Rusting is the slow electrochemical corrosion of iron that needs both water and oxygen. Iron is oxidised at an anodic region, oxygen is reduced at a cathodic region, and the product is hydrated iron(III) oxide, Fe2O3.xH2O, the brown flaky solid we call rust.
Rusting is the corrosion of iron, and the syllabus treats it as an electrochemical (redox) reaction rather than a simple burning. It appears in the Form 5 Redox chapter and links directly to the reactivity series, sacrificial protection and everyday rust prevention, so examiners like it because one context tests several ideas at once.
Rusting is a redox reaction
On a piece of damp iron, some regions behave as an anode and others as a cathode, with the water film acting as an electrolyte. Iron atoms lose electrons and dissolve, oxygen takes those electrons, and the two products meet to form rust.
- At the anode (oxidation): Fe(s) → Fe2+(aq) + 2e-
- At the cathode (reduction): O2(g) + 2H2O(l) + 4e- → 4OH−(aq)
The iron(II) ions are then oxidised further by dissolved oxygen to iron(III) ions, which combine with hydroxide ions and water to give hydrated iron(III) oxide. The overall change can be written as:
4Fe(s) + 3O2(g) + 2xH2O(l) → 2Fe2O3.xH2O(s)
The formula Fe2O3.xH2O shows that rust is a hydrated oxide, this is why both water and oxygen are essential.
Conditions required
- Water and oxygen must both be present. Remove either one and rusting stops.
- Rusting is speeded up by an electrolyte such as salt (why cars rust faster near the sea), by acids (acid rain), and by contact with a less reactive metal such as tin or copper.
- Rusting is slowed by contact with a more reactive metal (sacrificial protection) and by any barrier that keeps out air and water.
Observations
A brown or orange-brown flaky solid forms on the iron surface. The rust is porous, so it flakes off and exposes fresh metal underneath, letting corrosion continue until the object is weakened or eaten through.
The classic experiment and prevention
The standard experiment places clean iron nails in different tubes: ordinary water (rusts), boiled water sealed with oil so no oxygen is present (no rust), and dry air over anhydrous calcium chloride so no water is present (no rust). This proves both water and oxygen are needed. A further set wraps nails with zinc, magnesium, tin or copper: the more reactive metals (zinc, magnesium) protect the iron by corroding in its place, while the less reactive metals (tin, copper) make it rust faster. Prevention methods you should be able to explain include painting, oiling, galvanising with zinc, plating, alloying to stainless steel, and attaching sacrificial blocks of magnesium to ships and pipelines.
Common mistakes to avoid
Do not call rusting simply “iron reacting with oxygen”, the mark scheme wants water and oxygen. Do not write the product as iron(III) oxide, Fe2O3, alone; rust is the hydrated oxide, Fe2O3.xH2O. When explaining sacrificial protection, state that the more reactive metal loses electrons more readily and so is oxidised instead of the iron, which is why the iron is protected even where the coating is scratched.
How it appears in the SPM exam
In Paper 2 (4541/2) you may be asked to write the half-equation for iron at the anode, to explain a nail experiment, or to compare metals wrapped around iron using the reactivity series. Justify each answer with electron transfer, and rusting questions become straightforward marks.
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