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WAChemistrySyllabus dot point

Why does iron rust, and how can we use electrochemistry to prevent it?

Explain the corrosion of iron as an electrochemical process and evaluate methods of corrosion prevention

A focused answer to the WACE Year 12 Chemistry dot point on corrosion, explaining rusting as an electrochemical process with anodic and cathodic regions, and evaluating prevention methods including sacrificial anodes and protective coatings, with a worked example and common exam mistakes.

Generated by Claude Opus 4.88 min answerUpdated 2026-06-02

Reviewed by: AI editorial process; not yet individually human-reviewed

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What this dot point is asking

The corrosion of iron, commonly called rusting, is a spontaneous redox reaction in which iron is oxidised by oxygen in the presence of water. It is electrochemical: tiny galvanic cells form across the metal surface.

The electrochemistry of rusting

On a wet iron surface, regions of slightly different potential set up an electrochemical cell. At the anodic region, iron is oxidised:

\text{Fe}(s) \rightarrow \text{Fe}^{2+}(aq) + 2\text{e}^-

The electrons travel through the metal to a cathodic region, where oxygen dissolved in water is reduced:

\text{O}_2(g) + 2\text{H}_2\text{O}(l) + 4\text{e}^- \rightarrow 4\text{OH}^-(aq)

The iron(II) ions are then further oxidised by oxygen and combine with hydroxide and water to form hydrated iron(III) oxide, the familiar reddish-brown rust.

Methods of prevention

Prevention strategies fall into two categories.

Barrier methods (exclude oxygen and water):

Painting, oiling or coating with plastic forms a physical barrier. Effective until the coating is scratched.
Galvanising coats the iron with zinc, which is both a barrier and a sacrificial metal (see below).

Electrochemical methods (make the iron the cathode):

Sacrificial anode: a more reactive metal (such as zinc or magnesium) is connected to the iron. Because it is more easily oxidised (more negative electrode potential), it corrodes preferentially and protects the iron, which becomes the cathode. Used on ship hulls and pipelines.
Cathodic protection: an external power supply forces electrons onto the structure, keeping it the cathode so it cannot oxidise.

The chemistry of rust itself

The iron(II) ions produced at the anode do not stop there. In the oxygen-rich surface film they are oxidised further to iron(III), $4\text{Fe}^{2+}(aq) + \text{O}_2(g) + (4+2x)\text{H}_2\text{O}(l) \rightarrow 2\text{Fe}_2\text{O}_3 \cdot x\text{H}_2\text{O}(s) + 8\text{H}^+(aq)$ . Because the anodic and cathodic regions can be separated on the surface, the rust often deposits away from the actual pit where the metal is being eaten, which is why rusting can undermine a structure from beneath an apparently sound surface. Unlike the tightly adherent oxide layer that protects aluminium, hydrated iron(III) oxide is flaky and porous, so it does not seal the surface and corrosion continues underneath.

Why this matters

Corrosion costs economies billions each year, so understanding and preventing it is a major application of electrochemistry. This dot point ties together oxidation, half-equations and the electrode potential series in a real-world context that the examination often frames as an extended-response question.

Exam-style practice questions

Practice questions written in the style of SCSA exam questions on this dot point, with worked answer explainers. The year tag is the paper they imitate, not the source.

WACE 20226 marksA steel structure is protected by sacrificial magnesium anodes. (a) Using the half-equations, explain why magnesium and not the steel corrodes. (b) Write the overall ionic equation for the corrosion of the magnesium when oxygen and water are reduced. (c) Standard electrode potentials are

\text{Mg}^{2+}/\text{Mg} = -2.37\ \text{V}

and

\text{Fe}^{2+}/\text{Fe} = -0.44\ \text{V}

. Calculate the cell potential for the magnesium-iron couple and state what it tells you.

Show worked answer →

A 6 mark question rewards the half-equation reasoning, the overall equation, and the calculated potential.

(a) Magnesium has the more negative electrode potential, so it is oxidised more readily than iron: $\text{Mg}(s) \rightarrow \text{Mg}^{2+}(aq) + 2\text{e}^-$ . The electrons it releases flow to the steel, making the steel the cathode where $\text{O}_2 + 2\text{H}_2\text{O} + 4\text{e}^- \rightarrow 4\text{OH}^-$ occurs instead of iron oxidation, so the iron is protected.

(b) Combining $2\text{Mg} \rightarrow 2\text{Mg}^{2+} + 4\text{e}^-$ with $\text{O}_2 + 2\text{H}_2\text{O} + 4\text{e}^- \rightarrow 4\text{OH}^-$ :

2\text{Mg}(s) + \text{O}_2(g) + 2\text{H}_2\text{O}(l) \rightarrow 2\text{Mg}^{2+}(aq) + 4\text{OH}^-(aq).

(c) Treating magnesium as the anode (oxidation) and iron as the cathode:

E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}} = (-0.44) - (-2.37) = +1.93\ \text{V}.

The positive value confirms the reaction is spontaneous, so magnesium will preferentially corrode and protect the iron.

Markers reward the oxidation half-equation for $\text{Mg}$ , the overall equation, and $E^\circ_{\text{cell}} = +1.93\ \text{V}$ with a spontaneity statement.

WACE 20205 marksExplain, with reference to the electrochemistry of rusting, why (i) iron rusts faster in coastal salt-laden air than inland, and (ii) a tin coating accelerates rusting once it is scratched whereas a zinc coating does not.

Show worked answer →

A 5 mark answer needs the electrolyte effect plus the reactivity comparison for both coatings.

(i) Rusting is electrochemical and requires ions to carry charge through the surface water film. Dissolved salt (sodium chloride) greatly increases the conductivity of that electrolyte film, so the tiny corrosion cells operate faster and more iron is oxidised per unit time. Inland the film is closer to pure water, with low conductivity, so rusting is slower.

(ii) Tin is less reactive than iron (less easily oxidised). While the tin coat is intact it acts as a barrier, but once scratched the exposed iron becomes the anode (it is oxidised in preference to tin), so rusting is accelerated at the scratch. Zinc is more reactive than iron, so when a zinc coating is scratched the zinc becomes the anode and is sacrificially oxidised, keeping the iron as the protected cathode.

Markers reward the conductivity/electrolyte reasoning, and the anode-assignment for each metal based on relative reactivity.