Disproportionation Reaction
Definition and meaning of Disproportionation Reaction in chemistry.
A disproportionation reaction is a redox reaction in which a single chemical species is simultaneously oxidized and reduced, generating two different products that contain the same element in a higher and a lower oxidation state.
In more detail
This can only happen when the starting species holds an intermediate oxidation state, since some of its atoms must lose electrons while others gain electrons in the very same reaction. It differs from a typical redox reaction, where two distinct elements exchange electrons between separate species. Disproportionation is especially common among elements with several accessible oxidation states, such as the halogens, transition metals, and peroxide oxygen, and it underlies several important industrial and biological processes.
Key facts
| Field | General Chemistry |
|---|---|
| Also called | Dismutation |
| Requirement | Reactant must have an intermediate oxidation state |
| Example equation | Cl2 + 2NaOH → NaCl + NaOCl + H2O |
When chlorine gas is bubbled into cold, dilute sodium hydroxide, Cl2 (oxidation state 0) disproportionates into chloride (Cl⁻, -1) and hypochlorite (ClO⁻, +1): Cl2 + 2NaOH → NaCl + NaOCl + H2O. This reaction is the industrial basis for making household bleach.
Frequently asked questions
Is disproportionation the same as comproportionation?
No, comproportionation is the reverse process: two species containing the same element in different oxidation states react to form a single product with an intermediate oxidation state.
Why can't every redox reaction disproportionate?
Disproportionation requires the reactant's oxidation state to lie between two states it can adopt; an element already at its highest or lowest possible oxidation state cannot disproportionate.