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Some Important Compounds of Transition Elements


Table of Content

Ferrous sulphate (Green vitriol), FeSO4.7H­2Ferrous sulphate (Green vitriol), FeSO4.7H­2O

 It occurs in nature as copperas and commonly known as hara Kasis.

  • Preparation of Ferrous Sulphate
  1. By dissolving scrap Fe in dil. H2SO4

  2. From Kipp’s waste which contains ferrous sulphate with some free H2SO4; the latter is neutralised with scrap iron forming FeSO4 and hydrogen.

  3. By the action of air and water on iron pyrites. The solution is treated with scrap iron to remove H2SO4 and to reduce Fe2(SO4)3  to FeSO4.

  • Properties of Ferrous Sulphate
  1. Hydrated and anhydrous FeSO4 are green and white in colour respectively. It is isomorphous with epsom salt, MgSO4.7H2O and ZnSO4.7H2O. It effervesces on exposure to air.

  2. Light green crystals of FeSO4 lose water and turn brown on exposure to air, due to oxidation.

  3. On heating at 300°C it gives anhydrous FeSO4 which on further heating gives Fe2O3 and SO2.

  4. Like other ferrous salts, it takes up HNO3 forming brown coloured double compound, Fe(NO)SO4, nitroso ferrous sulphate (Ring test for nitrates).

  5. It decolourises acidified potassium permanganate and turns acidified dichromate green (reducing character).

  6. It forms double salts with sulphates of alkali metals with general formula R2SO4.FeSO4.6H2O. With ammonium sulphate, it forms a double salt known as ferrous ammonium sulphate or Mohr’s salt, FeSO4.(NH4)2SO4.6H2O. It does not effervesce. It ionises in solution to gives Fe2+, NH4+ and SO42– ions.

Ferric oxide, Fe2O3Ferric oxide, Fe2O3

  1. It occurs in nature as haematite.

  2. Fe­2O3 is a red powder, insoluble in H2O and not acted upon by air or H2O

  3. It is amphoteric in nature and reacts with acids and alkalies.

  4. It is reduced to iron by H2,C and CO.

  5. It is used as a catalyst in the oxidation of CO to CO2 in the Bosch process.

Ferric Chloride, FeCl3

  • Preparation of Ferric Chloride
  1. Hydrated ferric chloride (FeCl3.6H2O) can be prepared by dissolving iron, Fe(OH)3 or ferric oxide in dil. HCl.

  2. Reaction of Fe with dry Cl2 gives anhydrous FeCl3,

  • Properties  of Ferric Chloride

  1. Anhydrous salt is yellow, deliquescent compound and highly soluble in H2O.

  2. Its aqueous solution is acidic due to hydrolysis.

  3. On heating it gives FeCl2 and Cl2.

  4. It oxidizes H2S to S, SO2 to H2SO4, SnCl2 to SnCl4 and Na2S2O3 to Na2S4O6

Copper(II) Sulphate Pentahydrate or Blue Vitriol, CuSO4.5H2OCopper(II) Sulphate Pentahydrate or Blue Vitriol, CuSO4.5H2O

  • Preparation of Cupper (II) Sulphate:

In the laboratory, it is prepared by dissolving cupric oxide, cupric hydroxide or carbonate in dilute H2SO4.

CuO + H2SO4 → CuSO4 + H2O

Cu(OH)2 + H2SO4 → CuSO4 + 2H2O

CuCO3 + H2SO4 → CuSO4 + H2O + CO2

The solution of CuSO4 thus obtained is concentrated and cooled when crystals of blue vitriol, CuSO4.5H2O separates out.

Commercially it is prepared by the action of hot dilute sulphuric acid on scrap copper in the presence of air.

2Cu + 2H2SO4 + O2 → 2CuSO4 + 2H2O      

  • Properties of Cupper (II) Sulphate?

(1) Action of Heat

  1. It has 5 molecules of water of crystallisation; all of which can be removed on heating, to form colourless CuSO4 (again coloured with H2O).
    CuSO4.5H2\overset{373 K}{\rightarrow}CuSO4.H2O\overset{423 K}{\rightarrow} CuSO4 (white ppt)  \overset{Strong\ heating }{\rightarrow} CuO + SO3

  2. At high temperature it forms cupric oxide.

  3. It forms double salts with alkali sulphates, e.g. K2SO4.CuSO4.6H2O

  4. When treated with NH4OH, it first forms precipitate of cupric hydroxide copper (II) sulphate (Schweitzer’s reagent), used for dissolving cellulose in the manufacture of artificial silk.

  5. It reacts with KCN forming a complex compound K3[Cu(CN)4].

  6. It liberates iodine from soluble iodides.

(2) Action of Alkalis

CuSO4 + 2NaOH → Cu(OH)2 + Na2SO4

With NH4OH it forms tetraamminecopper (II) sulphate

CuSO4 + 4NH4OH → [Cu(NH3)4]SO4 + 4H2O

(3) Reaction with KI

CuSO4 + 2KI → CuI2 + K2SO4

2Cul2 → 2Cul + l2

The liberation of iodine in this reaction is quantitative. Therefore, this reaction is used to estimate copper volumetrically.

  • Uses of Cupper (II) Sulphate
  1. It is used as an electrolyte in electroplating, electrotyping and refining of copper.

  2. It is used in reservoirs and swimming pools to prevent the growth of weeds.

  3. It is used as a fungicide under the name Bordeaux mixture, which is a mixture of CuSO4 and slaked lime Ca(OH)2.

  4. Anhydrous CuSO4 is used for detection of moisture in organic liquids such as alcohol, ether etc.

Silver Nitrate, AgNO3Silver Nitrate, AgNO3

  • Preparation of Silver Nitrate

Silver nitrate is prepared by the action of dilute nitric acid on silver and then evaporating the solution to crystallization.

3Ag + 4HNO3 → 3AgNO3 + NO ↑ + 2H2O

  • Properties Silver Nitrate

(1) Action of Heat

It decomposes on heating.

2AgNO3 \overset{723 K}{\rightarrow} 2AgNO2\overset{980 K}{\rightarrow} 2Ag + 2NO2

On coming into contact with organic matter like skin or clothes, it is reduced to finely – divided silver, giving a black stain.

(2) Precipitation Reactions

Silver nitrat forms precipitates with some salt solutions which help in the detection of acid radicals.

Some of the precipitation reactions are:

  1. NaCl + AgNO3 → AgCl ↓ + NaNO3

  2. NaPO4 + 3AgNO3 → Ag3PO4 ↓ + 3NaNO3

  3. K2CrO4 + 3AgNO3 → AgCrO ↓ + 2KNO3

  4. Na2S + 2AgNO3 → Ag2S ↓ + 2NaNO3

  5. Na2S2O3 + 2AgNO3 → Ag2S2O3 ↓ + 2NaNO3

  6. Na2C2O4 + 2AgNO3 → Ag2C2O4 ↓ + 2NaNO3

  7. Na3BO3 + 3AgNO3 → Ag3BO3 ↓ + 3NaNO3

  • Uses of Silver Nitrate

It is used for  1. Preparing silver halides which are used in photography.  2. For making inks and hair dyes.  3. In qualitative and quantitative analysis.  4. For silvering of glass, i.e. preparation of mirrors.

It is used for

  1. Preparing silver halides which are used in photography.

  2. For making inks and hair dyes.

  3. In qualitative and quantitative analysis.

  4. For silvering of glass, i.e. preparation of mirrors.

Halides of Silver

  • Preparation of Silver Halides

Silver halides are prepared by the action of sodium or potassium halide on silver nitrate solution (except for AgF)

AgNO3 + NaX → AgX(s) + NaNO3   

Silver fluoride is prepared by the action of HF on silver (I) oxide. 

2HF + 2Ag2O → 2AgF + H2O

  • Properties of Silver Halides

  1. AgCl is white solid, AgBr is a pale yellow solid and AgI is a yellow solid.

  2. AgF is soluble in water whereas other halides are insoluble in water. AgCl dissolves in ammonia to form a complex.

  3. AgCl + 2NH4OH → [Ag(NH3)2]Cl + 2H2O

  4. AgBr is partially soluble and AgI is insoluble in NH4OH.

  5. All the silver halides dissolve in potassium cyanide and Na2S2O3 solution to form complexes.
    AgCl + 2KCN → K [Ag(CN)2] + KCl
    AgCl + 2Na2S2O3 → Na2[Ag(S2O3)2] + NaCl

  • Uses of Silver Halides

All silver halides (particularly AgBr) are photosensitive and hence are widely used in photography.

Mercury (I) Chloride / Mercurous Chloride /Calomel, (Hg2Cl2)

  • Preparation of Mercurous Chloride

It can be prepared by mixing a chloride solution with a mercury (I) salt solution.

Hg2(NO3)2 + 2NaCl → Hg2Cl2 ↓ + 2NaNO3 

It can also prepared by heating a mixture of mercuric chloride and mercury in an iron vessel.

HgCl2 + Hg \overset{Heat}{\rightarrow} Hg2Cl2
  • Properties of Mercurous Chloride

  1. It is a white power insoluble in water but soluble in chlorine water.
    Hg2Cl2 + Cl2 → 2HgCl2

  2. It decomposes on heating to HgCl2
    Hg2Cl2 → HgCl2 + Hg

  1. On treatment with ammonia, if turns black due to the formation of finely divided mercury.
    Hg2Cl2 + 2NH3→ Hg + Hg(NH2)Cl + NH4Cl

  • Uses of Mercurous Chloride

  1. In making standard calomel electrode and

  2. As a purgative in medicine.

Mercury (II) Chloride HgCl2

  • Preparation of Mercury (II) Chloride 
  1.  It is prepared by passing dry chlorine over heated mercury. 
    HgCl2 + Hg → Hg2Cl2
  1.  It is also obtained by treating HgO with HCl
    HgO + 2HCl → HgCl2 + H2O
  1. Commercially, it is prepared by heating a mixture of HgSO4 and NaCl in the presence of MnO2
    HgSO4 + 2NaCl \overset{MnO_{2}}{\rightarrow} HgCl2 + Na2SO4
  • Properties of Mercury (II) Chloride

It is a white crystalline solid sparingly soluble in cold water but soluble in hot water. Its solubility can be increased by adding Cl-.

HgCl2 + 2Cl → [HgCl4]2–

It is readily soluble in organic solvents suggesting its covalent nature.

When treated with SnCl2 it is reduced to mercury.

2HgCl2 + SnCl2 → SnCl4 + Hg2Cl2

Hg2Cl2 + SnCl2 → 2Hg + SnCl4

When Cu turnings are placed in its contact a shining grey film of mercury deposits over them.

HgCl2 + Cu → Hg + CuCl2

  • Uses of Mercury (II) Chloride 

It is used for preserving wood and hides and for making fungicides.

Mercury-II Iodide

  • Preparation of Mercury (II) Iodide

    Mercury (II)  Iodide is used to prepare Nessler’s reagent and for making ointments for treating skin infections.

It is prepared by treating HgCl2 with KI.

HgCl2 + 2Kl → Hgl2 + 2KCl

  • Properties of Mercury (II) Iodide

Mercuric iodide exists in two forms, i.e. red and yellow. The yellow form is stable above 400 K white the red form is stable below this temperature.

It readily dissolves in KI forming a complex

HgI2 + 2KI → K2[HgI4]

An alkaline solution of K2HgI4 is called Nessler’s reagent and is used to detect the presence of NH4+ with which it gives a brown precipitate due to the formation of iodide of Million’s base.

  • Uses of Mercury (II) Iodide

It is used to prepare Nessler’s reagent and for making ointments for treating skin infections.

Potassium Dichromate, K2Cr2O7

  • Preparation of Potassium Dichromate

It is prepared from the ore called chromate or ferrochrome or chrome iron, FeO.Cr2O3. The various steps involved arePotassium Dichromate, K2Cr2O7

(a) Preparation of sodium chromate

4FeO.Cr2O3 + O2 → Fe2O3 + 4Cr2O3

4Na2CO3 + 2Cr2O3 + 3O2 → 4Na2CrO4 + 4CO2

(b) Conversion of sodium chromate into sodium dichromate.

2Na2CrO4 + H2SO4 → Na2Cr2O7 + Na2SO4 + H2O

(c) Conversion of sodium dichromate into potassium dichromate.

Na2Cr2O7 + 2KCl → K2Cr2O7 + 2NaCl

  • Properties  of Potassium Dichromate

It forms orange red crystals. It is moderately soluble in cold water but freely soluble in hot water.

1. Action of heat

When heated, it decomposed to its chromate

4K2Cr2O7 \overset{Heat}{\rightarrow} 4K2CrO4 + 2Cr2O3 + 3O2

2. Action of alkalis

With alkalis it is converted into chromate which on acidifying gives back dichromate.

K2Cr2O7 + 2KOH → 2K2CrO4 + H2O

2K2Cr2O7 + H2SO4 → K2Cr2O7 + K2SO4 + H2O

In dichromate solution the Cr2O72– ions are in equilibrium with Cr2O72– ions at pH = 4.

Cr2O72– + H2\overset{PH =4}{\rightarrow}    2CrO42– + 2Hl

        orange red                            yellow

3. Action of conc. H2SO4 solution

  • In cold conditions
    K2Cr2O7 + 2H2SO4 ———→ 2CrO3 + 2KHSO4 + H2O

  • In hot conditions
    2K2Cr2O7 + 8H2SO4 ———→ 2K2SO4 + 2Cr2(SO4)3 + 8H2O + 3O2

4. Oxidising properties

It is a powerful oxidising agent. In the presence of dil. H2SO4 it furnishes 3 atoms of available oxygen.

K2Cr2O7 + 4H2SO4 → K2SO4 + Cr2(SO4)3 + 4H2O + 3O

Some of the oxidizing properties of K2Cr2O7 are

  1. It liberates I2 from KI
    K2Cr2O7 + 7H2SO4 + 6Kl → 4K2SO4 + Cr2(SO4)3 + 3l2 + 7H2O

  2. It oxidises ferrous salts to ferric salts
    K2Cr2O7 + 7H2SO4 + 6FeSO4 → K2SO4 + Cr2(SO4)3 + 3Fe2(SO4)3 + 2H2O

  3. It oxidises S-2 to S
    K2Cr2O7 + 4H2SO4 + 3H2S → K2SO4 + Cr2(SO4)3 + 7H2O + 3S

  4. It oxidises nitrites to nitrates
    K2Cr2O7 + 4H2SO4 + 3NaNO2 → K2SO4 + Cr2(SO4)3 + 3NaNO3 + 4H2O

  5. It oxidises SO2 to SO42–
    K2Cr2O7 + H2SO4 + 3SO2 → K2SO4 + Cr2(SO4)3 + 3H2O

  6. It oxidises ethyl alcohol to acetaldehyde and acetic acid.

5. Chromyl Chloride Test

When heated with conc. HCl or with a chloride in the presence of sulphuric acid, reddish brown vapours of chromyl chloride are obtained.

K2Cr2O7 + 4KCl + 6H2SO4 → 2CrO2Cl2 + 6KHSO4 + 3H2O

Thus reaction is used in the detection of chloride ions in qualitative analysis.

  • Uses of Potassium Dicromate

  1. In volumetric analysis for the estimation of Fe2+ and I-.

  2. In chrome tanning in leather industry.

  3. In photography and in hardening gelatin film. 

Potassium Permanganate, KMnO4

  • Potassium Permanganate,Preparation of Potassium Permanganate on a Large Scale

It is prepared from the mineral pyrolusite, MnO2. The preparation involves the following steps

Conversion of MnO2 into potassium manganate.

When finely powdered MnO2 is fused with KOH. K2MnO4 is obtained.

2MnO2 + 4KOH + O2 → 2K2MnO4 + 2H2O

Oxidation of potassium manganate into permanganate

(a) Chemical oxidation

K2MnO4 is oxidised to KMnO4 by bubbling CO2 or Cl2 or ozone into the former.

3K2MnO4 + 2CO2 → 2KMnO4 + MnO2 + 2K2CO3

(b) Electrolytic oxidation

The manganate solution is electrolysed between iron electrodes. The oxygen evolved at anode converts manganate into permanganate.

2K2MnO2 + H2O + O → 2K2MnO4 + 2KOH 

  • Properties of Potassium Permanganate

KMnO4 exists as deep purple prisms. It is moderately soluble in water at room temperature and its solubility in water increases with temperature. 

(i) Action of heat

When heated it decomposes to K2MnO4.

 2KMnO4 → K2MnO4 + MnO2 + O2

(ii) Action of conc. H­2SO4

With cold conc. H2SO4 it gives Mn2O7 which on warming decomposes to MnO2.

2MnO2 + 2H2SO4 → Mn2O7 + 2KHSO4 + 2H2O

2Mn2O7 \overset{Heat}{\rightarrow} 4MnO2 + 3O2

With hot Conc. H­2SO4 O2 is evolved

4KMnO4 + 6H2SO4 → 2K2SO4 + 4MnSO4 + 6H2O + 5O2

(iii) Oxidising properties

KMnO4 is a powerful oxidizing agent. The actual oxidizing action depends upon the medium i.e. acidic, basic or neutral.

In neutral solution, it acts as moderate oxidizing agent.

2KMnO4 + H2O → 2KOH + 2MnO2 + 3O

Some oxidizing properties of KMnO4 in neutral medium are

2KMnO4 + 3Na2S2O3 + H2O → 3K2SO4 + 8MnO2 + 3Na2SO4 + 2KOH

2KMnO4 + 4H2S → 2MnS + S + K2SO4 + 4H2O

In strong alkaline solution, it is converted into

2KMnO4 + 2KOH → 2K2MnO4 + H2O + O

2KMnO4 + H2O + Kl → 2MnO2 + 2KOH + KlO3

In acidic medium, Mn+7 is converted into Mn+2

2KMnO4 + 3H2SO4 → K2SO4 + 2MnSO4 + 3H2O + 5O

Some other reactions are

  1. 2KMnO4 + 3H2SO4 + 5H2S → K2SO4 + 2MnSO4 + 3H2O + 5S

  2. 2KMnO4 + 5SO2 + 2H2O → K2SO4 + 2MnSO4 + 2H2SO4

  3. 2KMnO4 + 3H2SO4 + 5KNO2 → K2SO4 + 2MnSO4 + 3H2O + 5KNO3

  4. 2KMnO4 + 3H2SO4 + 5C2H2O4 → K2SO4 + 2MnSO4 + 8H2O + 10CO2

  5. 2KMnO4 + 8H2SO4 + 10FeSO4 → K2SO4 + 2MnSO4 + 5Fe2(SO4)3 + 8H2O

  6. 2KMnO4 + 3H2SO4 + 10Kl → K2SO4 + 2MnSO4 + 8H2O + 5l2

  • Uses of Potassium Permanganate

  1. It is used in volumetric analysis for the estimation of ferrous salts, oxalates, iodides and H2O2.

  2. It is used as oxidizing agent in the laboratory as well as in industry.

  3. It is also used as disinfectant and germicide.

Question 1: CuSO4.5H2\overset{373 K}{\rightarrow}

a. CuSO4.H2O

b. CuSO4

c. CuO +SO2

d. Cu +SO3

Question 2: When heated with conc. HCl or with a chloride in the presence of sulphuric acid, reddish brown vapours of 

a. potassium dichromate

b. chromyl chloride 

c. chromic acid

d. chlorine

Question 3: Potassium dicromate  oxidises nitrites to 

a. nitrates

b. dinitrogen

c. nitrogen dioxide

d. nitrogen trioxide

Question 4: Hydrated ferric chloride (FeCl3.6H2O) can be prepared by dissolving iron, Fe(OH)3 or ferric oxide in 

a. dil. HCl.

b. dil NaOH

c. H2O

d. NH3









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