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chlorine (Cl)





chlorine
A poisonous, greenish-yellow, gaseous element, with a choking smell. Chlorine is one of the halogens. It occurs widely in nature as sodium chloride in seawater and as halite (NaCl), carnalite (KCl.MgCl2.6H2O), and sylvite (KCl).

Chlorine is more than twice as dense as air and is quite soluble in water. It does not burn but is a very reactive substance, combining directly with many elements, both metals and non-metals (e.g., sodium and phosphorus) and compounds. In some reactions it acts a strong oxidizing agent by removing hydrogen from compounds. The hydrogen combines with the the chlorine to yield hydrogen chloride. The action of chlorine as a bleaching agent is one example of its oxidizing properties. Chlorine reacts with most organic compounds, replacing hydrogen atoms (see alkyl halides) and adding to double and triple bonds.

Chlorine was discovered by Carl Scheele in Uppsala, Sweden, in 1774. Its name comes from the Greek chloros, meaning "pale green".


atomic number 17
atomic mass 35.453
electronic configuration 1s22s22p63s23p5
first ionization energy 1,251 kJ/mol
electronegativity 3.2
atomic radius 99 pm
ionic radius 181 pm
density 3.214 g/dm3
melting point -101.5°C (171.6 K)
boiling point -34.04°C (239.1 K)


Preparation of chlorine

Chlorine may be prepared by the oxidation of hydrochloric acid – hydrogen ions are oxidized to water thus releasing the free element chlorine. The usual way of bringing about the reduction is to heat manganese dioxide with concentrated hydrochloric acid. Although the gas is soluble in water, it may be collected over a strong solution of sodium chloride.

Chlorine is also obtained by the electrolysis of sodium chloride solution using carbon electrodes. Chlorine is liberated at the anode (positive electrode) while sodium is set free at the cathode (negative electrode). However, sodium is so reactive that it reacts with the water – hydrogen is liberated at the same time as sodium hydroxide is formed.

industrial production of chlorine
The industrial preparation of chlorine involves the electrolysis of bring, the reaction being 2NaCl +2H2O → Cl2 + 2NaOH + H2. The diaphragm cell shown here is one of the methods used.

1 brine (salt solution) inlet; 2 chlorine produced at the graphite anodes bubbles up through the brine reservoir; 3 graphite anode; 4 pipeline connected to the cells collects this chlorine; pipeline is rubber-lined to prevent corrosion; 5 hydrogen gas and sodium hydroxide produced inside the iron screen cathode pockets; 6 sodium hydroxide outlet through which the sodium hydroxide produced inside the pockets flows; 7 hydrogen gas led off from the cell; 8 concrete cell (Hooker diaphragm type)


Uses of chlorine

Chlorine is used in large quantities as a bleach, as a disinfectant for drinking water and swimming pools, and in the manufacture of plastics, solvents, and other compounds.






Compounds of chlorine

Chlorides, the commonest chlorine compounds, are typical halides except for carbon tetrachloride (see carbon), which is inert. Other chlorine compounds include a series of oxides, unstable and highly oxidizing, and a series of oxyanions – hypochlorites, chlorite, chlorate, and perchlorate – with the corresponding oxy-acids, all powerful oxidizing agents. Calcium hypochlorite (see bleaching powder and sodium chlorite are used as bleaches; chlorates are used as weedkillers and to make matches and fireworks; perchlorates are used as explosives and rocket fuels.


Related category

   • INORGANIC CHEMISTRY