Question

Zinc is attached to a ship's steel propeller to prevent the steel from rusting. Write balanced equations for the corrosion reactions that occur (a) in the presence of \(\mathrm{Zn}\) and \((\mathrm{b})\) in the absence of \(\mathrm{Zn}\).

Short answer

Zinc prevents steel corrosion by oxidizing first. Without zinc, steel oxidizes forming rust.

Step-by-step solution

Understanding the Galvanic Protection

Zinc is used to prevent steel from rusting through a process known as galvanic protection. In this arrangement, zinc acts as a sacrificial anode, corroding instead of the steel.

Writing the Reaction Equation in the Presence of Zinc

Steel (mainly composed of iron) does not corrode because zinc undergoes oxidation first due to its higher reactivity in the galvanic series. The balanced equation for the reaction in the presence of zinc is:\[\mathrm{Zn} \to \mathrm{Zn}^{2+} + 2\mathrm{e}^-\]In this reaction, zinc oxidizes faster than iron, thus protecting steel by sacrificing itself.

Corrosion Reaction in the Absence of Zinc

In the absence of zinc, iron in the steel propeller is exposed to the environment and begins to corrode. The balanced equation for this reaction is:\[\mathrm{Fe} \to \mathrm{Fe}^{2+} + 2\mathrm{e}^-\]These electrons go on to reduce oxygen in the presence of water:\[\mathrm{O}_2 + 2\mathrm{H}_2\mathrm{O} + 4\mathrm{e}^- \to 4\mathrm{OH}^-\]

Complete Overall Equations

Combine the oxidation of iron and the reduction of oxygen to form the rust formation reaction:\[2\mathrm{Fe}^{2+} + \mathrm{O}_2 + 2\mathrm{H}_2\mathrm{O} \to 2\mathrm{Fe(OH)}_2\]\[\mathrm{Fe(OH)}_2 + \frac{1}{2}\mathrm{O}_2 + \mathrm{H}_2\mathrm{O} \to \mathrm{Fe(OH)}_3\]Over time, \(\mathrm{Fe(OH)}_3\) precipitates to form rust, \(\mathrm{Fe}_2\mathrm{O}_3\cdot \mathrm{nH}_2\mathrm{O}\).

Key concepts

Sacrificial Anode

The sacrificial anode is a crucial concept in galvanic protection, where a more easily oxidized metal is intentionally used to protect another metal from corrosion. Here, zinc serves as a great example of a sacrificial anode to protect iron or steel. In simple terms, zinc "sacrifices" itself by corroding faster than steel.
This is because zinc is higher in the electrochemical series, meaning it has a greater tendency to lose electrons and oxidize.

• This property makes zinc an effective protector for steel surfaces, such as ship propellers.
• When zinc oxidizes, it forms positively charged ions and releases electrons.
• These released electrons are then used to protect the steel from corroding.

The concept is widely applied in industries to prolong the life of metal structures.

Corrosion Reactions

Corrosion reactions describe the chemical process through which metals deteriorate due to environmental interactions. When metal comes into contact with moisture, air, and other corrosive elements, its structure is compromised as electrons are transferred from the metal to the environment.
This process often involves redox reactions:

• Oxidation - the metal loses electrons.
• Reduction - another substance gains those electrons.

Zinc, serving as a sacrificial anode, undergoes oxidation as shown in the equation: \[\mathrm{Zn} \to \mathrm{Zn}^{2+} + 2\mathrm{e}^- \]This reaction helps prevent the corrosion of iron in the presence of zinc, as electrons from oxidized zinc are used to neutralize potential corrosion interactants like oxygen.

Iron Oxidation

Iron oxidation is a part of the corrosion process and is the reason why steel structures rust over time. When iron is exposed to moist air without the protection of a sacrificial anode like zinc, it loses electrons and forms iron ions:\[\mathrm{Fe} \to \mathrm{Fe}^{2+} + 2\mathrm{e}^- \]This transformation is the initial step in rust formation.
These iron ions can further react with oxygen and water to create additional iron compounds.

• Iron oxides and hydroxides form as by-products.
• These substances further degrade the metal structure.

To reduce this deterioration, galvanic protection is implemented through sacrificial anodes, thereby saving the iron from these damaging reactions.

Zinc Corrosion

Zinc corrosion is intentionally utilized in applications where it acts as a sacrificial anode to protect iron and steel. This corrosion is actually beneficial, as it prevents the steel it is attached to from rusting.
The equations demonstrate this principle in a straightforward way. In the presence of zinc, zinc itself oxidizes much more readily than iron:

• Zinc sacrifices its own electrons, becoming positively charged zinc ions in the process.
• The produced electrons are used to counteract potential oxidizing agents, like oxygen.

In contrast, without zinc, the electrons from iron are used to complete these same reactions, leading to the formation of rust and ultimately the degradation of steel when moisture and oxygen are present.