[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 MgK2(SO4)2,6H2O lose all their water at 135°, whilst MgSO4,7H2O only parts with 6H2O.[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, MgSO4+ 2NaCl + H2O = Na2SO4+ 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 MgX2, it follows, according to Berthollet's teaching, that MgCl2is also present.[32]As the crystallo-hydrates of the salts of sodium often contain 10H2O, so many of the salts of magnesium contain 6H2O.[33]This decomposition is most simply defined as the result of the two reverse reactions, MgCl2÷ H2O = MgO + 2HCl and MgO + 2HCl = MgCl2+ H2O, or as a distribution between O and Cl2on the one hand and H2and Mg on the other. (With O, MgCl2gives chlorine,seeChapter X., Note33, and Chapter II., Note 3 bis and others, where the reactions and applications of MgCl2are 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, MgCl2,6H2O, 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 (seeChapter 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 MgBr2crystallises not only with 6H2O (temperature of fusion 152°), but also with 10H2O (temperature of fusion +12°, formed at -18°). (Panfiloff, 1894).[34]According to Thomsen, the combination of MgCl2with 6H2O evolves 33,000 calories, and its solution in an excess of water 36,000.[34 bis]Hence MgCl2may be employed for the preparation of chlorine and hydrochloric acid (ChaptersX. andXI.). In general magnesium chloride, which is obtained in large quantities from sea water and Stassfurt carnallite, may find numerous practical uses.[35]There are many other methods of separating calcium from magnesium besides that mentioned above (Note22). Among them it will be sufficient to mention the behaviour of these bases towards a solution of sugar; hydratedlimeis exceedinglysoluble 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 (Note50).[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, calledpoorlime in distinction from that obtained from purer limestone, which is calledrichlime. 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, CaN2O6, 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: CaN2O6= CaO + 2NO2+ 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.[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 oflimeorcement, or the so-calledslaked 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.[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, H2SO4(containing neither free water nor anhydride, SO3). 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, themodus operandiof 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 conceptionof the mechanical theory of heatis applicable to dissociationandevaporation.[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 acrystallo-hydrate, CaCO3,5H2O (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 anacid salt, CaH2(CO3)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.[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).[47]According to Le Chatelier (1888), 1½H2O is lost at 120°—that is, H2O,2CaSO4is 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 H2O,2CaSO4is 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 (Note47), 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 CaSO4,2H2O, which contain as much as 1 part of CaSO4to 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 ChapterX. 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 sulphate, 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 CaK2(SO4)2,H2O; 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 asglauberite. Fritzsche also obtainedgaylussite, Na2Ca(CO3)2,5H2O, 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(NO3)2Ca(OH)2,2½H2O by leaving powdered caustic lime in a saturated solution of Ca(NO3)2until it solidified. This salt is decomposed by water.[50]Calcium chloride has a specific gravity 2·20, or, when fused, 2·12, and the sp. gr. of the crystallised salt CaCl2,6H2O 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,seeNote27). The solution containing 50 p.c. CaCl2boils 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 CaCl2,6H2O 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 CaCl2,4H2Oβ are formed, which at temperatures above 38·4° are converted into the crystallo-hydrates CaCl2,2H2O, whilst at temperatures below 18° the β variety passes into the more stable CaCl2,4H2Oα, which process is aided by mechanical friction. Hence, as is the case with magnesium sulphate (Note27), one and the same crystallo-hydrate appears in two forms—the β, which is easily produced but is unstable, and the α, 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°CaCl2,6H2O6075100(102·8)CaCl2,4H2Oα—90101117(154·2)CaCl2,4H2Oβ—104114—CaCl2,2H2O——(308·3)128137The 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 4H2O and 2H2O. The crystals CaCl2,2H2O 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 CaCl2,H2O 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 CaCl2,4H2Oβ, we will give the temperaturestat which the passage of one hydrate into another takes place and at which the solution CaCl2+nH2O, the two solids A and B and aqueous vapour, whose tension is given aspin millimetres, are able to exist together in stable equilibrium, according to Roozeboom's determinations:tnABp-55°14·5iceCaCl2,6H2O0+29·8°6·1CaCl2,6H2OCaCl2,4H2O6·845·3°4·7CaCl2,4H2OCaCl2,2H2O11·8175·5°2·1CaCl2,2H2OCaCl2,H2O842200°1·8CaCl2,H2OCaCl2Several atmospheresSolutions 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 CaCl2,6H2O.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 CaCl2,6H2O separates, and that a solution of the above composition (CaCl2,14H2O requires 44·0 parts calcium chloride per 100 of water) solidifies as a cryohydrate at about -55°.[51]The action of barium sulphate on sodium and potassium carbonates is given on p. 437.[52]Barium sulphide is decomposed by water, BaS + 2H2O = H2S + 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 BaH2O2and SrH2O2are 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,K2S, 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 BaSO4after 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, MnCl2, 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 meansbarium sulphateis 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 BaCl2,2H2O and SrCl2,6H2O. The latter (which separates out at 40°) resembles the salts of Ca and Mg in composition, and Étard (1892) obtained SrCl2,2H2O from solutions at 90–130°. We may also observe that the crystallo-hydrates BaBr2,H2O and BaI2,7H2O are known.[54]The nitrates Sr(NO3)2(in the cold its solutions give a crystallo-hydrate containing 4H2O) and Ba(NO3)2are 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 Chapter I., Note65. 100 parts of water dissolve0°20°40°60°80°BaO1·53·57·418·890·8SrO0·30·71·439Supersaturated solutions are easily formed.The anhydrous oxide BaO fuses in the oxyhydrogen flame. When ignited in the vapour of potassium, the latter takes up the oxygen; whilst in chlorine, oxygen is separated and barium chloride formed.[55 bis]Brugellmann, by heating BaH2O2in a graphite or clay crucible, obtained BaO in needles, sp. gr. 5·32, and by heating in a platinum crucible—in crystals belonging to the cubical system, sp. gr. 5·74. SrO is obtained in the latter form from the nitrate. The following are the specific gravities of the oxides from different sources:—MgOCaOSrOfromRN2O63·383·254·75„RCO33·483·264·45„RH2O23·413·254·57[56]The property of barium oxide of absorbing oxygen when heated, and giving the peroxide, BaO2, is very characteristic for this oxide (seeChapter III., Note7). It only belongs to the anhydrous oxide. The hydroxide does not absorb oxygen. Peroxides of calcium and strontium may be obtained by means of hydrogen peroxide. Barium peroxide is insoluble in water, but is able to form a hydrate with it, and also to combine with hydrogen peroxide, forming a very unstable compound having the composition BaH2O4(obtained by Professor Schöne), which in course of time evolves oxygen (Chapter IV., Note21).[57]Even in solutions a gradual progression in the increase of the specific gravity shows itself, not only for equivalent solutions (for instance, RCl2+ 200H2O), but even with an equal percentage composition, as is seen from the curves giving the specific gravity (water 4° = 10,000) at 15° (for barium chloride, according to Bourdiakoff's determinations):BeCl2: S = 9,992 + 67·21p+ 0·111p2CaCl2: S = 9,992 + 80·24p+ 0·476p2SrCl2: S = 9,992 + 85·57p+ 0·733p2BaCl2: S = 9,992 + 86·56p+ 0·813p2[58]One part of calcium sulphate at the ordinary temperature requires about 500 parts of water for solution, strontium sulphate about 7,000 parts, barium sulphate about 400,000 parts, whilst beryllium sulphate is easily soluble in water.[59]We refer beryllium to the class of the bivalent metals of the alkaline earths—that is, we ascribe to its oxide the formula BeO, and do not consider it as trivalent (Be = 13·5, Chapter VII., Note21), although that view has been upheld by many chemists. The true atomic composition of beryllium oxide was first given by the Russian chemist, Avdéeff (1819), in his researches on the compounds of this metal. He compared the compounds of beryllium to those of magnesium, and refuted the notion prevalent at the time, of the resemblance between the oxides of beryllium and aluminium, by proving that beryllium sulphate presents a greater resemblance to magnesium sulphate than to aluminium sulphate. It was especially noticed that the analogues of alumina give alums, whilst beryllium oxide, although it is a feeble base, easily giving, like magnesia, basic and double salts, does not form true alums. The establishment of the periodic system of the elements (1869), considered in the following chapter, immediately indicated that Avdéeff's view corresponded with the truth—that is, that beryllium is bivalent, which therefore necessitated the denial of its trivalency. This scientific controversy resulted in a long series of researches (1870–80) concerning this element, and ended in Nilson and Pettersson—two of the chief advocates of the trivalency of beryllium—determining the vapour density of BeCl2= 40, (Chapter VII., Note21), which gave an undoubted proof of its bivalency (seealso Note3).[60]Beryllium oxide, like aluminium oxide, is precipitated from solutions of its salts by alkalis as a gelatinous hydroxide, BeH2O2, which, like alumina, is soluble in an excess of caustic potash or soda. This reaction may be taken advantage of for distinguishing and separating beryllium from aluminium, because when the alkaline solution is diluted with water and boiled, beryllium hydroxide is precipitated, whilst the alumina remains in solution. The solubility of the beryllium oxide at once clearly indicates its feeble basic properties, and, as it were, separates this oxide from the class of the alkaline earths. But on arranging the oxides of the above-described metals of the alkaline earths according to their decreasing atomic weights we have the seriesBaO, SrO, CaO, MgO, BeO,in which the basic properties and solubility of the oxides consecutively and distinctly decrease until we reach a point when, had we not known of the existence of the beryllium oxide, we should expect to find in its place an oxide insoluble in water and of feeble basic properties. If an alcoholic solution of caustic potash be saturated with the hydrate of BeO, and evaporated under the receiver of an air pump, it forms silky crystals BeK2O2.Another characteristic of the salts of beryllium is that they give with aqueous ammonia a gelatinous precipitate which is soluble in an excess of ammonium carbonate like the precipitate of magnesia; in this beryllium oxide differs from the oxide of aluminium. Beryllium oxide easily forms a carbonate which is insoluble in water, and resembles magnesium carbonate in many respects. Beryllium sulphate is distinguished by its considerable solubility in water—thus, at the ordinary temperature it dissolves in an equal weight of water; it crystallises out from its solutions in well-formed crystals, which do not change in the air, and contain BeSO4,4H2O. When ignited it leaves beryllium oxide, but this oxide, after prolonged ignition, is re-dissolved by sulphuric acid, whilst aluminium sulphate, after a similar treatment, leaves aluminium oxide, which is no longer soluble in acids. With a few exceptions, the salts of beryllium crystallise with great difficulty, and to a considerable extent resemble the salts of magnesium; thus, for instance, beryllium chloride is analogous to magnesium chloride. It is volatile in an anhydrous state, and in a hydrated state it decomposes, with the evolution of hydrochloric acid.
[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 MgK2(SO4)2,6H2O lose all their water at 135°, whilst MgSO4,7H2O only parts with 6H2O.
[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 MgK2(SO4)2,6H2O lose all their water at 135°, whilst MgSO4,7H2O only parts with 6H2O.
[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.
[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, MgSO4+ 2NaCl + H2O = Na2SO4+ 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.
[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, MgSO4+ 2NaCl + H2O = Na2SO4+ 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 MgX2, it follows, according to Berthollet's teaching, that MgCl2is also present.
[31]As sea-water contains many salts, MCl and MgX2, it follows, according to Berthollet's teaching, that MgCl2is also present.
[32]As the crystallo-hydrates of the salts of sodium often contain 10H2O, so many of the salts of magnesium contain 6H2O.
[32]As the crystallo-hydrates of the salts of sodium often contain 10H2O, so many of the salts of magnesium contain 6H2O.
[33]This decomposition is most simply defined as the result of the two reverse reactions, MgCl2÷ H2O = MgO + 2HCl and MgO + 2HCl = MgCl2+ H2O, or as a distribution between O and Cl2on the one hand and H2and Mg on the other. (With O, MgCl2gives chlorine,seeChapter X., Note33, and Chapter II., Note 3 bis and others, where the reactions and applications of MgCl2are 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, MgCl2,6H2O, 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 (seeChapter 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 MgBr2crystallises not only with 6H2O (temperature of fusion 152°), but also with 10H2O (temperature of fusion +12°, formed at -18°). (Panfiloff, 1894).
[33]This decomposition is most simply defined as the result of the two reverse reactions, MgCl2÷ H2O = MgO + 2HCl and MgO + 2HCl = MgCl2+ H2O, or as a distribution between O and Cl2on the one hand and H2and Mg on the other. (With O, MgCl2gives chlorine,seeChapter X., Note33, and Chapter II., Note 3 bis and others, where the reactions and applications of MgCl2are 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, MgCl2,6H2O, 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 (seeChapter 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 MgBr2crystallises not only with 6H2O (temperature of fusion 152°), but also with 10H2O (temperature of fusion +12°, formed at -18°). (Panfiloff, 1894).
[34]According to Thomsen, the combination of MgCl2with 6H2O evolves 33,000 calories, and its solution in an excess of water 36,000.
[34]According to Thomsen, the combination of MgCl2with 6H2O evolves 33,000 calories, and its solution in an excess of water 36,000.
[34 bis]Hence MgCl2may be employed for the preparation of chlorine and hydrochloric acid (ChaptersX. andXI.). In general magnesium chloride, which is obtained in large quantities from sea water and Stassfurt carnallite, may find numerous practical uses.
[34 bis]Hence MgCl2may be employed for the preparation of chlorine and hydrochloric acid (ChaptersX. andXI.). In general magnesium chloride, which is obtained in large quantities from sea water and Stassfurt carnallite, may find numerous practical uses.
[35]There are many other methods of separating calcium from magnesium besides that mentioned above (Note22). Among them it will be sufficient to mention the behaviour of these bases towards a solution of sugar; hydratedlimeis exceedinglysoluble 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.
[35]There are many other methods of separating calcium from magnesium besides that mentioned above (Note22). Among them it will be sufficient to mention the behaviour of these bases towards a solution of sugar; hydratedlimeis exceedinglysoluble 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 (Note50).
[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 (Note50).
[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.
[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, calledpoorlime in distinction from that obtained from purer limestone, which is calledrichlime. 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, CaN2O6, 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: CaN2O6= CaO + 2NO2+ 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.
[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, calledpoorlime in distinction from that obtained from purer limestone, which is calledrichlime. 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, CaN2O6, 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: CaN2O6= CaO + 2NO2+ 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.
[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.
[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 oflimeorcement, or the so-calledslaked 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).
[40]The greater quantity of lime is used in making mortar for binding bricks or stones together, in the form oflimeorcement, or the so-calledslaked 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.
[41]Professor Glinka measured the transparent bright crystals of calcium hydroxide which are formed in common hydraulic (Portland) cement.
[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, H2SO4(containing neither free water nor anhydride, SO3). 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.
[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, H2SO4(containing neither free water nor anhydride, SO3). 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.
[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, themodus operandiof 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 conceptionof the mechanical theory of heatis applicable to dissociationandevaporation.
[44]Before the introduction of Deville's theory of dissociation, themodus operandiof 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 conceptionof the mechanical theory of heatis applicable to dissociationandevaporation.
[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 acrystallo-hydrate, CaCO3,5H2O (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 anacid salt, CaH2(CO3)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.
[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 acrystallo-hydrate, CaCO3,5H2O (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 anacid salt, CaH2(CO3)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.
[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).
[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).
[47]According to Le Chatelier (1888), 1½H2O is lost at 120°—that is, H2O,2CaSO4is 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 H2O,2CaSO4is also formed when gypsum is heated with water in a closed vessel at 150° (Hoppe-Seyler).
[47]According to Le Chatelier (1888), 1½H2O is lost at 120°—that is, H2O,2CaSO4is 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 H2O,2CaSO4is 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]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 (Note47), and he explains by their relation to water the phenomena observed in the solidification of a mixture of burnt gypsum and water.
[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 (Note47), 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 CaSO4,2H2O, which contain as much as 1 part of CaSO4to 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 ChapterX. 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 sulphate, 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 CaK2(SO4)2,H2O; 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 asglauberite. Fritzsche also obtainedgaylussite, Na2Ca(CO3)2,5H2O, 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(NO3)2Ca(OH)2,2½H2O by leaving powdered caustic lime in a saturated solution of Ca(NO3)2until it solidified. This salt is decomposed by water.
[49]As Marignac showed, gypsum, especially when desiccated at 120°, easily gives supersaturated solutions with respect to CaSO4,2H2O, which contain as much as 1 part of CaSO4to 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 ChapterX. 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 sulphate, 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 CaK2(SO4)2,H2O; 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 asglauberite. Fritzsche also obtainedgaylussite, Na2Ca(CO3)2,5H2O, 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(NO3)2Ca(OH)2,2½H2O by leaving powdered caustic lime in a saturated solution of Ca(NO3)2until it solidified. This salt is decomposed by water.
[50]Calcium chloride has a specific gravity 2·20, or, when fused, 2·12, and the sp. gr. of the crystallised salt CaCl2,6H2O 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,seeNote27). The solution containing 50 p.c. CaCl2boils 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 CaCl2,6H2O 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 CaCl2,4H2Oβ are formed, which at temperatures above 38·4° are converted into the crystallo-hydrates CaCl2,2H2O, whilst at temperatures below 18° the β variety passes into the more stable CaCl2,4H2Oα, which process is aided by mechanical friction. Hence, as is the case with magnesium sulphate (Note27), one and the same crystallo-hydrate appears in two forms—the β, which is easily produced but is unstable, and the α, 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°CaCl2,6H2O6075100(102·8)CaCl2,4H2Oα—90101117(154·2)CaCl2,4H2Oβ—104114—CaCl2,2H2O——(308·3)128137The 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 4H2O and 2H2O. The crystals CaCl2,2H2O 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 CaCl2,H2O 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 CaCl2,4H2Oβ, we will give the temperaturestat which the passage of one hydrate into another takes place and at which the solution CaCl2+nH2O, the two solids A and B and aqueous vapour, whose tension is given aspin millimetres, are able to exist together in stable equilibrium, according to Roozeboom's determinations:tnABp-55°14·5iceCaCl2,6H2O0+29·8°6·1CaCl2,6H2OCaCl2,4H2O6·845·3°4·7CaCl2,4H2OCaCl2,2H2O11·8175·5°2·1CaCl2,2H2OCaCl2,H2O842200°1·8CaCl2,H2OCaCl2Several atmospheresSolutions 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 CaCl2,6H2O.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 CaCl2,6H2O separates, and that a solution of the above composition (CaCl2,14H2O requires 44·0 parts calcium chloride per 100 of water) solidifies as a cryohydrate at about -55°.
[50]Calcium chloride has a specific gravity 2·20, or, when fused, 2·12, and the sp. gr. of the crystallised salt CaCl2,6H2O 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,seeNote27). The solution containing 50 p.c. CaCl2boils 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 CaCl2,6H2O 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 CaCl2,4H2Oβ are formed, which at temperatures above 38·4° are converted into the crystallo-hydrates CaCl2,2H2O, whilst at temperatures below 18° the β variety passes into the more stable CaCl2,4H2Oα, which process is aided by mechanical friction. Hence, as is the case with magnesium sulphate (Note27), one and the same crystallo-hydrate appears in two forms—the β, which is easily produced but is unstable, and the α, 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:—
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 4H2O and 2H2O. The crystals CaCl2,2H2O 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 CaCl2,H2O 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 CaCl2,4H2Oβ, we will give the temperaturestat which the passage of one hydrate into another takes place and at which the solution CaCl2+nH2O, the two solids A and B and aqueous vapour, whose tension is given aspin millimetres, are able to exist together in stable equilibrium, according to Roozeboom's determinations:
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 CaCl2,6H2O.
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 CaCl2,6H2O separates, and that a solution of the above composition (CaCl2,14H2O requires 44·0 parts calcium chloride per 100 of water) solidifies as a cryohydrate at about -55°.
[51]The action of barium sulphate on sodium and potassium carbonates is given on p. 437.
[51]The action of barium sulphate on sodium and potassium carbonates is given on p. 437.
[52]Barium sulphide is decomposed by water, BaS + 2H2O = H2S + 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 BaH2O2and SrH2O2are 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,K2S, 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 BaSO4after being diluted with water.
[52]Barium sulphide is decomposed by water, BaS + 2H2O = H2S + 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 BaH2O2and SrH2O2are 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,K2S, 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 BaSO4after 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, MnCl2, 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 meansbarium sulphateis 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 BaCl2,2H2O and SrCl2,6H2O. The latter (which separates out at 40°) resembles the salts of Ca and Mg in composition, and Étard (1892) obtained SrCl2,2H2O from solutions at 90–130°. We may also observe that the crystallo-hydrates BaBr2,H2O and BaI2,7H2O are known.
[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, MnCl2, 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 meansbarium sulphateis 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 BaCl2,2H2O and SrCl2,6H2O. The latter (which separates out at 40°) resembles the salts of Ca and Mg in composition, and Étard (1892) obtained SrCl2,2H2O from solutions at 90–130°. We may also observe that the crystallo-hydrates BaBr2,H2O and BaI2,7H2O are known.
[54]The nitrates Sr(NO3)2(in the cold its solutions give a crystallo-hydrate containing 4H2O) and Ba(NO3)2are 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).
[54]The nitrates Sr(NO3)2(in the cold its solutions give a crystallo-hydrate containing 4H2O) and Ba(NO3)2are 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 Chapter I., Note65. 100 parts of water dissolve0°20°40°60°80°BaO1·53·57·418·890·8SrO0·30·71·439Supersaturated solutions are easily formed.The anhydrous oxide BaO fuses in the oxyhydrogen flame. When ignited in the vapour of potassium, the latter takes up the oxygen; whilst in chlorine, oxygen is separated and barium chloride formed.
[55]The dissociation of the crystallo-hydrate of baryta is given in Chapter I., Note65. 100 parts of water dissolve
Supersaturated solutions are easily formed.
The anhydrous oxide BaO fuses in the oxyhydrogen flame. When ignited in the vapour of potassium, the latter takes up the oxygen; whilst in chlorine, oxygen is separated and barium chloride formed.
[55 bis]Brugellmann, by heating BaH2O2in a graphite or clay crucible, obtained BaO in needles, sp. gr. 5·32, and by heating in a platinum crucible—in crystals belonging to the cubical system, sp. gr. 5·74. SrO is obtained in the latter form from the nitrate. The following are the specific gravities of the oxides from different sources:—MgOCaOSrOfromRN2O63·383·254·75„RCO33·483·264·45„RH2O23·413·254·57
[55 bis]Brugellmann, by heating BaH2O2in a graphite or clay crucible, obtained BaO in needles, sp. gr. 5·32, and by heating in a platinum crucible—in crystals belonging to the cubical system, sp. gr. 5·74. SrO is obtained in the latter form from the nitrate. The following are the specific gravities of the oxides from different sources:—
[56]The property of barium oxide of absorbing oxygen when heated, and giving the peroxide, BaO2, is very characteristic for this oxide (seeChapter III., Note7). It only belongs to the anhydrous oxide. The hydroxide does not absorb oxygen. Peroxides of calcium and strontium may be obtained by means of hydrogen peroxide. Barium peroxide is insoluble in water, but is able to form a hydrate with it, and also to combine with hydrogen peroxide, forming a very unstable compound having the composition BaH2O4(obtained by Professor Schöne), which in course of time evolves oxygen (Chapter IV., Note21).
[56]The property of barium oxide of absorbing oxygen when heated, and giving the peroxide, BaO2, is very characteristic for this oxide (seeChapter III., Note7). It only belongs to the anhydrous oxide. The hydroxide does not absorb oxygen. Peroxides of calcium and strontium may be obtained by means of hydrogen peroxide. Barium peroxide is insoluble in water, but is able to form a hydrate with it, and also to combine with hydrogen peroxide, forming a very unstable compound having the composition BaH2O4(obtained by Professor Schöne), which in course of time evolves oxygen (Chapter IV., Note21).
[57]Even in solutions a gradual progression in the increase of the specific gravity shows itself, not only for equivalent solutions (for instance, RCl2+ 200H2O), but even with an equal percentage composition, as is seen from the curves giving the specific gravity (water 4° = 10,000) at 15° (for barium chloride, according to Bourdiakoff's determinations):BeCl2: S = 9,992 + 67·21p+ 0·111p2CaCl2: S = 9,992 + 80·24p+ 0·476p2SrCl2: S = 9,992 + 85·57p+ 0·733p2BaCl2: S = 9,992 + 86·56p+ 0·813p2
[57]Even in solutions a gradual progression in the increase of the specific gravity shows itself, not only for equivalent solutions (for instance, RCl2+ 200H2O), but even with an equal percentage composition, as is seen from the curves giving the specific gravity (water 4° = 10,000) at 15° (for barium chloride, according to Bourdiakoff's determinations):
BeCl2: S = 9,992 + 67·21p+ 0·111p2CaCl2: S = 9,992 + 80·24p+ 0·476p2SrCl2: S = 9,992 + 85·57p+ 0·733p2BaCl2: S = 9,992 + 86·56p+ 0·813p2
[58]One part of calcium sulphate at the ordinary temperature requires about 500 parts of water for solution, strontium sulphate about 7,000 parts, barium sulphate about 400,000 parts, whilst beryllium sulphate is easily soluble in water.
[58]One part of calcium sulphate at the ordinary temperature requires about 500 parts of water for solution, strontium sulphate about 7,000 parts, barium sulphate about 400,000 parts, whilst beryllium sulphate is easily soluble in water.
[59]We refer beryllium to the class of the bivalent metals of the alkaline earths—that is, we ascribe to its oxide the formula BeO, and do not consider it as trivalent (Be = 13·5, Chapter VII., Note21), although that view has been upheld by many chemists. The true atomic composition of beryllium oxide was first given by the Russian chemist, Avdéeff (1819), in his researches on the compounds of this metal. He compared the compounds of beryllium to those of magnesium, and refuted the notion prevalent at the time, of the resemblance between the oxides of beryllium and aluminium, by proving that beryllium sulphate presents a greater resemblance to magnesium sulphate than to aluminium sulphate. It was especially noticed that the analogues of alumina give alums, whilst beryllium oxide, although it is a feeble base, easily giving, like magnesia, basic and double salts, does not form true alums. The establishment of the periodic system of the elements (1869), considered in the following chapter, immediately indicated that Avdéeff's view corresponded with the truth—that is, that beryllium is bivalent, which therefore necessitated the denial of its trivalency. This scientific controversy resulted in a long series of researches (1870–80) concerning this element, and ended in Nilson and Pettersson—two of the chief advocates of the trivalency of beryllium—determining the vapour density of BeCl2= 40, (Chapter VII., Note21), which gave an undoubted proof of its bivalency (seealso Note3).
[59]We refer beryllium to the class of the bivalent metals of the alkaline earths—that is, we ascribe to its oxide the formula BeO, and do not consider it as trivalent (Be = 13·5, Chapter VII., Note21), although that view has been upheld by many chemists. The true atomic composition of beryllium oxide was first given by the Russian chemist, Avdéeff (1819), in his researches on the compounds of this metal. He compared the compounds of beryllium to those of magnesium, and refuted the notion prevalent at the time, of the resemblance between the oxides of beryllium and aluminium, by proving that beryllium sulphate presents a greater resemblance to magnesium sulphate than to aluminium sulphate. It was especially noticed that the analogues of alumina give alums, whilst beryllium oxide, although it is a feeble base, easily giving, like magnesia, basic and double salts, does not form true alums. The establishment of the periodic system of the elements (1869), considered in the following chapter, immediately indicated that Avdéeff's view corresponded with the truth—that is, that beryllium is bivalent, which therefore necessitated the denial of its trivalency. This scientific controversy resulted in a long series of researches (1870–80) concerning this element, and ended in Nilson and Pettersson—two of the chief advocates of the trivalency of beryllium—determining the vapour density of BeCl2= 40, (Chapter VII., Note21), which gave an undoubted proof of its bivalency (seealso Note3).
[60]Beryllium oxide, like aluminium oxide, is precipitated from solutions of its salts by alkalis as a gelatinous hydroxide, BeH2O2, which, like alumina, is soluble in an excess of caustic potash or soda. This reaction may be taken advantage of for distinguishing and separating beryllium from aluminium, because when the alkaline solution is diluted with water and boiled, beryllium hydroxide is precipitated, whilst the alumina remains in solution. The solubility of the beryllium oxide at once clearly indicates its feeble basic properties, and, as it were, separates this oxide from the class of the alkaline earths. But on arranging the oxides of the above-described metals of the alkaline earths according to their decreasing atomic weights we have the seriesBaO, SrO, CaO, MgO, BeO,in which the basic properties and solubility of the oxides consecutively and distinctly decrease until we reach a point when, had we not known of the existence of the beryllium oxide, we should expect to find in its place an oxide insoluble in water and of feeble basic properties. If an alcoholic solution of caustic potash be saturated with the hydrate of BeO, and evaporated under the receiver of an air pump, it forms silky crystals BeK2O2.Another characteristic of the salts of beryllium is that they give with aqueous ammonia a gelatinous precipitate which is soluble in an excess of ammonium carbonate like the precipitate of magnesia; in this beryllium oxide differs from the oxide of aluminium. Beryllium oxide easily forms a carbonate which is insoluble in water, and resembles magnesium carbonate in many respects. Beryllium sulphate is distinguished by its considerable solubility in water—thus, at the ordinary temperature it dissolves in an equal weight of water; it crystallises out from its solutions in well-formed crystals, which do not change in the air, and contain BeSO4,4H2O. When ignited it leaves beryllium oxide, but this oxide, after prolonged ignition, is re-dissolved by sulphuric acid, whilst aluminium sulphate, after a similar treatment, leaves aluminium oxide, which is no longer soluble in acids. With a few exceptions, the salts of beryllium crystallise with great difficulty, and to a considerable extent resemble the salts of magnesium; thus, for instance, beryllium chloride is analogous to magnesium chloride. It is volatile in an anhydrous state, and in a hydrated state it decomposes, with the evolution of hydrochloric acid.
[60]Beryllium oxide, like aluminium oxide, is precipitated from solutions of its salts by alkalis as a gelatinous hydroxide, BeH2O2, which, like alumina, is soluble in an excess of caustic potash or soda. This reaction may be taken advantage of for distinguishing and separating beryllium from aluminium, because when the alkaline solution is diluted with water and boiled, beryllium hydroxide is precipitated, whilst the alumina remains in solution. The solubility of the beryllium oxide at once clearly indicates its feeble basic properties, and, as it were, separates this oxide from the class of the alkaline earths. But on arranging the oxides of the above-described metals of the alkaline earths according to their decreasing atomic weights we have the series
BaO, SrO, CaO, MgO, BeO,
in which the basic properties and solubility of the oxides consecutively and distinctly decrease until we reach a point when, had we not known of the existence of the beryllium oxide, we should expect to find in its place an oxide insoluble in water and of feeble basic properties. If an alcoholic solution of caustic potash be saturated with the hydrate of BeO, and evaporated under the receiver of an air pump, it forms silky crystals BeK2O2.
Another characteristic of the salts of beryllium is that they give with aqueous ammonia a gelatinous precipitate which is soluble in an excess of ammonium carbonate like the precipitate of magnesia; in this beryllium oxide differs from the oxide of aluminium. Beryllium oxide easily forms a carbonate which is insoluble in water, and resembles magnesium carbonate in many respects. Beryllium sulphate is distinguished by its considerable solubility in water—thus, at the ordinary temperature it dissolves in an equal weight of water; it crystallises out from its solutions in well-formed crystals, which do not change in the air, and contain BeSO4,4H2O. When ignited it leaves beryllium oxide, but this oxide, after prolonged ignition, is re-dissolved by sulphuric acid, whilst aluminium sulphate, after a similar treatment, leaves aluminium oxide, which is no longer soluble in acids. With a few exceptions, the salts of beryllium crystallise with great difficulty, and to a considerable extent resemble the salts of magnesium; thus, for instance, beryllium chloride is analogous to magnesium chloride. It is volatile in an anhydrous state, and in a hydrated state it decomposes, with the evolution of hydrochloric acid.