Section 5: Chemistry in Industry - IGCSE Chemistry

Chemistry > Section 5: Chemistry in Industry

a) Extraction and uses of metals

5.1 explain how the methods of extraction of the metals in this section are related to their positions in the reactivity series

Metals can be extracted in three different ways: (i)With Carbon and (ii) By electrolysis & (iii)By more reactive material

Metals that are below zinc in reactivity series can be extracted using carbon and carbon monoxide. Carbon is cheap and can also be as the source of heat.

Metals above zinc in reactivity series are usually extracted by electrolysis. Unfortunately, the large amounts of electricity involved make this an expensive process.

If a metal is more reactive than a metal in a compound, the reactive material will displace it and our targeted metal will be extracted.

5.2 describe and explain the extraction of aluminium from purified aluminium oxide by electrolysis, including:

i. the use of molten cryolite as a solvent and to decrease the required operating temperature

ii. the need to replace the positive electrodes

iii. the cost of the electricity as a major factor

Aluminium is extracted from aluminium oxide which comes from bauxite. Aluminium oxide has very high melting point,

and it isn’t practical toSo aluminiumelectrolyseoxideisdissolved inmoltencryolite toalumini decrease melting point.

The electrolyte is a solution of aluminium oxide in molten cryolite at a temperature of about 1000oC. The molten aluminium is siphoned off from time to time and fresh aluminium oxide is added to the cell. The cell operates at 5-6 volts and with current 100,000amps. The heat generated by the huge current keeps the electrolyte molten. The large amounts of electricity needed are a major expense.


Aluminium ions are attracted to the cathode and are reduced to aluminium by gaining electrons.

Al3+(l) + 3e- ==> Al(l)

The molten aluminium produced sinks to the bottom of the cell.


The oxide ions are attracted to the anode and lose electrons to from oxygen gas.

2O2-(l) ==> O2(g) + 4e-

This creates a problem. Because of the high temperatures, the carbon anodes burn in the oxygen to from carbon dioxide. The anodes have to be replaced regularly, and this also adds to the expense of the process.

5.3 write ionic half-equations for the reactions at the electrodes in aluminium extraction

At the anode: 2O2-(l) ==> O2(g) + 4e-

At the cathode: Al3+(l) + 3e- ==> Al(l)

5.4 describe and explain the main reactions involved in the extraction of iron from iron ore (haematite), using coke, limestone and air in a blast furnace

Coke is used as the starting material. It is an impure carbon and it burns in how air blast to form carbon dioxide. This is strongly exothermic reaction.

C(s) + O2(g) ==> CO2(g)

At the high temperatures in the furnace, the carbon dioxide is reduced by more carbon to give carbon monoxide.

CO2(g) + C(s) ==> 2CO(g)

Carbon monoxide is the reducing agent. Iron is gained from haematite(Fe2O3)

Fe2O3(s) + 3CO(g) ==> 2Fe(l) +3CO2(g)

The iron melts and flows to the bottom of the furnace, where it can be tapped off.

In the hotter parts of the furnace, some of the iron oxide is also reduced by carbon itself.

Fe2O3(s) + 3C(s) ==> 2Fe(l) + 3CO(g)

Limestone is added to the furnace to remove impurities in the ore. Limestone thermally decomposed to calcium oxide and carbon dioxide. It is an endothermic reaction.

CaCO3(s) ==> CaO(s) + CO2(g)

Silicon dioxide occurs naturally as quartz and it is a form of impurities that needs to be removed. Calcium oxide react with it to form calcium silicate. This melts and trickles to the bottom of the furnace as a molten slag, which floats on top of the molten iron and can be tapped off separately.

CaO(s) + SiO2(s) ==> CaSiO3(l)

5.5 explain the uses of aluminium and iron, in terms of their properties.

Uses of aluminium:

Pure aluminium isn’t very strong, so aluminiumrrosionalloy and strong, it has various uses. Like it is used in aero planes, pans etc. For its good conductivity of electricity, it is used as cables.

Uses of iron:

Types of iron

Iron mixed with


Some uses

Wrought iron

(pure iron)


Decorative work such as gates and railings

Mild steel

Up to 0.25% carbon


Nails, car bodies, ship building, girders

High-carbon steel

0.25-1.5% carbon

Very hard, sometimes brittle

Cutting tools, masonry nails

Cast iron

About 4% carbon

Hard but brittle

Manhole covers, guttering, engine blocks

Stainless steel

Chromium and nickel

Resistant to corrosion

Cutlery, cooking utensils, kitchen sinks

b) Crude Oil

5.6 understand that crude oil is a mixture of hydrocarbons

Crude oil is a mixture of hydrocarbons - compounds containing carbon and hydrogen only.

5.7 describe and explain how the industrial process of fractional distillation separates crude oil into fractions

The process of refining involves separating the hydrocarbons into fractions or batches using a technique called fractional distillation. Each fraction separates as they have different boiling points. The crude oil is heated in a furnace to around 400oC. This allows all of the hydrocarbons in the crude oil to move into the bottom of the fractionating tower. The tower is hottest at the bottom and coolest at the top

The smallest molecules contained in the crude oil have lower boiling points and so move to the top of the tower. This is because the forces between these molecules are weak, so little energy is required to vaporise them. Larger molecules will remain lower down the tower as they have higher boiling points. This is because forces between the molecules are stronger.

5.8 recall the names and uses of the main fractions obtained from crude oil: refinery gases, gasoline, kerosene, diesel, fuel oil and bitumen

Refinery gases: refinery gases are a mixture of methane, ethane, propane and butane, which can be separated into individual gases if required. These gases are commonly used as LPG (liquefied petroleum gas) for domestic heating and cooking.

Gasoline (petrol): As with all other fractions, petrol is a mixture of hydrocarbons with similar boiling points. It is used in cars mainly.

Kerosine: Kerosine is used as fuel for jet aircraft, as domestic heating oil and as as 'paraffin' for small heaters and lamps.

Diesel oil (gas oil): This is used for buses, Iorries, some cars, and railway engines where the line hasn't been electrified. Some is also cracked to make other organic chemicals and produced more petrol

Fuel oil: This is used for ships’ boilers and for industrial heating.

Bitumen: Bitumen is a thick black material, which is melted and mixed with rock chippings to make the top surfaces of roads.

5.9 describe the trend in boiling point and viscosity of the main fractions

As the molecules get bigger, the following changes occur:

* boiling point increases

* The liquids become less volatile.

* The liquid flow less easily (they become more vicous)

* Bigger hydrocarbons do not burn as easily as smaller ones.

5.10 understand that incomplete combustion of fuels may produce carbon monoxide and explain that carbon monoxide is poisonous because it reduces the capacity of the blood to carry oxygen

If therenough air(or isn't enough oxygen), you get incomplete combustion. This leads to the formation of carbon or carbon monoxide instead of carbon dioxide.

2CH4(g) + 3O2(g) ==> 2CO(g) + 4H2O(l)

Carbon monoxide is colourless, odourless and is very poisonous. Carbon monoxide is poisonous because it combines with hemoglobin, preventing it from carrying oxygen. People can be made ill or even die, because of lack of oxygen in your body.

5.11 understand that, in car engines, the temperature reached is high enough to allow nitrogen and oxygen from air to react, forming nitrogen oxides

In car engines there is a high enough temperature to cause a reaction between oxygen and nitrogen in the air.

N2(g) + O2(g) ==> NO(g)

5.12 understand that nitrogen oxides and sulfur dioxide are pollutant gases which contribute to acid rain, and describe the problems caused by acid rain

Acid rain is formed when acidic air pollutants such as sulphur dioxide and nitrogen dissolve in rainwater. Sulphur dioxide dissolves in water to form sulphurous acid (H2SO3).

SO2(g) + H2O(l) ==> H2SO3(aq)

In the presence of oxygen in the air, the acid is slowly oxidized to sulphuric acid (H2SO4).

Oxides of nitrogen also contribute to acid rain. In the presence of oxygen and water, nitrogen dioxide is converted to nitric acid.

4NO2(g) + 2H2O(l) + O2(g) ==> 4HNO3(aq)

Carbon dioxide in the air dissolves in rainwater to form carbonic acid, which is a weak acid.

CO2(g) + H2O(l) ==> H2CO3(aq)

The pH value of normal rain is slightly below 7. The pH value of acid rain is approximately 3.4.

Problems of acid rain:

Acid rain reacts with metals and with carbonates in marble and limestone. When this happens, metal bridges and stone buildings are damaged.

It can reduce the pH value of natural water bodies from 6.5 and 8.5 to below 4. This will kill fish and other aquatic life.

It also leaches important nutrients from the soil and destroys plants. Without these nutrients, plant growth is stunted. In some cases, acid rain dissolves aluminium hydroxide in the soil to produce aluminium ions, which are toxic to plants.

5.13 understand that fractional distillation of crude oil produces more long-chain hydrocarbons than can be used directly and fewer short-chain hydrocarbons than required and explain why this makes cracking necessary

The amounts of each fraction you get will depend on the proportions of the various hydrocarbons in the original crude oil, not in the amount in which they are needed.

Long chain hydrocarbons which can't be used directly are become less flammable, more viscous and therefore less useful.

Short chain hydrocarbons burn well and flow well. Therefore, they are useful but these are produced less in the fractional distillation of crude oil. Thus chemists convert these large, less useful, heavy fractions into smaller, more useful ones by means of cracking.

5.14 describe how long-chain alkanes are converted to alkenes and shorter-chain alkanes by catalytic cracking, using silica or alumina as the catalyst and a temperature in the range of 600-700

o C.

The gas oil fraction is heated to give a gas and then passed over a catalyst of mixed silicon dioxide and aluminium oxide at about 600-700o C. Cracking can also be carried out at higher temperature without a catalyst.

Cracking is simply splitting of larger molecules to simpler ones. The molecules are broken up in an random way which produce a mixture of alkanes and alkenes.

c) Synthetic polymers

5.15 understand that an addition polymer is formed by joining up many small molecules called monomers

Polymerisation: The process of joining together a large number of small molecules (monomers) to from a macromolecule is called polymerisation.

Monomer : Monomers are small units which are joined to form polymer.

Polymer : When repeating monomers are joined together by polymerisation, them form a macromolecule called polymer.

There are two basic types of reactions for forming polymers:

i. Addition polymerisation

ii. Condensation polymerisation

Addition polymerisation occurs when monomer units join together without losing any molecules or atoms.

5.16 draw the repeat unit of addition polymers, including poly(ethene), poly(propene) and poly(chloroethene)

5.17 deduce the structure of a monomer from the repeat unit of an addition polymer

i. Indentify the repeat unit in the polymer.

ii. Write down the formula of the repeat unit.

iii. Convert the carbon-carbon single bond into a carbon-carbon double bond.

5.18 describe some uses for polymers, including poly(ethene), poly(propene) and poly(chloroethene)

Poly(ethene): Low density poly(ethene) is mainly used to make polythene bags. It is very flexible but not very strong.

High-density poly(ethene) is used where rather greater strength and rigidity is needed - for example to make plastic bottles.

Poly(propene): it is used to make ropes and crates.

Poly(chloroethene): It is quite strong and rigid and so can be used for drainpipes, or replacement windows. It can also be made flexible by adding 'plasticisers'. That makes it useful for sheet floor coverings, and even clothing. These polymers don't conduct electricity and PVC is used for electrical insulation.

5.19 explain that addition polymers are hard to dispose of as their inertness means that they do not easily biodegrade

Addition polymers are unreactive. So they don't easily biodegrade.

5.20 understand that some polymers, such as nylon, form by a different process called condensation polymerisation

Some polymers are made by reacting two different types of monomers. Each of the monomers involved has a functional group at each end of the molecule. When these monomers react, a polymer is produced, and a small molecule such as water is also produced as a by-product of the reaction. This type of reaction is called condensation polymerisation.

5.21 understand that condensation polymerisation produces a small molecule, such as water, as well as the polymer.

(Follow 5.20)

d) Industrial manufacture of chemicals

5.22 understand that nitrogen from air, and hydrogen from natural gas or the cracking of hydrocarbons, are used in the manufacture of ammonia

Ammonia (NH3) is manufactured by using nitrogen from air and hydrogen from natural gas.

N2(g) + H2(g) ==> NH3 (g)

5.23 describe the manufacture of ammonia by the Haber process, including the essential conditions: i a temperature of about 450°C

ii a pressure of about 200 atmospheres

iii an iron catalyst

Ammonia is manufactured by combining nitrogen and hydrogen in an important industrial process called the Haber process.

nitrogen + hydrogen symbol ammonia

N 2 (g) + 3H 2 (g) symbol 2NH 3 (g)

The Haber process:

  • Raw materials: nitrogen (from the air) & hydrogen :(made from natural gas)
  • The proportions: 1 volume of nitrogen to 3 volumes of hydrogen
  • The temperature: 450oC
  • The pressure: 200 atmospheres
  • The catalyst: iron

Each time the gases pass through the reaction vessel, only about 15% of the nitrogen and hydrogen combine to make ammonia. The reaction mixture is cooled and the ammonia condenses as a liquid. The unreacted nitrogen and hydrogen can simply be recycled through the reactor.

5.24 understand how the cooling of the reaction mixture liquefies the ammonia produced and allows the unused hydrogen and nitrogen to be recirculated

The products from the reactant are sent through a cooling mechanism, this is at a temperature that condenses ammonia, but not hydrogen and nitrogen. Liquid ammonia is then collected but hydrogen and nitrogen float right back into the reactor.

5.25 describe the use of ammonia in the manufacture of nitric acid and fertilisers

Ammonia is produced in large quantities because it has many important uses.

Manufacture of Nitric acid:

Industrially, nitric acid is made by the catalytic oxidation of ammonia over heated platinum. Oxidising ammonia produces oxides of nitrogen which can then be dissolved in water to produce nitric acid.

Initially nitrogen(II) oxide will be formed from the catalytic oxidation of ammonia using the transition metal platinum. 4NH3 (g) + 5O3 (g) ==> 4NO(g) + 6H2O (g)

The nitrogen(II) oxide is rapidly cooled before combining with oxygen (from excess air) to form nitrogen(IV) oxide. 2NO(g) + O2(g) ⇌ 2NO2 (g)

The nitrogen(IV) oxide is now allowed to react with water to form nitric acid. 4NO2 (g) + O2 (g) + 2H2O (l) ==> 4HNO3 (aq)

Most of the nitric acid made is used to make the all-important fertilisers such as ammonium nitrate.

Manufacture of Fertilizers:

The main use of ammonia is in the manufacture of fertilizers. Approximately 75% of all ammonia produced is converted into various ammonium compounds like ammonia sulphate, ammonium nitrate and urea (NH2CONH2). These compounds are called nitrogenous fertilizers. They are solids for ease in handling and water soluble so that they seep into the soil to be absorbed by the roots of the plant.

5.26 recall the raw materials used in the manufacture of sulfuric acid

Sulphur: (sulphur is found in rocks and some natural gasses)

Oxygen: from the air.

5.27 describe the manufacture of sulfuric acid by the contact process, including the essential conditions:

i. a temperature of about 450°C

ii. a pressure of about 2 atmospheres

iii. a vanadium(V) oxide catalyst

I) First Sulphur dioxide is produce:

Either burn sulfur in air:

S(s) + O2(g) ==> SO2(g)

Or heat sulphide ores strongly in air:

4FeS2(s) + 11O2(g) ==> 2Fe2O3(s) +8SO2(g)

II) Sulphur trioxide is produced:

Now sulfur dioxide is converted into sulfur trioxide using an excess of air from the previous processes.

2SO2(g) + O2(g) ⇌ 2SO3(g)

Because of the forward reaction is exothermic, there would be a higher percentage conversion of sulfur dioxide into sulfur trioxide at a low temperature. However, at a low temperature the rate of reaction would be very slow. 450oC is a compromise.

There are 3 gas molecules on the left-hand side of the equation, but only 2 on the right. Reactions in which the numbers of gas molecules decrease are favored by high pressures. In this case, though, the conversion is good at low pressures

that it isn’t economically worthwhile to use higher

The catalyst, vanadium(V) oxide, has no effect on the percentage conversion, but helps to speed up the reaction. Without the catalyst, the reaction would be extremely slow.

III) Making the sulphuric acid

In principal you can react sulfur trioxide with water to make sulphuric acid. In practice, this produce an uncontrolled fog of concentrated sulfuric acid. Instead the sulfur trioxide is absorbed in concentrated sulfuric acid to give fuming sulfuric acid.

H2SO4(l) + SO3(g) ==> H2S2O7(l)

This is converted into twice as much concentrated sulfuric acid by careful addition of water.

H2S2O7(l) + H2O(l) ==> 2H2SO4(l)

5.28 describe the use of sulfuric acid in the manufacture of detergents, fertilisers and paints

  • the production of fertilizers such as ammonium sulphate, potassium sulphate, calcium superphosphate, etc.
  • the manufacture of non-soapy detergents: modern detergents are organic compounds ‘ concentrated sulphuric acid.
  • In paint manufacture, sulfuric acid is used in extracting the white pigment titanium dioxide, TiO2, from titanium ores.
  • The making of artificial silks like rayon
  • The cleaning of metals by removing the surface oxide coating.
  • It's used as an electrolyte inside batteries for cars

5.29 describe the manufacture of sodium hydroxide and chlorine by the electrolysis of concentrated sodium chloride solution (brine) in a diaphragm cell

The concentrated salt solution can be electrolysed to produce three useful chemicals - sodium hydroxide, chlorine and hydrogen. The electrolysis can be carried out in a diaphragm cell.

At the titanium anode, chloride ions are discharged to produce chlorine gas. 2Cl-(aq) ==> Cl2(g) + 2e-

At the steel cathode, it is too difficult sodium ions, so hydrogen ions from the water are discharged instead to produce hydrogen gas.

2H+(aq) + 2e- ==> H2(g)

More and more water keeps splitting up to replace the hydrogen ions as soon as they discharged. Each time a water molecule splits up it produces a hydroxide ion as well. That means that there will be a build-up of sodium ions and hydroxide ions in the right-hand compartment-sodium hydroxide solution is formed. This is contaminated with uncharged sodium chloride.

The sodium hydroxide solution is concentrated by evaporating it. During this process, most of the sodium chloride crystallizes out as solid salt. This can be separated, redissolved in water and recycled back through the cell again.


What happens if chlorine comes into contact with sodium hydroxide or hydrogen?

The cell is designed to keep the products apart. If chlorine comes into contact with sodium hydroxide solution, it reacts to make bleach - a mixture of sodium chloride and sodium chlorate(I) solution. If chlorine comes into contact with hydrogen it produces a mixture which would explode violently on exposure to sunlight or heat to give hydrogen chloride.

5.30 write ionic half-equations for the reactions at the electrodes in the diaphragm cell

(Follow 5.29)

5.31 describe important uses of sodium hydroxide, including the manufacture of bleach, paper and soap; and of chlorine, including sterilising water supplies and in the manufacture of bleach and hydrochloric acid.

Uses of sodium hydroxide include:

  • The purification of bauxite to make aluminium oxi
  • Paper making - the sodium hydroxide helps break the wood down into pulp.
  • Soap making - sodium hydroxide reacts with animal and vegetable fats and oils to make compounds, such as sodium stearate, that are present in soap.
  • Making bleach - bleach is formed when sodium hydroxide and chlorine react together in the cold; it is a mixture of sodium chloride and sodium chlorate(I) solution. 2NaOH(aq) + Cl2(g) ==> NaCl(aq) + NaOCl(aq) + H2O(l)

Uses of chlorine include:

  • Sterilising water to make it safe to drink
  • Making hydrochloric acid (by controlled reaction with hydrogen)
  • Making bleach