SHAPES OF
MOLECULES
REMINDER ABOUT ELECTRONS
 Electrons
have negative charges
 Negative charges “repel” each
other
 In molecules, electrons want to get
as far away from each other as
possible
 As a result, this repulsion of electrons
leads to the shape of the molecule
SHAPE OF MOLECULES
 There
is a simple model used to determine
the shape of molecules.
 VSEPR (Valence Shell Electron Pair
Repulsion)

A simple model that predicts the general
shape of a molecule based on the
repulsion between both the bonding and
nonbonding electron clouds
 Always
based on the CENTRAL atom
DEFINITION
Electron
cloud: Any type of
bond (single, double or triple)
or any set of unshared pairs
of electrons.
Unshared pair of electron:
any pair of electrons not
involved in a covalent bond.
2 ELECTRON CLOUDS
 Example:
CO2
 Count the number of electron clouds
surrounding the central atom
..
..
O=C=O
··
··
 There are 2 double bonds around the
central Carbon (C)

Thus, there are 2 electron clouds
2 ELECTRON CLOUDS
 Electrons
have a negative charge
and repel each other
 Thus, the electrons that maintain
the bond will move as far away
from each other as possible
 If the central atom has only two
electron clouds, it will have a linear
shape
LINEAR
View
the clip
NOTE:
That the attached
atoms are 180° apart from
each other
3 ELECTRON CLOUDS
 Again
count the number of electron
clouds around the central atom (SO3)
..
: O:
..
| ..
:0-S=O
··
··
3 ELECTRON CLOUDS
Again
the electron clouds want
to move as far from each other
as possible
When the central atom has 3
electron clouds surrounding it,
the molecule has a trigonal
planar shape
TRIGONAL PLANAR
View
the clip
NOTE:
The attached atoms are
120° apart from each other
3 ELECTRON CLOUDS
What
happens if one of the
electron clouds is an
unshared pair of electrons
O 3
Do the Lewis dot structure
for this in your notes
3 ELECTRON CLOUDS
 Notice
the following:
 You have one double bond
 You have one single bond
 You have one unshared pair of
electrons
 When you view a molecule, you can’t see
the unshared pair of electrons
 This creates a bent shape
BENT
 View
the clip
 NOTE:
That the attached atoms and
unshared pair are 120° apart from each
other

Since you can’t see the unshared pair, the
molecule looks bent
4 ELECTRON CLOUDS
Again
count the number
of electron clouds
CCl4
Draw the Lewis dot
structure in your notes
4 ELECTRON CLOUDS
Same
as before, the electron
clouds want to get as far from
each other as possible
When the central atom has 4
electron clouds surrounding it,
you get a tetrahedral
TETRAHEDRAL
View
the clip
NOTE:
That the attached
atoms are 109.5° apart from
each other
4 ELECTRON CLOUDS
What
happens if you have 3
atoms bound to a central atom
with one unshared pair
NH3
Do the Lewis dot structure for
this in your notes
4 ELECTRON CLOUDS
 Notice
the following:
 You have three single bonds
 You have one unshared pair of
electrons
 When you view a molecule, you
can’t see the unshared pair of
electrons
 This creates a pyramidal shape
PYRAMIDAL
 View
the clip
 NOTE:
That the attached atoms
and unshared pair are 109.5° apart
from each other
 Since you can’t see the unshared
pair, the molecule looks like a
pyramid
4 ELECTRON CLOUDS
What
happens if you have 2
atoms bound to a central atom
with two unshared pairs
H2O
Do the Lewis dot structure for
this in your notes
4 ELECTRON CLOUDS
 Notice
the following:
 You have two single bonds
 You have two unshared pairs of
electrons
 When you view a molecule, you
can’t see the unshared pair of
electrons
 This creates a bent shape
BENT
 View
the clip
 NOTE:
That the attached atoms
and unshared pair are 109.5° apart
from each other
 Since you can’t see the unshared
pair, the molecule looks like a
pyramid
TRY THESE


HANDOUT
For each of the following, write the Lewis
structure and indicate the shape:
1. CBr4
2. CS2
POLARITY OF MOLECULES
 We’ve
already discussed the difference
between nonpolar, polar and ionic bonds
(electronegativity difference)
 Molecular shape is important for
determining the polarity of a molecule.
 Covalently bonded molecules can be
polar or nonpolar based on the shape of
the molecule
EXAMPLE
 Let’s


H2 O
CF4
 First


look at the shapes of
of all:
H-O bond of water has an electronegativity
difference of 1.4 (polar covalent)
C-F bond of CF4 has an electronegativity
difference of 1.5 (polar covalent)
EXAMPLE
 Since
both H2O and CF4 have polar
covalent bonds, we would expect both
molecules to be polar covalent
 This is not the case
 IN YOUR NOTES: Draw the molecular
shape of both:


H2 O
CF4
ANSWER – THINK OF TUG-O-WAR
 Water


As a result water is a polar molecule
You have a partial charge of σ- for O and
σ+ for H
 CF4


is a bent molecule
is a tetrahedral molecule
Because of the shape you have a nonpolar
molecule
Even though you have partial charges, the
charges cancel out because of the shape
WRITE WHICH SHAPES ARE
POLAR/NONPOLAR?
Take
a look at your handout
that shows the shape of
different molecules.
Which shapes do you think are
polar?
Which shapes do you think are
nonpolar?
ANSWER
 Polar
shapes (have lone pairs):
 Bent
 Pyramidal
 Nonpolar shapes (do not have lone
pairs):
 Linear
 Trigonal planar
 Tetrahedral
TRY THE FOLLOWING

Determine the Lewis structure of the following.
Are the molecules polar or nonpolar?
1.
2.
3.
Cl2O
CO2
NF3
ANSWER
Bent, polar molecule
2. Linear, nonpolar
molecule
3. Pyramidal, polar
molecule
1.