Comproportionation and disproportionation reactions

What are comproportionation and disprortionation reactions?

Disproportionation reaction:

A disproportionation reaction is a type of redox reaction in which a single chemical species undergoes simultaneous oxidation and reduction, forming two different products with distinct oxidation states.

In these reactions, the same element in the reactant is both oxidized and reduced.

General Form:

2A^p+⟶Aⁿ⁺+A^m+

Where n>p>m, meaning one part of the reactant is oxidized, and another part is reduced.

Examples:

1. Decomposition of Hydrogen Peroxide:

2H₂O₂⟶2H₂O+O₂

Oxygen in H₂O₂ ​ has an oxidation state of −1. It is reduced to H₂O (oxidation state −2) and oxidized to O₂​ (oxidation state 0).

2. Chlorine in Water:

Cl₂+H₂O⟶HCl+HClO

Chlorine (Cl₂​, oxidation state 0) is reduced to HCl (−1) and oxidized to HClO (+1).

3. Reaction of Sodium Hypochlorite:

3Cl₂+6OH−⟶5Cl−+ClO₃−+3H₂O

Chlorine (0) is reduced to chloride (−1) and oxidized to chlorate (+5).

Characteristics:

Oxidation and Reduction in One Element: The same element in the reactant undergoes both oxidation and reduction.

Specific to Certain Compounds: Elements or compounds with intermediate oxidation states, such as Cl₂, H₂O₂ and NO₂​, are prone to disproportionation.

Thermodynamic Favorability: The reaction occurs when the resulting products are more stable than the original compound.

Disproportionation is the opposite of comproportionation, where two species of the same element combine to form a single product with an intermediate oxidation state.

Comproportionation reactions:

A comproportionation reaction (also known as a symproportionation reaction) is a type of redox reaction in which two species of the same element, but with different oxidation states, react to form a single product with an intermediate oxidation state.

This reaction is essentially the reverse of a disproportionation reaction, where a single species is both oxidized and reduced. In comproportionation, two species are balanced toward an intermediate state.

General Form:

Aⁿ⁺+A^m+⟶2A^p+

Where n>p>m, meaning the higher oxidation state is reduced and the lower oxidation state is oxidized to produce the intermediate oxidation state.

Examples:

1. The redox reaction in lead storage cell:

Pb(s) + PbO₂(s) + 2H₂SO₄(aq) → 2PbSO₄(s) + 2H₂O(l)

Here the lower oxidation state 0 in Pb is oxidized to +2 in PbSO₄ and higher oxidation state +4 in PbO₂ is reduced to +2 in the same PbSO₄^.

2. Reaction in Iodine Chemistry:

5I^- + IO₃^- + 6H⁺ → 3I₂ + 3H₂O. 

Here I⁻ is oxidised to I₂ and I⁶⁺ is reduced to I₂.

 3. Reaction in Copper Chemistry:

Cu⁺+Cu²⁺⟶2Cu^1.5+

4. Sulfur Chemistry:

H₂S+S₈⟶2HSn

Characteristics:

Oxidation State Balance: The reaction balances the oxidation states of the same element.

Common in Multivalent Elements: Often observed in elements with multiple stable oxidation states, such as sulfur, iodine, nitrogen, and transition metals.

Thermodynamic Feasibility: The reaction occurs when the intermediate oxidation state is more stable than the starting states.