Metals react in a predictable order and acids react in predictable patterns. Electrolysis decomposes ionic compounds and is used to extract the most reactive metals.
Metals can be ordered by how readily they lose electrons to form positive ions. Potassium, sodium, lithium, calcium, magnesium, carbon, zinc, iron, hydrogen, copper, silver, gold.
A more reactive metal displaces a less reactive one from its compound, and this is used to establish the order experimentally. Carbon and hydrogen are included because their positions decide how a metal is extracted and whether it reacts with acid.
Most metals occur naturally in compounds such as oxides, sulfides and carbonates. Extracting the metal means reduction: the metal ions gain electrons. Where the ore is an oxide, that can also be described as removing the oxygen.
Metals below carbon can be reduced by heating with carbon, which takes the oxygen. Metals above carbon must be extracted by electrolysis. Gold is unreactive enough to occur naturally as the element rather than only in compounds.
In terms of oxygen, oxidation is gain and reduction is loss. In terms of electrons it is the other way round: oxidation is loss of electrons and reduction is gain.
The mnemonic is OIL RIG. Oxidation Is Loss, Reduction Is Gain. In a displacement reaction the ions of the less reactive metal are reduced, while the more reactive metal is oxidised.
acid + metal → salt + hydrogen. With dilute hydrochloric or sulfuric acid, metals above hydrogen in the reactivity series generally react this way; nitric acid behaves differently. Under the same conditions, the more reactive the metal the faster the fizzing.
This is a redox reaction: the metal loses electrons and is oxidised, while hydrogen ions gain them and are reduced.
Acids react with bases to give a salt and water. acid + metal oxide → salt + water and acid + metal hydroxide → salt + water; with a carbonate, carbon dioxide is also produced.
The acid determines the salt formed: hydrochloric gives chlorides, sulfuric gives sulfates, nitric gives nitrates.
To make a pure dry sample of a soluble salt, add the insoluble base to warm acid until no more dissolves, filter off the excess, then crystallise the filtrate by evaporating some of the water and leaving it to form crystals.
pH is a measure of the concentration of hydrogen ions in solution. Below 7 is acidic, 7 neutral, above 7 alkaline. Acids produce H+ ions in solution; alkalis produce OH− ions, and neutralisation is H+ + OH− → H2O.
The scale is logarithmic: each step of one on the pH scale is a factor of ten in hydrogen ion concentration, so a pH 3 solution has ten times the hydrogen-ion concentration of a pH 4 solution, and a hundred times that of pH 5.
A strong acid is fully ionised in solution: hydrochloric, sulfuric and nitric. A weak acid is only partially ionised: ethanoic, citric and carbonic.
Strong and weak are about the degree of ionisation; concentrated and dilute are about how much acid there is in a given volume. A concentrated weak acid is entirely possible.
A titration finds the volume of one solution that exactly reacts with a known volume of another, and from that the unknown concentration.
Use a pipette for the known solution, a burette for the one being added, and a single indicator with a sharp end point, such as phenolphthalein or methyl orange, not universal indicator, which changes gradually. Repeat until concordant results are obtained and take the mean.
Electrolysis breaks a compound down using electricity. The compound must be molten or dissolved so its ions are free to move.
Positive ions move to the cathode (negative electrode) and gain electrons; negative ions move to the anode (positive electrode) and lose them. For a molten binary compound the products are the metal at the cathode and the non-metal at the anode.
In solution, water also provides H+ and OH− ions, so there is a competition. The rules below are for inert electrodes and the solutions met at GCSE; the electrode material can change the products.
At the cathode, hydrogen is produced unless the metal is less reactive than hydrogen, in which case the metal is deposited. At the anode, oxygen is produced unless a halide is present, in which case the halogen is produced.
A half equation shows what happens at one electrode. For molten sodium chloride, electrons are gained at the cathode: Na+ + e− → Na, and lost at the anode: 2Cl− → Cl2 + 2e−. From an aqueous copper salt the cathode reaction is Cu2+ + 2e− → Cu.
Both charge and atoms must balance.
Metals more reactive than carbon are extracted by electrolysis of the molten compound. Aluminium is obtained from aluminium oxide this way.
The oxide is mixed with cryolite to lower the melting point and save energy, and the oxygen produced at the anode reacts with the hot carbon anodes, mainly forming carbon dioxide, so they wear away and must be replaced. The process is expensive because melting the ore and supplying the current both need large amounts of energy.
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