36

The Group Va Elements (N, P, As, Sb, Bi) andTheir Principal Anions and Cations

Nitrogen (N2) is a colourless, inert diatomic gas (boiling point: -196 °C). Nitrogenoccurs in Nature mainly as dinitrogen (N2) that comprises 78% by volume of theearth's atmosphere.

Phosphorus is solid at room temperature. There are three main forms of phosphorus:white, black, and red.White phosphorus is in the solid and liquid forms, and in the vapour phase below800 °C consists of tetrahedral P4 molecules. White phosphorus is the least stable solidallotrope, but all others revert to it when melted. It is highly reactive and toxic, and iscommonly stored under water to protect it from air.Orthorhombic black phosphorus, the most thermodynamically stable and leastreactive form, is obtained by heating white phosphorus under pressure. It has agraphitic appearance and consists of polymeric double layers. When subjected topressures above 12 kbar the orthorhombic form transforms successively to therhombohedral and cubic forms.Red phosphorus is of intermediate reactivity and is used commercially. Ordinarily itis amorphous. It is easily obtained by heating white phosphorus in a sealed vessel at ∼400 °C.

Arsenic, Antimony, and Bismuth. These elements have fewer allotropic forms thanphosphorus. For As and Sb unstable yellow allotropes comparable to whitephosphorus are obtainable by rapid condensation of vapours. They readily transformto the bright, "metallic" rhombohedral forms similar to rhombohedral blackphosphorus. This is also the commonest form for bismuth.Arsenic is a steel-grey, brittle solid with a metallic lustre. It sublimes on heating, anda characteristic garlic-like odour is apparent.Antimony is a lustrous, silver-white metal, which melts at 630 °C.Bismuth is a brittle, crystalline, reddish-white metal. It melts at 272 °C.

Solubility in water, aqueous acids and aqueous alkaliNitrogen is little soluble in water, but P, As, Sb, and Bi are not soluble.Phosphorus, arsenic, antimony, and bismuth are not affected by nonoxidizing

acids (e.g. HCl), but they are soluble in oxidising acids (e.g. in nitric acid, to produceH3PO4, H3AsO4, Sb2O3, and Bi(NO3)3).

Arsenic is insoluble in hydrochloric acid and in dilute sulphuric acid, but itdissolves readily in dilute nitric acid yielding arsenite ions and in concentrated nitricacid, aqua regia or sodium hypochlorite solution forming arsenate:

As + 4 H+ + NO3- → As3+ + NO ↑ + 2 H2O3 As + 5 HNO3 (conc) + 2 H2O → 3 AsO43- + 5 NO ↑ + 9 H+2 As + 5 OCl- + 3 H2O → 2 AsO43- + 5 Cl- + 6 H+

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Antimony is not soluble in dilute sulphuric acid, but it dissolves slowly in hot ,concentrated sulphuric acid forming antimony(III) ions:

2 Sb + 3 H2SO4 + 6 H+ → 2 Sb3+ + 3 SO2 ↑ + 6 H2ONitric acid oxidises antimony to an insoluble product, which can be regarded as amixture of Sb2O3 and Sb2O5. These anhydrides, in turn, can be dissolved in tartaricacid. A mixture of nitric acid and tartaric acid dissolves antimony easily.

Aqua regia dissolves antimony, when antimony(III) ions are formed:Sb + HNO3 + 3 HCl → Sb3+ + 3 Cl- + NO ↑ + 2 H2O

Bismuth dissolves in oxidising acids such as concentrated nitric acid, aquaregia, or hot, concentrated sulphuric acid:

2 Bi + 8 HNO3 → 2 Bi3+ + 6 NO3- + 2 NO ↑ + 4 H2OBi + HNO3 + 3 HCl → Bi3+ + 3 Cl- + NO ↑ + 2 H2O2 Bi + 6 H2SO4 → 2 Bi3+ + 3 SO42- + 3 SO2 ↑ + 6 H2O

Only white phosphorus is soluble in aq. alkali with disproportionation:

P4 + 3 NaOH + 3 H2O → PH3 ↑ + 3 NaH2PO2

Trihydrides of the Group Va elements (XH3) The gases XH3 can be obtained by treating ammonium salts with alkalies, by

treating phosphides or arsenides of electropositive metals with acids, by reduction ofsulphuric acid solutions of arsenic, antimony, or bismuth with an electropositive metalor electrolytically. The stability falls rapidly down in the group, so the SbH3 andBiH3 are very unstable thermally, the latter having been obtained only in traces.PH3, AsH3, and SbH3 are extremely poisonous.Ammonia NH3

Ammonia is a colourless pungent gas with a normal boiling point of -33.4 °C.The liquid has a large heat of evaporation and is therefore fairly easily handled inordinary laboratory equipment. Liquid ammonia resembles water in its physicalbehaviour, being highly associated because of the polarity of the molecules and stronghydrogen bonding.Liquid ammonia has lower reactivity than H2O toward electropositive metals, whichmay dissolve physically giving blue solutions (e.g. Na).

Nitric acid (HNO3)The pure acid is a colourless liquid. The normal concentrated aqueous acid (∼

70% by weight) is colourless but often becomes yellow as a result of photochemicaldecomposition, which gives NO2: 4 HNO3 → 4 NO2 + 2 H2O + O2The acid has the highest self-ionisation of the pure liquid acids, and the overall selfdissociation is

2 HNO3 ↔ NO2+ + NO3- + H2O

Nitric acid of concentration below 2 M has little oxidising power.

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The concentrated acid is a powerful oxidising agent and, of the metals, only Au, Pt, Ir,and Re are unattacked, although a few others such as Al, Fe, Cr are rendered passive,probably owing to formation of an oxide film. The attack on metals generallyinvolves reduction of nitrate. Nonmetals are usually oxidised by HNO3 to oxo acidsor oxides.The ability of HNO3, especially in the presence of concentrated H2SO4, to nitratemany organic compounds, is attributable to the formation of the nitronium ion, NO2+.The so-called fuming nitric acid contains dissolved NO2 in excess of the amount thatcan be hydrated to HNO3 + NO. Red fuming nitric acid contains N2O4.Aqua regia (∼3 vol. of conc. HCl + 1 vol. of conc. HNO3) contains free chlorine andClNO, and it attacks gold and platinum metals, its action being more effective thanthat of HNO3 mainly because of the complexing function of chloride ion.

Most important oxides of the Group Va elementsN P As Sb Bi

NO colourlessNO2 red

P4O10 (P2O5)white

very hygroscopic

As4O6 (As2O3)white

Sb4O6 (Sb2O3)white

Bi2O3yellow

Principal anions and cations of N, P, As, Sb, and Bi:N P As Sb Bi

NH4+ ammoniumNO2+ nitronium

NO2- nitriteNO3- nitrate

H2PO2-

hypophosphiteHPO32- phosphitePO43- phosphate

(As3+)*

(As5+)*

AsO33- arseniteAsO43- arsenate

Sb3+

Sb5+

SbO+

[Sb(OH)4]−

[Sb(OH)6]−

Bi3+

BiO+

* They exist in aqueous solution in the form of arsenite and arsenate.

Characteristic reactions of NH4+, NO2-, NO3-, PO43-, AsO33-, AsO43-,Sb3+, and Bi3+ ions

Ammonium ion, NH4+

Ammonium salts are generally water-soluble compounds, forming colourlesssolutions (unless the anion is coloured).

The reactions of ammonium ions are in general similar to those of potassium,because the sizes of the two ions are almost identical.

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Use a 0.5 M solution of ammonium chloride to study the reactions ofammonium ions.

1. Sodium hydroxide solution: ammonia gas is evolved on warming.

NH4+ + OH- → NH3 ↑ + H2O

Ammonia gas may be identifieda.) by its odour (cautiously smell the vapour after removing the test-tube from

the flame);b.) by the formation of white fumes of ammonium chloride when a glass rod

moistened with concentrated hydrochloric acid is held in the vapour;c.) by its turning moistened pH paper blue;d.) by its ability to turn filter paper moistened with mercury(I) nitrate solution

black (this is a very trustworthy test; arsine, however, must be absent):

2 NH3 + Hg22+ + NO3- → Hg(NH2)NO3 ↓ + Hg ↓ + NH4+

e.) filter paper moistened with a solution of manganese(II) chloride andhydrogen peroxide gives a brown colour, due to the oxidation of manganese by thealkaline solution thus formed:

2 NH3 + Mn2+ + H2O2 + H2O → MnO(OH)2 ↓ + 2 NH4+

2. Nessler's reagent (alkaline solution of potassium tetraiodomercurate(II)):brown precipitate or brown or yellow coloration is produced according to the

amount of ammonia or ammonium ions present. The precipitate is a basic mercury(II)amido-iodide:

NH4+ + 2 [HgI4]2- + 4 OH- → HgO.Hg(NH2)I ↓ + 7 I- + 3 H2O

The test is an extremely delicate one and will detect traces of ammonia present indrinking water. All metals except sodium or potassium, must be absent.

3. Sodium hexanitritocobaltate(III), Na3[Co(NO2)6]:yellow precipitate of ammonium hexanitritocobaltate(III), (NH4)3[Co(NO2)6]

, similar to that produced by potassium ions:

3 NH4+ + [Co(NO2)6]3- → (NH4)3[Co(NO2)6] ↓

4. Saturated sodium hydrogen tartrate solution, NaHC4H4O6:white precipitate of ammonium acid tartarate NH4HC4H4O6, similar to but

slightly more soluble than the corresponding potassium salt, from which it isdistinguished by the evolution of ammonia gas on being heated with sodiumhydroxide solution.

NH4+ + HC4H4O6- → NH4HC4H4O6 ↓

5. Perchloric acid or sodium perchlorate solution: no precipitate (distinction frompotassium).

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Compare the characteristic reactions of K+ and NH4+ ions:

Flame test cc HClO4 Tartaric acid [Co(NO2)6]3− Nessler`sreagent

K+

NH4+

Nitrites, NO2−

Silver nitrite is sparingly soluble in water (1.363 g AgNO2/100 ml water at 60°C). All other nitrites are soluble in water.

Use a 0.1M solution of potassium nitrite to study the reactions of nitrites.

1. Hydrochloric acid: Cautious addition of the acid to a solid nitrite in the cold yieldsa transient, pale-blue liquid (due to the presence of free nitrous acid, HNO2, or itsanhydride, N2O3) and the evolution of brown fumes of nitrogen dioxide, the latterbeing largely produced by combination of nitric oxide with the oxygen of the air.Similar results are obtained with the aqueous solution.

NO2- + H+ → HNO23 HNO2 → HNO3 + 2 NO ↑ + H2O2 NO ↑ + O2 ↑ → 2 NO2 ↑

2. Barium chloride solution: no precipitate.

3. Silver nitrate solution: white crystalline precipitate of silver nitrite only fromconcentrated solutions.

NO2- + Ag+ → AgNO2 ↓

4. Potassium iodide solution: the addition of a nitrite solution to a solution ofpotassium iodide, followed by acidification with acetic acid or with dilute sulphuricacid, results in the liberation of iodine, which may be identified by the blue colourproduced with starch solution. An alternative method is to extract the liberated iodinewith carbon tetrachloride.

2 NO2- + 2 I- + 4 H+ → I2 + 2 NO ↑ + 2 H2O

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5. Ammonium chloride. By boiling a solution of a nitrite with excess of solidammonium chloride, nitrogen is evolved and the nitrite is completely destroyed:

NO2- + NH4+ → N2 ↑ + 2 H2O

6. Urea: the nitrite is decomposed, and nitrogen and carbon dioxide are evolved,when a solution of a nitrite is treated with urea, CO(NH2)2, and the mixture isacidified with dilute hydrochloric acid.

2 NO2- + CO(NH2)2 + 2 H+ → 2 N2 ↑ + CO2 ↑ + 3 H2O

7. Sulphamic acid (H2N-SO3H). When a solution of a nitrite is treated withsulphamic acid, it is completely decomposed:

H2NSO3H + NO2- + H+ → N2 ↑ + 2 H+ + SO42- + H2O

8. Acidified potassium permanganate solution: decolourized by a solution of anitrite, but no gas is evolved.

5 NO2- + 2 MnO4- + 6 H+ → 5 NO3- + 2 Mn2+ + 3 H2O

9. Sulphanilic acid α-naphthylamine reagent. (Griess-Ilosvay test)This test depends upon the diazotization of sulphanilic acid by nitrous acid,

followed by coupling with α-naphthylamine to form a red azo dye:

NH 2

SO H3 NH2

NO2 H++ ++ 2 H O2+NH 2N NHSO3

The test solution must be very dilute, otherwise the reaction does not go beyond thediazotation stage.

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Nitrates, NO3−

All nitrates are soluble in water.The nitrates of mercury and bismuth yield basic salts on treatment with water;

these are soluble in dilute nitric acid.Use a 0.1M solution of potassium nitrate to study the reactions of nitrates.

1. Reduction to nitrite test. Nitrates are reduced to nitrites by metallic zinc in aceticacid solutions; the nitrite can be readily detected by means of the Griess-Ilosvay test.Nitrites, of course, interfere and are best removed with sulphamic acid.

2. Reduction of nitrates in alkaline medium. Ammonia is evolved when a solutionof a nitrate is boiled with zinc dust or gently warmed with aluminium powder andsodium hydroxide. solution. Ammonia is detected (i) by its odour, (ii) by its actionupon pH paper and (iii) upon mercury(I) nitrate paper.

NO3- + 4 Zn + 7 OH- + 6 H2O → NH3 ↑ + 4 [Zn(OH)4]2-3 NO3- + 8 Al + 5 OH- + 18 H2O → NH3 ↑ + 8 [Al(OH)4]-

Ammonium ions interfere and must be absent. Nitrites give similar reaction and maybe removed with sulphamic acid.

3. Iron(II) sulphate solution and concentrated sulphuric acid (brown ring test):This test is carried out in either of two ways:

a.) Add 3 ml freshly prepared saturatedsolution of iron(II) sulphate to 2 mlnitrate solution, and pour 3-5 ml concent-rated sulphuric acid slowly down the sideof the test tube so that the acid forms alayer beneath the mixture. A brown ringforms where the liquids meet.b.) Add 4 ml concentrated sulphuricacid slowly to 2 ml nitrate solution, mixthe liquids thoroughly and cool themixture under a stream of cold waterfrom the tap, or ice-water. Pour asaturated solution of iron(II) sulphateslowly down side of the tube so that itforms a layer on top of the liquid. Set thetube aside for 2-3 minutes. A brown ringwill form at the zone of contact of thetwo liquids.

The brown ring is due to the formation of the [Fe(NO)]2+. On shaking and warmingthe mixture the brown colour disappears, nitric oxide is evolved, and a yellowsolution of iron(III) ions remains.

2 NO3− + 4 H2SO4 + 6 Fe2+ → 6 Fe3+ + 2 NO ↑ + 4 SO4

2− + 4 H2OFe2+ + NO ↑ → [Fe(NO)]2+

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Orthophosphates, PO43-

Orthophosphoric acid (often referred to simply as phosphoric acid) is atriprotic acid giving rise to three series of salts:

primary orthophosphates, e.g. NaH2PO4;secondary orthophosphates, e.g. Na2HPO4:tertiary orthophosphates, e.g. Na3PO4.

Ordinary 'sodium phosphate' is disodium hydrogen phosphate, Na2HPO4.12H2O .

Solubility. The phosphates (primary, secondary, and tertiary) of the alkalimetals, with the exception of lithium and of ammonium, are soluble in water.The primary phosphates of the alkaline earth metals are also soluble in water.All the phosphates of the other metals, and also the secondary and tertiary phosphatesof the alkaline earth metals are sparingly soluble or insoluble in water.

To study the reactions of phosphates use a 0.1 M solution of disodiumhydrogen phosphate, Na2HPO4.12H2O .

1. Dilute hydrochloric acid. No apparent change.

2. Silver nitrate solution: yellow precipitate of normal silver orthophosphate,solubility product: Ksp(Ag3PO4, 25°C)= 8.88x10−17:

HPO42- + 3 Ag+ → Ag3PO4 ↓ + H+

The precipitate is soluble in dilute ammonia solution and in dilute nitric acid:

Ag3PO4 ↓ + 6 NH3 → 3 [Ag(NH3)2]+ + PO43−

Ag3PO4 ↓ + 2 H+ → H2PO4− + 3 Ag+

3. Barium chloride solution: white, amorphous precipitate of secondary bariumphosphate from neutral solutions, soluble in dilute mineral acids and in acetic acid.

HPO42- + Ba2+ → BaHPO4 ↓

In the presence of dilute ammonia solution, the less soluble tertiary phosphate isprecipitated:

2 HPO42- + 3 Ba2+ + 2 NH3 → Ba3(PO4)2 ↓ + 2 NH4+

4. Magnesium nitrate reagent or magnesia mixture: The former is a solutioncontaining Mg(NO3)2, NH4NO3, and a little aqueous NH3, and the latter is asolution containing MgCl2, NH4Cl, and a little aqueous NH3. With either reagent awhite, crystalline precipitate of magnesium ammonium phosphate,Mg(NH4)PO4.6H2O, is produced:

HPO42- + Mg2+ + NH3 → Mg(NH4)PO4 ↓

The precipitate is soluble in acetic acid and in mineral acids, but practically insolublein ammonia solution.

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5. Iron(III) chloride solution: yellowish-white precipitate of iron(III) phosphate:

HPO42− + Fe3+ → FePO4 ↓ + H+

The precipitate is soluble in dilute mineral acids, but insoluble in dilute acetic acid.

6. Ammonium molybdate reagent: The addition of a large excess of this reagent to asmall volume of a phosphate solution produces a yellow, crystalline precipitate ofammonium phosphomolybdate, (NH4)3[P(Mo3O10)4]. The resulting solution shouldbe strongly acid with nitric acid; the latter is usually present in the reagent andaddition is therefore unnecessary.(Prerare a clear solution of the reagent from (NH4)2MoO4 and concentrated HNO3 !)

HPO42- + 12 MoO42- + 3 NH4+ + 23 H+ → (NH4)3[P(Mo3O10)4] + 12 H2O

Reactions of arsenic(III) ions Arsenic(III) compounds can be derived from the amphoteric arsenic trioxide

As2O3. In strongly acid solutions the only detectable species is the pyramidalAs(OH)3. In strongly basic solutions the arsenite ion, AsO33-, appears to be present.

A 0.1 M solution of arsenic(III) oxide, As2O3, or sodium arsenite, Na3AsO3,can be used for studying the reactions of arsenic(III) ions.Arsenic(III) oxide does not dissolve in cold water, but by boiling the mixture for halfan hour, dissolution is complete. The mixture can be cooled without the danger ofprecipitating the oxide.

1. Hydrogen sulphide: yellow precipitate of arsenic(III) sulphide:

2 AsO33− + 6 H+ + 3 H2S → As2S3 ↓ + 6 H2O

The solution must be strongly acidic; if there is not enough acid present a yellowcoloration is visible only, owing to the formation of colloidal As2S3.The precipitate is insoluble in concentrated hydrochloric acid.The precipitate dissolves in hot concentrated nitric acid, alkali hydroxides, orammonia:

3 As2S3 ↓ + 28 HNO3 + 4 H2O → 6 AsO43− + 9 SO4

2− + 36 H+ + 28 NO ↑As2S3 ↓ + 6 OH− → AsO3

3− + AsS33− + 3 H2O

Ammonium sulphide and ammonium polysulphide also dissolves the precipitate,when thioarsenite (AsS3

3−) and thioarsenate (AsS43−) ions are formed, respectively:

As2S3 ↓ + 3 S2− → 2 AsS33−

As2S3 ↓ + 4 S22− → 2 AsS4

3− + S32−

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On reacidifying these both decompose, when arsenic(III) sulphide or arsenic(V)sulphide, and hydrogen sulphide are formed. The excess polysulphide reagent alsodecomposes and the precipitate is contaminated with sulphur:

2 AsS33− + 6 H+ → As2S3 ↓ + 3 H2S ↑

2 AsS43− + 6 H+ → As2S5 ↓ + 3 H2S ↑S2

2− + 2 H+ → H2S ↑ + S ↓

2. Silver nitrate: yellow precipitate of silver arsenite in neutral solution (distinctionfrom arsenates):

AsO33− + 3 Ag+ → Ag3AsO3 ↓

3. Magnesia mixture (a solution containing MgCl2, NH4Cl, and a little NH3): no precipitate (distinction from arsenate).

4. Copper sulphate solution: green precipitate of copper arsenite, variouslyformulated as CuHAsO3 and Cu3(AsO3)2.xH2O, from neutral solutions:

AsO33− + Cu2+ + H+ → CuHAsO3 ↓

The precipitate soluble in acids, and also in ammonia solution. The precipitate alsodissolves in sodium hydroxide solution; upon boiling, copper(I) oxide is precipitated.

5. Potassium tri-iodide (solution of iodine in potassium iodide): oxidizes arseniteions while becoming decolourized:

AsO33− + I3

− + H2O ↔ AsO43− + 3 I− + 2 H+

The reaction is reversible, and an equilibrium is reached.

6. Bettendorff's test (tin(II) chloride solution and concentrated hydrochloric acid): a few drops of the arsenite solution are added to a solution made of 0.5 ml saturatedtin(II) chloride solution and 2 ml concentrated hydrochloric acid, and the solution isgently warmed; the solution becomes dark brown and finally black, due to theseparation of elementary arsenic:

2 AsO33− + 12 H+ + 3 Sn2+ → 2 As ↓ + 3 Sn4+ + 6 H2O

7. Marsh's test.This test is based upon the fact that all soluble compounds of arsenic are reduced by'nascent' hydrogen in acid solution to arsine (AsH3), a colourless, extremelypoisonous gas with a garlic-like odour.If the gas, mixed with hydrogen, is conducted through a heated glass tube, it isdecomposed into hydrogen and metal arsenic, which is deposited as a brownish-black'mirror' just beyond the heated part of the tube, particularly if it is cooled.The deposit is soluble in sodium hypochlorite (distinction from antimony).

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AsO33− + 3 Zn + 9 H+ →

AsH3 ↑ + 3 Zn2+ + 3 H2O

4 AsH3 ↑ → heat → 4 As ↓ + 6 H2 ↑

2 As + 5 OCl− + 3 H2O →2 AsO4

3− + 5 Cl− + 6 H+

Reactions of arsenate ions, AsO43-Arsenic(V) compounds are derived from arsenic pentoxide, As2O5. This is the

anhydride of arsenic acid, H3AsO4, which forms salts such as sodium arsenate.Arsenic(V) therefore exists in solutions predominantly as the arsenate AsO43- ion.

A 0.1 M solution of disodium hydrogen arsenate Na2HAsO4 can be used forthe study of these reactions.

1. Hydrogen sulphide: no immediate precipitate in the presence of dilutehydrochloric acid. If the passage of the gas is continued, a mixture of arsenic(III)sulphide and sulphur is slowly precipitated. Precipitation is more rapid in hot solution.

AsO43− + H2S → AsO3

3− + S ↓ + H2O2 AsO3

3− + 6 H+ + 3 H2S → As2S3 ↓ + 6 H2O

If a large excess of concentrated hydrochloric acid is present and hydrogen sulphide ispassed rapidly into the cold solution, yellow arsenic pentasulphide is precipitated:

2 AsO43− + 5 H2S + 6 H+ → As2S5 ↓ + 8 H2O

Arsenic pentasulphide, like the trisulphide, is readily soluble in alkali hydroxides orammonia, ammonium sulphide, ammonium polisulphide, sodium or ammoniumcarbonate:

As2S5 ↓ + 6 OH− → AsS43− + AsO3S3− + 3 H2O

As2S5 ↓ + 3 S2− → 2 AsS43−

As2S5 ↓ + 6 S22− → 2 AsS4

3− + 3 S32−

As2S5 ↓ + 3 CO32− → AsS4

3− + AsO3S3− + 3 CO2

Upon acidifying these solutions with hydrochloric acid, arsenic pentasulphide isreprecipitated:

2 AsS43− + 6 H+ → As2S5 ↓ + 3 H2S ↑

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2. Silver nitrate solution: brownish-red precipitate of silver arsenate, Ksp(Ag3AsO4,25°C)= 1.03x10−22, from neutral solutions:

AsO43− + 3 Ag+ → Ag3AsO4 ↓

Soluble in acids and ammonia solution, but insoluble in acetic acid.

3. Magnesia mixture: white, crystalline precipitate of magnesium ammoniumarsenate Mg(NH4)AsO4.6H2O from neutral or ammoniacal solution:

AsO43− + Mg2+ + NH4

+ → MgNH4AsO4 ↓

Upon treating the white precipitate with silver nitrate solution containing a few dropsof acetic acid, red silver arsenate is formed:

MgNH4AsO4 ↓ + 3 Ag+ → Ag3AsO4 ↓ + Mg2+ + NH4+

4. Ammonium molybdate solution: when the reagent and nitric acid are added inconsiderable excess to a solution of an arsenate, a yellow crystalline precipitate isobtained on boiling:

AsO43− + 12 MoO4

2− + 3 NH4+ + 24 H+ → (NH4)3[As(Mo3O10)4] ↓ + 12 H2O

The precipitate is insoluble in nitric acid, but dissolves in ammonia solution and insolutions of caustic alkalis.

5. Potassium iodide solution: in the presence of concentrated hydrochloric acid,iodine is precipitated; upon shaking the mixture with 1-2 ml of carbon tetrachloride orof chloroform, the latter is coloured violet by the iodine.

AsO43− + 2 I− + 2 H+ ↔ AsO3

3− + I2 ↓ + H2O

The reaction is reversible.

Redox systems: I−/I2 H3AsO3/H3AsO4

Concentrations: [I−]= 0.1 M; [H3AsO4]= 0.1 Ma) [H3AsO3]= 0.001 Mb) [H3AsO3]= 0.01 Mc) [H3AsO3]= 0.1 M 0 1 2 3 4 5

0.2

0.3

0.4

0.5

0.6

0.7

cba

H3AsO3/ H3AsO4

I-/ I2

Redox potential (V)

pH

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6. Bettendorff's test (tin(II) chloride solution and concentrated hydrochloric acid): the solution becomes dark brown and finally black, due to the separation ofelementary arsenic:

2 AsO43− + 16 H+ + 5 Sn2+ → 2 As ↓ + 5 Sn4+ + 8 H2O

7. Marsh's test: The soluble compounds of arsenic are reduced by 'nascent' hydrogenin acid solution to arsine (AsH3). If the gas, mixed with hydrogen, is conductedthrough a heated glass tube, it is decomposed into hydrogen and metal arsenic, whichis deposited as a brownish-black 'mirror' just beyond the heated part of the tube.

AsO43− + 4 Zn + 11 H+ → AsH3 ↑ + 4 Zn2+ + 4 H2O

4 AsH3 ↑ → heat → 4 As ↓ + 6 H2 ↑2 As + 5 OCl− + 3 H2O → 2 AsO4

3− + 5 Cl− + 6 H+

Reactions of antimony(III) ions, Sb3+Antimony(III) compounds are easily dissolved in acids, when the ion Sb3+ is

stable. If the solution is made alkaline, or the concentration of hydrogen ions isdecreased by dilution, hydrolysis occurs when antimonyl, SbO+, ions are formed:

Sb3+ + H2O ↔ SbO+ + 2 H+

Use 0.1 M solution of antimony(III) chloride, SbCl3, to study the reactions ofantimony(III) ions.

1. Hydrogen sulphide: orange-red precipitate of antimony trisulphide, Sb2S3, fromsolutions which are not too acidic.

2 Sb3+ + 3 H2S → Sb2S3 ↓ + 6 H+

The precipitate is soluble in warm concentrated hydrochloric acid (distinction andmethod of separation from arsenic(III) and mercury(II) sulphide), in ammoniumsulphide (forming a thioantimonite) and polysulphide (forming a thioantimonate), andin alkali hydroxide solutions (forming antimonite and thioantimonite).

Sb2S3 ↓ + 6 HCl → 2 Sb3+ + 6 Cl− + 3 H2S ↑Sb2S3 ↓ + 3 S2− → 2 SbS3

3−

Sb2S3 ↓ + 4 S22− → 2 SbS4

3− + S32−

Sb2S3 ↓ + 4 OH− → [Sb(OH)4]− + SbS33−

Upon acidification of the thioantimonate solution with hydrochloric acid, antimonypentasulphide is precipitated initially but usually decomposes partially into thetrisulphide and sulphur:

2 SbS43− + 6 H+ → Sb2S5 ↓ + 3 H2S ↑ Sb2S5 ↓ → Sb2S3 ↓ + 2 S ↓

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Acidification of the thioantimonite solution or the antimonite-tioantimonite mixtureleads to the precipitation of the trisulphide:

2 SbS33− + 6 H+ → Sb2S3 ↓ + 3 H2S ↑

[Sb(OH)4]− + SbS33− + 4 H+ → Sb2S3 ↓ + 4 H2O

2. Water: when the solution is poured into water, a white precipitate of antimonylchloride SbOCl is formed, soluble in hydrochloric acid and in tartaric acid solution(difference from bismuth). With a large excess of water the hydrated oxideSb2O3.xH2O is produced.

Sb3+ + Cl− + H2O ↔ SbOCl ↓ + 2 H+

SbOCl ↓ + HOOC-CH(OH)-CH(OH)-COOH →→ [OOC-CH(OH)-CH(OH)-COOSbO]− + 2 H+ + Cl−

3. Sodium hydroxide or ammonia solution: white precipitate of the hydratedantimony(III) oxide Sb2O3.xH2O soluble in concentrated solutions of alkalis formingantimonites.

2 Sb3+ + 6 OH− → Sb2O3 ↓ + 3 H2OSb2O3 ↓ + 2 OH− + 3 H2O → 2 [Sb(OH)4]−

4. Zinc, tin, or iron: a black precipitate of antimony is produced.

Standard redox potentials:

Sb/Sb3+: +0.24 V

Zn/Zn2+: −0.76 V; Fe/Fe2+: −0.44 V; Sn/Sn2+: −0.14 V.

2 Sb3+ + 3 Zn ↓ → 2 Sb ↓ + 3 Zn2+

2 Sb3+ + 3 Sn ↓ → 2 Sb ↓ + 3 Sn2+

2 Sb3+ + 3 Fe ↓ → 2 Sb ↓ + 3 Fe2+

5. Potassium iodide solution: yellow coloration owing to the formation of a complexsalt:

Sb3+ + 6 I− → [SbI6]3−

6. Marsh's test: This test is based upon the fact that soluble compounds of antimonyare reduced by 'nascent' hydrogen in acid solution to stibine (SbH3), a colourless,thermally unstable, and extremely poisonous gas.If the gas, mixed with hydrogen, is conducted through a heated glass tube, it isdecomposed into hydrogen and metal antimony, which is deposited as a brownish-black 'mirror' on both sides of the heated part of the tube.The deposit is not soluble in sodium hypochlorite (distinction from arsenic).

50

Sb3+ + 3 Zn + 3 H+ → SbH3 ↑ + 3 Zn2+

4 SbH3 ↑ → heat → 4 Sb ↓ + 6 H2 ↑

Reactions of bismuth(III) ions, Bi3+The trivalent bismuth ion Bi3+ is the most common ion of bismuth.The hydroxide Bi(OH)3 is a weak base; bismuth salts therefore hydrolyse

readily, when the following process occurs:

Bi3+ + H2O ↔ BiO+ + 2 H+

The bismuthyl ion, BiO+, forms insoluble salts, like BiOCl, with most ions. If wewant to keep bismuth ions in solution, we must acidify the solution, when the aboveequilibrium shifts towards the left.

Use a 0.1 M solution of Bi(NO3)3, which contains 3-4 per cent nitric acid, tostudy these reactions.

1. Hydrogen sulphide: black precipitate of bismuth sulphide:

2 Bi3+ + 3 H2S → Bi2S3 ↓ + 6 H+

The precipitate is insoluble in cold, dilute acid and in ammonium sulphide.Boiling concentrated hydrochloric acid dissolves the precipitate, when hydrogensulphide gas is liberated.

Bi2S3 ↓ + 6 HCl → 2 Bi3+ + 6 Cl− + 3 H2S ↑

Hot dilute nitric acid dissolves bismuth sulphide, leaving behind sulphur in the formof a white precipitate.

Bi2S3 ↓ + 8 H+ + 2 NO3− → 2 Bi3+ + 3 S ↓ + 2 NO ↑ + 4 H2O

2. Ammonia solution: white basic salt of variable composition. The approximatechemical reaction is:

51

Bi3+ + NO3− + 2 NH3 + 2 H2O → Bi(OH)2NO3 ↓ + 2 NH4

+

The precipitate is insoluble in excess reagent.

3. Sodium hydroxide: white precipitate of bismuth(III) hydroxide:

Bi3+ + 3 OH− → Bi(OH)3 ↓

When boiled, the precipitate loses water and turns yellowish-white:

Bi(OH)3 ↓ → BiO.OH ↓ + H2O

4. Potassium iodide when added dropweise: black precipitate of bismuth(III) iodide.The precipitate dissolves readily in excess reagent, when orange-coloured tetraiodo-bismuthate ions are formed:

Bi3+ + 3 I− → BiI3 ↓BiI3 ↓ + I− ↔ [BiI4]−

Heating the BiI3 precipitate with water, it turns orange, owing to the formation ofbismuthyl iodide:

BiI3 ↓ + H2O → BiOI ↓ + 2 H+ + 2 I−

5. Sodium tetrahydroxostannate(II) (freshly prepared): in cold solution reducesbismuth(III) ions to bismuth metal which separates in the form of a black precipitate.First the sodium hydroxide present in the reagent reacts with bismuth(III) ions, andthen bismuth(III) hydroxide is reduced by tetrahydroxostannate(II) ions when bismuthmetal and hexahydroxostannate(IV) ions are formed:

Bi3+ + 3 OH− → Bi(OH)3 ↓2 Bi(OH)3 ↓ + 3 [Sn(OH)4]2− → 2 Bi ↓ + 3 [Sn(OH)6]2−

Standard redox potentials:in HCl solution: in NaOH solution: Bi/[BiCl4]

−: +0.16 V Bi/BiO+: +0.32 V Sn2+/Sn4+: +0.15 V [Sn(OH)4]2−/[Sn(OH)6]2−: −0.93 V

6. Water: when a solution of a bismuth salt is poured into a large volume of water. awhite precipitate of the corresponding basic salt is produced, which is soluble indilute mineral acids, but is insoluble in alkali hydroxides.

Bi3+ + NO3− + H2O → BiO(NO3) ↓ + 2 H+

Bi3+ + Cl− + H2O → BiO.Cl ↓ + 2 H+

52

Compare the characteristic reactions of arsenites, arsenates, and phosphates

PO43− AsO3

3− AsO43−

H2S

Magnesia mixture

+ AgNO3 soln.

Ag+

(NH4)2MoO4

Summarise the reactions of AsO33−, Sb3+, and Bi3+ ions with various reagents:

AsO33− Sb3+ Bi3+

+ H2Sprecipitate + NH4OH + 1:1 HCl + (NH4)2CO3

+ (NH4)2SMarsh's test + NaOClBettendorff's test

AgNO3

KI

Fe

Na2[Sn(OH)4]

53

The Group VIa Elements (O, S, Se, Te) andTheir Anions

Oxygen occurs in two allotropic forms: O2 and O3 (ozone).Oxygen (O2) is a colourless gas (boiling point: -183 °C). O2 is paramagnetic in thegaseous, liquid, and solid states. Both liquid and solid O2 are pale blue. Oxygenforms compounds with all the elements except He, Ne, and possibly Ar, and itcombines directly with all the other elements except the halogens, a few noble metals,and the noble gases, either at room or elevated temperatures.Oxygen is moderately soluble in water, and the neutral water saturated with O2 is afairly good oxidising agent.Ozone (O3) is a diamagnetic triatomic molecule. The gas is blue, the liquid is deepblue (boiling point: -112 °C), and the solid is black violet (melting point: -193 °C).

Sulphur has a strong tendency to catenation. It forms open and cyclic Sn speciesfrom n=2 to n=20 for cycles and higher for chains. This leads to enormous complexityin the physical and chemical behaviour of the element. There are several allotropicforms of sulphur.The common, stable form of sulphur at room temperature is the solid orthorhombic αsulphur, containing cyclo-S8 molecules.

Selenium has a smaller tendency to catenation than sulphur. It has three forms, the αand β red selenium, containing sulphur type cyclo-Se8 molecules, and the greyselenium.Grey selenium (metallic) is the stable form. The structure, which has no sulphuranalogue, contains infinite, spiral chains of selenium atoms.

Tellurium: the one form of tellurium is silvery-white, semimetallic, andisomorphous with grey selenium. Like the latter it is virtually insoluble in all liquidsexcept those with which it reacts.

Solubility of chalcogens in water, aqueous acids and aqueous alkaliOxygen is moderately soluble in water, but S, Se, and Te are not soluble.Sulphur, selenium, and tellurium are not affected by nonoxidizing acids, but

they are soluble in hot concentrated oxidising acids to produce H2SO4, H2SeO4, andH6TeO6, respectively, e.g.:

S + 2 HNO3 → H2SO4 + 2 NO

Sulphur is soluble in aq. alkali with disproportionation:

4 S + 6 NaOH → 2 Na2S + Na2S2O3 + 3 H2O

54

Dihydrides of chalcogensExcept water, the dihydrides, H2S, H2Se, and H2Te are extremely poisonous

gases with revolting odors; the toxicity of H2S far exceeds that of HCN. All behaveas very weak acids in aqueous solution. H2S dissolves in water to give a solution ca.0.1 M under 1 atm pressure. The dissociation constants are

H2S + H2O ↔ H3O+ + HS- K1= 1.3x10-7

HS- + H2O ↔ H3O+ + S2- K2= 7.1x10-15

Hydrogen peroxide H2O2Pure H2O2 is a colourless liquid (bp 150.2 °C, mp -0.43 °C) that resembles

water in many of its physical properties, although it is denser (1.44 g/cm3 at 25 °C).Dilute or 30% H2O2 solutions are widely used as oxidants. In acid solutionoxidations with H2O2 are most often slow, whereas in basic solution they are usuallyfast.H2O2 can also be oxidised in aqueous solution e.g. with MnO4-.Test the redox behaviour of H2O2 by performing the following tests:

1. Potassium iodide and starch. If potassium iodide and starch are added tohydrogen peroxide, acidified previously by dilute sulphuric acid, iodine is formedslowly and the solution turns gradually to deeper and deeper blue:

H2O2 + 2 H+ + 2 I- → I2 + 2 H2O

2. Potassium permanganate solution: decolourized in acid solution and oxygen isevolved:

2 MnO4- + 5 H2O2 + 6 H+ → 2 Mn2+ + 5 O2 ↑ + 8 H2O

Principal anions and oxoanions of chalcogens:O S Se Te

O2- oxideO22- peroxide

S2- sulphideSn2- polysulphide

SO32- sulphiteSO42- sulphate

S2O32- thiosulphateS2O42- dithioniteS2O52- disulphiteS2O62- dithionateS2O72- disulphate

Sn+2O62- polythionateSO52- peroxo-monosulphate

S2O82- peroxodisulphate

Se2- selenide

SeO32- seleniteSeO42- selenate

Te2- telluride

TeO32- telluriteTeO42- tellurate

55

Characteristic reactions of sulphide, sulphite, sulphate, and thiosulphate anions

Sulphides, S2-

Sulphides of alkali metals are soluble in water; their aqueous solutions reactalkaline because of hydrolysis:

S2- + H2O ↔ SH- + OH-SH- + H2O ↔ H2S + OH-

The sulphides of most other metals are insoluble in water. Those of alkalineearths are sparingly soluble, but are gradually changed by contact with water intosoluble hydrogen sulphides:

e.g.: CaS + H2O → Ca2+ + SH- + OH-

The sulphides of aluminium, chromium, and magnesium can only be preparedin the dry, as they are completely hydrolysed by water:

e.g.: Al2S3 + 6 H2O → 2 Al(OH)3 ↓ + 3 H2S ↑

For the study of the reactions of sulphides use a 2M solution of sodium sulphide.

1. Hydrochloric acid: hydrogen sulphide gas is evolved:

S2- + 2 H+ → H2S ↑

The gas may be identified by its characteristic odour, and by the blackening of filterpaper moistened with lead acetate solution:

H2S ↑ + Pb2+ → PbS ↓ + 2 H+

2. Silver nitrate solution: black precipitate of silver sulphide, solubility product:Ksp(Ag2S, 25°C)= 1.09x10−49:

S2- + 2 Ag+ → Ag2S ↓

3. Barium chloride solution: no precipitate.

4. Lead acetate solution: black precipitate of lead sulphide, solubility product:Ksp(PbS, 25°C)= 9.04x10−29:

S2- + Pb2+ → PbS ↓

56

5. Sodium nitroprusside solution (Na2[Fe(CN)5NO]): transient purple colour in thepresence of solutions of alkalis. No reactions occurs with solutions of hydrogensulphide or with the free gas: if, however, filter paper is moistened with a solution ofthe reagent made alkaline with sodium hydroxide or ammonia solution, a purplecolouration is produced with free hydrogen sulphide:

S2- + [Fe(CN)5NO]2- → [Fe(CN)5NOS]4-

Sulphites, SO32-

Only the sulphites of the alkali metals and of ammonium are soluble in water.Sulphites of the other metals are either sparingly soluble or insoluble in water.

The hydrogen sulphites of the alkali metals are soluble in water; the hydrogensulphites of the alkaline earth metals are known only in solution.

Use a 0.5M solution of sodium sulphite Na2SO3 to study the reactions of sulphites.

1. Hydrochloric acid (or dilute sulphuric acid): decomposition, more rapidly onwarming, with the evolution of sulphur dioxide:

SO32- + 2 H+ → SO2 ↑ + H2O

The gas may be identified (i) by its suffocating odour of burning sulphur, (ii) by theblue colouration, due to the formation of iodine, produced when a filter paper,moistened with potassium iodate and starch solution, is held over the mouth of thetest-tube.

5 SO2 ↑ + 2 IO3- + 4 H2O → I2 + 5 SO42- + 8 H+

2. Barium chloride or (strontium chloride) solution: white precipitate of barium ( orstrontium) sulphite:

SO32- + Ba2+ → BaSO3 ↓

the precipitate dissolves in dilute hydrochloric acid, when sulphur dioxide evolves.On standing, the precipitate is slowly oxidized to the sulphate and is then insoluble indilute mineral acids; this change is rapidly effected by warming with bromine wateror with hydrogen peroxide.

2 BaSO3 ↓ + O2 → 2 BaSO4 ↓BaSO3 ↓ + Br2 + H2O → BaSO4 ↓ + 2 Br- + 2 H+BaSO3 ↓ + H2O2 → BaSO4 ↓ + H2O

57

3. Silver nitrate solution: first, no visible change occurs because of the formation ofsulphitoargentate ions:

SO32- + Ag+ → [AgSO3]-

on the addition of more reagent, a white, crystalline precipitate of silver sulphite,solubility product: Ksp(Ag2SO3, 25°C)= 1.49x10−14, is formed:

[AgSO3]- + Ag+ → Ag2SO3 ↓

The precipitate dissolves if sulphite ions are added in excess:

Ag2SO3 ↓ + SO32- → 2 [AgSO3]-

On boiling the solution of the complex salt, or an aqueous suspension of theprecipitate, grey metallic silver is precipitated.The precipitate is soluble in dilute nitric acid, when sulphur dioxide gas is evolved.The precipitate also dissolves in ammonia.

4. Lead acetate or lead nitrate solution: white precipitate of lead sulphite:

SO32- + Pb2+ → PbSO3 ↓

The precipitate dissolves in dilute nitric acid. On boiling, the precipitate is oxidizedby atmospheric oxygen and white lead sulphate is formed:

2 PbSO3 ↓ + O2 → 2 PbSO4 ↓

This reaction can be used to distinguish sulphites and thiosulphates; the latter producea black precipitate on boiling.

5. Potassium dichromate solution (acidified with dilute sulphuric acid before the test): a green colouration, owing to the formation of chromium(III) ions:

3 SO32- + Cr2O72- + 8 H+ → 2 Cr3+ + 3 SO42- + 4 H2O

6. Lime water: This test is carried out by adding dilute hydrochloric acid to the solidsulphite, and bubbling the evolved sulphur dioxide through lime water; a whiteprecipitate of calcium sulphite is formed.

SO32- + Ca2+ → CaSO3 ↓

The precipitate dissolves on prolonged passage of the gas, due to the formation ofhydrogen sulphite ions:

CaSO3 ↓ + SO2 + H2O → Ca2+ + 2 HSO3-

A turbidity is also produced by carbonates; sulphur dioxide must therefore be firstremoved when testing for the latter. This may be effected by adding potassium

58

dichromate solution to the test-tube before acidifying. The dichromate oxidizes anddestroys the sulphur dioxide without affecting the carbon dioxide.

7. Zinc and sulphuric acid: hydrogen sulphide gas is evolved, which may bedetected by holding lead acetate paper to the mouth of the test-tube:

SO32- + 3 Zn + 8 H+ → H2S ↑ + 3 Zn2+ + 3 H2O

Compare the characteristic reactions of sulphites and carbonates

HCl Ba2+ Ag+ Pb2+K2Cr2O7

+acid

Zn+

H2SO4

CO32−

SO32−

Thiosulphates, S2O32-

Most of the thiosulphates that have been prepared are soluble in water;those of lead, silver, and barium are very sparingly soluble. Lead and silverthiosulphate dissolve in excess sodium thiosulphate solution forming complex salts.

To study the reactions of thiosulphates use a 0.5M solution of sodiumthiosulphate.

1. Hydrochloric acid: no immediate change in the cold with a solution of athiosulphate; the acidified liquid soon becomes turbid owing to the separation ofsulphur, and sulphur dioxide is evolved (especially by heating) which can berecognized by its odour and its action upon filter paper moistened with potassiumiodate and starch solution.

S2O32- + 2 H+ → S ↓ + SO2 ↑ + H2O

2. Barium chloride solution: white precipitate of barium thiosulphate frommoderately concentrated solutions.

S2O32- + Ba2+ → BaS2O3 ↓

59

3. Silver nitrate solution: white precipitate of silver thiosulphate:

S2O32- + 2 Ag+ → Ag2S2O3 ↓

At first no precipitation occurs because the soluble dithiosulphatoargentate(I)complex is formed. The precipitate is unstable, turning dark on standing, when silversulphide is formed:

Ag2S2O3 ↓ + H2O → Ag2S ↓ + 2 H+ + SO42-

The decomposition can be accelerated by warming.

4. Lead acetate or lead nitrate solution: first no change, but on further addition ofthe reagent a white precipitate of lead thiosulphate is formed:

S2O32- + Pb2+ → PbS2O3 ↓

The precipitate is soluble in excess thiosulphate.On boiling the suspension the precipitate darkens, forming finally a black precipitateof lead sulphide:

PbS2O3 ↓ + H2O → PbS ↓ + 2 H+ + SO42-

5. Iodine solution: decolourized when a colourless solution of tetrathionate ions isformed:

I2 + 2 S2O32- → 2 I- + S4O62-

Sulphates, SO42-

The sulphates of barium, strontium, and lead are practically insoluble in water,those of calcium and mercury(II) are slightly soluble.

Most of the remaining metallic sulphates are soluble in water.

Some basic sulphates such as those of mercury, bismuth, and chromium, arealso insoluble in water, but these dissolve in dilute hydrochloric or nitric acid.

To study the reactions of sulphates use a 0.1M solution of sodium sulphate.

1. Hydrochloric acid: no visible change.

2. Barium chloride solution: white precipitate of barium sulphate, solubility product:Ksp(BaSO4, 25°C)= 1.07x10−10, insoluble in warm dilute hydrochloric acid and indilute nitric acid.

SO42- + Ba2+ → BaSO4 ↓

60

3. Lead acetate solution: white precipitate of lead sulphate, PbSO4 (solubilityproduct: Ksp(PbSO4, 25°C)= 1.82x10−8), soluble in hot concentrated sulphuric acid, insolutions of ammonium acetate and of ammonium tartarate, and in sodium hydroxidesolution.

SO42- + Pb2+ → PbSO4 ↓

4. Silver nitrate solution: no precipitate in dilute solutions, but white, crystallineprecipitate of silver sulphate (Ksp(Ag2SO4, 25°C)= 1.20x10−5) from concentratedsolutions:

SO42- + 2 Ag+ → Ag2SO4 ↓

5. Mercury(II) nitrate solution: yellow precipitate of basic mercury(II) sulphate:

SO42- + 3 Hg2+ + 2 H2O → HgSO4.2HgO ↓ + 4 H+

This is a sensitive test, given even by suspensions of barium or lead sulphate.

61

Summarise the reactions of sulphide, sulphite, sulphate, and thiosulphate anionsby filling in the following tables:

HCl Ag+ Ba2+ Pb2+ Sr2+

S2-

SO32-

S2O32-

SO42-

Cl2-water I2(soln. is

decolourised)

KI(I2

precipitates)

KMnO4acidic media(decolourised

)

S2-

SO32-

S2O32-

SO42-

62

Summarise the reactions of As3+, Sb3+, Bi3+, Sn2+, Sn4+, and Pb2+ ions:

As3+ Sb3+ Bi3+ Sn2+ Sn4+ Pb2+

HCl

H2S

precipitate+

(NH4)2S

(NH4)2Sx

HNO3

cc HCl

NaOH

NH4OH

KI

SO42-

PO43-

CO32-

NaOH

Zn +H2SO4

Fe

63

The Group VIIa Elements (F, Cl, Br, I) andTheir Anions

Because of their reactivity, none of the halogens occurs in the elemental statein Nature. Since the atoms are only one electron short of the noble gas configuration,the elements readily form the anion X- or a single covalent bond.

Fluorine (F2) is a yellow gas. It is the chemically most reactive of all the elementsand combines directly at ordinary or elevated temperatures with all he elements otherthan nitrogen, oxygen, and the lighter noble gases, and also attacks many othercompounds. It reacts with water (chemically soluble in water):

2 F2 + 2 H2O → 2 H2F2 + O2

Chlorine (Cl2) is a greenish gas. It is moderately soluble in water, with which itpartially reacts (physical and chemical solubility):

Cl2 + H2O → HCl + HOCl

Bromine (Br2) is a dense, mobile, dark red liquid at room temperature. It ismoderately soluble in water and miscible with non-polar solvents such as CCl4 andCS2.Iodine (I2) is a black solid with a slight metallic lustre. At atmospheric pressure itsublimes (violet vapour) without melting. It is only slightly soluble in water (thesolubility in water is two order smaller than that of bromine), but it is readily solublein non polar solvents such as CCl4 and CS2 to give violet solutions. Iodine solutionsare brown in solvents such as unsaturated hydrocarbons, liquid SO2, alcohols, andketones, and pinkish brown in benzene.Although iodine is slightly soluble in water, it is very soluble in aqueous solution ofpotassium iodide, because iodide ions have a pronounced tendency to interact withone (or more) molecules of I2 to form polyiodide anions:

I- + I2 → I3-

Solubility of halogens in aqueous alkaliFluorine reacts fast with aqueous alkali forming fluorides and oxygen gas.By dissolving the halogens in cold base, halide (X-) and hipohalite (XO-) ions

are produced in principle according to the general reaction:

X2 + 2 OH- → X- + XO- + H2O

This situation, however, is complicated by the tendency of the hypohalite ions todisproportionate further in basic solution to produce the halate (XO3-) ions:

3 XO- → 2 X- + XO3-

The disproportionation of ClO- is slow at and below room temperature, thuswhen chlorine reacts with base " in the cold," reasonably pure solutions of Cl- and

64

ClO- are obtained. In hot solutions (70-90 °C) the rate of disproportionation is fairlyrapid and under proper conditions, good yields of ClO3- can be secured.

The disproportionation of BrO- is moderately fast even at room temperature.Consequently solutions of BrO- can only be made and kept at around 0 °C. Attemperatures of 50 to 80 °C quantitative yields of BrO3- are obtained:

3 Br2 + 6 OH- → 5 Br- + BrO3- + 3 H2O

The rate of disproportionation of IO- is very fast at all temperatures, so that itis unknown in solutions. Reaction of iodine with base gives IO3- quantitatively:

3 I2 + 6 OH- → 5 I- + IO3- + 3 H2O

Principal anions and oxoanions of halogensF Cl Br I

F− fluoride Cl− chlorideClO− hypochlorite

ClO2− chlorite

ClO3− chlorate

ClO4− perchlorate

Br− bromide

BrO3− bromate

I− iodide

IO3− iodate

Characteristic reactions of halide ionsTo study the reactions of fluorides, chlorides, bromides, and iodides

use a 0.1 M solution of sodium fluoride (NaF), sodium chloride (NaCl), sodiumbromide (NaBr), and potassium iodide (KI), respectively.

Fluorides, F-

The fluorides of the common alkali metals and of silver, mercury, aluminium,and nickel are readily soluble in water.

The fluorides of lead, copper, iron(III), barium, and lithium are slightlysoluble in water.

The fluorides of calcium, strontium, and magnesium are insoluble in water.

E.g. at 20 °C: Compound Solubility ( g / 100 ml H2O )CaF2PbF2

0.00160.064

KFAgF

92.3185

1. Silver nitrate solution: no precipitate, since silver fluoride is soluble in water.

65

2 Calcium chloride solution: white precipitate of calcium fluoride, solubilityproduct: Ksp(CaF2, 25°C)= 1.46x10−10:

2 F- + Ca2+ → CaF2 ↓

3. Barium chloride: white precipitate of barium fluoride, solubility product:Ksp(BaF2, 25°C)= 1.84x10−7:

2 F- + Ba2+ → BaF2 ↓

4. Concentrated sulphuric acid: with the solid fluoride, a colourless, corrosive gas,hydrogen fluoride, H2F2, is evolved on warming; the gas fumes on moist air, and thetest-tube acquires a greasy appearance as a result of the corrosive action of the vapouron the silica in the glass, which liberates the gas, silicon tetrafluoride, SiF4.Gelatinous silicic acid H2SiO3 is deposited on the glass as a product of thedecomposition of the silicon tetrafluoride.

2 F- + H2SO4 → H2F2 ↑ + SO42-SiO2 + 2 H2F2 → SiF4 ↑ + 2 H2O3 SiF4 ↑ + 3 H2O → 2 [SiF6]2- + H2SiO3 ↓ + 4 H+

Chlorides, Cl-

Most chlorides are soluble in water.

Mercury(I) chloride, Hg2Cl2, silver chloride, AgCl, lead chloride, PbCl2 (thisis sparingly soluble in cold but readily soluble in boiling water), copper(I) chloride,CuCl, bismuth oxychloride, BiOCl, antimony oxychloride, SbOCl, and mercury(II)oxychloride, Hg2OCl2, are insoluble in water.

1. Silver nitrate solution: white, curdy precipitate of silver chloride, solubilityproduct: Ksp(AgCl, 25°C)= 1.77x10−10:

Cl- + Ag+ → AgCl ↓

The precipitate is insoluble in water and dilute nitric acid, but soluble in diluteammonia solution, in conc. ammonium carbonate solution, and in potassium cyanideand sodium thiosulphate solutions:

AgCl ↓ + 2 NH3 → [Ag(NH3)2]+ + Cl-AgCl ↓ + (NH4)2CO3 → [Ag(NH3)2]+ + Cl- + CO2 + H2OAgCl ↓ + 2 Na2S2O3 → 4 Na+ + [Ag(S2O3)2]3- + Cl-AgCl ↓ + 2 KCN → 2 K+ + [Ag(CN)2]- + Cl-

66

2. Lead acetate solution: white precipitate of lead chloride from concentratedsolutions, solubility product: Ksp(PbCl2, 25°C)= 1.17x10−5:

2 Cl- + Pb2+ → PbCl2 ↓

3. Concentrated sulphuric acid: decomposition of chloride occurs with theevolution of hydrogen chloride:

Cl- + H2SO4 → HCl + HSO4-

4. Manganese dioxide and concentrated sulphuric acid: chlorine evolves, when thechloride is mixed with manganese dioxide, concentrated sulphuric acid added and themixture gently warmed. Chlorine is identified by its suffocating odour, yellowish-green colour, its bleaching of moistened litmus paper, and turning of potassiumiodide-starch paper blue.

2 Cl- + MnO2 + 2 H2SO4 → Mn2+ + Cl2 ↑ + 2 SO42- + 2 H2O

5. Potassium dichromate and sulphuric acid (chromyl chloride test): the solid chloride is mixed with three times itsweight of powdered potassium dichromate in atest tube, an equal bulk of concentratedsulphuric acid is added and the mixture gentlywarmed. The deep-red vapours of chromylchloride, CrO2Cl2, which are evolved arepassed into sodium hydroxide solutioncontained in a test tube. Chromyl chloride is areadily volatile liquid (b.p. 116.5 oC) anddecompose in sodium hydroxide solution toform sodium chromate, thus a yellow solutionis obtained. The formation of chromate can be confirmed by acidifying the solution with sulphuric acid, adding 1-2 ml amyl alcoholfollowed by a little hydrogen peroxide solution. The organic layer is coloured blue(formation of CrO5).

4 Cl- + Cr2O72- + 6 H+ → 2 CrO2Cl2 ↑ + 3 H2OCrO2Cl2 + 4 OH- → CrO42- + 2 Cl- + 2 H2O

Bromides and chlorides give rise to the free halogens, which yield colourlesssolutions with sodium hydroxide. If the ratio of iodide and chloride exceeds 1:15, thechromyl chloride formation is largely prevented and chlorine is evolved.Fluorides give rise to the volatile chromyl fluoride, which is decomposed by water,and hence should be absent or removed, if this test is used for identifying chlorides.

67

Bromides, Br-Silver, mercury(I), and copper(I) are insoluble in water. Lead bromide is

sparingly soluble in cold, but more soluble in boiling water.All other bromides are soluble in water.

1. Silver nitrate solution: pale-yellow precipitate of silver bromide, solubilityproduct: Ksp(AgBr, 25°C)= 5.35x10−13:

Br- + Ag+ → AgBr ↓

The precipitate is sparingly soluble in dilute, but readily soluble in concentratedammonia solution. The precipitate is also soluble in potassium cyanide and sodiumthiosulphate solutions, but insoluble in dilute nitric acid.

AgBr ↓ + 2 NH3 → [Ag(NH3)2]+ + Br-AgBr ↓ + 2 Na2S2O3 → 4 Na+ + [Ag(S2O3)2]3- + Br-AgBr ↓ + 2 KCN → 2 K+ + [Ag(CN)2]- + Br-

2. Lead acetate solution: white crystalline precipitate of lead bromide solubilityproduct: Ksp(PbBr2, 25°C)= 6.60x10−6:

2 Br- + Pb2+ → PbBr2 ↓

The precipitate is soluble in boiling water, and reappears by cooling down the hotsolution.

3. Concentrated sulphuric acid: first a reddish-brown solution is formed, laterreddish-brown vapours containing bromine and hydrogen bromide is evolved:

Br- + H2SO4 → HBr ↑ + HSO4-2 Br- + 2 H2SO4 → Br2 ↑ + SO2 ↑ + SO42- + 2 H2O

These reactions are accelerated by warming.

4. Manganese dioxide and concentrated sulphuric acid: reddish-brown vapours ofbromine are evolved. Bromine is recognised by its irritating odour.

2 Br- + MnO2 + 2 H2SO4 → Br2 ↑ + Mn2+ + 2 SO42- + 2 H2O

5. Chlorine water: the addition of this reagent dropwise to a solution of a bromideliberates free bromine, which colours the solution orange-red; if carbon tetrachlorideis added and the liquid is shaken, the bromine dissolves in the organic solvent and,after allowing to stand, forms a reddish-brown solution below the colourless aqueouslayer.With excess chlorine water, the bromine is converted into yellow brominemonochloride (or partially into colourless hypobromous or bromic acid) and a pale-yellow solution results (difference from iodide).

2 Br- + Cl2 → Br2 + 2 Cl-Br2 + Cl2 → 2 BrCl

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6. Nitric acid: hot, fairly concentrated (1:1) nitric acid oxidises bromides tobromine:

6 Br- + 8 HNO3 → 3 Br2 ↑ + 2 NO ↑ + 6 NO3- + 4 H2O

7. Potassium dichromate and sulphuric acid: on gently warming a mixture of abromide, concentrated sulphuric acid, and potassium dichromate, and passing theevolved vapours into water, a yellowish-brown solution, containing free bromine butno chromium, is produced. A colourless solution is obtained on treatment withsodium hydroxide solution; this does not give the chromate reaction with sulphuricacid, hydrogen peroxide and amyl alcohol (distinction from chloride).

6 Br- + Cr2O72- + 7 H2SO4 → 3 Br2 ↑ + 2 Cr3+ + 7 SO42- + 7 H2O

Iodides, I-

Silver, mercury, and copper(I) iodides are insoluble in water. Lead iodide issparingly soluble in cold, but more soluble in boiling water.

All other iodides are soluble in water.

E.g.: Compound Solubility ( g / 100 ml water) 20 °C 100 °C

AgIPbI2

----- ----- 0,063 0,41

CaI2 209 426

1. Silver nitrate solution: yellow, curdy precipitate of silver iodide, solubilityproduct: Ksp(AgI, 25°C)= 8.51x10−17, readily soluble in potassium cyanide and insodium thiosulphate solutions, very slightly soluble in concentrated ammoniasolution, and insoluble in dilute nitric acid:

I- + Ag+ → AgI ↓AgI ↓ + 2 Na2S2O3 → 4 Na+ + [Ag(S2O3)2]3- + I-AgI ↓ + 2 KCN → 2 K+ + [Ag(CN)2]- + I-

2. Lead acetate solution: yellow precipitate of lead iodide, solubility product:Ksp(PbI2, 25°C)= 8.49x10−9, soluble in much hot water forming a colourless solution,and yielding golden-yellow plates on cooling:

2 I- + Pb2+ → PbI2 ↓

3. Concentrated sulphuric acid: iodine is liberated; on warming, violet vapours areevolved, which turn starch paper blue. Some hydrogen iodide is formed, but most of it

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reduces the sulphuric acid to sulphur dioxide, hydrogen sulphide, and sulphur, therelative proportions of which depend upon the concentrations of the reagents:

2 I- + 2 H2SO4 → I2 ↑ + SO42- + 2 H2O + SO2 ↑

4. Manganese dioxide and concentrated sulphuric acid: only iodine is formed andthe sulphuric acid does not get reduced:

2 I- + MnO2 + 2 H2SO4 → I2 ↑ + Mn2+ + 2 SO42- + 2 H2O

5. Chlorine water: when this reagent is added dropwise to a solution of an iodide,iodine is liberated, which colours the solution brown; on shaking with carbontetrachloride or chloroform it dissolves in the organic phase forming a violet solution.The free iodine may also be identified by the characteristic blue colour it forms withstarch solution. If excess chlorine water is added, the iodine is oxidised to colourlessiodic acid:

2 I- + Cl2 → I2 + 2 Cl-I2 + 5 Cl2 + 6 H2O → 2 IO3- + 10 Cl- + 12 H+

6. Bromine water: iodine is liberated:

2 I- + Br2 → I2 + 2 Br-

7. Potassium dichromate and sulphuric acid: only iodine is liberated, and nochromate is present in the distillate. Difference from chloride (see chlorides).

8. Sodium nitrite solution: Iodine is liberated when this reagent is added to aniodide solution acidified with dilute acetic or sulphuric acid (difference from bromideand chloride). The iodine may be identified by colouring starch paste blue, or carbontetrachloride violet:

2 I- + 2 NO2- + 4 H+ → I2 + 2 NO ↑ + 2 H2O

9. Copper sulphate solution: brown precipitate consisting of a mixture of copper(I)iodide, CuI (solubility product: Ksp(CuI, 25°C)= 1.27x10−12), and iodine. The iodinemay be removed by the addition of sodium thiosulphate solution, and a nearly whiteprecipitate of copper(I) iodide obtained:

4 I- + 2 Cu2+ → 2 CuI ↓ + I2I2 + 2 S2O32- → 2 I- + S4O62-

10. Mercury(II) chloride solution: scarlet (red) precipitate of mercury(II) iodide,solubility product: Ksp(HgI2, 25°C)= 2.82x10−29:

2 I- + HgCl2 → HgI2 ↓ + 2 Cl-

The mercury(II) iodide precipitate dissolves in excess potassium iodide, forming atetraiodomercurate(II) complex:

HgI2 ↓ + 2 I- → [HgI4]2-

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11. Bismuth(III) nitrate solution: black precipitate of bismuth(III) iodide:

3 I- + Bi3+ → BiI3 ↓

The bismuth(III) iodide precipitate dissolves in excess potassium iodide, whenorange-coloured tetraiodo-bismuthate ions are formed:

BiI3 ↓ + I- → [BiI4]-

12. Starch test: iodides are readily oxidised in acid solution to free iodine by anumber of oxidising agents; the free iodine may then be identified by the deep-bluecoloration produced with starch solution. One of the best oxidising agent to employ isacidified potassium nitrite solution:

2 I- + 2 NO2- + 4 H+ → I2 + 2 NO ↑ + 2 H2O

or acidified hydrogen peroxide: H2O2 + 2 I- + 2 H+ → I2 + 2 H2O

Chlorates, ClO3−

All chlorates are soluble in water.

To study these reactions use a 0.1 M solution of potassium chlorate.

1. Concentrated sulphuric acid. All chlorates are decomposed with the formation ofthe greenich-yellow gas, chlorine dioxide, ClO2, which dissolves in the sulphuric acidto give an orange-yellow solution. On warming gently an explosive crackling occurs,which may develop into a violent explosion (Danger!!!!). In carrying out this test oneor two small crystalls of potassium chlorate are treated with 1 ml concentratedsulphuric acid in the cold; the yellow explosive chlorine dioxide can be seen onshaking the solution. The test-tube should not be warmed.

3 KClO3 + 3 H2SO4 → 2 ClO2 ↑ + ClO4− + 3 SO4

2− + 4 H+ + 3 K+ + H2O

2. Concentrated hydrochloric acid. All chlorates are decomposed by this acid, andchlorine, together with varying quantities of the explosive chlorine dioxide, isevolved; chlorine dioxide imparts a yellow colour to the acid. The experiment shouldbe conducted on a very small scale. The following two chemical reactions probablyoccur simultaneously:

2 KClO3 + 4 HCl → 2 ClO2 + Cl2 ↑ + 2 K+ + 2 Cl− + 2 H2O KClO3 + 6 HCl → 3 Cl2 ↑ + K+ + Cl− + 3 H2O

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3. Sodium nitrite solution. On warming this reagent with a solution of the chlorate,the latter is reduced to a chloride, which may be identified by adding silver nitratesolution after acidification with dilute nitric acid.

ClO3− + 3 NO2

− → Cl− + 3 NO3−

A solution of sulphurous acid acts similarly:

ClO3− + 3 H2SO3 → Cl− + 3 SO4

2− + 6 H+

4. Zinc and sodium hydroxide solution: the chlorate is reduced to a chloride.

ClO3− + 3 Zn + 6 OH− + 3 H2O → Cl− + 3 [Zn(OH)4]2−

The solution is acidified with dilute nitric acid after several minutes boiling andremoving the unreacted zinc with filtration, and silver nitrate is added:

Ag+ + Cl− → AgCl ↓

5. Potassium iodide solution: iodine is liberated if a mineral acid is present. If aceticacid is used, no iodine separates even on long standing.

ClO3− + 6 I− + 6 H+ → 3 I2 ↓ + Cl− + 3 H2O

6. Iron(II) sulphate solution: reduction to chloride upon boiling in the presence ofmineral acid.

ClO3− + 6 Fe2+ + 6 H+ → Cl− + 6 Fe3+ + 3 H2O

7. Silver nitrate solution: no precipitate in neutral solution or in the presence ofdilute nitric acid.

8. Barium chloride solution: no precipitate is obtained.

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Summarise the reactions of halides by filling in the following tables:

Ag+ Pb2+ Hg22+ Hg2+ Bi3+ Ba2+

F-

Cl-

Br-

I-

Cl2-water Br2-water cc H2SO4or HNO3

FeCl3 H2O2

Cl-

Br-

I-

Summarise the solubility of silver halide precipitates:

diluteHNO3 (NH4)2CO3 NH4OH Na2S2O3 KCN

AgCl

AgBr

AgI

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Pseudohalogens and Pseudohalides

A number of nitrogen compounds are known which undergo many reactionssuggestive of the halogens. These are called the "pseudohalogens". They include suchcompounds as cyanogen (CN)2, thiocyanogen (SCN)2, oxycyanogen (OCN)2, andselenocyanogen (SeCN)2. The pseudohalogens, like the halogens, are dimeric,oxidising, and capable of forming many stable salts. They also possess outer lonepairs which become available for coordination as the atoms acquire negative charge,enabling them to form very numerous and stable complexes. Having electron donorpairs on more than one atom, they are especially able to bridge acceptors together.Like the halogens they can be prepared by electrolytic or chemical oxidation of theirsimple salts. The pseudohalides are derived from pseudohalogens, like halides fromhalogens. Azides are sometimes included also even though no corresponding dimerexists.

Pseudohalogen Pseudohalide

(CN)2 cyanogen (OCN)2 oxycyanogen (SCN)2 tiocyanogen (SeCN)2 selenocyanogen ----

CN− cyanide OCN− cyanate SCN− tiocyanate SeCN− selenocyanate NNN− azide

Characteristic reactions of thiocyanates, SCN− To study the reactions of thiocyanates use a 0.1 M solution of

potassium thiocyanate (KSCN).

The thiocyanates of silver and copper(I) are practically insoluble in water, thethiocyanates of mercury(II) and lead are sparingly soluble.

The thiocyanates of most other metals are soluble in water.

1. Silver nitrate solution: white, curdy precipitate of silver thiocyanate, solubilityproduct: Ksp(AgSCN, 25°C)= 1.03x10−12:

SCN− + Ag+ → AgSCN ↓

The precipitate is insoluble in dilute nitric acid, but soluble in ammonia solution:

AgSCN ↓ + 2 NH3 → [Ag(NH3)2]+ + SCN-

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2. Copper sulphate solution: first a green colouration, than a black precipitate ofcopper(II) thiocyanate is observed:

2 SCN− + Cu2+ → Cu(SCN)2 ↓

3. Mercury(II) nitrate solution: white precipitate of mercury(II) thiocyanate,Hg(SCN)2, readily soluble in excess of the thiocyanate solution:

2 SCN− + Hg2+ → Hg(SCN)2 ↓Hg(SCN)2 ↓ + 2 SCN- → [Hg(SCN)4]2-

4. Iron(III) chloride solution: blood-red colouration, due to the formation of acomplex:

3 SCN− + Fe3+ ↔ Fe(SCN)3

The Fe(SCN)3 complex can be extracted by shaking with ether.The red colour is removed by fluorides, when colourless, more stable fluoro-complexes are formed:

Fe(SCN)3 + 6 F− → [FeF6]3− + 3 SCN−

5. Cobalt nitrate solution: blue colouration, due to the formation oftetrathiocyanatocobaltate(II) ions:

Co2+ + 4 SCN− → [Co(SCN)4]2−

If amyl alcohol or ether is added the free acid H2[Co(SCN)4] is formed and dissolvedby the organic solvent.

2 H+ + [Co(SCN)4]2− ↔ H2[Co(SCN)4]

The test is rendered more sensitive if the solution is acidified with concentratedhydrochloric acid, when the equilibrium shifts towards the formation of the free acid.

6. Hydrochloric acid (distillation test): Free isothiocyanic acid, HNCS, can beliberated by hydrochloric acid, distilled into ammonia solution, where it can beidentified with iron(III) chloride.

Place a few drops of the test solution in atest tube, acidify with dilute hydrochloricacid, boil the solution in the test tubegently so as to distil any HNCS into theammonia solution. After distillation,acidify slightly the ammonia solutionwith dilute hydrochloric acid and add adrop of iron(III) chloride solution.A red colouration is observed.

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7. Zinc and dilute hydrochloric acid: hydrogen sulphide and hydrogen cyanide areevolved:

SCN− + Zn + 3 H+ → H2S ↑ + HCN ↑ + Zn2+

The gas may be identified by the blackening of filter paper moistened with leadacetate solution:

H2S ↑ + Pb2+ → PbS ↓ + 2 H+

8. Sodium nitrite solution and dilute hydrochloric acid: the addition of a nitritesolution to a solution of potassium thiocyanate, followed by acidification with dilutehydrochloric acid, produces a red colour (similar to that of Fe(SCN)3) due to nitrosylthiocyanate:

SCN− + NO2− + 2 H+ → ONSCN + H2O

If carbon tetrachloride or ether is added, ONSCN is extracted by the organic solvent.Nitrosyl thiocyanate is not very stable at room temperature, especially not inconcentrate solutions or at elevated temperatures. It decomposes on heating with theformation of nitrogen oxide. The colourless nitrogen oxide gas, when mixed with air,is oxidised to red nitrogen dioxide:

2 ONSCN → 2 NO + (SCN)22 NO ↑ (colourless) + O2 ↑ → 2 NO2 ↑ (red)

9. Dilute nitric acid: decomposition upon warming, a red colouration is produced,and nitrogen oxyde and hydrogen cyanide are evolved:

SCN− + 2 NO3− + H+ → 2 NO ↑ + HCN ↑ + SO4

2−

With concentrated nitric acid a more vigorous reaction takes place, with the formationof nitrogen oxide and carbon dioxide.

10. Sulphuric acid: With the concentrated acid a yellow colouration is produced inthe cold: upon warming a violent reaction occurs:

SCN− + H2SO4 + H2O → COS ↑ + NH4+ + SO4

2−

The reaction gets slower and slower with dilution of the acid. With the 2.5 Macid no reaction occurs in the cold, but on boiling a yellow solution is formed, sulphurdioxide and a little carbonyl sulphide are evolved.

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Compare the reactions of thiocyanate ion to those of halides:

Ag+ Pb2+ Cu2+ Hg2+ Fe3+ NO2−

SCN-

Cl-

Br-

I-

Pharaoh's SnakeThe famous German chemist Friedrich Wöhler invented the trick known as

Pharaoh's snake in the XIXth century.If a pellet of Hg(SCN)2 is placed on a fireproof support or in a porcelain dish andlight is put to it, a porous mass emerges, which resembles a snake and grows everlarger.

Hg(SCN)2 may be prepared by dissolvingmercury metal in nitric acid, andprecipitating the mercury ions withpotassium thiocyanate solution as a whitepowderlike precipitate:

3 Hg + 8 HNO3 → 3 Hg(NO3)2 + 2 NO + 4 H2OHg2+ + 2 SCN− → Hg(SCN)2 ↓

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Summarise the reactions of anions:

HCl Ba2+ Ca2+ Na+ Ag+ Pb2+

CO32-

SiO32-

S2-

SO32-

S2O32-

SO42-

BO33-

PO43-

F-

Cl-

Br-

I-

NO2-

NO3-