AQA GCSE Chemistry 8462: Chemical changes – Study Guide

This guide teaches section 4.4 Chemical changes of the AQA GCSE Chemistry (8462) specification, for teaching from September 2016 and exams from June 2018 onwards. It covers every point from 4.4.1.1 to 4.4.3.5: the reactivity of metals, reactions of acids and electrolysis. The topic is assessed on Paper 1 (topics 1 to 5), set at Foundation and Higher Tier. Content the specification marks (HT only) is labelled Higher tier only. Required practicals 1, 2 and 3 belong to this topic.

For quick recall, use the Chemical changes revision notes. To test yourself, use the Chemical changes practice questions. The course hub is AQA GCSE Chemistry, and the printable checklist lists every specification point. Titration calculations use the mole ideas from Quantitative chemistry.

What this unit covers

Spec point	What you must be able to do	Tier
4.4.1.1 Metal oxides	Explain oxidation and reduction as gain or loss of oxygen	Both
4.4.1.2 The reactivity series	Describe reactions of K, Na, Li, Ca, Mg, Zn, Fe, Cu with water or dilute acid; link reactivity to forming positive ions; deduce an order from results; displacement	Both
4.4.1.3 Extraction of metals and reduction	Explain reduction with carbon; interpret extraction data; identify what is oxidised or reduced	Both
4.4.1.4 Oxidation and reduction in terms of electrons	Write ionic equations for displacement; identify species oxidised and reduced	Higher tier only
4.4.2.1 Acids with metals	Acid + metal → salt + hydrogen; explain as redox in terms of electrons (HT)	Both / HT
4.4.2.2 Neutralisation and salts	Predict products; deduce salt formulae from ions	Both
4.4.2.3 Soluble salts	Describe making a pure, dry salt (required practical 1)	Both
4.4.2.4 pH and neutralisation	pH scale, indicators, H⁺ + OH⁻ → H₂O	Both
4.4.2.5 Titrations	Describe a titration (required practical 2); calculate concentrations (HT)	Both / HT
4.4.2.6 Strong and weak acids	Ionisation, dilute vs concentrated, factor of 10 per pH unit	Higher tier only
4.4.3.1–4.4.3.4 Electrolysis	Process; molten compounds; aluminium extraction; aqueous solutions (required practical 3)	Both
4.4.3.5 Half equations	Write and balance half equations at electrodes	Higher tier only

4.4.1.1 Metal oxides: oxidation and reduction

Metals react with oxygen to make metal oxides. The metal gains oxygen, so this is oxidation. Reduction is the loss of oxygen. For example, 2Mg + O₂ → 2MgO: magnesium is oxidised.

4.4.1.2 The reactivity series

When metals react, their atoms form positive ions. The more easily a metal forms its positive ion, the more reactive it is. The non-metals hydrogen and carbon are often included for comparison.

Metal	With water (room temperature)	With dilute acid
Potassium	Very vigorous; fizzes, melts, moves about, lilac flame	Too dangerous to try
Sodium	Vigorous; fizzes, melts into a ball, moves about	Too dangerous to try
Lithium	Fizzes steadily, floats	Too dangerous to try
Calcium	Fizzes; cloudy solution forms	Very vigorous
Magnesium	Very slow, a few bubbles	Vigorous fizzing
Zinc	No visible reaction	Steady fizzing
Iron	No visible reaction	Slow fizzing
Copper	No reaction	No reaction

Order: potassium > sodium > lithium > calcium > magnesium > (carbon) > zinc > iron > (hydrogen) > copper.

Metals react with water to give a metal hydroxide and hydrogen, and with acids to give a salt and hydrogen. Room temperature only; reactions with steam are not needed.

Displacement: a more reactive metal displaces a less reactive metal from a compound. Magnesium in iron(II) sulfate solution becomes coated with iron; copper does nothing.

Worked example 1: deducing an order

Metals P, Q and R are added to dilute hydrochloric acid. P fizzes quickly. Q gives no bubbles. R fizzes slowly. R also displaces Q from a solution of Q sulfate.

1. Fizzing rate: P is more reactive than R, and R more than Q.
2. The displacement agrees: R is above Q.
3. Order: P > R > Q.

4.4.1.3 Extraction of metals and reduction

Unreactive metals such as gold are found in the Earth as the metal itself. Most metals are found as compounds and must be extracted by a chemical reaction. Metals less reactive than carbon can be extracted from their oxides by reduction with carbon. Details of industrial processes are not required, but you may be given information about an extraction to interpret or evaluate.

Worked example 2: oxidised or reduced?

2CuO + C → 2Cu + CO₂

1. Copper oxide loses oxygen to become copper: copper oxide is reduced.
2. Carbon gains oxygen to become carbon dioxide: carbon is oxidised.

4.4.1.4 Oxidation and reduction in terms of electrons (Higher tier only)

Oxidation is the loss of electrons. Reduction is the gain of electrons. A useful memory aid is OIL RIG.

In a displacement reaction the ions that do not change are spectator ions; leave them out of the ionic equation.

Worked example 3: ionic equation for displacement

Magnesium reacts with iron(II) sulfate solution: Mg + FeSO₄ → MgSO₄ + Fe

1. The sulfate ions are unchanged, so remove them.
2. Ionic equation: Mg + Fe²⁺ → Mg²⁺ + Fe
3. Half equations: Mg → Mg²⁺ + 2e⁻ (loses electrons, oxidised); Fe²⁺ + 2e⁻ → Fe (gains electrons, reduced).

4.4.2.1 Reactions of acids with metals

Acids react with some metals to produce a salt and hydrogen. You need the reactions of magnesium, zinc and iron with hydrochloric and sulfuric acids, for example Zn + H₂SO₄ → ZnSO₄ + H₂.

Higher tier only: these are redox reactions. In Zn + 2H⁺ → Zn²⁺ + H₂, zinc atoms lose electrons (oxidised) and hydrogen ions gain electrons (reduced).

4.4.2.2 Neutralisation and salt production

• acid + alkali (soluble metal hydroxide) → salt + water
• acid + base (insoluble metal hydroxide or metal oxide) → salt + water
• acid + metal carbonate → salt + water + carbon dioxide

The salt depends on the acid (hydrochloric → chlorides, nitric → nitrates, sulfuric → sulfates) and on the positive ion in the base, alkali or carbonate.

Worked example 4: salt formulae from ions

Balance the charges so the total is zero.

Salt	Ions	Formula
Magnesium nitrate	Mg²⁺, NO₃⁻	Mg(NO₃)₂
Potassium sulfate	K⁺, SO₄²⁻	K₂SO₄
Calcium chloride	Ca²⁺, Cl⁻	CaCl₂

4.4.2.3 Soluble salts (required practical 1)

A soluble salt can be made from an acid and a solid insoluble metal, metal oxide, hydroxide or carbonate. Method for copper sulfate crystals from copper oxide and dilute sulfuric acid:

1. Warm the dilute acid gently in a beaker using a Bunsen burner.
2. Add copper oxide a little at a time, stirring, until no more reacts and some black solid is left. The excess makes sure all the acid is used up.
3. Filter to remove the excess solid. The filtrate is copper sulfate solution.
4. Heat the solution in an evaporating basin on a water bath or electric heater to evaporate some water.
5. Leave it to cool so crystals form, then filter them off and dry them between filter paper.

4.4.2.4 The pH scale and neutralisation

Acids produce hydrogen ions, H⁺, in aqueous solution. Alkalis produce hydroxide ions, OH⁻. The pH scale runs from 0 to 14: 7 is neutral, below 7 acidic, above 7 alkaline. Measure pH with universal indicator (or a wide range indicator) by matching the colour to a chart, or with a pH probe.

In neutralisation between an acid and an alkali: H⁺(aq) + OH⁻(aq) → H₂O(l)

4.4.2.5 Titrations (required practical 2)

A titration measures the volumes of acid and alkali that react exactly, using a suitable indicator. Use strong acids and alkalis only (sulfuric, hydrochloric and nitric acids).

1. Use a pipette to put 25.0 cm³ of alkali into a conical flask. Add a few drops of indicator.
2. Fill a burette with acid and record the starting reading.
3. Add acid while swirling. Near the end point, add it drop by drop until the indicator changes colour.
4. Record the final reading. The titre is final minus start.
5. Repeat until you have close (concordant) results and find the mean.

Worked example 5: concentration from a titration (Higher tier only)

25.0 cm³ of sodium hydroxide solution is neutralised by 20.00 cm³ of 0.150 mol/dm³ sulfuric acid. H₂SO₄ + 2NaOH → Na₂SO₄ + 2H₂O

moles H2SO4 = 0.150 × 20.00 ÷ 1000 = 0.00300 mol
moles NaOH  = 2 × 0.00300          = 0.00600 mol
[NaOH]      = 0.00600 ÷ 0.0250     = 0.240 mol/dm³
in g/dm³    = 0.240 × 40           = 9.60 g/dm³

4.4.2.6 Strong and weak acids (Higher tier only)

• A strong acid is completely ionised in aqueous solution: hydrochloric, nitric and sulfuric acids.
• A weak acid is only partially ionised: ethanoic, citric and carbonic acids.
• Strong/weak describes the degree of ionisation. Dilute/concentrated describes the amount of acid in a given volume. A dilute strong acid and a concentrated weak acid are both possible.
• For the same concentration, the stronger the acid, the lower the pH.
• When pH falls by one unit, the hydrogen ion concentration rises by a factor of 10. From pH 5 to pH 2 is three units, so [H⁺] increases by 10 × 10 × 10 = 1000 times.

4.4.3 Electrolysis

4.4.3.1 The process

When an ionic compound is melted or dissolved in water, its ions are free to move. The liquid or solution conducts electricity and is called an electrolyte. Positive ions move to the negative electrode (cathode); negative ions move to the positive electrode (anode). The ions are discharged and form elements.

4.4.3.2 Molten ionic compounds

With inert electrodes, a molten binary compound gives the metal at the cathode and the non-metal at the anode. Molten lead bromide gives lead and bromine. Molten zinc chloride is a safer alternative in practical work.

4.4.3.3 Extracting metals by electrolysis

Electrolysis is used when a metal is too reactive to be reduced by carbon, or reacts with carbon. It uses large amounts of energy to melt the compounds and to produce the current. Aluminium is made by electrolysing a molten mixture of aluminium oxide and cryolite, with carbon anodes.

• The mixture melts at a lower temperature than aluminium oxide alone, so less energy is needed.
• Oxygen forms at the carbon anode and reacts with it to make carbon dioxide, so the anode must be replaced.

4.4.3.4 Aqueous solutions (required practical 3)

Water molecules break down into H⁺ and OH⁻ ions, which can also be discharged. With inert electrodes:

• Cathode: hydrogen forms if the metal is more reactive than hydrogen; otherwise the metal forms.
• Anode: oxygen forms, unless halide ions are present, when the halogen forms.

Worked example 6: predicting products

Solution	Cathode	Anode
Sodium chloride	Hydrogen (Na more reactive than H)	Chlorine (chloride present)
Copper sulfate	Copper (Cu less reactive than H)	Oxygen (no halide)
Potassium bromide	Hydrogen	Bromine

Required practical 3 is an investigation of this, starting from a hypothesis.

4.4.3.5 Half equations (Higher tier only)

At the cathode, positive ions gain electrons: reduction. At the anode, negative ions lose electrons: oxidation. Examples: 2H⁺ + 2e⁻ → H₂ and 4OH⁻ → O₂ + 2H₂O + 4e⁻ (also written 4OH⁻ − 4e⁻ → O₂ + 2H₂O).

Worked example 7: molten lead bromide

1. Cathode: Pb²⁺ + 2e⁻ → Pb (reduction).
2. Anode: 2Br⁻ → Br₂ + 2e⁻ (oxidation).
3. Check that charge balances in each half equation.

Common errors

• Mixing up which substance is oxidised in a carbon reduction: the metal oxide is reduced, carbon is oxidised.
• Writing a salt formula without brackets, e.g. MgNO₃₂ for Mg(NO₃)₂.
• Filtering before the acid is fully used up in required practical 1, so the product is contaminated with acid.
• Heating the salt solution to dryness instead of letting crystals form.
• Saying chlorine forms at the cathode, or that sodium forms from sodium chloride solution.

Next steps

Use the revision notes, then the practice questions. Titration moles are in the Quantitative chemistry revision notes. The free 10-minute diagnostics show where to focus next.

Official syllabus

AQA GCSE Chemistry (8462) specification, Version 1.1 (October 2019), for teaching from September 2016 and exams from June 2018 onwards (AQA), section 4.4 Chemical changes.