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