[88] By the action of water on iodine monochloride and trichloride a compound IHCl2 is obtained, which does not seem to be altered by water. Besides this compound, iodine and iodic acid are always formed, 10ICl + 3H2O = HIO3 + 5IHCl2 + 2I2; and in this respect iodine trichloride may be regarded as a mixture, ICl + ICl5 = 2ICl3, but ICl5 + 3H2O = IHO3 + 5HCl; hence iodic acid, iodine, the compound IHCl2, and hydrochloric acid are also formed by the action of water.

CHAPTER XII
SODIUM

The neutral salt, sodium sulphate, Na2SO4, obtained when a mixture of sulphuric acid and common salt is strongly heated (Chapter [X].),[1] forms a colourless saline mass consisting of fine crystals, soluble in water. It is the product of many other double decompositions, sometimes carried out on a large scale; for example, when ammonium sulphate is heated with common salt, in which case the sal-ammoniac is volatilised, &c. A similar decomposition also takes place when, for instance, a mixture of lead sulphate and common salt is heated; this mixture easily fuses, and if the temperature be further raised heavy vapours of lead chloride appear. When the disengagement of these vapours ceases, the remaining mass, on being treated with water, yields a solution of sodium sulphate mixed with a solution of undecomposed common salt. A considerable quantity, however, of the lead sulphate remains unchanged during this reaction, PbSO4 + 2NaCl = PbCl2 + Na2SO4, the vapours will contain lead chloride, and the residue will contain the mixture of the three remaining salts. The cause and nature of the reaction are just the same as were pointed out when considering the action of sulphuric acid upon NaCl. Here too it may be shown that the double decomposition is determined by the removal of PbCl2 from the sphere of the action of the remaining substances. This is seen from the fact that sodium sulphate, on being dissolved in water and mixed with a solution of any lead salt (and even with a solution of lead chloride, although this latter is but sparingly soluble in water), immediately gives a white precipitate of lead sulphate. In this case the lead takes up the elements of sulphuric acid from the sodium sulphate in the solutions. On heating, the reverse phenomenon is observed. The reaction in the solution depends upon the insolubility of the lead sulphate, and the decomposition which takes place on heating is due to the volatility of the lead chloride. Silver sulphate, Ag2SO4, in solution with common salt, gives silver chloride, because the latter is insoluble in water, Ag2SO4 + 2NaCl = Na2SO4 + 2AgCl. Sodium carbonate, mixed in solution with the sulphates of iron, copper, manganese, magnesium, &c., gives in solution sodium sulphate, and in the precipitate a carbonate of the corresponding metal, because these salts of carbonic acid are insoluble in water; for instance, MgSO4 + Na2CO3 = Na2SO4 + MgCO3. In precisely the same way sodium hydroxide acts on solutions of the majority of the salts of sulphuric acid containing metals, the hydroxides of which are insoluble in water—for instance, CuSO4 + 2NaHO = Cu(HO)2 + Na2SO4. Sulphate of magnesium, MgSO4, on being mixed in solution with common salt, forms, although not completely, chloride of magnesium, and sodium sulphate. On cooling the mixture of such (concentrated) solutions sodium sulphate is deposited, as was shown in Chapter [X]. This is made use of for preparing it on the large scale in works where sea-water is treated. In this case, on cooling, the reaction 2NaCl + MgSO4 = MgCl2 + Na2SO4 takes place.

Thus where sulphates and salts of sodium are in contact, it may be expected that sodium sulphate will be formed and separated if the conditions are favourable; for this reason it is not surprising that sodium sulphate is often found in the native state. Some of the springs and salt lakes in the steppes beyond the Volga, and in the Caucasus, contain a considerable quantity of sodium sulphate, and yield it by simple evaporation of the solutions. Beds of this salt are also met with; thus at a depth of only 5 feet, about 38 versts to the east of Tiflis, at the foot of the range of the ‘Wolf's mane'’ (Voltchia griva) mountains, a deep stratum of very pure Glauber's salt, Na2SO4,10H2O, has been found.[2] A layer two metres thick of the same salt lies at the bottom of several lakes (an area of about 10 square kilometres) in the Kouban district near Batalpaschinsk, and here its working has been commenced (1887). In Spain, near Arangoulz and in many parts of the Western States of North America, mineral sodium sulphate has likewise been found, and is already being worked.

The methods of obtaining salts by means of double decomposition from others already prepared are so general, that in describing a given salt there is no necessity to enumerate the cases hitherto observed of its being formed through various double decompositions.[3] The possibility of this occurrence ought to be foreseen according to Berthollet's doctrine from the properties of the salt in question. On this account it is important to know the properties of salts; all the more so because up to the present time those very properties (solubility, formation of crystallo-hydrates, volatility, &c.) which may be made use of for separating them from other salts have not been generalised.[4] These properties as yet remain subjects for investigation, and are rarely to be foreseen. The crystallo-hydrate of the normal sodium sulphate, Na2SO4,10H2O, very easily parts with water, and may be obtained in an anhydrous state if it be carefully heated until the weight remains constant; but if heated further, it partly loses the elements of sulphuric anhydride. The normal salt fuses at 843° (red heat), and volatilises to a slight extent when very strongly heated, in which case it naturally decomposes with the evolution of SO3. At 0° 100 parts of water dissolve 5 parts of the anhydrous salt, at 10° 9 parts, at 20° 19·4, at 30° 40, and at 34° 55 parts, the same being the case in the presence of an excess of crystals of Na2SO4,10H2O.[5] At 34° the latter fuses, and the solubility decreases at higher temperatures.[6] A concentrated solution at 34° has a composition nearly approaching to Na2SO4 + 14H2O, and the decahydrated salt contains 78·9 of the anhydrous salt combined with 100 parts of water. From the above figures it is seen that the decahydrated salt cannot fuse without decomposing,[7] like hydrate of chlorine, Cl2,8H2O (Chapter XI., Note [10]). Not only the fused decahydrated salt, but also the concentrated solution at 34° (not all at once, but gradually), yields the monohydrated salt, Na2SO4,H2O. The heptahydrated salt, Na2SO4,7H2O, also splits up, even at low temperatures, with the formation of this monohydrated salt, and therefore from 35° the solubility can be given only for the latter. For 100 parts of water this is as follows: at 40° 48·8, at 50° 46·7, at 80° 43·7, at 100° 42·5 parts of the anhydrous salt. If the decahydrated salt be fused, and the solution allowed to cool in the presence of the monohydrated salt, then at 30° 50·4 parts of anhydrous salt are retained in the solution, and at 20° 52·8 parts. Hence, with respect to the anhydrous and monohydrated salts, the solubility is identical, and falls with increasing temperature, whilst with respect to decahydrated salt, the solubility rises with increasing temperature. So that if in contact with a solution of sodium sulphate there are only crystals of that heptahydrated salt (Chapter I., Note [54]), Na2SO4,7H2O, which is formed from saturated solutions, then saturation sets in when the solution has the following composition per 100 parts of salt: at 0° 19·6, at 10° 30·5, at 20° 44·7, and at 25° 52·9 parts of anhydrous salt. Above 27° the heptahydrated salt, like the decahydrated salt at 34°, splits up into the monohydrated salt and a saturated solution. Thus sodium sulphate has three curves of solubility: one for Na2SO4,7H2O (from 0° to 26°), one for Na2SO4,10H2O (from 0° to 34°), and one for Na2SO4,H2O (a descending curve beginning at 26°), because there are three of these crystallo-hydrates, and the solubility of a substance only depends upon the particular condition of that portion of it which has separated from the solution or is present in excess.[8]

Thus solutions of sodium sulphate may give crystallo-hydrates of three kinds on cooling the saturated solution: the unstable heptahydrated salt is obtained at temperatures below 26°, the decahydrated salt forms under ordinary conditions at temperatures below 34°, and the monohydrated salt at temperatures above 34°. Both the latter crystallo-hydrates present a stable state of equilibrium, and the heptahydrated salt decomposes into them, probably according to the equation 3Na2SO4,7H20 = 2Na2SO4,10H2O + Na2SO4,H2O. The ordinary decahydrated salt is called Glauber's salt. All forms of these crystallo-hydrates lose their water entirely, and give the anhydrous salt when dried over sulphuric acid.[9]

Sodium sulphate, Na2SO4, only enters into a few reactions of combination with other salts, and chiefly with salts of the same acid, forming double sulphates. Thus, for example, if a solution of sodium sulphate be mixed with a solution of aluminium, magnesium, or ferrous sulphate, it gives crystals of a double salt when evaporated. Sulphuric acid itself forms a compound with sodium sulphate, which is exactly like these double salts. It is formed with great ease when sodium sulphate is dissolved in sulphuric acid and the solution evaporated. On evaporation, crystals of the acid salt separate, Na2SO4 + H2SO4 = 2NaHSO4. This separates from hot solutions, whilst the crystallo-hydrate, NaHSO4,H2O,[10] separates from cold solutions. The crystals when exposed to damp air decompose into H2SO4, which deliquesces, and Na2SO4 (Graham, Rose); alcohol also extracts sulphuric acid from the acid salt. This shows the feeble force which holds the sulphuric acid to the sodium sulphate.[11] Both acid sodium sulphate and all mixtures of the normal salt and sulphuric acid lose water when heated, and are converted into sodium pyrosulphate, Na2S2O7, at a low red heat.[11 bis] This anhydrous salt, at a bright red heat, parts with the elements of sulphuric anhydride, the normal sodium sulphate remaining behind—Na2S2O7 = Na2SO4 + SO3. From this it is seen that the normal salt is able to combine with water, with other sulphates, and with sulphuric anhydride or acid, &c.

Sodium sulphate may by double decomposition be converted into a sodium salt of any other acid, by means of heat and taking advantage of the volatility, or by means of solution and taking advantage of the different degree of solubility of the different salts. Thus, for instance, owing to the insolubility of barium sulphate, sodium hydroxide or caustic soda may be prepared from sodium sulphate, if barium hydroxide be added to its solution, Na2SO4 + Ba(HO)2 = BaSO4 + 2NaHO. And by taking any salt of barium, BaX2, the corresponding salt of sodium may be obtained, Na2SO4 + BaX2 = BaSO4 + 2NaX. Barium sulphate thus formed, being a very sparingly-soluble salt, is obtained as a precipitate, whilst the sodium hydroxide, or salt, NaX, is obtained in solution, because all salts of sodium are soluble. Berthollet's doctrine permits all such cases to be foreseen.