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

Question

Which of the following chemical reactions depicts the oxidizing behaviour of \(\mathrm{H}_{2} \mathrm{SO}_{4}\) ? (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

The correct reaction depicting oxidizing behaviour is (a).

Step 1: Understanding Oxidizing Behavior

Recognize that an oxidizing agent is a substance that can accept electrons during a chemical reaction. To determine which reaction involves the oxidizing behavior of \( \mathrm{H}_{2} \mathrm{SO}_{4} \), identify where \( \mathrm{H}_{2} \mathrm{SO}_{4} \) is reduced by gaining electrons from a reducing agent.

Step 2: Inspect Reaction (a) for Oxidation-Reduction

In the reaction \( 2 \mathrm{HI} + \mathrm{H}_{2} \mathrm{SO}_{4} \rightarrow \mathrm{I}_{2} + \mathrm{SO}_{2} + 2 \mathrm{H}_{2} \mathrm{O} \), \( \mathrm{I}^{-} \) ions are oxidized to \( \mathrm{I}_{2} \), and \( \mathrm{H}_{2} \mathrm{SO}_{4} \) is reduced to \( \mathrm{SO}_{2} \). This shows that \( \mathrm{H}_{2} \mathrm{SO}_{4} \) is acting as an oxidizing agent.

Step 3: Analyze the other reactions

- **Reaction (b):** In \( \mathrm{Ca(OH)}_{2} + \mathrm{H}_{2} \mathrm{SO}_{4} \rightarrow \mathrm{CaSO}_{4} + 2 \mathrm{H}_{2} \mathrm{O} \), \( \mathrm{H}_{2} \mathrm{SO}_{4} \) acts as an acid, not an oxidizer.- **Reaction (c):** In \( \mathrm{NaCl} + \mathrm{H}_{2} \mathrm{SO}_{4} \rightarrow \mathrm{NaHSO}_{4} + \mathrm{HCl} \), \( \mathrm{H}_{2} \mathrm{SO}_{4} \) acts as an acid, not an oxidizer.- **Reaction (d):** In \( 2 \mathrm{PCl}_{5} + \mathrm{H}_{2} \mathrm{SO}_{4} \rightarrow 2 \mathrm{POCl}_{3} + 2 \mathrm{HCl} + \mathrm{SO}_{2} \mathrm{Cl}_{2} \), \( \mathrm{H}_{2} \mathrm{SO}_{4} \) acts predominantly as a dehydrating agent, not as an oxidizer.

Step 4: Conclusion

Among the provided reactions, only reaction (a) involves an oxidation-reduction process where \( \mathrm{H}_{2} \mathrm{SO}_{4} \) is reduced, thereby displaying oxidizing behavior. Therefore, the answer is reaction (a).

Key Concepts

sulfuric acid

Sulfuric acid, often represented by the chemical formula \( \mathrm{H}_{2} \mathrm{SO}_{4} \), is a dense, oily liquid with a wide range of applications. It's renowned for its strong acidic nature and powerful oxidizing ability. In chemical reactions, it can behave as an oxidizing agent and a dehydrating agent.
- **Primary Role**: Often used in the production of fertilizers, sulfuric acid's oxidizing property can enhance reactions by accepting electrons from other substances. - **Versatility**: This makes it a key component in the chemical industry, both for processes requiring acid and ones requiring the compound to act as an agent that facilitates the release or transfer of electrons in redox reactions.
Sulfuric acid can break down into sulfur dioxide (\( \mathrm{SO}_{2} \)), especially under high temperatures, which illustrates its effective role as an oxidizer. As we investigate reactions, we observe sulfuric acid taking on this role in various environments.

redox reactions

Redox reactions, short for reduction-oxidation reactions, are fundamental to chemistry, involving the movement of electrons between substances. A redox reaction is characterized by two half-reactions, oxidation and reduction.
- **Oxidation**: This is the process where a substance loses electrons, increasing its oxidation state. - **Reduction**: Here, a substance gains electrons, decreasing its oxidation state.
In any redox process, you'll always find an oxidizing agent and a reducing agent. The oxidizing agent gains electrons, whereas the reducing agent loses them.
Redox reactions are vital in various processes, including energy production through combustion, cellular respiration, and even the rusting of metals. They underscore the interactions in many biochemical and industrial processes.

chemical equations

Chemical equations are symbolic representations of chemical reactions. They provide a quick way to visualize what's happening at a molecular level during a reaction.
- **Components**: The equations list reactants and products, often accompanied by physical states (e.g., \( g \), \( l \), \( s \), \( aq \) for gas, liquid, solid, and aqueous).- **Balancing**: A crucial part of chemical equations is ensuring they are balanced. This means the number of atoms for each element is the same on both sides of the equation.
Consider the reaction: \[ 2 \mathrm{HI} + \mathrm{H}_{2} \mathrm{SO}_{4} \rightarrow \mathrm{I}_{2} + \mathrm{SO}_{2} + 2 \mathrm{H}_{2} \mathrm{O} \]Here, all elements have been balanced, maintaining mass conservation and indicating the transformation of substances involved.

oxidation-reduction

Oxidation-reduction, or redox, processes describe reactions where electron transfer occurs between molecules or atoms, altering their oxidation states.
- **Identifying Redox**: To find if a reaction includes redox processes, check if there's a change in oxidation numbers.- **Example Insight**: In the reaction \( 2 \mathrm{HI} + \mathrm{H}_{2} \mathrm{SO}_{4} \rightarrow \mathrm{I}_{2} + \mathrm{SO}_{2} + 2 \mathrm{H}_{2} \mathrm{O} \), the iodine (\( \mathrm{I}^- \)) is oxidized to \( \mathrm{I}_{2} \), and sulfuric acid is reduced to \( \mathrm{SO}_{2} \).
Redox reactions are intrinsic to both natural and industrial processes, serving a pivotal role in chemistry. Understanding these transformations can elucidate many practical applications, from metallurgy to energy production.