med with variable composition Cl2 · nH2O (where n ranges from 6 to 8), which are in the form of yellow crystals of the isometric system that decompose with increasing temperature into chlorine and water. Chlorine is readily soluble in TiCl4, SiCl4, SnCl4, and some organic solvents, especially, hexane, C6H14, and carbon tetrachloride, CCl4. The chlorine molecule is diatomic (Cl2). The degree of thermal dissociation of Cl2 + 243 kJ ⇄ 2Cl is 2.07 × 10–4 percent at 1000°K and 0.909 percent at 2500°K.
The outer electron configuration of the chlorine atom is 3s23p5. Consequently, in its compounds, chlorine may have oxidation states of –1, +1, +3, +4, +5, +6, and +7. The covalent radius of the chlorine atom is 0.99 Å, while the ionic radius of Cl– is 1.82 Å. The electron affinity of the chlorine atom is 3.65 eV, while the ionization energy is 12.97 eV.
Chemically, chlorine is very reactive and combines directly with almost all metals (with some metals, it reacts only in the presence of moisture or upon heating) and with nonmetals (except carbon, nitrogen, oxygen, and the inert gases), forming the corresponding chlorides. It reacts with many compounds, replaces hydrogen in saturated hydrocarbons, and combines with unsaturated compounds. Chlorine replaces bromine and iodine from their compounds with hydrogen and metals and is itself replaced by fluorine from its compounds with these elements.
In the presence of tiny amounts of moisture, alkali metals react with chlorine by combustion. Most metals react with dry chlorine only upon heating. Steel, as well as some metals, are stable in the presence of dry chlorine at moderate temperatures and thus are used for the construction of equipment used with dry chlorine and tanks for storing dry chlorine. Phosphorus ignites in chlorine, forming PCl3, and upon further chlorination, PCl5. Sulfur reacts with chlorine to yield S2Cl2, SCl2, and other compounds with the general formula SnClm. Arsenic, antimony, bismuth, strontium, and tellurium react vigorously with chlorine.
A mixture of chlorine and hydrogen burns with a colorless or yellowish green flame, producing hydrogen chloride by a chain reaction. The maximum temperature of a hydrogen-chlorine flame is 2200°C. Mixtures of chlorine and hydrogen containing 5.8 to 88.5 percent hydrogen are explosive.
With oxygen, chlorine forms the oxides Cl2O, ClO2, O2O6, Cl2O7, and Cl2O8, as well as hypochlorites (salts of hypochlorous acid), chlorites, chlorates, and perchlorates. All the oxygen compounds of chlorine perchlorates. All the oxygen compounds of chlorine form explosive mixtures with readily oxidizable compounds. Chlorine oxides have low stability and may explode spontaneously. Hypochlorites upon storage decompose slowly, while chlorates and perchlorates may explode under the action of initiators.
Chlorine hydrolizes in water, forming hypochlorous and hydrochloric acids: Cl2 + H2O ⇆ HClO + HCl. Hypochlorites and chlorides are formed upon the chlorination of cold alkaline aqueous solutions: 2NaOH + Cl2 = NaClO + NaCl + H2. Chlorates are formed upon heating. Chlorinated lime is formed by the chlorination of dry calcium hydroxide (see).
Nitrogen trichloride is formed in the reaction between ammonia and chlorine. In the chlorination of organic compounds, chlorine either replaces hydrogen, for example, R—H + Cl2 = RCl + HCl, or attaches through multiple bonds, for example,
forming various chlorine-containing organic compounds (organic chlorides).
With other halogens, chlorine forms interhalogen compounds. The fluorides ClF, ClF3, and ClF5 are very reactive; for example, glass wool ingnites spontaneously in the presence of ClF3. Compounds of chlorine with oxygen and fluorine include chlorine oxyfluorides, such as ClO3F, ClO2F3, ClOF, and ClOF3, and fluorine perchlorate, FClO4.
Production. The industrial production of chlorine was begun in 1785 based on the reaction between hydrochloric acid and manganese dioxide, or pyrolusite. In 1867 the British chemist H. Deacon developed a method for the production of chlorine by the oxidation of HCl using atmo