4.4 Chemical changes — notes on one page

4.4.1.1 · Metal oxides

• A metal reacting with oxygen forms a metal oxide; the metal is oxidised because it gains oxygen.
• Reduction is the loss of oxygen from a substance, while oxidation is the gain of oxygen.
• For example, heating copper in oxygen forms copper oxide: $2\mathrm{Cu}+\mathrm{O_2}\rightarrow2\mathrm{CuO}$.
• Do not decide oxidation from the presence of oxygen alone: track whether oxygen is gained or lost by the named substance.
• Exam answers should link the observed product or property to the relevant particles, electron transfer or position in the reactivity series.

Exam tip. For oxidation or reduction in terms of oxygen, state explicitly whether oxygen is gained or lost.

4.4.1.2 · The reactivity series

Definition. The reactivity of a metal is its tendency to form positive ions. The reactivity series orders potassium, sodium, lithium, calcium, magnesium, zinc, iron and copper by that tendency, with the non-metals hydrogen and carbon included for comparison.

• When metals react they lose electrons and form positive ions, so the more easily a metal forms its positive ion, the more reactive it is. The specification places the order on reactions with water and with dilute acids at room temperature only — reactions with steam are outside 8462.
• Potassium, sodium, lithium and calcium react with cold water; magnesium, zinc and iron react with dilute acids but not appreciably with cold water; copper reacts with neither, because it sits below hydrogen.
• A more reactive metal displaces a less reactive metal from a compound, and displacement is the standard experimental route to an unknown metal's position.
• Carbon's position decides the extraction method: metals below carbon can be extracted by reduction with carbon, because carbon removes the oxygen; metals above carbon, such as aluminium, sodium and potassium, must be extracted by electrolysis.
• You must also deduce a reactivity order from experimental results.

What the examiner wants. A comparison, not a label. Every reactivity mark scheme awards the converse, so 'carbon is more reactive than tin' and 'tin is less reactive than carbon' both score — but 'tin is unreactive' scores nothing.

Exam tip. A displacement answer should compare both metals in the reactivity series and state which forms positive ions more readily.

4.4.1.3 · Extraction of metals and reduction

• Very unreactive metals can occur native, but most metals occur as compounds and must be extracted by chemical reactions.
• A metal below carbon in the reactivity series can be extracted from its oxide by reduction with carbon.
• In oxygen-transfer language, the metal oxide is reduced because it loses oxygen; carbon is oxidised because it gains oxygen.
• Use supplied evidence to evaluate an extraction route; detailed industrial processes beyond reduction of oxides with carbon are not required here.
• Exam answers should link the observed product or property to the relevant particles, electron transfer or position in the reactivity series.

Exam tip. For extraction, first locate the metal relative to carbon and then name the reduction method.

4.4.1.4 · Oxidation and reduction in terms of electrons (HT only)

• Higher tier: oxidation is loss of electrons and reduction is gain of electrons: OIL RIG. A half equation shows electrons explicitly and must balance both atoms and total charge.
• Add state symbols when the question asks for them; they can be part of the mark for a complete ionic equation.
• In a metal displacement, the more reactive metal loses electrons while the displaced metal ions gain electrons.
• Do not identify redox from charge signs alone: compare the same species before and after the reaction and track electron transfer.
• Exam answers should link the observed product or property to the relevant particles, electron transfer or position in the reactivity series.

Exam tip. Check both atom count and total charge after writing an ionic or half equation.

4.4.2.1 · Reactions of acids with metals

• Acids react with magnesium, zinc and iron to produce a salt and hydrogen gas. Hydrochloric acid forms chloride salts, while sulfuric acid forms sulfate salts.
• For example, $\mathrm{Zn}+2\mathrm{HCl}\rightarrow\mathrm{ZnCl_2}+\mathrm{H_2}$; the metal has replaced hydrogen from the acid.
• Do not add water as a product: acid plus metal gives salt plus hydrogen.
• Higher tier: AQA restricts "explain in terms of gain or loss of electrons, that these are redox reactions" and "identify which species are oxidised and which are reduced in given chemical equations" to Higher tier; naming the salt and the hydrogen produced is examined on both tiers.
• Exam answers should link the observed product or property to the relevant particles, electron transfer or position in the reactivity series.

Exam tip. For acid–metal reactions, identify the salt from the acid and metal, then add hydrogen.

4.4.2.2 · Neutralisation of acids and salt production

Definition. Acids are neutralised by alkalis (soluble metal hydroxides) and by bases (insoluble metal hydroxides and metal oxides) to produce a salt and water, and by metal carbonates to produce a salt, water and carbon dioxide.

• The salt produced is decided by two things: the acid used supplies the negative ion, and the positive ion comes from the base, alkali or carbonate. Hydrochloric acid produces chlorides, nitric acid produces nitrates, and sulfuric acid produces sulfates.
• So sulfuric acid plus zinc oxide gives zinc sulfate and water; hydrochloric acid plus copper carbonate gives copper chloride, water and carbon dioxide. The rule also works backwards: asked which acid makes zinc nitrate, the nitrate tells you nitric acid.
• To write the formula of the salt, combine the formulae of the common ions so that the charges cancel — one Zn(2+) needs two Cl(-) to give ZnCl₂, but only one SO₄(2-) to give ZnSO₄.
• Three general equations are worth knowing verbatim: acid + alkali gives salt + water; acid + metal oxide or hydroxide gives salt + water; acid + metal carbonate gives salt + water + carbon dioxide.
• Acid + metal works only for metals above hydrogen.

What the examiner wants. The salt named from the acid's anion and the base's cation, in that order. Where a question asks for 'two other substances', the accepted list is the metal, the metal oxide and the metal hydroxide — a named carbonate is usually excluded by the stem.

Exam tip. Product prediction earns marks for the salt name and every additional product required by the reactant type.

4.4.2.3 · Soluble salts

Definition. Soluble salts are made by reacting an acid with an insoluble solid — a metal, metal oxide, hydroxide or carbonate. The solid is added to the acid until no more reacts, the excess solid is filtered off, and the salt solution is crystallised to give the solid salt.

• This is Required practical 1, and the six-mark method question that follows it has an unusually fixed indicative-content list. The steps in order are: add the insoluble solid to the warm dilute acid and stir;
• keep adding until the solid is in excess, shown by solid remaining and no more effervescence; filter the mixture to remove the excess solid; heat the filtrate gently, using a water bath rather than a naked flame, until the crystallisation point is reached; leave the solution to cool and crystallise; and pat the crystals dry with filter paper.
• Each step has a reason the examiner may ask for separately. Warming the acid increases the rate of reaction.
• Using excess solid ensures all the acid is neutralised, so no acid contaminates the product; excess solid is used rather than excess acid because a solid can be filtered off and excess acid cannot.
• Gentle heating avoids decomposing the salt, and slow cooling gives larger, purer crystals.

What the examiner wants. A logically sequenced method that actually ends with a pure, dry solid. Level 3 needs the excess test, the filtration and the drying step; a method that stops at 'evaporate the water' loses the top band.

Exam tip. A required-practical method must explain excess solid, filtration, concentration, crystallisation and drying in order.

4.4.2.4 · The pH scale and neutralisation

Definition. Acids produce hydrogen ions, H⁺, in aqueous solution and aqueous solutions of alkalis contain hydroxide ions, OH⁻. The pH scale runs from 0 to 14: pH 7 is neutral, below 7 is acidic and above 7 is alkaline.

• pH can be measured with universal indicator, which gives an approximate value from a colour, or with a pH probe, which gives a numerical reading and is more precise. Universal indicator turns red at low pH, green at pH 7 and blue or purple at high pH.
• In a neutralisation between an acid and an alkali the hydrogen ions react with the hydroxide ions to produce water, which the specification writes as H⁺ + OH⁻ gives H₂O.
• That ionic equation is what neutralisation means at particle level, and it explains why the conductivity of a mixture falls as the free ions are removed.
• Higher tier candidates must also handle the logarithmic scale: as pH falls by one unit the hydrogen ion concentration increases tenfold, so pH 2.5 has one hundred times the hydrogen ion concentration of pH 4.5.
• A single suitable indicator such as methyl orange or phenolphthalein is used in titrations; universal indicator is explicitly not accepted.

What the examiner wants. The indicator named first and its colour change second, in the direction the question specifies. The second mark is dependent on the first, so a colour change with no named indicator scores zero.

Exam tip. For neutralisation, use hydrogen ions plus hydroxide ions forming water and keep the charges balanced.

4.4.2.5 · Titrations (chemistry only)

Definition. A titration measures the volumes of acid and alkali solutions that react exactly with each other, using a suitable indicator to show the end point. Required practical 2 uses a strong acid — sulfuric, hydrochloric or nitric — against a strong alkali.

• This is Required practical 2. Measure a fixed volume of one solution into a conical flask with a volumetric pipette and add a few drops of a single suitable indicator such as methyl orange or phenolphthalein; universal indicator is explicitly not accepted.
• Fill the burette with the other solution, check for bubbles below the tap, and take the initial reading at eye level from the bottom of the meniscus. Stand the flask on a white tile, add from the burette while swirling, and add dropwise near the end point until the first permanent colour change.
• Record the final reading and subtract to get the titre.
• Repeat until you have concordant titres — usually within 0.10 cm³ — and take the mean of the concordant results only.
• Higher tier: AQA restricts "calculate the chemical quantities in titrations involving concentrations in $\mathrm{mol/dm^3}$ and in $\mathrm{g/dm^3}$" to Higher tier; describing how to carry out a titration accurately is examined on both tiers.

What the examiner wants. Named apparatus and the concordance rule. 'Repeat and take a mean' only scores fully when you also say the rough and any non-concordant run are excluded from that mean.

Exam tip. A titration method should name the apparatus, endpoint, dropwise addition and concordant repeats.

4.4.2.6 · Strong and weak acids (HT only)

• Higher tier: a strong acid ionises completely in water, whereas a weak acid ionises only partially.
• Hydrochloric, nitric and sulfuric acids are strong; ethanoic, citric and carbonic acids are weak.
• Strength describes degree of ionisation, while concentration describes the amount of solute per volume; either a strong or weak acid may be dilute.
• For whole-number pH values, a decrease of one pH unit means a tenfold increase in hydrogen ion concentration, not an increase of one unit.
• Exam answers should link the observed product or property to the relevant particles, electron transfer or position in the reactivity series.

Exam tip. At equal concentration, link stronger acid → greater ionisation → more hydrogen ions → lower pH.

4.4.3.1 · The process of electrolysis

• An electrolyte is a molten ionic compound or ionic solution whose mobile ions allow it to conduct electricity. Positive ions move to the negative cathode, while negative ions move to the positive anode.
• When ions are discharged at the electrodes, they form elements; the electric current drives this decomposition process.
• In the electrolyte, charge is carried by moving ions rather than free electrons; a solid ionic compound does not conduct because its ions are fixed.
• Higher tier: throughout Section 4.4.3 AQA restricts "write half equations for the reactions occurring at the electrodes during electrolysis, and may be required to complete and balance supplied half equations" to Higher tier; describing how ions move and are discharged is examined on both tiers.
• Exam answers should link the observed product or property to the relevant particles, electron transfer or position in the reactivity series.

Exam tip. Label the cathode negative and anode positive before tracing ion movement.

4.4.3.2 · Electrolysis of molten ionic compounds

• A molten binary ionic compound contains only its metal ions and non-metal ions, so its products are predictable from those ions. The metal forms at the negative cathode and the non-metal forms at the positive anode when inert electrodes are used.
• For molten lead bromide, lead forms at the cathode and bromine forms at the anode.
• Do not apply the aqueous-solution rules to a molten compound: there is no water supplying competing hydrogen or hydroxide ions.
• Higher tier: write, complete and balance half equations for electrode reactions.
• Exam answers should link the observed product or property to the relevant particles, electron transfer or position in the reactivity series.

Exam tip. For a molten binary compound, predict the metal at the cathode and non-metal at the anode.

4.4.3.3 · Using electrolysis to extract metals

• Electrolysis extracts metals that are too reactive for carbon reduction or that react with carbon. Extraction uses much energy because the ionic compound must be molten and an electric current must be maintained.
• Aluminium is extracted from a molten mixture of aluminium oxide and cryolite; the mixture has a lower melting point than pure aluminium oxide.
• Oxygen produced at the carbon anode reacts with carbon to form carbon dioxide, so the positive electrodes wear away and must be replaced.
• Higher tier: write, complete and balance half equations for electrode reactions.
• Exam answers should link the observed product or property to the relevant particles, electron transfer or position in the reactivity series.

Exam tip. An aluminium-extraction explanation should cover both the lowered melting point and carbon-anode reaction.

4.4.3.4 · Electrolysis of aqueous solutions

• An aqueous electrolyte contains ions from the dissolved compound as well as $\mathrm{H^+}$ and $\mathrm{OH^-}$ originating from water. At the cathode, hydrogen forms if the metal is more reactive than hydrogen; otherwise the metal forms.
• At the anode, a halide ion forms its halogen; if no halide is present, oxygen forms.
• Apply both electrode rules separately and assume inert electrodes unless told otherwise; do not simply name the elements in the solute.
• Higher tier: write, complete and balance half equations for electrode reactions.
• Exam answers should link the observed product or property to the relevant particles, electron transfer or position in the reactivity series.

Exam tip. For aqueous electrolysis, apply the cathode and anode selection rules separately.

4.4.3.5 · Representation of reactions at electrodes as half equations (HT only)

• Higher tier: at the cathode, positive ions gain electrons, so cathode reactions are reductions.
• At the anode, negative ions lose electrons, so anode reactions are oxidations.
• Balance a half equation by conserving atoms and total charge; electrons go on the side needed to balance charge.
• Check electron direction: electrons are reactants in a reduction half equation and products in an oxidation half equation.
• Exam answers should link the observed product or property to the relevant particles, electron transfer or position in the reactivity series.

Exam tip. A half equation must balance atoms, charge and electrons, with reduction at the cathode and oxidation at the anode.