Problem 100 - A Complete Resource Book in Chemistry for JEE Main

Question

Corrosion of iron is essentially an electrochemical 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

Option (a) is correct: Fe oxidized to Fe^{2+} and O_2 reduced to OH^{-}.

Step 1: Understand the Corrosion Process

Corrosion of iron is an electrochemical process involving the oxidation of iron and the reduction of oxygen in water. It typically occurs in the presence of moisture and oxygen.

Step 2: Analyze the Oxidation Reaction

In corrosion, iron (Fe) is oxidized to ferrous ions (Fe^{2+}). This can be represented as the reaction: \[ Fe \rightarrow Fe^{2+} + 2e^{-} \] This is characteristic of iron corrosion and relates to the oxidation reaction.

Step 3: Identify the Reduction Reaction

In atmospheric conditions, the dissolved oxygen in water is typically reduced to hydroxide ions (OH^{-}). This is shown as: \[ 0.5 O_2 + H_2O + 2e^{-} \rightarrow 2 OH^{-} \] This is a standard reduction process in aqueous environments during iron corrosion.

Step 4: Match Reactions to Options

The correct option lists iron being oxidized to Fe^{2+} and oxygen in water being reduced to OH^{-}. A review of the given choices reveals that option (a) matches these criteria, as it involves the oxidation of iron to Fe^{2+} and the reduction of oxygen to OH^{-}.

Key Concepts

Electrochemical Reactions

Electrochemical reactions are chemical processes in which electrons are transferred between chemical species. This transfer leads to the creation of electric current. In the case of metal corrosion, such as iron, these reactions are pivotal. They involve a change in oxidation states of the elements involved. Specifically, iron loses electrons during the corrosion process, which is a hallmark of oxidation. Meanwhile, substances like oxygen gain these electrons, characterizing reduction. Understanding this transfer of electrons is key to grasping why corrosion can both cause damage and create electrical energy.

Oxidation of Iron

The oxidation of iron is a fundamental part of its corrosion. When iron oxidizes, it loses electrons and transforms from iron metal into iron ions, often forming ferrous ions \(Fe^{2+}\). This can be expressed in a simple chemical equation:

• \(Fe \rightarrow Fe^{2+} + 2e^{-}\)

Iron is naturally prone to oxidation when exposed to air and moisture because of its reactivity with oxygen. As the iron atom gives up electrons, it becomes positively charged, contributing to the overall electrochemical cell reaction. This phenomenon explains why metals like iron rust over time, especially in wet environments.

Reduction of Oxygen

Reduction of oxygen is the counterpart of iron oxidation in the corrosion process. In the presence of water, oxygen undergoes reduction. This means it gains electrons freed up by oxidizing iron. The typical reduction reaction in corrosion is:

• \(0.5 \, O_2 + H_2O + 2e^{-} \rightarrow 2 \, OH^{-}\)

This reaction occurs at the cathode of the electrochemical cell formed during corrosion. Oxygen, found dissolved in the water, turns into hydroxide ions \(OH^{-}\). These ions can further react with iron ions, forming rust. The reduction process is essential for completing the electrochemical circuit and making the corrosion process continue.

Iron Chemistry

In iron chemistry, iron can occur in multiple oxidation states, though \(Fe^{2+}\) and \(Fe^{3+}\) are the most common in corrosion contexts. Iron's tendency to convert to these ions comes from its electron configuration, making it reactive, especially in the presence of environmental oxidizing agents. As iron transitions to \(Fe^{2+}\), it loses two electrons. Further oxidation could lead to the conversion to \(Fe^{3+}\). The chemistry of iron is complex and involves not only these ionic transformations but also the interactions with water and air that contribute to rust formation. Understanding iron's chemical properties is critical to mastering materials science, especially in contexts of metal deterioration.

Electrochemical Corrosion Process

The electrochemical corrosion process is a sequence of interconnected reactions that lead to metal degradation. This involves both oxidation and reduction reactions. Initiated when iron interacts with oxygen and moisture, a galvanic cell is effectively created.

• The anode, where oxidation occurs, is the site on the iron surface converting iron to \(Fe^{2+}\).
• The cathode is where dissolved oxygen reduces to hydroxide ions \(OH^{-}\).
• These hydroxide ions then react with iron ions to form iron compounds, such as rust, further depleting the metal.

This process is insidious over time, as it gradually weakens metal structures. The corrosive process is markedly influenced by environmental factors like humidity and temperature. Preventative measures, such as coatings and cathodic protection, are vital in mitigating the damaging impacts of electrochemical corrosion.