Acids, alkalis, indicators, neutralisation and preparing salts.
Revise the key ideas
Acids, bases and alkalis
Acids dissolved in water (in aqueous solution) provide hydrogen ions, H⁺. Hydrochloric acid produces H⁺ and chloride ions, Cl⁻. Chloride ions are different from chlorine atoms or chlorine gas.
Alkalis provide hydroxide ions, OH⁻, in aqueous solution. Sodium hydroxide solution contains Na⁺ and OH⁻.
A base neutralises an acid to produce salt and water; metal oxides and metal hydroxides are common examples. An alkali is a base that dissolves in water.
All alkalis are bases, but not all bases are alkalis. Copper oxide is an insoluble base; sodium hydroxide is a soluble base, an alkali.
For ordinary GCSE aqueous solutions, acidic means pH below 7, neutral pH 7, and alkaline pH above 7. The familiar 0–14 scale is useful but not an absolute limit for every solution.
Wear eye protection and use the specified dilute reagents in school practicals; hazard and risk depend on both the substance and its concentration.
Indicators and pH
Universal indicator gives an approximate pH by comparison with a colour chart: red/orange/yellow for acidic solutions, green near neutral, blue/purple for alkaline solutions.
Litmus is red in acid and blue in alkali. It indicates acidity or alkalinity rather than providing an exact pH.
Methyl orange is red in sufficiently acidic solution and yellow in neutral or alkaline solution, with orange in its transition range.
Phenolphthalein is colourless in acid and neutral solution and pink in sufficiently alkaline solution. Use a suitable indicator for the expected endpoint.Colours in suitably acidic/alkaline solutions; indicator transition ranges differ.
A calibrated pH probe can measure pH more precisely than matching an indicator colour. Rinse it between samples to reduce contamination.
Lower pH means higher hydrogen-ion concentration. A decrease of one pH unit corresponds to ten times the H⁺ concentration; a decrease of two means a hundred times.
Strong, weak, dilute and concentrated
A strong acid ionises almost completely in water: almost all its acid molecules form ions. A weak acid ionises only partly. Its molecules and ions are in equilibrium. This splitting into ions is also called dissociation.
Concentrated means a large amount of dissolved acid per volume of solution; dilute means a smaller amount per volume. This is different from strength.
A dilute strong acid and a concentrated weak acid are both possible. Do not use “strong” simply to mean “lots of acid”.
At the same concentration, a strong acid generally supplies more hydrogen ions and has lower pH than a weak acid. Comparisons at different concentrations need more information.
Diluting an acidic solution lowers H⁺ concentration and moves pH towards 7. It does not change a strong acid into a weak acid.
For example, pH 2 has ten times the H⁺ concentration of pH 3, and a hundred times that of pH 4. Describe concentration rather than saying it is “a hundred times stronger”.Each unit decrease gives tenfold greater H⁺ concentration.
Acid reactions and salt names
A sufficiently reactive metal with a suitable dilute acid gives salt + hydrogen. For magnesium: Mg + 2HCl → MgCl₂ + H₂. Water is not an additional product in this equation.
Metal oxide + acid → salt + water. Example: CuO + H₂SO₄ → CuSO₄ + H₂O.
Metal hydroxide + acid → salt + water. Example: NaOH + HCl → NaCl + H₂O.
Metal carbonate + acid → salt + water + carbon dioxide. Example: CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂.
Hydrochloric acid makes chlorides, sulfuric acid sulfates, and nitric acid nitrates. The metal or ammonium ion supplies the first part of the salt name.
Hydrogen makes a squeaky pop with a lighted splint. Carbon dioxide turns limewater cloudy. Use small gas samples and the teacher's safe practical procedure.
The ionic equation for acid–alkali neutralisation is H⁺(aq) + OH⁻(aq) → H₂O(l). Ions such as Na⁺ and Cl⁻ do not change during the reaction. These spectator ions are left out of the ionic equation.The water-forming ionic equation is independent of spectator ions.
Copper sulfate crystals: core practical
To make a soluble salt using an insoluble base, warm dilute acid gently and add the base in small portions until some remains unreacted. Excess base ensures the acid is used up.
For hydrated copper sulfate crystals, react copper oxide with dilute sulfuric acid. Stir while adding the black solid until an excess remains.
Filter the mixture to remove excess copper oxide as residue; copper sulfate solution is the filtrate. Do not use the solid residue as the desired salt.
Concentrate the filtrate gently using a water bath, then leave it to cool and crystallise. Avoid heating to dryness, which can remove water of crystallisation or cause spitting.
Filter the crystals, wash with a little cold distilled water and dry between filter papers. Hydrated copper sulfate crystals are blue. They contain water molecules in their crystal structure, called water of crystallisation.Keep the blue hydrated crystals; use a water bath and avoid heating to dryness.
An insoluble excess base can be filtered off; a dissolved excess alkali cannot. A soluble acid–alkali salt preparation therefore needs measured reacting volumes.
Titration to prepare a soluble salt
Use a pipette and filler to transfer a measured alkali volume to a conical flask. Add a few drops of a suitable indicator; never pipette by mouth.
Fill a burette with acid, ensure its tip is full and read the initial level at eye height. Swirl the flask while adding acid, then add dropwise near the colour change.Read burette at eye level; swirl and add dropwise near the endpoint.
The titre is the volume delivered from the burette: final reading − initial reading. Repeat until the titres are close together (concordant). Follow the agreement required by your course or the question.
The endpoint is the indicator colour change indicating suitable reacting proportions. It is not automatically pH 7 for every acid–base combination.
To obtain pure salt, repeat with the measured reacting volumes but without indicator. Then gently concentrate, cool, crystallise and dry the salt.
Do not add excess soluble reactant and attempt to filter it out: it remains dissolved and contaminates the salt. Titration avoids excess acid or alkali.
Solubility and precipitation
All common sodium, potassium and ammonium salts, and all nitrates, are soluble in water under the GCSE solubility rules.
Common chlorides are soluble except silver chloride and lead chloride. Common sulfates are soluble except lead, barium and calcium sulfates.
Common carbonates and hydroxides are insoluble except those of sodium, potassium and ammonium. Calcium hydroxide is slightly soluble; use the specified rules for the question.
A precipitate is an insoluble solid formed when suitable solutions react. Predict possible ion combinations and check their solubility.
Silver nitrate plus sodium chloride solutions form silver chloride precipitate: Ag⁺(aq) + Cl⁻(aq) → AgCl(s). Sodium and nitrate ions remain dissolved.
To obtain pure dry insoluble salt, mix suitable solutions, filter off the precipitate, wash it with distilled water to remove dissolved salts, and dry it.Here the desired salt is the filter residue, unlike the copper sulfate filtrate method.
Changing pH: core practical
Measure a fixed volume of dilute hydrochloric acid and its initial pH. Add small measured portions of powdered calcium hydroxide or calcium oxide, stirring each time.
Measure pH after each addition with a suitable probe or indicator and record the added amount. pH rises as the acid is neutralised; excess base can make the mixture alkaline.
Control acid volume and concentration, powder choice and mixing method when comparing runs. Avoid splashing and wear eye protection.
Plot pH against amount added and interpret the trend. A pH rise reflects falling H⁺ concentration, not the acid being converted into a weaker acid type.
Watch CC8 · Acids and alkalis · Topic 3 — Chemical changes
Revise acids and alkalis with this narrated video. Use the player controls to pause, seek, adjust the volume or mute. Turn English captions on or off using the captions menu.
Mr Kidman's Classroom