chemistry xl-14a molecular shape and structure
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CHEMISTRY XL-14A MOLECULAR SHAPE AND STRUCTURE. July 23, 2011Robert Iafe. Valence shell electron pair repulsion theory (VSEPR). Lewis Theory Connectivity, electron tracking VSEPR Theory 3-D Structure around an atom Valence Bond Theory Molecular Orbital Theory. - PowerPoint PPT PresentationTRANSCRIPT
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CHEMISTRY XL-14A
MOLECULAR SHAPE
AND STRUCTU
REJuly 23, 2011 Robert Iafe
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Valence shell electron pair repulsion theory (VSEPR)
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Lewis Theory Connectivity, electron tracking
VSEPR Theory 3-D Structure around an atom
Valence Bond Theory Molecular Orbital Theory
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Valence shell electron pair repulsion theory (VSEPR)
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Valence-Shell Electron-Pair Repulsion (VSEPR) Accounts for 3D shapes of molecules Based on electron-electron repulsion Determine bond angles shape
Rules are based on experimental observation:
1. Areas of electron concentration (bonds and lone pairs) around the central atom repel each other.
2. Bonds and lone pairs stay as far away from each other as possible (without changing distance)
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Valence shell electron pair repulsion theory (VSEPR)
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Valence shell electron pair repulsion theory (VSEPR)
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Valence-Shell Electron-Pair Repulsion (VSEPR) – Accounts for 3D shapes of molecules in terms of electron-electron repulsion
Rules are based on experimental observation:
1. Regions of high electron concentration around the central atom (bonds and lone pairs) repel each other.
2. Single and Multiple bonds treated the same
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Coordination Number6
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Ethanol7
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AXE Model8
A = Central Atom X = Atomic Substituents E = Lone Pairs (If none, not shown)
AXnEm
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VESPR Base Shapes9
Linear: AX2 – (2 e- groups)
Cl ClBe
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VESPR Base Shapes10
Trigonal Planar: AX3
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VESPR Base Shapes11
Tetrahedral: AX4
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VESPR Base Shapes12
Trigonal Bipyramidal: AX5
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VESPR Base Shapes13
Octahedral: AX6
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Octahedral – cis/trans14
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VESPR Base Shapes15
Pentagonal Bipyramidal: AX7
F FI
FFF
FF
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Bond Angles16
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AX3 Derivatives17
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AX4 Derivatives18
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AX5 Derivatives19
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AX6 Derivatives20
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2 3 4 5 6
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Benzene – C6H622
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Cyclohexane – C6H1223
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Polar Molecules24
A polar covalent bond has a non-zero dipole moment
Polar molecules – molecules with non-zero dipole moment: ≠ 0 D
Examples: HCl HF HBr
A non-polar molecule has no dipole moment: = 0 D
Examples: O2 CH4 SF6
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Polar Molecules25
Polar bonds, Non-polar molecule?
If the dipole moments cancel out (Vector sum = 0):
Each C-O bond is polar
Two equally polar bonds, with opposing dipole moments
CO2
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Polar Molecules26
H2OEach O-H bond is polar
Overall molecule is ??????
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Polar Molecules27
C2H2Cl2
cis-CH2Cl2 trans-CH2Cl2
non-polar
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Polar Molecules28
non-polar polar
CCl4 CHCl3
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Valence Bond (VB) Theory29
Lewis Theory Connectivity, electron tracking
VSEPR Theory 3-D Structure around an atom
Valence Bond Theory Extended 3-D Structure Information Delocalization in Molecules Illustrates Multiple Bonding Prediction of Reactivity
Molecular Orbital Theory
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Answers to HΨ=EΨ30
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Answers to HΨ=EΨ31
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Bonding With Orbitals32
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Bonding with Orbitals33
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Valence-Bond Theory34
So far, we’ve been thinking of molecules using Lewis’ Theory:
Bonding electrons are located in between bonded atoms – electrons are localized
But….
We learned in Ch. 1 that we learned to think of electrons as wavefunctions, which are described by atomic orbitals
Valence Bond Theory – quantum mechanical view of bonding
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‘Types’ of Bonds – σ bond35
Lets start with H2, the simplest molecule:
Ground state H has one 1s electron
When the 2 H atoms bond, the atomic orbitals merge, forming a σ-bond
σbond (sigma bond) – along bond axis.
We say the atomic orbitals overlap
More overlap = Stronger bond
All single covalent bonds are σ-bonds
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‘Types’ of Bonds – σ bond36
All single covalent bonds are σ-bonds
Can have σ-bonds between any types
of orbitals:
Two s orbitals
Two p orbitals
An s and a p orbital
Etc….
Example: σ-bond between 1s orbital of H and 2pz orbital of F
H –– F
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‘Types’ of Bonds – Multiple bonds
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What happens in N2?
Remember the bond in N2 is a triple bond…
Each N atom has 3 unpaired 2p electrons
Lets look at the atomic orbitals of N:
pzpypx
N –– N
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‘Types’ of Bonds – Multiple bonds
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What happens in N2?
Remember the bond in N2 is a triple bond…
Each N atom has 3 unpaired 2p electrons
How do the atomic orbitals bond?
pz pz
Between the 2pz and 2pz orbitals, we have bonding on the bond axis
This is a σ-bond!
N –– N
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‘Types’ of Bonds – Multiple bonds
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What happens in N2?
Remember the bond in N2 is a triple bond…
Each N atom has 3 unpaired 2p electrons
How do the atomic orbitals bond?
Between 2px and 2px orbitals, the bonding is not on the bond axis
A different type of bond!
-bond –nodal plane along bond axis
Bonding occurs above and below the bond axis
px px
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Bonding of N240
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Bonding of N241
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Hybridization of Orbitals42
What about methane (CH4)?
According to Valence Bond Theory:
C should only make 2 bonds!
But we know that C can make 4 bonds
And CH4 has a tetrahedral shape according to VSEPR:
How do get a tetrahedral shape from the 2s and 2p orbitals?
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Experiments43
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Hybridization of Orbitals44
We solve the 4 bond problem by promoting an electron:
It takes energy to promote an electron to a higher E orbital
But, overall the energy is lower if C can make 4 bonds instead of 2.
Ok, what about the geometry problem?
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Hybridization of Orbitals45
By promoting an electron, we can now make 4 bonds
But the geometry of the px, py, and pz orbitals don’t match the tetrahedral shape of CH4
s and p orbitals are described by a wave-like model of the e
If we think of the orbitals as interfering with each other, we can define new hybrid orbitals:
h1 = s + px + py + pz
h2 = s – px – py + pz
h3 = s – px + py – pz
h4 = s + px – py – pz
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sp3 Hybridization46
C in CH4 uses 4 hybrid orbitals:
We took 1 s orbital and 3 p orbitals to make 4 sp3 orbitals.
The sum of the atomic orbitals = the sum of the hybrid orbitals
The 4 sp3 orbitals point in the 4 directions of the tetrahedral bonds
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sp3 Hybridization47
Back to CH4:
The 4 2sp3 hybrid orbitals on carbon make σ-bonds with the 1s orbitals of the 4 H atoms
NH3 also uses sp3 hybrid orbitals:
Whenever an atom has a tetrahedral structure, it is sp3 hybridized
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sp3 Hybridization48
Whenever an atom has a tetrahedral structure, we say it is sp3 hybridized
This includes molecules with multiple central atoms:
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Atomic Orbitals Hybrid Atomic Orbitals
49
psh1 psh2
sp3 bonding works for tetrahedrally shaped molecules
What about the other VSEPR shapes?
Linear bonding can be described by sp hybridization:
2 of the p orbitals remain as they were
Unused p orbitals are available for bonding
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Other Types of Hybridization
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sp3 bonding works for tetrahedrally shaped molecules
What about the other VSEPR shapes?
Trigonal planar can be described by sp2 hybridization:
The pz orbital is not used and remains as it was.
y1 p2sh
yx2 p2
1p2
3sh
yx3 p2
1p2
3sh
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Hybrid Orbitals51
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Hybrid Orbital Shapes52
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Look Familiar?53
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2 3 4 5 6 sp sp2 sp3 sp3d sp3d2
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Other Types of Hybridization
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N atomic orbitals always produce N hybrid orbitals
Spectroscopic data suggests terminal atoms use hybrid orbitals as well
A terminal Cl uses sp3 hybridization in the arrangement of its lone pairs?
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Examples56
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π Bond57
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Ethylene58
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Ethylene59
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Characteristics of Multiple Bonds60
double bond = 1 σ-bond + 1 -bond
triple bond = 1 σ-bond + 2 -bonds
σ-bonds result from head-on overlap of orbitals
-bonds results from side-by-side overlap
Atoms in a single bond can rotate freely, whereas atoms in a double bond are much less likely to:
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Characteristics of Multiple Bonds
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The shape of ethylene (C2H4)
Experimental evidence:
All six atoms lie in the same plane with 120º bond angles
Suggests trigonal planar structure sp2 hybridization
Each C: sp2 + p
σ-bonds with the sp2 orbitals
C –– C
4 C –– H
-bond with the leftover p orbital
C –– C
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Characteristics of Multiple Bonds62
The shape of acetylene (C2H2) is linear sp hybridization
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Hybridization and Benzene63
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Molecular Orbital Theory64
Lewis Theory Connectivity, electron tracking
VSEPR Theory 3-D Structure around an atom
Valence Bond Theory Extended 3-D Structure Information Delocalization in Molecules (LIMITED) Illustrates Multiple Bonding Prediction of Reactivity
Molecular Orbital Theory Orbitals as a function of the whole molecule Delocalization much more thorough Anti-bonding and Non-bonding electrons Behavior in a magnetic field Spectral Data
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Why MO Theory?65
O OB
H
H
BH
H H
H
O2 Bond Dissociation Energy = 498 kJ/mol O2
+ Bond Dissociation Energy = 623 kJ/mol
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Paramagnetic / Diamagnetic
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O O
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Constructive and Destructive Interference
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Molecular Orbitals68
MO Theory – electrons occupy molecular orbitals spread over entire molecule
All valence electrons are delocalized Molecular orbitals are built by adding together atomic orbitals Linear combination of atomic orbitals (LCAO)
H – H bond in H2 with two 1s electrons:
A1sB1s
• LCAO-MO shows constructive interference
• ELCAO-MO < EAO
• LCAO-MO are bigger than AO
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Molecular Orbitals69
“Conservation of orbitals”
H2 bonds with two 1s orbitals – it must have 2 LCAO-MOs
A1sB1s
1st LCAO-MO is lower in energy than Atomic orbitals
This is called a bonding orbital
2nd LCAO-MO - destructive interference, higher PE than atomic orbitals
Nodal plane! – much less e- density “internuclearly”
This is an antibonding orbital
A1sB1s
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Sigma Bonds70
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Molecular Orbital Energy Diagram71
E
Shows:
Relative Energies of
Atomic Orbitals
Bonding MOs
Antibonding MOs
Shapes of LCAO-MOs
Type of MO – σ or
Molecular Electron Configuration
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Electron Configurations of Diatomic Molecules
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Shows ALL valence electrons using the Building-Up Principle:
Electron configuration of H2 σ1s2
Based on MO theory – a single electron can hold a bond together
Use “Hund’s Rule just like in atomic orbitals!
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Bond Order73
Valence Bond Theory: Bond order was #
bonding pairs MO Theory, bond order
is: b = ½ (N – N*)
N = # e –s in bonding orbitals
N* = # e–s in antibonding orbitals
H2 b = 1 He2 b = 0
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2nd row Diatomic Molecules74
Must include both 2s and 2p orbitals in making our MOs
Only orbitals that are close in energy will form MOs
We have 2 (2s) orbitals and 6 (2p) orbitals 8 atomic orbitals
2s orbitals overlap in the same way as the 1s orbitals of H2
2p orbitals overlap in same orientations as in Valence bond theory:
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2nd row Diatomic Molecules75
We have 2 (2s) orbitals and 6 (2p) orbitals 8 atomic orbitals
2s orbitals overlap in the same way as the 1s orbitals of H2
2p orbitals overlap in same orientations as in Valence bond theory:
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Pi Bonds76
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2nd row Homonuclear Diatomic Molecules
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2nd row Homonuclear Diatomic Molecules
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Homonuclear Diatomic MO Energy diagram for all elements right of (and including) oxygen
O2 F2 ‘Ne2’
And their combinations
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2nd row Homonuclear Diatomic Molecules
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Homonuclear Diatomic MO Energy diagram for all elements left of oxygen
Li2 Be2 B2 C2 N2
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Relative MO Energy Levels80
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MO Diagram for Li281
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MO Diagram for Be282
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MO Diagram for N283
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MO Diagram for O284
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MO Diagram for F285
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86
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Relative AO Energy Levels87
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Heteronuclear Diatomic Molecules88
Account for ΔEN between AO’s
More electronegative atoms will have lower energy atomic orbitals
Contributions are dependent on energy level
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MO diagram for HF89
2p
2s
1sE
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Orbitals in Polyatomic Molecules90
An electron pair in a bonding orbital helps to bind together the whole molecule – can become very complex
Occasionally some MOs are neither bonding nor antibonding –these are called nonbonding orbitals
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MO Diagram for Carbon Monoxide
91
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Organic Molecules92
Chemists commonly mix VB and MO theories when discussing organic molecules:
VB theory is good for talking about σ bonds
MO theory is better for considering bonds
Consider Benzene (C6H6):
VB perspective – each C is sp2 hybridized
σ framework is overlap of sp2 orbitals
MO perspective – conjugation of π bonds extremely important
MOs account for delocalization of electrons
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Orbitals of Benzene93
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HOMOLUMO Transitions94
HOMO – highest occupied MO
LUMO – lowest unoccupied MO