[61]In general, 2HI + O = I2+ H2O, if the oxygen proceed from a substance from which it is easily evolved. For this reason compounds corresponding with the higher stages of oxidation or chlorination frequently give a lower stage when treated with hydriodic acid. Ferric oxide, Fe2O3, is a higher oxide, and ferrous oxide, FeO, a lower oxide; the former corresponds with FeX3, and the latter with FeX2, and this passage from the higher to the lower takes place under the action of hydriodic acid. Thus hydrogen peroxide and ozone (ChapterIV.) are able to liberate iodine from hydriodic acid. Compounds of copper oxide, CuO or CuX2, give compounds of the suboxide Cu2O, or CuX. Even sulphuric acid, which corresponds to the higher stage SO3, is able to act thus, forming the lower oxide SO2. The liberation of iodine from hydriodic acid proceeds with still greater ease under the action of substances capable of disengaging oxygen. In practice, many methods are employed for liberating iodine from acid liquids containing, for example, sulphuric acid and hydriodic acid. The higher oxides of nitrogen are most commonly used; they then pass into nitric oxide. Iodine may even be disengaged from hydriodic acid by the action of iodic acid, &c. But there is a limit in these reactions of the oxidation of hydriodic acid because, under certain conditions, especially in dilute solutions, the iodine set free is itself able to act as an oxidising agent—that is, it exhibits the character of chlorine, and of the halogens in general, to which we shall again have occasion to refer. In Chili, where a large quantity of iodine is extracted in the manufacture of Chili nitre, which contains NaIO3, it is mixed with the acid and normal sulphites of sodium in solution; the iodine is then precipitated according to the equation 2NaIO3+ 3Na2SO3+ 2NaHSO3= 5Na2SO4+I2+ H2O. The iodine thus obtained is purified by sublimation.[62]For the final purification of iodine, Stas dissolved it in a strong solution of potassium iodide, and precipitated it by the addition of water (seeNote58).[63]The solubility of iodine in solutions containing iodides, and compounds of iodine in general, may serve, on the one hand, as an indication that solution is due to a similarity between the solvent and dissolved substance, and, on the other hand, as an indirect proof of that view as to solutions which was cited in ChapterI., because in many instances unstable highly iodised compounds, resembling crystallo-hydrates, have been obtained from such solutions. Thus iodide of tetramethylammonium, N(CH3)4I, combines with I2, and I4. Even a solution of iodine in a saturated solution of potassium iodide presents indications of the formation of a definite compound KI3. Thus, an alcoholic solution of KI3does not give up iodine to carbon bisulphide, although this solvent takes up iodine from an alcoholic solution of iodine itself (Girault, Jörgensen, and others). The instability of these compounds resembles the instability of many crystallo-hydrates, for instance of HCl,2H2O.[64]The equality of the atomic volumes of the halogens themselves is all the more remarkable because in all the halogen compounds the volume augments with the substitution of fluorine by chlorine, bromine, and iodine. Thus, for example, the volume of sodium fluoride (obtained by dividing the weight expressed by its formula by its specific gravity) is about 15, of sodium chloride 27, of sodium bromide 32, and of sodium iodide 41. The volume of silicon chloroform, SiHCl3, is 82, and those of the corresponding bromine and iodine compounds are 108 and 122 respectively. The same difference also exists in solutions; for example, NaCl + 200H2O has a sp. gr. (at 15°/4°) of 1·0106, consequently the volume of the solution 3,658·5/1·0106 = 3,620, hence the volume of sodium chloride in solution = 3,620–3,603 (this is the volume of 200 H2O) = 17, and in similar solutions, NaBr = 26 and NaI = 35.[65]But the density (and also molecular volume, Note64) of a bromine compound is always greater than that of a chlorine compound, whilst that of an iodine compound is still greater. The order is the same in many other respects. For example, an iodine compound has a higher boiling point than a bromine compound, &c.[66]A. L. Potilitzin showed that in heating various metallic chlorides in a closed tube, with an equivalent quantity of bromine, a distribution of the metal between the halogens always occurs, and that the amounts of chlorine replaced by the bromine in the ultimate product are proportional to the atomic weights of the metals taken and inversely proportional to their equivalence. Thus, if NaCl + Br be taken, then out of 100 parts of chlorine, 5·54 are replaced by the bromine, whilst with AgCl + Br 27·28 parts are replaced. These figures are in the ratio 1 : 4·9, and the atomic weights Na : Ag = 1 : 4·7. In general terms, if a chloride MClnbe taken, it gives withnBr a percentage substitution = 4M/n2where M is the atomic weight of the metal. This law was deduced from observations on the chlorides of Li, K, Na, Ag (n= 1), Ca, Sr, Ba, Co, Ni, Hg, Pb (n= 2), Bi (n= 3), Sn (n= 4), and Fe2(n= 6).In these determinations of Potilitzin we see not only a brilliant confirmation of Berthollet's doctrine, but also the first effort to directly determine the affinities of elements by means of displacement. The chief object of these researches consisted in proving whether a displacement occurs in those cases where heat is absorbed, and in this instance it should be absorbed, because the formation of all metallic bromides is attended with the evolution of less heat than that of the chlorides, as is seen by the figures given in Note55.If the mass of the bromine be increased, then the amount of chlorine displaced also increases. For example, if masses of bromine of 1 and 4 equivalents act on a molecule of sodium chloride, then the percentages of the chlorine displaced will be 6·08 p.c. and 12·46 p.c.; in the action of 1, 4, 25, and 100 molecules of bromine on a molecule of barium chloride, there will be displaced 7·8, 17·6, 35·0, and 45·0 p.c. of chlorine. If an equivalent quantity of hydrochloric acid act on metallic bromides in closed tubes, and in the absence of water at a temperature of 300°, then the percentages of the substitution of the bromine by the chlorine in the double decomposition taking place between univalent metals are inversely proportional to their atomic weights. For example, NaBr + HCl gives at the limit 21 p.c. of displacement, KCl 12 p.c. and AgCl 4¼ p.c. Essentially the same action takes place in an aqueous solution, although the phenomenon is complicated by the participation of the water. The reactions proceed spontaneously in one or the other direction at the ordinary temperature but at differentrates. In the action of a dilute solution (1 equivalent per 5 litres) of sodium chloride on silver bromide at the ordinary temperature the amount of bromine replaced in six and a half days is 2·07 p.c., and with potassium chloride 1·5 p.c. With an excess of the chloride the magnitude of the substitution increases. These conversions also proceed with the absorption of heat. The reverse reactions evolving heat proceed incomparably more rapidly, but also to a certain limit; for example, in the reaction AgCl + RBr the following percentages of silver bromide are formed in different times:hours232296120K79·8287·488·22—94·21Na83·6390·7491·7095·49—That is, the conversions which are accompanied by an evolution of heat proceed with very much greater rapidity than the reverse conversions.[67]The dissociation of hydriodic acidhas been studied in detail by Hautefeuille and Lemoine, from whose researches we extract the following information. The decomposition of hydriodic acid is decided, but proceeds slowly at 180°; the rate and limit of decomposition increase with a rise of temperature. The reverse action—that is, I2+ H2= 2HI—proceeds not only under the influence of spongy platinum (Corenwinder), which also accelerates the decomposition of hydriodic acid, but also by itself, although slowly. The limit of the reverse reaction remains the same with or without spongy platinum. An increase of pressure has a very powerful accelerative effect on the rate of formation of hydriodic acid, and therefore spongy platinum by condensing gases has the same effect as increase of pressure. At the atmospheric pressure the decomposition of hydriodic acid reaches the limit at 250° in several months, and at 440° in several hours. The limit at 250° is about 18 p.c. of decomposition—that is, out of 100 parts of hydrogen previously combined in hydriodic acid, about 18 p.c. may be disengaged at this temperature (this hydrogen may be easily measured, and the measure of dissociation determined), but not more; the limit at 440° is about 26 p.c. If the pressure under which 2HI passes into H2+ I2be 4½ atmospheres, then the limit is 24 p.c.; under a pressure of ⅕ atmosphere the limit is 29 p.c. The small influence of pressure on the dissociation of hydriodic acid (compared with N2O4, Chapter VI. Note46) is due to the fact that the reaction 2HI = I2+ H2is not accompanied by a change of volume. In order to show the influence of time, we will cite the following figures referring to 350°: (1) Reaction H2+ I2; after 3 hours, 88 p.c. of hydrogen remained free; 8 hours, 69 p.c.; 34 hours, 48 p.c.; 76 hours, 29 p.c.; and 327 hours, 18·5 p.c. (2) The reverse decomposition of 2HI; after 9 hours, 3 p.c. of hydrogen was set free, and after 250 hours 18·6 p.c.—that is, the limit was reached. The addition of extraneous hydrogen diminishes the limit of the reaction of decomposition, or increases the formation of hydriodic acid from iodine and hydrogen, as would be expected from Berthollet's doctrine (ChapterX.). Thus at 440° 26 p.c. of hydriodic acid is decomposed if there be no admixture of hydrogen, while if H2be added, then at the limit only half as large a mass of HI is decomposed. Therefore, if an infinite mass of hydrogen be added there will be no decomposition of the hydriodic acid. Light aids the decomposition of hydriodic acid very powerfully. At the ordinary temperature 80 p.c. is decomposed under the influence of light, whilst under the influence of heat alone this limit corresponds with a very high temperature. The distinct action of light, spongy platinum, and of impurities in glass (especially of sodium sulphate, which decomposes hydriodic acid), not only render the investigations difficult, but also show that in reactions like 2HI = I2+ H2, which are accompanied by slight heat effects, all foreign and feeble influences may strongly affect the progress of the action (Note47).[68]The thermal determinations of Thomsen (at 18°) gave in thousands of calories, Cl + H = +22, HCl + Aq (that is, on dissolving HCl in a large amount of water) = +17·3, and therefore H + Cl + Aq = +39·3. In taking molecules, all these figures must be doubled. Br + H = +8·4; HBr + Aq = 19·9; H + Br + Aq = +28·3. According to Berthelot 7·2 are required for the vaporisation of Br2, hence Br2+ H2= 16·8 + 7·2 = +24, if Br2be taken as vapour for comparison with Cl2. H + I = -6·0, HI + Aq = 19·2; H + I + Aq= +13·2, and, according to Berthelot, the heat of fusion of I2= 3·0, and of vaporisation 6·0 thousand heat units, and therefore I2+ H2= -2(6·0) + 3 + 6 = -3·0, if the iodine be taken as vapour. Berthelot, on the basis of his determinations, gives, however, +0·8 thousand heat units. Similar contradictory results are often met with in thermochemistry owing to the imperfection of the existing methods, and particularly the necessity of depending on indirect methods for obtaining the fundamental figures. Thus Thomsen decomposed a dilute solution of potassium iodide by gaseous chlorine; the reaction gave +26·2, whence, having first determined the heat effects of the reactions KHO + HCl, KHO + HI and Cl + H in aqueous solutions, it was possible to find H + I + Aq; then, knowing HI + Aq, to find I + H. It is evident that unavoidable errors may accumulate.[69]One can believe, however, on the basis of Berthollet's doctrine, and the observations of Potilitzin (Note66), that a certain slow decomposition of water by iodine takes place. On this view the observations of Dossios and Weith on the fact that the solubility of iodine in water increases after the lapse of several months will be comprehensible. Hydriodic acid is then formed, and it increases the solubility. If the iodine be extracted from such a solution by carbon bisulphide, then, as the authors showed, after the action of nitrous anhydride iodine may be again detected in the solution by means of starch. It can easily be understood that a number of similar reactions, requiring much time and taking place in small quantities, have up to now eluded the attention of investigators, who even still doubt the universal application of Berthollet's doctrine, or only see the thermochemical side of reactions, or else neglect to pay attention to the element of time and the influence of mass.[70]On the basis of the data in Note68.[71]A number of similar cases confirm what has been said in ChapterX.[72]This is prevented by the reducibility of sulphuric acid. If volatile acids be taken they pass over, together with the hydrobromic and hydriodic acids, when distilled; whilst many non-volatile acids which are not reduced by hydrobromic and hydriodic acids only act feebly (like phosphoric acid), or do not act at all (like boric acid).[73]This is in agreement with the thermochemical data, because if all the substances be taken in the gaseous state (for sulphur the heat of fusion is 0·3, and the heat of vaporisation 2·3) we have H2+ S = 4·7; H2+ Cl2= 44; H2+ Br2= 24, and H2+ I2= -3 thousand heat units; hence the formation of H2S gives less heat than that of HCl and HBr, but more than that of HI. In dilute solutions H2+ S + Aq = 9·3, and consequently less than the formation of all the halogen acids, as H2S evolves but little heat with water, and therefore in dilute solutions chlorine, bromine, and iodine decompose hydrogen sulphide.[74]Here there are three elements, hydrogen, sulphur, and iodine, each pair of which is able to form a compound, HI, H2S, and SI, besides which the latter may unite in various proportions. The complexity of chemical mechanics is seen in such examples as these. It is evident that only the study of the simplest cases can give the key to the more complex problems, and on the other hand it is evident from the examples cited in the last pages that, without penetrating into the conditions of chemical equilibria, it would be impossible to explain chemical phenomena. By following the footsteps of Berthollet the possibility of unravelling the problems will be reached; but work in this direction has only been begun during the last ten years, and much remains to be done in collecting experimental material, for which occasions present themselves at every step. In speaking of the halogens I wished to turn the reader's attention to problems of this kind.[74 bis]The same essentially takes place when sulphurous anhydride, in a dilute solution, gives hydriodic acid and sulphuric acid with iodine. On concentration a reverse reaction takes place. The equilibrated systems and the part played by water are everywhere distinctly seen.[75]Methods of formation and preparation are nothing more than particular cases of chemical reaction. If the knowledge of chemical mechanics were more exact and complete than it now is it would be possible to foretell all cases of preparationwith every detail(of the quantity of water, temperature, pressure, mass, &c.) The study of practical methods of preparation is therefore one of the paths for the study of chemical mechanics. The reaction of iodine on phosphorus and water is a case like that mentioned in Note74, and the matter is here further complicated by the possibility of the formation of the compound PH3with HI, as well as the production of PI2, PI3, and the affinity of hydriodic acid and the acids of phosphorus for water. The theoretical interest of equilibria in all their complexity is naturally very great, but it falls into the background in presence of the primary interest of discovering practical methods for the isolation of substances, and the means of employing them for the requirements of man. It is only after the satisfaction of these requirements that interests of the other order arise, which in their turn must exert an influence on the former. For these reasons, whilst considering it opportune to point out the theoretical interest of chemical equilibria, the chief attention of the reader is directed in this work to questions of practical importance.[76]Hydrobromic acid is also obtained by the action of bromine on paraffin heated to 180°. Gustavson proposed to prepare it by the action of bromine (best added in drops together with traces of aluminium bromide) on anthracene (a solid hydrocarbon from coal tar). Balard prepared it by passing bromine vapour over moist pieces of common phosphorus. The liquid tribromide of phosphorus, directly obtained from phosphorus and bromine, also gives hydrobromic acid when treated with water. Bromide of potassium or sodium, when treated with sulphuric acid in the presence of phosphorus, also gives hydrobromic acid, but hydriodic acid is decomposed by this method. In order to free hydrobromic acid from bromine vapour it is passed over moist phosphorus and dried either by phosphoric anhydride or calcium bromide (calcium chloride cannot be used, as hydrochloric acid would be formed). Neither hydrobromic nor hydriodic acids can be collected over mercury, on which they act, but they may be directly collected in a dry vessel by leading the gas-conducting tube to the bottom of the vessel, both gases being much heavier than air. Merz and Holtzmann (1889) proposed to prepare HBr directly from bromine and hydrogen. For this purpose pure dry hydrogen is passed through a flask containing boiling bromine. The mixture of gas and vapour then passes through a tube provided with one or two bulbs, which is heated moderately in the middle. Hydrobromic acid is formed with a series of flashes at the part heated. The resultant HBr, together with traces of bromine, passes into a Woulfe's bottle into which hydrogen is also introduced, and the mixture is then carried through another heated tube, after which it is passed through water which dissolves the hydrobromic acid. According to the method proposed by Newth (1892) a mixture of bromine and hydrogen is led through a tube containing a platinum spiral, which is heated to redness after the air has been displaced from the tube. If the vessel containing the bromine be kept at 60°, the hydrogen takes up almost the theoretical amount of bromine required for the formation of HBr. Although the flame which appears in the neighbourhood of the platinum spiral does not penetrate into the vessel containing the bromine, still, for safety, a tube filled with cotton wool may be interposed.Hydriodic acid is obtained in the same manner as hydrobromic. The iodine is heated in a small flask, and its vapour is carried over by hydrogen into a strongly heated tube, The gas passing from the tube is found to contain a considerable amount of HI, together with some free iodine. At a low red heat about 17 p.c. of the iodine vapour enters into combination; at a higher temperature, 78 p.c. to 79 p.c.; and at a strong heat about 82 p.c.[77]But generally more phosphorus is taken than is required for the formation of PI3, because otherwise a portion of the iodine distils over. If less than one-tenth part of iodine be taken, much phosphonium iodide, PH4I, is formed. This proportion was established by Gay-Lussac and Kolbe. Hydriodic acid is also prepared in many other ways. Bannoff dissolves two parts of iodine in one part of a previously prepared strong (sp. gr. 1·67) solution of hydriodic acid, and pours it on to red phosphorus in a retort. Personne takes a mixture of fifteen parts of water, ten of iodine, and one of red phosphorus, which, when heated, disengages hydriodic acid mixed with iodine vapour; the latter is removed by passing it over moist phosphorus (Note76). It must be remembered however that reverse reaction (Oppenheim) may take place between the hydriodic acid and phosphorus, in which the compounds PH4I and PI2are formed.It should be observed that the reaction between phosphorus, iodine and water must be carried out in the above proportions and with caution, as they may react with explosion. With red phosphorus the reaction proceeds quietly, but nevertheless requires care.L. Meyer showed that with an excess of iodine the reaction proceeds without the formation of bye-products (PH4I), according to the equation P + 5I + 4H2O = PH3O4+ 5HI. For this purpose 100 grams of iodine and 10 grams of water are placed in a retort, and a paste of 5 grams of red phosphorus and 10 grams of water is added little by little (at first with great care). The hydriodic acid may be obtained free from iodine by directing the neck of the retort upwards and causing the gas to pass through a shallow layer of water (respecting the formation of HI,seealso Note75).[78]The specific gravities of their solutions as deduced by me on the basis of Topsöe and Berthelot's determinations for 15°/4° are as follows:—102030405060 p.c.HBr1·0711·1561·2581·3741·5051·650HI1·0751·1641·2671·3991·5671·769Hydrobromic acid forms two hydrates, HBr,2H2O and HBr,H2O, which have been studied by Roozeboom with as much completeness as the hydrate of hydrochloric acid (Chapter X. Note37).With metallic silver, solutions of hydriodic acid give hydrogen with great ease, forming silver iodide. Mercury, lead, and other metals act in a similar manner.[78 bis]Iodide of nitrogen, NHI2is obtained as a brown pulverulent precipitate on adding a solution of iodine (in alcohol, for instance) to a solution of ammonia. If it be collected on a filter-paper, it does not decompose so long as the precipitate is moist; but when dry it explodes violently, so that it can only be experimented upon in small quantities. Usually the filter-paper is torn into bits while moist, and the pieces laid upon a brick; on drying an explosion proceeds not only from friction or a blow, but even spontaneously. The more dilute the solution of ammonia, the greater is the amount of iodine required for the formation of the precipitate of NHI2. A low temperature facilitates its formation. NHI2dissolves in ammonia water, and when heated the solution forms HIO3and iodine. With KI, iodide of nitrogen gives iodine, NH3and KHO. These reactions (Selivanoff) are explained by the formation of HIO from NHI2+ 2H2O = NH3+ 2HIO—and then KI + HIO = I2+ KHO. Selivanoff (seeNote29) usually observed a temporary formation of hypoiodous acid, HIO, in the reaction of ammonia upon iodine, so that here the formation of NHI2is preceded by that of HIO—i.e.first I2+ H2O = HIO + HI, and then not only the HI combines with NH3, but also 2HIO + NH3= NHI2+ 2H2O. With dilute sulphuric acid iodide of nitrogen (like NCl3) forms hypoiodous acid, but it immediately passes into iodic acid, as is expressed by the equation 5HIO = 2I2+ HIO3+ 2H2O (first 3HIO = HIO3+ 2HI, and then HI + HIO = I2+ H2O). Moreover, Selivanoff found that iodide of nitrogen, NHI2, dissolves in an excess of ammonia water, and that with potassium iodide the solution gives the reaction for hypoiodous acid (the evolution of iodine in an alkaline solution). This shows that HIO participates in the formation and decomposition of NHI2, and therefore the condition of the iodine (its metaleptic position) in them is analogous, and differs from the condition of the halogens in the haloid-anhydrides (for instance, NO2Cl). The latter are tolerably stable, while (the haloid being designated by X) NHX2, NX3, XOH, RXO (seeChapter XIII. Note43), &c., are unstable, easily decomposed with the evolution of heat, and, under the action of water, the haloid is easily replaced by hydrogen (Selivanoff), as would be expected in true products of metalepsis.[79]Hypoiodous acid, HIO, is not known, but organic compounds, RIO, of this type are known. To illustrate the peculiarities of their properties we will mention one of these compounds, namely,iodosobenzol, C6H5IO. This substance was obtained by Willgerodt (1892), and also by V. Meyer, Wachter, and Askenasy, by the action of caustic alkalis upon phenoldiiodochloride, C6H5ICl2(according to the equation, C6H5ICl2+ 2MOH = C6H5IO + 2MCl + H2O). Iodosobenzol is an amorphous yellow substance, whose melting point could not be determined because it explodes at 210°, decomposing with the evolution of iodine vapour. This substance dissolves in hot water and alcohol, but is not soluble in the majority of other neutral organic solvents. If acids do not oxidise C6H5IO, they give saline compounds in which iodosobenzol appears as a basic oxide of a diatomic metal, C6H5I. Thus, for instance, when an acetic acid solution of iodosobenzol is treated with a solution of nitric acid, it gives large monoclinic crystals of a nitric acid salt having the composition C6H5I(NO3)2(like Ca(NO3)2). In appearing as the analogue of basic oxides, iodosobenzol displaces iodine from potassium iodide (in a solution acidulated with acetic or hydrochloric acid)—i.e.it acts with its oxygen like HClO. The action of peroxide of hydrogen, chromic acid, and other similar oxidising agents gives iodoxybenzol, C6H5IO2, which is a neutral substance—i.e.incapable of giving salts with acids (compare Chapter XIII. Note43).[79 bis]The oxidation of iodine by strong nitric acid was discovered by Connell; Millon showed that it is effected, although more slowly, by the action of the hydrates of nitric acid up to HNO3,H2O, but that the solution HNO3,2H2O, and weaker solutions, do not oxidise, but simply dissolve, iodine. The participation of water in reactions is seen in this instance. It is also seen, for example, in the fact that dry ammonia combines directly with iodine—for instance, at 0° forming the compound I2,4NH3—whilst iodide of nitrogen is only formed in presence of water.[80]Bromine also displaces chlorine—for instance, from chloric acid, directly forming bromic acid. If a solution of potassium chlorate be taken (75 parts per 400 parts of water), and iodine be added to it (80 parts), and then a small quantity of nitric acid, chlorine is disengaged on boiling, and potassium iodate is formed in the solution. In this instance the nitric acid first evolves a certain portion of the chloric acid, and the latter, with the iodine, evolves chlorine. The iodic acid thus formed acts on a further quantity of the potassium chlorate, sets a portion of the chloric acid free, and in this manner the action is kept up. Potilitzin (1887) remarked, however, that not only do bromine and iodine displace the chlorine from chloric acid and potassium chlorate, but also chlorine displaces bromine from sodium bromate, and, furthermore, the reaction does not proceed as a direct substitution of the halogens, but is accompanied by the formation of free acids; for example, 5NaClO3+ 3Br2+ 3H2O = 5NaBr + 5HClO3+ HBrO3.[81]If iodine be stirred up in water, and chlorine passed through the mixture, the iodine is dissolved; the liquid becomes colourless, and contains, according to the relative amounts of water and chlorine, either IHCl2, or ICl3, or HIO3. If there be a small amount of water, then the iodic acid may separate out directly as crystals, but a complete conversion (Bornemann) only occurs when not less than ten parts of water are taken to one part of iodine—ICl + 3H2O + 2Cl2= IHO3+ 5HCl.[82]Schönbein and Ogier proved this. Ogier found that at 45° ozone immediately oxidises iodine vapour, forming first of all the oxide I2O3, which is decomposed by water or on heating into iodic anhydride and iodine. Iodic acid is formed at the positive pole when a solution of hydriodic acid is decomposed by a galvanic current (Riche). It is also formed in the combustion of hydrogen mixed with a small quantity of hydriodic acid (Salet).[83]Kämmerer showed that a solution of sp. gr. 2·127 at 14°, containing 2HIO3,9H2O, solidified completely in the cold. On comparing solutions HI +mH2O with HIO3+mH2O, we find that the specific gravity increases but the volume decreases, whilst in the passage of solutions HCl +mH2O to HClO3+mH2O both the specific gravity and the volume increase, which is also observed in certain other cases (for example, H3PO3and H3PO4).[83 bis]Ditte (1890) obtained many iodates of great variety. A neutral salt, 2(LiIO3)H2O, is obtained by saturating a solution of lithia with iodic acid. There is an analogous ammonium salt, 2(NH4IO3)H2O. He also obtained hydrates of a more complex composition, such as 6(NH4IO3)H2O and 6(NH4IO3)2H2O. Salts of the alkaline earths, Ba(IO3)2H2O and Sr(IO3)2H2O, may be obtained by a reaction of double decomposition from the normal salts of the type 2(MeIO3)H2O. When evaporated at 70° to 80° with nitric acid these salts lose water. A mixture of solutions of nitrate of zinc and an alkaline iodate precipitates Zn(IO3)22H2O. An anhydrous salt is thrown out if nitric acid be added to the solutions. Analogous salts of cadmium, silver, and copper give compounds of the type 2Me′IO34NH3and Me″(IO3)24NH3, with gaseous ammonia (Me′ and Me″ being elements of the first (Ag) and second (Cd, Zn, Cu) groups). With an aqueous solution of ammonia the above salts give substances of a different composition, such as Zn(IO3)2(NH4)2O, Cd(IO3)2(NH4)2O. Copper gives Cu(IO3)24(NH4)2O and Cu(IO3)2(NH4)2O. These salts may be regarded as compounds of I2O5, and MeO and (NH4)2O; for example, Zn(IO3)2(NH4)2O may be regarded as ZnO(NH4)2OI2O5, or, as derived from the hydrate, I2O52H2O = 2(HIO3)H2O.[84]If sodium iodate be mixed with a solution of sodium hydroxide, heated, and chlorine passed through the solution, a sparingly soluble salt separates out, which corresponds with periodic acid, and has the composition Na4I2O9,3H2O.6NaHO + 2NaIO3+ 4Cl = 4NaCl + Na4I2O9+ 3H2O.This compound is sparingly soluble in water, but dissolves easily in a very dilute solution of nitric acid. If silver nitrate be added to this solution a precipitate is formed which contains the corresponding compound of silver, Ag4I2O9,3H2O. If this sparingly soluble silver compound be dissolved in hot nitric acid, orange crystals of a salt having the composition AgIO4separate on evaporation. This salt is formed from the preceding by the nitric acid taking up silver oxide—Ag4I2O_9 + 2HNO3= 2AgNO3+ 2AgIO4+ H2O. The silver salt is decomposed by water, with the re-formation of the preceding salt, whilst iodic acid remains in solution—4AgIO4+ H2O = Ag4I2O9+ 2HIO4.The structure of the first of these salts, Na4I2O9,3H2O, presents itself in a simpler form if the water of crystallisation is regarded as an integral portion of the salt; the formula is then divided in two, and takes the form of IO(OH)3(ONa)2—that is, it answers to the type IOX5, or IX7, like AgIO4which is IO3(OAg). The composition of all the salts of periodic acids are expressed by this type IX7. Kimmins (1889) refers all the salts of periodic acid to four types—the meta-salts of HIO4(salts of Ag, Cu, Pb), the meso-salts of H3IO5(PbH, Ag2H, CdH), the para-salts of H5IO6(Na2H3, Na3H2), and the di-salts of H4I2O9(K4, Ag4, Ni2). The three first are direct compounds of the type IX7, namely, IO3(OH), IO2(OH)3, and IO(OH)5, and the last are types of diperiodic salts, which correspond with the type of the meso-salts, as pyrophosphoric salts correspond with orthophosphoric salts—i.e.2H3IO5-H2O = H4I2O9.[85]Periodic acid, discovered by Magnus and Ammermüller, and whose salts were afterwards studied by Langlois, Rammelsberg, and many others, presents an example of hydrates in which it is evident that there is not that distinction between the water of hydration and of crystallisation which was at first considered to be so clear. In HClO,2H2O the water, 2H2O, is not displaced by bases, and must be regarded as water of crystallisation, whilst in HIO4,2H2O it must be regarded as water of hydration. We shall afterwards see that the system of the elements obliges us to consider the halogens as substances giving a highest saline type,GX7, whereGsignifies a halogen, andXoxygen (O =X2), OH, and other like elements. The hydrate IO(OH)5corresponding with many of the salts of periodic acid (for example, the salts of barium, strontium, mercury) does not exhaust all the possible forms. It is evident that various other pyro-, meta-, &c., forms are possible by the loss of water, as will be more fully explained in speaking of phosphoric acid, and as was pointed out in the preceding note.[86]With respect to hydrogen, oxygen, chlorine, and other elements, bromine occupies an intermediate position between chlorine and iodine, and therefore there is no particular need for considering at length the compounds of bromine. This is the great advantage of a natural grouping of the elements.[87]They were both obtained by Gay-Lussac and many others. Recent data respecting iodine monochloride, ICl, entirely confirm the numerous observations of Trapp (1854), and even confirm his statement as to the existence of two isomeric (liquid and crystalline) forms (Stortenbeker). With a small excess of iodine, iodine monochloride remains liquid, but in the presence of traces of iodine trichloride it easily crystallises. Tanatar (1893) showed that of the two modifications of ICl, one is stable, and melts at 27°; while the other, which easily passes into the first, and is formed in the absence of ICl3, melts at 14°. Schützenberger amplified the data concerning the action of water on the chlorides (Note88), and Christomanos gave the fullest data regarding the trichloride.After being kept for some time, the liquid monochloride of iodine yields red deliquescent octahedra, having the composition ICl4, which are therefore formed from the monochloride with the liberation of free iodine, which dissolves in the remaining quantity of the monochloride. This substance, however, judging by certain observations, is impure iodine trichloride. If 1 part of iodine be stirred up in 20 parts of water, and chlorine be passed through the liquid, then all the iodine is dissolved, and a colourless liquid is ultimately obtained which contains a certain proportion of chlorine, because this compound gives a metallic chloride and iodate with alkalis without evolving any free iodine: ICl5+ 6KHO = 5KCl + KIO3+ 3H2O. The existence of a pentachloride ICl5is, however, denied, because this substance has not been obtained in a free state.Stortenbeker (1888) investigated the equilibrium of the system containing the molecules I2, ICl, ICl3, and Cl2, in the same way that Roozeboom (Chapter X. Note38) examined the equilibrium of the molecules HCl, HCl,2H2O, and H2O. He found that iodine monochloride appears in two states, one (the ordinary) is stable and melts at 27°·2, whilst the other is obtained by rapid cooling, and melts at 13°·9, and easily passes into the first form. Iodine trichloride melts at 101° only in a closed tube under a pressure of 16 atmospheres.[88]By the action of water on iodine monochloride and trichloride a compound IHCl2is obtained, which does not seem to be altered by water. Besides this compound, iodine and iodic acid are always formed, 10ICl + 3H2O = HIO3+ 5IHCl2+ 2I2; and in this respect iodine trichloride may be regarded as a mixture, ICl + ICl5= 2ICl3, but ICl5+ 3H2O = IHO3+ 5HCl; hence iodic acid, iodine, the compound IHCl2, and hydrochloric acid are also formed by the action of water.
[61]In general, 2HI + O = I2+ H2O, if the oxygen proceed from a substance from which it is easily evolved. For this reason compounds corresponding with the higher stages of oxidation or chlorination frequently give a lower stage when treated with hydriodic acid. Ferric oxide, Fe2O3, is a higher oxide, and ferrous oxide, FeO, a lower oxide; the former corresponds with FeX3, and the latter with FeX2, and this passage from the higher to the lower takes place under the action of hydriodic acid. Thus hydrogen peroxide and ozone (ChapterIV.) are able to liberate iodine from hydriodic acid. Compounds of copper oxide, CuO or CuX2, give compounds of the suboxide Cu2O, or CuX. Even sulphuric acid, which corresponds to the higher stage SO3, is able to act thus, forming the lower oxide SO2. The liberation of iodine from hydriodic acid proceeds with still greater ease under the action of substances capable of disengaging oxygen. In practice, many methods are employed for liberating iodine from acid liquids containing, for example, sulphuric acid and hydriodic acid. The higher oxides of nitrogen are most commonly used; they then pass into nitric oxide. Iodine may even be disengaged from hydriodic acid by the action of iodic acid, &c. But there is a limit in these reactions of the oxidation of hydriodic acid because, under certain conditions, especially in dilute solutions, the iodine set free is itself able to act as an oxidising agent—that is, it exhibits the character of chlorine, and of the halogens in general, to which we shall again have occasion to refer. In Chili, where a large quantity of iodine is extracted in the manufacture of Chili nitre, which contains NaIO3, it is mixed with the acid and normal sulphites of sodium in solution; the iodine is then precipitated according to the equation 2NaIO3+ 3Na2SO3+ 2NaHSO3= 5Na2SO4+I2+ H2O. The iodine thus obtained is purified by sublimation.
[61]In general, 2HI + O = I2+ H2O, if the oxygen proceed from a substance from which it is easily evolved. For this reason compounds corresponding with the higher stages of oxidation or chlorination frequently give a lower stage when treated with hydriodic acid. Ferric oxide, Fe2O3, is a higher oxide, and ferrous oxide, FeO, a lower oxide; the former corresponds with FeX3, and the latter with FeX2, and this passage from the higher to the lower takes place under the action of hydriodic acid. Thus hydrogen peroxide and ozone (ChapterIV.) are able to liberate iodine from hydriodic acid. Compounds of copper oxide, CuO or CuX2, give compounds of the suboxide Cu2O, or CuX. Even sulphuric acid, which corresponds to the higher stage SO3, is able to act thus, forming the lower oxide SO2. The liberation of iodine from hydriodic acid proceeds with still greater ease under the action of substances capable of disengaging oxygen. In practice, many methods are employed for liberating iodine from acid liquids containing, for example, sulphuric acid and hydriodic acid. The higher oxides of nitrogen are most commonly used; they then pass into nitric oxide. Iodine may even be disengaged from hydriodic acid by the action of iodic acid, &c. But there is a limit in these reactions of the oxidation of hydriodic acid because, under certain conditions, especially in dilute solutions, the iodine set free is itself able to act as an oxidising agent—that is, it exhibits the character of chlorine, and of the halogens in general, to which we shall again have occasion to refer. In Chili, where a large quantity of iodine is extracted in the manufacture of Chili nitre, which contains NaIO3, it is mixed with the acid and normal sulphites of sodium in solution; the iodine is then precipitated according to the equation 2NaIO3+ 3Na2SO3+ 2NaHSO3= 5Na2SO4+I2+ H2O. The iodine thus obtained is purified by sublimation.
[62]For the final purification of iodine, Stas dissolved it in a strong solution of potassium iodide, and precipitated it by the addition of water (seeNote58).
[62]For the final purification of iodine, Stas dissolved it in a strong solution of potassium iodide, and precipitated it by the addition of water (seeNote58).
[63]The solubility of iodine in solutions containing iodides, and compounds of iodine in general, may serve, on the one hand, as an indication that solution is due to a similarity between the solvent and dissolved substance, and, on the other hand, as an indirect proof of that view as to solutions which was cited in ChapterI., because in many instances unstable highly iodised compounds, resembling crystallo-hydrates, have been obtained from such solutions. Thus iodide of tetramethylammonium, N(CH3)4I, combines with I2, and I4. Even a solution of iodine in a saturated solution of potassium iodide presents indications of the formation of a definite compound KI3. Thus, an alcoholic solution of KI3does not give up iodine to carbon bisulphide, although this solvent takes up iodine from an alcoholic solution of iodine itself (Girault, Jörgensen, and others). The instability of these compounds resembles the instability of many crystallo-hydrates, for instance of HCl,2H2O.
[63]The solubility of iodine in solutions containing iodides, and compounds of iodine in general, may serve, on the one hand, as an indication that solution is due to a similarity between the solvent and dissolved substance, and, on the other hand, as an indirect proof of that view as to solutions which was cited in ChapterI., because in many instances unstable highly iodised compounds, resembling crystallo-hydrates, have been obtained from such solutions. Thus iodide of tetramethylammonium, N(CH3)4I, combines with I2, and I4. Even a solution of iodine in a saturated solution of potassium iodide presents indications of the formation of a definite compound KI3. Thus, an alcoholic solution of KI3does not give up iodine to carbon bisulphide, although this solvent takes up iodine from an alcoholic solution of iodine itself (Girault, Jörgensen, and others). The instability of these compounds resembles the instability of many crystallo-hydrates, for instance of HCl,2H2O.
[64]The equality of the atomic volumes of the halogens themselves is all the more remarkable because in all the halogen compounds the volume augments with the substitution of fluorine by chlorine, bromine, and iodine. Thus, for example, the volume of sodium fluoride (obtained by dividing the weight expressed by its formula by its specific gravity) is about 15, of sodium chloride 27, of sodium bromide 32, and of sodium iodide 41. The volume of silicon chloroform, SiHCl3, is 82, and those of the corresponding bromine and iodine compounds are 108 and 122 respectively. The same difference also exists in solutions; for example, NaCl + 200H2O has a sp. gr. (at 15°/4°) of 1·0106, consequently the volume of the solution 3,658·5/1·0106 = 3,620, hence the volume of sodium chloride in solution = 3,620–3,603 (this is the volume of 200 H2O) = 17, and in similar solutions, NaBr = 26 and NaI = 35.
[64]The equality of the atomic volumes of the halogens themselves is all the more remarkable because in all the halogen compounds the volume augments with the substitution of fluorine by chlorine, bromine, and iodine. Thus, for example, the volume of sodium fluoride (obtained by dividing the weight expressed by its formula by its specific gravity) is about 15, of sodium chloride 27, of sodium bromide 32, and of sodium iodide 41. The volume of silicon chloroform, SiHCl3, is 82, and those of the corresponding bromine and iodine compounds are 108 and 122 respectively. The same difference also exists in solutions; for example, NaCl + 200H2O has a sp. gr. (at 15°/4°) of 1·0106, consequently the volume of the solution 3,658·5/1·0106 = 3,620, hence the volume of sodium chloride in solution = 3,620–3,603 (this is the volume of 200 H2O) = 17, and in similar solutions, NaBr = 26 and NaI = 35.
[65]But the density (and also molecular volume, Note64) of a bromine compound is always greater than that of a chlorine compound, whilst that of an iodine compound is still greater. The order is the same in many other respects. For example, an iodine compound has a higher boiling point than a bromine compound, &c.
[65]But the density (and also molecular volume, Note64) of a bromine compound is always greater than that of a chlorine compound, whilst that of an iodine compound is still greater. The order is the same in many other respects. For example, an iodine compound has a higher boiling point than a bromine compound, &c.
[66]A. L. Potilitzin showed that in heating various metallic chlorides in a closed tube, with an equivalent quantity of bromine, a distribution of the metal between the halogens always occurs, and that the amounts of chlorine replaced by the bromine in the ultimate product are proportional to the atomic weights of the metals taken and inversely proportional to their equivalence. Thus, if NaCl + Br be taken, then out of 100 parts of chlorine, 5·54 are replaced by the bromine, whilst with AgCl + Br 27·28 parts are replaced. These figures are in the ratio 1 : 4·9, and the atomic weights Na : Ag = 1 : 4·7. In general terms, if a chloride MClnbe taken, it gives withnBr a percentage substitution = 4M/n2where M is the atomic weight of the metal. This law was deduced from observations on the chlorides of Li, K, Na, Ag (n= 1), Ca, Sr, Ba, Co, Ni, Hg, Pb (n= 2), Bi (n= 3), Sn (n= 4), and Fe2(n= 6).In these determinations of Potilitzin we see not only a brilliant confirmation of Berthollet's doctrine, but also the first effort to directly determine the affinities of elements by means of displacement. The chief object of these researches consisted in proving whether a displacement occurs in those cases where heat is absorbed, and in this instance it should be absorbed, because the formation of all metallic bromides is attended with the evolution of less heat than that of the chlorides, as is seen by the figures given in Note55.If the mass of the bromine be increased, then the amount of chlorine displaced also increases. For example, if masses of bromine of 1 and 4 equivalents act on a molecule of sodium chloride, then the percentages of the chlorine displaced will be 6·08 p.c. and 12·46 p.c.; in the action of 1, 4, 25, and 100 molecules of bromine on a molecule of barium chloride, there will be displaced 7·8, 17·6, 35·0, and 45·0 p.c. of chlorine. If an equivalent quantity of hydrochloric acid act on metallic bromides in closed tubes, and in the absence of water at a temperature of 300°, then the percentages of the substitution of the bromine by the chlorine in the double decomposition taking place between univalent metals are inversely proportional to their atomic weights. For example, NaBr + HCl gives at the limit 21 p.c. of displacement, KCl 12 p.c. and AgCl 4¼ p.c. Essentially the same action takes place in an aqueous solution, although the phenomenon is complicated by the participation of the water. The reactions proceed spontaneously in one or the other direction at the ordinary temperature but at differentrates. In the action of a dilute solution (1 equivalent per 5 litres) of sodium chloride on silver bromide at the ordinary temperature the amount of bromine replaced in six and a half days is 2·07 p.c., and with potassium chloride 1·5 p.c. With an excess of the chloride the magnitude of the substitution increases. These conversions also proceed with the absorption of heat. The reverse reactions evolving heat proceed incomparably more rapidly, but also to a certain limit; for example, in the reaction AgCl + RBr the following percentages of silver bromide are formed in different times:hours232296120K79·8287·488·22—94·21Na83·6390·7491·7095·49—That is, the conversions which are accompanied by an evolution of heat proceed with very much greater rapidity than the reverse conversions.
[66]A. L. Potilitzin showed that in heating various metallic chlorides in a closed tube, with an equivalent quantity of bromine, a distribution of the metal between the halogens always occurs, and that the amounts of chlorine replaced by the bromine in the ultimate product are proportional to the atomic weights of the metals taken and inversely proportional to their equivalence. Thus, if NaCl + Br be taken, then out of 100 parts of chlorine, 5·54 are replaced by the bromine, whilst with AgCl + Br 27·28 parts are replaced. These figures are in the ratio 1 : 4·9, and the atomic weights Na : Ag = 1 : 4·7. In general terms, if a chloride MClnbe taken, it gives withnBr a percentage substitution = 4M/n2where M is the atomic weight of the metal. This law was deduced from observations on the chlorides of Li, K, Na, Ag (n= 1), Ca, Sr, Ba, Co, Ni, Hg, Pb (n= 2), Bi (n= 3), Sn (n= 4), and Fe2(n= 6).
In these determinations of Potilitzin we see not only a brilliant confirmation of Berthollet's doctrine, but also the first effort to directly determine the affinities of elements by means of displacement. The chief object of these researches consisted in proving whether a displacement occurs in those cases where heat is absorbed, and in this instance it should be absorbed, because the formation of all metallic bromides is attended with the evolution of less heat than that of the chlorides, as is seen by the figures given in Note55.
If the mass of the bromine be increased, then the amount of chlorine displaced also increases. For example, if masses of bromine of 1 and 4 equivalents act on a molecule of sodium chloride, then the percentages of the chlorine displaced will be 6·08 p.c. and 12·46 p.c.; in the action of 1, 4, 25, and 100 molecules of bromine on a molecule of barium chloride, there will be displaced 7·8, 17·6, 35·0, and 45·0 p.c. of chlorine. If an equivalent quantity of hydrochloric acid act on metallic bromides in closed tubes, and in the absence of water at a temperature of 300°, then the percentages of the substitution of the bromine by the chlorine in the double decomposition taking place between univalent metals are inversely proportional to their atomic weights. For example, NaBr + HCl gives at the limit 21 p.c. of displacement, KCl 12 p.c. and AgCl 4¼ p.c. Essentially the same action takes place in an aqueous solution, although the phenomenon is complicated by the participation of the water. The reactions proceed spontaneously in one or the other direction at the ordinary temperature but at differentrates. In the action of a dilute solution (1 equivalent per 5 litres) of sodium chloride on silver bromide at the ordinary temperature the amount of bromine replaced in six and a half days is 2·07 p.c., and with potassium chloride 1·5 p.c. With an excess of the chloride the magnitude of the substitution increases. These conversions also proceed with the absorption of heat. The reverse reactions evolving heat proceed incomparably more rapidly, but also to a certain limit; for example, in the reaction AgCl + RBr the following percentages of silver bromide are formed in different times:
That is, the conversions which are accompanied by an evolution of heat proceed with very much greater rapidity than the reverse conversions.
[67]The dissociation of hydriodic acidhas been studied in detail by Hautefeuille and Lemoine, from whose researches we extract the following information. The decomposition of hydriodic acid is decided, but proceeds slowly at 180°; the rate and limit of decomposition increase with a rise of temperature. The reverse action—that is, I2+ H2= 2HI—proceeds not only under the influence of spongy platinum (Corenwinder), which also accelerates the decomposition of hydriodic acid, but also by itself, although slowly. The limit of the reverse reaction remains the same with or without spongy platinum. An increase of pressure has a very powerful accelerative effect on the rate of formation of hydriodic acid, and therefore spongy platinum by condensing gases has the same effect as increase of pressure. At the atmospheric pressure the decomposition of hydriodic acid reaches the limit at 250° in several months, and at 440° in several hours. The limit at 250° is about 18 p.c. of decomposition—that is, out of 100 parts of hydrogen previously combined in hydriodic acid, about 18 p.c. may be disengaged at this temperature (this hydrogen may be easily measured, and the measure of dissociation determined), but not more; the limit at 440° is about 26 p.c. If the pressure under which 2HI passes into H2+ I2be 4½ atmospheres, then the limit is 24 p.c.; under a pressure of ⅕ atmosphere the limit is 29 p.c. The small influence of pressure on the dissociation of hydriodic acid (compared with N2O4, Chapter VI. Note46) is due to the fact that the reaction 2HI = I2+ H2is not accompanied by a change of volume. In order to show the influence of time, we will cite the following figures referring to 350°: (1) Reaction H2+ I2; after 3 hours, 88 p.c. of hydrogen remained free; 8 hours, 69 p.c.; 34 hours, 48 p.c.; 76 hours, 29 p.c.; and 327 hours, 18·5 p.c. (2) The reverse decomposition of 2HI; after 9 hours, 3 p.c. of hydrogen was set free, and after 250 hours 18·6 p.c.—that is, the limit was reached. The addition of extraneous hydrogen diminishes the limit of the reaction of decomposition, or increases the formation of hydriodic acid from iodine and hydrogen, as would be expected from Berthollet's doctrine (ChapterX.). Thus at 440° 26 p.c. of hydriodic acid is decomposed if there be no admixture of hydrogen, while if H2be added, then at the limit only half as large a mass of HI is decomposed. Therefore, if an infinite mass of hydrogen be added there will be no decomposition of the hydriodic acid. Light aids the decomposition of hydriodic acid very powerfully. At the ordinary temperature 80 p.c. is decomposed under the influence of light, whilst under the influence of heat alone this limit corresponds with a very high temperature. The distinct action of light, spongy platinum, and of impurities in glass (especially of sodium sulphate, which decomposes hydriodic acid), not only render the investigations difficult, but also show that in reactions like 2HI = I2+ H2, which are accompanied by slight heat effects, all foreign and feeble influences may strongly affect the progress of the action (Note47).
[67]The dissociation of hydriodic acidhas been studied in detail by Hautefeuille and Lemoine, from whose researches we extract the following information. The decomposition of hydriodic acid is decided, but proceeds slowly at 180°; the rate and limit of decomposition increase with a rise of temperature. The reverse action—that is, I2+ H2= 2HI—proceeds not only under the influence of spongy platinum (Corenwinder), which also accelerates the decomposition of hydriodic acid, but also by itself, although slowly. The limit of the reverse reaction remains the same with or without spongy platinum. An increase of pressure has a very powerful accelerative effect on the rate of formation of hydriodic acid, and therefore spongy platinum by condensing gases has the same effect as increase of pressure. At the atmospheric pressure the decomposition of hydriodic acid reaches the limit at 250° in several months, and at 440° in several hours. The limit at 250° is about 18 p.c. of decomposition—that is, out of 100 parts of hydrogen previously combined in hydriodic acid, about 18 p.c. may be disengaged at this temperature (this hydrogen may be easily measured, and the measure of dissociation determined), but not more; the limit at 440° is about 26 p.c. If the pressure under which 2HI passes into H2+ I2be 4½ atmospheres, then the limit is 24 p.c.; under a pressure of ⅕ atmosphere the limit is 29 p.c. The small influence of pressure on the dissociation of hydriodic acid (compared with N2O4, Chapter VI. Note46) is due to the fact that the reaction 2HI = I2+ H2is not accompanied by a change of volume. In order to show the influence of time, we will cite the following figures referring to 350°: (1) Reaction H2+ I2; after 3 hours, 88 p.c. of hydrogen remained free; 8 hours, 69 p.c.; 34 hours, 48 p.c.; 76 hours, 29 p.c.; and 327 hours, 18·5 p.c. (2) The reverse decomposition of 2HI; after 9 hours, 3 p.c. of hydrogen was set free, and after 250 hours 18·6 p.c.—that is, the limit was reached. The addition of extraneous hydrogen diminishes the limit of the reaction of decomposition, or increases the formation of hydriodic acid from iodine and hydrogen, as would be expected from Berthollet's doctrine (ChapterX.). Thus at 440° 26 p.c. of hydriodic acid is decomposed if there be no admixture of hydrogen, while if H2be added, then at the limit only half as large a mass of HI is decomposed. Therefore, if an infinite mass of hydrogen be added there will be no decomposition of the hydriodic acid. Light aids the decomposition of hydriodic acid very powerfully. At the ordinary temperature 80 p.c. is decomposed under the influence of light, whilst under the influence of heat alone this limit corresponds with a very high temperature. The distinct action of light, spongy platinum, and of impurities in glass (especially of sodium sulphate, which decomposes hydriodic acid), not only render the investigations difficult, but also show that in reactions like 2HI = I2+ H2, which are accompanied by slight heat effects, all foreign and feeble influences may strongly affect the progress of the action (Note47).
[68]The thermal determinations of Thomsen (at 18°) gave in thousands of calories, Cl + H = +22, HCl + Aq (that is, on dissolving HCl in a large amount of water) = +17·3, and therefore H + Cl + Aq = +39·3. In taking molecules, all these figures must be doubled. Br + H = +8·4; HBr + Aq = 19·9; H + Br + Aq = +28·3. According to Berthelot 7·2 are required for the vaporisation of Br2, hence Br2+ H2= 16·8 + 7·2 = +24, if Br2be taken as vapour for comparison with Cl2. H + I = -6·0, HI + Aq = 19·2; H + I + Aq= +13·2, and, according to Berthelot, the heat of fusion of I2= 3·0, and of vaporisation 6·0 thousand heat units, and therefore I2+ H2= -2(6·0) + 3 + 6 = -3·0, if the iodine be taken as vapour. Berthelot, on the basis of his determinations, gives, however, +0·8 thousand heat units. Similar contradictory results are often met with in thermochemistry owing to the imperfection of the existing methods, and particularly the necessity of depending on indirect methods for obtaining the fundamental figures. Thus Thomsen decomposed a dilute solution of potassium iodide by gaseous chlorine; the reaction gave +26·2, whence, having first determined the heat effects of the reactions KHO + HCl, KHO + HI and Cl + H in aqueous solutions, it was possible to find H + I + Aq; then, knowing HI + Aq, to find I + H. It is evident that unavoidable errors may accumulate.
[68]The thermal determinations of Thomsen (at 18°) gave in thousands of calories, Cl + H = +22, HCl + Aq (that is, on dissolving HCl in a large amount of water) = +17·3, and therefore H + Cl + Aq = +39·3. In taking molecules, all these figures must be doubled. Br + H = +8·4; HBr + Aq = 19·9; H + Br + Aq = +28·3. According to Berthelot 7·2 are required for the vaporisation of Br2, hence Br2+ H2= 16·8 + 7·2 = +24, if Br2be taken as vapour for comparison with Cl2. H + I = -6·0, HI + Aq = 19·2; H + I + Aq= +13·2, and, according to Berthelot, the heat of fusion of I2= 3·0, and of vaporisation 6·0 thousand heat units, and therefore I2+ H2= -2(6·0) + 3 + 6 = -3·0, if the iodine be taken as vapour. Berthelot, on the basis of his determinations, gives, however, +0·8 thousand heat units. Similar contradictory results are often met with in thermochemistry owing to the imperfection of the existing methods, and particularly the necessity of depending on indirect methods for obtaining the fundamental figures. Thus Thomsen decomposed a dilute solution of potassium iodide by gaseous chlorine; the reaction gave +26·2, whence, having first determined the heat effects of the reactions KHO + HCl, KHO + HI and Cl + H in aqueous solutions, it was possible to find H + I + Aq; then, knowing HI + Aq, to find I + H. It is evident that unavoidable errors may accumulate.
[69]One can believe, however, on the basis of Berthollet's doctrine, and the observations of Potilitzin (Note66), that a certain slow decomposition of water by iodine takes place. On this view the observations of Dossios and Weith on the fact that the solubility of iodine in water increases after the lapse of several months will be comprehensible. Hydriodic acid is then formed, and it increases the solubility. If the iodine be extracted from such a solution by carbon bisulphide, then, as the authors showed, after the action of nitrous anhydride iodine may be again detected in the solution by means of starch. It can easily be understood that a number of similar reactions, requiring much time and taking place in small quantities, have up to now eluded the attention of investigators, who even still doubt the universal application of Berthollet's doctrine, or only see the thermochemical side of reactions, or else neglect to pay attention to the element of time and the influence of mass.
[69]One can believe, however, on the basis of Berthollet's doctrine, and the observations of Potilitzin (Note66), that a certain slow decomposition of water by iodine takes place. On this view the observations of Dossios and Weith on the fact that the solubility of iodine in water increases after the lapse of several months will be comprehensible. Hydriodic acid is then formed, and it increases the solubility. If the iodine be extracted from such a solution by carbon bisulphide, then, as the authors showed, after the action of nitrous anhydride iodine may be again detected in the solution by means of starch. It can easily be understood that a number of similar reactions, requiring much time and taking place in small quantities, have up to now eluded the attention of investigators, who even still doubt the universal application of Berthollet's doctrine, or only see the thermochemical side of reactions, or else neglect to pay attention to the element of time and the influence of mass.
[70]On the basis of the data in Note68.
[70]On the basis of the data in Note68.
[71]A number of similar cases confirm what has been said in ChapterX.
[71]A number of similar cases confirm what has been said in ChapterX.
[72]This is prevented by the reducibility of sulphuric acid. If volatile acids be taken they pass over, together with the hydrobromic and hydriodic acids, when distilled; whilst many non-volatile acids which are not reduced by hydrobromic and hydriodic acids only act feebly (like phosphoric acid), or do not act at all (like boric acid).
[72]This is prevented by the reducibility of sulphuric acid. If volatile acids be taken they pass over, together with the hydrobromic and hydriodic acids, when distilled; whilst many non-volatile acids which are not reduced by hydrobromic and hydriodic acids only act feebly (like phosphoric acid), or do not act at all (like boric acid).
[73]This is in agreement with the thermochemical data, because if all the substances be taken in the gaseous state (for sulphur the heat of fusion is 0·3, and the heat of vaporisation 2·3) we have H2+ S = 4·7; H2+ Cl2= 44; H2+ Br2= 24, and H2+ I2= -3 thousand heat units; hence the formation of H2S gives less heat than that of HCl and HBr, but more than that of HI. In dilute solutions H2+ S + Aq = 9·3, and consequently less than the formation of all the halogen acids, as H2S evolves but little heat with water, and therefore in dilute solutions chlorine, bromine, and iodine decompose hydrogen sulphide.
[73]This is in agreement with the thermochemical data, because if all the substances be taken in the gaseous state (for sulphur the heat of fusion is 0·3, and the heat of vaporisation 2·3) we have H2+ S = 4·7; H2+ Cl2= 44; H2+ Br2= 24, and H2+ I2= -3 thousand heat units; hence the formation of H2S gives less heat than that of HCl and HBr, but more than that of HI. In dilute solutions H2+ S + Aq = 9·3, and consequently less than the formation of all the halogen acids, as H2S evolves but little heat with water, and therefore in dilute solutions chlorine, bromine, and iodine decompose hydrogen sulphide.
[74]Here there are three elements, hydrogen, sulphur, and iodine, each pair of which is able to form a compound, HI, H2S, and SI, besides which the latter may unite in various proportions. The complexity of chemical mechanics is seen in such examples as these. It is evident that only the study of the simplest cases can give the key to the more complex problems, and on the other hand it is evident from the examples cited in the last pages that, without penetrating into the conditions of chemical equilibria, it would be impossible to explain chemical phenomena. By following the footsteps of Berthollet the possibility of unravelling the problems will be reached; but work in this direction has only been begun during the last ten years, and much remains to be done in collecting experimental material, for which occasions present themselves at every step. In speaking of the halogens I wished to turn the reader's attention to problems of this kind.
[74]Here there are three elements, hydrogen, sulphur, and iodine, each pair of which is able to form a compound, HI, H2S, and SI, besides which the latter may unite in various proportions. The complexity of chemical mechanics is seen in such examples as these. It is evident that only the study of the simplest cases can give the key to the more complex problems, and on the other hand it is evident from the examples cited in the last pages that, without penetrating into the conditions of chemical equilibria, it would be impossible to explain chemical phenomena. By following the footsteps of Berthollet the possibility of unravelling the problems will be reached; but work in this direction has only been begun during the last ten years, and much remains to be done in collecting experimental material, for which occasions present themselves at every step. In speaking of the halogens I wished to turn the reader's attention to problems of this kind.
[74 bis]The same essentially takes place when sulphurous anhydride, in a dilute solution, gives hydriodic acid and sulphuric acid with iodine. On concentration a reverse reaction takes place. The equilibrated systems and the part played by water are everywhere distinctly seen.
[74 bis]The same essentially takes place when sulphurous anhydride, in a dilute solution, gives hydriodic acid and sulphuric acid with iodine. On concentration a reverse reaction takes place. The equilibrated systems and the part played by water are everywhere distinctly seen.
[75]Methods of formation and preparation are nothing more than particular cases of chemical reaction. If the knowledge of chemical mechanics were more exact and complete than it now is it would be possible to foretell all cases of preparationwith every detail(of the quantity of water, temperature, pressure, mass, &c.) The study of practical methods of preparation is therefore one of the paths for the study of chemical mechanics. The reaction of iodine on phosphorus and water is a case like that mentioned in Note74, and the matter is here further complicated by the possibility of the formation of the compound PH3with HI, as well as the production of PI2, PI3, and the affinity of hydriodic acid and the acids of phosphorus for water. The theoretical interest of equilibria in all their complexity is naturally very great, but it falls into the background in presence of the primary interest of discovering practical methods for the isolation of substances, and the means of employing them for the requirements of man. It is only after the satisfaction of these requirements that interests of the other order arise, which in their turn must exert an influence on the former. For these reasons, whilst considering it opportune to point out the theoretical interest of chemical equilibria, the chief attention of the reader is directed in this work to questions of practical importance.
[75]Methods of formation and preparation are nothing more than particular cases of chemical reaction. If the knowledge of chemical mechanics were more exact and complete than it now is it would be possible to foretell all cases of preparationwith every detail(of the quantity of water, temperature, pressure, mass, &c.) The study of practical methods of preparation is therefore one of the paths for the study of chemical mechanics. The reaction of iodine on phosphorus and water is a case like that mentioned in Note74, and the matter is here further complicated by the possibility of the formation of the compound PH3with HI, as well as the production of PI2, PI3, and the affinity of hydriodic acid and the acids of phosphorus for water. The theoretical interest of equilibria in all their complexity is naturally very great, but it falls into the background in presence of the primary interest of discovering practical methods for the isolation of substances, and the means of employing them for the requirements of man. It is only after the satisfaction of these requirements that interests of the other order arise, which in their turn must exert an influence on the former. For these reasons, whilst considering it opportune to point out the theoretical interest of chemical equilibria, the chief attention of the reader is directed in this work to questions of practical importance.
[76]Hydrobromic acid is also obtained by the action of bromine on paraffin heated to 180°. Gustavson proposed to prepare it by the action of bromine (best added in drops together with traces of aluminium bromide) on anthracene (a solid hydrocarbon from coal tar). Balard prepared it by passing bromine vapour over moist pieces of common phosphorus. The liquid tribromide of phosphorus, directly obtained from phosphorus and bromine, also gives hydrobromic acid when treated with water. Bromide of potassium or sodium, when treated with sulphuric acid in the presence of phosphorus, also gives hydrobromic acid, but hydriodic acid is decomposed by this method. In order to free hydrobromic acid from bromine vapour it is passed over moist phosphorus and dried either by phosphoric anhydride or calcium bromide (calcium chloride cannot be used, as hydrochloric acid would be formed). Neither hydrobromic nor hydriodic acids can be collected over mercury, on which they act, but they may be directly collected in a dry vessel by leading the gas-conducting tube to the bottom of the vessel, both gases being much heavier than air. Merz and Holtzmann (1889) proposed to prepare HBr directly from bromine and hydrogen. For this purpose pure dry hydrogen is passed through a flask containing boiling bromine. The mixture of gas and vapour then passes through a tube provided with one or two bulbs, which is heated moderately in the middle. Hydrobromic acid is formed with a series of flashes at the part heated. The resultant HBr, together with traces of bromine, passes into a Woulfe's bottle into which hydrogen is also introduced, and the mixture is then carried through another heated tube, after which it is passed through water which dissolves the hydrobromic acid. According to the method proposed by Newth (1892) a mixture of bromine and hydrogen is led through a tube containing a platinum spiral, which is heated to redness after the air has been displaced from the tube. If the vessel containing the bromine be kept at 60°, the hydrogen takes up almost the theoretical amount of bromine required for the formation of HBr. Although the flame which appears in the neighbourhood of the platinum spiral does not penetrate into the vessel containing the bromine, still, for safety, a tube filled with cotton wool may be interposed.Hydriodic acid is obtained in the same manner as hydrobromic. The iodine is heated in a small flask, and its vapour is carried over by hydrogen into a strongly heated tube, The gas passing from the tube is found to contain a considerable amount of HI, together with some free iodine. At a low red heat about 17 p.c. of the iodine vapour enters into combination; at a higher temperature, 78 p.c. to 79 p.c.; and at a strong heat about 82 p.c.
[76]Hydrobromic acid is also obtained by the action of bromine on paraffin heated to 180°. Gustavson proposed to prepare it by the action of bromine (best added in drops together with traces of aluminium bromide) on anthracene (a solid hydrocarbon from coal tar). Balard prepared it by passing bromine vapour over moist pieces of common phosphorus. The liquid tribromide of phosphorus, directly obtained from phosphorus and bromine, also gives hydrobromic acid when treated with water. Bromide of potassium or sodium, when treated with sulphuric acid in the presence of phosphorus, also gives hydrobromic acid, but hydriodic acid is decomposed by this method. In order to free hydrobromic acid from bromine vapour it is passed over moist phosphorus and dried either by phosphoric anhydride or calcium bromide (calcium chloride cannot be used, as hydrochloric acid would be formed). Neither hydrobromic nor hydriodic acids can be collected over mercury, on which they act, but they may be directly collected in a dry vessel by leading the gas-conducting tube to the bottom of the vessel, both gases being much heavier than air. Merz and Holtzmann (1889) proposed to prepare HBr directly from bromine and hydrogen. For this purpose pure dry hydrogen is passed through a flask containing boiling bromine. The mixture of gas and vapour then passes through a tube provided with one or two bulbs, which is heated moderately in the middle. Hydrobromic acid is formed with a series of flashes at the part heated. The resultant HBr, together with traces of bromine, passes into a Woulfe's bottle into which hydrogen is also introduced, and the mixture is then carried through another heated tube, after which it is passed through water which dissolves the hydrobromic acid. According to the method proposed by Newth (1892) a mixture of bromine and hydrogen is led through a tube containing a platinum spiral, which is heated to redness after the air has been displaced from the tube. If the vessel containing the bromine be kept at 60°, the hydrogen takes up almost the theoretical amount of bromine required for the formation of HBr. Although the flame which appears in the neighbourhood of the platinum spiral does not penetrate into the vessel containing the bromine, still, for safety, a tube filled with cotton wool may be interposed.
Hydriodic acid is obtained in the same manner as hydrobromic. The iodine is heated in a small flask, and its vapour is carried over by hydrogen into a strongly heated tube, The gas passing from the tube is found to contain a considerable amount of HI, together with some free iodine. At a low red heat about 17 p.c. of the iodine vapour enters into combination; at a higher temperature, 78 p.c. to 79 p.c.; and at a strong heat about 82 p.c.
[77]But generally more phosphorus is taken than is required for the formation of PI3, because otherwise a portion of the iodine distils over. If less than one-tenth part of iodine be taken, much phosphonium iodide, PH4I, is formed. This proportion was established by Gay-Lussac and Kolbe. Hydriodic acid is also prepared in many other ways. Bannoff dissolves two parts of iodine in one part of a previously prepared strong (sp. gr. 1·67) solution of hydriodic acid, and pours it on to red phosphorus in a retort. Personne takes a mixture of fifteen parts of water, ten of iodine, and one of red phosphorus, which, when heated, disengages hydriodic acid mixed with iodine vapour; the latter is removed by passing it over moist phosphorus (Note76). It must be remembered however that reverse reaction (Oppenheim) may take place between the hydriodic acid and phosphorus, in which the compounds PH4I and PI2are formed.It should be observed that the reaction between phosphorus, iodine and water must be carried out in the above proportions and with caution, as they may react with explosion. With red phosphorus the reaction proceeds quietly, but nevertheless requires care.L. Meyer showed that with an excess of iodine the reaction proceeds without the formation of bye-products (PH4I), according to the equation P + 5I + 4H2O = PH3O4+ 5HI. For this purpose 100 grams of iodine and 10 grams of water are placed in a retort, and a paste of 5 grams of red phosphorus and 10 grams of water is added little by little (at first with great care). The hydriodic acid may be obtained free from iodine by directing the neck of the retort upwards and causing the gas to pass through a shallow layer of water (respecting the formation of HI,seealso Note75).
[77]But generally more phosphorus is taken than is required for the formation of PI3, because otherwise a portion of the iodine distils over. If less than one-tenth part of iodine be taken, much phosphonium iodide, PH4I, is formed. This proportion was established by Gay-Lussac and Kolbe. Hydriodic acid is also prepared in many other ways. Bannoff dissolves two parts of iodine in one part of a previously prepared strong (sp. gr. 1·67) solution of hydriodic acid, and pours it on to red phosphorus in a retort. Personne takes a mixture of fifteen parts of water, ten of iodine, and one of red phosphorus, which, when heated, disengages hydriodic acid mixed with iodine vapour; the latter is removed by passing it over moist phosphorus (Note76). It must be remembered however that reverse reaction (Oppenheim) may take place between the hydriodic acid and phosphorus, in which the compounds PH4I and PI2are formed.
It should be observed that the reaction between phosphorus, iodine and water must be carried out in the above proportions and with caution, as they may react with explosion. With red phosphorus the reaction proceeds quietly, but nevertheless requires care.
L. Meyer showed that with an excess of iodine the reaction proceeds without the formation of bye-products (PH4I), according to the equation P + 5I + 4H2O = PH3O4+ 5HI. For this purpose 100 grams of iodine and 10 grams of water are placed in a retort, and a paste of 5 grams of red phosphorus and 10 grams of water is added little by little (at first with great care). The hydriodic acid may be obtained free from iodine by directing the neck of the retort upwards and causing the gas to pass through a shallow layer of water (respecting the formation of HI,seealso Note75).
[78]The specific gravities of their solutions as deduced by me on the basis of Topsöe and Berthelot's determinations for 15°/4° are as follows:—102030405060 p.c.HBr1·0711·1561·2581·3741·5051·650HI1·0751·1641·2671·3991·5671·769Hydrobromic acid forms two hydrates, HBr,2H2O and HBr,H2O, which have been studied by Roozeboom with as much completeness as the hydrate of hydrochloric acid (Chapter X. Note37).With metallic silver, solutions of hydriodic acid give hydrogen with great ease, forming silver iodide. Mercury, lead, and other metals act in a similar manner.
[78]The specific gravities of their solutions as deduced by me on the basis of Topsöe and Berthelot's determinations for 15°/4° are as follows:—
Hydrobromic acid forms two hydrates, HBr,2H2O and HBr,H2O, which have been studied by Roozeboom with as much completeness as the hydrate of hydrochloric acid (Chapter X. Note37).
With metallic silver, solutions of hydriodic acid give hydrogen with great ease, forming silver iodide. Mercury, lead, and other metals act in a similar manner.
[78 bis]Iodide of nitrogen, NHI2is obtained as a brown pulverulent precipitate on adding a solution of iodine (in alcohol, for instance) to a solution of ammonia. If it be collected on a filter-paper, it does not decompose so long as the precipitate is moist; but when dry it explodes violently, so that it can only be experimented upon in small quantities. Usually the filter-paper is torn into bits while moist, and the pieces laid upon a brick; on drying an explosion proceeds not only from friction or a blow, but even spontaneously. The more dilute the solution of ammonia, the greater is the amount of iodine required for the formation of the precipitate of NHI2. A low temperature facilitates its formation. NHI2dissolves in ammonia water, and when heated the solution forms HIO3and iodine. With KI, iodide of nitrogen gives iodine, NH3and KHO. These reactions (Selivanoff) are explained by the formation of HIO from NHI2+ 2H2O = NH3+ 2HIO—and then KI + HIO = I2+ KHO. Selivanoff (seeNote29) usually observed a temporary formation of hypoiodous acid, HIO, in the reaction of ammonia upon iodine, so that here the formation of NHI2is preceded by that of HIO—i.e.first I2+ H2O = HIO + HI, and then not only the HI combines with NH3, but also 2HIO + NH3= NHI2+ 2H2O. With dilute sulphuric acid iodide of nitrogen (like NCl3) forms hypoiodous acid, but it immediately passes into iodic acid, as is expressed by the equation 5HIO = 2I2+ HIO3+ 2H2O (first 3HIO = HIO3+ 2HI, and then HI + HIO = I2+ H2O). Moreover, Selivanoff found that iodide of nitrogen, NHI2, dissolves in an excess of ammonia water, and that with potassium iodide the solution gives the reaction for hypoiodous acid (the evolution of iodine in an alkaline solution). This shows that HIO participates in the formation and decomposition of NHI2, and therefore the condition of the iodine (its metaleptic position) in them is analogous, and differs from the condition of the halogens in the haloid-anhydrides (for instance, NO2Cl). The latter are tolerably stable, while (the haloid being designated by X) NHX2, NX3, XOH, RXO (seeChapter XIII. Note43), &c., are unstable, easily decomposed with the evolution of heat, and, under the action of water, the haloid is easily replaced by hydrogen (Selivanoff), as would be expected in true products of metalepsis.
[78 bis]Iodide of nitrogen, NHI2is obtained as a brown pulverulent precipitate on adding a solution of iodine (in alcohol, for instance) to a solution of ammonia. If it be collected on a filter-paper, it does not decompose so long as the precipitate is moist; but when dry it explodes violently, so that it can only be experimented upon in small quantities. Usually the filter-paper is torn into bits while moist, and the pieces laid upon a brick; on drying an explosion proceeds not only from friction or a blow, but even spontaneously. The more dilute the solution of ammonia, the greater is the amount of iodine required for the formation of the precipitate of NHI2. A low temperature facilitates its formation. NHI2dissolves in ammonia water, and when heated the solution forms HIO3and iodine. With KI, iodide of nitrogen gives iodine, NH3and KHO. These reactions (Selivanoff) are explained by the formation of HIO from NHI2+ 2H2O = NH3+ 2HIO—and then KI + HIO = I2+ KHO. Selivanoff (seeNote29) usually observed a temporary formation of hypoiodous acid, HIO, in the reaction of ammonia upon iodine, so that here the formation of NHI2is preceded by that of HIO—i.e.first I2+ H2O = HIO + HI, and then not only the HI combines with NH3, but also 2HIO + NH3= NHI2+ 2H2O. With dilute sulphuric acid iodide of nitrogen (like NCl3) forms hypoiodous acid, but it immediately passes into iodic acid, as is expressed by the equation 5HIO = 2I2+ HIO3+ 2H2O (first 3HIO = HIO3+ 2HI, and then HI + HIO = I2+ H2O). Moreover, Selivanoff found that iodide of nitrogen, NHI2, dissolves in an excess of ammonia water, and that with potassium iodide the solution gives the reaction for hypoiodous acid (the evolution of iodine in an alkaline solution). This shows that HIO participates in the formation and decomposition of NHI2, and therefore the condition of the iodine (its metaleptic position) in them is analogous, and differs from the condition of the halogens in the haloid-anhydrides (for instance, NO2Cl). The latter are tolerably stable, while (the haloid being designated by X) NHX2, NX3, XOH, RXO (seeChapter XIII. Note43), &c., are unstable, easily decomposed with the evolution of heat, and, under the action of water, the haloid is easily replaced by hydrogen (Selivanoff), as would be expected in true products of metalepsis.
[79]Hypoiodous acid, HIO, is not known, but organic compounds, RIO, of this type are known. To illustrate the peculiarities of their properties we will mention one of these compounds, namely,iodosobenzol, C6H5IO. This substance was obtained by Willgerodt (1892), and also by V. Meyer, Wachter, and Askenasy, by the action of caustic alkalis upon phenoldiiodochloride, C6H5ICl2(according to the equation, C6H5ICl2+ 2MOH = C6H5IO + 2MCl + H2O). Iodosobenzol is an amorphous yellow substance, whose melting point could not be determined because it explodes at 210°, decomposing with the evolution of iodine vapour. This substance dissolves in hot water and alcohol, but is not soluble in the majority of other neutral organic solvents. If acids do not oxidise C6H5IO, they give saline compounds in which iodosobenzol appears as a basic oxide of a diatomic metal, C6H5I. Thus, for instance, when an acetic acid solution of iodosobenzol is treated with a solution of nitric acid, it gives large monoclinic crystals of a nitric acid salt having the composition C6H5I(NO3)2(like Ca(NO3)2). In appearing as the analogue of basic oxides, iodosobenzol displaces iodine from potassium iodide (in a solution acidulated with acetic or hydrochloric acid)—i.e.it acts with its oxygen like HClO. The action of peroxide of hydrogen, chromic acid, and other similar oxidising agents gives iodoxybenzol, C6H5IO2, which is a neutral substance—i.e.incapable of giving salts with acids (compare Chapter XIII. Note43).
[79]Hypoiodous acid, HIO, is not known, but organic compounds, RIO, of this type are known. To illustrate the peculiarities of their properties we will mention one of these compounds, namely,iodosobenzol, C6H5IO. This substance was obtained by Willgerodt (1892), and also by V. Meyer, Wachter, and Askenasy, by the action of caustic alkalis upon phenoldiiodochloride, C6H5ICl2(according to the equation, C6H5ICl2+ 2MOH = C6H5IO + 2MCl + H2O). Iodosobenzol is an amorphous yellow substance, whose melting point could not be determined because it explodes at 210°, decomposing with the evolution of iodine vapour. This substance dissolves in hot water and alcohol, but is not soluble in the majority of other neutral organic solvents. If acids do not oxidise C6H5IO, they give saline compounds in which iodosobenzol appears as a basic oxide of a diatomic metal, C6H5I. Thus, for instance, when an acetic acid solution of iodosobenzol is treated with a solution of nitric acid, it gives large monoclinic crystals of a nitric acid salt having the composition C6H5I(NO3)2(like Ca(NO3)2). In appearing as the analogue of basic oxides, iodosobenzol displaces iodine from potassium iodide (in a solution acidulated with acetic or hydrochloric acid)—i.e.it acts with its oxygen like HClO. The action of peroxide of hydrogen, chromic acid, and other similar oxidising agents gives iodoxybenzol, C6H5IO2, which is a neutral substance—i.e.incapable of giving salts with acids (compare Chapter XIII. Note43).
[79 bis]The oxidation of iodine by strong nitric acid was discovered by Connell; Millon showed that it is effected, although more slowly, by the action of the hydrates of nitric acid up to HNO3,H2O, but that the solution HNO3,2H2O, and weaker solutions, do not oxidise, but simply dissolve, iodine. The participation of water in reactions is seen in this instance. It is also seen, for example, in the fact that dry ammonia combines directly with iodine—for instance, at 0° forming the compound I2,4NH3—whilst iodide of nitrogen is only formed in presence of water.
[79 bis]The oxidation of iodine by strong nitric acid was discovered by Connell; Millon showed that it is effected, although more slowly, by the action of the hydrates of nitric acid up to HNO3,H2O, but that the solution HNO3,2H2O, and weaker solutions, do not oxidise, but simply dissolve, iodine. The participation of water in reactions is seen in this instance. It is also seen, for example, in the fact that dry ammonia combines directly with iodine—for instance, at 0° forming the compound I2,4NH3—whilst iodide of nitrogen is only formed in presence of water.
[80]Bromine also displaces chlorine—for instance, from chloric acid, directly forming bromic acid. If a solution of potassium chlorate be taken (75 parts per 400 parts of water), and iodine be added to it (80 parts), and then a small quantity of nitric acid, chlorine is disengaged on boiling, and potassium iodate is formed in the solution. In this instance the nitric acid first evolves a certain portion of the chloric acid, and the latter, with the iodine, evolves chlorine. The iodic acid thus formed acts on a further quantity of the potassium chlorate, sets a portion of the chloric acid free, and in this manner the action is kept up. Potilitzin (1887) remarked, however, that not only do bromine and iodine displace the chlorine from chloric acid and potassium chlorate, but also chlorine displaces bromine from sodium bromate, and, furthermore, the reaction does not proceed as a direct substitution of the halogens, but is accompanied by the formation of free acids; for example, 5NaClO3+ 3Br2+ 3H2O = 5NaBr + 5HClO3+ HBrO3.
[80]Bromine also displaces chlorine—for instance, from chloric acid, directly forming bromic acid. If a solution of potassium chlorate be taken (75 parts per 400 parts of water), and iodine be added to it (80 parts), and then a small quantity of nitric acid, chlorine is disengaged on boiling, and potassium iodate is formed in the solution. In this instance the nitric acid first evolves a certain portion of the chloric acid, and the latter, with the iodine, evolves chlorine. The iodic acid thus formed acts on a further quantity of the potassium chlorate, sets a portion of the chloric acid free, and in this manner the action is kept up. Potilitzin (1887) remarked, however, that not only do bromine and iodine displace the chlorine from chloric acid and potassium chlorate, but also chlorine displaces bromine from sodium bromate, and, furthermore, the reaction does not proceed as a direct substitution of the halogens, but is accompanied by the formation of free acids; for example, 5NaClO3+ 3Br2+ 3H2O = 5NaBr + 5HClO3+ HBrO3.
[81]If iodine be stirred up in water, and chlorine passed through the mixture, the iodine is dissolved; the liquid becomes colourless, and contains, according to the relative amounts of water and chlorine, either IHCl2, or ICl3, or HIO3. If there be a small amount of water, then the iodic acid may separate out directly as crystals, but a complete conversion (Bornemann) only occurs when not less than ten parts of water are taken to one part of iodine—ICl + 3H2O + 2Cl2= IHO3+ 5HCl.
[81]If iodine be stirred up in water, and chlorine passed through the mixture, the iodine is dissolved; the liquid becomes colourless, and contains, according to the relative amounts of water and chlorine, either IHCl2, or ICl3, or HIO3. If there be a small amount of water, then the iodic acid may separate out directly as crystals, but a complete conversion (Bornemann) only occurs when not less than ten parts of water are taken to one part of iodine—ICl + 3H2O + 2Cl2= IHO3+ 5HCl.
[82]Schönbein and Ogier proved this. Ogier found that at 45° ozone immediately oxidises iodine vapour, forming first of all the oxide I2O3, which is decomposed by water or on heating into iodic anhydride and iodine. Iodic acid is formed at the positive pole when a solution of hydriodic acid is decomposed by a galvanic current (Riche). It is also formed in the combustion of hydrogen mixed with a small quantity of hydriodic acid (Salet).
[82]Schönbein and Ogier proved this. Ogier found that at 45° ozone immediately oxidises iodine vapour, forming first of all the oxide I2O3, which is decomposed by water or on heating into iodic anhydride and iodine. Iodic acid is formed at the positive pole when a solution of hydriodic acid is decomposed by a galvanic current (Riche). It is also formed in the combustion of hydrogen mixed with a small quantity of hydriodic acid (Salet).
[83]Kämmerer showed that a solution of sp. gr. 2·127 at 14°, containing 2HIO3,9H2O, solidified completely in the cold. On comparing solutions HI +mH2O with HIO3+mH2O, we find that the specific gravity increases but the volume decreases, whilst in the passage of solutions HCl +mH2O to HClO3+mH2O both the specific gravity and the volume increase, which is also observed in certain other cases (for example, H3PO3and H3PO4).
[83]Kämmerer showed that a solution of sp. gr. 2·127 at 14°, containing 2HIO3,9H2O, solidified completely in the cold. On comparing solutions HI +mH2O with HIO3+mH2O, we find that the specific gravity increases but the volume decreases, whilst in the passage of solutions HCl +mH2O to HClO3+mH2O both the specific gravity and the volume increase, which is also observed in certain other cases (for example, H3PO3and H3PO4).
[83 bis]Ditte (1890) obtained many iodates of great variety. A neutral salt, 2(LiIO3)H2O, is obtained by saturating a solution of lithia with iodic acid. There is an analogous ammonium salt, 2(NH4IO3)H2O. He also obtained hydrates of a more complex composition, such as 6(NH4IO3)H2O and 6(NH4IO3)2H2O. Salts of the alkaline earths, Ba(IO3)2H2O and Sr(IO3)2H2O, may be obtained by a reaction of double decomposition from the normal salts of the type 2(MeIO3)H2O. When evaporated at 70° to 80° with nitric acid these salts lose water. A mixture of solutions of nitrate of zinc and an alkaline iodate precipitates Zn(IO3)22H2O. An anhydrous salt is thrown out if nitric acid be added to the solutions. Analogous salts of cadmium, silver, and copper give compounds of the type 2Me′IO34NH3and Me″(IO3)24NH3, with gaseous ammonia (Me′ and Me″ being elements of the first (Ag) and second (Cd, Zn, Cu) groups). With an aqueous solution of ammonia the above salts give substances of a different composition, such as Zn(IO3)2(NH4)2O, Cd(IO3)2(NH4)2O. Copper gives Cu(IO3)24(NH4)2O and Cu(IO3)2(NH4)2O. These salts may be regarded as compounds of I2O5, and MeO and (NH4)2O; for example, Zn(IO3)2(NH4)2O may be regarded as ZnO(NH4)2OI2O5, or, as derived from the hydrate, I2O52H2O = 2(HIO3)H2O.
[83 bis]Ditte (1890) obtained many iodates of great variety. A neutral salt, 2(LiIO3)H2O, is obtained by saturating a solution of lithia with iodic acid. There is an analogous ammonium salt, 2(NH4IO3)H2O. He also obtained hydrates of a more complex composition, such as 6(NH4IO3)H2O and 6(NH4IO3)2H2O. Salts of the alkaline earths, Ba(IO3)2H2O and Sr(IO3)2H2O, may be obtained by a reaction of double decomposition from the normal salts of the type 2(MeIO3)H2O. When evaporated at 70° to 80° with nitric acid these salts lose water. A mixture of solutions of nitrate of zinc and an alkaline iodate precipitates Zn(IO3)22H2O. An anhydrous salt is thrown out if nitric acid be added to the solutions. Analogous salts of cadmium, silver, and copper give compounds of the type 2Me′IO34NH3and Me″(IO3)24NH3, with gaseous ammonia (Me′ and Me″ being elements of the first (Ag) and second (Cd, Zn, Cu) groups). With an aqueous solution of ammonia the above salts give substances of a different composition, such as Zn(IO3)2(NH4)2O, Cd(IO3)2(NH4)2O. Copper gives Cu(IO3)24(NH4)2O and Cu(IO3)2(NH4)2O. These salts may be regarded as compounds of I2O5, and MeO and (NH4)2O; for example, Zn(IO3)2(NH4)2O may be regarded as ZnO(NH4)2OI2O5, or, as derived from the hydrate, I2O52H2O = 2(HIO3)H2O.
[84]If sodium iodate be mixed with a solution of sodium hydroxide, heated, and chlorine passed through the solution, a sparingly soluble salt separates out, which corresponds with periodic acid, and has the composition Na4I2O9,3H2O.6NaHO + 2NaIO3+ 4Cl = 4NaCl + Na4I2O9+ 3H2O.This compound is sparingly soluble in water, but dissolves easily in a very dilute solution of nitric acid. If silver nitrate be added to this solution a precipitate is formed which contains the corresponding compound of silver, Ag4I2O9,3H2O. If this sparingly soluble silver compound be dissolved in hot nitric acid, orange crystals of a salt having the composition AgIO4separate on evaporation. This salt is formed from the preceding by the nitric acid taking up silver oxide—Ag4I2O_9 + 2HNO3= 2AgNO3+ 2AgIO4+ H2O. The silver salt is decomposed by water, with the re-formation of the preceding salt, whilst iodic acid remains in solution—4AgIO4+ H2O = Ag4I2O9+ 2HIO4.The structure of the first of these salts, Na4I2O9,3H2O, presents itself in a simpler form if the water of crystallisation is regarded as an integral portion of the salt; the formula is then divided in two, and takes the form of IO(OH)3(ONa)2—that is, it answers to the type IOX5, or IX7, like AgIO4which is IO3(OAg). The composition of all the salts of periodic acids are expressed by this type IX7. Kimmins (1889) refers all the salts of periodic acid to four types—the meta-salts of HIO4(salts of Ag, Cu, Pb), the meso-salts of H3IO5(PbH, Ag2H, CdH), the para-salts of H5IO6(Na2H3, Na3H2), and the di-salts of H4I2O9(K4, Ag4, Ni2). The three first are direct compounds of the type IX7, namely, IO3(OH), IO2(OH)3, and IO(OH)5, and the last are types of diperiodic salts, which correspond with the type of the meso-salts, as pyrophosphoric salts correspond with orthophosphoric salts—i.e.2H3IO5-H2O = H4I2O9.
[84]If sodium iodate be mixed with a solution of sodium hydroxide, heated, and chlorine passed through the solution, a sparingly soluble salt separates out, which corresponds with periodic acid, and has the composition Na4I2O9,3H2O.
6NaHO + 2NaIO3+ 4Cl = 4NaCl + Na4I2O9+ 3H2O.
This compound is sparingly soluble in water, but dissolves easily in a very dilute solution of nitric acid. If silver nitrate be added to this solution a precipitate is formed which contains the corresponding compound of silver, Ag4I2O9,3H2O. If this sparingly soluble silver compound be dissolved in hot nitric acid, orange crystals of a salt having the composition AgIO4separate on evaporation. This salt is formed from the preceding by the nitric acid taking up silver oxide—Ag4I2O_9 + 2HNO3= 2AgNO3+ 2AgIO4+ H2O. The silver salt is decomposed by water, with the re-formation of the preceding salt, whilst iodic acid remains in solution—
4AgIO4+ H2O = Ag4I2O9+ 2HIO4.
The structure of the first of these salts, Na4I2O9,3H2O, presents itself in a simpler form if the water of crystallisation is regarded as an integral portion of the salt; the formula is then divided in two, and takes the form of IO(OH)3(ONa)2—that is, it answers to the type IOX5, or IX7, like AgIO4which is IO3(OAg). The composition of all the salts of periodic acids are expressed by this type IX7. Kimmins (1889) refers all the salts of periodic acid to four types—the meta-salts of HIO4(salts of Ag, Cu, Pb), the meso-salts of H3IO5(PbH, Ag2H, CdH), the para-salts of H5IO6(Na2H3, Na3H2), and the di-salts of H4I2O9(K4, Ag4, Ni2). The three first are direct compounds of the type IX7, namely, IO3(OH), IO2(OH)3, and IO(OH)5, and the last are types of diperiodic salts, which correspond with the type of the meso-salts, as pyrophosphoric salts correspond with orthophosphoric salts—i.e.2H3IO5-H2O = H4I2O9.
[85]Periodic acid, discovered by Magnus and Ammermüller, and whose salts were afterwards studied by Langlois, Rammelsberg, and many others, presents an example of hydrates in which it is evident that there is not that distinction between the water of hydration and of crystallisation which was at first considered to be so clear. In HClO,2H2O the water, 2H2O, is not displaced by bases, and must be regarded as water of crystallisation, whilst in HIO4,2H2O it must be regarded as water of hydration. We shall afterwards see that the system of the elements obliges us to consider the halogens as substances giving a highest saline type,GX7, whereGsignifies a halogen, andXoxygen (O =X2), OH, and other like elements. The hydrate IO(OH)5corresponding with many of the salts of periodic acid (for example, the salts of barium, strontium, mercury) does not exhaust all the possible forms. It is evident that various other pyro-, meta-, &c., forms are possible by the loss of water, as will be more fully explained in speaking of phosphoric acid, and as was pointed out in the preceding note.
[85]Periodic acid, discovered by Magnus and Ammermüller, and whose salts were afterwards studied by Langlois, Rammelsberg, and many others, presents an example of hydrates in which it is evident that there is not that distinction between the water of hydration and of crystallisation which was at first considered to be so clear. In HClO,2H2O the water, 2H2O, is not displaced by bases, and must be regarded as water of crystallisation, whilst in HIO4,2H2O it must be regarded as water of hydration. We shall afterwards see that the system of the elements obliges us to consider the halogens as substances giving a highest saline type,GX7, whereGsignifies a halogen, andXoxygen (O =X2), OH, and other like elements. The hydrate IO(OH)5corresponding with many of the salts of periodic acid (for example, the salts of barium, strontium, mercury) does not exhaust all the possible forms. It is evident that various other pyro-, meta-, &c., forms are possible by the loss of water, as will be more fully explained in speaking of phosphoric acid, and as was pointed out in the preceding note.
[86]With respect to hydrogen, oxygen, chlorine, and other elements, bromine occupies an intermediate position between chlorine and iodine, and therefore there is no particular need for considering at length the compounds of bromine. This is the great advantage of a natural grouping of the elements.
[86]With respect to hydrogen, oxygen, chlorine, and other elements, bromine occupies an intermediate position between chlorine and iodine, and therefore there is no particular need for considering at length the compounds of bromine. This is the great advantage of a natural grouping of the elements.
[87]They were both obtained by Gay-Lussac and many others. Recent data respecting iodine monochloride, ICl, entirely confirm the numerous observations of Trapp (1854), and even confirm his statement as to the existence of two isomeric (liquid and crystalline) forms (Stortenbeker). With a small excess of iodine, iodine monochloride remains liquid, but in the presence of traces of iodine trichloride it easily crystallises. Tanatar (1893) showed that of the two modifications of ICl, one is stable, and melts at 27°; while the other, which easily passes into the first, and is formed in the absence of ICl3, melts at 14°. Schützenberger amplified the data concerning the action of water on the chlorides (Note88), and Christomanos gave the fullest data regarding the trichloride.After being kept for some time, the liquid monochloride of iodine yields red deliquescent octahedra, having the composition ICl4, which are therefore formed from the monochloride with the liberation of free iodine, which dissolves in the remaining quantity of the monochloride. This substance, however, judging by certain observations, is impure iodine trichloride. If 1 part of iodine be stirred up in 20 parts of water, and chlorine be passed through the liquid, then all the iodine is dissolved, and a colourless liquid is ultimately obtained which contains a certain proportion of chlorine, because this compound gives a metallic chloride and iodate with alkalis without evolving any free iodine: ICl5+ 6KHO = 5KCl + KIO3+ 3H2O. The existence of a pentachloride ICl5is, however, denied, because this substance has not been obtained in a free state.Stortenbeker (1888) investigated the equilibrium of the system containing the molecules I2, ICl, ICl3, and Cl2, in the same way that Roozeboom (Chapter X. Note38) examined the equilibrium of the molecules HCl, HCl,2H2O, and H2O. He found that iodine monochloride appears in two states, one (the ordinary) is stable and melts at 27°·2, whilst the other is obtained by rapid cooling, and melts at 13°·9, and easily passes into the first form. Iodine trichloride melts at 101° only in a closed tube under a pressure of 16 atmospheres.
[87]They were both obtained by Gay-Lussac and many others. Recent data respecting iodine monochloride, ICl, entirely confirm the numerous observations of Trapp (1854), and even confirm his statement as to the existence of two isomeric (liquid and crystalline) forms (Stortenbeker). With a small excess of iodine, iodine monochloride remains liquid, but in the presence of traces of iodine trichloride it easily crystallises. Tanatar (1893) showed that of the two modifications of ICl, one is stable, and melts at 27°; while the other, which easily passes into the first, and is formed in the absence of ICl3, melts at 14°. Schützenberger amplified the data concerning the action of water on the chlorides (Note88), and Christomanos gave the fullest data regarding the trichloride.
After being kept for some time, the liquid monochloride of iodine yields red deliquescent octahedra, having the composition ICl4, which are therefore formed from the monochloride with the liberation of free iodine, which dissolves in the remaining quantity of the monochloride. This substance, however, judging by certain observations, is impure iodine trichloride. If 1 part of iodine be stirred up in 20 parts of water, and chlorine be passed through the liquid, then all the iodine is dissolved, and a colourless liquid is ultimately obtained which contains a certain proportion of chlorine, because this compound gives a metallic chloride and iodate with alkalis without evolving any free iodine: ICl5+ 6KHO = 5KCl + KIO3+ 3H2O. The existence of a pentachloride ICl5is, however, denied, because this substance has not been obtained in a free state.
Stortenbeker (1888) investigated the equilibrium of the system containing the molecules I2, ICl, ICl3, and Cl2, in the same way that Roozeboom (Chapter X. Note38) examined the equilibrium of the molecules HCl, HCl,2H2O, and H2O. He found that iodine monochloride appears in two states, one (the ordinary) is stable and melts at 27°·2, whilst the other is obtained by rapid cooling, and melts at 13°·9, and easily passes into the first form. Iodine trichloride melts at 101° only in a closed tube under a pressure of 16 atmospheres.
[88]By the action of water on iodine monochloride and trichloride a compound IHCl2is obtained, which does not seem to be altered by water. Besides this compound, iodine and iodic acid are always formed, 10ICl + 3H2O = HIO3+ 5IHCl2+ 2I2; and in this respect iodine trichloride may be regarded as a mixture, ICl + ICl5= 2ICl3, but ICl5+ 3H2O = IHO3+ 5HCl; hence iodic acid, iodine, the compound IHCl2, and hydrochloric acid are also formed by the action of water.
[88]By the action of water on iodine monochloride and trichloride a compound IHCl2is obtained, which does not seem to be altered by water. Besides this compound, iodine and iodic acid are always formed, 10ICl + 3H2O = HIO3+ 5IHCl2+ 2I2; and in this respect iodine trichloride may be regarded as a mixture, ICl + ICl5= 2ICl3, but ICl5+ 3H2O = IHO3+ 5HCl; hence iodic acid, iodine, the compound IHCl2, and hydrochloric acid are also formed by the action of water.