concentric copper electrodes and rotated by an electromagnet;
it is worked at Vevey, Switzerland. The Rankin process, of
which very little is known, produces the arc with much lower
current.
The conversion of nitrogen into ammonia by electricity has received much attention, but the commercial aspect appears to have been first worked out by de Hemptinne in 1900, who used both the spark and silent discharge on mixtures of hydrogen and nitrogen, and found that the pressure and temperature must be kept low and the spark gap narrow. J. Schlutius in 1903 employed Dowson gas as a source of hydrogen, and induced combination by means of platinum and the silent discharge. Several non-electrical processes have been devised. In 1862 Fleck passed a mixture of steam, nitrogen and carbon monoxide over red-hot lime, whilst in 1904 Woltereck induced combination by passing steam and air over red-hot iron oxide (peat is used in practice). In de Lambilly's process air and steam is led over white-hot coke. and carbon dioxide or monoxide removed from the escaping gases according as ammonium formate or carbonate is wanted. The residual gas is then passed through a tube containing porous materials, such as wood- or bone-charcoal, platinized pumice or spongy platinum, then mixed with steam and again forced through the tube. The reactions are represented as
- (1) N2+3H2+2CO+2H2O＝2H⋅CO2NH4 (Ammonium formate).
- (2) N2+3H2+2CO2+2H2O＝2HO⋅CO2NH4 (Ammonium carbonate).
The best temperature for the first reaction is between 80° C. and 130° C. and for the second between 40° C. and 60° C. In another process, which originated with C. Kaiser (Abst. J.C.S., 1907, ii. p. 862), calcium is heated in a current of hydrogen, and nitrogen passed over the hydride so formed; this gives ammonia and calcium nitride, the latter of which gives up its nitrogen as ammonia and reforms the hydride when heated in a current of hydrogen.
The fixation of nitrogen as a nitride has not been attended with commercial success. H. Mehner patented heating the oxides of silicon, boron or magnesium with coal or coke in an electric furnace, and then passing in nitrogen, which forms, with the metal liberated by the action of the carbon, a readily decomposable nitride.
For an extended bibliography see Bulletin No. 63 of the Bureau of Soils, U.S. Department of Agriculture (Washington, 1910).
Nitrogen is a colourless, tasteless and odourless gas, which is only very slightly soluble in water. It is slightly lighter than air. Lord Rayleigh in 1894 found that the density of atmospheric nitrogen was about 12% higher than that of chemically prepared nitrogen, a discovery which led to the isolation of the rare gases of the atmosphere (see Argon). The values obtained are shown below.,
| Atmospheric Nitrogen. |
Chemical Nitrogen. | |
| 0·97209 | 0·96727 | Lord Rayleigh, Chem. News, 1897, 76, p. 315. |
| 0·9720 | 0·9671 | A. Leduc, Comptes rendus, 1896, 123, p. 805. |
Nitrogen is a very inert gas: it will neither burn nor support the combustion of ordinary combustibles. It combines directly with lithium, calcium and magnesium when heated, whilst nitrides of the rare earth metals are also produced when their oxides are mixed with magnesium and heated in a current of nitrogen (C. Matignon, Comptes rendus, 1900, 131, p. 837). Nitrogen has been liquefied, the critical temperature being −149° C. and the critical pressure 27·54 atmospheres. The liquefied gas boils at −195·5° C., and its specific gravity at its boiling point is 0·8103 (E. C. C. Baly and F. G. Donnan, Jour. Chem. Soc., 1902, 81, p. 912).
Compounds.
Nitrogen combines with hydrogen to form ammonia, NH3, hydrazine, N2H4, and azoimide, N3H (qq.v.); the other known hydrides, N4H4 and N5H5, are salts of azoimide, viz. NH4⋅N3 and N2H4N3H.
Nitrogen trichloride, NCl3, discovered by P. L. Dulong in 1811 (Schweigg. Journ., 1811, 8, p. 302), and obtained by the action of chlorine or sodium hypochlorite on ammonium chloride, or by the electrolysis of ammonium chloride solution, is a very volatile yellow oil. It possesses an extremely pungent smell, and its vapour is extremely irritating to the eyes. It is a most dangerous explosive (see D. L. Chapman and L. Vodden, Jour. Chem. Soc., 1909, 95, p. 138). Chlorine azide, Cl⋅N3, was discovered by F. Raschig in 1908 (see Azoamide); the corresponding iodine compound had been obtained in 1900 by A. Hantzsch (Ber., 33, p. 522). For the so-called nitrogen iodide see Ammonia.
Nitrogen forms five oxides, viz. nitrous oxide, NH2O, nitric oxide, NO, nitrogen trioxide, N2O3, nitrogen peroxide, NO2, and nitrogen pentoxide, N2O2, whilst three oxyacids of nitrogen are known: hyponitrous acid, H2N2O2, nitrous acid, HNO2, and nitric acid. HNO3 (q.v.). The first four oxides are gases, the fifth is a solid. Nitrous oxide, N2O, isolated in 1776 by J. Priestley, who obtained it by reducing nitrogen peroxide with iron, may be prepared by heating ammonium nitrate at 170-260° C., or by reducing a mixture of nitric and sulphuric acid with zinc. It is a colourless gas, which is practically odourless, but possesses a sweetish taste. It is somewhat soluble in water. When liquefied it boils at −89·8° C., and by further cooling may be solidified, the solid melting at −102·3° (W. Ramsay, Chem. News, 1893, 67, p. 140). It does not burn, but Supports the combustion of heated substances almost as well as oxygen. It is used as an anaesthetic, principally in dentistry, producing when inhaled a condition of hysterical excitement often accompanied by loud laughter, whence it is sometimes called “laughing gas.”
Nitric oxide, NO, first obtained by Van Helmont, is usually prepared by the action of dilute nitric acid (sp. gr. 1·2) on copper. This method does not give a pure gas, varying amounts of nitrous oxide and nitrogen being present (see Nitric Acid). In a purer condition it may be obtained by the action of sulphuric acid on a mixture of potassium nitrate and ferrous sulphate, or of hydrochloric acid on a mixture of potassium nitrate and ferric chloride. It is also formed by the action of concentrated sulphuric acid on sodium nitrite in the presence of mercury. It is a colourless gas which is only sparingly soluble in water. It may be liquefied, its critical temperature being −93·5°, and the liquid boils at −153·6° C. It is not a supporter of combustion, unless the substance introduced is at a sufficiently high temperature to decompose the gas, when combustion will continue at the expense of the liberated oxygen. If the gas be mixed with the vapour of carbon disulphide, the mixture burns with a vivid lavender-coloured flame. Nitric oxide is soluble in solutions of ferrous salts, a dark brown solution being formed, which is readily decomposed by heat, with evolution of nitric oxide. It combines with oxygen to form nitrogen peroxide. Nascent hydrogen reduces it to hydroxylamine (q.v.), whilst solutions of hypochlorites oxidize it to nitric acid. In some instances it reacts as a reducing agent, e.g. silver oxide is reduced to metallic silver at 170° C., lead dioxide to the monoxide and manganese dioxide to sesquioxide.
Nitrogen trioxide, N2O3, was first mentioned by J. R. Glauber in 1648 as a product of the reaction between nitric acid and arsenious oxide. Sir W. Ramsay (Jour. Chem. Soc., 1890, 5, p. 590), by distilling arsenious oxide with nitric acid and cooling the distillate, obtained a green liquid which consisted of nitrogen trioxide and peroxide in varying proportions, and concluded that the trioxide could not be obtained pure. He then tried the direct combination of nitric oxide with liquid nitrogen peroxide. A dark blue liquid is produced, and the first portions of gas boiling off from the mixture correspond fairly closely in composition with nitrogen trioxide. H. B. Baker (Jour. Chem. Soc., 1907, 91, p. 1862) obtained nitrogen trioxide in the gaseous form by volatilizing the liquid under special conditions. L. Francesconi and N. Sciacca (Gazz., 1904, 34 (i.), p. 447) have shown that liquid nitric oxide and oxygen, or gaseous nitric oxide and liquid oxygen, mixed in all proportions and yielded nitrogen trioxide, whilst gaseous nitric oxide mixed with excess of oxygen always gave the trioxide if the mixture was kept below −110° C. They also state that nitrogen trioxide is stable at ordinary pressure up to −21° C. N. M. v. Wittorf (Zeit. anorg. Chem., 1904, 41, p. 85) obtained blue crystals of the trioxide (melting at −103° C.) on saturating liquid nitrogen peroxide with nitric oxide and cooling the mixture. The liquid prepared by Baker is green in colour, and has a specific gravity 1·11 at ordinary temperature, but below −2° C. becomes of a deep indigo blue colour. It forms a mass of deep blue crystals at the temperature of liquid air. It is exceedingly soluble in concentrated sulphuric acid.
Nitrogen peroxide, NO2 or N2O4, may be obtained by mixing oxygen with nitric oxide and passing the red gas so obtained through a freezing mixture. The production of this red gas when air is mixed with nitric oxide was mentioned by R. Boyle in 1671. Nitrogen peroxide is also prepared by heating lead nitrate and passing the products of decomposition through a tube surrounded by a freezing mixture, when the gas liquefies. At low temperatures it is a colourless crystalline solid which melts at −10·14° C. (W. Ramsay, Chem. News, 1900, 61, p. 91). As the temperature increases the liquid becomes yellowish, the colour deepening with rise of temperature until at +15° C. it has a deep orange tint. The liquid boils at about 22° C. This change of colour is accompanied by a change in the vapour density, and is explained by the fact that nitrogen peroxide consists of a mixture of a colourless compound N2O4, and a red-brown gas NO2, the latter increasing in amount at the expense of the former as the temperature is raised (G. Salet, Comptes rendus, 1868, 67, p. 488; see also E. and L. Natanson, Wied. Ann., 1885, 24,
