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chlorine (klōrˈēn, klôrˈ–) [Gr.,=green], gaseous chemical element; symbol Cl; at. no. 17; interval in which at. wt. ranges 35.446–35.457; m.p. −100.98℃; b.p. −34.6℃; density 3.2 grams per liter at STP; valence −1, +1, +3, +5, +7. Chlorine is a greenish-yellow poisonous gas with a disagreeable, suffocating odor; it is about two and one-half times as dense as air. Only fluorine among the nonmetals is more chemically active. Chlorine belongs to the halogen family of elements, found in Group 17 of the periodic table. The gas is composed of diatomic molecules (Cl2) with molecular weight 70.906. Chlorine was discovered in 1774 by K. W. Scheele, who thought it was a compound of oxygen; it was named and identified as an element by Sir Humphry Davy in 1810.
Chlorine is soluble in water; its aqueous solution, called chlorine water, consists of a mixture of chlorine, hydrochloric acid, and hypochlorous acid; only a part of the chlorine introduced actually goes into solution, the major part reacting chemically with the water. Chlorine water has strong oxidizing properties resulting from the oxygen set free when the unstable hypochlorous acid decomposes. Chlorine reacts readily with hydrogen to form hydrogen chloride. It burns if ignited in a hydrogen atmosphere and, if unignited, can form explosive mixtures with hydrogen; it also unites with the hydrogen in compounds such as turpentine, a hydrocarbon. In the presence of moisture it combines directly with certain metals, such as copper and iron, to form chlorides. Iron ignites when heated in a chlorine atmosphere. With metals and oxygen, chlorine forms several chlorates; it also combines with many nonmetals and certain radicals.
Because of its activity chlorine does not occur uncombined in nature, but its compounds are numerous and abundant. Sodium chloride (common salt) is present in seawater, salt wells, and large salt deposits, often in association with other chlorides. Chlorine is produced commercially chiefly by the electrolysis of sodium chloride, either molten or in solution. Other chlorides are sometimes employed. Chlorine can also be prepared from hydrochloric acid by oxidation of the hydrogen chloride (Deacon's process) and from bleaching powder.
Chlorine is used in water purification; as a disinfectant and as an antiseptic (mercuric chloride); and in the manufacture of bleaching powder (chloride of lime), dyes, and explosives. Chlorinated hydrocarbons have been used extensively as pesticides; some examples are DDT, dieldrin, aldrin, endrin, lindane, chlordane, and heptachlor. These compounds resist degradation and have become very troublesome environmental pollutants. Carbon tetrachloride and trichloroethylene are used as solvents. The Freon refrigerants are hydrocarbons that have been reacted with chlorine and fluorine. Chlorine is an important constituent of many poison gases. It is used in such compounds as calomel, chloroform, and chloral hydrate, which are used in medicine. It is also employed in the extraction of bromine from seawater. It is used in preparing some synthetic rubbers, in petroleum refining, and to prepare pure hydrochloric acid (see hydrogen chloride).
Cl, a chemical element in Group VII of Mendeleev’s periodic system. Atomic number, 17; atomic weight, 35.453. A member of the halogen family.
Under normal conditions (0°C and 0.1 meganewton/m2, or 1 kilogram-force/cm2), chlorine is a yellowish green gas, with a pungent, irritating odor. It occurs naturally in the form of two stable isotopes: 35Cl (75.77 percent) and 37Cl (24.23 percent). A number of radioactive chlorine isotopes have been obtained artificially, with mass numbers 32, 33, 34, 36, 38, 39, and 40, which have a half-life of 0.31 sec, 2.5 sec, 1.56 sec, 3.1 × 105 yr, 37.3 min, 55.5 min, and 1.4 min, respectively. 36Cl and 38Cl are used as isotopic tracers.
History. Chlorine was first obtained in 1774 by K. Scheele by the reaction of hydrochloric acid with pyrolusite (manganese dioxide). However, it was only in 1810 that H. Davy established it as an element and named it chlorine (from Greek chloros, “yellowish green”). In 1813, J. L. Gay-Lussac proposed the French name chlore for this element from which the Russian name khlor is derived.
Distribution in nature. Chlorine is found in nature only as a component of compounds. The average content of chlorine in the earth’s crust (clarke) is 1.7 × 10–2 percent by weight. The average content in acidic igneous rocks, such as granites, is 2.4 × l0–2, and in basic and ultrabasic rocks, 5 × 10–3. Water migration plays a major role in the history of chlorine in the earth’s crust. In the form of the Cl– ion, chlorine is a component of the earth’s oceans (1.93 percent), underground brines, and salt lakes. There are 97 chlorine minerals, primarily natural chlorides, the most important of which is halite, NaCl (see and ROCK SALT). Numerous extensive deposits of potassium and magnesium chlorides and mixed chlorides are known: sylvite, KCl, sylvinite, (Na,K)Cl, carnallite, KCl · MgCl2 · 6H2Ó, kainite, KCl · MgSO4 · 3H2O, and bischofite MgCl2 · 6H2O. The migration of HCl contained in volcanic gases into the upper portions of the earth’s crust was of great significance in the earth’s geological history.
Physical and chemical properties. The boiling point of chlorine is –34.05°C, and the melting point –101°C. The density of chlorine gas at normal conditions is 3.214 g/liter (g/l), while the density of the saturated vapor at 0°C is 12.21 g/l. The density of liquid chlorine at its boiling point is 1.557 g/cm3, while the density of solid chlorine at –102°C is 1.9 g/cm2. The pressure of saturated chlorine vapor is 0.369 meganewton/m2 (MN/m2), or 3.69 kilograms-force (kgf/cm2), at 0”C, 0.772 MN/m2 (7.72 kgf/cm2) at 25°C, and 3.814 MN/m2 (38.14 kgf/cm2) at 100°C. The heat of fusion is 90.3 kilojoules/kg (kJ/kg), or 21.5 cal/g, while the heat of evaporation is 288 kJ/kg (68.8 cal/g). The heat capacity of chlorine gas at constant pressure is 0.48 kJ/(kg · °K), or 0.11 cal/(g · °C). The critical constants of chlorine are as follows: critical temperature, 144°C; critical pressure, 7.72 MN/m2 (77.2 kgf/cm2); critical density, 573 g/l; and critical volume, 1.745 × 10–3l/g. The solubility of chlorine at a partial pressure of 0.1 MN/m2 (1 kgf/cm2) in water is 14.8 g/l at 0°C, 5.8 g/l at 30°C, and 2.8 g/l at 70°C, while in a 300 g/l NaCl solution, its solubility is 1.42 g/l at 30°C and 0.64 g/l at 70°C.
Below 9.6°C in aqueous solutions, hydrates of chlorine are formed 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 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 compounds. Chlorine is stored and transported in liquid form in tanks, cylinders, railroad tank cars, or specially equipped ships. The following consumption of chlorine is characteristic of industrial countries: 60–75 percent is used for the production of chlorine-containing organic compounds, 10–20 percent for the production of chlorine-containing inorganic compounds, 5–15 percent for the bleaching of pulp and fabrics, and 2–6 percent for sanitary purposes and water chlorination.
Chlorine is also used for the chlorination of some ores in order to extract titanium, niobium, and zirconium.
Various chlorine-containing organic and inorganic compounds are discussed in separate articles (see Index).
L. M. IAKIMENKO
Chlorine in organisms. Chlorine is a biogenic element and a component of plant and animal tissues. The content of chlorine in plants ranges from thousandths of 1 percent to several percent (halophytes contain large amounts of chlorine), while the chlorine content in animals ranges from hundredths to tenths of 1 percent. The daily chlorine requirement of an adult human is 2–4 g and is met simply with the ingestion of food. In food, chlorine is usually present in excess in the form of sodium chloride and potassium chloride. Bread, meat, and milk products are especially rich in chlorine.
In animal organisms, chlorine is a major osmotically active substance of blood plasma, lymph, spinal fluid, and some tissues. It is important in water-salt metabolism, facilitating the retention of water by tissues. The regulation of the acid-base equilibrium in tissues is accomplished, in addition to other processes, by a change in the chlorine distribution between blood and other tissues.
In plants, chlorine participates in energy exchange, activating both oxidative phosphorylation and photophosphorylation. It also affects the absorption of oxygen by roots and is necessary for the formation of oxygen in photosynthesis by isolated chloroplasts. Chlorine is not included in the composition of most nutritive media for artificial plant cultivation. It is possible that very low concentrations of chlorine are sufficient for the development of plants.
M. IA. SHKOL’NIK
Poisoning. Chlorine poisoning is possible in the chemical, pulp-and-paper, textile, and pharmaceutical industries. Chlorine irritates the mucous membranes of the eyes and respiratory tract. Secondary infections usually follow the primary inflammatory changes. Acute poisoning develops almost immediately. Among the symptoms noted upon the inhalation of medium and low concentrations of chlorine are tightening and pain in the chest, dry cough, rapid breathing, burning sensation in the eyes and tearing, increased content of leukocytes in the blood, and increased body temperature. Bronchial pneumonia, toxic pulmonary edema, depression, and convulsions are possible. In light cases, recovery occurs after three to seven days. Catarrh of the upper respiratory tract and recurring bronchitis and pneumosclerosis are long-term sequellae; the activation of pulmonary tuberculosis is also possible. Upon prolonged breathing of low concentrations of chlorine, similar but slowly developing disorders are observed.
Safety measures to prevent chlorine poisoning include the hermetic sealing of production equipment, good ventilation, and, when necessary, the use of gas masks. The maximum allowable concentration of chlorine in the air at production sites is 1 mg/m3. The production of chlorine, chlorinated lime, and other chlorine-containing compounds is considered potentially harmful and consequently Soviet law restricts the use of female and juvenile labor.
A. A. KASPAROV
REFERENCESIakimenko, L. M. Proizvodstvo khlora, kausticheskoi sody i neorganicheskikh khlorproduktov. Moscow, 1974. Nekrasov, B. V. Osnovy obshchei khimii, 3rd ed. [vol.] 1. Moscow, 1973.
Vrednye veshchestva v promyshlennosti, 6th ed., vol. 2. Edited by N. V. Lazarev. Leningrad, 1971.
Comprehensive Inorganic Chemistry, vols. 1–5. Edited by J. C. Bailar et al. Oxford, 1973.