Cation Analysis
 

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Sodium Na⁺

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Flame test: covered earlier.  There are no simple reagents that will precipitate sodium compounds.
 

Inorganic analysis introduction   
Preliminary tests   
Anion analysis
The nature of hydroxide precipitates 
 

Potassium K⁺

Flame test: covered earlier.  There are no simple reagents that will precipitate potassium compounds.
 

Magnesium Mg²⁺

Ions

Calcium Ca²⁺

Dilute sulphuric acid gives a white precipitate of calcium sulphate if the original solution is fairly concentrated:

Ca²⁺(aq)  +  SO₄^2-(aq)  "   CaSO₄(s)

Calcium sulphate solution is permanently hard water; its solubility is about 0.05 mol dm^-3 at 25^oC.


Sodium or potassium carbonate solution precipitates white calcium carbonate:

Ca²⁺(aq)  +  CO₃^2-(aq)  "  CaCO₃(s)
 

Ammonium ethanedioate solution precipitates white calcium ethanedioate from neutral or alkaline solutions. The precipitate dissolves readily in dilute acid:

Ca²⁺(aq)  +  C₂O₄^2-(aq)  "   CaC₂O₄(s)

Calcium ethanedioate (calcium oxalate) is found in rhubarb leaves - it is what makes them poisonous.
 

Chromium(III) [Cr(H₂O)₆]³⁺

Sodium hydroxide solution precipitates grey-green chromium(III) hydroxide, which reacts with excess NaOH to give a deep green solution of the chromite ion:

[Cr(H₂O)₆]³⁺(aq)  +   3OH^-(aq)  "   Cr(OH)₃(s)  +  6H₂O(l)

Cr(OH)₃(s)  +  3OH^-(aq)  "   [Cr(OH)₆]^3-(aq)

Ammonia solution precipitates grey-green chromium(III) hydroxide, which with excess ammonia very slowly (see below) forms pinkish solutions of ammines. Ammines are transition metal complexes with ammonia, NH₃. Do not confuse them with amines RNH₂.The initial equation is as the first one above; then

Cr(OH)₃(s)  +  4NH₃(aq)  +  2H₂O(l)  "  [Cr(NH₃)₄(H₂O)₂]³⁺(aq)  +  3OH^-(aq)

The exact composition of the ammine will depend on the amount of ammonia used. However, the reaction is so slow that several days are needed to see anything at all.

Oxidising agents convert chromium(III) to yellow chromate(VI) in the presence of alkali. Hydrogen peroxide is favoured:

2[Cr(H₂O)₆]³⁺(aq)  +   3H₂O₂(aq)  +  10OH^-(aq)  "  2CrO₄^2-(aq) +  14H₂O(l)

The test solution is made alkaline with NaOH solution, then an equal volume of ‘20 volume’ hydrogen peroxide solution is added and the mixture boiled. 20-volume hydrogen peroxide solution gives 20cm³ of oxygen per cm³ of solution that is decomposed to oxygen and water.
 

Manganese(II) [Mn(H₂O)₆]²⁺

Sodium hydroxide solution precipitates beige manganese(II) hydroxide. This is insoluble in excess reagent, and rapidly darkens owing to oxidation to hydrated manganese(IV) oxide:

[Mn(H₂O)₆]²⁺(aq)  +   2OH^-(aq)  "  Mn(OH)₂(s)  +  6H₂O(l)

4Mn(OH)₂(s)  +  2H₂O(l)  +  O₂(aq)  "  4MnO₂.H₂O(s)

Sodium or potassium carbonate solution gives a white precipitate of manganese(II) carbonate:

[Mn(H₂O)₆]²⁺(aq)  +   CO₃^2-(aq)  "  MnCO₃(s)  +  6H₂O(l)

Lead(IV) oxide in the presence of nitric acid converts manganese(II) into purple manganate(VII); an alternative oxidising agent is sodium bismuthate(V), NaBiO₃:

5PbO₂(s) + 2Mn²⁺(aq) + 6H⁺(aq)  "  2MnO₄^-(aq) + 5Pb²⁺(aq)  +  2H₂O(l)

5BiO₃^-(aq)  +  14H⁺(aq)  +  2Mn²⁺(aq)  "  5Bi³⁺(aq)  +  2MnO₄^-(aq)  +  7H₂O(l)

A small amount of lead(IV) oxide or of sodium bismuthate(V) is added to the test solution, 6 or so drops of concentrated nitric acid added, and the mixture boiled. Filtration will give a purple filtrate if Mn²⁺ was present.
 

Iron(III) [Fe(H₂O)₆]³⁺

Aqueous iron(III) ions are not [Fe(H₂O)₆]³⁺; this ion is an amethyst (pale purple) colour, and is found only in solids such as aluminium iron(III) sulphate. In solution the hexaaqua ion is readily deprotonated by solvent water; the solution is acidic, and the yellow ion is [Fe(H₂O)₅OH]²⁺:

[Fe(H₂O)₆]³⁺ + H₂O  "   [Fe(H₂O)₅OH]²⁺ + H₃O⁺.

Sodium hydroxide solution precipitates foxy-red iron(III) hydroxide, insoluble in excess alkali:

[Fe(H₂O)₆]³⁺(aq)  +  3OH^-(aq)  "  Fe(OH)₃(s)  +  6H₂O(l)

Ammonia solution reacts in the same way as sodium hydroxide – iron does not form complexes with ammonia.

Potassium hexacyanoferrate(II) precipitates dark blue potassium iron(III) hexacyanoferrate(II), Prussian Blue:

K⁺(aq)  +  Fe³⁺(aq)  +  [Fe(CN)₆]^4-(aq)  "  KFe^III[Fe^II(CN)₆](s)

Prussian Blue is the pigment that is used to print 'blueprints'. It was also used for Prussian army uniforms, and for the locomotives and rolling stock of the Somerset & Dorset Joint Railway, pre-1923.

Iron(II) [Fe(H₂O)₆]²⁺

Sodium hydroxide solution precipitates dirty green iron(II) hydroxide; on standing the surface of the precipitate turns foxy-red owing to air oxidation to iron(III) hydroxide:

[Fe(H₂O)₆]²⁺(aq)  +  2OH^-(aq)  "   Fe(OH)₂(s)  +  6H₂O(l)

Ammonia solution behaves similarly to sodium hydroxide – again there are no ammine complexes.

Potassium hexacyanoferrate(III) precipitates dark blue potassium iron(II) hexacyanoferrate(III), Turnbull’s Blue:

K⁺(aq)  +  Fe²⁺(aq)  +  [Fe(CN)₆]^3-(aq)  "   KFe^II[Fe^III(CN)₆](s)

Although Turnbull's Blue and Prussian Blue don't have quite the same colour, they are in fact the same substance.

Potassium manganate(VII) oxidises iron(II) to iron(III) in acidic solution. The purple manganate colour is lost and the resulting solution of iron(III) is yellow:

MnO₄^-(aq)  +  5Fe²⁺(aq)  +  8H⁺(aq)  "   Mn²⁺(aq)  +  5Fe³⁺(aq)  +  4H₂O(aq)

Potassium dichromate(VI) oxidises iron(II) to iron(III) in acidic solution. The orange dichromate colour is lost and the resulting solution is green owing to presence of chromium(III) which masks the iron(III) colour:

Cr₂O₇^2-(aq)  +  6Fe²⁺(aq)  +  14H⁺(aq)  "  2Cr³⁺(aq) +  6Fe³⁺(aq)  +  7H₂O(l)

Copper(II) [Cu(H₂O)₆]²⁺

Sodium hydroxide solution gives a pale blue precipitate usually described as copper(II) hydroxide:

[Cu(H₂O)₆]²⁺(aq)  +  2OH^-(aq)  "  Cu(OH)₂(s)  +  6H₂O(l)

In fact this precipitate is the basic salt – thus from copper(II) sulphate solution the product is basic copper sulphate:

2[Cu(H₂O)₆]²⁺(aq)  +  2OH^-(aq)  +  SO₄^2-(aq)  "  Cu(OH)₂.CuSO₄(s)  +  12H₂O(l)

Basic copper sulphate or basic copper carbonate is responsible for the green patina - verdigris - found on weathered copper roofs. The sulphate is present in industrial surroundings. To get true copper(II) hydroxide, the solution of the copper(II) salt must be added to excess sodium hydroxide; the precipitate is then blue rather than turquoise. In both cases CAUTIOUS warming of the mixture causes the precipitate to lose water and turn black as CuO(s) is formed. The loss of water by heating in an aqueous medium is unusual.

Ammonia solution initially gives a blue precipitate as for sodium hydroxide. Further addition of ammonia gives deep blue soluble cuprammine complexes whose composition depends on the amount of ammonia present. The reaction is usually represented:

Cu(OH)₂(s)  +  4NH₃(aq)  +  2H₂O(l) "  [Cu(NH₃)₄(H₂O)₂]²⁺(aq)  +2OH^-(aq)

The cuprammine solution has the bizarre property of being able to dissolve cellulose - bits of filter paper, for e example. The resulting viscous substance can be extruded into a bath of dilute sulphuric acid to give viscose fibre, which is a form of artificial silk.

Concentrated HCl or a saturated solution of sodium chloride gives a bright green solution containing CuCl₄^2- ions; dilution with water causes the solution to turn pale blue as the [Cu(H₂O)₆]²⁺ ion re-forms.

[Cu(H₂O)₆]²⁺(aq)  +  4Cl^-(aq)  "  CuCl₄^2-(aq)  +  6H₂O(l)

Sodium carbonate solution precipitates greenish-blue basic carbonates of indefinite composition.

Potassium iodide solution precipitates iodine and copper(I) iodide; this reaction is used volumetrically for the estimation of copper, the liberated iodine being titrated with sodium thiosulphate solution:

2[Cu(H₂O)₆]²⁺(aq)  +  4I^-(aq)  "  2CuI(s)  +  I₂(aq)  +  6H₂O(l)

The mixture turns a sludgy brown colour; addition of sodium thiosulphate solution leaves a creamy-pink precipitate of copper(I) iodide.

Zinc(II) [Zn(H₂O)₆]²⁺

Sodium hydroxide solution precipitates white zinc(II) hydroxide, easily soluble in excess sodium hydroxide to give sodium zincate(II), because the hydroxide is amphoteric:

[Zn(H₂O)₆]²⁺(aq)  +  2OH^-(aq)  "  Zn(OH)₂(s)  +  6H₂O(l)

Zn(OH)₂(s)  +  2OH^-(aq)  "  Zn(OH)₄^2-(aq)

Ammonia solution initially precipitates zinc(II) hydroxide, as with sodium hydroxide. Excess ammonia causes the precipitate to disappear owing to the formation of ammine complexes – a different reason from that in 10.1:

Zn(OH)₂(s)  +  4NH₃(aq)  "  [Zn(NH₃)₄]²⁺(aq)  +  2OH^-(aq)

Sodium carbonate solution precipitates a white basic carbonate:

5Zn²⁺(aq)  +  2CO₃^2-(aq)  +  6OH^-(aq)  "  2ZnCO₃.3Zn(OH)₂(s)

This latter substance is what you get when you buy zinc carbonate – which is why on heating it gives off water vapour as well as carbon dioxide.
 

Aluminium Al[H₂O] ³⁺

Sodium hydroxide solution precipitates white gelatinous aluminium hydroxide. This reacts with excess NaOH to give a colourless solution of sodium aluminate:

[Al(H₂O)₆] ³⁺(aq)  +  3OH^-(aq)  "  Al(OH)₃(s) + 6H₂O(l)

 Al(OH)₃(s)  +  3OH^-(aq)  "  [Al(OH)₆]^3-(aq)

Ammonia solution precipitates white aluminium hydroxide, as with NaOH; however it does not react further with excess ammonia.

Lead(II) Pb²⁺

Dilute hydrochloric acid or other soluble chlorides precipitate white lead(II) chloride from moderately concentrated solutions of lead(II) salts.

Pb²⁺(aq)  +  2Cl^-(aq)  "   PbCl₂(s)

The precipitate dissolves in an excess of concentrated hydrochloric acid to give yellow solutions containing various species including PbCl₃^- and PbCl₄^2-.

Sodium hydroxide solution gives a white precipitate of lead(II) hydroxide which is soluble in excess NaOH to give colourless sodium plumbate(II)

Pb²⁺(aq)  +  2OH^-(aq)  "   Pb(OH)₂(s)

Pb(OH)₂(s) +  2OH^-(aq)  "   [Pb(OH)₄]^2-(aq)

Ammonia solution precipitates white lead(II) hydroxide; this does not dissolve in excess ammonia, because no complexes are formed and ammonia is not a strong enough base to bring out the amphoteric properties of lead(II) hydroxide.

Dilute sulphuric acid or other soluble sulphates give a white precipitate of lead(II) sulphate:

Pb²⁺(aq)  +  SO₄^2-(aq)  "  PbSO₄(s)

Potassium chromate(VI) solution gives a bright yellow precipitate of lead(II) chromate(VI):

Pb²⁺(aq)  +  CrO₄^2-(aq)  "  PbCrO₄(s)

Potassium iodide solution gives a bright yellow precipitate of lead(II) iodide. This will dissolve in boiling water to give a colourless solution – the colour comes from interactions in the crystal lattice rather than from coloured ions:

Pb²⁺(aq)  +  2I^-(aq)  "   PbI₂(s)

Ammonium NH₄⁺

Heat: all ammonium salts decompose on heating; ammonium nitrate may explode. Ammonium chloride and sulphate give products that recombine on cooling, so that the salts apparently sublime.

NH₄Cl(s)  "   NH₃(g)  +  HCl(g)

NH₄NO₃(s)  " N₂O(g)  +  2H₂O(g)

2(NH₄)₂SO₄(s)  D 2NH₃(g)  +  SO₃(g)  +   H₂O(g)

Ammonium dichromate(VI) decomposes spectacularly on ignition in a reaction that is oxidation of the cation by the anion; the initially orange solid gives a fluffy green product of much larger volume:

(NH₄)₂Cr₂O₇(s)  "  N₂(g)  +  Cr₂O₃(s)  +  4H₂O(g)

Alkalis (sodium hydroxide, calcium hydroxide) liberate ammonia from ammonium salts on warming with the solution, or even from a mixture of the solids. This is because OH^- is a stronger base than ammonia, so removes a hydrogen ion from the ammonium ion:

NH₄⁺(aq)  +  OH^-(aq)  "   NH₃(g)  +  H₂O(l)

The test solution is warmed with sodium hydroxide solution and the vapours tested with moist red litmus paper. It is important to test the vapours immediately heating begins, since the ammonia is lost very quickly and by the time the solution boils it may well have all gone.

Cobalt(II) [Co(H₂O)₆]²⁺

The (+2) state is the most stable for simple cobalt salts; they are coloured pink or blue. The pink colour of the hydrated ion [Co(H₂O)₆]²⁺ changes to blue on heating, on dehydration, or in the presence of concentrated acids.

Sodium hydroxide solution precipitates a blue basic salt which on warming with excess sodium hydroxide forms solid pink cobalt(II) hydroxide. With a solution of cobalt(II) nitrate the equations are:

Basic salts, usually hydroxides or carbonates, contain not only the expected hydroxide or carbonate ion but also the anion from the original salt. The precipitate from cobalt(II) nitrate solution is therefore different from that which would be given from cobalt(II) chloride. The latter would be Co(OH)Cl.

If the pink precipitate is boiled for some time (or is warmed with hydrogen peroxide solution) it is converted to brownish-black cobalt(III) hydroxide:

4Co(OH)₂ (s)  +  2H₂O (l)  +  O₂ (aq) "  4Co(OH)₃ (s)

Ammonia solution gives a blue basic precipitate as with sodium hydroxide; excess ammonia converts this to a brown solution of a cobaltammine which turns red on exposure to air giving an ammine of cobalt(III). The overall reaction is:

Co(OH)NO₃ (s) + 28NH₃(aq) + 6H₂O(l) + O₂(aq)   "  4[Co(NH₃)₆](OH)₃(aq) + 4NH₄NO₃(aq)

Nickel(II) [Ni(H₂O)₆]²⁺

Sodium hydroxide solution gives a green gelatinous precipitate of nickel(II) hydroxide, insoluble in excess of the reagent:

[Ni(H₂O)₆]²⁺ (aq)  +  2OH ^– (aq)  "  Ni(OH)₂ (s)  +  6H₂O(l)

Ammonia solution gives a green precipitate of a basic salt which dissolves readily in excess ammonia solution to form a blue-violet solution of complex nickel ammines. With nickel(II) chloride the reactions are:

[Ni(H₂O)₆]²⁺(aq) + Cl ^– (aq) + NH₃(aq) "  Ni(OH)Cl(s) + NH₄⁺(aq) + 5H₂O(l)

Ni(OH)Cl(s)  + 5NH₃(aq) + H₂O(l) "  [Ni(NH₃)₄](OH)₂(aq)  +  NH₄⁺(aq) + Cl ^– (aq)

Ni(OH)Cl(s)  + 7NH₃(aq) + H₂O(l) "  [Ni(NH₃)₆](OH)₂(aq)  +  NH₄⁺(aq) + Cl ^– (aq)