Corrosion of iron and its prevention: WACE Year 12 Chemistry

Syllabus dot point

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

8 min readUpdated 2 Jun 2026

Quick answer

Rusting is an electrochemical process: part of the iron surface acts as an anode where iron is oxidised, and another part acts as a cathode where oxygen and water are reduced. Both water and oxygen are needed. Prevention works either by excluding oxygen and water (coatings) or by making the iron the cathode (sacrificial anodes, galvanising, cathodic protection).

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.

Corrosion

The deterioration of a metal by an electrochemical reaction with its environment. For iron the product is hydrated iron(III) oxide, $\text{Fe}_2\text{O}_3 \cdot x\text{H}_2\text{O}$ , which is rust.

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.

Both water and oxygen are required

Rusting needs the metal, water and oxygen all present. Remove any one and rusting stops, which is the basis of several prevention methods. Salt water and acids speed corrosion because they increase the conductivity of the electrolyte film.

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.

Why zinc protects even when scratched

Galvanising with zinc protects iron even if the coating is scratched, because zinc is more reactive and acts as a sacrificial anode. Tin plating, by contrast, accelerates rusting once scratched because tin is less reactive than iron, so the exposed iron becomes the anode.

Choosing a sacrificial metal

Question

A steel pipeline is to be protected by a sacrificial anode. Given the standard electrode potentials of iron ( $-0.44$ V), zinc ( $-0.76$ V) and copper ( $+0.34$ V), which metal should be used and why?

Solution

A sacrificial anode must be more easily oxidised than iron, that is, it must have a more negative electrode potential so that it corrodes in preference to the iron.

Zinc ( $-0.76$ V) is more negative than iron ( $-0.44$ V), so zinc will be oxidised preferentially and protect the iron. Copper ( $+0.34$ V) is more positive, so it would make the iron the anode and accelerate corrosion. The correct choice is zinc.

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.

Saying only water (or only oxygen) causes rust

A common error is claiming rusting needs only water or only oxygen; both are required, along with the iron. Another is suggesting any coating metal protects iron; only a more reactive metal (like zinc) protects when scratched, while a less reactive coating (like tin) worsens corrosion once breached. Always justify a sacrificial metal using the electrode potential series.

Exam-style questions

Questions in the style of SCSA exam questions on this dot point, each with a worked answer. They are written by ExamExplained unless tagged "Past paper"; the year shows the paper a question is modelled on.

WACE 2022-style6 marks

A 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 2020-style5 marks

Explain, 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.

Practise this

Sources & how we know this

Chemistry ATAR course: syllabus and support materials — SCSA
Chemistry: past ATAR course examinations — SCSA