6. METHODS of MAKING SALTS - salt preparation procedures

6a-b. Two Methods of making Salts which are water soluble

6a. METHOD (a) Neutralising a soluble acid with a soluble base e.g. the hydroxide of an alkali metal like sodium hydroxide or ammonia solution. Steps (1) to (3) below is called a titration.

Examples of neutralisation equations are given in section 4.

Typical common soluble bases (alkalis) used for preparing soluble salts:

NaOH sodium hydroxide, KOH potassium hydroxide, NH₃ ammonia

(1) A known volume of acid is pipetted into a conical flask and universal indicator added. The acid is titrated with the alkali from the burette.

(2) The acid is added until the indicator turns green, pH 7 neutral. This means all the acid has been neutralised to form the salt

(3) The volume of alkali needed for neutralisation is then noted, this is called the endpoint volume. (1)-(3) are repeated with both known volumes mixed together BUT without the contaminating universal indicator.

(4) The solution is transferred to an evaporating dish and heated to partially evaporate the water causing crystallisation or can be left to slowly evaporate - which tends to give bigger and better crystals.

(5) The residual liquid can be decanted away and the crystals can be carefully collected and dried by 'dabbing' with a filter paper OR the crystals can be collected by filtration (below) and dried (as above).

Note (i) You can put the acid in the burette and the alkali in the flask.

(ii) Parts (1) to (3) are known specifically as an acid-base (alkali) titration, and the general method is known as a volumetric titration by which it possible to find out exactly what volume ratios are needed for neutralisation. So knowing one concentration, you can calculate the other.

(iii) Concentration calculations are on calculations pages sections 11. and 12.

(iv) Apparatus used: (1) pipette and conical flask; (2)-(3) burette and conical flask; (4) evaporating (crystallising) dish, bunsen burner, tripod and gauze; (5) filter paper.

(v) Other indicators e.g. phenolphthalein can be used instead (pink alkaline, colourless acid).

(vi) The burette and pipette are both used for the accurate measurement of volume.

(vii) The pH changes in this preparation are described in section 7.

6b. METHOD (b) Reacting an acid with a metal or with an insoluble base e.g. an insoluble metal oxide, hydroxide or carbonate, often of a Group 2 metal like calcium, magnesium or a Transition Metal like nickel, copper or zinc. Copper metal won't dissolve in acids, but its oxide and carbonate will.

Examples of neutralization equations are given in section 4.

Typical common insoluble bases used for preparing soluble salts:

MgO magnesium oxide, MgCO₃ magnesium carbonate

CaO Calcium oxide, CaCO₃ calcium carbonate, Ca(OH)₂ calcium hydroxide,

NiO nickel(II) oxide, ZnO zinc oxide, Zn(OH)₂, zinc hydroxide, ZnCO₃ zinc carbonate

(1) The required volume of acid is measured out into the beaker with a measuring cylinder. The insoluble metal, oxide, hydroxide or carbonate is weighed out and the solid added in small portions to the acid in the beaker with stirring.

(2) The mixture may be heated to speed up the reaction. When no more of the solid dissolves it means ALL the acid is neutralised and there should be a little excess solid.

(3)  The hot solution (with care!) is filtered to remove the excess solid metal/oxide/carbonate, into an evaporating dish.

(4) The hot solution is left to cool and crystallise. Then collect and dry the crystals with a filter paper.

Note (i) Apparatus used: (1) balance, measuring cylinder, beaker and glass stirring rod. (2) beaker/rod, bunsen burner, tripod and gauze; (3)-(4) filter funnel and filter paper, evaporating (crystallising) dish.

(ii) A measuring cylinder is adequate for measuring the acid volume, you do not need the accuracy of a pipette or burette required in method (a).

(iii) How to calculate amounts required and % yield is dealt with in Chemical Calculations Part 14.

Method 6c. Preparing an Insoluble Salt

• An insoluble salt can be made by mixing two solutions of soluble salts in a process is called precipitation. One solution contains the 1st required ion, and the other solution contains the 2nd required ion. The precipitated salt can then be filtered off with a filter funnel and paper. The collected solid is washed with distilled water to remove any remaining soluble salt impurities and removed from the filter paper to be dried. Examples ...

  • (i) Silver chloride is made by mixing solutions of solutions of silver nitrate and sodium chloride.

    • silver nitrate + sodium chloride ==> silver chloride + sodium nitrate

    • AgNO_3(aq) + NaCl_(aq) ==> AgCl_(s) + NaNO_3(aq)

    • in terms of ions it could be written as

    • Ag⁺NO₃^-_(aq) + Na⁺Cl^-_(aq) ==> AgCl_(s) + Na⁺NO₃^-_(aq)

    • or: Ag⁺_(aq) + NO₃^-_(aq) + Na⁺_(aq) + Cl^-_(aq) ==> AgCl_(s) + Na⁺_(aq) + NO₃^-_(aq)

    • but the spectator ions are nitrate NO₃^- and sodium Na⁺ which do not change at all,

    • so the ionic equation is simply: Ag⁺_(aq) + Cl^-_(aq) ==> AgCl_(s)

      • Note that ionic equations omit ions that do not change there chemical or physical state.

      • In this case the nitrate, NO₃^-_(aq) and sodium Na⁺_(aq) ions do not change physically or chemically and are called spectator ions,

      • BUT the aqueous silver ion, Ag⁺_(aq), combines with the aqueous chloride ion, Cl^-_(aq), to form the insoluble salt silver chloride, AgCl_(s), thereby changing their states both chemically and physically.

      • More Ionic equations explained with all spectator ions indicated.

    • If you use barium chloride the word and symbol equations are ...

    • barium chloride + silver nitrate ==> silver chloride + barium nitrate

    • BaCl_2(aq) + 2AgNO_3(aq) ==> 2AgCl_(s) + Ba(NO₃)_2(aq)

    • which can be written as

    • Ba²⁺_(aq) + 2Cl^-_(aq) + 2Ag⁺_(aq) + 2NO₃^-_(aq) ==> 2AgCl_(s) + Ba²⁺_(aq) + 2NO₃^-_(aq)

    • the spectator ions are Ba²⁺ and NO₃^-

    • so the ionic equation is: Ag⁺_(aq) + Cl^-_(aq) ==> AgCl_(s)

  • (ii) Lead(II) iodide, a yellow precipitate (insoluble in water!) can be made by mixing lead(II) nitrate solution with e.g. potassium iodide solution.

    • lead(II) nitrate + potassium iodide ==> lead(II) iodide + potassium nitrate

    • Pb(NO₃)_2(aq) + 2KI_(aq) ==> PbI_2(s) + 2KNO_3(aq)

    • which can be written as

    • Pb²⁺_(aq) + 2NO₃^-_(aq) + 2K⁺_(aq) + 2I^-_(aq) ==> PbI_2(s) + 2K⁺_(aq) + 2NO₃^-_(aq)

    • the ionic equation is: Pb²⁺_(aq) + 2I^-_(aq) ==> PbI_2(s)

    • because the spectator ions are nitrate NO₃^- and potassium K⁺.

    • In a similar way you can make lead(II) chloride by e.g. using dilute hydrochloric acid

      • lead(II) nitrate + hydrochloric acid ==> lead(II) chloride + nitric acid

      • Pb(NO₃)_2(aq) + 2HCl_(aq) ==> PbCl_2(s) + 2HNO_3(aq)

      • Pb²⁺_(aq) + 2NO₃^-_(aq) + 2H⁺_(aq) + 2Cl^-_(aq) ==> PbCl_2(s) + 2H⁺_(aq) + 2NO₃^-_(aq)

      • the ionic equation is: Pb²⁺_(aq) + 2Cl^-_(aq) ==> PbCl_2(s)

      • because the spectator ions are nitrate NO₃^- and hydrogen H⁺.

    • and you can make lead(II) bromide by e.g. using sodium bromide

      • lead(II) nitrate + sodium bromide ==> lead(II) bromide + sodium nitrate

      • Pb(NO₃)_2(aq) + 2NaBr_(aq) ==> PbBr_2(s) + 2NaNO_3(aq)

      • Pb²⁺_(aq) + 2NO₃^-_(aq) + 2Na⁺_(aq) + 2Br^-_(aq) ==> PbBr_2(s) + 2Na⁺_(aq) + 2NO₃^-_(aq)

      • the ionic equation is: Pb²⁺_(aq) + 2Br^-_(aq) ==> PbBr_2(s)

      • because the spectator ions are nitrate NO₃^- and sodium Na⁺.

  • (iii) Calcium carbonate, a white precipitate, forms on e.g. mixing calcium chloride and sodium carbonate solutions ...

    • calcium chloride + sodium carbonate ==> calcium carbonate + sodium chloride

    • CaCl_2(aq) + Na₂CO_3(aq) ==> CaCO_3(s) + 2NaCl_(aq)

    • Ca²⁺_(aq) + 2Cl^-_(aq) + 2Na⁺_{(aq) +} CO₃^2-_(aq) ==> CaCO_3(s) + 2Na⁺_(aq) + 2Cl^-_(aq)

    • ionically: Ca²⁺_(aq) + CO₃^2-_(aq) ==> CaCO_3(s)

    • because the spectator ions are chloride Cl^- and sodium Na⁺.

  • (iv) Barium sulphate, a white precipitate, forms on mixing e.g. barium chloride and dilute sulphuric acid ...

    • barium chloride + sulphuric acid ==> barium sulphate + hydrochloric acid

    • BaCl_2(aq) + H₂SO_4(aq) ==> BaSO_4(s) + 2HCl_(aq)

    • Ba²⁺_(aq) + 2Cl^-_(aq) + 2H⁺_{(aq) +} SO₄^2-_(aq) ==> BaSO_4(s) + 2H⁺_(aq) + 2Cl^-_(aq)

    • ionic equation: Ba²⁺_(aq) + SO₄^2-_(aq) ==> BaSO_4(s)

    • because the spectator ions are chloride Cl^- and hydrogen H⁺.

      • Or you can use sulphate salts like sodium sulphate, so the word and symbol equations are ..

      • barium chloride + sodium sulfate ==> barium sulfate + sodium chloride

      • BaCl_2(aq) + Na₂SO_4(aq) ==> BaSO_4(s) + 2NaCl_(aq)

      • The ionic equation is the same: Ba²⁺_(aq) + SO₄^2-_(aq) ==> BaSO_4(s)

      • because the spectator ions are sodium Na⁺ and chloride Cl^-

  • (v) Lead(II) sulphate, a white precipitate, forms in a similar way e.g.

    • lead(II) nitrate + sodium sulphate ==> lead(II) sulphate + sodium nitrate

    • Pb(NO₃)_{2 (aq)} + Na₂SO_{4 (aq)} ==> PbSO_{4 (s)} + 2NaNO_{3 (aq)}

    • ionically: Pb²⁺_(aq) + SO₄^2-_(aq) ==> PbSO_4(s)

    • because the spectator ions are sodium Na⁺ and nitrate NO₃^-

  • NOTE: A precipitation reaction is generally defined as 'the formation of an insoluble solid on mixing two solutions or bubbling a gas into a solution'.

• General rules which describe the solubility of common types of compounds in water:

  • All common sodium, potassium and ammonium salts are soluble e.g. NaCl, K₂SO₄, NH₄NO₃

  • All nitrate salts are soluble e.g. NaNO₃, Mg(NO₃)₂, Al(NO₃)₃, NH₄NO₃

  • Some ethanoate salts are soluble e.g. CH₃COONa

  • Common chloride salts are soluble except those of silver and lead e.g.

    • soluble: KCl, CaCl₂, AlCl₃ or insoluble AgCl, PbCl₂

  • Common sulfates are soluble except those of lead, barium and calcium: soluble e.g.

    • soluble: Na₂SO₄, MgSO₄, Al₂(SO₄)₃

    • insoluble: PbSO₄, BaSO₄, CaSO₄ is slightly soluble.

  • Common oxides, hydroxides and carbonates are usually insoluble (e.g. Group 2 and Transition Metals) except those of the Group 1 Alkali Metals sodium, potassium etc. and ammonium:

    • soluble: K₂O, KOH, NaOH, NH₄OH actually NH_3(aq), Na₂CO₃, (NH₄)₂CO₃  

    • insoluble: MgO, CuO, ZnO, Mg(OH)₂, Fe(OH)₂, Cu(OH)₂, CuCO₃, ZnCO₃, CaCO₃

Method 6d. Making a salt by direct combination of elements

• (5) By direct combination of the elements to form anhydrous salts e.g. if dry chlorine gas Cl₂ is passed over heated iron or aluminium, the chloride is produced. The experiment (shown above) should be done very carefully by the teacher in a fume cupboard.

  • 2Al_(s) + 3Cl_2(g) ==> 2AlCl_3(s)

  • The aluminium can burn intensely with a violet flame, white fumes of aluminium chloride sublime from the hot reacted aluminium and the white solid forms on the cold surface of the flask (its often discoloured yellow from the trace chlorides of copper or iron that may be formed).

  • 2Fe_(s) + 3Cl_2(g) ==> 2FeCl_3(s)

  • The iron (e.g. as steel wool) glows red and brown fumes of iron(III) chloride stream off, the brown solid collects on the cold flask surface.

  • Note (i): Both these chlorides react exothermically and hydrolyse with water to give the metal hydroxide and fumes of hydrogen chloride, and so dry conditions are needed.

  • Note (ii): Both these chlorides cannot be made in an anhydrous form from aqueous solution neutralisation. This is because on evaporation the compounds contain 'water of crystallisation'. On heating the hydrated salt  hydrolyses and decomposes into water, the oxide or hydroxide and fumes of hydrogen chloride, and maybe some impure anhydrous chloride, basically it a mess in terms of trying to make pure AlCl₃ and FeCl₃ in this way.

• Water of crystallisation is dealt with in section 8.