[48]The greatest part of the sulphuric acid is used in the soda manufacture, in the conversion of the common salt into sulphate. For this purpose an acid having a density of 60° Baumé is amply sufficient. Chamber acid has a density up to 1·57 = 50° to 51° Baumé; it contains about 35 per cent. of water. About 15 per cent. of this water can be removed in leaden stills, and nearly all the remainder may be expelled in glass or platinum vessels. Acid of 66° Baumé, = 1·847, contains about 96 per cent. of the hydrate H2SO4. The density falls with a greater or less proportion of water, the maximum density corresponding with 97½ per cent. of the hydrate H2SO4. The concentration of H2SO4in platinum retorts has the disadvantage that sulphuric acid, upwards of 90 per cent. in strength, does corrode platinum, although but slightly (a few grams per tens of tons of acid). The retorts therefore require repairing, and the cost of the platinum exceeds the price obtained for concentrating the acid from 90 per cent. to 98 per cent. (in factories the acid is not concentrated beyond this by evaporation in the air). This inconvenience has lately (1891, by Mathey) been eliminated by coating the inside of the platinum retorts with a thin (0·1 to 0·02 mm.) layer of gold which is 40 times less corroded by sulphuric acid than platinum. Négrier (1890) carries on the distillation in porcelain dishes, Blond by heating a thin platinum wire immersed in the acid by means of an electric current, but the most promising method is that of Kessler (1891), which consists in passing hot air over sulphuric acid flowing in a thin stream in stone vessels, so that there is no boiling but only evaporation at moderate temperatures: the transference of the heat is direct (and not through the sides of the vessels), which economises the fuel and prevents the distilling vessels being damaged.When, by evaporation of the water, sulphuric acid attains a density of 66° Baumé (sp. gr. 1·84), it is impossible to concentrate it further, because it then distils over unchanged.The distillation of sulphuric acidis not generally carried on on a large scale, but forms a laboratory process, employed when particularly pure acid is required. The distillation is effected either in platinum retorts furnished with corresponding condensers and receivers, or in glass retorts. In the latter case, great caution is necessary, because the boiling of sulphuric acid itself is accompanied by still more violent jerks and greater irregularity than even the evaporation of the last portions of water contained in the acid. If the glass retort which holds the strong sulphuric acid to be distilled be heated directly from below, it frequently jerks and breaks. For greater safety the heating is not effected from below, but at the sides of the retort. The evaporation then does not proceed in the whole mass, but only from the upper portions of the liquid, and therefore goes on much more quietly. The acid may be made to boil quietly also by surrounding the retort with good conductors of heat—for example, iron filings, or by immersing a bunch of platinum wires in the acid, as the bubbles of sulphuric acid vapour then form on the extremities of the wires.[49]Thus it appears that so common, and apparently so stable, a compound as sulphuric acid decomposes even at a low temperature with separation of the anhydride, but this decomposition is restricted by a limit, corresponding to the presence of about 1½ p.c. of water, or to a composition of nearly H2O,12H2SO4.Now there is no reason for thinking that this substance is a definite compound; it is an equilibrated system which does not decompose under ordinary circumstances below 338°. Dittmar carried on the distillation under pressures varying between 30 and 2,140 millimetres (of mercury), and he found that the composition of the residue hardly varies, and contains from 99·2 to 98·2 per cent. of the normal hydrate, although at 30 mm. the temperature of distillation is about 210° and at 2,140 mm. it is 382°. Furthermore, it is a fact of practical importance that under a pressure of two atmospheres the distillation of sulphuric acid proceeds very quietly.Sulphuric acid may bepurifiedfrom the majority of its impurities by distillation, if the first and last portions of the distillate be rejected. The first portions will contain the oxides of nitrogen, hydrochloric acid, &c., and the last portions the less volatile impurities. The oxides of nitrogen may be removed by heating the acid with charcoal, which converts them into volatile gases. Sulphuric acid may be freed from arsenic by heating it with manganese dioxide and then distilling. This oxidises all the arsenic into non-volatile arsenic acid. Without a preliminary oxidation it would partially remain as volatile arsenious acid, and might pass over into the distillate. The arsenic may also be driven off by first reducing it to arsenious acid, and then passing hydrochloric acid gas through the heated acid. It is then converted into arsenious chloride, which volatilises.[50]The amount of heat developed by the mixture of sulphuric acid with water is expressed in the diagram on p.77, Volume I., by the middle curve, whose abscissæ are the percentage amounts of acid (H2SO4) in the resultant solution, and ordinates the number of units of heat corresponding with the formation of 100 cubic centimetres of the solution (at 18°). The calculations on which the curve is designed are based on Thomsen's determinations, which show that 98 grams or a molecular amount of sulphuric acid, in combining withmmolecules of water (that is, withm=18 grams of water), develop the following number of units of heat, R:—m=123591949100200R =6379941811137131081495216256166841685917066c=0·4320·4700·5000·5760·7010·8210·9140·9540·975T =127°149°146°121°82°45°19°9°5°cstands for the specific heat of H2SO4mH2O (according to Marignac and Pfaundler), and T for the rise in temperature which proceeds from the mixture of H2SO4withmH2O. The diagram shows that contraction and rise of temperature proceed almost parallel with each other.[50 bis]Pickering (1890) showed (a) that dilute solutions of sulphuric acid containing up to H2SO4+ 10H2O deposit ice (at -0°·12 when there is 2,000H2O per H2SO4, at -0°·23 when there is 1,000H2O, at -1°·04 when there is 200H2O, at -2°·12 when there is 100H2O, at -4°·5 when there is 50H2O, at -15°·7 when there is 20H2O, and at -61° when the composition of the solution is H2SO4+ 10H2O); (b) that for higher concentrations crystals separate out at a considerable degree of cold, having the composition H2SO44H2O, which melt at -24°·5, and if either water or H2SO4be added to this compound the temperature of crystallisation falls, so that a solution of the composition 12H2SO4+ 100H2O gives crystals of the above hydrate at -70°, 15H2SO4+ 100H2O at -47°, 30H2SO4+ 100H2O at -32°, 40H2SO4+ 100H2O at -52°; (c) that if the amount of H2SO4be still greater, then a hydrate H2SO4H2O separates out and melts at +8°·5, while the addition of water or sulphuric acid to it lowers the temperature of crystallisation so that the crystallisation of H2SO4H2O from a solution of the composition H2SO4+ 1·73H2O takes place at -22°, H2SO4+ 1·5H2O at -6°·5, H2SO4+ 1·2H2O at +3°·7, H2SO4+ 0·75H2O at +2°·8, H2SO4+ 0·5H2O at -16°; (d) that when there is less than 40H2O per 100H2SO4, refrigeration separates out the normal hydrate H2SO4, which melts at +10°·35, and that a solution of the composition H2SO4+ 0·35H2O deposits crystals of this hydrate at -34°, H2SO4+ 0·1H2O at -4°·1, H2SO4+ 0·05H2O at +4°·9, while fuming acid of the composition H2SO4+ 0·05SO3deposits H2SO4at about +7°. Thus the temperature of the separation of crystals clearly distinguishes the above four regions of solutions, and in the space between H2SO4+ H2O and +25H2O a particular hydrate H2SO44H2O separates out, discovered by Pickering, the isolation of which deserves full attention and further research. I may add here that the existence of a hydrate H2SO44H2O was pointed out in my work,The Investigation of Aqueous Solutions, p. 120 (1887), upon the basis that it has at all temperatures a smaller value for the coefficient of expansionkin the formula St= S0/(1 -kt) than the adjacent (in composition) solutions of sulphuric acid. And for solutions approximating to H2SO410H2O in their composition,kis constant at all temperatures (for more dilute solutions the value ofkincreases withtand for more concentrated solutions it decreases). This solution (with 10H2O) forms the point of transition between more dilute solutions which deposit ice (water) when refrigerated and those which give crystals of H2SO44H2O. According to R. Pictet (1894) the solution H2SO410H2O freezes at -88° (but no reference is made as to what separates out),i.e.at a lower temperature than all the other solutions of sulphuric acid. However, in respect to these last researches of R. Pictet (for 88·88 p.c. H2SO4-55°, for H2SO4H2O +3·5°, for H2SO42H2O -70°, for H2SO44H2O -40°, &c.) it should be remarked that they offer some quite improbable data; for example, for H2SO475H2O they give the freezing point as 0°, for H2SO4300H2O +4°·5, and even for H2SO41000H2O +0°·5, although it is well known that a small amount of sulphuric acid lowers the temperature of the formation of ice. I have found by direct experiment that a frozen solidified solution of H2SO4+ 300H2O melted completely at 0°.[51]With an excess of snow, the hydrate H2SO4,H2O, like the normal hydrate, gives a freezing mixture, owing to the absorption of a large amount of heat (the latent heat of fusion). In melting, the molecule H2SO4absorbs 960 heat units, and the molecule H2SO4H2O 3,680 heat units. If therefore we mix one gram molecule of this hydrate with seventeen gram molecules of snow, there is an absorption of 18,080 heat units, because 17H2O absorbs 17 × 1,430 heat units, and the combination of the monohydrate with water evolves 9,800 heat units. As the specific heat of the resultant compound H2SO4,18H2O = 0·813, the fall of temperature will be -52°·6. And, in fact, a very low temperature may be obtained by means of sulphuric acid.[52]For example, if it be taken that at 19° the sp. gr. of pure sulphuric acid is 1·8330, then at 20° it is 1·8330 - (20 - 19)10·13 = 1·8320.[53]Unfortunately, notwithstanding the great number of fragmentary and systematic researches which have been made (by Parks, Ure, Bineau, Kolbe, Lunge, Marignac, Kremers, Thomsen, Perkin, and others) for determining the relation between the sp. gr. and composition of solutions of sulphuric acid, they contain discrepancies which amount to, and even exceed, 0·002 in the sp. gr. For instance, at 15°·4 the solution of composition H2SO43H2O has a sp. gr. 1·5493 according to Perkin (1886), 1·5501 according to Pickering (1890), and 1·5525 according to Lunge (1890). The cause of these discrepancies must be looked for in the methods employed for determining the composition of the solutions—i.e.in the inaccuracy with which the percentage amount of H2SO4is determined, for a difference of 1 p.c. corresponds to a difference of from 0·0070 (for very weak solutions) to 0·0118 (for a solution containing about 73 p.c.) in the specific gravity (that is the factords/dp) at 15°. As it is possible to determine the specific gravity with an accuracy even exceeding 0·0002, the specific gravities given in the adjoining tables are only averages and most probable data in which the error, especially for the 30–80 p.c. solutions cannot be less than 0·0010 (taking water at 4° as 1).[53 bis]Judging from the best existing determinations (of Marignac, Kremers, and Pickering) for solutions of sulphuric acid (especially those containing more than 5 p.c. H2SO4) within the limits of 0° and 30° (and even to 40°), the variation of the sp. gr. with the temperaturetmay (within the accuracy of the existing determinations) be perfectly expressed by the equation St= S0+ At+ Bt2. It must be added that (1) three specific gravities fully determine the variation of the density witht; (2)ds/dt= A + 2Bt—i.e.the factor of the temperature is expressed by a straight line; (3) the value of A (ifpbe greater than 5 p.c.) is negative, and numerically much greater than B; (4) the value of B for dilute solutions containing less than 25 p.c. is negative; for solutions approximating to H2SO43H2O in their composition it is equal to 0, and for solutions of greater concentration B is positive; (5) the factords/dpfor all temperatures attains a maximum value about H2SO4H2O; (6) on dividingds/dtby S0, and so obtaining the coefficient of expansionk(seeNote53), a minimum is obtained near H2SO4and H2SO44H2O, and a maximum at H2SO4H2O for all temperatures.[53 tri]These data (as well as those in the following table) have been recalculated by me chiefly upon the basis of Kremer's, Pickering's, Perkin's, and my own determinations; all the requisite corrections have been introduced, and I have reason for thinking that in each of them the probable error (or difference from the true figures, now unknown) of the specific gravity does not exceed ±0·0007 (if water at 4° = 1) for the 25–80 p.c. solutions, and ±0·0002 for the more dilute or concentrated solutions.[54]The factordS/dppasses through 0, that is, the specific gravity attains a maximum value at about 98 p.c. This was discovered by Kohlrausch, and confirmed by Chertel, Pickering, and others.[55]Naturally under the condition that there is no other ingredient besides water, which is sufficiently true. For commercial acid, whose specific gravity is usually expressed in degrees of Baumé's hydrometer, we may add that at 15°Specific gravity11·11·21·31·41·51·61·71·8Degree Baumé0132433·341·248·154·159·564·266° Baumé (the strongest commercial acid or oil of vitriol) corresponds to a sp. gr. 1·84.By employing the second table (by the method of interpolation) the specific gravity, at a given temperature (from 0° to 30°) can be found for any percentage amount of H2SO4, and therefore conversely the percentage of H2SO4can be found from the specific gravity.[55 bis]Whether similar (even small) breaks in the continuity of the factordS/dpexist or not, for other hydrates (for instance, for H2SO4H2O and H2SO44H2O) cannot as yet be affirmed owing to the want of accurate data (Note53). In my investigation of this subject (1887) I admit their possibility, but only conditionally; and now, without insisting upon a similar opinion, I only hold to the existence of a distinct break in the factor at H2SO4, being guided by C. Winkler's observations ond the specific gravities of fuming sulphuric acid.[56]In 1887, on considering all the existent observations for a temperature 0°, I gave the accompanying scheme (p.243) of the variation of the factords/dpat 0°.I did not then (1887) give this scheme an absolute value, and now after the appearance of two series of new determinations (Lunge and Pickering in 1890), which disagree in many points, I think it well to state quite clearly: (1) that Lunge's and Pickering's new determinations have not added to the accuracy of our data respecting the variation of the specific gravity of solutions of sulphuric acid; (2) that the sum total of existing data does not negative (within the limit of experimental accuracy) the possibility of a rectilinear and broken form for the factorsds/dp; (3) that the supposition of ‘special points’ inds/dp, indicating definite hydrates, finds confirmation in all the latest determinations; (4) that the supposition respecting the existence of hydrates determining a break of the factords/dpis in in way altered if, instead of a series of broken straight lines, there be a continuous series of curves, nearly approaching straight lines; and (5) that this subject deserves (as I mentioned in 1887) new and careful elaboration, because it concerns that foremost problem in our science—solutions—and introduces a special method into it—that is, the study of differential variations in a property which is so easily observed as the specific gravity of a liquid.[56 bis]These hydrates are: (a) H2SO4= SO3H2O (melts at + 10°·4); (b) H2SO4H2O = SO32H2O (crystallo-hydrate, melts at +8°·5); (c) H2SO42H2O (is apparently not crystallisable); (d) one of the hydrates between H2SO46H2O and H2SO43H2O, most probably H2SO44H2O = SO35H2O, for it crystallises at -24°·5 (Note50 bis); and (e) a certain hydrate with a large proportion of water, about H2SO4150H2O. The existence of the last is inferred from the fact that the factords/dpfirst falls, starting from water, and then rises, and this change takes place whenpis less than 5 p.c. Certainly a change in the variation ofds/dpords/dtdoes take place in the neighbourhood of these five hydrates (Pickering, 1890, recognised a far greater number of hydrates). I think it well to add that if the composition of the solutions be expressed by the percentage amount of molecules—r1SO3+ (100 -r1)H2O we find that for H2SO4,r1= 50, for H2SO42H2Or1= 25 = 50/2, for H2SO4H2O,r1= 33·333 = 50·⅔, while for H2SO44H2O,r1= 16·666 = 50·⅓—i.e.that the chief hydrates are distributed symmetrically between H2O and H2SO4. Besides which I may mention that my researches (1887) upon the abrupt changes in the factor for solutions of sulphuric acid, and upon the correspondence of the breaks ofds/dpwith definite hydrates, received an indirect confirmation not only in the solutions of HNO3, HCl, C2H6O, C3H8O, &c., which I investigated (in my work cited in Chapter I., Note19), but also in the careful observations made by Professor Cheltzoff on the solutions of FeCl3and ZnCl2(Chapter XVI., Note4) which showed the existence in these solutions of an almost similar change inds/dpas is found in sulphuric acid. The detailed researches (1893) made by Tourbaba on the solutions of many organic substances are of a similar nature. Besides which, H. Crompton (1888), in his researches on the electrical conductivity of solutions of sulphuric acid, and Tammann, in his observations on their vapour tension, found a correlation with the hydrates indicated as above by the investigation of their specific gravities. The influence of mixtures of a definite composition upon the chemical relations of solutions is even exhibited in such a complex process as electrolysis. V. Kouriloff (1891) showed that mixtures containing about 3 p.c., 47 p.c. and 73 p.c. of sulphuric acid—i.e.whose composition approaches that of the hydrates H2SO4150H2O, H2SO46H2O and H2SO42H2O—exhibit certain peculiarities in respect to the amount of peroxide of hydrogen formed during electrolysis. Thus a 3 p.c. solution gives a maximum amount of peroxide of hydrogen at the negative pole, as compared with that given by other neighbouring concentrations. Starting from 3 p.c., the formation of peroxide of hydrogen ceases until a concentration of 47 p.c. is reached.[57]Cellulose, for instance unsized paper or calico, is dissolved by strong sulphuric acid. Acid diluted with about half its volume of water converts it (if the action be of short duration) into vegetable parchment (Chapter I., Note18). The action of dilute solutions of sulphuric acid converts it into hydro-cellulose, and the fibre loses its coherent quality and becomes brittle. The prolonged action of strong sulphuric acid chars the cellulose while dilute acid converts it into glucose. If sulphuric acid be kept in an open vessel, the organic matter of the dust held in the atmosphere falls into it and blackens the acid. The same thing happens if sulphuric acid be kept in a bottle closed by a cork; the cork becomes charred, and the acid turns black. However, the chemical properties of the acid undergo only a very slight change when it turns black. Sulphuric acid which is considerably diluted with water does not produce the above effects, which clearly shows their dependence on the affinity of the sulphuric acid for water. It is evident from the preceding that strong sulphuric acid will act as a powerful poison; whilst, on the other hand, when very dilute it is employed in certain medicines and as a fertiliser for plants.[58]Weber (1884) obtained a series of salts R2O,8SO3nH2O for K, Rb, Cs, and Tl.[58 bis]Ditte (1890) divides all the metals into two groups with respect to sulphuric acid; the first group includes silver, mercury, copper, lead, and bismuth, which are only acted upon by hot concentrated acid. In this case sulphurous anhydride is evolved without any by-reactions. The second group contains manganese, nickel, cobalt, iron, zinc, cadmium, aluminium, tin, thallium, and the alkali metals. They react with sulphuric acid of any concentration at any temperature. At a low temperature hydrogen is disengaged, and at higher temperatures (and with very concentrated acid) hydrogen and sulphurous anhydride are simultaneously evolved.[59]For example, the action of hot sulphuric acid on nitrogenous compounds, as applied in Kjeldahl's method for the estimation of nitrogen (Volume I. p.249). It is obvious that when sulphuric acid acts as an oxidising agent it forms sulphurous anhydride.The action of sulphuric acid on the alcohols is exactly similar to its action on alkalis, because the alcohols, like alkalis, react on acids; a molecule of alcohol with a molecule of sulphuric acid separates water and forms anacidethereal salt—that is there is produced an ethereal compound corresponding with acid salts. Thus, for example, the action of sulphuric acid, H2SO4, on ordinary alcohol, C2H5OH, gives water and sulphovinic acid, C2H5HSO4—that is, sulphuric acid in which one atom of hydrogen is replaced by the radicle C2H5of ethyl alcohol, SO2(OH)(OC2H5), or, what is the same thing, the hydrogen in alcohol is replaced by the radicle (sulphoxyl) of sulphuric acid, C2H5O.SO2(OH).[60]We will mention the following difference between the sulphonic acids and the ethereal acid sulphates (Note59): the former re-form sulphuric acid with difficulty and the latter easily. Thus sulphovinic acid when heated with an excess of water is reconverted into alcohol and sulphuric acid. This is explained in the following manner. Both these classes of acids are produced by the substitution of hydrogen by SO3H, or the univalent radicle of sulphuric acid, but in the formation of ethereal acid sulphates the SO3H replaces the hydrogen of the hydroxyl in the alcohol, whilst in the formation of the sulphonic acids the SO3H replaces the hydrogen of a hydrocarbon. This difference is clearly evidenced in the existence of two acids of the composition SO4C2H6. The one, mentioned above, is sulphovinic acid or alcohol, C2H5.OH, in which the hydrogen of the hydroxyl is replaced by sulphoxyl = C2H5.OSO3H, whilst the other is alcohol, in which one atom of the hydrogen in ethyl, C2H5, is replaced by the sulphonic group—that is = (C2H4)SO3H·OH. The latter is called isethionic acid. It is more stable than sulphovinic acid. The details as to these interesting compounds must be looked for in works on organic chemistry, but I think it necessary to note one of the general methods of formation of these acids. The sulphites of the alkalis—for example, K2SO3—when heated with the halogen products of metalepsis, give a halogen salt and a salt of a sulphonic acid. Thus methyl iodide, CH3I, derived from marsh gas, CH4, when heated to 100° with a solution of potassium sulphite, K2SO3, gives potassium iodide, KI, and potassium methylsulphonate, CH3SO3K—that is a salt of the sulphonic acid. This shows that the sulphonic acid may be referred to sulphurous acid, and that there is a resemblance between sulphuric and sulphurous acid, which clearly reveals itself here in the formation of one product from them both.[61]The reaction BaO + O develops 12,000 heat units, whilst the reaction H2O + O absorbs 21,000 heat units.[62]Schöne obtained a compound of peroxide of barium with peroxide of hydrogen. If barium peroxide be dissolved in hydrochloric (or acetic) acid, or if a solution of hydrogen peroxide be diluted with a solution of barium hydroxide, a pure hydrate is precipitated having the composition BaO2,8H2O (sometimes the composition is taken as BaO2,6H2O). This fact was already known to Thénard. Schöne showed that if hydrogen peroxide be in excess, a crystalline compound of the two peroxides, BaO2H2O2, is precipitated. Schöne also obtained small well-formed crystals of the same composition by adding a solution of ammonia to an acid solution of barium peroxide (containing a barium salt and hydrogen peroxide or a compound of BaO2with the acid). Thus barium peroxide combines with both water and hydrogen peroxide. This is a very important fact for the comprehension of the composition of other peroxides. Moreover, if the peroxides are able to give hydrates they can also form corresponding salts,i.e.they can combine with bases and acids, as was afterwards found to be the case on further research into this subject.[63]Anhydroussulphuric peroxide, S2O7, is obtained by the prolonged (8 to 10 hours) action of a silent discharge of considerable intensity on a mixture of oxygen and sulphurous anhydride; the vapour of sulphuric peroxide, S2O7, condenses as liquid drops, or after being cooled to 0° in the form of long prismatic crystals, resembling those of sulphuric anhydride. The anhydrous compound S2O7(and also the hydrated compound) cannot be preserved long, as it splits up into oxygen and sulphuric anhydride. Direct experiment shows that a mixture of equal volumes of sulphurous anhydride and oxygen leaves a residue of a quarter of the oxygen taken, or half of the whole volume, which indicates the formula S2O7. This substance is soluble in water, and it then gives a hydrate, probably having the composition S2O7,H2O = 2SHO4. This solution oxidises the salts SnX2, potassium iodide, and others, which renders it possible to prove that the solution actually contains one atom of oxygen capable of effecting oxidation to two molecules of sulphuric anhydride.In order to fully demonstrate the reality of a peroxide form for acids, it should be mentioned that some years ago Brodie obtained the so-calledacetic peroxide, (C2H2O)2O2, by the action of barium peroxide on acetic anhydride, (C2H3O)2O. Its corresponding hydrate is also known. This shows that true peroxides and their hydrates, with reactions similar to those of hydrogen peroxide, are possible for acids. A similar higher oxide has long been known for chromium, and Berthelot obtained a like compound for nitric acid (Chapter VI., Note26).[64]When an acid of the strength H2SO46H2O is taken, at first only the hydrate of the sulphuric peroxide, S2O7H2O, is formed, but when the concentration at the positive pole reaches H2SO43H2O, a mixture of hydrogen peroxide and the hydrate of sulphuric peroxide begins to be formed. A state of equilibrium is ultimately arrived at when the amounts of these substances correspond to the proportion S2O7: 2H2O2, which, as it were, answers to a new hydrate, S2O92H2O. But its existence cannot be admitted because the sulphuric peroxide can be easily distinguished from the hydrogen peroxide in the solution owing to the fact that the former does not act on an acid solution of potassium permanganate, whilst the hydrogen peroxide disengages both its own oxygen and that of the permanganic acid, converting it into manganous oxide. Their common property of liberating iodine from an acid solution of the potassium iodide enables the sum of the active oxygen in them both to be determined.[65]If a solution of sulphuric acid which has been first subjected to electrolysis be neutralised with potash or baryta, the salt which is formed begins to decompose rapidly with the evolution of oxygen (Berthelot, 1890). On saturating with caustic baryta, the solution of the salt formed may be separated from the sulphate of barium, and then the composition of the resultant compound, BaS2O8, may be determined from the amount of oxygen disengaged. Marshall (1891) studied the formation of this class of compounds more fully; he subjected a saturated solution of bisulphate of potassium to electrolysis with a current of 3–3½ ampères; before electrolysis dilute sulphuric acid is added to the liquid surrounding the negative pole, and during electrolysis the solution at the anode is cooled. The electrolysis is continued without interruption for two days, and a white crystalline deposit separates at the anode. To avoid decomposition, the latter is not filtered through paper, but through a perforated platinum plate, and dried on a porous tile. The mother liquor, with the addition of a fresh solution of bisulphate of potassium, is again subjected to electrolysis and the crystals formed at the anode are again collected, &c. The salt so obtained may be recrystallised by dissolving it in hot water and rapidly cooling the solution after filtration; a small proportion of the salt is decomposed by this treatment. Rapid cooling is followed by the formation of small columnar crystals; slow cooling gives large prismatic crystals. The composition of the salt is determined either by igniting it, when it forms sulphate of potassium, or else by titrating the active oxygen with permanganate: its composition was found to correspond to the salt of persulphuric acid, K2S2O8. The solution of the salt has a neutral reaction, and does not give a precipitate with salts of other metals. K2S2O8is the most insoluble of the salts of persulphuric acid. With nitrate of silver it forms persulphate of silver, which gives peroxide of silver under the action of water according to the equation Ag2S2O8+ 2H2O = Ag2O2+ 2H2SO4. With an alkaline solution of a cupric salt (Fehling's solution) it forms a red precipitate of peroxide of copper. Manganese and cobalt salts give precipitates of MnO2and Co2O3. Ferrous salts are rapidly oxidised, potassium iodide slowly disengages iodine at the ordinary temperature. All these reactions indicate the powerful oxidising properties of K2S2O8. In oxidising in the presence of water it gives a residue of KHSO4. The decomposition of the dry salt begins at 100° but is not complete even at 250°. The freshly prepared salt is inodorous, but after being kept in a closed vessel it evolves a peculiar smell different from that of ozone. The ammonium salt of persulphuric acid, (NH4)2S2O8, is obtained in a similar manner. It is soluble to the extent of 58 parts per 100 parts by weight of water. The decomposition of the ammonium salt by the hydrated oxide of barium gives the barium salt, BaS2O84H2O, which is soluble to the extent of 52·2 parts in 100 parts of water at 0°. The crystals do not deliquesce in the air and decompose in the course of several days; they decompose most rapidly in perfectly dry air. Solutions of the pure salt decompose slowly at the ordinary temperature; on boiling barium sulphate is gradually precipitated, oxygen being liberated simultaneously. To completely decompose this salt it is necessary to boil its solution for a long time. Alcohol dissolves the solid salt; the anhydrous salt does not separate from the alcoholic solution, but a hydrate containing one molecule of water, BaS2O8H2O, which is soluble in water but insoluble in absolute alcohol. Solid barium persulphate decomposes even when slightly heated. The free acid, which may serve for the preparation of other salts, is obtained by treating the barium salt with sulphuric acid. The lead salt, PbS2O8, has been obtained from the free acid; it crystallises with two or three molecules of water. It is soluble in water, deliquesces in the air, and with alkalis gives a precipitate of the hydrated oxide which rapidly oxidises into the binoxide.Traube, before Marshall's researches, thought that the electrolysis of solutions of sulphuric acid did not give persulphuric acid but a persulphuric oxide having the composition SO4. On repeating his former researches (1892) Traube obtained a persulphuric oxide by the electrolysis of a 70 per cent. solution of sulphuric acid, and he separated it from the solution by means of barium phosphate. Analysis showed that this substance corresponded to the above composition SO4, and therefore Traube considers it very likely that the salts obtained by Marshall corresponded to an acid H2SO4+ SO4,i.e.that the indifferent oxide, SO4, can combine with sulphuric acid and form peculiar saline compounds.[65 bis]Or one of those supposed ions which appear at the positive pole in the decomposition of sulphuric acid by the action of a galvanic current.[66]If this be true one would expect the following peroxide hydrates: for phosphoric acid, (H2PO4)2= H4P2O8= 2H2O + 2PO3; for carbonic acid, (HCO3)2= H2C2O6= H2O + C2O5; and for lead the true peroxide will be also Pb2O5, &c. Judging from the example of barium peroxide (Note62), these peroxide forms will probably combine together. It seems to me that the compounds obtained by Fairley for uranium are very instructive as elucidating the peroxides. In the action of hydrogen peroxide in an acid solution on uranium oxide, UO3, there is formed a uranium peroxide, UO4,4H2O (U = 240), but hydrogen peroxide acts on uranium oxide in the presence of caustic soda; on the addition of alcohol a crystalline compound containing Na4UO8,4H2O is precipitated, which is doubtless a compound of the peroxides of sodium, Na2O2, and uranium, UO4. It is very possible that the first peroxide, UO4,4H2O, contains the elements of hydrogen peroxide and uranium peroxide, U2O7, or even U(OH)6,H2O2, just as the peroxide form lately discovered by Spring for tin perhaps contains Sn2O3,H2O2.[67]This view was communicated by me in 1870 to the Russian Chemical Society.[68]Dithionic acid, H2S2O6, is distinguished among the thionic acids as containing the least proportion of sulphur. It is also called hyposulphuric acid, because its supposed anhydride, S2O5, contains more O than sulphurous oxide, SO2or S2O4, and less than sulphuric anhydride, SO3or S2O6. Dithionic acid, discovered by Gay-Lussac and Welter, is known as a hydrate and as salts, but not as anhydride. The method for preparing dithionic acid usually employed is by the action of finely-powdered manganese dioxide on a solution of sulphurous anhydride. On shaking, the smell of the latter disappears, and the manganese salt of the acid in question passes into solution; MnO2+ 2SO2= MnS2O6. If the temperature be raised, the dithionate splits up into sulphurous anhydride and manganese sulphate, MnSO4. Generally owing to this a mixture of manganese sulphate and dithionate is obtained in the solution. They may be separated by mixing the solution of the manganese salts with a solution of barium hydroxide, when a precipitate of manganese hydroxide and barium sulphate is obtained. In this manner barium dithionate only is obtained in solution. It is purified by crystallisation, and separates as BaS2O6,2H2O; this is then dissolved in water, and decomposed with the requisite amount of sulphuric acid. Dithionic acid, H2S2O6, then remains in solution. By concentrating the resultant solution under the receiver of an air-pump it is possible to obtain a liquid of sp. gr. 1·347, but it still contains water, and on further evaporation the acid decomposes into sulphuric acid and sulphurous anhydride: H2S2O6= H2SO4+ SO2. The same decomposition takes place if the solution be slightly heated. Like all the thionic acids, dithionic acid is readily attacked by oxidising agents, and passes into sulphurous acid. No dithionate is able to withstand the action of heat, even when very slight, without giving off sulphurous anhydride: K2S2O6= K2SO4+ SO2. The alkali dithionates have a neutral reaction (which indicates the energetic nature of the acid) are soluble in water, and in this respect present a certain resemblance to the salts of nitric acid (their anhydrides are: N2O5and S2O5). Klüss (1888) described many of the salts of dithionic acid.Langlois, about 1840, obtained a peculiar thionic acid by heating a strong solution of acid potassium sulphite with flowers of sulphur to about 60°, until the disappearance of the yellow coloration first produced by the solution of the sulphur. On cooling, a portion of the sulphur was precipitated, and crystals of a salt oftrithionic acid, K2S3O6(partly mixed with potassium sulphate), separated out. Plessy afterwards showed that the action of sulphurous acid on a thiosulphate also gives sulphur and trithionic acid: 2K2S2O3+ 3SO2= 2K2S3O6+ S. A mixture of potassium acid sulphite and thiosulphate also gives a trithionate. It is very possible that a reaction of the same kind occurs in the formation of trithionic acid by Langloid's method, because potassium sulphite and sulphur yield potassium thiosulphate. The potassium thiosulphate may also be replaced by potassium sulphide, and on passing sulphurous anhydride through the solution thiosulphate is first formed and then trithionate: 4KHSO3+ K2S + 4SO2= 3K2S3O6+ 2H2O. The sodium salt is not formed under the same circumstances as the corresponding potassium salt. The sodium salt does not crystallise and is very unstable: the barium salt is, however, more stable. The barium and potassium salts are anhydrous, they give neutral solutions and decompose when ignited, with the evolution of sulphur and sulphurous anhydride, a sulphate being left behind, K2S3O6= K2SO4+ SO2+ S. If a solution of the potassium salt be decomposed by means of hydrofluosilicic or chloric acid, the insoluble salts of these acids are precipitated and trithionic acid is obtained in solution, which however very easily breaks up on concentration. The addition of salts of copper, mercury, silver, &c., to a solution of a trithionate is followed, either immediately or after a certain time, by the formation of a black precipitate of the sulphides whose formation is due to the decomposition of the trithionic acid with the transference of its sulphur to the metal.Tetrathionic acid, H2S4O6, in contradistinction to the preceding acids, is much more stable in the free state than in the form of salts. In the latter form it is easily converted into trithionate, with liberation of sulphur. Sodium tetrathionate was obtained by Fordos and Gélis, by the action of iodine on a solution of sodium thiosulphate. The reaction essentially consists in the iodine taking up half the sodium of the thiosulphate, inasmuch as the latter contains Na2S2O3, whilst the tetrathionate contains NaS2O3or Na2S4O6, so that the reaction is as follows: 2Na2S2O3+ I2= 2NaI + Na2S4O6. It is evident that tetrathionic acid stands to thiosulphuric acid in exactly the same relation as dithionic acid does to sulphurous acid; for the same amount of the other elements in dithionate, KSO3, and tetrathionate, KS2O3, there is half as much metal as in sulphite, K2SO3, and thiosulphate, K2S2O3. If in the above reaction the sodium thiosulphate be replaced by the lead salt PbS2O3, the sparingly-soluble lead iodide PbI2and the soluble salt PbS4O6are obtained. Moreover the lead salt easily gives tetrathionic acid itself (PbSO4is precipitated). The solution of tetrathionic acid may be evaporated over a water-bath, and afterwards in a vacuum, when it gives a colourless liquid, which has no smell and a very acid reaction. When dilute it may be heated to its boiling-point, but in a concentrated form it decomposes into sulphuric acid, sulphurous anhydride, and sulphur: H2S4O6= H2SO4+ SO2+ S2.Pentathionic acid, H2S5O6, also belongs to this series of acids. But little is known concerning it, either as hydrate or in salts. It is formed, together with tetrathionic acid, by the direct action of sulphurous acid on sulphuretted hydrogen in an aqueous solution; a large proportion of sulphur being precipitated at the same time: 5SO2+ 5H2S = H2S5O6+ 5S + 4H2O.If, as was shown above, the thionic acids are disulphonic acids, they may be obtained, like other sulphonic acids, by means of potassium sulphite and sulphur chloride. Thus Spring demonstrated the formation of potassium trithionate by the action of sulphur dichloride on a strong solution of potassium sulphite: 2KSO3K + SCl2= S(SO3K)2+ 2KCl. If sulphur chloride be taken, sulphur also is precipitated. The same trithionate is formed by heating a solution of double thiosulphates; for example, of AgKS2O3. Two molecules of the salts then form silver sulphide and potassium trithionate. If the thiosulphate be the potassium silver salt SO3K(AgS), then the structure of the trithionate must necessarily be (SO3K)2S. Previous to Spring's researches, the action of iodine on sodium thiosulphate was an isolated accidentally discovered reaction; he, however, showed its general significance by testing the action of iodine on mixtures of different sulphur compounds. Thus with iodine, I2, the mixture Na2S + Na2SO3forms 2NaI + Na2S2O3, whilst the mixture Na2S2O3+ Na2SO3+ I2gives 2NaI + Na2S3O6—that is, trithionic acid stands in the same relation to thiosulphuric acid as the latter does to sulphuretted hydrogen. We adopt the same mode of representation: by replacing one hydrogen in H2S by sulphuryl we obtain thiosulphuric acid, HSO3.HS, and by replacing a second hydrogen in the latter again by sulphuryl we obtain trithionic acid, (HSO3)2S. Furthermore, Spring showed that the action of sodium amalgam on the thionic acids causes reverse reactions to those above indicated for iodine. Thus sodium thiosulphate with Na2gives Na2S + Na2SO3, and Spring showed that the sodium here is not a simple element taking up sulphur, but itself enters into double decomposition, replacing sulphur; for on taking a potassium salt and acting on it with sodium, KSO3(SK) + NaNa = KSO3Na + (SK)Na. In a similar way sodium dithionate with sodium gives sodium sulphite: (NaSO3)2+ Na2= 2NaSO3Na; sodium trithionate forms NaSO3Na and NaSO3.SNa, and tetrathionate forms sodium thiosulphate, (NaSO3)S2(NaSO3) + Na2= 2(NaSO3)(NaS).In all the oxidised compounds of sulphur we may note the presence of the elements of sulphurous anhydride, SO2, the only product of the combustion of sulphur, and in this sense the compounds of sulphur containing one SO2are—fig_258_1while, according to this mode of representation, the thionic acids are—fig_258_2Hence it is evident that SO2has (whilst CO2has not) the faculty for combination, and aims at forming SO2X2. These X2can = O, and the question naturally suggests itself as to whether the O2which occurs in SO2is not of the same nature as this oxygen which adds itself to SO2—that is, whether SO2does not correspond with the more general type SX4, and its compounds with the type SX6? To this we may answer ‘Yes’ and ‘No’—‘Yes’ in the general sense which proceeds from the investigation of the majority of compounds, especially metals, where RO corresponds with RCl2, RX2; ‘No’ in the sense that sulphur does not give either SH4, SH6, or SCl6, and therefore the stages SX4and SX6are only observable in oxygen compounds. With reference to the type SX6a hydrate, S(HO)6, might be expected, if not SCl6. And we must recognise this hydrate from a study of the compounds of sulphuric acid with water. In addition to what has been already said respecting the complex acids formed by sulphur, I think it well to mention that, according to the above view, still more complex oxygen acids and salts of sulphur may be looked for. For instance, the salt Na2S4O8obtained by Villiers (1888) is of this kind. It is formed together with sodium trithionate and sulphur, when SO2is passed through a cold solution of Na2S2O3, which is then allowed to stand for several days at the ordinary temperature: 2Na2S2O3+ 4SO2= Na2S4O8+ Na2S3O6+ S. It may be assumed here, as in the thionic acids, that there are two sulphoxyls, bound together not only by S, but also by SO2, or what is almost the same thing, that the sulphoxyl is combined with the residue of trithionic acid,i.e.replaces one aqueous residue in trithionic acid.
[48]The greatest part of the sulphuric acid is used in the soda manufacture, in the conversion of the common salt into sulphate. For this purpose an acid having a density of 60° Baumé is amply sufficient. Chamber acid has a density up to 1·57 = 50° to 51° Baumé; it contains about 35 per cent. of water. About 15 per cent. of this water can be removed in leaden stills, and nearly all the remainder may be expelled in glass or platinum vessels. Acid of 66° Baumé, = 1·847, contains about 96 per cent. of the hydrate H2SO4. The density falls with a greater or less proportion of water, the maximum density corresponding with 97½ per cent. of the hydrate H2SO4. The concentration of H2SO4in platinum retorts has the disadvantage that sulphuric acid, upwards of 90 per cent. in strength, does corrode platinum, although but slightly (a few grams per tens of tons of acid). The retorts therefore require repairing, and the cost of the platinum exceeds the price obtained for concentrating the acid from 90 per cent. to 98 per cent. (in factories the acid is not concentrated beyond this by evaporation in the air). This inconvenience has lately (1891, by Mathey) been eliminated by coating the inside of the platinum retorts with a thin (0·1 to 0·02 mm.) layer of gold which is 40 times less corroded by sulphuric acid than platinum. Négrier (1890) carries on the distillation in porcelain dishes, Blond by heating a thin platinum wire immersed in the acid by means of an electric current, but the most promising method is that of Kessler (1891), which consists in passing hot air over sulphuric acid flowing in a thin stream in stone vessels, so that there is no boiling but only evaporation at moderate temperatures: the transference of the heat is direct (and not through the sides of the vessels), which economises the fuel and prevents the distilling vessels being damaged.When, by evaporation of the water, sulphuric acid attains a density of 66° Baumé (sp. gr. 1·84), it is impossible to concentrate it further, because it then distils over unchanged.The distillation of sulphuric acidis not generally carried on on a large scale, but forms a laboratory process, employed when particularly pure acid is required. The distillation is effected either in platinum retorts furnished with corresponding condensers and receivers, or in glass retorts. In the latter case, great caution is necessary, because the boiling of sulphuric acid itself is accompanied by still more violent jerks and greater irregularity than even the evaporation of the last portions of water contained in the acid. If the glass retort which holds the strong sulphuric acid to be distilled be heated directly from below, it frequently jerks and breaks. For greater safety the heating is not effected from below, but at the sides of the retort. The evaporation then does not proceed in the whole mass, but only from the upper portions of the liquid, and therefore goes on much more quietly. The acid may be made to boil quietly also by surrounding the retort with good conductors of heat—for example, iron filings, or by immersing a bunch of platinum wires in the acid, as the bubbles of sulphuric acid vapour then form on the extremities of the wires.
[48]The greatest part of the sulphuric acid is used in the soda manufacture, in the conversion of the common salt into sulphate. For this purpose an acid having a density of 60° Baumé is amply sufficient. Chamber acid has a density up to 1·57 = 50° to 51° Baumé; it contains about 35 per cent. of water. About 15 per cent. of this water can be removed in leaden stills, and nearly all the remainder may be expelled in glass or platinum vessels. Acid of 66° Baumé, = 1·847, contains about 96 per cent. of the hydrate H2SO4. The density falls with a greater or less proportion of water, the maximum density corresponding with 97½ per cent. of the hydrate H2SO4. The concentration of H2SO4in platinum retorts has the disadvantage that sulphuric acid, upwards of 90 per cent. in strength, does corrode platinum, although but slightly (a few grams per tens of tons of acid). The retorts therefore require repairing, and the cost of the platinum exceeds the price obtained for concentrating the acid from 90 per cent. to 98 per cent. (in factories the acid is not concentrated beyond this by evaporation in the air). This inconvenience has lately (1891, by Mathey) been eliminated by coating the inside of the platinum retorts with a thin (0·1 to 0·02 mm.) layer of gold which is 40 times less corroded by sulphuric acid than platinum. Négrier (1890) carries on the distillation in porcelain dishes, Blond by heating a thin platinum wire immersed in the acid by means of an electric current, but the most promising method is that of Kessler (1891), which consists in passing hot air over sulphuric acid flowing in a thin stream in stone vessels, so that there is no boiling but only evaporation at moderate temperatures: the transference of the heat is direct (and not through the sides of the vessels), which economises the fuel and prevents the distilling vessels being damaged.
When, by evaporation of the water, sulphuric acid attains a density of 66° Baumé (sp. gr. 1·84), it is impossible to concentrate it further, because it then distils over unchanged.The distillation of sulphuric acidis not generally carried on on a large scale, but forms a laboratory process, employed when particularly pure acid is required. The distillation is effected either in platinum retorts furnished with corresponding condensers and receivers, or in glass retorts. In the latter case, great caution is necessary, because the boiling of sulphuric acid itself is accompanied by still more violent jerks and greater irregularity than even the evaporation of the last portions of water contained in the acid. If the glass retort which holds the strong sulphuric acid to be distilled be heated directly from below, it frequently jerks and breaks. For greater safety the heating is not effected from below, but at the sides of the retort. The evaporation then does not proceed in the whole mass, but only from the upper portions of the liquid, and therefore goes on much more quietly. The acid may be made to boil quietly also by surrounding the retort with good conductors of heat—for example, iron filings, or by immersing a bunch of platinum wires in the acid, as the bubbles of sulphuric acid vapour then form on the extremities of the wires.
[49]Thus it appears that so common, and apparently so stable, a compound as sulphuric acid decomposes even at a low temperature with separation of the anhydride, but this decomposition is restricted by a limit, corresponding to the presence of about 1½ p.c. of water, or to a composition of nearly H2O,12H2SO4.Now there is no reason for thinking that this substance is a definite compound; it is an equilibrated system which does not decompose under ordinary circumstances below 338°. Dittmar carried on the distillation under pressures varying between 30 and 2,140 millimetres (of mercury), and he found that the composition of the residue hardly varies, and contains from 99·2 to 98·2 per cent. of the normal hydrate, although at 30 mm. the temperature of distillation is about 210° and at 2,140 mm. it is 382°. Furthermore, it is a fact of practical importance that under a pressure of two atmospheres the distillation of sulphuric acid proceeds very quietly.Sulphuric acid may bepurifiedfrom the majority of its impurities by distillation, if the first and last portions of the distillate be rejected. The first portions will contain the oxides of nitrogen, hydrochloric acid, &c., and the last portions the less volatile impurities. The oxides of nitrogen may be removed by heating the acid with charcoal, which converts them into volatile gases. Sulphuric acid may be freed from arsenic by heating it with manganese dioxide and then distilling. This oxidises all the arsenic into non-volatile arsenic acid. Without a preliminary oxidation it would partially remain as volatile arsenious acid, and might pass over into the distillate. The arsenic may also be driven off by first reducing it to arsenious acid, and then passing hydrochloric acid gas through the heated acid. It is then converted into arsenious chloride, which volatilises.
[49]Thus it appears that so common, and apparently so stable, a compound as sulphuric acid decomposes even at a low temperature with separation of the anhydride, but this decomposition is restricted by a limit, corresponding to the presence of about 1½ p.c. of water, or to a composition of nearly H2O,12H2SO4.
Now there is no reason for thinking that this substance is a definite compound; it is an equilibrated system which does not decompose under ordinary circumstances below 338°. Dittmar carried on the distillation under pressures varying between 30 and 2,140 millimetres (of mercury), and he found that the composition of the residue hardly varies, and contains from 99·2 to 98·2 per cent. of the normal hydrate, although at 30 mm. the temperature of distillation is about 210° and at 2,140 mm. it is 382°. Furthermore, it is a fact of practical importance that under a pressure of two atmospheres the distillation of sulphuric acid proceeds very quietly.
Sulphuric acid may bepurifiedfrom the majority of its impurities by distillation, if the first and last portions of the distillate be rejected. The first portions will contain the oxides of nitrogen, hydrochloric acid, &c., and the last portions the less volatile impurities. The oxides of nitrogen may be removed by heating the acid with charcoal, which converts them into volatile gases. Sulphuric acid may be freed from arsenic by heating it with manganese dioxide and then distilling. This oxidises all the arsenic into non-volatile arsenic acid. Without a preliminary oxidation it would partially remain as volatile arsenious acid, and might pass over into the distillate. The arsenic may also be driven off by first reducing it to arsenious acid, and then passing hydrochloric acid gas through the heated acid. It is then converted into arsenious chloride, which volatilises.
[50]The amount of heat developed by the mixture of sulphuric acid with water is expressed in the diagram on p.77, Volume I., by the middle curve, whose abscissæ are the percentage amounts of acid (H2SO4) in the resultant solution, and ordinates the number of units of heat corresponding with the formation of 100 cubic centimetres of the solution (at 18°). The calculations on which the curve is designed are based on Thomsen's determinations, which show that 98 grams or a molecular amount of sulphuric acid, in combining withmmolecules of water (that is, withm=18 grams of water), develop the following number of units of heat, R:—m=123591949100200R =6379941811137131081495216256166841685917066c=0·4320·4700·5000·5760·7010·8210·9140·9540·975T =127°149°146°121°82°45°19°9°5°cstands for the specific heat of H2SO4mH2O (according to Marignac and Pfaundler), and T for the rise in temperature which proceeds from the mixture of H2SO4withmH2O. The diagram shows that contraction and rise of temperature proceed almost parallel with each other.
[50]The amount of heat developed by the mixture of sulphuric acid with water is expressed in the diagram on p.77, Volume I., by the middle curve, whose abscissæ are the percentage amounts of acid (H2SO4) in the resultant solution, and ordinates the number of units of heat corresponding with the formation of 100 cubic centimetres of the solution (at 18°). The calculations on which the curve is designed are based on Thomsen's determinations, which show that 98 grams or a molecular amount of sulphuric acid, in combining withmmolecules of water (that is, withm=18 grams of water), develop the following number of units of heat, R:—
cstands for the specific heat of H2SO4mH2O (according to Marignac and Pfaundler), and T for the rise in temperature which proceeds from the mixture of H2SO4withmH2O. The diagram shows that contraction and rise of temperature proceed almost parallel with each other.
[50 bis]Pickering (1890) showed (a) that dilute solutions of sulphuric acid containing up to H2SO4+ 10H2O deposit ice (at -0°·12 when there is 2,000H2O per H2SO4, at -0°·23 when there is 1,000H2O, at -1°·04 when there is 200H2O, at -2°·12 when there is 100H2O, at -4°·5 when there is 50H2O, at -15°·7 when there is 20H2O, and at -61° when the composition of the solution is H2SO4+ 10H2O); (b) that for higher concentrations crystals separate out at a considerable degree of cold, having the composition H2SO44H2O, which melt at -24°·5, and if either water or H2SO4be added to this compound the temperature of crystallisation falls, so that a solution of the composition 12H2SO4+ 100H2O gives crystals of the above hydrate at -70°, 15H2SO4+ 100H2O at -47°, 30H2SO4+ 100H2O at -32°, 40H2SO4+ 100H2O at -52°; (c) that if the amount of H2SO4be still greater, then a hydrate H2SO4H2O separates out and melts at +8°·5, while the addition of water or sulphuric acid to it lowers the temperature of crystallisation so that the crystallisation of H2SO4H2O from a solution of the composition H2SO4+ 1·73H2O takes place at -22°, H2SO4+ 1·5H2O at -6°·5, H2SO4+ 1·2H2O at +3°·7, H2SO4+ 0·75H2O at +2°·8, H2SO4+ 0·5H2O at -16°; (d) that when there is less than 40H2O per 100H2SO4, refrigeration separates out the normal hydrate H2SO4, which melts at +10°·35, and that a solution of the composition H2SO4+ 0·35H2O deposits crystals of this hydrate at -34°, H2SO4+ 0·1H2O at -4°·1, H2SO4+ 0·05H2O at +4°·9, while fuming acid of the composition H2SO4+ 0·05SO3deposits H2SO4at about +7°. Thus the temperature of the separation of crystals clearly distinguishes the above four regions of solutions, and in the space between H2SO4+ H2O and +25H2O a particular hydrate H2SO44H2O separates out, discovered by Pickering, the isolation of which deserves full attention and further research. I may add here that the existence of a hydrate H2SO44H2O was pointed out in my work,The Investigation of Aqueous Solutions, p. 120 (1887), upon the basis that it has at all temperatures a smaller value for the coefficient of expansionkin the formula St= S0/(1 -kt) than the adjacent (in composition) solutions of sulphuric acid. And for solutions approximating to H2SO410H2O in their composition,kis constant at all temperatures (for more dilute solutions the value ofkincreases withtand for more concentrated solutions it decreases). This solution (with 10H2O) forms the point of transition between more dilute solutions which deposit ice (water) when refrigerated and those which give crystals of H2SO44H2O. According to R. Pictet (1894) the solution H2SO410H2O freezes at -88° (but no reference is made as to what separates out),i.e.at a lower temperature than all the other solutions of sulphuric acid. However, in respect to these last researches of R. Pictet (for 88·88 p.c. H2SO4-55°, for H2SO4H2O +3·5°, for H2SO42H2O -70°, for H2SO44H2O -40°, &c.) it should be remarked that they offer some quite improbable data; for example, for H2SO475H2O they give the freezing point as 0°, for H2SO4300H2O +4°·5, and even for H2SO41000H2O +0°·5, although it is well known that a small amount of sulphuric acid lowers the temperature of the formation of ice. I have found by direct experiment that a frozen solidified solution of H2SO4+ 300H2O melted completely at 0°.
[50 bis]Pickering (1890) showed (a) that dilute solutions of sulphuric acid containing up to H2SO4+ 10H2O deposit ice (at -0°·12 when there is 2,000H2O per H2SO4, at -0°·23 when there is 1,000H2O, at -1°·04 when there is 200H2O, at -2°·12 when there is 100H2O, at -4°·5 when there is 50H2O, at -15°·7 when there is 20H2O, and at -61° when the composition of the solution is H2SO4+ 10H2O); (b) that for higher concentrations crystals separate out at a considerable degree of cold, having the composition H2SO44H2O, which melt at -24°·5, and if either water or H2SO4be added to this compound the temperature of crystallisation falls, so that a solution of the composition 12H2SO4+ 100H2O gives crystals of the above hydrate at -70°, 15H2SO4+ 100H2O at -47°, 30H2SO4+ 100H2O at -32°, 40H2SO4+ 100H2O at -52°; (c) that if the amount of H2SO4be still greater, then a hydrate H2SO4H2O separates out and melts at +8°·5, while the addition of water or sulphuric acid to it lowers the temperature of crystallisation so that the crystallisation of H2SO4H2O from a solution of the composition H2SO4+ 1·73H2O takes place at -22°, H2SO4+ 1·5H2O at -6°·5, H2SO4+ 1·2H2O at +3°·7, H2SO4+ 0·75H2O at +2°·8, H2SO4+ 0·5H2O at -16°; (d) that when there is less than 40H2O per 100H2SO4, refrigeration separates out the normal hydrate H2SO4, which melts at +10°·35, and that a solution of the composition H2SO4+ 0·35H2O deposits crystals of this hydrate at -34°, H2SO4+ 0·1H2O at -4°·1, H2SO4+ 0·05H2O at +4°·9, while fuming acid of the composition H2SO4+ 0·05SO3deposits H2SO4at about +7°. Thus the temperature of the separation of crystals clearly distinguishes the above four regions of solutions, and in the space between H2SO4+ H2O and +25H2O a particular hydrate H2SO44H2O separates out, discovered by Pickering, the isolation of which deserves full attention and further research. I may add here that the existence of a hydrate H2SO44H2O was pointed out in my work,The Investigation of Aqueous Solutions, p. 120 (1887), upon the basis that it has at all temperatures a smaller value for the coefficient of expansionkin the formula St= S0/(1 -kt) than the adjacent (in composition) solutions of sulphuric acid. And for solutions approximating to H2SO410H2O in their composition,kis constant at all temperatures (for more dilute solutions the value ofkincreases withtand for more concentrated solutions it decreases). This solution (with 10H2O) forms the point of transition between more dilute solutions which deposit ice (water) when refrigerated and those which give crystals of H2SO44H2O. According to R. Pictet (1894) the solution H2SO410H2O freezes at -88° (but no reference is made as to what separates out),i.e.at a lower temperature than all the other solutions of sulphuric acid. However, in respect to these last researches of R. Pictet (for 88·88 p.c. H2SO4-55°, for H2SO4H2O +3·5°, for H2SO42H2O -70°, for H2SO44H2O -40°, &c.) it should be remarked that they offer some quite improbable data; for example, for H2SO475H2O they give the freezing point as 0°, for H2SO4300H2O +4°·5, and even for H2SO41000H2O +0°·5, although it is well known that a small amount of sulphuric acid lowers the temperature of the formation of ice. I have found by direct experiment that a frozen solidified solution of H2SO4+ 300H2O melted completely at 0°.
[51]With an excess of snow, the hydrate H2SO4,H2O, like the normal hydrate, gives a freezing mixture, owing to the absorption of a large amount of heat (the latent heat of fusion). In melting, the molecule H2SO4absorbs 960 heat units, and the molecule H2SO4H2O 3,680 heat units. If therefore we mix one gram molecule of this hydrate with seventeen gram molecules of snow, there is an absorption of 18,080 heat units, because 17H2O absorbs 17 × 1,430 heat units, and the combination of the monohydrate with water evolves 9,800 heat units. As the specific heat of the resultant compound H2SO4,18H2O = 0·813, the fall of temperature will be -52°·6. And, in fact, a very low temperature may be obtained by means of sulphuric acid.
[51]With an excess of snow, the hydrate H2SO4,H2O, like the normal hydrate, gives a freezing mixture, owing to the absorption of a large amount of heat (the latent heat of fusion). In melting, the molecule H2SO4absorbs 960 heat units, and the molecule H2SO4H2O 3,680 heat units. If therefore we mix one gram molecule of this hydrate with seventeen gram molecules of snow, there is an absorption of 18,080 heat units, because 17H2O absorbs 17 × 1,430 heat units, and the combination of the monohydrate with water evolves 9,800 heat units. As the specific heat of the resultant compound H2SO4,18H2O = 0·813, the fall of temperature will be -52°·6. And, in fact, a very low temperature may be obtained by means of sulphuric acid.
[52]For example, if it be taken that at 19° the sp. gr. of pure sulphuric acid is 1·8330, then at 20° it is 1·8330 - (20 - 19)10·13 = 1·8320.
[52]For example, if it be taken that at 19° the sp. gr. of pure sulphuric acid is 1·8330, then at 20° it is 1·8330 - (20 - 19)10·13 = 1·8320.
[53]Unfortunately, notwithstanding the great number of fragmentary and systematic researches which have been made (by Parks, Ure, Bineau, Kolbe, Lunge, Marignac, Kremers, Thomsen, Perkin, and others) for determining the relation between the sp. gr. and composition of solutions of sulphuric acid, they contain discrepancies which amount to, and even exceed, 0·002 in the sp. gr. For instance, at 15°·4 the solution of composition H2SO43H2O has a sp. gr. 1·5493 according to Perkin (1886), 1·5501 according to Pickering (1890), and 1·5525 according to Lunge (1890). The cause of these discrepancies must be looked for in the methods employed for determining the composition of the solutions—i.e.in the inaccuracy with which the percentage amount of H2SO4is determined, for a difference of 1 p.c. corresponds to a difference of from 0·0070 (for very weak solutions) to 0·0118 (for a solution containing about 73 p.c.) in the specific gravity (that is the factords/dp) at 15°. As it is possible to determine the specific gravity with an accuracy even exceeding 0·0002, the specific gravities given in the adjoining tables are only averages and most probable data in which the error, especially for the 30–80 p.c. solutions cannot be less than 0·0010 (taking water at 4° as 1).
[53]Unfortunately, notwithstanding the great number of fragmentary and systematic researches which have been made (by Parks, Ure, Bineau, Kolbe, Lunge, Marignac, Kremers, Thomsen, Perkin, and others) for determining the relation between the sp. gr. and composition of solutions of sulphuric acid, they contain discrepancies which amount to, and even exceed, 0·002 in the sp. gr. For instance, at 15°·4 the solution of composition H2SO43H2O has a sp. gr. 1·5493 according to Perkin (1886), 1·5501 according to Pickering (1890), and 1·5525 according to Lunge (1890). The cause of these discrepancies must be looked for in the methods employed for determining the composition of the solutions—i.e.in the inaccuracy with which the percentage amount of H2SO4is determined, for a difference of 1 p.c. corresponds to a difference of from 0·0070 (for very weak solutions) to 0·0118 (for a solution containing about 73 p.c.) in the specific gravity (that is the factords/dp) at 15°. As it is possible to determine the specific gravity with an accuracy even exceeding 0·0002, the specific gravities given in the adjoining tables are only averages and most probable data in which the error, especially for the 30–80 p.c. solutions cannot be less than 0·0010 (taking water at 4° as 1).
[53 bis]Judging from the best existing determinations (of Marignac, Kremers, and Pickering) for solutions of sulphuric acid (especially those containing more than 5 p.c. H2SO4) within the limits of 0° and 30° (and even to 40°), the variation of the sp. gr. with the temperaturetmay (within the accuracy of the existing determinations) be perfectly expressed by the equation St= S0+ At+ Bt2. It must be added that (1) three specific gravities fully determine the variation of the density witht; (2)ds/dt= A + 2Bt—i.e.the factor of the temperature is expressed by a straight line; (3) the value of A (ifpbe greater than 5 p.c.) is negative, and numerically much greater than B; (4) the value of B for dilute solutions containing less than 25 p.c. is negative; for solutions approximating to H2SO43H2O in their composition it is equal to 0, and for solutions of greater concentration B is positive; (5) the factords/dpfor all temperatures attains a maximum value about H2SO4H2O; (6) on dividingds/dtby S0, and so obtaining the coefficient of expansionk(seeNote53), a minimum is obtained near H2SO4and H2SO44H2O, and a maximum at H2SO4H2O for all temperatures.
[53 bis]Judging from the best existing determinations (of Marignac, Kremers, and Pickering) for solutions of sulphuric acid (especially those containing more than 5 p.c. H2SO4) within the limits of 0° and 30° (and even to 40°), the variation of the sp. gr. with the temperaturetmay (within the accuracy of the existing determinations) be perfectly expressed by the equation St= S0+ At+ Bt2. It must be added that (1) three specific gravities fully determine the variation of the density witht; (2)ds/dt= A + 2Bt—i.e.the factor of the temperature is expressed by a straight line; (3) the value of A (ifpbe greater than 5 p.c.) is negative, and numerically much greater than B; (4) the value of B for dilute solutions containing less than 25 p.c. is negative; for solutions approximating to H2SO43H2O in their composition it is equal to 0, and for solutions of greater concentration B is positive; (5) the factords/dpfor all temperatures attains a maximum value about H2SO4H2O; (6) on dividingds/dtby S0, and so obtaining the coefficient of expansionk(seeNote53), a minimum is obtained near H2SO4and H2SO44H2O, and a maximum at H2SO4H2O for all temperatures.
[53 tri]These data (as well as those in the following table) have been recalculated by me chiefly upon the basis of Kremer's, Pickering's, Perkin's, and my own determinations; all the requisite corrections have been introduced, and I have reason for thinking that in each of them the probable error (or difference from the true figures, now unknown) of the specific gravity does not exceed ±0·0007 (if water at 4° = 1) for the 25–80 p.c. solutions, and ±0·0002 for the more dilute or concentrated solutions.
[53 tri]These data (as well as those in the following table) have been recalculated by me chiefly upon the basis of Kremer's, Pickering's, Perkin's, and my own determinations; all the requisite corrections have been introduced, and I have reason for thinking that in each of them the probable error (or difference from the true figures, now unknown) of the specific gravity does not exceed ±0·0007 (if water at 4° = 1) for the 25–80 p.c. solutions, and ±0·0002 for the more dilute or concentrated solutions.
[54]The factordS/dppasses through 0, that is, the specific gravity attains a maximum value at about 98 p.c. This was discovered by Kohlrausch, and confirmed by Chertel, Pickering, and others.
[54]The factordS/dppasses through 0, that is, the specific gravity attains a maximum value at about 98 p.c. This was discovered by Kohlrausch, and confirmed by Chertel, Pickering, and others.
[55]Naturally under the condition that there is no other ingredient besides water, which is sufficiently true. For commercial acid, whose specific gravity is usually expressed in degrees of Baumé's hydrometer, we may add that at 15°Specific gravity11·11·21·31·41·51·61·71·8Degree Baumé0132433·341·248·154·159·564·266° Baumé (the strongest commercial acid or oil of vitriol) corresponds to a sp. gr. 1·84.By employing the second table (by the method of interpolation) the specific gravity, at a given temperature (from 0° to 30°) can be found for any percentage amount of H2SO4, and therefore conversely the percentage of H2SO4can be found from the specific gravity.
[55]Naturally under the condition that there is no other ingredient besides water, which is sufficiently true. For commercial acid, whose specific gravity is usually expressed in degrees of Baumé's hydrometer, we may add that at 15°
66° Baumé (the strongest commercial acid or oil of vitriol) corresponds to a sp. gr. 1·84.
By employing the second table (by the method of interpolation) the specific gravity, at a given temperature (from 0° to 30°) can be found for any percentage amount of H2SO4, and therefore conversely the percentage of H2SO4can be found from the specific gravity.
[55 bis]Whether similar (even small) breaks in the continuity of the factordS/dpexist or not, for other hydrates (for instance, for H2SO4H2O and H2SO44H2O) cannot as yet be affirmed owing to the want of accurate data (Note53). In my investigation of this subject (1887) I admit their possibility, but only conditionally; and now, without insisting upon a similar opinion, I only hold to the existence of a distinct break in the factor at H2SO4, being guided by C. Winkler's observations ond the specific gravities of fuming sulphuric acid.
[55 bis]Whether similar (even small) breaks in the continuity of the factordS/dpexist or not, for other hydrates (for instance, for H2SO4H2O and H2SO44H2O) cannot as yet be affirmed owing to the want of accurate data (Note53). In my investigation of this subject (1887) I admit their possibility, but only conditionally; and now, without insisting upon a similar opinion, I only hold to the existence of a distinct break in the factor at H2SO4, being guided by C. Winkler's observations ond the specific gravities of fuming sulphuric acid.
[56]In 1887, on considering all the existent observations for a temperature 0°, I gave the accompanying scheme (p.243) of the variation of the factords/dpat 0°.I did not then (1887) give this scheme an absolute value, and now after the appearance of two series of new determinations (Lunge and Pickering in 1890), which disagree in many points, I think it well to state quite clearly: (1) that Lunge's and Pickering's new determinations have not added to the accuracy of our data respecting the variation of the specific gravity of solutions of sulphuric acid; (2) that the sum total of existing data does not negative (within the limit of experimental accuracy) the possibility of a rectilinear and broken form for the factorsds/dp; (3) that the supposition of ‘special points’ inds/dp, indicating definite hydrates, finds confirmation in all the latest determinations; (4) that the supposition respecting the existence of hydrates determining a break of the factords/dpis in in way altered if, instead of a series of broken straight lines, there be a continuous series of curves, nearly approaching straight lines; and (5) that this subject deserves (as I mentioned in 1887) new and careful elaboration, because it concerns that foremost problem in our science—solutions—and introduces a special method into it—that is, the study of differential variations in a property which is so easily observed as the specific gravity of a liquid.
[56]In 1887, on considering all the existent observations for a temperature 0°, I gave the accompanying scheme (p.243) of the variation of the factords/dpat 0°.
I did not then (1887) give this scheme an absolute value, and now after the appearance of two series of new determinations (Lunge and Pickering in 1890), which disagree in many points, I think it well to state quite clearly: (1) that Lunge's and Pickering's new determinations have not added to the accuracy of our data respecting the variation of the specific gravity of solutions of sulphuric acid; (2) that the sum total of existing data does not negative (within the limit of experimental accuracy) the possibility of a rectilinear and broken form for the factorsds/dp; (3) that the supposition of ‘special points’ inds/dp, indicating definite hydrates, finds confirmation in all the latest determinations; (4) that the supposition respecting the existence of hydrates determining a break of the factords/dpis in in way altered if, instead of a series of broken straight lines, there be a continuous series of curves, nearly approaching straight lines; and (5) that this subject deserves (as I mentioned in 1887) new and careful elaboration, because it concerns that foremost problem in our science—solutions—and introduces a special method into it—that is, the study of differential variations in a property which is so easily observed as the specific gravity of a liquid.
[56 bis]These hydrates are: (a) H2SO4= SO3H2O (melts at + 10°·4); (b) H2SO4H2O = SO32H2O (crystallo-hydrate, melts at +8°·5); (c) H2SO42H2O (is apparently not crystallisable); (d) one of the hydrates between H2SO46H2O and H2SO43H2O, most probably H2SO44H2O = SO35H2O, for it crystallises at -24°·5 (Note50 bis); and (e) a certain hydrate with a large proportion of water, about H2SO4150H2O. The existence of the last is inferred from the fact that the factords/dpfirst falls, starting from water, and then rises, and this change takes place whenpis less than 5 p.c. Certainly a change in the variation ofds/dpords/dtdoes take place in the neighbourhood of these five hydrates (Pickering, 1890, recognised a far greater number of hydrates). I think it well to add that if the composition of the solutions be expressed by the percentage amount of molecules—r1SO3+ (100 -r1)H2O we find that for H2SO4,r1= 50, for H2SO42H2Or1= 25 = 50/2, for H2SO4H2O,r1= 33·333 = 50·⅔, while for H2SO44H2O,r1= 16·666 = 50·⅓—i.e.that the chief hydrates are distributed symmetrically between H2O and H2SO4. Besides which I may mention that my researches (1887) upon the abrupt changes in the factor for solutions of sulphuric acid, and upon the correspondence of the breaks ofds/dpwith definite hydrates, received an indirect confirmation not only in the solutions of HNO3, HCl, C2H6O, C3H8O, &c., which I investigated (in my work cited in Chapter I., Note19), but also in the careful observations made by Professor Cheltzoff on the solutions of FeCl3and ZnCl2(Chapter XVI., Note4) which showed the existence in these solutions of an almost similar change inds/dpas is found in sulphuric acid. The detailed researches (1893) made by Tourbaba on the solutions of many organic substances are of a similar nature. Besides which, H. Crompton (1888), in his researches on the electrical conductivity of solutions of sulphuric acid, and Tammann, in his observations on their vapour tension, found a correlation with the hydrates indicated as above by the investigation of their specific gravities. The influence of mixtures of a definite composition upon the chemical relations of solutions is even exhibited in such a complex process as electrolysis. V. Kouriloff (1891) showed that mixtures containing about 3 p.c., 47 p.c. and 73 p.c. of sulphuric acid—i.e.whose composition approaches that of the hydrates H2SO4150H2O, H2SO46H2O and H2SO42H2O—exhibit certain peculiarities in respect to the amount of peroxide of hydrogen formed during electrolysis. Thus a 3 p.c. solution gives a maximum amount of peroxide of hydrogen at the negative pole, as compared with that given by other neighbouring concentrations. Starting from 3 p.c., the formation of peroxide of hydrogen ceases until a concentration of 47 p.c. is reached.
[56 bis]These hydrates are: (a) H2SO4= SO3H2O (melts at + 10°·4); (b) H2SO4H2O = SO32H2O (crystallo-hydrate, melts at +8°·5); (c) H2SO42H2O (is apparently not crystallisable); (d) one of the hydrates between H2SO46H2O and H2SO43H2O, most probably H2SO44H2O = SO35H2O, for it crystallises at -24°·5 (Note50 bis); and (e) a certain hydrate with a large proportion of water, about H2SO4150H2O. The existence of the last is inferred from the fact that the factords/dpfirst falls, starting from water, and then rises, and this change takes place whenpis less than 5 p.c. Certainly a change in the variation ofds/dpords/dtdoes take place in the neighbourhood of these five hydrates (Pickering, 1890, recognised a far greater number of hydrates). I think it well to add that if the composition of the solutions be expressed by the percentage amount of molecules—r1SO3+ (100 -r1)H2O we find that for H2SO4,r1= 50, for H2SO42H2Or1= 25 = 50/2, for H2SO4H2O,r1= 33·333 = 50·⅔, while for H2SO44H2O,r1= 16·666 = 50·⅓—i.e.that the chief hydrates are distributed symmetrically between H2O and H2SO4. Besides which I may mention that my researches (1887) upon the abrupt changes in the factor for solutions of sulphuric acid, and upon the correspondence of the breaks ofds/dpwith definite hydrates, received an indirect confirmation not only in the solutions of HNO3, HCl, C2H6O, C3H8O, &c., which I investigated (in my work cited in Chapter I., Note19), but also in the careful observations made by Professor Cheltzoff on the solutions of FeCl3and ZnCl2(Chapter XVI., Note4) which showed the existence in these solutions of an almost similar change inds/dpas is found in sulphuric acid. The detailed researches (1893) made by Tourbaba on the solutions of many organic substances are of a similar nature. Besides which, H. Crompton (1888), in his researches on the electrical conductivity of solutions of sulphuric acid, and Tammann, in his observations on their vapour tension, found a correlation with the hydrates indicated as above by the investigation of their specific gravities. The influence of mixtures of a definite composition upon the chemical relations of solutions is even exhibited in such a complex process as electrolysis. V. Kouriloff (1891) showed that mixtures containing about 3 p.c., 47 p.c. and 73 p.c. of sulphuric acid—i.e.whose composition approaches that of the hydrates H2SO4150H2O, H2SO46H2O and H2SO42H2O—exhibit certain peculiarities in respect to the amount of peroxide of hydrogen formed during electrolysis. Thus a 3 p.c. solution gives a maximum amount of peroxide of hydrogen at the negative pole, as compared with that given by other neighbouring concentrations. Starting from 3 p.c., the formation of peroxide of hydrogen ceases until a concentration of 47 p.c. is reached.
[57]Cellulose, for instance unsized paper or calico, is dissolved by strong sulphuric acid. Acid diluted with about half its volume of water converts it (if the action be of short duration) into vegetable parchment (Chapter I., Note18). The action of dilute solutions of sulphuric acid converts it into hydro-cellulose, and the fibre loses its coherent quality and becomes brittle. The prolonged action of strong sulphuric acid chars the cellulose while dilute acid converts it into glucose. If sulphuric acid be kept in an open vessel, the organic matter of the dust held in the atmosphere falls into it and blackens the acid. The same thing happens if sulphuric acid be kept in a bottle closed by a cork; the cork becomes charred, and the acid turns black. However, the chemical properties of the acid undergo only a very slight change when it turns black. Sulphuric acid which is considerably diluted with water does not produce the above effects, which clearly shows their dependence on the affinity of the sulphuric acid for water. It is evident from the preceding that strong sulphuric acid will act as a powerful poison; whilst, on the other hand, when very dilute it is employed in certain medicines and as a fertiliser for plants.
[57]Cellulose, for instance unsized paper or calico, is dissolved by strong sulphuric acid. Acid diluted with about half its volume of water converts it (if the action be of short duration) into vegetable parchment (Chapter I., Note18). The action of dilute solutions of sulphuric acid converts it into hydro-cellulose, and the fibre loses its coherent quality and becomes brittle. The prolonged action of strong sulphuric acid chars the cellulose while dilute acid converts it into glucose. If sulphuric acid be kept in an open vessel, the organic matter of the dust held in the atmosphere falls into it and blackens the acid. The same thing happens if sulphuric acid be kept in a bottle closed by a cork; the cork becomes charred, and the acid turns black. However, the chemical properties of the acid undergo only a very slight change when it turns black. Sulphuric acid which is considerably diluted with water does not produce the above effects, which clearly shows their dependence on the affinity of the sulphuric acid for water. It is evident from the preceding that strong sulphuric acid will act as a powerful poison; whilst, on the other hand, when very dilute it is employed in certain medicines and as a fertiliser for plants.
[58]Weber (1884) obtained a series of salts R2O,8SO3nH2O for K, Rb, Cs, and Tl.
[58]Weber (1884) obtained a series of salts R2O,8SO3nH2O for K, Rb, Cs, and Tl.
[58 bis]Ditte (1890) divides all the metals into two groups with respect to sulphuric acid; the first group includes silver, mercury, copper, lead, and bismuth, which are only acted upon by hot concentrated acid. In this case sulphurous anhydride is evolved without any by-reactions. The second group contains manganese, nickel, cobalt, iron, zinc, cadmium, aluminium, tin, thallium, and the alkali metals. They react with sulphuric acid of any concentration at any temperature. At a low temperature hydrogen is disengaged, and at higher temperatures (and with very concentrated acid) hydrogen and sulphurous anhydride are simultaneously evolved.
[58 bis]Ditte (1890) divides all the metals into two groups with respect to sulphuric acid; the first group includes silver, mercury, copper, lead, and bismuth, which are only acted upon by hot concentrated acid. In this case sulphurous anhydride is evolved without any by-reactions. The second group contains manganese, nickel, cobalt, iron, zinc, cadmium, aluminium, tin, thallium, and the alkali metals. They react with sulphuric acid of any concentration at any temperature. At a low temperature hydrogen is disengaged, and at higher temperatures (and with very concentrated acid) hydrogen and sulphurous anhydride are simultaneously evolved.
[59]For example, the action of hot sulphuric acid on nitrogenous compounds, as applied in Kjeldahl's method for the estimation of nitrogen (Volume I. p.249). It is obvious that when sulphuric acid acts as an oxidising agent it forms sulphurous anhydride.The action of sulphuric acid on the alcohols is exactly similar to its action on alkalis, because the alcohols, like alkalis, react on acids; a molecule of alcohol with a molecule of sulphuric acid separates water and forms anacidethereal salt—that is there is produced an ethereal compound corresponding with acid salts. Thus, for example, the action of sulphuric acid, H2SO4, on ordinary alcohol, C2H5OH, gives water and sulphovinic acid, C2H5HSO4—that is, sulphuric acid in which one atom of hydrogen is replaced by the radicle C2H5of ethyl alcohol, SO2(OH)(OC2H5), or, what is the same thing, the hydrogen in alcohol is replaced by the radicle (sulphoxyl) of sulphuric acid, C2H5O.SO2(OH).
[59]For example, the action of hot sulphuric acid on nitrogenous compounds, as applied in Kjeldahl's method for the estimation of nitrogen (Volume I. p.249). It is obvious that when sulphuric acid acts as an oxidising agent it forms sulphurous anhydride.
The action of sulphuric acid on the alcohols is exactly similar to its action on alkalis, because the alcohols, like alkalis, react on acids; a molecule of alcohol with a molecule of sulphuric acid separates water and forms anacidethereal salt—that is there is produced an ethereal compound corresponding with acid salts. Thus, for example, the action of sulphuric acid, H2SO4, on ordinary alcohol, C2H5OH, gives water and sulphovinic acid, C2H5HSO4—that is, sulphuric acid in which one atom of hydrogen is replaced by the radicle C2H5of ethyl alcohol, SO2(OH)(OC2H5), or, what is the same thing, the hydrogen in alcohol is replaced by the radicle (sulphoxyl) of sulphuric acid, C2H5O.SO2(OH).
[60]We will mention the following difference between the sulphonic acids and the ethereal acid sulphates (Note59): the former re-form sulphuric acid with difficulty and the latter easily. Thus sulphovinic acid when heated with an excess of water is reconverted into alcohol and sulphuric acid. This is explained in the following manner. Both these classes of acids are produced by the substitution of hydrogen by SO3H, or the univalent radicle of sulphuric acid, but in the formation of ethereal acid sulphates the SO3H replaces the hydrogen of the hydroxyl in the alcohol, whilst in the formation of the sulphonic acids the SO3H replaces the hydrogen of a hydrocarbon. This difference is clearly evidenced in the existence of two acids of the composition SO4C2H6. The one, mentioned above, is sulphovinic acid or alcohol, C2H5.OH, in which the hydrogen of the hydroxyl is replaced by sulphoxyl = C2H5.OSO3H, whilst the other is alcohol, in which one atom of the hydrogen in ethyl, C2H5, is replaced by the sulphonic group—that is = (C2H4)SO3H·OH. The latter is called isethionic acid. It is more stable than sulphovinic acid. The details as to these interesting compounds must be looked for in works on organic chemistry, but I think it necessary to note one of the general methods of formation of these acids. The sulphites of the alkalis—for example, K2SO3—when heated with the halogen products of metalepsis, give a halogen salt and a salt of a sulphonic acid. Thus methyl iodide, CH3I, derived from marsh gas, CH4, when heated to 100° with a solution of potassium sulphite, K2SO3, gives potassium iodide, KI, and potassium methylsulphonate, CH3SO3K—that is a salt of the sulphonic acid. This shows that the sulphonic acid may be referred to sulphurous acid, and that there is a resemblance between sulphuric and sulphurous acid, which clearly reveals itself here in the formation of one product from them both.
[60]We will mention the following difference between the sulphonic acids and the ethereal acid sulphates (Note59): the former re-form sulphuric acid with difficulty and the latter easily. Thus sulphovinic acid when heated with an excess of water is reconverted into alcohol and sulphuric acid. This is explained in the following manner. Both these classes of acids are produced by the substitution of hydrogen by SO3H, or the univalent radicle of sulphuric acid, but in the formation of ethereal acid sulphates the SO3H replaces the hydrogen of the hydroxyl in the alcohol, whilst in the formation of the sulphonic acids the SO3H replaces the hydrogen of a hydrocarbon. This difference is clearly evidenced in the existence of two acids of the composition SO4C2H6. The one, mentioned above, is sulphovinic acid or alcohol, C2H5.OH, in which the hydrogen of the hydroxyl is replaced by sulphoxyl = C2H5.OSO3H, whilst the other is alcohol, in which one atom of the hydrogen in ethyl, C2H5, is replaced by the sulphonic group—that is = (C2H4)SO3H·OH. The latter is called isethionic acid. It is more stable than sulphovinic acid. The details as to these interesting compounds must be looked for in works on organic chemistry, but I think it necessary to note one of the general methods of formation of these acids. The sulphites of the alkalis—for example, K2SO3—when heated with the halogen products of metalepsis, give a halogen salt and a salt of a sulphonic acid. Thus methyl iodide, CH3I, derived from marsh gas, CH4, when heated to 100° with a solution of potassium sulphite, K2SO3, gives potassium iodide, KI, and potassium methylsulphonate, CH3SO3K—that is a salt of the sulphonic acid. This shows that the sulphonic acid may be referred to sulphurous acid, and that there is a resemblance between sulphuric and sulphurous acid, which clearly reveals itself here in the formation of one product from them both.
[61]The reaction BaO + O develops 12,000 heat units, whilst the reaction H2O + O absorbs 21,000 heat units.
[61]The reaction BaO + O develops 12,000 heat units, whilst the reaction H2O + O absorbs 21,000 heat units.
[62]Schöne obtained a compound of peroxide of barium with peroxide of hydrogen. If barium peroxide be dissolved in hydrochloric (or acetic) acid, or if a solution of hydrogen peroxide be diluted with a solution of barium hydroxide, a pure hydrate is precipitated having the composition BaO2,8H2O (sometimes the composition is taken as BaO2,6H2O). This fact was already known to Thénard. Schöne showed that if hydrogen peroxide be in excess, a crystalline compound of the two peroxides, BaO2H2O2, is precipitated. Schöne also obtained small well-formed crystals of the same composition by adding a solution of ammonia to an acid solution of barium peroxide (containing a barium salt and hydrogen peroxide or a compound of BaO2with the acid). Thus barium peroxide combines with both water and hydrogen peroxide. This is a very important fact for the comprehension of the composition of other peroxides. Moreover, if the peroxides are able to give hydrates they can also form corresponding salts,i.e.they can combine with bases and acids, as was afterwards found to be the case on further research into this subject.
[62]Schöne obtained a compound of peroxide of barium with peroxide of hydrogen. If barium peroxide be dissolved in hydrochloric (or acetic) acid, or if a solution of hydrogen peroxide be diluted with a solution of barium hydroxide, a pure hydrate is precipitated having the composition BaO2,8H2O (sometimes the composition is taken as BaO2,6H2O). This fact was already known to Thénard. Schöne showed that if hydrogen peroxide be in excess, a crystalline compound of the two peroxides, BaO2H2O2, is precipitated. Schöne also obtained small well-formed crystals of the same composition by adding a solution of ammonia to an acid solution of barium peroxide (containing a barium salt and hydrogen peroxide or a compound of BaO2with the acid). Thus barium peroxide combines with both water and hydrogen peroxide. This is a very important fact for the comprehension of the composition of other peroxides. Moreover, if the peroxides are able to give hydrates they can also form corresponding salts,i.e.they can combine with bases and acids, as was afterwards found to be the case on further research into this subject.
[63]Anhydroussulphuric peroxide, S2O7, is obtained by the prolonged (8 to 10 hours) action of a silent discharge of considerable intensity on a mixture of oxygen and sulphurous anhydride; the vapour of sulphuric peroxide, S2O7, condenses as liquid drops, or after being cooled to 0° in the form of long prismatic crystals, resembling those of sulphuric anhydride. The anhydrous compound S2O7(and also the hydrated compound) cannot be preserved long, as it splits up into oxygen and sulphuric anhydride. Direct experiment shows that a mixture of equal volumes of sulphurous anhydride and oxygen leaves a residue of a quarter of the oxygen taken, or half of the whole volume, which indicates the formula S2O7. This substance is soluble in water, and it then gives a hydrate, probably having the composition S2O7,H2O = 2SHO4. This solution oxidises the salts SnX2, potassium iodide, and others, which renders it possible to prove that the solution actually contains one atom of oxygen capable of effecting oxidation to two molecules of sulphuric anhydride.In order to fully demonstrate the reality of a peroxide form for acids, it should be mentioned that some years ago Brodie obtained the so-calledacetic peroxide, (C2H2O)2O2, by the action of barium peroxide on acetic anhydride, (C2H3O)2O. Its corresponding hydrate is also known. This shows that true peroxides and their hydrates, with reactions similar to those of hydrogen peroxide, are possible for acids. A similar higher oxide has long been known for chromium, and Berthelot obtained a like compound for nitric acid (Chapter VI., Note26).
[63]Anhydroussulphuric peroxide, S2O7, is obtained by the prolonged (8 to 10 hours) action of a silent discharge of considerable intensity on a mixture of oxygen and sulphurous anhydride; the vapour of sulphuric peroxide, S2O7, condenses as liquid drops, or after being cooled to 0° in the form of long prismatic crystals, resembling those of sulphuric anhydride. The anhydrous compound S2O7(and also the hydrated compound) cannot be preserved long, as it splits up into oxygen and sulphuric anhydride. Direct experiment shows that a mixture of equal volumes of sulphurous anhydride and oxygen leaves a residue of a quarter of the oxygen taken, or half of the whole volume, which indicates the formula S2O7. This substance is soluble in water, and it then gives a hydrate, probably having the composition S2O7,H2O = 2SHO4. This solution oxidises the salts SnX2, potassium iodide, and others, which renders it possible to prove that the solution actually contains one atom of oxygen capable of effecting oxidation to two molecules of sulphuric anhydride.
In order to fully demonstrate the reality of a peroxide form for acids, it should be mentioned that some years ago Brodie obtained the so-calledacetic peroxide, (C2H2O)2O2, by the action of barium peroxide on acetic anhydride, (C2H3O)2O. Its corresponding hydrate is also known. This shows that true peroxides and their hydrates, with reactions similar to those of hydrogen peroxide, are possible for acids. A similar higher oxide has long been known for chromium, and Berthelot obtained a like compound for nitric acid (Chapter VI., Note26).
[64]When an acid of the strength H2SO46H2O is taken, at first only the hydrate of the sulphuric peroxide, S2O7H2O, is formed, but when the concentration at the positive pole reaches H2SO43H2O, a mixture of hydrogen peroxide and the hydrate of sulphuric peroxide begins to be formed. A state of equilibrium is ultimately arrived at when the amounts of these substances correspond to the proportion S2O7: 2H2O2, which, as it were, answers to a new hydrate, S2O92H2O. But its existence cannot be admitted because the sulphuric peroxide can be easily distinguished from the hydrogen peroxide in the solution owing to the fact that the former does not act on an acid solution of potassium permanganate, whilst the hydrogen peroxide disengages both its own oxygen and that of the permanganic acid, converting it into manganous oxide. Their common property of liberating iodine from an acid solution of the potassium iodide enables the sum of the active oxygen in them both to be determined.
[64]When an acid of the strength H2SO46H2O is taken, at first only the hydrate of the sulphuric peroxide, S2O7H2O, is formed, but when the concentration at the positive pole reaches H2SO43H2O, a mixture of hydrogen peroxide and the hydrate of sulphuric peroxide begins to be formed. A state of equilibrium is ultimately arrived at when the amounts of these substances correspond to the proportion S2O7: 2H2O2, which, as it were, answers to a new hydrate, S2O92H2O. But its existence cannot be admitted because the sulphuric peroxide can be easily distinguished from the hydrogen peroxide in the solution owing to the fact that the former does not act on an acid solution of potassium permanganate, whilst the hydrogen peroxide disengages both its own oxygen and that of the permanganic acid, converting it into manganous oxide. Their common property of liberating iodine from an acid solution of the potassium iodide enables the sum of the active oxygen in them both to be determined.
[65]If a solution of sulphuric acid which has been first subjected to electrolysis be neutralised with potash or baryta, the salt which is formed begins to decompose rapidly with the evolution of oxygen (Berthelot, 1890). On saturating with caustic baryta, the solution of the salt formed may be separated from the sulphate of barium, and then the composition of the resultant compound, BaS2O8, may be determined from the amount of oxygen disengaged. Marshall (1891) studied the formation of this class of compounds more fully; he subjected a saturated solution of bisulphate of potassium to electrolysis with a current of 3–3½ ampères; before electrolysis dilute sulphuric acid is added to the liquid surrounding the negative pole, and during electrolysis the solution at the anode is cooled. The electrolysis is continued without interruption for two days, and a white crystalline deposit separates at the anode. To avoid decomposition, the latter is not filtered through paper, but through a perforated platinum plate, and dried on a porous tile. The mother liquor, with the addition of a fresh solution of bisulphate of potassium, is again subjected to electrolysis and the crystals formed at the anode are again collected, &c. The salt so obtained may be recrystallised by dissolving it in hot water and rapidly cooling the solution after filtration; a small proportion of the salt is decomposed by this treatment. Rapid cooling is followed by the formation of small columnar crystals; slow cooling gives large prismatic crystals. The composition of the salt is determined either by igniting it, when it forms sulphate of potassium, or else by titrating the active oxygen with permanganate: its composition was found to correspond to the salt of persulphuric acid, K2S2O8. The solution of the salt has a neutral reaction, and does not give a precipitate with salts of other metals. K2S2O8is the most insoluble of the salts of persulphuric acid. With nitrate of silver it forms persulphate of silver, which gives peroxide of silver under the action of water according to the equation Ag2S2O8+ 2H2O = Ag2O2+ 2H2SO4. With an alkaline solution of a cupric salt (Fehling's solution) it forms a red precipitate of peroxide of copper. Manganese and cobalt salts give precipitates of MnO2and Co2O3. Ferrous salts are rapidly oxidised, potassium iodide slowly disengages iodine at the ordinary temperature. All these reactions indicate the powerful oxidising properties of K2S2O8. In oxidising in the presence of water it gives a residue of KHSO4. The decomposition of the dry salt begins at 100° but is not complete even at 250°. The freshly prepared salt is inodorous, but after being kept in a closed vessel it evolves a peculiar smell different from that of ozone. The ammonium salt of persulphuric acid, (NH4)2S2O8, is obtained in a similar manner. It is soluble to the extent of 58 parts per 100 parts by weight of water. The decomposition of the ammonium salt by the hydrated oxide of barium gives the barium salt, BaS2O84H2O, which is soluble to the extent of 52·2 parts in 100 parts of water at 0°. The crystals do not deliquesce in the air and decompose in the course of several days; they decompose most rapidly in perfectly dry air. Solutions of the pure salt decompose slowly at the ordinary temperature; on boiling barium sulphate is gradually precipitated, oxygen being liberated simultaneously. To completely decompose this salt it is necessary to boil its solution for a long time. Alcohol dissolves the solid salt; the anhydrous salt does not separate from the alcoholic solution, but a hydrate containing one molecule of water, BaS2O8H2O, which is soluble in water but insoluble in absolute alcohol. Solid barium persulphate decomposes even when slightly heated. The free acid, which may serve for the preparation of other salts, is obtained by treating the barium salt with sulphuric acid. The lead salt, PbS2O8, has been obtained from the free acid; it crystallises with two or three molecules of water. It is soluble in water, deliquesces in the air, and with alkalis gives a precipitate of the hydrated oxide which rapidly oxidises into the binoxide.Traube, before Marshall's researches, thought that the electrolysis of solutions of sulphuric acid did not give persulphuric acid but a persulphuric oxide having the composition SO4. On repeating his former researches (1892) Traube obtained a persulphuric oxide by the electrolysis of a 70 per cent. solution of sulphuric acid, and he separated it from the solution by means of barium phosphate. Analysis showed that this substance corresponded to the above composition SO4, and therefore Traube considers it very likely that the salts obtained by Marshall corresponded to an acid H2SO4+ SO4,i.e.that the indifferent oxide, SO4, can combine with sulphuric acid and form peculiar saline compounds.
[65]If a solution of sulphuric acid which has been first subjected to electrolysis be neutralised with potash or baryta, the salt which is formed begins to decompose rapidly with the evolution of oxygen (Berthelot, 1890). On saturating with caustic baryta, the solution of the salt formed may be separated from the sulphate of barium, and then the composition of the resultant compound, BaS2O8, may be determined from the amount of oxygen disengaged. Marshall (1891) studied the formation of this class of compounds more fully; he subjected a saturated solution of bisulphate of potassium to electrolysis with a current of 3–3½ ampères; before electrolysis dilute sulphuric acid is added to the liquid surrounding the negative pole, and during electrolysis the solution at the anode is cooled. The electrolysis is continued without interruption for two days, and a white crystalline deposit separates at the anode. To avoid decomposition, the latter is not filtered through paper, but through a perforated platinum plate, and dried on a porous tile. The mother liquor, with the addition of a fresh solution of bisulphate of potassium, is again subjected to electrolysis and the crystals formed at the anode are again collected, &c. The salt so obtained may be recrystallised by dissolving it in hot water and rapidly cooling the solution after filtration; a small proportion of the salt is decomposed by this treatment. Rapid cooling is followed by the formation of small columnar crystals; slow cooling gives large prismatic crystals. The composition of the salt is determined either by igniting it, when it forms sulphate of potassium, or else by titrating the active oxygen with permanganate: its composition was found to correspond to the salt of persulphuric acid, K2S2O8. The solution of the salt has a neutral reaction, and does not give a precipitate with salts of other metals. K2S2O8is the most insoluble of the salts of persulphuric acid. With nitrate of silver it forms persulphate of silver, which gives peroxide of silver under the action of water according to the equation Ag2S2O8+ 2H2O = Ag2O2+ 2H2SO4. With an alkaline solution of a cupric salt (Fehling's solution) it forms a red precipitate of peroxide of copper. Manganese and cobalt salts give precipitates of MnO2and Co2O3. Ferrous salts are rapidly oxidised, potassium iodide slowly disengages iodine at the ordinary temperature. All these reactions indicate the powerful oxidising properties of K2S2O8. In oxidising in the presence of water it gives a residue of KHSO4. The decomposition of the dry salt begins at 100° but is not complete even at 250°. The freshly prepared salt is inodorous, but after being kept in a closed vessel it evolves a peculiar smell different from that of ozone. The ammonium salt of persulphuric acid, (NH4)2S2O8, is obtained in a similar manner. It is soluble to the extent of 58 parts per 100 parts by weight of water. The decomposition of the ammonium salt by the hydrated oxide of barium gives the barium salt, BaS2O84H2O, which is soluble to the extent of 52·2 parts in 100 parts of water at 0°. The crystals do not deliquesce in the air and decompose in the course of several days; they decompose most rapidly in perfectly dry air. Solutions of the pure salt decompose slowly at the ordinary temperature; on boiling barium sulphate is gradually precipitated, oxygen being liberated simultaneously. To completely decompose this salt it is necessary to boil its solution for a long time. Alcohol dissolves the solid salt; the anhydrous salt does not separate from the alcoholic solution, but a hydrate containing one molecule of water, BaS2O8H2O, which is soluble in water but insoluble in absolute alcohol. Solid barium persulphate decomposes even when slightly heated. The free acid, which may serve for the preparation of other salts, is obtained by treating the barium salt with sulphuric acid. The lead salt, PbS2O8, has been obtained from the free acid; it crystallises with two or three molecules of water. It is soluble in water, deliquesces in the air, and with alkalis gives a precipitate of the hydrated oxide which rapidly oxidises into the binoxide.
Traube, before Marshall's researches, thought that the electrolysis of solutions of sulphuric acid did not give persulphuric acid but a persulphuric oxide having the composition SO4. On repeating his former researches (1892) Traube obtained a persulphuric oxide by the electrolysis of a 70 per cent. solution of sulphuric acid, and he separated it from the solution by means of barium phosphate. Analysis showed that this substance corresponded to the above composition SO4, and therefore Traube considers it very likely that the salts obtained by Marshall corresponded to an acid H2SO4+ SO4,i.e.that the indifferent oxide, SO4, can combine with sulphuric acid and form peculiar saline compounds.
[65 bis]Or one of those supposed ions which appear at the positive pole in the decomposition of sulphuric acid by the action of a galvanic current.
[65 bis]Or one of those supposed ions which appear at the positive pole in the decomposition of sulphuric acid by the action of a galvanic current.
[66]If this be true one would expect the following peroxide hydrates: for phosphoric acid, (H2PO4)2= H4P2O8= 2H2O + 2PO3; for carbonic acid, (HCO3)2= H2C2O6= H2O + C2O5; and for lead the true peroxide will be also Pb2O5, &c. Judging from the example of barium peroxide (Note62), these peroxide forms will probably combine together. It seems to me that the compounds obtained by Fairley for uranium are very instructive as elucidating the peroxides. In the action of hydrogen peroxide in an acid solution on uranium oxide, UO3, there is formed a uranium peroxide, UO4,4H2O (U = 240), but hydrogen peroxide acts on uranium oxide in the presence of caustic soda; on the addition of alcohol a crystalline compound containing Na4UO8,4H2O is precipitated, which is doubtless a compound of the peroxides of sodium, Na2O2, and uranium, UO4. It is very possible that the first peroxide, UO4,4H2O, contains the elements of hydrogen peroxide and uranium peroxide, U2O7, or even U(OH)6,H2O2, just as the peroxide form lately discovered by Spring for tin perhaps contains Sn2O3,H2O2.
[66]If this be true one would expect the following peroxide hydrates: for phosphoric acid, (H2PO4)2= H4P2O8= 2H2O + 2PO3; for carbonic acid, (HCO3)2= H2C2O6= H2O + C2O5; and for lead the true peroxide will be also Pb2O5, &c. Judging from the example of barium peroxide (Note62), these peroxide forms will probably combine together. It seems to me that the compounds obtained by Fairley for uranium are very instructive as elucidating the peroxides. In the action of hydrogen peroxide in an acid solution on uranium oxide, UO3, there is formed a uranium peroxide, UO4,4H2O (U = 240), but hydrogen peroxide acts on uranium oxide in the presence of caustic soda; on the addition of alcohol a crystalline compound containing Na4UO8,4H2O is precipitated, which is doubtless a compound of the peroxides of sodium, Na2O2, and uranium, UO4. It is very possible that the first peroxide, UO4,4H2O, contains the elements of hydrogen peroxide and uranium peroxide, U2O7, or even U(OH)6,H2O2, just as the peroxide form lately discovered by Spring for tin perhaps contains Sn2O3,H2O2.
[67]This view was communicated by me in 1870 to the Russian Chemical Society.
[67]This view was communicated by me in 1870 to the Russian Chemical Society.
[68]Dithionic acid, H2S2O6, is distinguished among the thionic acids as containing the least proportion of sulphur. It is also called hyposulphuric acid, because its supposed anhydride, S2O5, contains more O than sulphurous oxide, SO2or S2O4, and less than sulphuric anhydride, SO3or S2O6. Dithionic acid, discovered by Gay-Lussac and Welter, is known as a hydrate and as salts, but not as anhydride. The method for preparing dithionic acid usually employed is by the action of finely-powdered manganese dioxide on a solution of sulphurous anhydride. On shaking, the smell of the latter disappears, and the manganese salt of the acid in question passes into solution; MnO2+ 2SO2= MnS2O6. If the temperature be raised, the dithionate splits up into sulphurous anhydride and manganese sulphate, MnSO4. Generally owing to this a mixture of manganese sulphate and dithionate is obtained in the solution. They may be separated by mixing the solution of the manganese salts with a solution of barium hydroxide, when a precipitate of manganese hydroxide and barium sulphate is obtained. In this manner barium dithionate only is obtained in solution. It is purified by crystallisation, and separates as BaS2O6,2H2O; this is then dissolved in water, and decomposed with the requisite amount of sulphuric acid. Dithionic acid, H2S2O6, then remains in solution. By concentrating the resultant solution under the receiver of an air-pump it is possible to obtain a liquid of sp. gr. 1·347, but it still contains water, and on further evaporation the acid decomposes into sulphuric acid and sulphurous anhydride: H2S2O6= H2SO4+ SO2. The same decomposition takes place if the solution be slightly heated. Like all the thionic acids, dithionic acid is readily attacked by oxidising agents, and passes into sulphurous acid. No dithionate is able to withstand the action of heat, even when very slight, without giving off sulphurous anhydride: K2S2O6= K2SO4+ SO2. The alkali dithionates have a neutral reaction (which indicates the energetic nature of the acid) are soluble in water, and in this respect present a certain resemblance to the salts of nitric acid (their anhydrides are: N2O5and S2O5). Klüss (1888) described many of the salts of dithionic acid.Langlois, about 1840, obtained a peculiar thionic acid by heating a strong solution of acid potassium sulphite with flowers of sulphur to about 60°, until the disappearance of the yellow coloration first produced by the solution of the sulphur. On cooling, a portion of the sulphur was precipitated, and crystals of a salt oftrithionic acid, K2S3O6(partly mixed with potassium sulphate), separated out. Plessy afterwards showed that the action of sulphurous acid on a thiosulphate also gives sulphur and trithionic acid: 2K2S2O3+ 3SO2= 2K2S3O6+ S. A mixture of potassium acid sulphite and thiosulphate also gives a trithionate. It is very possible that a reaction of the same kind occurs in the formation of trithionic acid by Langloid's method, because potassium sulphite and sulphur yield potassium thiosulphate. The potassium thiosulphate may also be replaced by potassium sulphide, and on passing sulphurous anhydride through the solution thiosulphate is first formed and then trithionate: 4KHSO3+ K2S + 4SO2= 3K2S3O6+ 2H2O. The sodium salt is not formed under the same circumstances as the corresponding potassium salt. The sodium salt does not crystallise and is very unstable: the barium salt is, however, more stable. The barium and potassium salts are anhydrous, they give neutral solutions and decompose when ignited, with the evolution of sulphur and sulphurous anhydride, a sulphate being left behind, K2S3O6= K2SO4+ SO2+ S. If a solution of the potassium salt be decomposed by means of hydrofluosilicic or chloric acid, the insoluble salts of these acids are precipitated and trithionic acid is obtained in solution, which however very easily breaks up on concentration. The addition of salts of copper, mercury, silver, &c., to a solution of a trithionate is followed, either immediately or after a certain time, by the formation of a black precipitate of the sulphides whose formation is due to the decomposition of the trithionic acid with the transference of its sulphur to the metal.Tetrathionic acid, H2S4O6, in contradistinction to the preceding acids, is much more stable in the free state than in the form of salts. In the latter form it is easily converted into trithionate, with liberation of sulphur. Sodium tetrathionate was obtained by Fordos and Gélis, by the action of iodine on a solution of sodium thiosulphate. The reaction essentially consists in the iodine taking up half the sodium of the thiosulphate, inasmuch as the latter contains Na2S2O3, whilst the tetrathionate contains NaS2O3or Na2S4O6, so that the reaction is as follows: 2Na2S2O3+ I2= 2NaI + Na2S4O6. It is evident that tetrathionic acid stands to thiosulphuric acid in exactly the same relation as dithionic acid does to sulphurous acid; for the same amount of the other elements in dithionate, KSO3, and tetrathionate, KS2O3, there is half as much metal as in sulphite, K2SO3, and thiosulphate, K2S2O3. If in the above reaction the sodium thiosulphate be replaced by the lead salt PbS2O3, the sparingly-soluble lead iodide PbI2and the soluble salt PbS4O6are obtained. Moreover the lead salt easily gives tetrathionic acid itself (PbSO4is precipitated). The solution of tetrathionic acid may be evaporated over a water-bath, and afterwards in a vacuum, when it gives a colourless liquid, which has no smell and a very acid reaction. When dilute it may be heated to its boiling-point, but in a concentrated form it decomposes into sulphuric acid, sulphurous anhydride, and sulphur: H2S4O6= H2SO4+ SO2+ S2.Pentathionic acid, H2S5O6, also belongs to this series of acids. But little is known concerning it, either as hydrate or in salts. It is formed, together with tetrathionic acid, by the direct action of sulphurous acid on sulphuretted hydrogen in an aqueous solution; a large proportion of sulphur being precipitated at the same time: 5SO2+ 5H2S = H2S5O6+ 5S + 4H2O.If, as was shown above, the thionic acids are disulphonic acids, they may be obtained, like other sulphonic acids, by means of potassium sulphite and sulphur chloride. Thus Spring demonstrated the formation of potassium trithionate by the action of sulphur dichloride on a strong solution of potassium sulphite: 2KSO3K + SCl2= S(SO3K)2+ 2KCl. If sulphur chloride be taken, sulphur also is precipitated. The same trithionate is formed by heating a solution of double thiosulphates; for example, of AgKS2O3. Two molecules of the salts then form silver sulphide and potassium trithionate. If the thiosulphate be the potassium silver salt SO3K(AgS), then the structure of the trithionate must necessarily be (SO3K)2S. Previous to Spring's researches, the action of iodine on sodium thiosulphate was an isolated accidentally discovered reaction; he, however, showed its general significance by testing the action of iodine on mixtures of different sulphur compounds. Thus with iodine, I2, the mixture Na2S + Na2SO3forms 2NaI + Na2S2O3, whilst the mixture Na2S2O3+ Na2SO3+ I2gives 2NaI + Na2S3O6—that is, trithionic acid stands in the same relation to thiosulphuric acid as the latter does to sulphuretted hydrogen. We adopt the same mode of representation: by replacing one hydrogen in H2S by sulphuryl we obtain thiosulphuric acid, HSO3.HS, and by replacing a second hydrogen in the latter again by sulphuryl we obtain trithionic acid, (HSO3)2S. Furthermore, Spring showed that the action of sodium amalgam on the thionic acids causes reverse reactions to those above indicated for iodine. Thus sodium thiosulphate with Na2gives Na2S + Na2SO3, and Spring showed that the sodium here is not a simple element taking up sulphur, but itself enters into double decomposition, replacing sulphur; for on taking a potassium salt and acting on it with sodium, KSO3(SK) + NaNa = KSO3Na + (SK)Na. In a similar way sodium dithionate with sodium gives sodium sulphite: (NaSO3)2+ Na2= 2NaSO3Na; sodium trithionate forms NaSO3Na and NaSO3.SNa, and tetrathionate forms sodium thiosulphate, (NaSO3)S2(NaSO3) + Na2= 2(NaSO3)(NaS).In all the oxidised compounds of sulphur we may note the presence of the elements of sulphurous anhydride, SO2, the only product of the combustion of sulphur, and in this sense the compounds of sulphur containing one SO2are—fig_258_1while, according to this mode of representation, the thionic acids are—fig_258_2Hence it is evident that SO2has (whilst CO2has not) the faculty for combination, and aims at forming SO2X2. These X2can = O, and the question naturally suggests itself as to whether the O2which occurs in SO2is not of the same nature as this oxygen which adds itself to SO2—that is, whether SO2does not correspond with the more general type SX4, and its compounds with the type SX6? To this we may answer ‘Yes’ and ‘No’—‘Yes’ in the general sense which proceeds from the investigation of the majority of compounds, especially metals, where RO corresponds with RCl2, RX2; ‘No’ in the sense that sulphur does not give either SH4, SH6, or SCl6, and therefore the stages SX4and SX6are only observable in oxygen compounds. With reference to the type SX6a hydrate, S(HO)6, might be expected, if not SCl6. And we must recognise this hydrate from a study of the compounds of sulphuric acid with water. In addition to what has been already said respecting the complex acids formed by sulphur, I think it well to mention that, according to the above view, still more complex oxygen acids and salts of sulphur may be looked for. For instance, the salt Na2S4O8obtained by Villiers (1888) is of this kind. It is formed together with sodium trithionate and sulphur, when SO2is passed through a cold solution of Na2S2O3, which is then allowed to stand for several days at the ordinary temperature: 2Na2S2O3+ 4SO2= Na2S4O8+ Na2S3O6+ S. It may be assumed here, as in the thionic acids, that there are two sulphoxyls, bound together not only by S, but also by SO2, or what is almost the same thing, that the sulphoxyl is combined with the residue of trithionic acid,i.e.replaces one aqueous residue in trithionic acid.
[68]Dithionic acid, H2S2O6, is distinguished among the thionic acids as containing the least proportion of sulphur. It is also called hyposulphuric acid, because its supposed anhydride, S2O5, contains more O than sulphurous oxide, SO2or S2O4, and less than sulphuric anhydride, SO3or S2O6. Dithionic acid, discovered by Gay-Lussac and Welter, is known as a hydrate and as salts, but not as anhydride. The method for preparing dithionic acid usually employed is by the action of finely-powdered manganese dioxide on a solution of sulphurous anhydride. On shaking, the smell of the latter disappears, and the manganese salt of the acid in question passes into solution; MnO2+ 2SO2= MnS2O6. If the temperature be raised, the dithionate splits up into sulphurous anhydride and manganese sulphate, MnSO4. Generally owing to this a mixture of manganese sulphate and dithionate is obtained in the solution. They may be separated by mixing the solution of the manganese salts with a solution of barium hydroxide, when a precipitate of manganese hydroxide and barium sulphate is obtained. In this manner barium dithionate only is obtained in solution. It is purified by crystallisation, and separates as BaS2O6,2H2O; this is then dissolved in water, and decomposed with the requisite amount of sulphuric acid. Dithionic acid, H2S2O6, then remains in solution. By concentrating the resultant solution under the receiver of an air-pump it is possible to obtain a liquid of sp. gr. 1·347, but it still contains water, and on further evaporation the acid decomposes into sulphuric acid and sulphurous anhydride: H2S2O6= H2SO4+ SO2. The same decomposition takes place if the solution be slightly heated. Like all the thionic acids, dithionic acid is readily attacked by oxidising agents, and passes into sulphurous acid. No dithionate is able to withstand the action of heat, even when very slight, without giving off sulphurous anhydride: K2S2O6= K2SO4+ SO2. The alkali dithionates have a neutral reaction (which indicates the energetic nature of the acid) are soluble in water, and in this respect present a certain resemblance to the salts of nitric acid (their anhydrides are: N2O5and S2O5). Klüss (1888) described many of the salts of dithionic acid.
Langlois, about 1840, obtained a peculiar thionic acid by heating a strong solution of acid potassium sulphite with flowers of sulphur to about 60°, until the disappearance of the yellow coloration first produced by the solution of the sulphur. On cooling, a portion of the sulphur was precipitated, and crystals of a salt oftrithionic acid, K2S3O6(partly mixed with potassium sulphate), separated out. Plessy afterwards showed that the action of sulphurous acid on a thiosulphate also gives sulphur and trithionic acid: 2K2S2O3+ 3SO2= 2K2S3O6+ S. A mixture of potassium acid sulphite and thiosulphate also gives a trithionate. It is very possible that a reaction of the same kind occurs in the formation of trithionic acid by Langloid's method, because potassium sulphite and sulphur yield potassium thiosulphate. The potassium thiosulphate may also be replaced by potassium sulphide, and on passing sulphurous anhydride through the solution thiosulphate is first formed and then trithionate: 4KHSO3+ K2S + 4SO2= 3K2S3O6+ 2H2O. The sodium salt is not formed under the same circumstances as the corresponding potassium salt. The sodium salt does not crystallise and is very unstable: the barium salt is, however, more stable. The barium and potassium salts are anhydrous, they give neutral solutions and decompose when ignited, with the evolution of sulphur and sulphurous anhydride, a sulphate being left behind, K2S3O6= K2SO4+ SO2+ S. If a solution of the potassium salt be decomposed by means of hydrofluosilicic or chloric acid, the insoluble salts of these acids are precipitated and trithionic acid is obtained in solution, which however very easily breaks up on concentration. The addition of salts of copper, mercury, silver, &c., to a solution of a trithionate is followed, either immediately or after a certain time, by the formation of a black precipitate of the sulphides whose formation is due to the decomposition of the trithionic acid with the transference of its sulphur to the metal.
Tetrathionic acid, H2S4O6, in contradistinction to the preceding acids, is much more stable in the free state than in the form of salts. In the latter form it is easily converted into trithionate, with liberation of sulphur. Sodium tetrathionate was obtained by Fordos and Gélis, by the action of iodine on a solution of sodium thiosulphate. The reaction essentially consists in the iodine taking up half the sodium of the thiosulphate, inasmuch as the latter contains Na2S2O3, whilst the tetrathionate contains NaS2O3or Na2S4O6, so that the reaction is as follows: 2Na2S2O3+ I2= 2NaI + Na2S4O6. It is evident that tetrathionic acid stands to thiosulphuric acid in exactly the same relation as dithionic acid does to sulphurous acid; for the same amount of the other elements in dithionate, KSO3, and tetrathionate, KS2O3, there is half as much metal as in sulphite, K2SO3, and thiosulphate, K2S2O3. If in the above reaction the sodium thiosulphate be replaced by the lead salt PbS2O3, the sparingly-soluble lead iodide PbI2and the soluble salt PbS4O6are obtained. Moreover the lead salt easily gives tetrathionic acid itself (PbSO4is precipitated). The solution of tetrathionic acid may be evaporated over a water-bath, and afterwards in a vacuum, when it gives a colourless liquid, which has no smell and a very acid reaction. When dilute it may be heated to its boiling-point, but in a concentrated form it decomposes into sulphuric acid, sulphurous anhydride, and sulphur: H2S4O6= H2SO4+ SO2+ S2.
Pentathionic acid, H2S5O6, also belongs to this series of acids. But little is known concerning it, either as hydrate or in salts. It is formed, together with tetrathionic acid, by the direct action of sulphurous acid on sulphuretted hydrogen in an aqueous solution; a large proportion of sulphur being precipitated at the same time: 5SO2+ 5H2S = H2S5O6+ 5S + 4H2O.
If, as was shown above, the thionic acids are disulphonic acids, they may be obtained, like other sulphonic acids, by means of potassium sulphite and sulphur chloride. Thus Spring demonstrated the formation of potassium trithionate by the action of sulphur dichloride on a strong solution of potassium sulphite: 2KSO3K + SCl2= S(SO3K)2+ 2KCl. If sulphur chloride be taken, sulphur also is precipitated. The same trithionate is formed by heating a solution of double thiosulphates; for example, of AgKS2O3. Two molecules of the salts then form silver sulphide and potassium trithionate. If the thiosulphate be the potassium silver salt SO3K(AgS), then the structure of the trithionate must necessarily be (SO3K)2S. Previous to Spring's researches, the action of iodine on sodium thiosulphate was an isolated accidentally discovered reaction; he, however, showed its general significance by testing the action of iodine on mixtures of different sulphur compounds. Thus with iodine, I2, the mixture Na2S + Na2SO3forms 2NaI + Na2S2O3, whilst the mixture Na2S2O3+ Na2SO3+ I2gives 2NaI + Na2S3O6—that is, trithionic acid stands in the same relation to thiosulphuric acid as the latter does to sulphuretted hydrogen. We adopt the same mode of representation: by replacing one hydrogen in H2S by sulphuryl we obtain thiosulphuric acid, HSO3.HS, and by replacing a second hydrogen in the latter again by sulphuryl we obtain trithionic acid, (HSO3)2S. Furthermore, Spring showed that the action of sodium amalgam on the thionic acids causes reverse reactions to those above indicated for iodine. Thus sodium thiosulphate with Na2gives Na2S + Na2SO3, and Spring showed that the sodium here is not a simple element taking up sulphur, but itself enters into double decomposition, replacing sulphur; for on taking a potassium salt and acting on it with sodium, KSO3(SK) + NaNa = KSO3Na + (SK)Na. In a similar way sodium dithionate with sodium gives sodium sulphite: (NaSO3)2+ Na2= 2NaSO3Na; sodium trithionate forms NaSO3Na and NaSO3.SNa, and tetrathionate forms sodium thiosulphate, (NaSO3)S2(NaSO3) + Na2= 2(NaSO3)(NaS).
In all the oxidised compounds of sulphur we may note the presence of the elements of sulphurous anhydride, SO2, the only product of the combustion of sulphur, and in this sense the compounds of sulphur containing one SO2are—
fig_258_1
while, according to this mode of representation, the thionic acids are—
fig_258_2
Hence it is evident that SO2has (whilst CO2has not) the faculty for combination, and aims at forming SO2X2. These X2can = O, and the question naturally suggests itself as to whether the O2which occurs in SO2is not of the same nature as this oxygen which adds itself to SO2—that is, whether SO2does not correspond with the more general type SX4, and its compounds with the type SX6? To this we may answer ‘Yes’ and ‘No’—‘Yes’ in the general sense which proceeds from the investigation of the majority of compounds, especially metals, where RO corresponds with RCl2, RX2; ‘No’ in the sense that sulphur does not give either SH4, SH6, or SCl6, and therefore the stages SX4and SX6are only observable in oxygen compounds. With reference to the type SX6a hydrate, S(HO)6, might be expected, if not SCl6. And we must recognise this hydrate from a study of the compounds of sulphuric acid with water. In addition to what has been already said respecting the complex acids formed by sulphur, I think it well to mention that, according to the above view, still more complex oxygen acids and salts of sulphur may be looked for. For instance, the salt Na2S4O8obtained by Villiers (1888) is of this kind. It is formed together with sodium trithionate and sulphur, when SO2is passed through a cold solution of Na2S2O3, which is then allowed to stand for several days at the ordinary temperature: 2Na2S2O3+ 4SO2= Na2S4O8+ Na2S3O6+ S. It may be assumed here, as in the thionic acids, that there are two sulphoxyls, bound together not only by S, but also by SO2, or what is almost the same thing, that the sulphoxyl is combined with the residue of trithionic acid,i.e.replaces one aqueous residue in trithionic acid.