 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 atmospheric oxygen in the presence of a catalyst. At the turn of the 20th century, chlorine was produced by the electrolysis of aqueous solutions of the chlorides of alkali metals. About 90–95 percent of the world production of chlorine was obtained by these methods in the 1970’s. Small amounts of chlorine are obtained as a by-product in the production of magnesium, calcium, sodium, and lithium by the electrolysis of molten chlorides. In 1975 the world production of chlorine was about 25 million tons.
The two major methods for the electrolysis of aqueous solutions of NaCl are electrolysis in a diaphragm cell with a solid cathode and electrolysis in a mercury cathode cell. In both methods, chlorine gas is liberated at the graphite anode or titanium oxide-ruthenium oxide anode. In the first method, hydrogen is liberated at the cathode, and a solution of NaOH and NaCl is formed, from which commercial caustic soda is obtained by subsequent treatment. In the second method, sodium amalgam is formed at the cathode. A NaOH solution, hydrogen, and pure mercury are formed upon the decomposition of sodium amalgam by pure water in a separate apparatus. The pure mercury formed is reused in production. Both methods yield 1.125 tons of NaOH per ton of chlorine produced.
Electrolysis in a diaphragm cell is a less expensive process and yields cheaper NaOH. The mercury cathode method permits the production of very pure NaOH, although mercury losses in the course of production pollute the environment. In 1970, 62.2 percent of the world production of chlorine was by the mercury cathode method, while the method using the diaphragm cell accounted for 33.6 percent and other methods for 4.2 percent. Beginning in 1970, electrolysis using a solid cathode and an ion-exchange membrane was used, a method that made possible the production of pure NaOH without mercury.
Uses. Chlorine production is one of the leading branches of the chemical industry. Most of the chlorine produced is converted at the production site into chlorine-containing compoun