IB Docs (2) Team VSEPR
The theory
For some reason many students seem to find it very difficult to predict the shape and bond angles of a simple molecule or ion using Valence Shell Electron Pair Repulsion Theory (VSEPR). At first sight this seems strange because the theory really is very simple. Essentially there are only three basic ideas:
- Electrons arrange themselves in pairs around the central atom or ion so that there is maximum repulsion between the pairs. Strictly speaking it is electron domains rather than electron pairs as this would take account of an unpaired electron in a free radical but this is not necessary at this level. In fact in previous syllabi the term 'negative charge centres' was used but in sub-topic 4.3 : Covalent structures it clearly states that 'electron domain' should be used in place of 'negative charge centre'.
- Non-bonding pairs of electrons exert a greater repulsion than bonded pairs.
- A multiple bond (i.e. a double or triple bond) must count as one pair of electrons (one electron domain) as all the pairs of electrons must be aligned in the same direction.
Applying the first idea leads to the fact that there are only five basic shapes to be considered. In fact there are only three at Standard Level as sub-topic 4.3 limits the shape to two, three and four electron domains on the central atom. For Higher Level students sub-topic 14.1 extends this to five and six pairs. These basic shapes and bond angles are:
| Number of pairs of electrons | Shape | Bond angle(s) |
| 2 | Linear | 180o |
| 3 | Trigonal planar | 120o |
| 4 | Tetrahedral | 109.5o (actually 109o 28') |
| 5 | Trigonal bipyramidal | 90o, 120o, 180o |
| 6 | Octahedral | 90o, 180o |
If one or more pairs are non-bonded then the basic shape adapts to take account of this. An example is ammonia.
There are four pairs of electrons around the central nitrogen atom. Three of these are bonding and one is non-bonding. Four pairs gives a tetrahedral shape but as one of the pairs is a non-bonding pair the actual shape will be trigonal pyramid (with the non-bonded pair pointing to the apex of what would have been the tetrahedron). As the non-bonding pair exerts a greater repulsion than bonding pairs this forces the bonded pairs slightly closer together and the H-N-H bond angle is reduced from the 109.5o of a regular tetrahedron to approximately 107o. Other examples follow a similar logic and basically that is all there is to it - at least as far as the IB goes. In fact it is not nearly so simple when you try to extend it to other molecules and ions. You can read more about this in 'Something to think about' on the page covering shapes of molecules or ions in 4.3 Covalent structures (2).
The problem for students
So why do many students seem to find it so difficult? I think there are two fundamental reasons. The first is that they are not clear about how to determine the number of electron pairs around the central atom. To do this they need to understand about Lewis structures but then ignore all the electron pairs around all the outer atoms. At Standard Level they can use the octet rule to give them the number and almost all examples will be four pairs due to the eight electrons around the central atom. However, because some of the pairs might be a double or triple bond this will reduce the number of pairs according to VSEPR theory. So that, for example, a carbon dioxide molecule which has the eight electrons around the central carbon atom arranged in two double bonds will have two pairs according to VSEPR and is linear.
The second reason is that many students seem unable to visualize the shapes in three dimensions. There is an interesting article about this written by Ashley Jennings, a high school student from Pennsylvania, in Journal of Chemical Education, Vol. 87 No. 5 May 2010 pages 462-463. In this article Ashley talks about how difficult it was for students to equate two-dimensional drawings with the three-dimensional models the teacher had shown them. Ashley's solution to this was to develop her own video showing the models in three dimensions.
This is a well-constructed video and you may find it helpful. I think though that the real problem is that often students are only shown the shapes and even in this video the shapes are shown on a two-dimensional screen. I have found that you will gain a much better understanding and concept of the shapes if you make your own models. Ideally you should have access to a molecular modelling kit, e.g. molymod and you should ask your teacher to buy the kits if your school does not already have them. If you always have these kits at hand then whenever a shape is discussed you can make the shape for yourself and feel it in three-dimensions.
Finally, if you are still having problems, it might be worth just learning the shapes of some specific molecules and ions as the examples the IB can use are quite limited in number!
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