XI. THE HALOGENS: CHLORINE, BROMINE, IODINE, AND FLUORINE

XI. THE HALOGENS: CHLORINE, BROMINE, IODINE, AND FLUORINE

Although hydrochloric acid, like water, is one of the most stable substances, it is nevertheless decomposed not only by the action of a galvanic current,[1] but also by a high temperature. Sainte-Claire Deville showed that decomposition already occurs at 1,300°, because a cold tube (as with CO, Chapter IX.) covered with an amalgam of silver absorbs chlorine from hydrochloric acid in a red-hot tube, and the escaping gas contains hydrogen. V. Meyer and Langer (1885) observed the decomposition of hydrochloric acid at 1,690° in a platinum vessel; the decomposition in this instance was proved not only from the fact that hydrogen diffused through the platinum (p. 142), owing to which the volume was diminished, but also from chlorine being obtained in the residue (the hydrogen chloride was mixed with nitrogen), which liberated iodine from potassium iodide.[2] The usual method for the preparation of chlorine consists in the abstraction of the hydrogen by oxidising agents.[2 bis]

[1] The decomposition of fused sodium chloride by an electric current has been proposed in America and Russia (N. N. Beketoff) as a means for the preparation of chlorine and sodium. A strong solution of hydrochloric acid is decomposed into equal volumes of chlorine and hydrogen by the action of an electric current. If sodium chloride and lead be melted in a crucible, the former being connected with the cathode and a carbon anode immersed in the lead, then the lead dissolves sodium and chlorine is disengaged as gas. This electrolytic method has not yet been practised on a large scale, probably because gaseous chlorine has not many applications, and because of the difficulty there is in dealing with it.

[2] To obtain so high a temperature (at which the best kinds of porcelain soften) Langer and Meyer employed the dense graphitoidal carbon from gas retorts, and a powerful blast. They determined the temperature by the alteration of the volume of nitrogen in the platinum vessel, for this gas does not permeate through platinum, and is unaltered by heat.

[2 bis] The acid properties of hydrochloric acid were known when Lavoisier pointed out the formation of acids by the combination of water with the oxides of the non-metals, and therefore there was reason for thinking that hydrochloric acid was formed by the combination of water with the oxide of some element. Hence when Scheele obtained chlorine by the action of hydrochloric acid on manganese peroxide he considered it as the acid contained in common salt. When it became known that chlorine gives hydrochloric acid with hydrogen, Lavoisier and Berthollet supposed it to be a compound with oxygen of an anhydride contained in hydrochloric acid. They supposed that hydrochloric acid contained water and the oxide of a particular radicle, and that chlorine was a higher degree of oxidation of this radicle muvias (from the Latin neme of hydrochloric acid, acidum muriaticum). It was only in 1811 that Gay-Lussac and Thénard in France and Davy in England arrived at the conclusion that the substance obtained by Scheele does not contain oxygen, nor under any conditions give water with hydrogen, and that there is no water in hydrochloric acid gas, and therefore concluded that chlorine is an elementary substance. They named it 'chlorine' from the Greek word [Greek: chlôros], signifying a green colour, because of the peculiar colour by which this gas is characterised.

An aqueous solution of hydrochloric acid is generally employed for t he evolution of chlorine. The hydrogen has to be abstracted from the hydrochloric acid. This is accomplished by nearly all oxidising substances, and especially by those which are able to evolve oxygen at a red heat (besides bases, such as mercury and silver oxides, which are able to give salts with hydrogen chloride); for example, manganese peroxide, potassium chlorate, chromic acid, &c. The decomposition essentially consists in the oxygen of the oxidising substance displacing the chlorine from 2HCl, forming water, H_{2}O, and setting the chlorine free, 2HCl + O (disengaged by the oxidising substances) = H_{2}O + Cl_{2}. Even nitric acid partially produces a like reaction; but as we shall afterwards see its action is more complicated, and it is therefore not suitable for the preparation of pure chlorine.[3] But other oxidising substances which do not give any other volatile products with hydrochloric acid may be employed for the preparation of chlorine. Among these may be mentioned: potassium chlorate, acid potassium chromate, sodium manganate, manganese peroxide, &c. Manganese peroxide is commonly employed in the laboratory, and on a large scale, for the preparation of chlorine. The chemical process in this case may be represented as follows: an exchange takes place between 4HCl and MnO_{2}, in which the manganese takes the place of the four atoms of hydrogen, or the chlorine and oxygen exchange places--that is, MnCl_{4} and 2H_{2}O are produced. The chlorine compound, MnCl_{4}, obtained is very unstable; it splits up into chlorine, which as a gas passes from the sphere of action, and a lower compound containing less chlorine than the substance first formed, which remains in the apparatus in which the mixture is heated, MnCl_{4} = MnCl_{2} + Cl_{2}.[3 bis] The action of hydrochloric acid requires a temperature of about 100°. In the laboratory the preparation of chlorine is carried on in flasks, heated over a water-bath, by acting on manganese peroxide with hydrochloric acid or a mixture of common salt and sulphuric acid[4] and washing the gas with water to remove hydrochloric acid.[5] Chlorine cannot be collected over mercury, because it combines with it as with many other metals, and it is soluble in water; however, it is but slightly soluble in hot water or brine. Owing to its great weight, chlorine may be directly collected in a dry vessel by carrying the gas-conducting tube down to the bottom of the vessel. The chlorine will lie in a heavy layer at the bottom of the vessel, displace the air, and the extent to which it fills the vessel may be followed by its colour.[6]

[3] However, nitric acid has been proposed as a means for obtaining chlorine, but by methods which have the drawback of being very complicated

[3 bis] This representation of the process of the reaction is most natural. However, this decomposition is generally represented as if chlorine gave only one degree of combination with manganese, MnCl_{2}, and therefore directly reacts in the following manner--MnO_{2} + 4HCl = MnCl_{2} + 2H_{2}O + Cl_{2}, in which case it is supposed that manganese peroxide, MnO_{2}, breaks up, as it were, into manganous oxide, MnO and oxygen, both of which react with hydrochloric acid, the manganous oxide acting upon HCl as a base, giving MnCl_{2} and at the same time 2HCl + O = H_{2}O + Cl_{2}. In reality, a mixture of oxygen and hydrochloric acid does give chlorine at a red heat, and this reaction may also take place at the moment of its evolution in this case.

All the oxides of manganese (Mn_{2}O_{3}, MnO_{2}, MnO_{3}, Mn_{2}O_{7}), with the exception of manganous oxide, MnO, disengage chlorine from hydrochloric acid, because manganous chloride, MnCl_{2}, is the only compound of chlorine and manganese which exists as a stable compound, all the higher chlorides of manganese being unstable and evolving chlorine. Hence we here take note of two separate changes: (1) an exchange between oxygen and chlorine, and (2) the instability of the higher chlorine compounds. As (according to the law of substitution) in the substitution of oxygen by chlorine, Cl_{2} takes the place of O, the chlorine compounds will contain more atoms than the corresponding oxygen compounds. It is not surprising, therefore, that certain of the chlorine compounds corresponding with oxygen compounds do not exist, or if they are formed are very unstable. And furthermore, an atom of chlorine is heavier than an atom of oxygen, and therefore a given element would have to retain a large mass of chlorine if in the higher oxides the oxygen were replaced by chlorine. For this reason equivalent compounds of chlorine do not exist for all oxygen compounds. Many of the former are immediately decomposed, when formed, with the evolution of chlorine. From this it is evident that there should exist such chlorine compounds as would evolve chlorine as peroxides evolve oxygen, and indeed a large number of such compounds are known. Amongst them may be mentioned antimony pentachloride, SbCl_{5}, which splits up into chlorine and antimony trichloride when heated. Cupric chloride, corresponding with copper oxide, and having a composition CuCl_{2}, similar to CuO, when heated parts with half its chlorine, just as barium peroxide evolves half its oxygen. This method may even be taken advantage of for the preparation of chlorine and cuprous chloride, CuCl. The latter attracts oxygen from the atmosphere, and in so doing is converted from a colourless substance into a green compound whose composition is Cu_{2}Cl_{2}O. With hydrochloric acid this substance gives cupric chloride (Cu_{2}Cl_{2}O + 2HCl = H_{2}O + 2CuCl_{2}), which has only to be dried and heated in order again to obtain chlorine. Thus, in solution, and at the ordinary temperature, the compound CuCl_{2} is stable, but when heated it splits up. On this property is founded Deacon's process for the preparation of chlorine from hydrochloric acid with the aid of air and copper salts, by passing a mixture of air and hydrochloric acid at about 440° over bricks saturated with a solution of a copper salt (a mixture of solutions of CuSO_{4} and Na_{2}SO_{4}). CuCl_{2} is then formed by the double decomposition of the salt of copper and the hydrochloric acid; the CuCl_{2} liberates chlorine, and the CuCl forms Cu_{2}Cl_{2}O with the oxygen of the air, which again gives CuCl_{2} with 2HCl, and so on.

Magnesium chloride, which is obtained from sea-water, carnallite, &c., may serve not only as a means for the preparation of hydrochloric acid, but also of chlorine, because its basic salt (magnesium oxychloride) when heated in the air gives magnesium oxide and chlorine (Weldon-Pechiney's process, 1888). Chlorine is now prepared on a large scale by this method. Several new methods based upon this reaction have been proposed for procuring chlorine from the bye-products of other chemical processes. Thus, Lyte and Tattars (1891) obtained up to 67 p.c. of chlorine from CaCl_{2} in this manner. A solution of CaCl_{2}, containing a certain amount of common salt, is evaporated and oxide of magnesium added to it. When the solution attains a density of 1·2445 (at 15°), it is treated with carbonic acid, which precipitates carbonate of calcium, while chloride of magnesium remains in solution. After adding ammonium chloride, the solution is evaporated to dryness and the double chloride of magnesium and ammonium formed is ignited, which drives off the chloride of ammonium. The chloride of magnesium which remains behind is used in the Weldon-Pechiney process. The De Wilde-Reychler (1892) process for the manufacture of chlorine consists in passing alternate currents of hot air and hydrochloric acid gas through a cylinder containing a mixture of the chlorides of magnesium and manganese. A certain amount of sulphate of magnesium which does not participate in any way in the reaction, is added to the mixture to prevent its fusing. The reactions may be expressed by the following equations: (1) 3MgCl_{2} + 3MnCl_{2} + 8O = Mg_{3}Mn_{3}O_{8} + 12Cl; (2) Mg_{3}Mn_{3}O_{8} + 16HCl = 3MgCl_{2} + 3MnCl_{2} + 8H_{2}O + 4Cl. As nitric acid is able to take up the hydrogen from hydrochloric acid, a heated mixture of these acids is also employed for the preparation of chlorine. The resultant mixture of chlorine and lower oxides of nitrogen is mixed with air and steam which regenerates the HNO_{3}, while the chlorine remains as a gas together with nitrogen, in which form it is quite capable of bleaching, forming chloride of lime, &c. Besides these, Solvay and Mond's methods of preparing chlorine must be mentioned. The first is based upon the reaction CaCl_{2} + SiO_{2} + O(air) = CaOSiO_{2} + Cl_{2}, the second on the action of the oxygen of the air (heated) upon MgCl_{2} (and certain similar chlorides) MgCl_{2} + O = MgO + Cl_{2} The remaining MgO is treated with sal-ammoniac to re-form MgCl_{2} (MgO + 2NH_{4}Cl = MgCl_{2} + H_{2}O + 2NH_{3}) and the resultant NH_{3} again converted into sal-ammoniac, so that hydrochloric acid is the only substance consumed. The latter processes have not yet found much application.

[4] The following proportions are accordingly taken by weight: 5 parts of powdered manganese peroxide, 11 parts of salt (best fused, to prevent its frothing), and 14 parts of sulphuric acid previously mixed with an equal volume of water. The mixture is heated in a salt bath, so as to obtain a temperature above 100°. The corks in the apparatus must be soaked in paraffin (otherwise they are corroded by the chlorine), and black india-rubber tubing smeared with vaseline must be used, and not vulcanised rubber (which contains sulphur, and becomes brittle under the action of the chlorine).

The reaction which proceeds may be expressed thus: MnO_{2} + 2NaCl + 2H_{2}SO_{4} = MnSO_{4} + Na_{2}SO_{4} + 2H_{2}O + Cl_{2}. The method of preparation of Cl_{2} from manganese peroxide and hydrochloric acid was discovered by Scheele, and from sodium chloride by Berthollet.

[5] The reaction of hydrochloric acid upon bleaching powder gives chlorine without the aid of heat, CaCl_{2}O_{2} + 4HCl = CaCl_{2} + 2H_{2}O + 2Cl_{2} and is therefore also used for the preparation of chlorine. This reaction is very violent if all the acid be added at once; it should be poured in drop by drop (Mermé, Kämmerer). C. Winkler proposed to mix bleaching powder with one quarter of burnt and powdered gypsum, and having damped the mixture with water, to press and cut it up into cubes and dry at the ordinary temperature. These cubes can be used for the preparation of chlorine in the same apparatus as that used for the evolution of hydrogen and carbonic anhydride--the disengagement of the chlorine proceeds uniformly.

A mixture of potassium dichromate and hydrochloric acid evolves chlorine perfectly free from oxygen (V. Meyer and Langer).

[6] [Illustration: FIG. 66.--Clay retort for the preparation of chlorine on a large scale.]

Chlorine is manufactured on a large scale from manganese peroxide and hydrochloric acid. It is most conveniently prepared in the apparatus shown in fig. 66, which consists of a three-necked earthenware vessel whose central orifice is the largest. A clay or lead funnel, furnished with a number of orifices, is placed in the central wide neck of the vessel. Roughly-ground lumps of natural manganese peroxide are placed in the funnel, which is then closed by the cover N, and luted with clay. One orifice is closed by a clay stopper, and is used for the introduction of the hydrochloric acid and withdrawal of the residues. The chlorine disengaged passes along a leaden gas-conducting tube placed in the other orifice. A row of these vessels is surrounded by a water-bath to ensure their being uniformly heated. Manganese chloride is found in the residue. In Weldon's process lime is added to the acid solution of manganese chloride. A double decomposition takes place, resulting in the formation of manganous hydroxide and calcium chloride. When the insoluble manganous hydroxide has settled, a further excess of milk of lime is added (to make a mixture 2Mn(OH)_{2} + CaO + _x_CaCl_{2}, which is found to be the best proportion, judging from experiment), and then air is forced through the mixture. The hydroxide is thus converted from a colourless to a brown substance, containing peroxide, MnO_{2}, and oxide of manganese, Mn_{2}O_{3}. This is due to the manganous oxide absorbing oxygen from the air. Under the action of hydrochloric acid this mixture evolves chlorine, because of all the compounds of chlorine and manganese the chloride MnCl_{2} is the only one which is stable (see Note 3). Thus one and the same mass of manganese may be repeatedly used for the preparation of chlorine. The same result is attained in other ways. If manganous oxide be subjected to the action of oxides of nitrogen and air (Coleman's process), then manganese nitrate is formed, which at a red heat gives oxides of nitrogen (which are again used in the process) and manganese peroxide, which is thus renewed for the fresh evolution of chlorine.

Chlorine is a gas of a yellowish green colour, and has a very suffocating and characteristic odour. On lowering the temperature to -50° or increasing the pressure to six atmospheres (at 0°) chlorine condenses[7] into a liquid which has a yellowish-green colour, a density of 1·3, and boils at -34°. The density and atomic weight of chlorine is 35·5 times greater than that of hydrogen, hence the molecule contains Cl_{2}[8]. At 0° one volume of water dissolves about 1-1/2 volume of chlorine, at 10° about 3 volumes, at 50° again 1-1/2 volume.[9] Such a solution of chlorine is termed 'chlorine water;' and is employed in a diluted form in medicine and as a laboratory reagent. It is prepared by passing chlorine through a series of Woulfe's bottles or into an inverted retort filled with water. Under the action of light, chlorine water gives oxygen and hydrochloric acid. At 0° a saturated solution of chlorine yields a crystallo-hydrate, Cl_{2},8H_{2}O, which easily splits up into chlorine and water when heated, so that if it be sealed up in a tube and heated to 35°, two layers of liquid are formed--a lower stratum of chlorine containing a small quantity of water, and an upper stratum of water containing a small quantity of chlorine.[10]

[7] Davy and Faraday liquefied chlorine in 1823 by heating the crystallo-hydrate Cl_{2}8H_{2}O in a bent tube (as with NH_{3}), surrounded by warm water, while the other end of the tube was immersed in a freezing mixture. Meselan condensed chlorine in freshly-burnt charcoal (placed in a glass tube), which when cold absorbs an equal weight of chlorine. The tube was then fused up, the bent end cooled, and the charcoal heated, by which means the chlorine was expelled from the charcoal, and the pressure increased.

[8] Judging from Ludwig's observations (1868), and from the fact that the coefficient of expansion of gases increases with their molecular weight (Chapter II., Note 26, for hydrogen = 0·367, carbonic anhydride = 0·373, hydrogen bromide = 0·386), it might be expected that the expansion of chlorine would be greater than that of air or of the gases composing it. V. Meyer and Langer (1885) having remarked that at 1,400° the density of chlorine (taking its expansion as equal to that of nitrogen) = 29, consider that the molecules of chlorine split up and partially give molecules Cl, but it might be maintained that the decrease in density observed only depends on the increase of the coefficient of expansion.

[9] Investigations on the solubility of chlorine in water (the solutions evolve all their chlorine on boiling and passing air through them) show many different peculiarities. First Gay-Lussac, and subsequently Pelouze, determined that the solubility increases between 0° and 8°-10° (from 1-1/2 to 2 vols. of chlorine per 100 vols. of water at 0° up to 3 to 2-3/4 at 10°). In the following note we shall see that this is not due to the breaking-up of the hydrate at about 8° to 10°, but to its formation below 9°. Roscoe observed an increase in the solubility of chlorine in the presence of hydrogen--even in the dark. Berthelot determined an increase of solubility with the progress of time. Schönbein and others suppose that chlorine acts on water, forming hypochlorous and hypochloric acids, (HClO + HCl).

The equilibrium between chlorine and steam as gases and between water, liquid chlorine, ice, and the solid crystallo-hydrate of chlorine is evidently very complex. Gibbs, Guldberg (1870) and others gave a theory for similar states of equilibrium, which was afterwards developed by Roozeboom (1887), but it would be inopportune here to enter into its details. It will be sufficient in the first place to mention that there is now no doubt (according to the theory of heat, and the direct observations of Ramsay and Young) that the vapour tensions at one and the same temperature are different for the liquid and solid states of substances; secondly, to call attention to the following note; and, thirdly, to state that, in the presence of the crystallo-hydrate, water between O°·24 and +28°·7 (when the hydrate and a solution may occur simultaneously) dissolves a different amount of chlorine than it does in the absence of the crystallo-hydrate.

[10] According to Faraday's data the hydrate of chlorine contains Cl_{2},10H_{2}O, but Roozeboom (1885) showed that it is poorer in water and = Cl_{2},8H_{2}O. At first small, almost colourless, crystals are obtained, but they gradually form (if the temperature be below their critical point 28°·7, above which they do not exist) large yellow crystals, like those of potassium chromate. The specific gravity is 1·23. The hydrate is formed if there be more chlorine in a solution than it is able to dissolve under the dissociation pressure corresponding with a given temperature. In the presence of the hydrate the percentage amount of chlorine at 0° = 0·5, at 9° = 0·9, and at 20° = 1·82. At temperatures below 9° the solubility (determined by Gay-Lussac and Pelouze, see Note 9) is dependent on the formation of the hydrate; whilst at higher temperatures under the ordinary pressure the hydrate cannot be formed, and the solubility of chlorine falls, as it does for all gases (Chapter I.). If the crystallo-hydrate is not formed, then below 9° the solubility follows the same rule (6° 1·07 p.c. Cl, 9° 0·95 p.c.). According to Roozeboom, the chlorine evolved by the hydrate presents the following tensions of dissociation: at 0° = 249 mm., at 4° = 398, at 8° = 620, at 10° = 797, at 14° = 1,400 mm. In this case a portion of the crystallo-hydrate remains solid. At 9°·6 the tension of dissociation is equal to the atmospheric pressure. At a higher pressure the crystallo-hydrate may form at temperatures above 9° up to 28°·7, when the vapour tension of the hydrate equals the tension of the chlorine. It is evident that the equilibrium which is established is on the one hand a case of a complex heterogeneous system, and on the other hand a case of the solution of solid and gaseous substances in water.

The crystallo-hydrate or chlorine water must be kept in the dark, or the access of light be prevented by coloured glass, otherwise oxygen is evolved and hydrochloric acid formed.

Chlorine explodes with hydrogen, if a mixture of equal volumes be exposed to the direct action of the sun's rays[11] or brought into contact with spongy platinum, or a strongly heated substance, or when subjected to the action of an electric spark. The explosion in this case takes place for exactly the same reasons--i.e. the evolution of heat and expansion of the resultant product--as in the case of detonating gas (Chapter III.) Diffused light acts in the same way, but slowly, whilst direct sunlight causes an explosion.[12] The hydrochloric acid gas produced by the reaction of chlorine on hydrogen occupies (at the original temperature and pressure) a volume equal to the sum of the original volumes; that is, a reaction of substitution here takes place: H_{2} + Cl_{2} = HCl + HCl. In this reaction twenty-two thousand heat units are evolved for one part by weight [1 gram] of hydrogen.[13]

[11] The chemical action of light on a mixture of chlorine and hydrogen was discovered by Gay-Lussac and Thénard (1809). It has been investigated by many savants, and especially by Draper, Bunsen, and Roscoe. Electric or magnesium light, or the light emitted by the combustion of carbon bisulphide in nitric oxide, and actinic light in general, acts in the same manner as sunlight, in proportion to its intensity. At temperatures below -12° light no longer brings about reaction, or at all events does not give an explosion. It was long supposed that chlorine that had been subjected to the action of light was afterwards able to act on hydrogen in the dark, but it was shown that this only takes place with moist chlorine, and depends on the formation of oxides of chlorine. The presence of foreign gases, and even of excess of chlorine or of hydrogen, very much enfeebles the explosion, and therefore the experiment is conducted with a detonating mixture prepared by the action of an electric current on a strong solution (sp. gr. 1·15) of hydrochloric acid, in which case the water is not decomposed--that is, no oxygen becomes mixed with the chlorine.

[12] The quantity of chlorine and hydrogen which combine is proportional to the intensity of the light--not of all the rays, but only those so-termed chemical (actinic) rays which produce chemical action. Hence a mixture of chlorine and hydrogen, when exposed to the action of light in vessels of known capacity and surface, may be employed as an actinometer--that is, as a means for estimating the intensity of the chemical rays, the influence of the heat rays being previously destroyed, which may be done by passing the rays through water. Investigations of this kind (photo-chemical) showed that chemical action is chiefly limited to the violet end of the spectrum, and that even the invisible ultra-violet rays produce this action. A colourless gas flame contains no chemically active rays; the flame coloured green by a salt of copper evinces more chemical action than the colourless flame, but the flame brightly coloured yellow by salts of sodium has no more chemical action than that of the colourless flame.

As the chemical action of light becomes evident in plants, photography, the bleaching of tissues, and the fading of colours in the sunlight, and as a means for studying the phenomenon is given in the reaction of chlorine on hydrogen, this subject has been the most fully investigated in photo-chemistry. The researches of Bunsen and Roscoe in the fifties and sixties are the most complete in this respect. Their actinometer contains hydrogen and chlorine, and is surrounded by a solution of chlorine in water. The hydrochloric acid is absorbed as it forms, and therefore the variation in volume indicates the progress of the combination. As was to be expected, the action of light proved to be proportional to the time of exposure and intensity of the light, so that it was possible to conduct detailed photometrical investigations respecting the time of day and season of the year, various sources of light, its absorption, &c. This subject is considered in detail in special works, and we only stop to mention one circumstance, that a small quantity of a foreign gas decreases the action of light; for example, 1/330 of hydrogen by 38 p.c., 1/200 of oxygen by 10 p.c., 1/100 of chlorine by 60 p.c., &c. According to the researches of Klimenko and Pekatoros (1889), the photo-chemical alteration of chlorine water is retarded by the presence of traces of metallic chlorides, and this influence varies with different metals.

As much heat is evolved in the reaction of chlorine on hydrogen, and as this reaction, being exothermal, may proceed by itself, the action of light is essentially the same as that of heat--that is, it brings the chlorine and hydrogen into the condition necessary for the reaction--it, as we may say, disturbs the original equilibrium; this is the work done by the luminous energy. It seems to me that the action of light on the mixed gases should be understood in this sense, as Pringsheim (1877) pointed out.

[13] In the formation of steam (from one part by weight [1 gram] of hydrogen) 29,000 heat units are evolved. The following are the quantities of heat (thousands of units) evolved in the formation of various other corresponding compounds of oxygen and of chlorine (from Thomsen's, and, for Na_{2}O, Beketoff's results):

{2NaCl, 195; CaCl_{2}, 170; HgCl_{2}, 63; 2AgCl, 59. { Na_{2}O, 100; CaO, 131; HgO, 42; Ag_{2}O, 6. {2AsCl_{3}, 143; 2PCl_{5}, 210; CCl_{4}, 21; 2HCl, 44 (gas). { As_{2}O_{3}, 155; P_{2}O_{5}, 370; CO_{2}, 97; H_{2}O, 58 (gas).

With the first four elements the formation of the chlorine compound gives the most heat, and with the four following the formation of the oxygen compound evolves the greater amount of heat. The first four chlorides are true salts formed from HCl and the oxide, whilst the remainder have other properties, as is seen from the fact that they are not formed from hydrochloric acid and the oxide, but give hydrochloric acid with water.

These relations show that the affinity of chlorine for hydrogen is very great and analogous to the affinity between hydrogen and oxygen. Thus[14] on the one hand by passing a mixture of steam and chlorine through a red-hot tube, or by exposing water and chlorine to the sunlight, oxygen is disengaged, whilst on the other hand, as we saw above, oxygen in many cases displaces chlorine from its compound with hydrogen, and therefore the reaction H_{2}O + Cl_{2} = 2HCl + O belongs to the number of reversible reactions, and hydrogen will distribute itself between oxygen and chlorine. This determines the relation of Cl to substances containing hydrogen and its reactions in the presence of water, to which we shall turn our attention after having pointed out the relation of chlorine to other elements.

[14] This has been already pointed out in Chapter III., Note 5.

Many metals when brought into contact with chlorine immediately combine with it, and form those metallic chlorides which correspond with hydrogen chloride and with the oxide of the metal taken. This combination may proceed rapidly with the evolution of heat and light; that is, metals are able to burn in chlorine. Thus, for example, sodium[15] burns in chlorine, synthesising common salt. Metals in the form of powders burn without the aid of heat, and become highly incandescent in the process; for instance, antimony, which is a metal easily converted into a powder.[16] Even such metals as gold and platinum,[17] which do not combine directly with oxygen and give very unstable compounds with it, unite directly with chlorine to form metallic chlorides. Either chlorine water or aqua regia may be employed for this purpose instead of gaseous chlorine. These dissolve gold and platinum, converting them into metallic chlorides. Aqua regia is a mixture of 1 part of nitric acid with 2 to 3 parts of hydrochloric acid. This mixture converts into soluble chlorides not only those metals which are acted on by hydrochloric and nitric acids, but also gold and platinum, which are insoluble in either acid separately. This action of aqua regia depends on the fact that nitric acid in acting on hydrochloric acid evolves chlorine. If the chlorine evolved be transferred to a metal, then a fresh quantity is formed from the remaining acids and also combines with the metal.[18] Thus the aqua regia acts by virtue of the chlorine which it contains and disengages.

[15] Sodium remains unaltered in perfectly dry chlorine at the ordinary temperature, and even when slightly warmed; but the combination is exceedingly violent at a red heat.

[16] An instructive experiment on combustion in chlorine may be conducted as follows: leaves of Dutch metal (used instead of gold for gilding) are placed in a glass globe, and a gas-conducting tube furnished with a glass cock is placed in the cork closing it, and the air is pumped out of the globe. The gas-conducting tube is then connected with a vessel containing chlorine, and the cock opened; the chlorine rushes in, and the metallic leaves are consumed.

[17] The behaviour of platinum to chlorine at a high temperature (1,400°) is very remarkable, because platinous chloride, PtCl_{2}, is then formed, whilst this substance decomposes at a much lower temperature into chlorine and platinum. Hence, when chlorine comes into contact with platinum at such high temperatures, it forms fumes of platinous chloride, and they on cooling decompose, with the liberation of platinum, so that the phenomenon appears to be dependent on the volatility of platinum. Deville proved the formation of platinous chloride by inserting a cold tube inside a red-hot one (as in the experiment on carbonic oxide). However, V. Meyer was able to observe the density of chlorine in a platinum vessel at 1,690°, at which temperature chlorine does not exert this action on platinum, or at least only to an insignificant degree.

[18] When left exposed to the air aqua regia disengages chlorine, and afterwards it no longer acts on gold. Gay-Lussac, in explaining the action of aqua regia, showed that when heated it evolves, besides chlorine, the vapours of two chloranhydrides--that of nitric acid, NO_{2}Cl (nitric acid, NO_{2}OH, in which HO is replaced by chlorine; see Chapter on Phosphorus), and that of nitrous acid, NOCl--but these do not act on gold. The formation of aqua regia may therefore be expressed by 4NHO_{3} + 8HCl = 2NO_{2}Cl + 2NOCl + 6H_{2}O + 2Cl_{2}. The formation of the chlorides NO_{2}Cl and NOCl is explained by the fact that the nitric acid is deoxidised, gives the oxides NO and NO_{2}, and they directly combine with chlorine to form the above anhydrides.

The majority of non-metals also react directly on chlorine; hot sulphur and phosphorus burn in it and combine with it at the ordinary temperature. Only nitrogen, carbon, and oxygen do not combine directly with it. The chlorine compounds formed by the non-metals--for instance, phosphorus trichloride, PCl_{3}, and sulphurous chloride, &c., do not have the properties of salts, and, as we shall afterwards see more fully, correspond to acid anhydrides and acids; for example, PCl_{3}--to phosphorous acid, P(OH)_{3}:

NaCl FeCl_{2} SnCl_{4} PCl_{3} HCl Na(HO) Fe(HO){2} Sn(HO){4} P(HO)_{3} H(HO)

As the above-mentioned relation in composition--i.e. substitution of Cl by the aqueous residue--exists between many chlorine compounds and their corresponding hydrates, and as furthermore some (acid) hydrates are obtained from chlorine compounds by the action of water, for instance,

PCl_{3} + 3H_{2}0 = P(HO)_{3} + 3HCl Phosphorus Water Phosphorous Hydrochloric trichloride acid acid

whilst other chlorine compounds are formed from hydroxides and hydrochloric acid, with the liberation of water, for example,

NaHO + HCl = NaCl + H_{2}O

we endeavour to express this intimate connection between the hydrates and chlorine compounds by calling the latter chloranhydrides. In general terms, if the hydrate be basic, then,

M(HO) + HCl = MCl + H_{2}O hydrate + hydrochloric acid = chloranhydride + water

and if the hydrate ROH be acid, then,

RCl + H_{2}O = R(HO) + HCl Chloranhydride + water = hydrate + hydrochloric acid

The chloranhydrides MCl corresponding to the bases are evidently metallic chlorides or salts corresponding to HCl. In this manner a distinct equivalency is marked between the compounds of chlorine and the so-called hydroxyl radicle (HO), which is also expressed in the analogy existing between chlorine, Cl_{2}, and hydrogen peroxide, (HO)_{2}.

As regards the chloranhydrides corresponding to acids and non-metals, they bear but little resemblance to metallic salts. They are nearly all volatile, and have a powerful suffocating smell which irritates the eyes and respiratory organs. They react on water like many anhydrides of the acids, with the evolution of heat and liberation of hydrochloric acid, forming acid hydrates. For this reason they cannot usually be obtained from hydrates--that is, acids--by the action of hydrochloric acid, as in that case water would be formed together with them, and water decomposes them, converting them into hydrates. There are many intermediate chlorine compounds between true saline metallic chlorides like sodium chloride and true acid chloranhydrides, just as there are all kinds of transitions between bases and acids. Acid chloranhydrides are not only obtained from chlorine and non-metals, but also from many lower oxides, by the aid of chlorine. Thus, for example, CO, NO, NO_{2}, SO_{2}, and other lower oxides which are capable of combining with oxygen may also combine with a corresponding quantity of chlorine. Thus COCl_{2}, NOCl, NO_{2}Cl, SO_{2}Cl_{2}, &c., are obtained. They correspond with the hydrates CO(OH){2}, NO(OH), NO{2}(OH), SO_{2}(OH){2}, &c., and to the anhydrides CO{2}, N_{2}O_{3}, N_{2}O_{5}, SO_{3}, &c. Here we should notice two aspects of the matter: (1) chlorine combines with that with which oxygen is able to combine, because it is in many respects equally if not more energetic than oxygen and replaces it in the proportion Cl_{2} : O; (2) that highest limit of possible combination which is proper to a given element or grouping of elements is very easily and often attained by combination with chlorine. If phosphorus gives PCl_{3} and PCl_{5}, it is evident that PCl_{5} is the higher form of combination compared with PCl_{3}. To the form PCl_{5}, or in general PX_{5}, correspond PH_{4}I, PO(OH){3}, POCl{3}, &c. If chlorine does not always directly give compounds of the highest possible forms for a given element, then generally the lower forms combine with it in order to reach or approach the limit. This is particularly clear in hydrocarbons, where we see the limit C_{n}H_{2_n_+2} very distinctly. The unsaturated hydrocarbons are sometimes able to combine with chlorine with the greatest ease and thus reach the limit. Thus ethylene, C_{2}H_{4}, combines with Cl_{2}, forming the so-called Dutch liquid or ethylene chloride, C_{2}H_{4}Cl_{2}, because it then reaches the limit C_{n}X_{2_n_+2}. In this and all similar cases the combined chlorine is able by reactions of substitution to give a hydroxide and a whole series of other derivatives. Thus a hydroxide called glycol, C_{2}H_{4}(OH){2}, is obtained from C{2}H_{4}Cl_{2}.

Chlorine in the presence of water very often acts directly as an oxidising agent. A substance A combines with chlorine and gives, for example, ACl_{2}, and this in turn a hydroxide, A(OH){2}, which on losing water forms AO. Here the chlorine has oxidised the substance A. This frequently happens in the simultaneous action of water and chlorine: A + H{2}O + Cl_{2} = 2HCl + AO. Examples of this oxidising action of chlorine may frequently be observed both in practical chemistry and technical processes. Thus, for instance, chlorine in the presence of water oxidises sulphur and metallic sulphides. In this case the sulphur is converted into sulphuric acid, and the chlorine into hydrochloric acid, or a metallic chloride if a metallic sulphide be taken. A mixture of carbonic oxide and chlorine passed into water gives carbonic anhydride and hydrochloric acid. Sulphurous anhydride is oxidised by chlorine in the presence of water into sulphuric acid, just as it is by the action of nitric acid: SO_{2} + 2H_{2}O + Cl_{2} = H_{2}SO_{4} + 2HCl.

The oxidising action of chlorine in the presence of water is taken advantage of in practice for the rapid bleaching of tissues and fibres. The colouring matter of the fibres is altered by oxidation and converted into a colourless substance, but the chlorine afterwards acts on the tissue itself. Bleaching by means of chlorine therefore requires a certain amount of technical skill in order that the chlorine should not act on the fibres themselves, but that its action should be limited to the colouring matter only. The fibre for making writing paper, for instance, is bleached in this manner. The bleaching property of chlorine was discovered by Berthollet, and forms an important acquisition to the arts, because it has in the majority of cases replaced that which before was the universal method of bleaching--namely, exposure to the sun of the fabrics damped with water, which is still employed for linens, &c. Time and great trouble, and therefore money also, have been considerably saved by this change.[19]

[19] Ozone and peroxide of hydrogen also bleach tissues. As the action of peroxide of hydrogen is easily controlled by taking a weak solution, and as it has hardly any action upon the tissues themselves, it is replacing chlorine more and more as a bleaching agent. The oxidising property of chlorine is apparent in destroying the majority of organic tissues, and proves fatal to organisms. This action of chlorine is taken advantage of in quarantine stations. But the simple fumigation by chlorine must be carried on with great care in dwelling places, because chlorine disengaged into the atmosphere renders it harmful to the health.

The power of chlorine for combination is intimately connected with its capacity for substitution, because, according to the law of substitution, if chlorine combines with hydrogen, then it also replaces hydrogen, and furthermore the combination and substitution are accomplished in the same quantities. Therefore the atom of chlorine which combines with the atom of hydrogen is also able to replace the atom of hydrogen. We mention this property of chlorine not only because it illustrates the application of the law of substitution in clear and historically important examples, but more especially because reactions of this kind explain those indirect methods of the formation of many substances which we have often mentioned and to which recourse is had in many cases in chemistry. Thus chlorine does not act on carbon,[20] oxygen, or nitrogen, but nevertheless its compounds with these elements may be obtained by the indirect method of the substitution of hydrogen by chlorine.

[20] A certain propensity of carbon to attract chlorine is evidenced in the immense absorption of chlorine by charcoal (Note 7), but, so far as is at present known (if I am not mistaken, no one has tried the aid of light), no combination takes place between the chlorine and carbon.

As chlorine easily combines with hydrogen, and does not act on carbon, it decomposes hydrocarbons (and many of their derivatives) at a high temperature, depriving them of their hydrogen and liberating the carbon, as, for example, is clearly seen when a lighted candle is placed in a vessel containing chlorine. The flame becomes smaller, but continues to burn for a certain time, a large amount of soot is obtained, and hydrochloric acid is formed. In this case the gaseous and incandescent substances of the flame are decomposed by the chlorine, the hydrogen combines with it, and the carbon is disengaged as soot.[21] This action of chlorine on hydrocarbons, &c., proceeds otherwise at lower temperatures, as we will now consider.

[21] The same reaction takes place under the action of oxygen, with the difference that it burns the carbon, which chlorine is not able to do. If chlorine and oxygen compete together at a high temperature, the oxygen will unite with the carbon, and the chlorine with the hydrogen.

A very important epoch in the history of chemistry was inaugurated by the discovery of Dumas and Laurent that chlorine is able to displace and replace hydrogen. This discovery is important from the fact that chlorine proved to be an element which combines with great ease simultaneously with both the hydrogen and the element with which the hydrogen was combined. This clearly proved that there is no opposite polarity between elements forming stable compounds. Chlorine does not combine with hydrogen because it has opposite properties, as Dumas and Laurent stated previously, accounting hydrogen to be electro-positive and chlorine electro-negative; this is not the reason of their combining together, for the same chlorine which combines with hydrogen is also able to replace it without altering many of the properties of the resultant substance. This substitution of hydrogen by chlorine is termed metalepsis. The mechanism of this substitution is very constant. If we take a hydrogen compound, preferably a hydrocarbon, and if chlorine acts directly on it, then there is produced on the one hand hydrochloric acid and on the other hand a compound containing chlorine in the place of the hydrogen--so that the chlorine divides itself into two equal portions, one portion is evolved as hydrochloric acid, and the other portion takes the place of the hydrogen thus liberated. Hence this metalepsis is always accompanied by the formation of hydrochloric acid.[22] The scheme of the process is as follows:

C_{n}H_{m}X + Cl_{2} = C_{n}H_{m-1}ClX + HCl Hydrocarbon Free Product of Hydrochloric chlorine metalepsis acid

Or, in general terms--

RH + Cl_{2} = RCl + HCl.

[22] This division of chlorine into two portions may at the same time be taken as a clear confirmation of the conception of molecules. According to Avogadro-Gerhardt's law, the molecule of chlorine (p. 310) contains two atoms of this substance; one atom replaces hydrogen, and the other combines with it.

The conditions under which metalepsis takes place are also very constant. In the dark chlorine does not usually act on hydrogen compounds, but the action commences under the influence of light. The direct action of the sun's rays is particularly propitious to metalepsis. It is also remarkable that the presence of traces of certain substances,[23] especially of iodine, aluminium chloride, antimony chloride, &c., promotes the action. A trace of iodine added to the substance subjected to metalepsis often produces the same effect as sunlight.[24]

[23] Such carriers or media for the transference of chlorine and the halogens in general were long known to exist in iodine and antimonious chloride, and have been most fully studied by Gustavson and Friedel, of the Petroffsky Academy--the former with respect to aluminium bromide, and the latter with respect to aluminium chloride. Gustavson showed that if a trace of metallic aluminium be dissolved in bromine (it floats on bromine, and when combination takes place much heat and light are evolved), the latter becomes endowed with the property of entering into metalepsis, which it is not able to do of its own accord. When pure, for instance, it acts very slowly on benzene, C_{6}H_{6}, but in the presence of a trace of aluminium bromide the reaction proceeds violently and easily, so that each drop of the hydrocarbon gives a mass of hydrobromic acid, and of the product of metalepsis. Gustavson showed that the modus operandi of this instructive reaction is based on the property of aluminium bromide to enter into combination with hydrocarbons and their derivatives. The details of this and all researches concerning the metalepsis of the hydrocarbons must be looked for in works on organic chemistry.

[24] As small admixtures of iodine, aluminium bromide, &c., aid the metalepsis of large quantities of a substance, just as nitric oxide aids the reaction of sulphurous anhydride on oxygen and water, so the principle is essentially the same in both cases. Effects of this kind (which should also be explained by a chemical reaction proceeding at the surfaces) only differ from true contact phenomena in that the latter are produced by solid bodies and are accomplished at their surfaces, whilst in the former all is in solution. Probably the action of iodine is founded on the formation of iodine chloride, which reacts more easily than chlorine.

If marsh gas be mixed with chlorine and the mixture ignited, then the hydrogen is entirely taken up from the marsh gas and hydrochloric acid and carbon formed, but there is no metalepsis.[25] But if a mixture of equal volumes of chlorine and marsh gas be exposed to the action of diffused light, then the greenish yellow mixture gradually becomes colourless, and hydrochloric acid and the first product of metalepsis--namely, methyl chloride--are formed:

CH_{4} + Cl_{2} = CH_{3}Cl + HCl Marsh gas Chlorine Methyl chloride Hydrochloric acid

[25] Metalepsis belongs to the number of delicate reactions--if it may be so expressed--as compared with the energetic reaction of combustion. Many cases of substitution are of this kind. Reactions of metalepsis are accompanied by an evolution of heat, but in a less quantity than that evolved in the formation of the resulting quantity of the halogen acids. Thus the reaction C_{2}H_{6} + Cl_{2} = C_{2}H_{5}Cl + HCl, according to the data given by Thomsen, evolves about 20,000 heat units, whilst the formation of hydrochloric acid evolves 22,000 units.

The volume of the mixture remains unaltered. The methyl chloride which is formed is a gas. If it be separated from the hydrochloric acid (it is soluble in acetic acid, in which hydrochloric acid is but sparingly soluble) and be again mixed with chlorine, then it may be subjected to a further metalepsical substitution--the second atom of hydrogen may be substituted by chlorine, and a liquid substance, CH_{2}Cl_{2}, called methylene chloride, will be obtained. In the same manner the substitution may be carried on still further, and CHCl_{3}, or chloroform, and lastly carbon tetrachloride, CCl_{4}, will be produced. Of these substances the best known is chloroform, owing to its being formed from many organic substances (by the action of bleaching powder) and to its being used in medicine as an anæsthetic; chloroform boils at 62° and carbon tetrachloride at 78°. They are both colourless odoriferous liquids, heavier than water. The progressive substitution of hydrogen by chlorine is thus evident, and it can be clearly seen that the double decompositions are accomplished between molecular quantities of the substance--that is, between equal volumes in a gaseous state.

Carbon tetrachloride, which is obtained by the metalepsis of marsh gas, cannot be obtained directly from chlorine and carbon, but it may be obtained from certain compounds of carbon--for instance, from carbon bisulphide--if its vapour mixed with chlorine be passed through a red-hot tube. Both the sulphur and carbon then combine with the chlorine. It is evident that by ultimate metalepsis a corresponding carbon chloride may be obtained from any hydrocarbon--indeed, the number of chlorides of carbon C_{n}Cl_{2_m_} already known is very large.

As a rule, the fundamental chemical characters of hydrocarbons are not changed by metalepsis; that is, if a neutral substance be taken, then the product of metalepsis is also a neutral substance, or if an acid be taken the product of metalepsis also has acid properties. Even the crystalline form not unfrequently remains unaltered after metalepsis. The metalepsis of acetic acid, CH_{3}·COOH, is historically the most important. It contains three of the atoms of the hydrogen of marsh gas, the fourth being replaced by carboxyl, and therefore by the action of chlorine it gives three products of metalepsis (according to the amount of the chlorine and conditions under which the reaction takes place), mono-, di-, and tri-chloracetic acids--CH_{2}Cl·COOH, CHCl_{2}·COOH, and CCl_{3}·COOH; they are all, like acetic acid, monobasic. The resulting products of metalepsis, in containing an element which so easily acts on metals as chlorine, possess the possibility of attaining a further complexity of molecules of which the original hydrocarbon is often in no way capable. Thus on treating with an alkali (or first with a salt and then with an alkali, or with a basic oxide and water, &c.) the chlorine forms a salt with its metal, and the hydroxyl radicle takes the place of the chlorine--for example, CH_{3}·OH is obtained from CH_{3}Cl. By the action of metallic derivatives of hydrocarbons--for example, CH_{3}Na--the chlorine also gives a salt, and the hydrocarbon radicle--for instance, CH_{3}--takes the place of the chlorine. In this, or in a similar manner, CH_{3}·CH_{3}, or C_{2}H_{6} is obtained from CH_{3}Cl and C_{6}H_{5}·CH_{3} from C_{6}H_{6}. The products of metalepsis also often react on ammonia, forming hydrochloric acid (and thence NH_{4}Cl) and an amide; that is, the product of metalepsis, with the ammonia radicle NH_{2}, &c. in the place of chlorine. Thus by means of metalepsical substitution methods were found in chemistry for an artificial and general means of the formation of complex carbon compounds from more simple compounds which are often totally incapable of direct reaction. Besides which, this key opened the doors of that secret edifice of complex organic compounds into which man had up to then feared to enter, supposing the hydrocarbon elements to be united only under the influence of those mystic forces acting in organisms.[26]

[26] With the predominance of the representation of compound radicles (this doctrine dates from Lavoisier and Gay-Lussac) in organic chemistry, it was a very important moment in its history when it became possible to gain an insight into the structure of the radicles themselves. It was clear, for instance, that ethyl, C_{2}H_{5}, or the radicle of common alcohol, C_{2}H_{5}·OH, passes, without changing, into a number of ethyl derivatives, but its relation to the still simpler hydrocarbons was not clear, and occupied the attention of science in the 'forties' and 'fifties.' Having obtained ethyl hydride, C_{2}H_{5}H = C_{2}H_{6}, it was looked on as containing the same ethyl, just as methyl hydride, CH_{4} = CH_{3}H, was considered as existing in methane. Having obtained free methyl, CH_{3}CH_{3} = C_{2}H_{6}, from it, it was considered as a derivative of methyl alcohol, CH_{3}OH, and as only isomeric with ethyl hydride. By means of the products of metalepsis it was proved that this is not a case of isomerism but of strict identity, and it therefore became clear that ethyl is methylated methyl, C_{2}H_{5} = CH_{2}CH_{3}. In its time a still greater impetus was given by the study of the reactions of monochloracetic acid, CH_{2}Cl·COOH, or CO(CH_{2}Cl)(OH). It appeared that metalepsical chlorine, like the chlorine of chloranhydrides--for instance, of methyl chloride, CH_{3}Cl, or ethyl chloride, C_{2}H_{5}Cl--is capable of substitution; for example, glycollic acid, CH_{2}(OH)(CO_{2}H), or CO(CH_{2}·OH)(OH), was obtained from it, and it appeared that the OH in the group CH_{2}(OH) reacted like that in alcohols, and it became clear, therefore, that it was necessary to examine the radicles themselves by analysing them from the point of view of the bonds connecting the constituent atoms. Whence arose the present doctrine of the structure of the carbon compounds. (See Chapter VIII., Note 42.)

It is not only hydrocarbons which are subject to metalepsis. Certain other hydrogen compounds, under the action of chlorine, also give corresponding chlorine derivatives in exactly the same manner; for instance, ammonia, caustic potash, caustic lime, and a whole series of alkaline substances.[27] In fact, just as the hydrogen in marsh gas can be replaced by chlorine and form methyl chloride, so the hydrogen in caustic potash, KHO, ammonia, NH_{3}, and calcium hydroxide, CaH_{2}O_{2} or Ca(OH){2}, may be replaced by chlorine and give potassium hypochlorite, KClO, calcium hypochlorite, CaCl{2}O_{2}, and the so-called chloride of nitrogen, NCl_{3}. For not only is the correlation in composition the same as in the substitution in marsh gas, but the whole mechanism of the reaction is the same. Here also two atoms of chlorine act: one takes the place of the hydrogen whilst the other is evolved as hydrochloric acid, only in the former case the hydrochloric acid evolved remained free, and in the latter, in presence of alkaline substances, it reacts on them. Thus, in the action of chlorine on caustic potash, the hydrochloric acid formed acts on another quantity of caustic potash and gives potassium chloride and water, and therefore not only KHO + Cl_{2} = HCl + KClO, but also KHO + HCl = H_{2}O + KCl, and the result of both simultaneous phases will be 2KHO + Cl_{2} = H_{2}O + KCl + KClO. We will here discuss certain special cases.

[27] By including many instances of the action of chlorine under metalepsis we not only explain the indirect formation of CCl_{4}, NCl_{3}, and Cl_{2}O by one method, but we also arrive at the fact that the reactions of the metalepsis of the hydrocarbons lose that exclusiveness which was often ascribed to them. Also by subjecting the chemical representations to the law of substitution we may foretell metalepsis as a particular case of a general law.

The action of chlorine on ammonia may either result in the entire breaking up of the ammonia, with the evolution of gaseous nitrogen, or in a product of metalepsis (as with CH_{4}). With an excess of chlorine and the aid of heat the ammonia is decomposed, with the disengagement of free nitrogen.[28] This reaction evidently results in the formation of sal-ammoniac, 8NH_{3} + 3Cl_{2} = 6NH_{4}Cl + N_{2}. But if the ammonium salt be in excess, then the reaction takes the direction of the replacement of the hydrogen in the ammonia by chlorine. The principal result is that NH_{3} + 3Cl_{2} forms NCl_{3} + 3HCl.[29] The resulting product of metalepsis, or chloride of nitrogen, NCl_{3}, discovered by Dulong, is a liquid having the property of decomposing with excessive ease not only when heated, but even under the action of mechanical influences, as by a blow or by contact with certain solid substances. The explosion which accompanies the decomposition is due to the fact that the liquid chloride of nitrogen gives gaseous products, nitrogen and chlorine.[29 bis]

[28] This may be taken advantage of in the preparation of nitrogen. If a large excess of chlorine water be poured into a beaker, and a small quantity of a solution of ammonia be added, then, after shaking, nitrogen is evolved. If chlorine act on a dilute solution of ammonia, the volume of nitrogen does not correspond with the volume of the chlorine taken, because ammonium hypochlorite is formed. If ammonia gas be passed through a fine orifice into a vessel containing chlorine, the reaction of the formation of nitrogen is accompanied by the emission of light and the appearance of a cloud of sal-ammoniac. In all these instances an excess of chlorine must be present.

[29] The hydrochloric acid formed combines with ammonia, and therefore the final result is 4NH_{3} + 3Cl_{2} = NCl_{3} + 3NH_{4}Cl. For this reason, more ammonia must enter into the reaction, but the metalepsical reaction in reality only takes place with an excess of ammonia or its salt. If bubbles of chlorine be passed through a fine tube into a vessel containing ammonia gas, each bubble gives rise to an explosion. If, however, chlorine be passed into a solution of ammonia, the reaction at first brings about the formation of nitrogen, because chloride of nitrogen acts on ammonia like chlorine. But when sal-ammoniac has begun to form, then the reaction directs itself towards the formation of chloride of nitrogen. The first action of chlorine on a solution of sal-ammoniac always causes the formation of chloride of nitrogen, which then reacts on ammonia thus: NCl_{3} + 4NH_{3} = N_{2} + 3NH_{4}Cl. Therefore, so long as the liquid is alkaline from the presence of ammonia the chief product will be nitrogen. The reaction NH_{4}Cl + 3Cl_{2} = NCl_{3} + 4HCl is reversible; with a dilute solution it proceeds in the above-described direction (perhaps owing to the affinity of the hydrochloric acid for the excess of water), but with a strong solution of hydrochloric acid it takes the opposite direction (probably by virtue of the affinity of hydrochloric acid for ammonia). Therefore there must exist a very interesting case of equilibrium between ammonia, hydrochloric acid, chlorine, water, and chloride of nitrogen which has not yet been investigated. The reaction NCl_{3} + 4HCl = NH_{4}Cl + 3Cl_{2} enabled Deville and Hautefeuille to determine the composition of chloride of nitrogen. When slowly decomposed by water, chloride of nitrogen gives, like a chloranhydride, nitrous acid or its anhydride, 2NCl_{3} + 3H_{2}O = N_{2}O_{3} + 6HCl. From these observations it is evident that chloride of nitrogen presents great chemical interest, which is strengthened by its analogy with trichloride of phosphorus. The researches of F. F. Selivanoff (1891-94) prove that NCl_{3} may be regarded as an ammonium derivative of hypochlorous acid. Chloride of nitrogen is decomposed by dilute sulphuric acid in the following manner: NCl_{3} + 3H_{2}O + H_{2}SO_{4} = NH_{4}HSO_{4} + 3HClO. This reaction is reversible and is only complete when some substance, combining with HClO (for instance, succinimide) or decomposing it, is added to the liquid. This is easily understood from the fact that hypochlorous acid itself, HClO, may, according to the view held in this book, be regarded as the product of the metalepsis of water, and consequently bears the same relation to NCl_{3} as H_{2}O does to NH_{3}, or as RHO to RNH_{2}, R_{2}NH, and R_{3}N--that is to say, NCl_{3} corresponds as an ammonium derivative to ClOH and Cl_{2} in exactly the same manner as NR_{3} corresponds to ROH and R_{2}. The connection of NCl_{3} and other similar explosive chloro-nitrogen compounds (called chloryl compounds by Selivanoff; for example, the C_{2}H_{5}NCl_{2} of Wurtz is chloryl ethylamine), such as NRCl_{2} (as NC_{2}H_{5}Cl_{2}), and NR_{2}Cl (for instance, N(CH_{3}CO)HCl, chlorylacetamide, and N(C_{2}H_{5}){2}Cl, chloryl diethylamine) with HClO is evident from the fact that under certain circumstances these compounds give hypochlorous acid, with water, for instance, NR{2}Cl + H_{2}O = NR_{2}H + HClO, and frequently act (like NCl_{3} and HClO, or Cl_{2}) in an oxidising and chloridising manner. We may take chloryl succinimide, C_{2}H_{4}(CO){2}NCl for example. It was obtained by Bender by the action of HClO upon succinimide, C{2}H_{4}(CO){2}NH, and is decomposed by water with the re-formation of amide and HClO (the reaction is reversible). Selivanoff obtained, investigated, and classified many of the compounds NR{2}Cl and NRCl_{2}, where R is a residue of organic acids or alcohols, and showed their distinction from the chloranhydrides, and thus supplemented the history of chloride of nitrogen, which is the simplest of the amides containing chlorine, NR_{3}, where R is fully substituted by chlorine.

[29 bis] In preparing NCl_{3} every precaution must he used to guard against an explosion, and care should he taken that the NCl_{3} remains under a layer of water. Whenever an ammoniacal substance comes into contact with chlorine great care must be taken, because it may be a case of the formation of such products and a very dangerous explosion may ensue. The liquid product of the metalepsis of ammonia may be most safely prepared in the form of small drops by the action of a galvanic current on a slightly warm solution of sal-ammoniac; chlorine is then evolved at the positive pole, and this chlorine acting on the ammonia gradually forms the product of metalepsis which floats on the surface of the liquid (being carried up by the gas), and if a layer of turpentine be poured on to it these small drops, on coming into contact with the turpentine, give feeble explosions, which are in no way dangerous owing to the small mass of the substance formed. Drops of chloride of nitrogen may with great caution be collected for investigation in the following manner. The neck of a funnel is immersed in a basin containing mercury, and first a saturated solution of common salt is poured into the funnel, and above it a solution of sal-ammoniac in 9 parts of water. Chlorine is then slowly passed through the solutions, when drops of chloride of nitrogen fall into the salt water.

Chloride of nitrogen is a yellow oily liquid of sp. gr. 1·65, which boils at 71°, and breaks up into N + Cl_{3} at 97°. The contact of phosphorus, turpentine, india-rubber, &c. causes an explosion, which is sometimes so violent that a small drop will pierce through a thick board. The great ease with which chloride of nitrogen decomposes is dependent upon the fact that it is formed with an absorption of heat, which it evolves when decomposed, to the amount of about 38,000 heat units for NCl_{3}, as Deville and Hautefeuille determined.

Chlorine, when absorbed by a solution of caustic soda (and also of other alkalis) at the ordinary temperature, causes the replacement of the hydrogen in the caustic soda by the chlorine, with the formation of sodium chloride by the hydrochloric acid, so that the reaction may be represented in two phases, as described above. In this manner, sodium hypochlorite, NaClO, and sodium chloride are simultaneously formed: 2NaHO + Cl_{2} = NaCl + NaClO + H_{2}O. The resultant solution contains NaClO and is termed 'eau de Javelle.' An exactly similar reaction takes place when chlorine is passed over dry hydrate of lime at the ordinary temperature: 2Ca(HO){2} + 2Cl{2} = CaCl_{2}O_{2} + CaCl_{2} + 2H_{2}O. A mixture of the product of metalepsis with calcium chloride is obtained. This mixture is employed in practice on a large scale, and is termed 'bleaching powder,' owing to its acting, especially when mixed with acids, as a bleaching agent on tissues, so that it resembles chlorine in this respect. It is however preferable to chlorine, because the destructive action of the chlorine can be moderated in this case, and because it is much more convenient to deal with a solid substance than with gaseous chlorine. Bleaching powder is also called chloride of lime, because it is obtained from chlorine and hydrate of lime, and contains[30] both these substances. It may be prepared in the laboratory by passing a current of chlorine through a cold mixture of water and lime (milk of lime). The mixture must be kept cold, as otherwise 3Ca(ClO){2} passes into 2CaCl{2} + Ca(ClO_{3})_{2}. In the manufacture of bleaching powder in large quantities at chemical works, the purest possible slaked lime is taken and laid in a thin layer in large flat chambers, M (whose walls are made of Yorkshire flags or tarred wood, on which chlorine has no action), and into which chlorine gas is introduced by lead tubes. The distribution of the plant is shown in the annexed drawing (fig. 67).

[30] Quicklime, CaO (or calcium carbonate, CaCO_{3}), does not absorb chlorine when cold, but at a red heat, in a current of chlorine, it forms calcium chloride, with the evolution of oxygen. (This was confirmed in 1893 by Wells, at Oxford.) This reaction corresponds with the decomposing action of chlorine on methane, ammonia, and water. Slaked lime (calcium hydroxide, CaH_{2}O_{2}) also, when dry, does not absorb chlorine at 100°. The absorption proceeds at the ordinary temperature (below 40°). The dry mass thus obtained contains not less than three equivalents of calcium hydroxide to four equivalents of chlorine, so that its composition is [Ca(HO){2}]{5}Cl_{4}. In all probability a simple absorption of chlorine by the lime at first takes place in this case, as may be seen from the fact that even carbonic anhydride, when acting on the dry mass obtained as above, disengages all the chlorine from it, leaving only calcium carbonate. But if the bleaching powder be obtained by a wet method, or if it be dissolved in water (in which it is very soluble), and carbonic anhydride be passed into it, then chlorine is no longer disengaged, but chlorine oxide, Cl_{2}O, and only half of the chlorine is converted into this oxide, while the other half remains in the liquid as calcium chloride. From this it may be inferred that calcium chloride is formed by the action of water on bleaching powder, and this is proved to be the case by the fact that small quantities of water extract a considerable amount of calcium chloride from bleaching powder. If a large quantity of water act on bleaching powder an excess of calcium hydroxide remains, a portion of which is not subjected to change. The action of the water may be expressed by the following formulæ: From the dry mass Ca_{3}(HO){6}Cl{4} there is formed lime, Ca(HO){2}, calcium chloride, CaCl{2}, and a saline substance, Ca(ClO){2}. Ca{3}H_{6}O_{6}Cl_{4} = CaH_{2}O_{2} + CaCl_{2}O_{2} + CaCl_{2} + 2H_{2}O. The resulting substances are not equally soluble; water first extracts the calcium chloride, which is the most soluble, then the compound Ca(ClO){2} and ultimately calcium hydroxide is left. A mixture of calcium chloride and hypochlorite passes into solution. On evaporation there remains Ca{2}O_{2}Cl_{4}3H_{2}O. The dry bleaching powder does not absorb more chlorine, but the solution is able to absorb it in considerable quantity. If the liquid be boiled, a considerable amount of chlorine monoxide is evolved. After this calcium chloride alone remains in solution, and the decomposition may be expressed as follows: CaCl_{2} + CaCl_{2}O_{2} + 2Cl_{2} = 2CaCl_{2} + 2Cl_{2}O. Chlorine monoxide may be prepared in this manner.

It is sometimes said that bleaching powder contains a substance, Ca(OH){2}Cl{2}, that is calcium peroxide, CaO_{2}, in which one atom of oxygen is replaced by (OH){2}, and the other by Cl{2}; but, judging from what has been said above, this can only be the case in the dry state, and not in solutions.

On being kept for some time, bleaching powder sometimes decomposes, with the evolution of oxygen (because CaCl_{2}O_{2} = CaCl_{2} + O_{2}, see p. 163); the same takes place when it is heated.

The products of the metalepsis of alkaline hydrates, NaClO and Ca(ClO)_{2}, which are present in solutions of 'Javelle salt' and bleaching powder (they are not obtained free from metallic chlorides), must be counted as salts, because their metals are capable of substitution. But the hydrate HClO corresponding with these salts, or hypochlorous acid, is not obtained in a free or pure state, for two reasons: in the first place, because this hydrate, as a very feeble acid, splits up (like H_{2}CO_{3} or HNO_{3}) into water and the anhydride, or chlorine monoxide, Cl_{2}O = 2HClO-H_{2}O; and, in the second place, because, in a number of instances, it evolves oxygen with great facility, forming hydrochloric acid: HClO = HCl + O. Both hypochlorous acid and chlorine monoxide may be regarded as products of the metalepsis of water, because HOH corresponds with ClOH and ClOCl. Hence in many instances bleaching salts (a mixture of hypochlorites and chlorides) break up, with the evolution of (1) chlorine, under the action of an excess of a powerful acid capable of evolving hydrochloric acid from sodium or calcium chlorides, and this takes place most simply under the action of hydrochloric acid itself, because (p. 462) NaCl + NaClO + 3HCl = 2NaCl + HCl + Cl_{2} + H_{2}O; (2) oxygen, as we saw in Chapter III.--The bleaching properties and, in general, oxidising action of bleaching salts is based on this evolution of oxygen (or chlorine); oxygen is also disengaged on heating the dry salts--for instance, NaCl + NaClO = 2NaCl + O; (3) and, lastly, chlorine monoxide, which contains both chlorine and oxygen. Thus, if a little sulphuric, nitric, or similar acid (not enough to liberate hydrochloric acid from the CaCl_{2}) be added to a solution of a bleaching salt (which has an alkaline reaction, owing either to an excess of alkali or to the feeble acid properties of HClO), then the hypochlorous acid set free gives water and chlorine monoxide. If carbonic anhydride (or boracic or a similar very feeble acid) act on the solution of a bleaching salt, then hydrochloric acid is not evolved from the sodium or calcium chlorides, but the hypochlorous acid is displaced and gives chlorine monoxide,[31] because hypochlorous acid is one of the most feeble acids. Another method for the preparation of chlorine monoxide is based on these feeble acid properties of hypochlorous acid. Zinc oxide and mercury oxide, under the action of chlorine in the presence of water, do not give a salt of hypochlorous acid, but form a chloride and hypochlorous acid, which fact shows the incapacity of this acid to combine with the bases mentioned. Therefore, if such oxides as those of zinc or mercury be shaken up in water, and chlorine be passed through the turbid liquid,[32] a reaction occurs which may be expressed in the following manner: 2HgO + 2Cl_{2} = Hg_{2}OCl_{2} + Cl_{2}O. In this case, a compound of mercury oxide with mercury chloride, or the so-called mercury oxychloride, is obtained: Hg_{2}OCl_{2} = HgO + HgCl_{2}. This is insoluble in water, and is not affected by hypochlorous anhydride, so that the solution will contain hypochlorous acid only, but the greater part of it splits up into the anhydride and water.[32 bis]

[31] For this reason it is necessary that in the preparation of bleaching powder the chlorine should be free from hydrochloric acid, and even the lime from calcium chloride. An excess of chlorine, in acting on a solution of bleaching powder, may also give chlorine monoxide, because calcium carbonate also gives chlorine monoxide under the action of chlorine. This reaction may be brought about by treating freshly precipitated calcium carbonate with a stream of chlorine in water: 2Cl_{2} + CaCO_{3} = CO_{2} + CaCl_{2} + Cl_{2}O. From this we may conclude that, although carbonic anhydride displaces hypochlorous anhydride, it may be itself displaced by an excess of the latter.

[32] Dry red mercury oxide acts on chlorine, forming dry hypochlorous anhydride (chlorine monoxide) (Balard); when mixed with water, red mercury oxide acts feebly on chlorine, and when freshly precipitated it evolves oxygen and chlorine. An oxide of mercury which easily and abundantly evolves chlorine monoxide under the action of chlorine in the presence of water may be prepared as follows: the oxide of mercury, precipitated from a mercuric salt by an alkali, is heated to 300° and cooled (Pelouze). If a salt, MClO, be added to a solution of mercuric salt, HgX_{2}, mercuric oxide is liberated, because the hypochlorite is decomposed.

[32 bis] A solution of hypochlorous anhydride is also obtained by the action of chlorine on many salts; for example, in the action of chlorine on a solution of sodium sulphate the following reaction takes place: Na_{2}SO_{4} + H_{2}O + Cl_{2} = NaCl + HClO + NaHSO_{4}. Here the hypochlorous acid is formed, together with HCl, at the expense of chlorine and water, for Cl_{2} + H_{2}O = HCl + HClO. If the crystallo-hydrate of chlorine be mixed with mercury oxide, the hydrochloric acid formed in the reaction gives mercury chloride, and hypochlorous acid remains in solution. A dilute solution of hypochlorous acid or chlorine monoxide may be concentrated by distillation, and if a substance which takes up water (without destroying the acid)--for instance, calcium nitrate--be added to the stronger solution, then the anhydride of hypochlorous acid--i.e. chlorine monoxide--is disengaged.

Chlorine monoxide, which corresponds to bleaching and hypochlorous salts, containing as it does the two elements oxygen and chlorine, forms a characteristic example of a compound of elements which, in the majority of cases, act chemically in an analogous manner. Chlorine monoxide, as prepared from an aqueous solution by the abstraction of water or by the action of dry chlorine on cold mercury oxide, is, at the ordinary temperature, a gas or vapour which condenses into a red liquid boiling at +20° and giving a vapour whose density (43 referred to hydrogen) shows that 2 vols. of chlorine and 1 vol. of oxygen give 2 vols. of chlorine monoxide. In an anhydrous form the gas or liquid easily explodes, splitting up into chlorine and oxygen. This explosiveness is determined by the fact that heat is evolved in the decomposition to the amount of about 15,000 heat units for Cl_{2}O.[33] The explosion may even take place spontaneously, and also in the presence of many oxidisable substances (for instance, sulphur, organic compounds, &c.), but the solution, although unstable and showing a strong oxidising tendency, does not explode.[34] It is evident that the presence of hypochlorous acid, HClO, may be assumed in an aqueous solution of Cl_{2}O, since Cl_{2}O + H_{2}O = 2HClO.

[33] All explosive substances are of this kind--ozone, hydrogen peroxide, chloride of nitrogen, nitro-compounds, &c. Hence they cannot be formed directly from the elements or their simplest compounds, but, on the contrary, decompose into them. In a liquid state chlorine monoxide explodes even on contact with powdery substances, or when rapidly agitated--for instance, if a file be rasped over the vessel in which it is contained.

[34] A solution of chlorine monoxide, or hypochlorous acid, does not explode, owing to the presence of the mass of water. In dissolving, chlorine monoxide evolves about 9,000 heat units, so that its store of heat becomes less.

The capacity of hypochlorous acid (studied by Carius and others) for entering into combination with the unsaturated hydrocarbons is very often taken advantage of in organic chemistry. Thus its solution absorbs ethylene, forming the chlorhydrin C_{2}H_{4}ClOH.

The oxidising action of hypochlorous acid and its salts is not only applied to bleaching but also to many reactions of oxidation. Thus it converts the lower oxides of manganese into the peroxide.

Hypochlorous acid, its salts, and chlorine monoxide serve as a transition between hydrochloric acid, chlorides, and chlorine, and a whole series of compounds containing the same elements combined with a still greater quantity of oxygen. The higher oxides of chlorine, as their origin indicates, are closely connected with hypochlorous acid and its salts:

Cl_{2}, NaCl, HCl, hydrochloric acid. Cl_{2}O, NaClO, HClO, hypochlorous acid. Cl_{2}O_{3}, NaClO_{2}, HClO_{2}, chlorous acid.[35] Cl_{2}O_{5}, NaClO_{3}, HClO_{3}, chloric acid. Cl_{2}O_{7}, NaClO_{4}, HClO_{4}, perchloric acid.

When heated, solutions of hypochlorites undergo a remarkable change. Themselves so unstable, they, without any further addition, yield two fresh salts which are both much more stable; one contains more oxygen than MClO, the other contains none at all.

3MClO = MClO_{3} + 2MCl hypochlorite chlorate chloride

[35] Chlorous acid, HClO_{2} (according to the data given by Millon, Brandau, and others) in many respects resembles hypochlorous acid, HClO, whilst they both differ from chloric and perchloric acids in their degree of stability, which is expressed, for instance, in their bleaching properties; the two higher acids do not bleach, but both the lower ones do so (oxidise at the ordinary temperature). On the other hand, chlorous acid is analogous to nitrous acid, HNO_{2}. The anhydride of chlorous acid, Cl_{2}O_{3}, is not known in a pure state, but it probably occurs in admixture with chlorine dioxide, ClO_{2}, which is obtained by the action of nitric and sulphuric acids on a mixture of potassium chlorate with such reducing substances as nitric oxide, arsenious oxide, sugar, &c. All that is at present known is that pure chlorine dioxide ClO_{2} (see Notes 39-43) is gradually converted into a mixture of hypochlorous and chlorous acids under the action of water (and alkalis); that is, it acts like nitric peroxide, NO_{2} (giving HNO_{3} and HNO_{2}), or as a mixed anhydride, 2ClO_{2} + H_{2}O = HClO_{3} + HClO_{2}. The silver salt, AgClO_{2}, is sparingly soluble in water. The investigations of Garzarolli-Thurnlackh and others seem to show that the anhydride Cl_{2}O_{3} does not exist in a free state.

Part of the salt--namely, two-thirds of it--parts with its oxygen in order to oxidise the remaining third.[36] From an intermediate substance, RX, two extremes, R and RX_{3} are formed, just as nitrous anhydride splits up into nitric oxide and nitric anhydride (or nitric acid). The resulting salt, MClO_{3}, corresponds with chloric acid and potassium chlorate, KClO_{3}. It is evident that a similar salt may be obtained directly by the action of chlorine on an alkali if its solution be heated, because RClO will be first formed, and then RClO_{3}; for example, 6KHO + 3Cl_{2} = KClO_{3} + 5KCl + 3H_{2}O. Chlorates are so prepared; for instance, potassium chlorate, which is easily separated from potassium chloride, being sparingly soluble in cold water.[37]

[36] Hydrochloric acid, which is an example of compounds of this kind, is a saturated substance which does not combine directly with oxygen, but in which, nevertheless, a considerable quantity of oxygen may be inserted between the elements forming it. The same may be observed in a number of other cases. Thus oxygen may be added or inserted between the elements, sometimes in considerable quantities, in the saturated hydrocarbons; for instance, in C_{3}H_{8}, three atoms of oxygen produce an alcohol, glycerin or glycerol, C_{3}H_{5}(OH){3}. We shall meet with similar examples hereafter. This is generally explained by regarding oxygen as a bivalent element--that is, as capable of combining with two different elements, such as chlorine, hydrogen, &c. On the basis of this view, it may be inserted between each pair of combined elements; the oxygen will then be combined with one of the elements by one of its affinities and with the other element by its other affinity. This view does not, however, express the entire truth of the matter, even when applied to the compounds of chlorine. Hypochlorous acid, HOCl--that is, hydrochloric acid in which one atom of oxygen is inserted--is, as we have already seen, a substance of small stability; it might therefore be expected that on the addition of a fresh quantity of oxygen, a still less stable substance would be obtained, because, according to the above view, the chlorine and hydrogen, which form such a stable compound together, are then still further removed from each other. But it appears that chloric and perchloric acid, HClO{3} and HClO_{4}, are much more stable substances. Furthermore, the addition of oxygen has also its limit, it can only be added to a certain extent. If the above representation were true and not merely hypothetical, there would be no limit to the combination of oxygen, and the more it entered into one continuous chain the more unstable would be the resultant compound. But not more than four atoms of oxygen can be added to hydrogen sulphide, nor to hydrochloric acid, nor to hydrogen phosphide. This peculiarity must lie in the properties of oxygen itself; four atoms of oxygen seem to have the power of forming a kind of radicle which retains two or several atoms of various other substances--for example, chlorine and hydrogen, hydrogen and sulphur, sodium and manganese, phosphorus and metals, &c., forming comparatively stable compounds, NaClO_{4}, Na_{2}SO_{4}, NaMnO_{4}, Na_{3}PO_{4}, &c. See Chapter X. Note 1 and Chapter XV.

[37] If chlorine be passed through a cold solution of potash, a bleaching compound, potassium chloride and hypochlorite, KCl + KClO, is formed, but if it be passed through a hot solution potassium chlorate is formed. As this is sparingly soluble in water, it chokes the gas-conducting tube, which should therefore be widened out at the end.

Potassium chlorate is usually obtained on a large scale from calcium chlorate, which is prepared by passing chlorine (as long as it is absorbed) into water containing lime, the mixture being kept warm. A mixture of calcium chlorate and chloride is thus formed in the solution. Potassium chloride is then added to the warm solution, and on cooling a precipitate of potassium chlorate is formed as a substance which is sparingly soluble in cold water, especially in the presence of other salts. The double decomposition taking place is Ca(ClO_{3}){2} + 2KCl = CaCl{2} + 2KClO_{3}. On a small scale in the laboratory potassium chlorate is best prepared from a strong solution of bleaching powder by passing chlorine through it and then adding potassium chloride. KClO_{3} is always formed by the action of an electric current on a solution of KCl, especially at 80° (Häussermann and Naschold, 1894), so that this method is now used on a large scale.

Potassium chlorate crystallises easily in large colourless tabular crystals. Its solubility in 100 parts of water at 0° = 3 parts, 20° = 8 parts, 40° = 14 parts, 60° = 25 parts, 80° = 40 parts. For comparison we will cite the following figures showing the solubility of potassium chloride and perchlorate in 100 parts of water: potassium chloride at O° = 28 parts, 20° = 35 parts, 40° = 40 parts, 100° = 57 parts; potassium perchlorate at 0° about 1 part, 20° about 1-3/4 part, 100° about 18 parts. When heated, potassium chlorate melts (the melting point has been given as from 335°-376°; according to the latest determination by Carnelley, 359°) and decomposes with the evolution of oxygen, potassium perchlorate being at first formed, as will afterwards be described (see Note 47). A mixture of potassium chlorate and nitric and hydrochloric acids effects oxidation and chlorination in solutions. It deflagrates when thrown upon incandescent carbon, and when mixed with sulphur (1/3 by weight) it ignites it on being struck, in which case an explosion takes place. The same occurs with many metallic sulphides and organic substances. Such mixtures are also ignited by a drop of sulphuric acid. All these effects are due to the large amount of oxygen contained in potassium chlorate, and to the ease with which it is evolved. A mixture of two parts of potassium chlorate, one part of sugar, and one part of yellow prussiate of potash acts like gunpowder, but burns too rapidly, and therefore bursts the guns, and it also has a very strong oxidising action on their metal. The sodium salt, NaClO_{3}, is much more soluble than the potassium salt, and it is therefore more difficult to free it from sodium chloride, &c. The barium salt is also more soluble than the potassium salt; O° = 24 parts, 20° = 37 parts, 80° = 98 parts of salt per 100 of water.

If dilute sulphuric acid be added to a solution of potassium chlorate, chloric acid is liberated, but it cannot be separated by distillation, as it is decomposed in the process. To obtain the free acid, sulphuric acid must be added to a solution of barium chlorate.[38] The sulphuric acid gives a precipitate of barium sulphate, and free chloric acid remains in solution. The solution may be evaporated under the receiver of an air-pump. This solution is colourless, has no smell, and acts as a powerful acid (it neutralises sodium hydroxide, decomposes sodium carbonate, gives hydrogen with zinc, &c.); when heated above 40°, however, it decomposes, forming chlorine, oxygen, and perchloric acid: 4HClO_{3} = 2HClO_{4} + H_{2}O + Cl_{2} + O_{3}. In a concentrated condition the acid acts as an exceedingly energetic oxidiser, so that organic substances brought into contact with it burst into flame. Iodine, sulphurous acid, and similar oxidisable substances form higher oxidation products and reduce the chloric acid to hydrochloric acid. Hydrochloric acid gas gives chlorine with chloric acid (and consequently with KClO_{3} also) acting in the same manner as it acts on the lower acids: HClO_{3} + 5HCl = 3H_{2}O + 3Cl_{2}.

[38] Barium chlorate, Ba(ClO_{3}){2},H{2}O, is prepared in the following way: impure chloric acid is first prepared and saturated with baryta, and the barium salt purified by crystallisation. The impure free chloric acid is obtained by converting the potassium in potassium chlorate into an insoluble salt. This is done by adding tartaric or hydrofluosilicic acid to a solution of potassium chlorate, because potassium tartrate and potassium silicofluoride are very sparingly soluble in water. Chloric acid is easily soluble in water.

By cautiously acting on potassium chlorate with sulphuric acid, the dioxide (chloric peroxide), ClO_{2},[39] is obtained (Davy, Millon). This gas is easily liquefied in a freezing mixture, and boils at +10°. The vapour density (about 35 if H = 1) shows that the molecule of this substance is ClO_{2}.[40] In a gaseous or liquid state it very easily explodes (for instance, at 60°, or by contact with organic compounds or finely divided substances, &c.), forming Cl and O_{2}, and in many instances[41] therefore it acts as an oxidising agent, although (like nitric peroxide) it may itself be further oxidised.[42] In dissolving in water or alkalis chloric peroxide gives chlorous and hypochlorous acids--2ClO_{2} + 2KHO = KClO_{3} + KClO_{2} + H_{2}O--and therefore, like nitric peroxide, the dioxide may be regarded as an intermediate oxide between the (unknown) anhydrides of chlorous and chloric acids: 4ClO_{2} = Cl_{2}O_{3} + Cl_{2}O_{3}.[43]

[39] To prepare ClO_{2} 100 grams of sulphuric acid are cooled in a mixture of ice and salt, and 15 grams of powdered potassium chlorate are gradually added to the acid, which is then carefully distilled at 20° to 40°, the vapour given off being condensed in a freezing mixture. Potassium perchlorate is then formed: 3KClO_{3} + 2H_{2}SO_{4} = 2KHSO_{4} + KClO_{4} + 2ClO_{2} + H_{2}O. The reaction may result in an explosion. Calvert and Davies obtained chloric peroxide without the least danger by heating a mixture of oxalic acid and potassium chlorate in a test tube in a water-bath. In this case 2KClO_{3} + 3C_{2}H_{2}O_{4},2H_{2}O = 2C_{2}HKO_{4} + 2CO_{2} + 2ClO_{2} + 8H_{2}O. The reaction is still further facilitated by the addition of a small quantity of sulphuric acid. If a solution of HCl acts upon KClO_{3} at the ordinary temperature, a mixture of Cl_{2} and ClO_{2} is formed, but if the temperature be raised to 80° the greater part of the ClO_{2} decomposes, and when passed through a hot solution of MnCl_{2} it oxidises it. Gooch and Kreider proposed (1894) to employ this method for preparing small quantities of chlorine in the laboratory.

[40] By analogy with nitric peroxide it might be expected that at low temperatures a doubling of the molecule into Cl_{2}O_{4} would take place, as the reactions of ClO_{2} point to its being a mixed anhydride of HClO_{2} and HClO_{3}.

[41] Owing to the formation of this chlorine dioxide, a mixture of potassium chlorate and sugar is ignited by a drop of sulphuric acid. This property was formerly made use of for making matches, and is now sometimes employed for setting fire to explosive charges by means of an arrangement in which the acid is caused to fall on the mixture at the moment required. An interesting experiment on the combustion of phosphorus under water may be conducted with chlorine dioxide. Pieces of phosphorus and of potassium chlorate are placed under water, and sulphuric acid is poured on to them (through a long funnel); the phosphorus then burns at the expense of the chlorine dioxide.

[42] Potassium permanganate oxidises chlorine dioxide into chloric acid (Fürst).

[43] The euchlorine obtained by Davy by gently heating potassium chlorate with hydrochloric acid is (Pebal) a mixture of chlorine dioxide and free chlorine. The liquid and gaseous chlorine oxide (Note 35), which Millon considered to be Cl_{2}O_{3}, probably contains a mixture of ClO_{2} (vapour density 35), Cl_{2}O_{3} (whose vapour density should be 59), and chlorine (vapour density 35·5), since its vapour density was determined to be about 40.

As the salts of chloric acid, HClO_{3}, are produced by the splitting up of the salts of hypochlorous acid, so in the same way the salts of perchloric acid, HClO_{4}, are produced from the salts of chloric acid, HClO_{3}. But this is the highest form of the oxidation of HCl. Perchloric acid, HClO_{4}, is the most stable of all the acids of chlorine. When fused potassium chlorate begins to swell up and solidify, after having parted with one-third of its oxygen, potassium chloride and potassium perchlorate have been formed according to the equation 2KClO_{3} = KClO_{4} + KCl + O_{2}.

The formation of this salt is easily observed in the preparation of oxygen from potassium chlorate, owing to the fact that the potassium perchlorate fuses with greater difficulty than the chlorate, and therefore appears in the molten salt as solid grains (see Chapter III. Note 12). Under the action of certain acids--for instance, sulphuric and nitric--potassium chlorate also gives potassium perchlorate. This latter may be easily purified, because it is but sparingly soluble in water, although all the other salts of perchloric acid are very soluble and even deliquesce in the air. The perchlorates, although they contain more oxygen than the chlorates, are decomposed with greater difficulty, and even when thrown on ignited charcoal give a much feebler deflagration than the chlorates. Sulphuric acid (at a temperature not below 100°) evolves volatile and to a certain extent stable perchloric acid from potassium perchlorate. Neither sulphuric nor any other acid will further decompose perchloric acid as it decomposes chloric acid. Of all the acids of chlorine, perchloric acid alone can be distilled.[44] The pure hydrate HClO_{4}[45] is a colourless and exceedingly caustic substance which fumes in the air and has a specific gravity 1·78 at 15° (sometimes, after being kept for some time, it decomposes with a violent explosion). It explodes violently when brought into contact with charcoal, paper, wood, and other organic substances. If a small quantity of water be added to this hydrate, and it be cooled, a crystallo-hydrate, ClHO_{4},H_{2}O, separates out. This is much more stable, but the liquid hydrate HClO_{4},2H_{2}O is still more so. The acid dissolves in water in all proportions, and its solutions are distinguished for their stability.[46] When ignited both the acid and its salts are decomposed, with the evolution of oxygen.[47]

[44] If a solution of chloric acid, HClO_{3}, be first concentrated over sulphuric acid under the receiver of an air-pump and afterwards distilled, chlorine and oxygen are evolved and perchloric acid is formed: 4HClO_{3} = 2HClO_{4} + Cl_{2} + 3O + H_{2}O. Roscoe accordingly decomposed directly a solution of potassium chlorate by hydrofluosilicic acid, decanted it from the precipitate of potassium silicofluoride, K_{2}SiF_{6}, concentrated the solution of chloric acid, and then distilled it, perchloric acid being then obtained (see following footnote). That chloric acid is capable of passing into perchloric acid is also seen from the fact that potassium permanganate is decolorised, although slowly, by the action of a solution of chloric acid. On decomposing a solution of potassium chlorate by the action of an electric current, potassium perchlorate is obtained at the positive electrode (where the oxygen is evolved). Perchloric acid is also formed by the action of an electric current on solutions of chlorine and chlorine monoxide. Perchloric acid was obtained by Count Stadion and afterwards by Serullas, and was studied by Roscoe and others.

[45] Perchloric acid, which is obtained in a free state by the action of sulphuric acid on its salts, may be separated from a solution very easily by distillation, being volatile, although it is partially decomposed by distillation. The solution obtained after distillation may be concentrated by evaporation in open vessels. In the distillation the solution reaches a temperature of 200°, and then a very constant liquid hydrate of the composition HClO_{4},2H_{2}O is obtained in the distillate. If this hydrate be mixed with sulphuric acid, it begins to decompose at 100°, but nevertheless a portion of the acid passes over into the receiver without decomposing, forming a crystalline hydrate HClO_{4},H_{2}O which melts at 50°. On carefully heating this hydrate it breaks up into perchloric acid, which distills over below 100°, and into the liquid hydrate HClO_{4},2H_{2}O. The acid HClO_{4} may also be obtained by adding one-fourth part of strong sulphuric acid to potassium chlorate, carefully distilling and subjecting the crystals of the hydrate HClO_{4},H_{2}O obtained in the distillate to a fresh distillation. Perchloric acid, HClO_{4}, itself does not distil, and is decomposed on distillation until the more stable hydrate HClO_{4}, H_{2}O is formed; this decomposes into HClO_{4} and HClO_{4},2H_{2}O, which latter hydrate distils without decomposition. This forms an excellent example of the influence of water on stability, and of the property of chlorine of giving compounds of the type ClX_{7}, of which all the above hydrates, ClO_{3}(OH), ClO_{2}(OH){3}, and ClO(OH){5}, are members. Probably further research will lead to the discovery of a hydrate Cl(OH)_{7}.

[46] According to Roscoe the specific gravity of perchloric acid = 1·782 and of the hydrate HClO_{4},H_{2}O in a liquid state (50°) 1·811; hence a considerable contraction takes place in the combination of HClO_{4} with H_{2}O.

[47] The decomposition of salts analogous to potassium chlorate has been more fully studied in recent years by Potilitzin and P. Frankland. Professor Potilitzin, by decomposing, for example, lithium chlorate LiClO_{3}, found (from the quantity of lithium chloride and oxygen) that at first the decomposition of the fused salt (368°) takes place according to the equation, 3LiClO_{3} = 2LiCl + LiClO_{4} + 5O, and that towards the end the remaining salt is decomposed thus: 5LiClO_{3} = 4LiCl + LiClO_{4} + 10O. The phenomena observed by Potilitzin obliged him to admit that lithium perchlorate is capable of decomposing simultaneously with lithium chlorate, with the formation of the latter salt and oxygen; and this was confirmed by direct experiment, which showed that lithium chlorate is always formed in the decomposition of the perchlorate. Potilitzin drew particular attention to the fact that the decomposition of potassium chlorate and of salts analogous to it, although exothermal (Chapter III., Note 12), not only does not proceed spontaneously, but requires time and a rise of temperature in order to attain completion, which again shows that chemical equilibria are not determined by the heat effects of reactions only.

P. Frankland and J. Dingwall (1887) showed that at 448° (in the vapour of sulphur) a mixture of potassium chlorate and powdered glass is decomposed almost in accordance with the equation 2KClO_{3} = KClO_{4} + KCl + O_{2}, whilst the salt by itself evolves about half as much oxygen, in accordance with the equation, 8KClO_{3} = 5KClO_{4} + 3KCl + 2O_{2}. The decomposition of potassium perchlorate in admixture with manganese peroxide proceeds to completion, KClO_{4} = KCl + 2O_{2}. But in decomposing by itself the salt at first gives potassium chlorate, approximately according to the equation 7KClO_{4} = 2KClO_{3} + 5KCl + 11O_{2}. Thus there is now no doubt that when potassium chlorate is heated, the perchlorate is formed, and that this salt, in decomposing with evolution of oxygen, again gives the former salt.

In the decomposition of barium hypochlorite, 50 per cent. of the whole amount passes into chlorate, in the decomposition of strontium hypochlorite (Potilitzin, 1890) 12·5 per cent., and of calcium hypochlorite about 2·5 per cent. Besides which Potilitzin showed that the decomposition of the hypochlorites and also of the chlorates is always accompanied by the formation of a certain quantity of the oxides and by the evolution of chlorine, the chlorine being displaced by the oxygen disengaged. Spring and Prost (1889) represent the evolution of oxygen from KClO_{3} as due to the salt first splitting up into base and anhydride, thus (1) 2MClO_{3} = M_{2}O + Cl_{2}O_{5}; (2) Cl_{2}O_{5} = Cl_{2} + O_{3}; and (3) M_{2}O + Cl = 2MCl + O.

I may further remark that the decomposition of potassium chlorate as a reaction evolving heat easily lends itself for this very reason to the contact action of manganese peroxide and other similar admixtures; for such very feeble influences as those of contact may become evident either in those cases (for instance, detonating gas, hydrogen peroxide, &c.), when the reaction is accompanied by the evolution of heat, or when (for instance, H_{2} + I_{2}, &c.) little heat is absorbed or evolved. In these cases it is evident that the existing equilibrium is not very stable, and that a small alteration in the conditions at the surfaces of contact may suffice to upset it. In order to conceive the modus operandi of contact phenomena, it is enough to imagine, for instance, that at the surface of contact the movement of the atoms in the molecules changes from a circular to an elliptical path. Momentary and transitory compounds may he formed, but their formation cannot affect the explanation of the phenomena.

On comparing chlorine as an element not only with nitrogen and carbon but with all the other non-metallic elements (chlorine has so little analogy with the metals that a comparison with them would be superfluous), we find in it the following fundamental properties of the halogens or salt-producers. With metals chlorine gives salts (such as sodium chloride, &c.); with hydrogen a very energetic and monobasic acid HCl, and the same quantity of chlorine is able by metalepsis to replace the hydrogen; with oxygen it forms unstable oxides of an acid character. These properties of chlorine are possessed by three other elements, bromine, iodine, and fluorine. They are members of one natural family. Each representative has its peculiarities, its individual properties and points of distinction, in combination and in the free state--otherwise they would not be independent elements; but the repetition in all of them of the same chief characteristics of the family enables one more quickly to grasp all their various properties and to classify the elements themselves.

In order to have a guiding thread in forming comparisons between the elements, attention must however be turned not only to their points of resemblance but also to those of their properties and characters in which they differ most from each other. And the atomic weights of the elements must be considered as their most elementary property, since this is a quantity which is most firmly established, and must be taken account of in all the reactions of the element. The halogens have the following atomic weights--

F = 19, Cl = 35·5, Br = 80, I = 127.

All the properties, physical and chemical, of the elements and their corresponding compounds must evidently be in a certain dependence on this fundamental point, if the grouping in one family be natural.[47 bis] And we find in reality that, for instance, the properties of bromine, whose atomic weight is almost the mean between those of iodine and chlorine, occupy a mean position between those of these two elements. The second measurable property of the elements is their equivalence or their capacity for forming compounds of definite forms. Thus carbon or nitrogen in this respect differs widely from the halogens. Although the form ClO_{2} corresponds with NO_{2} and CO_{2}, yet the last is the highest oxide of carbon, whilst that of nitrogen is N_{2}O_{5}, and for chlorine, if there were an anhydride of perchloric acid, its composition would be Cl_{2}O_{7}, which is quite different from that of carbon. In respect to the forms of their compounds the halogens, like all elements of one family or group, are perfectly analogous to each other, as is seen from their hydrogen compounds:

HF, HCl, HBr, HI.

[47 bis] See, for example the melting point of NaCl, NaBr, NaI in Chapter II. Note 27. According to F. Freyer and V. Meyer (1892), the following are the boiling points of some of the corresponding compounds of chlorine and bromine:

BCl_{5} 17° BBr_{3} 90° SiCl_{3} 59° SiBr_{4} 153° PCl_{3} 76° PBr_{3} 175° SbCl_{3} 223° SbBr_{3} 275° BiCl_{3} 447° BiBr_{3} 453° SnCl_{4} 606° SnBr_{4} 619° ZnCl_{2} 730° ZnBr_{2} 650°

Thus for all the more volatile compounds the replacement of chlorine by bromine raises the boiling point, but in the ease of ZnX_{2} it lowers it (Chapter XV. Note 19).

Their oxygen compounds exhibit a similar analogy. Only fluorine does not give any oxygen compounds. The iodine and bromine compounds corresponding with HClO_{3} and HClO_{4} are HBrO_{3} and HBrO_{4}, HIO_{3} and HIO_{4}. On comparing the properties of these acids we can even predict that fluorine will not form any oxygen compound. For iodine is easily oxidised--for instance, by nitric acid--whilst chlorine is not directly oxidised. The oxygen acids of iodine are comparatively more stable than those of chlorine; and, generally speaking, the affinity of iodine for oxygen is much greater than that of chlorine. Here also bromine occupies an intermediate position. In fluorine we may therefore expect a still smaller affinity for oxygen than in chlorine--and up to now it has not been combined with oxygen. If any oxygen compounds of fluorine should be obtained, they will naturally be exceedingly unstable. The relation of these elements to hydrogen is the reverse of the above. Fluorine has so great an affinity for hydrogen that it decomposes water at the ordinary temperature; whilst iodine has so little affinity for hydrogen that hydriodic acid, HI, is formed with difficulty, is easily decomposed, and acts as a reducing agent in a number of cases.

From the form of their compounds the halogens are univalent elements with respect to hydrogen and septivalent with respect to oxygen, N being trivalent to hydrogen (it gives NH_{3}) and quinqui-valent to oxygen (it gives N_{2}O_{5}), and C being quadrivalent to both H and O as it forms CH_{4} and CO_{2}. And as not only their oxygen compounds, but also their hydrogen compounds, have acid properties, the halogens are elements of an exclusively acid character. Such metals as sodium, potassium, barium only give basic oxides. In the case of nitrogen, although it forms acid oxides, still in ammonia we find that capacity to give an alkali with hydrogen which indicates a less distinctly acid character than in the halogens. In no other elements is the acid-giving property so strongly developed as in the halogens.

In describing certain peculiarities characterising the halogens, we shall at every step encounter a confirmation of the above-mentioned general relations.

As fluorine decomposes water with the evolution of oxygen, F_{2} + H_{2}O = 2HF + O, for a long time all efforts to obtain it in free state by means of methods similar to those for the preparation of chlorine proved fruitless.[48] Thus by the action of hydrofluoric acid on manganese peroxide, or by decomposing a solution of hydrofluoric acid by an electric current, either oxygen or a mixture of oxygen and fluorine were obtained instead of fluorine. Probably a certain quantity of fluorine[48 bis] was set free by the action of oxygen or an electric current on incandescent and fused calcium fluoride, but at a high temperature fluorine acts even on platinum, and therefore it was not obtained. When chlorine acted on silver fluoride, AgF, in a vessel of natural fluor spar, CaF_{2}, fluorine was also liberated; but it was mixed with chlorine, and it was impossible to study the properties of the resultant gas. Brauner (1881) also obtained fluorine by igniting cerium fluoride, 2CeF_{4} = 2CeF_{3} + F_{2}; but this, like all preceding efforts, only showed fluorine to be a gas which decomposes water, and is capable of acting in a number of instances like chlorine, but gave no possibility of testing its properties. It was evident that it was necessary to avoid as far as possible the presence of water and a rise of temperature; this Moissan succeeded in doing in 1886. He decomposed anhydrous hydrofluoric acid, liquefied at a temperature of -23° and contained in a U-shaped tube (to which a small quantity of potassium fluoride had been added to make it a better conductor), by the action of a powerful electric current (twenty Bunsen's elements in series). Hydrogen was then evolved at the negative pole, and fluorine appeared at the positive pole (of iridium platinum) as a pale green gas which decomposed water with the formation of ozone and hydrofluoric acid, and combined directly with silicon (forming silicon fluoride, SiF_{4}), boron (forming BF_{3}), sulphur, &c. Its density (H = 1) is 18, so that its molecule is F_{2}. But the action of fluorine on metals at the ordinary temperature is comparatively feeble, because the metallic fluoride formed coats the remaining mass of the metals; it is, however, completely absorbed by iron. Hydrocarbons (such as naphtha), alcohol, &c., immediately absorb fluorine, with the formation of hydrofluoric acid. Fluorine when mixed with hydrogen can easily be made to explode violently, forming hydrofluoric acid.[49]

[48] Even before free fluorine was obtained (1886) it was evident from experience gained in the efforts made to obtain it, and from analogy, that it would decompose water (see first Russian edition of the Principles of Chemistry).

[48 bis] It is most likely that in this experiment of Fremy's, which corresponds with the action of oxygen on calcium chloride, fluorine was set free, but that a converse reaction also proceeded, CaO + F_{2} = CaF_{2} + O--that is, the calcium distributed itself between the oxygen and fluorine. MnF_{4}, which is capable of splitting up into MnF_{2} and F_{2}, is without doubt formed by the action of a strong solution of hydrofluoric acid on manganese peroxide, but under the action of water the fluorine gives hydrofluoric acid, and probably this is aided by the affinity of the manganese fluoride and hydrofluoric acid. In all the attempts made (by Davy, Knox, Louget, Fremy, Gore, and others) to decompose fluorides (those of lead, silver, calcium, and others) by chlorine, there were doubtless also cases of distribution, a portion of the metal combined with chlorine and a portion of the fluorine was evolved; but it is improbable that any decisive results were obtained. Fremy probably obtained fluorine, but not in a pure state.

[49] According to Moissan, fluorine is disengaged by the action of an electric current on fused hydrogen potassium fluoride, KHF_{2}. The present state of chemical knowledge is such that the knowledge of the properties of an element is much more general than the knowledge of the free element itself. It is useful and satisfactory to learn that even fluorine in the free state has not succeeded in eluding experiment and research, that the efforts to isolate it have been crowned with success, but the sum total of chemical data concerning fluorine as an element gains but little by this achievement. The gain will, however, be augmented if it be now possible to subject fluorine to a comparative study in relation to oxygen and chlorine. There is particular interest in the phenomena of the distribution of fluorine and oxygen, or fluorine and chlorine, competing under different conditions and relations. We may add that Moissan (1892) found that free fluorine decomposes H_{2}S, HCl, HBr, CS_{2}, and CNH with a flash; it does not act upon O_{2}, N_{2}, CO, and CO_{2}; Mg, Al, Ag, and Ni, when heated, burn in it, as also do S, Se, P (forms PF_{5}); it reacts upon H_{2} even in the dark, with the evolution of 366·00 units of heat. At a temperature of -95°, F_{2} still retains its gaseous state. Soot and carbon in general (but not the diamond) when heated in gaseous fluorine form fluoride of carbon, CF_{4} (Moissan, 1890); this compound is also formed at 300° by the double decomposition of CCl_{4} and AgF; it is a gas which liquefies at 10° under a pressure of 5 atmospheres. With an alcoholic solution of KHO, CF_{4} gives K_{2}CO_{3}, according to the equation CF_{4} + 6KHO = K_{2}CO_{3} + 4KF + 3H_{2}O. CF_{4} is not soluble in water, but it is easily soluble in CCl_{4} and alcohol.

In 1894 Brauner obtained fluorine directly by igniting the easily formed[49 bis] double lead salt HF,3KF,PbF_{4}, which first, at 230°, decomposes with the evolution of HF, and then splits up forming 3KF,PbF_{2} and fluorine F_{2}, which is recognised by the fact that it liberates iodine from KI and easily combines with silicon, forming SiF_{4}. This method gives chemically pure fluorine, and is based upon the breaking up of the higher compound--tetrafluoride of lead, PbF_{4}, corresponding to PbO_{2}, into free fluorine, F_{2}, and the lower more stable form--bifluoride of lead, PbF_{2}, which corresponds to PbO; that is, this method resembles the ordinary method of obtaining chlorine by means of MnO_{2}, as MnCl_{4} here breaks up into MnCl_{2} and chlorine, just as PbF_{4} splits up into PbF_{2} and fluorine.

[49 bis] T. Nikolukin (1885) and subsequently Friedrich and Classen obtained PbCl_{4} and a double ammonium salt of tetrachloride of lead (starting from the binoxide), PbCl_{4}2NH_{4}Cl; Hutchinson and Pallard obtained a similar salt of acetic acid (1893) corresponding to PbX_{4} by treating red lead with strong acetic acid; the composition of this salt is Pb(C_{2}H_{3}O_{2}){4}; it melts (and decomposes) at about 175°. Brauner (1894) obtained a salt corresponding to tetrafluoride of lead, PbF{4}, and the acid corresponding to it, H_{4}PbF_{8}. For example, by treating potassium plumbate (Chapter XVIII. Note 55) with strong HF, and also the above-mentioned tetra-acetate with a solution of KHF_{2}, Brauner obtained crystalline HK_{3}PbF_{8}--i.e. the salt from which he obtained fluorine.

Among the compounds of fluorine, calcium fluoride, CaF_{2}, is somewhat widely distributed in nature as fluor spar,[50] whilst cryolite, or aluminium sodium fluoride, Na_{3}AlF_{6}, is found more rarely (in large masses in Greenland). Cryolite, like fluor spar, is also insoluble in water, and gives hydrofluoric acid with sulphuric acid. Small quantities of fluorine have also in a number of cases been found in the bodies of animals, in the blood, urine, and bones. If fluorides occur in the bodies of animals, they must have been introduced in food, and must occur in plants and in water. And as a matter of fact river, and especially sea, water always contains a certain, although small, quantity of fluorine compounds.

[50] It is called spar because it very frequently occurs as crystals of a clearly laminar structure, and is therefore easily split up into pieces bounded by planes. It is called fluor spar because when used as a flux it renders ores fusible, owing to its reacting with silica, SiO_{2} + 2CaF_{2} = 2CaO + SiF_{4}; the silicon fluoride escapes as a gas and the lime combines with a further quantity of silica, and gives a vitreous slag. Fluor spar occurs in mineral veins and rocks, sometimes in considerable quantities. It always crystallises in the cubic system, sometimes in very large semi-transparent cubic crystals, which are colourless or of different colours. It is insoluble in water. It melts under the action of heat, and crystallises on cooling. The specific gravity is 3·1. When steam is passed over incandescent fluor spar, lime and hydrofluoric acid are formed: CaF_{2} + H_{2}O = CaO + 2HF. A double decomposition is also easily produced by fusing fluor spar with sodium or potassium hydroxides, or potash, or even with their carbonates; the fluorine then passes over to the potassium or sodium, and the oxygen to the calcium. In solutions--for example, Ca(NO_{3}){2} + 2KF = CaF{2} (precipitate) + 2KNO_{3} (in solution)--the formation of calcium fluoride takes place, owing to its very sparing solubility. 26,000 parts of water dissolve one part of fluor spar.

Hydrifluoric acid, HF, cannot be obtained from fluor spar in glass retorts, because glass is acted on by and destroys the acid. It is prepared in lead vessels, and when it is required pure, in platinum vessels, because lead also acts on hydrofluoric acid, although only very feebly on the surface, and when once a coating of fluoride and sulphate of lead is formed no further action takes place. Powdered fluor spar and sulphuric acid evolve hydrofluoric acid (which fumes in the air) even at the ordinary temperature, CaF_{2} + H_{2}SO_{4} = CaSO_{4} + 2HF. At 130° fluor spar is completely decomposed by sulphuric acid. The acid is then evolved as vapour, which may be condensed by a freezing mixture into an anhydrous acid. The condensation is aided by pouring water into the receiver of the condenser, as the acid is easily soluble in cold water.

In the liquid anhydrous form hydrofluoric acid boils at +19°, and its specific gravity at 12·8° = 0·9849.[51] It dissolves in water with the evolution of a considerable amount of heat, and gives a solution of constant boiling point which distils over at 120°; showing that the acid is able to combine with water. The specific gravity of the compound is 1·15, and its composition HF,2H_{2}O.[52] With an excess of water a dilute solution distils over first. The aqueous solution and the acid itself must be kept in platinum vessels, but the dilute acid may be conveniently preserved in vessels made of various organic materials, such as gutta-percha, or even in glass vessels having an interior coating of paraffin. Hydrofluoric acid does not act on hydrocarbons and many other substances, but it acts in a highly corrosive manner on metals, glass, porcelain, and the majority of rock substances.[53] It also attacks the skin, and is distinguished by its poisonous properties, so that in working with the acid a strong draught must be kept up, to prevent the possibility of the fumes being inhaled. The non-metals do not act on hydrofluoric acid, but all metals--with the exception of mercury, silver, gold, and platinum, and, to a certain degree, lead--decompose it with the evolution of hydrogen. With bases it gives directly metallic fluorides, and behaves in many respects like hydrochloric acid. There are, however, several distinct individual differences, which are furthermore much greater than those between hydrochloric, hydrobromic, and hydriodic acids. Thus the silver compounds of the latter are insoluble in water, whilst silver fluoride is soluble. Calcium fluoride, on the contrary, is insoluble in water, whilst calcium chloride, bromide, and iodide are not only soluble, but attract water with great energy. Neither hydrochloric, hydrobromic, nor hydriodic acid acts on sand and glass, whilst hydrofluoric acid corrodes them, forming gaseous silicon fluoride. The other halogen acids only form normal salts, KCl, NaCl, with Na or K, whilst hydrofluoric acid gives acid salts, for instance HKF_{2} (and by dissolving KF in liquid HF, KHF_{2}2HF is obtained). This latter property is in close connection with the fact that at the ordinary temperature the vapour density of hydrofluoric acid is nearly 20, which corresponds with a formula H_{2}F_{2}, as Mallet (1881) showed; but a depolymerisation occurs with a rise of temperature, and the density approaches 10, which answers to the formula HF.[54]

[51] According to Gore. Fremy obtained anhydrous hydrofluoric acid by decomposing lead fluoride at a red heat, by hydrogen, or by beating the double salt HKF_{2}, which easily crystallises (in cubes) from a solution of hydrofluoric acid, half of which has been saturated with potassium hydroxide. Its vapour density corresponds to the formula HF.

[52] This composition corresponds to the crystallo-hydrate HCl,2H_{2}O. All the properties of hydrofluoric acid recall those of hydrochloric acid, and therefore the comparative ease with which hydrofluoric acid is liquefied (it boils at +19°, hydrochloric acid at -35°) must be explained by a polymerisation taking place at low temperatures, as will be afterwards explained, H_{2}F_{2} being formed, and therefore in a liquid state it differs from hydrochloric acid, in which a phenomenon of a similar kind has not yet been observed.

[53] The corrosive action of hydrofluoric acid on glass and similar siliceous compounds is based upon the fact that it acts on silica, SiO_{2}, as we shall consider more fully in describing that compound, forming gaseous silicon fluoride, SiO_{2} + 4HF = SiF_{4} + 2H_{2}O. Silica, on the other hand, forms the binding (acid) element of glass and of the mass of mineral substances forming the salts of silica. When it is removed the cohesion is destroyed. This is made use of in the arts, and in the laboratory, for etching designs and scales, &c., on glass. In engraving on glass the surface is covered with a varnish composed of four parts of wax and one part of turpentine. This varnish is not acted on by hydrofluoric acid, and it is soft enough to allow of designs being drawn upon it whose lines lay bare the glass. The drawing is made with a steel point, and the glass is afterwards laid in a lead trough in which a mixture of fluor spar and sulphuric acid is placed. The sulphuric acid must be used in considerable excess, as otherwise transparent lines are obtained (owing to the formation of hydrofluosilicic acid). After being exposed for some time, the varnish is removed (melted) and the design drawn by the steel point is found reproduced in dull lines. The drawing may be also made by the direct application of a mixture of a silicofluoride and sulphuric acid, which forms hydrofluoric acid.

[54] Mallet (1881) determined the density at 30° and 100°, previous to which Gore (1869) had determined the vapour density at 100°, whilst Thorpe and Hambly (1888) made fourteen determinations between 26° and 88°, and showed that within this limit of temperature the density gradually diminishes, just like the vapour of acetic acid, nitrogen dioxide, and others. The tendency of HF to polymerise into H_{2}F_{2} is probably connected with the property of many fluorides of forming acid salts--for example, KHF_{2} and H_{2}SiF_{6}. We saw above that HCl has the same property (forming, for instance, H_{2}PtCl_{6}, &c., p. 457), and hence this property of hydrofluoric acid does not stand isolated from the properties of the other halogens.

The analogy between chlorine and the other two halogens, bromine and iodine, is much more perfect. Not only have their hydrates or halogen acids much in common, but they themselves resemble chlorine in many respects,[55] and even the properties of the corresponding metallic compounds of bromine and iodine are very much alike. Thus, the chlorides, bromides, and iodides of sodium and potassium crystallise in the cubic system, and are soluble in water; the chlorides of calcium, aluminium, magnesium, and barium are just as soluble in water as the bromides and iodides of these metals. The iodides and bromides of silver and lead are sparingly soluble in water, like the chlorides of these metals. The oxygen compounds of bromine and iodine also present a very strong analogy to the corresponding compounds of chlorine. A hypobromous acid is known corresponding with hypochlorous acid. The salts of this acid have the same bleaching property as the salts of hypochlorous acid. Iodine was discovered in 1811 by Courtois in kelp, and was shortly afterwards investigated by Clement, Gay-Lussac, and Davy. Bromine was discovered in 1826 by Balard in the mother liquor of sea water.

[55] For instance, the experiment with Dutch metal foil (Note 16) may be made with bromine just as well as with chlorine. A very instructive experiment on the direct combination of the halogens with metals maybe made by throwing a small piece (a shaving) of aluminium into a vessel containing liquid bromine; the aluminium, being lighter, floats on the bromine, and after a certain time reaction sets in accompanied by the evolution of heat, light, and fumes of bromine. The incandescent piece of metal moves rapidly over the surface of the bromine in which the resultant aluminium bromide dissolves. For the sake of comparison we will proceed to cite several thermochemical data (Thomsen) for analogous actions of (1) chlorine, (2) bromine, and (3) iodine, with respect to metals; the halogen being expressed by the symbol X, and the plus sign connecting the reacting substances. All the figures are given in thousands of calories, and refer to molecular quantities in grams and to the ordinary temperature:--

1 2 3 K_{2} + X_{2} 211 191 160 Na_{2} + X_{2} 195 172 138 Ag_{2} + X_{2} 59 45 28 Hg_{2} + X_{2} 83 68 48 Hg + X_{2} 63 51 34 Ca + X_{2} 170 141 -- Ba + X_{2} 195 170 -- Zn + X_{2} 97 76 49 Pb + X_{2} 83 64 40 Al + X_{2} 161 120 70

We may remark that the latent heat of vaporisation of the molecular weight Br_{2} is about 7·2, and of iodine 6·0 thousand heat units, whilst the latent heat of fusion of Br_{2} is about 0·3, and of I_{2} about 3·0 thousand heat units. From this it is evident that the difference between the amounts of heat evolved does not depend on the difference in physical state. For instance, the vapour of iodine in combining with Zn to form ZnI_{2} would give 48 + 8 + 3, or about sixty thousand heat units, or 1-1/2 times less than Zn + Cl_{2}.

Bromine and iodine, like chlorine, occur in sea water in combination with metals. However, the amount of bromides, and especially of iodides, in sea water is so small that their presence can only be discovered by means of sensitive reactions.[56] In the extraction of salt from sea water the bromides remain in the mother liquor. Iodine and bromine also occur combined with silver, in admixture with silver chloride, as a rare ore which is mainly found in America. Certain mineral waters (those of Kreuznach and Staro-rossüsk) contain metallic bromides and iodides, always in admixture with an excess of sodium chloride. Those upper strata of the Stassfurt rock salt (Chapter X.) which are a source of potassium salts also contain metallic bromides,[57] which collect in the mother liquors left after the crystallisation of the potassium salts; and this now forms the chief source (together with certain American springs) of the bromine in common use. Bromine may be easily liberated from a mixture of bromides and chlorides, owing to the fact that chlorine displaces bromine from its compounds with sodium, magnesium, calcium, &c. A colourless solution of bromides and chlorides turns an orange colour after the passage of chlorine, owing to the disengagement of bromine.[58] Bromine may be extracted on a large scale by a similar method, but it is simpler to add a small quantity of manganese peroxide and sulphuric acid to the mother liquid direct. This sets free a portion of the chlorine, and this chlorine liberates the bromine.

[56] One litre of sea-water contains about 20 grams of chlorine, and about 0·07 gram of bromine. The Dead Sea contains about ten times as much bromine.

[57] But there is no iodine in Stassfurt carnallite.

[58] The chlorine must not, however, be in large excess, as otherwise the bromine would contain chlorine. Commercial bromine not unfrequently contains chlorine, as bromine chloride; this is more soluble in water than bromine, from which it may thus be freed. To obtain pure bromine the commercial bromine is washed with water, dried by sulphuric acid, and distilled, the portion coming over at 58° being collected; the greater part is then converted into potassium bromide and dissolved, and the remainder is added to the solution in order to separate iodine, which is removed by shaking with carbon bisulphide. By heating the potassium bromide thus obtained with manganese peroxide and sulphuric acid, bromine is obtained quite free from iodine, which, however, is not present in certain kinds of commercial bromine (the Stassfurt, for instance). By treatment with potash, the bromine is then converted into a mixture of potassium bromide and bromate, and the mixture (which is in the proportion given in the equation) is distilled with sulphuric acid, bromine being then evolved: 5KBr + KBrO_{3} + 6H_{2}SO_{4} = 6KHSO_{4} + 3H_{2}O + 3Br_{2}. After dissolving the bromine in a strong solution of calcium bromide and precipitating with an excess of water, it loses all the chlorine it contained, because chlorine forms calcium chloride with CaBr_{2}.

Bromine is a dark brown liquid, giving brown fumes, and having a poisonous suffocating smell, whence its name (from the Greek [Greek: brômos], signifying evil smelling). The vapour density of bromine shows that its molecule is Br_{2}. In the cold bromine freezes into brown-grey scales like iodine. The melting point of pure bromine is -7°·05.[59] The density of liquid bromide at 0° is 3·187, and at 15° about 3·0. The boiling point of bromine is about 58°·7. Bromine, like chlorine, is soluble in water; 1 part of bromine at 5° requires 27 parts of water, and at 15° 29 parts of water. The aqueous solution of bromine is of an orange colour, and when cooled to -2° yields crystals containing 10 molecules of water to 1 molecule of bromine.[60] Alcohol dissolves a greater quantity of bromine, and ether a still greater amount. But after a certain time products of the action of the bromine on these organic substances are formed in the solutions. Aqueous solutions of the bromides also absorb a large amount of bromine.

[59] There has long existed a difference of opinion as to the melting point of pure bromine. By some investigators (Regnault, Pierre) it was given as between -7° and -8°, and by others (Balard, Liebig, Quincke, Baumhauer) as between -20° and -25°. There is now no doubt, thanks more especially to the researches of Ramsay and Young (1885), that pure bromine melts at about -7°. This figure is not only established by direct experiment (Van der Plaats confirmed it), but also by means of the determination of the vapour tensions. For solid bromine the vapour tension p in mm. at t was found to be--

p = 20 25 30 35 40 45 mm. t = -16°·6 -14° -12° -10° -8·5° -7°

For liquid bromine--

p = 50 100 200 400 600 760 mm. t = -5°·0 8°·2 23°·4 40°·4 51°·9 58°·7

These curves intersect at -7°·05. Besides which, in comparing the vapour tension of many liquids (for example, those given in Chapter II., Note 27), Ramsay and Young observed that the ratio of the absolute temperatures (t + 273) corresponding with equal tension varies for every pair of substances in rectilinear proportion in dependence upon t, and, therefore, for the above pressure p, Ramsay and Young determined the ratio of t + 273 for water and bromine, and found that the straight lines expressing these ratios for liquid and solid bromine intersect also at 7°·05; thus, for example, for solid bromine--

p = 20 25 30 35 40 45 273 + t = 256·4 259 261 263 264·6 266 273 + t´ = 295·3 299 302·1 304·8 307·2 309·3 c = 1·152 1·154 1·157 1·159 1·161 1·163

where t´ indicates the temperature of water corresponding with a vapour tension p, and where c is the ratio of 273 + t´ to 273 + t. The magnitude of c is evidently expressed with great accuracy by the straight line c = 1·1703 + 0·0011_t_. In exactly the same way we find the ratio for liquid bromine and water to be _c__{1} = 1·1585 + 0·00057t. The intersection of these straight lines in fact corresponds with -7°·06, which again confirms the melting point given above for bromine. In this manner it is possible with the existing store of data to accurately establish and verify the melting point of substances. Ramsay and Young established the thermal constants of iodine by exactly the same method.

[60] The observations made by Paterno and Nasini (by Raoult's method, Chapter I. Note 49) on the temperature of the formation of ice (-1°·115, with 1·391 gram of bromine in 100 grams of water) in an aqueous solution of bromine, showed that bromine is contained in solutions as the molecule Br_{2}. Similar experiments conducted on iodine (Kloboukoff 1889 and Beckmann 1890) show that in solution the molecule is I_{2}.

B. Roozeboom investigated the hydrate of bromine as completely as the hydrate of chlorine (Notes 9, 10). The temperature of the complete decomposition of the hydrate is +6°·2; the density of Br_{2},10H_{2}O = 1·49.

With respect to iodine, it is almost exclusively extracted from the mother liquors after the crystallisation of natural sodium nitrate (Chili saltpetre) and from the ashes of the sea-weed cast upon the shores of France, Great Britain, and Spain, sometimes in considerable quantities, by the high tides. The majority of these sea-weeds are of the genera Fucus, Laminaria, &c. The fused ashes of these sea-weeds are called 'kelp' in Scotland and 'varech' in Normandy. A somewhat considerable quantity of iodine is contained in these sea-weeds. After being burnt (or subjected to dry distillation) an ash is left which chiefly contains salts of potassium, sodium, and calcium. The metals occur in the sea-weed as salts of organic acids. On being burnt these organic salts are decomposed, forming carbonates of potassium and sodium. Hence, sodium carbonate is found in the ash of sea plants. The ash is dissolved in hot water, and on evaporation sodium carbonate and other salts separate, but a portion of the substances remains in solution. These mother liquors left after the separation of the sodium carbonate contain chlorine, bromine, and iodine in combination with metals, the chlorine and iodine being in excess of the bromine. 13,000 kilos of kelp give about 1,000 kilos of sodium carbonate and 15 kilos of iodine.

The liberation of the iodine from the mother liquor is effected with comparative ease, because chlorine disengages iodine from potassium iodide and its other combinations with the metals. Not only chlorine, but also sulphuric acid, liberates iodine from sodium iodide. Sulphuric acid, in acting on an iodide, sets hydriodic acid free, but the latter easily decomposes, especially in the presence of substances capable of evolving oxygen, such as chromic acid, nitrous acid, and even ferric salts.[61] Owing to its sparing solubility in water, the iodine liberated separates as a precipitate. To obtain pure iodine it is sufficient to distil it, and neglect the first and last portions of the distillate, the middle portion only being collected. Iodine passes directly from a state of vapour into a crystalline form, and settles on the cool portions of the apparatus in tabular crystals, having a black grey colour and metallic lustre.[62]

[61] In general, 2HI + O = I_{2} + H_{2}O, if the oxygen proceed from a substance from which it is easily evolved. For this reason compounds corresponding with the higher stages of oxidation or chlorination frequently give a lower stage when treated with hydriodic acid. Ferric oxide, Fe_{2}O_{3}, is a higher oxide, and ferrous oxide, FeO, a lower oxide; the former corresponds with FeX_{3}, and the latter with FeX_{2}, and this passage from the higher to the lower takes place under the action of hydriodic acid. Thus hydrogen peroxide and ozone (Chapter IV.) are able to liberate iodine from hydriodic acid. Compounds of copper oxide, CuO or CuX_{2}, give compounds of the suboxide Cu_{2}O, or CuX. Even sulphuric acid, which corresponds to the higher stage SO_{3}, is able to act thus, forming the lower oxide SO_{2}. The liberation of iodine from hydriodic acid proceeds with still greater ease under the action of substances capable of disengaging oxygen. In practice, many methods are employed for liberating iodine from acid liquids containing, for example, sulphuric acid and hydriodic acid. The higher oxides of nitrogen are most commonly used; they then pass into nitric oxide. Iodine may even be disengaged from hydriodic acid by the action of iodic acid, &c. But there is a limit in these reactions of the oxidation of hydriodic acid because, under certain conditions, especially in dilute solutions, the iodine set free is itself able to act as an oxidising agent--that is, it exhibits the character of chlorine, and of the halogens in general, to which we shall again have occasion to refer. In Chili, where a large quantity of iodine is extracted in the manufacture of Chili nitre, which contains NaIO_{3}, it is mixed with the acid and normal sulphites of sodium in solution; the iodine is then precipitated according to the equation 2NaIO_{3} + 3Na_{2}SO_{3} + 2NaHSO_{3} = 5Na_{2}SO_{4} +I_{2} + H_{2}O. The iodine thus obtained is purified by sublimation.

[62] For the final purification of iodine, Stas dissolved it in a strong solution of potassium iodide, and precipitated it by the addition of water (see Note 58).

The specific gravity of the crystals of iodine is 4·95. It melts at 114° and boils at 184°. Its vapour is formed at a much lower temperature, and is of a violet colour, whence iodine receives its name ([Greek: ioeidês], violet). The smell of iodine recalls the characteristic smell of hypochlorous acid; it has a sharp sour taste. It destroys the skin and organs of the body, and is therefore frequently employed for cauterising and as an irritant for the skin. In small quantities it turns the skin brown, but the coloration disappears after a certain time, partly owing to the volatility of the iodine. Water dissolves only 1/5000 part of iodine. A brown solution is thus obtained, which bleaches, but much more feebly than bromine and chlorine. Water which contains salts, and especially iodides, in solution dissolves iodine in considerable quantities, and the resultant solution is of a dark brown colour. Pure alcohol dissolves a small amount of iodine, and in so doing acquires a brown colour, but the solubility of iodine is considerably increased by the presence of a small quantity of an iodine compound--for instance, ethyl iodide--in the alcohol.[63] Ether dissolves a larger amount of iodine than alcohol; but iodine is particularly soluble in liquid hydrocarbons, in carbon bisulphide, and in chloroform. A small quantity of iodine dissolved in carbon bisulphide tints it rose-colour, but in a somewhat larger amount it gives a violet colour. Chloroform (quite free from alcohol) is also tinted rose colour by a small amount of iodine. This gives an easy means for detecting the presence of free iodine in small quantities. The blue coloration which free iodine gives with starch may also, as has already been frequently mentioned (see Chapter IV.), serve for the detection of iodine.

[63] The solubility of iodine in solutions containing iodides, and compounds of iodine in general, may serve, on the one hand, as an indication that solution is due to a similarity between the solvent and dissolved substance, and, on the other hand, as an indirect proof of that view as to solutions which was cited in Chapter I., because in many instances unstable highly iodised compounds, resembling crystallo-hydrates, have been obtained from such solutions. Thus iodide of tetramethylammonium, N(CH_{3}){4}I, combines with I{2}, and I_{4}. Even a solution of iodine in a saturated solution of potassium iodide presents indications of the formation of a definite compound KI_{3}. Thus, an alcoholic solution of KI_{3} does not give up iodine to carbon bisulphide, although this solvent takes up iodine from an alcoholic solution of iodine itself (Girault, Jörgensen, and others). The instability of these compounds resembles the instability of many crystallo-hydrates, for instance of HCl,2H_{2}O.

If we compare the four elements, fluorine, chlorine, bromine, and iodine, we see in them an example of analogous substances which arrange themselves by their physical properties in the same order as they stand in respect to their atomic and molecular weights. If the weight of the molecule be large, the substance has a higher specific gravity, a higher melting and boiling point, and a whole series of properties depending on this difference in its fundamental properties. Chlorine in a free state boils at about -35°, bromine boils at 60°, and iodine only above 180°. According to Avogadro-Gerhardt's law, the vapour densities of these elements in a gaseous state are proportional to their atomic weights, and here, at all events approximately, the densities in a liquid (or solid) state are also almost in the ratio of their atomic weights. Dividing the atomic weight of chlorine (35·5) by its specific gravity in a liquid state (1·3), we obtain a volume = 27, for bromine (80/3·1) 26, and for iodine also (127/4·9) 26.[64]

[64] The equality of the atomic volumes of the halogens themselves is all the more remarkable because in all the halogen compounds the volume augments with the substitution of fluorine by chlorine, bromine, and iodine. Thus, for example, the volume of sodium fluoride (obtained by dividing the weight expressed by its formula by its specific gravity) is about 15, of sodium chloride 27, of sodium bromide 32, and of sodium iodide 41. The volume of silicon chloroform, SiHCl_{3}, is 82, and those of the corresponding bromine and iodine compounds are 108 and 122 respectively. The same difference also exists in solutions; for example, NaCl + 200H_{2}O has a sp. gr. (at 15°/4°) of 1·0106, consequently the volume of the solution 3,658·5/1·0106 = 3,620, hence the volume of sodium chloride in solution = 3,620-3,603 (this is the volume of 200 H_{2}O) = 17, and in similar solutions, NaBr = 26 and NaI = 35.

The metallic bromides and iodides are in the majority of cases, in most respects analogous to the corresponding chlorides,[65] but chlorine displaces the bromine and iodine from them, and bromine liberates iodine from iodides, which is taken advantage of in the preparation of these halogens. However, the researches of Potilitzin showed that a reverse displacement of chlorine by bromine may occur both in solutions and in ignited metallic chlorides in an atmosphere of bromine vapour--that is, a distribution of the metal (according to Berthollet's doctrine) takes place between the halogens, although however the larger portion, still unites with the chlorine, which shows its greater affinity for metals as compared with that of bromine and iodine.[66] The latter, however, sometimes behave with respect to metallic oxides in exactly the same manner as chlorine. Gay-Lussac, by igniting potassium carbonate in iodine vapour, obtained (as with chlorine) an evolution of oxygen and carbonic anhydride, K_{2}CO_{3} + I_{2} = 2KI + CO_{2} + O, only the reactions between the halogens and oxygen are more easily reversible with bromine and iodine than with chlorine. Thus, at a red heat oxygen displaces iodine from barium iodide. Aluminium iodide burns in a current of oxygen (Deville and Troost), and a similar, although not so clearly marked, relation exists for aluminium chloride, and shows that the halogens have a distinctly smaller affinity for those metals which only form feeble bases. This is still more the case with the non-metals, which form acids and evolve much more heat with oxygen than with the halogens (Note 13). But in all these instances the affinity (and amount of heat evolved) of iodine and bromine is less than that of chlorine, probably because the atomic weights are greater. The smaller store of energy in iodine and bromine is seen still more clearly in the relation of the halogens to hydrogen. In a gaseous state they all enter, with more or less ease, into direct combination with gaseous hydrogen--for example, in the presence of spongy platinum, forming halogen acids, HX--but the latter are far from being equally stable; hydrogen chloride is the most stable, hydrogen iodide the least so, and hydrogen bromide occupies an intermediate position. A very high temperature is required to decompose hydrogen chloride even partially, whilst hydrogen iodide is decomposed by light even at the ordinary temperature and very easily by a red heat. Hence the reaction I_{2} + H_{2} = HI + HI is very easily reversible, and consequently has a limit, and hydrogen iodide easily dissociates.[67] Judging by the direct measurement of the heat evolved (22,000 heat units) in the formation of HCl, the conversion of 2HCl into H_{2} + Cl_{2} requires the expenditure of 44,000 heat units. The decomposition of 2HBr into H_{2} + Br_{2} only requires, if the bromine be obtained in a gaseous state, a consumption of about 24,000 units, whilst in the decomposition of 2HI into H_{2} + I_{2} as vapour about 3,000 heat units are evolved;[68] these facts, without doubt, stand in causal connection with the great stability of hydrogen chloride, the easy decomposability of hydrogen iodide, and the intermediate properties of hydrogen bromide. From this it would be expected that chlorine is capable of decomposing water with the evolution of oxygen, whilst iodine has not the energy to produce this disengagement,[69] although it is able to liberate the oxygen from the oxides of potassium and sodium, the affinity of these metals for the halogens being very considerable. For this reason oxygen, especially in compounds from which it can be evolved readily (for instance, ClHO, CrO_{3}, &c.), easily decomposes hydrogen iodide. A mixture of hydrogen iodide and oxygen burns in the presence of an ignited substance, forming water and iodine. Drops of nitric acid in an atmosphere of hydrogen iodide cause the disengagement of violet fumes of iodine and brown fumes of nitric peroxide. In the presence of alkalis and an excess of water, however, iodine is able to effect oxidation like chlorine--that is, it decomposes water; the action is here aided by the affinity of hydrogen iodide for the alkali and water, just as sulphuric acid helps zinc to decompose water. But the relative instability of hydriodic acid is best seen in comparing the acids in a gaseous state. If the halogen acids be dissolved in water, they evolve so much heat that they approach much nearer to each other in properties. This is seen from thermochemical data, for in the formation of HX in solution (in a large excess of water) from the gaseous elements there is evolved for HCl 39,000, for HBr 32,000, and for HI 18,000 heat units.[70] But it is especially evident from the fact that solutions of hydrogen bromide and iodide in water have many points in common with solutions of hydrogen chloride, both in their capacity to form hydrates and fuming solutions of constant boiling point, and in their capacity to form haloid salts, &c. by reacting on bases.

[65] But the density (and also molecular volume, Note 64) of a bromine compound is always greater than that of a chlorine compound, whilst that of an iodine compound is still greater. The order is the same in many other respects. For example, an iodine compound has a higher boiling point than a bromine compound, &c.

[66] A. L. Potilitzin showed that in heating various metallic chlorides in a closed tube, with an equivalent quantity of bromine, a distribution of the metal between the halogens always occurs, and that the amounts of chlorine replaced by the bromine in the ultimate product are proportional to the atomic weights of the metals taken and inversely proportional to their equivalence. Thus, if NaCl + Br be taken, then out of 100 parts of chlorine, 5·54 are replaced by the bromine, whilst with AgCl + Br 27·28 parts are replaced. These figures are in the ratio 1 : 4·9, and the atomic weights Na : Ag = 1 : 4·7. In general terms, if a chloride MCl_{n} be taken, it gives with _n_Br a percentage substitution = 4M/n^2 where M is the atomic weight of the metal. This law was deduced from observations on the chlorides of Li, K, Na, Ag (n = 1), Ca, Sr, Ba, Co, Ni, Hg, Pb (n = 2), Bi (n = 3), Sn (n = 4), and Fe_{2} (n = 6).

In these determinations of Potilitzin we see not only a brilliant confirmation of Berthollet's doctrine, but also the first effort to directly determine the affinities of elements by means of displacement. The chief object of these researches consisted in proving whether a displacement occurs in those cases where heat is absorbed, and in this instance it should be absorbed, because the formation of all metallic bromides is attended with the evolution of less heat than that of the chlorides, as is seen by the figures given in Note 55.

If the mass of the bromine be increased, then the amount of chlorine displaced also increases. For example, if masses of bromine of 1 and 4 equivalents act on a molecule of sodium chloride, then the percentages of the chlorine displaced will be 6·08 p.c. and 12·46 p.c.; in the action of 1, 4, 25, and 100 molecules of bromine on a molecule of barium chloride, there will be displaced 7·8, 17·6, 35·0, and 45·0 p.c. of chlorine. If an equivalent quantity of hydrochloric acid act on metallic bromides in closed tubes, and in the absence of water at a temperature of 300°, then the percentages of the substitution of the bromine by the chlorine in the double decomposition taking place between univalent metals are inversely proportional to their atomic weights. For example, NaBr + HCl gives at the limit 21 p.c. of displacement, KCl 12 p.c. and AgCl 4-1/4 p.c. Essentially the same action takes place in an aqueous solution, although the phenomenon is complicated by the participation of the water. The reactions proceed spontaneously in one or the other direction at the ordinary temperature but at different rates. In the action of a dilute solution (1 equivalent per 5 litres) of sodium chloride on silver bromide at the ordinary temperature the amount of bromine replaced in six and a half days is 2·07 p.c., and with potassium chloride 1·5 p.c. With an excess of the chloride the magnitude of the substitution increases. These conversions also proceed with the absorption of heat. The reverse reactions evolving heat proceed incomparably more rapidly, but also to a certain limit; for example, in the reaction AgCl + RBr the following percentages of silver bromide are formed in different times:

hours 2 3 22 96 120 K 79·82 87·4 88·22 -- 94·21 Na 83·63 90·74 91·70 95·49 --

That is, the conversions which are accompanied by an evolution of heat proceed with very much greater rapidity than the reverse conversions.

[67] The dissociation of hydriodic acid has been studied in detail by Hautefeuille and Lemoine, from whose researches we extract the following information. The decomposition of hydriodic acid is decided, but proceeds slowly at 180°; the rate and limit of decomposition increase with a rise of temperature. The reverse action--that is, I_{2} + H_{2} = 2HI--proceeds not only under the influence of spongy platinum (Corenwinder), which also accelerates the decomposition of hydriodic acid, but also by itself, although slowly. The limit of the reverse reaction remains the same with or without spongy platinum. An increase of pressure has a very powerful accelerative effect on the rate of formation of hydriodic acid, and therefore spongy platinum by condensing gases has the same effect as increase of pressure. At the atmospheric pressure the decomposition of hydriodic acid reaches the limit at 250° in several months, and at 440° in several hours. The limit at 250° is about 18 p.c. of decomposition--that is, out of 100 parts of hydrogen previously combined in hydriodic acid, about 18 p.c. may be disengaged at this temperature (this hydrogen may be easily measured, and the measure of dissociation determined), but not more; the limit at 440° is about 26 p.c. If the pressure under which 2HI passes into H_{2} + I_{2} be 4-1/2 atmospheres, then the limit is 24 p.c.; under a pressure of 1/5 atmosphere the limit is 29 p.c. The small influence of pressure on the dissociation of hydriodic acid (compared with N_{2}O_{4}, Chapter VI. Note 46) is due to the fact that the reaction 2HI = I_{2} + H_{2} is not accompanied by a change of volume. In order to show the influence of time, we will cite the following figures referring to 350°: (1) Reaction H_{2} + I_{2}; after 3 hours, 88 p.c. of hydrogen remained free; 8 hours, 69 p.c.; 34 hours, 48 p.c.; 76 hours, 29 p.c.; and 327 hours, 18·5 p.c. (2) The reverse decomposition of 2HI; after 9 hours, 3 p.c. of hydrogen was set free, and after 250 hours 18·6 p.c.--that is, the limit was reached. The addition of extraneous hydrogen diminishes the limit of the reaction of decomposition, or increases the formation of hydriodic acid from iodine and hydrogen, as would be expected from Berthollet's doctrine (Chapter X.). Thus at 440° 26 p.c. of hydriodic acid is decomposed if there be no admixture of hydrogen, while if H_{2} be added, then at the limit only half as large a mass of HI is decomposed. Therefore, if an infinite mass of hydrogen be added there will be no decomposition of the hydriodic acid. Light aids the decomposition of hydriodic acid very powerfully. At the ordinary temperature 80 p.c. is decomposed under the influence of light, whilst under the influence of heat alone this limit corresponds with a very high temperature. The distinct action of light, spongy platinum, and of impurities in glass (especially of sodium sulphate, which decomposes hydriodic acid), not only render the investigations difficult, but also show that in reactions like 2HI = I_{2} + H_{2}, which are accompanied by slight heat effects, all foreign and feeble influences may strongly affect the progress of the action (Note 47).

[68] The thermal determinations of Thomsen (at 18°) gave in thousands of calories, Cl + H = +22, HCl + Aq (that is, on dissolving HCl in a large amount of water) = +17·3, and therefore H + Cl + Aq = +39·3. In taking molecules, all these figures must be doubled. Br + H = +8·4; HBr + Aq = 19·9; H + Br + Aq = +28·3. According to Berthelot 7·2 are required for the vaporisation of Br_{2}, hence Br_{2} + H_{2} = 16·8 + 7·2 = +24, if Br_{2} be taken as vapour for comparison with Cl_{2}. H + I =-6·0, HI + Aq = 19·2; H + I + Aq= +13·2, and, according to Berthelot, the heat of fusion of I_{2} = 3·0, and of vaporisation 6·0 thousand heat units, and therefore I_{2} + H_{2} =-2(6·0) + 3 + 6 =-3·0, if the iodine be taken as vapour. Berthelot, on the basis of his determinations, gives, however, +0·8 thousand heat units. Similar contradictory results are often met with in thermochemistry owing to the imperfection of the existing methods, and particularly the necessity of depending on indirect methods for obtaining the fundamental figures. Thus Thomsen decomposed a dilute solution of potassium iodide by gaseous chlorine; the reaction gave +26·2, whence, having first determined the heat effects of the reactions KHO + HCl, KHO + HI and Cl + H in aqueous solutions, it was possible to find H + I + Aq; then, knowing HI + Aq, to find I + H. It is evident that unavoidable errors may accumulate.

[69] One can believe, however, on the basis of Berthollet's doctrine, and the observations of Potilitzin (Note 66), that a certain slow decomposition of water by iodine takes place. On this view the observations of Dossios and Weith on the fact that the solubility of iodine in water increases after the lapse of several months will be comprehensible. Hydriodic acid is then formed, and it increases the solubility. If the iodine be extracted from such a solution by carbon bisulphide, then, as the authors showed, after the action of nitrous anhydride iodine may be again detected in the solution by means of starch. It can easily be understood that a number of similar reactions, requiring much time and taking place in small quantities, have up to now eluded the attention of investigators, who even still doubt the universal application of Berthollet's doctrine, or only see the thermochemical side of reactions, or else neglect to pay attention to the element of time and the influence of mass.

[70] On the basis of the data in Note 68.

In consequence of what has been said above, it follows that hydrobromic and hydriodic acids, being substances which are but slightly stable, cannot be evolved in a gaseous state under many of those conditions under which hydrochloric acid is formed. Thus if sulphuric acid in solution acts on sodium iodide, all the same phenomena take place as with sodium chloride (a portion of the sodium iodide gives hydriodic acid, and all remains in solution), but if sodium iodide be mixed with strong sulphuric acid, then the oxygen of the latter decomposes the hydriodic acid set free, with liberation of iodine, H_{2}SO_{4} + 2HI = 2H_{2}O + SO_{2} + I_{2}. This reaction takes place in the reverse direction in the presence of a large quantity of water (2,000 parts of water per 1 part of SO_{2}), in which case not only the affinity of hydriodic acid for water is brought to light but also the action of water in directing chemical reactions in which it participates.[71] Therefore, with a halogen salt, it is easy to obtain gaseous hydrochloric acid by the action of sulphuric acid, but neither hydrobromic nor hydriodic acid can be so obtained in the free state (as gases).[72] Other methods have to be resorted to for their preparation, and recourse must not be had to compounds of oxygen, which are so easily able to destroy these acids. Therefore hydrogen sulphide, phosphorus, &c., which themselves easily take up oxygen, are introduced as means for the conversion of bromine and iodine into hydrobromic and hydriodic acids in the presence of water. For example, in the action of phosphorus the essence of the matter is that the oxygen of the water goes to the phosphorus, and the union of the remaining elements leads to the formation of hydrobromic or hydriodic acid; but the matter is complicated by the reversibility of the reaction, the affinity for water, and other circumstances which are understood by following Berthollet's doctrine. Chlorine (and bromine also) directly decomposes hydrogen sulphide, forming hydrochloric acid and liberating sulphur, both in a gaseous form and in solutions, whilst iodine only decomposes hydrogen sulphide in weak solutions, when its affinity for hydrogen is aided by the affinity of hydrogen iodide for water. In a gaseous state iodine does not act on hydrogen sulphide,[73] whilst sulphur is able to decompose gaseous hydriodic acid, forming hydrogen sulphide and a compound of sulphur and iodine which with water forms hydriodic acid.[74]

[71] A number of similar cases confirm what has been said in Chapter X.

[72] This is prevented by the reducibility of sulphuric acid. If volatile acids be taken they pass over, together with the hydrobromic and hydriodic acids, when distilled; whilst many non-volatile acids which are not reduced by hydrobromic and hydriodic acids only act feebly (like phosphoric acid), or do not act at all (like boric acid).

[73] This is in agreement with the thermochemical data, because if all the substances be taken in the gaseous state (for sulphur the heat of fusion is 0·3, and the heat of vaporisation 2·3) we have H_{2} + S = 4·7; H_{2} + Cl_{2} = 44; H_{2} + Br_{2} = 24, and H_{2} + I_{2} = -3 thousand heat units; hence the formation of H_{2}S gives less heat than that of HCl and HBr, but more than that of HI. In dilute solutions H_{2} + S + Aq = 9·3, and consequently less than the formation of all the halogen acids, as H_{2}S evolves but little heat with water, and therefore in dilute solutions chlorine, bromine, and iodine decompose hydrogen sulphide.

[74] Here there are three elements, hydrogen, sulphur, and iodine, each pair of which is able to form a compound, HI, H_{2}S, and SI, besides which the latter may unite in various proportions. The complexity of chemical mechanics is seen in such examples as these. It is evident that only the study of the simplest cases can give the key to the more complex problems, and on the other hand it is evident from the examples cited in the last pages that, without penetrating into the conditions of chemical equilibria, it would be impossible to explain chemical phenomena. By following the footsteps of Berthollet the possibility of unravelling the problems will be reached; but work in this direction has only been begun during the last ten years, and much remains to be done in collecting experimental material, for which occasions present themselves at every step. In speaking of the halogens I wished to turn the reader's attention to problems of this kind.

If hydrogen sulphide be passed through water containing iodine, the reaction H_{2}S + I_{2} = 2HI + S proceeds so long as the solution is dilute, but when the mass of free HI increases the reaction stops, because the iodine then passes into solution. A solution having a composition approximating to 2HI + 4I_{2} + 9H_{2}O (according to Bineau) does not react with H_{2}S, notwithstanding the quantity of free iodine. Therefore only weak solutions of hydriodic acid can be obtained by passing hydrogen sulphide into water with iodine.[74 bis]

[74 bis] The same essentially takes place when sulphurous anhydride, in a dilute solution, gives hydriodic acid and sulphuric acid with iodine. On concentration a reverse reaction takes place. The equilibrated systems and the part played by water are everywhere distinctly seen.

To obtain[75] gaseous hydrobromic and hydriodic acids it is most convenient to take advantage of the reactions between phosphorus, the halogens, and water, the latter being present in small quantity (otherwise the halogen acids formed are dissolved by it); the halogen is gradually added to the phosphorus moistened with water. Thus if red phosphorus be placed in a flask and moistened with water, and bromine be added drop by drop (from a tap funnel), hydrobromic acid is abundantly and uniformly disengaged.[76] Hydrogen iodide is prepared by adding 1 part of common (yellow) dry phosphorus to 10 parts of dry iodine in a glass flask. On shaking the flask, union proceeds quietly between them (light and heat being evolved), and when the mass of iodide of phosphorus which is formed has cooled, water is added drop by drop (from a tap funnel) and hydrogen iodide is evolved directly without the aid of heat. These methods of preparation will be at once understood when it is remembered (p. 468) that phosphorus chloride gives hydrogen chloride with water. It is exactly the same here--the oxygen of the water passes over to the phosphorus, and the hydrogen to the iodine, thus, PI_{3} + 3H_{2}O = PH_{3}O_{3} + 3HI.[77]

[75] Methods of formation and preparation are nothing more than particular cases of chemical reaction. If the knowledge of chemical mechanics were more exact and complete than it now is it would be possible to foretell all cases of preparation with every detail (of the quantity of water, temperature, pressure, mass, &c.) The study of practical methods of preparation is therefore one of the paths for the study of chemical mechanics. The reaction of iodine on phosphorus and water is a case like that mentioned in Note 74, and the matter is here further complicated by the possibility of the formation of the compound PH_{3} with HI, as well as the production of PI_{2}, PI_{3}, and the affinity of hydriodic acid and the acids of phosphorus for water. The theoretical interest of equilibria in all their complexity is naturally very great, but it falls into the background in presence of the primary interest of discovering practical methods for the isolation of substances, and the means of employing them for the requirements of man. It is only after the satisfaction of these requirements that interests of the other order arise, which in their turn must exert an influence on the former. For these reasons, whilst considering it opportune to point out the theoretical interest of chemical equilibria, the chief attention of the reader is directed in this work to questions of practical importance.

[76] Hydrobromic acid is also obtained by the action of bromine on paraffin heated to 180°. Gustavson proposed to prepare it by the action of bromine (best added in drops together with traces of aluminium bromide) on anthracene (a solid hydrocarbon from coal tar). Balard prepared it by passing bromine vapour over moist pieces of common phosphorus. The liquid tribromide of phosphorus, directly obtained from phosphorus and bromine, also gives hydrobromic acid when treated with water. Bromide of potassium or sodium, when treated with sulphuric acid in the presence of phosphorus, also gives hydrobromic acid, but hydriodic acid is decomposed by this method. In order to free hydrobromic acid from bromine vapour it is passed over moist phosphorus and dried either by phosphoric anhydride or calcium bromide (calcium chloride cannot be used, as hydrochloric acid would be formed). Neither hydrobromic nor hydriodic acids can be collected over mercury, on which they act, but they may be directly collected in a dry vessel by leading the gas-conducting tube to the bottom of the vessel, both gases being much heavier than air. Merz and Holtzmann (1889) proposed to prepare HBr directly from bromine and hydrogen. For this purpose pure dry hydrogen is passed through a flask containing boiling bromine. The mixture of gas and vapour then passes through a tube provided with one or two bulbs, which is heated moderately in the middle. Hydrobromic acid is formed with a series of flashes at the part heated. The resultant HBr, together with traces of bromine, passes into a Woulfe's bottle into which hydrogen is also introduced, and the mixture is then carried through another heated tube, after which it is passed through water which dissolves the hydrobromic acid. According to the method proposed by Newth (1892) a mixture of bromine and hydrogen is led through a tube containing a platinum spiral, which is heated to redness after the air has been displaced from the tube. If the vessel containing the bromine be kept at 60°, the hydrogen takes up almost the theoretical amount of bromine required for the formation of HBr. Although the flame which appears in the neighbourhood of the platinum spiral does not penetrate into the vessel containing the bromine, still, for safety, a tube filled with cotton wool may be interposed.

Hydroiodic acid is obtained in the same manner as hydrobromic. The iodine is heated in a small flask, and its vapour is carried over by hydrogen into a strongly heated tube, The gas passing from the tube is found to contain a considerable amount of HI, together with some free iodine. At a low red heat about 17 p.c. of the iodine vapour enters into combination; at a higher temperature, 78 p.c. to 79 p.c.; and at a strong heat about 82 p.c.

[77] But generally more phosphorus is taken than is required for the formation of PI_{3}, because otherwise a portion of the iodine distils over. If less than one-tenth part of iodine be taken, much phosphonium iodide, PH_{4}I, is formed. This proportion was established by Gay-Lussac and Kolbe. Hydriodic acid is also prepared in many other ways. Bannoff dissolves two parts of iodine in one part of a previously prepared strong (sp. gr. 1·67) solution of hydriodic acid, and pours it on to red phosphorus in a retort. Personne takes a mixture of fifteen parts of water, ten of iodine, and one of red phosphorus, which, when heated, disengages hydriodic acid mixed with iodine vapour; the latter is removed by passing it over moist phosphorus (Note 76). It must be remembered however that reverse reaction (Oppenheim) may take place between the hydriodic acid and phosphorus, in which the compounds PH_{4}I and PI_{2} are formed.

It should be observed that the reaction between phosphorus, iodine and water must be carried out in the above proportions and with caution, as they may react with explosion. With red phosphorus the reaction proceeds quietly, but nevertheless requires care.

L. Meyer showed that with an excess of iodine the reaction proceeds without the formation of bye-products (PH_{4}I), according to the equation P + 5I + 4H_{2}O = PH_{3}O_{4} + 5HI. For this purpose 100 grams of iodine and 10 grams of water are placed in a retort, and a paste of 5 grams of red phosphorus and 10 grams of water is added little by little (at first with great care). The hydriodic acid may be obtained free from iodine by directing the neck of the retort upwards and causing the gas to pass through a shallow layer of water (respecting the formation of HI, see also Note 75).

In a gaseous form hydrobromic and hydriodic acids are closely analogous to hydrochloric acid; they are liquefied by pressure and cold, they fume in the air, form solutions and hydrates, of constant boiling point, and react on metals, oxides and salts, &c.[78] Only the relatively easy decomposability of hydrobromic acid and especially of hydriodic acid, clearly distinguish these acids from hydrochloric acid. For this reason, hydriodic acid acts in a number of cases as a deoxidiser or reducer, and frequently even serves as a means for the transference of hydrogen. Thus Berthelot, Baeyer, Wreden, and others, by heating unsaturated hydrocarbons in a solution of hydriodic acid, obtained their compounds with hydrogen nearer to the limit C_{n}H_{2_n_ + 2} or even the saturated compounds. For example, benzene, C_{6}H_{6}, when heated in a closed tube with a strong solution of hydriodic acid, gives hexylene, C_{6}H_{12}. The easy decomposability of hydriodic acid accounts for the fact that iodine does not act by metalepsis on hydrocarbons, for the hydrogen iodide liberated with the product of metalepsis, RI, formed, gives iodine and the hydrogen compound, RH, back again. And therefore, to obtain the products of iodine substitution, either iodic acid, HIO_{3} (Kekulé), or mercury oxide, HgO (Weselsky), is added, as they immediately react on the hydrogen iodide, thus: HIO_{3} + 5HI = 3H_{2}O + 3I_{2}, or, HgO + 2HI = HgI_{2} + H_{2}O. From these considerations it will be readily understood that iodine acts like chlorine (or bromine) on ammonia and sodium hydroxide, for in these cases the hydriodic acid produced forms NH_{4}I and NaI. With tincture of iodine or even the solid element, a solution of ammonia immediately forms a highly-explosive solid black product of metalepsis, NHI_{2}, generally known as iodide of nitrogen, although it still contains hydrogen (this was proved beyond doubt by Szuhay 1893), which may be replaced by silver (with the formation of NAgI_{2}): 3NH_{3} + 2I_{2} = 2NH_{4}I + NHI_{2}. However, the composition of the last product is variable, and with an excess of water NI_{3} seems to be formed. Iodide of nitrogen is just as explosive as nitrogen chloride.[78 bis] In the action of iodine on sodium hydroxide no bleaching compound is formed (whilst bromine gives one), but a direct reaction is always accomplished with the formation of an iodate, 6NaHO + 3I_{2} = 5NaI + 3H_{2}O + NaIO_{3} (Gay-Lussac). Solutions of other alkalis, and even a mixture of water and oxide of mercury, act in the same manner.[79] This direct formation of iodic acid, HIO_{3} = IO_{2}(OH), shows the propensity of iodine to give compounds of the type IX_{5}. Indeed, this capacity of iodine to form compounds of a high type emphasises itself in many ways. But it is most important to turn attention to the fact that iodic acid is easily and directly formed by the action of oxidising substances on iodine. Thus, for instance, strong nitric acid directly converts iodine into iodic acid, whilst it has no oxidising action on chlorine.[79 bis] This shows a greater affinity in iodine for oxygen than in chlorine, and this conclusion is confirmed by the fact that iodine displaces chlorine from its oxygen acids,[80] and that in the presence of water chlorine oxidises iodine.[81] Even ozone or a silent discharge passed through a mixture of oxygen and iodine vapour is able to directly oxidise iodine[82] into iodic acid. It is disengaged from solutions as a hydrate, HIO_{3}, which loses water at 170°, and gives an anhydride, I_{2}O_{5}. Both these substances are crystalline (sp. gr. I_{2}O_{5} 5·037, HIO_{3} 4·869 at 0°), colourless and soluble in water;[83] both decompose at a red heat into iodine and oxygen, are in many cases powerfully oxidising--for instance, they oxidise sulphurous anhydride, hydrogen sulphide, carbonic oxide, &c.--form chloride of iodine and water with hydrochloric acid, and with bases form salts, not only normal MIO_{3}, but also acid; for example, KIO_{3}HIO_{3}, KIO_{3}2HIO_{3}.[83 bis] With hydriodic acid iodic acid immediately reacts, disengaging iodine, HIO_{3} + 5HI = 3H_{2}O + 3I_{2}.

[78] The specific gravities of their solutions as deduced by me on the basis of Topsöe and Berthelot's determinations for 15°/4° are as follows:--

10 20 30 40 50 60 p.c. HBr 1·071 1·156 1·258 1·374 1·505 1·650 HI 1·075 1·164 1·267 1·399 1·567 1·769

Hydrobromic acid forms two hydrates, HBr,2H_{2}O and HBr,H_{2}O, which have been studied by Roozeboom with as much completeness as the hydrate of hydrochloric acid (Chapter X. Note 37).

With metallic silver, solutions of hydriodic acid give hydrogen with great ease, forming silver iodide. Mercury, lead, and other metals act in a similar manner.

[78 bis] Iodide of nitrogen, NHI_{2} is obtained as a brown pulverulent precipitate on adding a solution of iodine (in alcohol, for instance) to a solution of ammonia. If it be collected on a filter-paper, it does not decompose so long as the precipitate is moist; but when dry it explodes violently, so that it can only be experimented upon in small quantities. Usually the filter-paper is torn into bits while moist, and the pieces laid upon a brick; on drying an explosion proceeds not only from friction or a blow, but even spontaneously. The more dilute the solution of ammonia, the greater is the amount of iodine required for the formation of the precipitate of NHI_{2}. A low temperature facilitates its formation. NHI_{2} dissolves in ammonia water, and when heated the solution forms HIO_{3} and iodine. With KI, iodide of nitrogen gives iodine, NH_{3} and KHO. These reactions (Selivanoff) are explained by the formation of HIO from NHI_{2} + 2H_{2}O = NH_{3} + 2HIO--and then KI + HIO = I_{2} + KHO. Selivanoff (see Note 29) usually observed a temporary formation of hypoiodous acid, HIO, in the reaction of ammonia upon iodine, so that here the formation of NHI_{2} is preceded by that of HIO--i.e. first I_{2} + H_{2}O = HIO + HI, and then not only the HI combines with NH_{3}, but also 2HIO + NH_{3} = NHI_{2} + 2H_{2}O. With dilute sulphuric acid iodide of nitrogen (like NCl_{3}) forms hypoiodous acid, but it immediately passes into iodic acid, as is expressed by the equation 5HIO = 2I_{2} + HIO_{3} + 2H_{2}O (first 3HIO = HIO_{3} + 2HI, and then HI + HIO = I_{2} + H_{2}O). Moreover, Selivanoff found that iodide of nitrogen, NHI_{2}, dissolves in an excess of ammonia water, and that with potassium iodide the solution gives the reaction for hypoiodous acid (the evolution of iodine in an alkaline solution). This shows that HIO participates in the formation and decomposition of NHI_{2}, and therefore the condition of the iodine (its metaleptic position) in them is analogous, and differs from the condition of the halogens in the haloid-anhydrides (for instance, NO_{2}Cl). The latter are tolerably stable, while (the haloid being designated by X) NHX_{2}, NX_{3}, XOH, RXO (see Chapter XIII. Note 43), &c., are unstable, easily decomposed with the evolution of heat, and, under the action of water, the haloid is easily replaced by hydrogen (Selivanoff), as would be expected in true products of metalepsis.

[79] Hypoiodous acid, HIO, is not known, but organic compounds, RIO, of this type are known. To illustrate the peculiarities of their properties we will mention one of these compounds, namely, iodosobenzol, C_{6}H_{5}IO. This substance was obtained by Willgerodt (1892), and also by V. Meyer, Wachter, and Askenasy, by the action of caustic alkalis upon phenoldiiodochloride, C_{6}H_{5}ICl_{2} (according to the equation, C_{6}H_{5}ICl_{2} + 2MOH = C_{6}H_{5}IO + 2MCl + H_{2}O). Iodosobenzol is an amorphous yellow substance, whose melting point could not be determined because it explodes at 210°, decomposing with the evolution of iodine vapour. This substance dissolves in hot water and alcohol, but is not soluble in the majority of other neutral organic solvents. If acids do not oxidise C_{6}H_{5}IO, they give saline compounds in which iodosobenzol appears as a basic oxide of a diatomic metal, C_{6}H_{5}I. Thus, for instance, when an acetic acid solution of iodosobenzol is treated with a solution of nitric acid, it gives large monoclinic crystals of a nitric acid salt having the composition C_{6}H_{5}I(NO_{3}){2} (like Ca(NO{3})_{2}). In appearing as the analogue of basic oxides, iodosobenzol displaces iodine from potassium iodide (in a solution acidulated with acetic or hydrochloric acid)--i.e. it acts with its oxygen like HClO. The action of peroxide of hydrogen, chromic acid, and other similar oxidising agents gives iodoxybenzol, C_{6}H_{5}IO_{2}, which is a neutral substance--i.e. incapable of giving salts with acids (compare Chapter XIII. Note 43).

[79 bis] The oxidation of iodine by strong nitric acid was discovered by Connell; Millon showed that it is effected, although more slowly, by the action of the hydrates of nitric acid up to HNO_{3},H_{2}O, but that the solution HNO_{3},2H_{2}O, and weaker solutions, do not oxidise, but simply dissolve, iodine. The participation of water in reactions is seen in this instance. It is also seen, for example, in the fact that dry ammonia combines directly with iodine--for instance, at 0° forming the compound I_{2},4NH_{3}--whilst iodide of nitrogen is only formed in presence of water.

[80] Bromine also displaces chlorine--for instance, from chloric acid, directly forming bromic acid. If a solution of potassium chlorate be taken (75 parts per 400 parts of water), and iodine be added to it (80 parts), and then a small quantity of nitric acid, chlorine is disengaged on boiling, and potassium iodate is formed in the solution. In this instance the nitric acid first evolves a certain portion of the chloric acid, and the latter, with the iodine, evolves chlorine. The iodic acid thus formed acts on a further quantity of the potassium chlorate, sets a portion of the chloric acid free, and in this manner the action is kept up. Potilitzin (1887) remarked, however, that not only do bromine and iodine displace the chlorine from chloric acid and potassium chlorate, but also chlorine displaces bromine from sodium bromate, and, furthermore, the reaction does not proceed as a direct substitution of the halogens, but is accompanied by the formation of free acids; for example, 5NaClO_{3} + 3Br_{2} + 3H_{2}O = 5NaBr + 5HClO_{3} + HBrO_{3}.

[81] If iodine be stirred up in water, and chlorine passed through the mixture, the iodine is dissolved; the liquid becomes colourless, and contains, according to the relative amounts of water and chlorine, either IHCl_{2}, or ICl_{3}, or HIO_{3}. If there be a small amount of water, then the iodic acid may separate out directly as crystals, but a complete conversion (Bornemann) only occurs when not less than ten parts of water are taken to one part of iodine--ICl + 3H_{2}O + 2Cl_{2} = IHO_{3} + 5HCl.

[82] Schönbein and Ogier proved this. Ogier found that at 45° ozone immediately oxidises iodine vapour, forming first of all the oxide I_{2}O_{3}, which is decomposed by water or on heating into iodic anhydride and iodine. Iodic acid is formed at the positive pole when a solution of hydriodic acid is decomposed by a galvanic current (Riche). It is also formed in the combustion of hydrogen mixed with a small quantity of hydriodic acid (Salet).

[83] Kämmerer showed that a solution of sp. gr. 2·127 at 14°, containing 2HIO_{3},9H_{2}O, solidified completely in the cold. On comparing solutions HI + _m_H_{2}O with HIO_{3} + _m_H_{2}O, we find that the specific gravity increases but the volume decreases, whilst in the passage of solutions HCl + _m_H_{2}O to HClO_{3} + _m_H_{2}O both the specific gravity and the volume increase, which is also observed in certain other cases (for example, H_{3}PO_{3} and H_{3}PO_{4}).

[83 bis] Ditte (1890) obtained many iodates of great variety. A neutral salt, 2(LiIO_{3})H_{2}O, is obtained by saturating a solution of lithia with iodic acid. There is an analogous ammonium salt, 2(NH_{4}IO_{3})H_{2}O. He also obtained hydrates of a more complex composition, such as 6(NH_{4}IO_{3})H_{2}O and 6(NH_{4}IO_{3})2H_{2}O. Salts of the alkaline earths, Ba(IO_{3}){2}H{2}O and Sr(IO_{3}){2}H{2}O, may be obtained by a reaction of double decomposition from the normal salts of the type 2(MeIO_{3})H_{2}O. When evaporated at 70° to 80° with nitric acid these salts lose water. A mixture of solutions of nitrate of zinc and an alkaline iodate precipitates Zn(IO_{3}){2}2H{2}O. An anhydrous salt is thrown out if nitric acid be added to the solutions. Analogous salts of cadmium, silver, and copper give compounds of the type 2Me´IO_{3}4NH_{3} and Me´´(IO_{3}){2}4NH{3}, with gaseous ammonia (Me´ and Me´´ being elements of the first (Ag) and second (Cd, Zn, Cu) groups). With an aqueous solution of ammonia the above salts give substances of a different composition, such as Zn(IO_{3}){2}(NH{4}){2}O, Cd(IO{3}){2}(NH{4}){2}O. Copper gives Cu(IO{3}){2}4(NH{4}){2}O and Cu(IO{3}){2}(NH{4}){2}O. These salts may be regarded as compounds of I{2}O_{5}, and MeO and (NH_{4}){2}O; for example, Zn(IO{3}){2}(NH{4}){2}O may be regarded as ZnO(NH{4}){2}OI{2}O_{5}, or, as derived from the hydrate, I_{2}O_{5}2H_{2}O = 2(HIO_{3})H_{2}O.

As with chlorine, so with iodine, a periodic acid, HIO_{4}, is formed. This acid is produced in the form of its salts, by the action of chlorine on alkaline solutions of iodates, and also by the action of iodine on chloric acid.[84] It crystallises from solutions as a hydrate containing 2H_{2}O (corresponding with HClO_{4},2H_{2}O), but as it forms salts containing up to 5 atoms of metals, this water must be counted as water of constitution. Therefore IO(OH){5} = HIO{4},2H_{2}O corresponds with the highest form of halogen compounds, IX_{7}.[85] In decomposing (at 200°) or acting as an oxidiser, periodic acid first gives iodic acid, but it may also be ultimately decomposed.

[84] If sodium iodate be mixed with a solution of sodium hydroxide, heated, and chlorine passed through the solution, a sparingly soluble salt separates out, which corresponds with periodic acid, and has the composition Na_{4}I_{2}O_{9},3H_{2}O.

6NaHO + 2NaIO_{3} + 4Cl = 4NaCl + Na_{4}I_{2}O_{9} + 3H_{2}O.

This compound is sparingly soluble in water, but dissolves easily in a very dilute solution of nitric acid. If silver nitrate be added to this solution a precipitate is formed which contains the corresponding compound of silver, Ag_{4}I_{2}O_{9},3H_{2}O. If this sparingly soluble silver compound be dissolved in hot nitric acid, orange crystals of a salt having the composition AgIO_{4} separate on evaporation. This salt is formed from the preceding by the nitric acid taking up silver oxide--Ag_{4}I_{2}O_{9} + 2HNO_{3} = 2AgNO_3 + 2AgIO_{4} + H_{2}O. The silver salt is decomposed by water, with the re-formation of the preceding salt, whilst iodic acid remains in solution--

4AgIO_{4} + H_{2}O = Ag_{4}I_{2}O_{9} + 2HIO_{4}.

The structure of the first of these salts, Na_{4}I_{2}O_{9},3H_{2}O, presents itself in a simpler form if the water of crystallisation is regarded as an integral portion of the salt; the formula is then divided in two, and takes the form of IO(OH){3}(ONa){2}--that is, it answers to the type IOX_{5}, or IX_{7}, like AgIO_{4} which is IO_{3}(OAg). The composition of all the salts of periodic acids are expressed by this type IX_{7}. Kimmins (1889) refers all the salts of periodic acid to four types--the meta-salts of HIO_{4} (salts of Ag, Cu, Pb), the meso-salts of H_{3}IO_{5} (PbH, Ag_{2}H, CdH), the para-salts of H_{5}IO_{6} (Na_{2}H_{3}, Na_{3}H_{2}), and the di-salts of H_{4}I_{2}O_{9} (K_{4}, Ag_{4}, Ni_{2}). The three first are direct compounds of the type IX_{7}, namely, IO_{3}(OH), IO_{2}(OH){3}, and IO(OH){5}, and the last are types of diperiodic salts, which correspond with the type of the meso-salts, as pyrophosphoric salts correspond with orthophosphoric salts--i.e. 2H_{3}IO_{5}-H_{2}O = H_{4}I_{2}O_{9}.

[85] Periodic acid, discovered by Magnus and Ammermüller, and whose salts were afterwards studied by Langlois, Rammelsberg, and many others, presents an example of hydrates in which it is evident that there is not that distinction between the water of hydration and of crystallisation which was at first considered to be so clear. In HClO,2H_{2}O the water, 2H_{2}O, is not displaced by bases, and must be regarded as water of crystallisation, whilst in HIO_{4},2H_{2}O it must be regarded as water of hydration. We shall afterwards see that the system of the elements obliges us to consider the halogens as substances giving a highest saline type, _GX__{7}, where G signifies a halogen, and X oxygen (O = _X__{2}), OH, and other like elements. The hydrate IO(OH)_{5} corresponding with many of the salts of periodic acid (for example, the salts of barium, strontium, mercury) does not exhaust all the possible forms. It is evident that various other pyro-, meta-, &c., forms are possible by the loss of water, as will be more fully explained in speaking of phosphoric acid, and as was pointed out in the preceding note.

Compounds formed between chlorine and iodine must be classed among the most interesting halogen bodies.[86] These elements combine together directly with evolution of heat, and form iodine monochloride, ICl, or iodine trichloride, ICl_{3}.[87] As water reacts on these substances, forming iodic acid and iodine, they have to be prepared from dry iodine and chlorine.[88] Both substances are formed in a number of reactions; for example, by the action of aqua regia on iodine, of chlorine on hydriodic acid, of hydrochloric acid on periodic acid, of iodine on potassium chlorate (with the aid of heat, &c.) Trapp obtained iodine monochloride, in beautiful red crystals, by passing a rapid current of chlorine into molten iodine. The monochloride then distils over and solidifies, melting at 27°. By passing chlorine over the crystals of the monochloride, it is easy to obtain iodine trichloride in orange crystals, which melt at 34° and volatilise at 47°, but in so doing decompose (into Cl_{2} and ClI). The chemical properties of these chlorides entirely resemble those of chlorine and iodine, as would be expected, because, in this instance, a combination of similar substances has taken place as in the formation of solutions or alloys. Thus, for instance, the unsaturated hydrocarbons (for example, C_{2}H_{4}), which are capable of directly combining with chlorine and iodine, also directly combine with iodine monochloride.

[86] With respect to hydrogen, oxygen, chlorine, and other elements, bromine occupies an intermediate position between chlorine and iodine, and therefore there is no particular need for considering at length the compounds of bromine. This is the great advantage of a natural grouping of the elements.

[87] They were both obtained by Gay-Lussac and many others. Recent data respecting iodine monochloride, ICl, entirely confirm the numerous observations of Trapp (1854), and even confirm his statement as to the existence of two isomeric (liquid and crystalline) forms (Stortenbeker). With a small excess of iodine, iodine monochloride remains liquid, but in the presence of traces of iodine trichloride it easily crystallises. Tanatar (1893) showed that of the two modifications of ICl, one is stable, and melts at 27°; while the other, which easily passes into the first, and is formed in the absence of ICl_{3}, melts at 14°. Schützenberger amplified the data concerning the action of water on the chlorides (Note 88), and Christomanos gave the fullest data regarding the trichloride.

After being kept for some time, the liquid monochloride of iodine yields red deliquescent octahedra, having the composition ICl_{4}, which are therefore formed from the monochloride with the liberation of free iodine, which dissolves in the remaining quantity of the monochloride. This substance, however, judging by certain observations, is impure iodine trichloride. If 1 part of iodine be stirred up in 20 parts of water, and chlorine be passed through the liquid, then all the iodine is dissolved, and a colourless liquid is ultimately obtained which contains a certain proportion of chlorine, because this compound gives a metallic chloride and iodate with alkalis without evolving any free iodine: ICl_{5} + 6KHO = 5KCl + KIO_{3} + 3H_{2}O. The existence of a pentachloride ICl_{5} is, however, denied, because this substance has not been obtained in a free state.

Stortenbeker (1888) investigated the equilibrium of the system containing the molecules I_{2}, ICl, ICl_{3}, and Cl_{2}, in the same way that Roozeboom (Chapter X. Note 38) examined the equilibrium of the molecules HCl, HCl,2H_{2}O, and H_{2}O. He found that iodine monochloride appears in two states, one (the ordinary) is stable and melts at 27°·2, whilst the other is obtained by rapid cooling, and melts at 13°·9, and easily passes into the first form. Iodine trichloride melts at 101° only in a closed tube under a pressure of 16 atmospheres.

[88] By the action of water on iodine monochloride and trichloride a compound IHCl_{2} is obtained, which does not seem to be altered by water. Besides this compound, iodine and iodic acid are always formed, 10ICl + 3H_{2}O = HIO_{3} + 5IHCl_{2} + 2I_{2}; and in this respect iodine trichloride may be regarded as a mixture, ICl + ICl_{5} = 2ICl_{3}, but ICl_{5} + 3H_{2}O = IHO_{3} + 5HCl; hence iodic acid, iodine, the compound IHCl_{2}, and hydrochloric acid are also formed by the action of water.

CHAPTERR XII

SODIUM

The neutral salt, sodium sulphate, Na_{2}SO_{4}, 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, PbSO_{4} + 2NaCl = PbCl_{2} + Na_{2}SO_{4}, 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 PbCl_{2} 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, Ag_{2}SO_{4}, in solution with common salt, gives silver chloride, because the latter is insoluble in water, Ag_{2}SO_{4} + 2NaCl = Na_{2}SO_{4} + 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, MgSO_{4} + Na_{2}CO_{3} = Na_{2}SO_{4} + MgCO_{3}. 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, CuSO_{4} + 2NaHO = Cu(HO){2} + Na{2}SO_{4}. Sulphate of magnesium, MgSO_{4}, 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 + MgSO_{4} = MgCl_{2} + Na_{2}SO_{4} takes place.

[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.

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, Na_{2}SO_{4},10H_{2}O, 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.

[2] Anhydrous (ignited) sodium sulphate, Na_{2}SO_{4}, 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 Na_{2}SO_{4}H_{2}O is obtained, together with a supersaturated solution.

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, Na_{2}SO_{4},10H_{2}O, 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 SO_{3}. 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 Na_{2}SO_{4},10H_{2}O.[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 Na_{2}SO_{4} + 14H_{2}O, 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, Cl_{2},8H_{2}O (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, Na_{2}SO_{4},H_{2}O. The heptahydrated salt, Na_{2}SO_{4},7H_{2}O, 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), Na_{2}SO_{4},7H_{2}O, 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 Na_{2}SO_{4},7H_{2}O (from 0° to 26°), one for Na_{2}SO_{4},10H_{2}O (from 0° to 34°), and one for Na_{2}SO_{4},H_{2}O (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]

[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, Na_{2}S, or sodium sulphite, Na_{2}SO_{3}, &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, Na_{2}SO_{4}, is obtained from NaCl without the preliminary manufacture of H_{2}SO_{4}. 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 SO_{2}. 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, CaSO_{4},2H_{2}O, 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 Na{2}SO_{4} 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 + 90·2_p_ + 0·35_p_^2 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 CaCl_{2}).

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 3Na_{2}SO_{4},7H_{2}0 = 2Na_{2}SO_{4},10H_{2}O + Na_{2}SO_{4},H_{2}O. 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]

[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, Na_{2}SO_{4},10H_{2}O, 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 Na_{2}SO_{4},10H_{2}O 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 Na_{2}SO_{4},10H_{2}O 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.

Sodium sulphate, Na_{2}SO_{4}, 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, Na_{2}SO_{4} + H_{2}SO_{4} = 2NaHSO_{4}. This separates from hot solutions, whilst the crystallo-hydrate, NaHSO_{4},H_{2}O,[10] separates from cold solutions. The crystals when exposed to damp air decompose into H_{2}SO_{4}, which deliquesces, and Na_{2}SO_{4} (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, Na_{2}S_{2}O_{7}, 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--Na_{2}S_{2}O_{7} = Na_{2}SO_{4} + SO_{3}. 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.

[10] The very large and well-formed crystals of this salt resemble the hydrate H_{2}SO_{4},H_{2}O, 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 + 77·92_p_ + 0·231_p_^2 (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, H_{2}SO_{4} + 25H_{2}O with Na_{2}SO_{4} + 25H_{2}O 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 NaHSO_{4},H_{2}SO_{4} are separated (Schultz, 1868). This compound fuses at about 100°; the ordinary acid salt, NaHSO_{4}, at 149°.

[11 bis] On decreasing the pressure, sodium hydrogen sulphate, NaHSO_{4}, dissociates much more easily than at the ordinary pressure; it loses water and forms the pyrosulphate, Na_{2}S_{2}O_{7}; this reaction is utilised in chemical works.

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, Na_{2}SO_{4} + Ba(HO){2} = BaSO{4} + 2NaHO. And by taking any salt of barium, BaX_{2}, the corresponding salt of sodium may be obtained, Na_{2}SO_{4} + BaX_{2} = BaSO_{4} + 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 SO_{3} 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, Na_{2}SO_{3} (for instance, in the formation of glass); or else the decomposition proceeds further, and sodium sulphide, Na_{2}S, is formed, according to the equation Na_{2}SO_{4} + 2C = 2CO_{2} + Na_{2}S.

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, Na_{2}CO_{3}, 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 Na_{2}SO_{4}, 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, Na_{2}SO_{4} + 2C = Na_{2}S + 2CO_{2}. The sodium sulphide thus formed then enters into double decomposition with the calcium carbonate taken, and gives calcium sulphide and sodium carbonate, Na_{2}S + CaCO_{3} = Na_{2}CO_{3} + CaS.

Besides which, under the action of the heat, a portion of the excess of calcium carbonate is decomposed into lime and carbonic anhydride, CaCO_{3} = CaO + CO_{2}, 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: 2Na_{2}SO_{4} + 3CaCO_{3} + 9C = 2Na_{2}CO_{3} + 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]

[12] Calcium sulphide, CaS, like many metallic sulphides which are soluble in water, is decomposed by it (Chapter X.), CaS + H_{2}O = CaO + H_{2}S, 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] [Illustration: 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 Na_{2}CO_{3},10H_{2}O 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 + MgCl_{2} + 2H_{2}O = CaCl_{2} + Mg(OH){2} + H{2}S. Sulphurous anhydride is passed through the residue in order to form the insoluble calcium sulphite: CaCl_{2} + Mg(OH){2} + SO{2} = CaSO_{3} + MgCl_{2} + H_{2}O. 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:

CaSO_{3} + 2H_{2}S + 2HCl = CaCl_{2} + 3H_{2}O + 3S.

But most efforts have been directed towards avoiding the formation of soda waste.

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]

[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, Na_{6}Fe_{4}S_{3}, 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 Na_{2}S from Na_{2}SO_{4} (by heating it with carbon), dissolves it in water and subjects the solution to the action of an excess of CO_{2} in coke towers, thus obtaining H_{2}S (a gas which gives SO_{2} under perfect combustion, or sulphur when incompletely burnt, Chapter XX., Note 6) and bicarbonate of sodium; Na_{2}S + 2CO_{2} + 2H_{2}O = H_{2}S + 2HNaCO_{3}. The latter gives soda and CO_{2} 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 Na_{2}SO_{4}, like that which is found in the Caucasus, all the more since H_{2}S 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.

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, NH_{4}HCO_{3}. 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, NaHCO_{3}, is precipitated and ammonium chloride, NH_{4}Cl, 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 NaHCO_{3} according to the equation NaCl + NH_{4}HCO_{3} = NH_{4}Cl + NaHCO_{3}. 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]

[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 NH_{4}Cl 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 NH_{3}, and forms MgCl_{2}: the former goes back for the manufacture of soda, while the latter is employed either for making HCl or Cl_{2}.

[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 Na_{2}SO_{4}. 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 + 104·5_p_ + 0·165_p_^2. 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 = _S__{0} + Ap + Bp^2, where _S__{0} 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.

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.

[18] The resemblance is so great that, notwithstanding the difference in the molecular composition of Na_{2}SO_{4} and Na_{2}CO_{3}, they ought to be classed under the type (NaO){2}R, where R = SO{2} or CO. Many other sodium salts also contain 10 mol. H_{2}O.

At a red heat superheated steam liberates carbonic anhydride from sodium carbonate and forms caustic soda, Na_{2}CO_{3} + H_{2}O = 2NaHO + CO_{2}. 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 (BaCl_{2} + Na_{2}CO_{3} = 2NaCl + BaCO_{3}) and with the aluminium salts it precipitates aluminium hydroxide, carbonic anhydride being disengaged: 3Na_{2}CO_{3} + Al_{2}(SO_{4}){3} + 3H{2}O = 3Na_{2}SO_{4} + 2Al(OH){3} + 3CO{2}. 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), NaHCO_{3}.[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. MgCO_{3}. No such precipitate is formed by the acid salt, because magnesium carbonate is soluble in the presence of an excess of carbonic anhydride.

[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, H_{2}CO_{3}, with the normal salt, Na_{2}CO_{3}, just as the latter also combines with water. Such a combination is all the more likely because (1) there exists another salt, Na_{2}CO_{3},2NaHCO_{3},2H_{2}O (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 NaHCO_{3},Na_{2}CO_{3},2H_{2}O 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 Na{2}CO_{3},H_{2}O, this salt will, when powdered, absorb CO_{2} 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: Na_{2}CO_{3},10H_{2}O + CO_{2} = Na_{2}CO_{3},H_{2}CO_{3} + 9H_{2}O, 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.

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: Na_{2}CO_{3} + Ca(HO){2} = CaCO{3} + 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,7H_{2}O; 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.

[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 CO_{2} move in a circle round the elements Na_{2}O, but at the points of contact with Fe_{2}O_{3} the motion becomes elliptic with a long axis, and at some distance from Na_{2}O the elements of CO_{2} are parted, not having the faculty of attaching themselves to Fe_{2}O_{3}.

[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,7H_{2}O 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 H_{2}O.

[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 + _n_H_{2}O, on being mixed with water, evolve heat if they contain less than 6H_{2}O, 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 Caustic Sodium Alcohol salt soda salt

Such a decomposition is termed saponification; similar reactions were known very long ago for the ethereal salts corresponding with glycerin, C_{3}H_{5}(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, C{16}H_{32}O_{2} and C_{18}H_{38}O_{2}, and the liquid oleic acid, C_{18}H_{34}O_{2}. 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).

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(NO_{3}){2} + 2NaHO = Cu(HO){2} + 2NaNO_{3}. 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, Cl_{2}, 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.

[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.

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, NaHSO_{4}, NaHCO_{3}), 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, H_{2}CO_{3}, or sulphuric, H_{2}SO_{4}--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, NaHSO_{4}, and Na_{2}SO_{4}, 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, CaCO_{3} and CaSO_{4}.[35 bis]

[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, H_{3}PO_{4}--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.

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.

[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 Na_{2}O; then, however, only one half is displaced, with the formation of NaHO.

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:

Na_{2}CO_{3} + 2C = Na_{2} + 3CO

[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 6-1/2 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 = Na_{2}CO_{3} + 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).

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.

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.

[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 NaHg_{5}, the amount of heat evolved increases, after which the heat of formation falls, and the heat evolved decreases. In the formation of NaHg_{5} about 18,500 calories are evolved; when NaHg_{3} is formed, about 14,000; and for NaHg about 10,000 calories. Kraft regarded the definite crystalline amalgam as having the composition of NaHg_{6}, but at the present time, in accordance with Grimaldi's results, it is thought to be NaHg_{5}. 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, KHg_{10}. 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 Na_{2}Sn), lead (NaPb), bismuth (Na_{3}Bi), &c. (Joannis 1892 and others) have been investigated.

It is easy to form an alloy of mercury and sodium having a crystalline structure, and a definite atomic composition, NaHg_{5}. The alloy of sodium with hydrogen or sodium hydride, Na_{2}H, 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, H_{2}O, H_{3}N, H_{4}C, or even in the gaseous metallic hydrides AsH_{3}, TeH_{2}. 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]

[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 Na_{2}H p = 28 57 447 910 for K_{2}H 45 72 548 1100

[43] In general, during the formation of alloys the volumes change very slightly, and therefore from the volume of Na_{2}H 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 Na_{2}H 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 Na_{2}H (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 Na_{2}H 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 NH_{3} and forms (NH_{3}Na)_{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).

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 NH_{2}Na, 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]

[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, NH_{2}Na, (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, NH_{2}Na, hydrogen being separated. This substance with water forms sodium hydroxide and ammonia; with carbonic oxide, CO, it forms sodium cyanide, NaCN, and water, H_{2}O; 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, NH_{2}Na, only partially decomposes, with evolution of hydrogen, the principal part of it giving ammonia and sodium nitride, Na_{3}N, according to the equation 3NH_{2}Na = 2NH_{3} + NNa_{3}. 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 NH_{3}. Pure NH_{2}Na 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 NH_{2}Na. It partially volatilises (sublimes) in vacuo at 200°, and breaks up into 2Na + N_{2} + 2H_{2} at 500°. The same amide is formed when oxide of sodium is heated in NH_{3}: Na_{2}O + 2NH_{3} = 2NaH_{2}N + H_{2}O. NaHO is also formed to some extent by the resultant H_{2}O. Potassium and lithium form similar amides. With water, alcohol, and acids, NH_{2}Na gives NH_{3} and NaHO, which react further. Anhydrous CaO absorbs NH_{2}Na when heated without decomposing it. When sodamide is heated with SiO_{2}, NH_{3} is disengaged, and silicon nitride formed. It acts still more readily upon boric anhydride when heated with it: 2NH_{2}Na + B_{2}O_{3} = 2BN + 2NaHO + H_{2}O. When slightly heated, NH_{2}Na + NOCl = NaCl + N_{2} + H_{2}O (NHNa_{2} and NNa_{3} 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(C_{2}H_{5}){2}, when treated with sodium, loses zinc and forms sodium ethyl, C{2}H_{5}Na, 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).

With oxygen sodium unites in three degrees of combination, forming a suboxide Na_{4}O,[46 bis] an oxide, Na_{2}O, and a peroxide, NaO. They are thus termed because Na_{2}O is a stable basic oxide (with water it forms a basic hydroxide), whilst Na_{4}O 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: Na_{4}O + 3H_{2}O = 4NaHO + H_{2}[47] Oxide: Na_{2}O + H_{2}O = 2NaHO[48] Peroxide: Na_{2}O_{2} + H_{2}O = 2NaHO + O[49]

All three oxides form sodium hydroxide with water, but only the oxide Na_{2}O 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 Na_{2}O, 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, Na_{2}O, is the only salt-forming oxide of this metal, as water is in the case of hydrogen. Although the peroxide H_{2}O_{2} is derived from hydrogen, and Na_{2}O_{2} 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, CO_{2}. Nitrogen and chlorine both give several salt-forming oxides and types of salts. But of the oxides of nitrogen, NO and NO_{2} do not form salts, as do N_{2}O_{3}, N_{2}O_{4}, and N_{2}O_{5}, although N_{2}O_{4} does not form special salts, and N_{2}O_{5} 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, Na_{2}O, 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.

[46 bis] It is even doubtful whether the suboxide exists (see Note 47).

[47] A compound, Na_{2}Cl, 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 Na_{2}Cl; the more so as the mass obtained gives hydrogen when treated with water: Na_{2}Cl + H_{2}O = H + NaHO + NaCl, that is, it acts like suboxide of sodium. If Na_{2}Cl really exists as a salt, then the corresponding base Na_{4}O, according to the rule with other bases of the composition M_{4}O, 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, K_{2}O, and of sodium, Na_{2}O, by heating the metals with their hydroxides, whence NaHO + Na = Na_{2}O + 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, Na_{2}O, when heated in an atmosphere of hydrogen, gave a mixture of sodium and its hydroxide: Na_{2}O + 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 Na_{2}O_{2}; this is obtained when sodium is burnt in an excess of oxygen. If NaNO_{3} be melted, it gives Na_{2}O_{2} 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: Na_{2}O_{2} +I_{2} = Na_{2}OI_{2} + O. The compound Na_{2}OI_{2} is akin to the compound Cu_{2}OCl_{2} obtained by oxidising CuCl. This reaction is one of the few in which iodine directly displaces oxygen. The substance Na_{2}OI_{2} 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: Na_{2}CO_{3} = Na_{2}O_{2} + CO, whilst carbonic anhydride liberates oxygen from it. With nitrous oxide it reacts thus: Na_{2}O_{2} +2N_{2}O = 2NaNO_{2} +N_{2}; with nitric oxide it combines directly, forming sodium nitrite, NaO + NO = NaNO_{2}. Sodium peroxide, when treated with water, does not give hydrogen peroxide, because the latter in the presence of the alkali formed (Na_{2}O_{2}+ 2H_{2}O = 2NaHO + H_{2}O_{2}) decomposes into water and oxygen. In the presence of dilute sulphuric acid it forms H_{2}O_{2} (Na_{2}O_{2} + H_{2}SO_{4} = Na_{2}SO_{4} + H_{2}O_{2}). 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 H_{2}O_{2} formed). The oxidising properties of Na_{2}O_{2} under the action of heat are seen, for instance, in the fact that when heated with I it forms sodium iodate; with PbO, Na_{2}PbO_{3}; with pyrites, sulphates, &c. When peroxide of sodium comes into contact with water, it evolves much heat, forming H_{2}O_{2}, and decomposing with the disengagement of oxygen; but, as a rule, there is no explosion. But if Na_{2}O_{2} 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 Na_{2}O_{2} 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.

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.