IGCSE Chemistry Metals notes

Chemically, metals are reducing agents: their atoms tend to lose electrons and form positive ions. How easily this happens decides everything else in this topic. Very reactive metals (potassium, sodium, calcium) react with cold water, giving a hydroxide and hydrogen. Moderately reactive metals such as magnesium react with steam, and most metals above hydrogen in the reactivity series react with dilute acids, giving a salt and hydrogen. Metals react with oxygen to form oxides; more reactive metals burn more brightly.

• Aluminium in aircraft: very low density.
• Aluminium in overhead cables: low density and good electrical conductivity.
• Aluminium food containers: resistant to corrosion (see the oxide layer below).
• Copper in electrical wiring: excellent conductivity and ductility — it can be drawn into wires.

An alloy is a mixture of a metal with other elements. Brass is copper mixed with zinc; stainless steel is iron mixed with chromium, nickel and carbon. Alloys can be harder and stronger than the pure metal, which is why stainless steel is used for cutlery — it keeps its polish, resists rusting and survives rough handling.

Supplement: in a pure metal the atoms are one uniform size and sit in neat layers that can slide over each other when a force is applied — that is why pure metals are soft and malleable. Adding atoms of a different size disturbs the regular layers, so they can no longer slide easily. The distorted structure makes the alloy harder and stronger.

The reactivity series lists metals in order of how readily they form positive ions:

1. 1potassium (most reactive)
2. 2sodium
3. 3calcium
4. 4magnesium
5. 5aluminium
6. 6carbon (non-metal, included as a benchmark)
7. 7zinc
8. 8iron
9. 9hydrogen (benchmark)
10. 10copper
11. 11silver
12. 12gold (least reactive)

Exam questions give observations — fizzing speeds, temperature rises, colour changes — and expect you to deduce the order. Supplement: displacement reactions are the cleanest evidence. Dip a more reactive metal into the aqueous ions of a less reactive one and a reaction happens: iron placed in copper(II) sulfate solution becomes coated with brown copper because iron ions form more readily, Fe + Cu²⁺ → Fe²⁺ + Cu. Reverse the pair and nothing happens.

Supplement: aluminium sits high in the series yet barely reacts in everyday life. A tough, airtight oxide layer forms instantly on its surface and shields the metal underneath — the reason aluminium resists corrosion in food containers and window frames even though the metal itself is reactive.

The classic three-test-tube experiment nails down the conditions. A nail in dry air stays shiny, a nail in boiled water sealed with oil stays shiny, and a nail in open tap water rusts. So rusting needs both oxygen and water, forming hydrated iron(III) oxide, the flaky brown substance that crumbles off and exposes fresh metal.

Supplement: zinc is more reactive than iron, so a galvanised surface is protected twice — the zinc layer physically blocks air and moisture, and if the coating is scratched the zinc still loses electrons preferentially and corrodes first, sparing the iron underneath. Ships’ hulls and underground pipelines carry blocks of zinc or magnesium for exactly this reason.

An ore is a rock containing enough metal compound to make extraction worthwhile. Metals below carbon (zinc, iron, copper) can be extracted by heating their oxide with carbon, because carbon displaces them. Metals above carbon (potassium to aluminium) need electrolysis, which is far more expensive — the position in the series sets the price of the metal.

1. 1Coke burns

   C + O₂ → CO₂, releasing the heat that keeps the furnace at about 1500 °C.

2. 2Carbon dioxide is reduced

   C + CO₂ → 2CO, making the reducing agent carbon monoxide.

3. 3Iron oxide is reduced

   Fe₂O₃ + 3CO → 2Fe + 3CO₂ — molten iron runs to the base of the furnace.

4. 4Limestone removes impurities

   CaCO₃ → CaO + CO₂, then CaO + SiO₂ → CaSiO₃; the molten calcium silicate slag floats on the iron and is run off.

The equations above are Supplement; at Core you need the chemistry in words. Iron is extracted from hematite, its iron(III) oxide ore, in the blast furnace. Coke (carbon) burns in hot air to reach furnace temperatures and produce carbon dioxide; more coke converts that to carbon monoxide, which reduces the iron(III) oxide to molten iron. Limestone decomposes in the heat, and the calcium oxide formed combines with the sandy impurity (silicon dioxide) to make slag.

Aluminium is too reactive for carbon, so its ore bauxite (aluminium oxide) is dissolved in molten cryolite, which lowers the operating temperature and saves energy. Supplement: at the cathode, Al³⁺ + 3e⁻ → Al; at the anode, 2O²⁻ → O₂ + 4e⁻. The oxygen attacks the hot carbon anodes, forming carbon dioxide, so the anodes burn away and must be replaced regularly.