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Chloride Cl^-
AgCl, PbCl₂, Hg₂Cl₂ and CuCl are insoluble in water. Concentrated sulphuric acid liberates steamy acidic fumes of HCl from solid chlorides: NaCl(s) + H₂SO₄(l) " NaHSO₄(s) + HCl(g) Silver nitrate solution added to a solution of a chloride that has been acidified (test with blue litmus paper) with dilute nitric acid gives a white precipitate of silver chloride. The precipitate is readily soluble in dilute ammonia or in sodium thiosulphate solution: Ag⁺(aq) + Cl^-(aq) " AgCl(s) AgCl(s) + 2NH₃(aq) " [Ag(NH₃)₂]⁺(aq) + Cl^-(aq) AgCl(s) + 2S₂O₃^2-(aq) " [Ag(S₂O₃)₂]^3-(aq) + Cl^-(aq) Acidification with nitric acid is necessary to eliminate carbonate or sulphite, both of which interfere with the test by giving spurious precipitates. Concentrated solutions of sulphates can give a precipitate of silver sulphate in this test; its appearance is wholly different from AgCl. The latter is truly white; the sulphate is a pearly white, rather like pearlescent nail varnish.


Bromide Br^-
AgBr, PbBr₂, Hg₂Br₂ and CuBr are insoluble in water. Concentrated sulphuric acid gives a mixture of hydrogen bromide, bromine and sulphur dioxide with solid bromides; the HBr produced is oxidised by sulphuric acid. The mixture evolves steamy brownish acidic fumes: NaBr(s) + H₂SO₄(l) " NaHSO₄(s) + HBr(g) 2HBr + H₂SO₄ " Br₂ + SO₂ + 2H₂O Silver nitrate solution added to a solution of a bromide that has been acidified (test with blue litmus paper) with dilute nitric acid gives a cream precipitate of silver bromide. The precipitate is readily soluble in concentrated ammonia: Ag⁺(aq) + Br^-(aq) " AgBr(s) AgBr(s) + 2NH₃(aq) " [Ag(NH₃)₂]⁺(aq) + Br^-(aq) Acidification with nitric acid is necessary to eliminate carbonate or sulphite, both of which interfere with the test by giving spurious precipitates. Oxidising agents oxidise bromide to bromine, which is yellow or orange in aqueous solution. Bromine can be extracted from the solution by shaking with an immiscible organic solvent, for example hexane, the organic layer then turning orange. A suitable oxidising agent is sodium chlorate(I); this is added to the test solution, followed by a little dilute hydrochloric acid and a few cm³ of hexane: OCl^-(aq) + 2H⁺(aq) + 2Br^-(aq) " Br₂(aq) + Cl^-(aq) + H₂O(l)
Iodide I^-
AgI, PbI₂, Hg₂I₂ and CuI are insoluble in water. Concentrated sulphuric acid gives a mixture of hydrogen iodide, iodine, hydrogen sulphide, sulphur and sulphur dioxide when added to solid iodides; the HI produced is oxidised by sulphuric acid. The mixture evolves purple acidic fumes, turns to a brown slurry, and is a mess: NaI(s) + H₂SO₄(l) " NaHSO₄(s) + HI(g) 2HI + H₂SO₄ " I₂ + SO₂ + 2H₂O 6HI + H₂SO₄ " 3I₂ + S + 4H₂O 8HI + H₂SO₄ " 4I₂ + H₂S + 4H₂O There are no state symbols in these equations because the mixture is such a mess, and is more sulphuric acid than water. Silver nitrate solution added to a solution of an iodide that has been acidified (test with blue litmus paper) with dilute nitric acid gives a yellow precipitate of silver iodide. The precipitate is insoluble even in concentrated ammonia: Ag⁺(aq) + I^-(aq) " AgI(s) Acidification with nitric acid is necessary to eliminate carbonate or sulphite, both of which interfere with the test by giving spurious precipitates. Oxidising agents oxidise iodide to iodine, which is yellow or orange in aqueous solution. Iodine can be extracted from the solution by shaking with an immiscible organic solvent, for example hexane, the organic layer then turning purple. A suitable oxidising agent is sodium chlorate(I); this is added to the test solution, followed by a little dilute hydrochloric acid and a few cm³ of hexane: OCl^-(aq) + 2H⁺(aq) + 2Ir^-(aq) " I₂(aq) + Cl^-(aq) + H₂O(l) Lead ethanoate or lead nitrate solutions give a bright yellow precipitate of lead(II) iodide with iodides: Pb²⁺(aq) + 2I^-(aq) " PbI₂(s) The colour of lead(II) iodide comes from interactions in the lattice; dissolving the salt in boiling water gives a colourless solution which deposits glittering yellow plates on cooling. Solutions of copper(II) salts give a brown mixture containing iodine and copper(I) iodide when added to solutions of iodides. Addition of sodium thiosulphate solution decolourises the iodine and leaves pinkish-cream copper(I) iodide as a precipitate. 2Cu²⁺(aq) + 4I-(aq) " 2CuI(s) + I₂(aq) 2S₂O₃^2-(aq) + I₂(aq) " 2I^-(aq) + S₄O₆^2-(aq) This reaction is the basis for the volumetric estimation of copper; the iodine liberated from a known amount of a copper(II) solution is titrated with standard sodium thiosulphate solution.
Sulphite SO₃^2-
Sulphurous acid is considerably stronger than carbonic acid, so sulphites do not give the effervescence that is characteristic of carbonates when dilute acid is added. Dilute hydrochloric acid on warming with a sulphite evolves sulphur dioxide; this turns acidified potassium dichromate(VI) solution (or paper) green: SO₃^2-(aq) + 2H⁺(aq) " H₂O(l) + SO₂(g) Barium chloride solution gives a white precipitate of barium sulphite; addition of dilute hydrochloric acid causes the precipitate to dissolve without effervescence: SO₃^2-(aq) + Ba²⁺(aq) " BaSO₃(s)
Sulphate SO₄^2-
BaSO₄, SrSO₄ and PbSO₄ are insoluble; CaSO₄ is sparingly soluble. Barium chloride solution added to the test solution acidified with dilute hydrochloric acid gives a white precipitate of barium sulphate: Ba²⁺(aq) + SO₄^2-(aq) " BaSO₄(s) HSO₄^- does the same thing with barium ions; however the original test solution would then be very acidic, so that should be tested for. The addition of HCl destroys any carbonate or sulphite ions present so prevents the spurious positive result due to the precipitation of these barium salts. Barium nitrate solution can be used instead of barium chloride. 18.2 Lead ethanoate solution gives a precipitate of white lead sulphate: Pb²⁺(aq) + SO₄^2-(aq) " PbSO₄(s)
Nitrate NO₃^-
Since all nitrates are water soluble, there is no precipitation reaction for this ion. Solid nitrates decompose on heating; those of group 1 (except Li) give the nitrite and oxygen; 2NaNO₃(s) " 2NaNO₂(s) + O₂(g) All others give the metal oxide, nitrogen dioxide, and oxygen. A brown gas is emitted that re-lights a glowing splint: 2Pb(NO₃)₂(s) " 2PbO(s) + O₂(g) + 2NO₂(g) Nitrate ions are reduced to ammonia by boiling with aluminium or with Devarda?s Alloy in sodium hydroxide solution. Devarda?s Alloy contains aluminium, zinc and copper. Since ammonium ions also give ammonia with NaOH, the test solution must be boiled with NaOH and the vapour tested for ammonia; if present heating must continue until all the ammonia has gone. The mixture is then cooled, Devarda?s Alloy (or a piece of aluminium foil) added, and the mixture re-heated. A gas that turns moist red litmus paper blue indicates nitrate in the original solution: 3NO₃^-(aq) + 8Al(s) + 18H₂O(l) + 21 OH^-(aq) " 8[Al(OH)₆]^3-(aq) + 3NH₃(g) Not an equation to be remembered!
Carbonate CO₃^2-
Only the alkali metal and ammonium carbonates are water soluble. Some carbonates (e.g. zinc, copper(II)) are basic carbonates and contain a proportion of the hydroxide in their structure. Heating decomposes all but the alkali and alkaline earth metal carbonates (at Bunsen temperatures) giving the oxide and carbon dioxide: CuCO₃(s) " CuO(s) + CO₂(g) Dilute hydrochloric acid gives vigorous effervescence with carbonates, evolving carbon dioxide: CO₃^2-(aq or s) + 2H⁺(aq) " H₂O(l) + CO₂(g) Bicarbonates also give this effervescence. The reaction of carbonates with acid is exothermic; bicarbonates react endothermically.
Hydrogen carbonate (bicarbonate) HCO₃^-
Only the alkali metal and ammonium bicarbonates are obtainable as solids; group 2 bicarbonates exist only in solution. Calcium chloride solution on addition to a bicarbonate solution gives no precipitate since calcium bicarbonate is soluble; this distinguishes it from carbonate, which does give a precipitate. On heating the calcium chloride/bicarbonate mixture a white precipitate appears since the bicarbonate decomposes to carbonate: Ca²⁺(aq) + 2HCO₃^-(aq) " CaCO₃(s) + CO₂(g) + H₂O(l)
Ethanoate CH₃COO^-
Ethanoates on heating with dilute hydrochloric acid give ethanoic acid, recognisable by its vinegary smell. Neutral iron(III) chloride solution added to neutral solutions of ethanoate ion give a deep red colouration owing to formation of iron(III) ethanoate.
Ethanedioate C₂O₄^2-
Concentrated sulphuric acid added to a solid ethanedioate salt gives a mixture of carbon monoxide and carbon dioxide from dehydration of the ethanedioic acid formed: HOOC-COOH " H₂O + CO₂ + CO Potassium manganate(VII) solution acidified with dilute sulphuric acid added to a solution of an ethanedioate causes the purple colour to disappear: 2MnO₄^-(aq) + 5C₂O₄^2-(aq) + 16H⁺(aq) " 2Mn²⁺(aq) + 10CO₂(g) + 8H₂O(l) Calcium chloride solution added to a solution of an ethanedioate gives a white precipitate of calcium ethanedioate: Ca²⁺(aq) + C₂O₄^2-(aq) " CaC₂O₄(s) The precipitate dissolves readily in dilute hydrochloric acid.
Chromate(VI) CrO₄^2- and dichromate(VI) Cr₂O₇^2-
These ions are related through the equilibrium Cr₂O₇^2-(aq) + 2OH^-(aq) D 2CrO₄^2-(aq) + H₂O(l) In alkaline solution the yellow chromate(VI) dominates, in acidic solution orange dichromate(VI). All dichromate(VI) salts are soluble; addition of dichromate(VI) ions to solutions of ions of metals which have insoluble chromate(VI) salts leads to the precipitation of chromates. This means that the only dichromates that can exist are those of group 1 metals, ammonium, magnesium, calcium and strontium. Barium chloride solution added to a chromate(VI) or dichromate(VI) solution precipitates bright yellow barium chromate(VI): Ba²⁺(aq) + CrO₄^2-(aq) " BaCrO₄(s) The addition of a heavy metal ion to potassium dichromate solution precipitates the chromate and therefore moves the equilibrium to the right hand side. If the chromate is sparingly soluble (e.g. strontium) the supernatant liquid will remain yellow. Very insoluble chromates, such as lead, remove all the colour from the supernatant liquid. Dichromate(VI) ion solution in sulphuric acid is an oxidising agent; oxidation is shown by the solution turning from orange to green (Cr(III)). The following can be oxidised: (a) iron(II) to iron(III): Cr₂O₇^2-(aq) + 14H⁺(aq) + 6Fe²⁺(aq) " 2Cr³⁺(aq) + 7H₂O(l) + 6Fe³⁺(aq) (b) iodide to iodine (the solution turns murky greenish-brown): Cr₂O₇^2-(aq) + 14H⁺(aq) + 6I^-(aq) " 2Cr³⁺(aq) + 7H₂O(aq) + 3I₂(aq) (c) sulphite to sulphate: Cr₂O₇^2-(aq) + 8H⁺(aq) + 3SO₃^2-(aq) " 2Cr³⁺(aq) + 4H₂O(l) + 3SO₄^2-(aq) (d) nitrite to nitrate: Cr₂O₇^2-(aq) + 8H⁺(aq) + 3NO₂^-(aq) " 2Cr³⁺(aq) + 4H₂O(l) + 3NO₃^-(aq) (e) hydrogen peroxide reacts with acidified dichromate(VI) solutions to give a blue compound that rapidly turns green and evolves oxygen. The blue compound can be extracted into an organic solvent such as butan-1-ol. The blue compound is CrO₅, which contains a peroxy structure. It is covalent, and is stable in organic solvents though not in water. Cr₂O₇^2-(aq) + 8H⁺(aq) + 3H₂O₂ (aq) " 2Cr³⁺(aq) + 7H₂O(l) + 3 O₂(g) Alcohols are oxidised by acidified potassium dichromate(VI) solution. Primary alcohols give aldehydes and then acids, secondary alcohols give ketones. Ethanol can be used to test for dichromate(VI) ions, therefore, the solution turning green and the apple smell of ethanal being evident.

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