AQA GCSE Chemistry (8462) · Paper 1
🧪 Chemical Changes
Revision notes written to the specification, with examiner tips and the required practicals. Every point here has flashcards in the Stickwise app.
Reactivity & extraction
In terms of oxygen, oxidation means gaining oxygen, while reduction means losing oxygen. When a metal burns, it is oxidised to form a metal oxide, for example 2Mg + O₂ → 2MgO.
Higher content: oxidation and reduction can also be defined in terms of electrons. Oxidation is the loss of electrons and reduction is the gain of electrons, remembered as OILRIG (Oxidation Is Loss, Reduction Is Gain). A reaction in which both happen at once is a redox reaction.
The reactivity series ranks metals by how readily they form positive ions, which means how readily they lose electrons: potassium, sodium, lithium, calcium, magnesium, (carbon), zinc, iron, (hydrogen), copper, silver, gold. Very reactive metals react with water to form a metal hydroxide and hydrogen, while metals above hydrogen in the series react with acid to form a salt and hydrogen. A more reactive metal will displace a less reactive one from its compound, for example Zn + CuSO₄ → ZnSO₄ + Cu.
The method used to extract a metal depends on its position in the reactivity series. Gold and silver are so unreactive that they are found native, meaning as the metal itself. Metals below carbon in the series, such as iron, zinc and copper, are extracted by smelting: the ore is heated with carbon, which removes the oxygen from the metal oxide in a reduction reaction, and this is a relatively cheap method. Metals above carbon, such as aluminium and magnesium, hold onto their oxygen too strongly for this to work, so they must be extracted by electrolysis, which uses a large amount of energy and is expensive.
For higher tier, in a displacement reaction only the metal ions actually change, while the other ions present take no part in the reaction. These spectator ions stay unchanged and are left out of the ionic equation: Zn + Cu²⁺ → Zn²⁺ + Cu.
Acids & salts
Neutralisation can be written as chemistry's simplest ionic equation: H⁺ + OH⁻ → H₂O.
Making salts
The general word equations for making salts are: acid + metal → salt + hydrogen; acid + metal oxide (or hydroxide) → salt + water; and acid + carbonate → salt + water + CO₂, which causes the fizzing. The acid used determines the name of the salt: hydrochloric acid produces chlorides, sulfuric acid produces sulfates, and nitric acid produces nitrates.
For higher tier, strong acids, such as HCl, H₂SO₄ and HNO₃, fully ionise in water, while weak acids, such as ethanoic, citric and carbonic acid, only partially ionise. This is different from concentrated and dilute, which describe the amount of acid dissolved per volume of water, so a dilute strong acid and a concentrated weak acid can both exist. Each step down the pH scale represents a tenfold increase in the concentration of H⁺ ions.
Electrolysis
Electrolysis uses electricity to split apart an ionic compound. The compound being split, called the electrolyte, must be molten or dissolved so that its ions are free to move. Positive ions, called cations, travel to the negative electrode, the cathode, while negative ions travel to the positive electrode, the anode.
Electrolysis of a molten compound is straightforward: the metal forms at the cathode and the non-metal forms at the anode. Aluminium is extracted this way, from molten aluminium oxide mixed with cryolite, which lowers the melting point and so saves energy. The carbon anodes gradually burn away, because the oxygen produced there reacts with the carbon, so they need to be replaced regularly.
In aqueous solutions, water molecules also take part, so the products are not always as simple. At the cathode, hydrogen forms unless the metal is less reactive than hydrogen, such as copper or silver, in which case the metal is deposited instead. At the anode, oxygen forms unless a halide ion is present, such as Cl⁻, Br⁻ or I⁻, in which case the halogen is produced instead. For example, electrolysis of salty water produces hydrogen and chlorine.
For higher tier, these changes can be written as half equations. At the cathode: 2H⁺ + 2e⁻ → H₂, which is a reduction reaction. At the anode: 2Cl⁻ → Cl₂ + 2e⁻, or, if no halide is present, 4OH⁻ → O₂ + 2H₂O + 4e⁻; both of these are oxidation reactions, since electrons are lost at the anode.