The power of magnesium salts to form double and basic salts is very frequently shown in reactions, and is specially marked as regards ammonium salts. If saturated solutions of magnesium and ammonium sulphates are mixed together, a crystalline double salt Mg(NH4)2(SO4)2,6H2O,[20] is immediately precipitated. A strong solution of ordinary ammonium carbonate dissolves magnesium oxide or carbonate, and precipitates crystals of a double salt, Mg(NH4)2(CO3)2,4H2O, from which water extracts the ammonium carbonate. With an excess of an ammonium salt the double salt passes into solution,[21] and therefore if a solution contain a magnesium salt and an excess of an ammonium salt—for instance, sal-ammoniac—then sodium carbonate will no longer precipitate magnesium carbonate. A mixture of solutions of magnesium and ammonium chlorides, on evaporation or refrigeration, gives a double salt, Mg(NH4)Cl3,6H2O.[22] The salts of potassium, like those of ammonium, are able to enter into combination with the magnesium salts.[23] For instance, the double salt, MgKCl3,6H2O, which is known as carnallite,[24] and occurs in the salt mines of Stassfurt, may be formed by freezing a saturated solution of potassium chloride with an excess of magnesium chloride. A saturated solution of magnesium sulphate dissolves potassium sulphate, and solid magnesium sulphate is soluble in a saturated solution of potassium sulphate. A double salt, K2Mg(SO4)2,6H2O, which closely resembles the above-mentioned ammonium salt, crystallises from these solutions.[25] The nearest analogues of magnesium are able to give exactly similar double salts, both in crystalline form (monoclinic system) and composition; they, like this salt (see Chapter XV.), are easily able (at 140°) to part with all their water of crystallisation, and correspond with the salts of sulphuric acid, whose type may be taken as magnesium sulphate, MgSO4.[26] It occurs at Stassfurt as kieserite, MgSO4,H2O, and generally separates from solutions as a heptahydrated salt, MgSO4,7H2O, and from supersaturated solutions as a hexahydrated salt, MgSO4,6H2O; at temperatures below 0° it crystallises out as a dodecahydrated salt, MgSO4,12H2O, and a solution of the composition MgSO4,2H2O solidifies completely at -5°.[27] Thus between water and magnesium sulphate there may exist several definite and more or less stable degrees of equilibrium; the double salt MgSO4K2SO4,6H2O may be regarded as one of these equilibrated systems, the more so since it contains 6H2O, whilst MgSO4 forms its most stable system with 7H2O, and the double salt may be considered as this crystallo-hydrate in which one molecule of water is replaced by the molecule K2SO4.[28]
The power of forming basic salts is a very remarkable peculiarity of magnesia and other feeble bases, and especially of those corresponding with polyvalent metals. The very powerful bases corresponding with univalent metals—like potassium and sodium—do not form basic salts, and, indeed, are more prone to give acid salts, whilst magnesium easily and frequently forms basic salts, especially with feeble acids, although there are some oxides—as, for example, copper and lead oxides—which still more frequently give basic salts. If a cold solution of magnesium sulphate be mixed with a solution of sodium carbonate there is formed a gelatinous precipitate of a basic salt, Mg(HO)2,4MgCO3,9H2O; but all the magnesia is not precipitated in this case, as a portion of it remains in solution as an acid double salt. If sodium carbonate be added to a boiling solution of magnesium sulphate a precipitate of a still more basic salt is formed, 4MgSO4 + 4Na2CO3 + 4H2O = 4Na2SO4 + CO2 + Mg(OH)2,3MgCO3,3H2O. This basic salt forms the ordinary drug magnesia (magnesia alba), in the form of light porous lumps. Other basic salts are formed under certain modifications of temperature and conditions of decomposition. But the normal salt, MgCO3, which occurs in nature as magnesite in the form of rhombohedra of specific gravity 3·056, cannot be obtained by such a method of precipitation. In fact, the formation of the different basic salts shows the power of water to decompose the normal salt. It is possible, however, to obtain this salt both in an anhydrous and hydrated state. A solution of magnesium carbonate in water containing carbonic acid is taken for this purpose. The reason for this is easily understood—carbonic anhydride is one of the products of the decomposition of magnesium carbonate in the presence of water. If this solution be left to evaporate spontaneously the normal salt separates in a hydrated form, but in the evaporation of a heated solution, through which a stream of carbonic anhydride is passed, the anhydrous salt is formed as a crystalline mass, which remains unaltered in the air, like the natural mineral.[29] The decomposing influence of water on the salts of magnesium, which is directly dependent on the feeble basic properties of magnesia,[30] is most clearly seen in magnesium chloride, MgCl2. This salt is contained[31] in the last mother-liquors of the evaporation of sea-water. On cooling a sufficiently concentrated solution, the crystallo-hydrate, MgCl2,6H2O, separates;[32] but if it be further heated (above 106°) to remove the water, then hydrochloric acid passes off together with the latter, so that there ultimately remains magnesia with a small quantity of magnesium chloride.[33] From what has been said it is evident that anhydrous magnesium chloride cannot be obtained by simple evaporation. But if sal-ammoniac or sodium chloride be added to a solution of magnesium chloride, then the evolution of hydrochloric acid does not take place, and after complete evaporation the residue is perfectly soluble in water. This renders it possible to obtain anhydrous magnesium chloride from its aqueous solution. Indeed the mixture with sal-ammoniac (in excess) may be dried (the residue consists of an anhydrous double salt, MgCl2,2NH4Cl) and then ignited (460°), when the sal-ammoniac is converted into vapour and a fused mass of anhydrous magnesium chloride remains behind. The anhydrous chloride evolves a very considerable amount of heat on the addition of water, which shows the great affinity the salt has for water.[34] Anhydrous magnesium chloride is not only obtained by the above method, but is also formed by the direct combination of chlorine and magnesium, and by the action of chlorine on magnesium oxide, oxygen being evolved; this proceeds still more easily by heating magnesia with charcoal in a stream of chlorine, when the charcoal serves to take up the oxygen. This latter method is also employed for the preparation of chlorides which are formed in an anhydrous condition with still greater difficulty than magnesium chloride. Anhydrous magnesium chloride forms a colourless, transparent mass, composed of flexible crystalline plates of a pearly lustre. It fuses at a low red heat (708°) into a colourless liquid, remains unchanged in a dry state, but under the action of moisture is partially decomposed even at the ordinary temperature, with formation of hydrochloric acid. When heated in the presence of oxygen (air) it gives chlorine and the basic salt, which is formed with even greater facility under the action of heat in the presence of steam, when HCl is formed, according to the equation 2MgCl2 + H2O = MgOMgCl2 + 2HCl.[34 bis]
Calcium (or the metal of lime) and its compounds in many respects present a great resemblance to magnesium compounds, but are also clearly distinguished from them by many properties.[35] In general, calcium stands to magnesium in the same relation as potassium occupies in respect to sodium. Davy obtained metallic calcium, like potassium, as an amalgam by the action of a galvanic current; but neither charcoal nor iron decomposes calcium oxide, and even sodium decomposes calcium chloride[36] with difficulty. But a galvanic current easily decomposes calcium chloride, and metallic sodium somewhat easily decomposes calcium iodide when heated. As in the case of hydrogen, potassium, and magnesium, the affinity of iodine for calcium is feebler than that of chlorine (and oxygen), and therefore it is not surprising that calcium iodide may be subjected to that decomposition, which the chloride and oxide undergo with difficulty.[37] Metallic calcium is of a yellow colour, and has a considerable lustre, which it preserves in dry air. Its specific gravity is 1·58. Calcium is distinguished by its great ductility; it melts at a red heat and then burns in the air with a very brilliant flame; the brilliancy is due to the formation of finely divided infusible calcium oxide. Judging from the fact that calcium in burning gives a very large flame, it is probable that this metal is volatile. Calcium decomposes water at the ordinary temperature, and is oxidised in moist air, but not so rapidly as sodium. In burning, it gives its oxide or lime, CaO, a substance which is familiar to every one, and of which we have already frequently had occasion to speak. This oxide is not met with in nature in a free state, because it is an energetic base which everywhere encounters acid substances forming salts with them. It is generally combined with silica, or occurs as calcium carbonate or sulphate. The carbonate and nitrate are decomposed, at a red heat, with the formation of lime. As a rule, the carbonate, which is so frequently met with in nature, serves as the source of the calcium oxide, both commercial and pure. When heated, calcium carbonate dissociates: CaCO3 = CaO + CO2. In practice the decomposition is conducted at a bright red heat, in the presence of steam, or a current of a foreign gas, in heaps or in special kilns.[38]
Fig. 78.—Continually-acting kiln for burning lime. The lime is charged from above and calcined by four lateral grates, R, M. D, fire-bars. B, space for withdrawing the burnt lime. K, stoke-house. M. fire grate. Q, R, under-grate.
Calcium oxide—that is, quicklime—is a substance (sp. gr. 3·15) which is unaffected by heat,[39] and may therefore serve as a fire-resisting material, and was employed by Deville for the construction of furnaces in which platinum was melted, and silver volatilised by the action of the heat evolved by the combustion of detonating gas. The hydrated lime, slaked lime, or calcium hydroxide, CaH2O2 (specific gravity 2·07) is a most common alkaline substance, employed largely in building for making mortars or cements, in which case its binding property is mainly due to the absorption of carbonic anhydride.[40] Lime, like other alkalis, acts on many animal and vegetable substances, and for this reason has many practical uses—for example, for removing fats, and in agriculture for accelerating the decomposition of organic substances in the so-called composts or accumulations of vegetable and animal remains used for fertilising land. Calcium hydroxide easily loses its water at a moderate heat (530°), but it does not part with water at 100°. When mixed with water, lime forms a pasty mass known as slaked lime and in a more dilute form as milk of lime, because when shaken up in water it remains suspended in it for a long time and presents the appearance of a milky liquid. But, besides this, lime is directly soluble in water, not to any considerable extent, but still in such a quantity that lime water is precipitated by carbonic anhydride, and has clearly distinguishable alkaline properties. One part of lime requires at the ordinary temperature about 800 parts of water for solution. At 100° it requires about 1500 parts of water, and therefore lime-water becomes cloudy when boiled. If lime-water be evaporated in a vacuum, calcium hydroxide separates in six-sided crystals.[41] If lime-water be mixed with hydrogen peroxide minute crystals of calcium peroxide, CaO2,8H2O, separate; this compound is very unstable and, like barium peroxide, is decomposed by heat. Lime, as a powerful base, combines with all acids, and in this respect presents a transition from the true alkalis to magnesia. Many of the salts of calcium (the carbonate, phosphate, borate, and oxalate) are insoluble in water; besides which the sulphate is only sparingly soluble. As a more energetic base than magnesia, lime forms salts, CaX2, which are distinguished by their stability in comparison with the salts MgX2; neither does lime so easily form basic and double salts as magnesia.
Anhydrous lime does not absorb dry carbonic anhydride at the ordinary temperature. This was already known by Scheele, and Prof. Schuliachenko showed that there is no absorption even at 360°. It only proceeds at a red heat,[42] and then only leads to the formation of a mixture of calcium oxide and carbonate (Rose). But if the lime be slaked or dissolved, the absorption of carbonic anhydride proceeds rapidly and completely. These phenomena are connected with the dissociation of calcium carbonate, studied by Debray (1867) under the influence of the conceptions of dissociation introduced into science by Henri Saint-Claire Deville. Just as there is no vapour tension for non-volatile substances, so there is no dissociation tension of carbonic anhydride for calcium carbonate at the ordinary temperature. Just as every volatile substance has a maximum possible vapour tension for every temperature, so also calcium carbonate has its corresponding dissociation tension; this at 770° (the boiling point of cadmium) is about 85 mm. (of the mercury column), and at 930° (the boiling point of Zn) it is about 520 mm. As, if the tension be greater, there will be no evaporation, so also there will he no decomposition. Debray took crystals of calc spar, and could not observe the least change in them at the boiling point of zinc (930°) in an atmosphere of carbonic anhydride taken at the atmospheric pressure (760 mm.), whilst on the other hand calcium carbonate may be completely decomposed at a much lower temperature if the tension of the carbonic anhydride be kept below the dissociation tension, which may be done either by directly pumping away the gas with an air-pump, or by mixing it with some other gas—that is, by diminishing the partial pressure of the carbonic anhydride,[43] just as an object may be dried at the ordinary temperature by removing the aqueous vapour or by carrying it off in a stream of another gas. Thus it is possible to obtain calcium carbonate from lime and carbonic anhydride at a certain temperature above that at which dissociation begins, and conversely to decompose calcium carbonate at the same temperature into lime and carbonic anhydride.[44] At the ordinary temperature the reaction of the first order (combination) cannot proceed because the second (decomposition, dissociation) cannot take place, and thus all the most important phenomena with respect to the behaviour of lime towards carbonic anhydride are explained by starting from one common basis.[45]
Calcium carbonate, CaCO3, is sometimes met with in nature in a crystalline form, and it forms an example of the phenomenon termed dimorphism—that is, it appears in two crystalline forms. When it exhibits combinations of forms belonging to the hexagonal system (six-sided prisms, rhombohedra, &c.) it is called calc spar. Calc spar has a specific gravity of 2·7, and is further characterised by a distinct cleavage along the planes of the fundamental rhombohedron having an angle of 105°. Perfectly transparent Iceland spar presents a clear example of double refraction (for which reason it is frequently employed in physical apparatus). The other form of calcium carbonate occurs in crystals belonging to the rhombic system, and it is then called aragonite; its specific gravity is 3·0. If calcium carbonate be artificially produced by slow crystallisation at the ordinary temperature, it appears in the rhombohedral form, but if the crystallisation be aided by heat it then appears as aragonite. It may therefore be supposed that calc spar presents the form corresponding with a low temperature, and aragonite with a higher temperature during crystallisation.[46]