The reaction
A mixture of aluminium powder and iron(III) oxide, once started, reacts to give molten iron and aluminium oxide. It is a displacement reaction and a redox reaction, and it is strongly exothermic.
Fe2O3(s) + 2Al(s)→2Fe(l) + Al2O3(s)
ΔH ≈ −850 kJ per mole of Fe2O3
Why aluminium can displace iron
Aluminium is higher than iron in the reactivity series. A more reactive metal bonds to oxygen more strongly, so aluminium takes the oxygen away from iron(III) oxide and leaves iron behind as the free metal.
KNaCaMgAlCZnFePbCuAgAu
More reactiveLess reactive
The reverse reaction does not happen: heating iron with aluminium oxide gives no reaction, because iron is less reactive than aluminium.
Oxidation and reduction
Aluminium
Gains oxygen and loses electrons, so it is oxidised.
It is the reducing agent.
Al → Al3+ + 3e−
Iron(III) oxide
Loses oxygen; Fe3+ gains electrons, so it is reduced.
It is the oxidising agent.
Fe3+ + 3e− → Fe
Memory aid, OIL RIG: Oxidation Is Loss of electrons, Reduction Is Gain.
Energy changes
The mixture is stable at room temperature. Burning magnesium ribbon supplies the activation energy. Once started, the reaction releases far more energy than it took to start, enough to keep itself going and to melt the iron.
Why the iron is molten and glowing
The reaction reaches about 2500 °C, well above the melting point of iron (1538 °C), so the iron forms as a liquid. Hot objects give out light: dull red at around 600 °C, then orange, yellow and finally white as the temperature rises.
600 °C1000 °C1500 °C2500 °C
Molten iron is denser than aluminium oxide, so the iron sinks to the bottom of the crucible and the oxide floats on top as slag. That is why the iron can be run off from a tap hole at the bottom.
Worked calculation
What mass of iron forms from 1000 g of iron(III) oxide? (Ar: Fe = 56, O = 16, Al = 27)
- Mr of Fe2O3 = (2 × 56) + (3 × 16) = 160
- Moles of Fe2O3 = 1000 ÷ 160 = 6.25 mol
- 1 mol Fe2O3 gives 2 mol Fe, so 12.5 mol Fe
- Mass of Fe = 12.5 × 56 = 700 g
The same charge needs 12.5 × 27 = 337.5 g of aluminium. This is the 700 g shown in the simulation.
Uses
- Welding railway tracks on site. It needs no electricity or furnace, so a joint can be made anywhere along the line.
- Repairing large steel parts such as cracked machinery castings.
- Extracting metals such as chromium and manganese from their oxides (aluminothermic reduction).
Where marks are lost
- Calling aluminium the oxidising agent. It is the reducing agent; it is itself oxidised.
- Saying the reaction is endothermic because it has to be started with heat. The magnesium only supplies the activation energy; overall, energy is released.
- Writing an unbalanced equation. One Fe2O3 needs two Al.
- Naming the products vaguely. Write iron and aluminium oxide.
Safety
This reaction is only ever demonstrated by trained staff, outdoors or in a fume cupboard, behind a safety screen. Students must never attempt it.
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