p. 454; 1886, 27, p. 606). M. Berthelot and J. Ogier (Bull. Soc.
Chim., 1882 [2], 37, p. 434; 38, p. 60) have also shown that the
specific heat of the gas decreases with increase of temperature until
it reaches a minimum at about 198-253° C. Cryoscopic determinations
of the molecular weight of nitrogen peroxide dissolved in
glacial acetic acid show that it corresponds to the molecular formula
N2O4 at low temperatures (W. Ramsay, Jour. Chem. Soc., 1888, 53,
p. 621). Nitrogen peroxide is the most stable oxide of nitrogen.
It is decomposed by water, giving at 0° C. a mixture of nitric and
nitrous acids: 2NO2+H2O＝HNO3+HNO2. It combines with
sulphuric acid to form nitro-sulphonic acid, SO2(OH)(NO2). It does
not support the combustion of a taper, but burning phosphorus and
red-hot carbon will continue to burn in the gas. It converts many
metallic oxides into mixtures of nitrates and nitrites, and attacks
many metals, forming nitrates and being itself reduced to nitric
oxide. It is an energetic oxidizing agent.
Nitrogen pentoxide, N2O5, was first obtained in 1849 by H. Sainte-Claire-Deville (Ann. Chim. Phys., 1850 [3], 28, p. 241) by the action of dry chlorine on silver nitrate: 4AgNO3+2Cl2＝4AgCl+2N2O5 +O2. It may also be obtained by distilling nitric acid over phosphorus pentoxide. It crystallizes in large prisms which melt at 29-30° C. to a yellowish liquid, which boils at 45-50° C. with rapid decomposition. It is very unstable, decomposing slowly even at ordinary temperatures. It dissolves in water, forming nitric acid.
Hyponitrous acid, H2N2O2, was first obtained in the form of its salts by E. Divers in 1871 (Chem. News, 23, p. 206) by reducing a solution of potassium nitrite with sodium amalgam, and subsequent precipitation as silver salt. Hyponitrites also result when hydroxyamido-sulphonates, e.g. HO⋅NH⋅SO3Na, are hydrolysed by caustic alkalis (E. Divers and T. Haga, Jour. Chem. Soc., 1889, 55, p. 760), or when benzsulphohydroxamic acid, C6H5SO2⋅NH⋅OH, is treated in the same manner (O. Piloty, Ber., 1896, 29, p. 1560). They may also be prepared by the action of mercuric or cupric oxides on alkaline solutions of hydroxylamine (A. Hantzsch, Ann., 1896, 292, p. 317); by the action of hydroxylamine sulphate on alkaline nitrites in the presence of lime or calcium carbonate, the mixture being rapidly heated to 60° C.; or by the hydrolysis of dimethyl nitroso-oxyurea, (CH3)2N⋅CO⋅N(NO)⋅OH (A. Hantzsch, Ber., 1897, 30, p. 2356). The free acid, which crystallizes in brilliant scales, is best prepared by decomposing the silver salt with an ethereal solution of hydrochloric acid. It is very explosive, dissolves readily in water and behaves as a dibasic acid. It does not liberate iodine from potassium iodide, neither does it decolorize iodine solution. Bromine oxidizes it to nitric acid, but the reaction is not quantitative. In acid solution, potassium permanganate oxidizes it to nitric acid, but in alkaline solution only to nitrous acid. It decomposes slowly on standing, yielding water and nitrous oxide. The silver salt is a bright yellow solid, soluble in dilute sulphuric and nitric acids, and may be crystallized from concentrated solutions of ammonia. It slowly decomposes on exposure or on heating. The calcium salt, CaN2O2⋅4H2O, formed by the action of calcium chloride on the silver salt in the presence of a small quantity of nitric acid, is a lustrous crystalline powder, almost insoluble in water but readily soluble in dilute acids.” It is decomposed by sulphuric acid, with evolution of nitrous oxide.
Nitrous acid, HNO2, is found to some extent in the form of its salts in the atmosphere and in rain water. The pure acid has not yet been obtained, since in the presence of water it decomposes with formation of nitric acid and liberation of nitric oxide: 3HNO2＝HNO3+2NO+H2O. Its salts may be obtained in some cases by heating the corresponding nitrates, but the method does not give good results. Sodium nitrite, the most commonly used salt of the acid, is generally obtained by heating the nitrate with metallic lead; by heating sodium nitrate with sulphur and sodium hydroxide, the product then being fractionally crystallized (Read, Holliday & Sons): 3NaNO3+S+2NaOH＝Na2SO4+3NaNO2+H2O; by oxidizing atmospheric nitrogen in an electric arc, keeping the gases above 300° C., until absorption in alkaline hydroxide solution is effected (German Pat. 188188); or by passing air, or a mixture of oxygen and ammonia, over heated metallic oxides (ibid., 168272). The salts of the acid are colourless or faintly yellow. In aqueous solution the free acid acts as an oxidizing agent, bleaching indigo and liberating iodine from potassium iodide, or it may act as a reducing agent since it readily tends to pass into nitric acid: consequently it discharges the colour of acid solutions of permanganate's and chromates. The acid finds considerable use in organic chemistry, being employed to discriminate between the different types of alcohols and of amines, and also in the production of diazo, azo and diazo-amino compounds. It may be recognized by the blue colour it gives with diphenylamine sulphate and by its reaction with potassium iodide-starch paper.
Nitrosyl chloride, NOCl, is obtained by the direct union of nitric oxide with chlorine; or by distilling a mixture of concentrated nitric and hydrochloric acids, passing the resulting gases into concentrated sulphuric acid and heating the so-formed nitrosyl hydrogen sulphate with dry salt: HNO3+3HCl＝NOCl+C12 +H2O; NOCl + H2SO4＝HCl + NO⋅SO4H; NO⋅SO4H + NaCl＝NOCl+NaHSO4 (W. A. Tilden, Jour. Chem. Soc., 1860, p. 630). It is also prepared by the action of phosphorus pentachloride on potassium nitrite or on nitrogen peroxide. It is an orange-coloured gas which may be readily liquefied and by further cooling may be solidified. The liquid boils at −5° C. and the solid melts at −65° C. It forms double compounds with many metallic chlorides, and finds considerable application as a means of separating various members of the terpene group of compounds. It is readily decomposed by water and alkaline hydroxides, yielding a mixture of nitrite and chloride. On treatment with silver fluoride it yields nitrosyl fluoride, NOF (O. Ruff, Zeit. anorg. Chem., 1905, 47, p. 190). Nitroxyl fluoride, NO2F, is formed by the action of fluorine on nitric oxide at the temperature of liquid oxygen (H. Moissan and P. Lebeau, Comptes rendus, 1905, 140, pp. 1573, 1621). It is a gas at ordinary temperature; when liquefied it boils at −63·5° C. and on solidification melts at −139° C. Water decomposes it into nitric and hydrofluoric acids. Nitramide, NH2NO2, is obtained by the action of sulphuric and nitric acids on potassium imidosulphonate, or by the action of ice-cold sulphuric acid on potassium nitro-carbamate (J. Thiele and A. Lachmann, Ann., 1895, 288, p. 297): NO2⋅NK⋅CO2K+H2SO4＝NH2NO2+K2SO4+CO2. It crystallizes in prisms or leaflets which melt at 72-75° C. and are readily soluble in water and in all organic solvents except ligroin. It is somewhat volatile at ordinary temperature, and its aqueous solution possesses a strongly acid reaction. It is very unstable, decomposing into nitrous oxide and water when mixed with copper oxide, lead chromate or even powdered glass. On reduction it gives a strongly reducing substance, probably hydrazine. According to A. Hantzsch (Ann., 1896, 292, pp. 340 et seq.) hyponitrous acid and nitramide are to be regarded as stereo-isomers, being the anti- and syn- forms of the same compound. Thiele, however, regards nitramide as imidonitric acid, HN:NO(OH).
Nitrogen sulphide, N4S4, first obtained by W. Gregory (Jour. pharm., 1835, 21, p. 315) by the action of ammonia on sulphur chloride, has been investigated by O. Ruff and E. Geisel (Ber., 1904, 37, 1573; 1905, 38, p. 2659). who also obtained it by dissolving sulphur in liquid ammonia. It is a reddish-yellow crystalline solid, insoluble in water and melting at 178° C. It explodes readily when melted or subjected to shock. Dry hydrochloric acid gives ammonia but no nitrogen; with ammonia it gives N:SNH2 and S:S(NH2)2; and with secondary amines it forms thiodiamines, S(NR2)2, nitrogen and ammonia being liberated. When heated with CS2 to 100° C. under pressure, it forms liquid nitrogen sulphide, N2S5, a mobile red liquid which solidifies to an iodine-like mass of crystals which melt at 10-11° C. Water, alkalis and acids decompose it into sulphur and ammonia (W. Muthmann, Zeit. anorg. Chem., 1897, 13, p. 200).
For sulphonic acids containing nitrogen see Ammonia.
Numerous determinations of the atomic weight of nitrogen have been made by different observers, the values obtained varying somewhat according to the methods used. These methods have been purely chemical (either gravimetric or volumetric), physical (determinations of the density of nitrogen, nitric oxide, &c.), or physicochemical. P. A. Guye has given a critical discussion of the relative accuracy of the gravimetric and physico-chemical methods, and favours the latter, giving for the atomic weight a value less than 14·01. The more important papers dealing with the subject are: J. Stas, (Œuvres complètes, i. pp. 342 et seq.; Lord Rayleigh, Proc. Roy. Soc. (1894), 55, p. 340; (1904) 73, p. 153; G. Dean, Jour. Chem. Soc. (1901), 79, p. 147; R. W. Gray, Jour. Chem. Soc. (1906), 88, p. 1174; A. Scott, Proc. Chem. Soc. (1905), 21, p. 309; P. A. Guye, Chem. News (1905), 92, pp. 261 et seq.; (1906) 93, p. 13 et seq.; D. Berthelot, Comptes rendus (1907), 144, p. 269.
NITROGLYCERIN, C3H5(NO3)3 or CH2NO3⋅CH2NO3 glyceryl trinitrate, an explosive first obtained in 1846 by Ascanio Sobrero (Mem. Acad. Torino, 1847) by acting with a mixture of strong nitric and sulphuric acids on glycerin at the ordinary temperature. The reaction proceeds in several stages, mono-, di- and finally tri-nitrate being produced, the final stage requiring sulphuric acid as a dehydrator. When pure it is a very pale yellow oil of sp. gr. 1·614 at 4° C. and 1·60 at 15° C. One gram requires for solution between 800 and 1000 c.c. of water, 4 c.c. of absolute alcohol or 18 c.c. of wood spirit, and it is scarcely at all soluble in glycerin itself, but mixes in all proportions with ether, acetone, ethyl acetate and benzene.
In the manufacture glycerin is dropped in a very thin stream into a mixture of 3 parts of nitric (sp. gr. 1·5) and 5 parts of sulphuric acid (sp. gr. 1·84), the containing vessel being cooled by a water jacket and the acid mixture agitated by a stream of cooled air, the temperature being kept at about 15° C. A considerable excess of acids is necessary for the completion and safety of the reaction, usually about 8 parts of the acid mixture to 1 of glycerin. The higher the strength of the acids the higher the yield of nitroglycerin and the smaller the loss by solution in the waste acids. In recent practice some sulphin trioxide, or fuming sulphuric acid, is added, so that the mixture of acids contains less than 1% of water. The action is very rapid, and the product, which rises to the top of the acids, is separated and washed successively with cold and then tepid water, and finally with water made slightly alkaline with sodium carbonate or hydroxide, to remove all adhering or dissolved acids which would otherwise render the product very unstable. Nitroglycerin dissolves a little water and then appears thick or milky. Generally it is either
