1911 Encyclopædia Britannica/Caesium
CAESIUM (symbol Cs, atomic weight 132.9), one of the alkali metals. Its name is derived from the Lat. caesius, sky-blue, from two bright blue lines of its spectrum. It is of historical importance, since it was the first metal to be discovered by the aid of the spectroscope (R. Bunsen, Berlin Acad. Ber., 1860), although caesium salts had undoubtedly been examined before, but had been mistaken for potassium salts (see C.F. Plattner, Pog. Ann., 1846, p. 443, on the analysis of pollux and the subsequent work of F. Pisani, Comptes Rendus, 1864, 58, p. 714). Caesium is found in the mineral springs of Frankenhausen, Montecatini, di Val di Nievole, Tuscany, and Wheal Clifford near Redruth, Cornwall (W.A. Miller, Chem. News, 1864, 10, p. 181), and, associated with rubidium, at Dürkheim; it is also found in lepidolite, leucite, petalite, triphylline and in the carnallite from Stassfurt. The separation of caesium from the minerals which contain it is an exceedingly difficult and laborious process. According to R. Bunsen, the best source of rubidium and caesium salts is the residue left after extraction of lithium salts from lepidolite. This residue consists of sodium, potassium and lithium chlorides, with small quantities of caesium and rubidium chlorides. The caesium and rubidium are separated from this by repeated fractional crystallization of their double platinum chlorides, which are much less soluble in water than those of the other alkali metals (R. Bunsen, Ann., 1862, 122, p. 347; 1863, 125, p. 367). The platino-chlorides are reduced by hydrogen, and the caesium and rubidium chlorides extracted by water. See also A. Schrötter (Jour. prak. Chem., 1864, 93, p. 2075) and W. Heintz (Journ. prak. Chem., 1862, 87, p. 310). W. Feit and K. Kubierschky (Chem. Zeit., 1892, 16, p. 335) separate rubidium and caesium from the other alkali metals by converting them into double chlorides with stannic chloride; whilst J. Redtenbacher (Jour. prak. Chem., 1865, 94, p. 442) separates them from potassium by conversion into alums, which C. Setterberg (Ann., 1882, 211, p. 100) has shown are very slightly soluble in a solution of potash alum. In order to separate caesium from rubidium, use is made of the different solubilities of their various salts. The bitartrates RbHC4H406 and CsHC4H406 have been employed, as have also the alums (see above). The double chloride of caesium and antimony 3CsCl·2SbCl3 (R. Godeffroy, Ber., 1874, 7, p. 375; Ann., 1876, 181, p. 176) has been used, the corresponding compound not being formed by rubidium. The metal has been obtained by electrolysis of a mixture of caesium and barium cyanides (C. Setterberg, Ann., 1882, 211, p. 100) and by heating the hydroxide with magnesium or aluminium (N. Beketoff, Chem. Centralblatt, 1889, 2, p. 245). L. Hackspill (Comptes Rendus, 1905, 141, p. 101) finds that metallic caesium can be obtained more readily by heating the chloride with metallic calcium. A special V-shaped tube is used in the operation, and the reaction commences between 400°C. and 500°C. It is a silvery white metal which burns on heating in air. It melts at 26° to 27°C. and has a specific gravity of 1.88 (15°C.).
The atomic weight of caesium has been determined by the analysis of its chloride and bromide. Richards and Archibald (Zeit. anorg. Chem., 1903, 34, p. 353) obtained 132.879 (O=16).
Caesium hydroxide, Cs(OH)2, obtained by the decomposition of the sulphate with baryta water, is a greyish-white deliquescent solid, which melts at a red heat and absorbs carbon dioxide rapidly. It readily dissolves in water, with evolution of much heat. Caesium chloride, CsCl, is obtained by the direct action of chlorine on caesium, or by solution of the hydroxide in hydrochloric acid. It forms small cubes which melt at a red heat and volatilize readily. It deliquesces in moist air. Many double chlorides are known, and may be prepared by mixing solutions of the two components in the requisite proportions. The bromide, CsBr, and iodide, CsI, resemble the corresponding potassium salts. Many trihaloid salts of caesium are also known, such as CsBr3, CsClBr2, CsI3, CsBrI2, CsBr2I, &c. (H.L. Wells and S.L. Penfield, Zeit. fur anorg. Chem., 1892, i, p. 85). Caesium sulphate, Cs2SO4, may be prepared by dissolving the hydroxide or carbonate in sulphuric acid. It crystallizes in short hard prisms, which are readily soluble in water but insoluble in alcohol. It combines with many metallic sulphates (silver, zinc, cobalt, nickel, &c.) to form double sulphates of the type Cs2SO4·RSO4·6H2O. It also forms a caesium-alum Cs2SO4·Al2(SO4)3·24H2O. Caesium nitrate, CsNO3, is obtained by dissolving the carbonate in nitric acid, and crystallizes in glittering prisms, which melt readily, and on heating evolve oxygen and leave a residue of caesium nitrite. The carbonate, Cs2CO3, silicofluoride, Cs2SiF6, borate, Cs2O·3B2O3, and the sulphides Cs2S·4H2O, Cs2S2·H2O, Cs2S3·H2O, Cs2S4 and Cs2S6·H2O, are also known.
Caesium compounds can be readily recognized by the two bright blue lines (of wave length 4555 and 4593) in their flame spectrum, but these are not present in the spark spectrum. The other lines include three in the green, two in the yellow, and two in the orange.