Problem 71 - 43 Years JEE ADVANCED (1978-2020) + JEE MAIN Chapterwise & Topicwise Solved Papers Chemistry

Question

\(\mathrm{Fe}^{3+}\) is reduced to \(\mathrm{Fe}^{2+}\) by using (a) \(\mathrm{H}_{2} \mathrm{O}_{2}\) in presence of \(\mathrm{NaOH}\) (b) \(\mathrm{Na}_{2} \mathrm{O}_{2}\) in water (c) \(\mathrm{H}_{2} \mathrm{O}_{2}\) in presence of \(\mathrm{H}_{2} \mathrm{SO}_{4}\) (d) \(\mathrm{Na}_{2} \mathrm{O}_{2}\) in presence of \(\mathrm{H}_{2} \mathrm{SO}_{4}\)

Step-by-Step Solution

Answer

(c) \\mathrm{H}_{2} \\mathrm{O}_{2} in \\mathrm{H}_{2} \\mathrm{SO}_{4}

Step 1: Reaction with H2O2 in NaOH

To find if \( \mathrm{Fe}^{3+} \) is reduced by \( \mathrm{H}_{2} \mathrm{O}_{2} \) in an alkaline medium (\( \mathrm{NaOH} \)), examine the oxidation-reduction potential. In the alkaline condition, \( \mathrm{H}_2\mathrm{O}_2 \) acts as a reducing agent. Hence, \( \mathrm{Fe}^{3+} \) will be reduced to \( \mathrm{Fe}^{2+} \).

Step 2: Reaction with Na2O2 in Water

Consider \( \mathrm{Na}_{2} \mathrm{O}_{2} \) in water. Here, \( \mathrm{Na}_2\mathrm{O}_2 \) produces \( \mathrm{H}_2\mathrm{O}_2 \) in acidic conditions, which normally favors oxidation. Therefore, it is not effective for reducing \( \mathrm{Fe}^{3+} \) to \( \mathrm{Fe}^{2+} \) under these conditions.

Step 3: Reaction with H2O2 in H2SO4

Investigate \( \mathrm{H}_{2} \mathrm{O}_{2} \) in a strong acid like \( \mathrm{H}_{2} \mathrm{SO}_{4} \). In acidic conditions, \( \mathrm{H}_2\mathrm{O}_2 \) acts as an oxidizing agent and is able to reduce \( \mathrm{Fe}^{3+} \) to \( \mathrm{Fe}^{2+} \).

Step 4: Reaction with Na2O2 in H2SO4

Look at the behavior of \( \mathrm{Na}_{2} \mathrm{O}_{2} \) in the presence of \( \mathrm{H}_{2} \mathrm{SO}_{4} \). Similar to \( \mathrm{Na}_2\mathrm{O}_2 \) in water, the peroxide will produce \( \mathrm{H}_2\mathrm{O}_2 \), but due to the acidic environment, \( \mathrm{Fe}^{3+} \) will not be reduced as it tends to favor oxidation.

Key Concepts

Oxidation-Reduction Potential

Oxidation-reduction potential, or redox potential, is a measure that indicates the ability of a chemical species to acquire electrons and thereby be reduced. This potential is generally measured in volts and serves as a critical indicator when examining chemical reactions, especially redox reactions.
Consider it like a "pull" for electrons, similar to how gravity pulls an apple to the ground. If a substance has a high oxidation-reduction potential, it tends to pull electrons towards itself, making it more likely to be reduced.

• If the redox potential is positive, the species can act as an oxidizing agent.
• If the potential is negative, it serves as a reducing agent.

In the exercise, various agents were examined to see if they could reduce \(\mathrm{Fe}^{3+}\), converting it into \(\mathrm{Fe}^{2+}\). This process involves electron transfer, dictated by the redox potential of the agents involved.

Reducing Agents

Reducing agents are substances that donate electrons to another substance in a chemical reaction. By doing this, they themselves become oxidized, as they lose electrons.
Think of reducing agents as electron "givers"; they push electrons into other substances. When \(\mathrm{Fe}^{3+}\) is reduced, a reducing agent donates an electron, transforming it to \(\mathrm{Fe}^{2+}\).

• In alkaline conditions, such as when \(\mathrm{H}_2\mathrm{O}_2\) is combined with \(\mathrm{NaOH}\), \(\mathrm{H}_2\mathrm{O}_2\) acts as a reducing agent.
• In our problem, some conditions did not provide effective reducing agents, hence \(\mathrm{Fe}^{3+}\) remained unchanged.

This capacity to donate electrons and facilitate the reduction process, particularly in alkaline media, makes certain reagents highly valuable in redox reactions.

Oxidizing Agents

Oxidizing agents accept electrons from another substance. By accepting electrons, they become reduced while the substance donating the electrons becomes oxidized.
Consider them as electron "acceptors." In the exercise scenario, the role of the oxidizing agent is critical in understanding whether the \(\mathrm{Fe}^{3+}\) ions in solutions can be reduced or if the opposite happens, leading to oxidation instead.

• For instance, \(\mathrm{H}_2\mathrm{O}_2\) acts as an oxidizing agent in acidic conditions when combined with \(\mathrm{H}_2\mathrm{SO}_4\).
• This behavior shifts the balance depending on the surrounding environment.

Whether an agent acts as an oxidizer or a reducer greatly depends on the pH of the solution, and identifying this can help predict the result of complex chemical reactions.