Predicting the Shapes of Molecules

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A theory, called the valence shell electron pair repulsion theory, or more simply, the VSEPR theory, can predict the shape that a molecule can take. It is based on the idea that valence shell electron pairs want to stay as far apart from each other as possible (because they are negatively charged) so that their repulsions are at a mininium.
For example, take BeCl2 molecule. The Be is the central atom and the 2 Cl atoms extend out ward from it. But is the shape of the molecule linear or is it bent?
According to VSEPR, the electron pairs in the valence shell want to stay as far apart as possible. There are 4 valence electrons in Be, and therefore there are 2 pairs. The farthest apart that 2 electron pairs can stay is a 180° angle. In other words, BeCl2 is linear.
Now let's take BCl3. The B is the central atom, and it has 3 pairs of valence electrons. The 3 pairs try to stay as far apart as possible and form a planar triangular shape with 120° angles.
Here is a table showing the orientations expected for different numbers of valence electrons:

Electron Pairs

Shape

Example

2


linear

BeCl2

3


planar triangular

BCl3

4


tetrahedral

CH4

5


trigonal bipyramidal

PCl5

6


octahedral

SF6

When Some Electron Pairs Are Not in Bonds

Some molecules have a central atom with one or more pairs of electrons that are not shared with another atom. For example, SnCl2 has Sn as the center atom. Sn has 3 pairs but only 2 of those pairs are shared with the Cl. The left over pair is called the lone pair.
When this happens, you still consider that there are 3 pairs, therefore making the shape a planar triangular. But since one of the pairs are lone, you take one vertex off of the planar triangular. Now the remainder is a shape called bent, nonlinear, or V-shaped.

3 valence pairs

Planar Triangular--0 lone pairs
Bent--1 lone pair
In a molecule with four electron pairs, the shape is tetrahedral. But if there is one lone pair, one vertex is removed from the shape, making a trigonal pyramidal structure. If there are two lone pairs, another vertex is removed, making a bent shaped.

4 valence pairs

Tetrahedral--0 lone pairs
Trigonal Pyramidal--1 lone pair
Bent--2 lone pairs
Here are tables for the structures of 5 and 6 valence pairs:

5 valence pairs

Trigonal Bipyramidal--0 lone pairs
Unsymmetrical Tetrahedron--1 lone pair
T-shaped--2 lone pairs
Linear--3 lone pairs
For five valence pairs, the axial bonds are not removed because they have a larger bond angle from each other, 180°. The equatorial bonds are removed because they are only 120° apart.

6 valence pairs

Octahedral--0 lone pairs
Square Pyramidal--1 lone pair
Square Planar--2 lone pairs

Shapes of Molecules with Double or Triple Bonds

The presence of double or triple bonds does not complicate matters at all. Since in a double or triple bond, both electron pairs must stay together, we can treat them as one single bond. For example, in carbon dioxide (CO2) the oxygen atoms are double bonded to the central carbon atom. The carbon atom has no lone pairs. The two double bonds act as two single bonds, so therefore, there are two electron pairs. Thus, the molecule is linear.
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