Extraction of iron
Iron is the most widely used metal in the world, mainly in the form of steel. If you've used any
form of transport today, it probably contained plenty of iron. Steel is used to make most cars,
buses, and trains. In this unit you will see how we extract iron from its ore.
The main ore of iron is called haematite, figure 1. It is a mineral composed of iron(III)
Figure 1: Haematite.
Q: What is the charge on an ion of iron in haematite?
Q: What is the charge on the oxide ion in haematite?
To extract the iron from iron(lll) oxide, we remove the oxygen chemically. The iron
(III) oxide must be reduced. This is done in a blast furnace, figure 2.
The haematite is fed into the top of the furnace along with coke (a cheap form of
carbon, made from coal) and limestone.
Figure 2: A blast furnace.
Blast furnaces are huge structures that can be over 50 metres tall. They are run
continuously for years before their linings have to be replaced. The furnace is heated
by blasts of hot air and the combustion of coke inside the furnace raises the
temperature still further. Typically the temperature at the heart of a blast furnace is
about 1500 °C.
Q: Which of the following are raw materials used in the extraction of iron?
O Carbon dioxide,
Inside the blast furnace
The coke is added to the haematite to provide a reducing agent for the iron(III) oxide. Some
iron(III) oxide is reduced by carbon in contact with the haematite ore:
iron(III) oxide + carbon iron + carbon dioxide
2 Fe2O3(s) + 3 C(s) 4 Fe(l) + 3 CO2(g)
However, most of the iron(III) oxide is reduced by carbon monoxide gas. As a gas, this can
circulate freely in the blast furnace. It is made when carbon dioxide (made from the coke
burning in the blasts of hot air) reacts with more hot coke.
carbon + oxygen carbon dioxide
C(s) + O2(g) CO2(g)
carbon dioxide + carbon carbon monoxide
CO2(g) + C(s) 2 CO(g)
The carbon monoxide then reduces iron(III) oxide:
iron(III) oxide + carbon monoxide iron + carbon dioxide
Fe2O3(s) + 3 CO(g) 2 Fe(l) + 3 CO2(g)
The iron forms as a liquid in the high temperatures inside the blast furnace. It sinks to the
bottom of the furnace where it is tapped off.
Limestone (containing calcium carbonate) is added to get rid of the sandy bits of rock in the
ore, which are acidic. The limestone is broken down in the hot furnace:
calcium carbonate calcium oxide + carbon dioxide
CaCO3(s) CaO(s) + CO2(g)
Calcium oxide is a base, so it reacts with the acidic silicon dioxide (sand):
calcium oxide + silicon dioxide calcium silicate
CaO(s) + SiO 2 (s) CaSiO 3 (l)
The molten calcium silicate (known as slag) floats on top of the molten iron at the base
of the furnace. Here it is tapped off and cooled to use in building roads. Figure 3,
shows the stages in the extraction of iron using a blast furnace.
Figure 3: inside a blast furnace.
Q: Which two word equations show the reduction of iron as it happens in the blast
O iron(III) oxide + carbon iron + carbon dioxide,
O iron + carbon iron(III) oxide + carbon dioxide,
O iron (III) oxide + carbon monoxide iron + carbon dioxide.
Q: What is the main reducing agent in the blast furnace?
O carbon dioxide,
O carbon monoxide,
Q: Why is limestone added to the blast furnace?
O To reduce the iron(III) oxide,
O To remove sandy impurities,
O To produce carbon monoxide.
Iron is extracted from its ore, haematite, in a blast furnace. The ore is fed into the top of the
furnace along with coke and limestone. The main reducing agent is carbon monoxide:
iron(III) oxide + carbon monoxide iron + carbon dioxide
The molten iron is then tapped from the bottom of the furnace. The limestone decomposes in
the hot furnace, forming calcium oxide. This reacts with the sandy impurities (silicon dioxide)
to form a slag. The slag can be used in making roads.
1. Match the raw material used in a blast furnace to the substance it contains.
2. What temperature do we find in the heart of a blast furnace?
O 500 °C,
O 1500 °C,
O 5000 °C.
3. There are two reducing agents inside a blast furnace. Which of the following are the
O Iron(lll) oxide,
O Carbon monoxide,
O Carbon dioxide.
4. Were does molten iron collect in the blast furnace?
O At the bottom,
O Just above the molten slag,
O At the top of the furnace.
5. What role does limestone play in the extraction of iron?
O It provides an additional source of carbon for reducing the ore,
O It removes acidic (sandy) impurities present in the ore,
O Burning limestone helps to maintain the high temperature needed in the furnace,
O It acts as a catalyst in the conversion of carbon dioxide to carbon monoxide.
6. Which equation correctly describes the reduction of iron(III) oxide?
O Fe2O3(s) + CO(g) Fe(l) + CO2(g)
O Fe2O3(s) + CO(g) Fe2(l) + CO4(g)
O Fe2O3(s) + 3 CO(g) 2 Fe(l) + 3 CO2(g)
Extraction of Aluminium
When was the last time you used something containing aluminium metal? Perhaps it was drinking
from a can of cola or some other fizzy drink. Aluminium is a very useful metal. It is a good
conductor of heat and electricity. It resists corrosion - strange, you might think, for a metal so
high in the Reactivity Series - and has a low density for a metal.
Aluminium metal is extracted from its ore, bauxite. The brown colour in figure 4 comes from an
impurity - iron(III) oxide. Bauxite is often dug out of the ground in open-cast mines, figure 5. In
Europe we can find mines in Sardinia and Ireland.
Figure 4:Bauxite - an ore of aluminium.
Figure 5: Open-cast mining of bauxite ore.
The ore contains aluminium oxide. This will eventually be reduced (have its oxygen removed) by
electrolysis. But first we have to use chemical separation to extract the aluminium oxide from
the rest of the rock in the ore. For every 5 tonnes of bauxite we get only 1 tonne of aluminium.
The rest is liquid waste which is pumped into shallow storage ponds, where it gradually dries in
the sun. In Jamaica, this waste is now used as a low-cost raw material for bricks.
Q: What is the useful compound in bauxite?
O Aluminium chloride,
O Aluminium oxide,
O Iron(lll) oxide
Q: Make a list of the bad effects mining bauxite has on our environment.
Because aluminium is more reactive than carbon, it has to be extracted from its ore by
electrolysis. This is done industrially in reaction cells like the one below.
Figure 6: Extraction of aluminium.
Aluminium is extracted from aluminium oxide. The aluminium oxide is added to a reaction cell
filled with molten cryolite (a less common ore of aluminium), figure 6. The cryolite lowers the
melting point of the aluminium oxide to save energy.
Q: What are the positive electrodes in the cell made from?
O Carbon (graphite).
Q: Aluminium plants (Anglesey, Wales) are often sited near hydroelectric power stations. Why
O The bauxite is often found in mountains near the power station,
O You need a source of cheap electricity nearby because the electrolysis uses so much
O Water is needed in the electrolysis cells.
Q: At which electrode is the aluminium formed?
O Anode (+),
O Cathode (-).
Q: Which half-equation describes how aluminium ions are discharged in the cell?
O Al3+ + 3e- Al,
O Al+ + e- Al,
O Al3+ - 3e- Al.
Q: Why do the anodes in the cell have to be replaced frequently?
O They melt in the high temperature inside the cell,
O They are burned away by oxygen reacting with the carbon to form carbon dioxide gas,
O They get coated in a thick crust of solid aluminium oxide,
O The aluminium formed on them has to be removed.
Q: Why is molten cryolite used to dissolve the aluminium oxide the cell?
O To make the electrolyte richer in aluminium,
O To increase the resistance of the electrolyte,
O To increase the melting point of the electrolyte,
O To lower the melting point of the electrolyte.
Reduction and oxidation
We say that aluminium ions are reduced during electrolysis. Normally, we think of reduction as
the removal of oxygen from a compound. But we can define it more generally in terms of
Reduction is the gaining of electrons.
At the cathode, each Al3+ ion gains 3 electrons to form an aluminium atom. So the aluminium
ions are reduced: Al3+ + 3 e- Al
The chemical opposite of reduction is oxidation. Therefore:
Oxidation is the loss of electrons.
Oxidation always happens at the anode in electrolysis. In the case of the electrolysis of
aluminium oxide: 2 O 2- - 4e - O2
Here, the oxide ions, O2-, have been oxidized to form oxygen gas. This then reacts with
the carbon anodes to form carbon dioxide gas.
You can remember this by thinking of the word 'OILRIG'. This stands for: Oxidation Is
Loss, Reduction Is Gain (of electrons).
Decide whether the following half-equations represent oxidation or reduction reactions.
Na+ + e- Na
2 Cl- - 2 e- Cl2
2 O2- - 4 e- O2
Mg2+ + 2 e- Mg
Aluminium is extracted from its ore, bauxite. This contains aluminium oxide. The
aluminium oxide is dissolved in molten cryolite to lower the temperature needed to melt
When it is electrolysed, molten aluminium forms at the negative electrode (cathode):
Al3+ + 3 e- Al
The Al3+ ions are reduced at the negative electrode (cathode). They receive electrons and form
aluminium atoms. At the positive electrodes (anodes), the oxide ions (O2- ) are oxidized to form
oxygen gas (O2). The oxygen produced reacts with the carbon anodes. This makes carbon
dioxide gas and, in effect, the anodes are burned away. These have to be replaced frequently
in the cells.
In any reaction, reduction is the gain of electrons and oxidation is the loss of electrons.
1. From which ore is aluminium metal extracted?
2. Which aluminium compound is electrolysed in the extraction of aluminium?
O Aluminium chloride,
O Aluminium oxide,
O Aluminium fluoride.
3. Look at the diagram in figure 7 of the cell used in the extraction of aluminium, then match
the letters to the correct labels.
D_________________________ Figure 7
4. Why is aluminium formed at the cathode?
O Because aluminium ions are negatively charged, so they are attracted to the positively
O Because aluminium ions are positively charged, so they are attracted to the negatively
O Because aluminium ions are positively charged, so they are repelled by the positively charged
5. Explain what happens to an aluminium ion (Al3+) when it arrives at the cathode.
THE PURIFICATION OF COPPER
Copper must be very pure when used for making electrical wires. This is done using an
electrolysis cell :
• The ELECTROLYTE is a solution of copper sulphate.
• At the NEGATIVE electrode, copper ions (Cu2+) become copper atoms which are
deposited on the pure copper electrode :
copper ions + electrons copper atoms
• At the POSITIVE electrode, copper atoms in the electrode become
copper ions in solution :
copper atoms - electrons copper ions
• As the electrolysis proceeds, the negative pure copper electrode gets
thicker and the positive electrode gradually dissolves.
• Eventually, the negative electrode has to be replaced with a fresh thin
sheet of pure copper and the process is continued.
• Some of the impurities in the impure copper electrode (silver and gold)
collect as a sludge beneath the positive electrode and can be recovered
to help towards the costs of the process.
Copper can be purified by electrolysis.
Q: Copper metal in the impure anode becomes copper ions Cu2 + . Why do they travel towards
Q: What do the copper ions accept when they reach the cathode?
Q: Write an equation to show this.
Q: Write an equation to show what happens at the anode.
Q: Complete the following paragraph using the words. You can use them once, more than once
or not at all.
Purify, anode, positive, electricity, splitting, electrons, electrolyte, copper, sludge, anode,
electrolysis, copper metal
________ is the _______ of a compound by passing _________ through it. It is used to
________ metals. ______ can be purified in this way. Copper sulphate solution is the
________, which produces _______ ions and sulphate ions. The impure copper is
attached to the _______ electrode, the _______. This produces _______ ions which are
attracted to the negative cathode. Here they each gain _______ to become ________
metal. A _________ from the impure _______ forms underneath the _________.
26th October 2005.