Chlorine

Chlorine, Cl, is a poisonous greenish-yellow gaseous non-metallic element, found in Group VIIb (i.e. the Halogen Group of elements) of the periodic table.

It has two Isotopes

Atomic Number : 17
Atomic Mass : 35.453
Melting Point : -101 degC
Boilingg Point : -35 degC
Density : 0.003

Chlorine is easily liquefied under pressure.

Discovery

Chlorine was discovered by Scheele in 1774AD, by the action of Manganese Dioxide on Hydrochloric Acid.
MnO2 + 4 HCl ==> Cl2 + MnCl2 + 2 H2O
In 1774AD, Sir Humphery Davy demonstrated that the gas was an element and he suggested the name Chlorine (Greek, greenish yellow) for the gas.

Occurrence

Because of its reactivity, Chlorine does not exist in the free elemental state in nature, although it is widely distributed in combination with other elements. The most common chlorine compound is Common Salt (i.e. Sodium Chloride, NaCl) which occurs in seawater and in rock salt.

Extraction

Preparation

The most common laboratory method for preparation of Chlorine is to heat 100 gm. of Manganese Dioxide with 300 ml. of concentrated Hydrochloric Acid.
MnO2 + 4 HCl ==> MnCl2 + 2 H2O + Cl2
The gas is bubbled through water to remove any traces of hydrochloric gas that may be present and then it is dried by bubbling it through concentrated sulphuric acid.
Chlorine may also be prepared by dropping cold concentrated Hydrochloric Acid on crystals of Potassium Permanganate.
2 KMnO4 + 16 HCl
==> 2 MnCl2 + 2 KCl + 8 H2O + 5 Cl2
The gas is bubbled through water to remove any traces of Hydrochloric Acid gas that may be present and then it is dried by bubbling it through concentrated Sulphuric Acid.

Manufacture

Chlorine is manufactured industrially as a by-product in the manufacture of Caustic Soda by the electrolysis of brine.
2 NaCl + 2 H2O ==> Cl2 + H2 + 2 NaOH
This process was carried out in Kellner-Solvay Cells, using Mercury and Carbon as the electrodes. However, due to the toxicity of mercury, the modern version of the process uses metal electrodes with special membranes in the electrolytic cells.

Properties

Chlorine is

a highly toxic greenish yellow gas,
has a pungent odour, and
fumes in moist air.

Reactions

Chlorine is a highly reactive element, and undergoes reaction with a wide variety of other elements and compounds.

Chlorine is a good bleaching agent, due to its oxidising properties.

Chlorine is soluble in water (which solution is called Chlorine Water) and this loses its yellow colour on standing in sunlight, due to the formation of a mixture of Hypochlorous Acid and Hydrochloric Acid.

Cl2 + H2O ==> HOCl + Hcl

Chlorine gas supports the vigorous combustion of many elements to form their chlorides. For example, Sulphur and Phosphorus burn in the gas.

Cl2 + S ==> Scl2

Cl2 + P ==> PCl3 + PCl5

Reaction of Chlorine with Hydrogen

A mixture of Chlorine and Hydrogen explodes when exposed to sunlight to give Hydrogen Chloride. In the dark, no reaction occurs, so activation of the reaction by light energy is required.

Cl2 + H2 ==> 2 HCl

Reaction of Chlorine with Non-Metals

Chlorine combines directly with most non-metals (except with Nitrogen, Oxygen and Carbon, C).

Reaction of Chlorine with Metals

Chlorine combines directly with all metals forming metal chloride salts.

Oxidising Reaction of Chlorine

Chlorine is a strong oxidising agent. Chlorine oxidises Iron (II) Chloride, FeCl2, to the salt containing Iron in the higher oxidation state Iron (III) Chloride, FeCl3. This is possible because Iron has a variable valency.
2 FeCl2 + Cl2 ==> 2 FeCl3
Chlorine displaces the less electronegative Bromine and Iodine from their respective salts.
Cl2 + 2 KBr ==> 2 KCl + Br2 Chlorine removes Hydrogen from the hydrides of non-metals, forming Hydrogen Chloride, and leaving the non-metal element.
Cl2 + H2S ==> 2 HCl + S

Reaction of Chlorine with Water

When Chlorine Water (i.e. a solution of chlorine gas in water) in a flask inverted in a basin of the same liquid is exposed to bright sunlight, the Chlorine is decomposed and a solution of Hydrochloric Acid remains.
H2O + Cl2 ==> HCl + HClO
The Hydrochloric Acid, HClO, is not very stable and the solution readily decomcomposes, especially when exposed to sunlight, yielding Oxygen.
2 HClO ==> 2 HCl + O2

Reaction of Chlorine with Sodium Hydroxide

Chlorine reacts with a cold solution of Sodium Hydroxide to give a mixture of Sodium Chloride, NaCl, and Sodium Hypochlorite, NaOCl.

Cl2 + 2 NaOH ==> NaCl + NaOCl + H2O

Chlorine reacts with a hot solution of Sodium Hydroxide to give a mixture of Sodium Chloride and Sodium Chlorate.

3 Cl2 + 6 NaOH ==> 5 NaCl + NaClO3 + 3 H2O

Uses

Chlorine is used

for the manufacture of bleaching powder and liquid bleaches,
to bleach fabrics (e.g. linen and cotton), wood pulp and paper,
in the manufacture of a wide range of chloro-organic solvents, including Methylene Chloride, CH2Cl2, Chloroform, CHCl3, Carbon Tetrachloride, CCl4,
in the manufacture of a number of important inorganic chemicals, including Sulphur Chloride, S2Cl2, Thionyl Chloride, SOCl2, Phosgene (i.e. Carbonyl Chloride), COCl2, and inorganic Chlorates, (e.g. Sodium Chlorate, NaClO3),
for the direct manufacture of Hydrochloric Acid by the direct combination of its elements,
H2 + Cl2 ==> 2 HCl
for the extraction of Gold from its ores,
in the manufacture Sodium Hypochlorite (i.e. domestic bleach), disinfectants, insecticides, plastics and Hydrochloric Acid,
as a disinfectant used to kill bacteria in the preparation of drinking water.
Chlorine is also important in the manufacture of paints, aerosol propellants and plastics.

Detection and Analysis

Chlorine is detected by its action on Iodide salts (e.g. Sodium Iodide), which are oxidises to free elemental Iodine, I2.

2 NaI + Cl2 ==> 2 NaCl + I2

The Iodine liberated in this reaction turns starch indicator solution to a blue colour.

Chlorine bleaches litmus paper.

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