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Salt Analysis(XII) INORGANIC CHEMISTRY Salt Analysis(XII) INORGANIC CHEMISTRY
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  1. INORGANIC CHEMISTRY
  2. Compounds of Zinc Zinc, Oxide, ZnO: Zinc oxide is also called zinc white or Chinese white or philosopher’s wool. It occurs in nature as the mineral zincite or red zinc ore. Preparation: It is obtainedby the combustion of zinc orby the calcination ofzinc carbonate, zinc nitrate or zinc hydroxide. 2Zn + O2  2ZnO ZnCO3  ZnO + CO2 2Zn(NO3 )2  2ZnO + 4NO2 + O2 Zn(OH)2  ZnO + H2 O Very pure zinc oxideis preparedby mixinga solution of zincsulphate withsodium carbonate. The basic zinc carbonate thus, precipitated on heating gives pure zinc oxide. 4ZnSO4 + 4Na2 CO3 + 3H2 O  ZnCO3 .3Zn(OH)2 ppt. + 4Na2 SO4 + 3CO2 ZnCO3 .3Zn(OH)2   Heat 4ZnO + 3H2 O + CO2 Properties: (i) It is a white powder. It becomes yellow on heating and again turns white on cooling. (ii) It is very light. It is insoluble in water. It sublimes at 400o C. (iii) It is an amphoteric oxide and dissolves readily in acids forming corresponding zinc salts and alkalies forming zincates. ZnO + H2 SO4  ZnSO4 + H2 O ZnO + 2HCl  ZnCl2 + H2 O ZnO + 2NaOH  zincate Sodium 2 2ZnO Na + H2 O (iv) When heated in hydrogen above 400o C, it is reduced to metal. ZnO + H2  Zn + H2 O It is also reduced by carbon into zinc. ZnO + C  Zn +CO (v) When zinc oxide is heated with cobalt nitrate, a green mass is formed due to formation of cobalt zincate which is known as Riemann’s green. 2Co(NO3 )2  2CoO + 2NO2 + O2 ZnO + CoO  CoZnO2 or CoO.ZnO Uses: (i) Zinc oxide is used as a white pigment (paint). No doubt its covering power is less than but it is superior because it is not blackened in atmosphere of hydrogen sulphide. It can be used both as oil and water paint. (ii) It is used to prepare Rinmann’s green which is empolyed as a green pigment. (iii) It finds use as a catalyst along with Cr2 O3 in the manufacture of methyl alcohol from water gas. Zinc chloride, ZnCl2 .2H2 O Preparation: It is obtained by treating zinc oxide or zinc carbonate or zinc hydroxide with dilute hydrochloric acid. The solution on concentration and coolinggives hydrated zinc choloride crystals, ZnCl2 .2H2 O. ZnO + 2HCl  ZnCl2 + H2 O ZnCO3 + 2HCl  ZnCl2 + CO2 + H2 O Zn(OH)2 + 2HCl  ZnCl2 + 2H2 O
  3. Anhydrous zinc chloride cannot be obtained by heatingcrystals of hydrated zinc chloride as hydrolysis occurs and basic chloride (zinc hydroxy chloride) is formed which on further heating gives zinc oxide. ZnCl2 .2H2 O Zn(OH)Cl + HCl +H2 O Zn(OH)Cl  ZnO + HCl The anhydrous zinc chloride is obtained by heatingzinc in the atmosphere of dry chlorine or dry HCl gas. Zn + Cl2  ZnCl2 Zn + 2HCl  ZnCl2 + H2 This can also be formed by distilling zinc powder with mercuric chloride. Zn + HgCl2  ZnCl2 + Hg Properties: (a) Anhydrous zinc chloride is a white solid, deliquescent and soluble in water. It melts at 660o C and boils at 730o C. (b) Hydrated zinc chloride on heating forms zinc hydroxy chloride or zinc oxychloride. ZnCl2 .2H2 O  Zn(OH)Cl + HCl + H2 O 2ZnCl2 .2H2 O  Zn2 OCl2 + 2HCl + 3H2 O (c) When H2 Sis passed through the solution, a white precipitate of zinc sulphide, is formed. ZnCL2 + H2 S  ZnS+ 2HCl (d) When NaOH is added, a white precipitate of zinc hydroxide appears which dissolves in excess of sodium hydroxide forming sodium zincate. ZnCl2 + 2NaOH  Zn(OH)2 + 2NaCl Zn(OH)2 + 2NAOH  Na2 ZnO2 + 2H2 O (e) Onadding NH4 OHsolution, a white precipitate of zinchydroxide appearswhich dissolves in excess of ammonia forming a complex salt. ZnCl2 + 2NH4 OH  Zn(OH)2 + 2NH4 Cl Zn(OH)2 + 2NH4 OH + 2NH4 Cl    chloride zinc Tetrammine 2 4 3 Cl ) Zn(NH + 4H2 O (f) When the solution of zinc chloride is treated with a solution of sodium carbonate, a white precipitate of basic zinc carbonate is formed. 4ZnCl2 + 4Na2 CO3 + 3H2 O  carbonate zinc Basic 2 3.3Zn(OH) ZnCO + 8NaCl + 3CO2 Butwhena solutionofsodiumbicarbonate isused, awhiteprecipitateofnormalzinc carbonate is formed. ZnCl2 + 2NaHCo3  ZnCO3 + 2NaCl + H2 O + CO (g)Anhydrous zinc chloride absorbs ammonia gas and forms an addition compound. ZnCl2 + 4NH3  ZnCl2 .4NH3 (i) Its syrupy solution when mixed with zinc oxide, ZnO, sets to a hard mass forming an oxychloride, ZnCl2 .3ZnO. Zinc sulphate (White vitriol), ZnSO4 .7H2 O Preparation: It is prepared by reacting zinc with dilute sulphuric acid. It can also be prepared by dissolving zinc oxide or carbonate in dilute sulphuric acid. The solution on concentration and crystallisation below 39o C gives colourless crystals of zinc sulphate, ZnSO4 .7H2 O. Zn + H2 SO4  ZnSO4 + H2 ZnO + H2 SO4  ZnSO4 + H2 O ZnCO3 + H2 SO4  ZnSO4 + H2 O + CO2
  4. Properties: (a) It is a colourless, crystalline solid. It is an efflorescent substance. It is freely soluble in water. (b) On heatng, the following changes occur. ZnSO4 .7H2 O C o 70 Below C o 39 Above        ZnSO4 .6H2 O        C o 0 7 Above ZnSO4 .H2 O  C o 280 Above O2 + SO2 + ZnO      C o 800 ) (anhydrous 4 ZnSO ZnSO4     C o 800 ZnO + SO3  SO2 + ½O2 (c) When sodium hydroxide is added to the solution of zinc sulphate, a white precipitate of zinc hydroxide appears which dissolves in excess of NaOH forming sodium zincate. ZnSO4 + 2NaOH  Zn(OH)2 + Na2 SO4 Zn(OH)2 + 2NaOH  Na2 ZnO2 + 2H2 O (d) When sodium carbonate solution is added to the solution of zinc su;phate, a white precipitate of basic zinc carbonate is formed. 4ZnSO4 +4Na2 CO3 + 3H2 O  ZnCO3 .3Zn(OH)2 + 4Na2 SO4 + 3CO2 However, whenthe solutionof sodiumbecarbonate isadded, normalzinc carbonateis formed. ZnSO4 +2NaHCO3  ZnCO3 + Na2 SO4 +H2 O+CO2 (e) With alkali metal sulphates and (NH4 )2 SO4 , it forms double sulphates such as K2 SO4 .ZnSO4. 6H2 O. Compounds of Silver Silver nitrate (Lunar caustic), AgNO3 Silver nitrate is the most common and important salt of silver. Preparation: It is prepared by heating silver with dilute nitric acid. The solution is concentrated and cooled when the crystals of silver nitrate separate out. 3Ag + (Dilute) 3 4HNO   Heat 3AgNO3 + NO +2H2 O Properties: (a) It is a colourless crystalline compound, soluble in water and alcohol. It melts at 212o C.
  5. (b) In contact with organic substance it blackens due to decomposition into metallic silver. Thus, it leaves blackstains when comes in contact with skin and clothes. It produces burning sensation like caustic and leaves a white stain (usually a black stain) like the moon luna on skin and thus, called Lunar caustic. It is decomposed by light also and therefore stored in coloured bottles. (c) On heating above its melting point, it decomposes to silver nitrtie and oxygen. 2AgNO3  2AgNO2 + O2 When heated at red heat, it further decomposes to metallic silver. 2AgNO3  2Ag + 2NO2 + O2 (d) Solutions of halides, phosphates, sulphides, chromates, thiocyanates, sulphates and thiosulphates, allgive a precipitate ofthe correspondingsilver salt with silvernitrate solution. On account of these reactions, silver nitrate is an excellent laboratory regent for the identification of various acidic radical. (e) Solid AgNO3 absorbs ammonia gas with the formation of an addition compound, AgNO3 .3NH3 . (f) When treated with a solution of NaOH, it forms precipitate of silver oxide. Originally, it has brown colour but turns blackwhen dried. 2AgNO3 + 2NaOH  Ag2 O + 2NaNO3 + H2 O (g) When KCN is added to silver nitrate, a white precipitate of silver cyanide appears which dissolves in excess of KCN forming a cimplex salt, potassium argento cyanide. AgNO3 + KCN  AgCN + KNO3 AgCN + KCN  nide argentocya Potassium 2 KAg(CN)
  6. (h) When sodium thiosulphate is added to silver nitrate, a white precipitate of silver thiosulphate appears. This precipitate, however, dissolves in excess of sodium thioslphate forming a complex salt. 2AgNO3 + Na2 S2 O3  Ag2 S2 O3 + 2NaNO3 Ag2 S2 O3 + 3Na2 S2 O3  osulphate argentothi Sodium 2 3 2 3 ] ) O [Ag(S 2Na (i)AgNO3 reacts with iodine in two ways: (a) 6AgNO3 (excess) + 3I2 + 3H2 O  AgIO3 + 5AgI + 6HNO3 (b) 5AgNO3 + 3I2 (excess) + 3H2 O  HIO3 + 5AgI + 5HNO3 (j) Silver is readily displaced from as aqueous silver nitrate solution by the base metals, particularly, if the solution is somewhat acidic, 2AgNO3 + Cu  2Ag + Cu(NO3 )2 2AgNO3 + Zu  2Ag + Zn(NO3 )2 (k) Phosphine, arsine and stibine all precipitate silver from silver nitrate solution. PH3 + 6AgNO3 + 3H2 O  6Ag + 6HNO3 + H3 PO3 AsH3 + 6AgNO3 + 3H2 O  6Ag + 6HNO3 + H3 AsO3 (l)All halogen acids, except HF, precipitate silver halides fromaqueous solution ofAgNO3 . AgNO3 + HX  AgX + HNO3 (b) It converts glucose to gluconic acid. Ag2 O + C6 H12 O6  2Ag+ C6 H12 O7 (c) It oxidies formaldehyde to formic acid. Ag2 O + HCHO  2Ag+ HCOOH Uses: (i) It is used as a laboratory reagent for the identification ofvarious acidic radicals especially for chloride, bromide and iodide. The ammonical silver nitratesolution, i.e.,Tollen’s reagent sugars, etc. (ii) Silver nitrate is used for making silver halides which are used in photography. (v) It is used extensivley for the preparation of silver mirrors. The process of depositing a thin and uniform layer of silver on a clean glass surface is known as silvering of mirrors. It is employed for makinglookingglasses, concavemirrors and reflectingsurfaces. The process is based on the reduction of ammonical silver nitrate solution by some reducing agent like formaldehyde, gloucose, etc. The silver film deposited on the glass is first coated with a varnish and finally painted with red lead to prevent its being scraped off. Compounds of Copper Copper forms two series of compounds 1. Cuprous compounds: In which copper is monovalent. Most of the cuprous compounds are colourless and diamagnetic as 3d shell is completely filled. Cu2 O and Cu2 S are red and black, respectively, which are exceptions. Cuprous compounds are generally insoluble in water. The soluble compounds are unstable in aqueous solutions, sincethey disproportionate to Cu2+ and Cu. 2Cu+  Cu2+ + Cu Cuprouscompounds canbe obtainedby passingsulphur dioxidethrough asolution containing copper sulphate and sodium salt. Some examples are given below: (i) 2CuSO4 + 2NaCl + SO2 + 2H2 O  chloride Cuprous 2CuCl + Na2 SO4 + 2H2 SO4
  7. (ii) 2CuSO4 + 2NaBr + SO2 + 2H2  bromide Cuprous 2CuBr +Na2 SO4 + 2H2 SO4 (iii) 2CuSO4 + 2NaI + SO2 +2H2 O  iodide Cuprous 2CuI + Na2 SO4 + 2H2 SO4 (iv) 2CuSO4 + 2NaCN + SO2 + 2H2 O  cyandia Cuprous 2CuCN + Na2 SO4 + 2H2 SO4 (v) 2CuSO4 + 2NaCNS + SO2 + 2H2 O  te thiocyana Cuprous 2CuCNS + Na2 SO4 + 2H2 SO4 The reactions (iii) and (iv) can take place even in absence of sulphur dioxide. 2CuSO4 + 4NaI  2CuI+ 2Na2 SO4 + I2 2CuSO4 + 4NaCN  2CuCN + 2Na2 SO4 + (CN)2 The true molecular formula of cuprous compounds is still doubtful. There are expermental evidences for dimeric molecule. The most important compound of this class is cuprous chloride. Cuprous Chloride, Cu2 Cl2 Preparation: It is prepared (i) by heatingexcess of copper with concentrated hydrochloric acid in presence of a little potassium chlorate. Cu + 2HCl + O  CuCl2 + H2 O CuCl2 + Cu  Cu2 Cl2 (ii) By boiling copper sulphate solution with excess of copper turnings in presence of hydrochloric acid. CuSO4 + 2HCl  CuCl2 + H2 SO4 CuCl2 + Cu Cu2 CL2 (iii) By heating cupric chloride with zinc or sulphur dioside. 2CuCl2 + Zn  Cu2 Cl2 + ZnCL2 2CuCl2 + 2H2 O + SO2  Cu2 Cl2 + 2HCl + H2 SO4 (iv) By passing SO2 thruough the solution containing copper sulphate and sodium chloride. 2CuSO4 + 2NaCl + 2H2 O + SO2  Cu2 Cl2 + Na2 SO4 + 2H2 SO4 Properties:(i) It is a white solid. It is insoluble in water but soluble in excess of hydrochloric acid. Cu2 Cl2 + 4HCl  2H2 CuCl3 Cu2 Cl2 + 6HCl  2H3 CuCl4 (ii) It gradually turns green on exposure in air due to oxidation. 2Cu2 Cl2 + 2H2 O + O2  2[CuCl2 .Cu(OH)2 ] (iii) The solution of cuprous chloride in HCl is oxidised by air or oxidising agents into cupric chloride. Cu2 Cl2 + 2HCl + 2 O 2 1  2CuCl2 + H2 O (iv) The solution of cuprous chloride in HCl absorbs carbon monoxideand forms an addition compound. Cu2 Cl2 + 2CO  2CuCl. CO The addition compound decomposes on heating evolving carbon monoxide. The reaction is uitilised for the removel of carbon monoxide. (v) It dissolves in aqueous ammonia forming a colourless soultion due to the formation of the complex Cu(NH3 )2 Cl.
  8. (vi) The ammonicalcuprous chloridesolution absorbs acetylene to form bright red precipitate of cuprous acetylide, Cu2 C2 . 2Cu(NH3 )2 Cl + C2 H2  Cu2 C2 + 2NH3 + 2NH4 Cl Acetylene can be regenerated by treating the acetylide with strong HCl. The reaction is , therefore, used for the purification and separation of acetylene. Cu2 C2 + 2HCl  C2 H2 + Cu2 Cl2 (vii) Cuprous chloride with caustic alkalies gives a yellow precipitate of cuprous oxide which gradually changes to red. Cu2 Cl2 + 2NaOH  red to changing Yellow 2O Cu + 2NaCl + H2 O (viii) With H2 S, cuprous chloride forms a black precipitate of cuprous sulphide. Cu2 Cl2 + H2 S  Cu2 S + 2HCl (ix) With sodium chloride or potassium chloride solution cuprous chloride forms a soluble complex. Cu2 Cl2 + 6NaCl  2Na3 CuCl4 Cu2 Cl2 + 6KCl  2K3 CuCl4 (x) Dry cuprous chloride forms addition compounds with ammonia gas of the formula CuCl.nNH3 where n= 1, 2 1 1 , 3. Uses: (i) Ammonical solution of cuprous chloride is used for absorbing acetylene. (ii) HCl solution of cuprous chloride is used for absorption of carbon monooxide. (iii) It is also used for absorption of ammonia gas. Cupric compounds: In which copper is divalent, cupric compounds are more stable, more common and generally more stable. Most of the anhydrous cupric compounds are colourless while the hydrated compounds are generally blue due to the formation of blue hydrated ion, [Cu(H2 O)4 ]2+ or [Cu(H2 O)6 ]2+ . Compounds of Cu2+ ions are paramagenetic due to presence of one unpaired electron in 3d energy shell, i.e., configurationof Cu2+ is 3d9 . Some important cupric compounds are described here. Cupric Oxide, CuO It is called black oxide of copper. Preparation: It is prepared- (i) By heating Cu2 O in air or by heatingcopper for a longtime in air (The temperature should not exceed above 1100o C). Cu2 O + 2 O 2 1  2CuO 2Cu + O2  2CuO (ii) by heatingcupric hydroxide, Cu(OH)2  CuO + H2 O (iii) by heating copper nitrate, 2Cu(NO3 )2  2CuO + 4NO2 + O2 (iv) on a commercial scale, it is obtained by heating malachite which is found in nature. CuCO3 .Cu(OH)2  2CuO + CO2 + H2 O
  9. Properties: (a) It is black powder and stable to moderate heating. (b) The oxide is insoluble in water but dissolves in acids forming corresponding salts. CuO + 2HCl  CuCl2 + H2 O CuO + H2 SO4  CuSO4 + H2 O CuO + 2HNO3  Cu(NO3 )2 + H2 O (c) When heated to 1100 - 1200o C, it is converted into cuprous oxide with evolution of oxygen. 4CuO  2Cu2 O + O2 (d) It is reduced to metallic copper by reducing agents like hydrogen, carbon and carbon monoxide. CuO + H2  Cu + H2 O CuO + C  Cu + CO CuO + CO  Cu + CO2 Cupric Chloride, CuCl2 .2H2 O Preparation: (i) The metal or cupric oxide or cupric hydroxide or copper carbonate is dissolved in conc. HCl. The resulting solution on crystallisation gives green crystals ofhydrated cupric chloride. 2Cu + 4HCl +O2  2CuCl2 + 2H2 O CuO + 2HCl  CuCl2 + H2 O Cu(OH)2 CuCO3 + 4HCl  2CuCl2 + 3H2 O + CO2 (ii)Anhydrous cupric chloride is obtained as a dark brown mass whencopper metal is heated in excess of chlorine gas or by heating hydrated cupric chloride in HCl gas at 150o C. Cu + Cl2  CuCl2 CuCl2 .2H2 O gas HCl C 150      CuCl2 + 2H2 O Properties: (i) It is deliquescent compound and is reado;u soluble in water. The dilute solution is blue but concentrated solution is, however, green. It changes to yellow whne conc. HCl is added. The blue colour is due to complex cation [Cu(H2 O)4 ]2+ and yellow colour due to complex anion [CuCl4 ]2- and green when both are present. (ii) The aqueous solution is acidic due to its hydrolyis. CuCl2 + 2H2 O Cu(OH)2 + 2HCl (iii) The anhydrous salt on heatingforms Cu2 Cl2 and Cl2 . 2CuCl2  Cu2 Cl2 + Cl2 while the hydrated salt on strong heating gives CuO, Cu2 Cl2 , HCl and Cl2 . 3CuCl2 .2H2 O  CuO + Cu2 Cl2 + 2HCl + Cl2 5H2 O (iv) It is readily reduced to Cu2 Cl2 by copper turnings, or SO2 gas, or hydrogen (Nascent - obtained by the action of HCl on Zn) or SnCl2 . CuCl2 + Cu  Cu2 Cl2 2CuCl2 + SO2 + 2H2 O  CuCl2 + 2HCl + H2 SO4 2CuCl2 + 2H  CuCl2 + 2HCl 2CuCl2 + SnCl2  Cu2 Cl2 + SnCl4 (v)Apale blue precipitate of basic cupric chloride, CuCl2 .3Cu(OH)2 is obtainedwhen NaOH is added.
  10. CuCl2 2NaOH  Cu(OH)2 + 2NaCl CuCl2 +3Cu(OH)2  CuCl2 .3Cu(OH)2 It dissolves in ammonium hydroxide forming a deep blue solution. On evaporating of this solution deep blue crystals of tetrammine cupric chloride are obtained. CuCl2 + 4NH4 OH  Cu(NH3 )4 Cl2 .H2 O + 3H2 O Copper Sulphate (Blue Vitriol), CuSO4 .. 5H2 O Copper sulphate is the most common compound of copper. It is calles as blue vitriol or Nila Thotha. Preparation: (i) Copper sulphate is prepared in the laboratory by dissolving cupric oxide or hydroxide or carbonate in dilute sulphuric acid. The solution is evaporated and crystallised. CuO + H2 SO4  CuSO4 + H2 O Cu(OH)2 + H2 SO4  CuSO4 + 2H2 O Cu(OH)2 CuCO3 + 2H2 SO4  2CuSO4 + 3H2 O + CO2 (ii) On a commercial scale, it is prepared from scrap copper is placed in a perforated lead bucket which is dipped into hot dilute sulphuric acid.Air is blown thruough the acid. Copper sulphate is crystallised from the solution. Cu + H2 SO4 + (air) O 2 1 2  CuSO4 + H2 O Properties: (a) It is a blue crystalline compound and is fairly soluble in water. (b) Heating effect: CuSO4 .5H2 O crystals effloresce on exposure and converted into a pale blue powder, CuSO4 .3H2 Ois formed.The monohydrateloses last molecule of water at230o C giving the anhydrous salt, CuSO4 , which is white. CuSO4 .5H2 O blue Pale 2 4 O .3H CuSO te Bluish whi 2 4 O .H CuSO White 4 CuSO Anhydrous copper sulphate (white) regains its blue colour when moistened with a drop of water (test of water). If theanhydrous salt is heated at 720o C, it decomposes into cupricoxide andsulphur trioxide. CuSO4      C 720 CuO + (c) Action of NH4 OH: With ammonia solution, it forms the soluble blue complex. First it forms a precipitate of Cu(OH)2 which dissolves in excess of ammonia solution. CuSO4 + 2NH4 OH  Cu(OH)2 + (NH4 )2 SO4 Cu(OH)2 + 2NH4 OH + (NH4 )2 SO4  sulphate cupric Tetrammine 4 4 3 SO ) Cu(NH + 4H2 O The complex is known as Schwixer’s reagent which is used for dissolving cellulose in the manufacture of artificial silk. (d) Action of alkalies: Alkalies form a pale blue precipitate of copper hydroxide. CuSO4 + 2NaOH  Cu(OH)2 + Na2 SO4 (e) Action of potassium iodide: First cupric iodide is formed which decomposes to give white cuprous iodide and iodide. [CuSO4 + 2KI  CuI2 + K2 SO4 ] x 2 2CuI2  Cu2 I2 + I2 ––––––––––––––––––––––––––––––––––––––––––––––– 2CuSO4 + 4KI  Cu2 I2 + 2K2 SO4 + I2
  11. (f)Action ofpotassiumcyanide: First cupric cyanide is formed which decomposes to give cuprous cyanide and cyanogen gas. Cuprous cyanide dissolves in excess of potassium cyanide to form a complex, potassium cupro cyanide [K3 Cu(CN)4 ]. [CuSO4 + 2KCN  Cu(CN)2 + K2 SO4 ] x 2 2Cu(CN)2  Cu2 (CN)2 + (CN)2 Cu2 (CN)2 + 6KCN 2K3 Cu(CN)4 ___________________________________________________ 2CuSO4 + 10KCN  2K3 Cu(CN)4 + 2K2 SO4 + (CN)2 (g)Action ofpotassiumferrocyanide:Reddish brown precipitate of cupric ferrocyanideis formed. (test of Cu2+ ion) 2CuSO4 + K4 Fe(CN)6  Cu2 Fe(CN)6 + 2K2 SO4 (h) Addition of electropositive metals: Electropositive elements like zinc and iron precipitate copper from a solution of copper sulphate. CuSO4 + Fe  Cu + FeSO4 CuSO4 + Zn  Cu + ZnSO4 (i) Action of H2 S: When H2 S is passed through copper sulphate solution, a black precipitate of copper sulphide is formed. CuSO4 + H2 S  CuS + H2 SO4 The black precipitate dissolves in conc. HNO3 . 3CuS + 8HNO3  3Cu(NO3 )2 + 2NO + 3S + 4H2 O (j) Action of potassium sulphocyanide: Cupric sulphocyanide is formed. CuSO4 + 2KCNS  Cu(CNS)2 + K2 SO4 If SO2 ispassed through the solution, a white precipitate ofcuprous sulphocyanideis formed. 2CuSO4 + 2KCNS + SO2 +2H2 O  Cu2 (CNS)2 + K2 SO4 + 2H2 SSO4 [This is the general method for obtainingcuprous compounds]. (k) Double sulphates: Copper sulphate forms double salts with alkali sulphate to form cupric thiosulphate which is reduced by sodium thiosulphate. The cuprous compound thus formed dissolves in excess of sodium thiosulphate to form a complex. CuSO4 + Na2 S2 O3  CuS2 O3 + Na2 SO4 2CuS2 O3 + Na2 S2 O3 Cu2 S2 O3 + Na2 S4 O6 3Cu2 S2 O3 + 2Na2 S2 O3  Na4 [Cu6 (S2 O3 )5 ] Uses: (i) Copper sulphate is used for the preparation of other copper compounds. (ii) It finds use in electroplating, electrotyping, calicoprinting and dyeing. (iii) It is used in agriculture as afungicide and germicide. Bordeaux mixture consisting copper sulphate and lime is used to kill moulda and fungi on vines, trees, potatoes, etc. (iv) It is used as a laboratory reagent especially in the preparation of Fehling’s soluton. (v) It finds use as an antiseptic in medicine. (vi) It is extensivley used in electric batteries.
  12. Compounds of Iron Ferrous sulphate (Green witriol), FeSO4 .7H2 O Thisis thebest knownferrous salt. It occursin natureas copperand isformed bythe oxidation of purites under the action of water and atmospheric air. 2FeS2 + 7O2 + H2 O  2FeSO4 + 2H2 SO4 Preparation: (i) It is obtained by dissolving scrap iron in dilute sulphuric acid. Fe + H2 SO4  FeSO4 + H2 The solution is crystallised by the addition of alcohol as ferrous sulphate is sparingly soluble in it. Manufacture: Commercially, ferrous sulphate is obtained by the slow oxidation of iron pyrites in the presence of air and moisture. The pyrites are exposed to air in big heaps. 2FeS2 + 2H2 O + 7O2  2FeSO4 + 2H2 SO4 Properties: (i) Hydrated ferrous sulphate (FeSO4 .7H2 O) is green crystalline compound. Due to atmospheric oxidation, the crystals acquire brownish-yellow colour due to formation of basic ferric sulphate. 4FeSO4 + 2H2 O + O2  sulphate ferric Basic 4 4Fe(OH)SO (ii) Action of heat: At 300o , it becomes anhydrous. The anhydrous ferrous sulphate is colourless. The anhydrous salt when strongly heated, breaks up to form ferric oxide with the evolution of SO2 and SO3 . . O 2 7H C 300 White 4 Green 2 4 FeSO O .7H FeSO Temp High           Fe2 O3 + SO2 + SO3 (iii) The aqueous solution of ferrous sulphate is slightly acidic due to its hydrolysis. FeSO4 + 2H2 O base Weak 2 Fe(OH) + acid Strong 4 2SO H (iv) Ferrous sulphate is a strong reducing agent. (a) It decolourises acidified potassium permanganate. 2KMnO4 + 3H2 SO4  K2 SO4 + 2MnSO4 + 3H2 O + 5[O] [2FeSO4 + H2 SO4 + O  Fe2 (SO4 )3 + H2 O] x 5 __________________________________________________________ 10FeSO4 + 2KMnO4 + 8H2 SO4  5Fe2 (SO4 )3 + K2 SO4 +2MnSO4 +8H2 O
  13. (b) It turns potassium dichromate (acidified) green as dichromate is reduced to chromic salt (green). K2 Cr2 O7 + 4H2 SO4  K2 SO4 + Cr2 (SO4 )3 + 4H2 O + 3[O] [2FeSO4 | H2 SO4 + O Fe2 (SO4 )3 + H2 O x 3 __________________________________________________________ 6FeSO4 + K2 Cr2 O7 + 7H2 SO4  3Fe2 (SO4 )3 + K2 SO4 +Cr2 (SO4 )3 + 7H2 O (c) It reduces gold chloride to gold. AuCl3 + 3FeSO4  Au + Fe2 (SO4 )3 + FeCl3 (d) It reduces mercuric chloride to mercurous chloride. [2HgCl2  Hg2 Cl2 + 2Cl] x 3 [3FeSO4 + 3Cl  Fe2 (SO4 )3 + FeCl3 ] x 2 _________________________________________________________ 6HgCl2 + 6FeSO4  3Hg2 Cl2 + 2Fe2 (SO4 )3 + 2FeCl3 (v)Acold solution of ferrous sulphate absorbs nitirc oxide formingdark brown addition compound, nitroso ferrous sulphate. FeSO4 + NO  (Brown) sulphate ferrous Nitroso 4.NO FeSO The NO gas is evolved when the solution is heated. (vi) It forms double sulphates of the composition R2 SO4 .FeSO4 .6H2 O where R = an alkali metal metal or NH4 + radical. (NH4 )2 SO4 .FeSO4 .6H2 O (ferrous ammonium sulphate) is known as Mohr’s salt. (vii) It combines with potassium cyanide (excess) forming potassium ferrocyanide, K4 Fe(CN)6 . FeSO4 + 2KCN  Fe (CN)2 + K2 SO4 Fe(CN)2 + 4KCN  K4 Fe(CN)6 _____________________________ FeSO4 + 6KCN  K4 Fe(CN)6 + K2 SO4 Ferrous Ammonium Sulphate (Mohr’s salt). (NH4 )2 SO4 .FeSO4 .6H2 O Preparation: The double salt is best prepared by makingsaturated solutions of pure ferrous sulphte and pure ammonium sulphate in air free distilled water at 40o C. Both the solutions are mixed and allowed to cool. Generally, few drops of sulphuric acid and a little iron wire are added before crystallisation as to prevent oxidation of ferrous sulphate into ferric sulphate. The salt is obtained as pale green crystals. Properties: It is pale green crystalline compound which does not effloresce like ferrous sulphate. It is less readily oxidised in the solid state. Ferric chloride, FeCl3 This is the most important ferric salt. It is known in anhydrous and hydrated forms. The hydrated form consists of six water molecles, FeCl3 .6H2 O. Preparation: (i) Anhydrous ferric chloride is obtained by passing dry chlorine gas over heated iron fillingsAS shown in figure. The vapours are condensed in a bottle attached to the outlet of the tube. 2Fe + 3Cl2  2FeCl3 (ii) Hydrated ferric chloride is obtained by the action of hydrochloric acid on ferric carbonate, ferric hydroxide or ferric oxide. Fe2 (CO3 )3 + 6HCl  2FeCl3 + 3H2 O + 3CO2 Fe(OH)3 + 3HCl  FeCl3 + 3H2 O __________________________________ Fe2 O3 + 6HCl  2FeCl3 + 3H2 O
  14. The solution on evaporation and coolingdeposits yellow crystals of hydratedferric chloride, FeCl3 .6H2 O. Properties: (i) Anhydrous ferric chloride is a dark red deliquescent solid. It is sublimed at about 300o C and its vapour density corresponds to dimeric formula, Fe2 Cl6 . The dimer dissociates at high temperatures to FeCl3 . The dissociation into FeCl3 is complete at 750o C.Above this temperature it breaks into ferrous chloride and chlorine. Fe2 Cl6 2FeCl3 2FeCl2 + Cl2 (ii)Anhydrous ferric chloride behaves as a covalent compound as it is soluble in non-polar solvents like ether, alchol, etc. It is represented by chlorine bridge structure. (iii) It dissolves in water. The solution is acidic in nature due to its hydrolysis as shown below: FeCl3 + 3HOH Fe(OH)3 + 3HCl The solution is stabilised by the addition of hydrochloric acid to prevent hydrolysis. (iv)Anhydrous ferric chloride absorvs absorbs ammonia. FeCl3 + 6NH3  FeCl3 .6NH3 (v) Ferric chloride acts as an oxidising agent. (a) It oxidises stannous chloride to stannic chloride. 2FeCl3 + SnCl2  2FeCl2 + SnCl4 (b) It oxidises SO2 To H2 SO4 . 2FeCl3 + SO2 2H2 O  2FeCl2 + H2 SO4 + 2HCl (c) It oxidises H2 S to S. 2FeCl3 + H2 S  2FeCl2 + 2HCl + S (d) It liberates iodine from KI. 2FeCl3 + 2KI  2FeCl2 + 2KCl + I2 (e) Nascent hydrogen reduces FeCl3 into FeCl2 . FeCl3 + H  FeCl2 + HCl (vi) When ammonium hydroxide is added to the solution of ferric chloride, a reddish - brown precipitate of ferric hydroxide is formed. FeCl3 + 3NH4 OH  Fe(OH)3 + 3NH4 Cl (vii) When a solution of thiocyanate ions is added to ferric chloride solution, a deep red colouration is produced due to formation of a complex salt. FeCl3 + NH4 CNS  Fe(SCN)Cl2 + NH4 Cl Or FeCl3 + 3NH4 CNS  Fe(SCN)3 + 3NH4 Cl (viii) Ferric chloride forms a complex, prussian blue with potassium ferrocyanide. 4FeCl3 + 3K4 Fe(CN)6  de) ferrocyani (Ferri blue Prussian 3 6 4 ] [Fe(CN) Fe + 12KCl (ix) On heating hydrated ferric chloride FeCl3 .6H2 O, anhydrous ferric chloride is not obtained. It is changed to Fe2 O3 with evolution of H2 O and HCl. 2[FeCl3 .6H2 O]    Heat Fe2 O3 + 6HCl + 9H2 O Hydrated ferric chloride may be dehydrated by heating with thionyl chloride. FeCl3 .6H2 O + 6SOCl2  FeCl3 + 12HCl + 6SO2
  15. Corrosion of iron: Corrosion is defined as the gradual transformation of a metal into its combined state because of the reaction with the environment. Metals are usually extracted from their ores. Nature tries to convert them again into the ore form. The process by which the metals have the tendency to go back to their combined state, is termed corrosion. When iron is exposed to moist air, it is found covered with a reddish - brown coating which can easily be detached. The redish brown coatingis called ‘rust’. Thus, the corrosion of iron or fromation of the rust is called rusting. The composition of the rust is not certain but it mainly contains hydrated ferricoxide, 2Fe2 O3 .3H2 O, together witha small quantityof ferrous carbonate. The rust is formed by the action of water on iron in presence of dissolved oxygen and carbon dioxide. It has been observed that impure iron is more prone to rusting. The following are the favourable conditions for the rusting of iorn: (i) Presence of moisture (ii) Presence of a weakly acidic atmosphere (iii) presence of impurity in the iron. Various theories have been proposed to explain the phenomenon of rusting of iron but the accepted theory is the modern electrochemical theory. When impure iron comes in contact with water containing dissolved carbon dioxide, a voltaic cell is set up. The iron and other impurities act as electrodes while water having dissolved oxygen and carbon dioxide acts as an electrolyte. Iron atoms pass into solution as ferrous ions. Fe  Fe2+ + 2e Iron, thus, acts as anode. The impurities act as cathode.At the cathode, the cathode, the electrons are used in forming hydroxyl ions. H2 O + O + 2e  2OH- In presence of dissolved oxygen, ferrous ions are oxidised to ferric ions which combine with hydroxyl ions to form ferric hydroxide. Fe3+ + 3OH-  Fe(OH)3 [2Fe2+ + H2 O + O  2Fe3+ + 2OH- ] Corrosion or rustingis a surface phenomenon and thus, the protection of thesurface prevents the corrosion. Iron can be protected from the rusting by use of following methods: (i)Applying paints, lacquers and enamels on the surface of iron. (iii) By coating a thin film of zinc, tin, nickel, chromum, aluminium, etc. Ferric Oxide (Fe2 O3 ) Preparation: (i) Hydrolysis of FeCl3 actually give red-gelatinous ppt. of the hydrous oxide Fe2 O3 (H2 O)4 which on heating at 200o C give red-brown α- Fe2 O3 . (ii) It occur in haematite ore (Fe2 O3 ).
  16. (iii) On oxidation of Fe3 O4 , γ - Fe2 O3 is formed. (iv) 6Fe2 O3      CΔ 1400 4 Fe3 O4 + O2 Fe3 O4 isa mixedoxide FeO- Fe2 O3 (Occuras magnetite). Properties: (a) Freshly precipitate Fe2 O3 . (H2 O)4 dissolve in acid giving pale violet [Fe(H2 O)6 ]3+ ion. (b) [Fe2 O3 . (H2 O)4 ] also dissolve in concentrated NaOH forming [Fe(OH)6 ]3- . (v) Fusion of Fe2 O3 with Na2 CO3 give NaFeO2 (Sodium ferrites) which is hydeolysed to Fe2 O3 & NaOH Na2 CO3 + Fe2 O3  2NaFeO2 + CO2 2NaFeO2 + H2 O  2NaOH + Fe2 O3 (vi) If Cl2 gas is passed into an alkaline soultion of hydrated ferric oxide, a red purple soultion is formed, containing the ferrate ion [Feiv O4 ]2- . Fe2 O3 + 2NaOH  2NaFeO2 + H2 O 2NaFeO2 + H2 O  2Na2 FeO4 + 2NaCl + 2H2 (vii) Na2 FeO4 can also be obtained by oxidation of Fe2 O3 with NaOCl (Sod. hypochloride). (viii) Na2 FeO4 has Fe (+ VI) and is a strong oxidizing agent (like KMnO4 ). Potassium Permanganate, KMnO4 This is the most important and well known salt of permanganic acid. It is prepared from the pyrolusite ore. It is prepared by fusing pyrolusite ore either with KOH or K2 CO3 in presence of atmospheric oxygen or any other oxidising agent such as KNO3 . The mass turns green with the formation of potassiummanganate, K2 MnO4 . 2MnO2 + 4KOH + O2  2K2 MnO4 + 2H2 O 2MnO2 + 2K2 CO3 + O2  2K2 MnO4 + 2CO2 The fused mass is extractedwith water.The solution is now treated with a current of chlorine or ozone or carbon dioxide to convert manganate into permanganate. 2K2 MnO4 +Cl2  2KMnO4 + 2KCl 2K2 MnO4 = H2 O + O3  2KMnO4 + 2KOH + O2 3K2 MnO4 + 2CO2  2KMnO4 + MnO2 + 2K2 CO3 It is purple coloured crystalline compound. It is fairly soluble in water. When heated alone or with an alkali, it decomposes evolving oxygen. 2KMnO4  K2 MnO4 + MnO2 + O2 4KMnO4 + 4KOH  4K2 MnO4 + 2H2 O + O2 On treatment with conc. H2 SO4 , it forms manganese heptoxide via permanganyl sulphate which decomposes explosively on heating. 2KMnO4 + 3H2 SO4  2KHSO4 + (MnO3 )2 SO4 + 2H2 O (MnO3 )2 SO4 + H2 O  Mn2 O7 + H2 SO4 Mn2 O7  2MnO2 + 2 O 2 3 Potassium permanganate acts as an oxidising agent in alkaline, neutral or acidic solutions. (a) In alkaline solution: KMnO4 is first reduced to manganate and then to insoluble manganese dioxide. Colour changes first purple to green and finally becomes colourless. However brownish precipitate is formed. 2KMnO4 + 2KOH  2K2 MnO4 + H2 O + O
  17. 2K2 MnO4 + 2H2 O  2MnO2 + 4KOH + 2O _______________________________________ 2KMnO4 + H2 O     Alkaline 2MnO2 + 2KOH + 3[O] or  4 2MnO + H2 O 2MnO2 + 2OH- + 3[O] (b) In neutral solution: MnO2 is formed. Brownish ppt. is present. 2KMnO4 + H2 O  2MnO2 + 2KOH + 3[O] or  4 2MnO + H2 O  2MnO2 + 2OH- + 3[O] or  4 MnO + 2H2 O + 3e-  MnO2 + 4OH- (c) In acidic solution (in presence of dilute H2 SO4 ): Manganous sulphate is formed. The solution becomes colourless. 2KMnO4 + 3H2 SO4  K2 SO4 + 2MnSO4 + 3H2 O + 5[O] or 2  4 MnO + 6H+  2Mn2+ + 3H2 O + 5[O] or  4 MnO + 8H+ + 5e-  Mn2+ + 4H2 O Thismediumis used inquantitative (volumetric)estimations. Theequivalent mass of KMnO4 in acidic medium is = 5 mass Mol. . The oxidation reactions of acidified KMnO4 are catalysed by Mn (ii) ion. The important oxidation reactions are: (i) Ferrous salts are oxidised to ferric salts. 2KMnO4 + 3H2 SO4  K2 SO4 + 2MnSO4 + 3H2 O + 5[O] [2FeSO4 + H2 SO4 + [O]  Fe2 (SO4 )3 + H2 O] x 5 ________________________________________________________ 2KMnO4 + 10FeSO4 + 8H2 SO4  5Fe2 (SO4 )3 + K2 SO4 +2MnSO4 + 8H2 O or  4 2MnO + 10Fe2+ + 16H+  10Fe3+ + 2Mn2+ + 8H2 O (ii) Iodine is evolved from potassium iodide. 2KMnO4 + 3H2 SO4  K2 SO4 + 2MnSO4 + 3H2 O + 5[O] [2KI + H2 SO4 + [O]  K2 SO4 + I2 + H2 O] x 5 ________________________________________________ 2KMnO4 + 10 KI + 8H2 SO4  6K2 SO4 + 2MnSO4 + 5I2 + 8H2 O or  4 2MnO + 10I- + 16H+  2Mn2+ + 5I2 + 8H2 O (iii) H2 S is oxidised to sulphur. 2KMnO4 + 3H2 SO4 + 5H2 S  K2 SO4 + 2MnSO4 + 5S = 8H2 O (iv) SO2 is oxidised to H2 SO4 . 2KMnO4 + 5SO2 + 2H2O  K2 SO4 + 2MnSO4 + 2H2 SO4 (v) Nitrites are oxidised to nitrates. 2KMnO4 + 5KNO2 + 3H2 SO4  K2 SO4 + 2MnSO4 + 5KNO3 + 3H2 O (vi) Oxalic acid is oxidised to CO2 . + 2KMnO4 + 3H2 SO4  K2 SO4 + 2MnSO4 + 10CO2 + 8H2 O (vii) It oxidises hydrogen halides (HCl, HBr or HI) into X2 (halogen). 2KMnO4 + 3H2 SO4 + 10 HX  K2 SO4 + 2MnSO4 + 8H2 O + 5X2
  18. In neutral medium (i) H2 S is oxidised to sulphur. 2KMnO4 + H2 O  2MnO2 + 2KOH + 3[O] [H2 S + [O]  H2 O + S] x 3 ______________________________________ 2KMnO4 + 3H2 S  2KOH + 2MnO2 + 2H2 O + 3S (ii) Manganese sulpjate is oxidised to MnO2 . 2KMnO4 + H2 O  2MnO2 + 2KOH + 3 [O] [MnSO4 + H2 O + [O]  MnO2 + H2 SO4 ] x 3 2KOH + H2 SO4  K2 SO4 + 2H2 O ____________________________________________ 2KMnO4 + 3MnSO4 + 2H2 O K2 SO4 + 5MnO2 +2H2 SO4 (iii) Sodium thisoulphate is oxidised to sulphate and sulphur. 2KMnO4 + 3Na2 S2 O3 + H2 O  2KOH + 2MnO2 + 3Na2 SO4 + 3S In alkaline medium (i) It oxidises iodide to iodate. 2KMnO4 + H2 O  2KOH + 2MnO2 + 3[O] KI + 3[O]  KIO3 _____________________________________ 2KMnO4 + KI + H2 O  2KOH + 2MnO2 + KIO (ii) It oxidises ethylene to ethylene glyocl. + H2 O+[O]  In alkaline medium it is called Bayer’s reagent. Uses: (i) KMnO4 is used as an oxidisingagent inlaboratory andindustry. Involumetric esrimations, the solution is first standardised before use. (ii)Alkalinepotassiumpermanganateis calledBayer’s reagent.This reagentis usedin organic chemistry for the test of unsaturation. KMnO4 is used in the manufacture os saccharin, benxoic acid, acetaldehyde, etc. (iii) KMnO4 is used in qualitative analysis for detecting halides, sulphites, oxalates, etc. Potassium Dichromate, K2 Cr2 O7 It is the most important compound of Cr(VI). It is manufactured from chromite ore. is first converted into sodium dichromate . The hot saturated solution of sodium dichromate is mixed with KCl. Sodium chloride of sodium chloride precipitates out from the hot solution which is filtered off. On coolingthe mother liquor, crystals of potassiumdichromate separate out. It is orange-red coloured crystalline compound. It is moderately soluble in cold water but freely soluble in hot water. It melts at 398o C. On heating strongly, it decomposes liberating oxygen. 2K2 Cr2 O7  2K2 CrO4 + Cr2 O3 + 2 O 2 3 On heatingwith alkalies, it is converted to chromate, i.e., the colour changes from orange to yellow. On acidifying, yellow colour again changes to orange. K2 Cr2 O7 + 2KOH  2K2 CrO4 + H2 O
  19. Orange - 2 7 2O Cr + 2OH-  Yellow - 2 4 CrO + H2 O Yellow - 2 4 CrO + 2H+  Orange - 2 7 2O Cr + H2 O In alkaline solution, chromate ions are present while in acidic solution, dichromate ions are present. Potassium dichromate reacts with hydrochloric acid and evolves chlorine. K2 Cr2 O7 + 14HCl  2KCl + 2CrCl3 + 7H2 O + 3Cl2 It acts as a powerful oxidisingagent in acidic medium(dilute H2 SO4 ). - 2 7 2O Cr +14H+ 6e-  2Cr3+ + 7H2 O The oxidation state of Cr changes from +6 to +3. Some typical oxidation reactions are given bleow: (i) Iodine is liberated from potassium iodide. K2 Cr2 O7 + 4H2 SO4  K2 SO4 + Cr2 (SO4 )3 + 4H2 O + 3[O] [2KI + H2 SO4 + [O]  K2 SO4 + I2 + H2 O] x 3 _______________________________________________ K2 Cr2 O7 + 6KI + 7H2 SO4  4K2 SO4 + Cr2 (SO4 )3 + 7H2 O + 3I2 The equation in terms of electron method may also be written as : - 2 7 2O Cr + 14H+ + 6e- 2Cr3+ + 3I2 + 7H2 O 6I-  3I2 + 6e- _________________________________ - 2 7 2O Cr + 14H+ +6I- 2Cr3+ + 3I2 + 7H2 O (ii) Ferrous salts are oxidised to ferric salts. K2 Cr2 O7 + 4H2 SO4  K2 SO4 + Cr2 (SO4 )3 + 4H2 O + 3[O] [2FeSO4 + H2 SO4 + [O]  Fe2 (SO4 )3 + H2 O] x 3 ________________________________________________________ K2 Cr2 O7 + 6FeSO4 + 7H2 SO4  3Fe2 (SO4 )3 + Cr2 (SO4 )3 + 7H2 O + K2 SO4 or 6Fe2+ + - 2 7 2O Cr + 14H+  6Fe3+ + 2Cr3+ + 7H2 O (iii) Sulphites are oxidised to sulphates. K2 Cr2 O7 + 4H2 SO4  K2 SO4 + Cr2 (SO4 )3 + 4H2 O + 3[O] [Na2 SO3 + [O]  Na2 SO4 ]x 3 K2 Cr2 O7 + Na2 SO3 + 4H2 SO4 3Na2 SO4 + K2 SO4 + Cr2 )SO4 )3 + 4H2 O or - 2 7 2O Cr + - 2 3 3SO + 8H+  - 2 4 3SO + 2Cr3+ + 4H2 O (iv) H2 S is oxidised to sulphur. K2 Cr2 O7 + 4H2 SO4 + 3H2 S  K2 SO4 + Cr2 (SO4 )3 + 7H2 O + 3S or - 2 7 2O Cr + 3H2 S + 8H+  2Cr3+ + 7H2 O + 3S (v) So2 is oxidised to H2 SO4 . K2 Cr2 O7 + 4H2 SO4  K2 SO4 + Cr2 (SO4 )3 + 4H2 O + 3[O] [SO2 + [O] + H2 O  H2 SO4 ] x 3 ________________________________________ K2 Cr2 O7 + H2 SO4 3SO2  K2 SO4 + Cr2 (SO4 )3 + H2 O or - 2 7 2O Cr + 3SO2 + 2H+ 2Cr3+ + - 2 4 3SO + H2 O When the solution is evaporated, chrome-alum is obtained.
  20. (vi) It oxidises ethyl alcohal to acetaldehyde and acetaldehyde to acetic acid. alcohol Ethyl 5 2 OH H C  [O] de Acetaldehy 3CHO CH  [O] de Acetaldehy 3COOH CH It alsooxidises nitrites to nitrates, arsenites to arsenates, thiosulphate to sulphate and sulphur (  2 3 2O S + O   2 4 2O S + S), HBr to Br2 , HI to I2 etc. Chromyl chloride test:This is a test of chloride. When a mixture ofa metal chloride and potassium dichromate is heated with conc. H2 SO4 , orange red vapours of chrymyl chloride are evolved. K2 Cr2 O7 + 2H2 SO4  2KHSO4 + 2CrO3 + H2 O [NaCl + H2 SO4  NaHSO4 + HCl] x 4 [CrO3 + 2HCl  CrO2 Cl2 + H2 O] x 2 _______________________________________________________ K2 Cr2 O7 + 6H2 SO4 + 4NaCl  2KHSO4 + 4NaHSO4 + chloride Chromyl 2 2Cl CrO + 3H2 O When chromylchloride vapoursare passed through NaOHsolution, yellowcoloured solution is obtained. 4NaOH + CrO2 Cl2  soln. Yellow 4 2CrO Na + 2NaCl + 2H2 O Uses: Potassium dichromate is used: (i)As a volumetric reagent in teh estimation ofreducing agents such as oxalic acid, ferrous ions, iodide ions, etc. It is used as a primary standrad. Catalytic Properties: Transition metals and their compounds have catalytic propetries in most of the reactions due to surface adsorption.
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