Problem 207 - A Complete Resource Book in Chemistry for JEE Main

Question

Which of the following chemical reactions depicts the oxidizing behaviour of \(\mathrm{H}_{2} \mathrm{SO}_{4} ?\) [2006] (a) \(2 \mathrm{HI}+\mathrm{H}_{2} \mathrm{SO}_{4} \longrightarrow \mathrm{I}_{2}+\mathrm{SO}_{2}+2 \mathrm{H}_{2} \mathrm{O}\) (b) \(\mathrm{Ca}(\mathrm{OH})_{2}^{2}+\mathrm{H}_{2} \mathrm{SO}_{4} \longrightarrow \mathrm{CaSO}_{4}+2 \mathrm{H}_{2} \mathrm{O}\) (c) \(\mathrm{NaCl}+\mathrm{H}_{2} \mathrm{SO}_{4} \longrightarrow \mathrm{NaHSO}_{4}+\mathrm{HCl}\) (d) \(2 \mathrm{PCl}_{5}+\mathrm{H}_{2} \mathrm{SO}_{4} \longrightarrow 2 \mathrm{POCl}_{3}+2 \mathrm{HCl}+\mathrm{SO}_{2} \mathrm{Cl}_{2}\)

Step-by-Step Solution

Answer

Option (a) demonstrates the oxidizing behavior of \(\mathrm{H}_2\mathrm{SO}_4\).

Step 1: Understand Oxidizing Behavior

An oxidizing agent gains electrons and gets reduced in the process of oxidizing another substance. In the context of this problem, we need to identify in which reaction sulfur in \(\mathrm{H}_2\mathrm{SO}_4\) is being reduced.

Step 2: Analyze Option A

In the reaction \(2 \mathrm{HI}+\mathrm{H}_{2}\mathrm{SO}_{4} \to \mathrm{I}_{2}+\mathrm{SO}_{2}+2 \mathrm{H}_{2}\mathrm{O}\), sulfur goes from an oxidation state of +6 in \(\mathrm{H}_2\mathrm{SO}_4\) to +4 in \(\mathrm{SO}_2\). This indicates sulfur is reduced, thus \(\mathrm{H}_2\mathrm{SO}_4\) acts as an oxidizing agent.

Step 3: Analyze Option B

In the reaction \(\mathrm{Ca} (\mathrm{OH})_2 + \mathrm{H}_2\mathrm{SO}_4 \to \mathrm{CaSO}_4 + 2 \mathrm{H}_2\mathrm{O}\), there is no change in the oxidation state of sulfur, meaning no oxidizing behavior is demonstrated.

Step 4: Analyze Option C

In the reaction \(\mathrm{NaCl}+\mathrm{H}_{2}\mathrm{SO}_{4} \to \mathrm{NaHSO}_{4}+\mathrm{HCl}\), sulfur does not change its oxidation state; therefore, \(\mathrm{H}_2\mathrm{SO}_4\) is not acting as an oxidizing agent.

Step 5: Analyze Option D

In the reaction \(2\mathrm{PCl}_5 + \mathrm{H}_2\mathrm{SO}_4 \to 2 \mathrm{POCl}_3 + 2 \mathrm{HCl} + \mathrm{SO}_2\mathrm{Cl}_2\), sulfur's oxidation state changes from +6 in \(\mathrm{H}_2\mathrm{SO}_4\) to +6 in \(\mathrm{SO}_2\mathrm{Cl}_2\), so there's no reduction, indicating no oxidizing behavior.

Step 6: Conclusion: Identify the Correct Reaction

Considering the analysis, option (a) is the correct reaction as the sulfur in \(\mathrm{H}_2\mathrm{SO}_4\) is reduced from +6 to +4, demonstrating oxidizing behavior.

Key Concepts

Oxidation-Reduction Reactions

Oxidation-reduction reactions, commonly known as redox reactions, are fascinating processes where electrons are transferred between substances. These reactions involve two key components: the oxidation process and the reduction process. When a substance loses electrons, it is said to be oxidized, while the substance that gains electrons is reduced. This transfer is what differentiates a redox reaction from other types of chemical reactions.

• The oxidized substance acts as a reducing agent because it donates electrons to another substance.
• The reduced substance acts as an oxidizing agent because it accepts electrons from the oxidized substance.

In the context of the exercise, we are particularly interested in the reduction process of sulfuric acid \((\mathrm{H}_2\mathrm{SO}_4)\) and its role as an oxidizing agent. In reaction \((a)\), \(\mathrm{H}_2\mathrm{SO}_4\) oxidizes \(\mathrm{HI}\) to \(\mathrm{I}_2\) while being reduced to \(\mathrm{SO}_2\). Understanding this electron transfer process is crucial to identifying which reaction allows sulfuric acid to demonstrate its oxidizing behavior.

Oxidizing Agents

An oxidizing agent is a vital player in redox reactions. It facilitates the process by accepting electrons from the substance being oxidized. As a result, the oxidizing agent itself undergoes reduction.

• Common oxidizing agents include oxygen, halogens, and various acids, such as sulfuric acid \((\mathrm{H}_2\mathrm{SO}_4)\).
• Oxidizing agents are essential for processes ranging from metabolism in living organisms to industrial applications like bleaching and disinfection.

In reaction \((a)\) of the exercise, \(\mathrm{H}_2\mathrm{SO}_4\) acts as an oxidizing agent by reducing its sulfur atom from an oxidation state of +6 to +4. Understanding the role and mechanism of oxidizing agents helps demystify how substances interact and transform during chemical reactions.

Oxidation States

The concept of oxidation states, often referred to as oxidation numbers, helps us keep track of electron transfer in redox reactions. An oxidation state is an indicator of the degree of oxidation of an atom in a chemical compound. It is usually expressed as an integer representing the number of electrons lost or gained.

• A positive oxidation state indicates a loss of electrons, while a negative oxidation state indicates a gain of electrons.
• Understanding these states allows chemists to determine how electrons are being shared or transferred in a reaction.

For instance, in the exercise, the oxidation state of sulfur in \(\mathrm{H}_2\mathrm{SO}_4\) is +6. During the reaction in option \((a)\), sulfur is reduced to an oxidation state of +4 in \(\mathrm{SO}_2\). This change signifies the gaining of electrons by sulfur and confirms that sulfuric acid is acting as an oxidizing agent. A firm grasp of oxidation states is essential for comprehending the electron transfer processes in redox reactions.