INORGANIC CHEMISTRY

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

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

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.

(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)

(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

(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.

(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

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.

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

(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.

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

(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

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

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
).

(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

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

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

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.

(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.