• CHM 2045
• Molecular Geometry &
Chemical Bonding
Chapter 10
Using Lewis Theory to Predict
Molecular Shapes
• Lewis theory says that these regions of electron
groups should repel each other
• This idea can then be extended to predict the
shapes of molecules
– the position of atoms surrounding a central atom will be
determined by where the bonding electron groups are
– the positions of the electron groups will be determined by
trying to minimize repulsions between them
2
VSEPR Theory
• Electron groups around the central atom will be
most stable when they are as far apart as
possible – we call this valence shell electron
pair repulsion theory
• The resulting geometric arrangement will allow
us to predict the shapes and bond angles in the
molecule
Electron Group Geometry
• There are five basic arrangements of electron
groups around a central atom
• Each of these five basic arrangements results in
five different basic electron geometries
• For molecules that exhibit resonance, it doesn’t
matter which resonance form you use – the
electron geometry will be the same
4
Linear Electron Geometry
• When there are two electron groups around the
central atom, they will occupy positions on opposite
sides of the central atom
• This results in the electron groups taking a linear
geometry
• The bond angle is 180°
5
Trigonal Planar Electron Geometry
• When there are three electron groups around the
central atom, they will occupy positions in the shape
of a triangle around the central atom
• This results in the electron groups taking a trigonal
planar geometry
• The bond angle is 120°
6
Tetrahedral Electron Geometry
• When there are four electron groups around the
central atom, they will occupy positions in the shape
of a tetrahedron around the central atom
• This results in the electron groups taking a
tetrahedral geometry
• The bond angle is 109.5°
7
Tetrahedral Geometry
8
Trigonal Bipyramidal Electron Geometry
• When there are five electron groups around the
central atom, they will occupy positions in the
shape of two tetrahedra that are base-to-base
with the central atom in the center of the shared
bases.
• This results in the electron groups taking a
trigonal bipyramidal geometry
9
Trigonal Bipyramidal Geometry
10
Octahedral Electron Geometry
• When there are six electron groups around the
central atom, they will occupy positions in the shape
of two square-base pyramids that are base-to-base
with the central atom in the center of the shared
bases
• This results in the electron groups taking an
octahedral geometry
• All positions are equivalent
• The bond angle is 90°
11
Octahedral Geometry
12
Octahedral Geometry
13
Molecular Geometry
• The actual geometry of the molecule may be
different from the electron geometry
• When the electron groups are attached to
atoms of different size,
• when the bonding to one atom is different than
the bonding to another,
• When there are lone pairs since they occupy
space on the central atom, but are not “seen” as
points on the molecular geometry
14
The Effect of Lone Pairs
• Lone pair groups “occupy more space” on the
central atom
• Relative sizes of repulsive force interactions is
Lone Pair – Lone Pair > Lone Pair – Bonding Pair > Bonding Pair – Bonding Pair
• This affects the bond angles, making the bonding
pair – bonding pair angles smaller than expected
15
Bond Angle Distortion
from Lone Pairs
16
Bond Angle Distortion
from Lone Pairs
17
18
Molecular Geometry
Example 1
Which molecule or ion has a trigonal planar
molecular geometry?
a)PCl3
b)AsF3
c) HCN
d)HCCH
e)CO32–
Molecular Geometry
Example 2
Which molecule or ion has a trigonal pyramidal
molecular geometry?
a)CO32–
b)SO3
c) BF3
d)C2H4
e)SO32–
Molecular Geometry
Example 3
What is the bond angle in a linear molecule or
ion?
a)90°
b)109°
c) 120°
d)180°
Molecular Geometry
Example 4
The approximate H—C—C bond angle in ethane,
C2H6, is
a)60°.
b)90°.
c) 109°.
d)120°.
e)180°.
Molecular Geometry
Example 5
What is the H—O—H bond angle in water?
a)180°
b)120°
c) 90°
d)109°
e)slightly less than 109°
Dipole moment ,the measure of
charge separation
• A molecule has a dipole moment when there
is charge separation
• The strength of the dipole moment depends
on the difference in the electronegativity of
the atoms in the molecule
Dipole moment ,the measure of
charge separation
• Dipole moment can be shown with an arrow
with a cross on one end:
• An arrow indicates the direction in which the
electrons concentrate.
Molecule Polarity
The O─C bond is polar. The bonding
electrons are pulled equally toward both
O ends of the molecule. The net result is
a nonpolar molecule.
26
Molecule Polarity
The H─O bond is polar. Both sets of
bonding electrons are pulled toward the
O end of the molecule. The net result
is a polar molecule.
27
Predicting Polarity of Molecules
1. Draw the Lewis structure and determine the
molecular geometry
2. Determine whether the bonds in the
molecule are polar
a) if there are no polar bonds, the molecule
is nonpolar
3. Determine whether the polar bonds add
together to give a net dipole moment
28
Decide whether the following molecules Are polar
Trigonal
Bent
Trigonal
Planar
2.5
1. polar bonds, N-O
2. asymmetrical shape
polar
29
1. polar bonds, all S-O
2. symmetrical shape
nonpolar
Polar or Nonpolar
Example 6
Which of the following molecule is polar?
a)CF4
b)SO2
c) CS2
d)C2H4
e)C6H6
Polar or Nonpolar
Example 7
Which of the following compounds is nonpolar?
a)XeF2
b)HCl
c) SO2
d)H2S
e)N2O
Dipole Moment
Example 8
Which of the following molecules has a dipole
moment?
a)NF3
b)CCl4
c) SiCl4
d)SF6
e)BF3
Dipole Moment
Example 9
For which molecule or ion does the nitrogen
atom have the positive end of the dipole
moment?
a)NH4+
b)Ca3N2
c) HCN
d)AlN
e)NO
Valence Bond Theory
• Linus Pauling and others applied the principles
of quantum mechanics to molecules
• They reasoned that bonds between atoms
would occur when the orbitals on those atoms
interacted to make a bond
• The kind of interaction depends on whether the
orbitals align along the axis between the nuclei,
or outside the axis
34
Valence Bond Theory
• Linus Pauling and others applied the principles
of quantum mechanics to molecules
• They reasoned that bonds between atoms
would occur when the orbitals on those atoms
interacted to make a bond
• The kind of interaction depends on whether the
orbitals align along the axis between the nuclei,
or outside the axis
35
Valence Bond Theory – Hybridization
• One of the issues that arises is that the number
of partially filled or empty atomic orbitals did
not predict the number of bonds or orientation
of bonds
– C = 2s22px12py12pz0 would predict two or three
bonds that are 90° apart, rather than four bonds
that are 109.5° apart
36
Hybridization
• To adjust for these inconsistencies, it was
postulated that the valence atomic orbitals
could hybridize/mix before bonding took
place
– one hybridization of C is to mix all the 2s
and 2p orbitals to get four orbitals that
point at the corners of a tetrahedron
Unhybridized C Orbitals Predict the Wrong
Bonding & Geometry
38
Hybridization
• Some atoms hybridize their orbitals to maximize
bonding
• Hybridizing is mixing different types of orbitals
in the valence shell to make a new set of
degenerate orbitals
– sp, sp2, sp3, sp3d, sp3d2
• Same type of atom can have different types of
hybridization
– C = sp, sp2, sp3
39
Hybrid Orbitals
• The number of atomic orbitals combined = the
number of hybrid orbitals formed
• The number and type of standard atomic orbitals
combined determines the shape of the hybrid
orbitals
40
Orbital Diagram of the
sp3 Hybridization of C
41
Methane Formation with sp3 C
42
sp3 Hybridized Atoms
Orbital Diagrams
• Place electrons into hybrid and unhybridized valence orbitals
as if all the orbitals have equal energy
• Lone pairs generally occupy hybrid orbitals
sp3 hybridized atom
Unhybridized atom
2s

2s
43
 
2p
C

  
2p
N



  
2sp3
  
2sp3
Bonding with Valence Bond Theory
• According to valence bond theory, bonding
takes place between atoms when their atomic
or hybrid orbitals “overlap”
• To interact, the orbitals must either be aligned
along the axis between the atoms, or
• The orbitals must be parallel to each other and
perpendicular to the interatomic axis
44
Types of Bonds
• A sigma (s) bond results when the interacting atomic
orbitals point along the axis connecting the two bonding
nuclei
• A pi (p) bond results when the bonding atomic orbitals
are parallel to each other and perpendicular to the axis
connecting the two bonding nuclei
• The interaction between parallel orbitals is not as strong
as between orbitals that point at each other; therefore s
bonds are stronger than p bonds
45
46
Orbital Diagrams of Bonding
• “Overlap” between a hybrid orbital on one
atom with a hybrid or nonhybridized orbital
on another atom results in a s bond
• “Overlap” between unhybridized p orbitals on
bonded atoms results in a p bond
47
sp2 Hybridized Atoms
Orbital Diagrams

2s

2s
 
2p
  
2p
48
Tro: Chemistry: A Molecular Approach, 2/e
sp2 hybridized atom
C
3s
1p

N
2s
1p


Unhybridized atom
 
2sp2

2p
 
2sp2

2p
sp2 hybridization in Ethylene
• Two sp2-hybridized orbitals overlap to form a s
bond
• p orbitals overlap side-to-side to formation a pi
(p) bond
• sp2–sp2 s bond and 2p–2p p bond result in
sharing four electrons and formation of C-C
double bond
• Electrons in the p bond occupy regions on either
side of a line between nuclei
sp2 hybridization in Ethylene, C 2H4
sp hybridization in Acetylene, C2H2
Hybridization & Molecular Geometry
Hybridization
Example 11
The hybridization of the central atom in a
molecule is described as sp2. The arrangement
in space of the hybrid orbitals about that atom is
a) linear.
b) trigonal planar.
c)
tetrahedral.
d) trigonal bipyramidal.
Hybridization
Example 12
Which of the following statements is incorrect
regarding the water molecule?
a)The molecule is polar.
b)The hybridization of oxygen is sp3.
c) There are two lone pairs and two bonding
pairs on the central atom.
d)The molecular geometry is bent.
e)The hybridization of hydrogen is sp.
Hybridization
Example13
What is the hybridization of Se in SeF6?
a)sp
b)sp2
c) sp3
d)dsp3
e)d2sp3