Topic 4: Chemical changes

Chemical changes rearrange atoms to form new substances, and many of the most useful reactions involve metals, acids and electricity. This topic explains why some metals are more reactive than others, how we get pure metals out of their ores, and how acids react with metals, bases and carbonates to make salts. It finishes with the pH scale and electrolysis, two ideas that link reactivity to everyday chemistry and industry.

Metal oxides and the reactivity series

When a metal reacts with oxygen it forms a metal oxide; this is an example of oxidation, because oxidation is gain of oxygen. The reactivity series is a league table that lists metals in order of how vigorously they react, with the most reactive at the top. A common version, from most to least reactive, runs: potassium, sodium, lithium, calcium, magnesium, carbon, zinc, iron, hydrogen, copper. Carbon and hydrogen are non-metals but are placed in the list because we compare metals against them. A more reactive metal will displace a less reactive metal from a compound, which is the basis of displacement reactions. The reactivity of a metal is linked to how readily it forms positive ions by losing electrons: the more easily a metal loses electrons, the more reactive it is.

Reactions of metals with water and acids

Metals can be ranked by watching how they react with water and with dilute acids. Very reactive metals such as potassium, sodium, lithium and calcium react with cold water to give a metal hydroxide and hydrogen gas, for example sodium plus water gives sodium hydroxide and hydrogen. Less reactive metals such as magnesium, zinc and iron do not react much with cold water but do react with dilute acids to give a salt and hydrogen, for example magnesium plus hydrochloric acid gives magnesium chloride plus hydrogen. The vigour of bubbling (how fast hydrogen is produced) tells you the order of reactivity. Metals below hydrogen, such as copper, do not react with dilute acids at all. You can test the hydrogen produced with a lit splint, which gives a squeaky pop.

Extraction of metals by reduction with carbon

Most metals are found in the Earth combined with other elements in compounds called ores, so they must be extracted. The method used depends on the metal's position in the reactivity series. Metals less reactive than carbon, such as iron, zinc, tin and copper, can be extracted by reduction with carbon. Reduction is the loss of oxygen, so when carbon removes oxygen from a metal oxide the metal oxide is reduced. For example, iron(III) oxide is heated with carbon to produce iron and carbon dioxide; the carbon is oxidised because it gains oxygen. Metals more reactive than carbon, such as aluminium, cannot be extracted this way because carbon is not reactive enough to take their oxygen, so they are extracted using electrolysis instead. Unreactive metals like gold are found as the metal itself (native) and need no chemical extraction.

Oxidation and reduction in terms of electrons (Higher)

At Higher tier, oxidation and reduction are defined in terms of electrons rather than oxygen. Oxidation Is Loss of electrons and Reduction Is Gain of electrons, summarised by the mnemonic OIL RIG. A reaction in which one species is oxidised and another is reduced at the same time is called a redox reaction. In a displacement reaction, the more reactive metal is oxidised as it loses electrons to form positive ions, while the ions of the less reactive metal are reduced as they gain those electrons to become metal atoms. For example, when magnesium displaces copper from copper sulfate solution, magnesium atoms lose two electrons to become Mg2+ ions, and Cu2+ ions gain two electrons to become copper atoms. Writing ionic and half equations helps show clearly which species loses and which gains electrons.

Reactions of acids: metals, bases and carbonates

Acids take part in three key reaction types that all produce a salt. Acid plus metal gives a salt plus hydrogen. Acid plus a metal oxide or metal hydroxide (a base) gives a salt plus water. Acid plus a metal carbonate gives a salt plus water plus carbon dioxide, and the carbon dioxide can be tested by bubbling it through limewater, which turns cloudy. The name of the salt depends on the acid used: hydrochloric acid makes chlorides, sulfuric acid makes sulfates, and nitric acid makes nitrates. For example, copper oxide plus sulfuric acid gives copper sulfate plus water. A base is any substance that neutralises an acid; an alkali is a base that dissolves in water to give hydroxide ions.

Neutralisation and making soluble salts

Neutralisation is the reaction between an acid and a base in which hydrogen ions react with hydroxide ions to make water, shown by the ionic equation H+ plus OH- gives H2O. The salt forms from the leftover ions. To make a pure, dry sample of a soluble salt you react an acid with an excess of an insoluble base such as a metal oxide, carbonate or hydroxide. Warming the acid speeds the reaction, and adding the solid in excess ensures all the acid is used up. The leftover excess solid is removed by filtration, leaving a salt solution. The solution is then heated to evaporate some water and left to crystallise, and the crystals are dried. This produces hydrated salt crystals without any unreacted acid remaining.

The pH scale and indicators

The pH scale runs from 0 to 14 and measures how acidic or alkaline a solution is. A pH below 7 is acidic, a pH of 7 is neutral, and a pH above 7 is alkaline. The lower the pH the more acidic the solution; the higher the pH the more alkaline. Acidity is caused by H+ ions in solution, while alkalinity is caused by OH- ions. Indicators are dyes that change colour depending on pH. Universal indicator gives a range of colours and an approximate pH value, turning red in strong acid, green at neutral and purple in strong alkali. Litmus is red in acid and blue in alkali. For an accurate, numerical pH value a pH probe connected to a meter is used instead of an indicator.

Strong and weak acids (Higher)

At Higher tier, acids are classed as strong or weak. A strong acid is fully ionised in aqueous solution, meaning nearly all its molecules split up to release H+ ions; examples include hydrochloric, sulfuric and nitric acids. A weak acid is only partially ionised, so only a small fraction of its molecules release H+ ions at any moment; examples include ethanoic, citric and carbonic acids, and this ionisation is reversible. For a given concentration, a strong acid has a higher H+ concentration and therefore a lower pH than a weak acid. As the H+ concentration increases by a factor of ten, the pH decreases by one unit. Be careful to separate strength (degree of ionisation) from concentration (amount of acid dissolved in a given volume).

Electrolysis of molten compounds and aqueous solutions

Electrolysis uses an electric current to break down an ionic compound, called the electrolyte, into elements. The ions must be free to move, so the compound is either melted (molten) or dissolved in water (aqueous). Positive ions move to the negative electrode (the cathode) and negative ions move to the positive electrode (the anode). When a molten ionic compound is electrolysed, the metal forms at the cathode and the non-metal forms at the anode, for example molten lead bromide gives lead and bromine. Aluminium is extracted by electrolysing molten aluminium oxide mixed with cryolite to lower the melting point. In aqueous solutions, water also provides H+ and OH- ions, so the products can differ: hydrogen is produced at the cathode if the metal is more reactive than hydrogen, and oxygen is usually produced at the anode unless the solution contains halide ions, in which case the halogen forms.

Half equations at the electrodes (Higher)

At Higher tier you can represent each electrode reaction with a half equation that shows the transfer of electrons. At the cathode, positive ions gain electrons (reduction), for example Pb2+ plus 2e- gives Pb. At the anode, negative ions lose electrons (oxidation), for example 2Br- gives Br2 plus 2e-. The electrons appear on the left for reduction at the cathode and on the right for oxidation at the anode. For the electrolysis of aqueous sodium chloride, hydrogen forms at the cathode by 2H+ plus 2e- gives H2, and chlorine forms at the anode by 2Cl- gives Cl2 plus 2e-, leaving sodium hydroxide in solution. Balancing the number of electrons lost and gained confirms the overall reaction is consistent.