How has your IB year been so far? Whichever way your teacher(s) has chosen to deliver the deliver the course, chances are that round about now you will be covering the bonding part of the course (section 4).
Bonds are formed from electrons and section 4.3 deals with covalent bonds. In covalent bonds, pairs of electrons are shared between atoms. As a rule of thumb, a covalent bond is formed between two non-metals sharing electrons.
We can represent the bonds through diagrams called Lewis structures (or Lewis diagrams) – and I am sure you will have come across these in any pre IB course that you may have followed. Here in the UK we also refer to them as ‘dot and cross’ diagrams.
In Chemistry, we also conveniently move from one model to another in order to help to explain concepts. You may have come across the s, p, d notation to represent electrons in atoms to help explain the trends in ionisation energies of elements in a period while the Lewis structure model relies on a simpler understanding of electrons called the octet rule.
The octet rule involves the idea that when atoms form covalent bonds in molecules, the electrons will be shared so that each atom has a ‘full outer shell’ – in most elements this will be due to having eight electrons in the outer shell (hence, the phrase ‘octet’) although if there is hydrogen, only two electrons will need to be shared to form the ‘full outer shell’ – as you can see, the octet rule is breaking down (not working!) already – we could think of it as an anomaly.
That said, Chemistry is also full of anomalies (!) and two more stable molecules do not fit into the octet rule, BF3 and BeCl2.
The central atom in BF3 (ie, B) has only six electrons in its outer shell and in BeCl2 the central atom has only four electrons in its outer shell – the octet is incomplete. This may seem unusual but can be explained form the fact that both atoms are small, and due to electron repulsion they can only physical fit six in the case of B and four in the case of Be electrons around the central ion*
We use resonance structures help explain Lewis structures involving double bonds –in these structures it is possible to write more than one correct Lewis structure.
The classic example of this is Ozone (O3) and its resonance structure is shown below:
If there is a problem with the use of this image, please let me know and I will remove it.
The model indicates that the double bond is oscillating between the two states. This is a good way of showing things but in reality we believe something slightly different is happening. If you were to measure the O-O single and O=O double bond lengths you would find that the observed bond lengths in ozone are somewhere in between these expected length, indicating that a ‘true’ bond is not present.
We now tend to think that the bonding electrons are delocalised and oscillating very quickly between the resonance structures. Delocalised electrons are generally found when we have a series of alternating double and single bonds (for example, in benzene).
Can you find any other examples of resonance structures or delocalised electrons? There re lots out there? Please feel free to upload some names / formulas below!
*Not strictly true but the explanation fits the model being used (!)