Explaining Reactivity
Section: 9. Metals | Syllabus: Cambridge IGCSE Chemistry (0620)
Reactivity as a Tendency to Form Positive Ions
A metal's position in the reactivity series reflects how readily it loses electrons to form a positive ion. The more reactive the metal, the more strongly it "wants" to lose electrons - so a more reactive metal will always take the electrons away from a less reactive metal's ion, displacing it.
| Metal added | Mg²⁺(aq) | Zn²⁺(aq) | Fe²⁺(aq) | Cu²⁺(aq) | Ag⁺(aq) |
|---|---|---|---|---|---|
| Magnesium | — | displaces | displaces | displaces | displaces |
| Zinc | no reaction | — | displaces | displaces | displaces |
| Iron | no reaction | no reaction | — | displaces | displaces |
| Copper | no reaction | no reaction | no reaction | — | displaces |
| Silver | no reaction | no reaction | no reaction | no reaction | — |
This grid of results is itself evidence for the reactivity order Mg > Zn > Fe > Cu > Ag: whichever metal displaces every other metal's ions (magnesium, top row) must be the most reactive, and the one that never displaces anything (silver, bottom row) must be the least reactive.
Interpreting a Displacement Reaction as Electron Transfer
Take magnesium added to iron(II) sulfate solution:
Mg(s) + Fe²⁺(aq) → Mg²⁺(aq) + Fe(s)
- Mg → Mg²⁺ + 2e⁻ (magnesium is oxidised - loses electrons)
- Fe²⁺ + 2e⁻ → Fe (iron(II) ions are reduced - gain electrons)
Magnesium is more reactive than iron because it loses its electrons more readily - so in the mixture, magnesium always "wins" and gives its electrons to the iron(II) ions.
Worked Example
- Question: iron powder is added to silver nitrate solution and a grey coating forms on the iron. Explain this observation in terms of electron transfer. [2]
- Answer: iron is more reactive than silver, so iron atoms lose electrons (are oxidised) to form Fe²⁺ ions [1]; silver ions gain these electrons (are reduced) to form silver atoms, which coat the iron as a grey solid [1]
The Apparent Unreactivity of Aluminium
Aluminium sits just below magnesium in the reactivity series, so it should react about as vigorously as magnesium. In practice, aluminium objects seem far less reactive - they don't visibly react with air, water or dilute acids the way its position would suggest.
The reason is that aluminium reacts almost instantly with oxygen in the air to form a very thin layer of aluminium oxide (Al₂O₃) on its surface:
4Al(s) + 3O₂(g) → 2Al₂O₃(s)
This oxide layer is tough and impermeable, and sticks firmly to the metal underneath. It acts as a protective barrier that stops oxygen and water reaching the reactive aluminium beneath it - so the metal appears unreactive, even though the aluminium itself is genuinely very reactive.
Worked Example
- Question: which statement explains why aluminium has good resistance to corrosion? (A) aluminium is a highly reactive metal (B) aluminium has an oxide layer on its surface (C) aluminium requires heat to react (D) aluminium is an inert metal [1]
- Answer: B. The thin, tightly-adhering layer of aluminium oxide that forms on exposure to air acts as a barrier, preventing further reaction. Aluminium is not inert (ruling out D) - it is actually quite reactive (so A alone doesn't explain the corrosion resistance), and it reacts with oxygen at room temperature without needing heat (ruling out C).
Common Mistakes
- ❌ Concluding that aluminium must be unreactive, since it doesn't rust or corrode visibly
- ✓ aluminium is genuinely reactive (it sits above zinc and iron in the series) - it only appears unreactive because its own oxide layer protects it. If the oxide layer is removed or prevented from forming, aluminium reacts as vigorously as its position predicts.
Interactive revision notes, videos and practice questions load below.