Q.In a special type of redox reaction an element in one oxidation state is simultaneously oxidised and reduced.

Disproportionation Reactions: The Self-Oxidation-Reduction

Imagine you have a group of friends who are all equally tall. Now imagine that, for no external reason, some of them suddenly grow taller while others shrink shorter — all starting from the same height. That sounds strange, right? But that's exactly what happens in a disproportionation reaction: the same element, in the same starting oxidation state, simultaneously gets oxidised (loses electrons) and reduced (gains electrons).

In other words, one part of the substance acts as the oxidising agent, and another part acts as the reducing agent — on itself.

The Intuition First

Think of a chemical element like chlorine. In its elemental form ($Cl_2$), each chlorine atom has an oxidation state of 0. Now, if you put chlorine gas into water, something odd happens:

• Some chlorine atoms gain an electron (reduction) and become $Cl^-$ (oxidation state -1).
• Other chlorine atoms lose an electron (oxidation) and become $ClO^-$ (oxidation state +1).

The same starting material ($Cl_2$) produces two different products — one more reduced, one more oxidised. That's disproportionation.

The Precise Definition

The key conditions are:

1. One reactant — the same species is both oxidised and reduced.
2. Intermediate oxidation state — the starting element must be able to go both higher and lower.
3. Two products — one more oxidised, one more reduced.

The Classic Example: Chlorine in Water

Let's track the oxidation states:

The chlorine in $Cl_2$ (oxidation state 0) goes to -1 (gain of electron = reduction) and to +1 (loss of electron = oxidation). The same element, same starting state, two different directions.

How to Spot a Disproportionation Reaction

Look for these three clues:

1. One reactant, two products — especially if the reactant contains an element that can exist in multiple oxidation states.
2. The same element appears in two different oxidation states in the products — one higher, one lower than the starting state.
3. No external oxidising or reducing agent — the substance does it to itself.

More Examples

Hydrogen peroxide decomposition:

Oxygen in $H_2O_2$ has oxidation state -1. In $H_2O$, oxygen is -2 (reduced). In $O_2$, oxygen is 0 (oxidised). The -1 state is intermediate between -2 and 0.

Copper(I) chloride in aqueous solution:

Copper in $CuCl$ is +1. It goes to 0 (reduced) and +2 (oxidised).

Ammonium nitrite decomposition:

Nitrogen in $NH_4^+$ is -3, in $NO_2^-$ is +3. Both go to 0 in $N_2$ — this is actually the reverse (comproportionation), but it shows how the same element can meet in the middle.

The Opposite: Comproportionation …