(g) Acids, bases and salt preparations

Learning objectives

2.34Solubility rulesKnow the general rules for predicting the solubility of ionic compounds in water:• common sodium, potassium and ammonium compounds are soluble• all nitrates are soluble• common chlorides are soluble, except those of silver and lead(II)• common sulfates are soluble, except for those of barium, calcium and lead(II)• common carbonates are insoluble, except for those of sodium, potassium and ammonium• common hydroxides are insoluble except for those of sodium, potassium and calcium (calcium hydroxide is slightly soluble).2.35Acids and bases as proton transferUnderstand acids and bases in terms of proton transfer 2.36Proton donors and acceptorsUnderstand that an acid is a proton donor and a base is a proton acceptor 2.37Acid reactionsDescribe the reactions of hydrochloric acid, sulfuric acid and nitric acid with metals, bases and metal carbonates (excluding the reactions between nitric acid and metals) to form salts 2.38Bases and alkalisKnow that metal oxides, metal hydroxides and ammonia can act as bases, and that alkalis are bases that are soluble in water 2.39Preparing soluble salts from insoluble reactantsDescribe an experiment to prepare a pure, dry sample of a soluble salt, starting from an insoluble reactant 240C Preparing soluble salts by titrationDescribe an experiment to prepare a pure, dry sample of a soluble salt, starting from an acid and alkali 241C Preparing insoluble saltsDescribe an experiment to prepare a pure, dry sample of an insoluble salt, starting from two soluble reactants 2.42Soluble salt practicalPractical: prepare a sample of pure, dry hydrated copper(II) sulfate crystals starting from copper(II) oxide 243C Prepare a sample of pure, dry lead(II) sulfatePractical: prepare a sample of pure, dry lead(II) sulfate

An ionic compound is soluble if enough of it dissolves in water to form an aqueous solution. Apply the cation rule first when it gives an unconditional result, then check the anion and its exceptions.

Compound group	General rule	Exceptions or limits
sodium, potassium and ammonium compounds	soluble	none in the required rules
nitrates	soluble	none in the required rules
chlorides	soluble	silver chloride and lead(II) chloride are insoluble
sulfates	soluble	barium, calcium and lead(II) sulfates are insoluble
carbonates	insoluble	sodium, potassium and ammonium carbonates are soluble
hydroxides	insoluble	sodium and potassium hydroxides are soluble; calcium hydroxide is slightly soluble

Potassium carbonate is soluble because all common potassium compounds are soluble. Barium sulfate is insoluble because barium is a sulfate exception, so mixing solutions that supply $\ce{Ba^{2+}}$ and $\ce{SO4^{2-}}$ forms a precipitate.

Do not apply only the broad anion rule and ignore an exception. 'Slightly soluble' is the specified classification for calcium hydroxide; it is not grouped with fully soluble sodium and potassium hydroxides.

An acid–base reaction involves transfer of a proton, $\ce{H+}$ , from one particle to another. The proton is not released and left unaccounted for: one species loses it as another gains it.

\ce{H+(aq) + OH-(aq) -> H2O(l)}

In neutralisation, the acid supplies $\ce{H+}$ and the hydroxide ion gains that proton to become $\ce{H2O}$ . Tracking the same hydrogen and its positive charge shows exactly what has transferred.

Proton transfer is not electron transfer. A proton is $\ce{H+}$ ; oxidation and reduction instead track electrons, so an acid–base reaction is not automatically redox.

An acid is a proton donor and a base is a proton acceptor. Identify each role by comparing a species before and after the reaction.

\ce{HCl + NH3 -> NH4+ + Cl-}

Species	Change	Role
$\ce{HCl}$	loses $\ce{H+}$ and becomes $\ce{Cl-}$	acid: proton donor
$\ce{NH3}$	gains $\ce{H+}$ and becomes $\ce{NH4+}$	base: proton acceptor

Do not decide from a formula alone when a reaction is supplied: follow the proton. The acid loses $\ce{H+}$ ; the base does not donate it and does not need to contain $\ce{OH-}$ .

Hydrochloric, sulfuric and nitric acids form different salt families: chlorides, sulfates and nitrates. The other products depend on whether the acid reacts with a metal, a base or a metal carbonate.

Reactants	Products	Example
acid + metal	salt + hydrogen	$\ce{2HCl + Mg -> MgCl2 + H2}$
acid + base	salt + water	$\ce{H2SO4 + CuO -> CuSO4 + H2O}$
acid + metal carbonate	salt + water + carbon dioxide	$\ce{2HNO3 + CaCO3 -> Ca(NO3)2 + H2O + CO2}$

Metal reactions usually show hydrogen bubbles as the metal disappears. Carbonates effervesce because carbon dioxide escapes. With a solid base such as copper(II) oxide, the solid disappears as a salt solution forms when the base is not in excess.

Reactions between nitric acid and metals are explicitly excluded here, so do not apply the salt-plus-hydrogen pattern to them. Preserve formulas, coefficients and the acid-derived salt name.

A base neutralises an acid by accepting protons. Metal oxides, metal hydroxides and ammonia can act as bases; an alkali is specifically a base that is soluble in water.

Substance	Acts as a base?	Alkali?	Reason
copper(II) oxide	yes	no	neutralises acid but is insoluble in water
sodium hydroxide	yes	yes	dissolves in water and supplies $\ce{OH-}$
calcium hydroxide	yes	slightly soluble alkali	its limited solubility still gives an alkaline solution
ammonia	yes	aqueous ammonia is alkaline	accepts $\ce{H+}$ and produces an alkaline solution in water

All alkalis are bases because they neutralise acids, but not all bases are alkalis because many bases are insoluble. Solubility—not whether the name contains 'hydroxide'—controls the alkali label.

A metal oxide can be a base without containing $\ce{OH-}$ in its formula, and ammonia can be a base without being a metal hydroxide. Do not use 'base' and 'alkali' as exact synonyms.

Use an excess insoluble oxide, hydroxide or carbonate—or a suitable metal with hydrochloric or sulfuric acid—to consume all the acid, then remove the unused solid before crystallising the soluble salt.

Stage	Purpose
1. Warm dilute acid gently.	increases reaction rate without boiling away acid
2. Add the insoluble solid in small portions while stirring until some remains.	excess solid shows all acid has reacted
3. Filter.	removes the unreacted excess solid; the salt is in the filtrate
4. Heat the filtrate until near saturation.	evaporates some water without drying the salt completely
5. Cool to crystallise, filter the crystals and dry them.	obtains a pure, dry sample

Choose reactants that form the required soluble salt. For magnesium nitrate, warm nitric acid with excess magnesium oxide, then filter and crystallise the magnesium nitrate solution.

Do not evaporate the solution to dryness when hydrated crystals are required. Filtering before crystallisation is essential: otherwise excess reactant contaminates the crystals.

An acid and an alkali are both soluble, so excess reactant cannot be removed by filtration. First use a titration to find the exact reacting volumes, then remake the salt solution without indicator.

Stage	Action
1	Pipette a fixed alkali volume into a conical flask, add a suitable indicator and titrate with acid to the end-point.
2	Repeat to obtain a reliable reacting volume.
3	Mix the same measured acid and alkali volumes again, but add no indicator.
4	Heat the pure salt solution to near saturation, then leave it to cool and crystallise.
5	Filter the crystals, rinse with a little cold distilled water and dry them.

Sulfuric acid and sodium hydroxide form sodium sulfate solution: $\ce{H2SO4 + 2NaOH -> Na2SO4 + 2H2O}$ . The indicator-free repeat prevents coloured indicator contaminating the crystals.

Do not add excess acid or alkali: both remain dissolved and cannot be filtered off. The first titration finds the proportions; it is not normally the solution crystallised because it contains indicator.

Prepare an insoluble salt by mixing two soluble salt solutions whose ions combine to form the required precipitate.

Stage	Action and reason
1	Select two soluble compounds that supply the required cation and anion; dissolve them separately if starting from solids.
2	Mix and stir the solutions so the insoluble salt precipitates.
3	Filter to collect the precipitate as the residue.
4	Wash the residue with distilled water to remove soluble impurities.
5	Dry between filter papers, in a warm oven or in a desiccator.

\ce{BaCl2(aq) + Na2SO4(aq) -> BaSO4(s) + 2NaCl(aq)}

The desired insoluble salt is the filter residue, not the filtrate. Do not use evaporation or crystallisation: those methods recover a dissolved soluble salt, whereas the precipitate has already formed as a solid.

Prepare hydrated copper(II) sulfate by reacting warm dilute sulfuric acid with excess insoluble copper(II) oxide, then crystallising the blue solution.

\ce{CuO(s) + H2SO4(aq) -> CuSO4(aq) + H2O(l)}

Stage	Observation or purpose
1	Warm dilute sulfuric acid in a beaker; do not boil.
2	Add black copper(II) oxide a little at a time with stirring until black solid remains.
3	Filter off excess $\ce{CuO}$ ; collect the blue copper(II) sulfate filtrate.
4	Heat gently until a cooled drop forms crystals, then leave the solution to cool.
5	Filter the blue crystals, rinse with a little cold distilled water and dry between filter papers.

Excess copper(II) oxide removes all acid and is later filtered off. Do not strongly heat the final crystals or evaporate to dryness, because the target is hydrated copper(II) sulfate.

Lead(II) sulfate is insoluble, so prepare it as a precipitate by mixing two soluble solutions that supply $\ce{Pb^{2+}}$ and $\ce{SO4^{2-}}$ ions.

\ce{Pb(NO3)2(aq) + Na2SO4(aq) -> PbSO4(s) + 2NaNO3(aq)}

Stage	Action
1	Mix measured lead(II) nitrate and sodium sulfate solutions and stir to complete precipitation.
2	Filter the mixture; retain the solid lead(II) sulfate residue.
3	Wash the residue with distilled water to remove soluble sodium nitrate and excess ions.
4	Dry the solid between filter papers or in a warm oven.

Lead compounds are toxic: wear eye protection and gloves, avoid skin contact and dust, use small quantities, wash hands and place all lead-containing residues and liquids in the designated hazardous-waste container.

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—(h) Chemical testsRead

(g) Acids, bases and salt preparations

Learning objectives

Predict solubility from the required rules

See acid–base reactions as proton transfer

Identify proton donors and acceptors

Predict products of the three acid reaction patterns

Distinguish bases from alkalis

Prepare a soluble salt from an insoluble reactant

Prepare a soluble salt from an acid and alkali

Prepare an insoluble salt by precipitation

Prepare hydrated copper(II) sulfate crystals

Prepare pure, dry lead(II) sulfate

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