Molecular Shapes – VSEPR Theory
The more you study chemistry, the more you will
come to appreciate that it’s all about the electron.
You have probably already learned that when two
atoms form a covalent bond, they are in essence
sharing a pair of electrons. Molecules are made
up of atoms bonded together this way, and most
small molecules – in the 3-7 atom range – are
comprised of one central atom with all the other
atoms bonded to it, like the ones shown below:
Molecular Shapes – VSEPR Theory
But what exactly determines the shape of these
molecules? What causes some to be straight,
others to be bent? Some to be flat and others to
be more 3-dimensional? Again, the answer lies in
the electrons. Recall that electrons are negatively
charged and therefore repel one another. Since
the bonds are made up of electrons, these bonds
always repel one another and orient themselves
as far away from each other as possible.
Molecular Shapes – VSEPR Theory
This is known as “VSEPR” which stands for
“Valence Shell Electron Pair Repulsion Theory.”
Essentially it states that molecular shapes are
determined by the repulsive forces acting
between the electron pairs in the central atom’s
outermost (valence) level.
Molecular Shapes – VSEPR Theory
These electron pairs include not only the bonding
pairs that hold the surrounding atoms to the
central atom, but also the nonbonding pairs that
reside in the central atom’s valence level.
Lewis structures (AKA “electron dot structures”)
are diagrams that illustrate exactly how electrons
are being shared – and not shared – within a
molecule. Let’s look at some examples.
Molecular Shapes – VSEPR Theory
Below is a completed Lewis structure for BeF2.
Notice how the central atom (the Be) has just two
electron regions on it.
F
F Be
180°
These electron regions would naturally repel one
another and orient themselves as far away from
each other as they can get. This forces the
molecule into what is called a “linear” shape.
The bond angle is 180°.
Molecular Shapes – VSEPR Theory
Since a linear molecule is flat with all bonded
atoms existing in the same plane as the central
atom, drawing one is fairly easy:
F Be F
Drawing:
F Be F
Let’s look at another example:
Molecular Shapes – VSEPR Theory
Below is a completed Lewis structure for BF3.
Notice how the central atom (the B) has three
electron regions on it.
F
F B F
120°
These electron regions would naturally repel one
another and orient themselves as far away from
each other as they can get. This forces the
molecule into a “trigonal planar” shape.
The bond angles are now 120°.
Molecular Shapes – VSEPR Theory
Trigonal planar molecules are also completely flat,
so drawing one is rather simple.
Drawing:
F
B
F
F
F
B
F
F
Let’s look at another example:
Molecular Shapes – VSEPR Theory
Below is a completed Lewis structure for SO2.
Notice how the central atom (the S) again has
three electron regions on it.
O S O
These three regions repel one another, but with
no atom on top, the shape is simply called “bent”
or “V-shaped.” In fact, the nonbonding pair on
top repels the bonding pairs more than the
bonding pairs repel each other.
Molecular Shapes – VSEPR Theory
Below is a completed Lewis structure for SO2.
Notice how the central atom (the S) again has
three electron regions on it.
S
O O
<120°
These three regions repel one another, but with
no atom on top, the shape is simply called “bent”
or “V-shaped.” In fact, the nonbonding pair on
top repels the bonding pairs more than the
bonding pairs repel each other. This squeezes the
bond angle to something less than 120°.
Molecular Shapes – VSEPR Theory
A bent molecule like this is also flat and easy to
draw:
O
Drawing:
S
O
S
O
O
Let’s look at another example:
Molecular Shapes – VSEPR Theory
Below is a completed Lewis structure for CH4.
Notice how this time the central atom has four
electron regions on it.
H
HC H
H
You might be thinking that this “+ shape” with all
90° bond angles would be the most spread out
arrangement possible, but a “+ shape” doesn’t
take advantage of all the space available. Keep in
mind: no one said that molecules had to be flat.
Molecular Shapes – VSEPR Theory
The repelling force between the four bonding
electron pairs forces the hydrogen atoms into the
following arrangement:
H
HC H
H
109.5°
This shape is known as “tetrahedral,” since it has
the form of a symmetrical four-sided geometric
solid known as a tetrahedron. Its bond angles are
all 109.4712° – or simply 109.5°.
Molecular Shapes – VSEPR Theory
Since a tetrahedral molecule is not flat, drawing
Note this dotted line
one is a little trickier:
bond. Dotted lines
Note how these two
other bonds are drawn
as normal lines since
they are still in the same
plane as the central
atom.
H
C
H H
H
Drawing:
H
Let’s look at
another example:
H
C
H
H
are used to show that
a bond is receding
away from you –
behind the plane of
the central atom.
Note this wedge shaped
bond. Wedges are used
to show that a bond is
coming out toward you –
in front of the plane of
the central atom.
Also, drawing the closer “H”
a little larger and the more
distant “H” a little smaller
can help add to the threedimensional appearance of
the drawing.
Molecular Shapes – VSEPR Theory
Below is a completed Lewis structure for NH3.
Again, notice how the central atom has four
electron regions on it.
HNH
H
And though this Lewis structure makes the
molecule look T-shaped, remember: a Lewis
structure is not meant to convey the shape of a
molecule. These four regions spread themselves
out as they did before…
Molecular Shapes – VSEPR Theory
Below is a completed Lewis structure for NH3.
Again, notice how the central atom has four
electron regions on it.
HNH
H
And though this Lewis structure makes the
molecule look T-shaped, remember: a Lewis
structure is not meant to convey the shape of a
molecule. These four regions spread themselves
out as they did before…
Molecular Shapes – VSEPR Theory
And, like before, the nonbonding electron pair
repels the bonding electron pairs more than the
bonding electron pairs repel each other:
N
H H
<109.5°
H
This pushes the H’s down into a shape known as
“trigonal pyramidal.” This causes the bond angles
to be somewhat smaller than 109.5°.
Molecular Shapes – VSEPR Theory
Trigonal pyramidal molecules are drawn just like
tetrahedral one, but with a nonbonding electron
pair on the top
N
H H
H
Drawing:
H
Let’s look at
another example:
N
H
H
Again notice the
dotted line for the
bond receding away
from you and the
wedge shape for the
bond coming out
toward you.
Molecular Shapes – VSEPR Theory
Below is a completed Lewis structure for H2O.
Once again, notice how the central atom has four
electron regions on it.
H
HO
And these electron pairs again repel one another
as they did before, but with only two of the four
electron regions involved in bonds, the molecule
takes on a “bent” or “V-shaped” arrangement –
similar to the SO2 molecule we saw before.
Molecular Shapes – VSEPR Theory
And once again, the nonbonding electron pairs
repel the bonding pairs more than they repel one
another…
<109.5°
H
H
O
So the two H atoms get squeezed a little closer
together. And this gives this bent molecule a bond
angle that is significantly less than 109.5°.
Molecular Shapes – VSEPR Theory
This bent molecule is like the previous one, except
this one has two nonbonding electron pairs instead
of just one:
H
O
H
H
Drawing:
Let’s look at
another example:
O
H
It helps to put loops around
the nonbonding electron
pair and to have the closer
pair drawn a little larger and
overlapping the pair that is
farther away.
Molecular Shapes – VSEPR Theory
Below is a completed Lewis structure for PF5.
Notice how the central atom now has five electron
regions on it.
F
F P F
F F
90°
120°
These electron regions repel one another and
orient themselves as far away from each other as
possible. This gives rise to a shape known as
“trigonal bipyramidal.” And the bond angles are a
combination of 90° and 120°.
Molecular Shapes – VSEPR Theory
With a trigonal bipyramidal molecule, the top,
bottom and left hand atoms can all be drawn in
the same plane as the central atom:
F
F
PF
F
F
Drawing:
F
Let’s look at
another example:
F
F
P F
F
This shape may be thought
of as a combination of a
linear (going straight up and
down) and a trigonal planar
perpendicular to the linear.
Molecular Shapes – VSEPR Theory
Below is a completed Lewis structure for SF4.
Again, the central atom has a total of five electron
regions on it.
F
F
S
F F
These electron regions repel one another and
orient themselves as far away from each other as
they can get. This arrangement is commonly
known as “see-saw shaped.”
Molecular Shapes – VSEPR Theory
And, once again, the nonbonding electron pair on
the left repels the bonding pairs more than they
repel each other.
F
<90°
F
SF
F
<120°
This pushes the four F atoms to the right a bit.
This makes both the 90° and 120° bond angles a
little smaller.
Molecular Shapes – VSEPR Theory
See-saw molecules are drawn just like trigonal bipyarmidal molecules, but with a nonbonding
electron pair on the left:
F
F
SF
Drawing:
Let’s look at
another example:
F
F
F
S F
F
Molecular Shapes – VSEPR Theory
Below is a completed Lewis structure for BrF3.
Once again, the central atom has five electron
I pitty the fool who
regions on it.
can’t remember
F
F Br
F
T-shaped!
These five electron regions repel one another as
they did before. And this time the resulting shape
is completely flat and is known as “T-shaped.”
Molecular Shapes – VSEPR Theory
Again, the nonbonding electron pairs on the right
repel the bonding pairs more than the bonding
pairs repel each other.
<90°
F
F Br
F
That pushes the top and bottom F to the left.
And it squeezes the bond angles to a little less
than 90°.
Molecular Shapes – VSEPR Theory
A T-shaped molecule is completely flat and
therefore needs no dotted line or wedge-shaped
bonds. It is drawn like the trigonal bipyramidal but
with two nonbonding
pairs on the right:
F
Drawing:
Let’s look at
another example:
F Br
F
F
F Br
F
Molecular Shapes – VSEPR Theory
Below is a completed Lewis structure for KrF2.
Once again, the central atom has five electron
regions on it.
F
Kr
F
These five electron regions repel one another as
they did before. And because all three
nonbonding electron pairs arrange themselves
around the perimeter, the resulting molecule is
actually “linear.”
Molecular Shapes – VSEPR Theory
As before, the nonbonding electron pairs repel the
bonding pairs more than the bonding pairs repel
one another.
F
Kr
F
180°
But since they are pushing equally from all sides,
these extra repulsive forces have no effect on the
bond angle. It remains a precise 180°.
Molecular Shapes – VSEPR Theory
This linear molecule is drawn the same way as the
last one, only this one has three evenly spaced
nonbonding electron pairs wrapped around the
central atom:
F
Kr
Drawing:
F
F
Let’s look at
another example:
F
Kr
Molecular Shapes – VSEPR Theory
Below is a completed Lewis structure for SF6.
Notice how the central atom now has six electron
regions on it.
Notice how the
F
F
F
3
F
SF
F
These electron regions would naturally repel one
90°
4 angles
2 are
bond
1
all now7precisely
90°.8 5 6
another and orient themselves as far away from
each other as they can get. This pushes the atoms
into a shape known as “octahedral.” An octahedron
is a geometric solid with eight identical faces.
Molecular Shapes – VSEPR Theory
An octahedral molecule is best drawn with the top
and bottom atoms drawn in the same plane as the
central atom, then two atoms coming out toward
us and two others
receding away:
F
F
F
S
F FF
Drawing:
F
F
F
S
Let’s look at
F
F
another example:
F
Molecular Shapes – VSEPR Theory
Below is a completed Lewis structure for BrF5.
Once again, the central atom has six electron
regions on it.
F
F F
Br
F F
These regions repel one another as they did for
octahedral, but with one of the atoms replaced by
a non-bonding pair, the resulting shape is called
“square pyramidal” since it resembled a square
based pyramid”– like the ones in Egypt.
Molecular Shapes – VSEPR Theory
Once again, the nonbonding electron pair on
bottom repel the five bonding pairs more than
they repel one another.
F
F
Br F
<90°
F
F
This pushes the four nearby F atoms up a bit.
This makes all the bond angles a little less than 90°.
Molecular Shapes – VSEPR Theory
A square pyramidal molecule is drawn just like the
octahedral one, but with a nonbonding pair of
electrons in place of the bottom atom.
F
F
Drawing:
Let’s look at
one final example:
F
F
F
F
Br F
F
Br
F
F
Molecular Shapes – VSEPR Theory
Below is a completed Lewis structure for KrF4.
Once again, the central atom has six electron
regions on it.
F F
Kr
F F
These regions repel one another as they did for
octahedral, but with the top and bottom atoms
replaced by non-bonding pairs, the resulting
shape is called “square planar.”
Molecular Shapes – VSEPR Theory
Once again, the nonbonding electron pairs on top
and bottom repel the four bonding pairs more
than they repel one another.
90°
F
F
Kr
F
F
But since they are pushing equally from both
sides, these extra repulsive forces have no effect
on the bond angle. So it remains a precise 90°.
Molecular Shapes – VSEPR Theory
A square planar molecule is drawn just like the
octahedral one, but with nonbonding pairs of
electrons in place of the top and bottom atoms.
F
F
Drawing:
F
F
So now let’s try
some sample problems.
Kr
F
Kr
F
F
F
The next 15 slides will give you practice
in drawing these molecular shapes.
For each slide, draw and name the shape
and give the bond angle on a sheet of
scrap paper – or, if you prefer, directly on
the slide using the inking tool.
The answer will then be given.
If your answer does not precisely match
the one given, go back and try it again.
Remember: the more you practice, the
better you will get.
1. phosphorus trichloride:
Cl P Cl
Cl
Cl
P
Cl
Cl
trigonal pyramidal, <109.5°
Explanation: Since the
central atom has four
electron regions on it, think
tetrahedral. But one of
these regions is a NEP
(nonbonding electron pair),
so the shape is trigonal
pyramidal.
2. carbon disulfide:
S C S
S
C
S
linear, 180°
Explanation: Since the central
atom has two electron regions
on it, neither of which is a NEP,
the shape is linear. (It doesn’t
matter that these electron
regions happen to be double
bonds: a double bond counts
as just one region.)
3. sulfur trioxide:
O
O S O
O
O
S
O
trigonal planar, 120°
Explanation: Since the
central atom has three
electron regions on it, none
of which are NEP’s, the
shape is trigonal planar.
4. selenium tetrabromide:
Br
Se
Br
Br
Br
Br
Br
Se Br
Br
see-saw, <90°, <120°
Explanation: Since the central
atom has five electron regions
on it, think trigonal
bipyramidal. But one of these
regions is a NEP, so the
shape is see-saw.
5. sulfur difluoride:
F S F
F
S
F
Bent, <109.5°
Explanation: Since the
central atom has four
electron regions, think
tetrahedral. But two of
these regions are NEP’s so
the shape is bent.
6. xenon difluoride:
F Xe F
F
Xe
F
linear, 180°
Explanation: Since the
central atom has five
electron regions on it, think
trigonal bipyramidal. But
three of these regions are
NEP’s, so the shape is
linear.
7. boron trichloride:
Cl B Cl
Cl
Cl
B
Cl
Cl
trigonal planar, 120°
Explanation: Since the
central atom has three
electron regions on it, none
of which are NEP’s, the
shape is trigonal planar.
7. ammonium ion:
H
H 1+
H N H
H
H 1+
N H
H
tetrahedral, 109.5°
Explanation: Since the
central atom has four
electron regions on it, none
of which are NEP’s, the
shape is tetrahedral. The
fact that it is an ion rather
than a neutral molecule
does not have any impact
on its shape.
8. radon tetrafluoride:
F
F Rn F
F
F
F
Rn
F
F
square planar, 90°
Explanation: Since the
central atom has six
electron regions on it, think
octahedral. But two of
these regions are NEP’s, so
the shape is square planar.
9. arsenic pentaiodide:
I
I
I
I
I
As I
I
I
As I
I
trigonal bipyramidal, 90°, 120°
Explanation: Since the
central atom has five
electron regions on it, none
of which are NEP’s, the
shape is trigonal
bipyramidal.
10. tellurium hexachloride:
Cl
Cl
Cl
Te
Cl
Cl
Cl
Cl
Cl
Cl
Te
Cl
Cl
Cl
octahedral, 90°
Explanation: Since the
central atom has six
electron regions on it, none
of which are NEP’s, the
shape is octahedral.
11. selenium dioxide:
O Se O
Se
O
O
bent/V-shaped, <120°
Explanation: Since the
central atom has three
electron regions on it, think
trigonal planar. But one of
these regions is a NEP, so
the shape is bent/V-shaped.
12. iodine trichloride:
Cl I
Cl
Cl
Cl
Cl I
Cl
T-shaped, <90°
You really ought
to remember
T-shaped!
Explanation: Since the
central atom has five
electron regions on it,
think trigonal bipyramidal.
But two of these regions
are NEP’s, so the
molecule is T-shaped.
13. silicon tetrahydride (silane):
H
H
H Si H
H
H
Si H
H
tetrahedral, 109.5°
Explanation: Since the
central atom has four
electron regions on it, none
of which are NEP’s, the
shape is tetrahedral.
14. arsenic trifluoride:
F As F
F
F
As
F
F
trigonal pyramidal, <109.5°
Explanation: Since the
central atom has four
electron regions on it, think
tetrahedral. But one of
these regions is a NEP, so
the shape is trigonal
pyramidal.
15. bromine pentachloride:
Cl
Cl
Br
Cl
Cl
Cl
Cl
Cl
Cl
Br
Cl
Cl
square pyramidal, <90°
Explanation: Since the
central atom has six
electron regions on it, think
octahedral. But one of
these regions is a NEP, so
the shape is square
pyramidal.