Rusting and its prevention

This page covers one Form 5 content standard from Redox Equilibrium: rusting and its prevention. Rusting is the everyday redox reaction the exam loves, because it lets you apply oxidation, reduction, half-equations and the reactivity series to a real, familiar problem. Learn it as a redox story, not a list of facts.

Rusting is a redox reaction

Rust is hydrated iron(III) oxide, roughly Fe₂O₃.xH2O, the brown, flaky solid that forms on iron and steel. Rusting only happens when both water and oxygen are present; remove either one and iron does not rust. That is why the classic experiment uses three tubes, iron with water and air (rusts), iron with boiled water and oil to exclude air (does not rust), and iron with anhydrous calcium chloride and dry air (does not rust).

The process is redox. Iron is oxidised at one spot on the surface, losing electrons:

Fe → Fe²⁺ + 2e- (iron’s oxidation number rises from 0 to +2)

The electrons travel through the metal to another spot, where oxygen is reduced in the presence of water:

O₂ + 2H₂O + 4e- → 4OH⁻

The Fe²⁺ ions are then further oxidised to Fe³⁺, which combine with hydroxide and water to form the hydrated iron(III) oxide we call rust. So iron is the reducing agent and oxygen is the oxidising agent.

Factors that speed up rusting

• Electrolytes such as salt (sodium chloride) make water conduct better, so electrons and ions move more easily and rusting is faster. This is why cars and structures rust badly near the sea.
• Acidic conditions (for example acid rain) speed rusting for the same reason.
• Contact with a less reactive metal (for example iron touching copper or tin) makes iron rust faster, because iron loses electrons even more readily by comparison, iron becomes the one that is oxidised.

Preventing rust

There are three ideas, and you should be able to explain why each works.

1. Barrier methods, keep out water and oxygen by coating: painting, oiling/greasing, coating with plastic, or plating with an unreactive metal (for example chromium plating). Simple, but if the barrier is scratched, the exposed iron rusts.
2. Sacrificial protection, attach a more reactive metal such as zinc or magnesium. Because it is more reactive, it loses electrons instead of the iron, so it corrodes and is used up while the iron is protected. Crucially, this still works when the surface is scratched, because protection is by electron supply, not by a physical barrier. Ships’ hulls and underground pipes carry sacrificial zinc or magnesium blocks for this reason.
3. Galvanising, coating iron with zinc, is a clever double defence: the zinc is a barrier and, if scratched, gives sacrificial protection.
4. Alloying, mixing iron with other elements to make it rust-resistant, most famously stainless steel (iron with chromium and nickel), which forms a protective layer that does not flake off.

Worked example

Question. A steel gate is protected by attaching blocks of magnesium to it. Explain, using half-equations and the reactivity series, why the magnesium protects the iron even after the gate’s paint is scratched.

Step 1, compare reactivity. Magnesium is more reactive than iron, so magnesium loses electrons more readily than iron does.

Step 2, say what magnesium does. Magnesium is oxidised in place of the iron: Mg → Mg²⁺ + 2e-. The electrons it releases flow into the iron.

Step 3, say what this does for iron. With a supply of electrons pushed onto it, the iron is prevented from losing its own electrons, so Fe → Fe²⁺ + 2e- is suppressed and the iron does not rust. If anything, the extra electrons drive the oxygen-reduction half-reaction at the iron surface without consuming the iron.

Step 4, explain the “even if scratched” part. Because protection comes from the electron supply through the metal, not from a physical cover, a scratch in the paint does not matter, the magnesium keeps feeding electrons to the exposed iron.

Answer. Magnesium, being more reactive, is oxidised (Mg → Mg²⁺ + 2e-) and sacrificially supplies electrons to the iron, so the iron is not oxidised and does not rust; this works even when scratched because it depends on electron flow, not on a barrier.

Practice question

Two identical iron nails are half-dipped in agar jelly containing an indicator. Nail A is wrapped with a strip of zinc; nail B is wrapped with a strip of copper. Predict which nail rusts and explain in terms of the reactivity series.

Answer. Nail B (with copper) rusts; nail A (with zinc) does not. Zinc is more reactive than iron, so it is oxidised instead of the iron and protects it (sacrificial protection). Copper is less reactive than iron, so the iron loses electrons more readily than copper and is oxidised faster, the copper actually speeds up the rusting of nail B.

Exam tip

Two conditions, every time: rusting needs both water and oxygen, state both. For prevention questions, always name the method and give the redox reason: “zinc is more reactive, so it is oxidised in place of the iron and supplies electrons to it”. The highest-value point in sacrificial-protection answers is that it still works when the coating is scratched, examiners look for it. Do not confuse galvanising (zinc, gives sacrificial protection) with tin-plating (tin is less reactive, so a scratch makes the iron rust faster). These distinctions are common SPM Chemistry Paper 2 marks.

Where this fits in the chapter

Rusting ties the reactivity series to a real redox process and sets up the idea, used again in cells, that a more reactive metal gives up electrons more readily.

• Up to the chapter hub: Redox Equilibrium.
• Sideways: the chapter’s revision notes and the common mistakes list, which flags the tin-vs-zinc trap.

Our online 1-to-1 SPM Chemistry teachers, teaching in English from RM50 per hour, rehearse the exact wording examiners reward for sacrificial protection, so you never lose the “works even when scratched” mark.