Section 9.1-9.6
Reading: Chapter 9, sections 9.1-9.6
As you read these sections ask yourself:
How does the number of bonds and nonbonded pairs of electrons
affect the shape of a molecule?
Why is the repulsion between two domains of nonbonded pairs of
electrons greater than between two domains of bonded pairs of
electrons?
How can a molecule with polar bonds be nonpolar?
Why do we need theories of bonding that differ from VSEPR?
How does Valence Bond theory differ from the Lewis concept of
chemical bonding?
How does molecular orbital theory differ from valence bond
theory?
How does a hybrid orbital differ from a pure atomic orbital?
How are hybrid orbitals related to the VSEPR shapes you learned
earlier?
Chem 101
How do sigma and pi bonds differ from each other?
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1
Section 9.1-9.6
Chapter 9 Molecular Geometry and Bonding Theories
9.1 Molecular Shapes
AB3
AB2
linear
bent
trigonal
planar
trigonal
pyramidal
T-shaped
For molecules of the general form ABn there are 5 fundamental shapes:
180°
linear
109.5°
120°
90°
90°
120°
trigonal
planar
tetrahedral
trigonal
bipyramidal
90°
octahedral
We use the electron-domain geometry to help us predict the molecular geometry.
4
1
1. Draw Lewis structure, count electron domains
3
2
2. Arrange electron domains to minimize
repulsion
3. Inspect arrangement of atoms to determine
molecular geometry
2
2
0
3
3
0
2
1
tetrahedral
trigonal
pyramidal
2
Section 9.1-9.6
4
4
0
3
1
2
2
Effect of nonbonding electrons and multiple
bonds on bond angles
bonding
pair
non-bonding
pair
Electrons in nonbonding pairs and in multiple bonds repel more than electrons in
single bonds:
124.3°
111.4°
109.5°
107°
104.5°
124.3°
Molecules with Expanded Valence Shells
Atoms that have expanded octets have five electron domains (trigonal
bipyramidal) or six electron domains (octahedral) electron-domain geometries.
5
lone pair always in least
crowded position
5
0
PCl5
4
1
SF4
3
Section 9.1-9.6
5
note position
of lone pairs
6
3
2
ClF3
2
3
XeF2
6
0
SF6
5
1
BrF5
4
2
XeF4
lone pairs as far
apart as possible
Shapes of larger molecules
The interior atoms of more
complicated molecules can be
dealt with in turn using the
VSEPR model.
4
tetrahedral
4
bent
3
trigonal
planar
4
Section 9.1-9.6
9.3 Molecular Shape and Molecular Polarity
Polar molecules interact with electric fields. Binary compounds are polar if their
centers of negative and positive charge do not coincide.
Dipoles are a
vector quantity
=0
overall dipole
moment
electron density models
(red = high, blue = low)
Consider HCl, CCl4, NH3, BF3, CH3Cl
polar
polar
non-polar
polar
non-polar
5
Section 9.1-9.6
9.4 Covalent Bonding and Orbital Overlap
HCl
H
Cl2
H
overlap regions
H2
The change in potential energy
as two hydrogen atoms combine
to form the H2 molecule:
9.5 Hybrid Orbitals
s
p
2 × sp
large lobes of sp hybrid orbitals
F 2p
orbital
F 2p
orbital
overlap regions
6
Section 9.1-9.6
one s
orbital
hybridize
two p
orbitals
one s
orbital
three
sp2
hybrid
orbitals
sp2 hybrid orbitals
shown together
(large lobes only)
three p orbitals
sp3 hybrid orbitals
shown together
(large lobes only)
four sp3 hybrid orbitals
9.5 Hybrid Orbitals
For geometries involving expanded octets on the central atom, we use d orbitals in
our hybrids:
s, p, p, p, d
five sp 3d
trigonal
bipyramidal
octahedral
s, p, p, p, d, d
six sp 3d 2
7
Section 9.1-9.6
9.6 Multiple Bonds
one π bond
Bonding π molecular orbitals arising from p
atomic orbitals
8
Section 9.1-9.6
Delocalized π Bonding
When writing Lewis structures for species like the nitrate ion, we draw resonance
structures to more accurately reflect the structure of the molecule or ion
In reality, each of the four atoms in the nitrate
ion has a p orbital
The p orbitals on all three oxygens overlap with the
p orbital on the central nitrogen

the π electrons are delocalized throughout the ion
The organic molecule benzene has six σ bonds
and a p orbital on each carbon atom
9
Section 9.1-9.6
Molecular orbital (MO) theory
(only section 9.7)
10
Section 9.1-9.6
MO diagram (energy level diagram)
antibonding MO
raises energy
bonding MO
lowers energy
bonding electrons
Chem 101
21
Bond order
Bond order = ½ {no. bonding electrons – no. antibonding electrons}
antibonding electrons
11
Section 9.1-9.6
Metal bonding
(Sections 23.5 and 12.2)
MO model
an infinite chain of atoms
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Section 9.1-9.6
Metals, insulators and semiconductors
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