Upper secondary · 15 min

Why iron rusts: a corrosion cell on a nail

Compare bare iron, galvanised iron and painted iron in humid air to see the electrochemistry of rust.

Goal

Explain rusting as a short-circuited electrochemical cell and deduce two protection strategies.

Apparatus and reagents

A schematic rig: an iron bar under a film of humid air, with a mode switch for bare, zinc-coated or painted surfaces.

Procedure

Open the bare-iron mode: identify the anode region (Fe oxidised) and the cathode region (O₂ reduced).
Follow the arrows from the anode and cathode regions to the rust patch.
Switch to the zinc-coated mode: which metal corrodes now? Then to the painted mode: what is blocked?

What to observe

On bare iron, $\ce{Fe^{2+}}$ released at anodic spots migrates and meets $\ce{OH^-}$ formed at cathodic spots, ending as orange-brown $\ce{Fe2O3.xH2O}$ .
With a zinc coating the zinc corrodes instead of iron; under an intact paint film nothing happens because $\ce{O2}$ and water cannot reach the metal.

Explanation

Rusting needs both $\ce{O2}$ and $\ce{H2O}$ on the metal surface. A water film acts as electrolyte: at anodic spots $\ce{Fe -> Fe^{2+} + 2e^-}$ ; at cathodic spots $\ce{O2 + 2H2O + 4e^- -> 4OH^-}$ . The ions combine and oxidise further into hydrated iron(III) oxide — rust. Zinc protects iron sacrificially because it is more active ( $E°_{\ce{Zn^{2+}/Zn}} = -0{,}76$ V vs $-0{,}44$ V for iron), while paint simply isolates the metal from the electrolyte.

Chemists behind it

Michael Faraday Walther Nernst