1. It is easy to prove the possibility of the oxidation of ammonia into nitric acid by passing a mixture of ammonia and air over heated spongy platinum. This causes the oxidation of the ammonia, nitric acid being formed, which partially combines with the excess of ammonia.
The converse passage of nitric acid into ammonia is effected by the action of hydrogen at the moment of its evolution.[28] Thus metallic aluminium, evolving hydrogen from a solution of caustic soda, is able to completely convert nitric acid added to the mixture (as a salt, because the alkali gives a salt with the nitric acid) into ammonia, NHO3 + 8H = NH3 + 3H2O.
2. In 1890 Curtius in Germany obtained a gaseous substance of the composition HN3 (hydrogen trinitride), having the distinctive properties of an acid, and giving, like hydrochloric acid, salts; for example, a sodium salt, NaN3; ammonium salt, NH4N3 = N4H4; barium salt, Ba(N3)2, &c., which he therefore named hydronitrous acid, HN3.[28 bis] The extraordinary composition of the compound (ammonia, NH3, contains one N atom and three H atoms; in HN3, on the contrary, there are three N atoms and one H atom), the facile decomposition of its salts with an explosion, and above all its distinctly acid character (an aqueous solution shows a strong acid reaction to litmus), not only indicated the importance of this unexpected discovery, but at first gave rise to some perplexity as to the nature of the substance obtained, for the relations in which HN3 stood to other simple compounds of nitrogen which had long been known was not at all evident, and the scientific spirit especially requires that there should be a distinct bond between every innovation, every fresh discovery, and that which is already firmly established and known, for upon this basis is founded that apparently paradoxical union in science of a conservative stability with an irresistible and never-ceasing improvement. This missing, connection between the newly discovered hydronitrous acid, HN3, and the long known ammonia, NH3, and nitric acid, HNO3, may be found in the law of substitution, starting from the well-known properties and composition of nitric acid and ammonia, as I mentioned in the ‘Journal of the Russian Physico-Chemical Society’ (1890). The essence of the matter lies in the fact that to the hydrate of ammonium, or caustic ammonia, NH4OH, there should correspond, according to the law of substitution, an ortho-nitric acid (see Note [27]), H3NO4 = NO(OH)3, which equals NH4(OH) with the substitution in it of (a) two atoms of hydrogen by oxygen (O—H2) and (b) two atoms of hydrogen by the aqueous radicle (OH—H). Ordinary or meta-nitric acid is merely this ortho-nitric acid minus water. To ortho-nitric acid there should correspond the ammoniacal salts: mono-substituted, H2NH4NO4; bi-substituted, H(NH4)2NO4; and tri-substituted, (NH4)3NO4. These salts, containing as they do hydrogen and oxygen, like many similar ammoniacal salts (see, for instance, Chapter [IX].—Cyanides), are able to part with them in the form of water. Then from the first salt we have H2NH4NO4 - 4H2O = N2O—nitrous oxide, and from the second H(NH4)2NO4 - 4H2O = HN3—hydronitrous acid, and from the third (NH4)3NO - 4H2O = N4H4—the ammonium salt of the same acid. The composition of HN3 should be thus understood, whilst its acid properties are explained by the fact that the water (4H2O) from H(NH4)2N_O4 is formed at the expense of the hydrogen of the ammonium and oxygen of the nitric acid, so that there remains the same hydrogen as in nitric acid, or that which may be replaced by metals and give salts. Moreover, nitrogen undoubtedly belongs to that category of metalloids which give acids, like chlorine and carbon, and therefore, under the influence of three of its atoms, one atom of hydrogen acquires those properties which it has in acids, just as in HCN (hydrocyanic acid) the hydrogen has received these properties under the influence of the carbon and nitrogen (and HN3 may be regarded as HCN where C has been replaced by N2). Moreover, besides explaining the composition and acid properties of HN3, the above method gives the possibility of foretelling the closeness of the bond between hydronitrous acid and nitrous oxide, for N2O + NH3 = HN3 + H2O. This reaction, which was foreseen from the above considerations, was accomplished by Wislicenus (1892) by the synthesis of the sodium salt, by taking the amide of sodium, NH2Na (obtained by heating Na in a current of NH3), and acting upon it (when heated) with nitrous oxide, N2O, when 2NH2Na + N2O = NaN3 + NaHO + NH3. The resultant salt, NaN3, gives hydronitrous acid when acted upon by sulphuric acid, NaN3 + H2SO4 = NaHSO4 +HN3. The latter gives, with the corresponding solutions of their salts, the insoluble (and easily explosive) salts of silver, AgN3 (insoluble, like AgCl or AgCN), and lead, Pb(N3)2.
The compounds of nitrogen with oxygen present an excellent example of the law of multiple proportions, because they contain, for 14 parts by weight of nitrogen, 8, 16, 24, 32, and 40 parts respectively by weight of oxygen. The composition of these compounds is as follows:—
| N2O, nitrous oxide; hydrate NHO. |
| N2O2, nitric oxide, NO. |
| N2O3, nitrous anhydride; hydrate NHO2. |
| N2O4, peroxide of nitrogen, NO2. |
| N2O5, nitric anhydride; hydrate NHO3. |
Of these compounds,[29] nitrous and nitric oxides, peroxide of nitrogen, and nitric acid, NHO3, are characterised as being the most stable. The lower oxides, when coming into contact with the higher, may give the intermediate forms; for instance, NO and NO2 form N2O3, and the intermediate oxides may, in splitting up, give a higher and lower oxide. So N2O4 gives N2O3 and N2O5, or, in the presence of water, their hydrates.
We have already seen that, under certain conditions, nitrogen combines with oxygen, and we know that ammonia may he oxidised. In these cases various oxidation products of nitrogen are formed, but in the presence of water and an excess of oxygen they always give nitric acid. Nitric acid, as corresponding with the highest oxide, is able, in deoxidising, to give the lower oxides; it is the only nitrogen acid whose salts occur somewhat widely in nature, and it has many technical uses, for which reason we will begin with it.
Nitric acid, NHO3, is likewise known as aqua fortis. In a free state it is only met with in nature in small quantities, in the air and in rain-water after storms; but even in the atmosphere nitric acid does not long remain free, but combines with ammonia, traces of which are always found in air. On falling on the soil and into running water, &c., the nitric acid everywhere comes into contact with bases (or their carbonates), which easily act on it, and therefore it is converted into the nitrates of these bases. Hence nitric acid is always met with in the form of salts in nature. The soluble salts of nitric acid are called nitres. This name is derived from the Latin sal nitri. The potassium salt, KNO3, is common nitre, and the sodium salt, NaNO3, Chili saltpetre, or cubic nitre. Nitres are formed in the soil when a nitrogenous substance is slowly oxidised in the presence of an alkali by means of the oxygen of the atmosphere. In nature there are very frequent instances of such oxidation. For this reason certain soils and rubbish heaps—for instance, lime rubbish (in the presence of a base)—lime contain a more or less considerable amount of nitre. One of these nitres—sodium nitrate—is extracted from the earth in large quantities in Chili, where it was probably formed by the oxidation of animal refuse. This kind of nitre is employed in practice for the manufacture of nitric acid and the other oxygen compounds of nitrogen. Nitric acid is obtained from Chili saltpetre by heating it with sulphuric acid. The hydrogen of the sulphuric acid replaces the sodium in the nitre. The sulphuric acid then forms either an acid salt, NaHSO4, or a normal salt, Na2SO4, whilst nitric acid is formed from the nitre and is volatilised. The decomposition is expressed by the equations: (1) NaNO3 + H2SO4 = HNO3 + NaHSO4, if the acid salt be formed, and (2) 2NaNO3 + H2SO4 = Na2SO4 + 2HNO3, if the normal sodium sulphate is formed. With an excess of sulphuric acid, at a moderate heat, and at the commencement of the reaction, the decomposition proceeds according to the first equation; and on further heating with a sufficient amount of nitre according to the second, because the acid salt NaHSO4 itself acts like an acid (its hydrogen being replaceable as in acids), according to the equation NaNO3 + NaHSO4 = Na2SO4 + HNO3.
The sulphuric acid, as it is said, here displaces the nitric acid from its compound with the base.[29 bis] Thus, in the reaction of sulphuric acid on nitre there is formed a non-volatile salt of sulphuric acid, which remains, together with an excess of this acid, in the distilling apparatus, and nitric acid, which is converted into vapour, and may be condensed, because it is a liquid and volatile substance. On a small scale, this reaction may be carried on in a glass retort with a glass condenser. On a large scale, in chemical works, the process is exactly similar, only iron retorts are employed for holding the mixture of nitre and sulphuric acid, and earthenware three-necked bottles are used instead of a condenser,[30] as shown in fig. [47].
