
Core Core
Valence Valence
s(2s)
s(2s)
-1479
-1240 -1240
p p
s(2p)
-1155
p p
s(2p)
MO Diagram for N2
A ComplicationhellipMO Diagram for O2
(similar for F2 and Xe2)MO Diagram for B2
(similar for C2 and N2)O OO2
A ComplicationhellipMO Diagram for O2
(similar for F2 and Ne2)MO Diagram for B2
(similar for C2 and N2)
s-p mixing
No s-p mixing
Electron repulsion
Why does s-p mixing occur
s2s and s2p both have significant e- probability between the nuclei so e- in s2s
will repel e- in s2pEffect will decrease as you move across the Periodic Table increased nuclear charge pulls the s2s e-
closer making the s2s orbital smaller and decreasing the s2s and s2p interaction
Molecular Orbitals of X2 Molecules
sp orbital mixing (a little hybridization)bull lowers the energy of the s2s orbitals andbull raises the energy of the s2p orbitals
bull As a result E(s2p) gt E(p 2p) for B2 C2 and N2 bull As one moves right in Row 2 2s and 2p get
further apart in energy decreasing sndashp mixing E(s2p) lt E(p2p) for O2 F2 and Ne2 See text pages 680-681
bull Note that sndashp mixing does not affect bond order or magnetism in the common diatomics (N2 O2 and F2) Hence it is not of much practical importance
s-p mixing
No s-p mixing
B C and N all have 12 filled 2p orbitalsWhen does s-p mixing occur
O F and Xe all have gt 12 filled 2p orbitals
bull If 2 electrons are forced to be in the same orbital their energies go up
bull Electrons repel each other because they are negatively charged
bull Having gt 12 filled 2p orbitals raises the energies of these orbitals due to e- - e- repulsion
s-p mixing only occurs when the s and p atomic orbitals are close in energy ( 12 filled 2p orbitals)
Relating the MO Diagrams to Physical Properties
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11
Removing an electron from N2 decreases the bond energy of the resulting ion whereas removing an electron from O2 increases the bond energy of the resulting ion Explain these facts using MO diagrams
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11
Plan We first draw the MO energy levels for the four species recalling that they differ for N2 and O2 Then we determine the bond orders and compare them with the data bond order is related directly to bond energy and inversely to bond length
Sample Problem - ContinuedSolution The MO energy levels are
N2
sp
p2p
N2+
s2p
p2p
s2s
s2s
O2 O2+
sp
s2s
s2s
p2p
p2p
s2p
Bond Orders(8-2)2 = 3 (7-2)2 = 25 (8-4)2 = 2 (8-3)2 = 25
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11Bond Order 3 25 2 25
What have we learned so far1 Molecular orbitals (MO) explain the properties of valence electrons in molecules (Example O2)
2 s and p atomic orbitals can be mixed to form s s p and p molecular orbitals
3 Electrons in p or p molecular orbitals can have the same energies Degenerate orbitals
4 The ordering of s2p and p2p molecular orbitals depends on the electron occupancy s-p mixing
Bonding in Diatomic Molecules
Covalent
Ionic
IonicCovalent
HomonuclearH2
Nonpolar covalent bond(450 kJmol bond)
HeteronuclearHF
Polar covalent bond(565 kJmol bond)
Electronegativity
Electronegativity
Because F (EN = 40) is more electronegative than H (EN = 22) the electrons move closer to F
This gives rise to a polar bond
H F
Figure 1445
Electrons are not equally sharedin heteronuclear bondsHF
s Antibonding (s)Mostly H(1s)
s BondingMostly F(2p)
H F
H F
MOs of a Polar Covalent Bond HF
This approach simplifies model and only considers electrons involved in bond
MOs OF XY MOLECULES
Equal or unequal e sharing between 2 atoms is reflected in the composition of the MOsmiddot When 2 atoms X and Y have the same electronegativity (purely
covalent bond) their overlapping AOs have the same energy and the bonding and anti-bonding MOs are each half X and half Y AO All electrons spend equal time near X and Y Examples N2 O2 F2
middot If EN(Y) gt EN(X) (polar covalent X+Y) the Y AO has lower energy than the X AO The bonding MO is more like the Y AO and the anti-bonding MO more like the X AO Bonding e spend more time near Y than X vice versa for anti-bonding e Example CO
MOs OF XY MOLECULES ____ s ___ ___ p ___ ___ ___2p ___ ___ ___ 2p uarr ____ s
Energy ___ ___ p ____ s ___ 2s ___ 2s ____ s
middot CO Bond Order = 30 (same as N2)middot CO Bond Energy = 1076 kJmol (N2 = 945 kJmol)middot Isoelectronic to CO and N2 CNndash NO+middot NO has 1endash in p bond order = 25 this endash is more on N than O
NO NO+ easyhellip
C Atom (4endash) Cδ+Oδndash (10endash) O Atom (6endash)
___ ___ ___ 2p ___ 2s
Electronegativity
Bonding in NObull Two possible Lewis dot structures
for NObull The simplest structure minimizes
formal charges and places the lone (unpaired) electron on the nitrogen
bull The Lewis structure predicts a bond order of 2 but experimental evidence suggests a bond order between 2 and 3
bull How does MO theory help us understand bonding in NO
N=O
N=O
+1-1
When the electronegativities of the 2 atoms are more similar the bonding
becomes less polar
EN(N) = 30EN(O) = 34
N=O
Electronegativity
2s2s
2p2p
NON OBond order = 25 unpaired electron is in a N-like orbital
NO+
oxidation
NO
NO is easily oxidized to form NO+ Why What changes can we predict in the bonding and magnetism of the molecule
Bond Order = (8-3)2 = 25Paramagnetic
Bond Order = (8-2)2 = 3Diamagnetic
s2s
s2s
p2p p2p
p2p
-3320
-1835
-1444 -1374
s2p -1307
-597 p2p (empty)
MO diagram for NO
Key Points of MO Theory ndash Heteronuclear Molecules
bull The more electronegative atom has orbitals lower in energy than the more positive atom
bull Electrons in bonding orbitals are closer to the more electronegative atom anti-bonding electrons are closer to the more positive atom
bull For most diatomic molecules s-p mixing changes the orbital energy levels but since these orbitals are almost always fully occupied their order is less important to us
Combining the Localized Electron and Molecular Orbital Models (into a convenient working
model)
Figure 1447
Only the p bonding changes between these resonance structures - The MO model describes this p bonding more
effectively
Figure 1451
Atomic Orbitals Molecular Orbitals
Another example Benzene
p bonding
s bonding
p atomic orbitals p molecular orbital
MO Theory Expectationsbull You should be able to
ndash predict which atomic orbitals are higher or lower in energy (based on electronegativity differences)
ndash correctly fill a molecular orbital diagramndash correctly calculate bond orderndash predict molecular magnetic properties
based on orbital occupationndash understand how molecular properties
change upon ionization (oxidation or reduction) of molecules
- What Do Molecules Look Like
- Slide 2
- Slide 3
- Arranging Electron Pairs
- Arranging Electron Pairs (cont)
- VSEPR Structure Guidelines
- VSEPR 2 electron pairs
- VSEPR 3 electron pairs
- VSEPR 4 electron pairs (cont)
- Bonding vs Lone pairs
- VSEPR 4 electron pairs
- VSEPR 4 electron pairs (cont) (2)
- VSEPR 4 electron pairs (cont) (3)
- Slide 14
- Molecular vs Electron-Pair Geometry
- Slide 16
- VSEPR Beyond the Octet
- VSEPR 5 electron pairs
- VSEPR 5 electron pairs (2)
- 5-electron-pair geometries
- VSEPR 6 electron pairs
- Slide 22
- Molecular Dipole Moments
- Slide 24
- Molecular Dipoles (cont)
- Molecular Dipole Example
- Advanced VSEPR Application
- The VSEPR Table
- The VSEPR Table (2)
- Slide 30
- Slide 31
- Slide 32
- Slide 33
- Slide 34
- Slide 35
- Slide 36
- Slide 37
- Slide 38
- Slide 39
- Slide 40
- Slide 41
- Slide 42
- Slide 43
- Slide 44
- Slide 45
- Slide 46
- Slide 47
- Slide 48
- Hybridization
- Hybridization (2)
- Hybridization (3)
- General Notes
- What is important to know about hybridization
- What is important to know about hybridization (2)
- What is important to know about hybridization (3)
- Slide 56
- Slide 57
- Molecular Orbital Theory
- Slide 59
- Constructive and Destructive Interference
- Slide 61
- MO Diagrams
- Slide 63
- Slide 64
- Review of Orbital Filling
- Slide 66
- Slide 67
- Filling Molecular Orbitals with Electrons
- Slide 69
- Slide 70
- Slide 71
- Forming Bonds
- Anti-bonding Orbitals
- MO Theory in Bonding
- Slide 75
- Slide 76
- Slide 77
- Slide 78
- Slide 79
- Slide 80
- Slide 81
- Slide 82
- Slide 83
- Slide 84
- Slide 85
- MO Diagram for H2 vs N2
- Slide 87
- Slide 88
- Slide 89
- Slide 90
- Molecular Orbitals of X2 Molecules
- Slide 92
- Slide 93
- Slide 94
- Slide 95
- Slide 96
- Slide 97
- Slide 98
- Slide 99
- Slide 100
- Slide 101
- Slide 102
- Slide 103
- Slide 104
- Slide 105
- Bonding in NO
- Slide 107
- Slide 108
- Slide 109
- Key Points of MO Theory ndash Heteronuclear Molecules
- Slide 111
- Slide 112
- Slide 113
- MO Theory Expectations
-
1s(N) + 1s(N) 1s(N) ndash1s(N)-37875 -37871
-2965
Elec
tron
ener
gy (k
J mol
-1)
Core Core
Valence Valence
s(2s)
s(2s)
-1479
-1240 -1240
p p
s(2p)
-1155
p p
s(2p)
MO Diagram for N2
A ComplicationhellipMO Diagram for O2
(similar for F2 and Xe2)MO Diagram for B2
(similar for C2 and N2)O OO2
A ComplicationhellipMO Diagram for O2
(similar for F2 and Ne2)MO Diagram for B2
(similar for C2 and N2)
s-p mixing
No s-p mixing
Electron repulsion
Why does s-p mixing occur
s2s and s2p both have significant e- probability between the nuclei so e- in s2s
will repel e- in s2pEffect will decrease as you move across the Periodic Table increased nuclear charge pulls the s2s e-
closer making the s2s orbital smaller and decreasing the s2s and s2p interaction
Molecular Orbitals of X2 Molecules
sp orbital mixing (a little hybridization)bull lowers the energy of the s2s orbitals andbull raises the energy of the s2p orbitals
bull As a result E(s2p) gt E(p 2p) for B2 C2 and N2 bull As one moves right in Row 2 2s and 2p get
further apart in energy decreasing sndashp mixing E(s2p) lt E(p2p) for O2 F2 and Ne2 See text pages 680-681
bull Note that sndashp mixing does not affect bond order or magnetism in the common diatomics (N2 O2 and F2) Hence it is not of much practical importance
s-p mixing
No s-p mixing
B C and N all have 12 filled 2p orbitalsWhen does s-p mixing occur
O F and Xe all have gt 12 filled 2p orbitals
bull If 2 electrons are forced to be in the same orbital their energies go up
bull Electrons repel each other because they are negatively charged
bull Having gt 12 filled 2p orbitals raises the energies of these orbitals due to e- - e- repulsion
s-p mixing only occurs when the s and p atomic orbitals are close in energy ( 12 filled 2p orbitals)
Relating the MO Diagrams to Physical Properties
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11
Removing an electron from N2 decreases the bond energy of the resulting ion whereas removing an electron from O2 increases the bond energy of the resulting ion Explain these facts using MO diagrams
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11
Plan We first draw the MO energy levels for the four species recalling that they differ for N2 and O2 Then we determine the bond orders and compare them with the data bond order is related directly to bond energy and inversely to bond length
Sample Problem - ContinuedSolution The MO energy levels are
N2
sp
p2p
N2+
s2p
p2p
s2s
s2s
O2 O2+
sp
s2s
s2s
p2p
p2p
s2p
Bond Orders(8-2)2 = 3 (7-2)2 = 25 (8-4)2 = 2 (8-3)2 = 25
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11Bond Order 3 25 2 25
What have we learned so far1 Molecular orbitals (MO) explain the properties of valence electrons in molecules (Example O2)
2 s and p atomic orbitals can be mixed to form s s p and p molecular orbitals
3 Electrons in p or p molecular orbitals can have the same energies Degenerate orbitals
4 The ordering of s2p and p2p molecular orbitals depends on the electron occupancy s-p mixing
Bonding in Diatomic Molecules
Covalent
Ionic
IonicCovalent
HomonuclearH2
Nonpolar covalent bond(450 kJmol bond)
HeteronuclearHF
Polar covalent bond(565 kJmol bond)
Electronegativity
Electronegativity
Because F (EN = 40) is more electronegative than H (EN = 22) the electrons move closer to F
This gives rise to a polar bond
H F
Figure 1445
Electrons are not equally sharedin heteronuclear bondsHF
s Antibonding (s)Mostly H(1s)
s BondingMostly F(2p)
H F
H F
MOs of a Polar Covalent Bond HF
This approach simplifies model and only considers electrons involved in bond
MOs OF XY MOLECULES
Equal or unequal e sharing between 2 atoms is reflected in the composition of the MOsmiddot When 2 atoms X and Y have the same electronegativity (purely
covalent bond) their overlapping AOs have the same energy and the bonding and anti-bonding MOs are each half X and half Y AO All electrons spend equal time near X and Y Examples N2 O2 F2
middot If EN(Y) gt EN(X) (polar covalent X+Y) the Y AO has lower energy than the X AO The bonding MO is more like the Y AO and the anti-bonding MO more like the X AO Bonding e spend more time near Y than X vice versa for anti-bonding e Example CO
MOs OF XY MOLECULES ____ s ___ ___ p ___ ___ ___2p ___ ___ ___ 2p uarr ____ s
Energy ___ ___ p ____ s ___ 2s ___ 2s ____ s
middot CO Bond Order = 30 (same as N2)middot CO Bond Energy = 1076 kJmol (N2 = 945 kJmol)middot Isoelectronic to CO and N2 CNndash NO+middot NO has 1endash in p bond order = 25 this endash is more on N than O
NO NO+ easyhellip
C Atom (4endash) Cδ+Oδndash (10endash) O Atom (6endash)
___ ___ ___ 2p ___ 2s
Electronegativity
Bonding in NObull Two possible Lewis dot structures
for NObull The simplest structure minimizes
formal charges and places the lone (unpaired) electron on the nitrogen
bull The Lewis structure predicts a bond order of 2 but experimental evidence suggests a bond order between 2 and 3
bull How does MO theory help us understand bonding in NO
N=O
N=O
+1-1
When the electronegativities of the 2 atoms are more similar the bonding
becomes less polar
EN(N) = 30EN(O) = 34
N=O
Electronegativity
2s2s
2p2p
NON OBond order = 25 unpaired electron is in a N-like orbital
NO+
oxidation
NO
NO is easily oxidized to form NO+ Why What changes can we predict in the bonding and magnetism of the molecule
Bond Order = (8-3)2 = 25Paramagnetic
Bond Order = (8-2)2 = 3Diamagnetic
s2s
s2s
p2p p2p
p2p
-3320
-1835
-1444 -1374
s2p -1307
-597 p2p (empty)
MO diagram for NO
Key Points of MO Theory ndash Heteronuclear Molecules
bull The more electronegative atom has orbitals lower in energy than the more positive atom
bull Electrons in bonding orbitals are closer to the more electronegative atom anti-bonding electrons are closer to the more positive atom
bull For most diatomic molecules s-p mixing changes the orbital energy levels but since these orbitals are almost always fully occupied their order is less important to us
Combining the Localized Electron and Molecular Orbital Models (into a convenient working
model)
Figure 1447
Only the p bonding changes between these resonance structures - The MO model describes this p bonding more
effectively
Figure 1451
Atomic Orbitals Molecular Orbitals
Another example Benzene
p bonding
s bonding
p atomic orbitals p molecular orbital
MO Theory Expectationsbull You should be able to
ndash predict which atomic orbitals are higher or lower in energy (based on electronegativity differences)
ndash correctly fill a molecular orbital diagramndash correctly calculate bond orderndash predict molecular magnetic properties
based on orbital occupationndash understand how molecular properties
change upon ionization (oxidation or reduction) of molecules
- What Do Molecules Look Like
- Slide 2
- Slide 3
- Arranging Electron Pairs
- Arranging Electron Pairs (cont)
- VSEPR Structure Guidelines
- VSEPR 2 electron pairs
- VSEPR 3 electron pairs
- VSEPR 4 electron pairs (cont)
- Bonding vs Lone pairs
- VSEPR 4 electron pairs
- VSEPR 4 electron pairs (cont) (2)
- VSEPR 4 electron pairs (cont) (3)
- Slide 14
- Molecular vs Electron-Pair Geometry
- Slide 16
- VSEPR Beyond the Octet
- VSEPR 5 electron pairs
- VSEPR 5 electron pairs (2)
- 5-electron-pair geometries
- VSEPR 6 electron pairs
- Slide 22
- Molecular Dipole Moments
- Slide 24
- Molecular Dipoles (cont)
- Molecular Dipole Example
- Advanced VSEPR Application
- The VSEPR Table
- The VSEPR Table (2)
- Slide 30
- Slide 31
- Slide 32
- Slide 33
- Slide 34
- Slide 35
- Slide 36
- Slide 37
- Slide 38
- Slide 39
- Slide 40
- Slide 41
- Slide 42
- Slide 43
- Slide 44
- Slide 45
- Slide 46
- Slide 47
- Slide 48
- Hybridization
- Hybridization (2)
- Hybridization (3)
- General Notes
- What is important to know about hybridization
- What is important to know about hybridization (2)
- What is important to know about hybridization (3)
- Slide 56
- Slide 57
- Molecular Orbital Theory
- Slide 59
- Constructive and Destructive Interference
- Slide 61
- MO Diagrams
- Slide 63
- Slide 64
- Review of Orbital Filling
- Slide 66
- Slide 67
- Filling Molecular Orbitals with Electrons
- Slide 69
- Slide 70
- Slide 71
- Forming Bonds
- Anti-bonding Orbitals
- MO Theory in Bonding
- Slide 75
- Slide 76
- Slide 77
- Slide 78
- Slide 79
- Slide 80
- Slide 81
- Slide 82
- Slide 83
- Slide 84
- Slide 85
- MO Diagram for H2 vs N2
- Slide 87
- Slide 88
- Slide 89
- Slide 90
- Molecular Orbitals of X2 Molecules
- Slide 92
- Slide 93
- Slide 94
- Slide 95
- Slide 96
- Slide 97
- Slide 98
- Slide 99
- Slide 100
- Slide 101
- Slide 102
- Slide 103
- Slide 104
- Slide 105
- Bonding in NO
- Slide 107
- Slide 108
- Slide 109
- Key Points of MO Theory ndash Heteronuclear Molecules
- Slide 111
- Slide 112
- Slide 113
- MO Theory Expectations
-
A ComplicationhellipMO Diagram for O2
(similar for F2 and Xe2)MO Diagram for B2
(similar for C2 and N2)O OO2
A ComplicationhellipMO Diagram for O2
(similar for F2 and Ne2)MO Diagram for B2
(similar for C2 and N2)
s-p mixing
No s-p mixing
Electron repulsion
Why does s-p mixing occur
s2s and s2p both have significant e- probability between the nuclei so e- in s2s
will repel e- in s2pEffect will decrease as you move across the Periodic Table increased nuclear charge pulls the s2s e-
closer making the s2s orbital smaller and decreasing the s2s and s2p interaction
Molecular Orbitals of X2 Molecules
sp orbital mixing (a little hybridization)bull lowers the energy of the s2s orbitals andbull raises the energy of the s2p orbitals
bull As a result E(s2p) gt E(p 2p) for B2 C2 and N2 bull As one moves right in Row 2 2s and 2p get
further apart in energy decreasing sndashp mixing E(s2p) lt E(p2p) for O2 F2 and Ne2 See text pages 680-681
bull Note that sndashp mixing does not affect bond order or magnetism in the common diatomics (N2 O2 and F2) Hence it is not of much practical importance
s-p mixing
No s-p mixing
B C and N all have 12 filled 2p orbitalsWhen does s-p mixing occur
O F and Xe all have gt 12 filled 2p orbitals
bull If 2 electrons are forced to be in the same orbital their energies go up
bull Electrons repel each other because they are negatively charged
bull Having gt 12 filled 2p orbitals raises the energies of these orbitals due to e- - e- repulsion
s-p mixing only occurs when the s and p atomic orbitals are close in energy ( 12 filled 2p orbitals)
Relating the MO Diagrams to Physical Properties
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11
Removing an electron from N2 decreases the bond energy of the resulting ion whereas removing an electron from O2 increases the bond energy of the resulting ion Explain these facts using MO diagrams
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11
Plan We first draw the MO energy levels for the four species recalling that they differ for N2 and O2 Then we determine the bond orders and compare them with the data bond order is related directly to bond energy and inversely to bond length
Sample Problem - ContinuedSolution The MO energy levels are
N2
sp
p2p
N2+
s2p
p2p
s2s
s2s
O2 O2+
sp
s2s
s2s
p2p
p2p
s2p
Bond Orders(8-2)2 = 3 (7-2)2 = 25 (8-4)2 = 2 (8-3)2 = 25
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11Bond Order 3 25 2 25
What have we learned so far1 Molecular orbitals (MO) explain the properties of valence electrons in molecules (Example O2)
2 s and p atomic orbitals can be mixed to form s s p and p molecular orbitals
3 Electrons in p or p molecular orbitals can have the same energies Degenerate orbitals
4 The ordering of s2p and p2p molecular orbitals depends on the electron occupancy s-p mixing
Bonding in Diatomic Molecules
Covalent
Ionic
IonicCovalent
HomonuclearH2
Nonpolar covalent bond(450 kJmol bond)
HeteronuclearHF
Polar covalent bond(565 kJmol bond)
Electronegativity
Electronegativity
Because F (EN = 40) is more electronegative than H (EN = 22) the electrons move closer to F
This gives rise to a polar bond
H F
Figure 1445
Electrons are not equally sharedin heteronuclear bondsHF
s Antibonding (s)Mostly H(1s)
s BondingMostly F(2p)
H F
H F
MOs of a Polar Covalent Bond HF
This approach simplifies model and only considers electrons involved in bond
MOs OF XY MOLECULES
Equal or unequal e sharing between 2 atoms is reflected in the composition of the MOsmiddot When 2 atoms X and Y have the same electronegativity (purely
covalent bond) their overlapping AOs have the same energy and the bonding and anti-bonding MOs are each half X and half Y AO All electrons spend equal time near X and Y Examples N2 O2 F2
middot If EN(Y) gt EN(X) (polar covalent X+Y) the Y AO has lower energy than the X AO The bonding MO is more like the Y AO and the anti-bonding MO more like the X AO Bonding e spend more time near Y than X vice versa for anti-bonding e Example CO
MOs OF XY MOLECULES ____ s ___ ___ p ___ ___ ___2p ___ ___ ___ 2p uarr ____ s
Energy ___ ___ p ____ s ___ 2s ___ 2s ____ s
middot CO Bond Order = 30 (same as N2)middot CO Bond Energy = 1076 kJmol (N2 = 945 kJmol)middot Isoelectronic to CO and N2 CNndash NO+middot NO has 1endash in p bond order = 25 this endash is more on N than O
NO NO+ easyhellip
C Atom (4endash) Cδ+Oδndash (10endash) O Atom (6endash)
___ ___ ___ 2p ___ 2s
Electronegativity
Bonding in NObull Two possible Lewis dot structures
for NObull The simplest structure minimizes
formal charges and places the lone (unpaired) electron on the nitrogen
bull The Lewis structure predicts a bond order of 2 but experimental evidence suggests a bond order between 2 and 3
bull How does MO theory help us understand bonding in NO
N=O
N=O
+1-1
When the electronegativities of the 2 atoms are more similar the bonding
becomes less polar
EN(N) = 30EN(O) = 34
N=O
Electronegativity
2s2s
2p2p
NON OBond order = 25 unpaired electron is in a N-like orbital
NO+
oxidation
NO
NO is easily oxidized to form NO+ Why What changes can we predict in the bonding and magnetism of the molecule
Bond Order = (8-3)2 = 25Paramagnetic
Bond Order = (8-2)2 = 3Diamagnetic
s2s
s2s
p2p p2p
p2p
-3320
-1835
-1444 -1374
s2p -1307
-597 p2p (empty)
MO diagram for NO
Key Points of MO Theory ndash Heteronuclear Molecules
bull The more electronegative atom has orbitals lower in energy than the more positive atom
bull Electrons in bonding orbitals are closer to the more electronegative atom anti-bonding electrons are closer to the more positive atom
bull For most diatomic molecules s-p mixing changes the orbital energy levels but since these orbitals are almost always fully occupied their order is less important to us
Combining the Localized Electron and Molecular Orbital Models (into a convenient working
model)
Figure 1447
Only the p bonding changes between these resonance structures - The MO model describes this p bonding more
effectively
Figure 1451
Atomic Orbitals Molecular Orbitals
Another example Benzene
p bonding
s bonding
p atomic orbitals p molecular orbital
MO Theory Expectationsbull You should be able to
ndash predict which atomic orbitals are higher or lower in energy (based on electronegativity differences)
ndash correctly fill a molecular orbital diagramndash correctly calculate bond orderndash predict molecular magnetic properties
based on orbital occupationndash understand how molecular properties
change upon ionization (oxidation or reduction) of molecules
- What Do Molecules Look Like
- Slide 2
- Slide 3
- Arranging Electron Pairs
- Arranging Electron Pairs (cont)
- VSEPR Structure Guidelines
- VSEPR 2 electron pairs
- VSEPR 3 electron pairs
- VSEPR 4 electron pairs (cont)
- Bonding vs Lone pairs
- VSEPR 4 electron pairs
- VSEPR 4 electron pairs (cont) (2)
- VSEPR 4 electron pairs (cont) (3)
- Slide 14
- Molecular vs Electron-Pair Geometry
- Slide 16
- VSEPR Beyond the Octet
- VSEPR 5 electron pairs
- VSEPR 5 electron pairs (2)
- 5-electron-pair geometries
- VSEPR 6 electron pairs
- Slide 22
- Molecular Dipole Moments
- Slide 24
- Molecular Dipoles (cont)
- Molecular Dipole Example
- Advanced VSEPR Application
- The VSEPR Table
- The VSEPR Table (2)
- Slide 30
- Slide 31
- Slide 32
- Slide 33
- Slide 34
- Slide 35
- Slide 36
- Slide 37
- Slide 38
- Slide 39
- Slide 40
- Slide 41
- Slide 42
- Slide 43
- Slide 44
- Slide 45
- Slide 46
- Slide 47
- Slide 48
- Hybridization
- Hybridization (2)
- Hybridization (3)
- General Notes
- What is important to know about hybridization
- What is important to know about hybridization (2)
- What is important to know about hybridization (3)
- Slide 56
- Slide 57
- Molecular Orbital Theory
- Slide 59
- Constructive and Destructive Interference
- Slide 61
- MO Diagrams
- Slide 63
- Slide 64
- Review of Orbital Filling
- Slide 66
- Slide 67
- Filling Molecular Orbitals with Electrons
- Slide 69
- Slide 70
- Slide 71
- Forming Bonds
- Anti-bonding Orbitals
- MO Theory in Bonding
- Slide 75
- Slide 76
- Slide 77
- Slide 78
- Slide 79
- Slide 80
- Slide 81
- Slide 82
- Slide 83
- Slide 84
- Slide 85
- MO Diagram for H2 vs N2
- Slide 87
- Slide 88
- Slide 89
- Slide 90
- Molecular Orbitals of X2 Molecules
- Slide 92
- Slide 93
- Slide 94
- Slide 95
- Slide 96
- Slide 97
- Slide 98
- Slide 99
- Slide 100
- Slide 101
- Slide 102
- Slide 103
- Slide 104
- Slide 105
- Bonding in NO
- Slide 107
- Slide 108
- Slide 109
- Key Points of MO Theory ndash Heteronuclear Molecules
- Slide 111
- Slide 112
- Slide 113
- MO Theory Expectations
-
A ComplicationhellipMO Diagram for O2
(similar for F2 and Ne2)MO Diagram for B2
(similar for C2 and N2)
s-p mixing
No s-p mixing
Electron repulsion
Why does s-p mixing occur
s2s and s2p both have significant e- probability between the nuclei so e- in s2s
will repel e- in s2pEffect will decrease as you move across the Periodic Table increased nuclear charge pulls the s2s e-
closer making the s2s orbital smaller and decreasing the s2s and s2p interaction
Molecular Orbitals of X2 Molecules
sp orbital mixing (a little hybridization)bull lowers the energy of the s2s orbitals andbull raises the energy of the s2p orbitals
bull As a result E(s2p) gt E(p 2p) for B2 C2 and N2 bull As one moves right in Row 2 2s and 2p get
further apart in energy decreasing sndashp mixing E(s2p) lt E(p2p) for O2 F2 and Ne2 See text pages 680-681
bull Note that sndashp mixing does not affect bond order or magnetism in the common diatomics (N2 O2 and F2) Hence it is not of much practical importance
s-p mixing
No s-p mixing
B C and N all have 12 filled 2p orbitalsWhen does s-p mixing occur
O F and Xe all have gt 12 filled 2p orbitals
bull If 2 electrons are forced to be in the same orbital their energies go up
bull Electrons repel each other because they are negatively charged
bull Having gt 12 filled 2p orbitals raises the energies of these orbitals due to e- - e- repulsion
s-p mixing only occurs when the s and p atomic orbitals are close in energy ( 12 filled 2p orbitals)
Relating the MO Diagrams to Physical Properties
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11
Removing an electron from N2 decreases the bond energy of the resulting ion whereas removing an electron from O2 increases the bond energy of the resulting ion Explain these facts using MO diagrams
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11
Plan We first draw the MO energy levels for the four species recalling that they differ for N2 and O2 Then we determine the bond orders and compare them with the data bond order is related directly to bond energy and inversely to bond length
Sample Problem - ContinuedSolution The MO energy levels are
N2
sp
p2p
N2+
s2p
p2p
s2s
s2s
O2 O2+
sp
s2s
s2s
p2p
p2p
s2p
Bond Orders(8-2)2 = 3 (7-2)2 = 25 (8-4)2 = 2 (8-3)2 = 25
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11Bond Order 3 25 2 25
What have we learned so far1 Molecular orbitals (MO) explain the properties of valence electrons in molecules (Example O2)
2 s and p atomic orbitals can be mixed to form s s p and p molecular orbitals
3 Electrons in p or p molecular orbitals can have the same energies Degenerate orbitals
4 The ordering of s2p and p2p molecular orbitals depends on the electron occupancy s-p mixing
Bonding in Diatomic Molecules
Covalent
Ionic
IonicCovalent
HomonuclearH2
Nonpolar covalent bond(450 kJmol bond)
HeteronuclearHF
Polar covalent bond(565 kJmol bond)
Electronegativity
Electronegativity
Because F (EN = 40) is more electronegative than H (EN = 22) the electrons move closer to F
This gives rise to a polar bond
H F
Figure 1445
Electrons are not equally sharedin heteronuclear bondsHF
s Antibonding (s)Mostly H(1s)
s BondingMostly F(2p)
H F
H F
MOs of a Polar Covalent Bond HF
This approach simplifies model and only considers electrons involved in bond
MOs OF XY MOLECULES
Equal or unequal e sharing between 2 atoms is reflected in the composition of the MOsmiddot When 2 atoms X and Y have the same electronegativity (purely
covalent bond) their overlapping AOs have the same energy and the bonding and anti-bonding MOs are each half X and half Y AO All electrons spend equal time near X and Y Examples N2 O2 F2
middot If EN(Y) gt EN(X) (polar covalent X+Y) the Y AO has lower energy than the X AO The bonding MO is more like the Y AO and the anti-bonding MO more like the X AO Bonding e spend more time near Y than X vice versa for anti-bonding e Example CO
MOs OF XY MOLECULES ____ s ___ ___ p ___ ___ ___2p ___ ___ ___ 2p uarr ____ s
Energy ___ ___ p ____ s ___ 2s ___ 2s ____ s
middot CO Bond Order = 30 (same as N2)middot CO Bond Energy = 1076 kJmol (N2 = 945 kJmol)middot Isoelectronic to CO and N2 CNndash NO+middot NO has 1endash in p bond order = 25 this endash is more on N than O
NO NO+ easyhellip
C Atom (4endash) Cδ+Oδndash (10endash) O Atom (6endash)
___ ___ ___ 2p ___ 2s
Electronegativity
Bonding in NObull Two possible Lewis dot structures
for NObull The simplest structure minimizes
formal charges and places the lone (unpaired) electron on the nitrogen
bull The Lewis structure predicts a bond order of 2 but experimental evidence suggests a bond order between 2 and 3
bull How does MO theory help us understand bonding in NO
N=O
N=O
+1-1
When the electronegativities of the 2 atoms are more similar the bonding
becomes less polar
EN(N) = 30EN(O) = 34
N=O
Electronegativity
2s2s
2p2p
NON OBond order = 25 unpaired electron is in a N-like orbital
NO+
oxidation
NO
NO is easily oxidized to form NO+ Why What changes can we predict in the bonding and magnetism of the molecule
Bond Order = (8-3)2 = 25Paramagnetic
Bond Order = (8-2)2 = 3Diamagnetic
s2s
s2s
p2p p2p
p2p
-3320
-1835
-1444 -1374
s2p -1307
-597 p2p (empty)
MO diagram for NO
Key Points of MO Theory ndash Heteronuclear Molecules
bull The more electronegative atom has orbitals lower in energy than the more positive atom
bull Electrons in bonding orbitals are closer to the more electronegative atom anti-bonding electrons are closer to the more positive atom
bull For most diatomic molecules s-p mixing changes the orbital energy levels but since these orbitals are almost always fully occupied their order is less important to us
Combining the Localized Electron and Molecular Orbital Models (into a convenient working
model)
Figure 1447
Only the p bonding changes between these resonance structures - The MO model describes this p bonding more
effectively
Figure 1451
Atomic Orbitals Molecular Orbitals
Another example Benzene
p bonding
s bonding
p atomic orbitals p molecular orbital
MO Theory Expectationsbull You should be able to
ndash predict which atomic orbitals are higher or lower in energy (based on electronegativity differences)
ndash correctly fill a molecular orbital diagramndash correctly calculate bond orderndash predict molecular magnetic properties
based on orbital occupationndash understand how molecular properties
change upon ionization (oxidation or reduction) of molecules
- What Do Molecules Look Like
- Slide 2
- Slide 3
- Arranging Electron Pairs
- Arranging Electron Pairs (cont)
- VSEPR Structure Guidelines
- VSEPR 2 electron pairs
- VSEPR 3 electron pairs
- VSEPR 4 electron pairs (cont)
- Bonding vs Lone pairs
- VSEPR 4 electron pairs
- VSEPR 4 electron pairs (cont) (2)
- VSEPR 4 electron pairs (cont) (3)
- Slide 14
- Molecular vs Electron-Pair Geometry
- Slide 16
- VSEPR Beyond the Octet
- VSEPR 5 electron pairs
- VSEPR 5 electron pairs (2)
- 5-electron-pair geometries
- VSEPR 6 electron pairs
- Slide 22
- Molecular Dipole Moments
- Slide 24
- Molecular Dipoles (cont)
- Molecular Dipole Example
- Advanced VSEPR Application
- The VSEPR Table
- The VSEPR Table (2)
- Slide 30
- Slide 31
- Slide 32
- Slide 33
- Slide 34
- Slide 35
- Slide 36
- Slide 37
- Slide 38
- Slide 39
- Slide 40
- Slide 41
- Slide 42
- Slide 43
- Slide 44
- Slide 45
- Slide 46
- Slide 47
- Slide 48
- Hybridization
- Hybridization (2)
- Hybridization (3)
- General Notes
- What is important to know about hybridization
- What is important to know about hybridization (2)
- What is important to know about hybridization (3)
- Slide 56
- Slide 57
- Molecular Orbital Theory
- Slide 59
- Constructive and Destructive Interference
- Slide 61
- MO Diagrams
- Slide 63
- Slide 64
- Review of Orbital Filling
- Slide 66
- Slide 67
- Filling Molecular Orbitals with Electrons
- Slide 69
- Slide 70
- Slide 71
- Forming Bonds
- Anti-bonding Orbitals
- MO Theory in Bonding
- Slide 75
- Slide 76
- Slide 77
- Slide 78
- Slide 79
- Slide 80
- Slide 81
- Slide 82
- Slide 83
- Slide 84
- Slide 85
- MO Diagram for H2 vs N2
- Slide 87
- Slide 88
- Slide 89
- Slide 90
- Molecular Orbitals of X2 Molecules
- Slide 92
- Slide 93
- Slide 94
- Slide 95
- Slide 96
- Slide 97
- Slide 98
- Slide 99
- Slide 100
- Slide 101
- Slide 102
- Slide 103
- Slide 104
- Slide 105
- Bonding in NO
- Slide 107
- Slide 108
- Slide 109
- Key Points of MO Theory ndash Heteronuclear Molecules
- Slide 111
- Slide 112
- Slide 113
- MO Theory Expectations
-
Electron repulsion
Why does s-p mixing occur
s2s and s2p both have significant e- probability between the nuclei so e- in s2s
will repel e- in s2pEffect will decrease as you move across the Periodic Table increased nuclear charge pulls the s2s e-
closer making the s2s orbital smaller and decreasing the s2s and s2p interaction
Molecular Orbitals of X2 Molecules
sp orbital mixing (a little hybridization)bull lowers the energy of the s2s orbitals andbull raises the energy of the s2p orbitals
bull As a result E(s2p) gt E(p 2p) for B2 C2 and N2 bull As one moves right in Row 2 2s and 2p get
further apart in energy decreasing sndashp mixing E(s2p) lt E(p2p) for O2 F2 and Ne2 See text pages 680-681
bull Note that sndashp mixing does not affect bond order or magnetism in the common diatomics (N2 O2 and F2) Hence it is not of much practical importance
s-p mixing
No s-p mixing
B C and N all have 12 filled 2p orbitalsWhen does s-p mixing occur
O F and Xe all have gt 12 filled 2p orbitals
bull If 2 electrons are forced to be in the same orbital their energies go up
bull Electrons repel each other because they are negatively charged
bull Having gt 12 filled 2p orbitals raises the energies of these orbitals due to e- - e- repulsion
s-p mixing only occurs when the s and p atomic orbitals are close in energy ( 12 filled 2p orbitals)
Relating the MO Diagrams to Physical Properties
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11
Removing an electron from N2 decreases the bond energy of the resulting ion whereas removing an electron from O2 increases the bond energy of the resulting ion Explain these facts using MO diagrams
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11
Plan We first draw the MO energy levels for the four species recalling that they differ for N2 and O2 Then we determine the bond orders and compare them with the data bond order is related directly to bond energy and inversely to bond length
Sample Problem - ContinuedSolution The MO energy levels are
N2
sp
p2p
N2+
s2p
p2p
s2s
s2s
O2 O2+
sp
s2s
s2s
p2p
p2p
s2p
Bond Orders(8-2)2 = 3 (7-2)2 = 25 (8-4)2 = 2 (8-3)2 = 25
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11Bond Order 3 25 2 25
What have we learned so far1 Molecular orbitals (MO) explain the properties of valence electrons in molecules (Example O2)
2 s and p atomic orbitals can be mixed to form s s p and p molecular orbitals
3 Electrons in p or p molecular orbitals can have the same energies Degenerate orbitals
4 The ordering of s2p and p2p molecular orbitals depends on the electron occupancy s-p mixing
Bonding in Diatomic Molecules
Covalent
Ionic
IonicCovalent
HomonuclearH2
Nonpolar covalent bond(450 kJmol bond)
HeteronuclearHF
Polar covalent bond(565 kJmol bond)
Electronegativity
Electronegativity
Because F (EN = 40) is more electronegative than H (EN = 22) the electrons move closer to F
This gives rise to a polar bond
H F
Figure 1445
Electrons are not equally sharedin heteronuclear bondsHF
s Antibonding (s)Mostly H(1s)
s BondingMostly F(2p)
H F
H F
MOs of a Polar Covalent Bond HF
This approach simplifies model and only considers electrons involved in bond
MOs OF XY MOLECULES
Equal or unequal e sharing between 2 atoms is reflected in the composition of the MOsmiddot When 2 atoms X and Y have the same electronegativity (purely
covalent bond) their overlapping AOs have the same energy and the bonding and anti-bonding MOs are each half X and half Y AO All electrons spend equal time near X and Y Examples N2 O2 F2
middot If EN(Y) gt EN(X) (polar covalent X+Y) the Y AO has lower energy than the X AO The bonding MO is more like the Y AO and the anti-bonding MO more like the X AO Bonding e spend more time near Y than X vice versa for anti-bonding e Example CO
MOs OF XY MOLECULES ____ s ___ ___ p ___ ___ ___2p ___ ___ ___ 2p uarr ____ s
Energy ___ ___ p ____ s ___ 2s ___ 2s ____ s
middot CO Bond Order = 30 (same as N2)middot CO Bond Energy = 1076 kJmol (N2 = 945 kJmol)middot Isoelectronic to CO and N2 CNndash NO+middot NO has 1endash in p bond order = 25 this endash is more on N than O
NO NO+ easyhellip
C Atom (4endash) Cδ+Oδndash (10endash) O Atom (6endash)
___ ___ ___ 2p ___ 2s
Electronegativity
Bonding in NObull Two possible Lewis dot structures
for NObull The simplest structure minimizes
formal charges and places the lone (unpaired) electron on the nitrogen
bull The Lewis structure predicts a bond order of 2 but experimental evidence suggests a bond order between 2 and 3
bull How does MO theory help us understand bonding in NO
N=O
N=O
+1-1
When the electronegativities of the 2 atoms are more similar the bonding
becomes less polar
EN(N) = 30EN(O) = 34
N=O
Electronegativity
2s2s
2p2p
NON OBond order = 25 unpaired electron is in a N-like orbital
NO+
oxidation
NO
NO is easily oxidized to form NO+ Why What changes can we predict in the bonding and magnetism of the molecule
Bond Order = (8-3)2 = 25Paramagnetic
Bond Order = (8-2)2 = 3Diamagnetic
s2s
s2s
p2p p2p
p2p
-3320
-1835
-1444 -1374
s2p -1307
-597 p2p (empty)
MO diagram for NO
Key Points of MO Theory ndash Heteronuclear Molecules
bull The more electronegative atom has orbitals lower in energy than the more positive atom
bull Electrons in bonding orbitals are closer to the more electronegative atom anti-bonding electrons are closer to the more positive atom
bull For most diatomic molecules s-p mixing changes the orbital energy levels but since these orbitals are almost always fully occupied their order is less important to us
Combining the Localized Electron and Molecular Orbital Models (into a convenient working
model)
Figure 1447
Only the p bonding changes between these resonance structures - The MO model describes this p bonding more
effectively
Figure 1451
Atomic Orbitals Molecular Orbitals
Another example Benzene
p bonding
s bonding
p atomic orbitals p molecular orbital
MO Theory Expectationsbull You should be able to
ndash predict which atomic orbitals are higher or lower in energy (based on electronegativity differences)
ndash correctly fill a molecular orbital diagramndash correctly calculate bond orderndash predict molecular magnetic properties
based on orbital occupationndash understand how molecular properties
change upon ionization (oxidation or reduction) of molecules
- What Do Molecules Look Like
- Slide 2
- Slide 3
- Arranging Electron Pairs
- Arranging Electron Pairs (cont)
- VSEPR Structure Guidelines
- VSEPR 2 electron pairs
- VSEPR 3 electron pairs
- VSEPR 4 electron pairs (cont)
- Bonding vs Lone pairs
- VSEPR 4 electron pairs
- VSEPR 4 electron pairs (cont) (2)
- VSEPR 4 electron pairs (cont) (3)
- Slide 14
- Molecular vs Electron-Pair Geometry
- Slide 16
- VSEPR Beyond the Octet
- VSEPR 5 electron pairs
- VSEPR 5 electron pairs (2)
- 5-electron-pair geometries
- VSEPR 6 electron pairs
- Slide 22
- Molecular Dipole Moments
- Slide 24
- Molecular Dipoles (cont)
- Molecular Dipole Example
- Advanced VSEPR Application
- The VSEPR Table
- The VSEPR Table (2)
- Slide 30
- Slide 31
- Slide 32
- Slide 33
- Slide 34
- Slide 35
- Slide 36
- Slide 37
- Slide 38
- Slide 39
- Slide 40
- Slide 41
- Slide 42
- Slide 43
- Slide 44
- Slide 45
- Slide 46
- Slide 47
- Slide 48
- Hybridization
- Hybridization (2)
- Hybridization (3)
- General Notes
- What is important to know about hybridization
- What is important to know about hybridization (2)
- What is important to know about hybridization (3)
- Slide 56
- Slide 57
- Molecular Orbital Theory
- Slide 59
- Constructive and Destructive Interference
- Slide 61
- MO Diagrams
- Slide 63
- Slide 64
- Review of Orbital Filling
- Slide 66
- Slide 67
- Filling Molecular Orbitals with Electrons
- Slide 69
- Slide 70
- Slide 71
- Forming Bonds
- Anti-bonding Orbitals
- MO Theory in Bonding
- Slide 75
- Slide 76
- Slide 77
- Slide 78
- Slide 79
- Slide 80
- Slide 81
- Slide 82
- Slide 83
- Slide 84
- Slide 85
- MO Diagram for H2 vs N2
- Slide 87
- Slide 88
- Slide 89
- Slide 90
- Molecular Orbitals of X2 Molecules
- Slide 92
- Slide 93
- Slide 94
- Slide 95
- Slide 96
- Slide 97
- Slide 98
- Slide 99
- Slide 100
- Slide 101
- Slide 102
- Slide 103
- Slide 104
- Slide 105
- Bonding in NO
- Slide 107
- Slide 108
- Slide 109
- Key Points of MO Theory ndash Heteronuclear Molecules
- Slide 111
- Slide 112
- Slide 113
- MO Theory Expectations
-
Molecular Orbitals of X2 Molecules
sp orbital mixing (a little hybridization)bull lowers the energy of the s2s orbitals andbull raises the energy of the s2p orbitals
bull As a result E(s2p) gt E(p 2p) for B2 C2 and N2 bull As one moves right in Row 2 2s and 2p get
further apart in energy decreasing sndashp mixing E(s2p) lt E(p2p) for O2 F2 and Ne2 See text pages 680-681
bull Note that sndashp mixing does not affect bond order or magnetism in the common diatomics (N2 O2 and F2) Hence it is not of much practical importance
s-p mixing
No s-p mixing
B C and N all have 12 filled 2p orbitalsWhen does s-p mixing occur
O F and Xe all have gt 12 filled 2p orbitals
bull If 2 electrons are forced to be in the same orbital their energies go up
bull Electrons repel each other because they are negatively charged
bull Having gt 12 filled 2p orbitals raises the energies of these orbitals due to e- - e- repulsion
s-p mixing only occurs when the s and p atomic orbitals are close in energy ( 12 filled 2p orbitals)
Relating the MO Diagrams to Physical Properties
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11
Removing an electron from N2 decreases the bond energy of the resulting ion whereas removing an electron from O2 increases the bond energy of the resulting ion Explain these facts using MO diagrams
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11
Plan We first draw the MO energy levels for the four species recalling that they differ for N2 and O2 Then we determine the bond orders and compare them with the data bond order is related directly to bond energy and inversely to bond length
Sample Problem - ContinuedSolution The MO energy levels are
N2
sp
p2p
N2+
s2p
p2p
s2s
s2s
O2 O2+
sp
s2s
s2s
p2p
p2p
s2p
Bond Orders(8-2)2 = 3 (7-2)2 = 25 (8-4)2 = 2 (8-3)2 = 25
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11Bond Order 3 25 2 25
What have we learned so far1 Molecular orbitals (MO) explain the properties of valence electrons in molecules (Example O2)
2 s and p atomic orbitals can be mixed to form s s p and p molecular orbitals
3 Electrons in p or p molecular orbitals can have the same energies Degenerate orbitals
4 The ordering of s2p and p2p molecular orbitals depends on the electron occupancy s-p mixing
Bonding in Diatomic Molecules
Covalent
Ionic
IonicCovalent
HomonuclearH2
Nonpolar covalent bond(450 kJmol bond)
HeteronuclearHF
Polar covalent bond(565 kJmol bond)
Electronegativity
Electronegativity
Because F (EN = 40) is more electronegative than H (EN = 22) the electrons move closer to F
This gives rise to a polar bond
H F
Figure 1445
Electrons are not equally sharedin heteronuclear bondsHF
s Antibonding (s)Mostly H(1s)
s BondingMostly F(2p)
H F
H F
MOs of a Polar Covalent Bond HF
This approach simplifies model and only considers electrons involved in bond
MOs OF XY MOLECULES
Equal or unequal e sharing between 2 atoms is reflected in the composition of the MOsmiddot When 2 atoms X and Y have the same electronegativity (purely
covalent bond) their overlapping AOs have the same energy and the bonding and anti-bonding MOs are each half X and half Y AO All electrons spend equal time near X and Y Examples N2 O2 F2
middot If EN(Y) gt EN(X) (polar covalent X+Y) the Y AO has lower energy than the X AO The bonding MO is more like the Y AO and the anti-bonding MO more like the X AO Bonding e spend more time near Y than X vice versa for anti-bonding e Example CO
MOs OF XY MOLECULES ____ s ___ ___ p ___ ___ ___2p ___ ___ ___ 2p uarr ____ s
Energy ___ ___ p ____ s ___ 2s ___ 2s ____ s
middot CO Bond Order = 30 (same as N2)middot CO Bond Energy = 1076 kJmol (N2 = 945 kJmol)middot Isoelectronic to CO and N2 CNndash NO+middot NO has 1endash in p bond order = 25 this endash is more on N than O
NO NO+ easyhellip
C Atom (4endash) Cδ+Oδndash (10endash) O Atom (6endash)
___ ___ ___ 2p ___ 2s
Electronegativity
Bonding in NObull Two possible Lewis dot structures
for NObull The simplest structure minimizes
formal charges and places the lone (unpaired) electron on the nitrogen
bull The Lewis structure predicts a bond order of 2 but experimental evidence suggests a bond order between 2 and 3
bull How does MO theory help us understand bonding in NO
N=O
N=O
+1-1
When the electronegativities of the 2 atoms are more similar the bonding
becomes less polar
EN(N) = 30EN(O) = 34
N=O
Electronegativity
2s2s
2p2p
NON OBond order = 25 unpaired electron is in a N-like orbital
NO+
oxidation
NO
NO is easily oxidized to form NO+ Why What changes can we predict in the bonding and magnetism of the molecule
Bond Order = (8-3)2 = 25Paramagnetic
Bond Order = (8-2)2 = 3Diamagnetic
s2s
s2s
p2p p2p
p2p
-3320
-1835
-1444 -1374
s2p -1307
-597 p2p (empty)
MO diagram for NO
Key Points of MO Theory ndash Heteronuclear Molecules
bull The more electronegative atom has orbitals lower in energy than the more positive atom
bull Electrons in bonding orbitals are closer to the more electronegative atom anti-bonding electrons are closer to the more positive atom
bull For most diatomic molecules s-p mixing changes the orbital energy levels but since these orbitals are almost always fully occupied their order is less important to us
Combining the Localized Electron and Molecular Orbital Models (into a convenient working
model)
Figure 1447
Only the p bonding changes between these resonance structures - The MO model describes this p bonding more
effectively
Figure 1451
Atomic Orbitals Molecular Orbitals
Another example Benzene
p bonding
s bonding
p atomic orbitals p molecular orbital
MO Theory Expectationsbull You should be able to
ndash predict which atomic orbitals are higher or lower in energy (based on electronegativity differences)
ndash correctly fill a molecular orbital diagramndash correctly calculate bond orderndash predict molecular magnetic properties
based on orbital occupationndash understand how molecular properties
change upon ionization (oxidation or reduction) of molecules
- What Do Molecules Look Like
- Slide 2
- Slide 3
- Arranging Electron Pairs
- Arranging Electron Pairs (cont)
- VSEPR Structure Guidelines
- VSEPR 2 electron pairs
- VSEPR 3 electron pairs
- VSEPR 4 electron pairs (cont)
- Bonding vs Lone pairs
- VSEPR 4 electron pairs
- VSEPR 4 electron pairs (cont) (2)
- VSEPR 4 electron pairs (cont) (3)
- Slide 14
- Molecular vs Electron-Pair Geometry
- Slide 16
- VSEPR Beyond the Octet
- VSEPR 5 electron pairs
- VSEPR 5 electron pairs (2)
- 5-electron-pair geometries
- VSEPR 6 electron pairs
- Slide 22
- Molecular Dipole Moments
- Slide 24
- Molecular Dipoles (cont)
- Molecular Dipole Example
- Advanced VSEPR Application
- The VSEPR Table
- The VSEPR Table (2)
- Slide 30
- Slide 31
- Slide 32
- Slide 33
- Slide 34
- Slide 35
- Slide 36
- Slide 37
- Slide 38
- Slide 39
- Slide 40
- Slide 41
- Slide 42
- Slide 43
- Slide 44
- Slide 45
- Slide 46
- Slide 47
- Slide 48
- Hybridization
- Hybridization (2)
- Hybridization (3)
- General Notes
- What is important to know about hybridization
- What is important to know about hybridization (2)
- What is important to know about hybridization (3)
- Slide 56
- Slide 57
- Molecular Orbital Theory
- Slide 59
- Constructive and Destructive Interference
- Slide 61
- MO Diagrams
- Slide 63
- Slide 64
- Review of Orbital Filling
- Slide 66
- Slide 67
- Filling Molecular Orbitals with Electrons
- Slide 69
- Slide 70
- Slide 71
- Forming Bonds
- Anti-bonding Orbitals
- MO Theory in Bonding
- Slide 75
- Slide 76
- Slide 77
- Slide 78
- Slide 79
- Slide 80
- Slide 81
- Slide 82
- Slide 83
- Slide 84
- Slide 85
- MO Diagram for H2 vs N2
- Slide 87
- Slide 88
- Slide 89
- Slide 90
- Molecular Orbitals of X2 Molecules
- Slide 92
- Slide 93
- Slide 94
- Slide 95
- Slide 96
- Slide 97
- Slide 98
- Slide 99
- Slide 100
- Slide 101
- Slide 102
- Slide 103
- Slide 104
- Slide 105
- Bonding in NO
- Slide 107
- Slide 108
- Slide 109
- Key Points of MO Theory ndash Heteronuclear Molecules
- Slide 111
- Slide 112
- Slide 113
- MO Theory Expectations
-
s-p mixing
No s-p mixing
B C and N all have 12 filled 2p orbitalsWhen does s-p mixing occur
O F and Xe all have gt 12 filled 2p orbitals
bull If 2 electrons are forced to be in the same orbital their energies go up
bull Electrons repel each other because they are negatively charged
bull Having gt 12 filled 2p orbitals raises the energies of these orbitals due to e- - e- repulsion
s-p mixing only occurs when the s and p atomic orbitals are close in energy ( 12 filled 2p orbitals)
Relating the MO Diagrams to Physical Properties
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11
Removing an electron from N2 decreases the bond energy of the resulting ion whereas removing an electron from O2 increases the bond energy of the resulting ion Explain these facts using MO diagrams
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11
Plan We first draw the MO energy levels for the four species recalling that they differ for N2 and O2 Then we determine the bond orders and compare them with the data bond order is related directly to bond energy and inversely to bond length
Sample Problem - ContinuedSolution The MO energy levels are
N2
sp
p2p
N2+
s2p
p2p
s2s
s2s
O2 O2+
sp
s2s
s2s
p2p
p2p
s2p
Bond Orders(8-2)2 = 3 (7-2)2 = 25 (8-4)2 = 2 (8-3)2 = 25
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11Bond Order 3 25 2 25
What have we learned so far1 Molecular orbitals (MO) explain the properties of valence electrons in molecules (Example O2)
2 s and p atomic orbitals can be mixed to form s s p and p molecular orbitals
3 Electrons in p or p molecular orbitals can have the same energies Degenerate orbitals
4 The ordering of s2p and p2p molecular orbitals depends on the electron occupancy s-p mixing
Bonding in Diatomic Molecules
Covalent
Ionic
IonicCovalent
HomonuclearH2
Nonpolar covalent bond(450 kJmol bond)
HeteronuclearHF
Polar covalent bond(565 kJmol bond)
Electronegativity
Electronegativity
Because F (EN = 40) is more electronegative than H (EN = 22) the electrons move closer to F
This gives rise to a polar bond
H F
Figure 1445
Electrons are not equally sharedin heteronuclear bondsHF
s Antibonding (s)Mostly H(1s)
s BondingMostly F(2p)
H F
H F
MOs of a Polar Covalent Bond HF
This approach simplifies model and only considers electrons involved in bond
MOs OF XY MOLECULES
Equal or unequal e sharing between 2 atoms is reflected in the composition of the MOsmiddot When 2 atoms X and Y have the same electronegativity (purely
covalent bond) their overlapping AOs have the same energy and the bonding and anti-bonding MOs are each half X and half Y AO All electrons spend equal time near X and Y Examples N2 O2 F2
middot If EN(Y) gt EN(X) (polar covalent X+Y) the Y AO has lower energy than the X AO The bonding MO is more like the Y AO and the anti-bonding MO more like the X AO Bonding e spend more time near Y than X vice versa for anti-bonding e Example CO
MOs OF XY MOLECULES ____ s ___ ___ p ___ ___ ___2p ___ ___ ___ 2p uarr ____ s
Energy ___ ___ p ____ s ___ 2s ___ 2s ____ s
middot CO Bond Order = 30 (same as N2)middot CO Bond Energy = 1076 kJmol (N2 = 945 kJmol)middot Isoelectronic to CO and N2 CNndash NO+middot NO has 1endash in p bond order = 25 this endash is more on N than O
NO NO+ easyhellip
C Atom (4endash) Cδ+Oδndash (10endash) O Atom (6endash)
___ ___ ___ 2p ___ 2s
Electronegativity
Bonding in NObull Two possible Lewis dot structures
for NObull The simplest structure minimizes
formal charges and places the lone (unpaired) electron on the nitrogen
bull The Lewis structure predicts a bond order of 2 but experimental evidence suggests a bond order between 2 and 3
bull How does MO theory help us understand bonding in NO
N=O
N=O
+1-1
When the electronegativities of the 2 atoms are more similar the bonding
becomes less polar
EN(N) = 30EN(O) = 34
N=O
Electronegativity
2s2s
2p2p
NON OBond order = 25 unpaired electron is in a N-like orbital
NO+
oxidation
NO
NO is easily oxidized to form NO+ Why What changes can we predict in the bonding and magnetism of the molecule
Bond Order = (8-3)2 = 25Paramagnetic
Bond Order = (8-2)2 = 3Diamagnetic
s2s
s2s
p2p p2p
p2p
-3320
-1835
-1444 -1374
s2p -1307
-597 p2p (empty)
MO diagram for NO
Key Points of MO Theory ndash Heteronuclear Molecules
bull The more electronegative atom has orbitals lower in energy than the more positive atom
bull Electrons in bonding orbitals are closer to the more electronegative atom anti-bonding electrons are closer to the more positive atom
bull For most diatomic molecules s-p mixing changes the orbital energy levels but since these orbitals are almost always fully occupied their order is less important to us
Combining the Localized Electron and Molecular Orbital Models (into a convenient working
model)
Figure 1447
Only the p bonding changes between these resonance structures - The MO model describes this p bonding more
effectively
Figure 1451
Atomic Orbitals Molecular Orbitals
Another example Benzene
p bonding
s bonding
p atomic orbitals p molecular orbital
MO Theory Expectationsbull You should be able to
ndash predict which atomic orbitals are higher or lower in energy (based on electronegativity differences)
ndash correctly fill a molecular orbital diagramndash correctly calculate bond orderndash predict molecular magnetic properties
based on orbital occupationndash understand how molecular properties
change upon ionization (oxidation or reduction) of molecules
- What Do Molecules Look Like
- Slide 2
- Slide 3
- Arranging Electron Pairs
- Arranging Electron Pairs (cont)
- VSEPR Structure Guidelines
- VSEPR 2 electron pairs
- VSEPR 3 electron pairs
- VSEPR 4 electron pairs (cont)
- Bonding vs Lone pairs
- VSEPR 4 electron pairs
- VSEPR 4 electron pairs (cont) (2)
- VSEPR 4 electron pairs (cont) (3)
- Slide 14
- Molecular vs Electron-Pair Geometry
- Slide 16
- VSEPR Beyond the Octet
- VSEPR 5 electron pairs
- VSEPR 5 electron pairs (2)
- 5-electron-pair geometries
- VSEPR 6 electron pairs
- Slide 22
- Molecular Dipole Moments
- Slide 24
- Molecular Dipoles (cont)
- Molecular Dipole Example
- Advanced VSEPR Application
- The VSEPR Table
- The VSEPR Table (2)
- Slide 30
- Slide 31
- Slide 32
- Slide 33
- Slide 34
- Slide 35
- Slide 36
- Slide 37
- Slide 38
- Slide 39
- Slide 40
- Slide 41
- Slide 42
- Slide 43
- Slide 44
- Slide 45
- Slide 46
- Slide 47
- Slide 48
- Hybridization
- Hybridization (2)
- Hybridization (3)
- General Notes
- What is important to know about hybridization
- What is important to know about hybridization (2)
- What is important to know about hybridization (3)
- Slide 56
- Slide 57
- Molecular Orbital Theory
- Slide 59
- Constructive and Destructive Interference
- Slide 61
- MO Diagrams
- Slide 63
- Slide 64
- Review of Orbital Filling
- Slide 66
- Slide 67
- Filling Molecular Orbitals with Electrons
- Slide 69
- Slide 70
- Slide 71
- Forming Bonds
- Anti-bonding Orbitals
- MO Theory in Bonding
- Slide 75
- Slide 76
- Slide 77
- Slide 78
- Slide 79
- Slide 80
- Slide 81
- Slide 82
- Slide 83
- Slide 84
- Slide 85
- MO Diagram for H2 vs N2
- Slide 87
- Slide 88
- Slide 89
- Slide 90
- Molecular Orbitals of X2 Molecules
- Slide 92
- Slide 93
- Slide 94
- Slide 95
- Slide 96
- Slide 97
- Slide 98
- Slide 99
- Slide 100
- Slide 101
- Slide 102
- Slide 103
- Slide 104
- Slide 105
- Bonding in NO
- Slide 107
- Slide 108
- Slide 109
- Key Points of MO Theory ndash Heteronuclear Molecules
- Slide 111
- Slide 112
- Slide 113
- MO Theory Expectations
-
B C and N all have 12 filled 2p orbitalsWhen does s-p mixing occur
O F and Xe all have gt 12 filled 2p orbitals
bull If 2 electrons are forced to be in the same orbital their energies go up
bull Electrons repel each other because they are negatively charged
bull Having gt 12 filled 2p orbitals raises the energies of these orbitals due to e- - e- repulsion
s-p mixing only occurs when the s and p atomic orbitals are close in energy ( 12 filled 2p orbitals)
Relating the MO Diagrams to Physical Properties
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11
Removing an electron from N2 decreases the bond energy of the resulting ion whereas removing an electron from O2 increases the bond energy of the resulting ion Explain these facts using MO diagrams
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11
Plan We first draw the MO energy levels for the four species recalling that they differ for N2 and O2 Then we determine the bond orders and compare them with the data bond order is related directly to bond energy and inversely to bond length
Sample Problem - ContinuedSolution The MO energy levels are
N2
sp
p2p
N2+
s2p
p2p
s2s
s2s
O2 O2+
sp
s2s
s2s
p2p
p2p
s2p
Bond Orders(8-2)2 = 3 (7-2)2 = 25 (8-4)2 = 2 (8-3)2 = 25
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11Bond Order 3 25 2 25
What have we learned so far1 Molecular orbitals (MO) explain the properties of valence electrons in molecules (Example O2)
2 s and p atomic orbitals can be mixed to form s s p and p molecular orbitals
3 Electrons in p or p molecular orbitals can have the same energies Degenerate orbitals
4 The ordering of s2p and p2p molecular orbitals depends on the electron occupancy s-p mixing
Bonding in Diatomic Molecules
Covalent
Ionic
IonicCovalent
HomonuclearH2
Nonpolar covalent bond(450 kJmol bond)
HeteronuclearHF
Polar covalent bond(565 kJmol bond)
Electronegativity
Electronegativity
Because F (EN = 40) is more electronegative than H (EN = 22) the electrons move closer to F
This gives rise to a polar bond
H F
Figure 1445
Electrons are not equally sharedin heteronuclear bondsHF
s Antibonding (s)Mostly H(1s)
s BondingMostly F(2p)
H F
H F
MOs of a Polar Covalent Bond HF
This approach simplifies model and only considers electrons involved in bond
MOs OF XY MOLECULES
Equal or unequal e sharing between 2 atoms is reflected in the composition of the MOsmiddot When 2 atoms X and Y have the same electronegativity (purely
covalent bond) their overlapping AOs have the same energy and the bonding and anti-bonding MOs are each half X and half Y AO All electrons spend equal time near X and Y Examples N2 O2 F2
middot If EN(Y) gt EN(X) (polar covalent X+Y) the Y AO has lower energy than the X AO The bonding MO is more like the Y AO and the anti-bonding MO more like the X AO Bonding e spend more time near Y than X vice versa for anti-bonding e Example CO
MOs OF XY MOLECULES ____ s ___ ___ p ___ ___ ___2p ___ ___ ___ 2p uarr ____ s
Energy ___ ___ p ____ s ___ 2s ___ 2s ____ s
middot CO Bond Order = 30 (same as N2)middot CO Bond Energy = 1076 kJmol (N2 = 945 kJmol)middot Isoelectronic to CO and N2 CNndash NO+middot NO has 1endash in p bond order = 25 this endash is more on N than O
NO NO+ easyhellip
C Atom (4endash) Cδ+Oδndash (10endash) O Atom (6endash)
___ ___ ___ 2p ___ 2s
Electronegativity
Bonding in NObull Two possible Lewis dot structures
for NObull The simplest structure minimizes
formal charges and places the lone (unpaired) electron on the nitrogen
bull The Lewis structure predicts a bond order of 2 but experimental evidence suggests a bond order between 2 and 3
bull How does MO theory help us understand bonding in NO
N=O
N=O
+1-1
When the electronegativities of the 2 atoms are more similar the bonding
becomes less polar
EN(N) = 30EN(O) = 34
N=O
Electronegativity
2s2s
2p2p
NON OBond order = 25 unpaired electron is in a N-like orbital
NO+
oxidation
NO
NO is easily oxidized to form NO+ Why What changes can we predict in the bonding and magnetism of the molecule
Bond Order = (8-3)2 = 25Paramagnetic
Bond Order = (8-2)2 = 3Diamagnetic
s2s
s2s
p2p p2p
p2p
-3320
-1835
-1444 -1374
s2p -1307
-597 p2p (empty)
MO diagram for NO
Key Points of MO Theory ndash Heteronuclear Molecules
bull The more electronegative atom has orbitals lower in energy than the more positive atom
bull Electrons in bonding orbitals are closer to the more electronegative atom anti-bonding electrons are closer to the more positive atom
bull For most diatomic molecules s-p mixing changes the orbital energy levels but since these orbitals are almost always fully occupied their order is less important to us
Combining the Localized Electron and Molecular Orbital Models (into a convenient working
model)
Figure 1447
Only the p bonding changes between these resonance structures - The MO model describes this p bonding more
effectively
Figure 1451
Atomic Orbitals Molecular Orbitals
Another example Benzene
p bonding
s bonding
p atomic orbitals p molecular orbital
MO Theory Expectationsbull You should be able to
ndash predict which atomic orbitals are higher or lower in energy (based on electronegativity differences)
ndash correctly fill a molecular orbital diagramndash correctly calculate bond orderndash predict molecular magnetic properties
based on orbital occupationndash understand how molecular properties
change upon ionization (oxidation or reduction) of molecules
- What Do Molecules Look Like
- Slide 2
- Slide 3
- Arranging Electron Pairs
- Arranging Electron Pairs (cont)
- VSEPR Structure Guidelines
- VSEPR 2 electron pairs
- VSEPR 3 electron pairs
- VSEPR 4 electron pairs (cont)
- Bonding vs Lone pairs
- VSEPR 4 electron pairs
- VSEPR 4 electron pairs (cont) (2)
- VSEPR 4 electron pairs (cont) (3)
- Slide 14
- Molecular vs Electron-Pair Geometry
- Slide 16
- VSEPR Beyond the Octet
- VSEPR 5 electron pairs
- VSEPR 5 electron pairs (2)
- 5-electron-pair geometries
- VSEPR 6 electron pairs
- Slide 22
- Molecular Dipole Moments
- Slide 24
- Molecular Dipoles (cont)
- Molecular Dipole Example
- Advanced VSEPR Application
- The VSEPR Table
- The VSEPR Table (2)
- Slide 30
- Slide 31
- Slide 32
- Slide 33
- Slide 34
- Slide 35
- Slide 36
- Slide 37
- Slide 38
- Slide 39
- Slide 40
- Slide 41
- Slide 42
- Slide 43
- Slide 44
- Slide 45
- Slide 46
- Slide 47
- Slide 48
- Hybridization
- Hybridization (2)
- Hybridization (3)
- General Notes
- What is important to know about hybridization
- What is important to know about hybridization (2)
- What is important to know about hybridization (3)
- Slide 56
- Slide 57
- Molecular Orbital Theory
- Slide 59
- Constructive and Destructive Interference
- Slide 61
- MO Diagrams
- Slide 63
- Slide 64
- Review of Orbital Filling
- Slide 66
- Slide 67
- Filling Molecular Orbitals with Electrons
- Slide 69
- Slide 70
- Slide 71
- Forming Bonds
- Anti-bonding Orbitals
- MO Theory in Bonding
- Slide 75
- Slide 76
- Slide 77
- Slide 78
- Slide 79
- Slide 80
- Slide 81
- Slide 82
- Slide 83
- Slide 84
- Slide 85
- MO Diagram for H2 vs N2
- Slide 87
- Slide 88
- Slide 89
- Slide 90
- Molecular Orbitals of X2 Molecules
- Slide 92
- Slide 93
- Slide 94
- Slide 95
- Slide 96
- Slide 97
- Slide 98
- Slide 99
- Slide 100
- Slide 101
- Slide 102
- Slide 103
- Slide 104
- Slide 105
- Bonding in NO
- Slide 107
- Slide 108
- Slide 109
- Key Points of MO Theory ndash Heteronuclear Molecules
- Slide 111
- Slide 112
- Slide 113
- MO Theory Expectations
-
Relating the MO Diagrams to Physical Properties
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11
Removing an electron from N2 decreases the bond energy of the resulting ion whereas removing an electron from O2 increases the bond energy of the resulting ion Explain these facts using MO diagrams
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11
Plan We first draw the MO energy levels for the four species recalling that they differ for N2 and O2 Then we determine the bond orders and compare them with the data bond order is related directly to bond energy and inversely to bond length
Sample Problem - ContinuedSolution The MO energy levels are
N2
sp
p2p
N2+
s2p
p2p
s2s
s2s
O2 O2+
sp
s2s
s2s
p2p
p2p
s2p
Bond Orders(8-2)2 = 3 (7-2)2 = 25 (8-4)2 = 2 (8-3)2 = 25
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11Bond Order 3 25 2 25
What have we learned so far1 Molecular orbitals (MO) explain the properties of valence electrons in molecules (Example O2)
2 s and p atomic orbitals can be mixed to form s s p and p molecular orbitals
3 Electrons in p or p molecular orbitals can have the same energies Degenerate orbitals
4 The ordering of s2p and p2p molecular orbitals depends on the electron occupancy s-p mixing
Bonding in Diatomic Molecules
Covalent
Ionic
IonicCovalent
HomonuclearH2
Nonpolar covalent bond(450 kJmol bond)
HeteronuclearHF
Polar covalent bond(565 kJmol bond)
Electronegativity
Electronegativity
Because F (EN = 40) is more electronegative than H (EN = 22) the electrons move closer to F
This gives rise to a polar bond
H F
Figure 1445
Electrons are not equally sharedin heteronuclear bondsHF
s Antibonding (s)Mostly H(1s)
s BondingMostly F(2p)
H F
H F
MOs of a Polar Covalent Bond HF
This approach simplifies model and only considers electrons involved in bond
MOs OF XY MOLECULES
Equal or unequal e sharing between 2 atoms is reflected in the composition of the MOsmiddot When 2 atoms X and Y have the same electronegativity (purely
covalent bond) their overlapping AOs have the same energy and the bonding and anti-bonding MOs are each half X and half Y AO All electrons spend equal time near X and Y Examples N2 O2 F2
middot If EN(Y) gt EN(X) (polar covalent X+Y) the Y AO has lower energy than the X AO The bonding MO is more like the Y AO and the anti-bonding MO more like the X AO Bonding e spend more time near Y than X vice versa for anti-bonding e Example CO
MOs OF XY MOLECULES ____ s ___ ___ p ___ ___ ___2p ___ ___ ___ 2p uarr ____ s
Energy ___ ___ p ____ s ___ 2s ___ 2s ____ s
middot CO Bond Order = 30 (same as N2)middot CO Bond Energy = 1076 kJmol (N2 = 945 kJmol)middot Isoelectronic to CO and N2 CNndash NO+middot NO has 1endash in p bond order = 25 this endash is more on N than O
NO NO+ easyhellip
C Atom (4endash) Cδ+Oδndash (10endash) O Atom (6endash)
___ ___ ___ 2p ___ 2s
Electronegativity
Bonding in NObull Two possible Lewis dot structures
for NObull The simplest structure minimizes
formal charges and places the lone (unpaired) electron on the nitrogen
bull The Lewis structure predicts a bond order of 2 but experimental evidence suggests a bond order between 2 and 3
bull How does MO theory help us understand bonding in NO
N=O
N=O
+1-1
When the electronegativities of the 2 atoms are more similar the bonding
becomes less polar
EN(N) = 30EN(O) = 34
N=O
Electronegativity
2s2s
2p2p
NON OBond order = 25 unpaired electron is in a N-like orbital
NO+
oxidation
NO
NO is easily oxidized to form NO+ Why What changes can we predict in the bonding and magnetism of the molecule
Bond Order = (8-3)2 = 25Paramagnetic
Bond Order = (8-2)2 = 3Diamagnetic
s2s
s2s
p2p p2p
p2p
-3320
-1835
-1444 -1374
s2p -1307
-597 p2p (empty)
MO diagram for NO
Key Points of MO Theory ndash Heteronuclear Molecules
bull The more electronegative atom has orbitals lower in energy than the more positive atom
bull Electrons in bonding orbitals are closer to the more electronegative atom anti-bonding electrons are closer to the more positive atom
bull For most diatomic molecules s-p mixing changes the orbital energy levels but since these orbitals are almost always fully occupied their order is less important to us
Combining the Localized Electron and Molecular Orbital Models (into a convenient working
model)
Figure 1447
Only the p bonding changes between these resonance structures - The MO model describes this p bonding more
effectively
Figure 1451
Atomic Orbitals Molecular Orbitals
Another example Benzene
p bonding
s bonding
p atomic orbitals p molecular orbital
MO Theory Expectationsbull You should be able to
ndash predict which atomic orbitals are higher or lower in energy (based on electronegativity differences)
ndash correctly fill a molecular orbital diagramndash correctly calculate bond orderndash predict molecular magnetic properties
based on orbital occupationndash understand how molecular properties
change upon ionization (oxidation or reduction) of molecules
- What Do Molecules Look Like
- Slide 2
- Slide 3
- Arranging Electron Pairs
- Arranging Electron Pairs (cont)
- VSEPR Structure Guidelines
- VSEPR 2 electron pairs
- VSEPR 3 electron pairs
- VSEPR 4 electron pairs (cont)
- Bonding vs Lone pairs
- VSEPR 4 electron pairs
- VSEPR 4 electron pairs (cont) (2)
- VSEPR 4 electron pairs (cont) (3)
- Slide 14
- Molecular vs Electron-Pair Geometry
- Slide 16
- VSEPR Beyond the Octet
- VSEPR 5 electron pairs
- VSEPR 5 electron pairs (2)
- 5-electron-pair geometries
- VSEPR 6 electron pairs
- Slide 22
- Molecular Dipole Moments
- Slide 24
- Molecular Dipoles (cont)
- Molecular Dipole Example
- Advanced VSEPR Application
- The VSEPR Table
- The VSEPR Table (2)
- Slide 30
- Slide 31
- Slide 32
- Slide 33
- Slide 34
- Slide 35
- Slide 36
- Slide 37
- Slide 38
- Slide 39
- Slide 40
- Slide 41
- Slide 42
- Slide 43
- Slide 44
- Slide 45
- Slide 46
- Slide 47
- Slide 48
- Hybridization
- Hybridization (2)
- Hybridization (3)
- General Notes
- What is important to know about hybridization
- What is important to know about hybridization (2)
- What is important to know about hybridization (3)
- Slide 56
- Slide 57
- Molecular Orbital Theory
- Slide 59
- Constructive and Destructive Interference
- Slide 61
- MO Diagrams
- Slide 63
- Slide 64
- Review of Orbital Filling
- Slide 66
- Slide 67
- Filling Molecular Orbitals with Electrons
- Slide 69
- Slide 70
- Slide 71
- Forming Bonds
- Anti-bonding Orbitals
- MO Theory in Bonding
- Slide 75
- Slide 76
- Slide 77
- Slide 78
- Slide 79
- Slide 80
- Slide 81
- Slide 82
- Slide 83
- Slide 84
- Slide 85
- MO Diagram for H2 vs N2
- Slide 87
- Slide 88
- Slide 89
- Slide 90
- Molecular Orbitals of X2 Molecules
- Slide 92
- Slide 93
- Slide 94
- Slide 95
- Slide 96
- Slide 97
- Slide 98
- Slide 99
- Slide 100
- Slide 101
- Slide 102
- Slide 103
- Slide 104
- Slide 105
- Bonding in NO
- Slide 107
- Slide 108
- Slide 109
- Key Points of MO Theory ndash Heteronuclear Molecules
- Slide 111
- Slide 112
- Slide 113
- MO Theory Expectations
-
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11
Removing an electron from N2 decreases the bond energy of the resulting ion whereas removing an electron from O2 increases the bond energy of the resulting ion Explain these facts using MO diagrams
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11
Plan We first draw the MO energy levels for the four species recalling that they differ for N2 and O2 Then we determine the bond orders and compare them with the data bond order is related directly to bond energy and inversely to bond length
Sample Problem - ContinuedSolution The MO energy levels are
N2
sp
p2p
N2+
s2p
p2p
s2s
s2s
O2 O2+
sp
s2s
s2s
p2p
p2p
s2p
Bond Orders(8-2)2 = 3 (7-2)2 = 25 (8-4)2 = 2 (8-3)2 = 25
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11Bond Order 3 25 2 25
What have we learned so far1 Molecular orbitals (MO) explain the properties of valence electrons in molecules (Example O2)
2 s and p atomic orbitals can be mixed to form s s p and p molecular orbitals
3 Electrons in p or p molecular orbitals can have the same energies Degenerate orbitals
4 The ordering of s2p and p2p molecular orbitals depends on the electron occupancy s-p mixing
Bonding in Diatomic Molecules
Covalent
Ionic
IonicCovalent
HomonuclearH2
Nonpolar covalent bond(450 kJmol bond)
HeteronuclearHF
Polar covalent bond(565 kJmol bond)
Electronegativity
Electronegativity
Because F (EN = 40) is more electronegative than H (EN = 22) the electrons move closer to F
This gives rise to a polar bond
H F
Figure 1445
Electrons are not equally sharedin heteronuclear bondsHF
s Antibonding (s)Mostly H(1s)
s BondingMostly F(2p)
H F
H F
MOs of a Polar Covalent Bond HF
This approach simplifies model and only considers electrons involved in bond
MOs OF XY MOLECULES
Equal or unequal e sharing between 2 atoms is reflected in the composition of the MOsmiddot When 2 atoms X and Y have the same electronegativity (purely
covalent bond) their overlapping AOs have the same energy and the bonding and anti-bonding MOs are each half X and half Y AO All electrons spend equal time near X and Y Examples N2 O2 F2
middot If EN(Y) gt EN(X) (polar covalent X+Y) the Y AO has lower energy than the X AO The bonding MO is more like the Y AO and the anti-bonding MO more like the X AO Bonding e spend more time near Y than X vice versa for anti-bonding e Example CO
MOs OF XY MOLECULES ____ s ___ ___ p ___ ___ ___2p ___ ___ ___ 2p uarr ____ s
Energy ___ ___ p ____ s ___ 2s ___ 2s ____ s
middot CO Bond Order = 30 (same as N2)middot CO Bond Energy = 1076 kJmol (N2 = 945 kJmol)middot Isoelectronic to CO and N2 CNndash NO+middot NO has 1endash in p bond order = 25 this endash is more on N than O
NO NO+ easyhellip
C Atom (4endash) Cδ+Oδndash (10endash) O Atom (6endash)
___ ___ ___ 2p ___ 2s
Electronegativity
Bonding in NObull Two possible Lewis dot structures
for NObull The simplest structure minimizes
formal charges and places the lone (unpaired) electron on the nitrogen
bull The Lewis structure predicts a bond order of 2 but experimental evidence suggests a bond order between 2 and 3
bull How does MO theory help us understand bonding in NO
N=O
N=O
+1-1
When the electronegativities of the 2 atoms are more similar the bonding
becomes less polar
EN(N) = 30EN(O) = 34
N=O
Electronegativity
2s2s
2p2p
NON OBond order = 25 unpaired electron is in a N-like orbital
NO+
oxidation
NO
NO is easily oxidized to form NO+ Why What changes can we predict in the bonding and magnetism of the molecule
Bond Order = (8-3)2 = 25Paramagnetic
Bond Order = (8-2)2 = 3Diamagnetic
s2s
s2s
p2p p2p
p2p
-3320
-1835
-1444 -1374
s2p -1307
-597 p2p (empty)
MO diagram for NO
Key Points of MO Theory ndash Heteronuclear Molecules
bull The more electronegative atom has orbitals lower in energy than the more positive atom
bull Electrons in bonding orbitals are closer to the more electronegative atom anti-bonding electrons are closer to the more positive atom
bull For most diatomic molecules s-p mixing changes the orbital energy levels but since these orbitals are almost always fully occupied their order is less important to us
Combining the Localized Electron and Molecular Orbital Models (into a convenient working
model)
Figure 1447
Only the p bonding changes between these resonance structures - The MO model describes this p bonding more
effectively
Figure 1451
Atomic Orbitals Molecular Orbitals
Another example Benzene
p bonding
s bonding
p atomic orbitals p molecular orbital
MO Theory Expectationsbull You should be able to
ndash predict which atomic orbitals are higher or lower in energy (based on electronegativity differences)
ndash correctly fill a molecular orbital diagramndash correctly calculate bond orderndash predict molecular magnetic properties
based on orbital occupationndash understand how molecular properties
change upon ionization (oxidation or reduction) of molecules
- What Do Molecules Look Like
- Slide 2
- Slide 3
- Arranging Electron Pairs
- Arranging Electron Pairs (cont)
- VSEPR Structure Guidelines
- VSEPR 2 electron pairs
- VSEPR 3 electron pairs
- VSEPR 4 electron pairs (cont)
- Bonding vs Lone pairs
- VSEPR 4 electron pairs
- VSEPR 4 electron pairs (cont) (2)
- VSEPR 4 electron pairs (cont) (3)
- Slide 14
- Molecular vs Electron-Pair Geometry
- Slide 16
- VSEPR Beyond the Octet
- VSEPR 5 electron pairs
- VSEPR 5 electron pairs (2)
- 5-electron-pair geometries
- VSEPR 6 electron pairs
- Slide 22
- Molecular Dipole Moments
- Slide 24
- Molecular Dipoles (cont)
- Molecular Dipole Example
- Advanced VSEPR Application
- The VSEPR Table
- The VSEPR Table (2)
- Slide 30
- Slide 31
- Slide 32
- Slide 33
- Slide 34
- Slide 35
- Slide 36
- Slide 37
- Slide 38
- Slide 39
- Slide 40
- Slide 41
- Slide 42
- Slide 43
- Slide 44
- Slide 45
- Slide 46
- Slide 47
- Slide 48
- Hybridization
- Hybridization (2)
- Hybridization (3)
- General Notes
- What is important to know about hybridization
- What is important to know about hybridization (2)
- What is important to know about hybridization (3)
- Slide 56
- Slide 57
- Molecular Orbital Theory
- Slide 59
- Constructive and Destructive Interference
- Slide 61
- MO Diagrams
- Slide 63
- Slide 64
- Review of Orbital Filling
- Slide 66
- Slide 67
- Filling Molecular Orbitals with Electrons
- Slide 69
- Slide 70
- Slide 71
- Forming Bonds
- Anti-bonding Orbitals
- MO Theory in Bonding
- Slide 75
- Slide 76
- Slide 77
- Slide 78
- Slide 79
- Slide 80
- Slide 81
- Slide 82
- Slide 83
- Slide 84
- Slide 85
- MO Diagram for H2 vs N2
- Slide 87
- Slide 88
- Slide 89
- Slide 90
- Molecular Orbitals of X2 Molecules
- Slide 92
- Slide 93
- Slide 94
- Slide 95
- Slide 96
- Slide 97
- Slide 98
- Slide 99
- Slide 100
- Slide 101
- Slide 102
- Slide 103
- Slide 104
- Slide 105
- Bonding in NO
- Slide 107
- Slide 108
- Slide 109
- Key Points of MO Theory ndash Heteronuclear Molecules
- Slide 111
- Slide 112
- Slide 113
- MO Theory Expectations
-
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11
Plan We first draw the MO energy levels for the four species recalling that they differ for N2 and O2 Then we determine the bond orders and compare them with the data bond order is related directly to bond energy and inversely to bond length
Sample Problem - ContinuedSolution The MO energy levels are
N2
sp
p2p
N2+
s2p
p2p
s2s
s2s
O2 O2+
sp
s2s
s2s
p2p
p2p
s2p
Bond Orders(8-2)2 = 3 (7-2)2 = 25 (8-4)2 = 2 (8-3)2 = 25
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11Bond Order 3 25 2 25
What have we learned so far1 Molecular orbitals (MO) explain the properties of valence electrons in molecules (Example O2)
2 s and p atomic orbitals can be mixed to form s s p and p molecular orbitals
3 Electrons in p or p molecular orbitals can have the same energies Degenerate orbitals
4 The ordering of s2p and p2p molecular orbitals depends on the electron occupancy s-p mixing
Bonding in Diatomic Molecules
Covalent
Ionic
IonicCovalent
HomonuclearH2
Nonpolar covalent bond(450 kJmol bond)
HeteronuclearHF
Polar covalent bond(565 kJmol bond)
Electronegativity
Electronegativity
Because F (EN = 40) is more electronegative than H (EN = 22) the electrons move closer to F
This gives rise to a polar bond
H F
Figure 1445
Electrons are not equally sharedin heteronuclear bondsHF
s Antibonding (s)Mostly H(1s)
s BondingMostly F(2p)
H F
H F
MOs of a Polar Covalent Bond HF
This approach simplifies model and only considers electrons involved in bond
MOs OF XY MOLECULES
Equal or unequal e sharing between 2 atoms is reflected in the composition of the MOsmiddot When 2 atoms X and Y have the same electronegativity (purely
covalent bond) their overlapping AOs have the same energy and the bonding and anti-bonding MOs are each half X and half Y AO All electrons spend equal time near X and Y Examples N2 O2 F2
middot If EN(Y) gt EN(X) (polar covalent X+Y) the Y AO has lower energy than the X AO The bonding MO is more like the Y AO and the anti-bonding MO more like the X AO Bonding e spend more time near Y than X vice versa for anti-bonding e Example CO
MOs OF XY MOLECULES ____ s ___ ___ p ___ ___ ___2p ___ ___ ___ 2p uarr ____ s
Energy ___ ___ p ____ s ___ 2s ___ 2s ____ s
middot CO Bond Order = 30 (same as N2)middot CO Bond Energy = 1076 kJmol (N2 = 945 kJmol)middot Isoelectronic to CO and N2 CNndash NO+middot NO has 1endash in p bond order = 25 this endash is more on N than O
NO NO+ easyhellip
C Atom (4endash) Cδ+Oδndash (10endash) O Atom (6endash)
___ ___ ___ 2p ___ 2s
Electronegativity
Bonding in NObull Two possible Lewis dot structures
for NObull The simplest structure minimizes
formal charges and places the lone (unpaired) electron on the nitrogen
bull The Lewis structure predicts a bond order of 2 but experimental evidence suggests a bond order between 2 and 3
bull How does MO theory help us understand bonding in NO
N=O
N=O
+1-1
When the electronegativities of the 2 atoms are more similar the bonding
becomes less polar
EN(N) = 30EN(O) = 34
N=O
Electronegativity
2s2s
2p2p
NON OBond order = 25 unpaired electron is in a N-like orbital
NO+
oxidation
NO
NO is easily oxidized to form NO+ Why What changes can we predict in the bonding and magnetism of the molecule
Bond Order = (8-3)2 = 25Paramagnetic
Bond Order = (8-2)2 = 3Diamagnetic
s2s
s2s
p2p p2p
p2p
-3320
-1835
-1444 -1374
s2p -1307
-597 p2p (empty)
MO diagram for NO
Key Points of MO Theory ndash Heteronuclear Molecules
bull The more electronegative atom has orbitals lower in energy than the more positive atom
bull Electrons in bonding orbitals are closer to the more electronegative atom anti-bonding electrons are closer to the more positive atom
bull For most diatomic molecules s-p mixing changes the orbital energy levels but since these orbitals are almost always fully occupied their order is less important to us
Combining the Localized Electron and Molecular Orbital Models (into a convenient working
model)
Figure 1447
Only the p bonding changes between these resonance structures - The MO model describes this p bonding more
effectively
Figure 1451
Atomic Orbitals Molecular Orbitals
Another example Benzene
p bonding
s bonding
p atomic orbitals p molecular orbital
MO Theory Expectationsbull You should be able to
ndash predict which atomic orbitals are higher or lower in energy (based on electronegativity differences)
ndash correctly fill a molecular orbital diagramndash correctly calculate bond orderndash predict molecular magnetic properties
based on orbital occupationndash understand how molecular properties
change upon ionization (oxidation or reduction) of molecules
- What Do Molecules Look Like
- Slide 2
- Slide 3
- Arranging Electron Pairs
- Arranging Electron Pairs (cont)
- VSEPR Structure Guidelines
- VSEPR 2 electron pairs
- VSEPR 3 electron pairs
- VSEPR 4 electron pairs (cont)
- Bonding vs Lone pairs
- VSEPR 4 electron pairs
- VSEPR 4 electron pairs (cont) (2)
- VSEPR 4 electron pairs (cont) (3)
- Slide 14
- Molecular vs Electron-Pair Geometry
- Slide 16
- VSEPR Beyond the Octet
- VSEPR 5 electron pairs
- VSEPR 5 electron pairs (2)
- 5-electron-pair geometries
- VSEPR 6 electron pairs
- Slide 22
- Molecular Dipole Moments
- Slide 24
- Molecular Dipoles (cont)
- Molecular Dipole Example
- Advanced VSEPR Application
- The VSEPR Table
- The VSEPR Table (2)
- Slide 30
- Slide 31
- Slide 32
- Slide 33
- Slide 34
- Slide 35
- Slide 36
- Slide 37
- Slide 38
- Slide 39
- Slide 40
- Slide 41
- Slide 42
- Slide 43
- Slide 44
- Slide 45
- Slide 46
- Slide 47
- Slide 48
- Hybridization
- Hybridization (2)
- Hybridization (3)
- General Notes
- What is important to know about hybridization
- What is important to know about hybridization (2)
- What is important to know about hybridization (3)
- Slide 56
- Slide 57
- Molecular Orbital Theory
- Slide 59
- Constructive and Destructive Interference
- Slide 61
- MO Diagrams
- Slide 63
- Slide 64
- Review of Orbital Filling
- Slide 66
- Slide 67
- Filling Molecular Orbitals with Electrons
- Slide 69
- Slide 70
- Slide 71
- Forming Bonds
- Anti-bonding Orbitals
- MO Theory in Bonding
- Slide 75
- Slide 76
- Slide 77
- Slide 78
- Slide 79
- Slide 80
- Slide 81
- Slide 82
- Slide 83
- Slide 84
- Slide 85
- MO Diagram for H2 vs N2
- Slide 87
- Slide 88
- Slide 89
- Slide 90
- Molecular Orbitals of X2 Molecules
- Slide 92
- Slide 93
- Slide 94
- Slide 95
- Slide 96
- Slide 97
- Slide 98
- Slide 99
- Slide 100
- Slide 101
- Slide 102
- Slide 103
- Slide 104
- Slide 105
- Bonding in NO
- Slide 107
- Slide 108
- Slide 109
- Key Points of MO Theory ndash Heteronuclear Molecules
- Slide 111
- Slide 112
- Slide 113
- MO Theory Expectations
-
Sample Problem - ContinuedSolution The MO energy levels are
N2
sp
p2p
N2+
s2p
p2p
s2s
s2s
O2 O2+
sp
s2s
s2s
p2p
p2p
s2p
Bond Orders(8-2)2 = 3 (7-2)2 = 25 (8-4)2 = 2 (8-3)2 = 25
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11Bond Order 3 25 2 25
What have we learned so far1 Molecular orbitals (MO) explain the properties of valence electrons in molecules (Example O2)
2 s and p atomic orbitals can be mixed to form s s p and p molecular orbitals
3 Electrons in p or p molecular orbitals can have the same energies Degenerate orbitals
4 The ordering of s2p and p2p molecular orbitals depends on the electron occupancy s-p mixing
Bonding in Diatomic Molecules
Covalent
Ionic
IonicCovalent
HomonuclearH2
Nonpolar covalent bond(450 kJmol bond)
HeteronuclearHF
Polar covalent bond(565 kJmol bond)
Electronegativity
Electronegativity
Because F (EN = 40) is more electronegative than H (EN = 22) the electrons move closer to F
This gives rise to a polar bond
H F
Figure 1445
Electrons are not equally sharedin heteronuclear bondsHF
s Antibonding (s)Mostly H(1s)
s BondingMostly F(2p)
H F
H F
MOs of a Polar Covalent Bond HF
This approach simplifies model and only considers electrons involved in bond
MOs OF XY MOLECULES
Equal or unequal e sharing between 2 atoms is reflected in the composition of the MOsmiddot When 2 atoms X and Y have the same electronegativity (purely
covalent bond) their overlapping AOs have the same energy and the bonding and anti-bonding MOs are each half X and half Y AO All electrons spend equal time near X and Y Examples N2 O2 F2
middot If EN(Y) gt EN(X) (polar covalent X+Y) the Y AO has lower energy than the X AO The bonding MO is more like the Y AO and the anti-bonding MO more like the X AO Bonding e spend more time near Y than X vice versa for anti-bonding e Example CO
MOs OF XY MOLECULES ____ s ___ ___ p ___ ___ ___2p ___ ___ ___ 2p uarr ____ s
Energy ___ ___ p ____ s ___ 2s ___ 2s ____ s
middot CO Bond Order = 30 (same as N2)middot CO Bond Energy = 1076 kJmol (N2 = 945 kJmol)middot Isoelectronic to CO and N2 CNndash NO+middot NO has 1endash in p bond order = 25 this endash is more on N than O
NO NO+ easyhellip
C Atom (4endash) Cδ+Oδndash (10endash) O Atom (6endash)
___ ___ ___ 2p ___ 2s
Electronegativity
Bonding in NObull Two possible Lewis dot structures
for NObull The simplest structure minimizes
formal charges and places the lone (unpaired) electron on the nitrogen
bull The Lewis structure predicts a bond order of 2 but experimental evidence suggests a bond order between 2 and 3
bull How does MO theory help us understand bonding in NO
N=O
N=O
+1-1
When the electronegativities of the 2 atoms are more similar the bonding
becomes less polar
EN(N) = 30EN(O) = 34
N=O
Electronegativity
2s2s
2p2p
NON OBond order = 25 unpaired electron is in a N-like orbital
NO+
oxidation
NO
NO is easily oxidized to form NO+ Why What changes can we predict in the bonding and magnetism of the molecule
Bond Order = (8-3)2 = 25Paramagnetic
Bond Order = (8-2)2 = 3Diamagnetic
s2s
s2s
p2p p2p
p2p
-3320
-1835
-1444 -1374
s2p -1307
-597 p2p (empty)
MO diagram for NO
Key Points of MO Theory ndash Heteronuclear Molecules
bull The more electronegative atom has orbitals lower in energy than the more positive atom
bull Electrons in bonding orbitals are closer to the more electronegative atom anti-bonding electrons are closer to the more positive atom
bull For most diatomic molecules s-p mixing changes the orbital energy levels but since these orbitals are almost always fully occupied their order is less important to us
Combining the Localized Electron and Molecular Orbital Models (into a convenient working
model)
Figure 1447
Only the p bonding changes between these resonance structures - The MO model describes this p bonding more
effectively
Figure 1451
Atomic Orbitals Molecular Orbitals
Another example Benzene
p bonding
s bonding
p atomic orbitals p molecular orbital
MO Theory Expectationsbull You should be able to
ndash predict which atomic orbitals are higher or lower in energy (based on electronegativity differences)
ndash correctly fill a molecular orbital diagramndash correctly calculate bond orderndash predict molecular magnetic properties
based on orbital occupationndash understand how molecular properties
change upon ionization (oxidation or reduction) of molecules
- What Do Molecules Look Like
- Slide 2
- Slide 3
- Arranging Electron Pairs
- Arranging Electron Pairs (cont)
- VSEPR Structure Guidelines
- VSEPR 2 electron pairs
- VSEPR 3 electron pairs
- VSEPR 4 electron pairs (cont)
- Bonding vs Lone pairs
- VSEPR 4 electron pairs
- VSEPR 4 electron pairs (cont) (2)
- VSEPR 4 electron pairs (cont) (3)
- Slide 14
- Molecular vs Electron-Pair Geometry
- Slide 16
- VSEPR Beyond the Octet
- VSEPR 5 electron pairs
- VSEPR 5 electron pairs (2)
- 5-electron-pair geometries
- VSEPR 6 electron pairs
- Slide 22
- Molecular Dipole Moments
- Slide 24
- Molecular Dipoles (cont)
- Molecular Dipole Example
- Advanced VSEPR Application
- The VSEPR Table
- The VSEPR Table (2)
- Slide 30
- Slide 31
- Slide 32
- Slide 33
- Slide 34
- Slide 35
- Slide 36
- Slide 37
- Slide 38
- Slide 39
- Slide 40
- Slide 41
- Slide 42
- Slide 43
- Slide 44
- Slide 45
- Slide 46
- Slide 47
- Slide 48
- Hybridization
- Hybridization (2)
- Hybridization (3)
- General Notes
- What is important to know about hybridization
- What is important to know about hybridization (2)
- What is important to know about hybridization (3)
- Slide 56
- Slide 57
- Molecular Orbital Theory
- Slide 59
- Constructive and Destructive Interference
- Slide 61
- MO Diagrams
- Slide 63
- Slide 64
- Review of Orbital Filling
- Slide 66
- Slide 67
- Filling Molecular Orbitals with Electrons
- Slide 69
- Slide 70
- Slide 71
- Forming Bonds
- Anti-bonding Orbitals
- MO Theory in Bonding
- Slide 75
- Slide 76
- Slide 77
- Slide 78
- Slide 79
- Slide 80
- Slide 81
- Slide 82
- Slide 83
- Slide 84
- Slide 85
- MO Diagram for H2 vs N2
- Slide 87
- Slide 88
- Slide 89
- Slide 90
- Molecular Orbitals of X2 Molecules
- Slide 92
- Slide 93
- Slide 94
- Slide 95
- Slide 96
- Slide 97
- Slide 98
- Slide 99
- Slide 100
- Slide 101
- Slide 102
- Slide 103
- Slide 104
- Slide 105
- Bonding in NO
- Slide 107
- Slide 108
- Slide 109
- Key Points of MO Theory ndash Heteronuclear Molecules
- Slide 111
- Slide 112
- Slide 113
- MO Theory Expectations
-
Sample ProblemUsing MO Theory to Explain Bond Properties
Problem Consider the following data for these homonuclear diatomic species
N2 N2+ O2 O2
+
Bond energy (kJmol) 945 841 498 623Bond length (pm) 110 112 121 112No of valence electrons 10 9 12 11Bond Order 3 25 2 25
What have we learned so far1 Molecular orbitals (MO) explain the properties of valence electrons in molecules (Example O2)
2 s and p atomic orbitals can be mixed to form s s p and p molecular orbitals
3 Electrons in p or p molecular orbitals can have the same energies Degenerate orbitals
4 The ordering of s2p and p2p molecular orbitals depends on the electron occupancy s-p mixing
Bonding in Diatomic Molecules
Covalent
Ionic
IonicCovalent
HomonuclearH2
Nonpolar covalent bond(450 kJmol bond)
HeteronuclearHF
Polar covalent bond(565 kJmol bond)
Electronegativity
Electronegativity
Because F (EN = 40) is more electronegative than H (EN = 22) the electrons move closer to F
This gives rise to a polar bond
H F
Figure 1445
Electrons are not equally sharedin heteronuclear bondsHF
s Antibonding (s)Mostly H(1s)
s BondingMostly F(2p)
H F
H F
MOs of a Polar Covalent Bond HF
This approach simplifies model and only considers electrons involved in bond
MOs OF XY MOLECULES
Equal or unequal e sharing between 2 atoms is reflected in the composition of the MOsmiddot When 2 atoms X and Y have the same electronegativity (purely
covalent bond) their overlapping AOs have the same energy and the bonding and anti-bonding MOs are each half X and half Y AO All electrons spend equal time near X and Y Examples N2 O2 F2
middot If EN(Y) gt EN(X) (polar covalent X+Y) the Y AO has lower energy than the X AO The bonding MO is more like the Y AO and the anti-bonding MO more like the X AO Bonding e spend more time near Y than X vice versa for anti-bonding e Example CO
MOs OF XY MOLECULES ____ s ___ ___ p ___ ___ ___2p ___ ___ ___ 2p uarr ____ s
Energy ___ ___ p ____ s ___ 2s ___ 2s ____ s
middot CO Bond Order = 30 (same as N2)middot CO Bond Energy = 1076 kJmol (N2 = 945 kJmol)middot Isoelectronic to CO and N2 CNndash NO+middot NO has 1endash in p bond order = 25 this endash is more on N than O
NO NO+ easyhellip
C Atom (4endash) Cδ+Oδndash (10endash) O Atom (6endash)
___ ___ ___ 2p ___ 2s
Electronegativity
Bonding in NObull Two possible Lewis dot structures
for NObull The simplest structure minimizes
formal charges and places the lone (unpaired) electron on the nitrogen
bull The Lewis structure predicts a bond order of 2 but experimental evidence suggests a bond order between 2 and 3
bull How does MO theory help us understand bonding in NO
N=O
N=O
+1-1
When the electronegativities of the 2 atoms are more similar the bonding
becomes less polar
EN(N) = 30EN(O) = 34
N=O
Electronegativity
2s2s
2p2p
NON OBond order = 25 unpaired electron is in a N-like orbital
NO+
oxidation
NO
NO is easily oxidized to form NO+ Why What changes can we predict in the bonding and magnetism of the molecule
Bond Order = (8-3)2 = 25Paramagnetic
Bond Order = (8-2)2 = 3Diamagnetic
s2s
s2s
p2p p2p
p2p
-3320
-1835
-1444 -1374
s2p -1307
-597 p2p (empty)
MO diagram for NO
Key Points of MO Theory ndash Heteronuclear Molecules
bull The more electronegative atom has orbitals lower in energy than the more positive atom
bull Electrons in bonding orbitals are closer to the more electronegative atom anti-bonding electrons are closer to the more positive atom
bull For most diatomic molecules s-p mixing changes the orbital energy levels but since these orbitals are almost always fully occupied their order is less important to us
Combining the Localized Electron and Molecular Orbital Models (into a convenient working
model)
Figure 1447
Only the p bonding changes between these resonance structures - The MO model describes this p bonding more
effectively
Figure 1451
Atomic Orbitals Molecular Orbitals
Another example Benzene
p bonding
s bonding
p atomic orbitals p molecular orbital
MO Theory Expectationsbull You should be able to
ndash predict which atomic orbitals are higher or lower in energy (based on electronegativity differences)
ndash correctly fill a molecular orbital diagramndash correctly calculate bond orderndash predict molecular magnetic properties
based on orbital occupationndash understand how molecular properties
change upon ionization (oxidation or reduction) of molecules
- What Do Molecules Look Like
- Slide 2
- Slide 3
- Arranging Electron Pairs
- Arranging Electron Pairs (cont)
- VSEPR Structure Guidelines
- VSEPR 2 electron pairs
- VSEPR 3 electron pairs
- VSEPR 4 electron pairs (cont)
- Bonding vs Lone pairs
- VSEPR 4 electron pairs
- VSEPR 4 electron pairs (cont) (2)
- VSEPR 4 electron pairs (cont) (3)
- Slide 14
- Molecular vs Electron-Pair Geometry
- Slide 16
- VSEPR Beyond the Octet
- VSEPR 5 electron pairs
- VSEPR 5 electron pairs (2)
- 5-electron-pair geometries
- VSEPR 6 electron pairs
- Slide 22
- Molecular Dipole Moments
- Slide 24
- Molecular Dipoles (cont)
- Molecular Dipole Example
- Advanced VSEPR Application
- The VSEPR Table
- The VSEPR Table (2)
- Slide 30
- Slide 31
- Slide 32
- Slide 33
- Slide 34
- Slide 35
- Slide 36
- Slide 37
- Slide 38
- Slide 39
- Slide 40
- Slide 41
- Slide 42
- Slide 43
- Slide 44
- Slide 45
- Slide 46
- Slide 47
- Slide 48
- Hybridization
- Hybridization (2)
- Hybridization (3)
- General Notes
- What is important to know about hybridization
- What is important to know about hybridization (2)
- What is important to know about hybridization (3)
- Slide 56
- Slide 57
- Molecular Orbital Theory
- Slide 59
- Constructive and Destructive Interference
- Slide 61
- MO Diagrams
- Slide 63
- Slide 64
- Review of Orbital Filling
- Slide 66
- Slide 67
- Filling Molecular Orbitals with Electrons
- Slide 69
- Slide 70
- Slide 71
- Forming Bonds
- Anti-bonding Orbitals
- MO Theory in Bonding
- Slide 75
- Slide 76
- Slide 77
- Slide 78
- Slide 79
- Slide 80
- Slide 81
- Slide 82
- Slide 83
- Slide 84
- Slide 85
- MO Diagram for H2 vs N2
- Slide 87
- Slide 88
- Slide 89
- Slide 90
- Molecular Orbitals of X2 Molecules
- Slide 92
- Slide 93
- Slide 94
- Slide 95
- Slide 96
- Slide 97
- Slide 98
- Slide 99
- Slide 100
- Slide 101
- Slide 102
- Slide 103
- Slide 104
- Slide 105
- Bonding in NO
- Slide 107
- Slide 108
- Slide 109
- Key Points of MO Theory ndash Heteronuclear Molecules
- Slide 111
- Slide 112
- Slide 113
- MO Theory Expectations
-
What have we learned so far1 Molecular orbitals (MO) explain the properties of valence electrons in molecules (Example O2)
2 s and p atomic orbitals can be mixed to form s s p and p molecular orbitals
3 Electrons in p or p molecular orbitals can have the same energies Degenerate orbitals
4 The ordering of s2p and p2p molecular orbitals depends on the electron occupancy s-p mixing
Bonding in Diatomic Molecules
Covalent
Ionic
IonicCovalent
HomonuclearH2
Nonpolar covalent bond(450 kJmol bond)
HeteronuclearHF
Polar covalent bond(565 kJmol bond)
Electronegativity
Electronegativity
Because F (EN = 40) is more electronegative than H (EN = 22) the electrons move closer to F
This gives rise to a polar bond
H F
Figure 1445
Electrons are not equally sharedin heteronuclear bondsHF
s Antibonding (s)Mostly H(1s)
s BondingMostly F(2p)
H F
H F
MOs of a Polar Covalent Bond HF
This approach simplifies model and only considers electrons involved in bond
MOs OF XY MOLECULES
Equal or unequal e sharing between 2 atoms is reflected in the composition of the MOsmiddot When 2 atoms X and Y have the same electronegativity (purely
covalent bond) their overlapping AOs have the same energy and the bonding and anti-bonding MOs are each half X and half Y AO All electrons spend equal time near X and Y Examples N2 O2 F2
middot If EN(Y) gt EN(X) (polar covalent X+Y) the Y AO has lower energy than the X AO The bonding MO is more like the Y AO and the anti-bonding MO more like the X AO Bonding e spend more time near Y than X vice versa for anti-bonding e Example CO
MOs OF XY MOLECULES ____ s ___ ___ p ___ ___ ___2p ___ ___ ___ 2p uarr ____ s
Energy ___ ___ p ____ s ___ 2s ___ 2s ____ s
middot CO Bond Order = 30 (same as N2)middot CO Bond Energy = 1076 kJmol (N2 = 945 kJmol)middot Isoelectronic to CO and N2 CNndash NO+middot NO has 1endash in p bond order = 25 this endash is more on N than O
NO NO+ easyhellip
C Atom (4endash) Cδ+Oδndash (10endash) O Atom (6endash)
___ ___ ___ 2p ___ 2s
Electronegativity
Bonding in NObull Two possible Lewis dot structures
for NObull The simplest structure minimizes
formal charges and places the lone (unpaired) electron on the nitrogen
bull The Lewis structure predicts a bond order of 2 but experimental evidence suggests a bond order between 2 and 3
bull How does MO theory help us understand bonding in NO
N=O
N=O
+1-1
When the electronegativities of the 2 atoms are more similar the bonding
becomes less polar
EN(N) = 30EN(O) = 34
N=O
Electronegativity
2s2s
2p2p
NON OBond order = 25 unpaired electron is in a N-like orbital
NO+
oxidation
NO
NO is easily oxidized to form NO+ Why What changes can we predict in the bonding and magnetism of the molecule
Bond Order = (8-3)2 = 25Paramagnetic
Bond Order = (8-2)2 = 3Diamagnetic
s2s
s2s
p2p p2p
p2p
-3320
-1835
-1444 -1374
s2p -1307
-597 p2p (empty)
MO diagram for NO
Key Points of MO Theory ndash Heteronuclear Molecules
bull The more electronegative atom has orbitals lower in energy than the more positive atom
bull Electrons in bonding orbitals are closer to the more electronegative atom anti-bonding electrons are closer to the more positive atom
bull For most diatomic molecules s-p mixing changes the orbital energy levels but since these orbitals are almost always fully occupied their order is less important to us
Combining the Localized Electron and Molecular Orbital Models (into a convenient working
model)
Figure 1447
Only the p bonding changes between these resonance structures - The MO model describes this p bonding more
effectively
Figure 1451
Atomic Orbitals Molecular Orbitals
Another example Benzene
p bonding
s bonding
p atomic orbitals p molecular orbital
MO Theory Expectationsbull You should be able to
ndash predict which atomic orbitals are higher or lower in energy (based on electronegativity differences)
ndash correctly fill a molecular orbital diagramndash correctly calculate bond orderndash predict molecular magnetic properties
based on orbital occupationndash understand how molecular properties
change upon ionization (oxidation or reduction) of molecules
- What Do Molecules Look Like
- Slide 2
- Slide 3
- Arranging Electron Pairs
- Arranging Electron Pairs (cont)
- VSEPR Structure Guidelines
- VSEPR 2 electron pairs
- VSEPR 3 electron pairs
- VSEPR 4 electron pairs (cont)
- Bonding vs Lone pairs
- VSEPR 4 electron pairs
- VSEPR 4 electron pairs (cont) (2)
- VSEPR 4 electron pairs (cont) (3)
- Slide 14
- Molecular vs Electron-Pair Geometry
- Slide 16
- VSEPR Beyond the Octet
- VSEPR 5 electron pairs
- VSEPR 5 electron pairs (2)
- 5-electron-pair geometries
- VSEPR 6 electron pairs
- Slide 22
- Molecular Dipole Moments
- Slide 24
- Molecular Dipoles (cont)
- Molecular Dipole Example
- Advanced VSEPR Application
- The VSEPR Table
- The VSEPR Table (2)
- Slide 30
- Slide 31
- Slide 32
- Slide 33
- Slide 34
- Slide 35
- Slide 36
- Slide 37
- Slide 38
- Slide 39
- Slide 40
- Slide 41
- Slide 42
- Slide 43
- Slide 44
- Slide 45
- Slide 46
- Slide 47
- Slide 48
- Hybridization
- Hybridization (2)
- Hybridization (3)
- General Notes
- What is important to know about hybridization
- What is important to know about hybridization (2)
- What is important to know about hybridization (3)
- Slide 56
- Slide 57
- Molecular Orbital Theory
- Slide 59
- Constructive and Destructive Interference
- Slide 61
- MO Diagrams
- Slide 63
- Slide 64
- Review of Orbital Filling
- Slide 66
- Slide 67
- Filling Molecular Orbitals with Electrons
- Slide 69
- Slide 70
- Slide 71
- Forming Bonds
- Anti-bonding Orbitals
- MO Theory in Bonding
- Slide 75
- Slide 76
- Slide 77
- Slide 78
- Slide 79
- Slide 80
- Slide 81
- Slide 82
- Slide 83
- Slide 84
- Slide 85
- MO Diagram for H2 vs N2
- Slide 87
- Slide 88
- Slide 89
- Slide 90
- Molecular Orbitals of X2 Molecules
- Slide 92
- Slide 93
- Slide 94
- Slide 95
- Slide 96
- Slide 97
- Slide 98
- Slide 99
- Slide 100
- Slide 101
- Slide 102
- Slide 103
- Slide 104
- Slide 105
- Bonding in NO
- Slide 107
- Slide 108
- Slide 109
- Key Points of MO Theory ndash Heteronuclear Molecules
- Slide 111
- Slide 112
- Slide 113
- MO Theory Expectations
-
Bonding in Diatomic Molecules
Covalent
Ionic
IonicCovalent
HomonuclearH2
Nonpolar covalent bond(450 kJmol bond)
HeteronuclearHF
Polar covalent bond(565 kJmol bond)
Electronegativity
Electronegativity
Because F (EN = 40) is more electronegative than H (EN = 22) the electrons move closer to F
This gives rise to a polar bond
H F
Figure 1445
Electrons are not equally sharedin heteronuclear bondsHF
s Antibonding (s)Mostly H(1s)
s BondingMostly F(2p)
H F
H F
MOs of a Polar Covalent Bond HF
This approach simplifies model and only considers electrons involved in bond
MOs OF XY MOLECULES
Equal or unequal e sharing between 2 atoms is reflected in the composition of the MOsmiddot When 2 atoms X and Y have the same electronegativity (purely
covalent bond) their overlapping AOs have the same energy and the bonding and anti-bonding MOs are each half X and half Y AO All electrons spend equal time near X and Y Examples N2 O2 F2
middot If EN(Y) gt EN(X) (polar covalent X+Y) the Y AO has lower energy than the X AO The bonding MO is more like the Y AO and the anti-bonding MO more like the X AO Bonding e spend more time near Y than X vice versa for anti-bonding e Example CO
MOs OF XY MOLECULES ____ s ___ ___ p ___ ___ ___2p ___ ___ ___ 2p uarr ____ s
Energy ___ ___ p ____ s ___ 2s ___ 2s ____ s
middot CO Bond Order = 30 (same as N2)middot CO Bond Energy = 1076 kJmol (N2 = 945 kJmol)middot Isoelectronic to CO and N2 CNndash NO+middot NO has 1endash in p bond order = 25 this endash is more on N than O
NO NO+ easyhellip
C Atom (4endash) Cδ+Oδndash (10endash) O Atom (6endash)
___ ___ ___ 2p ___ 2s
Electronegativity
Bonding in NObull Two possible Lewis dot structures
for NObull The simplest structure minimizes
formal charges and places the lone (unpaired) electron on the nitrogen
bull The Lewis structure predicts a bond order of 2 but experimental evidence suggests a bond order between 2 and 3
bull How does MO theory help us understand bonding in NO
N=O
N=O
+1-1
When the electronegativities of the 2 atoms are more similar the bonding
becomes less polar
EN(N) = 30EN(O) = 34
N=O
Electronegativity
2s2s
2p2p
NON OBond order = 25 unpaired electron is in a N-like orbital
NO+
oxidation
NO
NO is easily oxidized to form NO+ Why What changes can we predict in the bonding and magnetism of the molecule
Bond Order = (8-3)2 = 25Paramagnetic
Bond Order = (8-2)2 = 3Diamagnetic
s2s
s2s
p2p p2p
p2p
-3320
-1835
-1444 -1374
s2p -1307
-597 p2p (empty)
MO diagram for NO
Key Points of MO Theory ndash Heteronuclear Molecules
bull The more electronegative atom has orbitals lower in energy than the more positive atom
bull Electrons in bonding orbitals are closer to the more electronegative atom anti-bonding electrons are closer to the more positive atom
bull For most diatomic molecules s-p mixing changes the orbital energy levels but since these orbitals are almost always fully occupied their order is less important to us
Combining the Localized Electron and Molecular Orbital Models (into a convenient working
model)
Figure 1447
Only the p bonding changes between these resonance structures - The MO model describes this p bonding more
effectively
Figure 1451
Atomic Orbitals Molecular Orbitals
Another example Benzene
p bonding
s bonding
p atomic orbitals p molecular orbital
MO Theory Expectationsbull You should be able to
ndash predict which atomic orbitals are higher or lower in energy (based on electronegativity differences)
ndash correctly fill a molecular orbital diagramndash correctly calculate bond orderndash predict molecular magnetic properties
based on orbital occupationndash understand how molecular properties
change upon ionization (oxidation or reduction) of molecules
- What Do Molecules Look Like
- Slide 2
- Slide 3
- Arranging Electron Pairs
- Arranging Electron Pairs (cont)
- VSEPR Structure Guidelines
- VSEPR 2 electron pairs
- VSEPR 3 electron pairs
- VSEPR 4 electron pairs (cont)
- Bonding vs Lone pairs
- VSEPR 4 electron pairs
- VSEPR 4 electron pairs (cont) (2)
- VSEPR 4 electron pairs (cont) (3)
- Slide 14
- Molecular vs Electron-Pair Geometry
- Slide 16
- VSEPR Beyond the Octet
- VSEPR 5 electron pairs
- VSEPR 5 electron pairs (2)
- 5-electron-pair geometries
- VSEPR 6 electron pairs
- Slide 22
- Molecular Dipole Moments
- Slide 24
- Molecular Dipoles (cont)
- Molecular Dipole Example
- Advanced VSEPR Application
- The VSEPR Table
- The VSEPR Table (2)
- Slide 30
- Slide 31
- Slide 32
- Slide 33
- Slide 34
- Slide 35
- Slide 36
- Slide 37
- Slide 38
- Slide 39
- Slide 40
- Slide 41
- Slide 42
- Slide 43
- Slide 44
- Slide 45
- Slide 46
- Slide 47
- Slide 48
- Hybridization
- Hybridization (2)
- Hybridization (3)
- General Notes
- What is important to know about hybridization
- What is important to know about hybridization (2)
- What is important to know about hybridization (3)
- Slide 56
- Slide 57
- Molecular Orbital Theory
- Slide 59
- Constructive and Destructive Interference
- Slide 61
- MO Diagrams
- Slide 63
- Slide 64
- Review of Orbital Filling
- Slide 66
- Slide 67
- Filling Molecular Orbitals with Electrons
- Slide 69
- Slide 70
- Slide 71
- Forming Bonds
- Anti-bonding Orbitals
- MO Theory in Bonding
- Slide 75
- Slide 76
- Slide 77
- Slide 78
- Slide 79
- Slide 80
- Slide 81
- Slide 82
- Slide 83
- Slide 84
- Slide 85
- MO Diagram for H2 vs N2
- Slide 87
- Slide 88
- Slide 89
- Slide 90
- Molecular Orbitals of X2 Molecules
- Slide 92
- Slide 93
- Slide 94
- Slide 95
- Slide 96
- Slide 97
- Slide 98
- Slide 99
- Slide 100
- Slide 101
- Slide 102
- Slide 103
- Slide 104
- Slide 105
- Bonding in NO
- Slide 107
- Slide 108
- Slide 109
- Key Points of MO Theory ndash Heteronuclear Molecules
- Slide 111
- Slide 112
- Slide 113
- MO Theory Expectations
-
HomonuclearH2
Nonpolar covalent bond(450 kJmol bond)
HeteronuclearHF
Polar covalent bond(565 kJmol bond)
Electronegativity
Electronegativity
Because F (EN = 40) is more electronegative than H (EN = 22) the electrons move closer to F
This gives rise to a polar bond
H F
Figure 1445
Electrons are not equally sharedin heteronuclear bondsHF
s Antibonding (s)Mostly H(1s)
s BondingMostly F(2p)
H F
H F
MOs of a Polar Covalent Bond HF
This approach simplifies model and only considers electrons involved in bond
MOs OF XY MOLECULES
Equal or unequal e sharing between 2 atoms is reflected in the composition of the MOsmiddot When 2 atoms X and Y have the same electronegativity (purely
covalent bond) their overlapping AOs have the same energy and the bonding and anti-bonding MOs are each half X and half Y AO All electrons spend equal time near X and Y Examples N2 O2 F2
middot If EN(Y) gt EN(X) (polar covalent X+Y) the Y AO has lower energy than the X AO The bonding MO is more like the Y AO and the anti-bonding MO more like the X AO Bonding e spend more time near Y than X vice versa for anti-bonding e Example CO
MOs OF XY MOLECULES ____ s ___ ___ p ___ ___ ___2p ___ ___ ___ 2p uarr ____ s
Energy ___ ___ p ____ s ___ 2s ___ 2s ____ s
middot CO Bond Order = 30 (same as N2)middot CO Bond Energy = 1076 kJmol (N2 = 945 kJmol)middot Isoelectronic to CO and N2 CNndash NO+middot NO has 1endash in p bond order = 25 this endash is more on N than O
NO NO+ easyhellip
C Atom (4endash) Cδ+Oδndash (10endash) O Atom (6endash)
___ ___ ___ 2p ___ 2s
Electronegativity
Bonding in NObull Two possible Lewis dot structures
for NObull The simplest structure minimizes
formal charges and places the lone (unpaired) electron on the nitrogen
bull The Lewis structure predicts a bond order of 2 but experimental evidence suggests a bond order between 2 and 3
bull How does MO theory help us understand bonding in NO
N=O
N=O
+1-1
When the electronegativities of the 2 atoms are more similar the bonding
becomes less polar
EN(N) = 30EN(O) = 34
N=O
Electronegativity
2s2s
2p2p
NON OBond order = 25 unpaired electron is in a N-like orbital
NO+
oxidation
NO
NO is easily oxidized to form NO+ Why What changes can we predict in the bonding and magnetism of the molecule
Bond Order = (8-3)2 = 25Paramagnetic
Bond Order = (8-2)2 = 3Diamagnetic
s2s
s2s
p2p p2p
p2p
-3320
-1835
-1444 -1374
s2p -1307
-597 p2p (empty)
MO diagram for NO
Key Points of MO Theory ndash Heteronuclear Molecules
bull The more electronegative atom has orbitals lower in energy than the more positive atom
bull Electrons in bonding orbitals are closer to the more electronegative atom anti-bonding electrons are closer to the more positive atom
bull For most diatomic molecules s-p mixing changes the orbital energy levels but since these orbitals are almost always fully occupied their order is less important to us
Combining the Localized Electron and Molecular Orbital Models (into a convenient working
model)
Figure 1447
Only the p bonding changes between these resonance structures - The MO model describes this p bonding more
effectively
Figure 1451
Atomic Orbitals Molecular Orbitals
Another example Benzene
p bonding
s bonding
p atomic orbitals p molecular orbital
MO Theory Expectationsbull You should be able to
ndash predict which atomic orbitals are higher or lower in energy (based on electronegativity differences)
ndash correctly fill a molecular orbital diagramndash correctly calculate bond orderndash predict molecular magnetic properties
based on orbital occupationndash understand how molecular properties
change upon ionization (oxidation or reduction) of molecules
- What Do Molecules Look Like
- Slide 2
- Slide 3
- Arranging Electron Pairs
- Arranging Electron Pairs (cont)
- VSEPR Structure Guidelines
- VSEPR 2 electron pairs
- VSEPR 3 electron pairs
- VSEPR 4 electron pairs (cont)
- Bonding vs Lone pairs
- VSEPR 4 electron pairs
- VSEPR 4 electron pairs (cont) (2)
- VSEPR 4 electron pairs (cont) (3)
- Slide 14
- Molecular vs Electron-Pair Geometry
- Slide 16
- VSEPR Beyond the Octet
- VSEPR 5 electron pairs
- VSEPR 5 electron pairs (2)
- 5-electron-pair geometries
- VSEPR 6 electron pairs
- Slide 22
- Molecular Dipole Moments
- Slide 24
- Molecular Dipoles (cont)
- Molecular Dipole Example
- Advanced VSEPR Application
- The VSEPR Table
- The VSEPR Table (2)
- Slide 30
- Slide 31
- Slide 32
- Slide 33
- Slide 34
- Slide 35
- Slide 36
- Slide 37
- Slide 38
- Slide 39
- Slide 40
- Slide 41
- Slide 42
- Slide 43
- Slide 44
- Slide 45
- Slide 46
- Slide 47
- Slide 48
- Hybridization
- Hybridization (2)
- Hybridization (3)
- General Notes
- What is important to know about hybridization
- What is important to know about hybridization (2)
- What is important to know about hybridization (3)
- Slide 56
- Slide 57
- Molecular Orbital Theory
- Slide 59
- Constructive and Destructive Interference
- Slide 61
- MO Diagrams
- Slide 63
- Slide 64
- Review of Orbital Filling
- Slide 66
- Slide 67
- Filling Molecular Orbitals with Electrons
- Slide 69
- Slide 70
- Slide 71
- Forming Bonds
- Anti-bonding Orbitals
- MO Theory in Bonding
- Slide 75
- Slide 76
- Slide 77
- Slide 78
- Slide 79
- Slide 80
- Slide 81
- Slide 82
- Slide 83
- Slide 84
- Slide 85
- MO Diagram for H2 vs N2
- Slide 87
- Slide 88
- Slide 89
- Slide 90
- Molecular Orbitals of X2 Molecules
- Slide 92
- Slide 93
- Slide 94
- Slide 95
- Slide 96
- Slide 97
- Slide 98
- Slide 99
- Slide 100
- Slide 101
- Slide 102
- Slide 103
- Slide 104
- Slide 105
- Bonding in NO
- Slide 107
- Slide 108
- Slide 109
- Key Points of MO Theory ndash Heteronuclear Molecules
- Slide 111
- Slide 112
- Slide 113
- MO Theory Expectations
-
Electronegativity
Because F (EN = 40) is more electronegative than H (EN = 22) the electrons move closer to F
This gives rise to a polar bond
H F
Figure 1445
Electrons are not equally sharedin heteronuclear bondsHF
s Antibonding (s)Mostly H(1s)
s BondingMostly F(2p)
H F
H F
MOs of a Polar Covalent Bond HF
This approach simplifies model and only considers electrons involved in bond
MOs OF XY MOLECULES
Equal or unequal e sharing between 2 atoms is reflected in the composition of the MOsmiddot When 2 atoms X and Y have the same electronegativity (purely
covalent bond) their overlapping AOs have the same energy and the bonding and anti-bonding MOs are each half X and half Y AO All electrons spend equal time near X and Y Examples N2 O2 F2
middot If EN(Y) gt EN(X) (polar covalent X+Y) the Y AO has lower energy than the X AO The bonding MO is more like the Y AO and the anti-bonding MO more like the X AO Bonding e spend more time near Y than X vice versa for anti-bonding e Example CO
MOs OF XY MOLECULES ____ s ___ ___ p ___ ___ ___2p ___ ___ ___ 2p uarr ____ s
Energy ___ ___ p ____ s ___ 2s ___ 2s ____ s
middot CO Bond Order = 30 (same as N2)middot CO Bond Energy = 1076 kJmol (N2 = 945 kJmol)middot Isoelectronic to CO and N2 CNndash NO+middot NO has 1endash in p bond order = 25 this endash is more on N than O
NO NO+ easyhellip
C Atom (4endash) Cδ+Oδndash (10endash) O Atom (6endash)
___ ___ ___ 2p ___ 2s
Electronegativity
Bonding in NObull Two possible Lewis dot structures
for NObull The simplest structure minimizes
formal charges and places the lone (unpaired) electron on the nitrogen
bull The Lewis structure predicts a bond order of 2 but experimental evidence suggests a bond order between 2 and 3
bull How does MO theory help us understand bonding in NO
N=O
N=O
+1-1
When the electronegativities of the 2 atoms are more similar the bonding
becomes less polar
EN(N) = 30EN(O) = 34
N=O
Electronegativity
2s2s
2p2p
NON OBond order = 25 unpaired electron is in a N-like orbital
NO+
oxidation
NO
NO is easily oxidized to form NO+ Why What changes can we predict in the bonding and magnetism of the molecule
Bond Order = (8-3)2 = 25Paramagnetic
Bond Order = (8-2)2 = 3Diamagnetic
s2s
s2s
p2p p2p
p2p
-3320
-1835
-1444 -1374
s2p -1307
-597 p2p (empty)
MO diagram for NO
Key Points of MO Theory ndash Heteronuclear Molecules
bull The more electronegative atom has orbitals lower in energy than the more positive atom
bull Electrons in bonding orbitals are closer to the more electronegative atom anti-bonding electrons are closer to the more positive atom
bull For most diatomic molecules s-p mixing changes the orbital energy levels but since these orbitals are almost always fully occupied their order is less important to us
Combining the Localized Electron and Molecular Orbital Models (into a convenient working
model)
Figure 1447
Only the p bonding changes between these resonance structures - The MO model describes this p bonding more
effectively
Figure 1451
Atomic Orbitals Molecular Orbitals
Another example Benzene
p bonding
s bonding
p atomic orbitals p molecular orbital
MO Theory Expectationsbull You should be able to
ndash predict which atomic orbitals are higher or lower in energy (based on electronegativity differences)
ndash correctly fill a molecular orbital diagramndash correctly calculate bond orderndash predict molecular magnetic properties
based on orbital occupationndash understand how molecular properties
change upon ionization (oxidation or reduction) of molecules
- What Do Molecules Look Like
- Slide 2
- Slide 3
- Arranging Electron Pairs
- Arranging Electron Pairs (cont)
- VSEPR Structure Guidelines
- VSEPR 2 electron pairs
- VSEPR 3 electron pairs
- VSEPR 4 electron pairs (cont)
- Bonding vs Lone pairs
- VSEPR 4 electron pairs
- VSEPR 4 electron pairs (cont) (2)
- VSEPR 4 electron pairs (cont) (3)
- Slide 14
- Molecular vs Electron-Pair Geometry
- Slide 16
- VSEPR Beyond the Octet
- VSEPR 5 electron pairs
- VSEPR 5 electron pairs (2)
- 5-electron-pair geometries
- VSEPR 6 electron pairs
- Slide 22
- Molecular Dipole Moments
- Slide 24
- Molecular Dipoles (cont)
- Molecular Dipole Example
- Advanced VSEPR Application
- The VSEPR Table
- The VSEPR Table (2)
- Slide 30
- Slide 31
- Slide 32
- Slide 33
- Slide 34
- Slide 35
- Slide 36
- Slide 37
- Slide 38
- Slide 39
- Slide 40
- Slide 41
- Slide 42
- Slide 43
- Slide 44
- Slide 45
- Slide 46
- Slide 47
- Slide 48
- Hybridization
- Hybridization (2)
- Hybridization (3)
- General Notes
- What is important to know about hybridization
- What is important to know about hybridization (2)
- What is important to know about hybridization (3)
- Slide 56
- Slide 57
- Molecular Orbital Theory
- Slide 59
- Constructive and Destructive Interference
- Slide 61
- MO Diagrams
- Slide 63
- Slide 64
- Review of Orbital Filling
- Slide 66
- Slide 67
- Filling Molecular Orbitals with Electrons
- Slide 69
- Slide 70
- Slide 71
- Forming Bonds
- Anti-bonding Orbitals
- MO Theory in Bonding
- Slide 75
- Slide 76
- Slide 77
- Slide 78
- Slide 79
- Slide 80
- Slide 81
- Slide 82
- Slide 83
- Slide 84
- Slide 85
- MO Diagram for H2 vs N2
- Slide 87
- Slide 88
- Slide 89
- Slide 90
- Molecular Orbitals of X2 Molecules
- Slide 92
- Slide 93
- Slide 94
- Slide 95
- Slide 96
- Slide 97
- Slide 98
- Slide 99
- Slide 100
- Slide 101
- Slide 102
- Slide 103
- Slide 104
- Slide 105
- Bonding in NO
- Slide 107
- Slide 108
- Slide 109
- Key Points of MO Theory ndash Heteronuclear Molecules
- Slide 111
- Slide 112
- Slide 113
- MO Theory Expectations
-
s Antibonding (s)Mostly H(1s)
s BondingMostly F(2p)
H F
H F
MOs of a Polar Covalent Bond HF
This approach simplifies model and only considers electrons involved in bond
MOs OF XY MOLECULES
Equal or unequal e sharing between 2 atoms is reflected in the composition of the MOsmiddot When 2 atoms X and Y have the same electronegativity (purely
covalent bond) their overlapping AOs have the same energy and the bonding and anti-bonding MOs are each half X and half Y AO All electrons spend equal time near X and Y Examples N2 O2 F2
middot If EN(Y) gt EN(X) (polar covalent X+Y) the Y AO has lower energy than the X AO The bonding MO is more like the Y AO and the anti-bonding MO more like the X AO Bonding e spend more time near Y than X vice versa for anti-bonding e Example CO
MOs OF XY MOLECULES ____ s ___ ___ p ___ ___ ___2p ___ ___ ___ 2p uarr ____ s
Energy ___ ___ p ____ s ___ 2s ___ 2s ____ s
middot CO Bond Order = 30 (same as N2)middot CO Bond Energy = 1076 kJmol (N2 = 945 kJmol)middot Isoelectronic to CO and N2 CNndash NO+middot NO has 1endash in p bond order = 25 this endash is more on N than O
NO NO+ easyhellip
C Atom (4endash) Cδ+Oδndash (10endash) O Atom (6endash)
___ ___ ___ 2p ___ 2s
Electronegativity
Bonding in NObull Two possible Lewis dot structures
for NObull The simplest structure minimizes
formal charges and places the lone (unpaired) electron on the nitrogen
bull The Lewis structure predicts a bond order of 2 but experimental evidence suggests a bond order between 2 and 3
bull How does MO theory help us understand bonding in NO
N=O
N=O
+1-1
When the electronegativities of the 2 atoms are more similar the bonding
becomes less polar
EN(N) = 30EN(O) = 34
N=O
Electronegativity
2s2s
2p2p
NON OBond order = 25 unpaired electron is in a N-like orbital
NO+
oxidation
NO
NO is easily oxidized to form NO+ Why What changes can we predict in the bonding and magnetism of the molecule
Bond Order = (8-3)2 = 25Paramagnetic
Bond Order = (8-2)2 = 3Diamagnetic
s2s
s2s
p2p p2p
p2p
-3320
-1835
-1444 -1374
s2p -1307
-597 p2p (empty)
MO diagram for NO
Key Points of MO Theory ndash Heteronuclear Molecules
bull The more electronegative atom has orbitals lower in energy than the more positive atom
bull Electrons in bonding orbitals are closer to the more electronegative atom anti-bonding electrons are closer to the more positive atom
bull For most diatomic molecules s-p mixing changes the orbital energy levels but since these orbitals are almost always fully occupied their order is less important to us
Combining the Localized Electron and Molecular Orbital Models (into a convenient working
model)
Figure 1447
Only the p bonding changes between these resonance structures - The MO model describes this p bonding more
effectively
Figure 1451
Atomic Orbitals Molecular Orbitals
Another example Benzene
p bonding
s bonding
p atomic orbitals p molecular orbital
MO Theory Expectationsbull You should be able to
ndash predict which atomic orbitals are higher or lower in energy (based on electronegativity differences)
ndash correctly fill a molecular orbital diagramndash correctly calculate bond orderndash predict molecular magnetic properties
based on orbital occupationndash understand how molecular properties
change upon ionization (oxidation or reduction) of molecules
- What Do Molecules Look Like
- Slide 2
- Slide 3
- Arranging Electron Pairs
- Arranging Electron Pairs (cont)
- VSEPR Structure Guidelines
- VSEPR 2 electron pairs
- VSEPR 3 electron pairs
- VSEPR 4 electron pairs (cont)
- Bonding vs Lone pairs
- VSEPR 4 electron pairs
- VSEPR 4 electron pairs (cont) (2)
- VSEPR 4 electron pairs (cont) (3)
- Slide 14
- Molecular vs Electron-Pair Geometry
- Slide 16
- VSEPR Beyond the Octet
- VSEPR 5 electron pairs
- VSEPR 5 electron pairs (2)
- 5-electron-pair geometries
- VSEPR 6 electron pairs
- Slide 22
- Molecular Dipole Moments
- Slide 24
- Molecular Dipoles (cont)
- Molecular Dipole Example
- Advanced VSEPR Application
- The VSEPR Table
- The VSEPR Table (2)
- Slide 30
- Slide 31
- Slide 32
- Slide 33
- Slide 34
- Slide 35
- Slide 36
- Slide 37
- Slide 38
- Slide 39
- Slide 40
- Slide 41
- Slide 42
- Slide 43
- Slide 44
- Slide 45
- Slide 46
- Slide 47
- Slide 48
- Hybridization
- Hybridization (2)
- Hybridization (3)
- General Notes
- What is important to know about hybridization
- What is important to know about hybridization (2)
- What is important to know about hybridization (3)
- Slide 56
- Slide 57
- Molecular Orbital Theory
- Slide 59
- Constructive and Destructive Interference
- Slide 61
- MO Diagrams
- Slide 63
- Slide 64
- Review of Orbital Filling
- Slide 66
- Slide 67
- Filling Molecular Orbitals with Electrons
- Slide 69
- Slide 70
- Slide 71
- Forming Bonds
- Anti-bonding Orbitals
- MO Theory in Bonding
- Slide 75
- Slide 76
- Slide 77
- Slide 78
- Slide 79
- Slide 80
- Slide 81
- Slide 82
- Slide 83
- Slide 84
- Slide 85
- MO Diagram for H2 vs N2
- Slide 87
- Slide 88
- Slide 89
- Slide 90
- Molecular Orbitals of X2 Molecules
- Slide 92
- Slide 93
- Slide 94
- Slide 95
- Slide 96
- Slide 97
- Slide 98
- Slide 99
- Slide 100
- Slide 101
- Slide 102
- Slide 103
- Slide 104
- Slide 105
- Bonding in NO
- Slide 107
- Slide 108
- Slide 109
- Key Points of MO Theory ndash Heteronuclear Molecules
- Slide 111
- Slide 112
- Slide 113
- MO Theory Expectations
-
MOs OF XY MOLECULES
Equal or unequal e sharing between 2 atoms is reflected in the composition of the MOsmiddot When 2 atoms X and Y have the same electronegativity (purely
covalent bond) their overlapping AOs have the same energy and the bonding and anti-bonding MOs are each half X and half Y AO All electrons spend equal time near X and Y Examples N2 O2 F2
middot If EN(Y) gt EN(X) (polar covalent X+Y) the Y AO has lower energy than the X AO The bonding MO is more like the Y AO and the anti-bonding MO more like the X AO Bonding e spend more time near Y than X vice versa for anti-bonding e Example CO
MOs OF XY MOLECULES ____ s ___ ___ p ___ ___ ___2p ___ ___ ___ 2p uarr ____ s
Energy ___ ___ p ____ s ___ 2s ___ 2s ____ s
middot CO Bond Order = 30 (same as N2)middot CO Bond Energy = 1076 kJmol (N2 = 945 kJmol)middot Isoelectronic to CO and N2 CNndash NO+middot NO has 1endash in p bond order = 25 this endash is more on N than O
NO NO+ easyhellip
C Atom (4endash) Cδ+Oδndash (10endash) O Atom (6endash)
___ ___ ___ 2p ___ 2s
Electronegativity
Bonding in NObull Two possible Lewis dot structures
for NObull The simplest structure minimizes
formal charges and places the lone (unpaired) electron on the nitrogen
bull The Lewis structure predicts a bond order of 2 but experimental evidence suggests a bond order between 2 and 3
bull How does MO theory help us understand bonding in NO
N=O
N=O
+1-1
When the electronegativities of the 2 atoms are more similar the bonding
becomes less polar
EN(N) = 30EN(O) = 34
N=O
Electronegativity
2s2s
2p2p
NON OBond order = 25 unpaired electron is in a N-like orbital
NO+
oxidation
NO
NO is easily oxidized to form NO+ Why What changes can we predict in the bonding and magnetism of the molecule
Bond Order = (8-3)2 = 25Paramagnetic
Bond Order = (8-2)2 = 3Diamagnetic
s2s
s2s
p2p p2p
p2p
-3320
-1835
-1444 -1374
s2p -1307
-597 p2p (empty)
MO diagram for NO
Key Points of MO Theory ndash Heteronuclear Molecules
bull The more electronegative atom has orbitals lower in energy than the more positive atom
bull Electrons in bonding orbitals are closer to the more electronegative atom anti-bonding electrons are closer to the more positive atom
bull For most diatomic molecules s-p mixing changes the orbital energy levels but since these orbitals are almost always fully occupied their order is less important to us
Combining the Localized Electron and Molecular Orbital Models (into a convenient working
model)
Figure 1447
Only the p bonding changes between these resonance structures - The MO model describes this p bonding more
effectively
Figure 1451
Atomic Orbitals Molecular Orbitals
Another example Benzene
p bonding
s bonding
p atomic orbitals p molecular orbital
MO Theory Expectationsbull You should be able to
ndash predict which atomic orbitals are higher or lower in energy (based on electronegativity differences)
ndash correctly fill a molecular orbital diagramndash correctly calculate bond orderndash predict molecular magnetic properties
based on orbital occupationndash understand how molecular properties
change upon ionization (oxidation or reduction) of molecules
- What Do Molecules Look Like
- Slide 2
- Slide 3
- Arranging Electron Pairs
- Arranging Electron Pairs (cont)
- VSEPR Structure Guidelines
- VSEPR 2 electron pairs
- VSEPR 3 electron pairs
- VSEPR 4 electron pairs (cont)
- Bonding vs Lone pairs
- VSEPR 4 electron pairs
- VSEPR 4 electron pairs (cont) (2)
- VSEPR 4 electron pairs (cont) (3)
- Slide 14
- Molecular vs Electron-Pair Geometry
- Slide 16
- VSEPR Beyond the Octet
- VSEPR 5 electron pairs
- VSEPR 5 electron pairs (2)
- 5-electron-pair geometries
- VSEPR 6 electron pairs
- Slide 22
- Molecular Dipole Moments
- Slide 24
- Molecular Dipoles (cont)
- Molecular Dipole Example
- Advanced VSEPR Application
- The VSEPR Table
- The VSEPR Table (2)
- Slide 30
- Slide 31
- Slide 32
- Slide 33
- Slide 34
- Slide 35
- Slide 36
- Slide 37
- Slide 38
- Slide 39
- Slide 40
- Slide 41
- Slide 42
- Slide 43
- Slide 44
- Slide 45
- Slide 46
- Slide 47
- Slide 48
- Hybridization
- Hybridization (2)
- Hybridization (3)
- General Notes
- What is important to know about hybridization
- What is important to know about hybridization (2)
- What is important to know about hybridization (3)
- Slide 56
- Slide 57
- Molecular Orbital Theory
- Slide 59
- Constructive and Destructive Interference
- Slide 61
- MO Diagrams
- Slide 63
- Slide 64
- Review of Orbital Filling
- Slide 66
- Slide 67
- Filling Molecular Orbitals with Electrons
- Slide 69
- Slide 70
- Slide 71
- Forming Bonds
- Anti-bonding Orbitals
- MO Theory in Bonding
- Slide 75
- Slide 76
- Slide 77
- Slide 78
- Slide 79
- Slide 80
- Slide 81
- Slide 82
- Slide 83
- Slide 84
- Slide 85
- MO Diagram for H2 vs N2
- Slide 87
- Slide 88
- Slide 89
- Slide 90
- Molecular Orbitals of X2 Molecules
- Slide 92
- Slide 93
- Slide 94
- Slide 95
- Slide 96
- Slide 97
- Slide 98
- Slide 99
- Slide 100
- Slide 101
- Slide 102
- Slide 103
- Slide 104
- Slide 105
- Bonding in NO
- Slide 107
- Slide 108
- Slide 109
- Key Points of MO Theory ndash Heteronuclear Molecules
- Slide 111
- Slide 112
- Slide 113
- MO Theory Expectations
-
MOs OF XY MOLECULES ____ s ___ ___ p ___ ___ ___2p ___ ___ ___ 2p uarr ____ s
Energy ___ ___ p ____ s ___ 2s ___ 2s ____ s
middot CO Bond Order = 30 (same as N2)middot CO Bond Energy = 1076 kJmol (N2 = 945 kJmol)middot Isoelectronic to CO and N2 CNndash NO+middot NO has 1endash in p bond order = 25 this endash is more on N than O
NO NO+ easyhellip
C Atom (4endash) Cδ+Oδndash (10endash) O Atom (6endash)
___ ___ ___ 2p ___ 2s
Electronegativity
Bonding in NObull Two possible Lewis dot structures
for NObull The simplest structure minimizes
formal charges and places the lone (unpaired) electron on the nitrogen
bull The Lewis structure predicts a bond order of 2 but experimental evidence suggests a bond order between 2 and 3
bull How does MO theory help us understand bonding in NO
N=O
N=O
+1-1
When the electronegativities of the 2 atoms are more similar the bonding
becomes less polar
EN(N) = 30EN(O) = 34
N=O
Electronegativity
2s2s
2p2p
NON OBond order = 25 unpaired electron is in a N-like orbital
NO+
oxidation
NO
NO is easily oxidized to form NO+ Why What changes can we predict in the bonding and magnetism of the molecule
Bond Order = (8-3)2 = 25Paramagnetic
Bond Order = (8-2)2 = 3Diamagnetic
s2s
s2s
p2p p2p
p2p
-3320
-1835
-1444 -1374
s2p -1307
-597 p2p (empty)
MO diagram for NO
Key Points of MO Theory ndash Heteronuclear Molecules
bull The more electronegative atom has orbitals lower in energy than the more positive atom
bull Electrons in bonding orbitals are closer to the more electronegative atom anti-bonding electrons are closer to the more positive atom
bull For most diatomic molecules s-p mixing changes the orbital energy levels but since these orbitals are almost always fully occupied their order is less important to us
Combining the Localized Electron and Molecular Orbital Models (into a convenient working
model)
Figure 1447
Only the p bonding changes between these resonance structures - The MO model describes this p bonding more
effectively
Figure 1451
Atomic Orbitals Molecular Orbitals
Another example Benzene
p bonding
s bonding
p atomic orbitals p molecular orbital
MO Theory Expectationsbull You should be able to
ndash predict which atomic orbitals are higher or lower in energy (based on electronegativity differences)
ndash correctly fill a molecular orbital diagramndash correctly calculate bond orderndash predict molecular magnetic properties
based on orbital occupationndash understand how molecular properties
change upon ionization (oxidation or reduction) of molecules
- What Do Molecules Look Like
- Slide 2
- Slide 3
- Arranging Electron Pairs
- Arranging Electron Pairs (cont)
- VSEPR Structure Guidelines
- VSEPR 2 electron pairs
- VSEPR 3 electron pairs
- VSEPR 4 electron pairs (cont)
- Bonding vs Lone pairs
- VSEPR 4 electron pairs
- VSEPR 4 electron pairs (cont) (2)
- VSEPR 4 electron pairs (cont) (3)
- Slide 14
- Molecular vs Electron-Pair Geometry
- Slide 16
- VSEPR Beyond the Octet
- VSEPR 5 electron pairs
- VSEPR 5 electron pairs (2)
- 5-electron-pair geometries
- VSEPR 6 electron pairs
- Slide 22
- Molecular Dipole Moments
- Slide 24
- Molecular Dipoles (cont)
- Molecular Dipole Example
- Advanced VSEPR Application
- The VSEPR Table
- The VSEPR Table (2)
- Slide 30
- Slide 31
- Slide 32
- Slide 33
- Slide 34
- Slide 35
- Slide 36
- Slide 37
- Slide 38
- Slide 39
- Slide 40
- Slide 41
- Slide 42
- Slide 43
- Slide 44
- Slide 45
- Slide 46
- Slide 47
- Slide 48
- Hybridization
- Hybridization (2)
- Hybridization (3)
- General Notes
- What is important to know about hybridization
- What is important to know about hybridization (2)
- What is important to know about hybridization (3)
- Slide 56
- Slide 57
- Molecular Orbital Theory
- Slide 59
- Constructive and Destructive Interference
- Slide 61
- MO Diagrams
- Slide 63
- Slide 64
- Review of Orbital Filling
- Slide 66
- Slide 67
- Filling Molecular Orbitals with Electrons
- Slide 69
- Slide 70
- Slide 71
- Forming Bonds
- Anti-bonding Orbitals
- MO Theory in Bonding
- Slide 75
- Slide 76
- Slide 77
- Slide 78
- Slide 79
- Slide 80
- Slide 81
- Slide 82
- Slide 83
- Slide 84
- Slide 85
- MO Diagram for H2 vs N2
- Slide 87
- Slide 88
- Slide 89
- Slide 90
- Molecular Orbitals of X2 Molecules
- Slide 92
- Slide 93
- Slide 94
- Slide 95
- Slide 96
- Slide 97
- Slide 98
- Slide 99
- Slide 100
- Slide 101
- Slide 102
- Slide 103
- Slide 104
- Slide 105
- Bonding in NO
- Slide 107
- Slide 108
- Slide 109
- Key Points of MO Theory ndash Heteronuclear Molecules
- Slide 111
- Slide 112
- Slide 113
- MO Theory Expectations
-
Bonding in NObull Two possible Lewis dot structures
for NObull The simplest structure minimizes
formal charges and places the lone (unpaired) electron on the nitrogen
bull The Lewis structure predicts a bond order of 2 but experimental evidence suggests a bond order between 2 and 3
bull How does MO theory help us understand bonding in NO
N=O
N=O
+1-1
When the electronegativities of the 2 atoms are more similar the bonding
becomes less polar
EN(N) = 30EN(O) = 34
N=O
Electronegativity
2s2s
2p2p
NON OBond order = 25 unpaired electron is in a N-like orbital
NO+
oxidation
NO
NO is easily oxidized to form NO+ Why What changes can we predict in the bonding and magnetism of the molecule
Bond Order = (8-3)2 = 25Paramagnetic
Bond Order = (8-2)2 = 3Diamagnetic
s2s
s2s
p2p p2p
p2p
-3320
-1835
-1444 -1374
s2p -1307
-597 p2p (empty)
MO diagram for NO
Key Points of MO Theory ndash Heteronuclear Molecules
bull The more electronegative atom has orbitals lower in energy than the more positive atom
bull Electrons in bonding orbitals are closer to the more electronegative atom anti-bonding electrons are closer to the more positive atom
bull For most diatomic molecules s-p mixing changes the orbital energy levels but since these orbitals are almost always fully occupied their order is less important to us
Combining the Localized Electron and Molecular Orbital Models (into a convenient working
model)
Figure 1447
Only the p bonding changes between these resonance structures - The MO model describes this p bonding more
effectively
Figure 1451
Atomic Orbitals Molecular Orbitals
Another example Benzene
p bonding
s bonding
p atomic orbitals p molecular orbital
MO Theory Expectationsbull You should be able to
ndash predict which atomic orbitals are higher or lower in energy (based on electronegativity differences)
ndash correctly fill a molecular orbital diagramndash correctly calculate bond orderndash predict molecular magnetic properties
based on orbital occupationndash understand how molecular properties
change upon ionization (oxidation or reduction) of molecules
- What Do Molecules Look Like
- Slide 2
- Slide 3
- Arranging Electron Pairs
- Arranging Electron Pairs (cont)
- VSEPR Structure Guidelines
- VSEPR 2 electron pairs
- VSEPR 3 electron pairs
- VSEPR 4 electron pairs (cont)
- Bonding vs Lone pairs
- VSEPR 4 electron pairs
- VSEPR 4 electron pairs (cont) (2)
- VSEPR 4 electron pairs (cont) (3)
- Slide 14
- Molecular vs Electron-Pair Geometry
- Slide 16
- VSEPR Beyond the Octet
- VSEPR 5 electron pairs
- VSEPR 5 electron pairs (2)
- 5-electron-pair geometries
- VSEPR 6 electron pairs
- Slide 22
- Molecular Dipole Moments
- Slide 24
- Molecular Dipoles (cont)
- Molecular Dipole Example
- Advanced VSEPR Application
- The VSEPR Table
- The VSEPR Table (2)
- Slide 30
- Slide 31
- Slide 32
- Slide 33
- Slide 34
- Slide 35
- Slide 36
- Slide 37
- Slide 38
- Slide 39
- Slide 40
- Slide 41
- Slide 42
- Slide 43
- Slide 44
- Slide 45
- Slide 46
- Slide 47
- Slide 48
- Hybridization
- Hybridization (2)
- Hybridization (3)
- General Notes
- What is important to know about hybridization
- What is important to know about hybridization (2)
- What is important to know about hybridization (3)
- Slide 56
- Slide 57
- Molecular Orbital Theory
- Slide 59
- Constructive and Destructive Interference
- Slide 61
- MO Diagrams
- Slide 63
- Slide 64
- Review of Orbital Filling
- Slide 66
- Slide 67
- Filling Molecular Orbitals with Electrons
- Slide 69
- Slide 70
- Slide 71
- Forming Bonds
- Anti-bonding Orbitals
- MO Theory in Bonding
- Slide 75
- Slide 76
- Slide 77
- Slide 78
- Slide 79
- Slide 80
- Slide 81
- Slide 82
- Slide 83
- Slide 84
- Slide 85
- MO Diagram for H2 vs N2
- Slide 87
- Slide 88
- Slide 89
- Slide 90
- Molecular Orbitals of X2 Molecules
- Slide 92
- Slide 93
- Slide 94
- Slide 95
- Slide 96
- Slide 97
- Slide 98
- Slide 99
- Slide 100
- Slide 101
- Slide 102
- Slide 103
- Slide 104
- Slide 105
- Bonding in NO
- Slide 107
- Slide 108
- Slide 109
- Key Points of MO Theory ndash Heteronuclear Molecules
- Slide 111
- Slide 112
- Slide 113
- MO Theory Expectations
-
When the electronegativities of the 2 atoms are more similar the bonding
becomes less polar
EN(N) = 30EN(O) = 34
N=O
Electronegativity
2s2s
2p2p
NON OBond order = 25 unpaired electron is in a N-like orbital
NO+
oxidation
NO
NO is easily oxidized to form NO+ Why What changes can we predict in the bonding and magnetism of the molecule
Bond Order = (8-3)2 = 25Paramagnetic
Bond Order = (8-2)2 = 3Diamagnetic
s2s
s2s
p2p p2p
p2p
-3320
-1835
-1444 -1374
s2p -1307
-597 p2p (empty)
MO diagram for NO
Key Points of MO Theory ndash Heteronuclear Molecules
bull The more electronegative atom has orbitals lower in energy than the more positive atom
bull Electrons in bonding orbitals are closer to the more electronegative atom anti-bonding electrons are closer to the more positive atom
bull For most diatomic molecules s-p mixing changes the orbital energy levels but since these orbitals are almost always fully occupied their order is less important to us
Combining the Localized Electron and Molecular Orbital Models (into a convenient working
model)
Figure 1447
Only the p bonding changes between these resonance structures - The MO model describes this p bonding more
effectively
Figure 1451
Atomic Orbitals Molecular Orbitals
Another example Benzene
p bonding
s bonding
p atomic orbitals p molecular orbital
MO Theory Expectationsbull You should be able to
ndash predict which atomic orbitals are higher or lower in energy (based on electronegativity differences)
ndash correctly fill a molecular orbital diagramndash correctly calculate bond orderndash predict molecular magnetic properties
based on orbital occupationndash understand how molecular properties
change upon ionization (oxidation or reduction) of molecules
- What Do Molecules Look Like
- Slide 2
- Slide 3
- Arranging Electron Pairs
- Arranging Electron Pairs (cont)
- VSEPR Structure Guidelines
- VSEPR 2 electron pairs
- VSEPR 3 electron pairs
- VSEPR 4 electron pairs (cont)
- Bonding vs Lone pairs
- VSEPR 4 electron pairs
- VSEPR 4 electron pairs (cont) (2)
- VSEPR 4 electron pairs (cont) (3)
- Slide 14
- Molecular vs Electron-Pair Geometry
- Slide 16
- VSEPR Beyond the Octet
- VSEPR 5 electron pairs
- VSEPR 5 electron pairs (2)
- 5-electron-pair geometries
- VSEPR 6 electron pairs
- Slide 22
- Molecular Dipole Moments
- Slide 24
- Molecular Dipoles (cont)
- Molecular Dipole Example
- Advanced VSEPR Application
- The VSEPR Table
- The VSEPR Table (2)
- Slide 30
- Slide 31
- Slide 32
- Slide 33
- Slide 34
- Slide 35
- Slide 36
- Slide 37
- Slide 38
- Slide 39
- Slide 40
- Slide 41
- Slide 42
- Slide 43
- Slide 44
- Slide 45
- Slide 46
- Slide 47
- Slide 48
- Hybridization
- Hybridization (2)
- Hybridization (3)
- General Notes
- What is important to know about hybridization
- What is important to know about hybridization (2)
- What is important to know about hybridization (3)
- Slide 56
- Slide 57
- Molecular Orbital Theory
- Slide 59
- Constructive and Destructive Interference
- Slide 61
- MO Diagrams
- Slide 63
- Slide 64
- Review of Orbital Filling
- Slide 66
- Slide 67
- Filling Molecular Orbitals with Electrons
- Slide 69
- Slide 70
- Slide 71
- Forming Bonds
- Anti-bonding Orbitals
- MO Theory in Bonding
- Slide 75
- Slide 76
- Slide 77
- Slide 78
- Slide 79
- Slide 80
- Slide 81
- Slide 82
- Slide 83
- Slide 84
- Slide 85
- MO Diagram for H2 vs N2
- Slide 87
- Slide 88
- Slide 89
- Slide 90
- Molecular Orbitals of X2 Molecules
- Slide 92
- Slide 93
- Slide 94
- Slide 95
- Slide 96
- Slide 97
- Slide 98
- Slide 99
- Slide 100
- Slide 101
- Slide 102
- Slide 103
- Slide 104
- Slide 105
- Bonding in NO
- Slide 107
- Slide 108
- Slide 109
- Key Points of MO Theory ndash Heteronuclear Molecules
- Slide 111
- Slide 112
- Slide 113
- MO Theory Expectations
-
NO+
oxidation
NO
NO is easily oxidized to form NO+ Why What changes can we predict in the bonding and magnetism of the molecule
Bond Order = (8-3)2 = 25Paramagnetic
Bond Order = (8-2)2 = 3Diamagnetic
s2s
s2s
p2p p2p
p2p
-3320
-1835
-1444 -1374
s2p -1307
-597 p2p (empty)
MO diagram for NO
Key Points of MO Theory ndash Heteronuclear Molecules
bull The more electronegative atom has orbitals lower in energy than the more positive atom
bull Electrons in bonding orbitals are closer to the more electronegative atom anti-bonding electrons are closer to the more positive atom
bull For most diatomic molecules s-p mixing changes the orbital energy levels but since these orbitals are almost always fully occupied their order is less important to us
Combining the Localized Electron and Molecular Orbital Models (into a convenient working
model)
Figure 1447
Only the p bonding changes between these resonance structures - The MO model describes this p bonding more
effectively
Figure 1451
Atomic Orbitals Molecular Orbitals
Another example Benzene
p bonding
s bonding
p atomic orbitals p molecular orbital
MO Theory Expectationsbull You should be able to
ndash predict which atomic orbitals are higher or lower in energy (based on electronegativity differences)
ndash correctly fill a molecular orbital diagramndash correctly calculate bond orderndash predict molecular magnetic properties
based on orbital occupationndash understand how molecular properties
change upon ionization (oxidation or reduction) of molecules
- What Do Molecules Look Like
- Slide 2
- Slide 3
- Arranging Electron Pairs
- Arranging Electron Pairs (cont)
- VSEPR Structure Guidelines
- VSEPR 2 electron pairs
- VSEPR 3 electron pairs
- VSEPR 4 electron pairs (cont)
- Bonding vs Lone pairs
- VSEPR 4 electron pairs
- VSEPR 4 electron pairs (cont) (2)
- VSEPR 4 electron pairs (cont) (3)
- Slide 14
- Molecular vs Electron-Pair Geometry
- Slide 16
- VSEPR Beyond the Octet
- VSEPR 5 electron pairs
- VSEPR 5 electron pairs (2)
- 5-electron-pair geometries
- VSEPR 6 electron pairs
- Slide 22
- Molecular Dipole Moments
- Slide 24
- Molecular Dipoles (cont)
- Molecular Dipole Example
- Advanced VSEPR Application
- The VSEPR Table
- The VSEPR Table (2)
- Slide 30
- Slide 31
- Slide 32
- Slide 33
- Slide 34
- Slide 35
- Slide 36
- Slide 37
- Slide 38
- Slide 39
- Slide 40
- Slide 41
- Slide 42
- Slide 43
- Slide 44
- Slide 45
- Slide 46
- Slide 47
- Slide 48
- Hybridization
- Hybridization (2)
- Hybridization (3)
- General Notes
- What is important to know about hybridization
- What is important to know about hybridization (2)
- What is important to know about hybridization (3)
- Slide 56
- Slide 57
- Molecular Orbital Theory
- Slide 59
- Constructive and Destructive Interference
- Slide 61
- MO Diagrams
- Slide 63
- Slide 64
- Review of Orbital Filling
- Slide 66
- Slide 67
- Filling Molecular Orbitals with Electrons
- Slide 69
- Slide 70
- Slide 71
- Forming Bonds
- Anti-bonding Orbitals
- MO Theory in Bonding
- Slide 75
- Slide 76
- Slide 77
- Slide 78
- Slide 79
- Slide 80
- Slide 81
- Slide 82
- Slide 83
- Slide 84
- Slide 85
- MO Diagram for H2 vs N2
- Slide 87
- Slide 88
- Slide 89
- Slide 90
- Molecular Orbitals of X2 Molecules
- Slide 92
- Slide 93
- Slide 94
- Slide 95
- Slide 96
- Slide 97
- Slide 98
- Slide 99
- Slide 100
- Slide 101
- Slide 102
- Slide 103
- Slide 104
- Slide 105
- Bonding in NO
- Slide 107
- Slide 108
- Slide 109
- Key Points of MO Theory ndash Heteronuclear Molecules
- Slide 111
- Slide 112
- Slide 113
- MO Theory Expectations
-
s2s
s2s
p2p p2p
p2p
-3320
-1835
-1444 -1374
s2p -1307
-597 p2p (empty)
MO diagram for NO
Key Points of MO Theory ndash Heteronuclear Molecules
bull The more electronegative atom has orbitals lower in energy than the more positive atom
bull Electrons in bonding orbitals are closer to the more electronegative atom anti-bonding electrons are closer to the more positive atom
bull For most diatomic molecules s-p mixing changes the orbital energy levels but since these orbitals are almost always fully occupied their order is less important to us
Combining the Localized Electron and Molecular Orbital Models (into a convenient working
model)
Figure 1447
Only the p bonding changes between these resonance structures - The MO model describes this p bonding more
effectively
Figure 1451
Atomic Orbitals Molecular Orbitals
Another example Benzene
p bonding
s bonding
p atomic orbitals p molecular orbital
MO Theory Expectationsbull You should be able to
ndash predict which atomic orbitals are higher or lower in energy (based on electronegativity differences)
ndash correctly fill a molecular orbital diagramndash correctly calculate bond orderndash predict molecular magnetic properties
based on orbital occupationndash understand how molecular properties
change upon ionization (oxidation or reduction) of molecules
- What Do Molecules Look Like
- Slide 2
- Slide 3
- Arranging Electron Pairs
- Arranging Electron Pairs (cont)
- VSEPR Structure Guidelines
- VSEPR 2 electron pairs
- VSEPR 3 electron pairs
- VSEPR 4 electron pairs (cont)
- Bonding vs Lone pairs
- VSEPR 4 electron pairs
- VSEPR 4 electron pairs (cont) (2)
- VSEPR 4 electron pairs (cont) (3)
- Slide 14
- Molecular vs Electron-Pair Geometry
- Slide 16
- VSEPR Beyond the Octet
- VSEPR 5 electron pairs
- VSEPR 5 electron pairs (2)
- 5-electron-pair geometries
- VSEPR 6 electron pairs
- Slide 22
- Molecular Dipole Moments
- Slide 24
- Molecular Dipoles (cont)
- Molecular Dipole Example
- Advanced VSEPR Application
- The VSEPR Table
- The VSEPR Table (2)
- Slide 30
- Slide 31
- Slide 32
- Slide 33
- Slide 34
- Slide 35
- Slide 36
- Slide 37
- Slide 38
- Slide 39
- Slide 40
- Slide 41
- Slide 42
- Slide 43
- Slide 44
- Slide 45
- Slide 46
- Slide 47
- Slide 48
- Hybridization
- Hybridization (2)
- Hybridization (3)
- General Notes
- What is important to know about hybridization
- What is important to know about hybridization (2)
- What is important to know about hybridization (3)
- Slide 56
- Slide 57
- Molecular Orbital Theory
- Slide 59
- Constructive and Destructive Interference
- Slide 61
- MO Diagrams
- Slide 63
- Slide 64
- Review of Orbital Filling
- Slide 66
- Slide 67
- Filling Molecular Orbitals with Electrons
- Slide 69
- Slide 70
- Slide 71
- Forming Bonds
- Anti-bonding Orbitals
- MO Theory in Bonding
- Slide 75
- Slide 76
- Slide 77
- Slide 78
- Slide 79
- Slide 80
- Slide 81
- Slide 82
- Slide 83
- Slide 84
- Slide 85
- MO Diagram for H2 vs N2
- Slide 87
- Slide 88
- Slide 89
- Slide 90
- Molecular Orbitals of X2 Molecules
- Slide 92
- Slide 93
- Slide 94
- Slide 95
- Slide 96
- Slide 97
- Slide 98
- Slide 99
- Slide 100
- Slide 101
- Slide 102
- Slide 103
- Slide 104
- Slide 105
- Bonding in NO
- Slide 107
- Slide 108
- Slide 109
- Key Points of MO Theory ndash Heteronuclear Molecules
- Slide 111
- Slide 112
- Slide 113
- MO Theory Expectations
-
Key Points of MO Theory ndash Heteronuclear Molecules
bull The more electronegative atom has orbitals lower in energy than the more positive atom
bull Electrons in bonding orbitals are closer to the more electronegative atom anti-bonding electrons are closer to the more positive atom
bull For most diatomic molecules s-p mixing changes the orbital energy levels but since these orbitals are almost always fully occupied their order is less important to us
Combining the Localized Electron and Molecular Orbital Models (into a convenient working
model)
Figure 1447
Only the p bonding changes between these resonance structures - The MO model describes this p bonding more
effectively
Figure 1451
Atomic Orbitals Molecular Orbitals
Another example Benzene
p bonding
s bonding
p atomic orbitals p molecular orbital
MO Theory Expectationsbull You should be able to
ndash predict which atomic orbitals are higher or lower in energy (based on electronegativity differences)
ndash correctly fill a molecular orbital diagramndash correctly calculate bond orderndash predict molecular magnetic properties
based on orbital occupationndash understand how molecular properties
change upon ionization (oxidation or reduction) of molecules
- What Do Molecules Look Like
- Slide 2
- Slide 3
- Arranging Electron Pairs
- Arranging Electron Pairs (cont)
- VSEPR Structure Guidelines
- VSEPR 2 electron pairs
- VSEPR 3 electron pairs
- VSEPR 4 electron pairs (cont)
- Bonding vs Lone pairs
- VSEPR 4 electron pairs
- VSEPR 4 electron pairs (cont) (2)
- VSEPR 4 electron pairs (cont) (3)
- Slide 14
- Molecular vs Electron-Pair Geometry
- Slide 16
- VSEPR Beyond the Octet
- VSEPR 5 electron pairs
- VSEPR 5 electron pairs (2)
- 5-electron-pair geometries
- VSEPR 6 electron pairs
- Slide 22
- Molecular Dipole Moments
- Slide 24
- Molecular Dipoles (cont)
- Molecular Dipole Example
- Advanced VSEPR Application
- The VSEPR Table
- The VSEPR Table (2)
- Slide 30
- Slide 31
- Slide 32
- Slide 33
- Slide 34
- Slide 35
- Slide 36
- Slide 37
- Slide 38
- Slide 39
- Slide 40
- Slide 41
- Slide 42
- Slide 43
- Slide 44
- Slide 45
- Slide 46
- Slide 47
- Slide 48
- Hybridization
- Hybridization (2)
- Hybridization (3)
- General Notes
- What is important to know about hybridization
- What is important to know about hybridization (2)
- What is important to know about hybridization (3)
- Slide 56
- Slide 57
- Molecular Orbital Theory
- Slide 59
- Constructive and Destructive Interference
- Slide 61
- MO Diagrams
- Slide 63
- Slide 64
- Review of Orbital Filling
- Slide 66
- Slide 67
- Filling Molecular Orbitals with Electrons
- Slide 69
- Slide 70
- Slide 71
- Forming Bonds
- Anti-bonding Orbitals
- MO Theory in Bonding
- Slide 75
- Slide 76
- Slide 77
- Slide 78
- Slide 79
- Slide 80
- Slide 81
- Slide 82
- Slide 83
- Slide 84
- Slide 85
- MO Diagram for H2 vs N2
- Slide 87
- Slide 88
- Slide 89
- Slide 90
- Molecular Orbitals of X2 Molecules
- Slide 92
- Slide 93
- Slide 94
- Slide 95
- Slide 96
- Slide 97
- Slide 98
- Slide 99
- Slide 100
- Slide 101
- Slide 102
- Slide 103
- Slide 104
- Slide 105
- Bonding in NO
- Slide 107
- Slide 108
- Slide 109
- Key Points of MO Theory ndash Heteronuclear Molecules
- Slide 111
- Slide 112
- Slide 113
- MO Theory Expectations
-
Combining the Localized Electron and Molecular Orbital Models (into a convenient working
model)
Figure 1447
Only the p bonding changes between these resonance structures - The MO model describes this p bonding more
effectively
Figure 1451
Atomic Orbitals Molecular Orbitals
Another example Benzene
p bonding
s bonding
p atomic orbitals p molecular orbital
MO Theory Expectationsbull You should be able to
ndash predict which atomic orbitals are higher or lower in energy (based on electronegativity differences)
ndash correctly fill a molecular orbital diagramndash correctly calculate bond orderndash predict molecular magnetic properties
based on orbital occupationndash understand how molecular properties
change upon ionization (oxidation or reduction) of molecules
- What Do Molecules Look Like
- Slide 2
- Slide 3
- Arranging Electron Pairs
- Arranging Electron Pairs (cont)
- VSEPR Structure Guidelines
- VSEPR 2 electron pairs
- VSEPR 3 electron pairs
- VSEPR 4 electron pairs (cont)
- Bonding vs Lone pairs
- VSEPR 4 electron pairs
- VSEPR 4 electron pairs (cont) (2)
- VSEPR 4 electron pairs (cont) (3)
- Slide 14
- Molecular vs Electron-Pair Geometry
- Slide 16
- VSEPR Beyond the Octet
- VSEPR 5 electron pairs
- VSEPR 5 electron pairs (2)
- 5-electron-pair geometries
- VSEPR 6 electron pairs
- Slide 22
- Molecular Dipole Moments
- Slide 24
- Molecular Dipoles (cont)
- Molecular Dipole Example
- Advanced VSEPR Application
- The VSEPR Table
- The VSEPR Table (2)
- Slide 30
- Slide 31
- Slide 32
- Slide 33
- Slide 34
- Slide 35
- Slide 36
- Slide 37
- Slide 38
- Slide 39
- Slide 40
- Slide 41
- Slide 42
- Slide 43
- Slide 44
- Slide 45
- Slide 46
- Slide 47
- Slide 48
- Hybridization
- Hybridization (2)
- Hybridization (3)
- General Notes
- What is important to know about hybridization
- What is important to know about hybridization (2)
- What is important to know about hybridization (3)
- Slide 56
- Slide 57
- Molecular Orbital Theory
- Slide 59
- Constructive and Destructive Interference
- Slide 61
- MO Diagrams
- Slide 63
- Slide 64
- Review of Orbital Filling
- Slide 66
- Slide 67
- Filling Molecular Orbitals with Electrons
- Slide 69
- Slide 70
- Slide 71
- Forming Bonds
- Anti-bonding Orbitals
- MO Theory in Bonding
- Slide 75
- Slide 76
- Slide 77
- Slide 78
- Slide 79
- Slide 80
- Slide 81
- Slide 82
- Slide 83
- Slide 84
- Slide 85
- MO Diagram for H2 vs N2
- Slide 87
- Slide 88
- Slide 89
- Slide 90
- Molecular Orbitals of X2 Molecules
- Slide 92
- Slide 93
- Slide 94
- Slide 95
- Slide 96
- Slide 97
- Slide 98
- Slide 99
- Slide 100
- Slide 101
- Slide 102
- Slide 103
- Slide 104
- Slide 105
- Bonding in NO
- Slide 107
- Slide 108
- Slide 109
- Key Points of MO Theory ndash Heteronuclear Molecules
- Slide 111
- Slide 112
- Slide 113
- MO Theory Expectations
-
Figure 1451
Atomic Orbitals Molecular Orbitals
Another example Benzene
p bonding
s bonding
p atomic orbitals p molecular orbital
MO Theory Expectationsbull You should be able to
ndash predict which atomic orbitals are higher or lower in energy (based on electronegativity differences)
ndash correctly fill a molecular orbital diagramndash correctly calculate bond orderndash predict molecular magnetic properties
based on orbital occupationndash understand how molecular properties
change upon ionization (oxidation or reduction) of molecules
- What Do Molecules Look Like
- Slide 2
- Slide 3
- Arranging Electron Pairs
- Arranging Electron Pairs (cont)
- VSEPR Structure Guidelines
- VSEPR 2 electron pairs
- VSEPR 3 electron pairs
- VSEPR 4 electron pairs (cont)
- Bonding vs Lone pairs
- VSEPR 4 electron pairs
- VSEPR 4 electron pairs (cont) (2)
- VSEPR 4 electron pairs (cont) (3)
- Slide 14
- Molecular vs Electron-Pair Geometry
- Slide 16
- VSEPR Beyond the Octet
- VSEPR 5 electron pairs
- VSEPR 5 electron pairs (2)
- 5-electron-pair geometries
- VSEPR 6 electron pairs
- Slide 22
- Molecular Dipole Moments
- Slide 24
- Molecular Dipoles (cont)
- Molecular Dipole Example
- Advanced VSEPR Application
- The VSEPR Table
- The VSEPR Table (2)
- Slide 30
- Slide 31
- Slide 32
- Slide 33
- Slide 34
- Slide 35
- Slide 36
- Slide 37
- Slide 38
- Slide 39
- Slide 40
- Slide 41
- Slide 42
- Slide 43
- Slide 44
- Slide 45
- Slide 46
- Slide 47
- Slide 48
- Hybridization
- Hybridization (2)
- Hybridization (3)
- General Notes
- What is important to know about hybridization
- What is important to know about hybridization (2)
- What is important to know about hybridization (3)
- Slide 56
- Slide 57
- Molecular Orbital Theory
- Slide 59
- Constructive and Destructive Interference
- Slide 61
- MO Diagrams
- Slide 63
- Slide 64
- Review of Orbital Filling
- Slide 66
- Slide 67
- Filling Molecular Orbitals with Electrons
- Slide 69
- Slide 70
- Slide 71
- Forming Bonds
- Anti-bonding Orbitals
- MO Theory in Bonding
- Slide 75
- Slide 76
- Slide 77
- Slide 78
- Slide 79
- Slide 80
- Slide 81
- Slide 82
- Slide 83
- Slide 84
- Slide 85
- MO Diagram for H2 vs N2
- Slide 87
- Slide 88
- Slide 89
- Slide 90
- Molecular Orbitals of X2 Molecules
- Slide 92
- Slide 93
- Slide 94
- Slide 95
- Slide 96
- Slide 97
- Slide 98
- Slide 99
- Slide 100
- Slide 101
- Slide 102
- Slide 103
- Slide 104
- Slide 105
- Bonding in NO
- Slide 107
- Slide 108
- Slide 109
- Key Points of MO Theory ndash Heteronuclear Molecules
- Slide 111
- Slide 112
- Slide 113
- MO Theory Expectations
-
Another example Benzene
p bonding
s bonding
p atomic orbitals p molecular orbital
MO Theory Expectationsbull You should be able to
ndash predict which atomic orbitals are higher or lower in energy (based on electronegativity differences)
ndash correctly fill a molecular orbital diagramndash correctly calculate bond orderndash predict molecular magnetic properties
based on orbital occupationndash understand how molecular properties
change upon ionization (oxidation or reduction) of molecules
- What Do Molecules Look Like
- Slide 2
- Slide 3
- Arranging Electron Pairs
- Arranging Electron Pairs (cont)
- VSEPR Structure Guidelines
- VSEPR 2 electron pairs
- VSEPR 3 electron pairs
- VSEPR 4 electron pairs (cont)
- Bonding vs Lone pairs
- VSEPR 4 electron pairs
- VSEPR 4 electron pairs (cont) (2)
- VSEPR 4 electron pairs (cont) (3)
- Slide 14
- Molecular vs Electron-Pair Geometry
- Slide 16
- VSEPR Beyond the Octet
- VSEPR 5 electron pairs
- VSEPR 5 electron pairs (2)
- 5-electron-pair geometries
- VSEPR 6 electron pairs
- Slide 22
- Molecular Dipole Moments
- Slide 24
- Molecular Dipoles (cont)
- Molecular Dipole Example
- Advanced VSEPR Application
- The VSEPR Table
- The VSEPR Table (2)
- Slide 30
- Slide 31
- Slide 32
- Slide 33
- Slide 34
- Slide 35
- Slide 36
- Slide 37
- Slide 38
- Slide 39
- Slide 40
- Slide 41
- Slide 42
- Slide 43
- Slide 44
- Slide 45
- Slide 46
- Slide 47
- Slide 48
- Hybridization
- Hybridization (2)
- Hybridization (3)
- General Notes
- What is important to know about hybridization
- What is important to know about hybridization (2)
- What is important to know about hybridization (3)
- Slide 56
- Slide 57
- Molecular Orbital Theory
- Slide 59
- Constructive and Destructive Interference
- Slide 61
- MO Diagrams
- Slide 63
- Slide 64
- Review of Orbital Filling
- Slide 66
- Slide 67
- Filling Molecular Orbitals with Electrons
- Slide 69
- Slide 70
- Slide 71
- Forming Bonds
- Anti-bonding Orbitals
- MO Theory in Bonding
- Slide 75
- Slide 76
- Slide 77
- Slide 78
- Slide 79
- Slide 80
- Slide 81
- Slide 82
- Slide 83
- Slide 84
- Slide 85
- MO Diagram for H2 vs N2
- Slide 87
- Slide 88
- Slide 89
- Slide 90
- Molecular Orbitals of X2 Molecules
- Slide 92
- Slide 93
- Slide 94
- Slide 95
- Slide 96
- Slide 97
- Slide 98
- Slide 99
- Slide 100
- Slide 101
- Slide 102
- Slide 103
- Slide 104
- Slide 105
- Bonding in NO
- Slide 107
- Slide 108
- Slide 109
- Key Points of MO Theory ndash Heteronuclear Molecules
- Slide 111
- Slide 112
- Slide 113
- MO Theory Expectations
-
MO Theory Expectationsbull You should be able to
ndash predict which atomic orbitals are higher or lower in energy (based on electronegativity differences)
ndash correctly fill a molecular orbital diagramndash correctly calculate bond orderndash predict molecular magnetic properties
based on orbital occupationndash understand how molecular properties
change upon ionization (oxidation or reduction) of molecules
- What Do Molecules Look Like
- Slide 2
- Slide 3
- Arranging Electron Pairs
- Arranging Electron Pairs (cont)
- VSEPR Structure Guidelines
- VSEPR 2 electron pairs
- VSEPR 3 electron pairs
- VSEPR 4 electron pairs (cont)
- Bonding vs Lone pairs
- VSEPR 4 electron pairs
- VSEPR 4 electron pairs (cont) (2)
- VSEPR 4 electron pairs (cont) (3)
- Slide 14
- Molecular vs Electron-Pair Geometry
- Slide 16
- VSEPR Beyond the Octet
- VSEPR 5 electron pairs
- VSEPR 5 electron pairs (2)
- 5-electron-pair geometries
- VSEPR 6 electron pairs
- Slide 22
- Molecular Dipole Moments
- Slide 24
- Molecular Dipoles (cont)
- Molecular Dipole Example
- Advanced VSEPR Application
- The VSEPR Table
- The VSEPR Table (2)
- Slide 30
- Slide 31
- Slide 32
- Slide 33
- Slide 34
- Slide 35
- Slide 36
- Slide 37
- Slide 38
- Slide 39
- Slide 40
- Slide 41
- Slide 42
- Slide 43
- Slide 44
- Slide 45
- Slide 46
- Slide 47
- Slide 48
- Hybridization
- Hybridization (2)
- Hybridization (3)
- General Notes
- What is important to know about hybridization
- What is important to know about hybridization (2)
- What is important to know about hybridization (3)
- Slide 56
- Slide 57
- Molecular Orbital Theory
- Slide 59
- Constructive and Destructive Interference
- Slide 61
- MO Diagrams
- Slide 63
- Slide 64
- Review of Orbital Filling
- Slide 66
- Slide 67
- Filling Molecular Orbitals with Electrons
- Slide 69
- Slide 70
- Slide 71
- Forming Bonds
- Anti-bonding Orbitals
- MO Theory in Bonding
- Slide 75
- Slide 76
- Slide 77
- Slide 78
- Slide 79
- Slide 80
- Slide 81
- Slide 82
- Slide 83
- Slide 84
- Slide 85
- MO Diagram for H2 vs N2
- Slide 87
- Slide 88
- Slide 89
- Slide 90
- Molecular Orbitals of X2 Molecules
- Slide 92
- Slide 93
- Slide 94
- Slide 95
- Slide 96
- Slide 97
- Slide 98
- Slide 99
- Slide 100
- Slide 101
- Slide 102
- Slide 103
- Slide 104
- Slide 105
- Bonding in NO
- Slide 107
- Slide 108
- Slide 109
- Key Points of MO Theory ndash Heteronuclear Molecules
- Slide 111
- Slide 112
- Slide 113
- MO Theory Expectations
-