Why does iron rust and how do we stop it?
Explain corrosion as an electrochemical process and evaluate methods of prevention.
The electrochemistry of rusting, the roles of oxygen and water, the effect of salt, and prevention methods including barriers, sacrificial anodes and cathodic protection, with worked TASC-style examples.
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What this dot point is asking
TASC expects you to explain rusting in electrochemical terms, including half-equations, and to evaluate the methods used to prevent it.
Corrosion as electrochemistry
Rusting is the oxidation of iron in the presence of water and oxygen. Different regions of the metal act as the electrodes of a tiny galvanic cell.
- At the anodic region, iron is oxidised: .
- At the cathodic region, oxygen is reduced in the presence of water: .
The iron(II) ions are then further oxidised by oxygen and combine with hydroxide and water to form hydrated iron(III) oxide, the flaky orange-brown solid we call rust.
Prevention methods
- Barrier methods keep water and oxygen away from the metal: painting, greasing, plastic coating and galvanising (a zinc layer) all form a physical barrier.
- Sacrificial protection uses a more reactive metal (zinc or magnesium) connected to the iron. The more reactive metal is preferentially oxidised, corroding instead of the iron. The block is a sacrificial anode and is replaced periodically.
- Cathodic protection forces the iron to be the cathode so it cannot be oxidised, either with a sacrificial anode or with an impressed current from an external supply.
Galvanising combines two effects: the zinc layer is a barrier, and because zinc is more reactive than iron, even a scratched layer still protects the exposed iron sacrificially.
Factors that speed corrosion
Several conditions accelerate the electrochemical process. Electrolytes such as dissolved salt or acid increase the conductivity of the surface water film, completing the circuit faster, which is why coastal and road-salted environments are so corrosive. Contact with a less reactive metal sets up a bimetallic (galvanic) couple in which the iron becomes the anode and corrodes faster, as in the copper-plated hull example. Mechanical stress and scratches expose fresh metal and concentrate anodic activity, and acidic rain or industrial pollutants lower the pH of the surface film, again speeding oxidation.
Stainless steel and passivation
Some metals resist corrosion by forming a thin, adherent oxide layer that seals the surface, a process called passivation. Aluminium forms a tough film, and stainless steel contains chromium that forms a protective chromium(III) oxide layer. Unlike rust, which is flaky and lets corrosion continue underneath, these oxide films are dense and self-repairing, so they protect the metal beneath. This contrasts with iron, whose oxide flakes away and exposes fresh metal to further attack.
In the exam, write the anode and cathode half-equations for rusting, state that both water and oxygen are required, note the accelerating effect of salt, and justify each prevention method by whether it forms a barrier or makes the iron a cathode.
Exam-style practice questions
Practice questions written in the style of TASC exam questions on this dot point, with worked answer explainers. The year tag is the paper they imitate, not the source.
TCE 20226 marksAn iron nail attaches a sign to a wooden post. (a) Explain how the nail corrodes over time, including relevant half-equations. (b) Explain how a zinc coating protects the iron even when the coating is scratched.Show worked answer →
(a) Where the iron contacts water and dissolved oxygen, iron is oxidised at anodic regions: . The electrons travel through the metal to cathodic regions where oxygen is reduced: . The and ions combine and are further oxidised by oxygen to hydrated iron(III) oxide (rust). Water acts as the electrolyte, so corrosion is fastest where the nail stays damp. (4 marks)
(b) Zinc is more easily oxidised than iron, so even if the coating is scratched the zinc acts as a sacrificial anode and is oxidised in preference; the iron becomes the cathode and is protected. This is sacrificial (galvanic) protection, the principle of galvanising. (2 marks)
TCE 20236 marksThree identical steel boats: Boat A on a freshwater beach at the water line, Boat B on a saltwater beach at the water line, Boat C underwater in a freshwater lake. (a) List them from fastest to slowest predicted corrosion and explain. (b) Explain how zinc attached to a hull would affect corrosion.Show worked answer →
(a) Fastest to slowest: B, then A, then C. Corrosion is electrochemical and needs water, oxygen and an electrolyte. Boat B (saltwater, water line) has the most dissolved ions (good electrolyte) plus ready oxygen, so it corrodes fastest. Boat A (freshwater, water line) has oxygen but a poorer electrolyte, so it is slower. Boat C is deep underwater where dissolved oxygen is very low, so the cathode reaction (oxygen reduction) is starved and corrosion is slowest. (4 marks)
(b) Zinc is more easily oxidised than iron, so it acts as a sacrificial anode: it is oxidised in preference to the steel, forcing the hull to behave as the cathode, protecting the iron even where the zinc is worn. (2 marks)
TCE 20213 marksAn 18th-century wooden hull was covered with copper plates attached by iron nails. Within a few years the copper detached because the iron nails were heavily corroded. Explain this, including appropriate half-equations.Show worked answer →
Iron and copper in contact in seawater form a galvanic (bimetallic) couple. Iron is the stronger reductant, so it becomes the anode and is oxidised, while copper acts as an inert cathode. Because the small iron nails are coupled to a large area of copper, the iron corrodes rapidly and the nails are eaten away, releasing the plates. (1 mark)
Anode (iron oxidised): . Cathode (oxygen reduced on the copper): . (2 marks)