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 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 ).

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 [2], 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 ). 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 (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 ). 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,