[60] We will mention the following difference between the sulphonic acids and the ethereal acid sulphates (Note [59]): 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 BaO2 with 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] Anhydrous sulphuric 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-called acetic 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., Note [26]).
[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. K2S2O8 is 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 MnO2 and 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 (Note [62]), 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.