Step 1: Understanding the Question:
The question asks about the electrochemical roles of the metals involved in galvanized iron.
Galvanization is the process of applying a protective zinc coating to steel or iron to prevent rusting.
Step 2: Detailed Explanation:
The Process of Galvanization:
In this process, iron (Fe) is coated with a layer of zinc (Zn). Zinc is used because it is more reactive than iron.
Electrochemical Series:
If we look at the standard reduction potentials, \(E^\circ_{Zn^{2+}/Zn} = -0.76\) V and \(E^\circ_{Fe^{2+}/Fe} = -0.44\) V.
Since Zinc has a more negative reduction potential, it has a greater tendency to undergo oxidation compared to iron.
Anodic and Cathodic Roles:
When the galvanized layer is intact, Zinc acts as a barrier. If the coating is scratched and both metals are exposed to moisture/air, a galvanic cell is formed.
Zinc, being more active, becomes the Anode and undergoes oxidation (corrosion): \(Zn \rightarrow Zn^{2+} + 2e^-\).
Iron becomes the Cathode. The electrons released by Zinc flow to the Iron, where they are used in the reduction of oxygen and water. This prevents the Iron from losing electrons (oxidizing), thereby protecting it from rust.
This is known as sacrificial protection or cathodic protection because the Zinc "sacrifices" itself to protect the Iron.
Tinning vs Galvanization:
Tinning involves coating iron with Tin (Sn). Since Tin is less active than iron, if the coating is scratched, iron becomes the anode and corrodes rapidly. Thus, galvanization is superior for protecting iron.
Step 3: Final Answer:
In galvanized iron, Zinc (\(Zn\)) acts as the anode and Iron (\(Fe\)) acts as the cathode, providing sacrificial protection to the iron.