see captionFig.87.—Concentration of sulphuric acid in glass retorts. The neck of each retort is attached to a bent glass tube, whose vertical arm is lowered into a glass or earthenware vessel acting as a receiver for the steam which comes over from the acid, as the former still contains a certain amount of acid.
Fig.87.—Concentration of sulphuric acid in glass retorts. The neck of each retort is attached to a bent glass tube, whose vertical arm is lowered into a glass or earthenware vessel acting as a receiver for the steam which comes over from the acid, as the former still contains a certain amount of acid.
Sulphuric acid, H2SO4, is formed by the combination of its anhydride, SO3, and water, with the evolution of a large amount of heat; the reaction SO3+ H2O develops 21,300 heat units. The method of its preparation on a large scale, and most of the methods employed for its formation, are dependent on the oxidation of sulphurous anhydride, and the formation of sulphuric anhydride, which forms sulphuric acid under the action of water. The technical method of its manufacture has been described in ChapterVI.The acid obtainedfrom the lead chamberscontains a considerable amount of water, and is also impure owing to the presence of oxides of nitrogen, lead compounds, and certain impurities from the burnt sulphur which have come over in a gaseous and vaporous state (for example, arsenic compounds). For practical purposes, hardly any notice is taken of the majority of these impurities, because they do not interfere with its general qualities. Most frequently endeavours are only made to remove, as far as possible, all the water which can be expelled.[45]That is, the objectis to obtain the hydrate, H2SO4, from the dilute acid (60 per cent.), and this is effected by evaporation by means of heat. Every given mixture of water and sulphuric acid begins to part with a certain amount of aqueous vapour when heated to a certain definite temperature. At a low temperature either there is no evaporation of water, or there can even be an absorption of moisture from the air. As the removal of the water proceeds, the vapour tension of the residue decreases for the same temperature, and therefore the more dilute the acid the lower the temperature at which it gives up a portion of its water. In consequence of this, the removal of water from dilute solutions of sulphuric acid may be easily carried on (up to 75 p.c. H2SO4) in lead vessels, because at low temperatures dilute sulphuric acid does not attack lead. But as the acid becomes more concentrated the temperature at which the water comes over becomes higher and higher, and then the acidbegins to act on lead (with the evolution of sulphuretted hydrogen and conversion of the lead into sulphate), and therefore lead vessels cannot be employed for the complete removal of the water. For this purpose the evaporation is generally carried on in glass or platinum retorts, like those depicted in figs. 87 and 88.
see captionFig.88.—Concentration of sulphuric acid in platinum retorts.
Fig.88.—Concentration of sulphuric acid in platinum retorts.
The concentration of sulphuric acidin glass retorts is not a continuous process, and consists of heating the dilute 75 per cent. acid until it ceases to give off aqueous vapour, and until acid containing 93–98 per cent. H2SO4(66° Baumé) is obtained—and this takes place when the temperature reaches 320° and the density of the residue reaches 1·847 (66° Baumé).[46]The platinum vessels designed for the continuous concentration of sulphuric acid consist of a stillB, furnished with a still headE, a connecting pipeE F, and a syphon tubeH R, which draws off the sulphuric acid concentrated in the boiler. A stream of sulphuric acid previously concentrated in lead retorts to a density of about 60° Baumé—i.e.to 75 per cent. or a sp. gr. of 1·7—runs continuously into the retort through a syphon funnelE′. The apparatus is fed from above, because the acid freshly supplied is lighter than that which has already lost water, and also because the water is more easily evaporated from the freshly supplied acid at the surface. The platinumretort is heated, and the steam coming off[47]is condensed in a wormF G, whilst as fresh dilute acid is supplied to the boiler the acid already concentrated is drawn off through the syphon tubeH B, which is furnished with a regulating cock by means of which the outflow of the concentrated acid from the bottom of the retort can be so regulated that it will always present one and the same specific gravity, corresponding with the strength required. For this purpose the acid flowing from the syphon is collected in a receiverR, in which a hydrometer, indicating its density, floats; if its density be less than 66° Baumé, the regulating cock is closed sufficiently to retard the outflow of sulphuric acid, so as to lengthen the time of its evaporation in the retort.[48]
Strictly speaking,sulphuric acid is not volatile, and at its so-called boiling-point it really decomposes into its anhydride and water; its boiling-point (338°) being nothing else but its temperature of decomposition. The products of this decomposition are substances boiling much below the temperature of the decomposition of sulphuric acid. This conclusion with regard to the process of the distillation of sulphuric acid may be deduced from Bineau's observations on the vapour-density of sulphuric acid. This density referred to hydrogen proved to be half that which sulphuric acid should have according to its molecular weight, H2SO4, in which case it should be 49, whilst the observed density was equal to 24·5. Besides which, Marignac showed that the first portions of the sulphuric acid distilling over contain less of the elements of water than the portion which remains behind, or which distils over towards the end. This is explained by the fact that on distillation the sulphuric acid is decomposed, but a portion of the water proceeding from its decomposition is retained by the remaining mass of sulphuric acid, and therefore at first a mixture of sulphuric acid and sulphuric anhydride—i.e.fuming sulphuric acid—is obtained in the distillate. It is possible by repeating the distillation several times and only collecting the first portions of the distillate, to obtain a distinctly fuming acid. To obtain the definite hydrate H2SO4it is necessary to refrigerate a highly concentrated acid, of as great a purity as possible, to which a small quantity of sulphuric anhydride has been previously added. Sulphuric acid containing a small quantity (a fraction of a per cent. by weight) of water only freezes at a very low temperature, while the pure normal acid, H2SO4, solidifies when it is cooled below 0°,and therefore the normal acid first crystallises out from the concentrated sulphuric acid. By repeating the refrigeration several times, and pouring off the unsolidified portion, it is possible to obtain a purenormal hydrate, H2SO4, which melts at 10°·4. Even at 40° it gives off distinct fumes—that is, it begins to evolve sulphuric anhydride, which volatilises, and therefore even in a dry atmosphere the hydrate H2SO4becomes weaker, until it contains 1½ p.c. of water.[49]
In a concentrated form sulphuric acid is commercially known asoil of vitriol, because for a long time it was obtained from green vitriol and because it has an oily appearance and flows from one vessel into another in a thick and somewhat sluggish stream, like the majority of oily substances, and in this clearly differs from such liquids as water, spirit, ether, and the like, which exhibit a far greater mobility. Among its reactions the first to be remarked is its faculty for the formation of many compounds. We already know that it combines with its anhydride, and with the sulphates of the alkali metals; that it is soluble in water, with which it forms more or less stable compounds. Sulphuric acid, when mixed with water, develops a very considerable amount of heat.[50]
Besides the normal hydrate H2SO4,another definite hydrate,H2SO4,H2O (84·48 per cent. of the normal hydrate, and 15·52 per cent. of water) is known; it crystallises[50 bis]extremely easily in large six-sidedprisms, which form above 0°—namely, at about +8°·5; when heated to 210° it loses water.[51]If the hydrates H2SO4and H2SO4,H2O exist at low temperatures as definite crystalline compounds, and if pyrosulphuric acid, H2SO4SO3, has the same property, and if they all decompose with more or less ease on a rise of temperature, with the disengagement of either SO3or H2O, and in their ordinary form present all the properties of simple solutions, it follows that between sulphuric anhydride, SO3, and water, H2O, there exists a consecutive series of homogeneous liquids or solutions, among which we must distinguishdefinite compounds, and therefore it is quite justifiable to look for other definite compounds between SO3and H2O, beyond the conditions for a change of state. In this respect we may be guided by the variation of properties of any kind, proceeding concurrently with a variation in the composition of a solution.
But only a few properties have been determined with sufficient accuracy. In those properties which have been determined for many solutions of sulphuric acid, it is actually seen that the above-mentioned definite compounds are distinguished by distinctive marks of change. As an example we may cite the variation of the specific gravity with a variation of temperature (namely K =ds/dt, ifsbe the sp. gr. andtthe temperature). For the normal hydrate, H2SO4, this factor is easily determined from the fact that—
s= 18528 - 10·65t+ 0·013t2,
wheresis the specific gravity att(degrees Celsius) if the sp. gr. of water at 4° = 10,000. Therefore K = 10·65 - 0·026t. This means that at 0° the sp. gr. of the acid H2SO4decreases by 10·65 for every rise of a degree of temperature, at 10° by 10·39, at 20° by 10·13, at 30° by 9·87.[52]And for solutions containing slightly more anhydride than the acid H2SO4(i.e.for fuming sulphuric acid), as well as for solutions containing more water, K is greater than for the acid H2SO4. Thus for the solution SO3,2H2SO4, at 10° K = 11·0. On diluting the acid H2SO4K again increases until the formation of the solution H2SO4,H2O (K = 11·1 at 10°), and then, on further dilution with water, it again decreases. Consequently both hydrates H2SO4and H2SO4,H2O are here expressed by an alteration of the magnitude of K.
see captionFig.89.—Diagram showing the variation of the factor (ds/dp) of the specific gravity of solutions of sulphuric acid. The percentage quantities of the acid, H2SO4, are laid out on the axes of abscissæ. The ordinates are the factors or rises in sp. gr. (water at 4 = 10,000) with the increase in the quantity of H2SO4.
Fig.89.—Diagram showing the variation of the factor (ds/dp) of the specific gravity of solutions of sulphuric acid. The percentage quantities of the acid, H2SO4, are laid out on the axes of abscissæ. The ordinates are the factors or rises in sp. gr. (water at 4 = 10,000) with the increase in the quantity of H2SO4.
This shows that in liquid solutions it is possible by studying the variation of their properties (without a change of physical state) to recognise the presence or formation of definite hydrate compounds, and therefore an exact investigation of the properties of solutions, of their specific gravity for instance, should give direct indications of such compounds.[53]The mean result of the most trustworthy determinationsof this nature is given in the following tables. The first of these tables gives the specific gravities (in vacuo, taking the sp. gr. of water at 4° = 1), at 0° (column 3), 15° (column 4), and 30° (column 5),[53 bis]for solutions having the composition H2SO4+nH2O (the value ofnis given in the first column), and containingp(column 2) per cent. (by weight in vacuo) of H2SO4.[53 tri]
In the second table the first column gives the percentage amountp(by weight) of H2SO4, the second column the weight in grams (S15) of a litre of the solution at 15° (at 4° the weight of a litre of water = 1,000 grams), the third column, the variation (dS/dt) of this weight for a rise of 1°, the fourth column, the variationdS/dpof this weight (at 15°) for a rise of 1 per cent. of H2SO4, the fifth column, the difference between the weight of a litre at 0° and 15° (S0- S15), and the sixth column, the difference between the weight of a litre at 15° and 30° (S15- S30).
The figures in these tables give the means of finding the amount of H2SO4contained in a solution from its specific gravity,[55]and also show that ‘special points’ in the lines of variation of the specific gravity with the temperature and percentage composition correspond to certain definite compounds of H2SO4with OH2. This is best seen in the variation of the factors (dS/dtanddS/dp) with the temperature andcomposition (columns 3, 4, second table). We have already mentioned how the factor of temperature points to the existence of hydrates, H2SO4and H2SO4,H2O. As regards the factordS/dp(giving the increase of sp. gr. with an increase of 1 per cent. H2SO4) the following are the three most salient points: (1) In passing from 98 per cent. to 100 per cent. the factor is negative, and at 100 per cent. about -0·0019 (i.e.at 99 per cent. the sp. gr. is about 1·8391, and at 100 per cent. about 1·8372, at 15°, the amount of H2SO4has increased whilst the sp. gr. has decreased), but as soon as a certain amount of SO3is added to the definite compound H2SO4(and ‘fuming’ acid formed) the specific gravity rises (for example, for H2SO40·136 SO3the sp. gr. at 15° = 1·866), that is the factor becomes positive (and, in fact, greater by +0·01), so that the formation of the definite hydrate H2SO4is accompanied by a distinct and considerable break in the continuity of the factor[55 bis]; (2) The factor (dS/dp) in increasing in its passage from dilute to concentrated solutions, attains a maximum value (at 15° about 0·012) about H2SO42H2O,i.e.at about the hydrate corresponding to the form SX6; proper to the compounds of sulphur, for S(OH)6= H2SO42H2O; the same hydrate corresponds to the composition of gypsum CaSO42H2O, and to it also corresponds the greatest contraction and rise of temperature in mixing H2SO4with H2O (seeChapter I., Note28); (3) The variation of the factor (dS/dp) under certain variations in the composition proceeds so uniformly and regularly, and is so different from the variation given under other proportions of H2SO4and H2O, that the sum of the variations ofdS/dpis expressed by a series of straight lines, if the values ofpbe laid along the axis of abscissæ and those ofdS/dpalong the ordinates.[56]Thus, for instance, for 15°, at10 per cent.dS/dp= 0·0071, at 20 per cent. = 0·0077, at 30 per cent. = 0·0082, at 40 per cent. = 0·0088, that is, for each 10 per cent. the factor increases by about 0·0006 for the whole of the above range, but beyond this it becomes larger, and then, after passing H2SO42H2O, it begins to fall rapidly. Such changes in the variation of the factor take place apparently about definite hydrates,[56 bis]and especially about H2SO44H2O, H2SO42H2O and H2SO4H2O. All this indicating as it does the special chemical affinity of sulphuric acid for water, although of no small significance for comprehending the nature of solutions (seeChapterI.and ChapterVII.), contains many special points which require detailed investigation, the chief difficulty being that it requires great accuracy in a large number of experimental data.
The great affinity of sulphuric acid for water is also seen fromthe fact that when the strong acid acts on the majority oforganic substancescontaining hydrogen and oxygen (especially on heating) it very frequentlytakes up these elements in the form of water. Thus strong sulphuric acid acting on alcohol, C2H6O, removes the elements of water from it, and converts it into olefiant gas, C2H4. It acts in a similar manner on wood and other vegetable tissues, which it chars. If a piece of wood be immersed in strong sulphuric acid it turns black. This is owing to the fact that the wood contains carbohydrates which give up hydrogen and oxygen as water to the sulphuric acid, leaving charcoal, or a black mass very rich in it. For example, cellulose, C6H10O5, acts in this manner.[57]
We have already had frequent occasion to notice the veryenergetic acid propertiesof sulphuric acid, and therefore we will now only consider a few of their aspects. First of all we must remember that, with calcium, strontium, and especially with barium and lead, sulphuric acid forms very slightly soluble salts, whilst with the majority of other metals it gives more easily soluble salts, which in the majority of cases are able, like sulphuric acid itself, to combine with water to form crystallo-hydrates. Normal sulphuric acid, containing two atoms of hydrogen in its molecule, is able for this reason alone to form two classes of salts,normalandacid, which it does with great facilitywith the alkali metals. The metals of the alkaline earths and the majority of other metals, if they do form acid sulphates, do so under exceptional conditions (with an excess of strong sulphuric acid), and these salts when formed are decomposable by water—that is, although having a certain degree of physical stability they have no chemical stability. Besides the acid salts RHSO4, sulphuric acid also gives other forms of acid salts. An entire series of salts having the composition RHSO4,H2SO4, or for bivalent metals RSO4,3H2SO4,[58]has been prepared. Such salts have been obtained for potassium, sodium, nickel, calcium, silver, magnesium, manganese. They are preparedby dissolving the sulphates in an excess of sulphuric acid and heating the solution until the excess of sulphuric acid is driven off; on cooling, the mass solidifies to a crystalline salt. Besides which, Rose obtained a salt having the composition Na2SO4,NaHSO4, and if HNaSO4be heated it easily forms a salt Na2S2O7= Na2SO4,SO3; hence it is clear that sulphuric anhydride combines with various proportions of bases, just as it combines with various proportions of water.
We have already learned that sulphuric acid displaces the acid from the salts of nitric, carbonic, and many other volatile acids. Berthollet's laws (ChapterX.) explain this by the small volatility of sulphuric acid; and, indeed, in an aqueous solution sulphuric acid displaces the much less soluble boric acid from its compounds—for instance, from borax, and it also displaces silica from its compounds with bases; but both boric anhydride and silica, when fused with sulphates, decompose them, displacing sulphuric anhydride, SO3, because they are less volatile than sulphuric anhydride. It is also well known that with metals, sulphuric acid forms salts giving off hydrogen (Fe, Zn, &c.), or sulphur dioxide (Cu, Hg, &c.).[58 bis]
The reactions of sulphuric acidwith respect to organic substancesare generally determined by its acid character, when the direct extraction of water, or oxidation at the expense of the oxygen of the sulphuric acid,[59]or disintegration does not take place. Thus the majority of the saturated hydrocarbons, CnH2m, form with sulphuric acid a special class ofsulphonic acids, CnH2m-1(HSO3); for example,benzene, C6H6, forms benzenesulphonic acid, C6H5.SO3H, water being separated, for the formation of which oxygen is taken up from the sulphuric acid, for the product contains less oxygen than the sulphuric acid. It is evident from the existence of these acids that the hydrogen in organic compounds is replaceable by the group SO3H, just as it may be replaced by the radicles Cl, NO2, CO2H and others. As the radicle of sulphuric acid orsulphoxyl, SO2OH or SHO3, contains, like carboxyl (Vol. I., p.395), one hydrogen (hydroxyl) of sulphuric acid, the resultant substances are acids whose basicity is equal to the number of hydrogens replaced by sulphoxyl. Since also sulphoxyl takes the place of hydrogen, and itself contains hydrogen, the sulpho-acids are equal to a hydrocarbon + SO3, just as every organic (carboxylic) acid is equal to a hydrocarbon + CO2. Moreover, here this relation corresponds with actual fact, because many sulphonic acids are obtained by the direct combination of sulphuric anhydride: C6H5,(SO3H) = C6H6+ SO3. The sulphonic acids give soluble barium salts, and are therefore easily distinguished from sulphuric acid. They are soluble in water, are not volatile, and when distilled give sulphurous anhydride (whilst the hydroxyl previously in combination with the sulphurous anhydride remains in the hydrocarbon group; thus phenol, C6H5.OH, is obtained from benzenesulphonic acid), and they are very energetic, because the hydrogen acting in them is of the same nature as in sulphuric acid itself.[60]
Sulphuric acid, as containing a large proportion of oxygen, is asubstance which frequently acts as an oxidising agent: in which case it isdeoxidised, forming sulphurous anhydrideand water (or even, although more rarely, sulphuretted hydrogen and sulphur). Sulphuric acid acts in this manner on charcoal, copper, mercury, silver, organic and other substances, which are unable to evolve hydrogen from it directly, as we saw in describing sulphurous anhydride.
Although the hydrate of a higher saline form of oxidation (ChapterXV.), sulphuric anhydride is capable of further oxidation, and forms a kind of peroxide, just as hydrogen gives hydrogen peroxide in addition to water, or as sodium and potassium, besides the oxides Na2O and K2O, give their peroxides, compounds which are in a chemical sense unstable, powerfully oxidising, and not directly able to enter into saline combinations. If the oxides of potassium, barium, &c., be compared to water, then their peroxides must in like manner correspond to hydrogen peroxide,[61]not only because the oxygen contained in them is very mobile and easily liberated, and because their reactions are similar, but also because they can be mutually transformed into each other, and are able to form compounds with each other, with bases and with water, and indeed form a kind of peroxide salts.[62]This is also the character ofpersulphuric acid, discovered in 1878 by Berthelot, and its corresponding anhydride or peroxide of sulphur S2O7. It is formed from 2SO3+ O with the absorption of heat (-27 thousand heat units), like ozone from O2+ O (-29 thousand units of heat), or hydrogen peroxide from H2O + O (-21 thousand heat units).
Peroxide of sulphur is produced by the action of a silent discharge upon a mixture of oxygen and sulphurous anhydride.[63]With waterS2O7gives persulphuric acid, H2S2O8. The latter is obtained more simply by mixing strong sulphuric acid (not weaker than H2SO4,2H2O) directly with hydrogen peroxide, or by the action of a galvanic current on sulphuric acid mixed with a certain amount of water, and cooled, the electrodes being platinum wires, when persulphuric acid naturally appears at the positive pole.[64]When an acid of the strength H2SO4,6H2O is taken, at first the hydrate of the sulphuric peroxide, S2O7,H2O only is formed; but when the concentration about the positive pole reaches H2SO4,3H2O, a mixture of hydrogen peroxide and the hydrate of sulphuric peroxide begins to be formed. Dilute solutions of sulphuric peroxide can be kept better than more concentrated solutions, but the latter may be obtained containing as much as 123 grams of the peroxide to a litre. It is a very instructive fact that hydrogen peroxide is always formed when strong solutions of persulphuric acid break up on keeping. So that the bond between the two peroxides is established both by analysis and synthesis: hydrogen peroxide is able to produce S2H2O8, and the latter to produce hydrogen peroxide. A mixture of sulphuric peroxide with sulphuric acid or water is immediately decomposed, with the evolution of oxygen, either when heated or under the action of spongy platinum. The same thingtakes place with a solution of baryta, although at first no precipitate is formed and the decomposition of the barium salt, BaS2O8, with the formation of BaSO4, only proceeds slowly, so that the solution may be filtered (the barium salt of persulphuric acid is soluble in water). Mercury, ferrous oxide, and the stannous salts, are oxidised by S2H2O8. These are all distinct signs of true peroxides. The same common properties (capacity for oxidising, property of forming peroxide of hydrogen, &c.) are possessed by the alkali salts of persulphuric acid, which are obtained by the action of an electric current upon certain sulphates, for instance ammonium or potassium sulphate. The ammonium salt of persulphuric acid, (NH4)2S2O8, is especially easily formed by this means, and is now prepared on a large scale and used (like Na2O2and H2O2) for bleaching tissues and fibres.[65]
In order to understand the relation of sulphuric peroxide to sulphuric acid we must first remark that hydrogen peroxide is to be considered, in accordance with the law of substitution, as water, H(OH), in which H is replaced by (OH). Now the relation of H2S2O8to H2SO4is exactly similar. The radicle of sulphuric acid, equivalent to hydrogen, is HSO4;[65 bis]it corresponds with the (OH) of water, and therefore sulphuric acid, H(SHO4), gives (SHO4)2or S2H2O8, in exactly the same manner as water gives (HO)2—i.e.H2O2.[66]
The largest partof the sulphuric acid madeis used for reacting on sodium chloride in the manufacture of sodium carbonate; for the manufacture of the volatile acids, like nitric, hydrochloric, &c., from their corresponding salts; for the preparation of ammonium sulphate, alums, vitriols (copper and iron), artificial manures, superphosphate (Chapter XIX., Note18) and other salts of sulphuric acid; in the treatment of bone ash for the preparation of phosphorus, and for the solution of metals—for example, of silver in its separation from gold—forcleaning metals from rust, &c. A large amount of oil of vitriol is also used in treatment of organic substances; it is used for the extraction of stearin, or stearic acid, from tallow, for refining petroleum and various vegetable oils, in the preparation of nitro-glycerine (Chapter VI., Notes37and37 bis), for dissolving indigo and other colouring matters, for the conversion of paper into vegetable parchment, for the preparation of ether from alcohol, for the preparation of various artificial scents from fusel oil, for the preparation of vegetable acids, such as oxalic, tartaric, citric, for the conversion of non-fermentable starchy substances into fermentable glucose, and in a number of other processes. It would be difficult to find another artificially-prepared substance which is so frequently applied in the arts as sulphuric acid. Where there are not works for its manufacture, the economical production of many other substances of great technical importance is impossible. In those localities which have arrived at a high technical activity the amount of sulphuric acid consumed is proportionally large; sulphuric acid, sodium carbonate, and lime are the most important of the artificially-prepared agents employed in factories.
Besides the normal acids of sulphur, H2SO3, H2SO3S, and H2SO4, corresponding with sulphuretted hydrogen, H2S, in the same way that the oxy-acids of chlorine correspond with hydrochloric acid, HCl, there exists a peculiar series of acids which are termedthionic acids. Their general composition is SnH2O6, wherenvaries from 2 to 5. Ifn= 2, the acid is called dithionic acid. The others are distinguished as trithionic, tetrathionic, and pentathionic acids. Their composition, existence, and reactions are very easily understood if they be referred to the class of the sulphonic acids—that is, if their relation to sulphuric acid be expressed in just the same manner as the relation of the organic acids to carbonic acid. The organic acids, as we saw (ChapterIX.), proceed from the hydrocarbons by the substitution of their hydrogen by carboxyl—that is, by the radicle of carbonic acid, CH2O3- HO = CHO2. The formation of the acids of sulphur by means of sulphoxyl may be represented in the same manner, HSO3= H2SO4- HO. Therefore to hydrogen H2, there should correspond the acids H.SHO3, sulphurous, and SHO3.SHO3= S2H2O6, or dithionic; to SH2there should correspond the acids SH(SHO3) = H2S2O3(thiosulphuric), and S(SHO3)2= H2S3O6(trithionic); to S2H2the acids S2H(SHO3) = H2S3O2(unknown), and S2(SHO3)2= H2S4O6(tetrathionic); to S3H2the acids S3H(SHO3) and S3(SHO3)2= H2S5O6(pentathionic). We know that iodine reacts directly with the hydrogen of sulphuretted hydrogen and combines with it, and if thiosulphuric acid contains the radicle of sulphuretted hydrogen (or hydrogen unitedwith sulphur) of the same nature as in sulphuretted hydrogen, it is not surprising that iodine reacts with sodium thiosulphate and forms sodium tetrathionate. Thus, thiosulphuric acid, HS(SHO3), when deprived of H, gives a radicle which immediately combines with another similar radicle, forming the tetrathionate S2(SO2HO)2. On this view[67]of the structure of the thionic acids and salts, it is also clear how all the thionic acids, like thiosulphuric acid, easily give sulphur and sulphides, with the exception only of dithionic acid, H2S2O6, which, judging from the above, stands apart from the series of the other thionic acids. Dithionic acid stands in the same relation to sulphuric acid as oxalic acid does to carbonic acid. Oxalic acid is dicarboxyl, (CHO2)2= C2H2O4, and so also dithionic acid is disulphoxyl, (SHO3)2= S2H2O6. Oxalic acid when ignited decomposes into carbonic anhydride and carbonic oxide, CO, and dithionic acid when heated decomposes into sulphuric anhydride and sulphurous anhydride, SO2, and SO2stands in the same relation to SO3as CO to CO2. This also explains the peculiarity of the calcium, barium, and lead, &c. salts of the thionic acids being easily soluble (although the corresponding salts of H2SO3, H2SO4, and H2S dissolve with difficulty), because the former are similar to the salts of the sulphonic acids, which are also soluble in water. Thus the thionic acids aredisulphonic acids, just as many dicarboxylic acids are known—for example, CH2(CO2H)2, C6H4(CO2H)2.[68]
Sulphur exhibits an acid character, not only in its compounds with hydrogen and oxygen, but also in those with other elements. The compound of sulphur and carbon has been particularly well investigated. It presents a great analogy to carbonic anhydride, both in its elementary composition and chemical character. This substance is the so-called carbon bisulphide, CS2, and corresponds with CO2.
The first endeavours to obtain a compound of sulphur with carbon were unsuccessful, for although sulphur does combine directly with carbon, yet the formation of this compound requires distinctly definite conditions. If sulphur be mixed with charcoal and heated, it is simply driven off from the latter, and not the smallest trace of carbon bisulphide is obtained. The formation of this compound requires that the charcoal should be first heated to a red heat, but not above, and then either the vapour of sulphur passed over it or lumps of sulphur thrown on to the red-hot charcoal, but in small quantities, so as not to lower the temperature of the latter. If the charcoal be heated to a white heat, the amount of carbon bisulphide formed is less. This depends, in the first place, on the carbon bisulphide dissociating at a high temperature.[69]In the second place, Favre and Silberman showed that in the combustion of one gram of carbon bisulphide (the products will be CO2+ 2SO2) 3,400 heat units are evolved—that is, the combustion of a molecular quantity of carbon bisulphide evolves 258,400 heat units (according to Berthelot, 246,000). From a molecule of carbon bisulphide in grams we may obtain 12 grams of carbon, whose combustion evolves 96,000 heat units, and 64 grams of sulphur, evolving by combustion (into SO2) 140,800 heat units. Hence we see that the component elements separately evolve less heat by their combustion (237,000 heat units) than carbon bisulphide itself—that is,that heat should be evolved (at the ordinary temperature) and not absorbed in its decomposition, and therefore that the formation of carbon bisulphide from charcoal and sulphur is in all probability accompanied by an absorption of heat.[70]It is therefore not surprising that, like other compounds produced with an absorption of heat (ozone, nitrous oxide, hydrogen peroxide, &c.), carbon bisulphide is unstable and easily converted into the original substances from which it is obtained. And indeed if the vapour of carbon bisulphide be passed through a red-hot tube, it is decomposed—that is, it dissociates—into sulphur and carbon. And this takes place at the temperature at which this substance is formed, just as water decomposes into hydrogen and oxygen at the temperature of its formation. In this absorption of heat in the formation of carbon bisulphide is explained the facility with which it suffers reactions of decomposition, which we shall see in the sequel, and its main difference from the closely analogous carbonic anhydride.