Problem 118 - A Complete Resource Book in Chemistry for JEE Main

Question

Which one of the following reactions represents the reducing property of \(\mathrm{H}_{2} \mathrm{O}_{2} ?\) (a) \(2 \mathrm{NaI}+\mathrm{H}_{2} \mathrm{SO}_{4}+\mathrm{H}_{2} \mathrm{O}_{2} \rightarrow \mathrm{Na}_{2} \mathrm{SO}_{4}+\mathrm{I}_{2}+2 \mathrm{H}_{2} \mathrm{O}\) (b) \(\mathrm{PbO}_{2}+\mathrm{H}_{2} \mathrm{O}_{2} \rightarrow \mathrm{PbO}+\mathrm{H}_{2} \mathrm{O}+\mathrm{O}_{2}\) (c) \(2 \mathrm{KMnO}_{4}+3 \mathrm{H}_{2} \mathrm{SO}_{4}+5 \mathrm{H}_{2} \mathrm{O}_{2} \rightarrow \mathrm{K}_{2} \mathrm{SO}_{4}+8 \mathrm{H}_{2} \mathrm{O}\) \(+5 \mathrm{O}_{2}+2 \mathrm{MnO}_{2}\) (d) \(2 \mathrm{~K}_{3}\left[\mathrm{Fe}(\mathrm{CN})_{6}\right]+2 \mathrm{KOH}+\mathrm{H}_{2} \mathrm{O}_{2} \rightarrow 2 \mathrm{~K}_{4}\left[\mathrm{Fe}(\mathrm{CN})_{6}\right]\) \(+2 \mathrm{H}_{2} \mathrm{O}+\mathrm{O}_{2}\)

Step-by-Step Solution

Answer

Reactions (c) and (d) both represent the reducing property of \( \mathrm{H}_{2} \mathrm{O}_{2} \).

Step 1: Understanding Reducing Property

A substance acts as a reducing agent when it donates electrons in a chemical reaction. In a redox reaction, the reducing agent is oxidized, meaning it loses electrons.

Step 2: Analyze Reaction (a)

The first reaction is: \( 2 \mathrm{NaI} + \mathrm{H}_{2} \mathrm{SO}_{4} + \mathrm{H}_{2} \mathrm{O}_{2} \rightarrow \mathrm{Na}_{2} \mathrm{SO}_{4} + \mathrm{I}_{2} + 2 \mathrm{H}_{2} \mathrm{O} \). In this reaction, \( \mathrm{H}_{2} \mathrm{O}_{2} \) is likely oxidized to \( \mathrm{O}_{2} \), however there is a reduction of \( \mathrm{I}^{-} \) to \( \mathrm{I}_{2} \). This indicates \( \mathrm{H}_{2} \mathrm{O}_{2} \) acts as an oxidizing agent, not reducing.

Step 3: Analyze Reaction (b)

The second reaction is: \( \mathrm{PbO}_{2} + \mathrm{H}_{2} \mathrm{O}_{2} \rightarrow \mathrm{PbO} + \mathrm{H}_{2} \mathrm{O} + \mathrm{O}_{2} \). Here, \( \mathrm{H}_{2} \mathrm{O}_{2} \) is decomposing and producing \( \mathrm{O}_{2} \), indicating oxidation, so it acts as an oxidizing agent rather than a reducing one.

Step 4: Analyze Reaction (c)

The third reaction is: \( 2 \mathrm{KMnO}_{4} + 3 \mathrm{H}_{2} \mathrm{SO}_{4} + 5 \mathrm{H}_{2} \mathrm{O}_{2} \rightarrow 2 \mathrm{K}_{2} \mathrm{SO}_{4} + 8 \mathrm{H}_{2} \mathrm{O} + 5 \mathrm{O}_{2} + 2 \mathrm{MnO}_{2} \). Here, \( \mathrm{KMnO}_{4} \) (a known oxidizing agent) itself is being reduced while \( \mathrm{H}_{2} \mathrm{O}_{2} \) is oxidized, showing its reducing property.

Step 5: Analyze Reaction (d)

The fourth reaction is: \( 2 \mathrm{~K}_{3}\left[\mathrm{Fe}(\mathrm{CN})_{6}\right] + 2 \mathrm{KOH} + \mathrm{H}_{2} \mathrm{O}_{2} \rightarrow 2 \mathrm{~K}_{4}\left[\mathrm{Fe}(\mathrm{CN})_{6}\right] + 2 \mathrm{H}_{2} \mathrm{O} + \mathrm{O}_{2} \). Here, \( \mathrm{H}_{2} \mathrm{O}_{2} \) causes the reduction of \( \mathrm{Fe} \), which indicates that \( \mathrm{H}_{2} \mathrm{O}_{2} \) is acting as a reducing agent.

Step 6: Conclusion

Compare the information from each analyzed reaction. Reactions (c) and (d) show the reducing property of \( \mathrm{H}_{2} \mathrm{O}_{2} \), where it helps another species reduce.

Key Concepts

Reducing Agents

Reducing agents are substances that facilitate the reduction of another material by donating electrons to it. During this exchange, reducing agents themselves undergo oxidation and become more positive in charge.

Understanding reducing agents is vital because they drive redox reactions by providing the electrons needed for reduction. Characteristics of reducing agents:

• They lose electrons in a reaction.
• They become oxidized in the process, often evidenced by an increase in their oxidation state.
• They are crucial in many chemical processes including combustion, metabolism, and industrial chemical reactions.

For example, hydrogen gas ( H₂) is a traditional reducing agent, frequently used to reduce metal oxides to pure metals. In the context of redox reactions, identifying the reducing agent reveals where electron donation occurs, crucial for understanding reaction mechanisms.

Oxidizing Agents

Oxidizing agents are substances that cause oxidation by accepting electrons from another species. During this process, the oxidizing agent itself gets reduced, which entails a decrease in its oxidation state.

Application of oxidizing agents is widespread, from industrial manufacturing to biological respiration. Key features of oxidizing agents:

• They gain electrons in a reaction.
• They undergo reduction as they lose positive charge.
• They are essential in applications like bleaching, disinfection, and energy production.

A common example of an oxidizing agent is oxygen ( O₂), which readily accepts electrons during combustion. In any redox reaction, identifying the oxidizing agent is critical to depicting the chemical changes and energy transfer.

Hydrogen Peroxide

Hydrogen Peroxide ( H₂O₂) is a versatile chemical often used in redox reactions. It can act as both a reducing and oxidizing agent depending on the circumstances.

In certain conditions, H₂O₂ donates oxygen, thus acting as an oxidizing agent, while in other situations, it can accept electrons and function as a reducing agent. Properties of H₂O₂:

• It is a pale blue liquid in its pure form, but usually appears as a clear solution in water.
• In dilute solutions, it is commonly used as a disinfectant and antiseptic.
• As a high-energy oxidizing agent, it decomposes into water and oxygen, releasing energy.

For example, in the reaction with KMnO₄, H₂O₂ demonstrates its reducing characteristics by causing the reduction of manganese while being oxidized itself. Understanding this dual nature helps in predicting and harnessing its behavior in various chemical and industrial processes.

Electrochemistry

Electrochemistry is the branch of chemistry that studies the relationship between electricity and chemical changes in a reaction. Redox reactions are central to electrochemical processes as they involve transfer of electrons.

Electrochemistry encompasses a variety of applications, from batteries and fuel cells to electroplating and sensors. Principles of electrochemistry:

• Involves oxidation and reduction reactions, where electrons are transferred between species.
• Governs the operation of electrochemical cells, where chemical energy is converted to electrical energy, or vice versa.
• Enables advancements in technology that rely on the flow of electrons for energy storage and synthesis of materials.

For students, mastering electrochemistry means understanding the electron exchange in reactions and applying these principles in both theoretical and practical contexts.