Chemical Elements - Chlorine - Physical and chemical properties.

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Chlorine is a greenish-yellow gas at room temperature and atmospheric pressure. It is con-siderably heavier than air. It has a choking smell, and inhalation causes suffocation, con-striction of the chest, tightness in the throat, and--after severe exposure--edema (filling with fluid) of the lungs. As little as 2.5 milligrams per litre in the atmosphere causes death within a few minutes, but less than 0.0001 percent by volume may be tolerated. Chlorine was the first gas used in chemical warfare in World War I. The gas is easily liquefied by cooling or by pressures of a few atmos-pheres at ordinary temperature.
Chlorine has a high electronegativity and a high electron affinity, the latter being even slightly higher than that of fluorine. The affinity of chlorine for hydrogen is so great that the reaction proceeds with explosive violence in light, as in the following equation:
In the presence of charcoal, the combination of chlorine and hydrogen takes place rapidly (but without explosion) in the dark. A jet of hydrogen will burn in chlorine with a silvery flame. Its high affinity for hydrogen allows chlorine to react with many compounds containing hydrogen. Chlorine reacts with hydrocarbons, for example, substituting chlorine atoms for the hydrogen atoms successively. If the hydrocarbon is saturated, however, chlorine atoms readily add to the double or triple bond.
Chlorine reacts with many elements of both metals and nonmetals to give chlorides. Only toward carbon, nitrogen, and oxygen is it fairly inert. The products of reaction with chlorine usually are chlorides with high oxidation numbers, such as iron(III) chloride (FeCl3), tin(IV) chloride (SnCl4), or antimony(V) chloride (SbCl5), but it should be noted that the chloride of highest oxidation number of a particular element is frequently in a lower oxidation state than the fluoride of highest oxidation number. Thus, vanadium forms a pentafluoride, whereas the pentachloride is unknown, and sulfur gives a hexafluoride but no hexachloride. With sulfur, even the tetrachloride is unstable.
Chlorine displaces the less electronegative halogens from compounds. The displacement of bromides, for example, occurs according to the following equation:
Furthermore, it converts several oxides into chlorides. An example is the conversion of iron(III) oxide to the corresponding chloride:
With carbon monoxide chlorine gives carbonyl chloride; and with sulfur dioxide, sulfuryl chloride:
Chlorine is moderately soluble in water, yielding chlorine water, and from this solution a solid hydrate of ideal composition Cl2 7.66H2O is obtained. This hydrate is characterized by a structure that is more open than that of ice; the unit cell contains 46 molecules of water and six cavities suitable for the chlorine molecules. When the hydrate stands, disproportionation takes place--that is, one chlorine atom in the molecule is oxidized and the other is reduced. At the same time, the solution becomes acidic, as shown in the following equation:
in which the oxidation numbers are written above the atomic symbols. Chlorine water loses its efficiency as an oxidizing agent on standing, because hypochlorous acid gradually decomposes. The reaction of chlorine with alkaline solutions yields salts of oxo acids.
The ionization potential of chlorine is high. Although ions in positive oxidation states are not very stable, high oxidation numbers are stabilized by coordination, mainly with oxygen and fluorine. In such compounds bonding is predominantly covalent, and chlorine is capable of exhibiting the oxidation numbers +1, +3, +4, +5, +6, and +7.

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Updated: 2/9/98
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