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 \(\mathrm{Fe}^{3+}\) (d) redox reaction

Step-by-Step Solution

Answer

Statement (c) is incorrect: metallic iron is oxidized, not reduced.

Step 1: Identify Oxidation and Reduction

Oxidation involves the loss of electrons, while reduction involves the gain of electrons. In this equation, the oxidation number of iron (798{Fe}) increases from 0 to +3, indicating that iron is oxidized. Oxygen (798{O_2}), on the other hand, decreases from 0 to -2, indicating that oxygen is reduced.

Step 2: Define Reducing and Oxidizing Agents

The reducing agent donates electrons, while the oxidizing agent accepts electrons. In our reaction, metallic iron (798{Fe}) loses electrons and is thus the reducing agent. As oxygen is reduced by gaining electrons, it acts as the oxidizing agent. Hence, 798{Fe^{3+}} is a product of oxidation but not an oxidizing agent.

Step 3: Evaluate the Statements for Accuracy

(a) Metallic iron is correctly identified as the reducing agent since it loses electrons. (b) 798{Fe^{3+}} cannot be an oxidizing agent because it is the result of oxidation, not the cause of reduction. (c) This statement is incorrect because metallic iron is oxidized, not reduced. (d) The reaction is indeed a redox reaction, as it involves both oxidation and reduction.

Key Concepts

Oxidation in Redox Reactions

In the realm of chemistry, oxidation is a core concept within redox reactions, where it involves the loss of electrons. This results in an increase in oxidation state. Consider the equation from the exercise, \(4 \text{Fe} + 3 \text{O}_{2} \rightarrow 4 \text{Fe}^{3+} + 6 \text{O}_{2}^{2-}\). Here, metallic iron \(\text{Fe}\) starts with an oxidation state of 0, and when it loses electrons, this oxidation state increases to +3. This process of losing electrons is what defines it as being oxidized.

Oxidation does not occur in isolation, as it is always coupled with another process: reduction. Together, they form the basis of any redox (reduction-oxidation) reaction. The term "oxidation" might bring to mind interactions with oxygen, but it applies to any process where electron loss occurs, fundamentally affecting the chemical properties of the element involved.

In real-world applications, oxidation is everywhere, from the rusting of iron to the biochemical processes in our own bodies. It's important to recognize how the reduction in one substance complements the oxidation in another, facilitating energy transfer and chemical changes that are pivotal in both industrial and natural contexts.

Reduction in Redox Reactions

Reduction is the counterpart to oxidation in a redox reaction, involving the gain of electrons and a decrease in oxidation state. In our example reaction, \(\text{O}_2\), initially with an oxidation state of 0, gains electrons, leading to an oxidation state of -2. This gain of electrons is what we define as reduction.

Just as with oxidation, reduction cannot happen in isolation. It is part of a redox reaction, balancing the changes occurring in both the oxidized and reduced elements. Reduction is crucial in many processes. For instance, photosynthesis is a large-scale natural event where carbon dioxide is reduced to glucose, facilitating life on Earth. In the context of this reaction, it helps in driving the transformation process and plays a vital role in maintaining the equilibrium of chemical reactions.

Understanding reduction can often be simplified by remembering the mnemonic "OIL RIG" - Oxidation Is Loss, Reduction Is Gain - referring to electrons. This simple concept helps demystify what is happening at the atomic level, making these reactions easier to analyze and understand.

Understanding Oxidizing and Reducing Agents

Oxidizing and reducing agents are key players in a redox reaction. The reducing agent is the substance that donates electrons, causing oxidation, while the oxidizing agent is the one that gains these electrons, leading to reduction.

In our example, metallic iron \(\text{Fe}\) serves as the reducing agent as it loses electrons in the reaction, subsequently being oxidized itself. Conversely, \(\text{O}_2\) acts as the oxidizing agent as it accepts electrons and is reduced. It's important to note that in terms of agents, the substance which gets oxidized is always the reducing agent, while the one that gets reduced is the oxidizing agent.

This aspect of chemistry is fundamental, as it helps in predicting the direction of reactions and understanding energy changes. Different substances have varying tendencies to act as oxidizing or reducing agents, largely influenced by their electron affinities and ionization energies. This concept is not only central to chemical reactions but also to technological applications like batteries and fuel cells, where the transfer of electrons is utilized for power generation.