A zinc-iron cell
Now pair iron with a metal above it: zinc and iron in dilute sulphuric acid. The same iron that was the anode a moment ago now becomes the cathode.
Figure to be added
Zinc is above iron in the activity series, so the more reactive zinc is oxidised and acts as the anode. Zinc atoms dissolve leaving electrons on it, so zinc is the negative terminal.
Zn (s) → Zn2+ (aq) + 2e
2H+ (aq) + 2e → H2 (g)
Zn (s) + 2H+ (aq) → Zn2+ (aq) + H2 (g)
(6) at the zinc anode; (7) at the iron cathode; (6) + (7) the cell reaction
Reduction happens at the iron, so iron is the cathode. Electrons flow to it along the wire, so it is the positive terminal. While the cell works, the zinc wears away and gas bubbles come off the iron — and the iron itself does not dissolve.
Why this matters for rust
Compare the two cells. Paired with copper, iron dissolves. Paired with zinc, iron is protected — the zinc dissolves instead, and the iron only has hydrogen bubbling off it. Whether iron is eaten away depends on its partner. This is exactly the idea behind galvanising iron with zinc, which you will meet at the end of the chapter: the zinc is sacrificed so that the iron survives.
The same metal can be either electrode
Iron is the anode against copper and the cathode against zinc. What decides is which metal of the pair is higher in the activity series.