Chapter 10 Notes
Chemical Bonding II
Molecular Shape,
Valence Bond Theory, and
Molecular Orbital Theory
Sections 10.1 – 10.8
Ch. 10 Problem Assignment
Pg. : 453 - 459
(Due test day)
(Sect. 10.2, 10.3, 10.4): 2, 3, 4, 5, 30, 31, 32, 34, 42
(Sect. 10.5): 7, 38 (ignore last part of directions), 40
(Sect. 10.6, 10.7): 11, 65, 66
I. Artificial Sweeteners: Fooled by Molecular Shape
Artificial sweeteners such as aspartame (Nutrasweet) taste sweet like
sugar but do not have the calories of sugar. This is because taste and
nutritional value are independent properties. The caloric value of
food depends on the amount of energy released when the food is
metabolized in the body. Many artificial sweeteners are not even
metabolized by the body- they just pass straight though.
The taste of a food begins with specialized cells on the tongue that
detect different molecules of food. These cells are so sensitive that
the tongue can detect one molecule of sugar out of thousands of
different molecules in a bite of food. We experience certain tastes
when molecules of food fit into a special part of a taste receptor on
our tongue, much in the same way a key fits into a lock. Artificial
sweeteners “trick” our tongues into tasting sweetness because these
molecules mimic the molecular shape of the sugar molecule and fit
snuggly into our taste receptors.
II. VSEPR Theory: The Five Basic Shapes
• VSEPR THEORY (VALENCE SHELL ELECTRON PAIR REPULSION):
- used to determine the 3- D shape of a molecule
- based on the idea that electrons groups (bonds & lone pair e─)
repel one another to varying degrees
- the combination of repulsions on the central atom of a molecule
determines its 3-D shape
*The basic idea behind VSEPR is that repulsions between electron
groups determine molecular geometry. The preferred geometry is
the one in which the electron groups have the maximum separation
(and therefore the minimum energy) possible.
Preview
We will first look at molecular geometries where there are two to
six electron groups around a central atom and all of the electron
groups are bonding groups (single or double bonds). We will then
look at what happens to the shapes when some of the electron
groups become lone pair electrons.
A. TWO ELECTRONS GROUPS: LINEAR GEOMETRY
 Bond Angles: 180°
Other examples:
SiO2
16
CS2
16
O
Si
O
S
C
S
B. THREE ELECTRONS GROUPS: TRIGONAL PLANAR
GEOMETRY
Other examples:
 Bond Angles: 120°
BH3
H
SO3 24
6
H
O
B
S
H
O
O
In CH2O, the bond angles deviate slightly from
120° because there is more electron density in a
double bond than in a single bond.
Different electron groups repel each other in
slightly different ways.
C. FOUR ELECTRONS GROUPS: TETRAHEDRAL GEOMETRY
 Bond Angles: 109.5°
Other examples:
SiCl4
Cl
CF4 32
32
Cl
F
Si
C
Cl
Cl
F
F
F
D. FIVE ELECTRONS GROUPS: TRIGONAL BIPYRAMIDAL
GEOMETRY
 Bond Angles: 90°, 120°
Other examples:
SOF4
ClO2F3 40
40
O
O
F
F
S
Cl
F
F
F
F
F
O
E. SIX ELECTRONS GROUPS: OCTAHEDRAL GEOMETRY
 Bond Angles: 90°
Other examples:
SCl6 48
Cl
Cl
Cl
S
Cl
Cl
Cl
III. VSEPR Theory: The Effect of Lone Pairs
Preview
We will now look at molecular geometries where there are four to
six electron groups around a central atom and some of the electron
groups are lone pairs. Keep in mind that these shapes are all
variations of the geometries from the last section, but now one or
more bonding groups have been replaced with lone pairs. The
difference in shapes results from the fact that lone pair electrons
repel other lone pair and bonding electrons to a greater extent than
bonding electrons repel one another.
THE ORDER OF ELECTRON PAIR REPULSION:
lone pair - lone pair > lone pair- bonding pair > bonding pair- bonding pair
Most repulsive
Least repulsive
* Want to be as far away
from each other as
possible
THE DIFFERENCE BETWEEN:
 Electron Geometry: The geometrical arrangement of electron groups
Molecular Geometry: The geometrical arrangement of the atoms
• In this section, don’t get confused between electron and molecular
geometries.
• In the last section, the electron and molecular geometries were the same
(same name for each).
• In this section, they are different.
• There are only 5 electron geometries, but there are 11 molecular
geometries.
A. FOUR ELECTRON GROUPS WITH LONE PAIRS
(TETRAHEDRAL ELECTRON GEOMETRY)
TRIGONAL PYRAMIDAL (MOLEC. GEO.): ONE LONE PAIR
 Bond Angles: 107°
Other examples:
PF3 26
NCl3 26
P
F
N
F
F
Cl
Cl
Cl
A. FOUR ELECTRON GROUPS WITH LONE PAIRS
(TETRAHEDRAL ELECTRON GEOMETRY)
BENT (MOLEC. GEO.): TWO LONE PAIRS
 Bond Angles: 104.5°
Other examples:
SCl2 18
Cl
H2S
S
Cl
H
8
S
H
Summarizing Tetrahedral Electron Geometries
B. FIVE ELECTRON GROUPS WITH LONE PAIRS
(BIPYRAMIDAL ELECTRON GEOMETRY)
SEESAW (MOLEC. GEO.): ONE LONE PAIR
 Bond Angles: 90°, 120°
Other examples:
SeCl4
IOF3 34
34
Cl
O
Cl
F
I
Se
Cl
Lone pair must go equatorial
Cl
F
F
B. FIVE ELECTRON GROUPS WITH LONE PAIRS
(BIPYRAMIDAL ELECTRON GEOMETRY)
T-SHAPED OR LINEAR (MOLEC. GEO.): TWO OR THREE
LONE PAIRS
 Bond Angles: 90°, 120°
Other examples:
I3 -
ClF3 28
F
Cl
I
F
F
Lone pairs must go equatorial
I
I
28
C. SIX ELECTRON GROUPS WITH LONE PAIRS
(OCTAHEDRAL ELECTRON GEOMETRY)
SQUARE PYRAMIDAL (MOLEC. GEO.): ONE LONE PAIR
SQUARE PLANAR (MOLEC. GEO.): TWO LONE PAIRS
 Bond Angles: 90°
Other examples:
BrF4-
XeOF4 42
36
O
F
F
F
Xe
F
F
Br
F
Lone pairs must go axial
F
F
Summary:
There are 5 ELECTRON Geometries:
LINEAR : 2 e- groups around central atom
TRIGONAL PLANAR : 3 e- groups around central atom
TETRAHEDRAL : 4 e- groups around central atom
TRIGONAL BIPYRAMIDAL : 5 e- groups around central atom
OCTAHEDRAL : 6 e- groups around
central atom
These electron geometries are also the MOLECULAR geometries
For molecules where all e- groups are bonding groups.
Summary:
There are 6 additional MOLECULAR Geometries.
These occur when one or more of the e- groups are lone-pair e-
Bent : 3 e- groups around central atom- 1 is a lone-pair
4 e- groups around central atom- 2 are a lone pair
Trigonal Pyramidal : 4 e- groups around central atom- 1 is a lone-pair
Seesaw: 5 e- groups around central atom- 1 is a lone-pair
T-Shaped: 5 e- groups around central atom- 2 are a lone-pair
Square Pyramidal: 6 e- groups around central atom- 1 is a lone-pair
Square Planar: 6 e- groups around central atom- 2 are a lone-pair
IV. VSEPR Theory: Predicting Molecular Geometries
Procedure:
1. Draw a Lewis structure for the molecule & total the valence e─
2. Determine the total number of electron groups around the
central atom: lone pairs, single, double, and triple bonds each
count as one electron group
3. Determine the number of (1) bonding groups and (2) the
number of lone pair groups around the central atom.
These should total the sum from Step 2
4. Use the VSEPR Table to determine the electron geometry and
the molecular geometry
Use pgs. 414-415 in your text to fill out the table
V. Molecular Shape and Polarity
Recall that in Ch 9, we said a bond is polar when there is an uneven
sharing of electrons between atoms of different electronegativities.
Now we will determine if an entire molecule is polar by looking at
both its bonds and shape. Just because a molecule contains polar
bonds does NOT mean it is a polar molecule. We must also look at
its molecular geometry.
To Determine if a Molecule is Polar:
1. Draw the molecule with the correct molecular geometry
2. Determine if each bond in the molecule is polar by considering
electronegativity differences. If the bond is polar ( ≥ 0.4 ∆EN),
draw an arrow over it pointing toward the more EN atom
3. Determine whether the polar bonds add together to form a net
dipole moment, which makes the molecule polar (See Table 10.2)
A
B
O
C
O
Dipoles cancel;
molecule is nonpolar
C
D
O
H
E
Others become polar
when bonds are NOT
identical
H
Dipoles do NOT cancel;
molecule is polar
F
N
H
C
H
Dipoles do NOT cancel;
molecule is polar
H
H
F: 4.0
C: 2.5
H: 2.1
Dipoles do NOT
H cancel;
molecule is polar
H
• In Lewis Theory, we use “dots” to represent valence electrons.
Although this theory helps us to understand bonding, it is not an
actual representation of how bonding between atoms truly occurs.
• In the next three sections we will learn about Valence Bond Theory
and Molecular Orbital Theory. These theories attempt to explain
how bonding actually occurs at the atomic level.
VI. Valence Bond Theory
• VALENCE BOND THEORY:
• e─ reside in atomic orbitals and/or hybridized atomic orbitals
•
bonds result when half-filled orbitals overlap
•
bonds are localized between atoms
Let’s apply the general concepts of valence bond theory to underStand the bonding in H2S:
VII. Valence Bond Theory: Hybridization of Orbitals
Although the overlap of half-filled orbitals adequately explains
the bonding in some molecules such as H2S, it cannot explain
the bonding in many other molecules. Let’s see how this idea
doesn’t predict the bonding of C and H in CH4, a stable
molecule that we know forms.
Because C only has two half-filled orbitals,
this model would suggest that carbon only
bonds with 2 H’s. Again, we know it bonds with 4.
To better understand bonding in Valence Bond Theory, we must
consider hybridization.
• HYBRIDIZATION: a mathematical procedure in which atomic
orbitals are combined to form new atomic
orbitals called hybrid orbitals
Some General Statements Regarding Hybridization:
• The number of standard atomic orbitals added together always
equals the number of hybrid orbitals formed. In order words, the
numbers of orbitals is preserved.
• The combination of orbitals determines the shapes and energies of
the orbitals formed.
• The specific type of hybridization that forms for a molecule is the
type that lowers the overall energy of the molecule.
A. SP3 HYBRIDIZATION
• One s orbital and three p orbitals hybridize to form four sp3 hybrid orbitals
of equal energy.
• Head-on overlap of sp3 orbitals only forms single bonds called
σ (sigma) bonds.
Now we get a better idea of how CH4 bonds.
Using sp3 hybrid orbitals, C now has four half-filled orbitals
to overlap with 4 H 1s orbitals, each of which is half-filled.
Remember: according to valence bond theory, bonds form
when half – filled orbitals overlap
B. SP2 HYBRIDIZATION
• One s orbital and two p orbitals hybridize to form three sp2 hybrid orbitals
of equal energy.
• One p orbital remains unhybridized  used to form a double bond
• Head-on overlap of sp2 orbitals only forms single bonds called
σ (sigma) bonds.
• Sideways overlap of two unhybridized p orbitals forms pi (π) bonds.
A DOUBLE BOND CONSISTS OF ONE SIGMA AND ONE PI BOND
- Double bonds consist of 1 sigma σ and 1 pi π bond.
- A σ bond forms from head-on overlap of two orbitals.
- A π bond forms from sideways overlap of two p orbitals.
C. SP HYBRIDIZATION
• One s orbital and one p orbital hybridize to form two sp hybrid orbitals
of equal energy.
• Two p orbitals remains unhybridized  used to form triple bonds
• Head-on overlap of sp orbitals only forms single bonds called
σ (sigma) bonds.
• Sideways overlap of four unhybridized p orbitals forms 2 π bonds.
DIFFERENCE IN BETWEEN DOUBLE AND TRIPLE BONDS:
•A double bond contains one sigma and one pi bond
•A triple bond contains one sigma and two pi bonds
Double bond
Triple bond
Summary of bond types:
Bond Type
# of σ
# of π
Single
1
0
Double
1
1
Triple
1
2
HOW TO DETERMINE HYBRIDIZATION and MOLECULAR
GEOMETRY for A MOLECULE:
Note: Each of the following count as ONE electron group:
A lone pair of e −, a single bond, a double bond, or a triple bond
Number of electron
groups around
Hybridization type
interior* atom
Electron geometry
around interior* atom
4
sp3
tetrahedral
3
sp2
trigonal planar
2
sp
linear
*An “interior atom” is an atom bonded to more than one other atom
Ex. Determine the hybridization type of the interior atoms in each
of the following molecules:
S
C
S
O
S
O
O
F
C
F
F
F
VIII. Molecular Orbital Theory
Molecular orbital theory is different from valence bond theory but
also attempts to explain how compounds actually bond at the atomic
level. In valence bond theory, we treated orbitals, whether hybridized or not, as being centered on individual atoms. In molecular
orbital theory, we will treat orbitals as overlapping the entire
molecule, not as “belonging” to individual atoms.
• MOLECULAR ORBITAL THEORY:
- e- reside in orbitals that span the entire molecule
 - molecular orbitals are made from atomic orbitals
To form our molecular orbitals, we will take a Linear Combination
of Atomic Orbitals (LCAO). This is a mathematical procedure that
is often performed via computers.
As in valence bond theory, when atomic orbitals are combined to
form molecular orbitals, the total number of orbitals does not change.
• Molecular orbital theory is a complicated theory that you can
understand at its basic level.
• It can help us to understand if bonds will form between atoms and
which type (double, triple, etc)
• Because we can often come to the same conclusion via simpler
theories, and because of time constraints, we will not study this
theory in more detail at this point.
• The key is to realize that more complicated bonding theories do exist
and they can more accurately explain and predict how bonding
occurs
Let’s first examine how to take a LCAO for s orbitals:
Whenever two atomic orbitals overlap, we get:
• a bonding molecular orbital which is lower in energy than the
atomic orbitals
• an antibonding molecular orbital which is higher in energy than the
atomic orbitals
• each orbital can hold up to two electrons
We can write molecular orbital energy diagrams to understand
bonding in molecules.
• Here is a molecular orbital diagram for H2:
Only valence e- go here
•Lets draw a molecular orbital diagram for He2:
He atom
He atom
We know that He2 does NOT form. We can see why by looking at
the bond order:
Bond Order = (# of e- in bonding MO’s) - (# of e- in antibonding MO’s)
2
• A positive bond order means there are more electrons in
bonding than in antibonding orbitals and the overall energy of
the electrons is lower in the molecule than in the individual atoms
• A negative or zero bond order indicates that the energy of the
electrons is overall higher than it would be if the atoms did not
bond, and therefore no bond forms.
Now let’s examine how to take a LCAO for p orbitals to understand
how they form molecular orbitals. Remember that p orbitals exist in
sets of three.
When three p orbitals from each atom combine, we get six molecular
orbitals. Three are bonding and three are anitbonding.
Of the three bonding orbitals, one overlaps head-on and form a
sigma bond and two overlap sideways to form two pi bonds.
Sigma bonding and
antibonding molecular
orbitals From p atomic
orbitals
Two pi bonding and antibonding molecular orbitals made from
four p atomic orbitals
Here’s the part that’s slightly tricky:
• For B, C, and N the π MO is lower in energy than the σ
MO. The π* MO is lower in energy than the σ* MO.
• For O, F, and Ne, it’s opposite. The σ MO is lower in
energy than the π MO are lower in energy. BUT the π*
MO is lower in energy than the σ* MO.
DISCLAIMER:
MO theory is a very useful method that adequately explains
bonding and other chemical and physical properties better than
most other bonding theories combined.
With the aid of computers, MO theory can be used to calculate
molecular orbitals for entire molecules (as seen below for some
simple examples).
Keep in mind that we are just touching on the very basics of
MO theory here just so that you have been introduced to the
concept.
Neither Lewis Theory nor Valence Bond Theory adequately conveys
the fact that all of the bonds in ozone (O3) are equal in length and energy.
However, a computer-generated molecular orbital is able to clearly show
that all O-O bonds are equivalent.