Working out the shapes of molecules is a fascinating topic in chemistry, stereochemistry is the term which is often used when discussing the shapes of
molecules. Molecules
come in all sorts of shapes and sizes; from small molecules such as
water to very large biological
polymers
such as proteins and DNA. The shape of molecules is perhaps most important in the chemical reactions that
take place in living organisms where a small alteration in the shape of an enzyme or a drug molecule can have dramatic
effects on its action or its side effects.
The overall shape of a molecule is determined by its bond angles; these are the angles between the lines representing the bonds between the nuclei of the atoms that make up the molecule. The molecule shown opposite has a central black atom surrounded by 4 other atoms. The shape of this molecule is tetrahedral. All the bond angles are all 109.50. To draw 3d molecules such as this tetrahedral one we use dotted or dashed lines to represent bonds which are behind the plane and a dark shaded triangular line to represent bonds sticking out in front of the plane. Solid lines show bonds which are in the plane. You can see in this tetrahedral molecule that 2 of the bonds are in the plane; one is behind and one is in front. This is shown in the image opposite.
In the example shown opposite we have a molecule with the formula AB3 where there are 3
covalent bonds
around the central blue atom. Each of these covalent bonds contains a pair of
electrons and since they
are all negatively charged they will repel each other and try to get as far apart in 3d space as possible.
So how can three electron pairs arrange themselves in space to get as far apart as possible? The answer
is to form a flat (planar) molecule with bond angles of 1200. The shape of this
molecule is
described as trigonal planar.
The shapes of molecules with single bonds between the atoms can be easily predicted using the VSEPR model. The shape of a molecule is basically determined by the number of bonding pairs of electrons and lone pairs of electrons around the central atom. Now being negatively charged these bonding pairs and lone pairs of electrons will obviously repel each other as much as possible in 3d space. Working out the shape of a particular molecule then is simply working out how these bonding pairs and lone pairs of electrons orientate themselves to reduce this repulsion between them. The image below shows three molecules you are probably very familiar with; ammonia, water and methane.
For molecules that contain more than one lone pair then this will have a further effect on the bond angles present, we can summarise the repulsion between bonding pairs and lone pairs of electrons as:
| Methane lone pair ↔ lone pair | is greater than | lone pair ↔ bonding pair | is greater than | bonding pair ↔ bonding pair |
However before we start to work out the shapes of various molecules it would help if we looked at the common shapes which are found in many molecules. You should make yourself familiar with these basic shapes and bond angles as they are the basis for all of the molecular shapes you are likely to be asked to work out. The basic shapes of common molecules are based on those shown in the table below:
The basic shape of a molecule will be determined largely by the number of electron pairs or covalent bonds the central atom in a molecule makes. These electrons pairs will repel each other and try to get as far apart as possible in 3d space. The table below shows the basic shapes of molecules with up to four electron pairs around the central atom in a molecule and also the bond angles between these electron pairs.
| Number of electron pairs | Shape of molecule | Name of molecular shape |
|---|---|---|
| 2 |  |
linear |
| 3 |  |
trigonal planar |
| 4 |  |
tetrahedral |
The tetrahedral molecule has 4 bonding pairs or 8 electrons around the central atom. In gcse chemistry we learned that 8 electrons around a central atom was the maximum number allowed to give a stable octet of electrons. However elements in period 3 and above can have an expanded valence shell which allows them to five or six pairs of electrons around the central atom. The shapes and the bond angles in these molecules are shown in the table below.
| Number of electron pairs | Shape of molecule | Name of molecular shape |
|---|---|---|
| 5 |  |
trigonal bipyramidal |
| 6 |  |
Octahedral |
In small molecules with up four electron pairs and also larger octahedral molecules all the bond angles in these molecules are the same. If you built models of these molecules and then flipped then on your desk no matter which way the molecule landed it would always look the same, simply because the bond angles are the same in all these molecules. However this is not true of trigonal bipyramidal molecules. In these molecules there are two different positions; called the axial and equatorial positions. These two different positions are shown in the image opposite:
In a molecule with a trigonal bipyramidal shape there are two different environments in which the bonded atoms can be in when they
surround the central atom. These two positions are called the equatorial and the axial positions. There are 3 atoms around the middle part of the trigonal bypyramidal molecule shown. They are separated from each other by bond angles of 1200. These atoms are described as being in an equatorial position.
There are also two atoms shown in green which are both at 900 to the equatorial atoms. These atoms are said to be in an axial position.
