Metals

Properties of metals

Metals share a distinctive set of physical and chemical properties that follow directly from the way their atoms are held together — a lattice of positive ions in a sea of delocalised electrons.

Physical properties

Each follows from metallic bonding. Electrons free to move carry both electric current and thermal energy. Layers of positive ions slide past each other under stress, so the metal bends rather than snaps. And the strong electrostatic attraction between the ions and the electron sea holds them together tightly — hence the high melting points.

Chemical properties

Metals react with three things Cambridge focuses on: water, dilute acid and oxygen. Each turns the metal into a compound — metal + water → metal hydroxide + hydrogen; metal + acid → salt + hydrogen; metal + oxygen → metal oxide. The vigour of these reactions is what changes from metal to metal, and it is exactly what the reactivity series tabulates.

Uses of metals & alloys

The right metal for a job is the one whose properties match. Aluminium in aircraft skins because it’s light; copper in wires because it conducts; stainless steel in cutlery because it doesn’t rust. And when no pure metal quite fits, an alloy usually does.

Match property to use

What makes an alloy stronger

A pure metal is soft because its identical, evenly-spaced atoms form neat layers that slide over each other under stress — the same thing that makes it malleable. In an alloy, atoms of a different size are mixed in among the original ones; those larger or smaller atoms sit awkwardly, disrupting the tidy layers.

With the layers no longer able to slide freely, the alloy is harder and stronger than the pure metal. Steel is the canonical example: adding just 0.1% carbon to iron transforms a soft, easily-bent metal into one strong enough for skyscrapers. Brass (copper + zinc) is harder than either metal alone and resists tarnishing. Stainless steel (iron + chromium + nickel) resists corrosion completely — the chromium oxide layer that forms is impermeable.

The reactivity series

The reactivity series is a ranked list of metals from most reactive at the top to least reactive at the bottom. It is built from experiment — watching each metal react with cold water, steam or dilute acid — and once assembled it explains almost everything else in this chapter.

Reactions with water and acid

A useful rule falls out: any metal above hydrogen in the series reacts with dilute acid; any metal below it does not. Copper, silver and gold sit below hydrogen and so are safe to touch acid. Example: Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g).

ExtendedThe underlying reason

Every reaction in the table happens because a metal atom loses electrons. Potassium’s outer electron sits far from the nucleus and is easy to lose, so potassium reacts explosively; gold’s outer electrons are tightly bound, so gold sits in jewellery unchanged for millennia. The reactivity series is really a series of “how easily does this metal lose its electrons”.

ExtendedDisplacement

Because a more reactive metal holds its electrons less strongly, it can force a less reactive one out of its compounds: Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s). The blue solution fades and pink-brown copper is deposited.

Corrosion of iron

Iron rusts. Left in the open air with any moisture around, it reddens and flakes as its structure is eaten by an oxide with the wrong shape and no strength. Preventing rust is a matter of blocking one of its two ingredients, or convincing something else to react in the iron’s place.

What rust is — and needs

Rust is hydrated iron(III) oxide. Its formation requires two reagents simultaneously: oxygen from the air and water. Remove either and rust does not form — shown by the classic three-tube experiment: iron in air only, iron in boiled-and-oil-sealed water, iron in tap water; only the last one rusts.

Prevention: two strategies

A barrier only works while it stays intact — a chip in the paint exposes the iron and rust restarts from that point. Galvanising also protects chemically: if the zinc coating is scratched, the zinc is still more reactive than iron and corrodes preferentially, leaving the iron untouched until every scrap of zinc is used up.

ExtendedSacrificial protection

The sacrificial principle also works without covering the iron. Blocks of zinc (or magnesium) bolted to steel ship hulls corrode preferentially — being above iron in the reactivity series, they lose their electrons first while the iron stays intact.

Iron extraction — the blast furnace

Iron sits below carbon in the reactivity series, so it can be extracted by heating with carbon. The industrial version happens in a blast furnace — a 30-metre tower fed with iron ore, coke and limestone at the top, and hot air blown in at the bottom.

Three zones, three reactions

The furnace is a continuous reactor; molten iron and slag come out of the bottom. Zone 1 — combustion: coke burns in the hot air, C(s) + O₂(g) → CO₂(g), highly exothermic, providing the heat for the whole furnace. Zone 2 — making the reducing agent: rising CO₂ meets more hot coke, CO₂(g) + C(s) → 2CO(g). Zone 3 — iron reduction: carbon monoxide meets iron(III) oxide from the descending ore, Fe₂O₃(s) + 3CO(g) → 2Fe(l) + 3CO₂(g); the molten iron collects at the bottom to be tapped off.

Slag — removing the sandy waste

Iron ore contains silicon dioxide (sand). The limestone deals with it in two steps: calcium carbonate thermally decomposes, CaCO₃(s) → CaO(s) + CO₂(g); then the calcium oxide reacts with the acidic silicon dioxide, CaO(s) + SiO₂(s) → CaSiO₃(l). The molten calcium silicate slag floats on top of the iron.

Aluminium extraction — electrolysis

Aluminium sits above carbon in the reactivity series, so its ore cannot be reduced with coke — electrolysis is the only industrial option. The result is a metal with extraordinary properties but a high energy cost that shows up in every recycling drive on the planet.

ExtendedWhy electrolysis

Any metal above carbon in the reactivity series holds its oxygen too tightly for carbon monoxide to remove. Electrolysis, which uses electrical energy directly to force the reduction, is the only viable route — aluminium, magnesium, calcium, sodium and potassium are all extracted this way.

ExtendedThe setup and electrode reactions

Purified aluminium oxide is dissolved in molten cryolite so the electrolyte flows at around 950 °C. The mixture is electrolysed in a steel container lined with carbon, which acts as the cathode; blocks of carbon dipped in from above are the anodes. At the cathode, aluminium ions gain three electrons and are deposited as molten metal: Al³⁺ + 3e⁻ → Al. At the anode, oxide ions lose electrons to form oxygen gas: 2O²⁻ → O₂ + 4e⁻.

ExtendedThe carbon-anode problem

The oxygen released reacts with the hot carbon anodes, burning them away to carbon dioxide: C(s) + O₂(g) → CO₂(g). This is why the anodes must be replaced regularly.