Hydrogen

Hydrogen [symbol H, atomic weight 1.008 (o = 16)], one of the chemical elements. Its name is derived from Gr. ὕδωρ, water, and γεννάειν, to produce, in allusion to the fact that water is produced when the gas burns in air. Hydrogen appears to have been recognized by Paracelsus in the 16th century; the combustibility of the gas was noticed by Turquet de Mayenne in the 17th century, whilst in 1700 N. Lémery showed that a mixture of hydrogen and air detonated on the application of a light. The first definite experiments concerning the nature of hydrogen were made in 1766 by H. Cavendish, who showed that it was formed when various metals were acted upon by dilute sulphuric or hydrochloric acids. Cavendish called it “inflammable air,” and for some time it was confused with other inflammable gases, all of which were supposed to contain the same inflammable principle, “phlogiston,” in combination with varying amounts of other substances. In 1781 Cavendish showed that water was the only substance produced when hydrogen was burned in air or oxygen, it having been thought previously to this date that other substances were formed during the reaction, A. L. Lavoisier making many experiments with the object of finding an acid among the products of combustion.

Hydrogen is found in the free state in some volcanic gases, in fumaroles, in the carnallite of the Stassfurt potash mines (H. Precht, Ber., 1886, 19, p. 2326), in some meteorites, in certain stars and nebulae, and also in the envelopes of the sun. In combination it is found as a constituent of water, of the gases from certain mineral springs, in many minerals, and in most animal and vegetable tissues. It may be prepared by the electrolysis of acidulated water, by the decomposition of water by various metals or metallic hydrides, and by the action of many metals on acids or on bases. The alkali metals and alkaline earth metals decompose water at ordinary temperatures; magnesium begins to react above 70° C., and zinc at a dull red heat. The decomposition of steam by red hot iron has been studied by H. Sainte-Claire Deville (Comptes rendus, 1870, 70, p. 1105) and by H. Debray (ibid., 1879, 88, p. 1341), who found that at about 1500° C. a condition of equilibrium is reached. H. Moissan (Bull. soc. chim., 1902, 27, p. 1141) has shown that potassium hydride decomposes cold water, with evolution of hydrogen, KH + H2O = KOH + H2. Calcium hydride or hydrolite, prepared by passing hydrogen over heated calcium, decomposes water similarly, 1 gram giving 1 litre of gas; it has been proposed as a commercial source (Prats Aymerich, Abst. J.C.S., 1907, ii. p. 543), as has also aluminium turnings moistened with potassium cyanide and mercuric chloride, which decomposes water regularly at 70°, 1 gram giving 1.3 litres of gas (Mauricheau-Beaupré, Comptes rendus, 1908, 147, p. 310). Strontium hydride behaves similarly. In preparing the gas by the action of metals on acids, dilute sulphuric or hydrochloric acid is taken, and the metals commonly used are zinc or iron. So obtained, it contains many impurities, such as carbon dioxide, nitrogen, oxides of nitrogen, phosphoretted hydrogen, arseniuretted hydrogen, &c., the removal of which is a matter of great difficulty (see E. W. Morley, Amer. Chem. Journ., 1890, 12, p. 460). When prepared by the action of metals on bases, zinc or aluminium and caustic soda or caustic potash are used. Hydrogen may also be obtained by the action of zinc on ammonium salts (the nitrate excepted) (Lorin, Comptes rendus, 1865, 60, p. 745) and by heating the alkali formates or oxalates with caustic potash or soda, Na2C2O4 \+ 2NaOH = H2 \+ 2Na2CO3. Technically it is prepared by the action of superheated steam on incandescent coke (see F. Hembert and Henry, Comptes rendus, 1885, 101, p. 797; A. Naumann and C. Pistor, Ber., 1885, 18, p. 1647), or by the electrolysis of a dilute solution of caustic soda (C. Winssinger, Chem. Zeit., 1898, 22, p. 609; “Die Elektrizitäts-Aktiengesellschaft,” Zeit. f. Elektrochem., 1901, 7, p. 857). In the latter method a 15% solution of caustic soda is used, and the electrodes are made of iron; the cell is packed in a wooden box, surrounded with sand, so that the temperature is kept at about 70° C.; the solution is replenished, when necessary, with distilled water. The purity of the gas obtained is about 97%.

Pure hydrogen is a tasteless, colourless and odourless gas of specific gravity 0.06947 (air = 1) (Lord Rayleigh, Proc. Roy. Soc., 1893, p. 319). It may be liquefied, the liquid boiling at −252.68° C. to −252.84° C., and it has also been solidified, the solid melting at −264° C. (J. Dewar, Comptes rendus, 1899, 129, p. 451; Chem. News, 1901, 84, p. 49; see also Liquid Gases). The specific heat of gaseous hydrogen (at constant pressure) is 3.4041 (water = 1), and the ratio of the specific heat at constant pressure to the specific heat at constant volume is 1.3852 (W. C. Röntgen, Pogg. Ann., 1873, 148, p. 580). On the spectrum see Spectroscopy. Hydrogen is only very slightly soluble in water. It diffuses very rapidly through a porous membrane, and through some metals at a red heat (T. Graham, Proc. Roy. Soc., 1867, 15, p. 223; H. Sainte-Claire Deville and L. Troost, Comptes rendus, 1863, 56, p. 977). Palladium and some other metals are capable of absorbing large volumes of hydrogen (especially when the metal is used as a cathode in a water electrolysis apparatus). L. Troost and P. Hautefeuille (Ann. chim. phys., 1874, (5) 2, p. 279) considered that a palladium hydride of composition Pd2H was formed, but the investigations of C. Hoitsema (Zeit. phys. Chem., 1895, 17, p. 1), from the standpoint of the phase rule, do not favour this view, Hoitsema being of the opinion that the occlusion of hydrogen by palladium is a process of continuous absorption. Hydrogen burns with a pale blue non-luminous flame, but will not support the combustion of ordinary combustibles. It forms a highly explosive mixture with air or oxygen, especially when in the proportion of two volumes of hydrogen to one volume of oxygen. H. B. Baker (Proc. Chem. Soc., 1902, 18, p. 40) has shown that perfectly dry hydrogen will not unite with perfectly dry oxygen. Hydrogen combines with fluorine, even at very low temperatures, with great violence; it also combines with carbon, at the temperature of the electric arc. The alkali metals when warmed in a current of hydrogen, at about 360° C., form hydrides of composition RH (R = Na, K, Rb, Cs), (H. Moissan, Bull. soc. chim., 1902, 27, p. 1141); calcium and strontium similarly form hydrides CaH2, SrH2 at a dull red heat (A. Guntz, Comptes rendus, 1901, 133, p. 1209). Hydrogen is a very powerful reducing agent; the gas occluded by palladium being very active in this respect, readily reducing ferric salts to ferrous salts, nitrates to nitrites and ammonia, chlorates to chlorides, &c.

For determinations of the volume ratio with which hydrogen and oxygen combine, see J. B. Dumas, Ann. chim. phys., 1843 (3), 8, p. 189; O. Erdmann and R. F. Marchand, ibid., p. 212; E. H. Keiser, Ber., 1887, 20, p. 2323; J. P. Cooke and T. W. Richards, Amer. Chem. Journ., 1888, 10, p. 191; Lord Rayleigh, Chem. News, 1889, 59, p. 147; E. W. Morley, Zeit. phys. Chem., 1890, 20, p. 417; and S. A. Leduc, Comptes rendus, 1899, 128, p. 1158.

Hydrogen combines with oxygen to form two definite compounds, namely, water (q.v.), H2O, and hydrogen peroxide, H2O2, whilst the existence of a third oxide, ozonic acid, has been indicated.

Hydrogen peroxide, H2O2, was discovered by L. J. Thénard in 1818 (Ann. chim. phys., 8, p. 306). It occurs in small quantities in the atmosphere. It may be prepared by passing a current of carbon dioxide through ice-cold water, to which small quantities of barium peroxide are added from time to time (F. Duprey, Comptes rendus, 1862, 55, p. 736; A. J. Balard, ibid., p. 758), BaO2 \+ CO2 \+ H2O = H2O2 \+ BaCO3. E. Merck (Abst. J.C.S., 1907, ii., p. 859) showed that barium percarbonate, BaCO4, is formed when the gas is in excess; this substance readily yields the peroxide with an acid. Or barium peroxide may be decomposed by hydrochloric, hydrofluoric, sulphuric or silicofluoric acids (L. Crismer, Bull. soc. chim., 1891 (3), 6, p. 24; Hanriot, Comptes rendus, 1885, 100, pp. 56, 172), the peroxide being added in small quantities to a cold dilute solution of the acid. It is necessary that it should be as pure as possible since the commercial product usually contains traces of ferric, manganic and aluminium oxides, together with some silica. To purify the oxide, it is dissolved in dilute hydrochloric acid until the acid is neatly neutralized, the solution is cooled, filtered, and baryta water is added until a faint permanent white precipitate of hydrated barium peroxide appears; the solution is now filtered, and a concentrated solution of baryta water is added to the filtrate, when a crystalline precipitate of hydrated barium peroxide, BaO2·H2O, is thrown down. This is filtered off and well washed with water. The above methods give a dilute aqueous solution of hydrogen peroxide, which may be concentrated somewhat by evaporation over sulphuric acid in vacuo. H. P. Talbot and H. R. Moody (Jour. Anal. Chem., 1892, 6, p. 650) prepared a more concentrated solution from the commercial product, by the addition of a 10% solution of alcohol and baryta water. The solution is filtered, and the barium precipitated by sulphuric acid. The alcohol is removed by distillation in vacuo, and by further concentration in vacuo a solution may be obtained which evolves 580 volumes of oxygen. R. Wolffenstein (Ber., 1894, 27, p. 2307) prepared practically anhydrous hydrogen peroxide (containing 99.1% H2O2) by first removing all traces of dust, heavy metals and alkali from the commercial 3% solution. The solution is then concentrated in an open basis on the water-bath until it contains 48% H2O2. The liquid so obtained is extracted with ether and the ethereal solution distilled under diminished pressure, and finally purified by repeated distillations. W. Staedel (Zeit. f. angew. Chem., 1902, 15, p. 642) has described solid hydrogen peroxide, obtained by freezing concentrated solutions.

Hydrogen peroxide is also found as a product in many chemical actions, being formed when carbon monoxide and cyanogen burn in air (H. B. Dixon); by passing air through solutions of strong bases in the presence of such metals as do not react with the bases to liberate hydrogen; by shaking zinc amalgam with alcoholic sulphuric acid and air (M. Traube, Ber., 1882, 15, p. 659); in the oxidation of zinc, lead and copper in presence of water, and in the electrolysis of sulphuric acid of such strength that it contains two molecules of water to one molecule of sulphuric acid (M. Berthelot, Comptes rendus, 1878, 86, p. 71).

The anhydrous hydrogen peroxide obtained by Wolffenstein boils at 84-85°C. (68 mm.); its specific gravity is 1.4996 (1.5° C.). It is very explosive (W. Spring, Zeit. anorg. Chem., 1895, 8, p. 424). The explosion risk seems to be most marked in the preparations which have been extracted with ether previous to distillation, and J. W. Brühl (Ber., 1895, 28, p. 2847) is of opinion that a very unstable, more highly oxidized product is produced in small quantity in the process. The solid variety prepared by Staedel forms colourless, prismatic crystals which melt at −2° C.; it is decomposed with explosive violence by platinum sponge, and traces of manganese dioxide. The dilute aqueous solution is very unstable, giving up oxygen readily, and decomposing with explosive violence at 100° C. An aqueous solution containing more than 1.5% hydrogen peroxide reacts slightly acid. Towards lupetidin [aa′ dimethyl piperidine, C5H9N(CH3)2] hydrogen peroxide acts as a dibasic acid (A. Marcuse and R. Wolffenstein, Ber., 1901, 34, p. 2430; see also G. Bredig, Zeit. Electrochem., 1901, 7, p. 622). Cryoscopic determinations of its molecular weight show that it is H2O2. [G. Carrara, Rend. della Accad. dei Lincei, 1892 (5), 1, ii. p. 19; W. R. Orndorff and J. White, Amer. Chem. Journ., 1893, 15, p. 347.] Hydrogen peroxide behaves very frequently as a powerful oxidizing agent; thus lead sulphide is converted into lead sulphate in presence of a dilute aqueous solution of the peroxide, the hydroxides of the alkaline earth metals are converted into peroxides of the type MO2·8H2O, titanium dioxide is converted into the trioxide, iodine is liberated from potassium iodide, and nitrites (in alkaline solution) are converted into acid-amides (B. Radziszewski, Ber., 1884, 17, p. 355). In many cases it is found that hydrogen peroxide will only act as an oxidant when in the presence of a catalyst; for example, formic, glycollic, lactic, tartaric, malic, benzoic and other organic acids are readily oxidized in the presence of ferrous sulphate (H. J. H. Fenton, Jour. Chem. Soc., 1900, 77, p. 69), and sugars are readily oxidized in the presence of ferric chloride (O. Fischer and M. Busch, Ber., 1891, 24, p. 1871). It is sought to explain these oxidation processes by assuming that the hydrogen peroxide unites with the compound undergoing oxidation to form an addition compound, which subsequently decomposes (J. H. Kastle and A. S. Loevenhart, Amer. Chem. Journ., 1903, 29, pp. 397, 517). Hydrogen peroxide can also react as a reducing agent, thus silver oxide is reduced with a rapid evolution of oxygen. The course of this reaction can scarcely be considered as definitely settled; M. Berthelot considers that a higher oxide of silver is formed, whilst A. Baeyer and V. Villiger are of opinion that reduced silver is obtained [see Comptes rendus, 1901, 133, p. 555; Ann. Chim. Phys., 1897 (7), 11, p. 217, and Ber., 1901, 34, p. 2769]. Potassium permanganate, in the presence of dilute sulphuric acid, is rapidly reduced by hydrogen peroxide, oxygen being given off, 2KMnO4 \+ 3H2SO4 \+ 5H2O2 = K2SO4 \+ 2MnSO4 \+ 8H2O + 5O2. Lead peroxide is reduced to the monoxide. Hypochlorous acid and its salts, together with the corresponding bromine and iodine compounds, liberate oxygen violently from hydrogen peroxide, giving hydrochloric, hydrobromic and hydriodic acids (S. Tanatar, Ber., 1899, 32, p. 1013).

On the constitution of hydrogen peroxide see C. F. Schönbein, Jour. prak. Chem., 1858-1868; M. Traube, Ber., 1882-1889; J. W. Brühl, Ber., 1895, 28, p. 2847; 1900, 33, p. 1709; S. Tanatar, Ber., 1903, 36, p. 1893.

Hydrogen peroxide finds application as a bleaching agent, as an antiseptic, for the removal of the last traces of chlorine and sulphur dioxide employed in bleaching, and for various quantitative separations in analytical chemistry (P. Jannasch, Ber., 1893, 26, p. 2908). It may be estimated by titration with potassium permanganate in acid solution; with potassium ferricyanide in alkaline solution, 2K3Fe(CN)6 \+ 2KOH + H2O2 = 2K4Fe(CN)6 \+ 2H2O + O2; or by oxidizing arsenious acid in alkaline solution with the peroxide and back titration of the excess of arsenious acid with standard iodine (B. Grützner, Arch. der Pharm., 1899, 237, p. 705). It may be recognized by the violet coloration it gives when added to a very dilute solution of potassium bichromate in the presence of hydrochloric acid; by the orange-red colour it gives with a solution of titanium dioxide in concentrated sulphuric acid; and by the precipitate of Prussian blue formed when it is added to a solution containing ferric chloride and potassium ferricyanide.

Ozonic Acid, H2O4. By the action of ozone on a 40% solution of potassium hydroxide, placed in a freezing mixture, an orange-brown substance is obtained, probably K2O4, which A. Baeyer and V. Villiger (Ber., 1902, 35, p. 3038) think is derived from ozonic acid, produced according to the reaction O3 \+ H2O = H2O4.

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