[2] Anhydrous (ignited) sodium sulphate, Na2SO4, is known in trade as ‘sulphate’ or salt-cake, in mineralogy thenardite. Crystalline decahydrated salt is termed in mineralogy mirabilite, and in trade Glauber's salt. On fusing it, the monohydrate Na2SO4H2O is obtained, together with a supersaturated solution.
[3] The salts may be obtained not only by methods of substitution of various kinds, but also by many other combinations. Thus sodium sulphate may be formed from sodium oxide and sulphuric anhydride, by oxidising sodium sulphide, Na2S, or sodium sulphite, Na2SO3, &c. When sodium chloride is heated in a mixture of the vapours of water, air, and sulphurous anhydride, sodium sulphate is formed. According to this method (patented by Hargreaves and Robinson), sodium sulphate, Na2SO4, is obtained from NaCl without the preliminary manufacture of H2SO4. Lumps of NaCl pressed into bricks are loosely packed into a cylinder and subjected, at a red heat, to the action of steam, air and SO2. Under these conditions, HCl, sulphate, and a certain amount of unaltered NaCl are obtained. This mixture is converted into soda by Gossage's process (see Note [15]) and may have some practical value.
[4] Many observations have been made, but little general information has been obtained from particular cases. In addition to which, the properties of a given salt are changed by the presence of other salts. This takes place not only in virtue of mutual decomposition or formation of double salts capable of separate existence, but is determined by the influence which some salts exert on others, or by forces similar to those which act during solution. Here nothing has been generalised to that extent which would render it possible to predict without previous investigation, if there be no close analogy to help us. Let us state one of these numerous cases: 100 parts of water at 20° dissolve 34 parts of potassium nitrate but on the addition of sodium nitrate the solubility of potassium nitrate increases to 48 parts in 10 of water (Carnelley and Thomson). In general, in all cases of which there are accurate observations it appears that the presence of foreign salts changes the properties of any given salt.
[5] The information concerning solubility (Chapter [I].) is given according to the determinations of Gay-Lussac, Lovell, and Mulder.
[6] In Chapter I., Note [24], we have already seen that with many other sulphates the solubility also decreases after a certain temperature is passed. Gypsum, CaSO4,2H2O, lime, and many other compounds present such a phenomenon. An observation of Tilden's (1884) is most instructive; he showed that on raising the temperature (in closed vessels) above 140° the solubility of sodium sulphate again begins to increase. At 100° 100 parts of water dissolve about 43 parts of anhydrous salt, at 140° 42 parts, at 160° 43 parts, at 180° 44 parts, at 230° 46 parts. According to Étard (1892) the solubility of 30 parts of Na2SO4 in 100 of solution (or 43 per 100 of water) corresponds to 80°, and above 240° the solubility again falls, and very rapidly, so that at 320° the solution contains 12 per 100 of solution (about 14 per 100 of water) and a further rise of temperature is followed by a further deposition of the salt. It is evident that the phenomenon of saturation, determined by the presence of an excess of the dissolved substance, is very complex, and therefore that for the theory of solutions considered as liquid indefinite chemical compounds, many useful statements can hardly be given.
[7] Already referred to in Chapter I., Note [56].
The example of sodium sulphate is historically very important for the theory of solutions. Notwithstanding the number of investigations which have been made, it is still insufficiently studied, especially from the point of the vapour tension of solutions and crystallo-hydrates, so that those processes cannot be applied to it which Guldberg, Roozeboom, Van't Hoff, and others applied to solutions and crystallo-hydrates. It would also be most important to investigate the influence of pressure on the various phenomena corresponding with the combinations of water and sodium sulphate, because when crystals are separated—for instance, of the decahydrated salt—an increase of volume takes place, as can be seen from the following data:—the sp. gr. of the anhydrous salt is 2·66, that of the decahydrated salt = 1·46, but the sp. gr. of solutions at 15°/4° = 9,992 + 90·2p + 0·35p2 where p represents the percentage of anhydrous salt in the solution, and the sp. gr. of water at 4° = 10,000. Hence for solutions containing 20 p.c. of anhydrous salt the sp. gr. = 1·1936; therefore the volume of 100 grams of this solution = 83·8 c.c., and the volume of anhydrous salt contained in it is equal to 20/2·66, or = 7·5 c.c., and the volume of water = 80·1 c.c. Therefore, the solution, on decomposing into anhydrous salt and water, increases in volume (from 83·8 to 87·6); but in the same way 83·8 c.c. of 20 p.c. solution are formed from (45·4/1·46 =) 31·1 c.c. of the decahydrated salt, and 54·6 c.c. of water—that is to say, that during the formation of a solution from 85·7 c.c., 83·8 c.c. are formed.
[8] From this example it is evident the solution remains unaltered until from the contact of a solid it becomes either saturated or supersaturated, crystallisation being determined by the attraction to a solid, as the phenomenon of supersaturation clearly demonstrates. This partially explains certain apparently contradictory determinations of solubility. The best investigated example of such complex relations is cited in Chapter XIV., Note [50] (for CaCl2).
[9] According to Pickering's experiments (1886), the molecular weight in grams (that is, 142 grams) of anhydrous sodium sulphate, on being dissolved in a large mass of water, at 0° absorbs (hence the - sign) -1,100 heat units, at 10°–700, at 15°–275, at 20° gives out +25, at 25° +300 calories. For the decahydrated salt, Na2SO4,10H2O, 5° - 4,225, 10° - 4,000, 15° - 3,570, 20° - 3,160, 25° - 2,775. Hence (just as in Chapter I., Note [56]) the heat of the combination Na2SO4,10H2O at 5° = +3,125, 10° = +3,250, 20° = +3,200, and 25° = +3,050.
It is evident that the decahydrated salt dissolving in water gives a decrease of temperature. Solutions in hydrochloric acid give a still greater decrease, because they contain the water of crystallisation in a solid state—that is, like ice—and this on melting absorbs heat. A mixture of 15 parts of Na2SO4,10H2O and 12 parts of strong hydrochloric acid produces sufficient cold to freeze water. During the treatment with hydrochloric acid a certain quantity of sodium chloride is formed.