Topic A: Planet Earth

Planet Earth introduces chemistry through the materials all around us: the air we breathe, the sea that covers most of the surface, and the rocks from which we win metals. It sets up core ideas of separating mixtures, writing balanced equations, and using chemical reactions to obtain useful substances from natural resources.

Composition of the atmosphere

Dry air is a mixture of gases, mainly nitrogen (about 78% by volume) and oxygen (about 21%), with the remaining 1% made up of argon, carbon dioxide and traces of other noble gases. Because air is a mixture, its components keep their own properties and can be separated by physical means. The most important industrial method is the fractional distillation of liquid air: air is cooled until it liquefies, then warmed slowly so that nitrogen (boiling point about -196 degrees Celsius), argon and oxygen (about -183 degrees Celsius) boil off at different temperatures and are collected separately. The fraction of carbon dioxide is small but chemically important, as it links the atmosphere to combustion, respiration and photosynthesis.

Oxygen, combustion and the oxygen cycle

Oxygen is the reactive component of air and supports both combustion and respiration. Burning is the rapid reaction of a substance with oxygen, releasing heat; for example carbon burns to give carbon dioxide, C + O2 -> CO2, and magnesium burns to give magnesium oxide, 2 Mg + O2 -> 2 MgO. Oxygen can be prepared in the laboratory by decomposing hydrogen peroxide with a manganese(IV) oxide catalyst, 2 H2O2 -> 2 H2O + O2, and tested with a glowing splint that relights. The oxygen cycle keeps the level of oxygen roughly constant: respiration and combustion remove oxygen and release carbon dioxide, while photosynthesis in green plants does the reverse, 6 CO2 + 6 H2O -> C6H12O6 + 6 O2.

Sea water and its components

Sea water is a complex aqueous solution containing many dissolved ionic salts, the most abundant being sodium chloride. Useful substances can be obtained from it by physical separation. Evaporating sea water in shallow ponds leaves common salt behind, since water is volatile and the salts are not. Pure water can be recovered by distillation, and in some regions desalination provides fresh water. Sea water is also a source of bromine and magnesium compounds. Because it is a mixture, no chemical change is needed to recover dissolved salt; only the input of energy to evaporate the water. The dissolved ions also make sea water a fair conductor of electricity, unlike pure water.

Rocks, minerals and ores

The solid crust is made of rocks, which are mixtures of minerals. A mineral is a naturally occurring compound or element with a definite composition, such as calcium carbonate in limestone and marble. An ore is a rock that contains enough of a metal compound to make extraction worthwhile. Limestone (calcium carbonate, CaCO3) is an especially useful rock: heating it strongly causes thermal decomposition, CaCO3 -> CaO + CO2, producing quicklime (calcium oxide). Calcium carbonate also reacts with dilute acid to give a salt, water and carbon dioxide, which is the standard test for a carbonate. These reactions show how a single abundant mineral can be a raw material for many industrial products.

Reactivity and methods of metal extraction

How a metal is extracted from its ore depends on its position in the reactivity series. Very reactive metals such as potassium, sodium, calcium, magnesium and aluminium hold their ores so strongly that they must be extracted by electrolysis of the molten compound. Moderately reactive metals such as zinc, iron and lead can be extracted by heating their oxides with carbon, which removes the oxygen by reduction; for iron, Fe2O3 + 3 CO -> 2 Fe + 3 CO2 in the blast furnace. Unreactive metals such as silver and gold occur native (uncombined) and need only physical separation. The more reactive the metal, the harder and more energy-demanding the extraction, which is why aluminium is far more expensive to produce than iron.

Extraction of iron in the blast furnace

Iron is extracted by reduction of its oxide ore (haematite, Fe2O3) in a blast furnace charged with ore, coke (carbon) and limestone. Hot air burns the coke to carbon dioxide, which reacts with more hot coke to form carbon monoxide, C + CO2 -> 2 CO; this carbon monoxide is the main reducing agent. Limestone decomposes to calcium oxide, which reacts with sandy impurities (silicon dioxide) to form molten slag, CaO + SiO2 -> CaSiO3. The molten iron sinks to the bottom and is tapped off, with the less dense slag floating on top. The process illustrates reduction, the role of a reducing agent, and the removal of impurities as slag, all central ideas in metal extraction.