7. Acids, bases and salts
- Syllabus
- 0620–2026–2027
- Section
- 7
- Level
- —
Exam analysis
7.1 The characteristic properties of acids and bases
- Syllabus
- 0620–2026–2027
- Topic
- 7.1
- Level
- —
Describe how acids react
| Acid reacts with | Products | Typical observation |
|---|---|---|
| a reactive metal | salt + hydrogen | effervescence; hydrogen gives a squeaky pop with a lighted splint |
| a base | salt + water | base dissolves; neutralisation occurs |
| a carbonate | salt + water + carbon dioxide | effervescence; CO₂ turns limewater milky |
Mg + H₂SO₄ → MgSO₄ + H₂; CuO + 2HNO₃ → Cu(NO₃)₂ + H₂O; CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂.
Choose the salt from the acid anion and the metal or ammonium ion supplied by the other reactant, then balance the equation.
Acids do not produce hydrogen with every metal. A metal must be reactive enough, and nitric acid may not behave like a typical acid with metals.
Recall indicator colours in acids
| Indicator | Colour in an acid |
|---|---|
| litmus | red; blue litmus turns red |
| thymolphthalein | colourless |
| methyl orange | red |
Add only a small amount of indicator and compare the resulting colour with its known acidic colour.
These indicators show an acidic or alkaline range but do not give a precise pH or acid strength.
Do not confuse methyl orange with universal indicator: methyl orange is red in acid, while universal indicator can show several acidic colours.
Distinguish bases and alkalis
Bases are oxides or hydroxides of metals. An alkali is a base that is soluble in water.
| Substance | Base? | Alkali? |
|---|---|---|
| copper(II) oxide | yes | no; insoluble |
| magnesium hydroxide | yes | no; only slightly soluble |
| sodium hydroxide | yes | yes; soluble |
| potassium hydroxide | yes | yes; soluble |
Every alkali is a base, but not every base is an alkali. Solubility is the deciding distinction.
Do not define every base as soluble. Insoluble metal oxides can still neutralise acids.
Describe how bases react
| Base reacts with | Products | Observation |
|---|---|---|
| acid | salt + water | neutralisation; an insoluble base may dissolve |
| ammonium salt, on warming | salt + water + ammonia | pungent alkaline gas released; damp red litmus turns blue |
CuO + 2HCl → CuCl₂ + H₂O. For an alkali and ammonium salt: NH₄⁺ + OH⁻ → NH₃ + H₂O.
Warm the mixture gently to release ammonia. Test the gas with damp indicator paper because ammonia must dissolve in water before showing alkalinity.
An ammonium salt does not release ammonia merely on its own; it must react with a base or alkali, usually on warming.
Recall indicator colours in alkalis
| Indicator | Colour in an alkali |
|---|---|
| litmus | blue; red litmus turns blue |
| thymolphthalein | blue |
| methyl orange | yellow |
Use a clean sample and a small amount of indicator so contamination does not shift the colour.
Indicator colour identifies an alkaline range, but universal indicator or a pH meter is needed to compare relative alkalinity.
Thymolphthalein is blue in alkali and colourless in acid; methyl orange is yellow in alkali and red in acid.
Link acids and alkalis to their ions
| Aqueous solution | Characteristic ion |
|---|---|
| acid | H⁺(aq) |
| alkali | OH⁻(aq) |
Acidic properties arise from mobile H⁺ ions in aqueous solution; alkaline properties arise from mobile OH⁻ ions.
HCl(aq) supplies H⁺ ions, while NaOH(aq) supplies OH⁻ ions. The state symbol (aq) matters because these ions are present in water.
A compound containing hydrogen is not automatically an acid, and a compound containing oxygen is not automatically an alkali.
Compare acidity and alkalinity using pH
| Approximate pH | Universal indicator colour | Interpretation |
|---|---|---|
| 0–2 | red | strongly acidic; highest H⁺ concentration |
| 3–4 | orange | acidic |
| 5–6 | yellow | weakly acidic |
| 7 | green | neutral |
| 8–10 | blue | alkaline |
| 11–14 | purple | strongly alkaline |
Touch universal indicator paper with a clean drop of solution and compare the colour with the chart. Lower pH means greater hydrogen ion concentration and greater acidity; higher pH means greater alkalinity.
Use equal measurement conditions when comparing samples. Neutrality is pH 7, not merely a pale or colourless solution.
Universal indicator gives an approximate pH range. It does not by itself distinguish acid strength from concentration unless concentrations are controlled.
Describe acid–alkali neutralisation
In neutralisation, H⁺ ions from an acid react with OH⁻ ions from an alkali to form water: H⁺(aq) + OH⁻(aq) → H₂O(l).
The remaining ions form the salt. They are spectator ions in the net ionic equation because they do not change.
HCl + NaOH → NaCl + H₂O. The full equation contains the salt; the ionic equation shows the chemical change common to every acid–alkali neutralisation.
Neutralisation describes acid reacting with base, but a neutral final pH occurs only when suitable reacting amounts are used; excess acid or alkali leaves the mixture non-neutral.
Define acids and bases by proton transfer
| Species role | Proton transfer |
|---|---|
| acid | donates H⁺ |
| base | accepts H⁺ |
In HCl + H₂O → H₃O⁺ + Cl⁻, HCl donates a proton and acts as the acid; water accepts it and acts as the base.
Compare each species before and after the reaction. The acid has lost H⁺; the base has gained H⁺.
Proton transfer is not electron transfer. A proton is H⁺, so acid–base roles are separate from oxidation and reduction roles.
Distinguish strong and weak acids
| Acid type | Dissociation in water | Particle picture |
|---|---|---|
| strong acid | complete | essentially all acid particles form ions |
| weak acid | partial | molecules and ions coexist at equilibrium |
At the same concentration, a strong acid has a greater H⁺ concentration, lower pH, greater electrical conductivity, and usually reacts faster than a weak acid.
Use a one-way arrow for complete dissociation and a reversible arrow for partial dissociation.
Strength is the extent of dissociation; concentration is the amount of acid per unit volume. A weak acid can be concentrated and a strong acid can be dilute.
Show that hydrochloric acid is strong
HCl(aq) → H⁺(aq) + Cl⁻(aq)
The one-way arrow shows that hydrochloric acid is completely dissociated into hydrogen ions and chloride ions in aqueous solution.
For this syllabus model, the aqueous solution is represented by H⁺ and Cl⁻ ions rather than an equilibrium containing a substantial amount of undissociated HCl.
Strong does not mean concentrated or dangerous by definition. It refers to complete dissociation in water.
Show that ethanoic acid is weak
CH₃COOH(aq) ⇌ H⁺(aq) + CH₃COO⁻(aq)
The reversible arrow shows partial dissociation: the aqueous mixture contains undissociated ethanoic acid molecules as well as H⁺ and ethanoate ions.
At the same total acid concentration, ethanoic acid produces fewer H⁺ ions and therefore has a higher pH than hydrochloric acid.
Weak does not mean that ethanoic acid cannot react with metals, bases, or carbonates. It describes partial dissociation, not absence of acidic reactions.
7.2 Oxides
- Syllabus
- 0620–2026–2027
- Topic
- 7.2
- Level
- —
Classify acidic and basic oxides
| Oxide type | Typical element character | Reaction pattern | Required examples |
|---|---|---|---|
| acidic | non-metal oxide | reacts with bases to form salt and water | CO₂, SO₂ |
| basic | metal oxide | reacts with acids to form salt and water | CuO, CaO |
CO₂ neutralises calcium hydroxide; SO₂ reacts with alkalis. CuO + 2HCl → CuCl₂ + H₂O, while CaO neutralises acids and acidic oxides.
Use metallic or non-metallic character as the first prediction, then use reaction evidence with acids and bases to confirm the class.
The metal/non-metal pattern is not universal: some oxides are amphoteric and some non-metal oxides such as CO are neutral. Use the specified examples and reaction evidence.
Describe amphoteric oxides
An amphoteric oxide reacts with both acids and bases to form a salt and water.
| Test | Evidence of amphoteric behaviour |
|---|---|
| add dilute acid | oxide reacts or dissolves; salt and water form |
| add aqueous alkali | oxide reacts or dissolves; a different salt and water form |
Place separate samples of the oxide in a named dilute acid and a named aqueous alkali. Reaction in both tests demonstrates amphoteric behaviour; use fresh oxide and controlled amounts for each test.
Amphoteric does not mean neutral or that the oxide forms a pH 7 solution. It means the oxide can react with both acidic and basic reagents.
Classify aluminium oxide and zinc oxide
Aluminium oxide, Al₂O₃, and zinc oxide, ZnO, are amphoteric oxides.
| Oxide | With acid | With aqueous alkali | Classification |
|---|---|---|---|
| Al₂O₃ | reacts to form an aluminium salt and water | reacts to form an aluminate salt and water | amphoteric |
| ZnO | reacts to form a zinc salt and water | reacts to form a zincate salt and water | amphoteric |
Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O; ZnO + 2HCl → ZnCl₂ + H₂O. Their additional ability to react with alkali distinguishes them from ordinary basic oxides.
Do not classify Al₂O₃ or ZnO as merely basic because they are metal oxides. Their reactions with both acids and bases make them amphoteric.
7.3 Preparation of salts
- Syllabus
- 0620–2026–2027
- Topic
- 7.3
- Level
- —
Prepare pure crystals of a soluble salt
| Reactants | How to make a neutral salt solution | Why this method fits |
|---|---|---|
| acid + alkali | titrate with indicator to find the exact reacting volumes; repeat those volumes without indicator | both reactants are soluble, so excess alkali cannot be filtered off |
| acid + excess metal | warm the acid, add metal until some remains, then filter | excess insoluble metal ensures no acid remains |
| acid + excess insoluble base | warm the acid, add base until some remains, then filter | excess insoluble base ensures no acid remains |
| acid + excess insoluble carbonate | warm the acid, add carbonate until fizzing stops and some remains, then filter | excess insoluble carbonate ensures no acid remains |
After obtaining the pure salt solution, heat it gently to evaporate some water until the solution is saturated. Leave it to cool so crystals form, then filter the crystals, wash them with a little cold distilled water and dry them between filter papers.
| Added solid | Useful observation | Other product |
|---|---|---|
| suitable metal | bubbling becomes slower and finally stops | hydrogen |
| insoluble base | solid eventually remains and no more dissolves | water |
| insoluble carbonate | effervescence becomes slower and finally stops | carbon dioxide and water |
Do not filter after using a soluble alkali: it passes through the filter. Do not evaporate the solution to dryness, because strong heating can damage the salt and prevents well-formed crystals from growing.
Apply salt solubility rules
| Compound family | Rule in water | Exceptions |
|---|---|---|
| sodium, potassium and ammonium salts | soluble | none required |
| nitrates | soluble | none required |
| chlorides | soluble | lead(II) chloride and silver chloride are insoluble |
| sulfates | soluble | barium sulfate, calcium sulfate and lead(II) sulfate are insoluble |
| carbonates | insoluble | sodium, potassium and ammonium carbonates are soluble |
| hydroxides | insoluble | sodium, potassium and ammonium hydroxides are soluble; calcium hydroxide is partially soluble |
Read the ions in the formula, then test the strongest general rule first: a sodium, potassium or ammonium ion, or a nitrate ion, makes the compound soluble. Otherwise apply the chloride, sulfate, carbonate or hydroxide rule and its named exceptions.
AgCl is insoluble because silver is a chloride exception. K₂CO₃ is soluble because potassium overrides the usual carbonate rule. BaSO₄ is insoluble because barium is a sulfate exception.
Partially soluble is not the same as fully soluble: calcium hydroxide dissolves only to a limited extent. State the correct qualification when applying the rule.
Distinguish hydrated and anhydrous substances
| Term | Meaning |
|---|---|
| hydrated substance | contains water chemically combined within its crystal structure |
| anhydrous substance | contains no water of crystallisation |
Heating a hydrated salt can remove its water of crystallisation and form the anhydrous salt. Adding water can reverse the change for suitable salts. For example: hydrated copper(II) sulfate ⇌ anhydrous copper(II) sulfate + water.
A formula with a dot followed by water, such as ZnSO₄·7H₂O, represents a hydrated salt. The corresponding anhydrous formula is ZnSO₄.
Hydrated does not mean merely wet. Water of crystallisation is present in a fixed chemical ratio inside the crystals; surface water can be removed by drying without changing the substance's formula.
Prepare an insoluble salt by precipitation
| Step | Action | Purpose |
|---|---|---|
| 1 | choose two soluble salts whose ions supply the required insoluble salt | ensures the reacting ions can meet in solution |
| 2 | dissolve each salt separately in distilled water | makes two aqueous solutions |
| 3 | mix the solutions and stir | forms the insoluble precipitate |
| 4 | filter the mixture | collects the precipitate as the residue |
| 5 | wash the residue with distilled water | removes soluble impurities |
| 6 | dry the residue between filter papers or in a warm place | obtains a pure, dry insoluble salt |
To prepare lead(II) iodide, mix aqueous lead(II) nitrate and aqueous potassium iodide: Pb(NO₃)₂(aq) + 2KI(aq) → PbI₂(s) + 2KNO₃(aq). The yellow PbI₂ precipitate is the wanted product.
The net ionic change is Pb²⁺(aq) + 2I⁻(aq) → PbI₂(s). Potassium and nitrate ions remain dissolved as spectator ions.
Keep the residue, not the filtrate. An insoluble salt is collected by filtration and drying; it is not obtained by evaporating the filtrate to crystallisation.
Define water of crystallisation
Water of crystallisation is water that is chemically combined in a hydrated crystal in a fixed ratio to the salt formula units.
| Hydrated formula | Meaning |
|---|---|
| CuSO₄·5H₂O | five water molecules for each CuSO₄ formula unit |
| CoCl₂·6H₂O | six water molecules for each CoCl₂ formula unit |
On heating, a hydrated salt can lose its water of crystallisation and become anhydrous. Hydrated copper(II) sulfate is blue and becomes white anhydrous copper(II) sulfate; hydrated cobalt(II) chloride is pink and becomes blue anhydrous cobalt(II) chloride.
The dot in a hydrated formula records the fixed number of water molecules in the crystal; it does not mean that the salt is simply dissolved in liquid water or has water on its surface.