Chapter 10: Chemical Bonding II: Molecular Shapes, Valence Bond Theory, and Molecular
Orbital Theory
1. Artificial Sweeteners: Fooled by Molecular Shape
a. In this chapter we will be looking at ways to predict and account for the shape of
molecules using the model called valence shell electron pair repulsion (VSEPR)
theory
2. VSEPR Theory: The Five Basic Shapes
a. Valence Shell Electron Pair Repulsion (VSEPR) theory is based on the idea that
electron groups (lone pairs, single bonds, multiple bonds, single electrons) repel
each other through Coulombic forces
i. This repulsion determines the geometry of the molecule
ii. The preferred geometry is where there is the least amount of repulsion
(groups have maximum separation)
b. Two electron groups
i. For example BeCl2
ii. There are two electron groups around the central atom so the maximum
separation is 180°
iii. This results in a linear geometry
c. Three electron groups
i. For example BF 3
ii. There are three electron groups around the central atom so the
maximum separation is 120°
iii. This results in a trigonal planar geometry
d. Four electron groups
i. For example CH4
ii. There are four electron groups around the central atom so the maximum
separation must be 109.5° (in 3D)
iii. This results in a tetrahedral geometry
e. Five electron groups
i. For example PCl5
ii. There are five electron groups around the central atom so the maximum
separation must be either 90° or 120° depending on the location of the
electron group
iii. This results in a trigonal bipyramidal geometry
f. Six electron groups
i. For example SF6
ii. There are six electron groups so the maximum separation must be 90
iii. This results in a octahedral geometry*
3. VSEPR Theory: The Effect of Lone Pairs
a. We can distinguish between the lone pairs and the bonding pairs by using what
we call molecular geometries
i. This is because electron repulsion strength is as follows:
1. 2 Lone pairs > lone pair – bonding pair > 2 bonding pairs
b. Four electron groups with lone pairs
i. For example NH3
1. One lone pair and three bonding pairs; we can call this trigonal
pyramidal
ii. For example H2O
1. Two lone pairs and two bonding pairs; we call this bent
Lone pairs make a difference in the bond angle as well:
c. Five electron groups with lone pairs
i. For example SF4
1. One lone pair and four bonding pairs; we call this seesaw
ii. For example BrF3
1. Two lone pairs and three bonding pairs; we call this T-shaped
iii. For example XeF2
1. Three lone pairs and two bonding pairs; we call this linear
d. Six electron groups with lone pairs
i. For example BrF5
1. One lone pair and five bonding pairs; we call this square
pyramidal
ii. For example XeF4
1. Two lone pairs and four bonding pairs; we call this square planar
4. VSEPR Theory: Predicting Molecular Geometries
a. To predict the geometry of a molecule, follow this procedure:
1. Draw a Lewis structure for the molecule
2. Determine the total number of electron groups around the central
atom
3. Determine the number of bonding groups and the number of lone
pairs around the central atom
4. Use table 10.1 (or handout) to determine the electron geometry
and the molecular geometry
b. When we write these structures on paper we use:
c. *Practice Worksheet
d. http://www.dartmouth.edu/~genchem/0102/spring/vsepr.html
e. Homework problems:
i. 1, 5, 6, 8, 9, 10, 12, 14, 16
5. Molecular Shape and Polarity
a. To determine if a molecule is polar (using geometries):
1. Draw a Lewis structure and determine the geometry
2. Determine if the molecule contains polar bonds (using EN)
3. Determine if the polar bonds add together to form a net dipole
moment
a. See table 10.2*
b. Homework problems:
i. 17, 18, 22
6. Valence Bond Theory: Orbital Overlap as a Chemical Bond
a. Valence bond theory is where electrons reside in quantum-mechanical orbits
(just like when doing electron configuration)
b. A chemical bond is the result from the overlap of two half-filled orbitals
c. The shape of the molecule is determined by the overlapping orbitals
i. For example H2S:
ii. Where the s orbitals from H overlap two of the p orbitals from S
7. Valence Bond Theory: Hybridization of Atomic Orbitals
a. This new theory is great, but there is something missing to be able to get the
correct shape for methane:
b. The reason for the difference is because these molecules use hybridization
i. Hybridization is a mathematical procedure in which the standard atomic
orbitals are combined to form new atomic orbitals called hybrid orbitals
c. Hybridization works by minimizing the energy of the molecule by maximizing the
orbital overlap in a bond
i. We will assume that only the outer orbitals (valence) get hybridized and:
1. The number of standard atomic orbitals (added together) always
equals the number of hybrid orbitals formed
2. The particular combinations of standard orbitals added together
determines the shape and energies of the hybrid orbitals
d. sp3 hybridization:
i. This is when one s orbital combines with three p orbitals to form 4 sp3
orbitals
e. sp2 hybridization:
i. This is when one s orbital combines with two p orbitals to form 3 sp2
orbitals
ii. If we looked at H2CO:
iii. When p orbitals overlap side by side we get a bond called a pi (π) bond
iv. When orbitals overlap end to end we get a bond called a sigma (σ) bond
f. sp hybridization:
i. This is when one s orbital combines with one p orbital to form 2 sp
orbitals
ii. You are NOT required to know any other kind of hybridization (although
it is theorized that they exist)
g. The bond angles and geometries for each type of hybridization are:
i. sp3
109°
tetrahedral
ii. sp2
120°
trigonal planar
iii. sp
180°
linear**
h. Homework problems:
i. 24, 26, 27, 29, 32, 33
8. Molecular Orbital Theory: Electron Delocalization
a. Valence bond theory has its limitations
b. If we used Schrödinger’s equation for molecules which is much better the math
gets insane, so we use something called Molecular Orbital (MO) Theory
i. Molecular Orbital Theory is where we use an educated “guess” as to
what the solution is
c. A bonding orbital is a molecular orbital that is lower in energy than any of the
atomic orbitals from which it was formed
i. When we put two or more orbitals together we get constructive
interference (bonding) and deconstructive interference (antibonding)
ii. We can assign a bond order (a number) that tells us if that molecule is
likely to form
1. We do NOT need to know how to do this 
2. BUT you should still look over what the orbitals look like on page
376 and 377
3. This is the reason that oxygen is paramagnetic when the electron
configuration is diamagnetic (see page 379 for further
explanation)
d. MO theory gives us a better idea as to how atoms bond, but it is just one of
many models that we have talked about and that you are responsible for
knowing
All chapter 10 homework problems: #1, 5, 6, 8, 9, 10, 12, 14, 16, 17, 18, 22, 24, 26, 27, 29, 32, 33
Review problems: #52, 55