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) \(\mathrm{Ca}(\mathrm{OH})_{2}+\mathrm{H}_{2} \mathrm{SO}_{4} \longrightarrow \mathrm{CaSO}_{4}+2 \mathrm{H}_{2} \mathrm{O}\) (c) \(\mathrm{NaCl}+\mathrm{H}_{2} \mathrm{SO}_{4} \stackrel{2}{\longrightarrow} \mathrm{NaHSO}_{4}+\mathrm{HCl}\) (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

(a) 2HI + H₂SO₄ → I₂ + SO₂ + 2H₂O depicts the oxidizing behavior of H₂SO₄.

Step 1: Define Oxidizing Behavior

The oxidizing behavior of a chemical compound is an ability to accept electrons in a redox reaction, leading to the increase in its oxidation number while causing the oxidation number of another element to decrease.

Step 2: Analyze Reaction (a)

In the reaction \(2 \mathrm{HI} + \mathrm{H}_2\mathrm{SO}_4 \longrightarrow \mathrm{I}_2 + \mathrm{SO}_2 + 2\mathrm{H}_2\mathrm{O}\), \(\mathrm{HI}\) is oxidized to \(\mathrm{I}_2\), with iodine's oxidation state increasing from \(-1\) in \(\mathrm{HI}\) to \(0\) in \(\mathrm{I}_2\). Simultaneously, sulfur is reduced from \(+6\) in \(\mathrm{H}_2\mathrm{SO}_4\) to \(+4\) in \(\mathrm{SO}_2\), showing \(\mathrm{H}_2\mathrm{SO}_4\)'s oxidizing behavior.

Step 3: Analyze Reaction (b)

The reaction \(\mathrm{Ca}( ext{OH})_2 + \mathrm{H}_2\mathrm{SO}_4 \rightarrow \mathrm{CaSO}_4 + 2 \mathrm{H}_2\mathrm{O}\) is a double displacement reaction without changes in oxidation states, indicating no oxidizing or reducing behavior from \(\mathrm{H}_2\mathrm{SO}_4\).

Step 4: Analyze Reaction (c)

In \(\mathrm{NaCl} + \mathrm{H}_2\mathrm{SO}_4 \rightarrow \mathrm{NaHSO}_4 + \mathrm{HCl}\), the reaction is primarily an acid-base reaction where no change in oxidation state occurs, signifying no oxidizing behavior is exhibited by \(\mathrm{H}_2\mathrm{SO}_4\).

Step 5: Analyze Reaction (d)

The reaction \(2 \mathrm{PCl}_5 + \mathrm{H}_2\mathrm{SO}_4 \rightarrow 2 \mathrm{POCl}_3 + 2 \mathrm{HCl} + \mathrm{SO}_2\mathrm{Cl}_2\) involves oxidizing and reducing, but mainly results in the formation of oxychloride and does not strongly depict \(\mathrm{H}_2\mathrm{SO}_4\)'s oxidizing behavior opposed to reaction (a).

Step 6: Conclusion

Analyzing all the options, reaction (a) distinctly demonstrates the oxidizing behavior of \(\mathrm{H}_2\mathrm{SO}_4\) as it oxidizes \(\mathrm{HI}\) to \(\mathrm{I}_2\) while itself being reduced.

Key Concepts

Redox Reactions

In chemical reactions, a redox process is where oxidation and reduction occur simultaneously. These types of reactions are pivotal in understanding how substances interact on a molecular level. To identify a redox reaction:

• Oxidation happens when a molecule, atom, or ion loses electrons. The element undergoing oxidation will see an increase in its oxidation state.
• Reduction is the gain of electrons, leading to a decrease in oxidation state of the participating substance.
• The substance that gets reduced is considered the "oxidizing agent," while the one that gets oxidized is the "reducing agent."

For example, in the reaction \(2\ \mathrm{HI} + \mathrm{H}_2\mathrm{SO}_4 \rightarrow \mathrm{I}_2 + \mathrm{SO}_2 + 2\mathrm{H}_2\mathrm{O}\), \(\mathrm{HI}\) is oxidized to \(\mathrm{I}_2\), while \(\mathrm{H}_2\mathrm{SO}_4\) is reduced, demonstrating a classic redox reaction.

Oxidation States

Understanding oxidation states helps us gauge how electrons are distributed in a molecule. It's a hypothetical charge that an atom would have if all the bonds were ionic. For better clarity, let's look at key points:

• Elemental state: Atoms in their elemental form have an oxidation state of zero, such as \(\mathrm{I}_2\) or \(\mathrm{O}_2\).
• Simple ions: For monoatomic ions, the oxidation state is equal to the charge of the ion. For instance, \(\mathrm{Na}^+\) has an oxidation state of \(+1\).
• Compounds: In molecules, the sum of oxidation states of all atoms equals the charge on the molecule. For instance, in \(\mathrm{H}_2\), each \(\mathrm{H}\) has an oxidation state of \(+1\), giving the molecule a net oxidation state of 0.

In the reaction \(2\ \mathrm{HI} + \mathrm{H}_2\mathrm{SO}_4 \rightarrow \mathrm{I}_2 + \mathrm{SO}_2 + 2\mathrm{H}_2\mathrm{O}\), iodine's oxidation state transitions from \(-1\) to 0, indicating oxidation, while sulfur changes from \(+6\) to \(+4\), indicating reduction.

Chemical Reactions Analysis

To effectively analyze chemical reactions, you'll need to follow a systematic approach that includes examining changes in the oxidation states to determine the kind of reaction taking place:

• Identify reactants and products: It's crucial to write down both sides of the equation to understand the starting materials and the resulting compounds.
• Determine oxidation states: Compare oxidation states of each element in reactants and products. A change signifies a redox process.
• Check for conservation of mass and charge: Ensure the number of atoms and.the overall charge are balanced on both sides, as these are conserved during reactions.
• Identify roles: Establish which substances act as oxidizing and reducing agents based on their role in gaining or losing electrons.

In our demonstrated reaction, \(\mathrm{H}_2\mathrm{SO}_4\) acts as an oxidizing agent, leading to the oxidation of \(\mathrm{HI}\). Understanding these steps allows us to deduce how and why certain reactions occur, giving us insight into the behavior and interaction of different chemicals.