Question

Which of the following equation depicts the oxidizing nature of $H _{2} O _{2} ?$

• A. $KIO _{4}+ H _{2} O _{2} \rightarrow KIO _{3}+ H _{2} O + O _{2}$
• B. $2 I ^{-}+ H _{2} O _{2}+2 H ^{+} \rightarrow I _{2}+2 H _{2} O$
• C. $I _{2}+ H _{2} O _{2}+2 OH ^{-} \rightarrow 2 I ^{-}+2 H _{2} O + O _{2}$
• D. $Cl _{2}+ H _{2} O _{2} \rightarrow 2 HCl + O _{2}$

Answer 1

The Correct Option is B

The question is asking which equation demonstrates the oxidizing nature of hydrogen peroxide ($H_{2}O_{2}$). To determine this, we need to identify the reaction where ($H_{2}O_{2}$) acts as an oxidizing agent by accepting electrons, thereby reducing other species.

Let's analyze each option:

1. $KIO_{4} + H_{2}O_{2} \rightarrow KIO_{3} + H_{2}O + O_{2}$
   • This reaction involves $KIO_{4}$ getting reduced to $KIO_{3}$. Here, ($H_{2}O_{2}$) is involved in a decomposition reaction, where it releases oxygen. ($H_{2}O_{2}$) does not act as an oxidizing agent in this reaction.
2. $2 I^{-} + H_{2}O_{2} + 2 H^{+} \rightarrow I_{2} + 2 H_{2}O$
   • In this reaction, $I^{-}$ ions (iodide ions) are oxidized to $I_{2}$. The gain of electrons by ($H_{2}O_{2}$) indicates its role as an oxidizing agent. Here, ($H_{2}O_{2}$) is reduced, making it the correct option that demonstrates its oxidizing nature.
3. $I_{2} + H_{2}O_{2} + 2 OH^{-} \rightarrow 2 I^{-} + 2 H_{2}O + O_{2}$
   • In this equation, $I_{2}$ is reduced to $I^{-}$, which means ($H_{2}O_{2}$) acts as a reducing agent rather than an oxidizing agent.
4. $Cl_{2} + H_{2}O_{2} \rightarrow 2 HCl + O_{2}$
   • This reaction involves chlorine ($Cl_{2}$) being reduced to hydrochloric acid ($HCl$), indicating ($H_{2}O_{2}$) acting as a reducing agent in this context.

Therefore, the correct answer is option 2: $2 I^{-} + H_{2}O_{2} + 2 H^{+} \rightarrow I_{2} + 2 H_{2}O$, which depicts the oxidizing nature of ($H_{2}O_{2}$). In this reaction, hydrogen peroxide oxidizes iodide ions to iodine.