BARIUM (symbol Ba, atomic weight 137.37 [O=16]), one of the metallic chemical elements included in the group of the alkaline earths. It takes its name from the Greek "barus" (heavy) on account of its presence in barytes or heavy spar which was first investigated in 1602 by V. Casciorolus, a shoemaker of Bologna, who found that after ignition with combustible substances it became phosphorescent, and on this account it was frequently called Bolognian phosphorus. In 1774 K. W. Scheele, in examining a specimen of pyrolusite, found a new substance to be present in the mineral, for on treatment with sulphuric acid it gave an insoluble salt which was afterwards shown to be identical with that contained in heavy spar. Barium occurs chiefly in the form of barytes or heavy spar, BaSO4, and witherite, BaCO3, and to a less extent in baryto-calcite, baryto-celestine, and various complex silicates. The metal is difficult to isolate, and until recently it may be doubted whether the pure metal had been obtained. Sir H. Davy tried to electrolyse baryta, but was unsuccessful; later attempts were made by him using barium chloride in the presence of mercury. In this way he obtained an amalgam, from which on distilling off the mercury the barium was obtained as a silver white residue. R. Bunsen in 1854 electrolysed a thick paste of barium chloride and dilute hydrochloric acid in the presence of mercury, at 100° C., obtaining a barium amalgam, from which the mercury was separated by a process of distillation. A. N. Guntz (Comptes rendus, 1901, 133, p. 872) electrolyses a saturated solution of barium chloride using a mercury cathode and obtains a 3% barium amalgam; this amalgam is transferred to an iron boat in a wide porcelain tube and the tube slowly heated electrically, a good yield of pure barium being obtained at about 1000° C. The metal when freshly cut possesses a silver white lustre, is a little harder than lead, and is extremely easily oxidized on exposure; it is soluble in liquid ammonia, and readily attacks both water and alcohol.
Three oxides of barium are known, namely, the monoxide, BaO, the dioxide, BaO2, and a suboxide, obtained by heating BaO with magnesium in a vacuum to 1100° (Guntz, loc. cit., 1906, p. 359). The monoxide is formed when the metal burns in air, but is usually prepared by the ignition of the nitrate, oxygen and oxides of nitrogen being liberated. It can also be obtained by the ignition of an intimate mixture of the carbonate and carbon, and in small quantities by the ignition of the iodate. It is a greyish coloured solid, which combines very energetically with water to form the hydroxide, much heat being evolved during the combination; on heating to redness in a current of oxygen it combines with the oxygen to form the dioxide, which at higher temperatures breaks up again into the monoxide and oxygen.
Barium hydroxide, Ba(OH)2, is a white powder that can be obtained by slaking the monoxide with the requisite quantity of water, but it is usually made on the large scale by heating heavy spar with small coal whereby a crude barium sulphide is obtained. This sulphide is then heated in a current of moist carbon dioxide, barium carbonate being formed, BaS + H2O + CO2 = BaCO3 + H2S, and finally the carbonate is decomposed by a current of superheated steam, BaCO3 + H2O = Ba(OH)2 + CO2, leaving a residue of the hydroxide. It is a white powder moderately soluble in cold water, readily soluble in hot water, the solution possessing an alkaline reaction and absorbing carbon dioxide readily. The solution, known as baryta-water, finds an extensive application in practical chemistry, being used in gas-analysis for the determination of the amount of carbon dioxide in the atmosphere; and also being used in organic chemistry as a hydrolysing agent for the decomposition of complex ureides and substituted aceto-acetic esters, while E. Fischer has used it as a condensing agent in the preparation of - and β-acrose from acrolein dibromide. A saturated solution of the hydroxide deposits on cooling a hydrated form Ba(OH)2 · 8H2O, as colourless quadratic prisms, which on exposure to air lose seven molecules of water of crystallization.
Barium dioxide, BaO2, can be prepared as shown above, or in the hydrated condition by the addition of excess of baryta-water to hydrogen peroxide solution, when it is precipitated in the crystalline condition as BaO2 · 8H2O. These crystals on heating to 130° C. lose the water of crystallization and leave a residue of the anhydrous peroxide. In the Brin process for the manufacture of oxygen, barium dioxide is obtained as an intermediate product by heating barium monoxide with air under pressure. It is a grey coloured powder which is readily decomposed by dilute acids with the production of hydrogen peroxide.
Barium chloride, BaCl2 · 2H2O, can be obtained by dissolving witherite in dilute hydrochloric acid, and also from heavy spar by ignition in a reverberatory furnace with a mixture of coal, limestone and calcium chloride, the barium chloride being extracted from the fused mass by water, leaving a residue of insoluble calcium sulphide. The chloride crystallizes in colourless rhombic tables of specific gravity 3.0 and is readily soluble in water, but is almost insoluble in concentrated hydrochloric acid and in absolute alcohol. It can be obtained in the anhydrous condition by heating it gently to about 120° C. It has a bitter taste and is a strong poison. Barium bromide is prepared by saturating baryta-water or by decomposing barium carbonate with hydrobromic acid. It crystallizes as BaBr2 · 2H2O isomorphous with barium chloride. Barium bromate, Ba(BrO3)2, can be prepared by the action of excess of bromine on baryta-water, or by decomposing a boiling aqueous solution of 100 parts of potassium bromate with a similar solution of 74 parts of crystallized barium chloride. It crystallizes in the monoclinic system, and separates from its aqueous solution as Ba(BrO3)2 · H2O. On heating, it begins to decompose at 260-265° C. Barium chlorate, Ba(ClO3)2, is obtained by adding barium chloride to sodium chlorate solution; on concentration of the solution sodium chloride separates first, and then on further evaporation barium chlorate crystallizes out and can be purified by recrystallization. It can also be obtained by suspending barium carbonate in boiling water and passing in chlorine. It crystallizes in monoclinic prisms of composition Ba(ClO3)2 · H2O, and begins to decompose on being heated to 250° C. Barium iodate, Ba(IO3)2, is obtained by the action of excess of iodic acid on hot caustic baryta solution or by adding sodium iodate to barium chloride solution. It crystallizes in monoclinic prisms of composition Ba(IO3)2 · H2O, and is only very sparingly soluble in cold water.
Barium carbide, BaC2, is prepared by a method similar to that in use for the preparation of calcium carbide (see Acetylene). L. Maquenne has also obtained it by distilling a mixture of barium amalgam and carbon in a stream of hydrogen. Barium sulphide, BaS, is obtained by passing sulphuretted hydrogen over heated barium monoxide, or better by fusion of the sulphate with a small coal. It is a white powder which is readily decomposed by water with the formation of the hydroxide and hydrosulphide. The phosphorescence of the sulphide obtained by heating the thiosulphate is much increased by adding uranium, bismuth, or thorium before ignition (J. pr. Chem., 1905, ii. p. 196).
Barium sulphate, BaSO4, is the most abundant of the naturally occurring barium compounds (see Barytes) and can be obtained artificially by the addition of sulphuric acid or any soluble sulphate to a solution of a soluble barium salt, when it is precipitated as an amorphous white powder of specific gravity 4.5. It is practically insoluble in water, and is only very slightly soluble in dilute acids; it is soluble to some extent, when freshly prepared, in hot concentrated sulphuric acid, and on cooling the solution, crystals of composition BaSO4 · H2SO4 are deposited. It is used as a pigment under the name of "permanent white" or blanc fixe.
Barium nitride, Ba3N2, is obtained as a brownish mass by passing nitrogen over heated barium amalgam. It is decomposed by water with evolution of hydrogen, and on heating in a current of carbonic oxide forms barium cyanide (L. Maquenne). Barium amide, Ba(NH2)2, is obtained from potassammonium and barium bromide.
Barium nitrate, Ba(NO3)2, is prepared by dissolving either the carbonate or sulphide in dilute nitric acid, or by mixing hot saturated solutions of barium chloride and sodium nitrate. It crystallizes in octahedra, having a specific gravity of 3.2, and melts at 597° C. (T. Carnelley). It is decomposed by heat, and is largely used in pyrotechny for the preparation of green fire. Barium carbonate, BaCO3, occurs rather widely distributed as witherite (q.v.), and may be prepared by the addition of barium chloride to a hot solution of ammonium carbonate, when it is precipitated as a dense white powder of specific gravity 4.3; almost insoluble in water.
Barium and its salts can be readily detected by the yellowish-green colour they give when moistened with hydrochloric acid and heated in the Bunsenflame, or by observation of their spectra, when two characteristic green lines are seen. In solution, barium salts may be detected by the immediate precipitate they give on the addition of calcium sulphate (this serves to distinguish barium salts from calcium salts), and by the yellow precipitate of barium chromate formed on the addition of potassium chromate. Barium is estimated quantitatively by conversion into the sulphate. The atomic weight of the element has been determined by C. Marignac by the conversion of barium chloride into barium sulphate, and also by a determination of the amount of silver required to precipitate exactly a known weight of the chloride; the mean value obtained being 136.84; T. W. Richards (Zeit. anorg. Chem., 1893, 6, p. 89), by determining the equivalent of barium chloride and bromide to silver, obtained the value 137.44. For the relation of barium to radium, see Radioactivity.
Note - this article incorporates content from Encyclopaedia Britannica, Eleventh Edition, (1910-1911)