Theories of Covalent Bonding
Theories of Covalent Bonding
Theories of Covalent Bonding
11-1
Theories of Covalent Bonding
11-2
The three models of chemical bonding
Figure 9.2
11-3
Covalent bond
formation in H2
Figure 9.11
11-4
Key Principles
shape = conformation
11-5
The Complementary Shapes of an Enzyme and Its Substrate
11-6
Valence-shell Electron-Pair Repulsion (VSEPR) Theory
11-7
A periodic table of partial ground-state electron configurations
Figure 8.12
11-8
The steps in determining a molecular shape
molecular Step 1
formula
molecular
Figure 10.12
shape
(AXmEn)
11-9
Steps to convert a molecular formula into a Lewis structure
11-10
Electron-group repulsions and the five basic molecular shapes
Figure 10.5
Examples:
CH4, SiCl4,
SO42-, ClO4-
Examples: Examples:
NH3 H 2O
PF3
OF2
ClO3
H 3O+ SCl2
Figure 10.8
11-12
The four molecular shapes of the trigonal bipyramidal
electron-group arrangement
Examples:
Examples:
SF4
PF5
XeO2F2
AsF5
IF4+
SOF4
IO2F2-
Examples:
Examples:
XeF2
ClF3
I3 -
BrF3
IF2-
Figure 10.10
11-13
VSEPR
(Valence Shell Electron Pair RepulsionTheory)
11-14
The Central Themes of Valence
Bond (VB) Theory
Basic Principle
A covalent bond forms when the orbitals of two atoms overlap
and are occupied by a pair of electrons that have the highest
probability of being located between the nuclei.
The greater the orbital overlap, the stronger (more stable) the
bond.
11-15
Orbital overlap and spin
pairing in three diatomic
molecules
hydrogen, H2
hydrogen fluoride, HF
Figure 11.1
fluorine, F2
11-16
Linus Pauling
11-17
Hybrid Orbitals
Key Points
11-18
The sp hybrid orbitals in gaseous BeCl2
atomic
orbitals
hybrid
VSEPR
orbitals
predicts a
linear
shape
Figure 11.2
Figure 11.2
11-20
The sp2 hybrid orbitals in BF3
VSEPR predicts
a trigonal planar
shape
Figure 11.3
11-21
The sp3 hybrid orbitals in CH4
VSEPR
predicts a
tetrahedral
shape
Figure 11.4
11-22
The sp3 hybrid orbitals in NH3
VSEPR predicts
a trigonal
pyramidal shape
Figure 11.5
11-23
The sp3 hybrid orbitals in H2O
VSEPR predicts
a bent (V) shape
Figure 11.5
11-24
The sp3d hybrid orbitals in PCl5
VSEPR predicts
a trigonal bipyramidal
Figure 11.6 shape
11-25
The sp3d2 hybrid orbitals in SF6
VSEPR predicts an
octahedral shape
Figure 11.7
11-26
11-27
Conceptual steps from molecular formula to the
hybrid orbitals used in bonding
molecular shape
molecular Lewis hybrid
and e- group
formula structure orbitals
arrangement
Figure 10.1 Figure 10.12 Table 11.1
Figure 11.8
11-28
SAMPLE PROBLEM 11.1 Postulating Hybrid Orbitals in a Molecule
11-29
SAMPLE PROBLEM 11.1 (continued)
2p 2p
sp3 sp3
hybridized C atom hybridized O atom
2s 2s
single C atom single O atom
(b) SF4 has a seesaw shape with four bonding and one non-bonding e- pairs.
F
F S
F 3d 3d
F
distorted 3p
trigonal
bipyramidal sp3d
hybridized
S atom
3s S atom
11-30
Covalent Bonds Between Carbon Atoms - Single Bonds
bonds in ethane, CH3-CH3
both carbons are sp3
hybridized s-sp3 overlaps to bonds
11-31
Covalent Bonds Between Carbon Atoms - Double Bonds
and bonds in ethylene, C2H4
overlap in one position -
p overlap -
hindered rotation
~120o
electron density
11-32 Figure 11.10
Covalent Bonds Between Carbon Atoms - Triple Bonds
and bonds in acetylene, C2H2
p overlap -
hindered rotation
180o
Figure 11.11
11-33
Video: Hybridization
11-34
SAMPLE PROBLEM 11.2 Describing bonding in molecules with
multiple bonds
SOLUTION:
sp2 sp2
sp3 hybridized
O O
O
H sp2
sp3 hybridized H C H
2
C C H sp sp3
H H H sp3 C
H sp2 C sp
2 3
C sp
sp2 hybridized C 3
H H3 C CH3
sp 3 sp 3H
sp3 H sp
H sp3
bond
bonds
11-35
Restricted rotation in -bonded molecules
cis trans
No spontaneous interconversion between
cis and trans forms (isomers) in solution at room temperature!
11-37
Molecular Orbital (MO) Theory
11-38
Central themes of molecular
orbital (MO) theory
11-39
An analogy between light waves and atomic wave
functions
Amplitudes of wave
functions are added
Figure 11.13
Amplitudes of wave
functions are
11-40 subtracted
Contours and energies of the bonding and antibonding
molecular orbitals in H2
Figure 11.14
11-41
number of AOs combined = number of MOs produced
Sigma () and pi () bonds are denoted as before; a star (asterick)
is used to denote antibonding MOs.
11-42
Molecular orbital diagram for
the H2 molecule
(aufbau and
exclusion
principles, Hund’s
rule)
Figure 11.15
11-43
The MO bond order
11-44
MO diagrams for He2+ and He2
*1s *1s
Energy
Energy
1s 1s 1s 1s
1s 1s
AO of MO of AO of AO of MO of AO of
He He+ He+ He He2 He
PROBLEM: Use MO diagrams to predict whether H2+ and H2- can exist.
Determine their bond orders and electron configurations.
PLAN: Use H2 as a model and accommodate the number of electrons in
bonding and antibonding orbitals. Calculate the bond order.
1s 1s 1s 1s
AO of H AO of H-
AO of H AO of H+
MO of H2+ MO of H2-
configuration is (1s)1 configuration is (1s)2(1s)1
11-46
Figure 11.17
*2s *2s
2s 2s 2s 2s
Energy
*1s *1s
1s 1s 1s 1s
1s 1s
11-47 Li2 bond order = 1 Be2 bond order = 0
Bonding and antibonding MOs for core
electrons cancel = no net contribution to bonding
11-48
Contours and energies of and MOs through
combinations of 2p atomic orbitals
end-to-end
overlap
side-to-side
overlap
Figure 11.18
11-49
Relative energies
11-50
Relative MO energy levels for Period 2 homonuclear
diatomic molecules
Figure 11.20
11-52
The paramagnetic
properties of O2
Explained by
MO diagram
Figure 11.21
11-53
SAMPLE PROBLEM 11.4 Using MO theory to explain bond properties
11-54
SAMPLE PROBLEM 11.4 (continued)
N2 N2+ O2 O2 +
2p 2p
2p 2p
2s 2s
2s 2s
1s
nonbonding MOs
2p
2px 2py
lower in energy
than 1s of H!
AO MO of AO
of H HF of F
11-56
In polar covalent compounds, bonding MOs
are closer in energy to the AOs of the more
electronegative atom.
11-57
Figure 11.23
*2s The MO diagram for NO
2p
Energy
2p 2p
MO of NO
11-58
The lowest energy -bonding MOs in benzene and ozone
O
O O
resonance hybrid
Figure 11.24
11-59