MF111 Study Guide 3
Shapes of molecules
Learning outcomes
 Appreciate the biological significance of molecular shape
 Draw a Lewis structure for a given molecular formula
 Calculate the formal charges of each atom in a molecule
 Determine the most prevalent resonance structure of a molecule
 Use the VSEPR theory to predict electron geometry and molecular
shape
 Determine the overall polarity of molecules based on the
electronegativities
of component atoms and molecular shape
Overview
It is possible to predict a wide range of molecular properties based on
the molecular formula for a covalently bonded molecule. The Lewis
structure is applied to determine the distribution of electrons about the
central atom. Formal charges of atoms can be calculated to predict the
most prevalent resonance structure. The valence shell electron pair
repulsion (VSEPR) theory is used to predict the electron geometry and
molecular shape of the molecule. Molecular shape affects the overall
polarity of molecules.
Biological significance of molecular shapes
Biological receptors such as those shown in Figure 3.1 are precisely
shaped cavities, often within a membrane, designed to respond to
particular classes of molecules analogous to a “lock and key”. The same
“lock and key” analogy can also be used to explain enzyme specificity in
various biological systems.
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(Silberberg fig B10.1)
Figure 3.1 Biological receptors of the olfactory system
Figure 3.2 Molecules with different composition but similar molecular
shape
All of the molecules shown in Figure 3.2 have the same bark-like smell
as camphor. This is due to the fact that similar shape molecules can bind
to olfactory receptors regardless of their composition.
Lewis structures from molecular formulas
There are four steps in drawing Lewis structures from a given molecular
formula:
Below is an example of applying these procedures to a molecule of
carbonyl bromide (COBr2):
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Step 1: Carbon has the lowest electronegativity, so it will be placed in
the centre:
Occasionally, unique atom in a molecular formula is the central atom,
even when it is the most electronegative e.g. ammonia (NH3) and
chlorine oxide (Cl2O). Note: Hydrogen never is the central atom because
it can only form one bond.
Step 2: Count the valence shell electrons using the group number for
each atom:
Atoms
Group
C
O
Br
Br
Net charge
Total
4
6
7
7
0
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Valence
electrons
4
6
7
7
0
24
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Step 3: Connect the atoms with single bonds, and subtract a pair of
valences electrons from the total for each bond:
24 electrons − 3(2 electrons) = 18 electrons remaining
Step 4: Distribute the remaining valence electrons in pairs, starting with
the surrounding atoms, giving each atom 8 electrons, until all of the
valence electrons are used up:
If after distributing the remaining valence electrons, the central atom has
fewer than 8 valence electrons, bring in a pair of lone pair electrons from
one of the surrounding atoms and form a multiple bond. For carbonyl
bromide (COBr2), there are two possible structures:
OR
Lewis structures for organic molecules
The most common elements in organic molecules and the number of
bonds they form to get a full valence shell are shown as follows:




Hydrogen atom forms 1 bond
Carbon atom forms 4 bonds
Nitrogen atom forms 3 bonds
Oxygen atom forms 2 bonds
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Halogens are less commonly found in organic molecules but chlorinated
solvents such as chloroform and dichloromethane very common:
Dichloromethane
Chloroform
Figure 3.3 Chlorinated solvents of chloroform and dichloromethane.
Chlorine forms one bond when they are bonded to carbon.
Organic molecules with more than one central atom are also very
common. They are often formed around a “carbon backbone” as shown
by examples below:
Ethane
Ethanol
Acetone
Figure 3.4 Organic molecules with more than one central atom
Resonance structures
When a central bond has double or triple bond together with a single
bond, multiple resonance structures may be possible (Figure 3.5). The
mentioned COBr2 molecule has two possible resonance structures:
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Figure 3.5 Resonance structures for carbonyl bromide
Molecules and ions with resonance have the following basic
characteristics:
 They can be represented by several correct Lewis formulas i.e.
multiple resonance structures. However, the real structure is not a
rapid inter-conversion of different resonance structures. Instead, a
resonance hybrid would be a more accurate representation.
Electrons are delocalised and are shared freely among the adjacent
bonds. Figure 3.6 is an example of a resonance hybrid involving
benzene.
 The different resonance structures are not isomers. They differ
only in the position of electrons, not in the position of nuclei.
When transforming from one resonance structure into another, no
bonds are broken.
 Each Lewis formula must have the same number of valence
electrons and the same number of unpaired electrons.
 Bonds that have different bond orders in different contributing
structures do not have typical bond lengths (Recall: Study Guide 2).
Bonds with delocalised electrons have bond lengths that are
averaged among the possible single, double or triple bonds. The
length of C−C bonds in benzene has an intermediate length
between single and double bonds (Figure 3.6).
 Certain resonance structures are more stable (lower potential
energy) and would be the more prevalent structure.
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Resonance structures
Hybrid structure
Figure 3.6 Resonance structures and hybrid structure of benzene
Determining the more prevalent resonance structure
Both resonance structure of benzene are equally prevalent as they
involve identical atoms and bond orders. However, this is not always the
case as can be observed in the COBr2 example. The more prevalent
resonance structure is the most stable structure with the least potential
energy. The more prevalent structure can be predicted by the
determining the formal charges for each of the atoms in each structure:
Formal charge of atom
= No. of valence electrons − (no. of unshared
valence electrons
+ ½ no. of shared valence electrons)
The structure on the right has the smallest set of formal charges, making
it the more prevalent structure.
Other criteria used to predict prevalent resonance structures are:
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 Similar formal charges on adjacent atoms are unfavourable
 A more negative formal charge should reside on a more
electronegative ion
VSEPR theory
The valence shell electron pair repulsion (VSEPR) theory is a model in
chemistry used to predict the shape of individual molecules based on the
extent of electron-pair electrostatic repulsion. It is also named GillespieNyholm theory after its two main developers. The acronym “VSEPR” is
sometimes pronounced “vesper” for ease of pronunciation.
The premise of VSEPR is that the valence electron pairs surrounding an
atom mutually repel each other, and will therefore adopt an arrangement
that minimizes this repulsion. Bonds and lone pairs of electrons would
be evenly spaced around the atom (Figure 3.7).
Electron geometry and molecular shape
There are four steps in determining molecular shapes according to the
VSEPR theory:
Using COBr2 as example:
Step 1: Draw the Lewis structure:
This has been done above.
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Step 2: Count the electron groups:
Two single bonds, one double bond, and no lone pairs = 3 electron
groups
The 3 groups of electrons give a trigonal planar electron geometry as
shown in figure:
Figure 3.7 Five basic shapes based on electron geometry of atoms with
two to six electron groups (bonds and lone pairs). Mutually repelling
electron pairs are analogous to balloons in that each balloon occupies
space pushing away other balloons.
Step 3: Determine the bond angles:
For trigonal planar geometry, the ideal VSEPR bond angle is 120°. Lone
pairs and double bonds, which require more space, will distort the bond
angles from the ideal angle.
Figure 3.8 The most prevalent carbonyl bromide structure
Since COBr2 has a double bond, we expect the Br–C–Br angle to be
slightly less than 120° and the Br–C–O angles to be slightly greater than
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120° (Figure 3.8).
Step 4: Determine the molecular shape:
This is done ignoring the lone pairs of electrons and observing the
resulting shape. For trigonal planar electron geometry there are two
possibilities (Figure 3.9):
i. Trigonal planar if there are no lone pairs of electrons
ii. Bent, if there is one pair of lone electrons
Since COBr2 has no lone pairs of electrons, it has a trigonal planar
molecular geometry. Other atoms with different electron geometry are
shown in Figure 3.10.
Figure 3.9 Possible shapes for molecules with trigonal planar electron
geometry
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Figure 3.10 Electron geometry and common molecular shapes
with two to six electron groups
Shapes of molecules with more than one central atom
Many molecules, especially organic molecules, have more than one
central atom. Molecular shapes in these molecules vary according to the
electron geometry of each central atom. Figure 3.11 shows the
tetrahedral centres of ethane and ethanol.
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Figure 3.11 Tetrahedral centres of ethane and ethanol
The carbon atoms in Figure 3.11 have no lone pair and therefore have a
tetrahedral molecular shape. The oxygen atom in ethanol has two lone
pairs, giving it a bent molecular shape.
Exercise: How many central atoms are there in acetone? What are the
shapes of the different central atoms?
Acetone
Molecular shape and polarity
As mentioned briefly in Study Guide 2, the presence of polar covalent
bonds does not always result in a polar molecule. To determine the
polarity of a given molecule, first determine if there are any polar
covalent bonds by comparing electronegativities. The following are the
electronegativities from the COBr2 example:
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Atom
EN
C
2.5
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O
3.5
Br
2.8
The bond polarities are determined by observing the differences in
electronegativity for each bond, ΔEN.
Bond
ΔEN
C=O
1
C−Br
0.3
The C=O bond is polar, whereas the C–Br bonds are only slightly polar.
Even though the C–Br bonds are in arrangement that allows them then
partially reduce the polarity of the C=O bond, the molecule will have a
net dipole:
Polarity of different molecules
Carbon dioxide is a non-polar molecule (Figure 3.12). Oxygen (EN 3.5)
when bonded to carbon (EN 2.5) forms a relatively polar dipole (∆EN
1.0). However, the linear shape of the molecules causes the two oxygen
induced dipoles to completely cancel each other.
..
.O.
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C
..
.O.
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Figure 3.12 Opposing dipoles in carbon dioxide
Ammonia is fairly polar molecule, even more so if the highly
electronegative nitrogen attracts a hydrogen ion (Figure 3.13). Nitrogen
(EN 3.0) bonded to hydrogen (EN 2.1) forms a relatively polar dipole
(∆EN 0.9). As a trigonal pyramidal molecule, all the electrons from the
three hydrogen atoms are pulled by the central nitrogen atom. The three
dipoles reinforce each other resulting in a net negative charge on the
nitrogen atom.
H
N
H
H
H
N
H
H
Figure 3.13 Three dipoles in ammonia reinforcing one another
Boron trifluoride is a non-polar molecule (Figure 3.14). Fluorine (EN
4.0) and boron (EN 2.0) forms a very polar bond with ∆EN of 2.0.
However, because of the trigonal planar shape of the molecule, all the
bonds are of the same angle and magnitude. This causes the three
dipoles to completely cancel each other.
F
B
F
F
120°
bond
angle
Figure 3.14 Opposing dipoles in boron trifluoride
Carbonyl sulphide is a relatively polar molecule (Figure 3.15). Sulphur
(EN 2.5) and carbon (EN 2.5) have equal electronegativities and no
dipole is formed between the two atoms. Oxygen and carbon however
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has a relatively large difference in electronegativity (∆EN 1.0). This
causes most of the electrons in the molecule to be attracted to the oxygen
atom, resulting in a net negative charge.
S
C
O
Figure 3.15 Net negative charge on the oxygen atom of carbonyl
sulphide
Exercise: Which is more polar, tetrachloromethane, chloroform or
dichloromethane?
Hint: Most electronegative atom is chlorine.
Tetrachlorometh
ane
Chloroform
Dichloromethan
e
Trivia: Why is monochloromethane never used as a solvent?
Reading Material
 Chapters 10.1 to 10.3
Silberberg, M.S. (2006). Chemistry: The Molecular Nature of
Matter and Change. 4th Ed. McGraw Hill.
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