Question

An example of disproportionation reaction is :

• A. \( 2H_2O_2 \rightarrow 2H_2O + O_2 \)
• B. \( 2KMnO_4 \rightarrow K_2MnO_4 + MnO_2 + O_2 \)
• C. \( 2MnO_4^- + 10I^- + 16H^+ \rightarrow 2Mn^{2+} + 5I_2 + 8H_2O \)
• D. \( 2NaI + Cl_2 \rightarrow 2NaCl + I_2 \)

Hint:

In disproportionation reactions, always check if the same element is getting oxidized and reduced simultaneously by comparing oxidation states before and after the reaction.

Answer 1

The Correct Option is A

Step 1: Understand disproportionation reaction.
A disproportionation reaction is a type of redox reaction in which the same substance undergoes both oxidation and reduction simultaneously.

Step 2: Analyze option (A).
In \( H_2O_2 \), oxygen has oxidation state \( -1 \).
After reaction:
In \( H_2O \), oxygen is \( -2 \) (reduction)
In \( O_2 \), oxygen is \( 0 \) (oxidation)
Thus, the same element oxygen is both oxidized and reduced.

Step 3: Check other options.
(B) \( KMnO_4 \) decomposes but manganese does not undergo both oxidation and reduction in the same way.
(C) This is a redox reaction between two different species, not disproportionation.
(D) This is a displacement reaction, not disproportionation.

Step 4: Conclusion.
Only option (A) shows the same substance undergoing both oxidation and reduction.
Therefore:
\[ \boxed{2H_2O_2 \rightarrow 2H_2O + O_2} \]