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 reaction: \(\mathrm{Fe}(s) \rightarrow \mathrm{Fe}^{2+}(aq) + 2e^{-}\) and cathode reaction: \( \frac{1}{2} \mathrm{O}_{2}(g) + \mathrm{H}_{2}\mathrm{O}(l) + 2e^{-} \rightarrow 2\mathrm{OH}^{-}(aq)\) (b) Balanced overall air oxidation reaction: \(4 \mathrm{Fe}^{2+}(aq) + \mathrm{O}_{2}(g) + 4\mathrm{H}^{+}(aq) \rightarrow 4 \mathrm{Fe}^{3+}(aq) + 2\mathrm{H}_{2}\mathrm{O}(l)\) and subsequent reaction: \( 4 \mathrm{Fe}^{3+}(aq) + 12\mathrm{OH}^{-}(aq) \rightarrow 2 \mathrm{Fe}_{2} \mathrm{O}_{3}(s) \cdot 3 \mathrm{H}_{2} \mathrm{O}(l)\)

Step 1: Anode Reaction

In the anode reaction, oxidation occurs. Iron loses two electrons to form iron(II) ions: \[\mathrm{Fe}(s) \rightarrow \mathrm{Fe}^{2+}(aq) + 2e^{-}\]

Step 2: Cathode Reaction

In the cathode reaction, reduction occurs. Some other species, typically present in the aqueous environment, will gain electrons. In case of corrosion of iron, oxygen dissolved in water is usually the electron acceptor; forming hydroxide ions: \( \frac{1}{2} \mathrm{O}_{2}(g) + \mathrm{H}_{2}\mathrm{O}(l) + 2e^{-} \rightarrow 2\mathrm{OH}^{-}(aq)\) (b)

Step 3: Reaction of Iron(II) with Oxygen

In air oxidation, the reaction of \(\mathrm{Fe}^{2+}(aq)\) with oxygen in the presence of water forms iron(III) hydroxide, which further dehydrates to form iron(III) oxide: \[ 4 \mathrm{Fe}^{2+}(aq) + \mathrm{O}_{2}(g) + 6\mathrm{H}_{2}\mathrm{O}(l) \rightarrow 4 \mathrm{Fe(OH)}_{3}(s) \]

Step 4: Dehydration of Iron(III) Hydroxide to Iron(III) Oxide

Iron(III) hydroxide dehydrates to form iron(III) oxide and water: \[ 2 \mathrm{Fe(OH)}_{3}(s) \rightarrow \mathrm{Fe}_{2} \mathrm{O}_{3}(s) + 3\mathrm{H}_{2}\mathrm{O}(l)\] To write the balanced half-reactions, we can split the air oxidation reaction into two parts - the oxidation of \(\mathrm{Fe}^{2+}\) ions and the reduction of oxygen.

Step 5: Oxidation Half-Reaction

In the oxidation half-reaction, the iron(II) ions lose their electrons to form iron(III) ions: \[2 \mathrm{Fe}^{2+}(aq) \rightarrow 2 \mathrm{Fe}^{3+}(aq) + 2e^{-}\]

Step 6: Reduction Half-Reaction

In the reduction half-reaction, oxygen gains electrons to form water: \[\mathrm{O}_{2}(g) + 4e^{-} + 4\mathrm{H}^{+}(aq) \rightarrow 2\mathrm{H}_{2}\mathrm{O}(l)\] Combine these half-reactions, ensuring that the total number of electrons lost and gained is the same. Solution for part (b):

Step 7: Balanced Overall Air Oxidation Reaction

By combining the oxidation and reduction half-reactions (multiplied as necessary to have equal numbers of electrons), we obtain the balanced overall air oxidation reaction: \[4 \mathrm{Fe}^{2+}(aq) + \mathrm{O}_{2}(g) + 4\mathrm{H}^{+}(aq) \rightarrow 4 \mathrm{Fe}^{3+}(aq) + 2\mathrm{H}_{2}\mathrm{O}(l)\] The subsequent step, where iron(III) ions react with water to form iron(III) hydroxide and finally dehydrate to \(\mathrm{Fe}_{2} \mathrm{O}_{3} \cdot 3 \mathrm{H}_{2} \mathrm{O}\), can be represented as: \[ 4 \mathrm{Fe}^{3+}(aq) + 12\mathrm{OH}^{-}(aq) \rightarrow 2 \mathrm{Fe}_{2} \mathrm{O}_{3}(s) \cdot 3 \mathrm{H}_{2} \mathrm{O}(l)\]

Key Concepts

Corrosion

Corrosion is a chemical process that breaks down metals through interactions with environmental elements, weakening them over time. It occurs when metal atoms lose electrons, a process known as oxidation. For iron, corrosion happens primarily when it comes into contact with water and oxygen, forming rust. This process is critical in electrochemistry because it illustrates how oxidation-reduction reactions can transform materials.

When iron is exposed to moisture, it starts losing electrons, turning into iron(II) ions. These electrons are then accepted by another species, typically oxygen dissolved in the water, forming hydroxide ions. This interaction is an electrochemical reaction, emphasizing the importance of understanding both oxidation at the anode (iron losing electrons) and reduction at the cathode (oxygen gaining electrons).

Corrosion is not just a surface issue; it affects the integrity of structures, leading to potential failures if unchecked. Prevention strategies often involve controlling the exposure of iron to these corrosive elements or using coatings or treatments that prevent electron transfer.

Oxidation-Reduction Reactions

Oxidation-reduction reactions, often called redox reactions, are a type of chemical reaction where electrons are transferred between two substances. In electrochemistry, these reactions are crucial as they describe how elements change oxidation states, signifying either a gain or loss of electrons.

In the context of corrosion and iron oxidation, this concept helps explain the transformation phases. A redox reaction involves two key parts: the oxidation half-reaction and the reduction half-reaction. During oxidation, a substance loses electrons, while in reduction, a different substance gains them.

For iron corrosion, the oxidation half-reaction involves iron atoms losing electrons to form iron(II) ions. Meanwhile, in the typical reduction half, the oxygen gains these electrons and forms hydroxide ions. By balancing these half-reactions, you get an understanding of the overall process, which is essential in practical applications like preventing corrosion.

• Oxidation: Iron loses electrons.
• Reduction: Oxygen gains electrons.

Understanding the delicate balance of these electron exchanges is vital for industries relying on iron and similar metals, influencing material sciences and engineering.

Iron(II) to Iron(III) Transformation

The transformation of iron(II) to iron(III) is a fascinating step in the corrosion process. Initially, when iron corrodes, iron(II) ions are formed from the oxidation of iron metal. These ions can undergo further chemical changes in the presence of oxygen.

The next stage involves the reaction of iron(II) ions with oxygen, leading to the formation of iron(III) hydroxide. This compound eventually dehydrates to form iron(III) oxide, more commonly known as rust. This transformation is a clear example of a redox process in action, aligning with electrochemistry principles.

In essence, this is about controlling the electron flow to transition iron from a +2 oxidation state in iron(II) to a +3 state in iron(III). This change explains why rust appears and the way it behaves as a brittle material. Such changes are significant because they dictate how different structures built from iron require maintenance over time.

• Initial oxidation: Formation of iron(II) ions.
• Further oxidation: Transition to iron(III) hydroxide and then iron(III) oxide.

Considering these transformations gives insight into how materials degrade over time and the chemistry involved in prolonging their lifespan.