Disproportionering

Disproportionering, also known as disproportionation, is a redox process in which a single substance in an intermediate oxidation state is simultaneously oxidized to a higher oxidation state and reduced to a lower oxidation state. The reaction yields two species of the same element in different oxidation states.

In general, disproportionation occurs when the resulting products are thermodynamically more stable than the intermediate state under the reaction conditions. The process can be influenced by factors such as pH, solvent, temperature, and the presence of catalysts or complexing ligands. Conceptually, the element acts as both oxidant and reductant in the same reaction mixture.

Common examples include:

- Hydrogen peroxide in aqueous solution: 2 H2O2 → 2 H2O + O2. Here oxygen is reduced from an average oxidation state of −1 in H2O2 to −2 in H2O and oxidized to 0 in O2.

- Copper(I) ions in solution: 2 Cu+ → Cu2+ + Cu(s). Copper(I) is simultaneously oxidized to copper(II) and reduced to metallic copper.

- Chlorite in neutral or basic solution: 2 ClO− → Cl− + ClO3−, where chlorine changes from +3 in ClO− to −1 in Cl− and +5 in ClO3−.

Disproportionation is the reverse process of comproportionation (synproportionation). It plays a role in redox chemistry, environmental chemistry, and certain synthetic contexts where the stabilization of multiple oxidation states is important.

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