Redox I: Disproportionation and Redox Classifications

Full Notes

What Is Disproportionation?

A disproportionation reaction is a specific type of redox reaction where one element is both oxidised and reduced in the same reaction. This results in the same element appearing in two different oxidation states in the products.

To recognise disproportionation:

• Look for a single element that appears once on the left-hand side (reactant).
• Then appears in two different forms on the right-hand side with different oxidation numbers.

Example of Disproportionation

Reaction of chlorine with water:

Cl₂ + H₂O → HCl + HClO

Oxidation number changes:

• Cl in Cl₂ = 0
• Cl in HCl = −1 = reduction
• Cl in HClO = +1 = oxidation

So chlorine is simultaneously oxidised and reduced, meaning this is a disproportionation reaction.

Reaction of chlorine with cold, dilute sodium hydroxide:

Cl₂ + 2NaOH → NaCl + NaClO + H₂O

• Cl in Cl₂ = 0
• Cl in NaCl = −1 (reduction)
• Cl in NaClO = +1 (oxidation)

Using Oxidation Numbers to Classify Reactions

Oxidation numbers can be used to identify redox processes (see oxidation numbers):

• If any oxidation number changes, the reaction is a redox reaction.
• If one species is both oxidised and reduced, the reaction is a disproportionation.

Summary

• Disproportionation involves one element being both oxidised and reduced in the same reaction.
• The element ends up in two different oxidation states in the products.
• Oxidation numbers help to identify disproportionation reactions.
• Examples include chlorine reacting with water or sodium hydroxide.