Iron nails are coated with zinc to prevent rusting. Even if the zinc layer is scratched, rusting does not occur because —

The correct option is Zinc acts as a sacrificial anode and corrodes instead of iron.

Here is a step-by-step explanation of why this happens:

1. Understanding Galvanization:
Coating iron nails with a layer of zinc is a process known as galvanization. The primary purpose of this coating is to protect the underlying iron from coming into contact with oxygen and moisture, which are the main requirements for rusting (oxidation of iron).

2. Activity Series of Metals:
Zinc is chemically more active (more reactive) than iron. In the reactivity series of metals, zinc lies above iron. This means that zinc has a higher tendency to lose electrons and undergo oxidation compared to iron. We can compare their standard reduction potentials ($E_{\mathrm{red}}°$):
For Zinc: ${\mathrm{Zn}}^{2+}+2e^{-}\to \mathrm{Zn},E°=-0.76V$
For Iron: ${\mathrm{Fe}}^{2+}+2e^{-}\to \mathrm{Fe},E°=-0.44V$
Since zinc has a more negative reduction potential, it is more easily oxidized than iron.

3. Sacrificial Protection:
Even if the zinc coating is scratched and the underlying iron is exposed to the atmosphere, the zinc and iron form a miniature electrochemical cell in the presence of moisture. Because zinc is more electropositive than iron, it behaves as the anode (undergoing oxidation) and sacrifices itself by corroding instead of the iron. The iron acts as the cathode and remains protected.

4. Electrochemical Reactions:
Zinc corrodes by losing electrons:
$\mathrm{Zn}\to {\mathrm{Zn}}^{2+}+2e^{-}$
The electrons released by the oxidation of zinc flow to the iron, preventing the iron from losing electrons and oxidizing to form rust (${\mathrm{Fe}}^{2+}$). Thus, zinc acts as a sacrificial anode.