Chapter XIII. These deposits also contain much _kainite_,

KMgCl(SO_{4}),3H_{2}O (sp. gr. 2·13; 100 parts of water dissolve 79·6 parts at 18°). This double salt contains two metals and two haloids. Feit (1889) also obtained a bromide corresponding to carnallite.

[25] The component parts of certain double salts diffuse at different rates, and as the diffused solution contains a different proportion of the component salts than the solution taken of the double salt, it shows that such salts are decomposed by water. According to Rüdorff, the double salts, like carnallite, MgK_{2}(SO_{4})_{2},6H_{2}O, and the alums, all belong to this order (1888). But such salts as tartar emetic, the double oxalates, and double cyanides are not separated by diffusion, which in all probability depends both on the relative rate of the diffusion of the component salts and on the degree of affinity acting between them. Those complex states of equilibrium which exist between water, the individual salts MX and NY, and the double salt MNXY, have been already partially analysed (as will be shown hereafter) in that case when the system is heterogeneous (that is, when something separates out in a solid state from the liquid solution), but in the case of equilibria in a homogeneous liquid medium (in a solution) the phenomenon is not so clear, because it concerns that very theory of solution which cannot yet be considered as established (Chapter I., Note 9, and others). As regards the heterogeneous decomposition of double salts, it has long been known that such salts as carnallite and K_{2}Mg(SO_{4})_{2} give up the more soluble salt if an insufficient quantity of water for their complete solution be taken. The complete saturation of 100 parts of water requires at 0° 14·1, at 20° 25, and at 60° 50·2 parts of the latter double salt (anhydrous), while 100 parts of water dissolve 27 parts of magnesium sulphate at 0°, 36 parts at 20°, and 55 parts at 60°, of the anhydrous salt taken. Of all the states of equilibrium exhibited by double salts the most fully investigated as yet is the system containing water, sodium sulphate, magnesium sulphate, and their double salt, Na_{2}Mg(SO_{4})_{2}, which crystallises with 4 and 6 mol. OH_{2}. The first crystallo-hydrate, MgNa_{2}(SO_{4})_{2},4H_{2}O, occurs at Stassfurt, and as a sedimentary deposit in many of the salt lakes near Astrakhan, and is therefore called _astrakhanite_. The specific gravity of the monoclinic prisms of this salt is 2·22. If this salt, in a finely divided state, be mixed with the necessary quantity of water (according to the equation MgNa_{2}(SO_{4})_{2},4H_{2}O + 13H_{2}O = Na_{2}SO_{4},10H_{2}O + MgSO_{4},7H_{2}O), the mixture solidifies like plaster of Paris into a homogeneous mass if the temperature be _below_ 22° (Van't Hoff und Van Deventer, 1886; Bakhuis Roozeboom, 1887); but if the temperature be above this _transition-point_ the water and double salt do not react on each other: that is, they do not solidify or give a mixture of sodium and magnesium sulphates. If a mixture (in equivalent quantities) of solutions of these salts be evaporated, and crystals of astrakhanite and of the individual salts capable of proceeding from it be added to the concentrated solution to avoid the possibility of a supersaturated solution, then at temperatures above 22° astrakhanite is exclusively formed (this is the method of its production), but at lower temperatures the individual salts are alone produced. If equivalent amounts of Glauber's salt and magnesium sulphate be mixed together in a solid state, there is no change at temperatures below 22°, but at higher temperatures astrakhanite and water are formed. The volume occupied by Na_{2}SO_{4},10H_{2}O in grams = 322/1·46 = 220·5 cubic centimetres, and by MgSO_{4},7H_{2}O = 246/1·68 = 146·4 c.c.; hence their mixture in equivalent quantities occupies a volume of 366·9 c.c. The volume of astrakhanite = 334/2·22 = 150·5 c.c., and the volume of 13H_{2}O = 234 c.c., hence their sum = 380·5 c.c., and therefore it is easy to follow the formation of the astrakhanite in a suitable apparatus (a kind of thermometer containing oil and a powdered mixture of sodium and magnesium sulphates), and to see by the variation in volume that below 22° it remains unchanged, and at higher temperatures proceeds the more quickly the higher the temperature. At the transition temperature the solubility of astrakhanite and of the mixture of the component salts is one and the same, whilst at higher temperatures a solution which is saturated for a mixture of the individual salts would be supersaturated for astrakhanite, and at lower temperatures the solution of astrakhanite will be supersaturated for the component salts, as has been shown with especial detail by Karsten, Deacon, and others. Roozeboom showed that there are two limits to the composition of the solutions which can exist for a double salt; these limits are respectively obtained by dissolving a mixture of the double salt with each of its component simple salts. Van't Hoff demonstrated, besides this, that the tendency towards the formation of double salts has a distinct influence on the progress of double decomposition, for at temperatures above 31° the mixture 2MgSO_{4},7H_{2}O + 2NaCl passes into MgNa_{2}(SO_{4})_{2},4H_{2}O + MgCl_{2},6H_{2}O + 4H_{2}O, whilst below 31° there is not this double decomposition, but it proceeds in the opposite direction, as may be demonstrated by the above-described methods. Van der Heyd obtained a potassium astrakhanite, K_{2}SO_{4}MgSO_{4},4H_{2}O, from solutions of the component salts at 100°.

From these experiments on double salts we see that there is as close a dependence between the temperature and the formation of substances as there is between the temperature and a change of state. It is a case of Deville's principles of dissociation, extended in the direction of the passage of a solid into a liquid. On the other hand, we see here how essential a _rôle_ water plays in the formation of compounds, and how the affinity for water of crystallisation is essentially analogous to the affinity between salts, and hence also to the affinity of acids for bases, because the formation of double salts does not differ in any essential point (except the degree of affinity--that is, from a quantitative aspect) from the formation of salts themselves. When sodium hydroxide with nitric acid gives sodium nitrate and water the phenomenon is essentially the same as in the formation of astrakhanite from the salts Na_{2}SO_{4},10H_{2}O and MgSO_{4},7H_{2}O. Water is disengaged in both cases, and hence the volumes are altered.

[26] This salt, and especially its crystallo-hydrate with 7H_{2}O, is generally known as Epsom salts. It has long been used as a purgative. It is easily obtained from magnesia and sulphuric acid, and it separates on the evaporation of sea water and of many saline springs. When carbonic anhydride is obtained by the action of sulphuric acid on magnesite, magnesium sulphate remains in solution. When dolomite--that is, a mixture of magnesium and calcium carbonates--is subjected to the action of a solution of hydrochloric acid until about half of the salt remains, the calcium carbonate is mostly dissolved and magnesium carbonate is left, which by treatment with sulphuric acid gives a solution of magnesium sulphate.

[27] The anhydrous salt, MgSO_{4} (sp. gr. 2·61), attracts moisture (7 mol. H_{2}O) from moist air; when heated in steam or hydrogen chloride it gives sulphuric acid, and when heated with carbon it is decomposed according to the equation 2MgSO_{4} + C = 2SO_{2} + CO_{2} + 2MgO. The monohydrated salt (kieserite), MgSO_{4},H_{2}O (sp. gr. 2·56), dissolves in water with difficulty; it is formed by heating the other crystallo-hydrates to 135°. The hexahydrated salt is dimorphous. If a solution, saturated at the boiling-point, be prepared, and cooled without access of crystals of the heptahydrated salt, then MgSO_{4},6H_{2}O crystallises out in _monoclinic_ prisms (Loewel, Marignac), which are quite as unstable as the salt, Na_{2}SO_{4},7H_{2}O; but if prismatic crystals of the cubic system of the copper-nickel salts of the composition MSO_{4},6H_{2}O be added, then crystals of MgSO_{4},6H_{2}O are deposited on them as prisms of the _cubic_ system (Lecoq de Boisbaudran). The common crystallo-hydrate, MgSO_{4},7H_{2}O, Epsom salts, belongs to the _rhombic_ system, and is obtained by crystallisation below 30°. Its specific gravity is 1·69. In a vacuum, or at 100°, it loses 5H_{2}O, at 132° 6H_{2}O, and at 210° all the 7H_{2}O (Graham). If crystals of ferrous or cobaltic sulphate be placed in a saturated solution, _hexagonal_ crystals of the heptahydrated salt are formed (Lecoq de Boisbaudran); they present an unstable state of equilibrium, and soon become cloudy, probably owing to their transformation into the more stable common form. Fritzsche, by cooling saturated solutions below 0°, obtained a mixture of crystals of ice and of a dodecahydrated salt, which easily split up at temperatures above 0°. Guthrie showed that dilute solutions of magnesium sulphate, when refrigerated, separate ice until the solution attains a composition MgSO_{4},24H_{2}O, which will completely freeze into a crystallo-hydrate at -5·3°. According to Coppet and Rüdorff, the temperature of the formation of ice falls by 0·073° for every part by weight of the heptahydrated salt per 100 of water. This figure gives (Chapter I., Note 49) _i_ = 1 for both the heptahydrated and the anhydrous salt, from which it is evident that it is impossible to judge the state of combination in which a dissolved substance occurs by the temperature of the formation of ice.

The solubility of the different crystallo-hydrates of magnesium sulphate, according to Loewel, also varies, like those of sodium sulphate or carbonate (_see_ Chapter XII., Notes 7 and 18). At 0° 100 parts of water dissolves 40·75 MgSO_{4} in the presence of the hexahydrated salt, 34·67 MgSO_{4} in the presence of the hexagonal heptahydrated salt, and only 26 parts of MgSO_{4} in the presence of the ordinary heptahydrated salt--that is, solutions giving the remaining crystallo-hydrates will be supersaturated for the ordinary heptahydrated salt.

All this shows how many diverse aspects of more or less stable equilibria may exist between water and a substance dissolved in it; this has already been enlarged on in Chapter I.

Carefully purified magnesium sulphate in its aqueous solution gives, according to Stcherbakoff, an alkaline reaction with litmus, and an acid reaction with phenolphthalein.

The specific gravity of solutions of certain salts of magnesium and calcium reduced to 15°/4° (see my work cited, Chapter I., Note 19), are, if water at 4° = 10,000,

MgSO_{4}: _s_ = 9,992 + 99·89_p_ + 0·553_p_^2 MgCl_{2}: _s_ = 9,992 + 81·31_p_ + 0·372_p_^2 CaCl_{2}: _s_ = 9,992 + 80·24_p_ + 0·476_p_^2

[28] Graham even distinguished the last equivalent of the water of crystallisation of the heptahydrated salt as that which is replaced by other salts, pointing out that double salts like MgK_{2}(SO_{4})_{2},6H_{2}O lose all their water at 135°, whilst MgSO_{4},7H_{2}O only parts with 6H_{2}O.

_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},4MgCO_{3},9H_{2}O; 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, 4MgSO_{4} + 4Na_{2}CO_{3} + 4H_{2}O = 4Na_{2}SO_{4} + CO_{2} + Mg(OH)_{2},3MgCO_{3},3H_{2}O. 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_, MgCO_{3}, 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_, MgCl_{2}. 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, MgCl_{2},6H_{2}O, 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, MgCl_{2},2NH_{4}Cl) 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 2MgCl_{2} + H_{2}O = MgOMgCl_{2} + 2HCl.[34 bis]

[29] The crystalline form of the anhydrous salt obtained in this manner is not the same as that of the natural salt. The former gives rhombohedra, like those in which calcium carbonate appears as calc spar, whilst the natural salt appears as rhombic prisms, like those sometimes presented by the same carbonate as aragonite, which will shortly be described.

[30] Magnesium sulphate enters into certain reactions which are proper to sulphuric acid itself. Thus, for instance, if a carefully prepared mixture of equivalent quantities of hydrated magnesium sulphate and sodium chloride be heated to redness, the evolution of hydrochloric acid is observed just as in the action of sulphuric acid on common salt, MgSO_{4} + 2NaCl + H_{2}O = Na_{2}SO_{4} + MgO + 2HCl. Magnesium sulphate acts in a similar manner on nitrates, with the evolution of nitric acid. A mixture of it with common salt and manganese peroxide gives chlorine. Sulphuric acid is sometimes replaced by magnesium sulphate in galvanic batteries--for example, in the well-known Meidinger battery. In the above-mentioned reactions we see a striking example of the similarity of the reactions of acids and salts, especially of salts which contain such feeble bases as magnesia.

[31] As sea-water contains many salts, MCl and MgX_{2}, it follows, according to Berthollet's teaching, that MgCl_{2} is also present.

[32] As the crystallo-hydrates of the salts of sodium often contain 10H_{2}O, so many of the salts of magnesium contain 6H_{2}O.

[33] This decomposition is most simply defined as the result of the two reverse reactions, MgCl_{2} ÷ H_{2}O = MgO + 2HCl and MgO + 2HCl = MgCl_{2} + H_{2}O, or as a distribution between O and Cl_{2} on the one hand and H_{2} and Mg on the other. (With O, MgCl_{2} gives chlorine, _see_ Chapter X., Note 33, and Chapter II., Note 3 bis and others, where the reactions and applications of MgCl_{2} are given.) It is then clear that, according to Berthollet's doctrine, the mass of the hydrochloric acid converts the magnesium oxide into chloride, and the mass of the water converts the magnesium chloride into oxide. The crystallo-hydrate, MgCl_{2},6H_{2}O, forms the limit of the reversibility. But an intermediate state of equilibrium may exist in the form of basic salts. On mixing ignited magnesia with a solution of magnesium chloride of specific gravity about 1·2, a solid mass is obtained which is scarcely decomposed by water at the ordinary temperature (_see_ Chapter XVI., Note 4). A similar means is employed for cementing sawdust into a solid mass, called cylolite, used for flooring, &c.

We may remark that MgBr_{2} crystallises not only with 6H_{2}O (temperature of fusion 152°), but also with 10H_{2}O (temperature of fusion +12°, formed at -18°). (Panfiloff, 1894).

[34] According to Thomsen, the combination of MgCl_{2} with 6H_{2}O evolves 33,000 calories, and its solution in an excess of water 36,000.

[34 bis] Hence MgCl_{2} may be employed for the preparation of chlorine and hydrochloric acid (Chapters X. and XI.). In general magnesium chloride, which is obtained in large quantities from sea water and Stassfurt carnallite, may find numerous practical uses.

_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: CaCO_{3} = CaO + CO_{2}. 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]

[35] There are many other methods of separating calcium from magnesium besides that mentioned above (Note 22). Among them it will be sufficient to mention the behaviour of these bases towards a solution of sugar; hydrated _lime_ is exceedingly _soluble in an aqueous solution of sugar_, whilst magnesia is but little soluble. All the lime may be extracted from dolomite by burning it, slaking the mixture of oxides thus obtained, and adding a 10 p.c. solution of sugar. Carbonic anhydride precipitates calcium carbonate from this solution. The addition of sugar (molasses) to the lime used for building purposes powerfully increases the binding power of the mortar, as I have myself found. I have been told that in the East (India, Japan) the addition of sugar to cement has long been practised.

[36] Moreover Caron obtained an alloy of calcium and zinc by fusing calcium chloride with zinc and sodium. The zinc distilled from this alloy at a white heat, leaving calcium behind (Note 50).

[37] Calcium iodide may be prepared by saturating lime with hydriodic acid. It is a very soluble salt (at 20° one part of the salt requires 0·49 part and at 43° 0·35 part of water for solution), is deliquescent in the air, and resembles calcium chloride in many respects. It changes but little when evaporated, and like calcium chloride fuses when heated, and therefore all the water may be driven off by heat. If anhydrous calcium iodide be heated with an equivalent quantity of sodium in a closely covered iron crucible, sodium iodide and metallic calcium are formed (Liés-Bodart). Dumas advises carrying on this reaction in a closed space under pressure.

[38] Kilns which act either intermittently or continuously are built for this purpose. Those of the first kind are filled with alternate layers of fuel and limestone; the fuel is lighted, and the heat developed by its combustion serves for decomposing the limestone. When the process is completed the kiln is allowed to cool somewhat, the lime raked out, and the same process repeated. In the continuously acting furnaces, constructed like that shown in fig. 78, the kiln itself only contains limestone, and there are lateral hearths for burning the fuel, whose flame passes through the limestone and serves for its decomposition. Such furnaces are able to work continuously, because the unburnt limestone may be charged from above and the burnt lime raked out from below. It is not every limestone that is suitable for the preparation of lime, because many contain impurities, principally clay, dolomite, and sand. Such limestones when burnt either fuse partially or give an impure lime, called _poor_ lime in distinction from that obtained from purer limestone, which is called _rich_ lime. The latter kind is characterised by its disintegrating into a fine powder when treated with water, and is suitable for the majority of uses to which lime is applied, and for which the poor lime is sometimes quite unfit. However, certain kinds of poor lime (as we shall see in Chapter XVIII., Note 25) are used in the preparation of hydraulic cements, which solidify into a hard mass under water.

In order to obtain perfectly pure lime it is necessary to take the purest possible materials. In the laboratory, marble or shells are used for this purpose as a pure form of calcium carbonate. They are first burnt in a furnace, then put in a crucible and moistened with a small quantity of water, and finally strongly ignited, by which means a pure lime is obtained. Pure lime may be more rapidly prepared by taking calcium nitrate, CaN_{2}O_{6}, which is easily obtained by dissolving limestone in nitric acid. The solution obtained is boiled with a small quantity of lime in order to precipitate the foreign oxides which are insoluble in water. The oxides of iron, aluminium, &c., are precipitated by this means. The salt is then crystallised and ignited: CaN_{2}O_{6} = CaO + 2NO_{2} + O.

In the decomposition of calcium carbonate the lime preserves the form of the lumps subjected to ignition; this is one of the signs distinguishing quicklime when it is freshly burnt and unaltered by air. It attracts moisture from the air and then disintegrates to a powder; if left long exposed in the air, it also attracts carbonic anhydride and increases in volume; it does not entirely pass into carbonate, but forms a compound of the latter with caustic lime.

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, CaH_{2}O_{2} (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_, CaO_{2},8H_{2}O, 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, CaX_{2}, which are distinguished by their stability in comparison with the salts MgX_{2}; neither does lime so easily form basic and double salts as magnesia.

[39] Lime, when raised to a white heat in the vapour of potassium, gives calcium, and in chlorine it gives off oxygen. Sulphur, phosphorus, &c., when heated with lime, are absorbed by it.

[40] The greater quantity of lime is used in making mortar for binding bricks or stones together, in the form of _lime_ or _cement_, or the so-called _slaked lime_. For this purpose the lime is mixed with water and sand, which serves to separate the particles of lime from each other. If only lime paste were put between two bricks they would not hold firmly together, because after the water had evaporated the lime would occupy a smaller space than before, and therefore cracks and powder would form in its mass, so that it would not at all produce that complete cementation of the bricks which it is desired to attain. Pieces of stone--that is, sand--mixed with the lime hinder this process of disintegration, because the lime binds together the individual grains of sand mixed with it, and forms one concrete mass, in consequence of a process which proceeds after the desiccation or removal of the water. The process of the solidification of lime, taken as slaked lime, consists first in the direct evaporation of the water and crystallisation of the hydrate, so that the lime binds the stones and sand mixed with it, just as glue binds two pieces of wood. But this preliminary binding action of lime is feeble (as is seen by direct experiment) unless there be further alteration of the lime leading to the formation of carbonates, silicates, and other salts of calcium which are distinguished by their great cohesiveness. With the progress of time the cement is partially subjected to the action of the carbonic anhydride in the air, owing to which calcium carbonate is formed, but not more than half the lime is thus converted into carbonate. Besides which, the lime partially acts on the silica of the bricks, and it is owing to these new combinations simultaneously forming in the cement that it gradually becomes stronger and stronger. Hence the binding action of the lime becomes stronger with the lapse of time. This is the reason (and not, as is sometimes said, because the ancients knew how to build stronger than we do) why buildings which have stood for centuries possess a very strongly binding cement. Hydraulic cements will be described later (Chapter XVIII., Note 25).

[41] Professor Glinka measured the transparent bright crystals of calcium hydroxide which are formed in common hydraulic (Portland) cement.

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]

[42] The act of heating brings the substance into that state of internal motion which is required for reaction. It should be considered that by the act of heating not only is the bond between the parts, or cohesion of the molecules, altered (generally diminished), not only is the motion or store of energy of the whole molecule increased, but also that in all probability the motion of the atoms themselves in molecules undergoes a change. The same kind of change is accomplished by the act of solution, or of combination in general, judging from the fact that a dissolved or combined substance--for instance, lime with water--reacts on carbonic anhydride as it does under the action of heat. For the comprehension of chemical phenomena it is exceedingly useful to recognise clearly this parallelism. Rose's observation on the formation (by the slow diffusion of solutions of calcium chloride and sodium carbonate) of aragonite from dilute, and of calc spar from strong, solutions is easily understood from this point of view. As aragonite is always formed from hot solutions, it appears that dilution with water acts like heat. The following experiment of Kühlmann is particularly instructive in this sense. Anhydrous (perfectly dry) barium oxide does not react with monohydrated sulphuric acid, H_{2}SO_{4} (containing neither free water nor anhydride, SO_{3}). But if either an incandescent object or a moist substance is brought into contact with the mixture a violent reaction immediately begins (it is essentially the same as combustion), and the whole mass reacts.

The influence of solution on the process of reaction is instructively illustrated by the following experiment. Lime, or barium oxide, is placed in a flask or retort having an upper orifice and connected with a tube immersed in mercury. A funnel furnished with a stopcock and filled with water is fixed into the upper orifice of the retort, which is then filled with dry carbonic anhydride. There is no absorption. When a constant temperature is arrived at, the unslaked oxide is made to absorb all the carbonic anhydride by carefully admitting water. A vacuum is formed, as is seen by the mercury rising in the neck of the retort. With water the absorption goes on to the end, whilst under the action of heat there remains the dissociating tension of the carbonic anhydride. Furthermore, we here see that, with a certain resemblance, there is also a distinction, depending on the fact that at low temperatures calcium carbonate does not dissociate; this determines the complete absorption of the carbonic anhydride in the aqueous solution.

[43] Experience has shown that by moistening partially-burnt lime with water and reheating it, it is easy to drive off the last traces of carbonic anhydride from it, and that, in general, by blowing air or steam through the lime, and even by using moist fuel, it is possible to accelerate the decomposition of the calcium carbonate. The partial pressure is decreased by these means.

[44] Before the introduction of Deville's theory of dissociation, the _modus operandi_ of decompositions like that under consideration was understood in the sense that decomposition starts at a certain temperature, and that it is accelerated by a rise of temperature, but it was not considered possible that combination could proceed at the same temperature as that at which decomposition goes on. Berthollet and Deville introduced the conception of equilibrium into chemical science, and elucidated the question of reversible reactions. Naturally the subject is still far from being clear--the questions of the rate and completeness of reaction, of contact, &c., still intrude themselves--but an important step has been made in chemical mechanics, and we have started on a new path which promises further progress, towards which much has been done not only by Deville himself, but more especially by the French chemists Debray, Troost, Lemoine, Hautefeuille, Le Chatelier, and others. Among other things those investigators have shown the close resemblance between the phenomena of evaporation and dissociation, and pointed out that the amount of heat absorbed by a dissociating substance may be calculated according to the law of the variation of dissociation-pressure, in exactly the same manner as it is possible to calculate the latent heat of the evaporation of water, knowing the variation of the tension with the temperature, on the basis of the second law of the mechanical theory of heat. Details of this subject must be looked for in special works on physical chemistry. _One and the same conception_ of the mechanical theory of heat _is applicable to dissociation_ and _evaporation_.

[45] But the question as to the formation of a basic calcium carbonate with a rise of temperature still remains undecided. The presence of water complicates all the relations between lime and carbonic anhydride, all the more as the existence of an attraction between calcium carbonate and water is seen from its being able to give a _crystallo-hydrate_, CaCO_{3},5H_{2}O (Pelouze), which crystallises in rhombic prisms of sp. gr. about 1·77 and loses its water at 20°. These crystals are obtained when a solution of lime in sugar and water is left long exposed to the air and slowly attracts carbonic anhydride from it, and also by the evaporation of such a solution at a temperature of about 3°. On the other band, it is probable that an _acid salt_, CaH_{2}(CO_{3})_{2}, is formed in an aqueous solution, not only because water containing carbonic acid dissolves calcium carbonate, but more especially in view of the researches of Schloesing (1872), which showed that at 16° a litre of water in an atmosphere of carbonic anhydride (pressure 0·984 atmosphere) dissolves 1·086 gram of calcium carbonate and 1·778 gram of carbonic anhydride, which corresponds with the formation of calcium hydrogen carbonate, and the solution of carbonic anhydride in the remaining water. Caro showed that a litre of water is able to dissolve as much as 3 grams of calcium carbonate if the pressure be increased to 4 and more atmospheres. The calcium carbonate is precipitated when the carbonic anhydride passes off in the air or in a current of another gas; this also takes place in many natural springs. Tufa, stalactites, and other like formations from waters containing calcium carbonate and carbonic acid in solution are formed in this manner. The solubility of calcium carbonate itself at the ordinary temperature does not exceed 13 milligrams per litre of water.

_Calcium carbonate_, CaCO_{3}, 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]

[46] Dimorphous bodies differ from true isomers and polymers in that they do not differ in their chemical reactions, which are determined by a difference in the distribution (motion) of the atoms in the molecules, and therefore dimorphism is usually ascribed to a difference in the distribution of similar molecules, building up a crystal. Although such a hypothesis is quite admissible in the spirit of the atomic and molecular theory, yet, as in such a redistribution of the molecules a perfect conservation of the distribution of the atoms in them cannot be imagined, and in every effort of chemical reaction there must take place a certain motion among the atoms; so in my opinion there is no firm basis for distinguishing dimorphism from the general conception of isomerism, under which the cases of those organic bodies which are dextro and lævo rotatory (with respect to polarised light) have recently been brought with such brilliant success. When calcium carbonate separates out from solutions, it has at first a gelatinous appearance, which leads to the supposition that this salt appears in a colloidal state. It only crystallises with the progress of time. The colloidal state of calcium carbonate is particularly clear from the following observations made by Prof. Famintzin, who showed that when it separates from solutions it is obtained under certain conditions in the form of grains having the peculiar paste-like structure proper to starch, which fact has not only an independent interest, but presents an example of a mineral substance being obtained in a form until then only known in the organic substances elaborated in plants. This shows that the forms (cells, vessels, &c.) in which vegetable and animal substances occur in organisms do not present in themselves anything peculiar to organisms, but are only the result of those particular conditions in which these substances are formed. Traube and afterwards Monnier and Vogt (1882) obtained formations which, under the microscope, were in every respect identical in appearance with vegetable cells, by means of a similar slow formation of precipitates (by reacting on sulphates of different metals with sodium silicate or carbonate).

_Calcium sulphate_ in combination with two equivalents of water, CaSO_{4},2H_{2}O, is very widely distributed in nature, and is known as _gypsum_. Gypsum loses one and a half and two equivalents of water at a moderate temperature,[47] and anhydrous or burnt gypsum is then obtained, which is also known as plaster of Paris, and is employed in large quantities for modelling.[48] This use depends on the fact that burnt and finely-divided and sifted gypsum forms a paste when mixed with water; after a certain time this paste becomes slightly heated and solidifies, owing to the fact that the anhydrous calcium sulphate, CaSO_{4}, again combines with water. When the plaster of Paris and water are first made into a paste they form a mechanical mixture, but when the mass solidifies, then a compound of the calcium sulphate with two molecules of water is produced; and this may be regarded as derived from S(OH)_{6} by the substitution of two atoms of hydrogen by one atom of bivalent calcium. Natural gypsum sometimes appears as perfectly colourless, or variegated, marble-like, masses, and sometimes in perfectly colourless crystals, _selenite_, of specific gravity 2·33. The semi-transparent gypsum, or _alabaster_, is often carved into small statues. Besides which an anhydrous calcium sulphate, CaSO_{4}, called _anhydrite_ (specific gravity 2·97), occurs in nature. It sometimes occurs along with gypsum. It is no longer capable of combining directly with water, and differs in this respect from the anhydrous salt obtained by gently igniting gypsum. If gypsum be very strongly heated it shrinks and loses its power of combining with water.[48 bis] One part of calcium sulphate requires at 0° 525 parts of water for solution, at 38° 466 parts, and at 100° 571 parts of water. The maximum solubility of gypsum is at about 36°, which is nearly the same temperature as that at which sodium sulphate is most soluble.[49]

[47] According to Le Chatelier (1888), 1-1/2H_{2}O is lost at 120°--that is, H_{2}O,2CaSO_{4} is formed, but at 194° all the water is expelled. According to Shenstone and Cundall (1888) gypsum begins to lose water at 70° in dry air. The semi-hydrated compound H_{2}O,2CaSO_{4} is also formed when gypsum is heated with water in a closed vessel at 150° (Hoppe-Seyler).

[48] For stucco-work it is usual to add lime and sand, as the mass is then harder and does not solidify so quickly. For imitating marble, glue is added to the plaster, and the mass is polished when thoroughly dry. Re-burnt gypsum cannot be used over again, as that which has once solidified is, like the natural anhydride, not able to recombine with water. It is evident that the structure of the molecules in the crystallised mass, or in general in any dense mass, exerts an influence on the chemical action, which is more particularly evident in metals in their different forms (powder, crystalline, rolled, &c.)

[48 bis] According to MacColeb, gypsum dehydrated at 200° has a specific gravity 2·577, and heated to its point of fusion, 2·654. Potilitzin (1894) also admits the two above-named modifications of anhydrous gypsum, which, moreover, always contain the semi-hydrated hydrate (Note 47), and he explains by their relation to water the phenomena observed in the solidification of a mixture of burnt gypsum and water.

[49] As Marignac showed, gypsum, especially when desiccated at 120°, easily gives supersaturated solutions with respect to CaSO_{4},2H_{2}O, which contain as much as 1 part of CaSO_{4} to 110 parts of water. Boiling dilute hydrochloric acid dissolves gypsum, forming calcium chloride. The behaviour of gypsum towards the alkaline carbonates has been described in Chapter X. Alcohol precipitates gypsum from its aqueous solutions, because, like the sulphates in general, it is sparingly soluble in alcohol. Gypsum, like all the sulphates, when heated with charcoal, gives up its oxygen, forming the sulphide, CaS.

Calcium ulphate, like magnesium sulphate, is capable of forming double salts, but with difficulty, and they are chemically less stable. They contain, as is always the case with double salts, less water of crystallisation than the component salts. Rose, Struvé, and others obtained the salt CaK_{2}(SO_{4})_{2},H_{2}O; a mixture of gypsum with an equivalent amount of potassium sulphate and water solidifies into a homogeneous mass. Fritzsche obtained the corresponding sodium salt in a hydrated and anhydrous state, by heating a mixture of gypsum with a saturated solution of sodium sulphate. The anhydrous salt occurs in nature as _glauberite_. Fritzsche also obtained _gaylussite_, Na_{2}Ca(CO_{3})_{2},5H_{2}O, by pouring a saturated solution of sodium carbonate on to freshly-precipitated calcium carbonate. Calcium also forms basic salts, but only a few. Veeren (1892) obtained Ca(NO_{3})_{2}Ca(OH)_{2},2-1/2H_{2}O by leaving powdered caustic lime in a saturated solution of Ca(NO_{3})_{2} until it solidified. This salt is decomposed by water.

As lime is a more energetic base than magnesia, so _calcium chloride_, CaCl_{2}, is not so easily decomposed by water, and its solutions only disengage a small quantity of hydrochloric acid when evaporated, and when the evaporation is conducted in a stream of hydrochloric acid it easily gives an anhydrous salt which fuses at 719°; otherwise an aqueous solution yields a crystallo-hydrate, CaCl_{2},6H_{2}O, which melts at 30°.[50]

[50] Calcium chloride has a specific gravity 2·20, or, when fused, 2·12, and the sp. gr. of the crystallised salt CaCl_{2},6H_{2}O is 1·69. If the volume of the crystals at 0° = 1, then at 29° it is 1·020, and the volume of the fused mass at the same temperature is 1·118 (Kopp) (specific gravity of solutions, _see_ Note 27). The solution containing 50 p.c. CaCl_{2} boils at 130°, 70 p.c. at 158°. Superheated steam decomposes calcium chloride with more difficulty than magnesium chloride and with greater ease than barium chloride (Kuhnheim). Sodium does not decompose fused calcium chloride even on prolonged heating (Liés-Bodart), but an alloy of sodium with zinc, lead, and bismuth decomposes it, forming an alloy of calcium with one of the above-named metals (Caron). The zinc alloy may be obtained with as much as 15 p.c. of calcium. Calcium chloride is soluble in alcohol and absorbs ammonia.

A gram molecular weight of calcium chloride in dissolving in an excess of water evolves 18,723 calories, and in dissolving in alcohol 17,555 units of heat, according to Pickering.

Roozeboom made detailed researches on the crystallo-hydrates of calcium chloride (1889), and found that CaCl_{2},6H_{2}O melts at 30·2°, and is formed at low temperatures from solutions containing not more than 103 parts of calcium chloride per 100 parts of water; if the amount of salt (always to 100 parts of water) reaches 120 parts, then tabular crystals of CaCl_{2},4H_{2}O[Greek: b] are formed, which at temperatures above 38·4° are converted into the crystallo-hydrates CaCl_{2},2H_{2}O, whilst at temperatures below 18° the [Greek: b] variety passes into the more stable CaCl_{2},4H_{2}O[Greek: a], which process is aided by mechanical friction. Hence, as is the case with magnesium sulphate (Note 27), one and the same crystallo-hydrate appears in two forms--the [Greek: b], which is easily produced but is unstable, and the [Greek: a], which is stable. The solubility of the above-mentioned hydrates of chloride of calcium, or amount of calcium chloride per 100 parts of water, is as follows:--

0° 20° 30° 40° 60°

CaCl_{2},6H_{2}O 60 75 100 (102·8) CaCl_{2},4H_{2}O[Greek: a] -- 90 101 117} (154·2) CaCl_{2},4H_{2}O[Greek: b] -- 104 114 --} CaCl_{2},2H_{2}O -- -- (308·3) 128 137

The amount of calcium chloride to 100 parts of water in the crystallo-hydrate is given in brackets. The point of intersection of the curves of solubility lies at about 30° for the first two salts and about 45° for the salts with 4H_{2}O and 2H_{2}O. The crystals CaCl_{2},2H_{2}O may, however, be obtained (Ditte) at the ordinary temperature from solutions containing hydrochloric acid. The vapour tension of this crystallo-hydrate equals the atmospheric at 165°, and therefore the crystals may be dried in an atmosphere of steam and obtained without a mother liquor, whose vapour tension is greater. This crystallo-hydrate decomposes at about 175° into CaCl_{2},H_{2}O and a solution; this is easily brought about in a closed vessel when the pressure is greater than the atmosphere. This crystallo-hydrate is destroyed at temperatures above 260°, anhydrous calcium chloride being formed.

Neglecting the unstable modification CaCl_{2},4H_{2}O[Greek: b], we will give the temperatures _t_ at which the passage of one hydrate into another takes place and at which the solution CaCl_{2} + _n_H_{2}O, the two solids A and B and aqueous vapour, whose tension is given as _p_ in millimetres, are able to exist together in stable equilibrium, according to Roozeboom's determinations:

_t_ _n_ A B _p_ -55° 14·5 ice CaCl_{2},6H_{2}O 0 +29·8° 6·1 CaCl_{2},6H_{2}O CaCl_{2},4H_{2}O 6·8 45·3° 4·7 CaCl_{2},4H_{2}O CaCl_{2},2H_{2}O 11·8 175·5° 2·1 CaCl_{2},2H_{2}O CaCl_{2},H_{2}O 842 260° 1·8 CaCl_{2},H_{2}O CaCl_{2} Several atmospheres

Solutions of calcium chloride may serve as a convenient example for the study of the supersaturated state, which in this case easily occurs, because different hydrates are formed. Thus at 25° solutions containing more than 83 parts of anhydrous calcium chloride per 100 of water will be supersaturated for the hydrate CaCl_{2},6H_{2}O.

On the other hand, Hammerl showed that solutions of calcium chloride, when frozen, deposit ice if they contain less than 43 parts of salt per 100 of water, and if more the crystallo-hydrate CaCl_{2},6H_{2}O separates, and that a solution of the above composition (CaCl_{2},14H_{2}O requires 44·0 parts calcium chloride per 100 of water) solidifies as a cryohydrate at about -55°.

Just as for potassium, K = 39 (and sodium, Na = 23), there are the near analogues, Rb = 85 and Cs = 133, and also another, Li = 7, so in exactly the same manner for calcium, Ca = 40 (and magnesium, Mg = 24), there is another analogue of lighter atomic weight, beryllium, Be = 9, besides the near analogues strontium, Sr = 87, and barium, Ba = 137. As rubidium and cæsium are more rarely met with in nature than potassium, so also strontium and barium are rarer than calcium (in the same way that bromine and iodine are rarer than chlorine). Since they exhibit many points of resemblance with calcium, strontium and barium may be characterised after a very short acquaintance with their chief compounds; this shows the important advantages gained by distributing the elements according to their natural groups, to which matter we shall turn our attention in the next chapter.

Among the compounds of barium met with in nature the commonest is the _sulphate_, BaSO_{4}, which forms anhydrous crystals of the rhombic system, which are identical in their crystalline form with anhydrite, and generally occur as transparent and semi-transparent masses of tabular crystals having a high specific gravity, namely 4·45, for which reason this salt bears the name of _heavy spar_ or _barytes_. Analogous to it is _celestine_, SrSO_{4}, which is, however, more rarely met with. Heavy spar frequently forms the gangue separated on dressing metallic ores from the vein stuff; this mineral is the source of all other barium compounds; for the carbonate, although more easily transformed into the other compounds (because acids act directly on it, evolving carbonic anhydride), is a comparatively rare mineral (BaCO_{3} forms the mineral _witherite_; SrCO_{3}, _strontianite_; both are rare, the latter is found at Etna). The treatment of barium sulphate is rendered difficult from the fact that it is insoluble both in water and acids, and has therefore to be treated by a method of reduction.[51] Like sodium sulphate and calcium sulphate, heavy spar when heated with charcoal parts with its oxygen and forms barium sulphide, BaS. For this purpose a pasty mixture of powdered heavy spar, charcoal, and tar is subjected to the action of a strong heat, when BaSO_{4} + 4C = BaS + 4CO. The residue is then treated with water, in which the barium sulphide is soluble.[52] When boiled with hydrochloric acid, barium chloride, BaCl_{2}, is obtained in solution, and the sulphur is disengaged as gaseous sulphuretted hydrogen, BaS + 2HCl = BaCl_{2} + H_{2}S. In this manner barium sulphate is converted into barium chloride,[53] and the latter by double decomposition with strong nitric acid or nitre gives the less soluble barium nitrate, Ba(NO_{3})_{2},[54] or with sodium carbonate a precipitate of barium carbonate, BaCO_{3}. Both these salts are able to give _barium oxide_, or _baryta_, BaO, and the hydroxide, Ba(HO)_{2}, which differs from lime by its great solubility in water,[55] and by the ease with which it forms a crystallo-hydrate, BaH_{2}O_{2},8H_{2}O, from its solutions. Owing to its solubility, baryta is frequently employed in manufactures and in practical chemistry as an alkali which has the very important property that it may be always entirely removed from solution by the addition of sulphuric acid, which entirely separates it as the insoluble barium sulphate, BaSO_{4}. It may also be removed whilst it remains in an alkaline state (for example, the excess which may remain when it is used for saturating acids) by means of carbonic anhydride, which also completely precipitates baryta as a sparingly soluble, colourless, and powdery carbonate. Both these reactions show that baryta has such properties as would very greatly extend its use were its compounds as widely distributed as those of sodium and calcium, and were its soluble compounds not poisonous. Barium nitrate is directly decomposed by the action of heat, barium oxide being left behind. The same takes place with barium carbonate, especially that form of it precipitated from solutions, and when mixed with charcoal or ignited in an atmosphere of steam. Barium oxide combines with water with the development of a large amount of heat, and the resultant hydroxide is very stable in its retention of the water, although it parts with it when strongly ignited.[55 bis] With oxygen the anhydrous oxide gives, as already mentioned in Chapters III. and IV., a _peroxide_, BaO_{2}.[56] Neither calcium nor strontium oxides are able to give such a peroxide directly, but they form peroxides under the action of hydrogen peroxide.

[51] The action of barium sulphate on sodium and potassium carbonates is given on p. 437.

[52] Barium sulphide is decomposed by water, BaS + 2H_{2}O = H_{2}S + Ba(OH)_{2} (the reaction is reversible), but both substances are soluble in water, and their separation is complicated by the fact that barium sulphide absorbs oxygen and gives insoluble barium sulphate. The hydrogen sulphide is sometimes removed from the solution by boiling with the oxides of copper or zinc. If sugar be added to a solution of barium sulphide, barium saccharate is precipitated on heating; it is decomposed by carbonic anhydride, so that barium carbonate is formed. An equivalent mixture of sodium sulphate with barium or strontium sulphates when ignited with charcoal gives a mixture of sodium sulphide and barium or strontium sulphide, and if this mixture be dissolved in water and the solution evaporated, barium or strontium hydroxide crystallises out on cooling, and sodium hydrosulphide, NaHS, is obtained in solution. The hydroxides BaH_{2}O_{2} and SrH_{2}O_{2} are prepared on a large scale, being applied to many reactions; for example, strontium hydroxide is prepared for sugar works for extracting crystallisable sugar from molasses.

We may remark that Boussingault, by igniting barium sulphate in hydrochloric acid gas, obtained a complete decomposition, with the formation of barium chloride. Attention should also be turned to the fact that Grouven, by beating a mixture of charcoal and strontium sulphate with magnesium and potassium sulphates, showed the easy decomposability depending on the formation of double salts, such as SrS,K_{2}S, which are easily soluble in water, and give a precipitate of strontium carbonate with carbonic anhydride. In such examples as these we see that the force which binds double salts may play a part in directing the course of reactions, and the number of double salts of silica on the earth's surface shows that nature takes advantage of these forces in her chemical processes. It is worthy of remark that Buchner (1893), by mixing a 40 per cent. solution of barium acetate with a 60 per cent. solution of sulphate of alumina, obtained a thick glutinous mass, which only gave a precipitate of BaSO_{4} after being diluted with water.

[53] Barium sulphate is sometimes converted into barium chloride in the following manner: finely-ground barium sulphate is heated with coal and manganese chloride (the residue from the manufacture of chlorine). The mass becomes semi-liquid, and when it evolves carbonic oxide the heating is stopped. The following double decompositions proceed during this operation: first the carbon takes up the oxygen from the barium sulphate, and gives sulphide, BaS, which enters into double decomposition with the chloride of manganese, MnCl_{2}, forming manganese sulphide, MnS, which is insoluble in water, and soluble barium chloride. This solution is easily obtained pure because many foreign impurities, such as iron, remain in the insoluble portion with the manganese. The solution of barium chloride is chiefly used for the preparation of barium sulphate, which is precipitated by sulphuric acid, by which means _barium sulphate_ is re-formed as a powder. This salt is characterised by the fact that it is unacted on by the majority of chemical reagents, is insoluble in water, and is not dissolved by acids. Owing to this, artificial barium sulphate forms a permanent white paint which is used instead of (and mixed with) white lead, and has been termed 'blanc fixé' or 'permanent white.'

The solution of one part of calcium chloride at 20° requires 1·36 part of water, the solution of one part of strontium chloride requires 1·88 part of water at the same temperature, and the solution of barium chloride 2·88 parts of water. The solubility of the bromides and iodides varies in the same proportion. The chlorides of barium and strontium crystallise out from solution with great ease in combination with water; they form BaCl_{2},2H_{2}O and SrCl_{2},6H_{2}O. The latter (which separates out at 40°) resembles the salts of Ca and Mg in composition, and Étard (1892) obtained SrCl_{2},2H_{2}O from solutions at 90-130°. We may also observe that the crystallo-hydrates BaBr_{2},H_{2}O and BaI_{2},7H_{2}O are known.

[54] The nitrates Sr(NO_{3})_{2} (in the cold its solutions give a crystallo-hydrate containing 4H_{2}O) and Ba(NO_{3})_{2} are so very sparingly soluble in water that they separate in considerable quantities when a solution of sodium nitrate is added to a strong solution of either barium or strontium chloride. They are obtained by the action of nitric acid on the carbonates or oxides. 100 parts of water at 15° dissolve 6·5 parts of strontium nitrate and 8·2 parts of barium nitrate, whilst more than 300 parts of calcium nitrate are soluble at the same temperature. Strontium nitrate communicates a crimson coloration to the flame of burning substances, and is therefore frequently used for Bengal fire, fireworks, and signal lights, for which purpose the salts of lithium are still better fitted. Calcium nitrate is exceedingly hygroscopic. Barium nitrate, on the contrary, does not show this property in the least degree, and in this respect it resembles potassium nitrate, and is therefore used instead of the latter for the preparation of a gunpowder which is called 'saxifragin powder' (76 parts of barium nitrate, 2 parts of nitre, and 22 parts of charcoal).

[55] The dissociation of the crystallo-hydrate of baryta is given in