[49] Nitrogen peroxide as a mixed substance has no corresponding independent salts, but Sabatier and Senderens (1892) showed that under certain conditions NO2 combines directly with some metals—for instance, copper and cobalt—forming Cu2NO2 and CoNO2 as dark brown powders, which do not, however, exhibit the reactions of salts. Thus by passing gaseous nitrogen dioxide over freshly reduced (from the oxides by heating with hydrogen) copper at 25°–30°, Cu2NO2 is directly formed. With water it partly gives off NO2 and partly forms nitrite of copper, leaving metallic copper and its suboxide. The nature of these compounds has not yet been sufficiently investigated.
[50] Ammonium nitrite may be easily obtained in solution by a similar method of double decomposition (for instance, of the barium salt with ammonium sulphate) to the other salts of nitrous acid, but it decomposes with great ease when evaporated, with evolution of gaseous nitrogen, as already mentioned (Chapter [V].) If the solution, however, be evaporated at the ordinary temperature under the receiver of an air-pump, a solid saline mass is obtained, which is easily decomposed when heated. The dry salt even decomposes with an explosion when struck, or when heated to about 70°—NH4NO2 = 2H2O + N2. It is also formed by the action of aqueous ammonia on a mixture of nitric oxide and oxygen, or by the action of ozone on ammonia, and in many other instances. Zörensen (1894) prepared NH4NO2 by the action of a mixture of N2O3 and other oxides of nitrogen on lumps of ammonium carbonate, extracting the nitrite of ammonium formed with absolute alcohol, and precipitating it from this solution by ether. This salt is crystalline, dissolves in water with absorption of heat, and attracts moisture from the air. The solid salt and its concentrated solutions decompose with an explosion when heated to 50°–80°, especially in the presence of traces of foreign acids. Decomposition also proceeds at the ordinary temperature, but more slowly; and in order to preserve the salt it should be covered with a layer of pure dry ether.
[51] Silver nitrite, AgNO2, is obtained as a very slightly soluble substance, as a precipitate, on mixing solutions of silver nitrate, AgNO3, and potassium nitrite, KNO2. It is soluble in a large volume of water, and this is taken advantage of to free it from silver oxide, which is also present in the precipitate, owing to the fact that potassium nitrite always contains a certain amount of oxide, which with water gives the hydroxide, forming oxide of silver with silver nitrate. The solution of silver nitrite gives, by double decomposition with metallic chlorides (for instance, barium chloride), insoluble silver chloride and the nitrite of the metal taken (in this case, barium nitrite, Ba(NO2)2).
[51 bis] Leroy (1889) obtained KNO2 by mixing powdered KNO3 with BaS, igniting the mixture in a crucible and washing the fused salts; BaSO4 is then left as an insoluble residue, and KNO2 passes into solution: 4KNO3 + BaS = 4KNO2 + BaSO4.
[52] Probably potassium nitrite, KNO2, when strongly heated, especially with metallic oxides, evolves N and O, and gives potassium oxide, K2O, because nitre is liable to such a decomposition; but it has, as yet, been but little investigated.
[53] There are many researches which lead to the conclusion that the reaction N2O3 = NO2 - NO is reversible, i.e. resembles the conversion of N2O4 into NO2. The brown colour of the fumes of N2O3 is due to the formation of NO2.
If nitrogen peroxide be cooled to -20°, and half its weight of water be added to it drop by drop, then the peroxide is decomposed, as we have already said, into nitrous and nitric acids; the former does not then remain as a hydrate, but straightway passes into the anhydride, and, hence, if the resultant liquid be slightly warmed vapours of nitrous anhydride, N2O3, are evolved, and condense into a blue liquid, as Fritzsche showed. This method of preparing nitrous anhydride apparently gives the purest product, but it easily dissociates, forming NO and NO2 (and therefore also nitric acid in the presence of water).
[54] According to Thorpe, N2O3 boils at +18°. According to Geuther, at +3°·5, and its sp. gr. at 0° = 1·449.
[55] In its oxidising action nitrous anhydride gives nitric oxide, N2O3 = 2NO + O. Thus its analogy to ozone becomes still more marked, because in ozone it is only one-third of the oxygen that acts in oxidising; from O3 there is obtained O, which acts as an oxidiser, and common oxygen O2. In a physical aspect the relation between N2O3 and O3 is revealed in the fact that both substances are of a blue colour when in the liquid state.
[56] This reaction is taken advantage of for converting the amides, NH2R (where R is an element or a complex group) into hydroxides, RHO. In this case NH2R + NHO2 forms 2N + H2O + RHO; NH2, is replaced by HO, the radicle of ammonia by the radicle of water. This reaction is employed for transforming many nitrogenous organic substances having the properties of amides into their corresponding hydroxides. Thus aniline, C6H5·NH2, which is obtained from nitrobenzene, C6H5·NO2 (Note [37]), is converted by nitrous anhydride into phenol, C6H5·OH, which occurs in the creosote extracted from coal tar. Thus the H of the benzene is successively replaced by NO2, NH2, and HO; a method which is suitable for other cases also.