GCSE Chemistry Reactivity & metals revision

A metal’s reactivity measures how eagerly its atoms lose electrons and form positive ions. Three experiment families rank them. With cold water: potassium, sodium, lithium and calcium fizz visibly (metal hydroxide + hydrogen), while the rest sit silent. With dilute acid: magnesium, aluminium, zinc and iron fizz steadily (salt + hydrogen) — but copper, silver and gold refuse, because they sit below hydrogen in the series. With displacement: drop one metal into the salt solution of another; reaction only happens if the strip’s metal is the more reactive of the two.

Aluminium looks unreactive in these tests only because its surface is sealed by a thin, tough skin of aluminium oxide — scratch it away and the metal beneath reacts readily. Non-metals carbon and hydrogen wedge into the list too: carbon between aluminium and zinc, hydrogen between iron and copper — placements that decide extraction routes.

Iron nails dropped into blue copper sulfate solution come out coated in orange-brown copper, and the blue fades: iron, being more reactive, has displaced copper. No oxygen is involved here, so the oxygen-language definition of oxidation does not apply; the Higher-tier treatment counts electrons. Iron atoms lose electrons (oxidation) to become Fe²⁺; copper ions gain electrons (reduction) to become copper atoms: Fe + Cu²⁺ → Fe²⁺ + Cu. Every displacement is one metal donating electrons to another metal’s ions — and the more reactive metal is always the donor.

Metals occur in ores mostly as oxides (or compounds convertible to oxides). Extraction means reduction — stripping the oxygen away. Carbon (as coke) is cheap and, sitting between aluminium and zinc in the series, it can reduce the oxides of every metal below it: the carbon takes the oxygen because carbon is more reactive than those metals. For iron(III) oxide in the blast furnace: Fe₂O₃ + 3CO → 2Fe + 3CO₂ — iron(III) oxide is reduced (loses oxygen), carbon monoxide is oxidised (gains it).

1. 1Blast furnace in outline

   Hot air blasts in at the base; coke (carbon) burns to carbon dioxide, releasing intense heat.

2. 2Making the reducer

   CO₂ reacts with more coke to form carbon monoxide, the actual reducing agent.

3. 3Reduction

   CO strips oxygen from iron(III) oxide descending from the top; molten iron collects and is tapped at the base.

4. 4Slag

   Limestone removes sandy impurities as slag, which floats and is tapped separately.

Metals above carbon — potassium to aluminium — cannot be reduced by carbon (carbon is not reactive enough to displace them), so they are won by electrolysis of their molten compounds: costly in electricity, which is exactly why aluminium costs several times more than iron.

High-grade copper ores are running out, so chemists turn to low-grade ores and even mine spoil. Phytoextraction grows plants that absorb copper ions from the soil into their shoots; the harvest is burned, leaving ash rich in copper compounds that can be processed by the usual routes. Bioleaching piles low-grade ore with bacteria that convert copper compounds into a solution (leachate) of copper ions — a form of indirect extraction. Both are slow (seasons, not minutes) but need little energy, work on ores no furnace would touch, and disturb the landscape less than new mines. The copper they yield can be purified by electrolysis, closing the loop with traditional methods.