Footnotes:[1]Whilst describing in some detail the properties of sodium chloride, hydrochloric acid, and sodium sulphate, I wish to impart, by separate examples, an idea of the properties of saline substances, but the dimensions of this work and its purpose and aim do not permit of entering into particulars concerning every salt, acid, or other substance. The fundamental object of this work—an account of the characteristics of the elements and an acquaintance with the forces acting between atoms—has nothing to gain from the multiplication of the number of as yet ungeneralised properties and relations.[2]Anhydrous (ignited) sodium sulphate, Na2SO4, is known in trade as ‘sulphate’ or salt-cake, in mineralogythenardite. Crystalline decahydrated salt is termed in mineralogymirabilite, and in trade Glauber's salt. On fusing it, the monohydrate Na2SO4H2O is obtained, together with a supersaturated solution.[3]The salts may be obtained not only by methods of substitution of various kinds, but also by many other combinations. Thus sodium sulphate may be formed from sodium oxide and sulphuric anhydride, by oxidising sodium sulphide, Na2S, or sodium sulphite, Na2SO3, &c. When sodium chloride is heated in a mixture of the vapours of water, air, and sulphurous anhydride, sodium sulphate is formed. According to this method (patented by Hargreaves and Robinson), sodium sulphate, Na2SO4, is obtained from NaCl without the preliminary manufacture of H2SO4. Lumps of NaCl pressed into bricks are loosely packed into a cylinder and subjected, at a red heat, to the action of steam, air and SO2. Under these conditions, HCl, sulphate, and a certain amount of unaltered NaCl are obtained. This mixture is converted into soda by Gossage's process (seeNote15) and may have some practical value.[4]Many observations have been made, but little general information has been obtained from particular cases. In addition to which, the properties of a given salt are changed by the presence of other salts. This takes place not only in virtue of mutual decomposition or formation of double salts capable of separate existence, but is determined by the influence which some salts exert on others, or by forces similar to those which act during solution. Here nothing has been generalised to that extent which would render it possible to predict without previous investigation, if there be no close analogy to help us. Let us state one of these numerous cases: 100 parts of water at 20° dissolve 34 parts of potassium nitrate but on the addition of sodium nitrate the solubility of potassium nitrate increases to 48 parts in 10 of water (Carnelley and Thomson). In general, in all cases of which there are accurate observations it appears that the presence of foreign salts changes the properties of any given salt.[5]The information concerning solubility (ChapterI.) is given according to the determinations of Gay-Lussac, Lovell, and Mulder.[6]In Chapter I., Note24, we have already seen that with many other sulphates the solubility also decreases after a certain temperature is passed. Gypsum, CaSO4,2H2O, lime, and many other compounds present such a phenomenon. An observation of Tilden's (1884) is most instructive; he showed that on raising the temperature (in closed vessels) above 140° the solubility of sodium sulphate again begins to increase. At 100° 100 parts of water dissolve about 43 parts of anhydrous salt, at 140° 42 parts, at 160° 43 parts, at 180° 44 parts, at 230° 46 parts. According to Étard (1892) the solubility of 30 parts of Na2SO4in 100 of solution (or 43 per 100 of water) corresponds to 80°, and above 240° the solubility again falls, and very rapidly, so that at 320° the solution contains 12 per 100 of solution (about 14 per 100 of water) and a further rise of temperature is followed by a further deposition of the salt. It is evident that the phenomenon of saturation, determined by the presence of an excess of the dissolved substance, is very complex, and therefore that for the theory of solutions considered as liquid indefinite chemical compounds, many useful statements can hardly be given.[7]Already referred to in Chapter I., Note56.The example of sodium sulphate is historically very important for the theory of solutions. Notwithstanding the number of investigations which have been made, it is still insufficiently studied, especially from the point of the vapour tension of solutions and crystallo-hydrates, so that those processes cannot be applied to it which Guldberg, Roozeboom, Van't Hoff, and others applied to solutions and crystallo-hydrates. It would also be most important to investigate the influence of pressure on the various phenomena corresponding with the combinations of water and sodium sulphate, because when crystals are separated—for instance, of the decahydrated salt—an increase of volume takes place, as can be seen from the following data:—the sp. gr. of the anhydrous salt is 2·66, that of the decahydrated salt = 1·46, but the sp. gr. of solutions at 15°/4° = 9,992 + 90·2p+ 0·35p2where p represents the percentage of anhydrous salt in the solution, and the sp. gr. of water at 4° = 10,000. Hence for solutions containing 20 p.c. of anhydrous salt the sp. gr. = 1·1936; therefore the volume of 100 grams of this solution = 83·8 c.c., and the volume of anhydrous salt contained in it is equal to 20/2·66, or = 7·5 c.c., and the volume of water = 80·1 c.c. Therefore, the solution, on decomposing into anhydrous salt and water, increases in volume (from 83·8 to 87·6); but in the same way 83·8 c.c. of 20 p.c. solution are formed from (45·4/1·46 =) 31·1 c.c. of the decahydrated salt, and 54·6 c.c. of water—that is to say, that during the formation of a solution from 85·7 c.c., 83·8 c.c. are formed.[8]From this example it is evident the solution remains unaltered until from the contact of a solid it becomes either saturated or supersaturated, crystallisation being determined by the attraction to a solid, as the phenomenon of supersaturation clearly demonstrates. This partially explains certain apparently contradictory determinations of solubility. The best investigated example of such complex relations is cited in Chapter XIV., Note50(for CaCl2).[9]According to Pickering's experiments (1886), the molecular weight in grams (that is, 142 grams) of anhydrous sodium sulphate, on being dissolved in a large mass of water, at 0° absorbs (hence the - sign) -1,100 heat units, at 10°–700, at 15°–275, at 20° gives out +25, at 25° +300 calories. For the decahydrated salt, Na2SO4,10H2O, 5° - 4,225, 10° - 4,000, 15° - 3,570, 20° - 3,160, 25° - 2,775. Hence (just as in Chapter I., Note56) the heat of the combination Na2SO4,10H2O at 5° = +3,125, 10° = +3,250, 20° = +3,200, and 25° = +3,050.It is evident that the decahydrated salt dissolving in water gives a decrease of temperature. Solutions in hydrochloric acid give a still greater decrease, because they contain the water of crystallisation in a solid state—that is, like ice—and this on melting absorbs heat. A mixture of 15 parts of Na2SO4,10H2O and 12 parts of strong hydrochloric acid produces sufficient cold to freeze water. During the treatment with hydrochloric acid a certain quantity of sodium chloride is formed.[10]The very large and well-formed crystals of this salt resemble the hydrate H2SO4,H2O, or SO(OH)4. In general the replacement of hydrogen by sodium modifies many of the properties of acids less than its replacement by other metals. This most probably depends on the volumes being nearly equal.[11]In solution (Berthelot) the acid salt in all probability decomposes most in the greatest mass of water. The specific gravity (according to the determinations of Marignac) of solutions at 15°/4° = 9,992 + 77·92p+ 0·231p2(seeNote7). From these figures, and from the specific gravities of sulphuric acid, it is evident that on mixing solutions of this acid and sodium sulphateexpansionwill always take place; for instance, H2SO4+ 25H2O with Na2SO4+ 25H2O increases from 483 volumes to 486. In addition to which, in weak solutions heat is absorbed, as shown in Chapter X., Note27. Nevertheless, even more acid salts may be formed and obtained in a crystalline form. For instance, on cooling a solution of 1 part of sodium sulphate in 7 parts of sulphuric acid, crystals of the composition NaHSO4,H2SO4are separated (Schultz, 1868). This compound fuses at about 100°; the ordinary acid salt, NaHSO4, at 149°.[11 bis]On decreasing the pressure, sodium hydrogen sulphate, NaHSO4, dissociates much more easily than at the ordinary pressure; it loses water and forms the pyrosulphate, Na2S2O7; this reaction is utilised in chemical works.[12]Calcium sulphide, CaS, like many metallic sulphides which are soluble in water, is decomposed by it (ChapterX.), CaS + H2O = CaO + H2S, because hydrogen sulphide is a very feeble acid. If calcium sulphide be acted on by a large mass of water, lime may be precipitated, and a state of equilibrium will be reached, when the system CaO + 2CaS remains unchanged. Lime, being a product of the action of water on CaS, limits this action. Therefore, if in black ash the lime were not in excess, a part of the sulphide would be in solution (actually there is but very little). In this manner in the manufacture of sodium carbonate the conditions of equilibrium which enter into double decompositions have been made use of (see above), and the aim is to form directly the unchangeable product CaO,2CaS. This was first regarded as a special insoluble compound, but there is no evidence of its independent existence.[13]see captionFig.69.—Apparatus for the methodical lixiviation of black ash, &c. Water flows into the tanks from the pipesr,r, and the saturated liquid is drawn off fromc,c.Methodical lixiviationis the extraction, by means of water, of a soluble substance from the mass containing it. It is carried on so as not to obtain weak aqueous solutions, and in such a way that the residue shall not contain any of the soluble substance. This problem is practically of great importance in many industries. It is required to extract from the mass all that is soluble in water. This is easily effected if water be first poured on the mass, the strong solution thus obtained decanted, then water again poured on, time being allowed for it to act, then again decanted, and so on until fresh water does not take up anything. But then finally such weak solutions are obtained that it would be very disadvantageous to evaporate them. This is avoided by pouring the fresh hot water destined for the lixiviation, not onto the fresh mass, but upon a mass which has already been subjected to a first lixiviation by weak solutions. In this way the fresh water gives a weak solution. The strong solution which goes to the evaporating pan flows from those parts of the apparatus which contain the fresh, as yet unlixiviated, mass, and thus in the latter parts the weak alkali formed in the other parts of the apparatus becomes saturated as far as possible with the soluble substance. Generally several intercommunicating vessels are constructed (standing at the same level) into which in turn the fresh mass is charged which is intended for lixiviation; the water is poured in, the alkali drawn off, and the lixiviated residue removed. The illustration represents such an apparatus, consisting of four communicating vessels. The water poured into one of them flows through the two nearest and issues from the third. The fresh mass being placed in one of these boxes or vessels, the stream of water passing through the apparatus is directed in such a manner as to finally issue from this vessel containing the fresh unlixiviated mass. The fresh water is added to the vessel containing the material which has been almost completely exhausted. Passing through this vessel it is conveyed by the pipe (syphon passing from the bottom of the first box to the top of the second) communicating with the second; it finally passes (also through a syphon pipe) into the box (the third) containing the fresh material. The water will extract all that is soluble in the first vessel, leaving only an insoluble residue. This vessel is then ready to be emptied, and refilled with fresh material. The levels of the liquids in the various vessels will naturally be different, in consequence of the various strengths of the solutions which they contain.It must not, however, be thought that sodium carbonate alone passes into the solution; there is also a good deal of caustic soda with it, formed by the action of lime on the carbonate of sodium, and there are also certain sodium sulphur compounds with which we shall partly become acquainted hereafter. The sodium carbonate, therefore, is not obtained in a very pure state. The solution is concentrated by evaporation. This is conducted by means of the waste heat from the soda furnaces, together with that of the gases given off. The process in the soda furnaces can only be carried on at a high temperature, and therefore the smoke and gases issuing from them are necessarily very hot. If the heat they contain was not made use of there would be a great waste of fuel; consequently in immediate proximity to these furnaces there is generally a series of pans or evaporating boilers, under which the gases pass, and into which the alkali solution is poured. On evaporating the solution, first of all the undecomposed sodium sulphate separates, then the sodium carbonate or soda crystals. These crystals as they separate are raked out and placed on planks, where the liquid drains away from them. Caustic soda remains in the residue, and also any sodium chloride which was not decomposed in the foregoing process.Part of the sodium carbonate is recrystallised in order to purify it more thoroughly. In order to do this a saturated solution is left to crystallise at a temperature below 30° in a current of air, in order to promote the separation of the water vapour. The large transparent crystals (efflorescent in air) of Na2CO3,10H2O are then formed which have already been spoken of (ChapterI.).[14]The whole of the sulphur used in the production of the sulphuric acid employed in decomposing the common salt is contained in this residue. This is the great burden and expense of the soda works which use Leblanc's method. As an instructive example from a chemical point of view, it is worth while mentioning here two of the various methods of recovering the sulphur from the soda waste. Chance's process is treated in Chapter XX., Note 6.Kynaston (1885) treats the soda waste with a solution (sp. gr. l·21) of magnesium chloride, which disengages sulphuretted hydrogen: CaS + MgCl2+ 2H2O = CaCl2+ Mg(OH)2+ H2S. Sulphurous anhydride is passed through the residue in order to form the insoluble calcium sulphite: CaCl2+ Mg(OH)2+ SO2= CaSO3+ MgCl2+ H2O. The solution of magnesium chloride obtained is again used, and the washed calcium sulphite is brought into contact at a low temperature with hydrochloric acid (a weak aqueous solution) and hydrogen sulphide, the whole of the sulphur then separating:CaSO3+ 2H2S + 2HCl = CaCl2+ 3H2O + 3S.But most efforts have been directed towards avoiding the formation of soda waste.[15]Among the drawbacks of the Leblanc process are the accumulation of ‘soda waste’ (Note14) owing to the impossibility at the comparatively low price of sulphur (especially in the form of pyrites) of finding employment for the sulphur and sulphur compounds for which this waste is sometimes treated, and also the insufficient purity of the sodium carbonate for many purposes. The advantages of the Leblanc process, besides its simplicity and cheapness, are that almost the whole of the acids obtained as bye-products have a commercial value; for chlorine and bleaching powder are produced from the large amount of hydrochloric acid which appears as a bye-product; caustic soda also is very easily made, and the demand for it increases every year. In those places where salt, pyrites, charcoal, and limestone (the materials required for alkali works) are found side by side—as, for instance, in the Ural or Don districts—conditions are favourable to the development of the manufacture of sodium carbonate on an enormous scale; and where, as in the Caucasus, sodium sulphate occurs naturally, the conditions are still more favourable. A large amount, however, of the latter salt, even from soda works, is used in making glass. The most important soda works, as regards the quantity of products obtained from them, are the English works.As an example of the other numerous and varied methods of manufacturing soda from sodium chloride, the following may be mentioned: Sodium chloride is decomposed by oxide of lead, PbO, forming lead chloride and sodium oxide, which, with carbonic anhydride, yields sodium carbonate (Scheele's process). In Cornu's method sodium chloride is treated with lime, and then exposed to the air, when it yields a small quantity of sodium carbonate. In E. Kopp's process sodium sulphate (125 parts) is mixed with oxide of iron (80 parts) and charcoal (55 parts), and the mixture is heated in reverberatory furnaces. Here a compound, Na6Fe4S3, is formed, which is insoluble in water absorbs oxygen and carbonic anhydride, and then forms sodium carbonate and ferrous sulphide; this when roasted gives sulphurous anhydride, the indispensable material for the manufacture of sulphuric acid, and ferric oxide which is again used in the process. In Grant's method sodium sulphate is transformed into sodium sulphide, and the latter is decomposed by a stream of carbonic anhydride and steam, when hydrogen sulphide is disengaged and sodium carbonate formed. Gossage prepares Na2S from Na2SO4(by heating it with carbon), dissolves it in water and subjects the solution to the action of an excess of CO2in coke towers, thus obtaining H2S (a gas which gives SO2under perfect combustion, or sulphur when incompletely burnt, Chapter XX., Note 6) and bicarbonate of sodium; Na2S + 2CO2+ 2H2O = H2S + 2HNaCO3. The latter gives soda and CO2when ignited. This process quite eliminates the formation of soda-waste (seeNote3) and should in my opinion be suitable for the treatment of native Na2SO4, like that which is found in the Caucasus, all the more since H2S gives sulphur as a bye-product.Repeated efforts have been made in recent times to obtain soda (and chlorine,seeChapter II., Note1) from strong solutions of salt (Chapter X., Note23 bis) by the action of an electric current, but until now these methods have not been worked out sufficiently for practical use, probably partly owing to the complicated apparatus needed, and the fact that the chlorine given off at the anode corrodes the electrodes and vessels and has but a limited industrial application. We may mention that according to Hempel (1890) soda in crystals is deposited when an electric current and a stream of carbonic acid gas are passed through a saturated solution of NaCl.Sodium carbonate may likewise be obtained from cryolite (Chapter XVII., Note 23) the method of treating this will be mentioned under Aluminium.[16]This process (Chapter XVII.) was first pointed out by Turck, worked out by Schloesing, and finally applied industrially by Solvay. The first (1883) large soda factories erected in Russia for working this process are on the banks of the Kama at Berezniak, near Ousolia, and belong to Lubimoff. But Russia, which still imports from abroad a large quantity of bleaching powder and exports a large amount of manganese ore, most of all requires works carrying on the Leblanc process. In 1890 a factory of this kind was erected by P. K. Oushkoff, on the Kama, near Elagoubi.[16 bis]Mond (seeChapter XI., Note3 bis) separates the NH4Cl from the residual solutions by cooling (Chapter X., Note44); ignites the sal-ammoniac and passes the vapour over MgO, and so re-obtains the NH3, and forms MgCl2: the former goes back for the manufacture of soda, while the latter is employed either for making HCl or Cl2.[17]Commercial soda ash (calcined, anhydrous) is rarely pure; the crystallised soda is generally purer. In order to purify it further, it is best to boil a concentrated solution of soda ash until two-thirds of the liquid remain, collect the soda which settles, wash with cold water, and then shake up with a strong solution of ammonia, pour off the residue, and heat. The impurities will then remain in the mother liquors, &c.Some numerical data may be given for sodium carbonate. The specific gravity of the anhydrous salt is 2·48, that of the decahydrated salt 1·46. Two varieties are known of the heptahydrated salt (Löwel, Marignac, Rammelsberg), which are formed together by allowing a saturated solution to cool under a layer of alcohol; the one is less stable (like the corresponding sulphate) and at 0° has a solubility of 32 parts (of anhydrous salt) in 100 water; the other is more stable, and its solubility 20 parts (of anhydrous salt) per 100 of water. The solubility of the decahydrated salt in 100 water = at 0°, 7·0; at 20°, 21·7; at 30°, 37·2 parts (of anhydrous salt). At 80° the solubility is only 46·1, at 90° 45·7, at 100°, 45·4 parts (of anhydrous salt). That is, it falls as the temperature rises, like Na2SO4. The specific gravity (Note7) of the solutions of sodium carbonate, according to the data of Gerlach and Kohlrausch, at 15°/4° is expressed by the formula,s= 9,992 + 104·5p+ 0·165p2. Weak solutions occupy a volume not only less than the sum of the volumes of the anhydrous salt and the water, but even less than the water contained in them. For instance, 1,000 grams of a 1 p.c. solution occupy (at 15°) a volume of 990·4 c.c. (sp. gr. 1·0097), but contain 990 grams of water, occupying at 15° a volume of 990·8 c.c. A similar case, which is comparatively rare occurs also with sodium hydroxide, in those dilute solutions for which the factorAis greater than 100 if the sp. gr. of water at 4° = 100,000, and if the sp. gr. of the solution be expressed by the formulaS=S0+Ap+Bp2, whereS0is the specific gravity of the water. For 5 p.c. the sp. gr. 15°/4° = 1·0520; for 10 p.c. 1·1057; for 15 p.c. 1·1603. The changes in the sp. gr. with the temperature are here almost the same as with solutions of sodium chloride with an equal value ofp.[18]The resemblance is so great that, notwithstanding the difference in the molecular composition of Na2SO4and Na2CO3, they ought to be classed under the type (NaO)2R, where R = SO2or CO. Many other sodium salts also contain 10 mol. H2O.[19]According to the observations of Pickering. According to Rose, when solutions of sodium carbonate are boiled a certain amount of carbonic anhydride is disengaged.[20]The composition of this salt, however, may be also represented as a combination of carbonic acid, H2CO3, with the normal salt, Na2CO3, just as the latter also combines with water. Such a combination is all the more likely because (1) there exists another salt, Na2CO3,2NaHCO3,2H2O (sodium sesquicarbonate), obtained by cooling a boiling solution of sodium bicarbonate, or by mixing this salt with the normal salt; but the formula of this salt cannot be derived from that of normal carbonic acid, as the formula of the bicarbonate can. At the same time the sesqui-salt has all the properties of a definite compound; it crystallises in transparent crystals, has a constant composition, its solubility (at 0° in 100 of water, 12·6 of anhydrous salt) differs from the solubility of the normal and acid salts; it is found in nature, and is known by the names oftronaandurao. The observations of Watts and Richards showed (1886) that on pouring a strong solution of the acid salt into a solution of the normal salt saturated by heating, crystals of the salt NaHCO3,Na2CO3,2H2O may be easily obtained, as long as the temperature is above 35°. The natural urao (Boussingault) has, according to Laurent, the same composition. This salt is very stable in air, and may be used for purifying sodium carbonate on the large scale. Such compounds have been little studied from a theoretical point of view, although particularly interesting, since in all probability they correspond with ortho-carbonic acid, C(OH)4, and at the same time correspond with double salts like astrakhanite (Chapter XIV., Note25). (2) Water of crystallisation does not enter into the composition of the crystals of the acid salt, so that on its formation (occurring only at low temperatures, as in the formation of crystalline compounds with water) the water of crystallisation of the normal salt separates and the water is, as it were, replaced by the elements of carbonic acid. If anhydrous sodium carbonate be mixed with the amount of water requisite for the formation of Na2CO3,H2O, this salt will, when powdered, absorb CO2as easily at the ordinary temperature as it does water.[21]100 parts of water at 0° dissolve 7 parts of the acid salt, which corresponds with 4·3 parts of the anhydrous normal salt, but at 0° 100 parts of water dissolve 7 parts of the latter. The solubility of the bi- or acid salt varies with considerable regularity; 100 parts of water dissolves at 15° 9 parts of the salt, at 30° 11 parts.The ammonium, and more especially the calcium, salt, is much more soluble in water. The ammonia process (seep.524) is founded upon this. Ammonium bicarbonate (acid carbonate) at 0° has a solubility of 12 parts in 100 water, at 30° of 27 parts. The solubility therefore increases very rapidly with the temperature. And its saturated solution is more stable than a solution of sodium bicarbonate. In fact, saturated solutions of these salts have a gaseous tension like that of a mixture of carbonic anhydride and water—namely, at 15° and at 50°, for the sodium salt 120 and 750 millimetres, for the ammonium salt 120 and 563 millimetres. These data are of great importance in understanding the phenomena connected with the ammonia process. They indicate that with an increased pressure the formation of the sodium salt ought to increase if there be an excess of ammonium salt.[22]Crystalline sodium carbonate (broken into lumps) also absorbs carbonic anhydride, but the water contained in the crystals is then disengaged: Na2CO3,10H2O + CO2= Na2CO3,H2CO3+ 9H2O, and dissolves part of the carbonate; therefore part of the sodium carbonate passes into solution together with all the impurities. When it is required to avoid the formation of this solution, a mixture of ignited and crystalline sodium carbonate is taken. Sodium bicarbonate is prepared chiefly for medicinal use, and is then often termedcarbonate of soda, also, for instance, in the so-called soda powders, for preparing certain artificial mineral waters, for the manufacture of digestive lozenges like those made at Essentuki, Vichy, &c.[23]In chemistry, sodium oxide is termed ‘soda,’ which word must be carefully distinguished from the word sodium, meaning the metal.[24]With a small quantity of water, the reaction either does not take place, or even proceeds in the reverse way—that is, sodium and potassium hydroxides remove carbonic anhydride from calcium carbonate (Liebig, Watson, Mitscherlich, and others). The influence of the mass of water is evident. According to Gerberts, however, strong solutions of sodium carbonate are decomposed by lime, which is very interesting if confirmed by further investigation.[25]As long as any undecomposed sodium carbonate remains in solution, excess of acid added to the solution disengages carbonic anhydride, and the solution after dilution gives a white precipitate with a barium salt soluble in acids, showing the presence of a carbonate in solution (if there be sulphate present, it also forms a white precipitate, but this is insoluble in acids). For the decomposition of sodium carbonate, milk of lime—that is, slaked slime suspended in water—is employed. Formerly pure sodium hydroxide was prepared (according to Berthollet's process) by dissolving the impure substance in alcohol (sodium carbonate and sulphate are not soluble), but now that metallic sodium has become cheap and is purified by distillation,pure caustic sodais prepared by acting on a small quantity of water with sodium. Perfectly pure sodium hydroxide may also be obtained by allowing strong solutions to crystallise (in the cold) (Note27).In alkali works where the Leblanc process is used, caustic soda is prepared directly from the alkali remaining in the mother liquors after the separation of the sodium carbonate by evaporation (Note14). If excess of lime and charcoal have been used, much sodium hydroxide maybe obtained. After the removal as much as possible of the sodium carbonate, a red liquid (from iron oxide) is left, containing sodium hydroxide mixed with compounds of sulphur and of cyanogen (seeChapterIX.) and also containing iron. This red alkali is evaporated and air is blown through it, which oxidises the impurities (for this purpose sometimes sodium nitrate is added, or bleaching powder, &c.) and leaves fused caustic soda. The fused mass is allowed to settle in order to separate the ferruginous precipitate, and poured into iron drums, where the sodium hydroxide solidifies. Such caustic soda contains about 10 p.c. of water in excess and some saline impurities, but when properly manufactured is almost free from carbonate and from iron. The greater part of the caustic soda, which forms so important an article of commerce, is manufactured in this manner.[26]Löwig gave a method of preparing sodium hydroxide from sodium carbonate by heating it to a dull red heat with an excess of ferric oxide. Carbonic anhydride is given off, and warm water extracts the caustic soda from the remaining mass. This reaction, as experiment shows, proceeds very easily, and is an example of contact action similar to that of ferric oxide on the decomposition of potassium chlorate. The reason of this may be that a small quantity of the sodium carbonate enters into double decomposition with the ferric oxide, and the ferric carbonate produced is decomposed into carbonic anhydride and ferric oxide, the action of which is renewed. Similar explanations expressing thereasonfor a reaction really adds but little to that elementary conception of contact which, according to my opinion, consists in the change of motion of the atoms in the molecules under the influence of the substance in contact. In order to represent this clearly it is sufficient, for instance, to imagine that in the sodium carbonate the elements CO2move in a circle round the elements Na2O, but at the points of contact with Fe2O3the motion becomes elliptic with a long axis, and at some distance from Na2O the elements of CO2are parted, not having the faculty of attaching themselves to Fe2O3.[27]By allowing strong solutions of sodium hydroxide to crystallise in the cold, impurities—such as, for instance, sodium sulphate—may be separated from them. The fused crystallo-hydrate 2NaHO,7H2O forms a solution having a specific gravity of 1·405 (Hermes). The crystals on dissolving in water produce cold, while NaHO produces heat. Besides which Pickering obtained hydrates with 1, 2, 4, 5, and 7 H2O.[28]In solid caustic soda there is generally an excess of water beyond that required by the formula NaHO. The caustic soda used in laboratories is generally cast in sticks, which are broken into pieces. It must be preserved in carefully closed vessels, because it absorbs water and carbonic anhydride from the air.[29]By the way it changes in air it is easy to distinguish caustic soda from caustic potash, which in general resembles it. Both alkalis absorb water and carbonic anhydride from the air, but caustic potash forms a deliquescent mass of potassium carbonate, whilst caustic soda forms a dry powder of efflorescent salt.[30]As the molecular weight of NaHO = 40, the volume of its molecule = 40/2·13 = 18·5, which very nearly approaches the volume of a molecule of water. The same rule applies to the compounds of sodium in general—for instance, its salts have a molecular volume approaching the volume of the acids from which they are derived.[31]The molecular quantity of sodium hydroxide (40 grams), on being dissolved in a large mass (200 gram molecules) of water, develops, according to Berthelot 9,780, and according to Thomsen 9,940, heat-units, but at 100° about 13,000 (Berthelot). Solutions of NaHO +nH2O, on being mixed with water, evolve heat if they contain less than 6H2O, but if more they absorb beat.[32]The specific gravity of solutions of sodium hydroxide at 15°/4° is given in the short table below:—NaHO, p.c.51015203040Sp. gr.1·0571·1131·1691·2241·3311·4361,000 grams of a 5 p.c. solution occupies a volume of 946 c.c.; that is, less than the water serving to make the solution (seeNote18).[33]Sodium hydroxide and some other alkalis are capable of hydrolysing—saponifying, as it is termed—the compounds of acids with alcohols. If RHO (or R(HO)n) represent the composition of an alcohol—that is, of the hydroxide of a hydrocarbon radicle—and QHO an acid, then the compound of the acid with the alcohol or ethereal salt of the given acid will have the composition RQO. Ethereal salts, therefore, present a likeness to metallic salts, just as alcohols resemble basic hydroxides. Sodium hydroxide acts on ethereal salts in the same way that it acts on the majority of metallic salts—namely, it liberates alcohol, and forms the sodium salt of that acid which was in the ethereal salt. The reaction takes place in the following way:—RQO+NaHO=NaQO+RHOEthereal saltCaustic sodaSodium saltAlcoholSuch a decomposition is termed saponification; similar reactions were known very long ago for the ethereal salts corresponding with glycerin, C3H5(OH)3(ChapterIX.), found in animals and plants, and composing what are called fats or oils. Caustic soda, acting on fat and oil, forms glycerin, and sodium salts of those acids which were in union with the glycerin in the fat, as Chevreul showed at the beginning of this century. The sodium salts of the fatty acids are commonly known as soaps. That is to say, soap is made from fat and caustic soda, glycerin being separated and a sodium salt or soap formed. As glycerin is usually found in union with certain acids, so also are the sodium salts of the same acids found in soap. The greater part of the acids found in conjunction with glycerin in fats are the solid palmitic and stearic acids, C16H32O2and C18H38O2, and the liquid oleic acid, C18H34O2. In preparing soap the fatty substances are mixed with a solution of caustic soda until an emulsion is formed; the proper quantity of caustic soda is then added in order to produce saponification on heating, the soap being separated from the solution either by means of an excess of caustic soda or else by common salt, which displaces the soap from the aqueous solution (salt water does not dissolve soap, neither does it form a lather). Water acting on soap partly decomposes it (because the acids of the soap are feeble), and the alkali set free acts during the application of soap. Hence it may be replaced by a very feeble alkali. Strong solutions of alkali corrode the skin and tissues. They are not formed from soap, because the reaction is reversible, and the alkali is only set free by the excess of water. Thus we see how the teaching of Berthollet renders it possible to understand many phenomena which occur in every-day experience (seeChapter IX., Note15).[34]On this is founded the process of Henkoff and Engelhardt for treating bones. The bones are mixed with ashes, lime, and water; it is true that in this case more potassium hydroxide than sodium hydroxide is formed, but their action is almost identical.[35]As explained in Note33.[35 bis]It might be expected, from what has been mentioned above, that bivalent metals would easily form acid salts with acids containing more than two atoms of hydrogen—for instance, with tribasic acids, such as phosphoric acid, H3PO4—and actually such salts do exist; but all such relations are complicated by the fact that the character of the base very often changes and becomes weakened with the increase of valency and the change of atomic weight; the feebler bases (like silver oxide), although corresponding with univalent metals, do not form acid salts, while the feeblest bases (CuO, PbO, &c.) easily form basic salts, and notwithstanding their valency do not form acid salts which are in any degree stable—that is, which are undecomposable by water. Basic and acid salts ought to be regarded rather as compounds similar to crystallo-hydrates, because such acids as sulphuric form with sodium not only an acid and a normal salt, as might be expected from the valency of sodium, but also salts containing a greater quantity of acid. In sodium sesquicarbonate we saw an example of such compounds. Taking all this into consideration, we must say that the property of more or less easily forming acid salts depends more upon the energy of the base than upon its valency, and the best statement is thatthe capacity of a base for forming acid and basic salts is characteristic, just as the faculty of forming compounds with hydrogen is characteristic of elements.[36]Deville supposes that such a decomposition of sodium hydroxide by metallic iron depends solely on the dissociation of the alkali at a white heat into sodium, hydrogen, and oxygen. Here the part played by the iron is only that it retains the oxygen formed, otherwise the decomposed elements would again reunite upon cooling, as in other cases of dissociation. If it be supposed that the temperature at the commencement of the dissociation of the iron oxides is higher than that of sodium oxide, then the decomposition may be explained by Deville's hypothesis. Deville demonstrates his views by the following experiment:—An iron bottle, filled with iron borings, was heated in such a way that the upper part became red hot, the lower part remaining cooler; sodium hydroxide was introduced into the upper part. The decomposition was then effected—that is, sodium vapours were produced (this experiment was really performed with potassium hydroxide). On opening the bottle it was found that the iron in the upper part was not oxidised, but only that in the lower part. This may be explained by the decomposition of the alkali into sodium, hydrogen, and oxygen taking place in the upper part, whilst the iron in the lower part absorbed the oxygen set free. If the whole bottle be subjected to the same moderate heat as the lower extremity, no metallic vapours are formed. In that case, according to the hypothesis, the temperature is insufficient for the dissociation of the sodium hydroxide.[37]It has been previously remarked (Chapter II. Note9) that Beketoff showed the displacement of sodium by hydrogen, not from sodium hydroxide but from the oxide Na2O; then, however, only one half is displaced, with the formation of NaHO.[38]Since the close of the eighties in England, where the preparation of sodium is at present carried out on a large commercial scale (from 1860 to 1870 it was only manufactured in a few works in France), it has been the practice to add to Deville's mixture iron, or iron oxide which with the charcoal gives metallic and carburetted iron, which still further facilitates the decomposition. At present a kilogram of sodium may be purchased for about the same sum (2/-) as a gram cost thirty years ago. Castner, in England, greatly improved the manufacture of sodium in large quantities, and so cheapened it as a reducing agent in the preparation of metallic aluminium. He heated a mixture of 44 parts of NaHO, and 7 parts of carbide of iron in large iron retorts at 1,000° and obtained about 6½ parts of metallic sodium. The reaction proceeds more easily than with carbon or iron alone, and the decomposition of the NaHO proceeds according to the equation: 3NaHO + C = Na2CO3+ 3H + Na. Subsequently, in 1891, aluminium was prepared by electrolysis (seeChapter XVII.), and metallic sodium found two new uses; (1) for the manufacture of peroxide of sodium (see later on) which is used in bleaching works, and (2) in the manufacture of potassium and sodium cyanide from yellow prussiate (Chapter XIII., Note12).[38 bis]This is also shown by the fall in the temperature of solidification of tin produced by the addition of sodium (and also Al and Zn). Heycock and Neville (1889).[39]By dissolving sodium amalgams in water and acids, and deducting the heat of solution of the sodium, Berthelot found thatfor each atom of the sodiumin amalgams containing a larger amount of mercury than NaHg5, the amount of heat evolved increases, after which the heat of formation falls, and the heat evolved decreases. In the formation of NaHg5about 18,500 calories are evolved; when NaHg3is formed, about 14,000; and for NaHg about 10,000 calories. Kraft regarded the definite crystalline amalgam as having the composition of NaHg6, but at the present time, in accordance with Grimaldi's results, it is thought to be NaHg5. A similar amalgam is very easily obtained if a 3 p.c. amalgam be left several days in a solution of sodium hydroxide until a crystalline mass is formed, from which the mercury may be removed by strongly pressing in chamois leather. This amalgam with a solution of potassium hydroxide forms a potassium amalgam, KHg10. It may be mentioned here that the latent heat of fusion (of atomic quantities) of Hg = 360 (Personne), Na = 730 (Joannis), and K = 610 calories (Joannis).[40]Alloys are so similar to solutions (exhibiting such complete parallelism in properties) that they are included in the same class of so-called indefinite compounds. But in alloys, as substances passing from the liquid to the solid state, it is easier to discover the formation of definite chemical compounds. Besides the alloys of Na with Hg, those with tin (Bailey 1892 found Na2Sn), lead (NaPb), bismuth (Na3Bi), &c. (Joannis 1892 and others) have been investigated.[41]Potassium forms a similar compound, but lithium, under the same circumstances, does not.[42]The tension of dissociation of hydrogenp, in millimetres of mercury, is:—t=330°350°400°430°for Na2Hp=2857447910for K2H45725481100[43]In general, during the formation of alloys the volumes change very slightly, and therefore from the volume of Na2H some idea may be formed of the volume of hydrogen in a solid or liquid state. Even Archimedes concluded that there was gold in an alloy of copper and gold by reason of its volume and density. From the fact that the density of Na2H is equal to 0·959, it may be seen that the volume of 47 grams (the gram molecule) of this compound = 49·0 c.c. The volume of 46 grams of sodium contained in the Na2H (the density under the same conditions being 0·97) is equal to 47·4 c.c. Therefore the volume of 1 gram of hydrogen in Na2H is equal to 1·6 c.c., and consequently the density of metallic hydrogen, or the weight of 1 c.c., approaches 0·6 gram. This density is also proper to the hydrogen alloyed with potassium and palladium. Judging from the scanty information which is at present available, liquid hydrogen near its absolute boiling point (ChapterII.) has a much lower density.[43 bis]We may remark that at low temperatures Na absorbs NH3and forms (NH3Na)2(seeChapter VI., Note14); this substance absorbs CO and gives (NaCO)n (Chapter IX., Note31), although by itself Na does not combine directly with CO (but K does).[44]H. A. Schmidt remarked that perfectly dry hydrogen chloride is decomposed with great difficulty by sodium, although the decomposition proceeds easily with potassium and with sodium in moist hydrogen chloride. Wanklyn also remarked that sodium burns with great difficulty in dry chlorine. Probably these facts are related to other phenomena observed by Dixon, who found that perfectly dry carbonic oxide does not explode with oxygen on passing an electric spark.[44 bis]Sodamide, NH2Na, (Chapter IV., Note14), discovered by Gay-Lussac and Thénard, has formed the object of repeated research, but has been most fully investigated by A. W. Titherley (1894). Until recently the following was all that was known about this compound:—By heating sodium in dry ammonia, Gay-Lussac and Thénard obtained an olive-green, easily-fusible mass,sodamide, NH2Na, hydrogen being separated. This substance with water forms sodium hydroxide and ammonia; with carbonic oxide, CO, it forms sodium cyanide, NaCN, and water, H2O; and with dry hydrogen chloride it forms sodium and ammonium chlorides. These and other reactions of sodamide show that the metal in it preserves its energetic properties in reaction, and that this compound of sodium is more stable than the corresponding chlorine amide. When heated, sodamide, NH2Na, only partially decomposes, with evolution of hydrogen, the principal part of it giving ammonia and sodium nitride, Na3N, according to the equation 3NH2Na = 2NH3+ NNa3. The latter is an almost black powdery mass, decomposed by water into ammonia and sodium hydroxide.Titherley's researches added the following data:—Iron or silver vessels should be used in preparing this body, because glass and porcelain are corroded at 300°–400°, at which temperature ammonia gas acts upon sodium and forms the amide with the evolution of hydrogen. The reaction proceeds slowly, but is complete if there be an excess of NH3. Pure NH2Na is colourless (its colouration is due to various impurities), semi-transparent, shows traces of crystallisation, has a conchoidal fracture, and melts at 145°. Judging from the increase in weight of the sodium and the quantity of hydrogen which is disengaged, the composition of the amide is exactly NH2Na. It partially volatilises (sublimes) in vacuo at 200°, and breaks up into 2Na + N2+ 2H2at 500°. The same amide is formed when oxide of sodium is heated in NH3: Na2O + 2NH3= 2NaH2N + H2O. NaHO is also formed to some extent by the resultant H2O. Potassium and lithium form similar amides. With water, alcohol, and acids, NH2Na gives NH3and NaHO, which react further. Anhydrous CaO absorbs NH2Na when heated without decomposing it. When sodamide is heated with SiO2, NH3is disengaged, and silicon nitride formed. It acts still more readily upon boric anhydride when heated with it: 2NH2Na + B2O3= 2BN + 2NaHO + H2O. When slightly heated, NH2Na + NOCl = NaCl + N2+ H2O (NHNa2and NNa3are apparently not formed at a higher temperature). The halogen organic compounds react with the aid of heat, but with so much energy that the reaction frequently leads to the ultimate destruction of the organic groups and production of carbon.[45]As sodium does not displace hydrogen from the hydrocarbons,it may be preservedin liquid hydrocarbons. Naphtha is generally used for this purpose, as it consists of a mixture of various liquid hydrocarbons. However, in naphtha sodium usually becomes coated with a crust composed of matter produced by the action of the sodium on certain of the substances contained in the mixture composing naphtha. In order that sodium may retain its lustre in naphtha, secondary octyl alcohol is added. (This alcohol is obtained by distilling castor oil with caustic potash.) Sodium keeps well in a mixture of pure benzene and paraffin.[46]If sodium does not directly displace the hydrogen in hydrocarbons, still by indirect means compounds may be obtained which contain sodium and hydrocarbon groups. Some of these compounds have been produced, although not in a pure state. Thus, for instance, zinc ethyl, Zn(C2H5)2, when treated with sodium, loses zinc and forms sodium ethyl, C2H5Na, but this decomposition is not complete, and the compound formed cannot be separated by distillation from the remaining zinc ethyl. In this compound the energy of the sodium is clearly manifest, for it reacts with substances containing haloids, oxygen, &c., and directly absorbs carbonic anhydride, forming a salt of a carboxylic acid (propionic).[46 bis]It is even doubtful whether the suboxide exists (seeNote47).[47]A compound, Na2Cl, which corresponds with the suboxide, is apparently formed when a galvanic current is passed through fused common salt; the sodium liberated dissolves in the common salt, and does not separate from the compound either on cooling or on treatment with mercury. It is therefore supposed to be Na2Cl; the more so as the mass obtained gives hydrogen when treated with water: Na2Cl + H2O = H + NaHO + NaCl, that is, it acts like suboxide of sodium. If Na2Cl really exists as a salt, then the corresponding base Na4O, according to the rule with other bases of the composition M4O, ought to be called a quaternary oxide. According to certain evidence, a suboxide is formed when thin sheets or fine drops of sodium slowly oxidise in moist air.[48]According to observations easily made, sodium when fused in air oxidises but does not burn, the combustion only commencing with the formation of vapour—that is, when considerably heated. Davy and Karsten obtained the oxides of potassium, K2O, and of sodium, Na2O, by heating the metals with their hydroxides, whence NaHO + Na = Na2O + H, but N. N. Beketoff failed to obtain oxides by this means. He prepared them by directly igniting the metals in dry air, and afterwards heating with the metal in order to destroy any peroxide. The oxide produced, Na2O, when heated in an atmosphere of hydrogen, gave a mixture of sodium and its hydroxide: Na2O + H = NaHO + Na (seeChapter II., Note9). If both the observations mentioned are accurate, then the reaction is reversible. Sodium oxide ought to be formed during the decomposition of sodium carbonate by oxide of iron (seeNote26), and during the decomposition of sodium nitrite. According to Karsten, its specific gravity is 2·8, according to Beketoff 2·3. The difficulty in obtaining it is owing to an excess of sodium forming the suboxide, and an excess of oxygen the peroxide. The grey colour peculiar to the suboxide and oxide perhaps shows that they contain metallic sodium. In addition to this, in the presence of water it may contain sodium hydride and NaHO.[49]Of the oxides of sodium, that easiest to form is the peroxide, NaO or Na2O2; this is obtained when sodium is burnt in an excess of oxygen. If NaNO3be melted, it gives Na2O2with metallic Na. In a fused state the peroxide is reddish yellow, but it becomes almost colourless when cold. When heated with iodine vapour, it loses oxygen: Na2O2+I2= Na2OI2+ O. The compound Na2OI2is akin to the compound Cu2OCl2obtained by oxidising CuCl. This reaction is one of the few in which iodine directly displaces oxygen. The substance Na2OI2is soluble in water, and when acidified gives free iodine and a sodium salt. Carbonic oxide is absorbed by heated sodium peroxide with formation of sodium carbonate: Na2CO3= Na2O2+ CO, whilst carbonic anhydride liberates oxygen from it. With nitrous oxide it reacts thus: Na2O2+2N2O = 2NaNO2+N2; with nitric oxide it combines directly, forming sodium nitrite, NaO + NO = NaNO2. Sodium peroxide, when treated with water, does not give hydrogen peroxide, because the latter in the presence of the alkali formed (Na2O2+ 2H2O = 2NaHO + H2O2) decomposes into water and oxygen. In the presence of dilute sulphuric acid it forms H2O2(Na2O2+ H2SO4= Na2SO4+ H2O2). Peroxide of sodium is now prepared on a large scale (by the action of air upon Na at 300°) for bleaching wool, silk &c. (when it acts in virtue of the H2O2formed). The oxidising properties of Na2O2under the action of heat are seen, for instance, in the fact that when heated with I it forms sodium iodate; with PbO, Na2PbO3; with pyrites, sulphates, &c. When peroxide of sodium comes into contact with water, it evolves much heat, forming H2O2, and decomposing with the disengagement of oxygen; but, as a rule, there is no explosion. But if Na2O2be placed in contact with organic matter, such as sawdust, cotton, &c., it gives a violent explosion when heated, ignited, or acted on by water. Peroxide of sodium forms an excellent oxidising agent for the preparation of the higher product of oxidation of Mn, Cr, W, &c., and also for oxidising the metallic sulphides. It should therefore find many applications in chemical analysis. To prepare Na2O2on a large scale, Castner melts Na in an aluminium vessel, and at 300° passes first air deprived of a portion of its oxygen (having been already once used), and then ordinary dry air over it.
Footnotes:
[1]Whilst describing in some detail the properties of sodium chloride, hydrochloric acid, and sodium sulphate, I wish to impart, by separate examples, an idea of the properties of saline substances, but the dimensions of this work and its purpose and aim do not permit of entering into particulars concerning every salt, acid, or other substance. The fundamental object of this work—an account of the characteristics of the elements and an acquaintance with the forces acting between atoms—has nothing to gain from the multiplication of the number of as yet ungeneralised properties and relations.
[1]Whilst describing in some detail the properties of sodium chloride, hydrochloric acid, and sodium sulphate, I wish to impart, by separate examples, an idea of the properties of saline substances, but the dimensions of this work and its purpose and aim do not permit of entering into particulars concerning every salt, acid, or other substance. The fundamental object of this work—an account of the characteristics of the elements and an acquaintance with the forces acting between atoms—has nothing to gain from the multiplication of the number of as yet ungeneralised properties and relations.
[2]Anhydrous (ignited) sodium sulphate, Na2SO4, is known in trade as ‘sulphate’ or salt-cake, in mineralogythenardite. Crystalline decahydrated salt is termed in mineralogymirabilite, and in trade Glauber's salt. On fusing it, the monohydrate Na2SO4H2O is obtained, together with a supersaturated solution.
[2]Anhydrous (ignited) sodium sulphate, Na2SO4, is known in trade as ‘sulphate’ or salt-cake, in mineralogythenardite. Crystalline decahydrated salt is termed in mineralogymirabilite, and in trade Glauber's salt. On fusing it, the monohydrate Na2SO4H2O is obtained, together with a supersaturated solution.
[3]The salts may be obtained not only by methods of substitution of various kinds, but also by many other combinations. Thus sodium sulphate may be formed from sodium oxide and sulphuric anhydride, by oxidising sodium sulphide, Na2S, or sodium sulphite, Na2SO3, &c. When sodium chloride is heated in a mixture of the vapours of water, air, and sulphurous anhydride, sodium sulphate is formed. According to this method (patented by Hargreaves and Robinson), sodium sulphate, Na2SO4, is obtained from NaCl without the preliminary manufacture of H2SO4. Lumps of NaCl pressed into bricks are loosely packed into a cylinder and subjected, at a red heat, to the action of steam, air and SO2. Under these conditions, HCl, sulphate, and a certain amount of unaltered NaCl are obtained. This mixture is converted into soda by Gossage's process (seeNote15) and may have some practical value.
[3]The salts may be obtained not only by methods of substitution of various kinds, but also by many other combinations. Thus sodium sulphate may be formed from sodium oxide and sulphuric anhydride, by oxidising sodium sulphide, Na2S, or sodium sulphite, Na2SO3, &c. When sodium chloride is heated in a mixture of the vapours of water, air, and sulphurous anhydride, sodium sulphate is formed. According to this method (patented by Hargreaves and Robinson), sodium sulphate, Na2SO4, is obtained from NaCl without the preliminary manufacture of H2SO4. Lumps of NaCl pressed into bricks are loosely packed into a cylinder and subjected, at a red heat, to the action of steam, air and SO2. Under these conditions, HCl, sulphate, and a certain amount of unaltered NaCl are obtained. This mixture is converted into soda by Gossage's process (seeNote15) and may have some practical value.
[4]Many observations have been made, but little general information has been obtained from particular cases. In addition to which, the properties of a given salt are changed by the presence of other salts. This takes place not only in virtue of mutual decomposition or formation of double salts capable of separate existence, but is determined by the influence which some salts exert on others, or by forces similar to those which act during solution. Here nothing has been generalised to that extent which would render it possible to predict without previous investigation, if there be no close analogy to help us. Let us state one of these numerous cases: 100 parts of water at 20° dissolve 34 parts of potassium nitrate but on the addition of sodium nitrate the solubility of potassium nitrate increases to 48 parts in 10 of water (Carnelley and Thomson). In general, in all cases of which there are accurate observations it appears that the presence of foreign salts changes the properties of any given salt.
[4]Many observations have been made, but little general information has been obtained from particular cases. In addition to which, the properties of a given salt are changed by the presence of other salts. This takes place not only in virtue of mutual decomposition or formation of double salts capable of separate existence, but is determined by the influence which some salts exert on others, or by forces similar to those which act during solution. Here nothing has been generalised to that extent which would render it possible to predict without previous investigation, if there be no close analogy to help us. Let us state one of these numerous cases: 100 parts of water at 20° dissolve 34 parts of potassium nitrate but on the addition of sodium nitrate the solubility of potassium nitrate increases to 48 parts in 10 of water (Carnelley and Thomson). In general, in all cases of which there are accurate observations it appears that the presence of foreign salts changes the properties of any given salt.
[5]The information concerning solubility (ChapterI.) is given according to the determinations of Gay-Lussac, Lovell, and Mulder.
[5]The information concerning solubility (ChapterI.) is given according to the determinations of Gay-Lussac, Lovell, and Mulder.
[6]In Chapter I., Note24, we have already seen that with many other sulphates the solubility also decreases after a certain temperature is passed. Gypsum, CaSO4,2H2O, lime, and many other compounds present such a phenomenon. An observation of Tilden's (1884) is most instructive; he showed that on raising the temperature (in closed vessels) above 140° the solubility of sodium sulphate again begins to increase. At 100° 100 parts of water dissolve about 43 parts of anhydrous salt, at 140° 42 parts, at 160° 43 parts, at 180° 44 parts, at 230° 46 parts. According to Étard (1892) the solubility of 30 parts of Na2SO4in 100 of solution (or 43 per 100 of water) corresponds to 80°, and above 240° the solubility again falls, and very rapidly, so that at 320° the solution contains 12 per 100 of solution (about 14 per 100 of water) and a further rise of temperature is followed by a further deposition of the salt. It is evident that the phenomenon of saturation, determined by the presence of an excess of the dissolved substance, is very complex, and therefore that for the theory of solutions considered as liquid indefinite chemical compounds, many useful statements can hardly be given.
[6]In Chapter I., Note24, we have already seen that with many other sulphates the solubility also decreases after a certain temperature is passed. Gypsum, CaSO4,2H2O, lime, and many other compounds present such a phenomenon. An observation of Tilden's (1884) is most instructive; he showed that on raising the temperature (in closed vessels) above 140° the solubility of sodium sulphate again begins to increase. At 100° 100 parts of water dissolve about 43 parts of anhydrous salt, at 140° 42 parts, at 160° 43 parts, at 180° 44 parts, at 230° 46 parts. According to Étard (1892) the solubility of 30 parts of Na2SO4in 100 of solution (or 43 per 100 of water) corresponds to 80°, and above 240° the solubility again falls, and very rapidly, so that at 320° the solution contains 12 per 100 of solution (about 14 per 100 of water) and a further rise of temperature is followed by a further deposition of the salt. It is evident that the phenomenon of saturation, determined by the presence of an excess of the dissolved substance, is very complex, and therefore that for the theory of solutions considered as liquid indefinite chemical compounds, many useful statements can hardly be given.
[7]Already referred to in Chapter I., Note56.The example of sodium sulphate is historically very important for the theory of solutions. Notwithstanding the number of investigations which have been made, it is still insufficiently studied, especially from the point of the vapour tension of solutions and crystallo-hydrates, so that those processes cannot be applied to it which Guldberg, Roozeboom, Van't Hoff, and others applied to solutions and crystallo-hydrates. It would also be most important to investigate the influence of pressure on the various phenomena corresponding with the combinations of water and sodium sulphate, because when crystals are separated—for instance, of the decahydrated salt—an increase of volume takes place, as can be seen from the following data:—the sp. gr. of the anhydrous salt is 2·66, that of the decahydrated salt = 1·46, but the sp. gr. of solutions at 15°/4° = 9,992 + 90·2p+ 0·35p2where p represents the percentage of anhydrous salt in the solution, and the sp. gr. of water at 4° = 10,000. Hence for solutions containing 20 p.c. of anhydrous salt the sp. gr. = 1·1936; therefore the volume of 100 grams of this solution = 83·8 c.c., and the volume of anhydrous salt contained in it is equal to 20/2·66, or = 7·5 c.c., and the volume of water = 80·1 c.c. Therefore, the solution, on decomposing into anhydrous salt and water, increases in volume (from 83·8 to 87·6); but in the same way 83·8 c.c. of 20 p.c. solution are formed from (45·4/1·46 =) 31·1 c.c. of the decahydrated salt, and 54·6 c.c. of water—that is to say, that during the formation of a solution from 85·7 c.c., 83·8 c.c. are formed.
[7]Already referred to in Chapter I., Note56.
The example of sodium sulphate is historically very important for the theory of solutions. Notwithstanding the number of investigations which have been made, it is still insufficiently studied, especially from the point of the vapour tension of solutions and crystallo-hydrates, so that those processes cannot be applied to it which Guldberg, Roozeboom, Van't Hoff, and others applied to solutions and crystallo-hydrates. It would also be most important to investigate the influence of pressure on the various phenomena corresponding with the combinations of water and sodium sulphate, because when crystals are separated—for instance, of the decahydrated salt—an increase of volume takes place, as can be seen from the following data:—the sp. gr. of the anhydrous salt is 2·66, that of the decahydrated salt = 1·46, but the sp. gr. of solutions at 15°/4° = 9,992 + 90·2p+ 0·35p2where p represents the percentage of anhydrous salt in the solution, and the sp. gr. of water at 4° = 10,000. Hence for solutions containing 20 p.c. of anhydrous salt the sp. gr. = 1·1936; therefore the volume of 100 grams of this solution = 83·8 c.c., and the volume of anhydrous salt contained in it is equal to 20/2·66, or = 7·5 c.c., and the volume of water = 80·1 c.c. Therefore, the solution, on decomposing into anhydrous salt and water, increases in volume (from 83·8 to 87·6); but in the same way 83·8 c.c. of 20 p.c. solution are formed from (45·4/1·46 =) 31·1 c.c. of the decahydrated salt, and 54·6 c.c. of water—that is to say, that during the formation of a solution from 85·7 c.c., 83·8 c.c. are formed.
[8]From this example it is evident the solution remains unaltered until from the contact of a solid it becomes either saturated or supersaturated, crystallisation being determined by the attraction to a solid, as the phenomenon of supersaturation clearly demonstrates. This partially explains certain apparently contradictory determinations of solubility. The best investigated example of such complex relations is cited in Chapter XIV., Note50(for CaCl2).
[8]From this example it is evident the solution remains unaltered until from the contact of a solid it becomes either saturated or supersaturated, crystallisation being determined by the attraction to a solid, as the phenomenon of supersaturation clearly demonstrates. This partially explains certain apparently contradictory determinations of solubility. The best investigated example of such complex relations is cited in Chapter XIV., Note50(for CaCl2).
[9]According to Pickering's experiments (1886), the molecular weight in grams (that is, 142 grams) of anhydrous sodium sulphate, on being dissolved in a large mass of water, at 0° absorbs (hence the - sign) -1,100 heat units, at 10°–700, at 15°–275, at 20° gives out +25, at 25° +300 calories. For the decahydrated salt, Na2SO4,10H2O, 5° - 4,225, 10° - 4,000, 15° - 3,570, 20° - 3,160, 25° - 2,775. Hence (just as in Chapter I., Note56) the heat of the combination Na2SO4,10H2O at 5° = +3,125, 10° = +3,250, 20° = +3,200, and 25° = +3,050.It is evident that the decahydrated salt dissolving in water gives a decrease of temperature. Solutions in hydrochloric acid give a still greater decrease, because they contain the water of crystallisation in a solid state—that is, like ice—and this on melting absorbs heat. A mixture of 15 parts of Na2SO4,10H2O and 12 parts of strong hydrochloric acid produces sufficient cold to freeze water. During the treatment with hydrochloric acid a certain quantity of sodium chloride is formed.
[9]According to Pickering's experiments (1886), the molecular weight in grams (that is, 142 grams) of anhydrous sodium sulphate, on being dissolved in a large mass of water, at 0° absorbs (hence the - sign) -1,100 heat units, at 10°–700, at 15°–275, at 20° gives out +25, at 25° +300 calories. For the decahydrated salt, Na2SO4,10H2O, 5° - 4,225, 10° - 4,000, 15° - 3,570, 20° - 3,160, 25° - 2,775. Hence (just as in Chapter I., Note56) the heat of the combination Na2SO4,10H2O at 5° = +3,125, 10° = +3,250, 20° = +3,200, and 25° = +3,050.
It is evident that the decahydrated salt dissolving in water gives a decrease of temperature. Solutions in hydrochloric acid give a still greater decrease, because they contain the water of crystallisation in a solid state—that is, like ice—and this on melting absorbs heat. A mixture of 15 parts of Na2SO4,10H2O and 12 parts of strong hydrochloric acid produces sufficient cold to freeze water. During the treatment with hydrochloric acid a certain quantity of sodium chloride is formed.
[10]The very large and well-formed crystals of this salt resemble the hydrate H2SO4,H2O, or SO(OH)4. In general the replacement of hydrogen by sodium modifies many of the properties of acids less than its replacement by other metals. This most probably depends on the volumes being nearly equal.
[10]The very large and well-formed crystals of this salt resemble the hydrate H2SO4,H2O, or SO(OH)4. In general the replacement of hydrogen by sodium modifies many of the properties of acids less than its replacement by other metals. This most probably depends on the volumes being nearly equal.
[11]In solution (Berthelot) the acid salt in all probability decomposes most in the greatest mass of water. The specific gravity (according to the determinations of Marignac) of solutions at 15°/4° = 9,992 + 77·92p+ 0·231p2(seeNote7). From these figures, and from the specific gravities of sulphuric acid, it is evident that on mixing solutions of this acid and sodium sulphateexpansionwill always take place; for instance, H2SO4+ 25H2O with Na2SO4+ 25H2O increases from 483 volumes to 486. In addition to which, in weak solutions heat is absorbed, as shown in Chapter X., Note27. Nevertheless, even more acid salts may be formed and obtained in a crystalline form. For instance, on cooling a solution of 1 part of sodium sulphate in 7 parts of sulphuric acid, crystals of the composition NaHSO4,H2SO4are separated (Schultz, 1868). This compound fuses at about 100°; the ordinary acid salt, NaHSO4, at 149°.
[11]In solution (Berthelot) the acid salt in all probability decomposes most in the greatest mass of water. The specific gravity (according to the determinations of Marignac) of solutions at 15°/4° = 9,992 + 77·92p+ 0·231p2(seeNote7). From these figures, and from the specific gravities of sulphuric acid, it is evident that on mixing solutions of this acid and sodium sulphateexpansionwill always take place; for instance, H2SO4+ 25H2O with Na2SO4+ 25H2O increases from 483 volumes to 486. In addition to which, in weak solutions heat is absorbed, as shown in Chapter X., Note27. Nevertheless, even more acid salts may be formed and obtained in a crystalline form. For instance, on cooling a solution of 1 part of sodium sulphate in 7 parts of sulphuric acid, crystals of the composition NaHSO4,H2SO4are separated (Schultz, 1868). This compound fuses at about 100°; the ordinary acid salt, NaHSO4, at 149°.
[11 bis]On decreasing the pressure, sodium hydrogen sulphate, NaHSO4, dissociates much more easily than at the ordinary pressure; it loses water and forms the pyrosulphate, Na2S2O7; this reaction is utilised in chemical works.
[11 bis]On decreasing the pressure, sodium hydrogen sulphate, NaHSO4, dissociates much more easily than at the ordinary pressure; it loses water and forms the pyrosulphate, Na2S2O7; this reaction is utilised in chemical works.
[12]Calcium sulphide, CaS, like many metallic sulphides which are soluble in water, is decomposed by it (ChapterX.), CaS + H2O = CaO + H2S, because hydrogen sulphide is a very feeble acid. If calcium sulphide be acted on by a large mass of water, lime may be precipitated, and a state of equilibrium will be reached, when the system CaO + 2CaS remains unchanged. Lime, being a product of the action of water on CaS, limits this action. Therefore, if in black ash the lime were not in excess, a part of the sulphide would be in solution (actually there is but very little). In this manner in the manufacture of sodium carbonate the conditions of equilibrium which enter into double decompositions have been made use of (see above), and the aim is to form directly the unchangeable product CaO,2CaS. This was first regarded as a special insoluble compound, but there is no evidence of its independent existence.
[12]Calcium sulphide, CaS, like many metallic sulphides which are soluble in water, is decomposed by it (ChapterX.), CaS + H2O = CaO + H2S, because hydrogen sulphide is a very feeble acid. If calcium sulphide be acted on by a large mass of water, lime may be precipitated, and a state of equilibrium will be reached, when the system CaO + 2CaS remains unchanged. Lime, being a product of the action of water on CaS, limits this action. Therefore, if in black ash the lime were not in excess, a part of the sulphide would be in solution (actually there is but very little). In this manner in the manufacture of sodium carbonate the conditions of equilibrium which enter into double decompositions have been made use of (see above), and the aim is to form directly the unchangeable product CaO,2CaS. This was first regarded as a special insoluble compound, but there is no evidence of its independent existence.
[13]see captionFig.69.—Apparatus for the methodical lixiviation of black ash, &c. Water flows into the tanks from the pipesr,r, and the saturated liquid is drawn off fromc,c.Methodical lixiviationis the extraction, by means of water, of a soluble substance from the mass containing it. It is carried on so as not to obtain weak aqueous solutions, and in such a way that the residue shall not contain any of the soluble substance. This problem is practically of great importance in many industries. It is required to extract from the mass all that is soluble in water. This is easily effected if water be first poured on the mass, the strong solution thus obtained decanted, then water again poured on, time being allowed for it to act, then again decanted, and so on until fresh water does not take up anything. But then finally such weak solutions are obtained that it would be very disadvantageous to evaporate them. This is avoided by pouring the fresh hot water destined for the lixiviation, not onto the fresh mass, but upon a mass which has already been subjected to a first lixiviation by weak solutions. In this way the fresh water gives a weak solution. The strong solution which goes to the evaporating pan flows from those parts of the apparatus which contain the fresh, as yet unlixiviated, mass, and thus in the latter parts the weak alkali formed in the other parts of the apparatus becomes saturated as far as possible with the soluble substance. Generally several intercommunicating vessels are constructed (standing at the same level) into which in turn the fresh mass is charged which is intended for lixiviation; the water is poured in, the alkali drawn off, and the lixiviated residue removed. The illustration represents such an apparatus, consisting of four communicating vessels. The water poured into one of them flows through the two nearest and issues from the third. The fresh mass being placed in one of these boxes or vessels, the stream of water passing through the apparatus is directed in such a manner as to finally issue from this vessel containing the fresh unlixiviated mass. The fresh water is added to the vessel containing the material which has been almost completely exhausted. Passing through this vessel it is conveyed by the pipe (syphon passing from the bottom of the first box to the top of the second) communicating with the second; it finally passes (also through a syphon pipe) into the box (the third) containing the fresh material. The water will extract all that is soluble in the first vessel, leaving only an insoluble residue. This vessel is then ready to be emptied, and refilled with fresh material. The levels of the liquids in the various vessels will naturally be different, in consequence of the various strengths of the solutions which they contain.It must not, however, be thought that sodium carbonate alone passes into the solution; there is also a good deal of caustic soda with it, formed by the action of lime on the carbonate of sodium, and there are also certain sodium sulphur compounds with which we shall partly become acquainted hereafter. The sodium carbonate, therefore, is not obtained in a very pure state. The solution is concentrated by evaporation. This is conducted by means of the waste heat from the soda furnaces, together with that of the gases given off. The process in the soda furnaces can only be carried on at a high temperature, and therefore the smoke and gases issuing from them are necessarily very hot. If the heat they contain was not made use of there would be a great waste of fuel; consequently in immediate proximity to these furnaces there is generally a series of pans or evaporating boilers, under which the gases pass, and into which the alkali solution is poured. On evaporating the solution, first of all the undecomposed sodium sulphate separates, then the sodium carbonate or soda crystals. These crystals as they separate are raked out and placed on planks, where the liquid drains away from them. Caustic soda remains in the residue, and also any sodium chloride which was not decomposed in the foregoing process.Part of the sodium carbonate is recrystallised in order to purify it more thoroughly. In order to do this a saturated solution is left to crystallise at a temperature below 30° in a current of air, in order to promote the separation of the water vapour. The large transparent crystals (efflorescent in air) of Na2CO3,10H2O are then formed which have already been spoken of (ChapterI.).
[13]
see captionFig.69.—Apparatus for the methodical lixiviation of black ash, &c. Water flows into the tanks from the pipesr,r, and the saturated liquid is drawn off fromc,c.
Fig.69.—Apparatus for the methodical lixiviation of black ash, &c. Water flows into the tanks from the pipesr,r, and the saturated liquid is drawn off fromc,c.
Methodical lixiviationis the extraction, by means of water, of a soluble substance from the mass containing it. It is carried on so as not to obtain weak aqueous solutions, and in such a way that the residue shall not contain any of the soluble substance. This problem is practically of great importance in many industries. It is required to extract from the mass all that is soluble in water. This is easily effected if water be first poured on the mass, the strong solution thus obtained decanted, then water again poured on, time being allowed for it to act, then again decanted, and so on until fresh water does not take up anything. But then finally such weak solutions are obtained that it would be very disadvantageous to evaporate them. This is avoided by pouring the fresh hot water destined for the lixiviation, not onto the fresh mass, but upon a mass which has already been subjected to a first lixiviation by weak solutions. In this way the fresh water gives a weak solution. The strong solution which goes to the evaporating pan flows from those parts of the apparatus which contain the fresh, as yet unlixiviated, mass, and thus in the latter parts the weak alkali formed in the other parts of the apparatus becomes saturated as far as possible with the soluble substance. Generally several intercommunicating vessels are constructed (standing at the same level) into which in turn the fresh mass is charged which is intended for lixiviation; the water is poured in, the alkali drawn off, and the lixiviated residue removed. The illustration represents such an apparatus, consisting of four communicating vessels. The water poured into one of them flows through the two nearest and issues from the third. The fresh mass being placed in one of these boxes or vessels, the stream of water passing through the apparatus is directed in such a manner as to finally issue from this vessel containing the fresh unlixiviated mass. The fresh water is added to the vessel containing the material which has been almost completely exhausted. Passing through this vessel it is conveyed by the pipe (syphon passing from the bottom of the first box to the top of the second) communicating with the second; it finally passes (also through a syphon pipe) into the box (the third) containing the fresh material. The water will extract all that is soluble in the first vessel, leaving only an insoluble residue. This vessel is then ready to be emptied, and refilled with fresh material. The levels of the liquids in the various vessels will naturally be different, in consequence of the various strengths of the solutions which they contain.
It must not, however, be thought that sodium carbonate alone passes into the solution; there is also a good deal of caustic soda with it, formed by the action of lime on the carbonate of sodium, and there are also certain sodium sulphur compounds with which we shall partly become acquainted hereafter. The sodium carbonate, therefore, is not obtained in a very pure state. The solution is concentrated by evaporation. This is conducted by means of the waste heat from the soda furnaces, together with that of the gases given off. The process in the soda furnaces can only be carried on at a high temperature, and therefore the smoke and gases issuing from them are necessarily very hot. If the heat they contain was not made use of there would be a great waste of fuel; consequently in immediate proximity to these furnaces there is generally a series of pans or evaporating boilers, under which the gases pass, and into which the alkali solution is poured. On evaporating the solution, first of all the undecomposed sodium sulphate separates, then the sodium carbonate or soda crystals. These crystals as they separate are raked out and placed on planks, where the liquid drains away from them. Caustic soda remains in the residue, and also any sodium chloride which was not decomposed in the foregoing process.
Part of the sodium carbonate is recrystallised in order to purify it more thoroughly. In order to do this a saturated solution is left to crystallise at a temperature below 30° in a current of air, in order to promote the separation of the water vapour. The large transparent crystals (efflorescent in air) of Na2CO3,10H2O are then formed which have already been spoken of (ChapterI.).
[14]The whole of the sulphur used in the production of the sulphuric acid employed in decomposing the common salt is contained in this residue. This is the great burden and expense of the soda works which use Leblanc's method. As an instructive example from a chemical point of view, it is worth while mentioning here two of the various methods of recovering the sulphur from the soda waste. Chance's process is treated in Chapter XX., Note 6.Kynaston (1885) treats the soda waste with a solution (sp. gr. l·21) of magnesium chloride, which disengages sulphuretted hydrogen: CaS + MgCl2+ 2H2O = CaCl2+ Mg(OH)2+ H2S. Sulphurous anhydride is passed through the residue in order to form the insoluble calcium sulphite: CaCl2+ Mg(OH)2+ SO2= CaSO3+ MgCl2+ H2O. The solution of magnesium chloride obtained is again used, and the washed calcium sulphite is brought into contact at a low temperature with hydrochloric acid (a weak aqueous solution) and hydrogen sulphide, the whole of the sulphur then separating:CaSO3+ 2H2S + 2HCl = CaCl2+ 3H2O + 3S.But most efforts have been directed towards avoiding the formation of soda waste.
[14]The whole of the sulphur used in the production of the sulphuric acid employed in decomposing the common salt is contained in this residue. This is the great burden and expense of the soda works which use Leblanc's method. As an instructive example from a chemical point of view, it is worth while mentioning here two of the various methods of recovering the sulphur from the soda waste. Chance's process is treated in Chapter XX., Note 6.
Kynaston (1885) treats the soda waste with a solution (sp. gr. l·21) of magnesium chloride, which disengages sulphuretted hydrogen: CaS + MgCl2+ 2H2O = CaCl2+ Mg(OH)2+ H2S. Sulphurous anhydride is passed through the residue in order to form the insoluble calcium sulphite: CaCl2+ Mg(OH)2+ SO2= CaSO3+ MgCl2+ H2O. The solution of magnesium chloride obtained is again used, and the washed calcium sulphite is brought into contact at a low temperature with hydrochloric acid (a weak aqueous solution) and hydrogen sulphide, the whole of the sulphur then separating:
CaSO3+ 2H2S + 2HCl = CaCl2+ 3H2O + 3S.
But most efforts have been directed towards avoiding the formation of soda waste.
[15]Among the drawbacks of the Leblanc process are the accumulation of ‘soda waste’ (Note14) owing to the impossibility at the comparatively low price of sulphur (especially in the form of pyrites) of finding employment for the sulphur and sulphur compounds for which this waste is sometimes treated, and also the insufficient purity of the sodium carbonate for many purposes. The advantages of the Leblanc process, besides its simplicity and cheapness, are that almost the whole of the acids obtained as bye-products have a commercial value; for chlorine and bleaching powder are produced from the large amount of hydrochloric acid which appears as a bye-product; caustic soda also is very easily made, and the demand for it increases every year. In those places where salt, pyrites, charcoal, and limestone (the materials required for alkali works) are found side by side—as, for instance, in the Ural or Don districts—conditions are favourable to the development of the manufacture of sodium carbonate on an enormous scale; and where, as in the Caucasus, sodium sulphate occurs naturally, the conditions are still more favourable. A large amount, however, of the latter salt, even from soda works, is used in making glass. The most important soda works, as regards the quantity of products obtained from them, are the English works.As an example of the other numerous and varied methods of manufacturing soda from sodium chloride, the following may be mentioned: Sodium chloride is decomposed by oxide of lead, PbO, forming lead chloride and sodium oxide, which, with carbonic anhydride, yields sodium carbonate (Scheele's process). In Cornu's method sodium chloride is treated with lime, and then exposed to the air, when it yields a small quantity of sodium carbonate. In E. Kopp's process sodium sulphate (125 parts) is mixed with oxide of iron (80 parts) and charcoal (55 parts), and the mixture is heated in reverberatory furnaces. Here a compound, Na6Fe4S3, is formed, which is insoluble in water absorbs oxygen and carbonic anhydride, and then forms sodium carbonate and ferrous sulphide; this when roasted gives sulphurous anhydride, the indispensable material for the manufacture of sulphuric acid, and ferric oxide which is again used in the process. In Grant's method sodium sulphate is transformed into sodium sulphide, and the latter is decomposed by a stream of carbonic anhydride and steam, when hydrogen sulphide is disengaged and sodium carbonate formed. Gossage prepares Na2S from Na2SO4(by heating it with carbon), dissolves it in water and subjects the solution to the action of an excess of CO2in coke towers, thus obtaining H2S (a gas which gives SO2under perfect combustion, or sulphur when incompletely burnt, Chapter XX., Note 6) and bicarbonate of sodium; Na2S + 2CO2+ 2H2O = H2S + 2HNaCO3. The latter gives soda and CO2when ignited. This process quite eliminates the formation of soda-waste (seeNote3) and should in my opinion be suitable for the treatment of native Na2SO4, like that which is found in the Caucasus, all the more since H2S gives sulphur as a bye-product.Repeated efforts have been made in recent times to obtain soda (and chlorine,seeChapter II., Note1) from strong solutions of salt (Chapter X., Note23 bis) by the action of an electric current, but until now these methods have not been worked out sufficiently for practical use, probably partly owing to the complicated apparatus needed, and the fact that the chlorine given off at the anode corrodes the electrodes and vessels and has but a limited industrial application. We may mention that according to Hempel (1890) soda in crystals is deposited when an electric current and a stream of carbonic acid gas are passed through a saturated solution of NaCl.Sodium carbonate may likewise be obtained from cryolite (Chapter XVII., Note 23) the method of treating this will be mentioned under Aluminium.
[15]Among the drawbacks of the Leblanc process are the accumulation of ‘soda waste’ (Note14) owing to the impossibility at the comparatively low price of sulphur (especially in the form of pyrites) of finding employment for the sulphur and sulphur compounds for which this waste is sometimes treated, and also the insufficient purity of the sodium carbonate for many purposes. The advantages of the Leblanc process, besides its simplicity and cheapness, are that almost the whole of the acids obtained as bye-products have a commercial value; for chlorine and bleaching powder are produced from the large amount of hydrochloric acid which appears as a bye-product; caustic soda also is very easily made, and the demand for it increases every year. In those places where salt, pyrites, charcoal, and limestone (the materials required for alkali works) are found side by side—as, for instance, in the Ural or Don districts—conditions are favourable to the development of the manufacture of sodium carbonate on an enormous scale; and where, as in the Caucasus, sodium sulphate occurs naturally, the conditions are still more favourable. A large amount, however, of the latter salt, even from soda works, is used in making glass. The most important soda works, as regards the quantity of products obtained from them, are the English works.
As an example of the other numerous and varied methods of manufacturing soda from sodium chloride, the following may be mentioned: Sodium chloride is decomposed by oxide of lead, PbO, forming lead chloride and sodium oxide, which, with carbonic anhydride, yields sodium carbonate (Scheele's process). In Cornu's method sodium chloride is treated with lime, and then exposed to the air, when it yields a small quantity of sodium carbonate. In E. Kopp's process sodium sulphate (125 parts) is mixed with oxide of iron (80 parts) and charcoal (55 parts), and the mixture is heated in reverberatory furnaces. Here a compound, Na6Fe4S3, is formed, which is insoluble in water absorbs oxygen and carbonic anhydride, and then forms sodium carbonate and ferrous sulphide; this when roasted gives sulphurous anhydride, the indispensable material for the manufacture of sulphuric acid, and ferric oxide which is again used in the process. In Grant's method sodium sulphate is transformed into sodium sulphide, and the latter is decomposed by a stream of carbonic anhydride and steam, when hydrogen sulphide is disengaged and sodium carbonate formed. Gossage prepares Na2S from Na2SO4(by heating it with carbon), dissolves it in water and subjects the solution to the action of an excess of CO2in coke towers, thus obtaining H2S (a gas which gives SO2under perfect combustion, or sulphur when incompletely burnt, Chapter XX., Note 6) and bicarbonate of sodium; Na2S + 2CO2+ 2H2O = H2S + 2HNaCO3. The latter gives soda and CO2when ignited. This process quite eliminates the formation of soda-waste (seeNote3) and should in my opinion be suitable for the treatment of native Na2SO4, like that which is found in the Caucasus, all the more since H2S gives sulphur as a bye-product.
Repeated efforts have been made in recent times to obtain soda (and chlorine,seeChapter II., Note1) from strong solutions of salt (Chapter X., Note23 bis) by the action of an electric current, but until now these methods have not been worked out sufficiently for practical use, probably partly owing to the complicated apparatus needed, and the fact that the chlorine given off at the anode corrodes the electrodes and vessels and has but a limited industrial application. We may mention that according to Hempel (1890) soda in crystals is deposited when an electric current and a stream of carbonic acid gas are passed through a saturated solution of NaCl.
Sodium carbonate may likewise be obtained from cryolite (Chapter XVII., Note 23) the method of treating this will be mentioned under Aluminium.
[16]This process (Chapter XVII.) was first pointed out by Turck, worked out by Schloesing, and finally applied industrially by Solvay. The first (1883) large soda factories erected in Russia for working this process are on the banks of the Kama at Berezniak, near Ousolia, and belong to Lubimoff. But Russia, which still imports from abroad a large quantity of bleaching powder and exports a large amount of manganese ore, most of all requires works carrying on the Leblanc process. In 1890 a factory of this kind was erected by P. K. Oushkoff, on the Kama, near Elagoubi.
[16]This process (Chapter XVII.) was first pointed out by Turck, worked out by Schloesing, and finally applied industrially by Solvay. The first (1883) large soda factories erected in Russia for working this process are on the banks of the Kama at Berezniak, near Ousolia, and belong to Lubimoff. But Russia, which still imports from abroad a large quantity of bleaching powder and exports a large amount of manganese ore, most of all requires works carrying on the Leblanc process. In 1890 a factory of this kind was erected by P. K. Oushkoff, on the Kama, near Elagoubi.
[16 bis]Mond (seeChapter XI., Note3 bis) separates the NH4Cl from the residual solutions by cooling (Chapter X., Note44); ignites the sal-ammoniac and passes the vapour over MgO, and so re-obtains the NH3, and forms MgCl2: the former goes back for the manufacture of soda, while the latter is employed either for making HCl or Cl2.
[16 bis]Mond (seeChapter XI., Note3 bis) separates the NH4Cl from the residual solutions by cooling (Chapter X., Note44); ignites the sal-ammoniac and passes the vapour over MgO, and so re-obtains the NH3, and forms MgCl2: the former goes back for the manufacture of soda, while the latter is employed either for making HCl or Cl2.
[17]Commercial soda ash (calcined, anhydrous) is rarely pure; the crystallised soda is generally purer. In order to purify it further, it is best to boil a concentrated solution of soda ash until two-thirds of the liquid remain, collect the soda which settles, wash with cold water, and then shake up with a strong solution of ammonia, pour off the residue, and heat. The impurities will then remain in the mother liquors, &c.Some numerical data may be given for sodium carbonate. The specific gravity of the anhydrous salt is 2·48, that of the decahydrated salt 1·46. Two varieties are known of the heptahydrated salt (Löwel, Marignac, Rammelsberg), which are formed together by allowing a saturated solution to cool under a layer of alcohol; the one is less stable (like the corresponding sulphate) and at 0° has a solubility of 32 parts (of anhydrous salt) in 100 water; the other is more stable, and its solubility 20 parts (of anhydrous salt) per 100 of water. The solubility of the decahydrated salt in 100 water = at 0°, 7·0; at 20°, 21·7; at 30°, 37·2 parts (of anhydrous salt). At 80° the solubility is only 46·1, at 90° 45·7, at 100°, 45·4 parts (of anhydrous salt). That is, it falls as the temperature rises, like Na2SO4. The specific gravity (Note7) of the solutions of sodium carbonate, according to the data of Gerlach and Kohlrausch, at 15°/4° is expressed by the formula,s= 9,992 + 104·5p+ 0·165p2. Weak solutions occupy a volume not only less than the sum of the volumes of the anhydrous salt and the water, but even less than the water contained in them. For instance, 1,000 grams of a 1 p.c. solution occupy (at 15°) a volume of 990·4 c.c. (sp. gr. 1·0097), but contain 990 grams of water, occupying at 15° a volume of 990·8 c.c. A similar case, which is comparatively rare occurs also with sodium hydroxide, in those dilute solutions for which the factorAis greater than 100 if the sp. gr. of water at 4° = 100,000, and if the sp. gr. of the solution be expressed by the formulaS=S0+Ap+Bp2, whereS0is the specific gravity of the water. For 5 p.c. the sp. gr. 15°/4° = 1·0520; for 10 p.c. 1·1057; for 15 p.c. 1·1603. The changes in the sp. gr. with the temperature are here almost the same as with solutions of sodium chloride with an equal value ofp.
[17]Commercial soda ash (calcined, anhydrous) is rarely pure; the crystallised soda is generally purer. In order to purify it further, it is best to boil a concentrated solution of soda ash until two-thirds of the liquid remain, collect the soda which settles, wash with cold water, and then shake up with a strong solution of ammonia, pour off the residue, and heat. The impurities will then remain in the mother liquors, &c.
Some numerical data may be given for sodium carbonate. The specific gravity of the anhydrous salt is 2·48, that of the decahydrated salt 1·46. Two varieties are known of the heptahydrated salt (Löwel, Marignac, Rammelsberg), which are formed together by allowing a saturated solution to cool under a layer of alcohol; the one is less stable (like the corresponding sulphate) and at 0° has a solubility of 32 parts (of anhydrous salt) in 100 water; the other is more stable, and its solubility 20 parts (of anhydrous salt) per 100 of water. The solubility of the decahydrated salt in 100 water = at 0°, 7·0; at 20°, 21·7; at 30°, 37·2 parts (of anhydrous salt). At 80° the solubility is only 46·1, at 90° 45·7, at 100°, 45·4 parts (of anhydrous salt). That is, it falls as the temperature rises, like Na2SO4. The specific gravity (Note7) of the solutions of sodium carbonate, according to the data of Gerlach and Kohlrausch, at 15°/4° is expressed by the formula,s= 9,992 + 104·5p+ 0·165p2. Weak solutions occupy a volume not only less than the sum of the volumes of the anhydrous salt and the water, but even less than the water contained in them. For instance, 1,000 grams of a 1 p.c. solution occupy (at 15°) a volume of 990·4 c.c. (sp. gr. 1·0097), but contain 990 grams of water, occupying at 15° a volume of 990·8 c.c. A similar case, which is comparatively rare occurs also with sodium hydroxide, in those dilute solutions for which the factorAis greater than 100 if the sp. gr. of water at 4° = 100,000, and if the sp. gr. of the solution be expressed by the formulaS=S0+Ap+Bp2, whereS0is the specific gravity of the water. For 5 p.c. the sp. gr. 15°/4° = 1·0520; for 10 p.c. 1·1057; for 15 p.c. 1·1603. The changes in the sp. gr. with the temperature are here almost the same as with solutions of sodium chloride with an equal value ofp.
[18]The resemblance is so great that, notwithstanding the difference in the molecular composition of Na2SO4and Na2CO3, they ought to be classed under the type (NaO)2R, where R = SO2or CO. Many other sodium salts also contain 10 mol. H2O.
[18]The resemblance is so great that, notwithstanding the difference in the molecular composition of Na2SO4and Na2CO3, they ought to be classed under the type (NaO)2R, where R = SO2or CO. Many other sodium salts also contain 10 mol. H2O.
[19]According to the observations of Pickering. According to Rose, when solutions of sodium carbonate are boiled a certain amount of carbonic anhydride is disengaged.
[19]According to the observations of Pickering. According to Rose, when solutions of sodium carbonate are boiled a certain amount of carbonic anhydride is disengaged.
[20]The composition of this salt, however, may be also represented as a combination of carbonic acid, H2CO3, with the normal salt, Na2CO3, just as the latter also combines with water. Such a combination is all the more likely because (1) there exists another salt, Na2CO3,2NaHCO3,2H2O (sodium sesquicarbonate), obtained by cooling a boiling solution of sodium bicarbonate, or by mixing this salt with the normal salt; but the formula of this salt cannot be derived from that of normal carbonic acid, as the formula of the bicarbonate can. At the same time the sesqui-salt has all the properties of a definite compound; it crystallises in transparent crystals, has a constant composition, its solubility (at 0° in 100 of water, 12·6 of anhydrous salt) differs from the solubility of the normal and acid salts; it is found in nature, and is known by the names oftronaandurao. The observations of Watts and Richards showed (1886) that on pouring a strong solution of the acid salt into a solution of the normal salt saturated by heating, crystals of the salt NaHCO3,Na2CO3,2H2O may be easily obtained, as long as the temperature is above 35°. The natural urao (Boussingault) has, according to Laurent, the same composition. This salt is very stable in air, and may be used for purifying sodium carbonate on the large scale. Such compounds have been little studied from a theoretical point of view, although particularly interesting, since in all probability they correspond with ortho-carbonic acid, C(OH)4, and at the same time correspond with double salts like astrakhanite (Chapter XIV., Note25). (2) Water of crystallisation does not enter into the composition of the crystals of the acid salt, so that on its formation (occurring only at low temperatures, as in the formation of crystalline compounds with water) the water of crystallisation of the normal salt separates and the water is, as it were, replaced by the elements of carbonic acid. If anhydrous sodium carbonate be mixed with the amount of water requisite for the formation of Na2CO3,H2O, this salt will, when powdered, absorb CO2as easily at the ordinary temperature as it does water.
[20]The composition of this salt, however, may be also represented as a combination of carbonic acid, H2CO3, with the normal salt, Na2CO3, just as the latter also combines with water. Such a combination is all the more likely because (1) there exists another salt, Na2CO3,2NaHCO3,2H2O (sodium sesquicarbonate), obtained by cooling a boiling solution of sodium bicarbonate, or by mixing this salt with the normal salt; but the formula of this salt cannot be derived from that of normal carbonic acid, as the formula of the bicarbonate can. At the same time the sesqui-salt has all the properties of a definite compound; it crystallises in transparent crystals, has a constant composition, its solubility (at 0° in 100 of water, 12·6 of anhydrous salt) differs from the solubility of the normal and acid salts; it is found in nature, and is known by the names oftronaandurao. The observations of Watts and Richards showed (1886) that on pouring a strong solution of the acid salt into a solution of the normal salt saturated by heating, crystals of the salt NaHCO3,Na2CO3,2H2O may be easily obtained, as long as the temperature is above 35°. The natural urao (Boussingault) has, according to Laurent, the same composition. This salt is very stable in air, and may be used for purifying sodium carbonate on the large scale. Such compounds have been little studied from a theoretical point of view, although particularly interesting, since in all probability they correspond with ortho-carbonic acid, C(OH)4, and at the same time correspond with double salts like astrakhanite (Chapter XIV., Note25). (2) Water of crystallisation does not enter into the composition of the crystals of the acid salt, so that on its formation (occurring only at low temperatures, as in the formation of crystalline compounds with water) the water of crystallisation of the normal salt separates and the water is, as it were, replaced by the elements of carbonic acid. If anhydrous sodium carbonate be mixed with the amount of water requisite for the formation of Na2CO3,H2O, this salt will, when powdered, absorb CO2as easily at the ordinary temperature as it does water.
[21]100 parts of water at 0° dissolve 7 parts of the acid salt, which corresponds with 4·3 parts of the anhydrous normal salt, but at 0° 100 parts of water dissolve 7 parts of the latter. The solubility of the bi- or acid salt varies with considerable regularity; 100 parts of water dissolves at 15° 9 parts of the salt, at 30° 11 parts.The ammonium, and more especially the calcium, salt, is much more soluble in water. The ammonia process (seep.524) is founded upon this. Ammonium bicarbonate (acid carbonate) at 0° has a solubility of 12 parts in 100 water, at 30° of 27 parts. The solubility therefore increases very rapidly with the temperature. And its saturated solution is more stable than a solution of sodium bicarbonate. In fact, saturated solutions of these salts have a gaseous tension like that of a mixture of carbonic anhydride and water—namely, at 15° and at 50°, for the sodium salt 120 and 750 millimetres, for the ammonium salt 120 and 563 millimetres. These data are of great importance in understanding the phenomena connected with the ammonia process. They indicate that with an increased pressure the formation of the sodium salt ought to increase if there be an excess of ammonium salt.
[21]100 parts of water at 0° dissolve 7 parts of the acid salt, which corresponds with 4·3 parts of the anhydrous normal salt, but at 0° 100 parts of water dissolve 7 parts of the latter. The solubility of the bi- or acid salt varies with considerable regularity; 100 parts of water dissolves at 15° 9 parts of the salt, at 30° 11 parts.
The ammonium, and more especially the calcium, salt, is much more soluble in water. The ammonia process (seep.524) is founded upon this. Ammonium bicarbonate (acid carbonate) at 0° has a solubility of 12 parts in 100 water, at 30° of 27 parts. The solubility therefore increases very rapidly with the temperature. And its saturated solution is more stable than a solution of sodium bicarbonate. In fact, saturated solutions of these salts have a gaseous tension like that of a mixture of carbonic anhydride and water—namely, at 15° and at 50°, for the sodium salt 120 and 750 millimetres, for the ammonium salt 120 and 563 millimetres. These data are of great importance in understanding the phenomena connected with the ammonia process. They indicate that with an increased pressure the formation of the sodium salt ought to increase if there be an excess of ammonium salt.
[22]Crystalline sodium carbonate (broken into lumps) also absorbs carbonic anhydride, but the water contained in the crystals is then disengaged: Na2CO3,10H2O + CO2= Na2CO3,H2CO3+ 9H2O, and dissolves part of the carbonate; therefore part of the sodium carbonate passes into solution together with all the impurities. When it is required to avoid the formation of this solution, a mixture of ignited and crystalline sodium carbonate is taken. Sodium bicarbonate is prepared chiefly for medicinal use, and is then often termedcarbonate of soda, also, for instance, in the so-called soda powders, for preparing certain artificial mineral waters, for the manufacture of digestive lozenges like those made at Essentuki, Vichy, &c.
[22]Crystalline sodium carbonate (broken into lumps) also absorbs carbonic anhydride, but the water contained in the crystals is then disengaged: Na2CO3,10H2O + CO2= Na2CO3,H2CO3+ 9H2O, and dissolves part of the carbonate; therefore part of the sodium carbonate passes into solution together with all the impurities. When it is required to avoid the formation of this solution, a mixture of ignited and crystalline sodium carbonate is taken. Sodium bicarbonate is prepared chiefly for medicinal use, and is then often termedcarbonate of soda, also, for instance, in the so-called soda powders, for preparing certain artificial mineral waters, for the manufacture of digestive lozenges like those made at Essentuki, Vichy, &c.
[23]In chemistry, sodium oxide is termed ‘soda,’ which word must be carefully distinguished from the word sodium, meaning the metal.
[23]In chemistry, sodium oxide is termed ‘soda,’ which word must be carefully distinguished from the word sodium, meaning the metal.
[24]With a small quantity of water, the reaction either does not take place, or even proceeds in the reverse way—that is, sodium and potassium hydroxides remove carbonic anhydride from calcium carbonate (Liebig, Watson, Mitscherlich, and others). The influence of the mass of water is evident. According to Gerberts, however, strong solutions of sodium carbonate are decomposed by lime, which is very interesting if confirmed by further investigation.
[24]With a small quantity of water, the reaction either does not take place, or even proceeds in the reverse way—that is, sodium and potassium hydroxides remove carbonic anhydride from calcium carbonate (Liebig, Watson, Mitscherlich, and others). The influence of the mass of water is evident. According to Gerberts, however, strong solutions of sodium carbonate are decomposed by lime, which is very interesting if confirmed by further investigation.
[25]As long as any undecomposed sodium carbonate remains in solution, excess of acid added to the solution disengages carbonic anhydride, and the solution after dilution gives a white precipitate with a barium salt soluble in acids, showing the presence of a carbonate in solution (if there be sulphate present, it also forms a white precipitate, but this is insoluble in acids). For the decomposition of sodium carbonate, milk of lime—that is, slaked slime suspended in water—is employed. Formerly pure sodium hydroxide was prepared (according to Berthollet's process) by dissolving the impure substance in alcohol (sodium carbonate and sulphate are not soluble), but now that metallic sodium has become cheap and is purified by distillation,pure caustic sodais prepared by acting on a small quantity of water with sodium. Perfectly pure sodium hydroxide may also be obtained by allowing strong solutions to crystallise (in the cold) (Note27).In alkali works where the Leblanc process is used, caustic soda is prepared directly from the alkali remaining in the mother liquors after the separation of the sodium carbonate by evaporation (Note14). If excess of lime and charcoal have been used, much sodium hydroxide maybe obtained. After the removal as much as possible of the sodium carbonate, a red liquid (from iron oxide) is left, containing sodium hydroxide mixed with compounds of sulphur and of cyanogen (seeChapterIX.) and also containing iron. This red alkali is evaporated and air is blown through it, which oxidises the impurities (for this purpose sometimes sodium nitrate is added, or bleaching powder, &c.) and leaves fused caustic soda. The fused mass is allowed to settle in order to separate the ferruginous precipitate, and poured into iron drums, where the sodium hydroxide solidifies. Such caustic soda contains about 10 p.c. of water in excess and some saline impurities, but when properly manufactured is almost free from carbonate and from iron. The greater part of the caustic soda, which forms so important an article of commerce, is manufactured in this manner.
[25]As long as any undecomposed sodium carbonate remains in solution, excess of acid added to the solution disengages carbonic anhydride, and the solution after dilution gives a white precipitate with a barium salt soluble in acids, showing the presence of a carbonate in solution (if there be sulphate present, it also forms a white precipitate, but this is insoluble in acids). For the decomposition of sodium carbonate, milk of lime—that is, slaked slime suspended in water—is employed. Formerly pure sodium hydroxide was prepared (according to Berthollet's process) by dissolving the impure substance in alcohol (sodium carbonate and sulphate are not soluble), but now that metallic sodium has become cheap and is purified by distillation,pure caustic sodais prepared by acting on a small quantity of water with sodium. Perfectly pure sodium hydroxide may also be obtained by allowing strong solutions to crystallise (in the cold) (Note27).
In alkali works where the Leblanc process is used, caustic soda is prepared directly from the alkali remaining in the mother liquors after the separation of the sodium carbonate by evaporation (Note14). If excess of lime and charcoal have been used, much sodium hydroxide maybe obtained. After the removal as much as possible of the sodium carbonate, a red liquid (from iron oxide) is left, containing sodium hydroxide mixed with compounds of sulphur and of cyanogen (seeChapterIX.) and also containing iron. This red alkali is evaporated and air is blown through it, which oxidises the impurities (for this purpose sometimes sodium nitrate is added, or bleaching powder, &c.) and leaves fused caustic soda. The fused mass is allowed to settle in order to separate the ferruginous precipitate, and poured into iron drums, where the sodium hydroxide solidifies. Such caustic soda contains about 10 p.c. of water in excess and some saline impurities, but when properly manufactured is almost free from carbonate and from iron. The greater part of the caustic soda, which forms so important an article of commerce, is manufactured in this manner.
[26]Löwig gave a method of preparing sodium hydroxide from sodium carbonate by heating it to a dull red heat with an excess of ferric oxide. Carbonic anhydride is given off, and warm water extracts the caustic soda from the remaining mass. This reaction, as experiment shows, proceeds very easily, and is an example of contact action similar to that of ferric oxide on the decomposition of potassium chlorate. The reason of this may be that a small quantity of the sodium carbonate enters into double decomposition with the ferric oxide, and the ferric carbonate produced is decomposed into carbonic anhydride and ferric oxide, the action of which is renewed. Similar explanations expressing thereasonfor a reaction really adds but little to that elementary conception of contact which, according to my opinion, consists in the change of motion of the atoms in the molecules under the influence of the substance in contact. In order to represent this clearly it is sufficient, for instance, to imagine that in the sodium carbonate the elements CO2move in a circle round the elements Na2O, but at the points of contact with Fe2O3the motion becomes elliptic with a long axis, and at some distance from Na2O the elements of CO2are parted, not having the faculty of attaching themselves to Fe2O3.
[26]Löwig gave a method of preparing sodium hydroxide from sodium carbonate by heating it to a dull red heat with an excess of ferric oxide. Carbonic anhydride is given off, and warm water extracts the caustic soda from the remaining mass. This reaction, as experiment shows, proceeds very easily, and is an example of contact action similar to that of ferric oxide on the decomposition of potassium chlorate. The reason of this may be that a small quantity of the sodium carbonate enters into double decomposition with the ferric oxide, and the ferric carbonate produced is decomposed into carbonic anhydride and ferric oxide, the action of which is renewed. Similar explanations expressing thereasonfor a reaction really adds but little to that elementary conception of contact which, according to my opinion, consists in the change of motion of the atoms in the molecules under the influence of the substance in contact. In order to represent this clearly it is sufficient, for instance, to imagine that in the sodium carbonate the elements CO2move in a circle round the elements Na2O, but at the points of contact with Fe2O3the motion becomes elliptic with a long axis, and at some distance from Na2O the elements of CO2are parted, not having the faculty of attaching themselves to Fe2O3.
[27]By allowing strong solutions of sodium hydroxide to crystallise in the cold, impurities—such as, for instance, sodium sulphate—may be separated from them. The fused crystallo-hydrate 2NaHO,7H2O forms a solution having a specific gravity of 1·405 (Hermes). The crystals on dissolving in water produce cold, while NaHO produces heat. Besides which Pickering obtained hydrates with 1, 2, 4, 5, and 7 H2O.
[27]By allowing strong solutions of sodium hydroxide to crystallise in the cold, impurities—such as, for instance, sodium sulphate—may be separated from them. The fused crystallo-hydrate 2NaHO,7H2O forms a solution having a specific gravity of 1·405 (Hermes). The crystals on dissolving in water produce cold, while NaHO produces heat. Besides which Pickering obtained hydrates with 1, 2, 4, 5, and 7 H2O.
[28]In solid caustic soda there is generally an excess of water beyond that required by the formula NaHO. The caustic soda used in laboratories is generally cast in sticks, which are broken into pieces. It must be preserved in carefully closed vessels, because it absorbs water and carbonic anhydride from the air.
[28]In solid caustic soda there is generally an excess of water beyond that required by the formula NaHO. The caustic soda used in laboratories is generally cast in sticks, which are broken into pieces. It must be preserved in carefully closed vessels, because it absorbs water and carbonic anhydride from the air.
[29]By the way it changes in air it is easy to distinguish caustic soda from caustic potash, which in general resembles it. Both alkalis absorb water and carbonic anhydride from the air, but caustic potash forms a deliquescent mass of potassium carbonate, whilst caustic soda forms a dry powder of efflorescent salt.
[29]By the way it changes in air it is easy to distinguish caustic soda from caustic potash, which in general resembles it. Both alkalis absorb water and carbonic anhydride from the air, but caustic potash forms a deliquescent mass of potassium carbonate, whilst caustic soda forms a dry powder of efflorescent salt.
[30]As the molecular weight of NaHO = 40, the volume of its molecule = 40/2·13 = 18·5, which very nearly approaches the volume of a molecule of water. The same rule applies to the compounds of sodium in general—for instance, its salts have a molecular volume approaching the volume of the acids from which they are derived.
[30]As the molecular weight of NaHO = 40, the volume of its molecule = 40/2·13 = 18·5, which very nearly approaches the volume of a molecule of water. The same rule applies to the compounds of sodium in general—for instance, its salts have a molecular volume approaching the volume of the acids from which they are derived.
[31]The molecular quantity of sodium hydroxide (40 grams), on being dissolved in a large mass (200 gram molecules) of water, develops, according to Berthelot 9,780, and according to Thomsen 9,940, heat-units, but at 100° about 13,000 (Berthelot). Solutions of NaHO +nH2O, on being mixed with water, evolve heat if they contain less than 6H2O, but if more they absorb beat.
[31]The molecular quantity of sodium hydroxide (40 grams), on being dissolved in a large mass (200 gram molecules) of water, develops, according to Berthelot 9,780, and according to Thomsen 9,940, heat-units, but at 100° about 13,000 (Berthelot). Solutions of NaHO +nH2O, on being mixed with water, evolve heat if they contain less than 6H2O, but if more they absorb beat.
[32]The specific gravity of solutions of sodium hydroxide at 15°/4° is given in the short table below:—NaHO, p.c.51015203040Sp. gr.1·0571·1131·1691·2241·3311·4361,000 grams of a 5 p.c. solution occupies a volume of 946 c.c.; that is, less than the water serving to make the solution (seeNote18).
[32]The specific gravity of solutions of sodium hydroxide at 15°/4° is given in the short table below:—
1,000 grams of a 5 p.c. solution occupies a volume of 946 c.c.; that is, less than the water serving to make the solution (seeNote18).
[33]Sodium hydroxide and some other alkalis are capable of hydrolysing—saponifying, as it is termed—the compounds of acids with alcohols. If RHO (or R(HO)n) represent the composition of an alcohol—that is, of the hydroxide of a hydrocarbon radicle—and QHO an acid, then the compound of the acid with the alcohol or ethereal salt of the given acid will have the composition RQO. Ethereal salts, therefore, present a likeness to metallic salts, just as alcohols resemble basic hydroxides. Sodium hydroxide acts on ethereal salts in the same way that it acts on the majority of metallic salts—namely, it liberates alcohol, and forms the sodium salt of that acid which was in the ethereal salt. The reaction takes place in the following way:—RQO+NaHO=NaQO+RHOEthereal saltCaustic sodaSodium saltAlcoholSuch a decomposition is termed saponification; similar reactions were known very long ago for the ethereal salts corresponding with glycerin, C3H5(OH)3(ChapterIX.), found in animals and plants, and composing what are called fats or oils. Caustic soda, acting on fat and oil, forms glycerin, and sodium salts of those acids which were in union with the glycerin in the fat, as Chevreul showed at the beginning of this century. The sodium salts of the fatty acids are commonly known as soaps. That is to say, soap is made from fat and caustic soda, glycerin being separated and a sodium salt or soap formed. As glycerin is usually found in union with certain acids, so also are the sodium salts of the same acids found in soap. The greater part of the acids found in conjunction with glycerin in fats are the solid palmitic and stearic acids, C16H32O2and C18H38O2, and the liquid oleic acid, C18H34O2. In preparing soap the fatty substances are mixed with a solution of caustic soda until an emulsion is formed; the proper quantity of caustic soda is then added in order to produce saponification on heating, the soap being separated from the solution either by means of an excess of caustic soda or else by common salt, which displaces the soap from the aqueous solution (salt water does not dissolve soap, neither does it form a lather). Water acting on soap partly decomposes it (because the acids of the soap are feeble), and the alkali set free acts during the application of soap. Hence it may be replaced by a very feeble alkali. Strong solutions of alkali corrode the skin and tissues. They are not formed from soap, because the reaction is reversible, and the alkali is only set free by the excess of water. Thus we see how the teaching of Berthollet renders it possible to understand many phenomena which occur in every-day experience (seeChapter IX., Note15).
[33]Sodium hydroxide and some other alkalis are capable of hydrolysing—saponifying, as it is termed—the compounds of acids with alcohols. If RHO (or R(HO)n) represent the composition of an alcohol—that is, of the hydroxide of a hydrocarbon radicle—and QHO an acid, then the compound of the acid with the alcohol or ethereal salt of the given acid will have the composition RQO. Ethereal salts, therefore, present a likeness to metallic salts, just as alcohols resemble basic hydroxides. Sodium hydroxide acts on ethereal salts in the same way that it acts on the majority of metallic salts—namely, it liberates alcohol, and forms the sodium salt of that acid which was in the ethereal salt. The reaction takes place in the following way:—
Such a decomposition is termed saponification; similar reactions were known very long ago for the ethereal salts corresponding with glycerin, C3H5(OH)3(ChapterIX.), found in animals and plants, and composing what are called fats or oils. Caustic soda, acting on fat and oil, forms glycerin, and sodium salts of those acids which were in union with the glycerin in the fat, as Chevreul showed at the beginning of this century. The sodium salts of the fatty acids are commonly known as soaps. That is to say, soap is made from fat and caustic soda, glycerin being separated and a sodium salt or soap formed. As glycerin is usually found in union with certain acids, so also are the sodium salts of the same acids found in soap. The greater part of the acids found in conjunction with glycerin in fats are the solid palmitic and stearic acids, C16H32O2and C18H38O2, and the liquid oleic acid, C18H34O2. In preparing soap the fatty substances are mixed with a solution of caustic soda until an emulsion is formed; the proper quantity of caustic soda is then added in order to produce saponification on heating, the soap being separated from the solution either by means of an excess of caustic soda or else by common salt, which displaces the soap from the aqueous solution (salt water does not dissolve soap, neither does it form a lather). Water acting on soap partly decomposes it (because the acids of the soap are feeble), and the alkali set free acts during the application of soap. Hence it may be replaced by a very feeble alkali. Strong solutions of alkali corrode the skin and tissues. They are not formed from soap, because the reaction is reversible, and the alkali is only set free by the excess of water. Thus we see how the teaching of Berthollet renders it possible to understand many phenomena which occur in every-day experience (seeChapter IX., Note15).
[34]On this is founded the process of Henkoff and Engelhardt for treating bones. The bones are mixed with ashes, lime, and water; it is true that in this case more potassium hydroxide than sodium hydroxide is formed, but their action is almost identical.
[34]On this is founded the process of Henkoff and Engelhardt for treating bones. The bones are mixed with ashes, lime, and water; it is true that in this case more potassium hydroxide than sodium hydroxide is formed, but their action is almost identical.
[35]As explained in Note33.
[35]As explained in Note33.
[35 bis]It might be expected, from what has been mentioned above, that bivalent metals would easily form acid salts with acids containing more than two atoms of hydrogen—for instance, with tribasic acids, such as phosphoric acid, H3PO4—and actually such salts do exist; but all such relations are complicated by the fact that the character of the base very often changes and becomes weakened with the increase of valency and the change of atomic weight; the feebler bases (like silver oxide), although corresponding with univalent metals, do not form acid salts, while the feeblest bases (CuO, PbO, &c.) easily form basic salts, and notwithstanding their valency do not form acid salts which are in any degree stable—that is, which are undecomposable by water. Basic and acid salts ought to be regarded rather as compounds similar to crystallo-hydrates, because such acids as sulphuric form with sodium not only an acid and a normal salt, as might be expected from the valency of sodium, but also salts containing a greater quantity of acid. In sodium sesquicarbonate we saw an example of such compounds. Taking all this into consideration, we must say that the property of more or less easily forming acid salts depends more upon the energy of the base than upon its valency, and the best statement is thatthe capacity of a base for forming acid and basic salts is characteristic, just as the faculty of forming compounds with hydrogen is characteristic of elements.
[35 bis]It might be expected, from what has been mentioned above, that bivalent metals would easily form acid salts with acids containing more than two atoms of hydrogen—for instance, with tribasic acids, such as phosphoric acid, H3PO4—and actually such salts do exist; but all such relations are complicated by the fact that the character of the base very often changes and becomes weakened with the increase of valency and the change of atomic weight; the feebler bases (like silver oxide), although corresponding with univalent metals, do not form acid salts, while the feeblest bases (CuO, PbO, &c.) easily form basic salts, and notwithstanding their valency do not form acid salts which are in any degree stable—that is, which are undecomposable by water. Basic and acid salts ought to be regarded rather as compounds similar to crystallo-hydrates, because such acids as sulphuric form with sodium not only an acid and a normal salt, as might be expected from the valency of sodium, but also salts containing a greater quantity of acid. In sodium sesquicarbonate we saw an example of such compounds. Taking all this into consideration, we must say that the property of more or less easily forming acid salts depends more upon the energy of the base than upon its valency, and the best statement is thatthe capacity of a base for forming acid and basic salts is characteristic, just as the faculty of forming compounds with hydrogen is characteristic of elements.
[36]Deville supposes that such a decomposition of sodium hydroxide by metallic iron depends solely on the dissociation of the alkali at a white heat into sodium, hydrogen, and oxygen. Here the part played by the iron is only that it retains the oxygen formed, otherwise the decomposed elements would again reunite upon cooling, as in other cases of dissociation. If it be supposed that the temperature at the commencement of the dissociation of the iron oxides is higher than that of sodium oxide, then the decomposition may be explained by Deville's hypothesis. Deville demonstrates his views by the following experiment:—An iron bottle, filled with iron borings, was heated in such a way that the upper part became red hot, the lower part remaining cooler; sodium hydroxide was introduced into the upper part. The decomposition was then effected—that is, sodium vapours were produced (this experiment was really performed with potassium hydroxide). On opening the bottle it was found that the iron in the upper part was not oxidised, but only that in the lower part. This may be explained by the decomposition of the alkali into sodium, hydrogen, and oxygen taking place in the upper part, whilst the iron in the lower part absorbed the oxygen set free. If the whole bottle be subjected to the same moderate heat as the lower extremity, no metallic vapours are formed. In that case, according to the hypothesis, the temperature is insufficient for the dissociation of the sodium hydroxide.
[36]Deville supposes that such a decomposition of sodium hydroxide by metallic iron depends solely on the dissociation of the alkali at a white heat into sodium, hydrogen, and oxygen. Here the part played by the iron is only that it retains the oxygen formed, otherwise the decomposed elements would again reunite upon cooling, as in other cases of dissociation. If it be supposed that the temperature at the commencement of the dissociation of the iron oxides is higher than that of sodium oxide, then the decomposition may be explained by Deville's hypothesis. Deville demonstrates his views by the following experiment:—An iron bottle, filled with iron borings, was heated in such a way that the upper part became red hot, the lower part remaining cooler; sodium hydroxide was introduced into the upper part. The decomposition was then effected—that is, sodium vapours were produced (this experiment was really performed with potassium hydroxide). On opening the bottle it was found that the iron in the upper part was not oxidised, but only that in the lower part. This may be explained by the decomposition of the alkali into sodium, hydrogen, and oxygen taking place in the upper part, whilst the iron in the lower part absorbed the oxygen set free. If the whole bottle be subjected to the same moderate heat as the lower extremity, no metallic vapours are formed. In that case, according to the hypothesis, the temperature is insufficient for the dissociation of the sodium hydroxide.
[37]It has been previously remarked (Chapter II. Note9) that Beketoff showed the displacement of sodium by hydrogen, not from sodium hydroxide but from the oxide Na2O; then, however, only one half is displaced, with the formation of NaHO.
[37]It has been previously remarked (Chapter II. Note9) that Beketoff showed the displacement of sodium by hydrogen, not from sodium hydroxide but from the oxide Na2O; then, however, only one half is displaced, with the formation of NaHO.
[38]Since the close of the eighties in England, where the preparation of sodium is at present carried out on a large commercial scale (from 1860 to 1870 it was only manufactured in a few works in France), it has been the practice to add to Deville's mixture iron, or iron oxide which with the charcoal gives metallic and carburetted iron, which still further facilitates the decomposition. At present a kilogram of sodium may be purchased for about the same sum (2/-) as a gram cost thirty years ago. Castner, in England, greatly improved the manufacture of sodium in large quantities, and so cheapened it as a reducing agent in the preparation of metallic aluminium. He heated a mixture of 44 parts of NaHO, and 7 parts of carbide of iron in large iron retorts at 1,000° and obtained about 6½ parts of metallic sodium. The reaction proceeds more easily than with carbon or iron alone, and the decomposition of the NaHO proceeds according to the equation: 3NaHO + C = Na2CO3+ 3H + Na. Subsequently, in 1891, aluminium was prepared by electrolysis (seeChapter XVII.), and metallic sodium found two new uses; (1) for the manufacture of peroxide of sodium (see later on) which is used in bleaching works, and (2) in the manufacture of potassium and sodium cyanide from yellow prussiate (Chapter XIII., Note12).
[38]Since the close of the eighties in England, where the preparation of sodium is at present carried out on a large commercial scale (from 1860 to 1870 it was only manufactured in a few works in France), it has been the practice to add to Deville's mixture iron, or iron oxide which with the charcoal gives metallic and carburetted iron, which still further facilitates the decomposition. At present a kilogram of sodium may be purchased for about the same sum (2/-) as a gram cost thirty years ago. Castner, in England, greatly improved the manufacture of sodium in large quantities, and so cheapened it as a reducing agent in the preparation of metallic aluminium. He heated a mixture of 44 parts of NaHO, and 7 parts of carbide of iron in large iron retorts at 1,000° and obtained about 6½ parts of metallic sodium. The reaction proceeds more easily than with carbon or iron alone, and the decomposition of the NaHO proceeds according to the equation: 3NaHO + C = Na2CO3+ 3H + Na. Subsequently, in 1891, aluminium was prepared by electrolysis (seeChapter XVII.), and metallic sodium found two new uses; (1) for the manufacture of peroxide of sodium (see later on) which is used in bleaching works, and (2) in the manufacture of potassium and sodium cyanide from yellow prussiate (Chapter XIII., Note12).
[38 bis]This is also shown by the fall in the temperature of solidification of tin produced by the addition of sodium (and also Al and Zn). Heycock and Neville (1889).
[38 bis]This is also shown by the fall in the temperature of solidification of tin produced by the addition of sodium (and also Al and Zn). Heycock and Neville (1889).
[39]By dissolving sodium amalgams in water and acids, and deducting the heat of solution of the sodium, Berthelot found thatfor each atom of the sodiumin amalgams containing a larger amount of mercury than NaHg5, the amount of heat evolved increases, after which the heat of formation falls, and the heat evolved decreases. In the formation of NaHg5about 18,500 calories are evolved; when NaHg3is formed, about 14,000; and for NaHg about 10,000 calories. Kraft regarded the definite crystalline amalgam as having the composition of NaHg6, but at the present time, in accordance with Grimaldi's results, it is thought to be NaHg5. A similar amalgam is very easily obtained if a 3 p.c. amalgam be left several days in a solution of sodium hydroxide until a crystalline mass is formed, from which the mercury may be removed by strongly pressing in chamois leather. This amalgam with a solution of potassium hydroxide forms a potassium amalgam, KHg10. It may be mentioned here that the latent heat of fusion (of atomic quantities) of Hg = 360 (Personne), Na = 730 (Joannis), and K = 610 calories (Joannis).
[39]By dissolving sodium amalgams in water and acids, and deducting the heat of solution of the sodium, Berthelot found thatfor each atom of the sodiumin amalgams containing a larger amount of mercury than NaHg5, the amount of heat evolved increases, after which the heat of formation falls, and the heat evolved decreases. In the formation of NaHg5about 18,500 calories are evolved; when NaHg3is formed, about 14,000; and for NaHg about 10,000 calories. Kraft regarded the definite crystalline amalgam as having the composition of NaHg6, but at the present time, in accordance with Grimaldi's results, it is thought to be NaHg5. A similar amalgam is very easily obtained if a 3 p.c. amalgam be left several days in a solution of sodium hydroxide until a crystalline mass is formed, from which the mercury may be removed by strongly pressing in chamois leather. This amalgam with a solution of potassium hydroxide forms a potassium amalgam, KHg10. It may be mentioned here that the latent heat of fusion (of atomic quantities) of Hg = 360 (Personne), Na = 730 (Joannis), and K = 610 calories (Joannis).
[40]Alloys are so similar to solutions (exhibiting such complete parallelism in properties) that they are included in the same class of so-called indefinite compounds. But in alloys, as substances passing from the liquid to the solid state, it is easier to discover the formation of definite chemical compounds. Besides the alloys of Na with Hg, those with tin (Bailey 1892 found Na2Sn), lead (NaPb), bismuth (Na3Bi), &c. (Joannis 1892 and others) have been investigated.
[40]Alloys are so similar to solutions (exhibiting such complete parallelism in properties) that they are included in the same class of so-called indefinite compounds. But in alloys, as substances passing from the liquid to the solid state, it is easier to discover the formation of definite chemical compounds. Besides the alloys of Na with Hg, those with tin (Bailey 1892 found Na2Sn), lead (NaPb), bismuth (Na3Bi), &c. (Joannis 1892 and others) have been investigated.
[41]Potassium forms a similar compound, but lithium, under the same circumstances, does not.
[41]Potassium forms a similar compound, but lithium, under the same circumstances, does not.
[42]The tension of dissociation of hydrogenp, in millimetres of mercury, is:—t=330°350°400°430°for Na2Hp=2857447910for K2H45725481100
[42]The tension of dissociation of hydrogenp, in millimetres of mercury, is:—
[43]In general, during the formation of alloys the volumes change very slightly, and therefore from the volume of Na2H some idea may be formed of the volume of hydrogen in a solid or liquid state. Even Archimedes concluded that there was gold in an alloy of copper and gold by reason of its volume and density. From the fact that the density of Na2H is equal to 0·959, it may be seen that the volume of 47 grams (the gram molecule) of this compound = 49·0 c.c. The volume of 46 grams of sodium contained in the Na2H (the density under the same conditions being 0·97) is equal to 47·4 c.c. Therefore the volume of 1 gram of hydrogen in Na2H is equal to 1·6 c.c., and consequently the density of metallic hydrogen, or the weight of 1 c.c., approaches 0·6 gram. This density is also proper to the hydrogen alloyed with potassium and palladium. Judging from the scanty information which is at present available, liquid hydrogen near its absolute boiling point (ChapterII.) has a much lower density.
[43]In general, during the formation of alloys the volumes change very slightly, and therefore from the volume of Na2H some idea may be formed of the volume of hydrogen in a solid or liquid state. Even Archimedes concluded that there was gold in an alloy of copper and gold by reason of its volume and density. From the fact that the density of Na2H is equal to 0·959, it may be seen that the volume of 47 grams (the gram molecule) of this compound = 49·0 c.c. The volume of 46 grams of sodium contained in the Na2H (the density under the same conditions being 0·97) is equal to 47·4 c.c. Therefore the volume of 1 gram of hydrogen in Na2H is equal to 1·6 c.c., and consequently the density of metallic hydrogen, or the weight of 1 c.c., approaches 0·6 gram. This density is also proper to the hydrogen alloyed with potassium and palladium. Judging from the scanty information which is at present available, liquid hydrogen near its absolute boiling point (ChapterII.) has a much lower density.
[43 bis]We may remark that at low temperatures Na absorbs NH3and forms (NH3Na)2(seeChapter VI., Note14); this substance absorbs CO and gives (NaCO)n (Chapter IX., Note31), although by itself Na does not combine directly with CO (but K does).
[43 bis]We may remark that at low temperatures Na absorbs NH3and forms (NH3Na)2(seeChapter VI., Note14); this substance absorbs CO and gives (NaCO)n (Chapter IX., Note31), although by itself Na does not combine directly with CO (but K does).
[44]H. A. Schmidt remarked that perfectly dry hydrogen chloride is decomposed with great difficulty by sodium, although the decomposition proceeds easily with potassium and with sodium in moist hydrogen chloride. Wanklyn also remarked that sodium burns with great difficulty in dry chlorine. Probably these facts are related to other phenomena observed by Dixon, who found that perfectly dry carbonic oxide does not explode with oxygen on passing an electric spark.
[44]H. A. Schmidt remarked that perfectly dry hydrogen chloride is decomposed with great difficulty by sodium, although the decomposition proceeds easily with potassium and with sodium in moist hydrogen chloride. Wanklyn also remarked that sodium burns with great difficulty in dry chlorine. Probably these facts are related to other phenomena observed by Dixon, who found that perfectly dry carbonic oxide does not explode with oxygen on passing an electric spark.
[44 bis]Sodamide, NH2Na, (Chapter IV., Note14), discovered by Gay-Lussac and Thénard, has formed the object of repeated research, but has been most fully investigated by A. W. Titherley (1894). Until recently the following was all that was known about this compound:—By heating sodium in dry ammonia, Gay-Lussac and Thénard obtained an olive-green, easily-fusible mass,sodamide, NH2Na, hydrogen being separated. This substance with water forms sodium hydroxide and ammonia; with carbonic oxide, CO, it forms sodium cyanide, NaCN, and water, H2O; and with dry hydrogen chloride it forms sodium and ammonium chlorides. These and other reactions of sodamide show that the metal in it preserves its energetic properties in reaction, and that this compound of sodium is more stable than the corresponding chlorine amide. When heated, sodamide, NH2Na, only partially decomposes, with evolution of hydrogen, the principal part of it giving ammonia and sodium nitride, Na3N, according to the equation 3NH2Na = 2NH3+ NNa3. The latter is an almost black powdery mass, decomposed by water into ammonia and sodium hydroxide.Titherley's researches added the following data:—Iron or silver vessels should be used in preparing this body, because glass and porcelain are corroded at 300°–400°, at which temperature ammonia gas acts upon sodium and forms the amide with the evolution of hydrogen. The reaction proceeds slowly, but is complete if there be an excess of NH3. Pure NH2Na is colourless (its colouration is due to various impurities), semi-transparent, shows traces of crystallisation, has a conchoidal fracture, and melts at 145°. Judging from the increase in weight of the sodium and the quantity of hydrogen which is disengaged, the composition of the amide is exactly NH2Na. It partially volatilises (sublimes) in vacuo at 200°, and breaks up into 2Na + N2+ 2H2at 500°. The same amide is formed when oxide of sodium is heated in NH3: Na2O + 2NH3= 2NaH2N + H2O. NaHO is also formed to some extent by the resultant H2O. Potassium and lithium form similar amides. With water, alcohol, and acids, NH2Na gives NH3and NaHO, which react further. Anhydrous CaO absorbs NH2Na when heated without decomposing it. When sodamide is heated with SiO2, NH3is disengaged, and silicon nitride formed. It acts still more readily upon boric anhydride when heated with it: 2NH2Na + B2O3= 2BN + 2NaHO + H2O. When slightly heated, NH2Na + NOCl = NaCl + N2+ H2O (NHNa2and NNa3are apparently not formed at a higher temperature). The halogen organic compounds react with the aid of heat, but with so much energy that the reaction frequently leads to the ultimate destruction of the organic groups and production of carbon.
[44 bis]Sodamide, NH2Na, (Chapter IV., Note14), discovered by Gay-Lussac and Thénard, has formed the object of repeated research, but has been most fully investigated by A. W. Titherley (1894). Until recently the following was all that was known about this compound:—
By heating sodium in dry ammonia, Gay-Lussac and Thénard obtained an olive-green, easily-fusible mass,sodamide, NH2Na, hydrogen being separated. This substance with water forms sodium hydroxide and ammonia; with carbonic oxide, CO, it forms sodium cyanide, NaCN, and water, H2O; and with dry hydrogen chloride it forms sodium and ammonium chlorides. These and other reactions of sodamide show that the metal in it preserves its energetic properties in reaction, and that this compound of sodium is more stable than the corresponding chlorine amide. When heated, sodamide, NH2Na, only partially decomposes, with evolution of hydrogen, the principal part of it giving ammonia and sodium nitride, Na3N, according to the equation 3NH2Na = 2NH3+ NNa3. The latter is an almost black powdery mass, decomposed by water into ammonia and sodium hydroxide.
Titherley's researches added the following data:—
Iron or silver vessels should be used in preparing this body, because glass and porcelain are corroded at 300°–400°, at which temperature ammonia gas acts upon sodium and forms the amide with the evolution of hydrogen. The reaction proceeds slowly, but is complete if there be an excess of NH3. Pure NH2Na is colourless (its colouration is due to various impurities), semi-transparent, shows traces of crystallisation, has a conchoidal fracture, and melts at 145°. Judging from the increase in weight of the sodium and the quantity of hydrogen which is disengaged, the composition of the amide is exactly NH2Na. It partially volatilises (sublimes) in vacuo at 200°, and breaks up into 2Na + N2+ 2H2at 500°. The same amide is formed when oxide of sodium is heated in NH3: Na2O + 2NH3= 2NaH2N + H2O. NaHO is also formed to some extent by the resultant H2O. Potassium and lithium form similar amides. With water, alcohol, and acids, NH2Na gives NH3and NaHO, which react further. Anhydrous CaO absorbs NH2Na when heated without decomposing it. When sodamide is heated with SiO2, NH3is disengaged, and silicon nitride formed. It acts still more readily upon boric anhydride when heated with it: 2NH2Na + B2O3= 2BN + 2NaHO + H2O. When slightly heated, NH2Na + NOCl = NaCl + N2+ H2O (NHNa2and NNa3are apparently not formed at a higher temperature). The halogen organic compounds react with the aid of heat, but with so much energy that the reaction frequently leads to the ultimate destruction of the organic groups and production of carbon.
[45]As sodium does not displace hydrogen from the hydrocarbons,it may be preservedin liquid hydrocarbons. Naphtha is generally used for this purpose, as it consists of a mixture of various liquid hydrocarbons. However, in naphtha sodium usually becomes coated with a crust composed of matter produced by the action of the sodium on certain of the substances contained in the mixture composing naphtha. In order that sodium may retain its lustre in naphtha, secondary octyl alcohol is added. (This alcohol is obtained by distilling castor oil with caustic potash.) Sodium keeps well in a mixture of pure benzene and paraffin.
[45]As sodium does not displace hydrogen from the hydrocarbons,it may be preservedin liquid hydrocarbons. Naphtha is generally used for this purpose, as it consists of a mixture of various liquid hydrocarbons. However, in naphtha sodium usually becomes coated with a crust composed of matter produced by the action of the sodium on certain of the substances contained in the mixture composing naphtha. In order that sodium may retain its lustre in naphtha, secondary octyl alcohol is added. (This alcohol is obtained by distilling castor oil with caustic potash.) Sodium keeps well in a mixture of pure benzene and paraffin.
[46]If sodium does not directly displace the hydrogen in hydrocarbons, still by indirect means compounds may be obtained which contain sodium and hydrocarbon groups. Some of these compounds have been produced, although not in a pure state. Thus, for instance, zinc ethyl, Zn(C2H5)2, when treated with sodium, loses zinc and forms sodium ethyl, C2H5Na, but this decomposition is not complete, and the compound formed cannot be separated by distillation from the remaining zinc ethyl. In this compound the energy of the sodium is clearly manifest, for it reacts with substances containing haloids, oxygen, &c., and directly absorbs carbonic anhydride, forming a salt of a carboxylic acid (propionic).
[46]If sodium does not directly displace the hydrogen in hydrocarbons, still by indirect means compounds may be obtained which contain sodium and hydrocarbon groups. Some of these compounds have been produced, although not in a pure state. Thus, for instance, zinc ethyl, Zn(C2H5)2, when treated with sodium, loses zinc and forms sodium ethyl, C2H5Na, but this decomposition is not complete, and the compound formed cannot be separated by distillation from the remaining zinc ethyl. In this compound the energy of the sodium is clearly manifest, for it reacts with substances containing haloids, oxygen, &c., and directly absorbs carbonic anhydride, forming a salt of a carboxylic acid (propionic).
[46 bis]It is even doubtful whether the suboxide exists (seeNote47).
[46 bis]It is even doubtful whether the suboxide exists (seeNote47).
[47]A compound, Na2Cl, which corresponds with the suboxide, is apparently formed when a galvanic current is passed through fused common salt; the sodium liberated dissolves in the common salt, and does not separate from the compound either on cooling or on treatment with mercury. It is therefore supposed to be Na2Cl; the more so as the mass obtained gives hydrogen when treated with water: Na2Cl + H2O = H + NaHO + NaCl, that is, it acts like suboxide of sodium. If Na2Cl really exists as a salt, then the corresponding base Na4O, according to the rule with other bases of the composition M4O, ought to be called a quaternary oxide. According to certain evidence, a suboxide is formed when thin sheets or fine drops of sodium slowly oxidise in moist air.
[47]A compound, Na2Cl, which corresponds with the suboxide, is apparently formed when a galvanic current is passed through fused common salt; the sodium liberated dissolves in the common salt, and does not separate from the compound either on cooling or on treatment with mercury. It is therefore supposed to be Na2Cl; the more so as the mass obtained gives hydrogen when treated with water: Na2Cl + H2O = H + NaHO + NaCl, that is, it acts like suboxide of sodium. If Na2Cl really exists as a salt, then the corresponding base Na4O, according to the rule with other bases of the composition M4O, ought to be called a quaternary oxide. According to certain evidence, a suboxide is formed when thin sheets or fine drops of sodium slowly oxidise in moist air.
[48]According to observations easily made, sodium when fused in air oxidises but does not burn, the combustion only commencing with the formation of vapour—that is, when considerably heated. Davy and Karsten obtained the oxides of potassium, K2O, and of sodium, Na2O, by heating the metals with their hydroxides, whence NaHO + Na = Na2O + H, but N. N. Beketoff failed to obtain oxides by this means. He prepared them by directly igniting the metals in dry air, and afterwards heating with the metal in order to destroy any peroxide. The oxide produced, Na2O, when heated in an atmosphere of hydrogen, gave a mixture of sodium and its hydroxide: Na2O + H = NaHO + Na (seeChapter II., Note9). If both the observations mentioned are accurate, then the reaction is reversible. Sodium oxide ought to be formed during the decomposition of sodium carbonate by oxide of iron (seeNote26), and during the decomposition of sodium nitrite. According to Karsten, its specific gravity is 2·8, according to Beketoff 2·3. The difficulty in obtaining it is owing to an excess of sodium forming the suboxide, and an excess of oxygen the peroxide. The grey colour peculiar to the suboxide and oxide perhaps shows that they contain metallic sodium. In addition to this, in the presence of water it may contain sodium hydride and NaHO.
[48]According to observations easily made, sodium when fused in air oxidises but does not burn, the combustion only commencing with the formation of vapour—that is, when considerably heated. Davy and Karsten obtained the oxides of potassium, K2O, and of sodium, Na2O, by heating the metals with their hydroxides, whence NaHO + Na = Na2O + H, but N. N. Beketoff failed to obtain oxides by this means. He prepared them by directly igniting the metals in dry air, and afterwards heating with the metal in order to destroy any peroxide. The oxide produced, Na2O, when heated in an atmosphere of hydrogen, gave a mixture of sodium and its hydroxide: Na2O + H = NaHO + Na (seeChapter II., Note9). If both the observations mentioned are accurate, then the reaction is reversible. Sodium oxide ought to be formed during the decomposition of sodium carbonate by oxide of iron (seeNote26), and during the decomposition of sodium nitrite. According to Karsten, its specific gravity is 2·8, according to Beketoff 2·3. The difficulty in obtaining it is owing to an excess of sodium forming the suboxide, and an excess of oxygen the peroxide. The grey colour peculiar to the suboxide and oxide perhaps shows that they contain metallic sodium. In addition to this, in the presence of water it may contain sodium hydride and NaHO.
[49]Of the oxides of sodium, that easiest to form is the peroxide, NaO or Na2O2; this is obtained when sodium is burnt in an excess of oxygen. If NaNO3be melted, it gives Na2O2with metallic Na. In a fused state the peroxide is reddish yellow, but it becomes almost colourless when cold. When heated with iodine vapour, it loses oxygen: Na2O2+I2= Na2OI2+ O. The compound Na2OI2is akin to the compound Cu2OCl2obtained by oxidising CuCl. This reaction is one of the few in which iodine directly displaces oxygen. The substance Na2OI2is soluble in water, and when acidified gives free iodine and a sodium salt. Carbonic oxide is absorbed by heated sodium peroxide with formation of sodium carbonate: Na2CO3= Na2O2+ CO, whilst carbonic anhydride liberates oxygen from it. With nitrous oxide it reacts thus: Na2O2+2N2O = 2NaNO2+N2; with nitric oxide it combines directly, forming sodium nitrite, NaO + NO = NaNO2. Sodium peroxide, when treated with water, does not give hydrogen peroxide, because the latter in the presence of the alkali formed (Na2O2+ 2H2O = 2NaHO + H2O2) decomposes into water and oxygen. In the presence of dilute sulphuric acid it forms H2O2(Na2O2+ H2SO4= Na2SO4+ H2O2). Peroxide of sodium is now prepared on a large scale (by the action of air upon Na at 300°) for bleaching wool, silk &c. (when it acts in virtue of the H2O2formed). The oxidising properties of Na2O2under the action of heat are seen, for instance, in the fact that when heated with I it forms sodium iodate; with PbO, Na2PbO3; with pyrites, sulphates, &c. When peroxide of sodium comes into contact with water, it evolves much heat, forming H2O2, and decomposing with the disengagement of oxygen; but, as a rule, there is no explosion. But if Na2O2be placed in contact with organic matter, such as sawdust, cotton, &c., it gives a violent explosion when heated, ignited, or acted on by water. Peroxide of sodium forms an excellent oxidising agent for the preparation of the higher product of oxidation of Mn, Cr, W, &c., and also for oxidising the metallic sulphides. It should therefore find many applications in chemical analysis. To prepare Na2O2on a large scale, Castner melts Na in an aluminium vessel, and at 300° passes first air deprived of a portion of its oxygen (having been already once used), and then ordinary dry air over it.
[49]Of the oxides of sodium, that easiest to form is the peroxide, NaO or Na2O2; this is obtained when sodium is burnt in an excess of oxygen. If NaNO3be melted, it gives Na2O2with metallic Na. In a fused state the peroxide is reddish yellow, but it becomes almost colourless when cold. When heated with iodine vapour, it loses oxygen: Na2O2+I2= Na2OI2+ O. The compound Na2OI2is akin to the compound Cu2OCl2obtained by oxidising CuCl. This reaction is one of the few in which iodine directly displaces oxygen. The substance Na2OI2is soluble in water, and when acidified gives free iodine and a sodium salt. Carbonic oxide is absorbed by heated sodium peroxide with formation of sodium carbonate: Na2CO3= Na2O2+ CO, whilst carbonic anhydride liberates oxygen from it. With nitrous oxide it reacts thus: Na2O2+2N2O = 2NaNO2+N2; with nitric oxide it combines directly, forming sodium nitrite, NaO + NO = NaNO2. Sodium peroxide, when treated with water, does not give hydrogen peroxide, because the latter in the presence of the alkali formed (Na2O2+ 2H2O = 2NaHO + H2O2) decomposes into water and oxygen. In the presence of dilute sulphuric acid it forms H2O2(Na2O2+ H2SO4= Na2SO4+ H2O2). Peroxide of sodium is now prepared on a large scale (by the action of air upon Na at 300°) for bleaching wool, silk &c. (when it acts in virtue of the H2O2formed). The oxidising properties of Na2O2under the action of heat are seen, for instance, in the fact that when heated with I it forms sodium iodate; with PbO, Na2PbO3; with pyrites, sulphates, &c. When peroxide of sodium comes into contact with water, it evolves much heat, forming H2O2, and decomposing with the disengagement of oxygen; but, as a rule, there is no explosion. But if Na2O2be placed in contact with organic matter, such as sawdust, cotton, &c., it gives a violent explosion when heated, ignited, or acted on by water. Peroxide of sodium forms an excellent oxidising agent for the preparation of the higher product of oxidation of Mn, Cr, W, &c., and also for oxidising the metallic sulphides. It should therefore find many applications in chemical analysis. To prepare Na2O2on a large scale, Castner melts Na in an aluminium vessel, and at 300° passes first air deprived of a portion of its oxygen (having been already once used), and then ordinary dry air over it.