Step 1: Understanding the Concept:
Galvanization is the process of applying a protective zinc coating to steel or iron to prevent rusting. This is a form of "sacrificial protection."
Step 2: Key Formula or Approach:
Compare the standard reduction potentials ($E^\circ$):
* $E^\circ (Fe^{2+}/Fe) = -0.44 \text{ V}$
* $E^\circ (Zn^{2+}/Zn) = -0.76 \text{ V}$
Step 3: Detailed Explanation:
Since Zinc has a more negative reduction potential than Iron, it is more "active." This means Zinc will oxidize (lose electrons) more readily than Iron. In the presence of moisture and air, a galvanic cell is formed where Zinc acts as the Anode and undergoes corrosion, while the Iron acts as the Cathode and remains protected. Even if the coating is scratched, the Zinc continues to protect the exposed Iron.
Step 4: Final Answer:
In galvanized iron, Zn acts as the anode and Fe acts as the cathode.