Problem 62 - A Complete Resource Book in Chemistry for JEE Main

Question

In the reaction \(4 \mathrm{Fe}+3 \mathrm{O}_{2} \longrightarrow 4 \mathrm{Fe}^{3+}+6 \mathrm{O}_{2}^{2-}\) which of the following statements is incorrect? (a) metallic iron is reducing agent (b) \(\mathrm{Fe}^{3+}\) is an oxidizing agent (c) metallic iron is reduced to \(\mathrm{Fe}^{31}\) (d) redox reaction

Step-by-Step Solution

Answer

Statement (c), metallic iron is reduced to \(\mathrm{Fe}^{31}\), is incorrect.

Step 1: Identify the Oxidation and Reduction Processes

In the given reaction, identify which element is oxidized and which is reduced. Iron (Fe) starts with an oxidation state of 0 in its metallic form and changes to +3 when it becomes \(\mathrm{Fe}^{3+}\). Oxygen (\(\mathrm{O}_2\)) starts with an oxidation state of 0 and changes to -2 in \(\mathrm{O}_2^{2-}\). This indicates that Iron is oxidized and Oxygen is reduced.

Step 2: Define the Reducing and Oxidizing Agents

Determine which species act as reducing and oxidizing agents. A substance that donates electrons is the reducing agent, and one that gains electrons is the oxidizing agent. Since iron (Fe) loses electrons, it acts as the reducing agent. Conversely, since oxygen (\(\mathrm{O}_2\)) gains electrons, it acts as the oxidizing agent.

Step 3: Analyze Given Statements about the Reaction

Check each statement against the identified oxidation and reduction processes: (a) "Metallic iron is reducing agent" is true.(b) "\(\mathrm{Fe}^{3+}\) is an oxidizing agent" is false because iron was oxidized, hence cannot act as an oxidizing agent.(c) "Metallic iron is reduced to \(\mathrm{Fe}^{31}\)" is incorrect as iron was oxidized, not reduced.(d) "Redox reaction" is true as both oxidation and reduction processes occur.

Step 4: Determine the Incorrect Statement

According to the analysis, both statements (b) and (c) are incorrect. However, the incorrect statement specifically that Fe is reduced should be highlighted, i.e., option (c).

Key Concepts

Oxidation

When talking about redox reactions, oxidation is one of the essential processes at play. In simple terms, oxidation involves the loss of electrons from an atom or molecule. This process often results in an increase in the oxidation state of the element involved. For example, in the reaction we are examining, iron (Fe) starts with an oxidation state of 0 in its metallic form and changes to +3 when it becomes \( \text{Fe}^{3+} \).
This change happens because iron loses electrons during the reaction. When an atom loses electrons, it becomes more positively charged, hence the increase in the oxidation number. Remember, oxidation is often associated with the addition of oxygen, but at its core, it's about losing electrons.

• Iron goes from Fe to \(\text{Fe}^{3+}\)
• This involves a loss of three electrons per iron atom

Reduction

Reduction is the complementary process to oxidation in a redox reaction. It involves the gain of electrons by an atom or molecule, which leads to a decrease in its oxidation state. In the concept of redox reactions, reduction plays hand in hand with oxidation.
In the reaction example provided, oxygen \((\text{O}_2)\) undergoes reduction. It starts with an oxidation state of 0 and changes to -2 when it forms \(\text{O}^{2-}\). This means oxygen gains electrons, becoming more negatively charged.

• Oxygen goes from \(\text{O}_2\) to \(\text{O}^{2-}\)
• This involves a gain of two electrons per oxygen atom

Reduction can be remembered by the phrase "gaining electrons is reduced" because it results in a more negative oxidation state.

Oxidizing Agent

An oxidizing agent, or oxidant, is a substance that causes another substance to undergo oxidation; in doing so, it itself is reduced. In simpler terms, the oxidizing agent is the substance that gains electrons in the reaction.
In the iron and oxygen reaction, oxygen \((\text{O}_2)\) acts as the oxidizing agent. It accepts the electrons that iron \((\text{Fe})\) loses, thus getting reduced itself. The capability of accepting electrons makes the oxidizing agent crucial in facilitating the redox process. Some points to note are:

• Oxygen acts as the oxidizing agent
• The oxidizing agent undergoes reduction (gains electrons)

This role is essential as it drives the transfer of electrons that characterizes redox reactions.

Reducing Agent

The reducing agent, or reductant, is the opposite of the oxidizing agent. It is the substance that donates electrons in a redox reaction. By losing electrons, it causes another substance to be reduced.
For example, iron \((\text{Fe})\) in its metallic form acts as the reducing agent in our specific reaction. It gives up its electrons to oxygen \((\text{O}_2)\), allowing the latter to be reduced. By serving as an electron donor, the reducing agent plays a crucial role in enabling the oxido-reduction process.
Some key points about reducing agents include:

• Iron acts as the reducing agent by losing electrons
• A reducing agent undergoes oxidation (loses electrons)

Thus, the reducing agent and oxidizing agent work in tandem to facilitate the unique electron exchange that defines a redox reaction.