see captionFig.77.—Preparation of lithium by the action of a galvanic current on fused lithium chloride.
Fig.77.—Preparation of lithium by the action of a galvanic current on fused lithium chloride.
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 sodiumand potassium; but thecarbonateis sparingly soluble in cold water, which fact is taken advantage of for separating lithium from potassium and sodium. This salt, Li2CO3, 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,H2O. Metalliclithiumis 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 When a somewhat considerable quantity of the metal has accumulated on the wire it is withdrawn, the metal is collected from it, and the experiment is then carried on as before.[39 bis]Lithium is the lightest of all metals, its specific gravity is 0·59,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 thered colorationwhich they impart to a colourless flame.[40]
Bunsen in 1860 tried to determine by means of the spectroscopewhether 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 namedrubidium, fromrubidius(dark red), and the other is calledcæsium, 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 calledpollux, 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, K2PtCl6,[41]Bunsen succeeded in separating both metals from each other and from potassium, and demonstrated the great resemblancethey 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 CsAlO2with Mg([42 bis]).
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 theirhydroxides RHO and carbonates R2CO3are soluble in water, whilst the hydroxides and carbonates of nearly all other metals are insoluble, shows that these metals form a natural group ofalkali 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 trueacid elements, the latter clearly-definedbasic 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 quantitativeresemblance 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 elementsreplace 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, K2O, HClO, and Cl2O, with water. It must not be forgotten, however, that the halogens give, with oxygen, besides compounds of the type R2O, 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]
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; thoseof 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.