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GCSE Chemistry Notes: Preparing a salt by a precipitation reaction

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method describe preparation of insoluble salt by precipitation apparatus reactants equation double decomposition gcse chemistry igcse international gcse O level

6c. MAKING an INSOLUBLE SALT BY A PRECIPITATION REACTION

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Sub-index for this page

Introduction and solubilities of salts

The method of making a salt by precipitation

The preparation of silver chloride

The preparation of lead(II) iodide

The preparation of calcium carbonate

The preparation of barium sulfate

The preparation of lead(II) sulfate

The preparation of calcium sulfate

The uses of some insoluble salts

Method (a) Making a salt by neutralising a soluble acid with a soluble base (alkali) – neutralisation reaction

Method (b) preparing a salt by reacting an acid with a metal or an insoluble base – oxide, hydroxide or carbonate

Method (c) Preparing an insoluble salt by mixing solutions of two soluble compounds (this page)

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

Doc Brown's chemistry revision notes: basic school chemistry science GCSE chemistry, IGCSE  chemistry, O level & ~US grades 8, 9, 10 school science courses for ~14-16 year old science students for national examinations in chemistry topics including acids bases alkalis salts preparations reactions

6. METHODS of MAKING SALTS – salt preparation procedures

6c. Method (c) Preparing an Insoluble Salt by a precipitation reaction

The procedure involves making an insoluble salt my mixing solutions of soluble compounds to form a precipitate.

The two soluble compounds must each provide one of the constituent ions of the desired insoluble salt which precipitates out when the solutions are mixed.

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

Note that several precipitation reactions are used as simple tests for e.g. for sulfate, chloride, bromide, iodide

Salt solubility affects the method you choose to make a salt, the table below will help you decide on the method

A solubility guide for salts and other compounds

Information required to decide on the method used to prepare a salt
salts and other compounds

solubility?
common salts of sodium, potassium and ammonium ions usually soluble in water
common sulfates (sulfates) usually quite soluble except for calcium sulfate (slightly soluble), lead sulfate and barium sulfate are both insoluble
common chlorides (similar rule for bromides and iodides) usually soluble except for insoluble lead(II) chloride and silver chloride
common nitrates all soluble
common carbonates most metal carbonates are insoluble apart from sodium & potassium carbonate.  Ammonium carbonate is also soluble.
common hydroxides most metal hydroxides are insoluble apart from sodium, potassium and ammonium hydroxide

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, 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

THE METHOD

How to make an insoluble salt from two soluble compounds (which may be a soluble salt or acid)

MAKING AN INSOLUBLE SALT - by mixing solutions of two soluble substances

These reactions are sometimes described as a DOUBLE DECOMPOSITION reaction - the term is explained in examples.

  • 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 AgCl, lead(II) chloride (lead chloride) PbCl2, lead(II) iodide (lead iodide) PbI2, calcium carbonate CaCO3, barium sulfate (barium sulfate) BaSO4, lead(II) sulfate (lead sulfate, lead sulfate)PbSO4, and 'slightly soluble' calcium sulfate (calcium sulfate) CaSO4, which can all be made by a precipitation reaction.

    • All of the above insoluble salts are white (as in diagram), except lead(II) iodide which is yellow.

  • Many of the salt precipitate are WHITE, but lead iodide is pale yellow.

method describe preparation of insoluble salt by precipitation apparatus reactants equation double decomposition gcse chemistry igcse international gcse O level

  • METHOD (c) Preparation of a insoluble salt using a precipitation reaction

  • Don't forget to wear safety glasses or goggles when doing this preparation.

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

    • These reactions are often described as double decomposition reactions - in the reaction, the ions swap who they are combined with in terms of the two ions precipitated as the salt and the two ions left in solution.

    • 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.

    • STEP 1. So you must prepare two solutions of soluble compounds, each of which provides one of the two ions required to combine and precipitate out as the insoluble salt.

      • Each soluble compound is weighed out into its own beaker and dissolved in a suitable volume of water until the solutions are both quite clear.

      • One solution is then  poured into the other, order doesn't really matter.

      • The two solutions of SOLUBLE compounds must be thoroughly mixed together to ensure all the reactants are used up, so the maximum amount of INSOLUBLE salt precipitate is formed.

      • You see the two original clear solutions on mixing forming a cloudy mixture as the insoluble compound is formed, known as the precipitate.

    • STEP 2.The mixture is then carefully poured into a funnel holding a filter paper.

      • The precipitated salt can then be filtered off with the filter funnel and paper.

    • STEP 3.While still in the filter paper and funnel, the collected solid precipitate is washed with distilled/deionised water to remove any remaining soluble salt impurities and just the damp, but otherwise pure, insoluble salt is left.

    • STEP 4.The precipitate is then carefully removed from the filter paper into a clean dish or basin to be dried e.g. left out in a dry room or warmed in a pre–heated oven.

  • This four step procedure applies to the precipitation reactions below, to prepare a variety of insoluble salts.

  • Examples ...

    • (i) Silver chloride is made by mixing solutions of solutions of silver nitrate and sodium chloride - but it can be any soluble chloride salt solution or even dilute hydrochloric acid (HCl). Silver nitrate is one of the few common soluble salts of silver.

    • method describe preparation of insoluble salt silver chloride AgCl by precipitation apparatus reactants equation double decomposition gcse chemistry igcse international gcse O level

      • The silver nitrate provides the silver ion and the sodium chloride provides the chloride ion to prepare the insoluble salt silver chloride which forms as a white precipitate.

      • 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,

        • The 'active ions' and resulting precipitate are highlighted in yellow.

        • Describing this as a double decomposition reaction:

        • In the reaction, the ions swap who they are combined with in terms of the two ions precipitated as the salt (silver Ag+ and chloride Cl-) and the two ions left in solution (sodium Na+ and nitrate NO3-) - the silver ion is no longer with the nitrate ion and the hydrogen ion is no longer with the chloride ion.

      • so the proper ionic equation is simply: Ag+(aq) + Cl–(aq) ==> AgCl(s)

        • Note (i) the use of state symbols (aq) and (s) AND

        • (ii) that ionic equations omit ions that do not change there chemical or physical state.

          • You must NOT include the spectator ions in an ionic equation!

        • 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

      • You can just use dilute hydrochloric acid and silver nitrate solution to make insoluble silver chloride by precipitation.

      • The silver nitrate provides the silver ion and the hydrochloric acid provides the chloride ion to prepare the insoluble salt silver chloride which forms as a white precipitate.

      • silver nitrate + hydrochloric acid ==> silver chloride + nitric acid

      • AgNO3(aq) + HCl(aq) ==> AgCl(s) + HNO3(aq)

        •

      • Describing this as a double decomposition reaction:

        • In the reaction, the ions swap who they are combined with in terms of the two ions precipitated as the salt (silver Ag+ and chloride Cl-) and the two ions left in solution (hydrogen H+ and nitrate NO3-) - the silver ion is no longer with the nitrate ion and the hydrogen ion is no longer with the chloride ion.

        • You can apply this idea of double decomposition to all the reactions described below.

      • In terms of ions the reaction can be written as

      • Ag+NO3–(aq) + H+Cl–(aq) ==> AgCl(s) + H+NO3–(aq)

      • or: Ag+(aq) + NO3–(aq) + H+(aq) + Cl–(aq) ==> AgCl(s) + H+(aq) + NO3–(aq)

        •

      • BUT, the spectator ions are nitrate NO3– and hydrogen H+ which do not change at all - they remain in solution,

        • The 'active ions' and resulting precipitate are highlighted in yellow.

      • so the proper ionic equation is simply: Ag+(aq) + Cl–(aq) ==> AgCl(s)

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

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

        • Can you describe this as a double decomposition reaction?

      • 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)

        • You can prepare the insoluble salts silver bromide and silver iodide in a similar way by precipitation.

          •

        • You can make silver bromide by mixing solutions of silver nitrate and potassium bromide.

        • The silver nitrate provides the silver ion and the potassium bromide provides the bromide ion to prepare the insoluble salt silver bromide which forms as a cream precipitate.

        • silver nitrate + potassium bromide ==> silver bromide + potassium nitrate

        • AgNO3(aq) + KBr(aq) ==> AgBr(s) + KNO3(aq)

          •

        • the spectator ions are nitrate NO3– and potassium K+ which do not change at all,

        • so the proper ionic equation is simply: Ag+(aq) + Br–(aq) ==> AgBr(s)

        • Similarly you can make silver iodide by mixing solutions of silver nitrate and potassium iodide.

        • The silver nitrate provides the silver ion and the potassium iodide provides the iodide ion to prepare the insoluble salt silver iodide which forms as a pale yellow precipitate.

        • silver nitrate + potassium iodide ==> silver iodide + potassium nitrate

        • AgNO3(aq) + KI(aq) ==> AgI(s) + KNO3(aq)

          •

        • the spectator ions are nitrate NO3– and potassium K+ which do not change at all,

        • so the proper ionic equation is simply: Ag+(aq) + I–(aq) ==> AgI(s)

        • TOP OF PAGE and sub-index for page

    • (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 - but it can be any soluble iodide salt solution mixed with lead nitrate - one of the few common soluble salts of lead.

    • method describe preparation of insoluble salt lead iodide PbI2 lead(II) iodide by precipitation apparatus reactants equation double decomposition gcse chemistry igcse international gcse O level

      • The

        • lead nitrate supplies the lead ion and the potassium iodide the iodide ion to make the insoluble salt lead iodide, which is a yellow crystalline solid.

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

        • Can you describe this as a double decomposition reaction?

      • 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

        • The hydrochloric acid supplies the chloride ion.

        • 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 proper ionic equation is: Pb2+(aq) + 2Cl–(aq) ==> PbCl2(s)

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

      • Similarly you can make lead(II) chloride by e.g. using dilute sodium chloride solution

        • The silver nitrate provides the silver ion and the sodium chloride provides the chloride ion to prepare the insoluble salt lead chloride which forms as a white precipitate.

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

        • Pb(NO3)2(aq) + 2NaCl(aq) ==> PbCl2(s) + 2NaNO3(aq)

        • Pb2+(aq) + 2NO3–(aq) + 2Na+(aq) + 2Cl–(aq) ==> PbCl2(s) + 2Na+(aq) + 2NO3–(aq)

          •

        • the proper ionic equation is: Pb2+(aq) + 2Cl–(aq) ==> PbCl2(s)

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

      • You can make lead(II) bromide by e.g. using sodium bromide solution

        • 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 proper ionic equation is: Pb2+(aq) + 2Br–(aq) ==> PbBr2(s)

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

        • The silver nitrate provides the silver ion and the sodium bromide provides the bromide ion to prepare the insoluble salt silver bromide which forms as a cream precipitate.

        • You can also use potassium bromide, just swap the Na for a K.

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

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

        • Pb(NO3)2(aq) + 2NaCl(aq) ==> PbCl2(s) + 2NaNO3(aq)

        • Pb2+(aq) + 2NO3–(aq) + 2Na+(aq) + 2Cl–(aq) ==> PbCl2(s) + 2Na+(aq) + 2NO3–(aq)

          •

        • the proper ionic equation is: Pb2+(aq) + 2Cl–(aq) ==> PbCl2(s)

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

        • TOP OF PAGE and sub-index for page

          •

    • (iii) Calcium carbonate, a white precipitate, forms on e.g. mixing calcium chloride solution and sodium carbonate solutions - but it can be any soluble calcium salt solution mixed with a solution of any soluble carbonate.

    • method describe preparation of insoluble salt calcium carbonate CaCO3 by precipitation apparatus reactants equation double decomposition gcse chemistry igcse international gcse O level

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

        • Can you describe this as a double decomposition reaction?

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

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

        •

      • the ionic equation is: Ca2+(aq) + CO32–(aq) ==> CaCO3(s)

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

      • TOP OF PAGE and sub-index for page

    • (iv) Barium sulfate (barium sulphate), a white precipitate, forms on mixing e.g. barium chloride solution and dilute sulfuric acid - but it can be any soluble barium salt solution mixed with a solution of any soluble sulfate salt solution.

    • method describe preparation of insoluble salt barium sulfate BaSO4 by precipitation apparatus reactants equation double decomposition gcse chemistry igcse international gcse O level

      • barium chloride + sulfuric acid ==> barium sulfate + hydrochloric acid

        • Can you describe this as a double decomposition reaction?

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

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

        •

      • so the ionic equation is: Ba2+(aq) + SO42–(aq) ==> BaSO4(s)

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

        • Or you can use sulfate salts like sodium sulfate, 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–

        • You can do exactly the same preparation using potassium sulfate and barium nitrate or barium chloride, basically any solutions of a soluble barium salt and a soluble sulfate salt can be used to prepare barium sulfate.

        • e.g.     barium nitrate + potassium sulfate ==> barium sulfate + potassium nitrate

        • or        barium chloride + potassium sulfate ==> barium sulfate + potassium chloride

        • See at the end of the page the uses of barium sulfate.

        • TOP OF PAGE and sub-index for page

    • (v) Lead(II) sulfate (lead sulfate), a white precipitate, forms in a similar way e.g. mixing lead(II) nitrate solution with sodium sulfate solution - but it can be any soluble lead salt mixed with a solution of any soluble sulfate salt or dilute sulfuric acid solution.

    • method describe preparation of insoluble salt lead sulfate PbSO4 lead(II) sulfate by precipitation apparatus reactants equation double decomposition gcse chemistry igcse international gcse O level

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

      • Pb(NO3)2 (aq) + Na2SO4(aq) ==> PbSO4(s) + 2NaNO3 (aq)

        •

      • ionic equation: Pb2+(aq) + SO42–(aq) ==> PbSO4(s)

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

      • Lead(II) sulfate can also be precipitated using dilute sulfuric acid.

        • lead(II) nitrate + sodium sulfate ==> lead(II) sulfate + nitric acid

        • Pb(NO3)2 (aq) + H2SO4(aq) ==> PbSO4(s) + 2HNO3 (aq)

        • ionic equation: Pb2+(aq) + SO42–(aq) ==> PbSO4(s)

          •

        • because the spectator ions are hydrogen H+ and nitrate NO3– ions

        • TOP OF PAGE and sub-index for page

    • (vi) Calcium sulfate (calcium sulfate), a white precipitate, forms on mixing e.g. calcium chloride solution and dilute sulfuric acid - but it can be any soluble calcium salt solution mixed with a solution of any soluble sulfate salt.

    • method describe preparation of insoluble salt calcium sulfate CaSO4 by precipitation apparatus reactants equation double decomposition gcse chemistry igcse international gcse O level

      • calcium chloride + sulfuric acid ==> calcium sulfate + hydrochloric acid

      • CaCl2(aq) + H2SO4(aq) ==> CaSO4(s) + 2HCl(aq)

      • Ca2+(aq) + 2Cl–(aq) + 2H+(aq) + SO42–(aq) ==> CaSO4(s) + 2H+(aq) + 2Cl–(aq)

      • so the ionic equation is: Ca2+(aq) + SO42–(aq) ==> CaSO4(s)

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

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

        • calcium chloride + sodium sulfate ==> calcium sulfate + sodium chloride

        • CaCl2(aq) + Na2SO4(aq) ==> CaSO4(s) + 2NaCl(aq)

        • The ionic equation is the same: Ca2+(aq) + SO42–(aq) ==> CaSO4(s)

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

          •

        • You can do exactly the same preparation using potassium sulfate and calcium nitrate or calcium chloride, basically any solutions of a soluble calcium salt and a soluble sulfate salt can be used to prepare and precipitate calcium sulfate.

          • The yield is a little lower than the theoretical because calcium sulfate is slightly soluble in water.

        • See at the end of the page the uses of barium sulfate.

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

    • TOP OF PAGE and sub-index for page

  • Some uses of insoluble salts (their preparation has been described above)

    • Calcium sulfate CaSO4 (calcium sulfate) is used in plaster of Paris and plaster for domestic wall covering.

    • Barium sulfate BaSO4 (barium sulfate)

      • Barium sulfates quite a dense material is used in with X-rays for particular medical examinations.

      • Barium sulfate, like all barium salts is toxic, BUT, because it is insoluble, non of it is absorbed by the body into the bloodstream and eventually it will pass right through the gut system and be expelled in the normal faeces.

      • X-rays are used to investigate bone structure e.g. fractured or broken bones and the technique works because bone material is too dense to let weak X-rays through, so on shooting an X-ray photograph you get the bone structure as a sort of shadow effect where the X-rays have been absorbed by the bone.

      • Therefore, normally you can't use X-rays to examine soft tissue areas like the gut.

      • However, if the patient takes a 'barium meal', a thick harmless fluid containing a suspension of the insoluble barium sulfate, this can pass through the stomach and into the gut.

      • Therefore it is then possible to X-ray the intestinal gut system because the barium sulfate absorbs the X-rays just like bone.

      • So, because barium sulfate is opaque to X-rays, the X-ray photograph via the barium meal will highlight the structure of the tissue lining of the gut and show up any structural abnormalities and blockages.

GCSE/IGCSE Acids & Alkalis revision notes sub–index: Index of all pH, Acids, Alkalis, Salts Notes 1. Examples of everyday acids, alkalis, salts, pH of solution, hazard warning signs : 2. pH scale, indicators, ionic theory of acids–alkali neutralisation : 4. Reactions of acids with metals/oxides/hydroxides/carbonates, neutralisation reactions : 5. Reactions of bases–alkalis like ammonia & sodium hydroxide : 6. Four methods of making salts : 7. Changes in pH in a neutralisation, choice and use of indicators : 8. Important formulae of compounds, salt solubility and water of crystallisation : 10. More on Acid–Base Theory and Weak and Strong Acids

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

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