What is disproportionation in redox chemistry?

Disproportionation in redox chemistry is a reaction where an element in a single species is both oxidised and reduced.

In more detail, disproportionation is a specific type of redox reaction (reduction-oxidation reaction) that involves one species undergoing both oxidation and reduction. This means that the same element in a single reactant is simultaneously oxidised and reduced, leading to two different products.

For instance, consider the reaction of chlorine with water. In this reaction, chlorine (in the zero oxidation state) is both reduced to chloride ions (in the -1 oxidation state) and oxidised to chlorate ions (in the +5 oxidation state). This is a classic example of a disproportionation reaction.

2Cl2 + 2H2O → 4HCl + O2

In the above reaction, chlorine is both oxidised and reduced. It is oxidised to form oxygen gas (O2) and reduced to form hydrochloric acid (HCl).

It's important to note that disproportionation reactions only occur with elements that have at least three different oxidation states. The element in the reactant is in an intermediate oxidation state, and in the products, it is in a higher and a lower oxidation state.

Understanding disproportionation reactions is crucial in redox chemistry as they are common in many chemical and biological processes. For example, they play a key role in the metabolism of many organisms, in the operation of batteries, and in the industrial production of many chemicals.