Lesson 3 · 30 min

Extracting metals

Gold can be picked out of river sand, iron must be smelted with coke in a blast furnace, and aluminium needs electrolysis. Thermodynamics decides the reducing agent; the energy bill decides the price.

In this lesson

Choose an extraction method from the stability of a metal’s oxide or sulfide: native metal, heating, reduction by carbon, reduction by another metal, or electrolysis
Describe the main reactions in each zone of the blast furnace, and the roles of slag and steelmaking
Calculate the decomposition voltage, output and energy use of a Hall–Héroult cell
Compare the energy cost of extraction routes and the value of recycling

Gold can be panned straight out of river sand; people learned to smelt iron with charcoal more than three thousand years ago; yet aluminium — the most abundant metal in the Earth’s crust — was still more precious than gold in the 1850s, when Napoleon III reputedly kept aluminium cutlery for his most honoured guests. Only in 1886, when Hall and Héroult independently invented the electrolysis of molten alumina, did aluminium become cheap. All of these are reductions of a metal compound — so why do the difficulties differ so wildly?

Predict first

Why can aluminium not be won like iron, by reducing its oxide with coke in a furnace?

Pick an answer and lock it in before you look. A wrong guess teaches more than no guess.

Al₂O₃ is too stable: at any practical furnace temperature, ΔG for its reduction by carbon is positive Aluminium melts too low and would boil away in the furnace Carbon and aluminium never touch, so they cannot react Electrolysis is cheaper than a blast furnace

From ore to metal

Ore is first concentrated; sulfide ores are usually roasted in air to oxides ( $\ce{2ZnS} + \ce{3O2} \mathrel{\ce{->}} \ce{2ZnO} + \ce{2SO2}$ — the $\ce{SO2}$ goes to make sulfuric acid). Which reduction to use then depends on how stable the metal oxide is (Ellingham diagrams turn this into a map):

Native metals: gold and platinum barely combine with oxygen and occur as the elements.
Heat alone: compounds of mercury and silver give up the metal on heating, e.g. $\ce{HgS} + \ce{O2} \mathrel{\ce{->}} \ce{Hg} + \ce{SO2}$ .
Carbon or carbon monoxide: the oxides of iron, zinc, lead, tin and copper are reduced by coke or CO at high temperature.
A more reactive metal: the thermite route $\ce{Cr2O3} + \ce{2Al} \mathrel{\ce{->}} \ce{2Cr} + \ce{Al2O3}$ ; titanium is first turned into $\ce{TiCl4}$ and reduced with magnesium (the Kroll process), because carbon would give titanium carbide.
Electrolysis: sodium, magnesium, aluminium and other very reactive metals form compounds too stable for anything else.

Try the combinations yourself in the smelter lab.

Iron: the blast furnace and steel

A blast furnace is charged at the top with iron ore, coke and limestone, and hot air is blown in through tuyeres near the bottom. Coke burns in front of the tuyeres (about 2000 °C), and the $\ce{CO2}$ formed is reduced by more hot coke to CO ( $\ce{CO2} + \ce{C} \mathrel{\ce{->}} \ce{2CO}$ , with ΔG° < 0 above about 710 °C). The rising CO reduces the iron oxides step by step in the stack: $\ce{Fe2O3} \mathrel{\ce{->}} \ce{Fe3O4} \mathrel{\ce{->}} \ce{FeO} \mathrel{\ce{->}} \ce{Fe}$ . Limestone breaks down to CaO, which combines with the $\ce{SiO2}$ of the ore to form molten slag, $\ce{CaSiO3}$ , tapped off above the molten iron.

Blast-furnace iron carries about 4 % carbon and is hard and brittle. In steelmaking pure oxygen is blown through the molten iron to burn the excess carbon off as CO, while silicon and phosphorus go into the slag, leaving steel with less than 2 % carbon. Add more than 10.5 % chromium and a dense $\ce{Cr2O3}$ film forms on the surface: stainless steel.

ΔG of each reaction at the zone’s temperature, from ΔH° and S°

The blast furnace, layer by layer

Reaction	ΔG / kJ
$\ce{C + O2 -> CO2}$ Coke burns in the hot blast	−397
$\ce{CO2 + C -> 2CO}$ CO₂ is reduced to CO by coke	−43
$\ce{3Fe2O3 + CO -> 2Fe3O4 + CO2}$ Haematite → magnetite	−104
$\ce{Fe3O4 + CO -> 3FeO + CO2}$ Magnetite → iron(II) oxide	−44
$\ce{FeO + CO -> Fe + CO2}$ FeO reduced to iron by CO	10
$\ce{FeO + C -> Fe + CO}$ Direct reduction by coke	−32
$\ce{CaCO3 -> CaO + CO2}$ Limestone decomposes	−17
$\ce{CaO + SiO2 -> CaSiO3}$ Slag forms	−92

Main reactions here

At 950 °C, FeO + CO ⇌ Fe + CO₂ has K = 0.37: FeO is reduced only if CO makes up more than 73 % of the CO + CO₂.

ΔG° against temperature for key reactions

CO₂ is reduced to CO by coke
Direct reduction by coke
FeO reduced to iron by CO
Limestone decomposes

CO₂ is reduced to CO by coke: ΔG° < 0 above about 708 °C
Direct reduction by coke: ΔG° < 0 above about 746 °C
FeO reduced to iron by CO: ΔG° < 0 above about 360 °C
Limestone decomposes: ΔG° < 0 above about 845 °C

Step down from the top zone by zone: the main reactions of each zone are highlighted, and ΔG is computed at that zone’s temperature. Notice that FeO + CO has a slightly positive ΔG when hot — the gas has to be mostly CO for it to go — while direct reduction of FeO by coke only becomes favourable above about 750 °C.

Aluminium: electrolysis of molten alumina

Bauxite is first dissolved in hot NaOH solution (the Bayer process — which relies on the amphoterism of $\ce{Al(OH)3}$ , see boron, aluminium, carbon and silicon) and pure alumina is precipitated back out. $\ce{Al2O3}$ melts only at 2072 °C, but it dissolves in molten cryolite, $\ce{Na3AlF6}$ , and is electrolysed at about 960 °C: aluminium forms at the cathode and pools at the bottom of the cell, while the carbon anodes are consumed — overall $\ce{2Al2O3} + \ce{3C} \mathrel{\ce{->}} \ce{4Al} + \ce{3CO2}$ .

Thermodynamic data put the reversible decomposition voltage of this reaction at only about 1.19 V at 960 °C, yet real cells run at about 4.2 V: the extra overcomes electrode overvoltages and the resistance of the bath and ends up as heat — exactly what keeps the bath molten. Each kilogram of aluminium takes 13–15 kWh of electricity; recycling needs only remelting, about a twentieth of that energy.

m = \frac{I\,t\,\eta\,M}{zF}\qquad w = \frac{zFU}{M\eta}

Faraday’s laws: the mass deposited, m, and the electrical energy per kilogram of product, w.

$I,\ t$: current and time
$\eta$: current efficiency (the fraction of the current that deposits metal)
$M,\ z$: molar mass and electrons per ion (3 for aluminium)
$U$: cell voltage
$F$: Faraday constant, 96 485 C/mol

Decomposition voltage from thermodynamic data, output from Faraday’s laws

The energy budget of a Hall–Héroult cell

Current · 400 kA Current density · 0.80 A/cm² Anode–cathode distance · 4.5 cm Current efficiency · 94 %

Cell cross-section (schematic)

$\ce{2Al2O3 + 3C -> 4Al + 3CO2}$

Reversible decomposition 1.19 V
Anode overvoltage 0.55 V
Cathode overvoltage 0.08 V
Ohmic drop in the bath 1.64 V
Gas bubbles 0.25 V
Conductors 0.60 V
Cell voltage 4.30 V

Electricity per kg Al 13.6 kWh/kg Modern smelters: 13–15 kWh/kg

Aluminium per cell per day 3.03 t

Direct CO₂ from the anode 1.30 kg/kg

Only 28 % of the voltage goes into splitting alumina; the rest becomes heat that keeps the bath molten. Recycling only has to remelt the metal, about 0.7 kWh/kg — less than a twentieth of smelting.

Adjust the anode–cathode distance: the ohmic drop in the bath is the largest “loss”, and a shorter gap saves power — but too close and the metal pad shorts to the anode and current efficiency falls. Switch to an inert anode: the decomposition voltage rises from about 1.2 V to about 2.2 V, but no CO₂ comes directly from the anode.

Worked example

How much aluminium does one cell make in a day?

A cell runs at 400 kA with 94 % current efficiency for 24 hours. What mass of aluminium does it make (molar mass 26.98 g/mol)?

Concept check

How is titanium produced industrially?

Reducing TiO₂ with coke in a blast furnace Electrolysing an aqueous solution of TiO₂ Converting TiO₂ into TiCl₄ and reducing it with magnesium Simply heating TiO₂ until it decomposes

Try the calculation

At 960 °C, $\ce{2Al2O3} + \ce{3C} \mathrel{\ce{->}} \ce{4Al} + \ce{3CO2}$ has $\Delta_\mathrm{r}G \approx +1374$ kJ with 12 electrons transferred. What is the reversible decomposition voltage?

Key takeaways

The extraction route follows the compound’s stability: gold is native, mercury needs only heat, iron, zinc and copper use carbon, titanium and chromium a reactive metal, and sodium, magnesium and aluminium electrolysis.
In the blast furnace burning coke supplies heat and CO, which reduces the iron oxides step by step; CaO and SiO₂ form slag; blowing oxygen through the iron removes excess carbon to make steel.
Aluminium is electrolysed from alumina dissolved in cryolite at about 960 °C: decomposition voltage about 1.2 V, cell voltage about 4.2 V, 13–15 kWh per kilogram.
Recycled aluminium needs about a twentieth of the energy of primary metal.

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Data and references

K. Grjotheim, H. Kvande (eds.), Introduction to Aluminium Electrolysis, 2nd ed., Aluminium-Verlag (1993).
A. Geerdes, H. Toxopeus, C. van der Vliet, Modern Blast Furnace Ironmaking, 2nd ed., IOS Press (2009).
J. D. Cox, D. D. Wagman, V. A. Medvedev, CODATA Key Values for Thermodynamics (1989); D. D. Wagman et al., NBS tables (1982).
N. N. Greenwood, A. Earnshaw, Chemistry of the Elements, 2nd ed. (1997), ch. 7 (aluminium), 21 (titanium), 25 (iron).
D. R. Gaskell, Introduction to the Thermodynamics of Materials, 5th ed., Taylor & Francis (2008) (Ellingham diagrams).