Question

In which of the following reactions the hydrogen peroxide acts as a reducing agent?

• A. \(PbS +4 H _2 O _2 \rightarrow PbSO _4+4 H _2 O\)
• B. \(2 Fe ^{2+}+ H _2 O _2 \rightarrow 2 Fe ^{3+}+2 OH ^{-}\)
• C. \(HOCl + H _2 O _2 \rightarrow H _3 O ^{+}+ Cl ^{-}+ O _2\)
• D. \(Mn ^{2+}+ H _2 O _2 \rightarrow Mn ^{4+}+2 OH ^{-}\)

Hint:

\(H_2O_2\) acts as a reducing agent when it gets oxidized to \(O_2 \)in the reaction.

Answer 1

The Correct Option is C

To determine in which of the given reactions hydrogen peroxide \((H_2O_2)\) acts as a reducing agent, we must analyze the changes in oxidation states that occur during each reaction.

1. \(PbS + 4 H_2O_2 \rightarrow PbSO_4 + 4 H_2O\):
   • In this reaction, the oxidation state of lead (Pb) remains unchanged, but sulfur (S) in PbS is oxidized from -2 to +6 in \(PbSO_4\). Hence, \(H_2O_2\) acts as an oxidizing agent here.
2. \(2 Fe^{2+} + H_2O_2 \rightarrow 2 Fe^{3+} + 2 OH^-\):
   • The iron ions \(Fe^{2+}\) are oxidized to \(Fe^{3+}\), indicating that \(H_2O_2\) is acting as an oxidizing agent rather than a reducing agent.
3. \(HOCl + H_2O_2 \rightarrow H_3O^+ + Cl^- + O_2\):
   • In this reaction, chlorine in \(HOCl\) is reduced from an oxidation state of +1 to -1 in Cl\(^-\). Meanwhile, the oxygen in \(H_2O_2\) is oxidized from -1 to 0 in \(O_2\). Thus, hydrogen peroxide is acting as a reducing agent here.
4. \(Mn^{2+} + H_2O_2 \rightarrow Mn^{4+} + 2 OH^-\):
   • The manganese ion \(Mn^{2+}\) is oxidized to \(Mn^{4+}\), so \(H_2O_2\) is acting as an oxidizing agent in this reaction as well.

Based on the analysis, in the reaction \(HOCl + H_2O_2 \rightarrow H_3O^+ + Cl^- + O_2\), hydrogen peroxide acts as a reducing agent, which is the correct answer.