Problem 53 - A Complete Resource Book in Chemistry for JEE Main

Question

Corrosion of iron is essentially an electro-chemical phenomenon, where the cell reactions are (a) \(\mathrm{Fe}\) is oxidized to \(\mathrm{Fe}^{2+}\) and dissolved oxygen in water is reduced to \(\mathrm{OH}\) (b) \(\mathrm{Fe}\) is oxidized to \(\mathrm{Fe}^{3+}\) and \(\mathrm{H}_{2} \mathrm{O}\) is reduced to \(\mathrm{O}_{2}^{2-}\) (c) \(\mathrm{Fe}\) is oxidized to \(\mathrm{Fe}^{2+}\) and \(\mathrm{H}_{2} \mathrm{O}\) is reduced to \(\mathrm{O}_{2}\) (d) \(\mathrm{Fe}\) is oxidized to \(\mathrm{Fe}^{2+}\) and \(\mathrm{H}_{2} \mathrm{O}\) is reduced to \(\mathrm{O}_{2}\)

Step-by-Step Solution

Answer

The correct answer is (a): \( \mathrm{Fe} \) is oxidized to \( \mathrm{Fe}^{2+} \) and dissolved oxygen is reduced to \( \mathrm{OH} \).

Step 1: Understanding Corrosion of Iron

Corrosion of iron involves an electrochemical process where iron reacts with water and oxygen, leading to its oxidation.

Step 2: Identify Oxidation Process

During corrosion, iron undergoes oxidation. The typical oxidation reaction for iron in a corrosion process is: \( \mathrm{Fe} \rightarrow \mathrm{Fe}^{2+} + 2\mathrm{e}^- \). This indicates that iron is oxidized to \( \mathrm{Fe}^{2+} \).

Step 3: Identify the Reduction Process

In the presence of water and air (oxygen), the dissolved oxygen is typically reduced to hydroxide ions \( \mathrm{OH}^- \). The general reduction reaction is: \( \mathrm{O}_2 + 2\mathrm{H}_2\mathrm{O} + 4\mathrm{e}^- \rightarrow 4\mathrm{OH}^- \).

Step 4: Matching the Reactions

From the given options, option (a) matches both the oxidation of \( \mathrm{Fe} \) to \( \mathrm{Fe}^{2+} \) and the reduction of dissolved oxygen to \( \mathrm{OH}^- \).

Step 5: Conclusion

The reactions that align with the described electrochemical process of iron corrosion involve iron's oxidation to \( \mathrm{Fe}^{2+} \) and the reduction of dissolved oxygen to \( \mathrm{OH}^- \).

Key Concepts

Electrochemical Process

Corrosion of iron is a fascinating example of an electrochemical process. In simple terms, an electrochemical process involves chemical changes accompanied by the transfer of electrons between two substances. This process occurs at the interface of an electrode and an electrolyte. Here, iron acts as the electrode, while water and oxygen function as electrolytes. As iron interacts with moisture and air, electrons are transferred from iron to oxygen, causing chemical reactions that lead to rust formation.
This electron transfer is key, as it causes iron to lose electrons – a process also known as oxidation – while oxygen gains electrons, which is called reduction. Understanding this exchange is crucial to comprehending how and why iron corrodes. The overall electrochemical process not only forms rust but also leads to structural damage and material degradation over time.

Oxidation Reaction

In the context of iron corrosion, oxidation refers to the chemical reaction where iron loses electrons. The process of oxidation is what transforms metallic iron (Fe) into its ionized form. Specifically, during iron oxidation, the reaction looks like this:

• The iron atom (\( Fe \)) lreleases two electrons to form the iron ion (\( Fe^{2+} \)).

Symbolically, it can be represented as:\( Fe \rightarrow Fe^{2+} + 2e^- \).
This reaction is critical in the entire corrosion process because it signifies the first step in deteriorating iron's integrity. The loss of electrons makes iron vulnerable, allowing it to interact and form compounds with other elements. Essentially, oxidation is the driving force that transforms strong, durable iron into brittle rust.

Reduction Reaction

Simultaneously, as iron undergoes oxidation, a reduction reaction takes place. Reduction involves the gain of electrons by another element. In the case of iron corrosion, oxygen in the presence of water gains electrons. This reaction can be expressed as:

• Dissolved oxygen (\( O_2 \)) gains electrons to form hydroxide ions (\( OH^- \)).

The balanced equation for this reduction reaction is:\( O_2 + 2H_2O + 4e^- \rightarrow 4OH^- \).
Reduction complements oxidation in the electrochemical process, forming a complete redox (reduction-oxidation) cycle. Without this, electrons released by iron would remain unaccepted and incomplete reactions would stall. Reduction is vital for the continuation and sustainability of the corrosion process.

Iron Oxidation

Iron oxidation is the core process that leads to the weakening of metal. During this chemical change, metallic iron turns into ferrous ions, which are more reactive and prone to further chemical reactions. Iron's transition from its metallic state to its oxidized form is central in understanding why iron decays in the presence of water and air.
As \( Fe \) is oxidized into \( Fe^{2+} \), it becomes destabilized and susceptible. This susceptibility is what typically causes structural failures in iron-based constructions over time. Iron's oxidation could lead to different compounds, including rust (\( Fe_2O_3 \)), which results from prolonged exposure to atmospheric conditions.

Chemical Reactions in Corrosion

Corrosion of iron is not just a single chemical reaction but a series of reactions occurring together. These reactions collectively contribute to the breakdown of iron. The major reactions include:

• electrochemical reactions, which cover both oxidation of iron and reduction of oxygen.
• chemical reactions that convert iron oxides and hydroxides into rust, the end product.

As iron oxidizes, it interacts with air and water to form hydrated iron(III) oxides (\( Fe_2O_3 \cdot H_2O \)). This iron rust is what you often see as the reddish-brown flakey layer on corroded metal.
Understanding these reactions helps in predicting the lifespan of iron structures and aids in developing strategies to prevent or slow down corrosion, such as coatings or galvanization.