Problem 83 - Chemistry: The Central Science

Question

(a) Write the anode and cathode reactions that cause the corrosion of iron metal to aqueous iron(II). (b) Write the balanced half-reactions involved in the air oxidation of \(\mathrm{Fe}^{2+}(a q)\) to \(\mathrm{Fe}_{2} \mathrm{O}_{3} \cdot 3 \mathrm{H}_{2} \mathrm{O}\),

Step-by-Step Solution

Answer

(a) Anode half-reaction: \[\mathrm{Fe(s)} \rightarrow \mathrm{Fe}^{2+}(a q)+2 \mathrm{e}^{-}\] Cathode half-reaction: \[\mathrm{O}_{2}(g)+2\mathrm{H}_{2}\mathrm{O}(l)+4\mathrm{e}^{-}\rightarrow 4\mathrm{OH}^{-}(a q)\] (b) Oxidation half-reaction: \[4 \mathrm{Fe}^{2+}(a q)+\mathrm{O}_{2}(g) \rightarrow 4 \mathrm{Fe}^{3+}(a q)+4\mathrm{e}^{-}\] Precipitation half-reaction: \[6 \,\mathrm{Fe}^{3+}(a q) + 9\,\mathrm{H}_{2}\mathrm{O}(l) \rightarrow 2\,\mathrm{Fe}_{2} \mathrm{O}_{3}\cdot 3 \mathrm{H}_{2} \mathrm{O}(s)\]

Step 1: Identifying the Oxidation and Reduction Reactions in the Corrosion Process

The corrosion of iron in iron metal can be seen as a galvanic cell where the iron undergoes oxidation at the anode and the released electrons reduce oxygen in the presence of water at the cathode. This process leads to the formation of aqueous iron(II) ions and hydroxyl ions.

Step 2: Write the Anode and Cathode Reactions

For the anode reaction, iron will be oxidized to form iron(II) ions, releasing 2 electrons: Anode half-reaction: \[\mathrm{Fe(s)} \rightarrow \mathrm{Fe}^{2+}(a q)+2 \mathrm{e}^{-}\] For the cathode reaction, oxygen in the presence of water will be reduced using the released electrons from iron oxidation to form hydroxyl ions: Cathode half-reaction: \[\mathrm{O}_{2}(g)+2\mathrm{H}_{2}\mathrm{O}(l)+4\mathrm{e}^{-}\rightarrow 4\mathrm{OH}^{-}(a q)\] For the second part of the exercise:

Step 3: Identify the Half-reactions involved in the Air Oxidation Process

In the air oxidation process, the presence of oxygen will oxidize \(\mathrm{Fe}^{2+}(a q)\) to form \(\mathrm{Fe}^{3+}(a q)\). Then, \(\mathrm{Fe}^{3+}(a q)\) can precipitate with hydroxyl ions to eventually form \(\mathrm{Fe}_{2} \mathrm{O}_{3} \cdot 3 \mathrm{H}_{2} \mathrm{O}\).

Step 4: Write the Balanced Half-reactions for Air Oxidation

For the oxidation of \(\mathrm{Fe}^{2+}(a q)\) to \(\mathrm{Fe}^{3+}(a q)\) involving oxygen, the half-reaction can be written as: \[4 \mathrm{Fe}^{2+}(a q)+\mathrm{O}_{2}(g) \rightarrow 4 \mathrm{Fe}^{3+}(a q)+4\mathrm{e}^{-}\] For the reaction of \(\mathrm{Fe}^{3+}(a q)\) with water forming \(\mathrm{Fe}_{2} \mathrm{O}_{3} \cdot 3 \mathrm{H}_{2} \mathrm{O}\), the half-reaction is given by: \[6 \,\mathrm{Fe}^{3+}(a q) + 9\,\mathrm{H}_{2}\mathrm{O}(l) \rightarrow 2\,\mathrm{Fe}_{2} \mathrm{O}_{3}\cdot 3 \mathrm{H}_{2} \mathrm{O}(s)\] These are the balanced half-reactions for the air oxidation of \(\mathrm{Fe}^{2+}(a q)\) to \(\mathrm{Fe}_{2} \mathrm{O}_{3} \cdot 3 \mathrm{H}_{2} \mathrm{O}\).

Key Concepts

Anode Reaction

In the corrosion of iron, the anode reaction is a crucial part of the oxidation-reduction process. Here, iron serves as the anode, undergoing oxidation. This involves the loss of electrons, transforming solid iron (\( ext{Fe(s)} \)) into aqueous iron(II) ions (\( ext{Fe}^{2+}(aq) \)). The reaction is described by the anode half-reaction:

• anode reaction: \( ext{Fe(s)} \rightarrow ext{Fe}^{2+}(aq) + 2 ext{e}^- \)

This process not only generates electrons but also initiates the series of reactions that results in iron's degradation. Losing electrons causes the iron to dissolve into its ionic form, making it susceptible to further reaction with environmental substances. Understanding this reaction is key in addressing issues related to iron corrosion and developing protective measures.

Cathode Reaction

The cathode reaction complements the anode reaction in the corrosion process. After iron loses electrons at the anode, these electrons are utilized in the reduction of oxygen, which acts as the cathode in this system. With the presence of water, oxygen undergoes reduction, forming hydroxyl ions (\( ext{OH}^- \)) in an aqueous solution:

• cathode reaction: \( ext{O}_2(g) + 2 ext{H}_2 ext{O}(l) + 4 ext{e}^- \rightarrow 4 ext{OH}^- (aq) \)

This reaction is critical because it neutralizes the electrons produced at the anode. It also contributes to the formation of compounds that facilitate the corrosion of iron. Understanding the cathode reaction helps us comprehend how oxygen plays a role in the degradation of metals.

Oxidation-Reduction Reactions

Oxidation-reduction reactions, also known as redox reactions, involve the transfer of electrons between substances. In the corrosion of iron, a redox reaction occurs where iron undergoes oxidation, losing electrons, while oxygen undergoes reduction, gaining electrons. Each piece of the reaction is concurrent and indispensable:

• Oxidation occurs at the anode: electron loss and production of ions.
• Reduction occurs at the cathode: electron gain and involvement of oxygen.

This transfer of electrons between iron and oxygen effectively drives the corrosion process. By studying redox reactions, we can learn about the energetics and dynamics of such processes, which is fundamental for innovations in corrosion protection and management.

Galvanic Cell

A galvanic cell is an essential framework for understanding the process of iron corrosion. It depicts how chemical energy is converted into electrical energy through spontaneous redox reactions. In this setup, the anode and cathode reactions occur within the same system but at different locations:

• The anode, where oxidation occurs, releases electrons.
• The cathode, where reduction takes place, consumes electrons.

These reactions form a complete circuit allowing electron flow between them, much like a simple battery. In the context of corrosion, the surrounding environment facilitates this electron exchange, promoting metal degradation. Recognizing iron corrosion as a galvanic activity is instrumental in developing practical anti-corrosion technologies like sacrificial anodes or coatings.

Iron Oxidation

Iron oxidation is at the heart of the corrosion process and exemplifies the transformation of iron from its metallic form to an oxidized state. It is characterized by the loss of electrons from the iron atoms:

• Oxidized iron ions (\( ext{Fe}^{2+} \)) are produced.

This process starts at the anode and sets the stage for the complex formation of compounds like rust, which includes iron oxides and hydroxides. The trend in iron oxidation also explains why metals corrode faster under certain environmental conditions, like high humidity or salinity. It is a fundamental concept for understanding material durability and longevity in architectural and industrial contexts.