The standard reduction potential of zinc ($E^\circ_{\text{Zn}^{2+}/\text{Zn}} = -0.76 \text{ V}$) is more negative than that of iron ($E^\circ_{\text{Fe}^{2+}/\text{Fe}} = -0.44 \text{ V}$). This means zinc has a higher negative electrode potential than iron, making it a stronger reducing agent. When zinc is coated on iron (galvanisation), zinc acts as a sacrificial anode and undergoes oxidation preferentially, protecting the iron from corrosion. The reverse is not possible because iron, being a weaker reducing agent with a lower negative electrode potential, cannot protect zinc.