Fig. 48.—The method of decomposition of nitrous anhydride, also applicable to the other oxides of nitrogen, and to their analysis. NO2 is generated from nitrate of lead in the retort A. Nitric acid and other less volatile products are condensed in B. The tube C C contains copper, and is heated from below. Undecomposed volatile products (if any are formed) are condensed in D, which is cooled. If the decomposition be incomplete, brown fumes make their appearance in this receiver. The gaseous nitrogen is collected in the cylinder E.
If the vapour of nitric acid is passed through an even moderately heated glass tube, the formation of dark-brown fumes of the lower oxides of nitrogen and the separation of free oxygen may be observed, 2NHO3 = H2O + 2NO2 + O. The decomposition is complete at a white heat—that is, nitrogen is formed, 2NHO3 = H2O + N2 + O5. Hence it is easily understood that nitric acid may part with its oxygen to a number of substances capable of being oxidised.[39] It is consequently an oxidising agent. Charcoal, as we have already seen, burns in nitric acid; phosphorus, sulphur, iodine, and the majority of metals also decompose nitric acid, some on heating and others even at the ordinary temperature: the substances taken are oxidised and the nitric acid is deoxidised, yielding compounds containing less oxygen. Only a few metals, such as gold and platinum, do not act on nitric acid, but the majority decompose it; in so doing, an oxide of the metal is formed, which, if it has the character of a base, acts on the remaining nitric acid; hence, with the majority of metals the result of the reaction is usually not an oxide of the metal, but the corresponding salt of nitric acid, and, at the same time, one of the lower oxides of nitrogen. The resulting salts of the metals are soluble, and hence it is said that nitric acid dissolves nearly all metals.[40] This case is termed the solution of metals by acids, although it is not a case of simple solution, but a complex chemical change of the substances taken. When treated with this acid, those metals whose oxides do not combine with nitric acid yield the oxide itself, and not a salt; for example, tin acts in this manner on nitric acid, forming a hydrated oxide, SnH2O3, which is obtained in the form of a white powder, Sn + 4NHO3 = H2SnO3 + 4NO2 + H2O. Silver is able to take up still more oxygen, and to convert a large portion of nitric acid into nitrous anhydride, 4Ag + 6HNO3 = 4AgNO3 + N2O3 + 3H2O. Copper takes up still more oxygen from nitric acid, converting it into nitric oxide, and, by the action of zinc, nitric acid is able to give up a still further quantity of nitrogen, forming nitrous oxide, 4Zn + 10NHO3 = 4Zn(NO3)2 + N2O + 5H2O.[41] Sometimes, and especially with dilute solutions of nitric acid, the deoxidation proceeds as far as the formation of hydroxylamine and ammonia, and sometimes it leads to the formation of nitrogen itself. The formation of one or other nitrogenous substance from nitric acid is determined, not only by the nature of the reacting substances, but also by the relative mass of water and nitric acid, and also by the temperature and pressure, or the sum total of the conditions of reaction; and as in a given mixture even these conditions vary (the temperature and the relative mass vary), it not unfrequently happens that a mixture of different products of the deoxidation of nitric acid is formed.
Thus the action of nitric acid on metals consists in their being oxidised, whilst the acid itself is converted, according to the temperature, concentration in which it is taken, and the nature of the metal, &c., into lower oxides, ammonia, or even into nitrogen.[42] Many compounds are oxidised by nitric acid like metals and other elements; for instance, lower oxides are converted into higher oxides. Thus, arsenious acid is converted into arsenic acid, suboxide of iron into oxide, sulphurous acid into sulphuric acid, the sulphides of the metals, M2S, into sulphates, M2SO4, &c.; in a word, nitric acid brings about oxidation, its oxygen is taken up and transferred to many other substances. Certain substances are oxidised by strong nitric acid so rapidly and with so great an evolution of heat that they deflagrate and burst into flame. Thus turpentine, C10H16, bursts into flame when poured into fuming nitric acid. In virtue of its oxidising property, nitric acid removes the hydrogen from many substances. Thus it decomposes hydriodic acid, separating the iodine and forming water; and if fuming nitric acid be poured into a flask containing gaseous hydriodic acid, then a rapid reaction takes place, accompanied by flame and the separation of violet vapours of iodine and brown fumes of oxides of nitrogen.[43]
As nitric acid is very easily decomposed with the separation of oxygen, it was for a long time supposed that it was not capable of forming the corresponding nitric anhydride, N2O5; but Deville first and subsequently Weber and others, discovered the methods of its formation. Deville obtained nitric anhydride by decomposing silver nitrate by chlorine under the influence of a moderate heat. Chlorine acts on the above salt at a temperature of 95° (2AgNO3 + Cl2 = 2AgCl + N2O5 + O), and when once the reaction is started, it continues by itself without further heating. Brown fumes are given off, which are condensed in a tube surrounded by a freezing-mixture. A portion condenses in this tube and a portion remains in a gaseous state. The latter contains free oxygen. A crystalline mass and a liquid substance are obtained in the tube; the liquid is poured off, and a current of dry carbonic acid gas is passed through the apparatus in order to remove all traces of volatile substances (liquid oxides of nitrogen) adhering to the crystals of nitric anhydride. These form a voluminous mass of rhombic crystals (density 1·64), which sometimes are of rather large size; they melt at about 30° and distil at about 47°. In distilling, a portion of the substance is decomposed. With water these crystals give nitric acid. Nitric anhydride is also obtained by the action of phosphoric anhydride, P2O5, on cold pure nitric acid (below 0°). During the very careful distillation of equal parts by weight of these two substances a portion of the acid decomposes, giving a liquid compound, H2O,2N2O5 = N2O5,2HNO3, whilst the greater part of the nitric acid gives the anhydride according to the equation 2NHO3 + P2O5 = 2PHO3 + N2O5. On heating, nitric anhydride decomposes with an explosion, or gradually, into nitric peroxide and oxygen, N2O5 = N2O4 + O.
Nitrogen peroxide, N2O4, and nitrogen dioxide, NO2, express one and the same composition, but they should be distinguished like ordinary oxygen and ozone, although in this case their mutual conversion is more easily effected and takes place on vaporisation; also, O3 loses heat in passing into O2, whilst N2O4 absorbs heat in forming NO2.
Nitric acid in acting on tin and on many organic substances (for example, starch) gives brown vapours, consisting of a mixture of N2O3 and NO2. A purer product is obtained by the decomposition of lead nitrate by heat, Pb(NO3)2 = 2NO2 + O + PbO, when non-volatile lead oxide, oxygen gas, and nitrogen peroxide are formed. The latter condenses, in a well-cooled vessel, to a brown liquid, which boils at about 22°. The purest peroxide of nitrogen, solidifying at -9°, is obtained by mixing dry oxygen in a freezing-mixture with twice its volume of dry nitric oxide, NO, when transparent prisms of nitrogen peroxide are formed in the receiver: they melt into a colourless liquid at about -10°. When the temperature of the receiver is above -9°, the crystals melt,[44] and at 0° give a reddish yellow liquid, like that obtained in the decomposition of lead nitrate. The vapours of nitrogen peroxide have a characteristic odour, and at the ordinary temperature are of a dark-brown colour, but at lower temperatures the colour of the vapour is much fainter. When heated, especially above 50°, the colour becomes a very dark brown, so that the vapours almost lose their transparency.
The causes of these peculiarities of nitrogen peroxide were not clearly understood until Deville and Troost determined the density and dissociation of the vapour of this substance at different temperatures, and showed that the density varies. If the density be referred to that of hydrogen at the same temperature and pressure, then it is found to vary from 38 at the boiling point, or about 27°, to 23 at 135°, after which the density remains constant up to those high temperatures at which the oxides of nitrogen are decomposed. As on the basis of the laws enunciated in the [following chapter], the density 23 corresponds with the compound NO2 (because the weight corresponding with this molecular formula = 46, and the density referred to hydrogen as unity is equal to half the molecular weight); therefore at temperatures above 135° the existence of nitrogen dioxide only must be recognised. It is this gas which is of a brown colour. At a lower temperature it forms nitrogen peroxide, N2O4, whose molecular weight, and therefore density, is twice that of the dioxide. This substance, which is isomeric with nitrogen dioxide, as ozone is isomeric with oxygen, and has twice as great a vapour density (46 referred to hydrogen), is formed in greater quantity the lower the temperature, and crystallises at -10°. The reasons both of the variation of the colour of the gas (N2O4 gives colourless and transparent vapours, whilst those of NO2 are brown and opaque) and the variation of the vapour density with the variation of temperature are thus made quite clear; and as at the boiling point a density 38 was obtained, therefore at that temperature the vapours consist of a mixture of 79 parts by weight of N2O4 with 21 parts by weight of NO2.[45] It is evident that a decomposition here takes place the peculiarity of which consists in the fact that the product of decomposition, NO2, is polymerised (i.e. becomes denser, combines with itself) at a lower temperature; that is, the reaction
N2O4 = NO2 + NO2
is a reversible reaction, and consequently the whole phenomenon represents a dissociation in a homogeneous gaseous medium, where the original substance, N2O4, and the resultant, NO2, are both gases. The measure of dissociation will be expressed if we find the proportion of the quantity of the substance decomposed to the whole amount of the substance. At the boiling point, therefore, the measure of the decomposition of nitrogen peroxide will be 21 p.c.; at 135° it = 1, and at 10° it = 0; that is, the N2O4 is not then decomposable. Consequently the limits of dissociation here are -10° and 135° at the atmospheric pressure.[46] Within the limits of these temperatures the vapours of nitrogen peroxide have not a constant density, but, on the other hand, above and below these limits definite substances exist. Thus above 135° N2O4 has ceased to exist and NO2 alone remains. It is evident that at the ordinary temperature there is a partially dissociated system or mixture of nitrogen peroxide, N2O4, and nitrogen dioxide, NO2. In the brown liquid boiling at 22° probably a portion of the N2O4 has already passed into NO2, and it is only the colourless liquid and crystalline substance at -10° that can be considered as pure nitrogen peroxide.[47]
The above explains the action of nitrogen peroxide on water at low temperatures. N2O4 then acts on water like a mixture of the anhydrides of nitrous and nitric acids. The first, N2O3, may be looked on as water in which each of the two atoms of hydrogen is replaced by the radicle NO, while in the second each hydrogen is replaced by the radicle NO2, proper to nitric acid; and in nitrogen peroxide one atom of the hydrogen of water is replaced by NO and the other by NO2, as is seen from the formulæ—