As the acids derived from chlorine, phosphorus, and carbon are the oxidised hydrogen compounds of these elements, so also we can form an idea of the acid hydrates of sulphur, or of the normal acids of sulphur, by representing them as the oxidised products of sulphuretted hydrogen—
| HCl | H2S | H3P | H4C |
| HClO | H2SO(?) | H3PO(?) | H4CO |
| HClO2 | H2SO2(?) | H3PO2 | H4CO2 |
| HClO3 | H2SO3 | H3PO3 | H4CO3 |
| HClO4 | H2SO4 | H3PO4 | H4CO4[30] |
In the case of chlorine, if not all the hydrates, at all events salts of all the normal hydrates are known, whilst in the case of sulphur only the acids H2S, H2SO3 and H2SO4 are known. But, on the other hand, the latter are obtained not only as hydrates but also as stable anhydrides, SO2 and SO3, which are formed with the evolution of heat from sulphur and oxygen; 32 parts of sulphur in combining with 32 parts of oxygen—that is, in forming SO2—evolve 71,000 heat units,[31] and if the oxidation proceeds to the formation of SO3, 103,000 heat units are evolved. These figures may be compared with those which correspond with the passage of carbon into CO and CO2, when 29,000 and 97,000 units of heat are evolved. This determines the stability of the higher oxides of sulphur, and also expresses the peculiarity of sulphur as an element which, although an analogue of oxygen, forms stable compounds with it, and thus fundamentally differs from chlorine. The higher and lower oxides of chlorine are powerful oxidising agents, whilst the higher oxide of sulphur, SO3, has but feeble oxidising powers, and the lower oxide, SO2, frequently acts as a reducing agent, and is formed by the direct combustion of sulphur, just as carbonic anhydride, CO2, proceeds from the combustion of carbon. In the combustion of sulphur, and also in the oxidation (roasting) of the sulphides and polysulphides by their ignition in air, sulphurous oxide, or sulphurous anhydride, or sulphur dioxide, SO2,[31 bis] is exclusively formed. It is prepared on a large scale by burning sulphur or roasting iron pyrites or other sulphides[32] for the manufacture of sulphuric acid (Chapter [VI.]), and for direct application in the manufacture of wine or for bleaching tissues and other purposes. In the latter instances its application is based on the fact that sulphurous anhydride acts on certain vegetable matters, and has the property of a reducing and feeble acid.[32 bis]
In the laboratory—that is, on a small scale—sulphurous anhydride is best prepared by deoxidising sulphuric acid by heating it with charcoal, or copper, sulphur, mercury, &c. Charcoal produces this decomposition of sulphuric acid at but moderately high temperatures; it is itself converted into carbonic anhydride,[32 tri] and therefore when sulphuric acid is heated with charcoal it evolves a mixture of sulphurous and carbonic anhydrides: C + 2H2SO4 = CO2 + 2SO2 + 2H2O. The metals which are unable to decompose water, and which do not, therefore, expel hydrogen from sulphuric acid, are frequently capable of decomposing sulphuric acid, with the evolution of sulphurous anhydride, just as they decompose nitric acid, forming the lower oxides of nitrogen. These metals are silver, mercury, copper, lead, and others. Thus, for example, the action of copper on sulphuric acid may be expressed by the following equation: Cu + 2H2SO4 = CuSO4 + SO2 + 2H2O. In the laboratory this reaction is carried on in a flask with a gas-conducting tube, and does not take place unless aided by heat.[33]
In its physical and chemical properties sulphurous anhydride presents a great resemblance to carbonic anhydride. It is a heavy gas, somewhat considerably soluble in water, very easily condensed into a liquid; it forms normal and acid salts, does not evolve oxygen under the direct action of heat,[34] although such metals as sodium and magnesium burn in it, just as in carbonic anhydride. It has a suffocating odour, which is well known owing to its being evolved when sulphur or sulphur matches are burnt. In characterising the properties of sulphurous anhydride, it is very important to remember (Chapter [II.]) also that it is more easily liquefied (at -10°, or at 0° under two atmospheres pressure) than carbonic anhydride (thirty-six atmospheres at 0°),[35] that it is more soluble than carbonic anhydride (Vol. I. p. [79]); at 0°, 100 vols. of water dissolve 180 vols. of carbonic anhydride and 688 vols. of sulphuric anhydride), that the molecular weight of SO2 = 64 and of CO2 = 44, and that the density of liquid sulphurous anhydride at 0° = 1·43 (molecular volume = 45) and of carbonic anhydride = 0·95 (molecular volume = 49). Although sulphur dioxide is the anhydride of an acid, nevertheless, like carbonic anhydride, it does not form any stable compounds with water, but gives a solution from which it may be entirely expelled by the action of heat.[36] The acid character of sulphurous anhydride is clearly expressed by the fact that it is entirely absorbed by alkalis, with which it forms acid and normal salts easily soluble in water. With salts of barium, calcium, and the heavy metals, the normal salts of the alkalis, M2SO3, give precipitates exactly like those formed by the carbonates. In general, the salts of sulphurous acid are closely analogous to the corresponding carbonates.
Acid sodium sulphite, NaHSO3, may be obtained by passing sulphurous anhydride into a solution of sodium hydroxide. It is also formed by saturating a solution of sodium carbonate with the gas (carbonic anhydride is then given off), and as the solubility of the acid sulphite is much greater than that of the carbonate, a further quantity of the latter may be dissolved after the passage of the sulphurous anhydride, so that ultimately a very strong solution of the sulphite may be formed in this manner, from which it may be obtained in a crystalline form, either by cooling and evaporating (without heating, for then the salt would give off sulphurous anhydride) or by adding alcohol to the solution. When exposed to the air this salt loses sulphurous anhydride and attracts oxygen, which converts it into sodium sulphate. The acid sulphites of the alkali metals are able to combine not only with oxygen, but also with many other substances—for example, a solution of the sodium salt dissolves sulphur, forming sodium thiosulphate, gives crystalline compounds with the aldehydes and ketones, and dissolves many bases, converting them into double sulphites. Having the faculty of attracting or absorbing oxygen, acid sodium sulphite is also able to absorb chlorine, and is therefore employed, like sodium thiosulphate, for the removal of chloride (as an antichlor), especially in the bleaching of fabrics, when it is necessary to remove the last traces of the chlorine held in the tissues, which might otherwise have an injurious effect on them. If a solution of an alkali hydroxide be divided into two parts, and one half is saturated with sulphurous anhydride, and then the other half added to it, a normal salt will be obtained in the solution, having an alkaline reaction, like a solution of sodium carbonate. The acid salt has a neutral reaction.[36 bis] Like sodium carbonate, normal sodium sulphite has the composition Na2SO3,10H2O, and its maximum solubility is at 33°—in a word, it very closely resembles sodium carbonate. Although this salt does not give off sulphurous anhydride from its solution, it is able, like the acid salt, to absorb oxygen from the air, and is then converted into sodium sulphate.[37]
Besides the acid character we must also point out the reducing character of sulphurous anhydride. The reducing action of sulphurous acid, its anhydride and salts, is due to their faculty of passing into sulphuric acid and sulphates. The reducing action of the sulphites is particularly energetic, so that they even convert nitric oxide into nitrous oxide: K2SO3 + 2NO = K2SO4 + N2O. The salts of many of the higher oxides are converted into those of the lower—for example, FeX3 into FeX2, CuX2 into CuX, HgX2 into HgX; thus 2FeX3 + SO2 + 2H2O = 2FeX2 + H2SO4 + 2HX. In the presence of water, sulphurous anhydride is oxidised by chlorine (SO2 + 2H2O + Cl2 = H2SO4 + 2HCl), iodine, nitrous acid, hydrogen peroxide, hypochlorous acid, chloric acid, and other oxygen compounds of the halogens, chromic, manganic, and many other metallic acids and higher oxides, as well as all peroxides. Free oxygen in the presence of spongy platinum is able to oxidise sulphurous anhydride even in the absence of water, in which case sulphuric anhydride SO3 is formed, so that the latter may be prepared by passing a mixture of sulphurous anhydride and oxygen over incandescent spongy platinum, or, as it is now prepared on a large scale in chemical works, by passing this mixture over asbestos or pumice stone moistened with a solution of platinum salt and ignited. Sulphurous anhydride is completely absorbed by certain higher oxides—for instance, by barium peroxide and lead dioxide (PbO2 + SO2 = PbSO4).[38]
There are, however, cases where sulphurous anhydride acts as an oxidising agent—that is, it is deoxidised in the presence of substances which are capable of absorbing oxygen with still greater energy than the sulphurous anhydride itself. This oxidising action proceeds with the formation of sulphuretted hydrogen or of sulphides, while the reducing agent is oxidised at the expense of the oxygen of the sulphurous anhydride. In this respect, the action of stannous salts is particularly remarkable. Stannous chloride, SnCl2, in an aqueous solution gives a precipitate of stannic sulphide, SnS2, with sulphurous anhydride—that is, the latter is deoxidised to sulphuretted hydrogen, while SnX2 is oxidised into SnX4. A solution of sulphurous anhydride has also an oxidising action on zinc. The zinc passes into solution, but no hydrogen is evolved,[39] because a salt of hydrosulphurous acid, ZnS2O4, is formed. The free acid is still less stable than the salt.