Chapter X., Note 2.

[2] Carnallite belongs to the number of double salts which are directly decomposed by water, and it only crystallises from solutions which contain an excess of magnesium chloride. It may be prepared artificially by mixing strong solutions of potassium and magnesium chlorides, when colourless crystals of sp. gr. 1·60 separate, whilst the Stassfurt salt is usually of a reddish tint, owing to traces of iron. At the ordinary temperature sixty-five parts of carnallite are soluble in one hundred parts of water in the presence of an excess of the salt. It deliquesces in the air, forming a solution of magnesium chloride and leaving potassium chloride. The quantity of carnallite produced at Stassfurt is now as much as 100,000 tons a year.

[3] The method of separating sodium chloride from potassium chloride has been described in Chapter I. On evaporation of a mixture of the saturated solutions, sodium chloride separates; and then, on cooling, potassium chloride separates, owing to the difference of rate of variation of their solubilities with the temperature. The following are the most trustworthy figures for the solubility of _potassium chloride_ in one hundred parts of water (for sodium chloride, _see_ Chapter X., Note 13):--

10° 20° 40° 60° 100° 32 35 40 46 57

When mixed with solutions of other salts the solubility of potassium chloride naturally varies, but not to any great extent.

[3 bis] The specific gravity of the solid salt is 1·99--that is, less than that of sodium chloride. All the salts of sodium are specifically heavier than the corresponding salts of potassium, as are also their solutions for equal percentage compositions. If the specific gravity of water at 4° = 10,000, then at 15° the specific gravity of a solution of _p_ p.c. potassium chloride = 9,992 + 63·29_p_ + 0·226_p_^2, and therefore for 10 p.c. = 1·0647, 20 p.c. = 1·1348, &c.

Potassium chloride combines with iodine trichloride to form a compound KCl + ICl_{3} = KICl_{4}, which has a yellow colour, is fusible, loses iodine trichloride at a red heat, and gives potassium iodate and hydrochloric acid with water. It is not only formed by direct combination, but also by many other methods; for instance, by passing chlorine into a solution of potassium iodide so long as the gas is absorbed, KI + 2Cl_{2} = KCl,ICl_{3}. Potassium iodide, when treated with potassium chlorate and strong hydrochloric acid, also gives this compound; another method for its formation is given by the equation KClO_{3} + I + 6HCl = KCl,ICl_{3} + 3Cl + 3H_{2}O. This is a kind of salt corresponding with KIO_{2} (unknown) in which the oxygen is replaced by chlorine. If valency be taken as the starting-point in the study of chemical compounds, and the elements considered as having a constant atomicity (number of bonds)--that is, if K, Cl, and I be taken as univalent elements--then it is impossible to explain the formation of such a compound because, according to this view, univalent elements are only able to form dual compounds with each other; such as, KCl, ClI, KI, &c., whilst here they are grouped together in the molecule KICl_{4}. Wells, Wheeler, and Penfield (1892) obtained a large number of such poly-haloid salts. They may all be divided into two large classes: the tri-haloid and the penta-haloid salts. They have been obtained not only for K but also for Rb and Cs, and partially also for Na and Li. The general method of their formation consists in dissolving the ordinary halogen salt of the metal in water, and treating it with the requisite amount of free halogen. The poly-haloid salt separates out after evaporating the solution at a more or less low temperature. In this manner, among the tri-haloid salts, may be obtained: KI_{3}, KBr_{2}I, KCl_{2}I, and the corresponding salts of rubidium and cæsium, for instance, CsI_{3}, CsBrI_{2}, CsBr_{2}I, CsClBrI,CsCl_{2}I, CsBr_{3}, CsClBr_{2}, CsCl_{2}Br, and in general MX_{3} where X is a halogen. The colour of the crystals varies according to the halogen, thus CsI_{3} is black, CrBr_{3} yellowish red, CrBrI_{2} reddish brown, CsBr_{2}I red, CsCl_{2}Br yellow. The cæsium salts are the most stable, and those of potassium least so, as also those which contain Br and I separately or together; for cæsium no compounds containing Cl and I were obtained. The penta-haloid salts form a smaller class; among these salts potassium forms KCl_{4}I, rubidium RbCl_{4}I, cæsium CsI_{5}, CsBr, CsCl_{4}I, lithium LiCl_{4}I (with 4H_{2}O) and sodium NaCl_{4}I (with 2H_{2}O). The most stable are those salts containing the metal with the greatest atomic weight--cæsium (_see_ Chapter XI., Note 63).

[4] It is possible to extract the compounds of potassium directly from the primary rocks which are so widely distributed over the earth's surface and so abundant in some localities. From a chemical point of view this problem presents no difficulty; for instance, by fusing powdered orthoclase with lime and fluor spar (Ward's method) and then extracting the alkali with water (on fusion the silica gives an insoluble compound with lime), or by treating the orthoclase with hydrofluoric acid (in which case silicon fluoride is evolved as a gas) it is possible to transfer the alkali of the orthoclase to an aqueous solution, and to separate it in this manner from the other insoluble oxides. However, as yet there is no profit in, nor necessity for, recourse to this treatment, as carnallite and potash form abundant materials for the extraction of potassium compounds by cheaper methods. Furthermore, the salts of potassium are now in the majority of chemical reactions replaced by salts of sodium, especially since the preparation of sodium carbonate has been facilitated by the Leblanc process. The replacement of potassium compounds by sodium compounds not only has the advantage that the salts of sodium are in general cheaper than those of potassium, but also that a smaller quantity of a sodium salt is needed for a given reaction than of a potassium salt, because the combining weight of sodium (23) is less than that of potassium (39).

[5] It has been shown by direct experiment on the cultivation of plants in artificial soils and in solutions that under conditions (physical, chemical, and physiological) otherwise identical plants are able to thrive and become fully developed in the entire absence of sodium salts, but that their development is impossible without potassium salts.

The primary rocks contain an almost equal proportion of potassium and sodium. But in sea water the compounds of the latter metal predominate. It may be asked, what became of the compounds of potassium in the disintegration of the primary rocks, if so small a quantity went to the sea water? They remained with the other products of the decomposition of the primary rocks. When granite or any other similar rock formation is disintegrated, there are formed, besides the soluble substances, also insoluble substances--sand and finely-divided clay, containing water, alumina, and silica. This clay is carried away by the water, and is then deposited in strata. It, and especially its admixture with vegetable remains, retain compounds of potassium in a greater quantity than those of sodium. This has been proved with absolute certainty to be the case, and is due to the _absorptive power of the soil_. If a dilute solution of a potassium compound be filtered through common mould used for growing plants, containing clay and the remains of vegetable decomposition, this mould will be found to have retained a somewhat considerable percentage of the potassium compounds. If a salt of potassium be taken, then during the filtration an equivalent quantity of a salt of calcium--which is also found, as a rule, in soils--is set free. Such a process of filtration through finely divided earthy substances proceeds in nature, and the compounds of potassium are everywhere retained by the friable earth in considerable quantity. This explains the presence of so small an amount of potassium salts in the water of rivers, lakes, streams, and oceans, where the lime and soda have accumulated. The compounds of potassium retained by the friable mass of the earth are absorbed as an aqueous solution by the roots of _plants_. Plants, as everyone knows, when burnt leave an ash, and this ash, besides various other substances, without exception contains compounds of potassium. Many land plants contain a very small amount of sodium compounds,[6] whilst potassium and its compounds occur in all kinds of vegetable ash. Among the generally cultivated plants, grass, potatoes, the turnip, and buckwheat are particularly rich in potassium compounds. The ash of plants, and especially of herbaceous plants, buckwheat straw, sunflower and potato leaves are used in practice for the extraction of potassium compounds. There is no doubt that potassium occurs in the plants themselves in the form of complex compounds, and often as salts of organic acids. In certain cases such salts of potassium are even extracted from the juice of plants. Thus, sorrel and oxalis, for example, contain in their juices the acid oxalate of potassium, C_{2}HKO_{4}, which is employed for removing ink stains. Grape juice contains the so-called cream of tartar, which is the acid tartrate of potassium, C_{4}H_{5}KO_{6}.[7] This salt also separates as a sediment from wine. When the plants, containing one or more of the salts of potassium, are burnt, the carbonaceous matter is oxidised, and in consequence the potassium is obtained in the ash as carbonate, K_{2}CO_{3}, which is generally known as _potashes_. Hence potashes occur ready prepared in the ash of plants, and therefore the ash of land plants is employed as a source for the extraction of potassium compounds. Potassium carbonate is extracted by lixiviating the ash with water.[8] Potassium carbonate may also be obtained from the chloride by a method similar to that by which sodium carbonate is prepared from sodium chloride.[8 bis] There is no difficulty in obtaining any salt of potassium--for example, the sulphate,[9] bromide, and iodide[10]--by the action of the corresponding acid on KCl and especially on the carbonate, whilst the hydroxide, _caustic potash_, KHO, which is in many respects analogous to caustic soda, is easily obtained by means of lime in exactly the same manner in which sodium hydroxide is prepared from sodium carbonate.[11] Therefore, in order to complete our knowledge of the alkali metals, we will only describe two salts of potassium which are of practical importance, and whose analogues have not been described in the preceding chapter, potassium cyanide and potassium nitrate.

[6] If herbaceous plants contain much sodium salts, it is evident that these salts mainly come from the sodium compounds in the water absorbed by the plants.

[7] As plants always contain mineral substances and cannot thrive in a medium which does not contain them, more especially in one which is free from the salts of the four basic oxides, K_{2}O, CaO, MgO, and Fe_{2}O_{3}, and of the four acid oxides, CO_{2}, N_{2}O_{5}, P_{2}O_{5}, and SO_{3}, and as the amount of ash-forming substances in plants is small, the question inevitably arises as to what part these play in the development of plants. With the existing chemical data only one answer is possible to this question, and it is still only a hypothesis. This answer was particularly clearly expressed by Professor Gustavson of the Petroffsky Agricultural Academy. Starting from the fact (Chapter XI., Note 55) that a small quantity of aluminium renders possible or facilitates the reaction of bromine on hydrocarbons at the ordinary temperature, it is easy to arrive at the conclusion, which is very probable and in accordance with many data respecting the reactions of organic compounds, that the addition of mineral substances to organic compounds lowers the temperature of reaction and in general facilitates chemical reactions in plants, and thus aids the conversion of the most simple nourishing substances into the complex component parts of the plant organism. The province of chemical reactions proceeding in organic substances in the presence of a small quantity of mineral substances has as yet been but little investigated, although there are already several disconnected data concerning reactions of this kind, and although a great deal is known with regard to such reactions among inorganic compounds. The essence of the matter may be expressed thus--two substances, A and B, do not react on each other of their own accord, but the addition of a small quantity of a third particularly active substance, C, produces the reaction of A on B, because A combines with C, forming AC, and B reacts on this new compound, which has a different store of chemical energy, forming the compound AB or its products, and setting C free again or retaining it.

It may here be remarked that all the mineral substances necessary for plants (those enumerated at the beginning of the note) are the highest saline compounds of their elements, that they enter into the plants as salts, that the lower forms of oxidation of the same elements (for instance, sulphites and phosphites) are harmful to plants (poisonous), and that strong solutions of the salts assimilated by plants (their osmotic pressure being great and contracting the cells, as De Vries showed, (_see_ Chapter I., Note 19)) not only do not enter into the plants but kill them (poison them).

[8] Besides which, it will be understood from the preceding paragraph that the salts of potassium may become exhausted from the soil by long cultivation, and that there may therefore be cases when the direct fertilisation by salts of potassium may be profitable. But manure and animal excrements, ashes, and, in general, nearly all refuse which may serve for fertilising the soil, contain a considerable quantity of potassium salts, and therefore, as regards the natural salts of potassium (Stassfurt), and especially potassium sulphate, if they often improve the crops, it is in all probability due to their action on the properties of the soil. The agriculturist cannot therefore be advised to add potassium salts, without making special experiments showing the advantage of such a fertiliser on a given kind of soil and plant.

The animal body also contains potassium compounds, which is natural, since animals consume plants. For example, milk, and especially human milk, contains a somewhat considerable quantity of potassium compounds. Cow's milk, however, does not contain much potassium salt. Sodium compounds generally predominate in the bodies of animals. The excrement of animals, and especially of herbivorous animals, on the contrary, often contains a large proportion of potassium salts. Thus sheep's dung is rich in them, and in washing sheep's wool salts of potassium pass into the water.

The ash of tree stems, as the already dormant portion of the plant (Chapter VIII., Note 1), contains little potash. For the extraction of potash, which was formerly carried on extensively in the east of Russia (before the discovery of the Stassfurt salt), the ash of grasses, and the green portions of potatoes, buckwheat, &c., are taken and treated with water (lixiviated), the solution is evaporated, and the residue ignited in order to destroy the organic matter present in the extract. The residue thus obtained is composed of raw potash. It is refined by a second dissolution in a small quantity of water, for the potash itself is very soluble in water, whilst the impurities are sparingly soluble. The solution thus obtained is again evaporated, and the residue ignited, and this potash is then called refined potash, or pearlash. This method of treatment cannot give chemically pure potassium carbonate. A certain amount of impurities remain. To obtain chemically pure potassium carbonate, some other salt of potassium is generally taken and purified by crystallisation. Potassium carbonate crystallises with difficulty, and it cannot therefore be purified by this means, whilst other salts, such as the tartrate, acid carbonate, sulphate, or nitrate, &c., crystallise easily and may thus be directly purified. The tartrate is most frequently employed, since it is prepared in large quantities (as a sediment from wine) for medicinal use under the name of cream of tartar. When ignited without access of air, it leaves a mixture of charcoal and potassium carbonate. The charcoal so obtained being in a finely-divided condition, the mixture (called 'black flux'), is sometimes used for reducing metals from their oxides with the aid of heat. A certain quantity of nitre is added to burn the charcoal formed by heating the cream of tartar. Potassium carbonate thus prepared is further purified by converting it into the acid salt, by passing a current of carbonic anhydride through a strong solution. KHCO_{3} is then formed, which is less soluble than the normal salt (as is also the case with the corresponding sodium salts), and therefore crystals of the acid salt separate from the solution on cooling. When ignited, they part with their water and carbonic anhydride, and pure potassium carbonate remains behind. The physical properties of potassium carbonate distinguish it sufficiently from sodium carbonate; it is obtained from solutions as a powdery white mass, having an alkaline taste and reaction, and, as a rule, shows only traces of crystallisation. It also attracts the moisture of the air with great energy. The crystals do not contain water, but absorb it from the air, deliquescing into a saturated solution. It melts at a red heat (1045°), and at a still higher temperature is even converted into vapour, as has been observed at glass works where it is employed. It is very soluble. At the ordinary temperature, water dissolves an equal weight of the salt. Crystals containing two equivalents of water separate from such a saturated solution when strongly cooled (Morel obtained K_{2}CO_{3}3H_{2}O in well-formed crystals at +10°). There is no necessity to describe its reactions, because they are all analogous to those of sodium carbonate. When manufactured sodium carbonate was but little known, the consumption of potassium carbonate was very considerable, and even now washing soda is frequently replaced for household purposes by 'lye'--_i.e._ an aqueous solution obtained from ashes. It contains potassium carbonate, which acts like the sodium salt in washing tissues, linen, &c.

A mixture of potassium and sodium carbonates fuses with much greater ease than the separate salts, and a mixture of their solutions gives well-crystallised salts--for instance (Marguerite's salt), K_{2}CO_{3},6H_{2}O,2Na_{2}CO_{3},6H_{2}O. Crystallisation also occurs in other multiple proportions of K and Na (in the above case 1 : 2, but 1 : 1 and 1 : 3 are known), and always with 6 mol. H_{2}O. This is evidently a combination _by similarity_, as in alloys, solutions, &c.

[8 bis] About 25,000 tons of potash annually are now prepared from KCl by this method at Stassfurt.

[9] _Potassium sulphate_, K_{2}SO_{4}, crystallises from its solutions in an anhydrous condition, in which respect it differs from the corresponding sodium salt, just as potassium carbonate differs from sodium carbonate. In general, it must be observed that the majority of sodium salts combine more easily with water of crystallisation than the potassium salts. The solubility of _potassium sulphate_ does not show the same peculiarities as that of sodium sulphate, because it does not combine with water of crystallisation; at the ordinary temperature 100 parts of water dissolve about 10 parts of the salt, at 0° 8·3 parts, and at 100° about 26 parts. _The acid sulphate_, KHSO_{4}, obtained easily by heating crystals of the normal salt with sulphuric acid, is frequently employed in chemical practice. On heating the mixture of acid and salt, fumes of sulphuric acid are at first given off; when they cease to be evolved, the acid salt is contained in the residue. At a higher temperature (of above 600°) the acid salt parts with all the acid contained in it, the normal salt being re-formed. The definite composition of this acid salt, and the ease with which it decomposes, render it exceedingly valuable for certain chemical transformations accomplished by means of sulphuric acid at a high temperature, because it is possible to take, in the form of this salt, a strictly definite quantity of sulphuric acid, and to cause it to act on a given substance at a high temperature, which it is often necessary to do, more especially in chemical analysis. In this case, the acid salt acts in exactly the same manner as sulphuric acid itself, but the latter is inefficient at temperatures above 400°, because it all evaporates, while at that temperature the acid salt still remains in a fused state, and acts with the elements of sulphuric acid on the substance taken. Hence by its means the boiling-point of sulphuric acid is raised. Thus the acid potassium sulphate is employed, where for conversion of certain oxides, such as those of iron, aluminium, and chromium, into salts, a high temperature is required.

Weber, by heating potassium sulphate with an excess of sulphuric acid at 100°, observed the formation of a lower stratum, which was found to contain a definite compound containing eight equivalents of SO_{3} per equivalent of K_{2}O. The salts of rubidium, cæsium, and thallium give a similar result, but those of sodium and lithium do not. (_See_ Note 1.)

[10] The _bromide_ and _iodide_ of potassium are used, like the corresponding sodium compounds, in medicine and photography. Potassium iodide is easily obtained in a pure state by saturating a solution of hydriodic acid with caustic potash. In practice, however, this method is rarely had recourse to, other more simple processes being employed although they do not give so pure a product. They aim at the direct formation of hydriodic acid in the liquid in the presence of potassium hydroxide or carbonate. Thus iodine is thrown into a solution of pure potash, and hydrogen sulphide passed through the mixture, the iodine being thus converted into hydriodic acid. Or a solution is prepared from phosphorus, iodine, and water, containing hydriodic and phosphoric acid; lime is then added to this solution, when calcium iodide is obtained in solution, and calcium phosphate as a precipitate. The solution of calcium iodide gives, with potassium carbonate, insoluble calcium carbonate and a solution of potassium iodide. If iodine is added to a slightly-heated solution of caustic potash (free from carbonate--that is, freshly prepared), so long as the solution is not coloured from the presence of an excess of iodine, there is formed (as in the action of chlorine on a solution of caustic potash) a mixture of potassium iodide and iodate. On evaporating the solution thus obtained and igniting the residue, the iodate is destroyed and converted into iodide, the oxygen being disengaged, and potassium iodide only is left behind. On dissolving the residue in water and then evaporating, cubical crystals of the anhydrous salt are obtained, which are soluble in water and alcohol, and on fusion give an alkaline reaction, owing to the fact that when ignited a portion of the salt decomposes, forming potassium oxide. The neutral salt may be obtained by adding hydriodic acid to this alkaline salt until it gives an acid reaction. It is best to add some finely-divided charcoal to the mixture of iodate and iodide before igniting it, as this facilitates the evolution of the oxygen from the iodate. The iodate may also be converted into iodide by the action of certain reducing agents, such as zinc amalgam, which when boiled with a solution containing an iodate converts it into iodide. Potassium iodide may also be prepared by mixing a solution of ferrous iodide (it is best if the solution contain an excess of iodine) and potassium carbonate, in which case ferrous carbonate FeCO_{3}, is precipitated (with an excess of iodine the precipitate is granular, and contains a compound of the suboxide and oxide of iron), while potassium iodide remains in solution. Ferrous iodide, FeI_{2}, is obtained by the direct action of iodine on iron in water. Potassium iodide considerably lowers the temperature (by 24°), when it dissolves in water, 100 parts of the salt dissolve in 73·5 parts of water at 12·5°, in 70 parts at 18°, whilst the saturated solution which boils at 120° contains 100 parts of salt per 45 parts of water. Solutions of potassium iodide dissolve a considerable amount of iodine; strong solutions even dissolving as much or more iodine than they contain as potassium iodide (_see_ Note 3 bis and Chapter XI., Note 64).

[11] Caustic potash is not only formed by the action of lime on dilute solutions of potassium carbonate (as sodium hydroxide is prepared from sodium carbonate), but by igniting potassium nitrate with finely-divided copper (_see_ Note 15), and also by mixing solutions of potassium sulphate (or even of alum, KAlS_{2}O_{8}) and barium hydroxide, BaH_{2}O_{2}. It is sometimes purified by dissolving it in alcohol (the impurities, for example, potassium sulphate and carbonate, are not dissolved) and then evaporating the alcohol.

The specific gravity of potassium hydroxide is 2·04, but that of its solutions (see Chapter XII., Note 18) at 15° S = 9,992 + 90·4_p_ + 0·28_p_^2 (here _p_^2 is +, and for sodium hydroxide it is-). Strong solutions, when cooled, yield a crystallo-hydrate, KHO,4H_{2}O, which dissolves in water, producing cold (like 2NaHO,7H_{2}O), whilst potassium hydroxide in solution develops a considerable amount of heat.

_Potassium cyanide_, which presents in its chemical relations a certain analogy with the halogen salts of potassium, is not only formed according to the equation, KHO + HCN = H_{2}O + KCN, but also whenever a nitrogenous carbon compound--for instance, animal matter--is heated in the presence of metallic potassium, or of a compound of potassium, and even when a mixture of potash and carbon is heated in a stream of nitrogen. Potassium cyanide is obtained from yellow prussiate, which has been already mentioned in Chapter IX., and whose preparation on a large scale will be described in Chapter XXII. If the yellow prussiate be ground to a powder and dried, so that it loses its water of crystallisation, it then melts at a red heat, and decomposes into carbide of iron, nitrogen, and potassium cyanide, FeK_{4}C_{6}N_{6} = 4KCN + FeC_{2} + N_{2}. After the decomposition it is found that the yellow salt has been converted into a white mass of potassium cyanide. The carbide of iron formed collects at the bottom of the vessel. If the mass thus obtained be treated with water, the potassium cyanide is partially decomposed by the water, but if it be treated with alcohol, then the cyanide is dissolved, and on cooling separates in a crystalline form.[12] A solution of potassium cyanide has a powerfully alkaline reaction, a smell like that of bitter almonds, peculiar to prussic acid, and acts as a most powerful poison. Although exceedingly stable in a fused state, potassium cyanide easily changes when in solution. Prussic acid is so very feebly energetic that even water decomposes potassium cyanide. A solution of the salt, even without access of air, easily turns brown and decomposes, and when heated evolves ammonia and forms potassium formate; this is easily comprehensible from the representation of the cyanogen compounds which was developed in Chapter IX., KCN + 2H_{2}O = CHKO_{2} + NH_{3}. Furthermore, as carbonic anhydride acts on potassium cyanide with evolution of prussic acid, and as potassium cyanate, which is also unstable, is formed by the action of air, it will be easily seen that solutions of potassium cyanide are very unstable. Potassium cyanide, containing as it does carbon and potassium, is a substance which can act in a very vigorously reducing manner, especially when fused; it is therefore used as a powerful reducing agent at a red heat.[13] The property of potassium cyanide of giving double salts with other cyanides is very clearly shown by the fact that many metals dissolve in a solution of potassium cyanide, with the evolution of hydrogen. For example, iron, copper, and zinc act in this manner. Thus--

4KCN + 2H_{2}O + Zn = K_{2}ZnC_{4}N_{4} + 2KHO + H_{2}

[12] When the yellow prussiate is heated to redness, all the cyanogen which was in combination with the iron is decomposed into nitrogen, which is evolved as gas, and carbon, which combines with the iron. In order to avoid this, potassium carbonate is added to the yellow prussiate while it is being fused. A mixture of 8 parts of anhydrous yellow prussiate and 3 parts of pure potassium carbonate is generally taken. Double decomposition then takes place, resulting in the formation of ferrous carbonate and potassium cyanide. But by this method, as by the first, a pure salt is not obtained, because a portion of the potassium cyanide is oxidised at the expense of the iron carbonate and forms potassium cyanate, FeCO_{3} + KCN = CO_{2} + Fe + KCNO; and the potassium cyanide very easily forms oxide, which acts on the sides of the vessel in which the mixture is heated (to avoid this iron vessels should be used). By adding one part of charcoal powder to the mixture of 8 parts of anhydrous yellow prussiate and 3 parts of potassium carbonate a mass is obtained which is free from cyanate, because the carbon absorbs the oxygen, but in that case it is impossible to obtain a colourless potassium cyanide by simple fusion, although this may be easily done by dissolving it in alcohol. Cyanide of potassium may also be obtained from potassium thiocyanate, which is formed from ammonium thiocyanate obtained by the action of ammonia upon bisulphide of carbon (_see_ works upon Organic Chemistry). Potassium cyanide is now prepared in large quantities from yellow prussiate for gilding and silvering. When fused in large quantities the action of the oxygen of the air is limited, and with great care the operation may be successfully conducted, and therefore, on a large scale, very pure salt is sometimes obtained. When slowly cooled, the fused salt separates in cubical crystals like potassium chloride.

Pure KCN is obtained by passing CNH gas into an alcoholic solution of KHO. The large amount of potassium cyanide which is now required for the extraction of gold from its ores, is being replaced by a mixture (Rossler and Gasslaker, 1892) of KCN and NaCN, prepared by heating powdered and dried yellow prussiate with metallic sodium: K_{4}Fe(CN)_{6} + 2Na = 4KCN + 2NaCN + Fe. This method offers two advantages over the above methods: (1) the whole of the cyanide is obtained, and does not decompose with the formation of N_{2}; and (2) no cyanates are formed, as is the case when carbonate of potash is heated with the prussiate.

[13] A considerable quantity of potassium cyanide is used in the arts, more particularly for the preparation of metallic solutions which are decomposed by the action of a galvanic current; thus it is very frequently employed in electro-silvering and gilding. An alkaline solution is prepared, which is moderately stable owing to the fact that potassium cyanide in the form of certain double salts--that is, combined with other cyanides--is far more stable than when alone (yellow prussiate, which contains potassium cyanide in combination with ferrous cyanide, is an example of this).

Gold and silver are soluble in potassium cyanide in the presence of air, in which case the hydrogen, which would otherwise be evolved in the reaction, combines with the oxygen of the air, forming water (Eissler, MacLaurin, 1893), for example, 4Au + 4KCN + O + H_{2}O = 2AuKC_{2}N_{2} + 2KHO, which is taken advantage of for extracting gold from its ores (Chapter XXIV.).[13 bis] Platinum, mercury, and tin are not dissolved in a solution of potassium cyanide, even with access of air.

[13 bis] A dilute solution of KCN is taken, not containing more than 1 per cent. KCN. MacLaurin explains this by the fact that strong solutions dissolve gold less rapidly, owing to their dissolving less air, whose oxygen is necessary for the reaction.

_Potassium nitrate_, or common _nitre_ or _saltpetre_, KNO_{3}, is chiefly used as a component part of gunpowder, in which it cannot be replaced by the sodium salt, because the latter is deliquescent. It is necessary that the nitre in gunpowder should be perfectly pure, as even small traces of sodium, magnesium, and calcium salts, especially chlorides, render the nitre and the gunpowder capable of attracting moisture. Nitre may easily be obtained pure, owing to its great disposition to form crystals both large and small, which aids its separation from other salts. The considerable differences between the solubility of nitre at different temperatures aids this crystallisation. A solution of nitre saturated at its boiling point (116°) contains 335 parts of nitre to 100 parts of water, whilst at the ordinary temperature--for instance, 20°--the solution is only able to retain 32 parts of the salt. Therefore, in the preparation and refining of nitre, its solution, saturated at the boiling point, is cooled, and nearly all the nitre is obtained in the form of crystals. If the solution be quietly and slowly cooled in large quantities then large crystals are formed, but if it be rapidly cooled and agitated then small crystals are obtained. In this manner, if not all, at all events the majority, of the impurities present in small quantities remain in the mother liquor. If an unsaturated solution of nitre be rapidly cooled, so as to prevent the formation of large crystals (in whose crevices the mother liquor, together with the impurities, would remain), the very minute crystals of nitre known as saltpetre flour are obtained.

Common nitre occurs in nature, but only in small quantities in admixture with other nitrates, and especially with sodium, magnesium, and calcium nitrates. Such a mixture of salts of nitric acid is formed in nature in fertile earth, and in those localities where, as in _the soil_, nitrogenous organic remains are decomposed in the presence of alkalis or alkaline bases with free access of air. This method of the formation of nitrates requires moisture, besides the free access of air, and takes place principally during warm weather.[14] In warm countries, and in temperate climates during the summer months, fertile soils produce a small quantity of nitre. In this respect India is especially known as affording a considerable supply of nitre extracted from the soil. The nitre-bearing soil after the rainy season sometimes becomes covered during the summer with crystals of nitre, formed by the evaporation of the water in which it was previously dissolved. This soil is collected, subjected to repeated lixiviations, and treated for nitre as will be presently described. In temperate climates nitrates are obtained from the lime rubbish of demolished buildings which have stood for many years, and especially from those portions which have been in contact with the ground. The conditions there are very favourable for the formation of nitre, because the lime used as a cement in buildings contains the base necessary for the formation of nitrates, while the excrement, urine, and animal refuse are sources of nitrogen. By the methodical lixiviation of this kind of rubbish a solution of nitrogenous salts is formed similar to that obtained by the lixiviation of fertile soil. A similar solution is also obtained by the lixiviation of the so-called _nitre plantations_. They are composed of manure interlaid with brushwood, and strewn over with ashes, lime, and other alkaline rubbish. These nitre plantations are set up in those localities where the manure is not required for the fertilisation of the soil, as, for example, in the south-eastern 'black earth' Governments of Russia. The same process of oxidation of nitrogenous matter freely exposed to air and moisture during the warm season in the presence of alkalis takes place in nitre plantations as in fertile soil and in the walls of buildings. From all these sources there is obtained a solution containing various salts of nitric acid mixed with soluble organic matter. The simplest method of treating this impure solution of nitre is to add a solution of potassium carbonate, or to simply treat it with ashes containing this substance. The potassium carbonate enters into double decomposition with the calcium and magnesium salts, forming insoluble carbonates of these bases and leaving the nitre in solution. Thus, for instance, K_{2}CO_{3} + Ca(NO_{3})_{2} = 2KNO_3 + CaCO_{3}. Both calcium and magnesium carbonates are insoluble, and therefore after treatment with potassium carbonate the solution no longer contains salts of these metals but only the salts of sodium and potassium together with organic matter. The latter partially separates on heating in an insoluble form, and is entirely destroyed by heating the nitre to a low red heat. The nitre thus obtained is easily purified by repeated crystallisation. The greater part of the nitre used for making gunpowder is now obtained from the sodium salt _Chili saltpetre_ or _cubic nitre_, which occurs in nature, as already mentioned. The conversion of this salt into common nitre is also carried on by means of a double decomposition. This is done either by adding potassium carbonate (when, on mixing the strong and hot solutions, sodium carbonate is directly obtained as a precipitate), or, as is now most frequent, potassium chloride. When a mixture of strong solutions of potassium chloride and sodium nitrate is evaporated, sodium chloride first separates, because this salt, which is formed by the double decomposition KCl + NaNO_{3} = KNO_{3} + NaCl, is almost equally soluble in hot and cold water; on cooling, therefore, a large amount of potassium nitrate separates from the saturated solution, while the sodium chloride remains dissolved. The nitre is ultimately purified by recrystallisation and by washing with a saturated solution of nitre, which cannot dissolve a further quantity of nitre but only the impurities.

[14] Besides which Schloesing and Müntz, by employing similar methods to Pasteur, showed that the formation of nitre in the decomposition of nitrogenous substances is accomplished by the aid of peculiar micro-organisms (ferments), without which the simultaneous action of the other necessary conditions (alkalis, moisture, a temperature of 37°, air, and nitrogenous substances) cannot give nitre.

Nitre is a colourless salt having a peculiar cool taste. It crystallises easily in long striated six-sided rhombic prisms terminating in rhombic pyramids. Its crystals (sp. gr. 1·93) do not contain water, but their cavities generally contain a certain quantity of the solution from which they have crystallised. For this reason in refining nitre, the production of large crystals is prevented, _saltpetre flour_ being prepared. At a low red heat (339°) nitre melts to a colourless liquid.[14 bis] Potassium nitrate at the ordinary temperature and in a solid form is inactive and stable, but _at a high temperature_ it acts as a powerful _oxidising agent_, giving up a considerable amount of oxygen to substances mixed with it.[15] When thrown on to incandescent charcoal it brings about its rapid combustion, and a mechanical mixture of powdered charcoal and nitre ignites when brought into contact with a red-hot substance, and continues to burn by itself. In this action, nitrogen is evolved, and the oxygen oxidises the charcoal, in consequence of which potassium carbonate and carbonic anhydride are formed: 4KNO_{3} + 5C = 2K_{2}CO_{3} + 3CO_{2} + 2N_{2}. This phenomenon depends on the fact that oxygen in combining with carbon evolves more heat than it does in combining with nitrogen. Hence, when once the combustion has been started at the expense of the nitre, it is able to go on without requiring the aid of external heat. A similar oxidation or combustion at the expense of the contained oxygen takes place when nitre is heated with different combustible substances. If a mixture of sulphur and nitre be thrown upon a red-hot surface, the sulphur burns, forming potassium sulphate and sulphurous anhydride. In this case, also, the nitrogen of the nitre is evolved as gas: 2KNO_{3} + 2S = K_{2}SO_{4} + N_{2} + SO_{2}. A similar phenomenon occurs when nitre is heated with many metals. The oxidation of those metals which are able to form acid oxides with an excess of oxygen is especially remarkable. In this case they remain in combination with potassium oxide as potassium salts. Manganese, antimony, arsenic, iron, chromium, &c. are instances of this kind. These elements, like carbon and sulphur, displace free nitrogen. The lower oxides of these metals when fused with nitre pass into the higher oxides. Organic substances are also oxidised when heated with nitre--that is, they burn at the expense of the nitre. It will be readily understood from this that nitre is frequently used in practical chemistry and the arts as an oxidising agent at high temperatures. Its application in _gunpowder_ is based on this property; gunpowder consists of a mechanical mixture of finely-ground sulphur, nitre, and charcoal. The relative proportion of these substances varies according to the destination of the powder and to the kind of charcoal employed (a friable, incompletely-burnt charcoal, containing therefore hydrogen and oxygen, is employed). Gases are formed in its combustion, chiefly nitrogen and carbonic anhydride, which create a considerable pressure if their escape be in any way impeded. This action of gunpowder may be expressed by the equation: 2KNO_{3} + 3C + S = K_{2}S + 3CO_{2} + N_{2}.

[14 bis] Before fusing, the crystals of potassium nitrate change their form, and take the same form as sodium nitrate--that is, they change into rhombohedra. Nitre crystallises from hot solutions, and in general under the influence of a rise of temperature, in a different form from that given at the ordinary or lower temperatures. Fused nitre solidifies to a radiated crystalline mass; but it does not exhibit this structure if metallic chlorides be present, so that this method may be taken advantage of to determine the degree of purity of nitre.

Carnelley and Thomson (1888) determined the fusing point of mixtures of potassium and sodium nitrates. The first salt fuses at 339° and the second at 316°, and if _p_ be the percentage amount of potassium nitrate, then the results obtained were--

_p_ = 10 20 30 40 50 60 70 80 90 298° 283° 268° 242° 231° 231° 242° 284° 306°

which confirms Shaffgotsch's observation (1857) that the lowest fusing point (about 231°) is given by mixing molecular quantities (_p_ = 54·3) of the salts--that is, in the formation of the alloy, KNO_{3},NaNO_{3}.

A somewhat similar result was discovered by the same observers for the solubility of mixtures of these salts at 20° in 100 parts of water. Thus, if _p_ be the weight of potassium nitrate mixed with 100-_p_ parts by weight of sodium nitrate taken for solution, and _c_ be the quantity of the mixed salts which dissolves in 100, the solubility of sodium nitrate being 85, and of potassium nitrate 34, parts in 100 parts of water, then--

_p_ = 10 20 30 40 50 60 70 80 90 _c_ = 110 136 136 138 106 81 73 54 41

The maximum solubility proved not to correspond with the most fusible mixture, but to one much richer in sodium nitrate.

Both these phenomena show that in homogeneous liquid mixtures the chemical forces that act between substances are the same as those that determine the molecular weights of substances, even when the mixture consists of such analogous substances as potassium and sodium nitrates, between which there is no direct chemical interchange. It is instructive to note also that the maximum solubility does not correspond with the minimum fusing point, which naturally depends on the fact that in solution a third substance, namely water, plays a part, although an attraction between the salts, like that which exists between sodium and potassium carbonates (Note 8), also partially acts.

[15] Fused nitre, with a further rise of temperature, disengages oxygen and then nitrogen. The nitrite KNO_{2} is first formed and then potassium oxide. The admixture of certain metals--for example, of finely-divided copper--aids the last decomposition. The oxygen in this case naturally passes over to the metal.

It is found by this equation that gunpowder should contain thirty-six parts of charcoal (13·3 p.c.), and thirty-two parts (11·9 p.c.) of sulphur, to 202 parts (74·8 p.c.) of nitre, which is very near to its actual composition.[16]

[16] In China, where the manufacture of gunpowder has long been carried on, 75·7 parts of nitre, 14·4 of charcoal, and 9·9 of sulphur are used. Ordinary powder for sporting purposes contains 80 parts of nitre, 12 of charcoal, and 8 of sulphur, whilst the gunpowder used in heavy ordnance contains 75 of nitre, 15 of charcoal, and 10 of sulphur. Gunpowder explodes when heated to 300°, when struck, or by contact with a spark. A compact or finely-divided mass of gunpowder burns slowly and has but little disruptive action, because it burns gradually. To act properly the gunpowder must have a definite rate of combustion, so that the pressure should increase during the passage of the projectile along the barrel of the fire-arm. This is done by making the powder in large granules or in the shape of six-sided prisms with holes through them (prismatic powder).

The products of combustion are of two kinds: (1) gases which produce the pressure and are the cause of the dynamical action of gunpowder, and (2) a solid residue, usually of a black colour owing to its containing unburnt particles of charcoal. Besides charcoal, the residue generally contains potassium sulphide, K_{2}S, and a whole series of other salts--for instance, carbonate and sulphate. It is apparent from this that the combustion of gunpowder is not so simple as it appears to be from the above formula, and hence the weight of the residue is also greater than indicated by that formula. According to the formula, 270 parts of gunpowder give 110 parts of residue--that is, 100 parts of powder give 37·4 parts of residue, K_{2}S, whilst in reality the weight of the residue varies from 40 p.c. to 70 p.c. (generally 52 p.c.). This difference depends on the fact that so much oxygen (of the nitre) remains in the residue, and it is evident that if the residue varies the composition of the gases evolved by the powder will vary also, and therefore the entire process will be different in different cases. The difference in the composition of the gases and residue depends, as the researches of Gay-Lussac, Shishkoff and Bunsen, Nobel and Abel, Federoff, Debus, &c., show, on the conditions under which the combustion of the powder proceeds. When gunpowder burns in an open space, the gaseous products which are formed do not remain in contact with the residue, and then a considerable portion of the charcoal entering into the composition of the powder remains unburnt, because the charcoal burns after the sulphur at the expense of the oxygen of the nitre. In this extreme case the commencement of the combustion of the gunpowder may be expressed by the equation, 2KNO_{3} + 3C + S = 2C + K_{2}SO_{4} + CO_{2} + N_{2}. The residue in a blank cartridge often consists of a mixture of C, K_{2}SO_{4}, K_{2}CO_{3}, and K_{2}S_{2}O_{3}. If the combustion of the gunpowder be impeded--if it take place in a cartridge in the barrel of a gun--the quantity of potassium sulphate will first be diminished, then the amount of sulphite, whilst the amount of carbonic anhydride in the gases and the amount of potassium sulphide in the residue will increase. The quantity of charcoal entering into the action will then be also increased, and hence the amount in the residue will decrease. Under these circumstances the weight of the residue will be less--for example, 4K_{2}CO_{3} + 4S = K_{2}SO_{4} + 3K_{2}S + 4CO_{2}. Besides which, carbonic oxide has been found in the gases, and potassium bisulphide, K_{2}S_{2}, in the residue of gunpowder. The amount of potassium sulphide, K_{2}S, increases with the completeness of the combustion, and is formed in the residue at the expense of the potassium sulphite. In recent times the knowledge of the action of gunpowder and other explosives has made much progress, and has developed into a vast province of artillery science, which, guided by the discoveries of chemistry,

has worked out a 'smokeless powder' which burns without leaving a residue, and does not therefore give any 'powder smoke' (to hinder the rapidity of firing and aiming), and at the same time disengages a greater volume of gas and consequently gives (under proper conditions of combustion) the possibility of communicating to the charge a greater initial velocity, and therefore greater distance, force, and accuracy of aim. Such 'smokeless powder' is prepared either from the varieties of nitrocellulose (Chapter VI., Note 37) or from a mixture of them with nitro-glycerine (_ibid_). In burning they give, besides steam and nitrogen, generally a large amount of oxide of carbon (this is a very serious drawback in all the present forms of smokeless powder, because carbonic oxide is poisonous), and also CO_{2}, H_{2}, &c.

_Metallic potassium_ was obtained like sodium; first by the action of a galvanic current, then by reduction of the hydroxide by means of metallic iron, and lastly, by the action of charcoal on the carbonate at a high temperature. The behaviour of metallic potassium differs, however, from that of sodium, because it easily combines with carbonic oxide, forming an explosive and inflammable mass.[17]

[17] The substances obtained in this case are mentioned in Chapter IX., Note 31.

Potassium is quite as volatile as sodium, if not more so. At the ordinary temperature potassium is even softer than sodium; its freshly-cut surfaces present a whiter colour than sodium, but, like the latter, and with even greater ease, it oxidises in moist air. It is brittle at low temperatures, but is quite soft at 25°, and melts at 58°. At a low red heat (667°, Perkin) it distils without change, forming a green vapour, whose density,[18] according to A. Scott (1887), is equal to 19 (if that of hydrogen = 1). This shows that the molecule of potassium (like that of sodium, mercury, and zinc) contains but one atom. This is also the case with many other metals, judging by recent researches.[19] The specific gravity of potassium at 15° is 0·87, and is therefore less than that of sodium, as is also the case with all its compounds.[20] Potassium decomposes water with great ease at the ordinary temperature, evolving 45,000 heat units per atomic weight in grams. The heat evolved is sufficient to inflame the hydrogen, the flame being coloured violet from the presence of particles of potassium.[21]

[18] A. Scott (1887) determined the vapour densities of many of the alkali elements and their compounds in a platinum vessel heated in a furnace and previously filled with nitrogen. But these, the first data concerning a subject of great importance, have not yet been sufficiently fully described, nor have they received as much attention as could be desired. Taking the density of hydrogen as unity, Scott found the vapour densities of the following substances to be--

Na 12·75 (11·5). KI 92 (84). K 19 (19·5). RbCl 70 (60). CsCl 89·5 (84·2). CsI 133 (130). FeCl_{3} 68. AgCl 80 (71·7).

In brackets are given the densities corresponding with the formulæ, according to Avogadro-Gerhardt's law. This figure is not given for FeCl_{3}, because in all probability under these conditions (the temperature at which it was determined) a portion of the FeCl_{3} was decomposed. If it was not decomposed, then a density 81 would correspond with the formula FeCl_{3}, and if the decomposition were Fe_{2}Cl_{6} = 2FeCl_{2} + Cl_{2}, then the density should be 54. With regard to the silver chloride, there is reason to think that the platinum decomposed this salt. The majority of Scott's results so closely correspond with the formulæ that a better concord cannot be expected in such determinations. V. Meyer (1887) gives 93 as the density of KI.

[19] The molecules of non-metals are more complex--for instance, H_{2}, O_{3}, Cl_{2}, &c. But arsenic, whose superficial appearance recalls that of metals, but whose chemical properties approach more nearly to the non-metals, has a complex molecule containing As_{4}.

[20] As the atomic weight of potassium is greater than that of sodium, the volumes of the molecules, or the quotients of the molecular weight by the specific gravity, for potassium compounds are greater than those of sodium compounds, because both the denominator and numerator of the fraction increase. We cite for comparison the volumes of the corresponding compounds--

Na 24 NaHO 18 NaCl 28 NaNO_{3} 37 Na_{2}SO_{4} 54 K 45 KHO 27 KCl 39 KNO_{3} 48 K_{2}SO_{4} 66

[21] The same precautions must be taken in decomposing water by potassium as have to be observed with sodium (Chapter II., Note 8).

It must be observed that potassium decomposes carbonic anhydride and carbonic oxide when heated, the carbon being liberated and the oxygen taken up by the metal, whilst on the other hand charcoal takes up oxygen from potassium, as is seen from the preparation of potassium by heating potash with charcoal, hence the reaction K_{2}O + C = K_{2} + CO is reversible and the relation is the same in this case as between hydrogen and zinc.

With regard to the relation of potassium to hydrogen and oxygen, it is closely analogous to sodium in this respect. Thus, with hydrogen it forms potassium hydride, K_{2}H (between 200° and 411°), and with oxygen it gives a suboxide K_{4}O, oxide K_{2}O, and peroxide, only more oxygen enters into the composition of the latter than in sodium peroxide; potassium peroxide contains KO_{2}, but it is probable that in the combustion of potassium an oxide KO is also formed. Potassium, like sodium, is soluble in mercury.[22] In a word, the relation between sodium and potassium is as close as that between chlorine and bromine, or, better still, between fluorine and chlorine, as the atomic weight of sodium, 23, is as much greater than that of fluorine, 19, as that of potassium, 39, is greater than that of chlorine, 35·5.

[22] _Potassium_ forms _alloys with sodium_ in all proportions. The alloys containing 1 and 3 equivalents of potassium to one equivalent of sodium are _liquids_, like mercury at the ordinary temperature. Joannis, by determining the amount of heat developed by these alloys in decomposing water, found the evolution for Na_{2}K, NaK, NaK_{2} and NaK_{3} to be 44·5, 44·1, 43·8 and 44·4 thousand heat units respectively (for Na 42·6 and for K 45·4). The formation of the alloy NaK_{2} is therefore accompanied by the development of heat, whilst the other alloys may be regarded as solutions of potassium or sodium in this alloy. In any case a fall of the temperature of fusion is evident in this instance as in the alloys of nitre (Note 14). The liquid alloy NaK_{2} is now used for filling thermometers employed for temperatures above 360°, when mercury boils.

The resemblance between _potassium_ and _sodium_ is so great that _their compounds_ can only be easily _distinguished_ in the form of certain of their salts. For instance, the acid potassium tartrate, C_{4}H_{5}KO_{6} (cream of tartar), is distinguished by its sparing solubility in water and in alcohol, and in a solution of tartaric acid, whilst the corresponding sodium salt is easily soluble. Therefore, if a solution of tartaric acid be added in considerable excess to the solutions of the majority of potassium salts, a precipitate of the sparingly-soluble acid salt is formed, which does not occur with salts of sodium. The chlorides KCl and NaCl in solutions easily give double salts K_{2}PtCl_{6} and Na_{2}PtCl_{6}, with platinic chloride, PtCl_{4}, and the solubility of these salts is very different, especially in a mixture of alcohol and ether. The sodium salt is easily soluble, whilst the potassium salt is insoluble or almost so, and therefore the reaction with platinic chloride is that most often used for the separation of potassium from sodium, as is more fully described in works on analytical chemistry.

It is possible to discover the least traces of these metals in admixture together, by means of their property of imparting different colours to _a flame_. The presence of a salt of sodium in a flame is recognised by a brilliant yellow coloration, and a pure potassium salt colours a colourless flame violet. However, in the presence of a sodium salt the pale violet coloration given by a potassium salt is quite undistinguishable, and it is at first sight impossible in this case to discover the potassium salt in the presence of that of sodium. But by decomposing the light given by a flame coloured by these metals or a mixture of them, by means of a prism, they are both easily distinguishable, because the yellow light emitted by the sodium salt depends on a group of light rays having a definite index of refraction which corresponds with the yellow portion of the solar spectrum, having the index of refraction of the Fraunhofer line (strictly speaking, group of lines) D, whilst the salts of potassium give a light from which these rays are entirely absent, but which contain rays of a red and violet colour. Therefore, if a potassium salt occur in a flame, on decomposing the light (after passing it through a narrow slit) by means of a prism, there will be seen red and violet bands of light situated at a considerable distance from each other; whilst if a sodium salt be present a yellow line will also appear. If both metals simultaneously occur in a flame and emit light, the spectrum lines corresponding to the potassium and the sodium will appear simultaneously.

For convenience in carrying on this kind of testing, _spectroscopes_ (fig. 72) are constructed,[23] consisting of a refracting prism and three tubes placed in the plane of the refracting angle of the prism. One of the tubes, C, has a vertical slit at the end, giving access to the light to be tested, which then passes into the tube (collimator), containing a lens which gives the rays a parallel direction. The rays of light having passed through the slit, and having become parallel, are refracted and dispersed in the prism, and the spectrum formed is observed through the eye-piece of the other telescope B. The third tube D contains a horizontal transparent scale (at the outer end) which is divided into equal divisions. The light from a source such as a gas burner or candle placed before this tube, passes through the scale, and is reflected on that face of the prism which stands before the telescope B, so that the image of the scale is seen through this telescope simultaneously with the spectrum given by the rays passing through the slit of the tube C. In this manner the image of the scale and the spectrum given by the source of light under investigation are seen simultaneously. If the sun's rays be directed through the slit of the tube C, then the observer looking through the eye-piece of B will see the solar spectrum, and (if the aperture of the slit be narrow and the apparatus correctly adjusted) the dark Fraunhofer lines in it.[24] Small-sized spectroscopes are usually so adjusted that (looking through B) the violet portion of the spectrum is seen to the right and the red portion to the left, and the Fraunhofer line D (in the bright yellow portion of the spectrum) is situated on the 50th division of the scale.[25] If the light emitted by an incandescent solid--for example, the Drummond light--be passed through the spectroscope, then all the colours of the solar spectrum are seen, but not the Fraunhofer lines. To observe the result given by a flame coloured by various salts a Bunsen gas burner (or the pale flame of hydrogen gas issuing from a platinum orifice) giving so pale a flame that its spectrum will be practically invisible is placed before the slit. If any compound of sodium be placed in the flame of the gas burner (for which purpose a platinum wire on whose end sodium chloride is fused is fixed to the stand), then the flame is coloured yellow, and on looking through the spectroscope the observer will see a bright _yellow_ line falling upon the 50th division of the scale, which is seen together with the spectrum in the telescope. No yellow lines of other refractive index, nor any rays of any other colour, will be seen, and, therefore, the spectrum corresponding with sodium compounds consists of yellow rays of that index of refraction which belong to the Fraunhofer (black) line D of the solar spectrum. If a potassium salt be introduced into the flame instead of a sodium salt, then two bands will be seen which are much feebler than the bright sodium band--namely, one red line near the Fraunhofer line A and another violet line. Besides which, a pale, almost continuous, spectrum will be observed in the central portions of the scale. If a mixture of sodium and potassium salts be now introduced into the flame, three lines will be seen simultaneously--namely, the red and pale violet lines of potassium and the yellow line of sodium. In this manner it is possible, by the aid of the spectroscope, to determine the relation between the spectra of metals and known portions of the solar spectrum. The continuity of the latter is interrupted by dark lines (that is, by an absence of light of a definite index of refraction), termed the Fraunhofer lines of the solar spectrum. It has been shown by careful observations (by Fraunhofer, Brewster, Foucault, Ångstrom, Kirchhoff, Cornu, Lockyer, Dewar, and others) that there exists an exact _agreement between the spectra_ of certain _metals_ and certain of the _Fraunhofer_ lines. Thus the bright yellow sodium line exactly corresponds with the dark Fraunhofer line D of the solar spectrum. A similar agreement is observed in the case of many other metals. This is not an approximate or chance correlation. In fact, if a spectroscope having a large number of refracting prisms and a high magnifying power be used, it is seen that the dark line D of the solar spectrum consists of an entire system of closely adjacent but definitely situated fine and wide (sharp, distinct) dark lines,[26] and an exactly similar group of bright lines is obtained when the yellow sodium line is examined through the same apparatus, so that each bright sodium line exactly corresponds with a dark line in the solar spectrum.[26 bis] This conformity of the bright lines formed by sodium with the dark lines of the solar spectrum cannot be accidental. This conclusion is further confirmed by the fact that the bright lines of other metals correspond with dark lines of the solar spectrum. Thus, for example, a series of sparks passing between the iron electrodes of a Ruhmkorff coil gives 450 very distinct lines characterising this metal. All these 450 bright lines, constituting the whole spectrum corresponding with iron, are repeated, as Kirchhoff showed, in the solar spectrum as dark Fraunhofer lines which occur in exactly the same situations as the bright lines in the iron spectrum, just as the sodium lines correspond with the band D in the solar spectrum. Many observers have in this manner studied the solar spectrum and the spectra of different metals simultaneously, and discovered in the former lines which correspond not only with sodium and iron, but also with many other metals.[27] The spectra of such elements as hydrogen, oxygen, nitrogen, and other gases may be observed in the so-called Geissler's tubes--that is, in glass tubes containing rarefied gases, through which the discharge of a Ruhmkorff's coil is passed. Thus hydrogen gives a spectrum composed of three lines--a red line corresponding with the Fraunhofer line C, a green line corresponding with the line F, and a violet line corresponding with one of the lines between G and H. Of these rays the red is the brightest, and therefore the general colour of luminous hydrogen (with an electric discharge through a Geissler tube) is reddish.

[23] For accurate measurements and comparative researches more complicated spectroscopes are required which give a greater dispersion, and are furnished for this purpose with several prisms--for example, in Browning's spectroscope the light passes through six prisms, and then, having undergone an internal total reflection, passes through the upper portion of the same six prisms, and again by an internal total reflection passes into the ocular tube. With such a powerful dispersion the relative position of the spectral lines may be determined with accuracy. For the absolute and exact determination of the wave lengths it is particularly important that the spectroscope should be furnished with diffraction gratings. The construction of spectroscopes destined for special purposes (for example, for investigating the light of stars, or for determining the absorption spectra in microscopic preparations, &c.) is exceedingly varied. Details of the subject must be looked for in works on physics and on spectrum analysis. Among the latter the best known for their completeness and merit are those of Roscoe, Kayser, Vogel, and Lecoq de Boisbaudran.

[24] The arrangement of all the parts of the apparatus so as to give the clearest possible vision and accuracy of observation must evidently precede every kind of spectroscopic determination. Details concerning the practical use of the spectroscope must be looked for in special works on the subject. In this treatise the reader is supposed to have a certain knowledge of the physical data respecting the refraction of light, and its dispersion and diffraction, and the theory of light, which allows of the determination of the length of the waves of light in absolute measure on the basis of observations with diffraction gratings, the distance between whose divisions may be easily measured in fractions of a millimetre; by such means it is possible to determine the wave-length of any given ray of light.

[25] In order to give an idea of the size of the scale, we may observe that the ordinary spectrum extends from the zero of the scale (where the red portion is situated) to the 170th division (where the end of the visible violet portion of the spectrum is situated), and that the Fraunhofer line A (the extreme prominent line in the red) corresponds with the 17th division of the scale; the Fraunhofer line F (at the beginning of the blue, near the green colour) is situated on the 90th division, and the line G, which is clearly seen in the beginning of the violet portion of the spectrum, corresponds with the 127th division of the scale.

[26] The two most distinct lines of D, or of sodium, have wave-lengths of 589·5 and 588·9 millionths of a millimeter, besides which fainter and fainter lines are seen whose wave-lengths in millionths of a millimeter are 588·7 and 588·1, 616·0 and 615·4, 515·5 and 515·2, 498·3 and 498·2, &c., according to Liveing and Dewar.

[26 bis] In the ordinary spectroscopes which are usually employed in chemical research, one yellow band, which does not split up into thinner lines, is seen instead of the system of sodium lines, owing to the small dispersive power of the prism and the width of the slit of the object tube.

[27] The most accurate investigations made in this respect are carried on with spectra obtained by diffraction, because in this case the position of the dark and bright lines does not depend on the index of refraction of the material of the prism, nor on the dispersive power of the apparatus. The best--that is, the most general and accurate--method of expressing the results of such determinations consists in determining the lengths of the waves corresponding to the rays of a definite index of refraction. (Sometimes instead of this the fraction of 1 divided by the square of the wave-length is given.) We will express this _wave-length_ in _millionth parts of a millimetre_ (the ten-millionth parts are already doubtful, and fall within the limits of error). In order to illustrate the relation between the wave-lengths and the positions of the lines of the spectrum, we will cite the wave-lengths corresponding with the chief Fraunhofer lines and colours of the spectrum.

Fraunhofer line A B C D E b Wave-length 761·0 687·5 656·6 589·5-588·9 527·3 518·7 +-----------------+ +---------+ +----+ +---+ Colour red orange yellow green

Fraunhofer line F G H Wave-length 486·5 431·0 397·2 +----------+ +----+ Colour blue violet

In the following table are given the _wave-lengths_ of the light rays (the longest and most distinct, _see_ later) for certain elements, those in black type being the most clearly defined and distinct lines, which are easily obtained either in the flame of a Bunsen's burner, or in Geissler's tubes, or in general, by an electric discharge. These lines refer to the elements (the lines of compounds are different, as will be afterwards explained, but many compounds are decomposed by the flame or by an electric discharge), and moreover to the elements in an incandescent and rarefied gaseous state, for the spectra sometimes vary considerably with a variation of temperature and pressure.

It may be mentioned that the _red_ colour corresponds with lines having a wave-length of from 780 (with a greater wave-length the lines are hardly visible, and are ultra red) to 650, the _orange_ from 650 to 590, the _yellow_ from 590 to 520, the _green_ from 520 to 490, the _blue_ from 490 to 420, and the _violet_ from 420 to 380 millionth parts of a millimetre. Beyond 380 the lines are scarcely visible, and belong to the ultra-violet. For fluorine Moissan found as many as 13 bright lines from 744 to 623.

In the table (p. 565) which is arranged in conformity with the image of the spectrum as it is seen (the red lines on the left-hand and the violet on the right-hand side), the figures in black type correspond with lines which are so bright and distinctly visible that they may easily be made use of, both in determining the relation between the divisions of the scale and the wave-lengths, and in determining the admixture of a given element with another. Brackets join those lines between which several other lines are clearly visible if the dispersive power of the spectroscope permits distinguishing the neighbouring lines. In the ordinary laboratory spectroscopes with one prism, even with all possible precision of arrangement and with a brilliancy of light permitting the observations being made with a very narrow aperture, the lines whose wave-lengths only differ by 2-3 millionths of a millimetre, are blurred together; and with a wide aperture a series of lines differing by even as much as 20 millionths of a millimetre appear as one wide line. With a faint light (that is, with a small quantity of light entering into the spectroscope) only the most _brilliant_ lines are clearly visible. The _length_ of the lines does not always correspond with their brilliancy. According to Lockyer this length is determined by placing the carbon electrodes (between which the incandescent vapours of the metals are formed), not horizontally to the slit (as they are generally placed, to give more light), but vertically to it. Then certain lines appear long and others short. As a rule (Lockyer, Dewar, Cornu), the longest lines are those with which it is easiest to obtain _reversed_ spectra (_see_ later). Consequently, these lines are the most characteristic. Only the longest and most brilliant are given in our table, which is composed on the basis of a collection of the data at our disposal for _bright_ spectra of the _incandescent and rarefied vapours of the elements_. As the spectra change with great variations of temperature and vapour density (the faint lines become brilliant whilst the bright lines sometimes disappear), which is particularly clear from Ciamician's researches on the halogens, until the method of observation and the theory of the subject are enlarged, particular theoretical importance should not be given to the wave-lengths showing the maximum brilliancy, which only possess a practical significance in the common methods of spectroscopic observations. In general the spectra of metals are simpler than those of the halogens, and the latter are variable; at an increased pressure all spectral lines become broader.

+-------+-------+--------+--------+-------+-------+------+------+ | N_{2} | O_{2} | Cl_{2} | Br_{2} | I_{2} | Pb | Sn | Tl | +-------+-------+--------+--------+-------+-------+------+------+ | -- | -- | -- | -- | -- | -- | -- | -- | | -- | -- | -- | -- | -- | -- | -- | -- | | 662 | -- | -- | -- | -- | -- | -- | -- | | 632} | -- | -- | 636 | -- | -- | 645 | -- | | 620} | 615 | -- | -- | 621 | -- | -- | -- | | 585} | -- | -- | -- | 613 | 605·7 | 580} | -- | | 574} | -- | 546} | -- | 579} | 560·7 | 556} | -- | | 544} | 543 | 539} | 544} | 560} | 554·7 | -- | 549 | | 535} | 533 | 528} | 523} | 545} | 537 | -- | 535 | | 527} | 516 | 519} | 517 | 506} | -- | -- | -- | | 516} | 495} | 494} | 479 | -- | -- | -- | 489 | | -- | 494} | 480} | 470} | -- | -- | -- | -- | | 457 | 470} | -- | 462} | -- | -- | 452 | -- | | 442 | 465} | -- | 454 | 445 | -- | -- | -- | | 436} | 447} | 436} | 437 | -- | -- | -- | -- | | 426} | 432} | 431} | -- | -- | -- | -- | -- | | -- | -- | -- | -- | 421 | -- | -- | -- | | -- | -- | -- | -- | -- | -- | -- | -- | | 409 | -- | -- | -- | -- | 406 | -- | -- | | -- | -- | -- | -- | -- | -- | -- | -- | +-------+-------+--------+--------+-------+-------+------+------+

+-----+-----+-----+-------+-----+------+------+-----+-------+-----+ | In | Ga | Al | Ba | Sr | Ca | Mg | Zn | Cd | Hg | |-----+-----+-----+-------+-----|------+------|-----+-------+-----+ | -- | -- | -- | -- | -- | -- | -- | -- | -- | -- | | -- | -- | -- | -- | -- | -- | -- | -- | -- | -- | | -- | -- | -- | -- | -- | 646 | -- | -- | -- | -- | | -- | -- | 624 | 649·7 | 641 | 644 | -- | 636 | 643·8 | -- | | 619 | -- | 623 | 614 | 606 | -- | -- | -- | -- | 615 | | -- | -- | -- | -- | -- | 612 | -- | -- | -- | 579 | | -- | -- | 572 | 553·5 | -- | -- | -- | -- | -- | 577 | | -- | -- | 570 | 549 | 548 | 559 | -- | -- | 537·7 | 546 | | 525 | -- | -- | -- | 524 | -- | 518 | -- | 533·6 | -- | | -- | -- | -- | -- | -- | -- | 516 | 492 | 508·5 | -- | | -- | -- | -- | 493·3 | -- | -- | -- | 481 | -- | -- | | -- | -- | -- | -- | -- | -- | 471 | 472 | 479·9 | -- | | 451 | -- | 466 | 455 | 460 | -- | -- | 468 | 467·7 | -- | | -- | -- | -- | -- | -- | 445} | 448 | -- | -- | -- | | -- | -- | -- | -- | -- | 442} | -- | -- | -- | 436 | | -- | -- | -- | -- | 430 | -- | -- | -- | -- | -- | | -- | 417 | -- | -- | 421 | 423 | -- | -- | -- | -- | | -- | -- | -- | 413 | -- | -- | -- | -- | -- | -- | | 410 | 403 | 396 | -- | 408 | 397 | 384} | -- | -- | 404 | | -- | -- | 394 | -- | -- | 393 | 383} | - -| -- | -- | +-----+-----+-----+-------+-----+------+------+-----+-------+-----+

+------+------+-----+-------+-----+-------+------+-------+-------+ | Mn | Fe | Cu | Ag | Cs | Rb | K | Na | Li | |------+------+-----+-------+-----+-------+------+-------+-------+ | -- | -- | -- | -- | -- | 780 | 770 | -- | -- | | -- | -- | -- | -- | -- | -- | 766 | -- | -- | | -- | -- | -- | -- | -- | -- | -- | -- | 670·6 | | -- | 640 | -- | -- | -- | -- | -- | -- | -- | | 602} | -- | -- | -- | 622 | 629·6 | -- | -- | 610 | | 601} | 561} | 578 | -- | 600 | -- | 583} | 589·5 | -- | | 551} | 544} | 570 | 546·4 | -- | -- | 578} | 588·9 | -- | | 534} | 537 | 522 | -- | -- | -- | -- | -- | -- | | -- | 532} | 515 | 520·8 | -- | -- | 535 | -- | -- | | -- | 521} | 511 | -- | -- | -- | 532} | -- | -- | | 482} | 496} | -- | -- | -- | -- | -- | -- | 497 | | 471} | 489} | -- | -- | -- | -- | -- | -- | -- | | -- | -- | -- | -- | 459 | -- | -- | -- | 460·3 | | -- | 441} | -- | -- | 456 | -- | -- | -- | -- | | -- | 430} | -- | -- | -- | -- | -- | -- | -- | | -- | 427 | -- | -- | -- | -- | -- | -- | -- | | 424} | -- | -- | 421 | -- | 420 | -- | -- | -- | | 403} | 407} | -- | -- | -- | -- | -- | -- | -- | | -- | 404} | -- | -- | -- | -- | 404 | -- | -- | | -- | -- | -- | -- | -- | -- | -- | -- | -- | +------+------+-----+-------+-----+-------+------+-------+-------+

+-------+ | H_{2} | +-------+ | -- | | -- | | -- | | 656·2 | | -- | | -- | | -- | | -- | | -- | | -- | | 486·1 | | -- | | -- | | -- | | 434 | | -- | | -- | | -- | | -- | | -- | +-------+

The correlation of the Fraunhofer lines with the spectra of metals depends on the phenomenon of the so-called _reversal of the spectrum_. This phenomenon consists in this, that instead of the bright spectrum corresponding with a metal, under certain circumstances a similar dark spectrum in the form of Fraunhofer lines may be obtained, as will be explained directly. In order to clearly understand the phenomenon of reversed spectra, it must be known that when light passes through certain transparent substances these substances retain rays of a certain refrangibility. The colour of solutions is a proof of this. Light which has passed through a yellow solution of a uranium salt contains no violet rays, and after having passed through a red solution of a permanganate, does not contain many rays in the yellow, blue, and green portions of the spectrum. Solutions of copper salts absorb nearly all red rays. Sometimes colourless solutions also absorb rays of certain definite refractive indexes, and give _absorption spectra_. Thus solutions of salts of didymium absorb rays of a certain refrangibility, and therefore an impression of black lines is received,[28] as shown in fig. 73. Many vapours (iodine) and gases (nitric peroxide) give similar spectra. Light which has passed through a deep layer of aqueous vapour, oxygen, or nitrogen also gives an absorption spectrum. For this reason the peculiar (winter) dark lines discovered by Brewster are observed in sunlight, especially in the evening and morning, when the sun's rays pass through the atmosphere (containing these substances) by a longer path than at mid-day. It is evident that the Fraunhofer lines may be ascribed to the absorption of certain rays of light in its passage from the luminous mass of the sun to the earth. The remarkable progress made in all spectroscopic research dates from the investigations made by _Kirchhoff_ (1859) on the relation between absorption spectra and the spectra of luminous incandescent gases. It had already been observed long before (by Fraunhofer, Foucault, Ångstrom) that the bright spectrum of the sodium flame gives two bright lines which are in exactly the same position as two black lines known as D in the solar spectrum, which evidently belong to an absorption spectrum. When Kirchhoff caused diffused sunlight to fall upon the slit of a spectroscope, and placed a sodium flame before it, a perfect superposition was observed--the bright sodium lines completely covered the black lines D of the solar spectrum. When further the continuous spectrum of a Drummond light showed the black line D on placing a sodium flame between it and the slit of the spectroscope--that is, when the Fraunhofer line of the solar spectrum was artificially produced--then there was no doubt that its appearance in the solar spectrum was due to the light passing somewhere through incandescent vapours of sodium. Hence a new theory of _reversed spectra_[29] arose--that is, of the relation between the waves of light emitted and absorbed by a substance under given conditions of temperature; this is expressed by Kirchhoff's law, discovered by a careful analysis of the phenomena. This law may be formulated in an elementary way as follows: At a given temperature the relation between the intensity of the light emitted (of a definite wave-length) and the absorptive capacity with respect to the same colour (of the same wave-length) is a constant quantity.[30] As a black dull surface emits and also absorbs a considerable quantity of heat rays whilst a polished metallic surface both absorbs and emits but few, so a flame coloured by sodium emits a considerable quantity of yellow rays of a definite refrangibility, and has the property of absorbing a considerable quantity of the rays of the same refractive index. In general, the medium which emits definite rays also absorbs them.

[28] The method of observing absorption spectra consists in taking a continuous spectrum of white light (one which does not show either dark lines or particularly bright luminous bands--for instance, the light of a candle, lamp, or other source). The collimator (that is, the tube with the slit) is directed towards this light, and then all the colours of the spectrum are visible in the ocular tube. A transparent absorptive medium--for instance, a solution or tube containing a gas--is then placed between the source of light and the apparatus (or anywhere inside the apparatus itself in the path of the rays). In this case either the entire spectrum is uniformly fainter, or absorption bands appear on the bright field of the continuous spectrum in definite positions along it. These bands have different lengths and positions, and distinctness and intensity of absorption, according to the properties of the absorptive medium. Like the luminous spectra given by incandescent gases and vapours, the absorption spectra of a number of substances have already been studied, and some with great precision--as, for example, the spectrum of the brown vapours of nitrogen dioxide by Hasselberg (at Pulkowa), the spectra of colouring matters (Eder and others), especially of those applied to orthochromatic photography, the spectra of blood, chlorophyll (the green constituent of leaves), and other similar substances, all the more carefully as by the aid of their spectra the presence of these substances may be discovered in small quantities (even in microscopical quantities, by the aid of special appliances on the microscope), and the changes they undergo investigated.

The absorption spectra, obtained at the ordinary temperature and proper to substances in all physical states, offer a most extensive but as yet little studied field, both for the general theory of spectroscopy, and for gaining an insight into the structure of substances. The investigation of colouring matters has already shown that in certain cases a definite change of composition and structure entails not only a definite change of the colours but also a displacement of the absorption bands by a definite number of wave-lengths.

[29] A number of methods have been invented to demonstrate the reversibility of spectra; among these methods we will cite two which are very easily carried out. In Bunsen's method sodium chloride is put into an apparatus for evolving hydrogen (the spray of the salt is then carried off by the hydrogen and colours the flame with the yellow sodium colour), and the hydrogen is ignited in two burners--in one large one with a wide flame giving a bright yellow sodium light, and in another with a small fine orifice whose flame is pale: this flame will throw a dark patch on the large bright flame. In Ladoffsky's method the front tube (p. 561) is unscrewed from a spectroscope directed towards the light of a lamp (a continuous spectrum), and the flame of a spirit lamp coloured by a small quantity of NaCl is placed between the tube and the prism; a black band corresponding to sodium will then be seen on looking through the ocular tube. This experiment is always successful if only there be the requisite relation between the strength of light of the two lamps.

[30] The absorptive capacity is the relation between the intensity of the light (of a given wave-length) falling upon and retained by a substance. Bunsen and Roscoe showed by direct experiment that this ratio is a constant quantity for every substance. If _A_ stand for this ratio for a given substance at a given temperature--for instance, for a flame coloured by sodium--and _E_ be the intensity of the light of the same wave-length emitted at the same temperature by the same substance, then Kirchhoff's law, the explanation and deduction of which must be looked for in text-books of physics, states that the fraction _A/E_ is a constant quantity depending on the nature of a substance (as _A_ depends on it) and determined by the temperature and wave-length.

Thus the bright spectral rays characteristic of a given metal may be reversed--that is, converted into dark lines--by passing light which gives a continuous spectrum through a space containing the heated vapours of the given metal. A similar phenomenon to that thus artificially produced is observed in sunlight, which shows dark lines characteristic of known metals--that is, the Fraunhofer lines form an absorption spectrum or depend on a reversed spectrum; it being presupposed that the sun itself, like all known sources of artificial light, gives a continuous spectrum without Fraunhofer lines.[31] We must imagine that the sun, owing to the high temperature which is proper to it, emits a brilliant light which gives a continuous spectrum, and that this light, before reaching our eyes, passes through a space full of the vapours of different metals and their compounds. As the earth's atmosphere[32] contains very little, or no, metallic vapours, and as they cannot be supposed to exist in the celestial space,[32 bis] the only place in which the existence of such vapours can be admitted is in the _atmosphere surrounding the sun itself_. As the cause of the sun's luminosity must be looked for in its high temperature, the existence of an atmosphere containing metallic vapours is readily understood, because at that high temperature such metals as sodium, and even iron, are separated from their compounds and converted into vapour. The sun must be imagined as surrounded by an atmosphere of incandescent vaporous and gaseous matter,[33] including those elements whose reversed spectra correspond with the Fraunhofer lines--namely, sodium, iron, hydrogen, lithium, calcium, magnesium, &c. Thus in spectrum analysis we find a means of determining the composition of the inaccessible heavenly luminaries, and much has been done in this respect since Kirchhoff's theory was formulated. By observations on the spectra of many heavenly bodies, changes have been discovered going on in them,[34] and many of the elements known to us have been found with certainty in them.[35] From this it must be concluded that the same elements which exist on the earth occur throughout the whole universe, and that at that degree of heat which is proper to the sun those simple substances which we accept as the elements in chemistry are still undecomposed and remain unchanged. A high temperature forms one of those conditions under which compounds most easily decompose; and if sodium or a similar element were a compound, in all probability it would be decomposed into component parts at the high temperature of the sun. This may indeed be concluded from the fact that in ordinary spectroscopic experiments the spectra obtained often belong to the metals and not to the compounds taken; this depends on the decomposition of these compounds in the heat of the flame. If common salt be introduced into the flame of a gas-burner, a portion of it is decomposed, first forming, in all probability, with water, hydrochloric acid and sodium hydroxide, and the latter then becoming partially decomposed by the hydrocarbons, giving metallic sodium, whose incandescent vapour emits light of a definite refrangibility. This conclusion is arrived at from the following experiment:--If hydrochloric acid gas be introduced into a flame coloured by sodium it is observed that the sodium spectrum disappears, owing to the fact that metallic sodium cannot remain in the flame in the presence of an excess of hydrochloric acid. The same thing takes place on the addition of sal-ammoniac, which in the heat of the flame gives hydrochloric acid. If a porcelain tube containing sodium chloride (or sodium hydroxide or carbonate), and closed at both ends by glass plates, be so powerfully heated that the salt volatilises, then the sodium spectrum is not observable; but if the salt be replaced by sodium, then either the bright line or the absorption spectra is obtained, according to whether the light emitted by the incandescent vapour be observed, or light passing through the tube. Thus the above spectrum is not given by sodium chloride or other sodium compound, but is proper to the metal sodium itself. This is also the case with other analogous metals. The chlorides and other halogen _compounds_ of barium, calcium, copper, &c., give independent spectra which differ from those of the metals. If barium chloride be introduced into a flame, it gives a mixed spectrum belonging to metallic barium and barium chloride. If besides barium chloride, hydrochloric acid or sal-ammoniac be introduced into the flame, then the spectrum of the metal disappears, and that of the chloride remains, which differs distinctly from the spectrum of barium fluoride, barium bromide, or barium iodide. A certain common resemblance and certain common lines are observed in the spectra of two different compounds of one and the same element obtained in the above-described manner, and also in the spectrum of the metal, but they all have their peculiarities. The independent spectra of the compounds of copper are easily observed (fig. 75). Thus certain compounds which exist in a state of vapour, and are luminous at a high temperature, give their independent spectra. In the majority of cases the spectra of compounds are composed of indistinct luminous lines and complete bright bands, whilst metallic elements generally give a few clearly-defined spectral lines.[36] There is no reason for supposing that the spectrum of a compound is equal to the sum of the spectra of its elements--that is, _every compound_ which is not decomposed by heat _has its own proper spectrum_. This is best proved by absorption spectra, which are essentially only reversed spectra observed at low temperatures. If every salt of sodium, lithium, and potassium gives one and the same spectrum, this must be ascribed to the presence in the flame of the free metals liberated by the decomposition of their salts. Therefore _the phenomena of the spectrum are determined by molecules, and not by atoms_--that is, the molecules of the metal sodium, and not its atoms, produce those particular vibrations which determine the spectrum of a sodium salt. Where there is no free metallic sodium there is no sodium spectrum.

[31] Heated metals begin to emit light (only visible in the dark) at about 420° (varying with the metal). On further heating, solids first emit red, then yellow, and lastly white light. Compressed or heavy gases (_see_ Chapter III., Note 44), when strongly heated, also emit white light. Heated liquids (for example, molten steel or platinum) also give a white compound light. This is readily understood. In a dense mass of matter the collisions of the molecules and atoms are so frequent that waves of only a few definite lengths cannot appear; the reverse is possible in rarefied gases or vapours.

[32] Brewster, as is mentioned above, first distinguished the atmospheric, cosmical Fraunhofer lines from the solar lines. Janssen showed that the spectrum of the atmosphere contains lines which depend on the absorption produced by aqueous vapour. Egoreff, Olszewski, Janssen, and Liveing and Dewar showed by a series of experiments that the oxygen of the atmosphere gives rise to certain lines of the solar spectrum, especially the line A. Liveing and Dewar took a layer of 165 c.m. of oxygen compressed under a pressure of 85 atmospheres, and determined its absorption spectrum, and found that, besides the Fraunhofer lines A and B, it contained the following groups: 630-622, 581-568, 535, 480-475. The same lines were found for liquid oxygen.

[32 bis] If the material of the whole heavenly space formed the absorbent medium, the spectra of the stars would be the same as the solar spectrum; but Huyghens, Lockyer, and others showed not only that this is the case for only a few stars, but that the majority of stars give spectra of a different character with dark and bright lines and bands.

[33] Eruptions, like our volcanic eruptions, but on an incomparably larger scale, are of frequent occurrence on the sun. They are seen as protuberances visible during a total eclipse of the sun, in the form of vaporous masses on the edge of the solar disc and emitting a faint light. These protuberances of the sun are now observed at all times by means of the spectroscope (Lockyer's method), because they contain luminous vapours (giving bright lines) of hydrogen and other elements.

[34] The great interest and vastness of astro-physical observations concerning the sun, comets, stars, nebulæ, &c., render this new province of natural science very important, and necessitate referring the reader to special works on the subject.

The most important astro-physical data since the time of Kellner are those referring to the _displacement_ of the lines of the spectrum. Just as a musical note changes its pitch with the approach or withdrawal of the resonant object or the ear, so the pitch of the luminous note or wave-length of the light varies if the luminous (or absorbent) vapour and the earth from which we observe it approach or recede from each other; this expresses itself in a visible displacement of the spectral lines. The solar eruptions even give broken lines in the spectrum, because the rapidly moving eruptive masses of vapour and gases either travel in the direction of the eye or fall back towards the sun. As the earth travels with the solar system among the stars, so it is possible to determine the direction and velocity with which the sun travels in space by the displacement of the spectral lines and light of the stars. The changes proceeding on the sun in its mass, which must be pronounced as vaporous, and in its atmosphere, are now studied by means of the spectroscope. For this purpose, many special astro-physical observatories now exist where these investigations are carried on.

We may remark that if the observer or luminous object moves with a velocity ±_v_, the ray, whose wave-length is [Greek: l], has an apparent wave-length [Greek: l](_n_±_v_)/_n_, where _n_ is the velocity of light. Thus Tolon, Huyghens, and others proved that the star Aldebaran approaches the solar system with a velocity of 30 kilometres per second, while Arcturus is receding with a velocity of 45 kilometres. The majority of stars give a distinct hydrogen spectrum, besides which nebulæ also give the spectrum of nitrogen. Lockyer classes the stars from their spectra, according to their period of formation, showing that some stars are in a period of increasing temperature (of formation or aggregation), whilst others are in a period of cooling. Altogether, in the astro-physical investigation of the spectra of heavenly bodies we find one of the most interesting subjects of recent science.

[35] Spectrum analysis has proved the indubitable existence in the sun and stars of a number of elements known in chemistry. Huyghens, Secchi, Lockyer, and others have furnished a large amount of material upon this subject. A compilation of existing information on it has been given by Prof. S. A. Kleiber, in the Journal of the Russian Physico-chemical Society for 1885 (vol. xviii. p. 146). Besides which, a peculiar element called helium has been discovered, which is characterised by a line (whose wave-length is 587·5, situated near D), which is seen very brightly in the projections (protuberances) and spots of the sun, but which does not belong to any known element, and is not reproducible as a reversed, dark line. This may be a right conclusion--that is to say, it is possible that an element may be discovered to which the spectrum of helium corresponds--but it may be that the helium line belongs to one of the known elements, because spectra vary in the brilliancy and position of their lines with changes of temperature and pressure. Thus, for instance, Lockyer could only see the line 423, at the very end of the calcium spectrum, at comparatively low temperatures, whilst the lines 397 and 393 appear at a higher temperature, and at a still higher temperature the line 423 becomes quite invisible.

[36] Spectroscopic observations are still further complicated by the fact that one and the same substance gives different spectra at different temperatures. This is especially the case with gases whose spectra are obtained by an electric discharge in tubes. Plücker, Wüllner, Schuster, and others showed that at low temperatures and pressures the spectra of iodine, sulphur, nitrogen, oxygen, &c. are quite different from the spectra of the same elements at high temperatures and pressures. This may either depend on the fact that the elements change their molecular structure with a change of temperature, just as ozone is converted into oxygen (for instance, from N_{2} molecules are obtained containing only one atom of nitrogen), or else it may be because at low temperature certain rays have a greater relative intensity than those which appear at higher temperatures. If we suppose that the molecules of a gas are in continual motion, with a velocity dependent on the temperature, then it must be admitted that they often strike against each other and rebound, and thus communicate peculiar motions to each other and the supposed ether, which express themselves in luminiferous phenomena. A rise of the temperature or an increase in the density of a gas must have an influence on the collision of its molecules and luminiferous motions thus produced, and this may be the cause of the difference of the spectra under these circumstances. It has been shown by direct experiment that gases compressed by pressure, when the collision of the molecules must be frequent and varied, exhibit a more complex spectrum on the passage of an electric spark than rarefied gases, and that even a continuous spectrum appears. In order to show the variability of the spectrum according to the circumstances under which it proceeds, it may be mentioned that potassium sulphate fused on a platinum wire gives, on the passage of a series of sparks, a distinct system of lines, 583-578, whilst when a series of sparks is passed through a solution of this salt this system of lines is faint, and when Roscoe and Schuster observed the absorption spectrum of the vapour of metallic potassium (which is green) they remarked a number of lines of the same intensity as the above system in the red, orange, and yellow portions.

The spectra of solutions are best observed by means of Lecoq de Boisbaudran's arrangement, shown in fig. 76. A bent capillary tube, D F, inside which a platinum wire, A _a_ (from 0·3 to 0·5 mm. in diameter) is fused, is immersed in a narrow cylinder, C (in which it is firmly held by a cork). The projecting end, _a_, of the wire is covered by a fine capillary tube, _d_, which extends 1-2 mm. beyond the wire. Another straight capillary tube, E, with a platinum wire, B _b_, about 1 mm. in diameter (a finer wire soon becomes hot), is held (by a cork or in a stand) above the end of the tube, D. If the wire A be now connected with the positive, and the wire B with the negative terminal of a Ruhmkorff's coil (if the wires be connected in the opposite order, the spectrum of air is obtained), a series of sparks rapidly following each other appear between _a_ and _b_, and their light may be examined by placing the apparatus in front of the slit of a spectroscope. The variations to which a spectrum is liable may easily be observed by increasing the distance between the wires, altering the direction of the current or strength of the solution, &c.

_Spectrum analysis_ has not only endowed science with a knowledge of the composition of distant heavenly bodies (of the sun, stars, nebulæ, comets, &c.), but has also given a new _method_ for studying the matter of the earth's surface. With its help Bunsen discovered two new elements belonging to the group of the alkali metals, and thallium, indium, and gallium were afterwards discovered by the same means. The spectroscope is employed in the study of rare metals (which in solution often give distinct absorption spectra), of dyes, and of many organic substances, &c.[37] With respect to the metals which are analogous to sodium, they all give similar very volatile salts and such very characteristic spectra that the least traces of them[38] are discovered with great ease by means of the spectroscope. For instance, _lithium_ gives a very brilliant red coloration to a flame and a very bright red spectral line (wave-length, 670 millionths mm.), which indicates the presence of this metal in admixture with compounds of other alkali metals.

[37] The importance of the spectroscope for the purpose of chemical research was already shown by Gladstone in 1856, but it did not become an accessory to the laboratory until after the discoveries of Kirchhoff and Bunsen. It may be hoped that in time spectroscopic researches will meet certain wants of the theoretical (philosophical) side of chemistry, but as yet all that has been done in this respect can only be regarded as attempts which have not yet led to any trustworthy conclusions. Thus many investigators, by collating the wave-lengths of all the light vibrations excited by a given element, endeavour to find the law governing their mutual relations; others (especially Hartley and Ciamician), by comparing the spectra of analogous elements (for instance, chlorine, bromine, and iodine), have succeeded in noticing definite features of resemblance in them, whilst others (Grünwald) search for relations between the spectra of compounds and their component elements, &c.; but--owing to the multiplicity of the spectral lines proper to many elements, and (especially in the ultra-red and ultra-violet ends of the spectrum) the existence of lines which are undistinguishable owing to their faintness, and also owing to the comparative novelty of spectroscopic research--this subject cannot be considered as in any way perfected. Nevertheless, in certain instances there is evidently some relationship between the wave-lengths of all the spectral lines formed by a given element. Thus, in the hydrogen spectrum the wave-length = 364·542 _m_^2/(_m_^{2}-4), if _m_ varies as a series of whole numbers from 3 to 15 (Walmer, Hagebach, and others). For example, when _m_ = 3, the wave-length of one of the brightest lines of the hydrogen spectrum is obtained (656·2), when _m_ = 7, one of the visible violet lines (396·8), and when _m_ is greater than 9, the ultra-violet lines of the hydrogen spectrum.

[38] In order to show the degree of sensitiveness of spectroscopic reactions the following observation of Dr. Bence Jones may be cited: If a solution of 3 grains of a lithium salt be injected under the skin of a guinea-pig, after the lapse of four minutes, lithium can be discovered in the bile and liquids of the eye, and, after ten minutes, in all parts of the animal.

_Lithium_, Li, is, like potassium and sodium, somewhat widely spread in siliceous rocks, but only occurs in small quantities and as mere traces in considerable masses of potassium and sodium salts. Only a very few rather rare minerals contain more than traces of it,[39] for example, spodumene and lithia mica. Many compounds of lithium are in all respects closely analogous to the corresponding compounds of sodium and potassium; but the _carbonate_ is sparingly soluble in cold water, which fact is taken advantage of for separating lithium from potassium and sodium. This salt, Li_{2}CO_{3}, is easily converted into the other compounds of lithium. Thus, for instance, the lithium hydroxide, LiHO, is obtained in exactly the same way as caustic soda, by the action of lime on the carbonate, and it is soluble in water and crystallises (from its solution in alcohol) as LiHO,H_{2}O. Metallic _lithium_ is obtained by the action of a galvanic current on fused lithium chloride; for this purpose a cast-iron crucible, furnished with a stout cover, is filled with lithium chloride, heated until the latter fuses, and a strong galvanic current is then passed through the molten mass. The positive pole (fig. 77) consists of a dense carbon rod C (surrounded by a porcelain tube P fixed in an iron tube BB), and the negative pole of an iron wire, on which the metal is deposited after the current has passed through the molten mass for a certain length of time. Chlorine is evolved at the positive pole. When a somewhat considerable quantity of the metal has accumulated on the wire it is withdrawn, the metal is collected from it, and the experiment is then carried on as before.[39 bis] Lithium is the lightest of all metals, its specific gravity is 0·59, owing to which it floats even on naphtha; it melts at 180°, but does not volatilise at a red heat. Its appearance recalls that of sodium, and, like it, it has a yellow tint. At 200° it burns in air with a very bright flame, forming lithium oxide. In decomposing water it does not ignite the hydrogen. The characteristic test for lithium compounds is the _red coloration_ which they impart to a colourless flame.[40]

[39] Thus _spodumene_ contains up to 6 p.c. of lithium oxide, and _petolite_, and _lepidolite_ or lithia mica, about 3 p.c. of lithium oxide. This mica is met with in certain granites in a somewhat considerable quantity, and is therefore most frequently employed for the preparation of lithium compounds. The treatment of lepidolite is carried on on a large scale, because certain salts of lithium are employed in medicine as a remedy for certain diseases (stone, gouty affections), as they have the power of dissolving the insoluble uric acid which is then deposited. Lepidolite, which is unacted on by acids in its natural state, decomposes under the action of strong hydrochloric acid after it has been fused. After being subjected to the action of the hydrochloric acid for several hours all the silica is obtained in an insoluble form, whilst the metallic oxides pass into solution as chlorides. This solution is mixed with nitric acid to convert the ferrous salts into ferric, and sodium carbonate is then added until the liquid becomes neutral, by which means a precipitate is formed of the oxides of iron, alumina, magnesia, &c., as insoluble oxides and carbonates. The solution (with an excess of water) then contains the chlorides of the alkaline metals KCl, NaCl, LiCl, which do not give a precipitate with sodium carbonate in a dilute solution. It is then evaporated, and a strong solution of sodium carbonate added. This precipitates lithium carbonate, which, although soluble in water, is much less so than sodium carbonate, and therefore the latter precipitates lithium from strong solutions as carbonate, 2LiCl + Na_{2}CO_{3} = 2NaCl + Li_{2}CO_{3}. _Lithium carbonate_, which resembles sodium carbonate in many respects, is a substance which is very slightly soluble in cold water and is only moderately soluble in boiling water. In this respect lithium forms a transition between the metals of the alkalis and other metals, especially those of the alkaline earths (magnesium, barium), whose carbonates are only sparingly soluble. Oxide of lithium, Li_{2}O, may be obtained by heating lithium carbonate with charcoal. Lithium oxide in dissolving gives (per gram-molecule) 26,000 heat units; but the combination of Li_{2} with O evolves 140,000 calories--that is, more than Na_{2}O (100,000 calories) and K_{2}O (97,000 calories), as shown by Beketoff (1887). Oeuvrard (1892) heated lithium to redness in nitrogen, and observed the absorption of N and formation of Li_{3}N, like Na_{3}N (_see_ Chapter XII. Note 50).

LiCl, LiBr, and LiI form crystallo-hydrates with H_{2}O, 2H_{2}O, and 3H_{2}O. As a rule, LiBr,2H_{2}O crystallises out, but Bogorodsky (1894) showed that a solution containing LiBr + 3·7H_{2}O, cooled to -62°, separates out crystals LiBr,3H_{2}O, which decompose at +4° with the separation of H_{2}O. LiF is but slightly soluble (in 800 parts) in water (and still less so in a solution of NH_{4}F).

[39 bis] Guntz (1893) recommends adding KCl to the LiCl in preparing Li by this method, and to act with a current of 10 ampères at 20 volts, and not to heat above 450°, so as to avoid the formation of Li_{2}Cl.

[40] In determining the presence of lithium in a given compound, it is best to treat the material under investigation with acid (in the case of mineral silicon compounds hydrofluoric acid must be taken), and to treat the residue with sulphuric acid, evaporate to dryness, and extract with alcohol, which dissolves a certain amount of the lithium sulphate. It is easy to discover lithium in such an alcoholic solution by means of the coloration imparted to the flame on burning it, and in case of doubt by investigating its light in a spectroscope, because lithium gives a red line, which is very characteristic and is found as a dark line in the solar spectrum. Lithium was first discovered in 1817 in petolite by Arfvedson.

Bunsen in 1860 tried to determine by means of the spectroscope whether any other as yet unknown metals might not occur in different natural products together with lithium, potassium, and sodium, and he soon discovered two new alkali metals showing independent spectra. They are named after the characteristic coloration which they impart to the flame. One which gives a red and violet band is named _rubidium_, from _rubidius_ (dark red), and the other is called _cæsium_, because it colours a pale flame sky blue, which depends on its containing bright blue rays, which appear in the spectrum of cæsium as two blue bands (table on p. 565). Both metals accompany sodium, potassium, and lithium, but in small quantities; rubidium occurs in larger quantity than cæsium. The amount of the oxides of cæsium and rubidium in lepidolite does not generally exceed one-half per cent. Rubidium has also been found in the ashes of many plants, while the Stassfurt carnallite (the mother-liquor obtained after having been treated for KCl) forms an abundant source for rubidium and also partly for cæsium. Rubidium also occurs, although in very small quantities, in the majority of mineral waters. In a very few cases cæsium is not accompanied by rubidium; thus, in a certain granite on the Isle of Elba, cæsium has been discovered, but not rubidium. This granite contains a very rare mineral called _pollux_, which contains as much as 34 per cent. of cæsium oxide. Guided by the spectroscope, and aided by the fact that the double salts of platinic chloride and rubidium and cæsium chlorides are still less soluble in water than the corresponding potassium salt, K_{2}PtCl_{6},[41] Bunsen succeeded in separating both metals from each other and from potassium, and demonstrated the great resemblance they bear to each other. The isolated metals,[42] rubidium and cæsium, have respectively the specific gravities 1·52 and 2·366, and melting points 39° and 27° as N. N. Beketoff showed (1894), he having obtained cæsium by heating CsAlO_{2} with Mg([42 bis]).

[41] The salts of the majority of metals are precipitated as carbonates on the addition of ammonium carbonate--for instance, the salts of calcium, iron, &c. The alkalis whose carbonates are soluble are not, however, precipitated in this case. On evaporating the resultant solution and igniting the residue (to remove the ammonium salts), we obtain salts of the alkali metals. They may he separated by adding hydrochloric acid together with a solution of platinic chloride. The chlorides of lithium and sodium give easily soluble double salts with platinic chloride, whilst the chlorides of potassium, rubidium, and cæsium form double salts which are sparingly soluble. A hundred parts of water at 0° dissolve 0·74 part of the potassium platinochloride; the corresponding rubidium platinochloride is only dissolved to the amount of 0·134 part, and the cæsium salt, 0·024 part; at 100° 5·13 parts of potassium platinochloride, K_{2}PtCl_{6}, are dissolved, 0·634 part of rubidium platinochloride, and 0·177 part of cæsium platinochloride. From this it is clear how the salts of rubidium and cæsium may be isolated. The separation of cæsium from rubidium by this method is very tedious. It can be better effected by taking advantage of the difference of the solubility of their carbonates in alcohol; cæsium carbonate, Cs_{2}CO_{3}, is soluble in alcohol, whilst the corresponding salts of rubidium and potassium are almost insoluble. Setterberg separated these metals as alums, but the best method, that given by Scharples, is founded on the fact that from a mixture of the chlorides of potassium, sodium, cæsium, and rubidium in the presence of hydrochloric acid, stannic chloride precipitates a double salt of cæsium, which is very slightly soluble. The salts of Rb and Cs are closely analogous to those of potassium.

[42] Bunsen obtained rubidium by distilling a mixture of the tartrate with soot, and Beketoff (1888) by heating the hydroxide with aluminium, 2RbHO + Al = RbAlO_{2} + H_{2} + Rb. By the action of 85 grams of rubidium on water, 94,000 heat units are evolved. Setterberg obtained cæsium (1882) by the electrolysis of a fused mixture of cyanide of cæsium and of barium. Winkler (1890) showed that metallic magnesium reduces the hydrates and carbonates of Rb and Cs like the other alkaline metals. N. N. Beketoff obtained them with aluminium (see following note).

[42 bis] Beketoff (1888) showed that metallic aluminium reduces the hydrates of the alkaline metals at a red heat (they should be perfectly dry) with the formation of aluminates (Chapter XVII.), RAlO_{2}--for example, 2KHO + Al = KAlO_{2} + K + H_{2}. It is evident that in this case only half of the alkaline metal is obtained free. On the other hand, K. Winkler (1889) showed that magnesium powder is also able to reduce the alkaline metals from their hydrates and carbonates. N. N. Beketoff and Tscherbacheff (1894) prepared cæsium upon this principle by heating its aluminate CsAlO_{2} with magnesium powder. In this case aluminate of magnesium is formed, and the whole of the cæsium is obtained as metal: 2CsAlO_{2} + Mg = MgOAl_{2}O_{5} + 2Cs. A certain excess of alumina was taken (in order to obtain a less hygroscopic mass of aluminate), and magnesium powder (in order to decompose the last traces of water); the CsAlO_{2} was prepared by the precipitation of cæsium alums by caustic baryta, and evaporating the resultant solution. We may add that N. N. Beketoff (1887) prepared oxide of potassium, K_{2}O, by heating the peroxide, KO, in the vapour of potassium (disengaged from its alloy with silver), and showed that in dissolving in an excess of water it evolves (for the above-given molecular weight) 67,400 calories (while 2KHO in dissolving in water evolves 24,920 cal.; so that K_{2}O + H_{2}O gives 42,480 cal.), whence (knowing that K_{2} + O+H_{2}O in an excess of water evolves 164,500) it follows that K_{2} + O evolves 97,100 cal. This quantity is somewhat less than that (100,260 cal.) which corresponds to sodium, and the energy of the action of potassium upon water is explained by the fact that K_{2}O evolves more heat than Na_{2}O in combining with water (_see_ Chapter II. Note 9). Just as hydrogen displaces half the Na from Na_{2}O forming NaHO, so also N. N. Beketoff found from experiment and thermochemical reasonings that hydrogen displaces half the potassium from K_{2}O forming KHO and evolving 7,190 calories. Oxide of lithium, Li_{2}O, which is easily formed by igniting Li_{2}CO_{3} with carbon (when Li_{2}O + 2CO is formed), disengages 26,000 cals. with an excess of water, while the reaction Li_{2} + O gives 114,000 cals. and the reaction Li_{2} + H_{2}O gives only 13,000 cals., and metallic lithium cannot be liberated from oxide of lithium with hydrogen (nor with carbon). Thus in the series Li, Na, K, the formation of R_{2}O gives most heat with Li and least with K, while the formation of RCl evolves most heat with K (105,000 cals.) and least of all with Li (93,500 cals.). Rubidium, in forming Rb_{2}O, gives 94,000 cals. (Beketoff). Cæsium, in acting upon an excess of water, evolves 51,500 cals., and the reaction Cs_{2} + O evolves about 100,000 cals.--_i.e._ more than K and Rb, and almost as much as Na--and oxide of cæsium reacts with hydrogen (according to the equation Cs_{2}O + H = CsHO + Cs) more easily than any of the oxides of the alkali metals, and this reaction takes place at the ordinary temperature (the hydrogen is absorbed), as Beketoff showed (1893). He also obtained a mixed oxide, AgCsO, which was easily formed in the presence of silver, and absorbed hydrogen with the formation of CsHO.

Judging by the properties of the free metals, and of their corresponding and even very complex compounds, lithium, sodium, potassium, rubidium, and cæsium present an indubitable chemical resemblance. The fact that the metals easily decompose water, and that their hydroxides RHO and carbonates R_{2}CO_{3} are soluble in water, whilst the hydroxides and carbonates of nearly all other metals are insoluble, shows that these metals form a natural group of _alkali metals_. The halogens and the alkali metals form, by their character, the two extremes of the elements. Many of the other elements are metals approaching the alkali metals, both in their capacity of forming salts and in not forming acid compounds, but are not so energetic as the alkali metals, that is, they form less energetic bases. Such are the common metals, silver, iron, copper, &c. Some other elements, in the character of their compounds, approach the halogens, and, like them, combine with hydrogen, but these compounds do not show the energetic property of the halogen acids; in a free state they easily combine with metals, but they do not then form such saline compounds as the halogens do--in a word, the halogen properties are less sharply defined in them than in the halogens themselves. Sulphur, phosphorus, arsenic, &c. belong to this order of elements. The clearest distinction of the properties of the halogens and alkali metals is expressed in the fact that the former give acids and do not form bases, whilst the latter, on the contrary, only give bases. The first are true _acid elements_, the latter clearly-defined _basic or metallic elements_. On combining together, the halogens form, in a chemical sense, unstable compounds, and the alkali metals alloys in which the character of the metals remains unaltered, just as in the compound ICl the character of the halogens remains undisguised; thus both classes of elements on combining with members of their own class form non-characteristic compounds, which have the properties of their components. On the other hand, the halogens on combining with the alkali metals form compounds which are, in all respects, stable, and in which the original characters of the halogens and alkali metals have entirely disappeared. The formation of such compounds is accompanied by evolution of a large amount of heat, and by an entire change of both the physical and chemical properties of the substances originally taken. The alloy of sodium and potassium, although liquid at the ordinary temperature, is perfectly metallic, like both its components. The compound of sodium and chlorine has neither the appearance nor the properties of the original elements; sodium chloride melts at a higher temperature, and is more difficultly volatile, than either sodium or chlorine.

With all these qualitative differences there is, however, an important quantitative _resemblance between the halogens and the alkali metals_. This resemblance is clearly expressed by stating that both orders of elements belong to those which are univalent with respect to hydrogen. It is thus correct to say that both the above-named orders of elements replace hydrogen atom for atom. Chlorine is able to take the place of hydrogen by metalepsis, and the alkali metals take the place of hydrogen in water and acids. As it is possible to consecutively replace every equivalent of hydrogen in a hydrocarbon by chlorine, so it is possible in an acid containing several equivalents of hydrogen to replace the hydrogen consecutively equivalent after equivalent by an alkali metal; hence an atom of these elements is analogous to an atom of hydrogen, which is taken, in all cases, as the unit for the comparison of the other elements. In ammonia, and in water, chlorine and sodium are able to bring about a direct replacement. According to the law of substitution, the formation of sodium chloride, NaCl, at once shows the equivalence of the atoms of the alkali metals and the halogens. The halogens and hydrogen and the alkali metals combine with such elements as oxygen, and it is easily proved that in such compounds one atom of oxygen is able to retain two atoms of the halogens, of hydrogen, and of the alkali metals. For this purpose it is enough to compare the compounds KHO, K_{2}O, HClO, and Cl_{2}O, with water. It must not be forgotten, however, that the halogens give, with oxygen, besides compounds of the type R_{2}O, higher acid grades of oxidation, which the alkali metals and hydrogen are not capable of forming. We shall soon see that these relations are also subject to a special law, showing a gradual transition of the properties of the elements from the alkali metals to the halogens.[43]

[43] We may here observe that the halogens, and especially iodine, may play the part of metals (hence iodine is more easily replaced by metals than the other halogens, and it approaches nearer to the metals in its physical properties than the other halogens). Schützenberger obtained a compound C_{2}H_{3}O(OCl), which he called chlorine acetate, by acting on acetic anhydride, (C_{2}H_{3}O)_{2}O, with chlorine monoxide, Cl_{2}O. With iodine this compound gives off chlorine and forms iodine acetate, C_{2}H_{3}O(OI), which also is formed by the action of iodine chloride on sodium acetate, C_{2}H_{3}O(ONa). These compounds are evidently nothing else than mixed anhydrides of hypochlorous and hypoiodous acids, or the products of the substitution of hydrogen in RHO by a halogen (_see_ Chapter XI., Notes 29 and 78 bis). Such compounds are very unstable, decompose with an explosion when heated, and are changed by the action of water and of many other reagents, which is in accordance with the fact that they contain very closely allied elements, as does Cl_{2}O itself, or ICl or KNa. By the action of chlorine monoxide on a mixture of iodine and acetic anhydride, Schützenberger also obtained the compound I(C_{2}H_{3}O_{2})_{3}, which is analogous to ICl_{3}, because the group C_{2}H_{3}O_{2} is, like Cl, a halogen, forming salts with the metals. Similar properties are found in iodosobenzene (Chapter XI., Note 79).

The atomic weights of the alkali metals, lithium 7, sodium 23, potassium 39, rubidium 85, and cæsium 133, show that here, as in the class of halogens, the elements may be arranged according to their atomic weights in order to compare the properties of the analogous compounds of the members of this group. Thus, for example, the platinochlorides of lithium and sodium are soluble in water; those of potassium, rubidium, and cæsium sparingly soluble, and the greater the atomic weight of the metal the less soluble is the salt. In other cases the reverse is observed--the greater the atomic weight the more soluble are the corresponding salts. The variation of properties with the variation in atomic weights even shows itself in the metals themselves; thus lithium volatilises with difficulty, whilst sodium is obtained by distillation, potassium volatilises more easily than sodium, and rubidium and cæsium as we have seen, are still more volatile.