chapter 8 – covalent bonding 8.1the covalent bond 8.2 naming molecules 8.3molecular structures...
TRANSCRIPT
Chapter 8 – Covalent Bonding
8.1 The Covalent Bond
8.2 Naming Molecules
8.3 Molecular Structures
8.4 Molecular Shape
8.5 Electronegativity and Polarity
Topic is Lewis Structures (combination of material found in 2 sections)
Sections 8.1/8.3 Covalent Bond/Molecular Structures
• Apply the octet rule to atoms that form covalent bonds.
• Describe the similarities and differences between ionic and covalent bonding.
• Describe the forces that act and energy changes that occur when atoms form a molecule.
Atoms gain stability when they share electrons and form covalent bonds. The sharing can be described by the Lewis structure of the compound.
• Categorize valance electrons as bonding or non-bonding.
• List the basic steps used to draw Lewis structures.
• Describe the formation of single, double, and triple covalent bonds using Lewis structures.
• Explain why resonance occurs, and identify resonance structures.
• Draw valid Lewis structures for molecules, including those involving multiple bonds, resonance, expanded octets, and electron deficient molecules.
Sections 8.1/8.3 Covalent Bond/Molecular Structures
Why Atoms Bond
Noble gas configuration especially stable
• ns2np6 (except for He)• Full outer energy level• Octet of electrons
Atoms bond to achieve a lower energy state (more stability)
Ionic vs Covalent Bonding
In ionic bonding, electrons transferred to achieve octet for each ion
• Number of ionic compounds small compared to total number of known compounds
In covalent bonding, electrons shared to achieve octet (mostly) for each atom
Covalent Bonding
In covalent bonding, electrons shared to achieve octet (mostly) for each atom
• Sharing occurs when electronegativities of atoms same or similar
• Majority of covalent bonds formed between nonmetallic elements
• Electronegativity difference < 1.7 (see next slide) – bond will have more covalent character than ionic character
EN Difference & Bond Character
0 1.0 2.0 3.0
Electronegativity Difference
% Io
nic
Cha
ract
er
25
50
75
100
Ionic Bonds
Covalent Bonds
Covalent Bonding
In covalent bonding, electrons shared to achieve octet (mostly) for each atom
Molecule formed when 2 or more atoms bond covalently
Covalent Bonding – Forces
Covalent Bonding – Forces
No interaction
Nucleus attracted to other atoms
electrons –
Not optimum distance
Net repulsion
from positive nuclei
Nucleus attracted to other atoms
electrons –
optimum distance
Covalent Bonding – Energy for H2
Internuclear Distance (pm)
100 200Pot
entia
l Ene
rgy
(kJ/
mol
)
-432 kJ/mol
Lewis Structures - Atoms (Electron Dot Diagrams)
Way of keeping track of valence electrons
To write for atom• Write symbol for element X• Put one dot for each
valence electron• Don’t pair up until you
have to (Hund’s rule)
Valence electrons of each element in molecule are divided into 2 categories:
• Bonding – pair of electrons shared by two atoms to form the covalent bond
Shared pair represented by a line connecting the element symbols H—H
• Nonbonding – called lone pairs A few molecules have odd # total
electrons – have unpaired nonbonding electron
Lewis Structure - Covalent Molecules
Lewis Structure - Covalent Molecules
Example – formation of H2 molecule
Bond = shared electron pair
H does not form octet
Space-Filling Model View Formation of H2
+
Bond = shared electron pair H_H
•HH••HH•
Ways of Representing Molecules: H2O
Structural Formula Space-Filling Model
Ball-and-Stick ModelOrbital Model
Ways of Representing Molecules: PH3
Covalent Bonding – F2
F 1s22s22p5
7 valence electrons
Forms F2 molecule
Each F shares 1 valence electron
Molecule is more stable than individual atoms
Lewis Structure - Covalent MoleculesExample – formation of F2 molecule
Bond = shared electron pairOctet
formed
Lewis Structure - Covalent MoleculesExample – formation of F2 molecule
Octet formed
Lewis Structure - Covalent MoleculesExample – formation of H2O molecule
Bonds = shared electron pairsOctet
formed
Two lone pairs
Shape of molecule
Lewis Structure - Covalent MoleculesExample – formation of ammonia, NH3
+ + +
Shape of molecule
Lone pair
Bonds = shared electron
pairs
Octet formed
Multiple Covalent Bonds
C, N, O, S often form multiple bonds
Double bond – O2 (6 valence e per O)
+
Triple bond – N2 (5 valence e per N)
+
Guide for Writing Lewis Structures
Step 1 – Write skeletal structure• least electronegative atom usually
occupies central position
Step 2 – Count total number of valence electrons
• polyatomic anions, add # of - charges e.g. CO3
2- add 2 electrons to total• polyatomic cations, subtract # of +
charges
Similar to procedure on p. 254, but without # of bonding pairs
Step 3 – Place single bond between central atom and surrounding atoms
Step 4 – Complete octet for terminal atoms (not for H)
Step 5 – Add remaining to central atom
Step 6 – If octet rule not satisfied for central atom, add multiple (double, triple) bonds between terminal and central atom, using the lone pairs from the terminal atoms
Lewis Structures – Common Bonding Patterns
C 4 bonds & 0 lone pairs
4 single (CH4), or 2 double (CO2), or single + triple (HCCH), or 2 single + double (CH2CH2)
N 3 bonds & 1 lone pair (NH3)
O 2 bonds & 2 lone pairs (H2O)
H & halogen 1 bond (CH4, CF4)
Be 2 bonds & 0 lone pairs (BeH2, electron def.)
B 3 bonds & 0 lone pairs (BH3, electron def.)
B C N O F
Lewis Structure Examples
Draw Single Bonds
Total Valence
Electrons
Calculate Number of Electrons
Remaining
Use Remaining
Electrons to Achieve
Noble Gas Configuration
Check Number of Electrons
a, HF 1 + 7 = 8 H-F 6
b, N2 5 + 5 = 10 N-N 8 c, NH3 5 + 3(1) = 8 2
d, CH4 4 + 4(1) = 8 0
e, CF4 4 + 4(7) = 32 24
f, NO+ 5 + 6 - 1 = 10 N-O 8
NH H
H
CH H
H
H
CF F
F
F
H, 2
F, 8
N,8
N, 8
O, 8
H, 2
N, 8
H, 2
C, 8
F, 8
C, 8
NH H
H
N N
FH
CH H
H
H
CF F
F
F
N O+
Practice
Problems 1-5 page 244
Problems 37-38, page 255
Problems 39-40, page 256 (mult bonds)
Problems 41-42 page 257 (ions)
Problems 104(a-d), page 275
Problems 1(a-d), page 979
Problems 4(a-e) page 980
Lewis Structure Example: NO3─
1. Write skeletal structure
N central because it is least electronegative
2. Count valence electrons
ONO
O
N = 5
3O = 3 x 6 = 18
(-) = 1
Total = 24 e-
Example NO3─ , Continued
3. Attach atoms with single bonds (pairs of electrons) & subtract from total ONO
O
——
Electrons
Start 24
Used 6
Left 18
4. Complete octets, outside-in
Keep going until all atoms have an octet or you run out of electrons
::
::
—— ONO
O
Electrons
Start18
Used18
Left 0
Example NO3─ , Continued
5. If central atom does not have octet, bring in electron pairs from outside atoms to shareIf structure is an ion, use brackets and indicate the charge
6. For this ion an extra step is needed – draw resonance structures
::
::
— ONO|
O
::
::
—— ONO
O
Example NO3─ , Continued
-1
Can have more than one correct Lewis structure for molecules or ions with double and single bonds
-
Example NO3─ , Continued
Resonance StructuresResonance structures differ only in position of electron pairs, never the atom positions
Molecule behaves as if it had only one structure (the average one)
• NO3- has all bond lengths
identical
-
Practice (Resonance Structures)
Problems 43-46 page 258
Problems 101,103 pages 274-5
Problems 5, 6 page 980
Practice—Lewis Structures
NClO
H3BO3
NO2-1
H3PO4
SO3-2
P2H4
NClO
H3BO3
NO2-1
H3PO4
SO3-2
P2H4
O P
O
O
O
HH
H
••
••
••
••
••
••
••
••
••
O S
O
O
••
••
•• •
•••
••
••
••
••
••
O N O ••
••
••
••
••••
18 e-
26 e-
32 e-
14 e- H P P H
HH
•• ••
O B
O
OH H
H••
••
••
••
••
••
24 e-
O N Cl ••
••
••
••
••••18 e-
Practice—Lewis Structures
*
*
* Has resonance structures
-1
-2
Exceptions to Octet RuleMolecules with odd number of total valence electrons
NO2 – 17 valence electrons
Also ClO2, NO
Exceptions to Octet RuleElectron deficient – form with fewer than 8 electrons around atom
• Be, B • Rare
+
Tend to form coordinate covalent bonds – both electrons in shared pair donated by single atom
Exceptions to Octet Rule
BeH2 – 4 electrons
BF3 – 6 electrons
Exceptions to Octet RuleMore than 8 valence electrons = expanded octet
PCl5 SF6
d orbitals involved• Only can occur for period 3 and higher,
not periods 1 or 2
Practice (Octet Exceptions)
Problems 47 - 49 page 260
Problems 102 (a-d), 104(a-d) page 273
Problem 7, page 980
Chapter 8 – Covalent Bonding
8.1 The Covalent Bond
8.2 Naming Molecules
8.3 Molecular Structures
8.4 Molecular Shape
8.5 Electronegativity and Polarity
Section 8.2 Naming Molecules
• Translate molecular formulas into binary molecular compound names and also the reverse process.
• Name acidic solutions
Specific rules are used when naming binary molecular compounds, binary acids, and oxyacids.
Naming Binary Covalent CompoundsFirst element named first, using entire element name
Second element named using same procedure as for ionic compounds – root of element name + ide ending
Use prefixes except if first element = 1• Drop final letter in prefix if precedes
vowel• Carbon monoxide , not monooxide
Prefixes in Covalent CompoundsTable 9-1, page 248
# Atoms Prefix # Atoms Prefix
1 mono- 6 hexa-
2 di- 7 hepta-
3 tri- 8 octa-
4 tetra- 9 nona-
5 penta- 10 deca-
Naming Binary Covalent Compounds
Name of AlCl3 ?Aluminum chloride
Name of PCl3 ?Phosphorus trichloride
Name of Al2O3?Aluminum oxide
Name of P2O5 ?Diphosphorus pentoxide
The naming systems for ionic and covalent compounds are different!!!
Common NamesTable 9-2, page 249
Formula
Common Name
Molecular Compound Name
H2O Water Dihydrogen monoxide
NH3 Ammonia Nitrogen trihydride
N2H4 Hydrazine Dinitrogen tetrahydride
N2O Nitrous oxide (laughing gas)
Dinitrogen monoxide
NO Nitric oxide Nitrogen monoxide
Naming Acids
For our purposes, acids are what result when molecules dissolved in water produce H+ (hydrogen ions)
• HCl(g) in water H+(aq) + Cl-(aq)• Product is hydrochloric acid
Two common types• Binary – H and one other element• Oxyacid – H and an oxyanion
Naming True Binary Acids
Use prefix hydro- to name hydrogen part of compound
For remainder, use a “form of the root” of 2d element plus suffix –ic followed by word acid
HCl – hydrochloric acid
H2S – hydrosulfuric acid• Root of S for acid name not “sulf” as in
Na2S (sodium sulfide)
Naming Acids Similar to Binary Acids (Rare)
If second part of compound is a polyatomic anion that does not contain oxygen (rare), use same system as for a true binary acid employing the root name for the anion
CN- – cyanide anion
HCN – hydrocyanic acid
Naming OxyacidsName is based solely on the anion
“A form of the root name of the anion” + suffix + acid
Anion suffix Acid Suffix
-ate -ic
-ite -ous
HNO3 Nitric acid NO3- = nitrate
HNO2 Nitrous acid NO2- = nitrite
Naming Molecular CompoundsFlow Chart, Fig 9-9, page 251
Naming Molecular CompoundsFlow Chart, Fig 9-9, page 251
Acidic Not Acidic
Practice
Problems 13-17 page 249 (binary covalent)Problems 18-22 page 250 (acids)Problems 27-29 page 251 (mixed)Problems 94-96(all a-d) page 273Problems 97-98(all a-d) page 273Problems 2 (a-f) page 874 (binary cov)Problem 3 page 875 (acids)
Chapter 8 – Covalent Bonding
8.1 The Covalent Bond – Bond Strength
8.2 Naming Molecules
8.3 Molecular Structures
8.4 Molecular Shape
8.5 Electronegativity and Polarity
Section 8.1 The Covalent Bond
• Relate the strength of a covalent bond to its bond length, bond order, and bond dissociation energy.
• Describe how the overall energy of a reaction (i.e., whether it is an endo- or exothermic reaction) is related to the bond energies of the reactant and product molecules.
Covalent Bonding – Energy for H2
Internuclear Distance (pm)
100 200Pot
entia
l Ene
rgy
(kJ/
mol
)
-432 kJ/mol
Bond Strength & Bond Length/Order
Distance between bonding nuclei at position of max attraction = bond length
Scale of bond length: ~10-10 m =100 pmBond order: Single 1 Double 2 Triple 3
Bond Strength & Bond Length/Order
Strength of bond related to bond lengthBond dissociation energy = energy needed to break bond Triple bond > double bond > single bond
Molecule Bond Length (pm)
Dissoc. Energy kJ/mol
F2 143 159
O2 121 498
N2 110 945
Bond Strength & Bond Length/Order
Reaction Energies & Bond Energies
Chemical reaction
Bonds in reactant molecules broken
New bonds formed in product molecules
CH4 + 2O2 2H2O + CO2
Breaking C-H bonds and O=O bonds
Making O-H bonds and C=O bonds
Reaction Energies & Bond Energies
CH4(g) + 2O2(g) 2H2O(g) + CO2(g)
Total energy change determined by difference of energy of bonds broken (reactant side) and formed (product side)• Endothermic – need more energy to
break than get back in formation• Exothermic – bond formation energy
larger than energy needed to break bonds
Reaction Energies & Bond EnergiesE
ntha
lpy
-SBE (products)
-SBE (products)
SBE (reactants)
SBE (reactants)
BE = Bond energy
Chapter 8 – Covalent Bonding8.1 The Covalent Bond
8.2 Naming Molecules
8.3 Molecular Structures
8.4 Molecular Shape
8.5 Electronegativity and Polarity
Section 8.4 Molecular Shapes
• Summarize the VSEPR bonding theory, including the role of bonding and nonbonding pairs of electrons.
• Predict the shape of, and the bond angles in, a molecule using VSEPR theory.
The VSEPR model is used to determine molecular shape.
2 Simple Theories Related to Covalent Bonding
Valence Shell Electron Pair Repulsion Theory (VSEPR)• Use Lewis structures to predict shape
Valence Bond Theory• Extends Lewis bonding model to focus
on orbitals, particularly hybridized orbitals
VSEPR
Valence Shell Electron Pair Repulsion Theory - allows us to predict geometry
Lewis structures tell us how the atoms are connected to each other
Lewis structures don’t tell us anything about shape
Shape of a molecule can greatly affect its properties
Molecular Shape & Biological Sensors
For some biological systems, a response is generated or a chemical change is initiated when a molecular key fits into correspondingly shaped molecular lock• Key is typically small molecule• Lock is typically large molecule with a
shaped receptor siteOnly interacts with key of a specific shape
Lewis Structure (a) & Tetrahedral Geometry (b) for Methane (CH4)
VSEPR
Molecules take a shape that puts electron pairs as far away from each other as possible (electron pair repulsion)
Have to draw the Lewis structure to determine categories of electron pairs
• bonding• nonbonding lone pair
Lone pair take more space
Multiple bonds count as one pair
Balloon Analogy for the MutualRepulsion of Electron Groups
Two Three Four Five Six
Number of Electron Groups
VSEPR
The number of pairs determines• bond angles• underlying structure
The number and position of atoms determines
• actual molecular shape
VSEPR – Underlying Shapes
# Elec. pairs Bond Angles Shape2 180° Linear
3 120° Trigonal Planar
4 109.5° Tetrahedral
590° &120°
Trigonal Bipyramidal
6 90° Octahedral
Actual Molecular Shapes
3 3 0 trigonal planar
4 4 0 tetrahedral4 3 1 trigonal pyramidal
2 2 0 linear
3 2 1 bent
4 2 2 bent
ElectronPairs
BondingPairs
Non-Bonding
Pairs Shape
Actual Molecular Shapes
5 5 0 trigonal bipyrimidal
5 4 1 See-saw
5 3 2 T-shaped5 2 3 linear
ElectronPairs
BondingPairs
Non-Bonding
Pairs Shape
Actual Molecular Shapes
6 6 0 Octahedral
6 5 1 Square Pyramidal
6 4 2 Square Planar6 3 3 T-shaped6 2 4 linear
ElectronPairs
BondingPairs
Non-Bonding
Pairs Shape
Relative Sizes: Bonding Pairs vs Lone Pairs
CH4 NH3H2O
Molecular GeometryCan predict geometry around each atom center and build overall molecular geometry piece by piece
GlycineN
C1
C2
O1
O2
Chapter 8 – Covalent Bonding
8.1 The Covalent Bond8.2 Naming Molecules8.3 Molecular Structures8.4 Molecular Shape (extension of book) Valence Bond Theory - Orbital Overlap Hybrid Orbitals
Quantum mechanical calculations8.5 Electronegativity and Polarity
Section 8.4 Molecular Shapes
• Describe the valence bond model of bonding
• Explain the similarities and differences between the Lewis and valence bond models of chemical bonds.
• Describe sigma and pi bonds and identify these bonds within molecules.
• Define hybridization.
The Valence Bond model is used to determine molecular shape via the concept of overlap of orbitals, particularly hybrid orbitals.
Section 8.4 Molecular Shapes
• Relate the type of hybridization (sp3, sp2, etc.) to the VSEPR geometry of a molecule
• Identify the specific type of hybridization that occurs within a given molecule and identify the specific orbitals (hybrid or non-hybrid) that are involved in each sigma and pi bond.
• Explain how quantum mechanics and the wave function concept can be applied to a molecule.
Valence Bond Theory
Lewis structures indicate status of electrons• Shared in bond• Lone pair
No information about orbitals involved
Valence bond theory• Bonds are formed by overlap of half-filled
atomic orbitals• Orbital geometry can give direct information
about molecular shape
Sigma Bonds
Single covalent bonds = sigma bond• Symbol Greek letter
Occurs when electron pair shared in area centered between two atoms
Atomic orbitals overlap end to end, forming a bonding orbital
• Localized region where bonding electrons will most likely be found
Sigma Bond Formation by Orbital Overlap
Two s orbitals overlap
Two s orbitals overlap
Two p orbitals overlap
H2
HF
F2
Sigma Bond Formation
Sigma Bonding – F2
px1py
2pz2 px
1py2pz
2
F—F
Area of overlap for atomic
orbitals
Pi Bond ()Formed when parallel orbitals overlap to share electrons
Shared pair occupies space above and below a line connecting atoms
Multiple bonds always have one sigma and at least one pi bond
• Double: 1 , 1 bond• Triple: 1 , 2 bonds
Sigma & Pi Bonding
Sigma () and Pi () Bonds
Hybrid Orbitals
For correct geometry of polyatomic molecules using the valence bond model, have to use concept of hybrid orbitals
• CH4 has 109.5 angles, but atomic p orbitals are at right angles to each other
Hybrid Orbitals
Hybrid orbitals – orbitals obtained when 2 or more nonequivalent orbitals combine to form an equal number of identical, degenerate orbitals
Hybridization – mixing of atomic orbitals in an atom (usually a central atom) to generate a set of hybrid orbitals
Use VSEPR logic to determine geometry of hybrid orbitals formed
Valence Orbitals on a Free Carbon Atom: 2s, 2px, 2py, and 2pz
s py
px pz
Formation of sp3 Hybrid Orbitals From Original Valence Orbitals
Hybridization
Cross Section of sp3 Orbital
Energy-Level Diagram Showing Formation of Four sp3 Orbitals
C Orbitals in CH4 molecule
Orbitals in free C atom
Hybridization
2s 2p
C
1s 1s 1s 1s
4 H
C*
Valence Bond Theory Treatment of CH4
Overlap of sp3 hybrid orbitals on C with 1s orbitals on H atoms gives 4 C-H (sp3)-1s bonds oriented 109.47° from each otherHas tetrahedral geometry predicted by VSEPR
HC
H
HH
sp3
C* (sp3)
Tetrahedral Set of Four sp3 Orbitals Forming Sigma Bonds with s Orbitals of Four Hydrogen Atoms(CH4)
Formation of sp2 Hybrid Orbitals from s, px, and py Atomic Orbitals
Hybridization
Energy-Level Diagram Showing Formation of Three sp2 Orbitals
Orbitals in sp2 hybridized C
Orbitals in free C atom
Hybridization
Note: Inconsistent with actual bonding – 4 valence electrons populate only sp2 orbitals (Aufbau) leaving only 1 unpaired electron in sp2
An sp2 Hybridized C Atom
Formation of sp Hybrid Orbitals from s and px, Atomic Orbitals
Hybridization
Energy-Level Diagram Showing Formation of Two sp Hybrid Orbitals
Orbitals in sp hybridized C
Orbitals in free C atom
Hybridization
Note: Inconsistent with actual bonding – 4 valence electrons should populate only sp orbitals (Aufbau) leaving no unpaired electrons
Orbitals of sp Hybridized Carbon Atom
sp3d (dsp3) Hybrid Orbitals
3s 3p
P
P*
3d
P* (sp3d) 3d
Can only occur for periods 3 & higher (need d orbitals) – example shown is for PLinked to geometry with 5 pairs (trigonal bipyramid)
3dz23pz 3py 3px3s sp3dz2
Set of dsp3 Hybrid Orbitals on a Phosphorus Atom
sp3d2 (d2sp3) Hybrid Orbitals
Can only occur for periods 3 & higher (need d orbitals)Linked to geometry with 6 pairs (octahedral)Example on next slide for S
S - Octahedral Set of d2sp3 Orbitals
Relationship among the number of effective pairs, geometry, and the
hybrid orbital set required to obtain this geometry shown on the following two
slides
Linear sp2
3
4
Trigonal sp2
planar
Tetra sp3
hedral
# Geometry Hybridization
Trigonal sp3dbipyramidal
5
6
# Geometry Hybridization
Octa sp3d2
hedral
Geometry & Hybridization - Steps1. Draw Lewis structure
2. Determine # of effective electron pairs (count double & triple bonds as one pair)
3. Determine basic geometry from number of pairs (e.g., 5 pairs = trigonal bipyramid)
4. Determine hybridization type from number of pairs (e.g., 5 pairs = sp3d)
5. Form single (sigma) bonds from hybrid orbitals; lone pairs also go in hybrid orbitals
6. Form pi bonds using unhybridized orbitals
Geometry & Hybridization - StepsFollowing slides give examples of using the steps listed on previous slide for these molecules:
1. Ammonia
2. Ethylene
3. Diatomic nitrogen
4. Acetylene
5. Carbon dioxide
6. Phosphorus pentachloride
N in Ammonia
sp3 Hybridized (4 pairs)
N in Ammonia
Trigonal pyramidal molecule with lone pair occupying hybrid orbital
Sigma & Pi Bonds Using Hybrid Orbitals - Ethylene
Three electron pairs for C sp2 hybridization & trigonal planar geometry
C 1s22s22p2 1s22(sp2)32p1
Hydrogens have 1s1 orbitals (spherical)
s Bonds in Ethylene – Top View
Sigma () bonds
Sigma and Pi Bonds in Ethylene
Pi () bond
Sigma Bonds in Ethylene
Because each C has trigonal planar geometry, entire molecule is planar
N2 Bonding
Two sp hybrid orbitals and two normal p orbitals
sp
py
pz
sp
Two pi bonds
One sigma bond
lone pair sigma lone pair
sp hybridized (2 pairs)
Sigma and Pi Bonds in Acetylene
Two electron pairs for C sp hybridization, linear geometry (triple bond = single pair)
C 1s22s22p2 1s22(sp)22p2
Hydrogens have 1s1 orbitals (spherical)
Orbitals of sp Hybridized Carbon Atom
Sigma and Pi Bonds in Acetylene
sp hybrid orbitals on C form single (sigma) bond with H and other C
Remaining two unhybridized p orbitals overlap to form two pi bonds
Sigma and Pi Bonds in Acetylene
Sigma & Pi Bonds Using Hybrid Orbitals in CO2
Two electron pairs for C sp hybridization, linear geometry (double bond = single pair)
C 1s22s22p2 1s22(sp)22p2
Three electron pairs for O sp2 hybridization & trigonal planar geometry
O 1s22s22p4 1s22(sp2)5p1
Orbitals of sp Hybridized Carbon Atom
Orbital Arrangement for an sp2 Hybridized Oxygen Atom
Sigma Bonds using Hybrid Orbitals in CO2 Molecule
Sigma () bonds
Sigma and Pi Bonds Using Hybrid Orbitals in CO2
Sigma Bonds Using Hybrid Orbitals in PCl5
Five electron pairs for P sp3d hybridization & trigonal bipyramidal geometry
P [Ne]3s23p3 [Ne](sp3d)5
Four electron pairs for Cl sp3 hybridization & tetrahedral geometry
Cl [Ne]3s23p5 [Ne]3(sp3)7
Set of dsp3 Hybrid Orbitals on a Phosphorus Atom
Structure of PCI5 and Orbitals Used to Form Sigma Bonds
Sigma () bond
Lone pairs on Cl in sp3 orbitals
Geometry & HybridizationSupply for each indicated atom in structure
# of sigma & pi bonds in molecule?
HC
HC
C O
O
C
H
H
H
C
N
3 pairs
Trigonal planar
sp2
3 pairs
Trigonal planar
sp2
2 pairs
Linear
sp
4 pairs (2 lone)
Bent
sp3
4 pairs
Tetrahedral
sp3
12 , 4
Practice (Shape, Angles, Hybridization)
Problems 56 – 60, 65 - 67 page 264
Problems 108,110 - 112 page 275
Problem 8 page 980
Quantum Mechanics & Molecules
Quantum Mechanics & Molecules
Y (wave function) exists for entire molecule and can be obtained from solution to Schrodinger wave equation written for the molecule
Y2 - Square of Y gives probability of finding electron at particular position around molecule – defines what is called a molecular orbital (MO)
Quantum Mechanics & MoleculesUsing certain types of approximations and today’s computers, wave functions for molecules (not individual atoms) can be obtained and molecular properties calculated from this information
Energy, absorption spectrum, dipole moment, etc
Molecular orbital theory is most advanced way of describing covalent bonding
Chapter 8 – Covalent Bonding
8.1 The Covalent Bond
8.2 Naming Molecules
8.3 Molecular Structures
8.4 Molecular Shape
8.5 Electronegativity and Polarity
Section 8.5 Electronegativity and Polarity
• Describe how electronegativity is used to determine bond type and characterize bonds between given pairs of atoms as being polar or nonpolar.
• Compare and contrast polar and nonpolar covalent bonds and polar and nonpolar molecules.
• Describe the term “dipole moment” and relate it to the terms polar and nonpolar.
A chemical bond’s character is related to each atom’s attraction for the electrons in the bond.
Section 8.5 Electronegativity and Polarity
• Identify molecules as being polar or nonpolar.
• Describe how polarity affects the solubility of one substance in another substance.
• Describe how polarity can give rise to intermolecular forces.
Polar Covalent Bonds
Polarity of bond determined by electronegativity difference
Difference = 0 Nonpolar
Difference > 0 Polar
Very large differences • No longer covalent compound
EN Difference & Bond Character
0 1.0 2.0 3.0
Electronegativity Difference
Ionic Bonds
Covalent Bonds
% Io
nic
Cha
ract
er
25
50
75
100
Relationship Between EN Difference and Bond Type
Relationship Between EN Difference and Bond Type
EN=0
Nonpolar Covalent
EN=medium
Polar Covalent
EN=large
Ionic
Scalars & VectorsScalar
• Completely specified by magnitude and units
Vector• Has magnitude, direction, and units
v = 3.5 m/s (scalar)
v = 3.5 m/s to northeast (vector)
Trigonometric Functions
Pythagorean Theorem
Dipole MomentTwo equal and opposite charges +Q and -Q separated by a distance l have a dipole moment p:
(vector points from –Q to +Q) Q lp = p =
r
+Q -Q
l
pr
Polarity and Dipole Moment
Dipole moment is a vector pointing from center of - charge to center of + chargeMagnitude proportional to size of charges and to separation distanceAll polar covalent bonds have a dipole moment
+ _Dipole
Polarity and Dipole Moment
Units of p are Debye units (D)% ionic character of bond determined by size of measured dipole moment relative to value calculated from using full (ionic) charges as Q
Q lp = p = +Q -Q
l
pr
Dipole Moments of Gas Phase Molecules
Dipole Moment
moment) dipole(net 21 ppp += r rr
Dipole moments from bonds add as vectors to give dipole moment of molecule
Molecular Polarity – Linear Molecule
O=C bond polar; bonding electrons pulled equally toward both O ends of molecule
Net result is nonpolar molecule (dipole moments of bonds cancel each other) [note: red arrows are opposite dipole direction]
HO bond polar
Both sets of bonding electrons pulled toward O end; net result is polar molecule (y components of bond dipole moment add, x components cancel)
[note: red arrows are opposite dipole direction]
Molecular Polarity – Bent Molecule
Polar Molecules
Molecule can have polar bonds but be a nonpolar molecule
-+
+-
-
-
+
-
Polar Nonpolar
Polar Bonds in Nonpolar Molecules
In symmetric molecules, vector addition of bond dipole moments results in zero dipole moment for the molecule
All molecules having basic VSEPR shapes & equal bonds are nonpolar
Linear+
Trigonal Planar +
Practice (Polar Bonds & Polar Molecules)
Problems 74 – 77 page 270
Problems 117 – 123 page 275
Problem 9 page 980
Polarity Effects
Polarity of molecule determines solubility characteristics – “like dissolves like”
Oil (nonpolar) and water
(polar) don’t mix
Dipole in an Electric FieldThe + and – charges in an electric dipole are pulled in opposite directions in an electric field, producing a net torque on the dipole, and orienting it.
+
-δ + δ
F HFieldOff
FieldOn
Dipole in Electric Field – HF Molecule
Polar Molecule & Electric FieldPolar molecules affected by electric field in an EM wave
Oscillating field twists water molecule and energy transferred (heats up)
Basis for microwave oven operation
Properties of Covalent Compounds
Bonding types affect properties
Many properties controlled by intermolecular forces
• Forces between molecules• Also known as van der Waals forces
Intermolecular forces are weaker than chemical bonds
[Note: intermolecular forces treated in more depth in section 12.2 – Forces of Attraction]
Intermolecular Forces
Forces between nonpolar molecules relatively weak
• Tend to be gases or volatile liquids• O2, N2, small hydrocarbons
Forces between polar molecules are stronger due to dipole-dipole forces
• Hydrogen bonding a particular strong version - H and F, O, or N
Hydrogen Bonding – Water Pentamer
Hydrogen Bonds
Hydrogen Bonding in Nylon
Hydrogen bonding helps make nylon strong
End of Chapter 8