Unit 6
Molecular
Structure &
Covalent Bonding
Theories
Slides courtesy Brooks/Cole
Stereochemistry
 Stereochemistry is the study of the three
dimensional shapes of molecules.
 Some questions to examine in this chapter
are:
1.
2.
3.
4.
Why are we interested in shapes?
What role does molecular shape play in life?
How do we determine molecular shapes?
How do we predict molecular shapes?
2
Two Simple Theories of Covalent Bonding
 Valence Shell Electron Pair Repulsion Theory
 Commonly designated as VSEPR
 Helps us to predict the spatial arrangement of atoms in a
polyatomic molecule or ion
 It does not explain how bonding occurs just where it occurs &
where unshared pairs of valence e-s are directed.
 Valence Bond Theory
 Describes how the boding takes place in terms of overlapping
orbitals
 Involves the use of hybridized (mixed) atomic orbitals
 Used together they enable us to understand bonding,
molecular shapes and properties of polyatomic molecules
and ions.
3
VSEPR Theory

Regions of high electron density (electron
groups) around the central atom are arranged
as far apart as possible to minimize repulsions.


Central atom – any atom that is bonded to more than
one other atom
The number of electron groups around the
central atom are counted as follows:
1. Each bonded atom is counted as one e- group for VSEPR,
regardless of whether the bonding is single, double or triple.
2. Each lone pair of valence e-s on the central atom is counted
as one e- group for VSEPR.
4
VSEPR Theory
 There are five basic molecular shapes
based on the number of regions of high
electron density around the central atom.
 Some molecules may have more than one
central atom, in such a case, we determine the
arrangement/ shape around each central atom
to get an overall shape of the molecule.
 Several modifications of these five basic
shapes will also be examined.
5
VSEPR Theory
 Two regions of high electron density around
the central atom.
6
VSEPR Theory
 Three regions of high electron density
around the central atom.
7
VSEPR Theory
 Four regions of high electron density
around the central atom.
8
VSEPR Theory
 Five regions of high electron density around
the central atom.
9
VSEPR Theory
 Six regions of high electron density around
the central atom.
10
VSEPR Theory
Frequently, we will describe two geometries for
each molecule.
 Electronic geometry is determined by the
locations of regions of high electron density
around the central atom(s).
 Molecular geometry is determined by the
arrangement of atoms around the central atom(s).
 Electron pairs are not used in the molecular geometry
determination just the positions of the atoms in the
molecule are used.
11
VSEPR Theory
 An example of a molecule that has the
same electronic and molecular geometries
is methane, CH4.
 Electronic and molecular geometries are
tetrahedral.
12
VSEPR Theory
 An example of a molecule that has different
electronic and molecular geometries is
water, H2O.
 Electronic geometry is tetrahedral.
 Molecular geometry is bent or angular.
13
VSEPR Theory
 Lone pairs of electrons (unshared pairs)
require more volume than shared pairs.
 Consequently, there is an ordering of
repulsions of electrons around central atom.
 Criteria for the ordering of the repulsions:
14
VSEPR Theory
1. Lone pair to lone pair is the strongest repulsion.
2. Lone pair to bonding pair is intermediate
repulsion.
3. Bonding pair to bonding pair is weakest
repulsion.
 Mnemonic for repulsion strengths
lp/lp > lp/bp > bp/bp
 Lone pair to lone pair repulsion is why bond
angles in water are less than 109.5o.
15
Polar Molecules: The Influence of Molecular
Geometry
 Molecular geometry affects molecular
polarity.
 Due to the effect of the bond dipoles and how
they either cancel or reinforce each other.
16
Polar Molecules: The Influence of Molecular
Geometry

For a molecule to be polar, two conditions must
both be met:
1. There must be at least one polar bond or one lone pair of
electrons on central atom.
2. Neither bonds nor lone pairs can be symmetrically arranged
so that their polarities cancel.

In other words, if there are no polar bonds or unshared
e-s on the central atom  molecule is non-polar AND
A molecule can have individual bond dipoles but the
entire molecule may be non-polar  if bond dipoles
cancel.


E.g. compare CO2 with H2O
17
Polar Molecules: The Influence of Molecular
Geometry
18
A guide to determining whether a polyatomic molecule is polar
or nonpolar
Fig. 8-1, p. 292
Valence Bond (VB) Theory
 Covalent bonds are formed by the overlap of atomic orbitals.
 Atomic orbitals on the central atom can mix
 Process is called hybridization.
 Accounts for the observed geometries of molecules
 Explains how it is possible for larger molecules and polyatomic ions
can form
 The number of hybrid orbitals = number of AOs mixed
 The type of hybrid orbitals obtained varies with the types of
AOs mixed
 Example:
 2s and the three 2p orbitals  four sp3 hybrid orbitals
 Hybrid Orbitals have the same shapes as predicted by
VSEPR.
20
sp3 Hybrid Atomic Orbitals (Example: C atom)
21
Fig. 3-6, p. 69
Sometimes N and O atoms also have sp3
hybrid orbitals
22
sp2 Hybridization
23
sp Hybridization
24
Fig. 3-14, p. 76
Valence Bond (VB) Theory
25
Molecular Shapes and Bonding
 In the next sections we will use the
following terminology:
A = central atom
B = bonding pairs around central atom
U = lone pairs around central atom
 For example:
AB3U designates that there are 3 bonding pairs
and 1 lone pair around the central atom.
26
Linear Electronic Geometry:AB2 Species
(No Lone Pairs of Electrons on A)
 Some examples of molecules with this
geometry are:
 BeCl2, BeBr2, BeI2, HgCl2, CdCl2
 All of these examples are linear, nonpolar
molecules.
 Important exceptions occur when the two
substituents are not the same!
 BeClBr or BeIBr will be linear and polar!
27
Linear Electronic Geometry:AB2 Species
(No Lone Pairs of Electrons on A)
Dot Formula
Electronic Geometry
Linear
28
Linear Electronic Geometry:AB2 Species
(No Lone Pairs of Electrons on A)
Molecular Geometry
Polarity
Very polar bonds
Symmetrical dipole cancel
29
Linear Electronic Geometry:AB2 Species
(No Lone Pairs of Electrons on A)
Electronic Structures
Lewis Formulas
4Be
1s

3s
17Cl
[Ne] 
2s 2p

3p
  
30
Linear Electronic Geometry:AB2 Species
(No Lone Pairs of Electrons on A)
Valence Bond Theory (Hybridization)
31
Linear Electronic Geometry:AB2 Species
(No Lone Pairs of Electrons on A)
32
Trigonal Planar Electronic Geometry: AB3 Species
(No Lone Pairs of Electrons on A)
 Some examples of molecules with this
geometry are:
 BF3, BCl3
 All of these examples are trigonal planar,
nonpolar molecules.
 Important exceptions occur when the three
substituents are not the same!
 BF2Cl or BCI2Br will be trigonal planar and
polar!
33
Trigonal Planar Electronic Geometry: AB3 Species
(No Lone Pairs of Electrons on A)
34
p. 296
Trigonal Planar Electronic Geometry: AB3 Species
(No Lone Pairs of Electrons on A)
Electronic Structures
Lewis Formulas
B
1s

2s

2p

F [He]
3s
3p
   
·· .
B
F
35
Trigonal Planar Electronic Geometry: AB3 Species
(No Lone Pairs of Electrons on A)
 Again we can use VB theory to explain B-F
bonds
 The 2s and two of the 2p orbitals of B hybridize to
for a set of three equivalent sp2 hybrid orbitals
36
p. 297
 The sp2 hybrid orbitals point toward the corners of an
equilateral triangle.
 We can imagine that there is 1 e- in each hybrid orbital
 Each the F atoms has a 2p orbital with one unpaired e the 2p orbital can overlap with the sp2 hybrid orbitals on B
37
Trigonal Planar Electronic Geometry: AB3 Species
(No Lone Pairs of Electrons on A)
sp2 hybridization occurs at the
central atom whenever there are
3 electron groups around the
central atom
AB3 molecules & ions with no
lone pairs on the central atom
have trigonal planar electronic
AND molecular geometry as well
as sp2 hybridization on the
central atom
38
Trigonal Planar Electronic Geometry: AB3 Species
(No Lone Pairs of Electrons on A)
39
p. 297
Tetrahedral Electronic Geometry: AB4 Species
(No Lone Pairs of Electrons on A)
 Some examples of molecules with this
geometry are:
 CH4, CF4, CCl4, SiH4, SiF4
 All of these examples are tetrahedral,
nonpolar molecules.
 Important exceptions occur when the four
substituents are not the same!
 CF3Cl or CH2CI2 will be tetrahedral and polar!
40
p. 299
Tetrahedral Electronic Geometry: AB4 Species
(No Lone Pairs of Electrons on A)
Polarity
H
Molecular Geometry
H
H
C
H
H
tetrahedral
H
C
H
H
nonpolar molecule
C - H
Electronegativities 2.5
2.1
1
424
3
0.4
non polar bonds
42
Tetrahedral Electronic Geometry: AB4 Species
(No Lone Pairs of Electrons on A)
 All AB4 molecules in which there are no unshared
e- pairs on the central element and all 4 B atoms
are identical  will be non-polar
43
p. 300
Tetrahedral Electronic Geometry: AB4 Species
(No Lone Pairs of Electrons on A)
 If the atoms bonded to
the central atom are
not all identical then
such molecules are
usually polar
44
p. 300
Tetrahedral Electronic Geometry: AB4 Species
(No Lone Pairs of Electrons on A)
 According to VB theory for a tetrahedral
arrangement the central atom must make 4
equivalent orbitals
 Four sp3 hybrid orbitals are formed by mixing the s and
all three p orbitals in the outer shell of the central atom
 This results in 4 unpaired e-s
45
p. 301
Tetrahedral Electronic Geometry: AB4 Species
(No Lone Pairs of Electrons on A)
 The sp3 hybrid orbitals are directed toward
the corners of a regular tetrahedron
46
Tetrahedral Electronic Geometry: AB4 Species
(No Lone Pairs of Electrons on A)
 Each of the 4 atoms that bond to C has a half-filled atomic
orbital: these can overlap the half-filled sp3 hybrid orbital
47
Tetrahedral Electronic Geometry: AB4 Species
(No Lone Pairs of Electrons on A)
 sp3 hybridization occurs at
the central atom whenever
there are 4 electron groups
around the central atom
 AB4 molecules and ions with
no lone pairs on the central
atom have a tetrahedral
electronic AND molecular
geometry as well as sp3
hybridization
48
p. 302
Example of Molecules with More Than One
Central Atom
 It is difficult to assign
one geometry to
compounds with
more than one
central atom.
 Can get an overall
idea about shape by
examining the
geometry around
each central atom
H
CH4
C H
H
H
H
H C
H
H
H H
C2H6
H
H
C3H8
C C H
H H
H
H
H C
H
H
H
H
C H
C
H H H
C
H
H C
H
H
49
Example of Molecules with More Than One Central
Atom
The electronic and molecular geometry at each C atom of ethane
is tetrahedral
50
p. 303
Tetrahedral Electronic Geometry: AB3U Species
(One Lone Pair of Electrons on A)
 Some examples of molecules with this geometry
are:
 NH3, NF3, PH3, PCl3, AsH3
 These molecules are our first examples of central
atoms with one lone pair of electrons.
 Thus, the electronic and molecular geometries are
different.
 All three substituents are the same but molecule is
polar.
 NH3 and NF3 have a trigonal pyramidal molecular
geometry and are polar molecules.
51
Tetrahedral Electronic Geometry: AB3U Species
(One Lone Pair of Electrons on A)
 Both NH3 and NF3 have 4 e- groups around
the central atoms  tetrahedral electronic
geometry
52
p. 304
Tetrahedral Electronic Geometry: AB3U Species
(One Lone Pair of Electrons on A)
 The lone pair of e-s on the N atom repel the shared e-s of
the N-H and N-F bonds  bond angle reduced (as opposed
to 109.5o for tetrahedral shape
53
p. 305
Tetrahedral Electronic Geometry: AB3U Species
(One Lone Pair of Electrons on A)
54
p. 305
p. 305
Tetrahedral Electronic Geometry: AB3U Species
(One Lone Pair of Electrons on A)
Valence Bond Theory (Hybridization)
To figure the hybridization on the central atom we need to
look at the electronic geometry around the central atom
56
Tetrahedral Electronic Geometry: AB3U Species
(One Lone Pair of Electrons on A)
Valence Bond Theory (Hybridization)
57
p. 307
Tetrahedral Electronic Geometry: AB2U2 Species
(Two Lone Pairs of Electrons on A)
 Some examples of molecules with this geometry
are:
 H2O, OF2, H2S
 These molecules are our first examples of central
atoms with two lone pairs of electrons.
 Thus, the electronic and molecular geometries are
different.
 Both substituents are the same but molecule is polar.
 Molecules are angular, bent, or V-shaped and
polar.
58
Tetrahedral Electronic Geometry: AB2U2 Species
(Two Lone Pairs of Electrons on A)
Electronic Structures
Lewis Formulas
O [He]
2s
2p
   
H
1s

··
·· O .
.
H .
59
Tetrahedral Electronic Geometry: AB2U2 Species
(Two Lone Pairs of Electrons on A)
Polarity
O - H
Electroneg ativities 3.5
2.1
1
424
3
Molecular Geometry
··
1.4
2 lone pairs
H
ver y polar bonds
O
··
H
bent, angular
or V-shaped
bond dipoles
reinforce lone
pairs
··
H
O
H
··
asymetric dipoles
very polar molecule
1.7 D
60
Tetrahedral Electronic Geometry: AB2U2 Species
(Two Lone Pairs of Electrons on A)
Valence Bond Theory (Hybridization)
O [He]
2s
 
2p
four sp3 hybrids
  
61
Tetrahedral Electronic Geometry:
ABU3 Species
(Three Lone Pairs of Electrons on A)
 Some examples of molecules with this geometry
are:
 HF, HCl, HBr, HI, FCl, IBr
 These molecules are examples of central atoms
with three lone pairs of electrons.
 Again, the electronic and molecular geometries are
different.
 Molecules are linear and polar when the two
atoms are different.
 Cl2, Br2, I2 are nonpolar.
62
Tetrahedral Electronic Geometry:
ABU3 Species - (Three Lone Pairs of Electrons on A)
Dot Formula
Electronic Geometry
··
H ·· F ··
··
··
H
F
··
··
tetrahedral
63
Tetrahedral Electronic Geometry:
ABU3 Species
(Three Lone Pairs of Electrons on A)
Molecular Geometry
Polarity
HF is a polar molecule.
··
H
F
··
3 lone pairs
··
linear
64
Tetrahedral Electronic Geometry:
ABU3 Species
(Three Lone Pairs of Electrons on A)
Valence Bond Theory (Hybridization)
F [He]
four sp3 hybrids
   
2s
2p
  
··
H
F
··
··
tetrahedral
65
Trigonal Bipyramidal Electronic Geometry:
AB5, AB4U, AB3U2, and AB2U3
 Some examples of molecules with this geometry
are:
PF5, AsF5, PCl5, etc.
 These molecules are examples of central atoms
with five bonding pairs of electrons.
The electronic and molecular geometries are the same.
 Molecules are trigonal bipyramidal and nonpolar
when all five substituents are the same.
If the five substituents are not the same polar molecules
can result, AsF4Cl is an example.
66
Trigonal Bipyramidal Electronic Geometry:
AB5, AB4U, AB3U2, and AB2U3
Electronic Structures
Lewis Formulas
As [Ar] 3d10
F [He]
2s

4s
4p
   
2p
  
··
.. As
.
.
As
.
··
·· F
.
··
67
Trigonal Bipyramidal Electronic Geometry:
AB5, AB4U, AB3U2, and AB2U3
Dot Formula
Electronic Geometry
··
·· F
··
··
·· F ··
··
··
·· As
·· ··
·· F ·
·
··
··
··
F
··
··
F
··
··
··
··
··
As ·
·
··
trigonal bipyramidal
68
Trigonal Bipyramidal Electronic Geometry:
AB5, AB4U, AB3U2, and AB2U3
Molecular Geometry
··
·· F ·· ··
F ··
··
·· F As ··
·· ·
··
F·
··
·
·· F ·
··
trigonal bipyramid
Polarity
As - F
Electroneg ativities 2.1
4.0
1
424
3
1.9
··
· F · ·· ve ry polar bonds
· ·
F ··
··
· F As ··
·· ·
·
··
F·
··
· F ··
· ··
symmetric dipoles cancel
nonpolar molecule
69
Trigonal Bipyramidal Electronic Geometry:
AB5, AB4U, AB3U2, and AB2U3
Valence Bond Theory (Hybridization)
As [Ar] 3d10
______
4s
4p
   
five sp3 d hybrids
    
4d
_________

4d
____________
70
Trigonal Bipyramidal Electronic Geometry:
AB5, AB4U, AB3U2, and AB2U3
If lone pairs are incorporated into the trigonal
bipyramidal structure, there are three possible new
shapes.

1.
2.
3.
One lone pair - Seesaw shape
Two lone pairs - T-shape
Three lone pairs – linear
The lone pairs occupy equatorial positions because
they are 120o from two bonding pairs and 90o from
the other two bonding pairs.


Results in decreased repulsions compared to lone pair in
axial position.
71
Trigonal Bipyramidal Electronic Geometry:
AB5, AB4U, AB3U2, and AB2U3
 AB4U molecules have:
1. trigonal bipyramid electronic geometry
2. seesaw shaped molecular geometry
3. and are polar
 One example of an AB4U molecule is
SF4
 Hybridization of S atom is sp3d.
72
Trigonal Bipyramidal Electronic Geometry:
AB5, AB4U, AB3U2, and AB2U3
Molecular Geometry
73
Trigonal Bipyramidal Electronic Geometry:
AB5, AB4U, AB3U2, and AB2U3
 AB3U2 molecules have:
1. trigonal bipyramid electronic geometry
2. T-shaped molecular geometry
3. and are polar
 One example of an AB3U2 molecule is
IF3
 Hybridization of I atom is sp3d.
74
Trigonal Bipyramidal Electronic Geometry:
AB5, AB4U, AB3U2, and AB2U3
Molecular Geometry
75
Trigonal Bipyramidal Electronic Geometry:
AB5, AB4U, AB3U2, and AB2U3
 AB2U3 molecules have:
1.trigonal bipyramid electronic geometry
2.linear molecular geometry
3.and are nonpolar
 One example of an AB3U2 molecule is
XeF2
 Hybridization of Xe atom is sp3d.
76
Trigonal Bipyramidal Electronic Geometry:
AB5, AB4U, AB3U2, and AB2U3
Molecular Geometry
77
Octahedral Electronic Geometry:
AB6, AB5U, and AB4U2
 Some examples of molecules with this
geometry are:
 SF6, SeF6, SCl6, etc.
 These molecules are examples of central
atoms with six bonding pairs of electrons.
 Molecules are octahedral and nonpolar
when all six substituents are the same.
 If the six substituents are not the same polar
molecules can result, SF5Cl is an example.
78
Octahedral Electronic Geometry:
AB6, AB5U, and AB4U2
Electronic Structures
Lewis Formulas
Se [Ar] 3d10
F [He]
4s

2s

4p
  
2p
  
··
·· Se .
.
··
·· F .
··
79
Octahedral Electronic Geometry:
AB6, AB5U, and AB4U2
Molecular Geometry
Polarity
F
F
F
Electroneg ativities
F
Se - F
2.4
4.0
1
424
3
1.6
Se
very polar bonds
F
F
F
F
octahedral
F
F
Se
F
H
H C
H
H
F
symmetric dipoles cancel
nonpolar molecule
80
Octahedral Electronic Geometry:
AB6, AB5U, and AB4U2
Valence Bond Theory (Hybridization)
Se [Ar] 3d10
4s

4p
  
4d
__________

six sp3 d2 hybrids
     
4d
______
81
Octahedral Electronic Geometry:
AB6, AB5U, and AB4U2
 If lone pairs are incorporated into the octahedral
structure, there are two possible new shapes.
1. One lone pair - square pyramidal
2. Two lone pairs - square planar
 The lone pairs occupy axial positions because they are
90o from four bonding pairs.

Results in decreased repulsions compared to lone pairs in
equatorial positions.
82
Octahedral Electronic Geometry:
AB6, AB5U, and AB4U2
 AB5U molecules have:
1.octahedral electronic geometry
2.Square pyramidal molecular geometry
3.and are polar.
 One example of an AB5U molecule is
IF5
 Hybridization of I atom is sp3d2.
83
Octahedral Electronic Geometry:
AB6, AB5U, and AB4U2
Molecular Geometry
84
Octahedral Electronic Geometry:
AB6, AB5U, and AB4U2
 AB4U2 molecules have:
1.octahedral electronic geometry
2.square planar molecular geometry
3.and are nonpolar.
 One example of an AB4U2 molecule is
XeF4
 Hybridization of Xe atom is sp3d2.
85
Octahedral Electronic Geometry:
AB6, AB5U, and AB4U2
Molecular Geometry
86
Compounds Containing
Double Bonds
 Ethene or ethylene, C2H4, is the
simplest organic compound containing a
double bond.
Lewis dot formula
N = 2(8) + 4(2) = 24
A = 2(4) + 4(1) = 12
S
= 12
 Compound must have a double bond to
obey octet rule.
87
Compounds Containing
Double Bonds
Lewis Dot Formula
H·
H
·
·
·
C ·· ·· C
·
·· H
H·
H
H
C
or
C
H
H
88
Compounds Containing
Double Bonds
Valence Bond Theory (Hybridization)
C atom has four electrons.
Three electrons from each C atom are in sp2
hybrids.
One electron in each C atom remains in an
unhybridized p orbital
2s 2p
three sp2 hybrids 2p
C    


89
Compounds Containing
Double Bonds
 An sp2 hybridized C atom has this shape.
Remember there will be one electron in each of the three
sp2 lobes and one in the p orbital.
Top View
Side View
90
Compounds Containing
Double Bonds
 Two sp2 hybridized C atoms plus p orbitals in
proper orientation to form C=C double bond.
91
Compounds Containing
Double Bonds
 The portion of the double bond formed from the
head-on overlap of the sp2 hybrids is designated as
a s bond.
92
Compounds Containing
Double Bonds
 The other portion of the double bond, resulting
from the side-on overlap of the p orbitals, is
designated as a p bond.
93
Compounds Containing
Double Bonds
 Thus a C=C bond looks like this and is made of two
parts, one s and one p bond.
94
Compounds Containing
Triple Bonds
 Ethyne or acetylene, C2H2, is the simplest triple
bond containing organic compound.
Lewis Dot Formula
N = 2(8) + 2(2) = 20
A = 2(4) + 2(1) =10
S
= 10
 Compound must have a triple bond to obey octet
rule.
95
Compounds Containing
Triple Bonds
Lewis Dot Formula
H ·· C ·· ·· ··C ·· H
or
H C C H
VSEPR Theory suggests regions of high
electron density are 180o apart.
H
C
C
H
96
Compounds Containing
Triple Bonds
Valence Bond Theory (Hybridization)
Carbon has 4 electrons.
Two of the electrons are in sp hybrids.
Two electrons remain in unhybridized p
orbitals.
C [He]
2s

2p
two sp hybrids 2p
 


97
Compounds Containing
Triple Bonds
A s bond results from the head-on overlap of
two sp hybrid orbitals.
98
Compounds Containing
Triple Bonds
 The unhybridized p orbitals form two p bonds.

Note that a triple bond consists of one s and
two p bonds.
99
Summary of Electronic & Molecular
Geometries
100