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CHLORINE (symbol Cl, atomic weight 35...
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See also:CHLORINE (See also:symbol Cl, atomic See also:weight 35`46 (0=16) , a gaseous chemical See also:element of the halogen See also:group, taking its name from the See also:colour, greenish-yellow (Gr. xXwpos). It was discovered in 1774 by See also:Scheele, who called it dep`ilogisticated muriatic See also:acid; about 1785, C. L. Berthollet, regarding it as being a See also:compound of hydrochloric acid and See also:oxygen, termed it oxygenized muriatic acid. This view was generally held until about 1810-1811, when See also:Sir H. See also: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 See also:condition, but in See also:combination with the See also:alkali metals it occurs widely distributed in the See also:form of See also:rock-See also:salt (See also:sodium chloride); as sylvine and carnallite, at Stassfiirt; and to a smaller extent in various other minerals such as matlockite and See also:horn-See also:mercury. In the form of alkaline chlorides it is found in See also:sea-See also:water and various See also:spring See also:waters, and in the tissues of animals and See also:plants; while, as hydrochloric acid it is found in volcanic gases. The preparation of chlorine, both on the small See also:scale and commercially, depends on the oxidation of hydrochloric acid; the usual oxidizing See also:agent is See also:manganese dioxide, which, when heated with concentrated hydrochloric acid, forms manganese chloride, water and chlorine :—MnO2+4HC1--MnC12+2H20+ C12. The manganese dioxide may be replaced by various other substances, such as red See also:lead, lead dioxide, See also:potassium bichromate, and potassium permanganate. Instead of See also:heating hydrochloric acid with manganese dioxide, use is frequently made of a mixture of See also:common salt and manganese dioxide, to which concentrated sulphuric acid is added and the mixture is then heated:—MnO2 +2NaCl+3H2SO4 = MnSO4+2NaHSO4+2H20+C12, Chlorine may also be obtained by the See also:action of dilute sulphuric acid on See also:bleaching, See also:powder. Owing to the enormous quantities of chlorine required for various See also: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 See also:necessity of using manganese in any form (see ALKALI MANUFACTURE). Owing to the reduction in the See also:supply of available hydrochloric acid (on See also:account of the increasing use of the " See also:ammonia-soda " See also:process in See also:place of the "Leblanc" process for the manufacture of soda) See also:Weldon tried to adapt the former to the See also:production of chlorine or hydrochloric acid. His method consisted in using See also:magnesia instead of See also: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 See also:magnesium chloride See also:solution and heating the See also:residue in See also:steam, to condense the acid vapours and so obtain hydrochloric acid. One See also: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 See also:conjunction with Pechiney, of Salindres (near See also:Alais, See also: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 See also: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 See also:Soc. of Chem. See also:Industry, 1884, p. 387). Greater success attended the efforts of See also:Ludwig See also:Mond, of the See also:firm of See also:Brunner, Mond & Co. In this process the ammonium chloride is volatilized in large See also:iron retorts lined with See also:Doulton tiles, and then led into large upright wrought-iron cylinders lined with See also:fire-bricks. These cylinders are filled with pills; made of a mixture of magnesia, potassium chloride and fireclay, the See also: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 See also:gas, containing no See also: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 I8-2o% of chlorine. This percentage drops gradually, and when it is reduced to about 3 % the temperature of the apparatus is lowered, by the See also:admission of air, to about 350° C., and the air stream containing the small percentage of chlorine is led off to a second See also: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, See also:British Assoc. Reports, 1896, p. 734). More recently, owing to the production of See also: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 See also:smell. It can be liquefied at — 340 C. under atmospheric pressure, and at —roe° C. it solidifies and crystallizes. Its specific See also:heat at See also:constant pressure 1S o' 1155, and at constant See also:volume o•08731 (A. Strecker, Wied. See also:Ann., 1877 [2], 13, p. 20); and its refractive See also:index 1.000772, whilst in the liquid condition the refractive index is 1'367. The See also:density is 2.4885 (air =1) (Treadwell and See also:Christie, Zeit. anorg. Chem., 1905, 47, P. 446). Its See also: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 See also:ordinary temperatures it unites directly with many other elements; thus with See also:hydrogen, combination takes place in See also:direct sunlight with explosive violence; See also:arsenic, See also:antimony, thin See also:copper See also:foil and See also:phosphorus take fire in an See also:atmosphere of chlorine, forming the corresponding chlorides. Many compounds containing hydrogen are readily decomposed by the gas; for example, a piece of See also:paper dipped in See also:turpentine inflames in an atmosphere of chlorine, producing hydrochloric acid and a copious See also:deposit of See also:soot; a lighted See also:taper See also:burns in chlorine with a dull smoky See also: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 See also:property that its bleaching and disinfecting See also:power depends (see BLEACHING). Water saturated with chlorine at o° C. deposits crystals of a See also:hydrate C12.8H20, which is readily decomposed at a higher temperature into its constituents. Chlorine hydrate has an See also:historical importance, as by sealing it up in a See also:bent See also:tube, and heating the end containing the hydrate, whilst the other See also:limb of the tube was enclosed in a freezing mixture, M. See also:Faraday was first able to obtain liquid chlorine. Chlorine is used commercially for the extraction of See also:gold (q.v.) and for the manufacture of " bleaching powder " and of See also:chlorates. It also finds an extensive use in organic See also: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 See also:rule only See also:works slowly on saturated compounds, but the reaction may be accelerated by the action of sunlight or on warming, or by using a " See also:carrier." In these latter cases the reaction may proceed in different directions; thus, with the aromatic See also:hydrocarbons, chlorine in the See also:cold or in the presence of a carrier substitutes in the See also:benzene See also:nucleus, but in the presence of sunlight or on warming, substitution takes place in the See also:side See also:chain. See also:Iodine, antimony trichloride, See also:molybdenum pentachloride, ferric chloride, ferric See also:oxide, antimony, See also:tin, stannic oxide and ferrous sulphate have all been used as chlorine See also:carriers. The atomic weight of chlorine was determined by J. Berzellusand by F. See also:Penny (Phil. Trans,, 1839, 13). . S. See also:Stas, from the See also:synthesis of See also:silver chloride, obtained the value 35.457 (0 = i6), and C. See also:Marignac found the value 34.462. More See also:recent determinations are: H. B. See also:Dixon and E. C. See also:Edgar (Phil. Trans., 1905); T. W. See also:Richards and G. See also: A. Noyes and H. C. See also:Weber (ibid., 1908), and Edgar (ibid., 1908). Hydrochloric Acid.—Chlorine combines with hydrogen to form hydrochloric acid, HC1, the only known compound of these two elements. The acid itself was first obtained by J. R. See also:Glauber in about 1648, but J. See also:Priestley in 1772 was the first to isolate it in the gaseous condition, and Sir H. Davy in r8ro showed that it contained hydrogen and chlorine only, as up to that See also:time it was considered to contain oxygen. It may be pre-pared by the direct See also:union of its constituents (see See also:Burgess and See also: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 o concentrated sulphuric acid, NaCl+H2SO4=NaHSO4-I-HCI. 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 See also: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 552.3 ° 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 See also: The salts of hydrochloric acid, known as chlorides, can, in most cases, be prepared by dissolving either the See also:metal, its hydroxide, oxide, or carbonate in the acid; or by heating the metal in a current of chlorine, or by precipitation. The See also: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 See also: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. See also:Bismuth and antimony chlorides are decomposed by water with production of oxychlorides, whilst See also:titanium tetrachloride yields titanic acid under the same conditions. All the metallic chlorides, with the exception of those of the alkali and alkaline See also:earth metals, are reduced either to the metallic condition or to that of a See also: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 See also:distillation with a mixture of potassium bichromate and concentrated sulphuric acid they yield See also:chromium oxychloride, whereas bromides and iodides by the same treatment give See also:bromine and iodine respectively. Some metallic chlorides readily form See also: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 See also:low boiling-point, which can be distilled without decomposition and are de-composed 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 See also:conversion into silver chloride, or if in the form of alkaline chlorides (in the See also:absence of other metals, and of any free acids) by titration with See also:standard silver nitrate solution, using potassium chromate as an See also:indicator. Chlorine and oxygen do not combine directly, but compounds can be obtained indirectly. Three oxides are known: chlorine monoxide, C120, chlorine peroxide, C102, and chlorine heptoxide, C1207. Chlorine monoxide results on passing chlorine over dry precipitated mercuric oxide. It is a See also: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 See also:shock or disturbance, or on exposure to sunlight. It is readily soluble in water, with which it combines to form hypoclsloreus acid. See also:Sulphur, paospilorus, See also:carbon compounds, and the alkali metals react violently with the gas, taking fire with explosive decomposition. A. J. See also:Balard determined the volume See also:composition of the gas by decomposition over mercury on See also: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, See also: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 See also:retort. After solution the retort is gently heated by warm water when the gas is liberated :—3KC103+2H2SO4 = KC104+2KHSO4 +H20+C102. 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 7o° C. on the water See also: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 See also:influence of See also:light, on warming, or on application of shock. It is a very powerful oxidant; a mixture of potassium chlorate and See also:sugar in about equal proportions spontaneously inflames when touched with a See also: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 See also:running in a thin stream of concentrated sulphuric acid (see papers by See also:Bray, Zeit. phys. Chem., 1906, et seq.). Chlorine heptoxide was obtained by A. See also:Michael by slowly adding perchloric acid to phosphoric oxide below–1o° 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 See also: 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 See also:electrolysis of brine solution in the presence of carbon electrodes, having no See also:diaphragm in the electrolytic See also:cell, and mixing the anode and See also:cathode products by agitating the liquid. The temperature should be kept at about 15° C., and the concentration of the hypo-. See also:chlorite produced must not be allowed to become too great, in See also: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 :2C10,+Na202=2NaC102+02. The silver and lead salts are unstable, being decomposed with explosive violence at loo° 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 See also: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 See also: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 See also:analysis of pure silver chlorate. The salts of this acid are known as chlorates (4.v.). Perchloric acid is best prepared by distilling potassium'perchlorate with concentrated sulphuric acid. According to Sir H. See also:Roscoe, pure perchloric acid distils over at first, but if the distillation he continueda white crystalline See also:mass of hydrated perchloric acid, HC104•H20, 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 (4 °). The crystalline hydrate melts at 5o° C. The pure acid decomposes slowly on See also:standing, but is See also:stable in dilute aqueous solution. It is a very powerful oxidizing agent; See also: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 1s not reduced by sulphurous acid. The salts of the acid are known as the per-chlorates, and are all soluble in water; the potassium and See also:rubidium salts, however, are only soluble to a slight extent. Potassium Perchlorate, KC104, 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. Additional information and CommentsThere are no comments yet for this article.
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