1911 Encyclopædia Britannica/Chlorine
CHLORINE (symbol Cl), atomic weight 35.46 (O = 16), a gaseous chemical element of the halogen group, taking its name from the colour, greenish-yellow (Gr. χλωρός). It was discovered in 1774 by Scheele, who called it dephlogisticated muriatic acid; about 1785, C. L. Berthollet, regarding it as being a compound of hydrochloric acid and oxygen, termed it oxygenized muriatic acid. This view was generally held until about 1810–1811, when Sir H. Davy showed definitely that it was an element, and gave it the name which it now bears.
Chlorine is never found in nature in the uncombined condition, but in combination with the alkali metals it occurs widely distributed in the form of rock-salt (sodium chloride); as sylvine and carnallite, at Stassfürt; and to a smaller extent in various other minerals such as matlockite and horn-mercury. In the form of alkaline chlorides it is found in sea-water and various spring waters, and in the tissues of animals and plants; while, as hydrochloric acid it is found in volcanic gases.
The preparation of chlorine, both on the small scale and commercially, depends on the oxidation of hydrochloric acid; the usual oxidizing agent is manganese dioxide, which, when heated with concentrated hydrochloric acid, forms manganese chloride, water and chlorine:—MnO2＋4HCl = MnCl2＋2H2O + Cl2. The manganese dioxide may be replaced by various other substances, such as red lead, lead dioxide, potassium bichromate, and potassium permanganate. Instead of heating hydrochloric acid with manganese dioxide, use is frequently made of a mixture of common salt and manganese dioxide, to which concentrated sulphuric acid is added and the mixture is then heated:—MnO2 + 2NaCl＋3H2SO4 = MnSO4＋2NaHSO4＋2H2O＋Cl2. Chlorine may also be obtained by the action of dilute sulphuric acid on bleaching powder.
Owing to the enormous quantities of chlorine required for various industrial purposes, many processes have been devised, either for the recovery of the manganese from the crude manganese chloride of the chlorine stills, so that it can be again utilized, or for the purpose of preparing chlorine without the necessity of using manganese in any form (see Alkali Manufacture).
Owing to the reduction in the supply of available hydrochloric acid (on account of the increasing use of the “ammonia-soda” process in place of the “Leblanc” process for the manufacture of soda) Weldon tried to adapt the former to the production of chlorine or hydrochloric acid. His method consisted in using magnesia instead of lime for the recovery of the ammonia (which occurs in the form of ammonium chloride in the ammonia-soda process), and then by evaporating the magnesium chloride solution and heating the residue in steam, to condense the acid vapours and so obtain hydrochloric acid. One day before him E. Solvay had patented the same process, but neither of them was able to make the method a commercial success. However, in conjunction with Pechiney, of Salindres (near Alais, France), the Weldon-Pechiney process was worked out. The residual magnesium chloride of the ammonia-soda process is evaporated until it ceases to give off hydrochloric acid, and is then mixed with more magnesia: the magnesium oxychloride formed is broken into small pieces and heated in a current of air, when it gives up its chlorine, partly in the uncombined condition and partly in the form of hydrochloric acid, and leaves a residue of magnesia, which can again be utilized for the decomposition of more ammonium chloride (W. Weldon, Journ. of Soc. of Chem. Industry, 1884, p. 387). Greater success attended the efforts of Ludwig Mond, of the firm of Brunner, Mond & Co. In this process the ammonium chloride is volatilized in large iron retorts lined with Doulton tiles, and then led into large upright wrought-iron cylinders lined with fire-bricks. These cylinders are filled with pills, made of a mixture of magnesia, potassium chloride and fireclay, the object of the potassium chloride being to prevent any formation of hydrochloric acid, which might occur if the magnesia was not perfectly dry. At 300° C. the ammonium chloride is decomposed by the magnesia, with the formation of magnesium chloride and ammonia. The mixture is now heated to 600° C. in a current of hot dry gas, containing no free oxygen (the gas from the carbonating plant being used), and then a current of air at the same temperature is passed in. Decomposition takes place and the issuing gas contains 18-20% of chlorine. This percentage drops gradually, and when it is reduced to about 3% the temperature of the apparatus is lowered, by the admission of air, to about 350° C., and the air stream containing the small percentage of chlorine is led off to a second cylinder of pills, which have just been treated with ammonium chloride vapour and are ready for the hot air current. With four cylinders the process is continuous (L. Mond, British Assoc. Reports, 1896, p. 734).
More recently, owing to the production of caustic soda by electrolytic methods, much chlorine has consequently been produced in the same manner (see Alkali Manufacture).
Chlorine is a gas of a greenish-yellow colour, and possesses a characteristic unpleasant and suffocating smell. It can be liquefied at −34° C. under atmospheric pressure, and at −102° C. it solidifies and crystallizes. Its specific heat at constant pressure is 0.1155, and at constant volume 0.08731 (A. Strecker, Wied. Ann., 1877 , 13, p. 20); and its refractive index 1.000772, whilst in the liquid condition the refractive index is 1.367. The density is 2.4885 (air = 1) (Treadwell and Christie, Zeit. anorg. Chem., 1905, 47, p. 446). Its critical temperature is 146° C. Liquid and solid chlorine are both yellow in colour. The gas must be collected either by downward displacement, since it is soluble in water and also attacks mercury; or over a saturated salt solution, in which it is only slightly soluble. At ordinary temperatures it unites directly with many other elements; thus with hydrogen, combination takes place in direct sunlight with explosive violence; arsenic, antimony, thin copper foil and phosphorus take fire in an atmosphere of chlorine, forming the corresponding chlorides. Many compounds containing hydrogen are readily decomposed by the gas; for example, a piece of paper dipped in turpentine inflames in an atmosphere of chlorine, producing hydrochloric acid and a copious deposit of soot; a lighted taper burns in chlorine with a dull smoky flame. The solution of chlorine in water, when freshly prepared, possesses a yellow colour, but on keeping becomes colourless, on account of its decomposition into hydrochloric acid and oxygen. It is on this property that its bleaching and disinfecting power depends (see Bleaching). Water saturated with chlorine at 0° C. deposits crystals of a hydrate Cl2·8H2O, which is readily decomposed at a higher temperature into its constituents. Chlorine hydrate has an historical importance, as by sealing it up in a bent tube, and heating the end containing the hydrate, whilst the other limb of the tube was enclosed in a freezing mixture, M. Faraday was first able to obtain liquid chlorine.
Chlorine is used commercially for the extraction of gold (q.v.) and for the manufacture of “bleaching powder” and of chlorates. It also finds an extensive use in organic chemistry as a substituting and oxidizing agent, as well as for the preparation of addition compounds. For purposes of substitution, the free element as a rule only works slowly on saturated compounds, but the reaction may be accelerated by the action of sunlight or on warming, or by using a “carrier.” In these latter cases the reaction may proceed in different directions; thus, with the aromatic hydrocarbons, chlorine in the cold or in the presence of a carrier substitutes in the benzene nucleus, but in the presence of sunlight or on warming, substitution takes place in the side chain. Iodine, antimony trichloride, molybdenum pentachloride, ferric chloride, ferric oxide, antimony, tin, stannic oxide and ferrous sulphate have all been used as chlorine carriers.
The atomic weight of chlorine was determined by J. Berzelius and by F. Penny (Phil, Trans., 1839, 13). J. S. Stas, from the synthesis of silver chloride, obtained the value 35.457 (O = 16), and C. Marignac found the value 34.462. More recent determinations are: H. B. Dixon and E. C. Edgar (Phil. Trans., 1905); T. W. Richards and G. Jones (Abst. J.C.S., 1907); W. A. Noyes and H. C. Weber (ibid., 1908), and Edgar (ibid., 1908).
Hydrochloric Acid.—Chlorine combines with hydrogen to form hydrochloric acid, HCl, the only known compound of these two elements. The acid itself was first obtained by J. R. Glauber in about 1648, but J. Priestley in 1772 was the first to isolate it in the gaseous condition, and Sir H. Davy in 1810 showed that it contained hydrogen and chlorine only, as up to that time it was considered to contain oxygen. It may be prepared by the direct union of its constituents (see Burgess and Chapman, J.C.S., 1906, 89, p. 1399), but on the large scale and also for the preparation of small quantities it is made by the decomposition of salt by means of concentrated sulphuric acid, NaCl + H2SO4 = NaHSO4 + HCl. It is chiefly obtained as a by-product in the manufacture of soda-ash by the Leblanc process (see Alkali Manufacture). The commercial acid is usually yellow in colour and contains many impurities, such as traces of arsenic, sulphuric acid, chlorine, ferric chloride and sulphurous acid; but these do not interfere with its application to the preparation of bleaching powder, in which it is chiefly consumed. Without further purification it is also used for “souring” in bleaching, and in tin and lead soldering.
It is a colourless gas, which can be condensed by cold and pressure to a liquid boiling at −83.7° C., and can also be solidified, the solid melting at −112.5° C. (K. Olszewski). Its critical temperature is 52.3° C., and its critical pressure is 86 atmos. The gas fumes strongly in moist air, and it is rapidly dissolved by water, one volume of water at 0° C. absorbing 503 volumes of the gas. The gas does not obey Henry’s law, that is, its solubility in water is not proportional to its pressure. It is one of the “strong” acids, being ionized to the extent of about 91.4% in decinormal solution. The strongest aqueous solution of hydrochloric acid at 15° C. contains 42.9% of the acid, and has a specific gravity of 1.212. Perfectly dry hydrochloric acid gas has no action on metals, but in aqueous solution it dissolves many of them with evolution of hydrogen and formation of chlorides.
The salts of hydrochloric acid, known as chlorides, can, in most cases, be prepared by dissolving either the metal, its hydroxide, oxide, or carbonate in the acid; or by heating the metal in a current of chlorine, or by precipitation. The majority of the metallic chlorides are solids (stannic chloride, titanic chloride and antimony pentachloride are liquids) which readily volatilize on heating. Many are readily soluble in water, the chief exceptions being silver chloride, mercurous chloride, cuprous chloride and palladious chloride which are insoluble in water, and thallous chloride and lead chloride which are only slightly soluble in cold water, but are readily soluble in hot water. Bismuth and antimony chlorides are decomposed by water with production of oxychlorides, whilst titanium tetrachloride yields titanic acid under the same conditions. All the metallic chlorides, with the exception of those of the alkali and alkaline earth metals, are reduced either to the metallic condition or to that of a lower chloride on heating in a current of hydrogen; most are decomposed by concentrated sulphuric acid. They can be distinguished from the corresponding bromides and iodides by the fact that on distillation with a mixture of potassium bichromate and concentrated sulphuric acid they yield chromium oxychloride, whereas bromides and iodides by the same treatment give bromine and iodine respectively. Some metallic chlorides readily form double chlorides, the most important of these double salts being the platinochlorides of the alkali metals. The chlorides of the non-metallic elements are usually volatile fuming liquids of low boiling-point, which can be distilled without decomposition and are decomposed by water. Hydrochloric acid and its metallic salts can be recognized by the formation of insoluble silver chloride, on adding silver nitrate to their nitric acid solution, and also by the formation of chromium oxychloride (see above). Chlorides can be estimated quantitatively by conversion into silver chloride, or if in the form of alkaline chlorides (in the absence of other metals, and of any free acids) by titration with standard silver nitrate solution, using potassium chromate as an indicator.
Chlorine and oxygen do not combine directly, but compounds can be obtained indirectly. Three oxides are known: chlorine monoxide, Cl2O, chlorine peroxide, ClO2, and chlorine heptoxide, Cl2O7.
Chlorine monoxide results on passing chlorine over dry precipitated mercuric oxide. It is a pale yellow gas which can be condensed, on cooling, to a dark-coloured liquid boiling at 5° C. (under a pressure of 737.9 mm.). It is extremely unstable, decomposing with extreme violence on the slightest shock or disturbance, or on exposure to sunlight. It is readily soluble in water, with which it combines to form hypochlorous acid. Sulphur, phosphorus, carbon compounds, and the alkali metals react violently with the gas, taking fire with explosive decomposition. A. J. Balard determined the volume composition of the gas by decomposition over mercury on gentle warming, followed by the absorption of the chlorine produced with potassium hydroxide, and then measured the residual oxygen.
Chlorine peroxide was first obtained by Sir H. Davy in 1815 by the action of concentrated sulphuric acid on potassium chlorate. As this oxide is a dangerous explosive, great care must be taken in its preparation; the chlorate is finely powdered and added in the cold, in small quantities at a time, to the acid contained in a retort. After solution the retort is gently heated by warm water when the gas is liberated:—3KClO3＋2H2SO4 = KClO4＋2KHSO4＋H2O＋ClO2. A mixture of chlorine peroxide and chlorine is obtained by the action of hydrochloric acid on potassium chlorate, and similarly, on warming a mixture of potassium chlorate and oxalic acid to 70° C. on the water bath, a mixture of chlorine peroxide and carbon dioxide is obtained. Chlorine peroxide must be collected by displacement, as it is soluble in water and readily attacks mercury. It is a heavy gas of a deep yellow colour and possesses an unpleasant smell. It can be liquefied, the liquid boiling at 9.9° C., and on further cooling it solidifies at −79° C. It is very explosive, being resolved into its constituents by influence of light, on warming, or on application of shock. It is a very powerful oxidant; a mixture of potassium chlorate and sugar in about equal proportions spontaneously inflames when touched with a rod moistened with concentrated sulphuric acid, the chlorine peroxide liberated setting fire to the sugar, which goes on burning. Similarly, phosphorus can be burned under water by covering it with a little potassium chlorate and running in a thin stream of concentrated sulphuric acid (see papers by Bray, Zeit. phys. Chem., 1906, et seq.).
Chlorine heptoxide was obtained by A. Michael by slowly adding perchloric acid to phosphoric oxide below −10° C.; the mixture is allowed to stand for a day and then gently warmed, when the oxide distils over as a colourless very volatile oil of boiling-point 82° C. It turns to a greenish-yellow colour in two or three days and gives off a greenish gas; it explodes violently on percussion or in contact with a flame, and is gradually converted into perchloric acid by the action of water. On the addition of iodine to this oxide, chlorine is liberated and a white substance is produced, which decomposes, on heating to 380° C, into iodine and oxygen; bromine is without action (see A. Michael, Amer. Chem. Jour., 1900, vol. 23; 1901, vol. 25).
Several oxy-acids of chlorine are known, namely, hypochlorous acid, HClO, chlorous acid, HClO2 (in the form of its salts), chloric acid, HClO3, and perchloric acid, HClO4. Hypochlorous acid is formed when chlorine monoxide dissolves in water, and can be prepared (in dilute solution) by passing chlorine through water containing precipitated mercuric oxide in suspension. Precipitated calcium carbonate may be used in place of the mercuric oxide, or a hypochlorite may be decomposed by a dilute mineral acid and the resulting solution distilled. For this purpose a filtered solution of bleaching-powder and a very dilute solution of nitric acid may be employed. The acid is only known in aqueous solution, and only dilute solutions can be distilled without decomposition. The solution has a pale yellow colour, and is a strong oxidizing and bleaching agent; it is readily decomposed by hydrochloric acid, with evolution of oxygen. The salts of this acid are known as hypochlorites, and like the acid itself are very unstable, so that it is almost impossible to obtain them pure. A solution of sodium hypochlorite (Eau de Javel), which can be prepared by passing chlorine into a cold aqueous solution of caustic soda, has been extensively used for bleaching purposes. One of the most important derivatives of hypochlorous acid is bleaching powder. Sodium hypochlorite can be prepared by the electrolysis of brine solution in the presence of carbon electrodes, having no diaphragm in the electrolytic cell, and mixing the anode and cathode products by agitating the liquid. The temperature should be kept at about 15° C., and the concentration of the hypochlorite produced must not be allowed to become too great, in order to prevent reduction taking place at the cathode.
Chlorous acid is not known in the pure condition; but its sodium salt is prepared by the action of sodium peroxide on a solution of chlorine peroxide: 2ClO2＋Na2O2 = 2NaClO2＋O2. The silver and lead salts are unstable, being decomposed with explosive violence at 100° C. On adding a caustic alkali solution to one of chlorine peroxide, a mixture of a chlorite and a chlorate is obtained.
Chloric acid was discovered in 1786 by C. L. Berthollet, and is best prepared by decomposing barium chlorate with the calculated amount of dilute sulphuric acid. The aqueous solution can be concentrated in vacuo over sulphuric acid until it contains 40% of chloric acid. Further concentration leads to decomposition, with evolution of oxygen and formation of perchloric acid. The concentrated solution is a powerful oxidizing agent; organic matter being oxidized so rapidly that it frequently inflames. Hydrochloric acid, sulphuretted hydrogen and sulphurous acid are rapidly oxidized by chloric acid. J. S. Stas determined its composition by the analysis of pure silver chlorate. The salts of this acid are known as chlorates (q.v.).
Perchloric acid is best prepared by distilling potassium perchlorate with concentrated sulphuric acid. According to Sir H. Roscoe, pure perchloric acid distils over at first, but if the distillation be continued a white crystalline mass of hydrated perchloric acid, HClO4·H2O, passes over; this is due to the decomposition of some of the acid into water and lower oxides of chlorine, the water produced then combining with the pure acid to produce the hydrated form. This solid, on redistillation, gives the pure acid, which is a liquid boiling at 39° C. (under a pressure of 56 mm.) and of specific gravity 1.764 (22/4)°. The crystalline hydrate melts at 50° C. The pure acid decomposes slowly on standing, but is stable in dilute aqueous solution. It is a very powerful oxidizing agent; wood and paper in contact with the acid inflame with explosive violence. In contact with the skin it produces painful wounds. It may be distinguished from chloric acid by the fact that it does not give chlorine peroxide when treated with concentrated sulphuric acid, and that it is not reduced by sulphurous acid. The salts of the acid are known as the perchlorates, and are all soluble in water; the potassium and rubidium salts, however, are only soluble to a slight extent. Potassium perchlorate, KClO4, can be obtained by carefully heating the chlorate until it first melts and then nearly all solidifies again. The fused mass is then extracted with water to remove potassium chloride, and warmed with hydrochloric acid to remove unaltered chlorate, and finally extracted with water again, when a residue of practically pure perchlorate is obtained. The alkaline perchlorates are isomorphous with the permanganates.