It must be observed that all the complex nitrogenous substances of plants, animals, and soils are decomposed when heated with an excess of sulphuric acid, the whole of their nitrogen being converted into ammonium sulphate, from which it may be liberated by treatment with an excess of alkali. This reaction is so complete that it forms the basis of Kjeldahl's method for estimating the amount of nitrogen in its compounds.
Fig. 45.—The liquefaction of ammonia in a thick bent glass tube. A compound of chloride of silver and ammonia is placed in the end a, and the end c is then sealed up.
Ammonia is a colourless gas, resembling those with which we are already acquainted in its outward appearance, but clearly distinguishable from any other gas by its very characteristic and pungent smell. It irritates the eyes, and it is positively impossible to inhale it. Animals die in it. Its density, referred to hydrogen, is 8·5; hence it is lighter than air. It belongs to the class of gases which are easily liquefied.[7] Faraday employed the following method for liquefying ammonia. Ammonia when passed over dry silver chloride, AgCl, is absorbed by it to a considerable extent, especially at low temperatures.[8] The solid compound AgCl,3NH3 thus obtained is introduced into a bent tube (fig. [45]), whose open end c is then fused up. The compound is then slightly heated at a, and the ammonia comes off, owing to the easy dissociation of the compound. The other end of the tube is immersed in a freezing mixture. The pressure of the gas coming off, combined with the low temperature at one end of the tube, causes the ammonia evolved to condense into a liquid, in which form it collects at the cold end of the tube. If the heating be stopped, the silver chloride again absorbs the ammonia. In this manner one tube may serve for repeated experiments. Ammonia may also be liquefied by the ordinary methods—that is, by means of pumping dry ammonia gas into a refrigerated space. Liquefied ammonia is a colourless and very mobile liquid,[9] whose specific gravity at 0° is 0·63 (E. Andréeff). At the temperature (about -70°) given by a mixture of liquid carbonic anhydride and ether, liquid ammonia crystallises, and in this form its odour is feeble, because at so low a temperature its vapour tension is very inconsiderable. The boiling point (at a pressure of 760 mm.) of liquid ammonia is about -32°. Hence this temperature may be obtained at the ordinary pressure by the evaporation of liquefied ammonia.
Ammonia, containing, as it does, much hydrogen, is capable of combustion; it does not, however, burn steadily, and sometimes not at all, in ordinary atmospheric air. In pure oxygen it burns with a greenish-yellow flame,[10] forming water, whilst the nitrogen set free gives its oxygen compounds—that is, oxides of nitrogen. The decomposition of ammonia into hydrogen and nitrogen not only takes place at a red heat and under the action of electric sparks, but also by means of many oxidising substances; for instance, by passing ammonia through a tube containing red-hot copper oxide. The water thus formed may be collected by substances absorbing it, and the quantity of nitrogen may be measured in a gaseous form, and thus the composition of ammonia determined. In this manner it is very easy to prove that ammonia contains 3 parts by weight of hydrogen to 14 parts by weight of nitrogen; and, by volume, 3 vols. of hydrogen and 1 vol. of nitrogen form 2 vols. of ammonia.[11]
Ammonia is capable of combining with a number of substances, forming, like water, substances of various degrees of stability. It is more soluble than any of the gases yet described, both in water and in many aqueous solutions. We have already seen, in the [first chapter], that one volume of water, at the ordinary temperature, dissolves about 700 vols. of ammonia gas. The great solubility of ammonia enables it to be always kept ready for use in the form of an aqueous solution,[12] which is commercially known as spirits of hartshorn. Ammonia water is continually evolving ammoniacal vapour, and so has the characteristic smell of ammonia itself. It is a very characteristic and important fact that ammonia has an alkaline reaction, and colours litmus paper blue, just like caustic potash or lime; it is therefore sometimes called caustic ammonia (volatile alkali). Acids may be saturated by ammonia water or gas in exactly the same way as by any other alkali. In this process ammonia combines directly with acids, and this forms the most essential chemical reaction of this substance. If sulphuric, nitric, acetic, or any other acid be brought into contact with ammonia it absorbs it, and in so doing evolves a large amount of heat and forms a compound having all the properties of a salt. Thus, for example, sulphuric acid, H2SO4, in absorbing ammonia, forms (on evaporating the solution) two salts, according to the relative quantities of ammonia and acid. One salt is formed from NH3 + H2SO4, and consequently has the composition NH5SO4, and the other is formed from 2NH3 + H2SO4, and its composition is therefore N2H8SO4. The former has an acid reaction and the latter a neutral reaction, and they are called respectively acid ammonium sulphate (ammonium hydrogen sulphate), and normal ammonium sulphate, or simply ammonium sulphate. The same takes place in the action of all other acids; but certain of them are able to form normal ammonium salts only, whilst others give both acid and normal ammonium salts. This depends on the nature of the acid and not on the ammonia, as we shall afterwards see. Ammonium salts are very similar in appearance and in many of their properties to metallic salts; for instance, sodium chloride, or table salt, resembles sal-ammoniac, or ammonium chloride, not only in its outward appearance but even in crystalline form, in its property of giving precipitates with silver salts, in its solubility in water, and in its evolving hydrochloric acid when heated with sulphuric acid—in a word, a most perfect analogy is to be remarked in an entire series of reactions. An analogy in composition is seen if sal-ammoniac, NH4Cl, be compared with table salt, NaCl; and the ammonium hydrogen sulphate, NH4HSO4, with the sodium hydrogen sulphate, NaHSO4; or ammonium nitrate, NH4NO3, with sodium nitrate, NaNO3.[13] It is seen, on comparing the above compounds, that the part which sodium takes in the sodium salts is played in ammonium salts by a group NH4, which is called ammonium. If table salt be called ‘sodium chloride,’ then sal-ammoniac should be and is called ‘ammonium chloride.’
The hypothesis that ammoniacal salts correspond with a complex metal ammonium bears the name of the ammonium theory. It was enunciated by the famous Swedish chemist Berzelius after the proposition made by Ampère. The analogy admitted between ammonium and metals is probable, owing to the fact that mercury is able to form an amalgam with ammonium similar to that which it forms with sodium or many other metals. The only difference between ammonium amalgam and sodium amalgam consists in the instability of the ammonium, which easily decomposes into ammonia and hydrogen.[14] Ammonium amalgam may be prepared from sodium amalgam. If the latter be shaken up with a strong solution of sal-ammoniac, the mercury swells up violently and loses its mobility whilst preserving its metallic appearance. In so doing, the mercury dissolves ammonium—that is, the sodium in the mercury is replaced by the ammonium, and replaces it in the sal-ammoniac, forming sodium chloride, NH4Cl + HgNa = NaCl + HgNH4. Naturally, the formation of ammonium amalgam does not entirely prove the existence of ammonium itself in a separate state; but it shows the possibility of this substance existing, and its analogy with the metals, because only metals dissolve in mercury.[15] Ammonium amalgam crystallises in cubes, three times heavier than water; it is only stable in the cold, and particularly at very low temperatures. It begins to decompose at the ordinary temperature, evolving ammonia and hydrogen in the proportion of two volumes of ammonia and one volume of hydrogen, NH4 = NH3 + H. By the action of water, ammonium amalgam gives hydrogen and ammonia water, just as sodium amalgam gives hydrogen and sodium hydroxide; and therefore, in accordance with the ammonium theory, ammonia water must be looked on as containing ammonium hydroxide, NH4OH,[16] just as an aqueous solution of sodium hydroxide, contains NaOH. The ammonium hydroxide, like ammonium itself, is an unstable substance, which easily dissociates, and can only exist in a free state at low temperatures.[17] Ordinary solutions of ammonia must be looked on as the products of the dissociation of this hydroxide, inasmuch as NH4OH = NH3 + H2O.
All ammoniacal salts decompose at a red heat into ammonia and an acid, which, on cooling in contact with each other, re-combine together. If the acid be non-volatile, the ammoniacal salt, when heated, evolves the ammonia, leaving the non-volatile acid behind; if the acid be volatile, then, on heating, both the acid and ammonia volatilise together, and on cooling re-combine into the salt which originally served for the formation of their vapours.[18]