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Disproportionation is a redox reaction in which a single substance is simultaneously oxidized and reduced to form two products that differ in oxidation state but contain the same element. This type of reaction occurs when the element can adopt multiple oxidation states and the intermediate state is stabilized under the reaction conditions. Disproportionation is the reverse process of comproportionation, where two species with different oxidation states react to form a single intermediate state.

Common examples illustrate the concept. Hydrogen peroxide undergoes disproportionation to yield water and oxygen: 2 H2O2 -> 2 H2O + O2. In this reaction, oxygen in H2O2 has an oxidation state of -1, which becomes -2 in water and 0 in molecular oxygen. Another well-known example involves hypochlorite in aqueous solution: 3 ClO- -> ClO3- + 2 Cl-. Here chlorine is +1 in ClO-, +5 in ClO3-, and -1 in Cl-, showing simultaneous oxidation and reduction of chlorine.

Factors such as pH, concentration, and catalysts influence disproportionation. Catalysis is common in biological and industrial contexts; for instance, catalase accelerates the decomposition of hydrogen peroxide, a form of disproportionation of H2O2. In environmental and chemical processes, disproportionation can affect bleaching, disinfection, and redox balance.

Disproportionation is often contrasted with comproportionation, in which two species with different oxidation states react to produce an intermediate state. While both involve redox chemistry of the same element, they proceed in opposite directions and under different conditions.