CHEMISTRY

HALOGEN ACIDS OR HYDRACIDS
Preparation (HCl, HBr and HI)
(i) By direct combinations of elements: Hydrogen chloride (HCl) is formed by burning chlorine in
excess of hydrogen.
H2
+ Cl2  2HCl
HBr can be prepared by passing a mixture of H2
and Br2
over a platinum spiral heated to redness by an
electric current.
H2
+ Br2 
Pt 2HBr
HI can also be prepared in small quantities by passing hydrogen and iodine vapours over red hot fine
platinum.
H2
+ I2
2HI
(ii) By heating a halide with acid: Hydrogen chloride is prepared by heating sodium chloride with
concentrated sulphuric acid.
2NaCl + H2
SO4  Na2
SO4
+ 2HCl
HBr and HI cannot be prepared by heating bromides and iodides with concentrated H2
SO4
because
HBr and HI are strong reducing agents and reduce H2
SO4
to SO2
and get themselves oxidised to
bromine and iodine, respectively.
KBr + H2
SO4  KHSO4
+ HBr
KI + H2
SO4 KHSO4
+ HI
H2
SO4
+ 2HBr  SO2
+ Br2
+ 2H2
O
H2
SO4
+ 2HI  SO2
+ I2
+ 2H2
O
Thus, HBr and HI are prepared by heating bromides and iodides respectivelywith conc. orthophosphoric
acid.
3KBr + H3
PO4 K3
PO4
+ 3HBr
KI + H3
PO4  K3
PO4
+ 3HI
(iii) HBr and HI can be prepared by hydorlysis of coresponding phosphorus trihalides.
2P + 3Br2  2PBr3
2P + 3I2  2PI3
PBr3
+ 3H2
O  H3
PO3
+ 3HBr
PI3
+ 3H2
O  H3
PO3
+ 3HI
Hydrogen fluoride : It is an associated molecule and is represented by formula, H2
F2
.
Pure anhydrous hydrogen fluoride is obtained by heating potassium hydrogen fluoride (dry) in a platinum
retort.
2KHF2  K2
F2
+ H2
F2
The Vapours are condensed in a platinum condenser and collected in a platinum receiver.
The aqueous solution of H2
F2
can be obtained byheating calcium fluoride with 90% conc. sulphuric acid in
a lead retort. The vapours are condensed in a strongly colled lead receiver. Aqueous acid of about 40%
strength can be obtained by passing the vapours through water in a lead receiver.
CaF2
+ H2
SO4  CaSO4
+ H2
F2

The aqueous solution of H2
F2
is kept in quttapercha bottles when dilute and in wax bottles when
concentrated. The anhydrous acid is stored in well cooled platinum, gold or silver vessels or steelcylinders.
Properties
(i) Physical state: Except H2
F2
, other hydrogen halides are gases. They fume in air and have pungent
odour. These are colourless liquids.
Their melting and boiling points increase with increase in atomic mass of the halogen. The low values are
due to covalent nature.
Order of M.P. & B.P. HCl > HBr > HI
H2
F2
is a liquid with boiling point 19.5o
C. This behaviour is due to association of HF molecules through
hydrogen bonding.
HFHFHFHF
(ii) Stability: The bond strength HX decreases fromHF to HI. Thus, HF is most stable while HI is least
stable. The decrease in stability is due to decrease in electronegativity from F to I. This is also observed in
the values of dissociation energy of HX bond.
Order of Dissociation energy HF >HCl > HBr >HI
HBr and HCl are stable upt 1200o
C, HBr dissociates slightly and HI dissociated considerably (20%) at
440o
C.
(iii) Acid strength: HF is only slightly ionised but HCl, HBr and HI are almost completelyionised, i.e., the
relative strength increases from HF to HI. HF is the weakest acid and HI is the strongest acid. On the
basis ofelectronegativity, this appears to rather surprising. [The electronegativitydifference suggests that
HF should be strongest and HI should be weakest acid.]
The lower strength of HF is largely due to its high bond dissociation energyas the bond length is very small
(1.01 A
).
(iv) Reducing nature: The reducing nature increases fromHF to HI as the stability decreases form HF to
HI. HF does not show reducing nature. It cannot be oxidised even by strong oxidising agents. HI is the
strongest reducing agent. Its aqueous solution gets oxidised even byatmospheric oxygen.
4HI + O2  2H2
O + 2I2
The reducing action can also be explained on the basis of increasing size of the halide ions from F
to I
.
The bigger ion can lose electron easily.
HCl can be oxidised by strong oxidising agents like MnO2
, KMnO4
, K2
Cr2
O7
, PbO2
, Pb3
O4
, etc.
HBr acts stronger reducing agent than HCl. It can be oxidised by H2
SO4
and atmospheric oxygen.
H2
SO4
+ 2HBr  SO2
+ Br2
+ 2H2
O
4HBr + O2  2H2
O + Br2
HI is the strongest reducing agent.
It reduces H2
SO4
to SO2
, S and H2
S, nitric acid to NO2
, nitrous acid to NO, FeCl3
to FeCl2
, cupric salt to
cuprous salt, etc.
H2
SO4
+ 2HI  SO2
+ I2
+ 2H2
O
H2
SO4
+ 6HI  S + 3I2
+ 4H2
O
H2
SO4
+ 8HI  H2
S + 4I2
+ 4H2
O
2HNO3
+ 2HI  2NO2
+ 2H2
O + I2
2HNO2
+ 2HI  2NO + 2H2
O + I2
2FeCl3
+ 2HI  2FeCl2
+ I2
+ 2HCl
2CuSO4
+ 4HI  Cu2
I2
+ 2H2
SO4
+ I2
Potassium peroxy K2
S2
O8
+ 2HI  K2
SO4
+ H2
SO4
+ I2
disulphate

(vi) Precipitation reactions: HCl forms insoluble chlorides with the solube salts of Ag, Pb and Hg(ous).
AgNO3
+ HCl  AgCl + HNO3
(white)
Pb(NO3
)2
+ 2HCl  PbCl2
+ 2HNO3
(white)
Hg2
(NO3
)2
+ 2HCl  Hg2
Cl2
+ 2HNO3
(white)
AgCl is insoluble in HNO3
but soluble in NH4
OH.PbCl2
is soluble in hot water while Hg2
Cl2
is soluble in
aquaregia.
HBr reacts with soluble salts of Ag and Pb.
AgNO3
+ HBr  AgBr + HNO3
(Pale yellow)]
Pb(NO3
)2
+ 2HBr  PbBr2
+ 2HNO3
(white)
AgBr is insoluble in HNO3
but sparingly soluble in NH4
OH. PbBr2
is soluble in hot water.
HI forms insoluble iodides with the salts of Ag, Pb and Hg (ic),
AgNO3
+ HI  AgI + HNO3
(Yellow)
Pb(CH3
COO)2
+ 2HI  PbI2
+ 2CH3
COOH
(Yellow)
HgCl2
+ 2HI  HgI2
+ 2HCl
(Scarlet)
AgI is insoluble in HNO3
and NH4
OH. PbI2
is soluble in hot water.
(vii) Action of halogens: F2
can displace Cl2
, Br2
and I2
from HCl, HBr and HI.
2HCl + F2  2HF + Cl2
2HBr + F2  2HF + Br2
2HI + F2  2HF + I2
Cl2
can displace Br2
and I2
from HBr and HI and Br2
can displace only I2
from HI. Iodine can displace
none.
(viii) Reaction with ammonia: HCl, HBr and HI combine with ammonia giving white fumes of ammonium
halides.
NH3
+ HCl  NH4
Cl (Ammonium chloride)
NH3
+ HBr  NH4
Br (Ammonium bromide)
NH3
+ HI  NH4
I (Ammonium iodide)
(ix) Acidic properties: Allthe three (HCl, HBr and HI) react with certain metals, their oxides, carbonates,
hydroxides, etc., and form salts.
Mg + 2HX  MgX2
+ H2
MgO + 2HX  MgX2
+ H2
CaCO3
+ 2HX  CaX2
+ H2
O + CO2
NaHCO3
+ HX  NaX + H2
O + CO2
NaOH + HX  NaX + H2
O
[HX = HCl, HBr or HI]
(x) Action of salts: HCl, HBr or HI decompose the salts of weaker acids.
Na2
S + 2HX  2NaX + H2
S

Ca3
P2
+ 6HX  3CaX2
+ 2PH3
Na2
SO3
+ 2HX  2NaX + H2
O + SO2
2NaNO2
+ 2HX  2NaX + NO + NO2
+ H2
O
OXIDES OF CHLORINE
Chlorine forms a number of oxides such as Cl2
O, ClO2
, Cl2
O6
and Cl2
O7
. All of them are unstable and
highly reactive.
(i) Dichloro oxide, Cl2
O: It is prepared by passing dry chlorine over freshly precipitated yellow mercuric
oxide.
HgO + 2Cl2
(dry)  HgCl2
+ Cl2
O
It is a brownishyellow gas which condenses to an orange coloured liquid in a freezing mixture (b. pt. =
2o
C). It has characteristic penetrating odour. Liquid Cl2
O explodes readily on heating or sparking Cl2
and
O2
. It also undergoes photochemical decomposition.
2Cl2
O  2Cl2
+ O2
It dissolves inn water giving golden yellow solution of hypochlorous acid. HClO. It is thus anhydride of
hypochlorous acid.
Cl2
O + H2
O 2HClO
Its gaseous mixture with ammonia explodes violently.
3Cl2
O + 10NH3  2N2
+ 6NH4
Cl + 3H2
O
it is a strong oxidising agent. It oxidises HCl to Cl2
.
Cl2
O + 2HCl  2Cl2
+ H2
O
Cl2
O has Vshaped structure. The oxygen atom undergoes sp3
hybridization. The Cl O  Cl bond angle
is 110.9o
C.
(ii) Chlorine dioxide, ClO2
: Pure ClO2
is obtained by passing dry Cl2
over AgClO3
heated to 90o
C.
2AgClO3
+ Cl2
(dry)  2AgCl + 2ClO2
+ O2
It can also be obtained by the action of Cl2
on sodium chlorite.
2NaClO2
+ Cl2  2NaCl + 2ClO2
It can be condensed by cooling to a coloured liquid (b.pt. 11o
C). The gas explodes and is decomposed to
Cl2
and O2
by an electric spark. It dissolves in water giving a mixture of chlorous acid and chloric acid.
2ClO2
+ H2
O  HClO2
+ HClO3
With alkalies, it gives a mixture ofchlorite and chlorate.
2ClO2
+ 2KOH  KClO2
+ KClO3
+ H2
O
(iii) Dichlorine hexoxide, Cl2
O6
: It is obtained by mixing chlorine dioxide with ozonised air at 0o
C.
2ClO2
+ 2O3  Cl2
O6
+ 2O2
It is a dark red liquid. It is unstable and decomposed into ClO2
and O2
. It is a strong oxidising agent. It
reacts with water forming chloric and perchloric acids.
Cl2
O6
+ H2
O  HClO3
+ HClO4
It reacts with alkalies to give chlorates and perchlorates.
Cl2
O6
+ 2KOH  KClO3
+ KClO4
+ H2
O
With HF, it gives HClO4
and chloryl fluride.
Cl2
O6
+ HF  HClO4
+ ClO2
F
In the vapour state it exists as ClO3
molecule which has odd number of electrons and hence paramagnetic
while in liquid state it is in dimeric form having even number of electrons and hence diamagnetic innature.
2ClO3
Cl2
O6
The exact structure is unknown. However, the following structure containing ClCl linkage is proposed in
which each chlorine is sp3
hybridized.

(iv) Chlorine heptoxide, Cl2
O7
: It is formed by dehydration of perchloric acid with P2
O5
at 10o
C.
2HClO4 
 
 5
2O
P Cl2
O7
+ H2
O
It is a colourless oily liquid which is explosive in nature. It slowly reacts with water forming perchloric acid.
It is, thus, anhydride of perchloric acid.
Cl2
O7
+ H2
O  2HClO4
It is less reactive in comparison to lower oxides of chlorine. Cl2
O7
has the structure O3
Cl  O  ClO3
in
which two tetrahedral are sharing one Oatom.
OXYACIDS OF CHLORINE
(i) Hypochlorous acid, HClO: The acid is known only in solution. It is obtained by shaking precipitated
HgO with chlorine water.
2HgO + 2Cl2
+ H2
O  Hg2
OCl2
+ 2HClO
(Oxychloride of mercury)
Commerciall, it is obtained by passing CO2
through suspension of bleaching powder and then distilling.
2CaOCl2
+ H2
O + CO2  CaCl2
+ CaCO3
+ 2HClO
It is a weak acid. Its concentrated solution is yellow in colour while dilute solution is colourless. It is
unstable and decomposes.
2HClO  2HCl + O2
It dissolves magnesiumwith evilution of hydrogen.
Mg + 2HClO  Mg(ClO)2
+ H2
With alkalies, it forms salts called hypochlorites.
It acts as a powerful oxidising and bleaching agent. This is due to release of nascent oxygen easily.
HClO  HCl + O
Bleaching Powder, CaOCl2
.H2
O
Bleaching powder is also called calcium chlorohypochlorite because it is considered as a mixed salt of
hydrochloric acid and hypochlorous acid. It is represented as:
(Odling formula)
It is manufactured by the action of chlorine on dry slaked lime, Ca(OH)2
, at 40o
C.
Ca(OH)2
+ Cl2  Ca(OCl)Cl + H2
O
This is the Odling view about its formation. There is another view proposed by Clifford according towhich
bleaching powder is a mixture of calcium hypochlorite and basic calcium chloride.
2Ca(OH)2
+ 2Cl2  Ca(OCl)2
+ CaCl2
+ 2H2
O
O
H
O
H
)
OH
(
Ca
CaCl
)
OCl
(
Ca
)
OH
(
Ca
3
Cl
2
O
H
)
OH
(
Ca
.
CaCl
O
H
)
OH
(
Ca
CaCl
2
powder
Bleaching
2
2
2
2
2
2
2
2
2
2
2
2












 





 

The manufacture of bleaching powder is carried ot in any one ofthe following plants:
(i) Hasenclever’s plant (Old process),

(ii) Backmann’s plant (Modern process).
Properties: (a) It is a pale yellow powder. It has a strong smell of chlorine. It is soluble in water but a
clear solution is never formed due to the presence of impurities.
(b) On long standing, it undergoes autooxidation into calciumchlorate and calcium chloride.
6CaOCl2  Ca(ClO3
)2
+ 5CaCl2
(c) In presence of cobalt chloride, CoCl2
, it loses its oxygen.
2CaOCl2 
 
 2
CoCl 2CaCl2
+ O2
(d) In presence of a slight amount of a dilute acid, it loses oxygen.
2CaOCl2
+ H2
CO4  CaCl2
+ CaSO4
+ 2HClO
HClO  HCl + O
On account ofthe formation of nascent oxygen, it shows oxidising and bleaching properties.
(i) Oxidising properties:
CaOCl2
+ H2
S  CaCl2
+ H2
O + S
CaOCl2
+ 2FeSO4
+ H2
SO4  Fe2
(SO4
)3
+ CaCl2
+ H2
O
CaOCl2
+ KNO2  CaCl2
+ KNO3
3CaOCl2
+ 2NH3  3CaCl2
+ 3H2
O + N2
CaOCl2
+ 2KI + 2HCl  CaCl2
+ 2KCl + H2
O + I2
Na3
AsO3
+ CaOCl2  CaCl2
+ Na3
AsO4
(ii) Bleaching action:
Coloured matter + [O]  Colourless product
(e) It loses its chlorine by the action of dilute acids (in excess) or carbon dioxide.
CaOCl2
+ 2HCl  CaCl2
+ H2
O + Cl2
CaOCl2
+ H2
SO4  CaSO4
+ H2
O + Cl2
CaOCl2
+ CO2  CaCO3
+Cl2
The amount of chlorine obtained from a sample of bleaching powder by treatment with excess of
dilute acids or carbon dioxide is called available chlorine. Agood sample of bleaching powder contains
3538% of available chlorine.
(f) Bleaching powder converts acetone or ethylalcohol into chloroform.
CaOCl2
+ H2
O Ca(OH)2
+ Cl2
Uses of bleaching powder: It is used (i) as a disinfectant and germicide especially in the sterlisation of
drinking water. (ii) for manufacture of chloroform. (iii) for making wool unshrinkable. (iv) as an oxidising
agent in industry. (v) mainly as bleaching agent for cotton, linen and wood pulp. However, delicatearticles
like straw, silk, ivory, etc., are not bleached by bleaching powder.
Bleaching process: In textile industry, the cotton cloth is mainly bleached with the help of bleaching
powder. The following procedure is followed.
(i) Cloth is first made to pass through a very dilute hot solution of caustic soda as to remove the greasy
matter.
(ii) It is now passed through a bath containing bleaching powder solution.
(iii) The cloth is next passed through a bath containing a very dilute solution of HCl. Nascent oxygen is
given out which bleaches the cloth. Some chlorine remains sticking to the fibre and its removal is necessary

otherwise the cloth is likely to be damaged.
(iv) For the removalofchlorine, the cloth is passed through a bath containing the solutionofhypo(antichlor).
Na2
S2
O3
+ Cl2
+ H2
O Na2
SO4
+ 2HCl + S
(v) Finally the cloth is washed in running water and thus the club is freed from all the sticking chemical
substances.
Chlorous Acid, HClO2
: It is obtained in aqueous solution when barium chlorite suspension in water is
treated with H2
SO4
. The insoluble barium sulphate is filtered off.
Ba(ClO2
)2
+ H2
SO4  BaSO4
+ 2HClO2
The freshly prepared solution is colourless but it soon decomposes to ClO2
which makes the solution
yellow.
5HClO2  4ClO2
+ HCl + 2H2
O
The acid liberates iodine from KI
4KI + HClO2
+ 2H2
O  4KOH + HCl + 2I2
Chloric Acid, HClO3
: This acid is only known in solution. The acid is prepared by the action of dilute
H2
SO4
on barium chlorate.
Ba(ClO3
)2
+ H2
SO4  BaSO4
+ 2HClO3
BaSO4
is filtered off.
Concentrated acid is colourless and pungent smelling liquid. It decomposes in light. However, it isstable in
dark. It acts as a strong oxidising and bleaching agent in light. Organic substances like paper, cotton, wool,
etc. catch fire in contact with the acid.
Perchloric acid, HClO4
: It is the most stable oxy acid of chlorine. Anhydrous HClO4
is obtained by
doing distillation of KClO4
, potassium perchlorate, with 9697.5% H2
SO4
under low pressure at 90
160o
C.
KClO4
+ H2
SO4 KHSO4
+ HClO4
An aqueous solution of the acid is obtained by reaction Ba(ClO4
)2
, barium perchlorate with calculated
quantity of dilute H2
SO4
. The insoluble barium sulphate is removed by filtration.
Ba(ClO4
)2
+ H2
SO4  BaSO4
+ 2HClO4
Anhydrous HClO4
is a colourless oily liquid which turns dark on keeping. It fumes in moist air. It is one of
the strongest acid and ionises as follows:
HClO4
H+
+ ClO4

It dissolves most of the metals.
Zn + 2HClO4  Zn(ClO4
)2
+ H2
It is an oxidising agent and explodes in contact with organic matter.When dehydrated with P2
O5
, it yields,
Cl2
O7
.
2HClO4
+ P2
O5  2HPO3
+ Cl2
O7
FLUORINE
Occurence: Fluorine is extremely reactive and hence does not occur in free state. It is widelyfound in the
combined state as fluorides. The important minerals are:
(i) Feldspar (Fluorite) , CaF2
(ii) Cryolite , Na3
AlF6
(iii) Fluorapatite , CaF2
·3Ca3
(PO4
)2
Isolation
In modern methods, fluorine is prepared by electrolysis of a fused fluoride (usually potassium hydrogen

fluoride, KHF2
). The electrolytic cells are made of copper, nickel or monel metal. The anode is generally of
graphite and the fluorine set free contains some carbon tetrafluoride.
Reactions in the electrolytic cell
KHF2
 KF + HF
K+
+ e  K F–
 F + e
K + HF  KF + H 2F  F2
2H  H2
Dennis Method
Whytlaw Gray Method
Properties: (a) The gas is pale greenish yellow in colour. It can be condensed to yellow liquid at –1880
C
and pale yellow solid at –2230
C. (b) It has pungent odour resembling that of a mixture of ozone and
chlorine. (c) It is a poisonous gas but less poisonous than HF gas. (d) It is the most reactive element.
(i) It can displace all the halogens from their halides
F2
+ 2NaX  2NaF + X2
(X = Cl–
, Br–
or I–
) (X2
= Cl2
, Br2
or I2
)
(ii) It combines with most of the metals and non-metals to form fluorides.
2Ag + F2
 2AgF
2Al + 3F2
 2AlF3
Mg + F2
 MgF2
Hydrogen explodes violently in fluorine even in the dark.
H2
+ F2
 H2
F2
A mixture of oxygen and fluorine explodes in presence ofsilent electric discharge.
O2
+ F2
 O2
F2
(Dioxygen difluorides)
It reacts with wood, charcoal, phosphorus, arsenic, antimony, boron, silicon producing a flame.
C + 2F2
 CF4
; Si + 2F2
 SiF4
2P + 3F2
 2PF3
; 2As + F2
 2AsF3
2B + 3F2
 2BF3
However, fluorine does not react with oxygen, nitrogen and inert gases directly under ordinary
conditions.
(iii) Fluorine reacts vigorouslywith water giving O2
and O3
.
2H2
O + 2F2
 4HF + O2
3H2
O + 3F2
 6HF + O3
(iv) Fluorine reacts with dilute alkalies to form oxygen difluorides (OF2
) and with conc. alkalies O2
is
formed.
2F2
+ NaOH  OF2
+ 2NaF + H2
O
2F2
+ 4NaOH  4NaF + 2H2
O + O2
(v) It acts as a strong oxidising agent. It oxidises chlorates to perchlorates, iodates to periodates and
bisulphates to peroxysulphates.
KClO3
+ F2
+ H2
O  KClO4
+ H2
F2
KIO3
+ F2
+ H2
O  KIO4
+ H2
F2
2NaHSO4
+ F2
 a2
S2
O8
+ 2HF
(vi) It attacks glass at about 1000
C forming SiF4
. However, the attack of dry fluorine is slow. At low
temperatures, dry fluorine can be stored in dry galss vessel.

(vii) It reacts with hydrocarbons explosively and forms fluorinated hydrocarbons.
CH4
F2
 
 CH3
F F2
 
 CH2
F2
F2
 
 CHF3
F2
 
 CF4
Fluorine reacts with NH3
to form nitrogen and H2
S forming SF6
(H2
S burns)
2NH3
+ 3F2
 N2
+ 6HF
H2
S + 4F2
 SF6
+ 2HF
(viii) It combines with xenon to form XeF2
, XeF4
, XeF6
(ix) Fluorination ofmethanol
CH3
OH + 2F2
AgF2
 

 CF3
OF + 4HF
CF3
OF is verystrong oxidising agent.
CHLORINE
Occurrence: Chlorine is widely distributed in nature in the combined state in the form of chlorides of
various metals. Common salt (NaCl) is the most important chloride which occurs in seawater, lakes and in
rocks. The other important chloride minerals are:
(i) Sylvine (potassium chloride), KCl
(ii) Carnallite, KCl, MgCl2
.6H2
O
(iii) Chlorapatite, 3Ca3
(PO4
)2
.CaCl2
(iv) Horn silver, AgCl
Preparation: (i) Bythe oxidation of hydrochloric acid. The oxidising agents which can be used are MnO2
,
PbO2
, Pb3
O4
(red lead), KMnO4
(Potassium permanganate), K2
Cr2
O7
(Potassium dechromate), NaClO
(Sodium hypochlorite), O3
, etc.
(a) MnO2
+ 4HCl  MnCl2
+ 2H2
O + Cl2
(b) PbO2
+ 4HCl  PbCl2
+ 2H2
O + Cl2
(c) Pb3
O4
+ 8HCl  3PbCl2
+ 4H2
O + Cl2
(d) 2KMnO4
+ 16HCl  2KCl + 2MnCl2
+ 8H2
O + 5Cl2
(e) K2
Cr2
O7
+ 14HCl  2KCl + 2CrCl3
+ 7H2
O + 3Cl2
(f) NaClO + 2HCl  NaCl + H2
O + Cl2
(g) O3
+ 2HCl  O2
+ H2
O + Cl2
(ii) By the action of mineral acids or carbon dioxide on bleaching powder.
CaOCl2
+ H2
SO4  CaSO4
+ H2
O + Cl2
CaOCl2
+ 2HCl  CaCl2
+ H2
O + Cl2
CaOCl2
+ CO2  O2
+ H2
O + Cl2
(iii) By heating anychloride with concentrated H2
SO4
in presence of MnO2
.
2NaCl + MnO2
+ 3H2
SO4  2NaHSO4
+ MnSO4
+ 2H2
O + Cl2
Pure chlorine: It may be obtained by heating dry platinic chloride (PtCl4
) or gold chloride (AuCl3
) in a
hard glass tube.
PtCl4 
 
 C
374o
PtCl2
+ Cl2 
 
 C
582o
Pt + 2Cl2
2AuCl3 
 
 C
175o
2AuCl + Cl2 
 
 C
185o
2Au + 3Cl2
Manufacture: Bulk of chlorine required for commercial purposes is obtained as a by-product in the
manufacture ofcaustic soda byelectrolysis ofbrine solution or in the manufacture ofsodium by electrolysis
of fused sodium chloride.
The other methods used for its manufacture are:
(i) Weldon’s process
(ii) deacon’s process
(iii) Nitrosylchloride process
These chemical methods are bit costlier and only the nitrosyl chloride process is used these days to
manufacture chlorine.

1. Weldon’s process: This process involves the heating of pyrolusite mineral (MnO2
) with concentrated
hydrochloric acid in stone-ware stills.
MnCl2
+ Ca(OH)2  Mn(OH)2
+ CaCl2
2Mn(OH)2
+ O2
(Air)  2MnO2
+ 2H2
O
MnO2
+ Ca(OH)2  manganite
Calcium
3
CaMnO + H2
O
Calcium manganite settles down to the bottom as a dark coloured mud. This is used for oxidation offresh
quantities ofhydrochloric acid.
CaMnO3
+ 6HCl  CaCl2
+ MnCl2
+ Cl2
+ 3H2
O
(ii) Deacon’s process: This process is superior to Weldon’s process. In this process, chlorine is obtained
by the oxidation of HCl with air in presence of cupric chloride heated to 400o
C.
4HCl + O2 C
400
CaCl
o
2

 
 2Cl2
+ 2H2
O
The catalytic action of cupric chloride can be explained with the following mechanism.
2CuCl2 

 
 .
temp
High Cu2
Cl2
+ Cl2
2Cu2
Cl2
+ O2
(air)  e
oxychlorid
Copper
2
2OCl
Cu
2
Cu2
OCl2
+ 2HCl  2CuCl2
+ H2
O
(iii) Nitrosyl chloride process: In this process, common salt (NaCl) is treated with concentrated nitric
acid, when a mixture of chlorine and nitrosylchloride is evolved.
[NaCl + HNO3  NaNO3
+ HCl] x 3
O
H
2
Cl
NOCl
NaNO
3
HNO
4
NaCl
3
O
H
2
Cl
NOCl
HCl
3
HNO
2
2
3
3
2
2
3









The gaseous mixture is oxidised with oxygen.
2NOCl + O2  2NO2
+ Cl2
The gases are liquefied and chlorine distilled out. Nitrogen dioxide is absorbed in water in presence of
oxygen to form nitric acid which can be used again.
4NO2
+ O2
+ 2H2
O  4HNO3
Properties of chlorine
Physical:(a) Chlorine is a yellowish-green gas with a pungent suffocating smell.
Chemical
(a) Combustibility: Chlorine does not burn and also does not help in burning.
(b) Combination with elements: It combines with allthe nonmetals except nitrogen, oxygen, carbon and
inert gases.
It combines with hydrogen in light (but not in dark) with explosion to form HCl. However, in presence of
charcoal as a catalyst, the reaction is safe.
H2
+ Cl2  2HCl
Arsenic, boron, sulphur, phosphorus (red), etc., combine with chlorine when heated.
2As + 3Cl2  e
trichlorid
Arsenic
3
AsCl
2
2B + 3Cl2  e
trichlorid
Boron
3
BCl
2
2S + 3Cl2  dichloride
Disulphur
2
2Cl
S

)
d
(Re
P
2 + 3Cl2  e
trichlorid
Phosphorus
3
PCl
2
Yellow phosphorus readily catches fire in chlorine.
2Na + Cl2  2NaCl : 2Fe + 3Cl2  2FeCl3
2Al + 3Cl2  )
Cl
Al
or
(
3
6
2
AlCl
2 : Cu + Cl2  CuCl2
Mg + Cl2  MgCl2
: Zn + Cl2  ZnCl2
(c) Affinity for hydrogen: It has great affinity for hydrogen. It decopses several hydrogen compounds
forming HCl.
(i) Turpentine burns in chlorine forming HCl and carbon.
C10
H16
+ 8Cl2  10C + 16HCl
(ii) It reacts with water in the presence of sunlight and forms O2
.
2H2
O + 2Cl2  4HCl + O2
Chlorine water loses its yellow colour on exposure to sunlight. This is due to the formation of HCl.
(iii) It reacts with ammonia under two conditions:
1. When ammonia in excess: The products are nitrogen and ammonium chloride.
8NH3
+ 3Cl2  N2
+ 6NH4
Cl
2. When chlorine in excess: The products are nitrogen trichloride and HCl.
NH3
+ 3Cl2  NCl3
+ 3HCl
(iv) It removes hydrogen from H2
S
H2
S + Cl2  2HCl + S
(d) Oxidising and bleaching nature: In presence ofmoisture, chlorine acts as an oxidising and a bleaching
agent. Chlorine reacts with water forming HCl and HClO (hypochlorous acid). HClO is not so stable and
decoposes giving nascent oxygen which is responsible for oxidising and bleaching properties of chlorine.
Cl2
+ H2
O  HCl + HclO
O
HCl
2
O
H
Cl
O
HCl
HClO
2
2 




Coloured matter + Nascent oxygen  Colourless matter
The bleaching action is permanent. The delicate articles should not be bleached with chlorine.
1. It displaces bromine and iodine from bromides and iodides respectively.
2KBr + Cl2  2KCl + Br2
2KI + Cl2  2KCl + I2
2. Ferrous chloride/ferrous sulphate containing dilute H2
SO4
, is converted to ferric chloride/ferric sulphate.
2FeCl2
+ Cl2  2FeCl3
2FeSO4
+ H2
SO4
+ Cl2  Fe2
(SO4
)3
+ 2HCl
3. Moist sulphur dioxide is oxidised to H2
SO4
.
SO2
+ 2H2
O + Cl2  H2
SO4
+ 2HCl
4. Sulphites are oxidised to sulphates.
Na2
SO3
+ H2
O + Cl2  Na2
SO4
+ 2HCl
5. Sulphur is precipitated from sodium thiosulphate.
Na2
S2
O3
+ H2
O + Cl2  Na2
SO4
+ 2HCl + S
(e) Action of alkalies: Whenpassed into cold dilute solution ofalkalies, it forms chloridesand hypochlorites.
2NaOH + Cl2  NaCl + NaClO + H2
O
2KOH + Cl2  KCl + KClO + H2
O
dilute
and
Cold
2
)
OH
(
Ca
2 + 2Cl2  CaCl2
+ Ca(ClO)2
+ 2H2
O
With hot concentrated alkali solutions, chlorates are formed.

6NaOH + 3Cl2  5NaCl + NaClO3
+ 3H2
6KOH + 3Cl2  5KCl + KClO3
+ 3H2
O
.
conc
and
Hot
2
)
OH
(
Ca
6 + 6Cl2  CaCl2
+ Ca(ClO3
)2
+ 6H2
O
However, when chlorine is passed over dry slaked lime, bleaching powder is obtained.
e
lim
Slaked
2
)
OH
(
Ca + Cl2  powder
Bleaching
2
2 O
H
.
CaOCl
(f) Additionm reactions: Chlorine forms addition compounds with SO2
, CO and unsaturated hydrocarbons.
SO2
+ Cl2  chloride
Sulphuryl
2
2Cl
SO : CO + Cl2 
)
Phosgene
(
chloride
Carbonyl
2
COCl
Phosgene is highly poisonous gas.
C2
H4
+ Cl2  chloride
Ethylene
2
4
2 Cl
H
C
BROMINE
Occurrence: It is neve found in free state in nature. The chief sources of bromine are:
(i) Sea water contains bromides such as NaBr, KBr, MgBr2
, CaBr2
, etc., in minute amounts.
Preparation: The following reactions can be used for its preparation.
(i) Bypassing chlorine through the solution of potassium or magnesium bromide.
2KBr + Cl2  2KCl + Br2
MgBr2
+ Cl2  MgCl2
+ Br2
(ii) By adding hydrochloric acid to the mixture containing potassium bromide and potassium bromate.
5KBr + KBrO3
+ 6HCl  6KCl + 3Br2
+ 3H2
O
(iii) By heating mixture of potassium bromide and manganese dioxide with conc. H2
SO4
.
MnO2
+ H2
SO4  MnSO4
+ H2
O + O
[KBr + H2
SO4  KHSO4
+ HBr] x 2
O
H
2
Br
MnSO
KHSO
2
SO
H
3
MnO
KBr
2
Br
O
H
O
HBr
2
2
2
4
4
4
2
2
2
2









This reaction is commonly used in the laboratory for the preparation of bromine.
Manufacture
1. From carnallite mother-liquor
2. From sea water
Properties
Physical: (a) It is a reddish brown, heavy mobile liquid.
Chemical: In chemical behaviour, bromine closely resembles chlorine.
(a) Combination with elements:
(i) Bromine reacts directly with hydrogen, though less readily than chlorine. The reaction,
H2
+ Br2
2HBr
is slightlyreversible. It combines with hydrogen at 200o
C or at room temperature in presence of platinum
as catalyst.
(ii) It combines with sulphur, phosphorus, arsenic, etc.
2P + 3Br2  2PBr3
2As + 3Br2  2AsBr3
2B + 3Br2 

 
 heat
d
Re 2BBr3

(iii) It combines directly with many metals.
2Na + Br2  2NaBr (Slow reaction)
2K + Br2  2KBr (Fast reaction)
2Al + 3Br2  2AlBr3
(Violent reaction)
Zinc reacts in presence of moisture with bromine.
Zn + Br2  ZnBr2
(b) Oxidising action: Bromine is a good oxidising agent. It ordinarily does not react with water but in
presence of an oxidisable substance it forms hydrogen bromide and oxygen with water.
Br2
+ H2
O  2HBr + O
ce
tan
subs
Oxidisable
X + O  XO
(i) It oxidises sO2
into sulphuric acid
SO2
+ Br2
+ 2H2
O  2HBr + H2
SO4
]
(ii) It oxidises sulphites into sulphates
Na2
SO3
+ Br2
+ H2
O  Na2
SO4
+ 2HBr
(iii) It oxidises arsenites into arsenates.
Na3
AsO3
+ Br2
+ H2
O  Na3
AsO4
+ 2HBr
(iv) It oxidises thiosulphate into sulphate
Na2
S2
O3
+ Br2
+ H2
O  Na2
SO4
+ S + 2HBr
(v) Sulphur is formed by oxidation of hydrogen sulphide.
H2
S + Br2  2HBr + S
(vi) Iodine is liberated from iodides.
2KI + Br2  2KBr + I2
(c) Action with alkalies: The actiion of bromine on alkalies is similar to that of chlorine, a hypobromite
with a cold dilute solution and a bromate with a hot solution are formed.
)
cold
and
.
Dil
(
KOH
2 + Br2  KBr + e
hypobromit
.
Pot
KBrO + H2
O
)
hot
and
.
Conc
(
KOH
6 + 3Br2  5KBr + bromate
.
Pot
3
KBrO + 3H2
O
)
hot
and
.
Conc
(
2
)
OH
(
Ba
6 + 6Br2  5BaBr2
+ bromate
Barium
2
3)
BrO
(
Ba + 6H2
O
)
hot
and
.
Conc
(
3
2CO
Na
3 + 3Br2  5NaBr + NaBrO3
+ 3CO2
However, with slaked lime (dry) it forms a compound similar to bleaching powder.
)
dry
(
e
lim
Slaked
2
)
OH
(
Ca
+ Br2  CaOBr2
+ H2
O
(d) Action of ammonia: It is similar to that of chlorine.
8NH3
+ 3Br2  6NH4
Br + N2
(e) Action of organic compounds:It forms addition compoundswithunsturated compounds and substituted
products with the saturated compounds.
C2
H4
+ Br2  C2
H4
Br2
; CH4
+ Br2  CH3
Br + HBr
(f) Bleaching action: The bleaching actionn of bromine is weaker than that of chlorine. The bleaching
action is due to its oxidising property.
(g) Action of mercuric oxide: Bromine water reacts with mercuric oxide to formmercury oxybromide.
2HgO + 2Br2
+ H2
O 
oxybromide
Mercury
2 HgO
.
HgBr + 2HBrO

However, on passing bromine vapours over dry HgO at 60o
C, bromine monoxide is obtained.
HgO + 2Br2  Br2
O + HgBr2
Uses of bromine: (i) The chiefuse of bromine is in the manufacture of ethylene bromide (C2
H4
Br2
) which
is extensively employed as a component of anti-knock gasoline.
(ii) It is used in the manufacture of NaBr and KBr (used in medicine as sedative), and certain organic dyes.
(iii) Bromine water is used in organic chemistry to test the unsturation of organic compounds.
(iv) It is used as germicide.
(v) It is employed as an oxidising agent and in organic synthesis.
IODINE
Source: Caliche or crude chile saltpetre which contains about 0.2% of NaIO3
(sodium iodate).
Preparation: The following reactions can be applied for the preparation of iodine:
(i) Bypassing chlorine through potassiumiodide solution.
2KI + Cl2  2KCl + I2
(ii) By adding potassium iodide solution to copper sulphate solution.
CuSO4
+ 2KI  CuI2
+ K2
SO4
2CuI2
 Cu2
I2
+ I2
Cupric Cuprous
iodide iodide
(iii) By adding conc. sulphuric acid to a mixture of sodium/potassiumiodide and sodium/potassium
iodate.
NaIO3
+ 5NaI + 6H2
SO4
 6NaHSO4
+ 3H2
O + 3I2
(iv) By adding an oxidising agent such as hydrogen peroxide, ozone, acidified potassium permanganate
or acidified potassiumdichromate to potassium iodide.
2KI + H2
O2
 2KOH + I2
(v) Byreacting sodium iodate with sodium hydrogen sulphite.
2NaIO3
+ 5NaHSO3
 3NaHSO4
+ 2Na2
SO4
+ I2
+ H2
O
(vi) Byheating mixture of potassiumiodide and manganese dioxide with conc. H2
SO4
.
2KI + MnO2
+ 3H2
SO4
 2KHSO4
+ MnSO4
+ 2H2
O + I2
Manufacture:
(i) From Caliche: 2NaIO3
+ 5NaHSO3
*  3NaHSO4
+ 2Na2
SO4
+ H2
O + I2
*NaHSO3
if present in excess, HI is produced, NaHSO3
+ I2
+ H2
O  NaHSO4
+ 2HI
Chemical properties: Iodine is chemicallyless reactive in comparison to chlorine and bromine.
(a) Combination with element: It combines directly with many elements.
H2
+ I2
Heat
Pt
 

 2HI
2P + 3I2
Heat
 

 2PI3
2KI + I2
 2KI
2Fe + 3I2
Heat
 

 2FeI3
(b) Oxidising properties: It is a weak oxidising agent. The following reactions show its oxidising
nature.
(i) H2
S + I2
 2HI + S
(ii) SO2
+ I2
+ 2H2
O  H2
SO4
+ 2HI
(iii) Na2
SO3
+ I2
+ H2
O  Na2
SO4
+ 2HI
(iv) Na3
AsO3
+ I2
+ H2
O  Na3
AsO4
+ 2HI
(v) 2FeSO4
+ H2
SO4
+ I2
 Fe2
(SO4
)3
+ 2HI

(c) Reaction with alkalies: Iodine reacts with cold dilute solution of NaOH forming hypoiodite. The
colour ofthe solution becomes faint yellow. However, hypoiodite undergoes hydrolysis forming HIO.
2NaOH + I2
 NaIO + NaI + H2
O
NaIO + H2
O  NaOH + HIO
–––––––––––––––––––––––––––––––
NaOH + I2
 NaI + HIO
Hypoiodous acid
Iodine when heated with concentrated alkali solution, iodide and iodate are formed. The colour of iodine
disappears.
6NaOH + 3I2
 5 3
NaI NaIO
colourless

 
 
 
+ 3H2
O
(d) Iodine does not displace chlorine and bromine from clorides and bromides, respectively but it
displaces them from their oxy salts.
2KCIO3
+ I2
 2KIO3
+ Cl2
2KBrO3
+ I2
 2KIO3
+ Br2
(e) Reaction with sodium thiosulphate: Iodine solution is decolourised by sodium thiosulphate (hypo)
as sodiumtetrathionate and sodium iodide are formed. Both are colourless and soluble.
2Na2
S2
O3
+ I2
 2NaI + Na2
S4
O6
Sodium thiosulphate Sodium tetrathionate
This reaction is involved in iodometric titrations in presence ofstarch indicator.
(f) Action of NH3
: Iodine forms a mild explosive with ammonia.
2NH3
+ 3I2
 NI3
·NH3
+ 3HI
Explosive
8NI3
·NH3
 5N2
+9I2
+ 6NH4
I
(g) Reaction with strong oxidising agents: Strong oxidising agents convert iodine to iodic acid (HIO3
).
(i) with conc. HNO3
I2
+ 10HNO3
 2HIO3
+ 10NO2
+ 4H2
O
(ii) with ozone
I2
+ H2
O + 5O3
 2HIO3
+ 5O2
(iii) with chlorine water
I2
+ 5Cl2
+ 6H2
O  2HIO3
+ 10HCl
(h) Reaction with freshly precipitated mercuric oxide: Iodine forms hypoiodous acid with mercuric
oxide
2HgO + 2I2
+ H2
O  HgI2
·HgO + 2HIO
(i) Iodine when warmed with alcohol and alkali forms iodoform.
C2
H5
OH + 6KOH + 4I2
 CHI3
+ HCOOK + 5KI + + 5H2
O
(j) Reaction with chlorine: When chlorine is passed over iodine, at first iodine monochloride, ICI
(dark liquid) and finally iodine trichloride, ICI3
(yellow solid) is formed.
(k) Reaction with starch: Iodine turns starch solution blue. This is one of the tests of iodine.
(l) Action on unsaturated organic compounds: Addition products are formed with unsaturated organic
compounds.
C2
H4
+ I2
 C2
H4
I2
Uses of Iodine: (a) It is used in medicine, e.g. tincture of iodine [½OZ. iodine, 1/4OZ. KI and 1 pint
rectified spirit] and iodex as antiseptic and analgesic. Solutions of iodine in KI are used in the treatment of
goitre.
* * * * *