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Chemistry · Ch 7 — Elements of Groups 16, 17 and 18

Chlorine

7.11.1

Chlorine

Chlorine was discovered by Scheele, a German Swedish chemist, in 1774, by the action of HCl on MnO2\mathrm{MnO_2}. In 1810 Davy established its elementary nature and suggested the name chlorine on account of its colour (Greek, Chloros = greenish yellow).

a. Preparation :

1. Chlorine can be prepared by the oxidation of hydrochloric acid with any of the following oxidising agents:

i. Manganese dioxide :

MnO2+4HCl⟶MnCl2+Cl2+2H2O\mathrm{MnO_2 + 4HCl \longrightarrow MnCl_2 + Cl_2 + 2H_2O}

ii. Potassium permanganate :

2KMnO4+16HCl⟶2KCl+2MnCl2+8H2O+5Cl2\mathrm{2KMnO_4 + 16HCl \longrightarrow 2KCl + 2MnCl_2 + 8H_2O + 5Cl_2}

2. Chlorine can also be prepared by the action of concentrated sulfuric acid on a mixture of sodium chloride (common salt) and manganese dioxide. The reaction takes place in two steps:

4NaCl+4H2SO4⟶4NaHSO4+4HCl\mathrm{4NaCl + 4H_2SO_4 \longrightarrow 4NaHSO_4 + 4HCl}

MnO2+4HCl⟶MnCl2+2H2O+Cl2\mathrm{MnO_2 + 4HCl \longrightarrow MnCl_2 + 2H_2O + Cl_2}

Overall:

4NaCl+MnO2+4H2SO4⟶4NaHSO4+MnCl2+2H2O+Cl2\mathrm{4NaCl + MnO_2 + 4H_2SO_4 \longrightarrow 4NaHSO_4 + MnCl_2 + 2H_2O + Cl_2}

b. Manufacture of chlorine :

i. Deacon's process : Chlorine is manufactured by the oxidation of hydrogen chloride gas by atmospheric oxygen, in the presence of CuCl2\mathrm{CuCl_2} as catalyst, at 723 K.

4HCl+O2→CuCl22Cl2+2H2O\mathrm{4HCl + O_2 \xrightarrow{CuCl_2} 2Cl_2 + 2H_2O}

ii. Electrolytic process : By the electrolysis of brine (concentrated NaCl solution), chlorine is liberated at the anode.

NaCl⇌Na++Cl−\mathrm{NaCl \rightleftharpoons Na^{+} + Cl^{-}}

Cathode: 2H2O+2e−⟶H2+2OH− ;Na++OH−⟶NaOH\text{Cathode: } \mathrm{2H_2O + 2e^{-} \longrightarrow H_2 + 2OH^{-}}\,; \qquad \mathrm{Na^{+} + OH^{-} \longrightarrow NaOH}

Anode: Cl−⟶Cl+e− ;Cl+Cl⟶Cl2\text{Anode: } \mathrm{Cl^{-} \longrightarrow Cl + e^{-}}\,; \qquad \mathrm{Cl + Cl \longrightarrow Cl_2}

c. Physical Properties of Chlorine :

i. Chlorine is a greenish-yellow gas having a pungent and suffocating odour.

ii. It is poisonous in nature.

iii. It can be easily liquified into a greenish yellow liquid, which boils at 239 K.

iv. It dissolves in water to give chlorine water.

v. It is 2-5 times heavier than air.

d. Chemical properties of chlorine :

i. Reaction with metals : Chlorine reacts with metals to form chlorides.

2Al+3Cl2⟶2AlCl3\mathrm{2Al + 3Cl_2 \longrightarrow 2AlCl_3}

2Na+Cl2⟶2NaCl\mathrm{2Na + Cl_2 \longrightarrow 2NaCl}

2Fe+3Cl2⟶2FeCl3\mathrm{2Fe + 3Cl_2 \longrightarrow 2FeCl_3}

ii. Reaction with nonmetals : Chlorine reacts with nonmetals to form their chlorides.

P4+6Cl2⟶4PCl3\mathrm{P_4 + 6Cl_2 \longrightarrow 4PCl_3}

S8+4Cl2⟶4S2Cl2\mathrm{S_8 + 4Cl_2 \longrightarrow 4S_2Cl_2}

iii. Affinity for hydrogen : Chlorine has great affinity for hydrogen. It reacts with hydrogen and with compounds containing hydrogen to form HCl.

H2+Cl2⟶2HCl\mathrm{H_2 + Cl_2 \longrightarrow 2HCl}

H2S+Cl2⟶2HCl+S\mathrm{H_2S + Cl_2 \longrightarrow 2HCl + S}

iv. Reaction with NH3\mathrm{NH_3} : Chlorine, when reacted with excess of ammonia, gives ammonium chloride and nitrogen:

8NH3 (Excess)+3Cl2⟶6NH4Cl+N2\mathrm{8NH_3\,(Excess) + 3Cl_2 \longrightarrow 6NH_4Cl + N_2}

Excess of chlorine reacts with ammonia to give nitrogen trichloride (explosive):

NH3+3Cl2 (Excess)⟶NCl3+3HCl\mathrm{NH_3 + 3Cl_2\,(Excess) \longrightarrow NCl_3 + 3HCl}

v. Reaction with alkali : Chlorine reacts with cold and dilute alkali to produce a mixture of chloride and hypochlorite. When reacted with hot concentrated alkali, chloride and chlorate are produced.

2NaOH+Cl2⟶NaCl+NaOCl+H2O (Cold and dilute)\mathrm{2NaOH + Cl_2 \longrightarrow NaCl + NaOCl + H_2O}\ \text{(Cold and dilute)}

6NaOH+3Cl2⟶5NaCl+NaClO3+3H2O (Hot and conc.)\mathrm{6NaOH + 3Cl_2 \longrightarrow 5NaCl + NaClO_3 + 3H_2O}\ \text{(Hot and conc.)}

Chlorine, when reacted with dry slaked lime, gives bleaching powder:

2Ca(OH)2+2Cl2⟶Ca(OCl)2+CaCl2+2H2O\mathrm{2Ca(OH)_2 + 2Cl_2 \longrightarrow Ca(OCl)_2 + CaCl_2 + 2H_2O}

vi. Reaction with hydrocarbons : Chlorine reacts with saturated hydrocarbons to give substitution products, and with unsaturated hydrocarbons it gives addition products.

CH4Methane+Cl2→U.VCH3ClMethylchloride+HCl\mathrm{\underset{Methane}{CH_4} + Cl_2 \xrightarrow{U.V} \underset{Methylchloride}{CH_3Cl} + HCl}

H2C=CH2Ethene+Cl2⟶CH2−CH2∣ ∣ClCl1,2-Dichloroethane\mathrm{\underset{Ethene}{H_2C{=}CH_2} + Cl_2 \longrightarrow \underset{1,2\text{-}Dichloroethane}{\begin{array}{c}\mathrm{CH_2{-}CH_2}\\ \mid\qquad\ \mid\\ \mathrm{Cl}\qquad \mathrm{Cl}\end{array}}}

vii. Oxidising property : Chlorine oxidises ferrous salts to ferric salts, and sulfites to sulfates.

2FeSO4+H2SO4+Cl2⟶Fe2(SO4)3+2HCl\mathrm{2FeSO_4 + H_2SO_4 + Cl_2 \longrightarrow Fe_2(SO_4)_3 + 2HCl}

Na2SO3+Cl2+H2O⟶Na2SO4+2HCl\mathrm{Na_2SO_3 + Cl_2 + H_2O \longrightarrow Na_2SO_4 + 2HCl}

It oxidises sulfur dioxide to sulfur trioxide, and iodine to iodate. In the presence of water they form sulfuric acid and iodic acid respectively.

SO2+2H2O+Cl2⟶H2SO4+2HCl\mathrm{SO_2 + 2H_2O + Cl_2 \longrightarrow H_2SO_4 + 2HCl}

I2+6H2O+5Cl2⟶2HIO3+10HCl\mathrm{I_2 + 6H_2O + 5Cl_2 \longrightarrow 2HIO_3 + 10HCl}

viii. Bleaching Property : Chlorine requires the presence of moisture (water) for bleaching. It liberates nascent oxygen from water, which is responsible for its oxidising and bleaching property.

Cl2+H2O⟶HCl+HOCl\mathrm{Cl_2 + H_2O \longrightarrow HCl + HOCl}

HOCl⟶HCl+[O]\mathrm{HOCl \longrightarrow HCl + [O]}

Chlorine bleaches vegetable matter or coloured organic matter, in the presence of moisture, to colourless matter: …

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