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, and prevention methods including barriers, sacrificial anodes and cathodic protection.
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What this dot point is asking
You must explain rusting in electrochemical terms and evaluate methods of preventing it.
Corrosion as electrochemistry
Rusting is the oxidation of iron in the presence of water and oxygen. Different regions of the metal act as 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 (zinc coating) all form a physical barrier.
Sacrificial protection uses a more reactive metal connected to the iron. The more reactive metal, such as zinc or magnesium, is preferentially oxidised, so it corrodes instead of the iron. The block of reactive metal is called a sacrificial anode and is replaced periodically.
Cathodic protection forces the iron to be the cathode of a cell so it cannot be oxidised. This can be done with a sacrificial anode or by 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 if the layer is scratched the exposed iron is still protected sacrificially.
In the exam, write the anode and cathode half-equations for rusting, state that both water and oxygen are required, 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.
2022 TASC6 marksAn iron nail was used to attach a sign to a wooden fence post. Explain how the nail corrodes over time, including relevant half-equation(s). Then explain how a zinc coating protects an iron object from corrosion, even if the zinc coating is damaged.Show worked answer →
Corrosion of the nail (4 marks). Where the iron is in contact with water and dissolved oxygen, iron is oxidised at the anodic regions: Fe(s) -> Fe2+(aq) + 2e-. The electrons travel through the metal to cathodic regions where oxygen is reduced: O2(g) + 2H2O(l) + 4e- -> 4OH-(aq). The Fe2+ ions and OH- ions then combine and are further oxidised by oxygen to form hydrated iron(III) oxide (rust). Water acts as the electrolyte, so corrosion is fastest where the nail stays damp.
Zinc coating (2 marks). Zinc is a stronger reducer than iron (more easily oxidised), so even if the coating is scratched the zinc acts as a sacrificial anode and is oxidised in preference to the iron. The iron becomes the cathode and is protected. This is sacrificial (galvanic) protection, the principle behind galvanising.
2023 TASC6 marksThree identical steel boats washed up: Boat A on a freshwater beach at the water line, Boat B on a saltwater beach at the water line, and Boat C 100 m underwater in a freshwater lake. Assuming the steel hull is exposed in all three cases, list the boats in order of predicted rate of corrosion from fastest to slowest and explain your reasoning. Then explain how zinc attached to a hull would affect the corrosion of the boat.Show worked answer →
Order and reasoning (4 marks). Fastest to slowest: B, then A, then C. Corrosion is an electrochemical process needing water, oxygen and an electrolyte. Boat B (saltwater, water line) has the most dissolved ions (salt is a good electrolyte) plus ready access to atmospheric oxygen, so it corrodes fastest. Boat A (freshwater, water line) also has oxygen but fresh water is a poorer electrolyte, so it is slower than B. Boat C is deep underwater where dissolved oxygen is very low, so the cathode reaction (oxygen reduction) is starved and corrosion is slowest.
Zinc (2 marks). Zinc is more easily oxidised than iron, so it acts as a sacrificial anode. The zinc is oxidised in preference to the steel, forcing the hull to behave as the cathode, so the iron is protected even where the zinc is worn or damaged.
2021 TASC3 marksIn the 18th century the wooden hull of a ship was covered with copper plates attached with iron nails. A few years later the copper had become detached because the iron nails had been heavily corroded. Explain this observation, including appropriate half-equations.Show worked answer →
Iron and copper in contact in seawater form a galvanic (bimetallic) couple. Iron is the stronger reducer, so it becomes the anode and is oxidised, while the 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.
Anode (iron oxidised): Fe(s) -> Fe2+(aq) + 2e-.
Cathode (oxygen reduced on the copper): O2(g) + 2H2O(l) + 4e- -> 4OH-(aq). (3 marks for identifying the galvanic coupling and both half-equations.)