[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 Li2Cl.
[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.
[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, K2PtCl6, 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, Cs2CO3, 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 = RbAlO2 + H2 + 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.), RAlO2—for example, 2KHO + Al = KAlO2 + K + H2. 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 CsAlO2 with magnesium powder. In this case aluminate of magnesium is formed, and the whole of the cæsium is obtained as metal: 2CsAlO2 + Mg = MgOAl2O5 + 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 CsAlO2 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, K2O, 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 K2O + H2O gives 42,480 cal.), whence (knowing that K2 + O + H2O in an excess of water evolves 164,500) it follows that K2 + 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 K2O evolves more heat than Na2O in combining with water (see Chapter II. Note [9]). Just as hydrogen displaces half the Na from Na2O forming NaHO, so also N. N. Beketoff found from experiment and thermochemical reasonings that hydrogen displaces half the potassium from K2O forming KHO and evolving 7,190 calories. Oxide of lithium, Li2O, which is easily formed by igniting Li2CO3 with carbon (when Li2O + 2CO is formed), disengages 26,000 cals. with an excess of water, while the reaction Li2 + O gives 114,000 cals. and the reaction Li2 + H2O 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 R2O 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 Rb2O, gives 94,000 cals. (Beketoff). Cæsium, in acting upon an excess of water, evolves 51,500 cals., and the reaction Cs2 + 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 Cs2O + 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.
[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 C2H3O(OCl), which he called chlorine acetate, by acting on acetic anhydride, (C2H3O)2O, with chlorine monoxide, Cl2O. With iodine this compound gives off chlorine and forms iodine acetate, C2H3O(OI), which also is formed by the action of iodine chloride on sodium acetate, C2H3O(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 Cl2O 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(C2H3O2)3, which is analogous to ICl3, because the group C2H3O2 is, like Cl, a halogen, forming salts with the metals. Similar properties are found in iodosobenzene (Chapter XI., Note [79]).
CHAPTER XIV
THE VALENCY AND SPECIFIC HEAT OF THE METALS. MAGNESIUM. CALCIUM, STRONTIUM, BARIUM, AND BERYLLIUM
It is easy by investigating the composition of corresponding compounds, to establish the equivalent weights of the metals compared with hydrogen—that is, the quantity which replaces one part by weight of hydrogen. If a metal decomposes acids directly, with the evolution of hydrogen, the equivalent weight of the metal may be determined by taking a definite weight of it and measuring the volume of hydrogen evolved by its action on an excess of acid; it is then easy to calculate the weight of the hydrogen from its volume.[1] The same result may be arrived at by determining the composition of the normal salts of the metal; for instance, by finding the weight of metal which combines with 35·5 parts of chlorine or 80 parts of bromine.[2] The equivalent of a metal may be also ascertained by simultaneously (i.e. in one circuit) decomposing an acid and a fused salt of a given metal by an electric current and determining the relation between the amounts of hydrogen and metal separated, because, according to Faraday's law, electrolytes (conductors of the second order) are always decomposed in equivalent quantities.[2 bis] The equivalent of a metal may even be found by simply determining the relation between its weight and that of its salt-giving oxide, as by this we know the quantity of the metal which combines with 8 parts by weight of oxygen, and this will be the equivalent, because 8 parts of oxygen combine with 1 part by weight of hydrogen. One method is verified by another, and all the processes for the accurate determination of equivalents require the greatest care to avoid the absorption of moisture, further oxidation, volatility, and other accidental influences which affect exact weighings. The description of the methods necessary for the attainment of exact results belongs to the province of analytical chemistry.