The Principles of Chemistry, Volume I

CHAPTER XII SODIUM

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

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

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

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

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

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

The reactions of decomposition of sodium sulphate are above all noticeable by the separation of oxygen. Sodium sulphate by itself is very stable, and it is only at a temperature sufficient to melt iron that it is possible to separate the elements SO₃ from it, and then only partially. However, the oxygen may be separated from sodium sulphate, as from all other sulphates, by means of many substances which are able to combine with oxygen, such as charcoal and sulphur, but hydrogen is not able to produce this action. If sodium sulphate be heated with charcoal, then carbonic oxide and anhydride are evolved, and there is produced, according to the circumstances, either the lower oxygen compound, sodium sulphite, Na₂SO₃ (for instance, in the formation of glass); or else the decomposition proceeds further, and sodium sulphide, Na₂S, is formed, according to the equation Na₂SO₄ + 2C = 2CO₂ + Na₂S.

On the basis of this reaction the greater part of the sulphate of sodium prepared at chemical works is converted into soda ash—that is, sodium carbonate, Na₂CO₃, which is used for many purposes. In the form of carbonates, the metallic oxides behave in many cases just as they do in the state of oxides or hydroxides, owing to the feeble acid properties of carbonic acid. However, the majority of the salts of carbonic acid are insoluble, whilst sodium carbonate is one of the few soluble salts of this acid, and therefore reacts with facility. Hence sodium carbonate is employed for many purposes, in which its alkaline properties come into play. Thus, even under the action of feeble organic acids it immediately parts with its carbonic acid, and gives a sodium salt of the acid taken. Its solutions exhibit an alkaline reaction on litmus. It aids the passage of certain organic substances (tar, acids) into solution, and is therefore used, like caustic alkalis and soap (which latter also acts by virtue of the alkali it contains), for the removal of certain organic substances, especially in bleaching cotton and similar fabrics. Besides which a considerable quantity of sodium carbonate is used for the preparation of sodium hydroxide or caustic soda, which has also a very wide application. In large chemical works where sodium carbonate is manufactured from Na₂SO₄, it is usual first to manufacture sulphuric acid, and then by its aid to convert common salt into sodium sulphate, and lastly to convert the sodium sulphate thus obtained into carbonate and caustic soda. Hence these works prepare both alkaline substances (soda ash and caustic soda) and acid substances (sulphuric and hydrochloric acids), the two classes of chemical products which are distinguished for the greatest energy of their reactions and are therefore most frequently applied to technical purposes. Factories manufacturing soda are generally called alkali works.

The process of the conversion of sodium sulphate into sodium carbonate consists in strongly heating a mixture of the sulphate with charcoal and calcium carbonate. The following reactions then take place: the sodium sulphate is first deoxidised by the charcoal, forming sodium sulphide and carbonic anhydride, Na₂SO₄ + 2C = Na₂S + 2CO₂. The sodium sulphide thus formed then enters into double decomposition with the calcium carbonate taken, and gives calcium sulphide and sodium carbonate, Na₂S + CaCO₃ = Na₂CO₃ + CaS.

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Fig. 68.—Reverberatory furnace for the manufacture of sodium carbonate. F, grate. A, bridge. M, hearth for the ultimate calcination of the mixture of sodium sulphate, coal, and calcium carbonate, which is charged from above into the part of the furnace furthest removed from the fire F. P, P, doors for stirring and bringing the mass towards the grate F by means of stirrers R. At the end of the operation the semifused mass is charged into trucks C.

Besides which, under the action of the heat, a portion of the excess of calcium carbonate is decomposed into lime and carbonic anhydride, CaCO₃ = CaO + CO₂, and the carbonic anhydride with the excess of charcoal forms carbon monoxide, which towards the end of the operation shows itself by the appearance of a blue flame. Thus from a mass containing sodium sulphate we obtain a mass which includes sodium carbonate, calcium sulphide, and calcium oxide, but none of the sodium sulphide which was formed on first heating the mixture. The entire process, which proceeds at a high temperature, may be expressed by a combination of the three above-mentioned formulÆ, if it be considered that the product contains one equivalent of calcium oxide to two equivalents of calcium sulphide.[12] The sum of the reactions may then be expressed thus: 2Na₂SO₄ + 3CaCO₃ + 9C = 2Na₂CO₃ + CaO,2CaS + 10CO. Indeed, the quantities in which the substances are mixed together at chemical works approaches to the proportion required by this equation. The entire process of decomposition is carried on in reverberatory furnaces, into which a mixture of 1,000 parts of sodium sulphate, 1,040 parts of calcium carbonate (as a somewhat porous limestone), and 500 parts of small coal is introduced from above. This mixture is first heated in the portion of the furnace which is furthest removed from the fire-grate; it is then brought to the portion nearest to the fire-grate, when it is stirred during heating. The partially fused mass obtained at the end of the process is cooled, and then subjected to methodical lixiviation[13] to extract the sodium carbonate, the mixture of calcium oxide and sulphide forming the so-called ‘soda waste’ or ‘alkali waste.’[14]

The above-mentioned process for making soda was discovered in the year 1808 by the French doctor Leblanc, and is known as the Leblanc process. The particulars of the discovery are somewhat remarkable. Sodium carbonate, having a considerable application in industry, was for a long time prepared exclusively from the ash of marine plants (Chapter XI., page 497). Even up to the present time this process is carried on in Normandy. In France, where for a long time the manufacture of large quantities of soap (so-called Marseilles soap) and various fabrics required a large amount of soda, the quantity prepared at the coast was insufficient to meet the demand. For this reason during the wars at the beginning of the century, when the import of foreign goods into France was interdicted, the want of sodium carbonate was felt. The French Academy offered a prize for the discovery of a profitable method of preparing it from common salt. Leblanc then proposed the above-mentioned process, which is remarkable for its great simplicity.[15]

Of all other industrial processes for manufacturing sodium carbonate, the ammonia process is the most worthy of mention.[16] In this the vapours of ammonia, and then an excess of carbonic anhydride, are directly introduced into a concentrated solution of sodium chloride in order to form the acid ammonium carbonate, NH₄HCO₃. Then, by means of the double saline decomposition of this salt, sodium chloride is decomposed, and in virtue of its slight solubility acid sodium carbonate, NaHCO₃, is precipitated and ammonium chloride, NH₄Cl, is obtained in solution (with a portion of the sodium chloride and acid sodium carbonate). The reaction proceeds in the solution owing to the sparing solubility of the NaHCO₃ according to the equation NaCl + NH₄HCO₃ = NH₄Cl + NaHCO₃. The ammonia is recovered from the solution by heating with lime or magnesia,[16 bis] and the precipitated acid sodium carbonate is converted into the normal salt by heating. It is thus obtained in a very pure state.[17]

Sodium carbonate, like sodium sulphate, loses all its water on being heated, and when anhydrous fuses at a bright-red heat (1098°). A small quantity of sodium carbonate placed in the loop of a platinum wire volatilises in the heat of a gas flame, and therefore in the furnaces of glass works part of the soda is always transformed into the condition of vapour. Sodium carbonate resembles sodium sulphate in its relation to water.[18] Here also the greatest solubility is at the temperature of 37°; both salts, on crystallising at the ordinary temperature, combine with ten molecules of water, and such crystals of soda, like crystals of Glauber's salt, fuse at 34°. Sodium carbonate also forms a supersaturated solution, and, according to the conditions, gives various combinations with water of crystallisation (mentioned on page 108), &c.

At a red heat superheated steam liberates carbonic anhydride from sodium carbonate and forms caustic soda, Na₂CO₃ + H₂O = 2NaHO + CO₂. Here the carbonic anhydride is replaced by water; this depends on the feebly acid character of carbonic anhydride. By direct heating, sodium carbonate is only slightly decomposed into sodium oxide and carbonic anhydride; thus, when sodium carbonate is fused, about 1 per cent. of carbonic anhydride is disengaged.[19] The carbonates of many other metals—for instance, of calcium, copper, magnesium, iron, &c.—on being heated lose all their carbonic anhydride. This shows the considerable basic energy which sodium possesses. With the soluble salts of most metals, sodium carbonate gives precipitates either of insoluble carbonates of the metals, or else of the hydroxides (in this latter case carbonic anhydride is disengaged); for instance, with barium salts it precipitates an insoluble barium carbonate (BaCl₂ + Na₂CO₃ = 2NaCl + BaCO₃) and with the aluminium salts it precipitates aluminium hydroxide, carbonic anhydride being disengaged: 3Na₂CO₃ + Al₂(SO₄)₃ + 3H₂O = 3Na₂SO₄ + 2Al(OH)₃ + 3CO₂. Sodium carbonate, like all the salts of carbonic acid, evolves carbonic anhydride on treatment with all acids which are to any extent energetic. But if an acid diluted with water be gradually added to a solution of sodium carbonate, at first such an evolution does not take place, because the excess of the carbonic anhydride forms acid sodium carbonate (sodium bicarbonate), NaHCO₃.[20] The acid sodium carbonate is an unstable salt. Not only when heated alone, but even on being slightly heated in solution, and also at the ordinary temperature in damp air, it loses carbonic anhydride and forms the normal salt. And at the same time it is easy to obtain it in a pure crystalline form, if a strong solution of sodium carbonate be cooled and a stream of carbonic anhydride gas passed through it. The acid salt is less soluble in water than the normal,[21] and therefore a strong solution of the latter gives crystals of the acid salt if carbonic anhydride be passed through it. The acid salt may be yet more conveniently formed from effloresced crystals of sodium carbonate, which, on being considerably heated, very easily absorb carbonic anhydride.[22] The acid salt crystallises well, but not, however, in such large crystals as the normal salt; it has a brackish and not an alkaline taste like that of the normal salt; its reaction is feebly alkaline, nearly neutral. At 70° its solution begins to lose carbonic anhydride, and on boiling the evolution becomes very abundant. From the preceding remarks it is clear that in most reactions this salt, especially when heated, acts similarly to the normal salt, but has, naturally, some distinction from it. Thus, for example, if a solution of sodium carbonate be added to a normal magnesium salt, a turbidity (precipitate) is formed of magnesium carbonate. MgCO₃. No such precipitate is formed by the acid salt, because magnesium carbonate is soluble in the presence of an excess of carbonic anhydride.

Sodium carbonate is used for the preparation of caustic soda[23]—that is, the hydrate of sodium oxide, or the alkali which corresponds to sodium. For this purpose the action of lime on a solution of sodium carbonate is generally made use of. The process is as follows: a weak, generally 10 per cent., solution of sodium carbonate is taken,[24] and boiled in a cast-iron, wrought-iron, or silver boiler (sodium hydroxide does not act on these metals), and lime is added, little by little, during the boiling. This latter is soluble in water, although but very slightly. The clear solution becomes turbid on the addition of the lime because a precipitate is formed; this precipitate consists of calcium carbonate, almost insoluble in water, whilst caustic soda is formed and remains in solution. The decomposition is effected according to the equation: Na₂CO₃ + Ca(HO)₂ = CaCO₃ + 2NaHO. On cooling the solution the calcium carbonate easily settles as a precipitate, and the clear solution or alkali above it contains the easily soluble sodium hydroxide formed in the reaction.[25] After the necessary quantity of lime has been added, the solution is allowed to stand, and is then decanted off and evaporated in cast or wrought iron boilers, or in silver pans if a perfectly pure product is required.[26] The evaporation cannot be conducted in china, glass, or similar vessels, because caustic soda attacks these materials, although but slightly. The solution does not crystallise on evaporation, because the solubility of caustic soda when hot is very great, but crystals containing water of crystallisation may be obtained by cooling. If the evaporation of the alkali be conducted until the specific gravity reaches 1·38, and the liquid is then cooled to 0°, transparent crystals appear containing 2NaHO,7H₂O; they fuse at +6°.[27] If the evaporation be conducted so long as water is disengaged, which requires a considerable amount of heat, then, on cooling, the hydroxide, NaHO, solidifies in a semi-transparent crystalline mass,[28] which eagerly absorbs moisture and carbonic anhydride from the air.[29] Its specific gravity is 2·13;[30] it is easily soluble in water, with disengagement of a considerable quantity of heat.[31] A saturated solution at the ordinary temperature has a specific gravity of about 1·5, contains about 45 per cent. of sodium hydroxide, and boils at 130°; at 55° water dissolves an equal weight of it.[32] Caustic soda is not only soluble in water but in alcohol, and even in ether. Dilute solutions of sodium hydroxide produce a soapy feeling on the skin because the active base of soap consists of caustic soda.[33] Strong solutions have a corroding action.

The chemical reactions of sodium hydroxide serve as a type for those of a whole class of alkalis—that is, of soluble basic hydroxides, MOH. The solution of sodium hydroxide is a very caustic liquid—that is to say, it acts in a destructive way on most substances, for instance on most organic tissues—hence caustic soda, like all soluble alkalis, is a poisonous substance; acids, for example hydrochloric, serve as antidotes. The action of caustic soda on bones, fat, starch, and similar vegetable and animal substances explains its action on organisms. Thus bones, when plunged into a weak solution of caustic soda, fall to powder,[34] and evolve a smell of ammonia, owing to the caustic soda changing the gelatinous organic substance of the bones (which contains carbon, hydrogen, nitrogen, oxygen, and sulphur, like albumin), dissolving it and in part destroying it, whence ammonia is disengaged. Fats, tallow, and oils become saponified by a solution of caustic soda—that is to say, they form with it soaps soluble in water, or sodium salts of the organic acids contained in the fats.[35] The most characteristic reactions of sodium hydroxide are determined by the fact that it saturates all acids, forming salts with them, which are almost all soluble in water, and in this respect caustic soda is as characteristic amongst the bases as nitric acid is among the acids. It is impossible to detect sodium by means of the formation of precipitates of insoluble sodium salts, as may be done with other metals, many of whose salts are but slightly soluble. The powerful alkaline properties of caustic soda determine its capacity for combining with even the feeblest acids, its property of disengaging ammonia from ammonium salts, its faculty of forming precipitates from solutions of salts whose bases are insoluble in water, &c. If a solution of the salt of almost any metal be mixed with caustic soda, then a soluble sodium salt will be formed, and an insoluble hydroxide of the metal will be separated—for instance, copper nitrate yields copper hydroxide, Cu(NO₃)₂ + 2NaHO = Cu(HO)₂ + 2NaNO₃. Even many basic oxides precipitated by caustic soda are capable of combining with it and forming soluble compounds, and therefore caustic soda in the presence of salts of such metals first forms a precipitate of hydroxide, and then, employed in excess, dissolves this precipitate. This phenomenon occurs, for example, when caustic soda is added to the salts of aluminium. This shows the property of such an alkali as caustic soda of combining not only with acids, but also with feeble basic oxides. For this reason caustic soda acts on most elements which are capable of forming acids or oxides similar to them; thus the metal aluminium gives hydrogen with caustic soda in consequence of the formation of alumina, which combines with the caustic soda—that is, in this case, the caustic alkali acts on the metal just as sulphuric acid does on Fe or Zn. If caustic soda acts in this manner on a metalloid capable of combining with the hydrogen evolved (aluminium does not give a compound with hydrogen), then it forms such a hydrogen compound. Thus, for instance, phosphorus acts in this way on caustic soda, yielding hydrogen phosphide. When the hydrogen compound disengaged is capable of combining with the alkali, then, naturally, a salt of the corresponding acid is formed. For example, chlorine and sulphur act in this way on caustic soda. Chlorine, with the hydrogen of the caustic soda, forms hydrochloric acid, and the latter forms common salt with the sodium hydroxide, whilst the other atom in the molecule of chlorine, Cl₂, takes the place of the hydrogen, and forms the hypochlorite, NaClO. In the same way, by the action of sodium hydroxide on sulphur, hydrogen sulphide is formed, which acts on the soda forming sodium sulphide, in addition to which sodium thiosulphate is formed (see Chapter XX.) By virtue of such reactions, sodium hydroxide acts on many metals and non-metals. Such action is often accelerated by the presence of the oxygen of the air, as by this means the formation of acids and oxides rich in oxygen is facilitated. Thus many metals and their lower oxides, in the presence of an alkali, absorb oxygen and form acids. Even manganese peroxide, when mixed with caustic soda, is capable of absorbing the oxygen of the air, and forming sodium manganate. Organic acids when heated with caustic soda give up to it the elements of carbonic anhydride, forming sodium carbonate, and separating that hydrocarbon group which exists, in combination with carbonic anhydride, in the organic acid.

Thus sodium hydroxide, like the soluble alkalis in general, ranks amongst the most active substances in the chemical sense of the term, and but few substances are capable of resisting it. Even siliceous rocks, as we shall see further on, are transformed by it, forming when fused with it vitreous slags. Sodium hydroxide (like ammonium and potassium hydroxides), as a typical example of the basic hydrates, in distinction from many other basic oxides, easily forms acid salts with acids (for instance, NaHSO₄, NaHCO₃), and does not form any basic salts at all; whilst many less energetic bases, such as the oxides of copper and lead, easily form basic salts, but acid salts only with difficulty. This capability of forming acid salts, particularly with polybasic acids, may be explained by the energetic basic properties of sodium hydroxide, contrasted with the small development of these properties in the bases which easily form basic salts. An energetic base is capable of retaining a considerable quantity of acid, which a slightly energetic base would not have the power of doing. Also, as will be shown in the subsequent chapters, sodium belongs to the univalent metals, being exchangeable for hydrogen atom for atom—that is, amongst metals sodium may, like chlorine amongst the non-metals, serve as the representative of the univalent properties. Most of the elements which are not capable of forming acid salts are bivalent. Whence it may be understood that in a bibasic acid—for instance, carbonic, H₂CO₃, or sulphuric, H₂SO₄—the hydrogen may be exchanged, atom for atom, for sodium, and yield an acid salt by means of the first substitution, and a normal salt by means of the second—for instance, NaHSO₄, and Na₂SO₄, whilst such bivalent metals as calcium and barium do not form acid salts because one of their atoms at once takes the place of both hydrogen atoms, forming, for example, CaCO₃ and CaSO₄.[35 bis]

We have seen the transformation of common salt into sodium sulphate, of this latter into sodium carbonate, and of sodium carbonate into caustic soda. Lavoisier still regarded sodium hydroxide as an element, because he was unacquainted with its decomposition with the formation of metallic sodium, which separates the hydrogen from water, reforming caustic soda.

The preparation of metallic sodium was one of the greatest discoveries in chemistry, not only because through it the conception of elements became broader and more correct, but especially because in sodium, chemical properties were observed which were but feebly shown in the other metals more familiarly known. This discovery was made in 1807 by the English chemist Davy by means of the galvanic current. By connecting with the positive pole (of copper or carbon) a piece of caustic soda (moistened in order to obtain electrical conductivity), and boring a hole in it filled with mercury connected with the negative pole of a strong Volta's pile, Davy observed that on passing the current a peculiar metal dissolved in the mercury, less volatile than mercury, and capable of decomposing water, again forming caustic soda. In this way (by analysis and synthesis) Davy demonstrated the compound nature of alkalis. On being decomposed by the galvanic current, caustic soda disengages hydrogen and sodium at the negative pole and oxygen at the positive pole. Davy showed that the metal formed volatilises at a red heat, and this is its most important physical property in relation to its extraction, all later methods being founded on it. Besides this Davy observed that sodium easily oxidises, its vapour taking fire in air, and the latter circumstance was for a long time an obstacle to the easy preparation of this metal. The properties of sodium were subsequently more thoroughly investigated by Gay-Lussac and ThÉnard, who observed that metallic iron at a high temperature was capable of reducing caustic soda to sodium.[36] Brunner latterly discovered that not only iron, but also charcoal, has this property, although hydrogen has not.[37] But still the methods of extracting sodium were very troublesome, and consequently it was a great rarity. The principal obstacle to its production was that an endeavour was made to condense the easily-oxidising vapours of sodium in vacuo in complicated apparatus. For this reason, when Donny and Maresca, having thoroughly studied the matter, constructed a specially simple condenser, the production of sodium was much facilitated. Furthermore, in practice the most important epoch in the history of the production of sodium is comprised in the investigation of Sainte-Claire Deville, who avoided the complex methods in vogue up to that time, and furnished those simple means by which the production of sodium is now rendered feasible in chemical works.

For the production of sodium according to Deville's method, a mixture of anhydrous sodium carbonate (7 parts), charcoal (two parts), and lime or chalk (7 parts) is heated. This latter ingredient is only added in order that the sodium carbonate, on fusing, shall not separate from the charcoal.[38] The chalk on being heated loses carbonic anhydride, leaving infusible lime, which is permeated by the sodium carbonate and forms a thick mass, in which the charcoal is intimately mixed with the sodium carbonate. When the charcoal is heated with the sodium carbonate, at a white heat, carbonic oxide and vapours of sodium are disengaged, according to the equation:

Na₂CO₃ + 2C = Na₂ + 3CO

On cooling the vapours and gases disengaged, the vapours condense into molten metal (in this form sodium does not easily oxidise, whilst in vapour it burns) and the carbonic oxide remains as gas.

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Fig. 70.—Manufacture of sodium by Deville's process. A C, iron tube containing a mixture of soda, charcoal, and chalk. B, condenser.

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Fig. 71.—Donny and Maresca's sodium condenser, consisting of two cast-iron plates screwed together.

In sodium works an iron tube, about a metre long and a decimeter in diameter, is made out of boiler plate. The pipe is luted into a furnace having a strong draught, capable of giving a high temperature, and the tube is charged with the mixture required for the preparation of sodium. One end of the tube is closed with a cast-iron stopper A with clay luting, and the other with the cast-iron stopper C provided with an aperture. On heating, first of all the moisture contained in the various substances is given off, then carbonic anhydride and the products of the dry distillation of the charcoal, then the latter begins to act on the sodium carbonate, and carbonic oxide and vapours of sodium appear. It is easy to observe the appearance of the latter, because on issuing from the aperture in the stopper C they take fire spontaneously and burn with a very bright yellow flame. A pipe is then introduced into the aperture C, compelling the vapours and gases formed to pass through the condenser B. This condenser consists of two square cast-iron trays, A and A', fig. 71, with wide edges firmly screwed together. Between these two trays there is a space in which the condensation of the vapours of sodium is effected, the thin metallic walls of the condenser being cooled by the air but remaining hot enough to preserve the sodium in a liquid state, so that it does not choke the apparatus, but continually flows from it. The vapours of sodium, condensing in the cooler, flow in the shape of liquid metal into a vessel containing some non-volatile naphtha or hydrocarbon. This is used in order to prevent the sodium oxidising as it issues from the condenser at a somewhat high temperature. In order to obtain sodium of a pure quality it is necessary to distil it once more, which may even be done in porcelain retorts, but the distillation must be conducted in a stream of some gas on which sodium does not act, for instance in a stream of nitrogen; carbonic anhydride is not applicable, because sodium partially decomposes it, absorbing oxygen from it. Although the above described methods of preparing sodium by chemical means have proved very convenient in practice, still it is now (since 1893) found profitable in England to obtain it (to the amount of several tons a week) by Davy's classical method, i.e. by the action of an electric current at a moderately high temperature, because the means for producing an electric current (by motors and dynamos) now render this quite feasible. This may be regarded as a sign that in process of time many other technical methods for producing various substances by decomposition may be profitably carried on by electrolysis.

Pure sodium is a lustrous metal, white as silver, soft as wax; it becomes brittle in the cold. In ordinary moist air it quickly tarnishes and becomes covered with a film of hydroxide, NaHO, formed at the expense of the water in the air. In perfectly dry air sodium retains its lustre for an indefinite time. Its density at the ordinary temperature is equal to 0·98, so that it is lighter than water; it fuses very easily at a temperature of 95°, and distils at a bright red heat (742° according to Perman, 1889). Scott (1887) determined the density of sodium vapour and found it to be nearly 12 (if H = 1). This shows that its molecule contains one atom (like mercury and cadmium) Na.[38 bis] It forms alloys with most metals, combining with them, heat being sometimes evolved and sometimes absorbed. Thus, if sodium (having a clean surface) be thrown into mercury, especially when heated, there is a flash, and such a considerable amount of heat is evolved that part of the mercury is transformed into vapour.[39] Compounds or solutions of sodium in mercury, or amalgams of sodium, even when containing 2 parts of sodium to 100 parts of mercury, are solids. Only those amalgams which are the very poorest in sodium are liquid. Such alloys of sodium with mercury are often used instead of sodium in chemical investigations, because in combination with mercury sodium is not easily acted on by air, and is heavier than water, and therefore more convenient to handle, whilst at the same time it retains the principal properties of sodium,[40] for instance it decomposes water, forming NaHO.

It is easy to form an alloy of mercury and sodium having a crystalline structure, and a definite atomic composition, NaHg₅. The alloy of sodium with hydrogen or sodium hydride, Na₂H, which has the external appearance of a metal,[41] is a most instructive example of the characteristics of alloys. At the ordinary temperature sodium does not absorb hydrogen, but from 300° to 421° the absorption takes place at the ordinary pressure (and at an increased pressure even at higher temperatures), as shown by Troost and Hautefeuille (1874). One volume of sodium absorbs as much as 238 volumes of hydrogen. The metal increases in volume, and when once formed the alloy can be preserved for some time without change at the ordinary temperature. The appearance of sodium hydride resembles that of sodium itself; it is as soft as this latter, when heated it becomes brittle, and decomposes above 300°, evolving hydrogen. In this decomposition all the phenomena of dissociation are very clearly shown—that is, the hydrogen gas evolved has a definite tension[42] corresponding with each definite temperature. This confirms the fact that the formation of substances capable of dissociation can only be accomplished within the dissociation limits. Sodium hydride melts more easily than sodium itself, and then does not undergo decomposition if it is in an atmosphere of hydrogen. It oxidises easily in air, but not so easily as potassium hydride. The chemical reactions of sodium are retained in its hydride, and, if we may so express it, they are even increased by the addition of hydrogen. At all events, in the properties of sodium hydride[43] we see other properties than in such hydrogen compounds as HCl, H₂O, H₃N, H₄C, or even in the gaseous metallic hydrides AsH₃, TeH₂. Platinum, palladium, nickel, and iron, in absorbing hydrogen form compounds in which hydrogen is in a similar state. In them, as in sodium hydride, the hydrogen is compressed, absorbed, occluded (Chapter II.)[43 bis]

The most important chemical property of sodium is its power of easily decomposing water and evolving hydrogen from the majority of the hydrogen compounds, and especially from all acids, and hydrates in which hydroxyl must be recognised. This depends on its power of combining with the elements which are in combination with the hydrogen. We already know that sodium disengages hydrogen, not only from water, hydrochloric acid,[44] and all other acids, but also from ammonia,[44 bis] with the formation of sodamide NH₂Na, although it does not displace hydrogen from the hydrocarbons.[45] Sodium burns both in chlorine and in oxygen, evolving much heat. These properties are closely connected with its power of taking up oxygen, chlorine, and similar elements from most of their compounds. Just as it removes the oxygen from the oxides of nitrogen and from carbonic anhydride, so also does it decompose the majority of oxides at definite temperatures. Here the action is essentially the same as in the decomposition of water. Thus, for instance, when acting on magnesium chloride the sodium displaces the magnesium, and when acting on aluminium chloride it displaces metallic aluminium. Sulphur, phosphorus, arsenic and a whole series of other elements, also combine with sodium.[46]

With oxygen sodium unites in three degrees of combination, forming a suboxide Na₄O,[46 bis] an oxide, Na₂O, and a peroxide, NaO. They are thus termed because Na₂O is a stable basic oxide (with water it forms a basic hydroxide), whilst Na₄O and NaO do not form corresponding saline hydrates and salts. The suboxide is a grey inflammable substance which easily decomposes water, disengaging hydrogen; it is formed by the slow oxidation of sodium at the ordinary temperature. The peroxide is a greenish yellow substance, fusing at a bright red heat; it is produced by burning sodium in an excess of oxygen, and it yields oxygen when treated with water:

Suboxide:	Na₄O	+	3H₂O	=	4NaHO	+	H₂[47]
Oxide:	Na₂O	+	H₂O	=	2NaHO[48]
Peroxide:	Na₂O₂	+	H₂O	=	2NaHO	+	O[49]

All three oxides form sodium hydroxide with water, but only the oxide Na₂O is directly transformed into a hydrate. The other oxides liberate either hydrogen or oxygen; they also present a similar distinction with reference to many other agents. Thus carbonic anhydride combines directly with the oxide Na₂O, which when heated in the gas burns, forming sodium carbonate, whilst the peroxide yields oxygen in addition. When treated with acids, sodium and all its oxides only form the salts corresponding with sodium oxide—that is, of the formula or type NaX. Thus the oxide of sodium, Na₂O, is the only salt-forming oxide of this metal, as water is in the case of hydrogen. Although the peroxide H₂O₂ is derived from hydrogen, and Na₂O₂ from sodium, yet there are no corresponding salts known, and if they are formed they are probably as unstable as hydrogen peroxide. Although carbon forms carbonic oxide, CO, still it has only one salt-forming oxide—carbonic anhydride, CO₂. Nitrogen and chlorine both give several salt-forming oxides and types of salts. But of the oxides of nitrogen, NO and NO₂ do not form salts, as do N₂O₃, N₂O₄, and N₂O₅, although N₂O₄ does not form special salts, and N₂O₅ corresponds with the highest form of the saline compounds of nitrogen. Such distinctions between the elements, according to their power of giving one or several saline forms, is a radical property of no less importance than the basic or acid properties of their oxides. Sodium as a typical metal does not form any acid oxides, whilst chlorine, as a typical non-metal, does not form bases with oxygen. Therefore sodium as an element may be thus characterised: it forms one very stable salt-forming oxide, Na₂O, having powerful basic properties, and its salts are of the general formula, NaX, therefore in its compounds it is, like hydrogen, a basic and univalent element.

On comparing sodium and its analogues, which will be described later with other metallic elements, it will be seen that these properties, together with the relative lightness of the metal itself and its compounds, and the magnitude of its atomic weight comprise the most essential properties of this element, clearly distinguishing it from others, and enabling us easily to recognise its analogues.

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 mineralogy thenardite. Crystalline decahydrated salt is termed in mineralogy mirabilite, 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 (see Note 15) 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 (Chapter I.) is given according to the determinations of Gay-Lussac, Lovell, and Mulder.

[6] In Chapter I., Note 24, 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 Na₂SO₄ in 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., Note 56.

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 + 0·35p² where 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., Note 50 (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., Note 56) 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 (see Note 7). From these figures, and from the specific gravities of sulphuric acid, it is evident that on mixing solutions of this acid and sodium sulphate expansion will 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., Note 27. 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,H₂SO₄ are 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 (Chapter X.), 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]

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Fig. 69.—Apparatus for the methodical lixiviation of black ash, &c. Water flows into the tanks from the pipes r, r, and the saturated liquid is drawn off from c, c.

Methodical lixiviation is 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 (Chapter I.).

[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’ (Note 14) 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 CO₂ in coke towers, thus obtaining H₂S (a gas which gives SO₂ under 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 CO₂ when ignited. This process quite eliminates the formation of soda-waste (see Note 3) 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, see Chapter II., Note 1) from strong solutions of salt (Chapter X., Note 23 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 (see Chapter XI., Note 3 bis) separates the NH₄Cl from the residual solutions by cooling (Chapter X., Note 44); 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 (Note 7) 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 factor A is 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 formula S = S₀ + Ap + Bp², where S₀ is 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 of p.

[18] The resemblance is so great that, notwithstanding the difference in the molecular composition of Na₂SO₄ and Na₂CO₃, they ought to be classed under the type (NaO)₂R, where R = SO₂ or 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 of trona and urao. 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., Note 25). (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 CO₂ as 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 (see p. 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 termed carbonate 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 soda is 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) (Note 27).

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 (Note 14). 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 (see Chapter IX.) 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 the reason for 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 CO₂ move in a circle round the elements Na₂O, but at the points of contact with Fe₂O₃ the motion becomes elliptic with a long axis, and at some distance from Na₂O the elements of CO₂ are 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.	5	10	15	20	30	40
Sp. gr.	1·057	1·113	1·169	1·224	1·331	1·436

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 (see Note 18).

[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	+	RHO
Ethereal salt		Caustic soda		Sodium salt		Alcohol

Such a decomposition is termed saponification; similar reactions were known very long ago for the ethereal salts corresponding with glycerin, C₃H₅(OH)₃ (Chapter IX.), 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, C₁₆H₃₂O₂ and 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 (see Chapter IX., Note 15).

[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 Note 33.

[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 that the 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. Note 9) 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 (see Chapter 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., Note 12).

[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 that for each atom of the sodium in 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 NaHg₅ about 18,500 calories are evolved; when NaHg₃ is 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 hydrogen p, in millimetres of mercury, is:—

	t =	330°	350°	400°	430°
for Na₂H	p =	28	57	447	910
for K₂H		45	72	548	1100

[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 (Chapter II.) has a much lower density.

[43 bis] We may remark that at low temperatures Na absorbs NH₃ and forms (NH₃Na)₂ (see Chapter VI., Note 14); this substance absorbs CO and gives (NaCO)n (Chapter IX., Note 31), 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., Note 14), 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 + 2H₂ at 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 NH₃ and NaHO, which react further. Anhydrous CaO absorbs NH₂Na when heated without decomposing it. When sodamide is heated with SiO₂, NH₃ is 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 (NHNa₂ and NNa₃ are 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 preserved in 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 (see Note 47).

[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 (see Chapter II., Note 9). 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 (see Note 26), 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 NaNO₃ be melted, it gives Na₂O₂ with 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 Na₂OI₂ is akin to the compound Cu₂OCl₂ obtained by oxidising CuCl. This reaction is one of the few in which iodine directly displaces oxygen. The substance Na₂OI₂ is 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 H₂O₂ formed). The oxidising properties of Na₂O₂ under 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 Na₂O₂ be 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 Na₂O₂ on 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.

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