DOC BROWN'S Science-CHEMISTRY HOMEPAGE KS3 SCIENCE QUIZZES and WORKSHEETS (~US grades 6-8)
GCSE SCIENCE help links GCSE ADDITIONAL SCIENCE help links
KS3 BIOLOGY Quizzes KS3 CHEMISTRY Quizzes & Worksheets KS3 PHYSICS Quizzes
KS4 Science GCSE/IGCSE CHEMISTRY NOTES (~US grades 8-10) KS4 Science GCSE/IGCSE CHEMISTRY QUIZZES and WORKSHEETS (~US grades 8-10) ADVANCED LEVEL CHEMISTRY QUIZZES and WORKSHEETS (~US grades 11-12)
Custom Search

Doc Brown's Chemistry GCSE/IGCSE Science-Chemistry Revision Notes

The pH scale of acidity and alkalinity, acids, alkalis, salts and neutralisation

6. Methods of making salts

How do we make salts? What preparations are available to us?

Four basic methods for preparing salts are described on this page, with annotated diagrams.

Method (a) Making a salt by neutralising a soluble acid with a soluble base (alkali).

Method (b) preparing a salt by reacting an acid with a metal or with an insoluble base.

Method (c) Preparing an Insoluble Salt.

Method (d) Making a salt by directly combining its constituent elements

GCSE/IGCSE Sub-index: Index of all pH, Acids, Alkalis, Salts Notes 1. Examples of acid-alkali chemistry : 2. pH scale, indicators, ionic theory of acids-alkali neutralisation : 3. pH examples of acid, neutral or alkaline solutions : 4. Acid reactions with metals/oxides/hydroxides/carbonates and neutralisation reactions : 5. Reactions of bases-alkalis like sodium hydroxide : 6. Four methods of making salts : 7. Changes in pH in a neutralisation : 8. Important formulae, salt solubility and water of crystallisation : 9. Further examples of word/symbol equations for salt preparations : 10. More on Acid-Base Theory and Weak and Strong Acids

See also Advanced Level Chemistry Students Acid-Base Revision Notes - use index


6. METHODS of MAKING SALTS - salt preparation procedures

Salt solubility affects the method you choose to make a salt and so section 8. contains tables of information-data on salt solubility which will help you decide on the method to prepare a salt.

top

6a-b. Two Methods of Making Water Soluble Salts


6a. METHOD (a) Neutralising a soluble acid with a soluble base (alkali) to give a soluble salt

The extraction of lead metal from lead carbonate ores Extraction of Lead

soluble salt preparation from soluble base-acid neutralisation

e.g. the hydroxide of an alkali metal like sodium hydroxide or ammonia solution. Steps (1) to (3) below is called a titration.

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

NaOH sodium hydroxide, KOH potassium hydroxide

Typical examples shown by the word and symbol equations below include ...

sodium hydroxide + hydrochloric acid ==> sodium chloride + water

NaOH(aq) + HCl(aq) ==> NaCl(aq) + H2O(l)

or

potassium hydroxide + sulphuric acid ==> potassium sulphate + water

2KOH(aq) + H2SO4(aq) ==> K2SO4(aq) + 2H2O(l)

or

sodium hydroxide + nitric acid ==> sodium nitrate + water

NaOH(aq) + HNO3(aq) ==> NaNO3(aq) + 2H2O(l)

(c) doc b More examples of neutralization equations are given in section 4.

METHOD (a)

(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) (c) doc b The pH changes in this preparation are described in section 7

(viii) Salt solubility affects the method you choose to make a salt and so (c) doc b section 8. contains tables of information-data on salt solubility which will help you decide on the method to prepare a salt.

top


6b. METHOD (b) Reacting an acid with a metal or with an insoluble base to give a soluble salt

soluble salt preparation from insoluble base-acid neutralisatione.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.

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, CuO copper(II) oxide, CuCO3 copper(II) carbonate, PbCO3 lead(II) carbonate (with nitric acid to make lead(II) nitrate), FeCO3 iron(II) carbonate (to make iron(II) salts), MnCO3 manganese(II) carbonate

Typical examples shown by the word and symbol equations below include ...

copper(II) oxide + sulphuric acid ==> copper(II) sulphate + water

CuO + H2SO4 ==> CuSO+ H2O

CuO(s) + H2SO4(aq) ==> CuSO4(aq) + H2O(l)

magnesium hydroxide + sulfuric acid ==> magnesium sulfate + water

Mg(OH)2 + H2SO4 ==> MgSO+ 2H2O

Mg(OH)2(s) + H2SO4(aq) ==> MgSO4(aq) + 2H2O(l)

 

Zinc carbonate + nitric acid ==> zinc nitrate + water + carbon dioxide

ZnCO3 + 2HNO3 ==> Zn(NO3)2 + H2O + CO2

ZnCO3(s) + 2HNO3(aq) ==> Zn(NO3)2(aq) + H2O(l) + CO2 (g)

  • Carbonates are frequently used in this method of salt making, e.g. using copper carbonate

    • The equations are give with, and without sate symbols.

  • copper(II) carbonate + hydrochloric acid ==> Copper(II) chloride + water + carbon dioxide

    • CuCO3 + 2HCl ==> CuCl2 + H2O + CO2

      • CuCO3(s) + 2HCl(aq) ==> CuCl2(aq) + H2O(l) + CO2(g)

    • and with sulphuric acid a blue solution of copper(II) sulphate is formed.

  • copper(II) carbonate + sulphuric acid ==> Copper(II) sulphate + water + carbon dioxide

    • CuCO3 + H2SO4 ==> CuSO4 + H2O + CO2

      • CuCO3(s) + H2SO4(aq) ==> CuSO4(aq) + H2O(l) + CO2(g)

  • copper(II) carbonate + nitric acid ==> Copper(II) nitrate + water + carbon dioxide

    • CuCO3 + 2HNO3 ==> Cu(NO3)2 + H2O + CO2

      • CuCO3(s) + 2HNO3(aq) ==> CuSO4(aq) + H2O(l) + CO2(g)

  • Similar equations for other carbonates to give soluble salts which can be crystallised from solution e.g.

    • calcium carbonate CaCO3, to make two salts - calcium chloride/nitrate (calcium sulfate is not very soluble)

    • iron(II) carbonate FeCO3, to make three salts - iron(II) chloride/sulfate/nitrate

    • magnesium carbonate MgCO3, to make three salts - magnesium chloride/sulfate/nitrate

    • manganese(II) carbonate MnCO3, to make three salts - manganese(II) chloride/sulfate/nitrate

    • zinc carbonate ZnCO3, to make three salts - zinc chloride/sulfate/nitrate

    • lead(II) carbonate PbCO3, only nitric acid to make lead(II) nitrate, lead(II) chloride and lead(II) sulfate are insoluble and must be prepared by method (c)

 

(c) doc b More examples of neutralization equations are given in section 4.

METHOD (b)

(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 (c) doc b(c) doc bChemical Calculations Part 14

Salt solubility affects the method you choose to make a salt and so  section 8. contains tables of information-data on salt solubility which will help you decide on the method to prepare a salt.

top


6c. Method (c) Preparing an Insoluble Salt

  • How can we make an insoluble salt? How do we prepare an insoluble salt from two soluble compounds?

  • This section describes the preparation of insoluble salts like silver chloride, lead(II) iodide (lead iodide), calcium carbonate, barium sulfate (barium sulphate), lead(II) sulfate (lead sulphate),

  • METHOD (c)

  • An insoluble salt can be made by mixing two solutions of soluble salts in a process is called precipitation.

    • The method is quite simple - illustrated above, assuming in this case the insoluble salt is colourless-white.

    • One solution contains the 1st required ion, and the other solution contains the 2nd required ion.

    • The two solutions of SOLUBLE compounds are mixed together so the INSOLUBLE salt precipitate is formed.

    • 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 carefully removed from the filter paper to be dried e.g. left out in dry room or warmed in a pre-heated oven.

  • 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

      • AgNO3(aq) + NaCl(aq) ==> AgCl(s) + NaNO3(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.

        • (c) doc b 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

      • BaCl2(aq) + 2AgNO3(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) ==> PbI2(s) + 2KNO3(aq)

      • which can be written as

      • Pb2+(aq) + 2NO3-(aq) + 2K+(aq) + 2I-(aq) ==> PbI2(s) + 2K+(aq) + 2NO3-(aq)

      • the ionic equation is: Pb2+(aq) + 2I-(aq) ==> PbI2(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) ==> PbCl2(s) + 2HNO3(aq)

        • Pb2+(aq) + 2NO3-(aq) + 2H+(aq) + 2Cl-(aq) ==> PbCl2(s) + 2H+(aq) + 2NO3-(aq)

        • the ionic equation is: Pb2+(aq) + 2Cl-(aq) ==> PbCl2(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) ==> PbBr2(s) + 2NaNO3(aq)

        • Pb2+(aq) + 2NO3-(aq) + 2Na+(aq) + 2Br-(aq) ==> PbBr2(s) + 2Na+(aq) + 2NO3-(aq)

        • the ionic equation is: Pb2+(aq) + 2Br-(aq) ==> PbBr2(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

      • CaCl2(aq) + Na2CO3(aq) ==> CaCO3(s) + 2NaCl(aq)

      • Ca2+(aq) + 2Cl-(aq) + 2Na+(aq) + CO32-(aq) ==> CaCO3(s) + 2Na+(aq) + 2Cl-(aq)

      • ionically: Ca2+(aq) + CO32-(aq) ==> CaCO3(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

      • BaCl2(aq) + H2SO4(aq) ==> BaSO4(s) + 2HCl(aq)

      • Ba2+(aq) + 2Cl-(aq) + 2H+(aq) + SO42-(aq) ==> BaSO4(s) + 2H+(aq) + 2Cl-(aq)

      • ionic equation: Ba2+(aq) + SO42-(aq) ==> BaSO4(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

        • BaCl2(aq) + Na2SO4(aq) ==> BaSO4(s) + 2NaCl(aq)

        • The ionic equation is the same: Ba2+(aq) + SO42-(aq) ==> BaSO4(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) + Na2SO4 (aq) ==> PbSO4 (s) + 2NaNO3 (aq)

      • ionically: Pb2+(aq) + SO42-(aq) ==> PbSO4(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 bases-alkalis oxides, hydroxides or carbonates: K2O, KOH, NaOH, NH4OH actually NH3(aq), Na2CO3, (NH4)2CO3  

      • insoluble bases - basic oxides, hydroxides or insoluble carbonates: MgO, CuO, ZnO, Mg(OH)2, Fe(OH)2, Cu(OH)2, CuCO3, ZnCO3, CaCO3

  • Salt solubility affects the method you choose to make a salt and so  section 8. contains tables of information-data on salt solubility which will help you decide on the method to prepare a salt.

top


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

The apparatus for the preparation of aluminium chlorise (c) doc b

  • How can we make aluminium chloride? How do we prepare iron(III) chloride?

  • METHOD (d)

  • These compounds can be made 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. These experiment preparations (shown above) should be done very carefully by the teacher in a fume cupboard.

    • 2Al(s) + 3Cl2(g) ==> 2AlCl3(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) + 3Cl2(g) ==> 2FeCl3(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.

top


Revision KS4 Science GCSE/IGCSE/O level Chemistry Information Study Notes for revising for AQA GCSE Science, Edexcel 360Science/IGCSE Chemistry & OCR 21stC Science, OCR Gateway Science  WJEC gcse science chemistry CCEA/CEA gcse science chemistry O Level Chemistry (revise courses equal to US grade 8, grade 9 grade 10) tuition help science chemistry courses revision guides

top

ALL Website content copyright Dr Phil Brown 2000-2014 All rights reserved on revision notes, images, puzzles, quizzes, worksheets, x-words etc. * Copying of website material is not permitted * chemhelp@tiscali.co.uk

Describing how to make a salt by four different methods, how to do it, what laboratory equipment is needed and how to crystallise or collect the salt in the end

Teach yourself chemistry online ALPHABETICAL SITE INDEX for chemistry

Alphabetical Index for Science Pages Content A B C D E F G H I J K L M N O P Q R S T U V W X Y Z

SITE HELP SEARCH - ENTER SPECIFIC WORDS/FORMULA etc.