21 6 MELLOR: ON THE UNION OF X1X.-On the Union of Hydrogen and Oilorine. Parts I to 1.1. By J. W. MELLOB. THIS paper contains the preliminary results of an investigation, sugges- ted by Professor Dixon, on the mode of formation of hydrogen chloride. So far as I can gather, William Cruickshank,l of Woolwich, was the first t o observe the gradual combination of hydrogen and chlorine gases. In reference t o the action of chlorine on hydrogen, hydro- carbons, and carbon monoxide he said on August 10, 1801 : “ If the pure oxigenated muriatic acid, in the form of a gas, be mixed in certain proportiona with any of these inflammable gases and introduced into a bottle filled with and inverted over water, though no immediate action may be at first perceptible, yet, in twenty-four hours a complete decomposition will be found to have taken place, the producte varying according t o the nature of the gases employed., . “I introduced into a phial with a glass stopper, filled with and inverted over water, one measure of pure hydrogen and afterwards two mea- sures of very pure oxigenated muriatic gas, this nearly filled the bottle; the stopper was then introduced very tight under water, and before the stopper was introduced, a whitish cloud appeared in the mixture yet very little or no diminution could be observed . . . ; at the end of twenty-four hours when the stopper was withdrawn the whole of the gas instantly disappeared except about one-tenth of a measure, which was found to be azote, and must have been originally contained in the two measures of oxigenated muriatic acid gas.I n tbis case the products were manifestly common muriatic acid and water, for the water in the phial contained common muriatic acid, but did not in the least smell of the oxigenated.” The words I have italicised appear to be the first record of a, pheno- menon named, Iater, the period of induction. On February 27, 1809, Gay Lussac and Thenard announced that an Compare Desormes and Clement, Ann. Chim., 1801, 89, 26 ; Berthollet’s “Easai de Statique Chimique,” 1803, 1, 423. Cruickshank, Hicholson’s Journ., 1801, [i], 5, 202.HYDROGEN AND CHLORINE, PARTS I TO .III. 217 explosive combination occurs when a mixture of hydrogen and chlorine gases is exposed to direct sunlight.l I n June of the same year, John Dalton showed the influence of light in this reaction, and on repeating Cruickshank’s experiment, also observed that ( I the gases after being put together (over water) seemed to have no effect for one or two minutes, when suddenly the mixture began to diminish with rapidity.” I n a letter to Goethe in 1810, Seebeck intimated that a mixture of hydrogen and chlorine gases contained in a clear glass vessel detonated the sunshine, whilst under a dark blue glass combination occurred without explosion in one minute, and under a dark red glass theaction either took place very slowly or not a t all.3 This observation was more particularly investigated by Berard (1813),4 Draper (1843),5 Favre and Silbermann (1853),6 and finally by Bunsen and Roscoe (1 85 7).7 Draper took up the subject about 1840, and made an instrument, called the tithonometer, to measure the rate of combination of hydro- gen and chlorine under the influence of light.The action of his in- strument is based on the fact, that the hydrogen chloride formed is at once absorbed by the liquid in the same vessel. The resulting con- traction is measured on a suitable index. Draper believed that the first action was to induce a more active, allotropic modification of chlorine, for he found that insolated chlorine combines with hydrogen more readily, and even in the dark. The period of inertness, pre- viously noted by Cruickshank and Dalton, was then also suppressed. This allotropism was not confirmed by Bunsen and Roscoe (1855), or by Askenasy and Meyer (1892),S although Favre and Silbermann (1853), and Amato (1884) have given experimental evidence in favour of Draper’s original statement. Fremy and Becquerel believe it to be due to the presence of oxychlorine compounds formed by tbe action of chlorine on the water vapour present.9 Draper also records that if an intense light, such as that of a spark from a Leyden jar, be momentarily 1 Gay Lussac and Thenard, Mem.phys. Chim. SOC. d’drcueil, 1809, 2, 340, or Gilbert’s Ann., 1810, 35, 8. Alembic Club Reprints, No. 13, p. 43. Dalton’s ‘‘ A New System of Chemieal Philosophy,” 1811, 2, 189. Seebeck, “ Von der Cheniischen Action des Lichts und der farbigen Beleuchtung,” in Goethe’s “ Zur Farbenlehre,” Tubingen, 1810, quoted in Eder’s “Geschichte der Photochemie und Photographie,” 1891, 1, 73. Berard, Ann.Chim., 1813, 85, 309. Draper’s “Collected Memoirs,” 1878 ; Phil. Mag., 1843, [iii], 23, 401 ; 1845, Favre and Silbermann, Ann. Chim. Phys., 1853, [iii], 37, 479. [iii], 27, 327. 7 Bunsen and Roscoe, Pogg. Ann., 1855, 96, 373 ; 1857, 100, 43, 481 ; 1857, 101, 235; 1859, 108, 193 ; 1862, 117, 529 ; Phil. Trans., 1857, 146, 355, 601 ; 1859 148, 879 ; Ostwalcl’s “Klassiker,” Nos. 34 and 38. Askenasy and V. Meyer, AnnaZen, 1892, 269, 72. Becquerel and Fremy, Wurtz’s “Dictionnaire d0 Chimie,” 1879, 2, 255.218 MELLOR: ON THE UNION OF flashed on to the mixture, a sudden expansion, followed instantly by a return to the original volume, takes place (therefore no formation of hydrogen chloride). This phenomenon mill be named, after its first observer, the I‘ Draper effect.” Favre and Silbermannl found that the heat developed in the action of insolated chlorine on potash was greater than that of non- insolated chlorine by some 39 cal.The increase in the activity of the chlorine is not accompanied by a change in volume. The first part of Bunsen and Roscoe’s classical work appeared in 1855. The final result was the establishment of the more important laws of the chemical action of light. These investigators, by means of a perfected form of Draper’s tithonometer, found that an amount of actinic energy disappeared in the act of photochemical combination equivalent to the amount of light absorbed. This phenomenon was styled ‘‘ photochemical extinction.” After the period of inertness, observed by Cruickshank, the rate of combination of hydrogen and chlorine was found to gradually increase until a maximum steady state was attained (period of acceleration 2).The interval between the first impact of light and the periodof constancy was termed the period of “ photochemical induction.” Bunsen and Roscoe also found that the presence of minute traces of oxygen, or of an excess of either of the reacting components, considerably retarded the rate of formation of hydrogen chloride. Gautier and Helier,3 under somewhat different conditions, found an acceleration in the rate when either of the react- ing gases is present in excess. This is what the dynamicnl theory of mass action would lead us to expect. In 1871, Budde discovered that when chlorine is exposed to a source of actinic light it occupies a greater volume.This expansion cannot be attributed to the direct effects of heat. Recklinghausen 5 (1894) found that a photo-expansion also occurs when the chlorine is mixed with hydrogen, carbon monoxide, or ethylene. No change in volume occurs when perfectly pure and dry chlorine is exposed to light under somewhat similar conditions (Baker6 and Shenstone 7). Since Bunsen and Roscoe’s great work, the most important contribu- tion to the subject was published by Pringsheims in 1887. Prings- LOC. cit. Veley, Phil. Mag., 1894, [v], 37, 165. Budde, PhiZ. Mug., 1871, [v], 4 2 , 290; Pqg. Ann. Ergbd., 1873, 6, 477. See Recklinghausen, Zeit. physikal. Chem., 1894,14, 494. Baker, Brit. Assoc. Rep., 1894. 7 Shenstone, Trans., 1897, 71, 471. * Pringsheim, Wied.Ann., 1887, 32, 384 ; Dixon and Harker, Hem. and Proc. 3 Gautier and Helier, Compt. rend., 1897, 124, 1121. also Richardson, Proc. London Phys. Soc., 1891, 11, 186. Manchester Lit. Phil. Soc., 1889, [iv], 3, 118.HYDROaEN AND CHLORINE. PARTS I TO 111. 219 heim rediscovered the Draper effect, and also found that by drying the mixed electrolytic gases the rate of combination was retarded, for in sunlight the reaction was only accompanied by a feeble clicking sound (“ ein sehr schwaches knisterndes Gerausch ”). Pringsheim believes that during the period of induction some such intermediate compound as chlorine monoxide is formed by the action of chlorine on the water vapour present, A most interesting observation has just been recorded by Cordier’ somewhat t o the effect that dry chlorine is transparent, but moist chlorine opaque, to the actinic rays.It has been my purpose to investigate the mode in which light effects the combination of hydrogen with chlorine. With this object in view, I began by studying the electrolysis of hydrochloric acid, and the solution of chlorine in this acid, since an exact knowledge of these processes is necessary to interpret the work of previous investigators. I.-THE ELECTROLYSIS OF HYDROCHLORIC ACID. The general and most convenient mode of preparing a mixture of equal volumes of hydrogen and chlorine gases (Chlorknallgas) is by the electrolysis of hydrochloric acid, under the conditions Bunsen and Roscoe I have examined the gases prepared by this method with a view to finding what impurities, if any, are present, The highly successful experiment of Baker 3 in which a dried mixture of hydrogen and chlorine, prepared by another process, did not completely combine when exposed for two days to the direct rays of the sun, has not been equalled by the use of the electrolytic gases.Bunsen and Roscoe have also shown that the influence of impurities in modifying the rate of combination of electrolytic hydrogen and chlorine is most remarkable. The amount of foreign gas sufficient to materially disturb the normal rate cannot approach the billionth part of the total volume of thegas.” Draper has stated that the electrolysis of hydrochloric acid never yields equal volumes of hydrogen and chlorine. Bunsen and Roscoe, however, very carefully studied the action, and came to the conclusion that small variations from equality in the proportions of the two gases can be brought within the limits of analysis by taking found to be most favourable.Cordier, Monatsh., 1900, 21, 660. Bunsen and Roscoe, Pogg. Ann., 1855, 96, 373. Baker, Trans., 1894, 65, 611. It might also be pointed out that 26 per cent. of the mixture remained uncombined after four days exposure, two of which were of bright sunshine. * Draper, Zoc. cit.220 MELLOR: ON THE UNION OF suitable precautions and keeping the strength of the acid over 23 per cent. of hydrogen chloride.1 If, during the electrolysis of concentrated hydrochloric acid, 4HC1+ 2H,O are decomposed to form 2HOClf 01, + 3R,, the electro- lytic gases would consist of (m + n)H, and (mHOC1 + nCI,), neglecting the hydrogen chloride and steam present.If a cylinder containing these gases be opened under an aqueous solution of potassium iodide, then for every four volumes of HOCl present a quantity of iodine equivalent to two volumes of chlorine will be liberated. In Bunsen and Roscoe’s analyses there is a mean error of - 0.72 per cent., assuming that the mixture contained equal volumes of the gases. Therefore, unless the electrolytic gases contain as impurity an amount of HOCl vapour exceeding 0.18 per cent., it would have escaped detection.2 A slight excess of hydrogen would also occur if a very little water were decomposed in the electrolysis. A large amount of matter has been published on the electrolysis of the hypochlorites and chlorides. The following refers to hydrochloric acid.Riche (1858) observed that dilute hydrochloric acid yields per- chloric acid when a current from 10 Bunsen cells is passed through it. Tommasi4 (1882) found chlorine oxides at the positive pole even with concentrated acid, and pointed out the possibility of their forma- tion by the decomposition of the hydrate, HCI,GH,O, or by the action of oxygen or water on the electrolyte. Haber and Grinberg (1898), in a very complete investigation, con- firm Bunsen and Roscoe’s observation. Working on small quantities, they have shown that concentrated hydrochloric acid, with platinum electrodes, furnishes a 100 per cent. yield of chlorine, which falls to zero with increasing dilution. They also trace the presence of varying quantities of oxygen, perchloric, chloric, and hypochlorous acids t o the combination of C1 ions with the OH ions of water.6 The great solubility of the liberated chlorine and its diffusion over to the cathode is, no doubt, an important factor in the inducing of secondary action^.^ It is this that causes the electrical sign of the 1 It is interesting to observe in this connection that E.Morley (Zeit. phyGkaZ. Chem., 1896, 20, 430) was unable to obtain an electrolytic mixture of hydrogen and oxygen in the proportions H, : 0 by the electrolysis of water. Bunsen, Annulen, 1853, 86, 273. Riche, Compt. rend., 1858, 46, 350. Tommasi, Compt. rend., 1882, 95, 689. 5 Haber and Grinberg, Zeit. anorg. Chem., 1898, 16, 198 ; 18, 37. 6 Matteucci found that the more intense the current, the greater the amount of 7 Townsend, Proc.Camb. Phil. Xoc., 1897, 11, 245, or PhQ. Mag., 1898, [v], a, oxygen evolved (Gmelin’s ‘( Handbook of Chemistry,” Eng. ed., 1848, 1, 455). 25 ; Enright, Phil. Mug., 1890, [v], 49, 56.HYDROGEN AND CHLORINE. PARTS I TO 111. 221 liberated hydrogen to change from its initial positive value to a final negative one. The formation of oxygen compounds of chlorine by the dissociation of chlorine water was first indicated by Millon 1 in an equation subsequently developed by Jakowkin in the following form : + C1,AqS (HOCl + H + 6)Aq. The electrolysis of aqueous solutions of chlorine is said to lead to the production of hydrochloric acid at the negative pole, and of chloric acid at the positive Oette14 has shown that nascent hydrogen at the cathode recombines with the chlorine in the solution.Gautier and Helier say, L’Qlectrolyse de !’acid chlorhydrique ou dos chlorures fournissent, B chaud ou A froid, un gaz m$lh d’oxyde de chlore, comme ou peut s’en assurer en le faisant passer, aprbs dessicca- tion prhalable, dans un tube de porcelain chauffQ au rouge, en receuillant dans une Bprouvette pleine de potasse les gaz dhgaghs, il reste toujours de l’oxyghe ralluman t les corps en ignition.” On carefully repeating this experiment, my equivalent to their “ tube de porcelain chauff B au rouge ” was broken by regurgitation of the potash solution. The ex- periment succeeds equally well without the hot tube, The following experiments prove that variable quantities of oxygen are evolved during the electrolysis of concentrated hydrochloric acid, SERIES 1.-Pure cold hydrochloric acid saturated with dry hydrogen chloride gas was subjected to electrolysis i n a cell from which the spent acid could be removed and new acid introduced without ad- mission of air (see Fig.1, p. 222). The electrode gases were led off separately on account of the subsequent heating of the anode gases, otherwise explosive combination occurs at 430-4409 The pre pared carbon electrodes were cemented in while warm with a pulp of asbestos and sodium silicate (water glass). I n other respects Bunsen and Roscoe’s directions were closely followed. (a), The anode gases were led through a three-way cock and two Mitscherlich absorption bulbs, the one containing water the other con- centrated sulphuric acid of sp.gr. 1 *9 ; then over fragments of glass wet with the same acid in a V-tube 7; thence through a heated glass tube, 1 Millon, Compt. rend., 1849, 28, 42. 2 Jakowkin, Ber., 1897, 30, 518 ; E. Mnller, Zeit. EZektrochem., 1900, 6, 573. 3 Balard, J. pr. Chem., 1835, 4, 167. 4 Oettel, Chem. Centr., 1895, [iv], 7, ii, 3. 5 Gautier and Helier, Compt. rend., 1897, 124, 1129, 1267. 6 Meyer and Freyer, Ber., 1893, 26, 428. 7 Bailey and Fowler (Trans., 1888, 53, 755) have shown that if the gas contains traces of hydrogen chloride, the reaction 2P20, + 3HC1=POC13 + 3HP0, occurs. Chromic and copper compounds are objectionable for removing the hydrogen chloride on acconnt of the possible action of hypochlorous acid. Hence it appeared better not to rise phosphoric oxide (compare Gutmann, Annalen, 1898, 299, 267 ; Baker,222 MELLOR: ON THE UNION OF and finally collected in an eprouvette over an aqueous solution of‘ potassium hydroxide.Air was carefully swept out of the apparatus. by means of a current of chlorine, prepared by Gautier and Helier’s process, until a blank experiment gave no result. The gases evolved during the first two hours electrolysis escaped via the three-way cock. The different parts of the apparatus were sealed together before the blowpipe. In an average experiment, approximately 13-1 4 litres of the anode gases gave 1.2 C.C. of oxygen (at normal temperature and pressure) in successive measurements of 0.4, 0.1, 0.3, and 0.4 C.C. (b). The dried cathode gases, from which chlorine and hydrogen had The experiments were done in a dark cellar.FIG. l.--Electroly2ic vessel. 3 been removed, were passed through a temoirz tube and then over warm palladium asbestos.l The water was absorbed in a weighed phosphoric oxide tube. Air was, as before, swept out of the apparatus by a current of hydrogen previous to an experiment. No perceptible increase in weight was noticed. The large quantity of desiccating agents used in these experiments E. Morley 9 has shown that a slow current of is a serious objection. Trans., 1898, 73, 422). Calcium chloride cannot be used, for any chlorine monoxide present would form hypochlorites (Garzarolli-Thurnlackh and Schachal, Annalen, 1885, 230, 280). 1 Winkler’s “ Anleitung zur chemische Untersuchung der Industrie-Gase,” 258, (1877). Morley, Amer.J. Sci., 1885, [iii], 30, 140.HYDROGEN AND CHLORINE. PARTS I TO 111. 223 air, passed through strong sulphuric acid, failed to remove something like 0.002 milligram of water per litre. This amount of moisture passed along with chlorine through a red hot tube mould probably liberate oxygen, but not sufficient to account for that obtained, The absorption of oxychlorine compounds, as well as of hydrogen chloride, would take place in the first washing bulbs. Whatever the method by which Gautier and Helier performed their (‘ dessiccation prblable, ” their hot tube is unnecessary, for the experiment succeeds equally well without it. Unfortunately, there is no satisfactory means of detecting hypo. chlorous acid (or chlorine monoxide) under the conditions of these experiments, as was shown by Haber and Grinberg in a recent ex- amination of iche delicacy of the various methods proposed for the detection of hypochlorous acid in the presence of hydrochloric acid.They also point out that Wolter’s methodP2 used by Pedler 3 to establish an equation for the action of light on chlorine water, is quite unreliable for ‘ I nachweisbare Mengen unterchloriger Saure konnen selbst in &=& norm. Salzsaure nicht mehr bestehen.” Millon’s * manganous chloride test can, however, be used as a ‘(Vergleichsprobe” with chlorine water, hypochlorous acid giving a brown colour rapidly, chlorine water slowly. Jakowkin confirms this observation and says, I ‘ es est deshalb ganz unerklarlich, auf welche Weise Pedler die An- wesenheit von HClO in einer Chlorlosung, welche (nach der Belichtung) uberschussige Salzsaure enthielt, konstatiert hatte.” SERIES 11.-A double globe (Fig.2) had a t one end three necks- two wider ones for carbon electrodes, and a central one carrying a two- way cock (a)-on the other end was sealed a capillary tube, bent twice a t right angles, and carrying a three-way cock (d). The free end of the capillary had a piece of wider glass tubing sealed on, as shown in the figure, p. 224. Both globes were filled with concentrated hydrochloric acid (completely saturated with hydrogen chloride in the cold) as far as the three-way cock, which was lubricated with glacial phosphoric acid, This filling was easily done by dipping the tip of a in the acid and connectingd with an aspirator, or as shown in the figure.The upper cock d was then closed, and the lower one partially so. An electric current was switched on (2--4volts), and while the globe b was being filled with electrolytic gas, the capillary tube, on the side of the apparatus, was filled with concentrated sulphuric acid by pouring the acid in e and applying suction a t d. A drying tube was then fitted on to e. The lower globe was painted black. 1 LOC. cit. 6 Jakowkin, Zeit. physikak Chem., 1899, 29, 613. Wolters, J. pr. Chem., 1873, [ii], 7, 468. 3 Pedler, Trans., 1890, 57, 613. LOC. eit.224 MELLOR: ON THE UNION OF When the globe b was almost full of the electrolytic gas, the cock at a was closed, and d opened for some two hours. The right and left hand sides of the vessel were then put into communication, and the upper globe exposed to some source of artificial light (coal gas lamp).The current was so regulated by the introduction of a suitable resistance that the ratio of the rate of formation of the gases H2+C12 and the rate of solution of the reformed HCl was approximately cons tan t. After 14 days, the three-way cock d was joined to a Hempel burette containing an aqueous solution of manganous chloride, and the current FIG. 2. stopped. All the air between the capillary of the cock and the burette was driven over into the right hand side of the apparatus. Water was then run into the lower cock, until all the gases in the upper globe were transferred to the burette. Analysis, by Haber and Grinberg’s method, invariably showed that oxygen is formed during the electrolysis.From this and the preceding experiments, it follows that if x be the amount of oxygen in the vessel at the end of the time t , we have: x = +(t). The manganous chloride comparison test generally shows the presenceHYDROGEN AND CHLORINE, PARTS I TO 111. 225 of traces of hypochlorous acid, thus confirming the suspicion of Haber and Grinberg. The gases exposed to the light are, therefore, a mixture of chlorine, hydrogen, steam, hydrogen chloride, and oxygen, There is probably a condensation of water vapour on the glass. A reaction between chlorine and water is, therefore, quite possible. From the thermo- dynamical principle of maximum work, it can be shown that, while gaseous chlorine will not decompose steam at loo’, it will act slowly on water at atmospheric temperatures.Thus, in round numbers : (1). Water Vapour and Chlorine Gas. [H2,0] gas a t looo= + 58 Cal. ; [H,Cl] = + 22 Gal. H20 (gas) + 2Cl (gas) = 2HC1 (gas) + 0 (gas) - 14 Gal. (2). Liquid Water and Chlorine Gas. H,O (liquid) + 2C1 (gas) + Aq = 2HClAq + 0 (gas) -t- 10 Cal. [H2,0] liquid = 68.4 Cal. ; [HCl,Aq] = + 17.2 Cal. Under similar conditions, in a cool cellar and an atmosphere of It is thus evident that : (1). Oxygen is present among the gaseous products of the electro- lysis of hydrochloric acid. (2). Even though traces of the lower chlorine oxides may be formed during the electrolysis, it is unlikely that any escape a preliminary washing of the gases. moist chlorine in the globe 6, a negative result was 0btained.l 11.THE SOLUBILITY OF CHLORINE IN AQUEOUS HYDROCHLORIC ACID. When a mixture of equal volumes of hydrogen and chlorine in the presence of water saturated with the two gases is exposed to the action of light, hydrogen chloride is formed by the water a t a rate pro- portional to the intensity of the light. Thinking that this absorption of the hydrogen chloride might disturb the equilibrium of the gases in the insolation vessel of Bunsen and Roscoe’s actinometer, I have in- vestigated the solubility of chlorine in water containing varying quantities of hydrogen chloride at a constant temperature. See the various reports to the British Association collected by Richardson (23. A. Reports, 1888, 89 ; 1889, 69 ; 1890, 263). Bunsen and Roscoe (this Journ., 1856, 8, 190), investigating Wittwer’s proposal (Pogg. Ann., 1855, 94, 527) to measure the chemical action of light by the decomposition of chlorine water [say, 2C1, (liquid) + 2H20 (liquid) = 4HC1 (liquid) + 0, (gas)], found that the presence of hydrochloric acid greatly retarded the action.226 MELLOR: ON THE UNION OF No systematic work appears to have been done on this subject.Three isolated records were all I could find. I n 1856 Roscoe 1 found that the presence of 1/120th part of hydro- gen chloride lowered the value of the coefficient of absorption of chlorine in water from a=2*3911 to 1.9789 at 14’. (1880) found that a 38 per cent. solution of hydrogen chloride absorbs 17-3 grams of chlorine per litre ; a solution containing 1/3HC1, that is, 33 per cent., 3 absorbs 11.0 grams; whilst a 3 per cent.solution absorbs 6 grams of chlorine per litre (temperature not stated). Berthelot suspects the formation of hydrogen perchloride, HCl,, in strong solutions, and quotes thermochemical data in support of this view. Goodwin 4 (1882) investigated the influence of temperature on the solubility of the following different strengths of acid and found : Berthelot Hydrochloric acid of sp. gr. 1.046, a=2*5403 at 23.6’ (752 mm.). ?, 1, ,, 1.080, a=4*1433 at 15.5’ (763 mm.), 9, 9 , ,, 1.125, a=4-7631 at 20.7’ (762 mm.). In my preliminary work I found that Heidenhain and V. Meyef’s method for saturating the liquid by shaking with the gas did not work so satisfactorily as the one described below. Chlorine, evolved from the liquid, was washed in boiled distilled water, then in chromic acid solution, and again in water, The gas was then passed into a vessel containing the given solution until two titrations, with sodium thiosulphate, showed constant figures, The saturation vessel stood in a water-bath maintained at a temperature of 20-21’ by means of a This work was done in a dark cellar.Ten C.C. of the saturated solution were run 7 into an aqueous N/10 solution of potassium iodide. The free iodine was determined by Roscoe, this Journ., 1856, 8, 14. Berthelot, Ann. Chim. Phys., 1881, [v], 22, 462 ; or Compt. rend., 1880, 91, 194. 3 A. M. Comey, “Dictionary of Solubilities,” 105, 1896, translates this as one- third of the 38 per cent. solution, and, therefore, wrongly reads “ 12.7 per cent. HCl absorbs 11 grains of chlorine per litre.” Goodwin, Trans.Roy. SOC. Edin., 1882, 30, 597 ; or Ber., 1882, 15, 3039. Timofeeff, Zeit. physikal. Chem., 1890, 6, 141. This temperature was chosen t o eliminate, as far as possible, any disturbance due to the formation of hydrates (compare Roozeboom, Rec. trau. Pays-Bas, 1884, 3, 59 ; 1885, 4, 69 ; Isambert, Conzpt. rend., 1878, 86, 481). The slight loss of chlorine involved in the withdrawal of this liquid by a suction pipette is avoided by using a pipette similar to that described by Reid (Chem. News, 1892, 66, 167), or Rnting’s patent pipette (Zeit. physikal. Chem., 1899, 29, 626).HYDROGEN AND CHLORINE. PART8 I TO 111. 227 means of standardised sodium thiosulphate solution in the usual way. The loss due to the decomposing effect of hydrochloric acid on the thiosulphate was negligibly small, since the thiosulphate was added direct to the iodine.2 Let 0 be the temperatuqp the height of the barometer in mm.of mercury, u the coefficient of absorption, X the coefficient of solubility, n, the number of C.C. of the standard thiosulphate required in titrating, w the volume of chlorine in C.C. absorbed by the given solution. Since each C.C. of the thiosulphate solution was equivalent to 0*01386 gram of iodine, or 0*003848 gram or 1.2127 C.C. of chlorine, that is, v = 1.2127, 273 p v a=- 273+8.760 3' For chlorine in pure water a t 2l0, u=2*1167, Schijnfelds gives 2.1148. The variations in the barometer readings were so small that fheir influence on the results is well within the errors of experiment.Temperature variations were, for the same reason, neglected. Hence : 273+0=n. 273 x 1.2127. Absorption coeficients of chlorine. Grams HCI Per 1000 C.C. 313.401 282*060 250'720 2 19 '380 188'040 156.700 125 *3 60 94.020 62.680 31 *340 15'670 12.540 9 *402 6.248 3.134 nil 2%. 31.52 29 *57 27.77 25'82 24-01 22.30 20.18 18.73 16'60 14-87 13-27 12.61 12.38 12.87 1377 18-80 P. 761 761 759 759 761 761 761 760 760 762 759 759 760 760 760 760 ~ t". 21 *o 21.0 21'0 21.0 20'2 20 '2 20 *5 20 *5 20.0 31'0 21'0 21.0 21.0 20 '0 21'0 21 '0 Grams C1, Per 1000 C.C. 12.03 11-87 10.68 9 -93 9 '23 8-58 7 *76 7'19 6'38 5 '81 5-10 4-85 4-76 4'94 5 *30 i 7'23 a. 3'5492 3 '3278 3.1272 2,9243 2.7020 2'5095 2.2711 2'1044 1'8682 1'6736 1 '4933 1 '4200 1'3942 1'4483 1.5496 2.1157 ~ A.3.8224 3.6859 3.3677 3.1312 2.9117 2.7043 2.4473 2,2677 2.0131 1.8033 1.6092 1.5292 1.5013 1.5607 1.6698 2.2799 When the amounts of hydrogen chloride contained in the solution are plotted as abscissz against the amounts of chlorine absorbed, two n, Norton, Arner. J. Xci., 1899, [vii], 7, 237. 3 Schonield, Annalen, 1855, 93, 26. Pickering, this Journ., 1870, 37, 135.228 MELLOR: ON THE UNION OF distinct curves appear. The one is subsequently referred to as the ‘I curve of dissociation,” the other as the ‘‘ curve of association.” The first action of chlorine on dilute hydrochloric acid is apparently represented by some exponential curve which intersects another simpler linear curve represented by the equation : p=ap+b where a and 6 are constants approximately equal to 0.023 and 4.92 respectively, and p and p respectively denote the amounts of hydrogen chloride and of chlorine per 1000 C.C.of solution. Grams of HCl per 100 C.C. = p. Isothemnal curve of the solubility of chlorine in aqueous solutions of hydrochloric acid of varying concentration. Jakowkinl has proved that the action of chlorine on water is a reversible dissociation somewhat in the form of the equation : CI,Aq t (HC1+ HOC1)Aq. He takes advantage of the fact that undissociated chlorine divides in a known ratio between water and carbon tetrachloride, in order to Jakowkin, Zeit. physikal. Chem., 1899, 29, 613.HYDROGEN AND CHLORINE, PARTS I TO 111. 229 determine the amounts of dissociated and undissociated chlorine in aqueous solution, hydrochloric and hypocblorous acids being insoluble in carbon tetrachloride.According to the dissociation theory, HClAq t. (H + Cl)Aq, while the hypochlorous acid suffers little if any dissociation. The latter statement is confirmed by (1) its weak acidity; (2) its easy hydrolysis with strong bases; (3) the rapid increase of its molecular conductivity with dilution ; and (4) its normal molecular weight by cryoscopic methods. + - Even in darkness the action is termolecular, 3 . - C1,Aq (H + C1+ HOCl)Aq, and is therefore denoted by the formula where A represents the total number of gram-molecules of chlorine in the solution, c that of the undissociated chlorine, A - c the number of dissociated chlorine molecules in the solution. The total volume of a substance taken up by unit volume of solvent is often referred to as the apparent solubility, whilst the amount of substance which remains unchanged in unit volume of solution is termed the real solubitity. Applying Nernst’s distribution law,l A = c + ( A - c ) , + + and since for every C1 ion that goes to form HOCl one H ion is set free, we have, according to the mass law, c R = { A - c ) ~ .’ (2) where K has not necessarily its former vahe. If c2 is the concentration of a second electrolyte which has one ion in common with the ions already in solution, we have for H or C1 ions, as before, and But since A’ = d + (A’ - c’), c’R’ = (A‘ - c‘)(A’ - C’ + c2) . (3) ( A - c ) < ( A ‘ - c ’ ) cf > c, that is, the amount of undissociated chlorine will be increased by the addition of a second electrolyte containing either H or C1 ions ; hence Nernst, Zeit.physikal. C%enz., 1889, 4, 372 ; Noyes, i6id., 1890, 6, 241. VOL. LXXIX. R230 MELLOR: ON THE UNION OF it follows from the ‘‘ theorem of constant solubility ” that the soh- bility of chlorine will be diminished. Let X represent the apparent solubility, and x the real solubility of chlorine in hydrochloric acid, x the amount of chlorine or of disso- ciated (Ht-C1) in solution according to Jakowkin’s equation, y the amount of chlorine in solution assumed, for the present, to be in some way combined with HCl. Hence + - X = x + x +ye For dilute solutions we should have x = x + x . The value of x is easily calculated from the dissociation data compiled by Fitzgerald in the Reports to the British Association, 1893.l The results are not altogether in accord with experiment, showing that under these conditions x is either not constant, or y cannot be neglected.Beyond a certain limit, however, y becomes relatively large, while x becomes small. That is assuming x to be constant. This diminution in the solubility of chlorine can be readily shown in a qualitative way, by adding a few C.C. of concentrated hydrochloric acid to a quantity of saturated chlorine water contained in a narrow vessel. Bubbles of chlorine soon form and escape to the surface. It is evident from the termolecular action of chlorine on water : (1) The addition of an electrolyte capable of supplying C1 ions to the solution will cause a diminution in the solubility of chlorine.Thus Kumpf,2 G~odwin,~ and Jakowkin3 have shown this diminu- tion in the solubility of chlorine in saturated solutions of alkali chlorides. The separation of sodium chloride, when chlorine is passed into saturated aqueous solutions, is another consequence of the same law. Conversely, Engel has shown4 that the addition of hydrogen chloride diminishes the solubility of electrolytic chlorides. Non- electrolytic chlorides, however, do not influence tho solution unless molecular association occurs : for example, mercuric chloride. (2) The addition of an electrolyte capable of supplying H ions must also effect a reduction in the solubility of the chlorine. This has been proved, for the electrolytic acids, nitric acid, hydro- 1 Reprinted in Whetham’s ‘‘ Solution and Electrolysis,” 215, (1895).2 Kumpf (Innug. Dissert.) Wied. Biebl., 1882, 6, 276. LOG. cit. Engel, Bull. SOG. Chim., 1889, [iii], 1, 695 ; or ConLpt. rend., 1889, 104, 1710 ; Ditte, Aim. Chirn. Phys., 1897, [vii], 10, 556.HYDROGEN AND CHLORINE. PARTS I TO 111. 231 chloric acid, sulphuric acid, and acetic acid, and its three chloro-deriva- tives. Non-electrolytic acids have no influence on the result: for example, boric acid. (3) The addition of hypochlorous acid reduces the solubility of chlorine in water, since it acts in virtue of the change, C1,Aq = (HC1 + H0Cl)Aq. Hydrogen chloride, supplying as it does both H and C1 ions, has a very marked influence in dilute solution. This is shown in Fig. 3, p. 228, as an isothermal curve of dissociation. The more concentrated the solution of hydrogen chloride the less the dissociation, For very concentrated solutions of hydrogen chloride we should expect a bimole- cular action : C1,Aq (HC1+ HOCl)Aq, and from equation (1) - (4) ( A - c ) ~ K= ~ cK= ( A - c)(A - c + x), C or where x denotes the number of hydrogen chloride molecules added to the solution.When the amount of dissociated chlorine (z) in the solution is exactly equivalent to the amount in combination (y) x, or y = g(X - x). From Berthelot's original paper,l the conclusion may be drawn that there is a concentration of hydrogen chloride having a maximum - NO. - 1 4 5 2 6 3 7 8 7 C.C. Thio- sulphate. 17-7 18'9 21 '7 23'4 31.1 31 *8 34.9 36-2 Gram HCl Per 100 gram If solution. 2'90 3 -22 11 ti3 12.19 31'24 32-00 34.57 35.90 Gram C1 per litre.6-0 6.3 7 -2 7 '76 10'3 10.55 11-6 12.7 t. ~ 15-0 16.3 1 6 3 15'0 16'3 15.0 16.3 16'3 Berthelot. 'e&Et' Gram C1. 1 ---I I- 759 763 763 759 763 759 763 763 3 3 - - 33 33 38 38 6 -0 6-0 - - 11.0 11.0 17'3 17-3 The number 17.3 for the weight of chlorine in a 38 per cent. solution of hydro- chloric acid attributed above to Berthelot is given in his paper as 7-3. He has recently informed me that this is a misprint. The mistake also occurs in Comey's Dictionary. The experiments in the text were made a t the end of 1899 in the attempt to find a point of maximum solubility, before i t was found that 7.3 was not the true number. R 2232 MELLOR: ON THE UNION OF power of absorption for chlorine. Hydrochloric acid saturated with hydrogen chloride in the cold was used for the strongest solution.The chlorine was passed into the solutions as indicated above. Free chlorine was determined by the usual thiosulphate titration, total chlorine by boiling 10 C.C. of the saturated solution with ferrous sul- phate and aqueous potassium hydroxide. The chlorine was then pre- cipitated as chloride from the solution, acidified with nitric acid, and weighed in the usual way. One C.C. of thiosulphate = 0.003319 gram of chlorine. There is thus no indication of a point of a maximum followed by a diminishing solubility with increasing concentration. An objection to the preceding method of finding the solubility of chlorine in the stronger solutions might be pointed out. I f the current of gaseous chlorine occupies any considerable time, the acid will tend to attain that particular concentration which has a constant com- position, and at equilibrium, the chlorine and hydrogen chloride will be distributed according to their partial pressures and the phase rule. On the Zxistence of HCI, in Liquid Solution.The curve of association indicates the possibility of the existence of some combination of hydrogen chloride and chlorine, possibly stable only in the presence of a great excess of hydrogen chloride, just as the great quantity of chlorine retained by strong solutions led Draper 2 to believe in the existence of a, ‘‘ bichloride of hydrogen,” and Ber- thelot of a g‘ perchloride of hydrogen.” I n a quite analogous way, Engel found that the solubility of certain chlorides is increased if hydrochloric acid is present in the solution owing to formation of ‘( chlorhydrates.” Some of these were isolated ; for instance, those of Etannic, ferric, cupric, and mercuric chlorides.The following evidence for the existence of HCI, might be cited : (1). The existence of other well-established tri- and penta-halides.4 (2). The partition-coefficient of iodine, between aqueous solutions of potassium iodide and carbon disulphide, leads t o the formula KI.1, (Jak~wkin).~ 1 Perman, Trans., 1895, 67, 868. Draper, Phi2. Mag., 1843, [iii], 23, 431. 3 Millon (J. Pharm., 1841, 28, 299) regarded the yellow liquid remaining when lend chloride is removed by the cooling of the products of the interaction of lead peroxide and concentrated hydrochloric acid according to the equation PbO, + 6HC1= PbC], + 2H2O+2HC1,.It is now generally believed that the action is Pb02+4HC1= PbCl,+ 2H,O, although the action PbO,+ 5HC1=PbCl2+ 2H,O + HC13 appears to be equally probable. 4 Conipare Wells, Amer. J. Sci., 1892, [iii], 43, 17 ; Wells and Wheeler, ibid., 1892, 44, 42, 475. 5 Jakowkin, Zeit. physikal. Chm., 1894, 13, 539. Compare Wilderman’s BrHBr,, &c., ibid, 1893, 11, 407 ; Noyes and Seidenstraker, ibid., 1898, 27, 357.HYDROGEX AND CHLORINE. PARTS I TO 111. 233 (3). The increasing solubility of chlorine with increasing additions of hydrogen chloride, arid the analogy with Engel’s “ chlorhydrates.” (4). The heat disengaged by solutions of chlorine in concentrated hydrochloric acid approaches that required for HCl,, and resembles that required for KI, and KBr, (Berthelot).On the Existence of HCI, in Guseous Solution. Evidence for the existence of gaseous HCI, was sought by bringing hydrogen chloride and chlorine together in the dark by means of the apparatus devised by Dixon and Harker.l A slightly greater con- traction occurred, differing from that with hydrogen chloride and air, or with chlorine and air. This might be attributed either to a con- densation of the gases on the glass or else to some form of molecular attraction between hydrogen chloride and chlorine. The only evidence of chemical combination between certain gases is a slight difference between the total volume occupied by the separate and by the mixed gases. It is assumed that if no chemical combination takes place, the mixture will obey Dalton’s law of partial pressures, namely, ‘‘ the pressure exerted by a mixture of gases is equal to the sum of the pressures separately exerted by the several components.” The work of Regnault on mixtures of air with carbon dioxide and with sulphur dioxide ; of Andrews on mixtures of carbon dioxide with air and with nitrogen, and of Braun4 on mixtures of sulphur dioxide and carbon dioxide, sulphur dioxide and hydrogen, sulphur dioxide and nitrogen, hydrogen and carbon dioxide, hydrogen and air, hydrogen and nitrogen, and carbon dioxide and air, shows that Dalton’s simple law of addition is not strictly followed.Dalton’s law assumes : (I) That each component exerts the same pressure in, the mixture that it would if it occupied the space alone.To avoid this limitation, Sarrau5 has proposed to recast Dalton’s law somewhat in this form : the specific volume of a mixture of gases is equal t o the sum of the specific volumes of the several components. Leduc and Sacerdote6 find that in this form the law agrees better with the results of their experiments. 1 Dixon and Harker, M e m . and Proc. Maitchester Lit. Phil. Soc., 1890, [iv], 2 Regnault, Mem. de I’Acad., 1862, 26, 256. 8 Andrews, Phil, Mag., 1876, [ v ] , 1, 7 8 ; Phil. Trans., 1888, 178, 57. 4 Braun, Wied. Ann,, 1888, %> 943. G Sarrau’s “ Introduction h la Theorie des Explosifs,” 25, (1895). 3,118. Compt. rend., 1898, 126, 218, 1853 ; Leduc’s “ Recherches sur les gaz,” 106, (1898).234 MELLOR: ON THE UNION OF D. Berthelot 3 has deduced an expression from van der Waals’ equa- tion which gives results in close agreement with experiment for the change of pressure accompanying the mixing of gases, for example : Observed increase.0.0001 1 atm. (Sacerdote). CO, + so, ... ... 0*0019 ,, 1 0*0018 ,, CO, + N,O , . . . . . 0°00013 atm. 9 , Calculated increase. These calculations are based on the assumption that the change of pressure which accompanies the expansion of each component of the mixture obeys van der Waals’ modification of Boyle’s law. Some interesting examples in which the final pressure is obscured by the dissociation of one of the components of the mixture are treated in a recent paper by Professor Dixon and J. D. Peterkin., (2). That the molecules of the diferent gases exert nezther attractive nor repdsive forces on one another.According to the kinetic theory, intermolecular attraction will (1) increase the number of collisions between the molecules, (2) cause certain molecules to swerve from their normal rectilinear path, (3) diminish the outward pressure of the gas. The molecules of such a gas are only attracted from within, its volume will therefore be less than that of a gas containing the same number of non-attracting molecules subject t o the same external pressure. It is proved in works on the kinetic theory of gases that for every molecule that loses its motion by collision, another will acquire the same motion by another simultaneous collision. That is to say, unless the attracting molecules during a collision remain in contact a longer time than non-attracting molecules, their motion will go on just the same as if there were no collision at all.Sutherland 3 by assuming that this attractive force varies in- versely as the fourth power of the distance between the molecules, has deduced very satisfactory formulse to explain certain physical proper- ties of gases. For instance, Lord Kelvin and Joule found that the cooling effect produced when n, mixture of gases undergoes expansion is not exactly the value calculated on the assumption that there are no attractive forces between the molecules. Sutherland, applying his lam of the inverse fourth, obtains results in harmony with experiment. Similarly with the variation of viscosity with temperature, diffusion of gases, &c. The term av-2 in van der Waals’ equation is intended to allow for 1 D.Berthelot, Contpt. rend., 1898, 126, 954, 1030, 1415, 1703, 1857 ; 1899, 2 Dixon and Peterkin, Trans., 1899, 75, 613. 3 Sutherland, Phil. Mag., 1893, [v], 36, 507 ; also 1886, [v], 22, 81 ; 1895, [v], 128, 1159 ; Leduc, ibid., 1898, 126, 1859 ; Van der Waals, ibid., 126, 1856. 40, 433.HYDROGEN AND CHLORINE. PARTS I TO 111. 235 the effects of the attraction OF the molecules when gases undergo certain changes in volume under the influence of a varying pressure. In a private communication last June, M. D. Berthelot pointed out to me that there is no reason to suppose that the increase of pressure calculated for the mixture of hydrogen chloride and chlorine would differ very much from that for the mixture of carbon dioxide and sulphur dioxide, unless the chlorine exercised some action on the walls of the vessel.Of course this action may to some extent be allowed for by a preliminary saturation of the walls of the vessel with chlorine and comparative experiments with other gases. Any slight contraction, therefore, which might occur on mixing two gases (say hydrogen chloride and chlorine) cannot be taken as con- clusive evidence of a chemical combination (say, formation of HCIJ until it has been shown that intermolecular forces are inadequate to account for the discrepancy. 111. THERMODYNAMICS OF SOLUTIONS OF CHLORINE AND HYDROGEN CHLORIDE IN WATER. If hydrogen chloride be added to a saturated solution of chlorine water in equilibrium with its atmosphere, the chlorine will be redis- tributed until the potential energy of the system attains a minimum value.From the properties of the thermodynamical potential, it can be shown t h a t for an increase 6v in the amount of HCl present, there will be an increase or a decrease in the amount of chlorine retained by the solution, according as the van’t Hoff factor i is less or greater than unity. J. Willard Gibbsl (1876) has shown that the differential of the energy of any material system, subject to gain or loss of energy and of mass, is expressed by the equation : dU= ed+ -pdv + pldm, + . . . pndmn . . - (1) where U denotes the energy, 8 the absolute temperature, CP the en- tropy, p the pressure, IJ the volume, ,U the (‘ potential ” (Gibbs) or ‘‘ intensity ” (Helm) factor expressing the rate of increase of energy in a reversible increase of unit mass with constant volume energy, namely : where rn denotes the mass of the body, and as a suffix implies that all other m’s in the above formula are constant.See H. le Chatelier’s ‘‘ Equilibre des S y s t h e s Chimiques, par J. Willard Gibbs,” 54, (1899).236 MELLOR: ON THE UNION OF Putting, with Duhem, aj=U-OQ,+pv . . . . . (2) d@ = - + d O + v d p + S p h . . . . . (3) we get from (1) where CP is clearly a quantity depending only on the parameters describing the particular state of the system, that is, @ is R complete diff erentia1,l and a2Qr - a2Qr aX,aX, ax,ax,’ - - _ _ Consider now the work (W) gained during an isothermal compres- sion from an initial pressure po to a greater pressure pr and, integrating by the aid of the ordinary gas equation, v = RBlog,31 , .. . . ( 5 ) P O By differentiation of (2) d@ = vdp- +do, and since d@ is a perfect differential, as in Massieu’s well known functions, from (4), therefore da? = vdp It now remains to show tbat @ is a function of the amount of chlorine in solution, or that where X is the coefficient of solubility, The extension of the gas laws to dilute solutions by Arrhenius, van’t Hoff, and Nernst, enables the various components of the mixture to be expressed as functions of the parameters describing the thermo- dynamic state of the mixture. Nernst’s distribution law allows us to replace the uapowr phase in the extension of Gibb’s equation to Henry’s law, by a second liquid phase. Let the formula @ = f ( 0 pu = i K 8 .. . . . . . (7) 1 Duhem’s ‘‘ Le Potential Thermodynaniique,” 33, (189s) ; Trevor, J. Phy&Z Ch., 1897, 1, 205, 633.HYDROGEN AND CHLORINE. PARTS I TO 111. 237 be applicable to dilute solutions of electrolytes, p n o w representing the osmotic pressure of the dissolved substance, i the isotonic coeffici- ent greater than unity for dissociated substances, that is to say i = a gram-molecule of the substance in solution. From (2), ( 6 ) , and (7), if a, now represent the potential of a given mass, i, of chlorine in dilute solution of hydrogen chloride, Q2 that of a more dilute solution, we have at constant temperature, where p , and p , represents the osmotic pressures of the chlorine in. the two solutions, hence d@, = iRBdhgeg2. PI If the molecular weight i of chlorine in the two solutions is the Barnel then, by the properties of dilute solutions, where Cl and C, are the cdfiegntrations of the two solutions C =- . ( t) We have, therefore, the relation d@=iI&dlog$2 . . . I . - , (9) c, = df (A> I n words the potential energy of chlorine in dilute solutions of hydrogen chloride is increased by a further addition of the latter gas. By Helmholtz's law, any dynamical system behaves so that the decrease in the potential energy may be a maximum, therefore the solution under these conditions cannot dissolve so much chlorine. If otherwise, the gas is either present in a supersaturated state, that is, in a state of unstable equilibrium (I' faux Bquilibres," Duhem), or else the stipulation that the molecular weight, i, of chlorine is constant in the different solutions, no longer holds. In stronger solutions the same solution has phases in which the. van't Hoff factor may be b l , i"<l, i = l . From the earlier part of this paper it follows that there is a, considerable variation in the relative values of i for chlorine in the different solutions of hydrogen chloride. For the curve of dissociation i'>l, and for the curve of association i"<l. In the more dilute solu- tions of HCl, the relative proportion of i'> 1 is the greater, while in concentrated solutions i"<l is the larger. At the point of inter-238 DAWSON: ON THE NATURE OF POLYIODIDES AND THEIR section of these two curves, we can say no more than that the amount of chlorine (Sx) for which i’>1 may be equal to the amount 89 for which i < 1, or dx: dy dz dy +=@ dp dp or ---=O. where p is the amount of hydrogen chloride in the solution. It is then evident from an equation similar to (9) that if, corre- sponding to HCl,, iff < 1, the solution will dissolve more chlorine in order that the potential energy may be a minimum, an inference which may be deduced from (1) when the system includes another term p2dm,. There is, however, a considerable amount of uncertainty as to what actually takes place in these and all other concentrated solutions. No further progress can be made in a quantitative way until this has been determined. A consistent theory for concentrated solutions is wanting. THE OWENS COLLEGE, M ANCHESTER.