78

The Group Ib Elements (Cu, Ag, Au) andTheir Principle Ions

. Copper is reddish coloured metal, takes on a bright metallic luster, and is soft,malleable, ductile, and a good conductor of heat and electricity (second only to silverin electrical conductivity). It melts at 1083 °C. It is only superficially oxidised in air,sometimes acquiring a green coating of hydroxo carbonate and hydroxo sulphate.

Silver. Pure silver has a brilliant white metallic luster. It is a little harder thangold and is very ductile and malleable, being exceeded only by gold and perhapspalladium. Pure silver has the highest electrical and thermal conductivity of allmetals. It melts at 962 °C. It is stable in pure air and water, but tarnishes whenexposed to ozone, hydrogen sulphide, or air containing sulphur.

Gold is a heavy metal with its characteristic yellow colour when in a mass. Inpowderous form it is usually reddish-brown, but when finely divided it may be black,ruby, or purple. It melts at 1064 °C. Gold is the most malleable and ductile metal, it issoft and a good conductor of heat and electricity, and is unaffected by air and mostreagents.

Solubility in acids and alkalis Because of their positive standard electrode potential copper, silver, and gold

are insoluble in hydrochloric acid and in dilute sulphuric acid.

Hot, concentrated sulphuric acid dissolves copper and silver, but gold isresistant against it.

Cu + 2 H2SO4 → Cu2+ + SO42− + SO2 ↑ + 2 H2O

2 Ag + 2 H2SO4 → 2 Ag+ + SO42− + SO2 ↑ + 2 H2O

Medium-concentrated (8M) nitric acid also dissolves copper and silver, butgold is resistant.

3 Cu + 8 HNO3 → 3 Cu2+ + 6 NO3− + 2 NO ↑ + 4 H2O

6 Ag + 8 HNO3 → 6 Ag+ + 6 NO3− + 2 NO ↑ + 4 H2O

Aqua regia dissolves copper and gold:

3 Cu + 6 HCl + 2 HNO3 → 3 Cu2+ + 6 Cl− + 2 NO ↑ + 4 H2OAu + 4 HCl + HNO3 → H[AuCl4] + NO ↑ + 2 H2O

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Silver and gold are resistant against alkalis.Copper is hardly soluble in strong alkalis in the presence of oxygen, but it is

soluble in concentrated ammonia solution in the presence of oxygen:

4 Cu + 8 NH3 + 2 H2O + O2 → 4 [Cu(NH3)2]+ + 4 OH−

Copper, silver, and gold is soluble in alkali cyanide solutions in the presenceof oxygen:

4 Cu + 8 KCN + 2 H2O + O2 → 8 K+ + 4 [Cu(CN)2]− + 4 OH−

4 Ag + 8 KCN + 2 H2O + O2 → 8 K+ + 4 [Ag(CN)2]− + 4 OH−

4 Au + 8 KCN + 2 H2O + O2 → 8 K+ + 4 [Au(CN)2]− + 4 OH−

Compare thestandard redox potentials

Au-Au+= +1.69 VAu-Au3+= +1.50 V

NO-NO3−= +0.96 V

Au-[AuCl4]−= +1.00 VAg-Ag+= +0.80 VCu-Cu+= +0.52 V

OH−-O2= +0.40 VCu-Cu2+= +0.34 V

H2-H+= 0.0 VCu-[Cu(NH3)2]+= -0.12 VAg-[Ag(CN)2]−= -0.31 VCu-[Cu(CN)2]−= -0.43 VAu-[Au(CN)2]−= -0.60 V

Summarise the solubility of selected elements in cold concentrated nitric acid:

Au Be Al C Si Pb

cc HNO3

80

Principal cations of copper, silver, and gold (Cu+) *

Cu2+Ag+

(Ag2+)*** (Au+) **

Au3+

* Copper(I) ions are unstable in aqueous solution. Copper(I) compounds are colourless and most ofthem are insoluble in water.** Au+ ions are exceedingly unstable with respect to the disproportionation to Au and Au3+.*** The ion is unstable to water (reduced by water into Ag+).

Oxides of copper, silver and gold

oxide Cu2O CuO Ag2O AgO Au2O3

colour red black brown black brown

Copper oxides are insoluble in water. CuO is soluble in acids, in NH4Cl orKCN solutions. Cu2O is soluble in hydrochloric acid, in ammonia, and in NH4Clsolutions, slightly soluble in dilute nitric acid.

Silver oxides are soluble in nitric acid, in sulphuric acid, and also in ammoniasolution. Ag2O is slightly soluble in water (solubility at 20 °C is 0.0013 g/100 mlwater) and its aqueous suspensions are alkaline. It is more soluble in strongly alkalinesolutions than in water, and AgOH and Ag(OH)2

− are formed. AgO is of littleimportance, it is actually AgIAgIIIO2.

Gold(III) oxide is of little importance. It is obtained in hydrated form as anamorphous brown precipitate on addition of base to AuCl4

− solutions. It is weaklyacidic and dissolves in excess strong base, probably as Au(OH)4

−. It is soluble inhydrochloric acid, in concentrated nitric acid, and in alkali cyanide solutions.

Reactions of copper(II) ions, Cu2+

Solubility of the most common copper(II) compoundsCopper(II) chloride, chlorate, nitrate, and sulphate are soluble in water.

Copper(II) acetate and fluoride are slightly soluble in water, and all the othercompounds are practically insoluble.

E.g. at 20 °C Compound Solubility ( g / 100 ml H2O)CuCl2CuF2

70,6 (0 °C-on)4,7

CuS 0,000033

To study these reactions use a 0.1 M solution of copper(II) sulphate.

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1. Hydrogen sulphide gas: black precipitate of copper(II) sulphide. Solubilityproduct: Ksp(CuS, 25 °C)= 1.27x10−36.

Cu2+ + H2S → CuS ↓ + 2 H+

The solution must be acidic in order to obtain a crystalline, well-filterable precipitate.

The precipitate is insoluble in hydrochloric acid, in boiling dilute sulphuric acid, insodium hydroxide, in ammonium sulphide, in sodium sulphide, and only very slightlysoluble in polysulphides.

Hot, concentrated nitric acid dissolves the copper(II) sulphide, leaving behind sulphuras a white precipitate:

3 CuS ↓ + 8 HNO3 → 3 Cu2+ + 6 NO3− + 3 S ↓ + 2 NO ↑ + 4 H2O

When boiled for longer, sulphur is oxidised to sulphuric acid and a clear, bluesolution is obtained:

S ↓ + 2 HNO3 → 2 H+ + SO42− + 2 NO ↑

Potassium cyanide dissolves the precipitate, when colourless tetracyanocuprate(I)ions and disulphide ions are formed (copper is reduced, sulphur is oxidised):

2 CuS ↓ + 8 CN− → 2 [Cu(CN)4]3− + S22−

2. Ammonia solution: when added sparingly, a blue precipitate of basic coppersulphate is obtained:

2 Cu2+ + SO42− + 2 NH3 + 2 H2O → Cu(OH)2.CuSO4 ↓ + 2 NH4

+

the precipitate is soluble in excess reagent, when a deep blue coloration is obtainedowing to the formation of tetraamminocuprate(II) complex ions:

Cu(OH)2.CuSO4 ↓ + 8 NH3 → 2 [Cu(NH3)4]2+ + SO42− + 2 OH−

If the solution contains ammonium salts, precipitation does not occur at all, but theblue colour is formed right away. (The reaction is characteristic for copper(II) ions inthe absence of nickel.)

3. Sodium hydroxide in cold solution: blue precipitate of copper(II) hydroxide:

Cu2+ + 2 OH− → Cu(OH)2 ↓

The precipitate is insoluble in excess reagent, but soluble in ammonia solution when adeep blue coloration is obtained:

Cu(OH)2 ↓ + 4 NH3 → [Cu(NH3)4]2+ + 2 OH−

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When heated, the precipitate is converted to black copper(II) oxide bydehydration:

Cu(OH)2 ↓ → CuO ↓ + H2O

4. Potassium iodide solution: precipitates white copper(I) iodide (solubilityproduct: Ksp(CuI, 25 °C)= 1.27x10−12), but the solution is intensely brownbecause of the formation of tri-iodide ions (iodine):

2 Cu2+ + 5 I− → 2 CuI ↓ + I3−

Adding an excess of sodium thiosulphate to the solution, tri-iodide ions are reduced tocolourless iodide ions and the white colour of the precipitate becomes visible:

I3− + 2 S2O3

2− → 3 I− + S4O62−

5. Potassium cyanide: when added sparingly forms first a yellow precipitate ofcopper(II) cyanide:

Cu2+ + 2 CN− → Cu(CN)2 ↓

The precipitate quickly decomposes into white copper(I) cyanide and cyanogen(highly poisonous gas):

2 Cu(CN)2 ↓ → 2 CuCN ↓ + (CN)2 ↑

The precipitate dissolves in excess reagent, when colourless tetracyanocuprate(I)complex is formed:

CuCN ↓ + 3 CN− → [Cu(CN)4]3−

The complex is so stable (i.e. the concentration of copper(I) ions is so low) thathydrogen sulphide cannot precipitate copper(I) sulphide from this solution.(Solubility product: Ksp(Cu2S, 25 °C)= 2.26x10−48.)

6. Potassium thiocyanate: black precipitate of copper(II) thiocyanate.

Cu2+ + 2 SCN− → Cu(SCN)2 ↓

The precipitate decomposes slowly to form white copper(I) thiocyanate (Solubilityproduct: Ksp(CuSCN, 25 °C)= 1.77x10−13):

2 Cu(SCN)2 ↓ → 2 CuSCN ↓ + (SCN)2 ↑

Copper(II) thiocyanate can be transformed to copper(I) thiocyanate immediately byadding a suitable reducing agent; e.g. a saturated solution of sulphur dioxide:

2 Cu(SCN)2 ↓ + SO2 + 2 H2O → 2 CuSCN ↓ + 2 SCN− + SO42− + 4 H+

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8. Iron: if a clean iron nail is immersed in a solution of a copper salt, a red depositof copper is obtained and equivalent amount of iron dissolves:

Cu2+ + Fe → Cu + Fe2+

The electrode potential of the copper-copper(II) system is more positive than that ofthe iron-iron(II) system.

9. Flame test: green colour is imparted to the Bunsen flame.

Compare the characteristic reactions of copper(II) and bismuth(III):

H2S NH3 soln.in excess

NaOH KI Fe flametest

Cu2+

Bi3+

Reactions of silver(I) ions, Ag+

Solubility of the most common silver(I) compoundsSilver nitrate, fluoride, chlorate, and perchlorate are readily soluble in water,

silver acetate, nitrite, and sulphate are slightly soluble, while all the other silvercompounds are practically insoluble.

E.g. at 0 °C: Compound Solubility ( g / 100 ml H2O)AgNO3Ag2SO4

1220,57

AgI 0,0000002

To study these reactions use a 0.1 M solution of silver(I) nitrate.

1. Dilute hydrochloric acid (or soluble chlorides): white precipitate of silverchloride. Solubility product: Ksp(AgCl, 25 °C)= 1.77x10−10.

Ag+ + Cl− → AgCl ↓

With concentrated hydrochloric acid precipitation does not occur. Decanting theliquid from over the precipitate, it dissolves in concentrated hydrochloric acid, when adichloroargentate complex is formed:

AgCl ↓ + Cl− ↔ [AgCl2]−

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By diluting with water, the equilibrium shifts back to the left and the precipitatereappears.

Dilute ammonia solution dissolves the precipitate when diamminoargentatecomplex ion is formed:

AgCl ↓ + 2 NH3 ↔ [Ag(NH3)2]+ + Cl−

Dilute nitric acid or hydrochloric acid neutralises the excess ammonia, and theprecipitate reappears because of the equilibrium is shifted back towards the left.

Potassium cyanide or sodium thiosulphate dissolves the AgCl precipitate:

AgCl ↓ + 2 CN− → [Ag(CN)2]− + Cl−

AgCl ↓ + 2 S2O32− → [Ag(S2O3)2]3− + Cl−

2. Hydrogen sulphide gas: in neutral or acidic medium, black precipitate of silversulphide. Solubility product: Ksp(Ag2S, 25 °C)= 6.69x10−50.

2 Ag+ + H2S → Ag2S ↓ + 2 H+

The precipitate is insoluble in ammonium sulphide, ammonium polysulphide,ammonia, potassium cyanide, or sodium thiosulphate. (Silver sulphide can beprecipitated from solutions containing dicyanato- or dithiosulphato-argentatecomplexes with hydrogen sulphide.)

Hot, concentrated nitric acid decomposes the silver sulphide, and sulphurremains in the form of a white precipitate:

3 Ag2S ↓ + 8 HNO3 → 6 Ag+ + 6 NO3− + S ↓ + 2 NO ↑ + 4 H2O

If the mixture is heated with concentrated nitric acid for a considerable time, sulphuris oxidised to sulphate and the precipitate disappears:

S ↓ + 2 HNO3 → 2 H+ + SO42− + 2 NO ↑

3. Ammonia solution: brown precipitate of silver oxide:

2 Ag+ + 2 NH3 + H2O → Ag2O ↓ + 2 NH4+

The reaction reaches an equilibrium and therefore precipitation is incomplete at anystage. The precipitate dissolves in diluted nitric acid and also in excess of the reagent:

Ag2O ↓ + 2 HNO3 → 2 Ag+ + 2 NO3− + H2O

Ag2O ↓ + 4 NH3 + H2O → 2 [Ag(NH3)2]+ + 2 OH−

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4. Sodium hydroxide: brown precipitate of silver oxide:

2 Ag+ + OH− → Ag2O ↓ + H2O

The precipitate is insoluble in excess reagent.The precipitate dissolves in ammonia solution or in nitric acid:

Ag2O ↓ + 4 NH3 + H2O → 2 [Ag(NH3)2]+ + 2 OH−

Ag2O ↓ + 2 HNO3 → 2 Ag+ + 2 NO3− + H2O

A well-washed suspension of the precipitate shows a slight alkaline reaction owing tothe hydrolysis equilibrium:

Ag2O ↓ + H2O ↔ 2 AgOH ↓ ↔ 2 Ag+ + 2 OH−

5. Potassium iodide: yellow precipitate of silver iodide. Solubility product:Ksp(AgI, 25 °C)= 8.51x10−17.

Ag+ + I− → AgI ↓

The precipitate is insoluble in dilute or concentrated ammonia, butdissolves readily in potassium cyanide or in sodium thiosulphate solution:

AgI ↓ + 2 CN− → [Ag(CN)2]− + I−

AgI ↓ + 2 S2O32− → [Ag(S2O3)2]3− + I−

6. Potassium chromate in neutral solution: brownish-red precipitate of silverchromate. Solubility product: Ksp(Ag2CrO4, 25 °C)= 1.12x10−12.

2 Ag+ + CrO42− → Ag2CrO4 ↓

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

Ag2CrO4 ↓ + 4 NH3 → 2 [Ag(NH3)2]+ + CrO42−

2 Ag2CrO4 ↓ + 2 H+ ↔ 4 Ag+ + Cr2O72− + H2O

7. Potassium cyanide solution: when added dropwise to a neutral solution of silvernitrate, white precipitate of silver cyanide is formed. Solubility product:Ksp(AgCN, 25 °C)= 5.97x10−17.

Ag+ + CN− → AgCN ↓

When potassium cyanide is added in excess, the precipitate disappears owing to theformation of dicyanoargentate ions:

AgCN ↓ + CN− → [Ag(CN)2]−

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8. Sodium carbonate solution: yellowish-white precipitate of silver carbonate.Solubility product: Ksp(Ag2CO3, 25 °C)= 8.45x10−12.

2 Ag+ + CO32− → Ag2CO3 ↓

Nitric acid and ammonia solution dissolve the precipitate:

Ag2CO3 ↓ + 2 HNO3 → 2 Ag+ + 2 NO3− + CO2 ↑ + H2O

Ag2CO3 ↓ + 4 NH3 → 2 [Ag(NH3)2]+ + CO32−

When heating, the silver carbonate precipitate decomposes and brown silver oxideprecipitate is formed:

Ag2CO3 ↓ → Ag2O ↓ + CO2 ↑

9. Disodium hydrogen phosphate in neutral solution: yellow precipitate of silverphosphate. Ksp(Ag3PO4, 25 °C)= 8.88x10−17.

3 Ag+ + HPO42− → Ag3PO4 ↓ + H+

Nitric acid and ammonia solution dissolve the precipitate:

Ag3PO4 ↓ + 3 HNO3 → 3 Ag+ + 3 NO3− + H3PO4

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

Fill in the following table:

Fe Zn Sn Cu

Cu2+

Ag+

Standard redox potentials: Fe-Fe2+= -0.44 V; Zn-Zn2+= -0.76 V; Sn-Sn2+= -0.14 V;Cu-Cu2+= +0.34 V; Ag-Ag+= +0.80 V.

87

The Group IIb Elements (Zn, Cd, Hg) andTheir Principle Ions

Zinc is a bluish-white, lustrous metal. It is brittle at ordinary temperatures, butmalleable and ductile at 100 to 150 °C. It melts at 420 °C. It is a fair conductor ofelectricity, and burns in air at high red heat with evolution of white clouds of theoxide. Zinc is stable in air at ordinary conditions, because a protective zinc oxideand/or basic zinc carbonate layer forms on the surface.

Cadmium is a soft, malleable, ductile, and bluish-white metal which is easilycut with a knife. It melts at 321 °C. It is stable in air at ordinary conditions and it issimilar in many respects to zinc.

Mercury is a heavy, silvery-white, liquid metal; it is the only common metalliquid at ordinary temperatures. It melts at −39 °C. It easily forms alloys with manymetals, such as gold, silver, and tin, which are called amalgams.The chemistries of Zn and Cd are very similar, but that of Hg differs considerably andcannot be regarded as homologous.

Solubility in acids and alkalis The very pure zinc metal dissolves very slowly in acids and in alkalis; the

presence of impurities or contact with e.g. platinum or copper accelerates the reaction.This explains the good solubility of commercial zinc.

Zinc, owing to its negative standard electrode potential of −0.76 V, dissolvesreadily in dilute hydrochloric acid and in dilute sulphuric acid with the evolution ofhydrogen:

Zn + 2 H+ → Zn2+ + H2 ↑

With hot, concentrated sulphuric acid, sulphur dioxide is evolved:

Zn + 2 H2SO4 → Zn2+ + SO42− + SO2 ↑ + 2 H2O

Zinc dissolves in very dilute nitric acid, when no gas is evolved:

4 Zn + 10 H+ + NO3− → 4 Zn2+ + NH4

+ + 3 H2O

With increasing concentration of nitric acid, dinitrogen oxide (N2O) and nitric oxide(NO) are formed:

4 Zn + 10 H+ + 2 NO3− → 4 Zn2+ + N2O ↑ + 5 H2O

3 Zn + 8 HNO3 → 3 Zn2+ + 2 NO ↑ + 6 NO32− + 4 H2O

Concentrated nitric acid has little effect on zinc because of the low solubility of zincnitrate in such a medium.

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Cadmium dissolves slowly in dilute acids with the evolution of hydrogen,owing to its negative standard electrode potential of −0.40 V:

Cd + 2 H+ → Cd2+ + H2 ↑

It dissolves in hot, concentrated sulphuric acid with the evolution of sulphur dioxide,and in medium concentrated nitric acid with the evolution of nitrogen monoxide.

Mercury, owing to its positive standard potential (Hg/Hg2+= +0.85 V;Hg/Hg2

2+= +0.80 V), is unaffected when treated with hydrochloric acid or dilute (2M)sulphuric acid.

Hot, concentrated sulphuric acid dissolves mercury. The product is mercury(I)ion if mercury is in excess, while if the acid is in excess, mercury(II) ions are formed:

Hg in excess: 2 Hg + 2 H2SO4 → Hg22+ + SO4

2− + SO2 ↑ + 2 H2Oacid in excess: Hg + 2 H2SO4 → Hg2+ + SO4

2− + SO2 ↑ + 2 H2O

It reacts readily with nitric acid. Cold, medium concentrated (8M) nitric acidwith an excess of mercury yields mercury(I) ions, and with and excess of hotconcentrated nitric acid mercury(II) ions are formed:

Hg in excess: 6 Hg + 8 HNO3 → 3 Hg22+ + 2 NO ↑ + 6 NO3

− + 4 H2Oacid in excess: 3 Hg + 8 HNO3 → 3 Hg2+ + 2 NO ↑ + 6 NO3

− +4 H2O

Zinc dissolves in alkali hydroxides, when tetrahydroxozincate(II) is formed:

Zn + 2 OH− + 2 H2O → [Zn(OH)4]2− + H2 ↑

Cadmium and mercury is insoluble in alkalis.

Summarise the solubility of selected metals in nitric acid.

Ag Pb Zn

cold, concentratedHNO3

medium conc.(8M) HNO3

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Oxides of zinc, cadmium, and mercury:ZnO CdO Hg2O HgOwhite brown black yellow or red

The oxides are practically insoluble in water; ZnO and HgO are a very little soluble,and the solubility at 25 °C is 0.00016 g and 0.0053 g/100 ml of water, respectively.CdO and Hg2O are insoluble.They all are soluble in acids (Hg2O only in nitric acid).

Principal cations of zinc, cadmium, and mercury

Zn2+ Cd2+Hg2

2+

Hg2+

Reactions of zinc(II) ions, Zn2+

Solubility of the most common zinc compoundsZinc(II) chloride, bromide, iodide, chlorate, nitrate, sulphate, and acetate are

soluble in water. Zinc(II) fluoride is very little soluble in water, and all the othercompounds (e.g. sulphide, carbonate, phosphate) are practically insoluble.

E.g. at 20 °C: Compound Solubility ( g / 100 ml H2O)ZnBr2

Zn(NO3)2.6H2O447

184,3

ZnF2 1,62ZnS (β)Zn(CN)2

0,0000650,0005

To study these reactions use a 0.1 M solution of zinc(II) sulphate.

1. Hydrogen sulphide gas: no precipitation occurs in acidic solution (pH about 0−6),only partial precipitation of zinc sulphide in neutral solutions. In alkaline solution,e.g. adding alkali acetate, the precipitation of white zinc(II) sulphide (solubilityproduct: Ksp(ZnS, 25 °C)= 2.93x10−25) is almost complete.

Zn2+ + H2S → ZnS ↓ + 2 H+

Zinc sulphide is also precipitated from alkaline solutions of tetrahydroxozincate:

[Zn(OH)4]2− + H2S → ZnS ↓ + 2 OH− + 2 H2O

2. Ammonium sulphide: white precipitate of zinc sulphide from neutral or alkalinesolutions:

Zn2+ + S2− → ZnS ↓

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The precipitate is insoluble in excess reagent, in acetic acid, and in solutions ofcaustic alkalis, but dissolves in dilute mineral acids.

The precipitate obtained is partially colloidal; it is difficult to wash and tendsto run through the filter paper, particularly on washing. To obtain the zinc sulphide ina form which can be ready filtered, the precipitation is carried out in boiling solution. 3. Ammonia solution: white precipitate of zinc hydroxide (solubility product:

Ksp(Zn(OH)2, 25 °C)= 6.86x10−17), readily soluble in excess reagent and insolutions of ammonium salts owing to the formation of tetramminezincate(II):

Zn2+ + 2 NH3 + 2 H2O ↔ Zn(OH)2 ↓ + 2 NH4+

Zn(OH)2 ↓ + 4 NH3 ↔ [Zn(NH3)4]2+ + 2 OH−

4. Sodium hydroxide: white gelatinous precipitate of zinc(II) hydroxide:

Zn2+ + 2 OH− ↔ Zn(OH)2 ↓

The precipitate is soluble in acids and also in the excess of the reagent:

Zn(OH)2 ↓ + 2 H+ ↔ Zn2+ + 2 H2OZn(OH)2 ↓ + 2 OH− ↔ [Zn(OH)4]2−

5. Disodium hydrogen phosphate solution: white precipitate of zinc phosphate:

3 Zn2+ + 2 HPO42− ↔ Zn3(PO4)2 ↓ + 2 H+

In the presence of ammonium ions zinc ammonium phosphate is formed:

Zn2+ + NH4+ + HPO4

2− ↔ Zn(NH4)PO4 ↓ + H+

Both precipitates are soluble in dilute acids, when the reaction is reversed.Also, both precipitates are soluble in ammonia:

Zn3(PO4)2 ↓ + 12 NH3 → 3 [Zn(NH3)4]2+ + 2 PO43−

Zn(NH4)PO4 ↓ + 3 NH3 → [Zn(NH3)4]2+ + HPO42−

6. Potassium hexacyanoferrate(II) solution: white precipitate of variablecomposition; if the reagent is added in some excess the composition of theprecipitate is K2Zn3[Fe(CN)6]2:

3 Zn2+ + 2 K+ + 2 [Fe(CN)6]4− → K2Zn3[Fe(CN)6]2 ↓

The precipitate is insoluble in dilute acids, but dissolves readily in sodium hydroxide.This reaction can be used to distinguish zinc from aluminium.

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7. Dithizone test. Dithizone (diphenyl thiocarbazone) forms complexes with anumber of metal ions, which can be extracted with carbon tetrachloride. The zinccomplex, formed in neutral, alkaline, or acetic acid solutions, is red:

S CNH

N N

NHS C

NH

N N

NC S

NN

HNNZn

���Zn2+ +2 + + 2 H

Acidify the test solution with acetic acid, and add a few drops of the reagent(dithizone dissolved in carbon tetrachloride). The organic phase turns red in thepresence of zinc. (Cu2+, Hg2

2+, Hg2+, and Ag+ ions interfere.)

Summarise the reactions of selected metal ions with various anions.

F− Cl− Br− S2− SO42− CrO4

2− CO32− OH−

Zn2+

Al3+

Ca2+

Pb2+

Reactions of cadmium(II) ions, Cd2+

Solubility of the most common cadmium compoundsCadmium acetate, sulphate, nitrate, iodide, bromide, chloride, and

chlorate are readily soluble in water, cadmium fluoride is little soluble, while all theother cadmium compounds (e.g. sulphide, carbonate, phosphate) are practicallyinsoluble.

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

0 °C:CdCl2

Cd(NO3)2

140109

25 °C: CdF2 4,3526 °C: Cd(OH)2

Cd3(PO4)2

0,00026-----

To study these reactions use a 0.1 M solution of cadmium(II) sulphate.

1. Hydrogen sulphide gas: in acidic medium, characteristic yellow precipitate ofcadmium sulphide. Solubility product: Ksp(CdS, 25 °C)= 1.40x10−29.

Cd2+ + H2S → CdS ↓ + 2 H+

The reaction is reversible; if the concentration of a strong acid in the solution is above0.5M, precipitation is incomplete. Concentrated acids dissolve the precipitate for thesame reason.The precipitate is insoluble in potassium cyanide (difference from copper ions).

2. Ammonia solution when added dropwise: white precipitate of cadmium(II)hydroxide (Solubility product: Ksp(Cd(OH)2, 25 °C)= 5.27x10−15):

Cd2+ + 2 NH3 + 2 H2O ↔ Cd(OH)2 ↓ + 2 NH4+

The precipitate dissolves in acid when the equilibrium shifts towards the left.An excess of the reagent dissolves the precipitate, when colourlesstetramminecadminate(II) complex ions are formed:

Cd(OH)2 ↓ + 4 NH3 → [Cd(NH3)4]2+ + 2 OH−

3. Sodium hydroxide solution: white precipitate of cadmium(II) hydroxide:

Cd2+ + 2 OH− → Cd(OH)2 ↓

The precipitate is insoluble in excess reagent.Dilute acids dissolve the precipitate.

4. Potassium cyanide solution: white precipitate of cadmium cyanide, when addedslowly to the solution:

Cd2+ + 2 CN− → Cd(CN)2 ↓

An excess of the reagent dissolves the precipitate, when tetracyanocadminate(II) ionsare formed:

Cd(CN)2 ↓ + 2 CN− → [Cd(CN)4]2−

The colourless compound is not too stable; when hydrogen sulphide gas is introduced,cadmium sulphide is precipitated:

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[Cd(CN)4]2− + H2S → CdS ↓ + 2 H+ + 4 CN−

The marked difference in the stabilities of the copper and cadmium tetracyanocomplexes serves as the basis for the separation of copper and cadmium ions, and alsofor the identification of cadmium in the presence of copper.

5. Potassium iodide: forms no precipitate (difference from copper).

Summarise the reactions of Cu2+, Cd2+, and Bi3+ ions

Cu2+ Cd2+ Bi3+

H2Sin acidic solution

NaOH

NH3in excess

KI

Fe nail

Reactions of mercury(I) ions, Hg22+

Solubility of the most common mercury(I) compounds:Mercury(I) nitrate is soluble in water and tends to decompose. Other common

inorganic salts are very slightly soluble or insoluble.

E.g. at 25 °C: Compound Solubility ( g / 100 ml H2O)Hg2SO4 0,06Hg2CO3Hg2Br2

0,00000450,000004

To study these reactions use a 0.1 M solution of mercury(I) nitrate.1. Dilute hydrochloric acid or soluble chlorides: white precipitate of mercury(I)

chloride (calomel). Solubility product: Ksp(Hg2Cl2, 25 °C)= 1.45x10−18.

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Hg22+ + 2 Cl− → Hg2Cl2 ↓

The precipitate is insoluble in dilute acids.Ammonia solution converts the precipitate into a mixture of mercury(II)amidochloride and mercury metal, both insoluble precipitates; the mercury(II)amidochloride is a white precipitate, but the finely divided mercury makes it shinyblack (disproportionation takes place; mercury(I) is converted partly to mercury(II)and partly to mercury metal):

Hg2Cl2 ↓ + 2 NH3 → Hg ↓ + Hg(NH2)Cl ↓ + NH4+ + Cl−

This reaction can be used to differentiate mercury(I) ions from lead(II) and silver(I).

Mercury(I) chloride dissolves in aqua regia, forming undissociated but solublemercury(II) chloride:

3 Hg2Cl2 ↓ + 2 HNO3 + 6 HCl → 6 HgCl2 + 2 NO ↑ + 4 H2O

2. Hydrogen sulphide in neutral or dilute acid medium: black precipitate, which isa mixture of mercury(II) sulphide and mercury metal:

Hg22+ + H2S → Hg ↓ + HgS ↓ + 2 H+

Owing to the extremely low solubility product of mercury(II) sulphide (6.44x10−53)the reaction is very sensitive.Aqua regia dissolves the precipitate, yielding undissociated mercury(II) chloride andsulphur:

3 Hg ↓ + 3 HgS ↓ + 12 HCl + 4 HNO3 → 6 HgCl2 + 3 S ↓ + 4 NO ↑ + 8 H2O

When heated with aqua regia, sulphur is oxidised to sulphuric acid and the solutionbecomes clear:

S ↓ + 6 HCl + 2 HNO3 → SO42− + 6 Cl− + 8 H+ + 2 NO ↑

3. Ammonia solution: black precipitate which is a mixture of mercury metal andbasic mercury(II) amidonitrate, which itself is a white precipitate:

2 Hg22+ + NO3

− + 4 NH3 + H2O → 2 Hg ↓ + HgO.Hg(NH2)NO3 ↓ + 3 NH4+

This reaction can be used to differentiate between mercury(I) and mercury(II) ions.

4. Sodium hydroxide: black precipitate of mercury(I) oxide.

Hg22+ + 2 OH− → Hg2O ↓ + H2O

The precipitate is insoluble in excess reagent, but dissolves readily in dilute nitricacid.

When boiling, the colour of the precipitate turns to grey, owing todisproportionation, when mercury(II) oxide and mercury metal are formed:

Hg2O ↓ → HgO ↓ + Hg ↓

95

5. Potassium chromate in hot solution: red crystalline precipitate of mercury(I)chromate:

Hg22+ + CrO4

2− → Hg2CrO4 ↓

If the test is carried out in cold, a brown amorphous precipitate is formed with anundefined composition. When heated the precipitate turns to red, crystallinemercury(I) chromate.

Sodium hydroxide turns the precipitate into black mercury(I) oxide:

Hg2CrO4 ↓ + 2 OH− → Hg2O ↓ + CrO42− + H2O

6. Potassium iodide, added slowly in cold solution: green precipitate of mercury(I)iodide:

Hg22+ + 2 I− → Hg2I2 ↓

If excess reagent is added, a disproportionation reaction takes place, solubletetraiodomercurate(II) ions and a black precipitate of finely divided mercury beingformed:

Hg2I2 ↓ + 2 I− → [HgI4]2− + Hg ↓

When boiling the mercury(I) iodide precipitate with water, disproportionation takesplace again, and a mixture of red mercury(II) iodide precipitate and finely distributedblack mercury is formed:

Hg2I2 ↓ → HgI2 ↓ + Hg ↓

7. Sodium carbonate in cold solution: yellow precipitate of mercury(I) carbonate(solubility product: Ksp(Hg2CO3, 25 °C)= 3.67x10−17):

Hg22+ + CO3

2− → Hg2CO3 ↓

The precipitate turns slowly to blackish grey, when mercury(II) oxide and mercuryare formed:

Hg2CO3 ↓ → HgO ↓ + Hg ↓ + CO2 ↑

The decomposition can be speeded up by heating the mixture.

8. Disodium hydrogen phosphate: white precipitate of mercury(I) hydrogenphosphate:

Hg22+ + HPO4

2− → Hg2HPO4 ↓

9. Potassium cyanide solution: produces mercury(II) cyanide solution and mercuryprecipitate:

Hg22+ + 2 CN− → Hg ↓ + Hg(CN)2

Mercury(II) cyanide, though soluble, is practically undissociated.

96

10. Tin(II) chloride solution: reduces mercury(I) ions to mercury metal, whichappears in the form of a greyish-black precipitate:

Hg22+ + Sn2+ → 2 Hg ↓ + Sn4+

11. Copper sheet or copper coin: deposit of mercury metal is formed on the coppersurface.

Hg22+ + Cu → 2 Hg ↓ + Cu2+

Compare the reactions of mercury(I), silver(I), and lead(II) ions.

Hg22+ Ag+ Pb2+

HCl

H2S

NH3

NH3in excess

KI

KIin excess

NaOH

Na2CO3

K2CrO4

Cu

97

Reactions of mercury(II) ions, Hg2+

Solubility of the most common mercury(II) compounds:Mercury(II) nitrate is readily soluble in water; chloride, chlorate, cyanide, and

acetate are also soluble, but their solubility is much less than that of nitrate. All theother mercury(II) compounds (e.g. sulphide, carbonate, iodide) are practicallyinsoluble.

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

10 °C:Hg(ClO3)2Hg(Ac)2

2525

25 °C: HgI2 (α) 0,0118 °C: HgS (α) 0,000001

To study these reactions use a 0.1 M solution of mercury(II) nitrate.

1. Hydrogen sulphide gas: black precipitate of mercury(II) sulphide. Solubilityproduct constant: Ksp(HgS, 25 °C)= 6.44x10−53.

Hg2+ + H2S → HgS ↓ + 2 H+

In the presence of dilute hydrochloric acid, initially a white precipitate of mercury(II)chlorosulphide (Hg3S2Cl2 ↓), which decomposes when further amounts of hydrogensulphide are added and finally a black precipitate of mercury(II) sulphide is formed:

3 Hg2+ + 2 Cl− + 2 H2S → Hg3S2Cl2 ↓ + 4 H+

Hg3S2Cl2 ↓ + H2S → 3 HgS ↓ + 2 H+ + 2 Cl−

The HgS precipitate is insoluble in water, hot dilute nitric acid, alkali hydroxides, orammonium sulphide.Aqua regia dissolves the precipitate:

3 HgS ↓ + 6 HCl + 2 HNO3 → 3 HgCl2 + 3 S ↓ + 2 NO ↑ + 4 H2O

Sulphur remains as a white precipitate, which however dissolves readily if thesolution is heated, to form sulphuric acid.

Sodium sulphide (2M) dissolves the HgS precipitate when thedisulphomercurate(II) complex ion is formed:

HgS ↓ + S2− → [HgS2]2−

Adding ammonium chloride to the solution, mercury(II) sulphide precipitates again.

2. Ammonia solution: white precipitate with a mixed composition; essentially itconsists of mercury(II) oxide and mercury(II) amidonitrate:

2 Hg2+ + NO3− + 4 NH3 + H2O → HgO.Hg(NH2)NO3 ↓ + 3 NH4

+

98

3. Sodium hydroxide (added in small amounts): brownish-red precipitate withvarying composition; if added in stoichiometric amounts the precipitate turns toyellow when mercury(II) oxide is formed:

Hg2+ + 2 OH− → HgO ↓ + H2O

The precipitate is insoluble in excess sodium hydroxide. Acids dissolve the precipitatereadily.

4. Potassium iodide (added slowly to the solution): red precipitate of mercury(II)iodide, Ksp(HgI2, 25 °C)= 2.82x10−29:

Hg2+ + 2 I− → HgI2 ↓

The precipitate dissolves in excess reagent, when colourless tetraiodomercurate(II)ions are formed:

HgI2 ↓ + 2 I− → [HgI4]−

5. Tin(II) chloride: when added in moderate amounts, white, silky precipitate ofmercury(I) chloride (calomel), Ksp(Hg2Cl2, 25 °C)= 1.45x10−18, is formed:

2 Hg2+ + Sn2+ + 2 Cl− → Hg2Cl2 ↓ + Sn4+

If more reagent is added, mercury(I) chloride is further reduced and black precipitateof mercury is formed:

Hg2Cl2 ↓ + Sn2+ → 2 Hg ↓ + Sn4+ + 2 Cl−

6. Copper sheet or coin: reduces mercury(II) ions to the metal:

Cu + Hg2+ → Hg ↓ + Cu2+

(Standard reduction potentials: Cu/ Cu2+= +0.3419 V; Sn2+/ Sn4+= +0.151 V;Hg2

2+/ Hg2+= +0.920 V; Hg/ Hg2+= +0.851 V; Hg/ Hg22+= +0.7973 V)

Compare the characteristic reactions of mercury(I) and mercury(II).

HCl H2S NH3 NaOH KI Cu

Hg22+

Hg2+

99

Compare the characteristic reactions of arsenic(III)(arsenite), antimony(III),tin(II), tin(IV), copper(II), cadmium(II), bismuth(III), and mercury(II) ions.

As3+ Sb3+ Sn2+ Sn4+ Cu2+ Cd2+ Bi3+ Hg2+

HCl

H2S

precipitate

+(NH4)2SX

+HCl

NaOH

KI

KCN

NH3

SnCl2

Fe

100

Titanium (Group IVb) and Its Common Ions

Titanium, when pure, is a lustrous, white metal. It has a low density, good strength,is easily fabricated, and has excellent corrosion resistance (melting point: 1660 °C).The metal is not attacked by mineral acids at room temperature or even by hotaqueous alkali.

Solubility in acids The titanium metal is not soluble in mineral acids at room temperature, but

soluble in hot, concentrated hydrochloric acid and sulphuric acid, and in hydrogenfluoride:

2 Ti + 6 HCl → 2 Ti3+ + 6 Cl− + 3 H2 ↑

Ti + 4 H2SO4 → Ti4+ + 2 SO42− + 2 SO2 ↑ + 4 H2O

The best solvents of the metal are HF and acids to which fluoride ions have beenadded; such media dissolve titanium and hold it in solution as fluoro complexes.

Titanium is insoluble in hot, concentrated nitric acid, like tin, because of theformation of titanic acid (TiO2.xH2O) on the surface of the metal, which protects therest of the metal from the acid.

Common cations of titanium in aqueous solutionTi3+

violetTi4+

colourless* Titanium(II) ions are not stable in aqueous solution; they liberate hydrogen gas from water(Ti2+/Ti3+= −0.369 V; Ti2+/Ti(OH)2

2+= −0.135 V).** Titanium(III) ions are rather unstable and are readily oxidised to titanium(IV) in aqueoussolutions, e.g. by air oxygen if exposed to air (Ti3+/Ti(OH)2

2+= +0.099 V).

The white titanium(IV) oxide, TiO2, is used as a pigment and is far the mostimportant titanium oxide, which occurs also in the nature. The solubility of TiO2depends considerably on its chemical and thermal history; strongly roasted specimensare chemically inert.

Titanium(IV) ions exist only in strongly acid solutions; they tend to hydrolyse.In strong acid the Ti4+ (aquated) ions are in equilibrium with Ti(OH)2

2+, Ti(OH)3+ and

TiO2+ ions (aquated); the main species is Ti(OH)22+ and if the acidity of the solution is

lowered titanium(IV) hydroxide is precipitated.

101

Reactions of titanium(IV) ions, Ti4+

To study these reactions use a 0.1 M solution of titanium(IV) sulphate, whichis prepared by dissolving Ti(SO4)2 in 5 per cent sulphuric acid.

1. Solutions of sodium hydroxide, ammonia or ammonium sulphide solution:white gelatinous precipitate of titanium(IV) hydroxide, Ti(OH)4 (or orthotitanicacid, H4TiO4), in the cold; this is almost insoluble in excess reagent, but soluble inmineral acids:

Ti(OH)22+ + 2 OH− → Ti(OH)4 ↓

Ti(OH)4 ↓ + H2SO4 → Ti(OH)22+ + 2 H2O + SO4

2−

Ti(OH)4 ↓ + 3 HCl → Ti(OH)Cl2+ + 3 H2O + 2 Cl−

If precipitation takes place from hot solution, white TiO(OH)2 (or metatitanic acid,H2TiO3) is formed, which is sparingly soluble in dilute acids.

Ti(OH)22+ + 2 OH− → TiO(OH)2 ↓ + H2O

2. Water: a white precipitate of metatitanic acid is obtained on boiling a solution ofa titanic salt with excess water:

Ti(OH)22+ + 2 OH− → TiO(OH)2 ↓ + H2O

3. Sodium phosphate solution: white precipitate of titanium(IV) phosphate in dilutesulphuric acid solution:

Ti(OH)22+ + 2 H2PO4

− → Ti(HPO4)2 ↓ + 2 H2O

4. Zinc or tin metal: when any of these metals is added to an acid solution of atitanium(IV) salt, a violet coloration is produced, due to reduction to titanium(III)ions:

2 Ti4+ + Zn → 2 Ti3+ + Zn2+

5. Hydrogen peroxide. An intense orange coloration is produced (yellow with verydilute solution), due to formation of stable peroxo complexes:

Ti(OH)22+ + H2O2 + OH− → Ti(O2)OH+ + 2 H2O

Ti

OH

OH2

OH2

OH2H2O

HO

4+

--�������� �������� ���������

���������

��������������������

-4+

OH2

OH2

OH2

OH

TiO

O

����������

102

Vanadium (Group Vb) and Its Common Ions

Vanadium is a bright, white metal, and is soft and ductile (melting point: 1890 °C). Ithas good corrosion resistance to alkalis, sulphuric and hydrochloric acid, and saltwaters. The metal has good structural strength.

Solubility in acids The vanadium metal is not soluble in hydrochloric, nitric, or sulphuric acids

or in alkalis at room temperature due to passivation (thin protective oxide layerforms).

It dissolves readily in aqua regia, in hot nitric acid, hot and concentratedsulphuric acid, or in a mixture of concentrated nitric acid and hydrogen fluoride:

3 V + 4 HNO3 + 6 HCl → 3 VO2+ + 6 Cl− + 4 NO ↑ + 5 H2O3 V + 10 HNO3 → 3 VO2+ + 6 NO3

− + 4 NO ↑ + 5 H2OV + 3 H2SO4 → VO2+ + SO4

2− + 2 SO2 ↑ + 3 H2O

Common ions of vanadium in aqueous solutionsoxidation state cations anions

+2 V2+

+3 V3+

+4 VO2+

+5 VO2+ VO4

3− vanadate 13<pHHVO4

2− monovanadate 8<pH<13HV2O7

3− divanadateVO3

− metavanadateV10O28

6− decavanadate 2<pH<6

* The orange decavanadate ion can exist in several protonated form, and with increasingacidity of the solution rapidly gives the dioxovanadium(V) ion, VO2

+.

Vanadium(II) and vanadium(III) ions are instable in aqueous solution, and easilyoxidised to vanadium(IV), due to their small standard reduction potentials:

[VO2(H2O)4] [VO(H2O)5] [V(H2O)6] [V(H2O)6] V 2+3+2++���� ���� ���� ���� ���� �������� ����

+0.999 V +0.359 V -0.256 V -1.186 V

103

Vanadium(V) oxide, V2O5, is the most stable and is far the most important vanadiumoxide. It is an orange (or brick red) powder, which is insoluble in water, but solublein both mineral acids and alkalis.Vanadium(V) is moderately strong oxidising agent, thus if the oxide is dissolved inhydrochloric acid chlorine gas is evolved and vanadium(IV) is produced. The oxide isalso reduced by warm sulphuric acid.Vanadium pentoxide dissolves in sodium hydroxide to give colourless solutions andin the highly alkaline region, pH>13, the main ion is VO4

3−. As the basicity isreduced, a series of complicated reactions occurs with the formation of variousvanadates (mono, di, meta, deca, etc.).

Metavanadates, VO3−

To study these reactions use a 0.1 M solution of ammonium metavanadate,NH4VO3, or sodium metavanadate, NaVO3. The addition of some sulphuric acidkeeps these solutions stable.

1. Hydrogen sulphide. No precipitate is produced in acidic solution, but a bluesolution (due to the production of vanadium(IV) ions) is formed and sulphurseparates:

2 VO3− + H2S + 6 H+ → 2 VO2+ + S ↓ + 8 H2O

2. Zinc or aluminium acid solution. Zn and Al carry the reduction stillfurther than H2S. The solution turns atfirst blue (VO2+ ions),then green (V3+ ions) andfinally violet (V2+ ions).

VO2+/VO2+ = +0.999 V

V3+/ VO2+ = +0.359 V V2+/ V3+ = −0.256 V

Zn/ Zn2+ = −0.762 V

3. Ammonium sulphide solution: the solution is coloured claret-red, due to theformation of thiovanadates (VS4

3−).

VO3− + 4 S2− + 3 H2O → VS4

3− + 6 OH−

Upon acidification of the solution, brown vanadium sulphide, V2S5, is precipitated,and the filtrate usually has a blue colour:

2 VS43− + 6 H+ → V2S5 ↓ + 3 H2S ↑

The precipitate is soluble in solutions of alkalis, alkali carbonates, and sulphides.

104

4. Hydrogen peroxide: A red coloration is produced when a few drops of hydrogenperoxide solution are added dropwise to an acid (15-20 per cent sulphuric acid)solution of a vanadate; excess hydrogen peroxide should be avoided.

The red colour is due to the formation of themono- and diperoxovanadium(V) ions,VO(O2)+ and VO(O2)2

−:

VO3− + 2 H+ + H2O2 → VO(O2)+ + 2 H2O

VO3− + 2 H2O2 → VO(O2)2

− + 2 H2O

If the solution is made alkaline and morehydrogen peroxide is added, the colourchanges to yellow, due to the formation ofdiperoxoorthovanadate(V) ions:

VO(O2)+ + H2O2 + 4 OH− → VO2(O2)23− + 3

H2OVO(O2)2

− + 2 OH− ↔ VO2(O2)23− + H2O

The reaction is reversible; on acidification thesolution again turns red.

��������������������

������������ 5+OH2

OH2

VO

O

O

OV

OH2

O

5+�������������������������� ����������

����������

O

O

O

OH2

5. Lead acetate solution: yellow precipitate of lead vanadate, turning white or paleyellow on standing; the precipitate is insoluble in dilute acetic acid but soluble indilute nitric acid.

6. Barium chloride solution: yellow precipitate of barium vanadate; soluble in dilute hydrochloric acid.

7. Copper sulphate solution: green precipitate with metavanadates.

8. Iron(III) chloride:

VO3− + 4 H+ + Fe2+ ↔ VO2+ + Fe3+ + 2 H2O

The reaction proceeds from left to right in acid solution and in the reverse direction inalkaline solution. (ε°(Fe2+/Fe3+)= +0.771 V)

105

Chromium (Group VIb) and Its Common Ions

Chromium is a silver-white, lustrous, hard, and brittle metal that takes a high polish(melting point: 1857 °C). Chromium is extremely resistant to ordinary corrosiveagents, which accounts for its extensive use as an electroplated protective coating.

Solubility in acids The metal, if it is passivated (probably due to a thin protective oxide layer), is

not soluble in mineral acids, but the metal is rather active when not passivated.Redox potentials: ε°(Cr/Cr2+)= −0.913 V; ε°(Cr/Cr3+)= −0.744 V.

The chromium metal is soluble in dilute or concentrated hydrochloric acid.If air is excluded, chromium(II) ions are formed:

Cr + 2 HCl → Cr2+ + 2 Cl− + H2 ↑

In the presence of atmospheric oxygenchromium(II) gets wholly oxidised to thetervalent state:

4 Cr2+ + O2 + 4 H+ → 4 Cr3+ + 2 H2O1 2 3 4 5 6 7 8 9

-0.7

-0.6

-0.5

-0.4

-0.3

-0.2

-0.1

0.0

[Cr2+]= [Cr3+] = 0.05 M

Cr2+/ Cr3+

H2/ H+

Redox potential (V)

pH

Even if the solution is protected from air, chromium(II) ions decompose at ratesvarying with acidity, by reducing water with liberation of hydrogen (ε°(Cr2+/Cr3+)=−0.407 V).Chromium(II) ions are stable only in neutral and pure solutions at the exclusion of air.

Dilute sulphuric acid attacks chromium slowly, with the formation ofhydrogen.In hot, concentrated sulphuric acid chromium dissolves readily, when chromium(III)ions and sulphur dioxide are formed:

2 Cr + 6 H2SO4 → 2 Cr3+ + 3 SO42− + 3 SO2 ↑ + 6 H2O

Both dilute and concentrated nitric acid render chromium passive, as doescold, concentrated sulphuric acid and aqua regia.

Principal cations and anions of chromium in aqueous solution(Cr2+ chromous)*

Cr3+ chromicCrO4

2− chromateCr2O7

2− dichromate* chromium(II) ions are rather unstable, as they are strong reducing agents. Atmospheric oxygen oxidises them readily to chromium(III) ions.

106

The green chromium(III) oxide, Cr2O3, and its hydrous form, Cr2O3.nH2O, areamphoteric, dissolving readily in acids and in concentrated alkali, but if ignited toostrongly Cr2O3 becomes inert toward both acid and base.

Chromate (CrO42−) and dichromate (Cr2O7

2−) ionsIn basic solutions above pH 7, the yellow chromate ion CrO4

2− is the mainspecies; between pH 1 and 6, HCrO4

− and the orange-red dichromate ion Cr2O72− are

in equilibrium; and at pH<0 the main species are H2CrO4 and HCr2O7−.

The equilibria are the following:

H2CrO4 ↔ HCrO4− + H+

HCrO4− ↔ CrO4

2− + H+ HCr2O7− ↔ Cr2O7

2− + H+

2 HCrO4− ↔ H2O + Cr2O7

2−

0 2 4 6 8 10 12 14

0.00

0.01

0.02

0.03

0.04

0.05 co (K2CrO4)= 0.05 M

H2CrO4

HCr2O7-

HCrO4-

CrO42-

Cr2O72-

C (mol/ l)

pH

The chromates of the alkali metals and of magnesium and calcium are solublein water.Strontium chromate is sparingly soluble in water, and most other metallic chromatesare insoluble.

Sodium, potassium, and ammonium dichromates are well known and they aresoluble in water.

Reactions of chromium(III) ions, Cr3+

Chromium(III) sulphide, Cr2S3, like aluminium sulphide, can be prepared onlyin dry, because it hydrolyses readily with water to form chromium(III) hydroxide andhydrogen sulphide.

107

Hydrated chromium(III) sulphate, nitrate, chloride, bromide, iodide, andacetate are soluble in water.Chromium oxide, hydroxide, phosphate, and anhydrous halogenides (fluoride,chloride, bromide, iodide) are hardly soluble or not soluble in water.

E.g.: Compound Solubility ( g / 100 ml H2O)20 °C:25 °C:

Cr2(SO4)3.18H2OCrCl3.6H2O

12058,5

CrCl3 -----

To study these reactions use a 0.1 M solution of chromium(III) chloride CrCl3or chromium(III) sulphate Cr2(SO4)3.

1. Ammonia solution: grey-green to grey-blue gelatinous precipitate ofchromium(III) hydroxide, slightly soluble in excess of the reagent in the coldforming a violet or pink solution containing complex hexammine-chromate(III)ion; upon boiling the solution, chromium(III) hydroxide is precipitated. Hence forcomplete precipitation of chromium as the hydroxide, it is essential that thesolution be boiling and excess aqueous ammonia solution be avoided.

Cr3+ + 3 NH3 + 3 H2O → Cr(OH)3 ↓ + 3 NH4+

Cr(OH)3 ↓ + 6 NH3 → [Cr(NH3)6]3+ + 3 OH−

2. Sodium hydroxide solution: precipitate of chromium(III) hydroxide:

Cr3+ + 3 OH− → Cr(OH)3 ↓

In excess reagent the precipitate dissolves readily, when tetrahydroxochromate(III)ions are formed:

Cr(OH)3 ↓ + OH− ↔ [Cr(OH)4]−

The solution is green. On adding hydrogen peroxide to the alkaline solution, a yellowsolution is obtained, owing to the oxidation of chromium(III) to chromate:

2 [Cr(OH)4]− + 3 H2O2 + 2 OH− → 2 CrO42− + 8 H2O

After decomposing the excess of hydrogen peroxide by boiling, chromate ions may beidentified in the solution by one of their characteristic reactions.

3. Sodium carbonate solution: precipitate of chromium(III) hydroxide:

2 Cr3+ + 3 CO32− + 3 H2O → 2 Cr(OH)3 ↓ + 3 CO2 ↑

4. Ammonium sulphide solution: precipitate of chromium(III) hydroxide:

2 Cr3+ + 3 S2− + 6 H2O → 2 Cr(OH)3 ↓ + 3 H2S ↑

108

5. Chromate test.Chromium(III) ions can be oxidised to chromate, and than chromate ions can

be identified on the basis of their characteristic reactions.Oxidation of chromium(III): adding an excess of sodium hydroxide to a

solution of chromium(III) salt followed by a few ml of hydrogen peroxide.

2 [Cr(OH)4]− + 3 H2O2 + 2 OH− → 2 CrO42− + 8 H2O

The excess of H2O2 can be removed by boiling the mixture for a few minutes.

Identification of chromium after oxidation to chromate: a. Barium chloride test. After acidifying the solution with acetic acid and adding barium chloridesolution, a yellow precipitate of barium chromate is formed:

Ba2+ + CrO42− → BaCrO4 ↓

b. Chromium pentoxide test. Acidifying the solution with dilute sulphuric acid,adding 2-3 ml of ether or amyl alcohol to the mixture andfinally adding some hydrogen peroxide, a blue colorationis formed, which can be extracted into the organic phaseby gently shaking. Chromium pentoxide is formed duringthe reaction:

Cr O O

OOO

CrO42− + 2 H+ + 2 H2O2 → CrO5 + 3 H2O

In aqueous solution the blue colour fades rapidly, because chromiumpentoxide decomposes to chromium(III) and oxygen.

Reactions of chromate (CrO42−) and dichromate (Cr2O7

2−) ionsThe chromates of metal ions are usually coloured solids, yielding yellow

solutions when dissolved in water. In the presence of dilute mineral acids chromatesare partially converted into dichromates; the latter yield orange-red aqueous solutions.

2 CrO42− + 2 H+ ↔ Cr2O7

2− + H2O

To study the reactions of chromates and dichromates use a 0.1 M solution ofpotassium chromate and dichromate, respectively.

1. Barium chloride solution: pale-yellow precipitate of barium chromate, solubilityproduct constant Ksp(BaCrO4)= 1.17x10−10:

CrO42− + Ba2+ → BaCrO4 ↓

The precipitate is insoluble in water, sodium hydroxide, and acetic acid, but soluble inmineral acids.

109

Dichromate ions produce the same precipitate, but as strong acid is formed,precipitation is only partial:

Cr2O72− + 2 Ba2+ + H2O ↔ 2 BaCrO4 ↓ + 2 H+

2. Silver nitrate solution: brownish-red precipitate of silver chromate with asolution of a chromate, Ksp(Ag2CrO4)= 1.12x10−12:

CrO42− + 2 Ag+ → Ag2CrO4 ↓

The precipitate is soluble in dilute nitric acid and in ammonia solution, but isinsoluble in acetic acid. Hydrochloric acid converts the precipitate into silverchloride, Ksp(AgCl)= 1.77x10−10:

2 Ag2CrO4 ↓ + 2 H+ → 4 Ag+ + Cr2O72− + H2O

Ag2CrO4 ↓ + 4 NH3 → 2 [Ag(NH3)2]+ + CrO42−

Ag2CrO4 ↓ + 2 Cl− → 2 AgCl ↓ + CrO42−

A reddish-brown precipitate of silver dichromate is formed with aconcentrated solution of a dichromate; this passes into the less soluble silver chromateon boiling with water:

Cr2O72− + 2 Ag+ → Ag2Cr2O7 ↓

Ag2Cr2O7 ↓ + H2O → Ag2CrO4 ↓ + CrO42− + 2 H+

3. Lead acetate solution: yellow precipitate of lead chromate, Ksp(PbCrO4)=1.77x10−14:

CrO42− + Pb2+ → PbCrO4 ↓

The precipitate is insoluble in acetic acid, but soluble in dilute nitric acid and sodiumhydroxide solution:

2 PbCrO4 ↓ + 2 H+ ↔ 2 Pb2+ + Cr2O72− + H2O

PbCrO4 ↓ + 4 OH− ↔ [Pb(OH)4]2− + CrO42−

4. Hydrogen peroxide. (chromium pentoxide test; see above)

5. Hydrogen sulphide: an acid solution of a chromate is reduced by this reagent to agreen solution of chromium(III) ions:

2 CrO42− + 3 H2S + 10 H+ → 2 Cr3+ + 3 S ↓ + 8 H2O

6. Potassium iodide solution: chromate is reduced into chromium(III) in thepresence of dilute mineral acids. Iodine formed in the reaction can be extractedwith e.g. CCl4, yielding a violet organic phase.

2 CrO42− + 6 I− + 16 H+ → 2 Cr3+ + 3 I2 + 8 H2O

110

7. Iron(II) sulphate: reduces chromates or dichromates in the presence of mineralacid smoothly:

CrO42− + 3 Fe2+ + 8 H+ → Cr3+ + 3 Fe3+ + 4 H2O

8. Concentrated hydrochloric acid: on heating a solid chromate or dichromate withconcentrated hydrochloric acid, chlorine is evolved, and a solution containingchromium(III) ions is produced:

2 K2CrO4 + 16 HCl → 2 Cr3+ + 3 Cl2 ↑ + 4 K+ + 10 Cl− + 8 H2OK2Cr2O7 + 14 HCl → 2 Cr3+ + 3 Cl2 ↑ + 2 K+ + 8 Cl− + 7 H2O

9. Concentrated sulphuric acid and a chloride: (see chromyl chloride test)

O

OCr

OH

OH

+ HCl

+ HCl

��������

��������

+ HOH

+ HOH

Cl

ClCr

O

O

Summarise the reactions of chromates.

Ag+ Ba2+ Pb2+

CrO42−

solubility inacetic acid

HNO3 soln.

NaOH soln.

NH3 soln.

111

Manganese (Group VIIb) and Its Common Ions

Manganese is grey-white, resembling iron, but is harder and very brittle (meltingpoint: 1244 °C). Manganese is roughly similar to Fe in its physical and chemicalproperties, the chief difference being that it is harder and more brittle. The metal isreactive, and slowly reacts even with cold water.

Solubility in water and acids Due to its highly negative electrode potential (ε°(Mn/Mn(OH)2)= −1.56 V),

manganese reacts with water (the reaction is slow in the cold, but fast if the water iswarm) forming manganese(II) hydroxide and hydrogen (there is no protective oxidelayer, like in the case of chromium):

Mn + 2 H2O → Mn(OH)2 ↓ + H2 ↑

Dilute mineral acids and also acetic acid dissolve the metal with theproduction of manganese(II) salts and hydrogen (ε°(Mn/Mn2+)= −1.185 V):

Mn + 2 H+ → Mn2+ + H2 ↑

With hot, concentrated sulphuric acid, sulphur dioxide is evolved:

Mn + H2SO4 → Mn2+ + SO42− + SO2 ↑ + 2 H2O

Principal ions of manganese in aqueous solutionsoxidation state cations anions

+2 Mn2+

+3 (Mn3+)*+4 (Mn4+)* (MnO4

4− or MnO32−)*

+5 (MnO43−)**

+6 MnO42− ***

+7 MnO4−

* manganese(III) and manganese(IV) cations, and manganate(IV) anion are unstablein aqueous solutions, they are easily reduced to manganese(II).** unstable in aqueous solution, disproportionates to Mn(VII) and Mn(IV).*** stable in alkaline solutions, but upon neutralisation a disproportination reactiontakes place:

3 MnO42− + 2 H2O → MnO2 ↓ + 2 MnO4

− + 4 OH−

112

Permanganate ions, MnO4−, can be reduced step by step, e.g. with perborate

solution, to study the colour of the different oxidation state of manganese:

4+6+7+����� ����� ����� �����

����������

����������

+0.564 V +0.27 VMnO4 MnO4 MnO4 MnO4

2 3 45+

Alkali permanganates, MeMnO4 (Me= metal ion), are stable compounds,producing violet coloured solutions. They are all strong oxidising agents, and aresoluble in water.

Manganese(II) forms an extensive series of salts with all common anions.Most are soluble in water, although the phosphate and carbonate are only slightly so.

Five oxides of manganese are known so far:MnO Mn2O3 MnO2 Mn2O7 Mn3O4green brown black reddish oil reddish-brown

MnO is a grey-green to dark green powder, insoluble in water, but soluble inmineral acids.

Mn2O3 is insoluble in water; it produces Mn(II) ions if treated with mineralacids. If hydrochloric acid or sulphuric acid is used, chlorine and oxygen are evolved,respectively:

Mn2O3 ↓ + 6 HCl → 2 Mn2+ + Cl2 ↑ + 4 Cl− + 3 H2O2 Mn2O3 ↓ + 4 H2SO4 → 4 Mn2+ + O2 ↑ + 4 SO4

2− + 4 H2O

MnO2 is inert to most acids except when heated, but it does not dissolve togive Mn(IV) in solution; instead it functions as an oxidising agent, the exact mannerof this depending on the acid. With concentrated hydrochloric or sulphuric acid,chlorine and oxygen gas are evolved, respectively, and manganese(II) ions areproduced:

MnO2 ↓ + 4 HCl → Mn2+ + Cl2 ↑ + 2 Cl− + 2 H2O2 MnO2 ↓ + 2 H2SO4 → 2 Mn2+ + O2 ↑ + 2 SO4

2− + 2 H2O

Mn2O7 is an explosive oil, which can be extracted into CCl4 in which it isreasonably stable and safe. It solidifies at 5 to 9 °C to red crystals.

Mn3O4 is a spinel, MnIIMn2IIIO4 (MnO.Mn2O3). It is insoluble in water, but

soluble in mineral acids.

113

Reactions of manganese(II) ions, Mn2+

Manganese(II) sulphate, nitrate, chloride, bromide, and iodide are soluble inwater. Other common inorganic manganese(II) compounds (e.g. phosphate andcarbonate) are hardly soluble or insoluble in water.

E.g.: Compound Solubility ( g / 100 ml H2O)0 °C:

25 °C:MnBr2MnCl2

127,372,3

40 °C:18 °C:

MnF2Mn(OH)2

0,660,0002

To study these reactions use a 0.1 M solution of manganese(II) chloride ormanganese(II) sulphate.

1. Sodium hydroxide solution: an initially white precipitate of manganese(II)hydroxide; solubility product constant, Ksp(Mn(OH)2, 25 °C)= 2.06x10−13:

Mn2+ + 2 OH− → Mn(OH)2 ↓

The precipitate is insoluble in excess reagent, but soluble in dilute acids.The precipitate rapidly oxidises on exposure to air, becoming brown, when hydratedmanganese dioxide, MnO2.yH2O, is formed (ε°(Mn(OH)2/MnO2)= −0.05 V):

2 Mn(OH)2 ↓ + O2 → 2 MnO2.H2O ↓

Hydrogen peroxide converts manganese(II) hydroxide rapidly into hydratedmanganese dioxide:

Mn(OH)2 ↓ + H2O2 → MnO2.H2O ↓ + H2O

2. Ammonia solution: partial precipitation of (initially) white manganese(II)hydroxide:

Mn2+ + 2 NH3 + 2 H2O ↔ Mn(OH)2 ↓ + 2 NH4+

The precipitate is soluble in ammonium salts, when the reaction proceeds towards theleft.

3. Ammonium sulphide solution: pink precipitate of manganese(II) sulphide, solubility product constant, Ksp(MnS, 25 °C)= 4.65x10−14:

Mn2+ + S2− → MnS ↓

The precipitate is readily soluble in mineral acids and even in acetic acid.

MnS ↓ + 2 H+ → Mn2+ + H2S ↑

114

4. Disodium hydrogen phosphate solution: in the presence of ammonia (orammonium ions), pink precipitate of manganese ammonium phosphate:

Mn2+ + NH3 + HPO42− → Mn(NH4)PO4 ↓

If ammonium salts are absent, manganese(II) phosphate is formed:

3 Mn2+ + 2 HPO42− → Mn3(PO4)2 ↓ + 2 H+

The precipitates are soluble in acids.

5. Sodium carbonate solution: rose precipitate of manganese(II) carbonate, solubility product constant, Ksp(MnCO3, 25 °C)= 2.24x10−11:

Mn2+ + CO32− → MnCO3 ↓

The precipitate is soluble in dilute mineral acids and even in acetic acid.

6. Lead dioxide and concentrated nitric acid. On boiling a dilute solution ofmanganese(II) ions, free from hydrochloric acid and chlorides, with lead dioxideand a little concentrated nitric acid, and then diluting somewhat and allowing thesuspended solid containing unattacked lead dioxide to settle, the liquid acquires aviolet-red (or purple) colour due to permanganic acid formed.

5 PbO2 + 2 Mn2+ + 4 H+ → 2 MnO4− + 5 Pb2+ + 2 H2O

Permanganates, MnO4−

To study these reactions use a 0.01 M solution of potassium permanganate,KMnO4.

1. Hydrogen peroxide. The addition of this reagent to a solution of potassiumpermanganate, acidified with dilute sulphuric acid, results in decolourisation andthe formation of pure but moist oxygen: ε°(Mn2+/MnO4

−)= +1.507 V ε°(H2O2/O2)= +0.695 V

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

2. Concentrated hydrochloric acid. All permanganates on boiling with concentratedhydrochloric acid evolve chlorine. ε°(Cl−/Cl2)= +1.358 V

2 MnO4− + 16 HCl → 5 Cl2 ↑ + 2 Mn2+ + 6 Cl− + 8 H2O

115

3. Hydrogen sulphide: in the presence of dilute sulphuric acid the solutiondecolourizes and sulphur is precipitated: ε°(H2S/S)= +0.142 V

2 MnO4− + 5 H2S + 6 H+ → 5 S ↓ + 2 Mn2+ + 8 H2O

4. Iron(II) sulphate solution: in the presence of sulphuric acid permanganate isreduced to manganese(II). ε°(Fe2+/Fe3+)= +0.771 V

MnO4− + 5 Fe2+ + 8 H+ → 5 Fe3+ + Mn2+ + 4 H2O

The solution becomes yellow because of the formation of iron(III) ions. The yellowcolour disappears if potassium fluoride is added; it forms colourless complex withiron(III).

5. Potassium iodide solution: reduces permanganate with the formation of iodine,in the presence of sulphuric acid.

2 MnO4− + 10 I− + 16 H+ → 5 I2 + 2 Mn2+ + 8 H2O

In alkaline solution the permanganate is decolourized, but manganese dioxide isprecipitated. In the presence of sodium hydroxide solution, potassium iodide isconverted into potassium iodate.

2 MnO4− + I− + H2O → 2 MnO2 ↓ + IO3

− + 2 OH−

6. Sodium hydroxide solution. Upon warming a concentrated solution of potassiumpermanganate with concentrated sodium hydroxide solution, a green solution ofpotassium manganate is produced and oxygen is evolved.

ε°(MnO42−/MnO4

−)= +0.564 V ε°(OH−/O2)= +0.401 V

4 MnO4− + 4 OH− → 4 MnO4

2− + O2 ↑ + 2 H2O

When the manganate solution is poured into a large volume of water or is acidifiedwith dilute sulphuric acid, the purple colour of the potassium permanganate isrestored, and manganese dioxide is precipitated. ε°(MnO4

2−/MnO2)=+0.60 V

3 MnO42− + 2 H2O → 2 MnO4

− + MnO2 ↓ + 4 OH−

116

Summarise the reactions of CrO42−, Cr2O7

2−, and MnO4− ions.

CrO42− Cr2O7

2− MnO4−

colour(alkali metal salt)

H2Sacid solution

H2O2acid solution

KIacid solution

FeSO4acid solution

cc HCl

117

The Group VIIIb Elements (Fe, Co, Ni) andTheir Principle Ions

Iron is a relatively abundant element in the universe. Its nuclei are verystable. The metal is the fourth most abundant element on earth, by weight, making upthe crust of the earth. The use of iron is prehistoric. Iron is a vital constituent of plantand animal life, and appears in haemoglobin.The chemically pure iron is a silver-white, tenacious, and ductile metal. It melts at1535 °C. The pure metal is very reactive chemically, and rapidly corrodes, especiallyin moist air or at elevated temperatures. Iron can be magnetised. It has four allotropicforms, from which only the α-form is magnetic.The commercial iron is rarely pure and usually contains small quantities of carbide,silicide, phosphide, and sulphide of iron, and some graphite. These contaminants playan important role in the strength of iron structures. Other additives such as nickel,chromium, vanadium, etc. are also used to produce alloy steels. Iron is the cheapestand most abundant, useful, and important of all metals.

Cobalt is a steel-grey, slightly magnetic, brittle, and hard metal, closelyresembling iron and nickel in appearance. Melting point is 1495 °C.

Nickel is silvery white and takes on a high polish. It is hard, malleable,ductile, very tenacious, somewhat ferromagnetic, and a fair conductor of heat andelectricity. It melts at 1453 °C. It is quite resistant to attack by air or water at ordinarytemperatures when compact.

Solubility in acids Dilute or concentrated hydrochloric acid and dilute sulphuric acid dissolve

iron, cobalt, and nickel, when iron(II), cobalt(II), and nickel(II) salts and hydrogengas are produced.

Fe + 2 H+ → Fe2+ + H2 ↑ ε°(Fe/Fe2+)= −0.447 VCo + 2 H+ → Co2+ + H2 ↑ ε°(Co/Co2+)= −0.28 VNi + 2 H+ → Ni2+ + H2 ↑ ε°(Ni/Ni2+)= −0.257 V

Hot, concentrated sulphuric acid yields iron(III), cobalt(II), and nickel(II) ionsand sulphur dioxide:

2 Fe + 3 H2SO4 + 6 H+ → 2 Fe3+ + 3 SO2 ↑ + 6 H2OCo + H2SO4 + 2 H+ → Co2+ + SO2 ↑ + 2 H2ONi + H2SO4 + 2 H+ → Ni2+ + SO2 ↑ + 2 H2O

Cold, concentrated nitric acid and sulphuric acid renders iron passive.

118

Cold, dilute nitric acid, yields iron(II) and ammonium ions:

4 Fe + 10 H+ + NO3− → 4 Fe2+ + NH4

+ + 3 H2O

Medium concentrated nitric acid, or hot, concentrated nitric acid dissolves ironwith the formation of nitrogen oxide gas and iron(III) ions:

Fe + HNO3 + 3 H+ → Fe3+ + NO ↑ + 2 H2O

Dilute nitric acid dissolve cobalt and nickel readily in cold:

3 Co + 2 HNO3 + 6 H+ → 3 Co2+ + 2 NO ↑ + 4 H2O3 Ni + 2 HNO3 + 6 H+ → 3 Ni2+ + 2 NO ↑ + 4 H2O

Like iron, cobalt and nickel does not dissolve in concentrated nitric acidbecause it is rendered passive by this reagent.

Principal cations of iron, cobalt, and nickelFe2+

Fe3+Co2+

(Co3+)*Ni2+

* Cobalt(III) ions are unstable in water, but their complexes are stable both in solution and indry form.

Oxides of iron, cobalt, and nickel FeO* black Fe2O3 red Fe3O4 red-brown

CoO olive-green Co2O3 brown-black Co3O4 black

NiO green Ni2O3 ** Ni3O4 ***

* FeO is not stable under 560 °C; disproportionates to Fe and Fe3O4. It is also easilyoxidised by air oxygen.** There is no good evidence for Ni2O3. The black NiO(OH), however, is well known.*** Ni3O4 is not known, only a NiII-NiIII hydroxide of stoichiometry Ni3O2(OH)4.

Iron(III) oxide, Fe2O3, is soluble in dilute acids at room temperature, but if heated toostrongly it is almost insoluble even in hot concentrated hydrochloric acid.Iron(II,III) oxide, Fe3O4=FeIIFe2

IIIO4, is very resistant to attack by acids and alkalis.

Cobalt(II) oxide, CoO, and cobalt(III) oxide, Co2O3, are soluble in mineral acids.Cobalt(II,III) oxide, Co3O4=CoIICo2

IIIO4, is very slightly soluble in mineral acids.

Nickel(II) oxide, NiO, dissolves readily in acids.

119

Reactions of iron(II) ions, Fe2+ Solubility of the most common iron(II) compoundsThe chloride, bromide, iodide, chlorate, perchlorate, nitrate, sulphate, and

acetate are soluble in water.The fluoride is very slightly soluble, and oxide, carbonate, sulphide, and phosphateare practically insoluble.

E.g.: Compound Solubility ( g / 100 ml H2O)10 °C:10 °C:

FeCl2FeBr2

64,4109

18 °C:25 °C:

Fe(OH)2FeCO3

0,000150,0067

To study these reactions use a 0.1 M solution of iron(II) sulphate or iron(II)ammonium sulphate (Mohr’s salt).

1. Sodium hydroxide solution: white precipitate of iron(II) hydroxide in thecomplete absence of air, solubility product constant Ksp(Fe(OH)2, 25°C)=4.87x10−17, insoluble in excess of the reagent, but soluble in acids.

Upon exposure to air, iron(II) hydroxide is rapidly oxidised, yielding ultimatelyreddish-brown iron(III) hydroxide, Ksp(Fe(OH)3, 25°C)= 2.64x10−39 .Under ordinary conditions it appears as a dirty-green precipitate; the addition ofhydrogen peroxide immediately oxidises it to iron(III) hydroxide.

Fe2+ + 2 OH− → Fe(OH)2 ↓4 Fe(OH)2 ↓ + 2 H2O + O2 → 4 Fe(OH)3 ↓2 Fe(OH)2 ↓ + H2O2 → 2 Fe(OH)3 ↓

2. Ammonia solution: precipitation of iron(II) hydroxide occurs.

Fe2+ + 2 NH3 + 2 H2O → Fe(OH)2 ↓ + 2 NH4+

If, however, larger amounts of ammonium ions are present, precipitation does notoccur.

3. Hydrogen sulphide: no precipitation takes place in acid solution since thesulphide ion concentration is insufficient to exceed the solubility product ofiron(II) sulphide.

4. Ammonium sulphide solution: black precipitate of iron(II) sulphide FeS, Ksp(FeS, 25°C)= 1.59x10−19:

Fe2+ + S2− → FeS ↓

FeS is readily soluble in acids with evolution of hydrogen sulphide.The moist precipitate becomes brown upon exposure to air, due to its oxidation tobasic iron(III) sulphate Fe2O(SO4)2.

120

FeS ↓ + 2 H+ → Fe2+ + H2S ↑4 FeS ↓ + 9 O2 → 2 Fe2O(SO4)2 ↓

5. Potassium hexacyanoferrate(II) solution: in the complete absence of air a whiteprecipitate of potassium iron(II) hexacianoferrate(II) is formed:

Fe2+ + 2 K+ + [Fe(CN)6]4− → K2Fe[Fe(CN)6] ↓

under ordinary atmospheric conditions a pale-blue precipitate is obtained.

6. Potassium hexacyanoferrate(III) solution: a dark-blue precipitate is obtained.First hexacyanoferrate(II) and iron(III) is formed from hexacyanoferrate(III) andiron(II) ions according to the following equilibrium,

Fe2+ + [FeIII(CN)6]3− ↔ Fe3+ + [FeII(CN)6]4−

and than these ions combine to a precipitate called Turnbull’s blue:

4 Fe3+ + 3 [Fe(CN)6]4− → Fe4[Fe(CN)6]3 ↓

The structure of Turnbull's blue isbased on a three-dimensional cubicframework with FeII and FeIII atoms atthe corners of a cube and with FeII−N-C−FeIII links.

Note, that the composition of thisprecipitate is identical to that ofPrussian blue (see below at the reactionof iron(III) ions).

III

IIFe IIIFe IIFe

FeII IIIFe

Fe IIFeIIIFeII

Fe

FeII

Fe

III

Fe

II

Fe

Fe

IIIII

Fe

FeFeIIIFe II

III

Fe

IIFe IIIFe IIFe

Fe IIIII Fe

Fe IIFeIIIFeII

III

III

CN-CN-

CN-

CN-

CN-

CN-

CN-

CN-CN-

CN-

CN-CN-

CN-

CN-CN-

CN- CN-

CN-

CN-

CN-

CN-CN-

CN-

CN-

CN-

CN- CN-

CN-

CN-CN-

CN-

CN-CN-

CN-

CN- CN-CN-

CN-CN-

CN-

CN-

CN-

CN-

CN-

CN- CN-

CN-

CN-

CN-

CN-

CN-CN-

CN-

7. Ammonium thiocyanate solution: no coloration is obtained with pure iron(II)salts (distinction from iron(III) ions).

8. Dimethylglyoxime reagent: soluble red iron(II) dimethylglyoxime in ammoniacalsolution. Iron(III) salts give no coloration, but nickel, cobalt, and large quantities ofcopper salts interfere and must be absent.

121

Reactions of iron(III) ions, Fe3+ Solubility of the most common iron(III) compoundsThe chloride, bromide, iodide, chlorate, perchlorate, nitrate and sulphate are

soluble in water.The fluoride is very slightly soluble, and oxide, carbonate, sulphide, and phosphateare practically insoluble.

E.g.: Compound Solubility ( g / 100 ml H2O)0 °C:0 °C:

FeCl3Fe(NO3)3.6H2O

74,4150

Fe2O3 -----

To study these reactions use a 0.1 M solution of iron(III) chloride FeCl3.

1. Sodium hydroxide solution: reddish-brown, gelatinous precipitate of iron(III)hydroxide, Ksp(Fe(OH)3, 25°C)= 2.64x10−39, insoluble in excess of the reagent,but soluble in acids.

Fe3+ + 3 OH− → Fe(OH)3 ↓

Iron(III) hydroxide can be converted on strong heating to iron(III) oxide; the heatedoxide is soluble with difficulty in dilute acids, but dissolves on vigorous boiling withconcentrated hydrochloric acid.

2 Fe(OH)3 ↓ → Fe2O3 + 3 H2OFe2O3 + 6 H+ → 2 Fe3+ + 3 H2O

2. Ammonia solution: reddish-brown, gelatinous precipitate of iron(III) hydroxide,insoluble in excess of the reagent, but soluble in acids.

Fe3+ + 3 NH3 + 3 H2O → Fe(OH)3 ↓ + 3 NH4+

The solubility product of iron(III) hydroxide is so small (2.64x10−39) that completeprecipitation takes place even in the presence of ammonium salts.

3. Hydrogen sulphide: in acidic solution reduces iron(III) ions to iron(II) andsulphur is formed as a milky-white precipitate:

ε°(Fe2+/Fe3+)= +0.771 Vε°(H2S/ S)= +0.142 V

2 Fe3+ + H2S → 2 Fe2+ + 2 H+ + S ↓

The finely distributed sulphur cannot be readily filtered with ordinary filter papers.By boiling the solution with a few torn pieces of filter paper the precipitate coagulatesand can be filtered.

4. Ammonium sulphide solution: black precipitate, consisting of iron(II) sulphideand sulphur is formed:

2 Fe3+ + 3 S2− → 2 FeS ↓ + S ↓

122

The black iron(II) sulphide precipitate dissolves in hydrochloric acid and the whitecolour of sulphur becomes visible:

FeS ↓ + 2 H+ → Fe2+ + H2S ↑

From alkaline solutions black iron(III) sulphide is obtained:

2 Fe3+ + 3 S2− → Fe2S3 ↓

On acidification with hydrochloric acid iron(III) ions are reduced to iron(II) andsulphur is formed:

Fe2S3 ↓ + 4 H+ → 2 Fe2+ + 2 H2S ↑ + S ↓

6. Potassium hexacyanoferrate(II) solution: intense blue precipitate of iron(III)hexacyanoferrate(II) (Prussian blue):

4 Fe3+ + 3 [Fe(CN)6]4− → Fe4[Fe(CN)6]3

The precipitate is insoluble in dilute acids, but decomposes with concentratedhydrochloric acid. A large excess of the reagent dissolves it partly or entirely, whenan intense blue solution is obtained.

7. Potassium hexacyanoferrate(III) solution: a brown coloration is produced, dueto the formation of an undissociated complex, iron(III) hexacyanoferrate(III):

Fe3+ + [Fe(CN)6]3− → Fe[Fe(CN)6]

Upon adding some tin(II) chloride solution, the hexacyanoferrate(III) part of thecompound is reduced and Prussian blue is precipitated.

8. Ammonium thiocyanate solution: in slightly acidic solution a deep-redcolouration is produced (difference from iron(II) ions), due to the formation of anon-dissociated iron(III) thiocyanate:

Fe3+ + 3 SCN− → Fe(SCN)3

This molecule can be extracted by ether or amyl alcohol.Fluorides and phosphates bleach the colour because of the formation of the morestable hexafluoro and triphosphato complexes:

Fe(SCN)3 + 6 F− → [FeF6]3− + 3 SCN−

Fe(SCN)3 + 3 PO43− → [Fe(PO4)3]6− + 3 SCN−

Upon addition of SnCl2 solution in excess, the red colour disappears due to thereduction of iron(III) to iron(II):

2 Fe(SCN)3 + Sn2+ → 2 Fe2+ + Sn4+ + 3 SCN−

123

9. Disodium hydrogen phosphate solution: a yellowish-white precipitate ofiron(III) phosphate is formed, Ksp(FePO4.2H2O, 25°C)= 9.92x10−29:

Fe3+ + HPO42− → FePO4 ↓ + H+

Summarise the reactions of iron cations with hexacyanoferrates.

[Fe(CN)6]4− [Fe(CN)6]3−

Fe2+

Fe3+

Reactions of cobalt(II) ions, Co2+ Solubility of the most common cobalt(II) compoundsThe chloride, bromide, iodide, chlorate, perchlorate, nitrate, sulphate, and

acetate are soluble in water.The fluoride is very slightly soluble, and oxide, hidroxide, carbonate, sulphide, andphosphate are practically insoluble.

E.g.: Compound Solubility ( g / 100 ml H2O)7 °C:

20 °C:CoCl2CoSO4

4536,2

25 °C: CoF2 1,5Co(OH)2 0,00032

To study these reactions use a 0.1 M solution of cobalt(II) chloride orcobalt(II) nitrate.

1. Sodium hydroxyde solution: in cold a blue basic salt is precipitated:

Co2+ + OH− + NO3− → Co(OH)NO3 ↓

Upon warming with excess alkali (or sometimes merely upon addition of excessreagent) the basic salt is converted into pink cobalt(II) hydroxide precipitateKsp(Co(OH)2, 25°C)= 1.09x10−15:

124

Co(OH)NO3 ↓ + OH− → Co(OH)2 ↓ + NO3−

The hydroxide is slowly transformed into the brownish black cobalt(III) hydroxide onexposure to the air:

4 Co(OH)2 ↓ + O2 +2 H2O → 4 Co(OH)3 ↓

Cobalt(II) hydroxide precipitate is readily soluble in ammonia or concentratedsolutions of ammonium salts.

2. Ammonia solution: in the absence of ammonium salts small amounts ofammonia precipitate the basic salt:

Co2+ + NH3 + H2O + NO3− → Co(OH)NO3 ↓ + NH4

+

The excess of the reagent dissolves the precipitate, when hexamminocobaltate(II) ionsare formed:

Co(OH)NO3 ↓ + 6 NH3 → [Co(NH3)6]2+ + NO3− + OH−

The precipitation of the basic salt does not take place at all if larger amounts ofammonium ions are present, but the complex is formed in one step.

3. Ammonium sulphide solution: black precipitate of cobalt(II) sulphide fromneutral or alkaline solution:

Co2+ + S2− → CoS ↓

The precipitate is insoluble in hydrochloric or acetic acids.Hot, concentrated nitric acid or aqua regia dissolve the precipitate, when whitesulphur remaines:

3 CoS ↓ + 2 HNO3 + 6 H+ → 3 Co2+ + 3 S ↓ + 2 NO ↑ + 4 H2OCoS ↓ + HNO3 + 3 HCl → Co2+ + S ↓ + NOCl ↑ + 2 Cl− + 2 H2O

On longer heating the mixture becomes clear because sulphur gets oxidised tosulphate.

The CoS precipitate dissolves also in the 1+1 mixture of concentrated aceticacid and 30% hydrogen peroxide:

CoS ↓ + 4 H2O2 → Co2+ + SO42− + 4 H2O

4. Potassium nitrite solution: yellow precipitate of potassium hexanitrito-cobaltate(III), K3[Co(NO2)6].3H2O:

Co2+ + 7 NO2− + 2 H+ + 3 K+ → K3[Co(NO2)6] ↓ + NO ↑ + H2O

125

The test can be carried out most conviniently as follows: to a neutral solution ofcobalt(II) add acetic acid, than a freshly prepared saturated solution of potassiumnitrite.

5. Ammonium thiocyanate test: adding a few crystals of ammonium thiocyanate toa neutral or acidic solution of cobalt(II) a blue colour appears owing to theformation of tetrathiocyanatocobaltate(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.

Reactions of nickel(II) ions, Ni2+ Solubility of the most common nickel compoundsThe chloride, bromide, iodide, chlorate, perchlorate, nitrate, sulphate, and

acetate are soluble in water.The fluoride is very slightly soluble, and oxide, carbonate, sulphide, and phosphateare practically insoluble.

E.g.: Compound Solubility ( g / 100 ml H2O)20 °C:0 °C:

NiCl2NiI2

64,2124,2

25 °C: NiF2 425 °C: NiCO3 0,0093

To study these reactions use a 0.1 M solution of nickel(II) sulphate ornickel(II) chloride.

1. Sodium hydroxyde solution: green precipitate of nickel(II) hydroxide, solubilityproduct constant, Ksp(Ni(OH)2, 25 °C)= 5.47x10−16:

Ni2+ + 2 OH− → Ni(OH)2 ↓

The precipitate is insoluble in excess reagent.Ammonia solution dissolves the precipitate; in the presence of excess alkalihydroxide ammonium salts also dissolve the precipitate:

126

Ni(OH)2 ↓ + 6 NH3 → [Ni(NH3)6]2+ + 2 OH−

Ni(OH)2 ↓ + 6 NH4+ + 4 OH− → [Ni(NH3)6]2+ + 6 H2O

The green nickel(II) hydroxide precipitate can be oxidised to black nickel(III)hydroxide with sodium hypochlorite solution:

2 Ni(OH)2 ↓ + ClO− + H2O → 2 Ni(OH)3 ↓ + Cl−

Hydrogen peroxide solution, however, does not oxidise nickel(II) hydroxide, but theprecipitate catalyses the decomposition of hydrogen peroxide to oxygen and waterwithout any other visible change.

2. Ammonia solution: green precipitate of nickel(II) hydroxide:

Ni2+ + 2 NH3 + 2 H2O → Ni(OH)2 ↓ + 2 NH4+

which dissolves in excess reagent:

Ni(OH)2 ↓ + 6 NH3 → [Ni(NH3)6]2+ + 2 OH−

the solution turns deep blue. If ammonium salts are present, no precipitation occurs,but the complex is formed immediately.

3. Ammonium sulphide solution: black precipitate of nickel sulphide, fromneutral or slightly alkaline solutions, Ksp(NiS, 25 °C)= 1.07x10−21:

Ni2+ + S2− → NiS ↓

If the reagent is added in excess, a dark-brown colloidal solution is formed which runsthrough the filter paper. If the colloidal solution is boiled, the colloidal solution(hydrosol) is coagulated and can than be filtered.The precipitate is insoluble in hydrochloric or acetic acids.Hot, concentrated nitric acid or aqua regia dissolve the precipitate with the separationof white sulphur:

3 NiS ↓ + 2 HNO3 + 6 H+ → 3 Ni2+ + 3 S ↓ + 2 NO ↑ + 4 H2ONiS ↓ + HNO3 + 3 HCl → Ni2+ + S ↓ + NOCl ↑ + 2 Cl− + 2 H2O

On longer heating the mixture becomes clear because sulphur gets oxidised tosulphate.

S ↓ + 2 HNO3 → SO42− + 2 H+ + 2 NO ↑

S ↓ + 3 HNO3 + 9 HCl → SO42− + 6 Cl− + 3 NOCl ↑ + 8 H+ + 2 H2O

The NiS precipitate dissolves also in the 1+1 mixture of concentrated aceticacid and 30% hydrogen peroxide:

NiS ↓ + 4 H2O2 → Ni2+ + SO42− + 4 H2O

127

4. Potassium nitrite solution: no precipitate is produced in the presence of aceticacid (difference from cobalt).

5. Potassium cyanide solution: green precipitate of nickel(II) cyanide:

Ni2+ + 2 CN− → Ni(CN)2 ↓

The precipitate is readily soluble in excess reagent, when a yellow solution appearsowing to the formation of tetracyanonickelate(II) complex ions:

Ni(CN)2 ↓ + 2 CN− → [Ni(CN)4]2−

6. Dimethylglyoxime reagent: red precipitate of nickel dimethylglyoxime fromsolutions just alkaline with ammonia or acid solutions buffered with sodium acetate:

CH3 C N OH

CH3 C N OHNi

N N

N N

C

C

C

C

OH

O

O OH

H3C

H3C

CH3

CH3

����+ 2 H +Ni + 2+ 2

Iron(II) (red colouration), bismuth (yellow precipitate), and larger amount ofcobalt (brown colouration) interfere in ammoniakal solution.

Compare the characteristic reactions of copper(II) and nickel(II):

NH3 soln.in excess

H2S KCN NaOH flametest

dimethyl-glyoxime

Cu2+

Ni2+

128

Summarise the solubility of sulphides:

MnS FeS CoS NiS

colour

HCl

hot, cc. HNO3

aqua regia

acetic acid +H2O2

Summarise the reactions of various metal ions with NaOH and NH3 solutions.

Mn2+ Fe2+ Fe3+ Co2+ Ni2+ Al3+ Cr3+ Zn2+

NaOH

NaOHin excess

NH3

NH3in excess

Summarise the solubility of selected metals in cold concentrated and 1+1 dilutednitric acid:

Ti V Cr Fe Co Ni

dilutedHNO3

cc HNO3