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.
HYDROGRAPHY (Gr. ὕδωρ, water, and γράφειν, to write),
the science dealing with all the waters of the earth’s surface,
including the description of their physical features and conditions;
the preparation of charts and maps showing the position
of lakes, rivers, seas and oceans, the contour of the sea-bottom,
the position of shallows, deeps, reefs and the direction and
volume of currents; a scientific description of the position,
volume, configuration, motion and condition of all the waters
of the earth. See also Surveying (Nautical) and Ocean and Oceanography.
The Hydrographic Department of the British
Admiralty, established in 1795, undertakes the making of charts
for the admiralty, and is under the charge of the hydrographer to
the admiralty (see Chart).
HYDROLYSIS (Gr. ὕδωρ, water, λύειν, to loosen), in chemistry,
a decomposition brought about by water after the manner shown
in the equation R·X + H·OH = R·H + X·OH. Modern research
has proved that such reactions are not occasioned by water
acting as H2O, but really by its ions (hydrions and hydroxidions),
for the velocity is proportional (in accordance with the law of
chemical mass action) to the concentration of these ions. This
fact explains the so-called “catalytic” action of acids and bases
in decomposing such compounds as the esters. The term
“saponification” (Lat. sapo, soap) has the same meaning, but
it is more properly restricted to the hydrolysis of the fats, i.e.
glyceryl esters of organic acids, into glycerin and a soap (see
Chemical Action).
