CHAPTER VI

TTHE COMPOUNDS OF NITROGEN WITH HYDROGEN AND OXYGEN

In the last chapter we saw that nitrogen does not directly combine with hydrogen, but that a mixture of these gases in the presence of hydrochloric acid gas, HCl, forms ammonium chloride, NH₄Cl, on the passage of a series of electric sparks.[1] In ammonium chloride, HCl is combined with NH₃, consequently N with H₃ forms ammonia.[2] Almost all the nitrogenous substances of plants and animals evolve ammonia when heated with an alkali. But even without the presence of an alkali the majority of nitrogenous substances, when decomposed or heated with a limited supply of air, evolve their nitrogen, if not entirely, at all events partially, in the form of ammonia. When animal substances such as skins, bones, flesh, hair, horns, &c., are heated without access of air in iron retorts—they undergo what is termed dry distillation. A portion of the resultant substances remains in the retort and forms a carbonaceous residue, whilst the other portion, in virtue of its volatility, escapes through the tube leading from the retort. The vapours given off, on cooling, form a liquid which separates into two layers; the one, which is oily, is composed of the so-called animal oils (oleum animale): the other, an aqueous layer, contains a solution of ammonia salts. If this solution be mixed with lime and heated, the lime takes up the elements of carbonic acid from the ammonia salts, and ammonia is evolved as a gas.[3] In ancient times ammonia compounds were imported into Europe from Egypt, where they were prepared from the soot obtained in the employment of camels' dung as fuel in the locality of the temple of Jupiter Ammon (in Lybia), and therefore the salt obtained was called ‘sal-ammoniacale,’ from which the name of ammonia is derived. At the present time ammonia is obtained exclusively, on a large scale, either from the products of the dry distillation of animal or vegetable refuse, from urine, or from the ammoniacal liquors collected in the destructive distillation of coal for the preparation of coal gas. This ammoniacal liquor is placed in a retort with lime and heated; the ammonia is then evolved together with steam.[4] In the arts, only a small amount of ammonia is used in a free state—that is, in an aqueous solution; the greater portion of it is converted into different salts having technical uses, especially sal-ammoniac, NH₄Cl, and ammonium sulphate, (NH₄)₂SO₄. They are saline substances which are formed because ammonia, NH₃, combines with all acids, HX, forming ammonia salts, NH₄X. Sal-ammoniac, NH₄Cl, is a compound of ammonia with hydrochloric acid. It is prepared by passing the vapours of ammonia and water, evolved, as above described, from ammoniacal liquor, into an aqueous solution of hydrochloric acid, and on evaporating the solution sal-ammoniac is obtained in the form of soluble crystals[5] resembling common salt in appearance and properties. Ammonia may be very easily prepared from this sal-ammoniac, NH₄Cl, as from any other ammoniacal salt, by heating it with lime. Calcium hydroxide, CaH₂O₂, as an alkali takes up the acid and sets free the ammonia, forming calcium chloride, according to the equation 2NH₄Cl + CaH₂O₂ = 2H₂O + CaCl₂ + 2NH₃. In this reaction the ammonia is evolved as a gas.[6]

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]

Ammonia is not only capable of combining with acids, but also with many salts, as was seen from its forming definite compounds, AgCl,3NH₃ and 2AgCl,3NH₃, with silver chloride. Just as ammonia is absorbed by various oxygen salts of the metals, so also is it absorbed by the chlorine, iodine, and bromine compounds of many metals, and in so doing evolves heat. Certain of these compounds part with their ammonia even when left exposed to the air, but others only do so at a red heat; many give up their ammonia when dissolved, whilst others dissolve without decomposition, and when evaporated separate from their solutions unchanged. All these facts only indicate that ammoniacal, like aqueous, compounds dissociate with greater or lesser facility.[19] Certain metallic oxides also absorb ammonia and are dissolved in ammonia water. Such are, for instance, the oxides of zinc, nickel, copper, and many others; the majority of such compounds are unstable. The property of ammonia of combining with certain oxides explains its action on certain metals.[20] By reason of such action, copper vessels are not suitable for holding liquids containing ammonia. Iron is not acted on by such liquids.

The similarity between the relation of ammonia and water to salts and other substances is more especially marked in those cases in which the salt is capable of combining with both ammonia and water. Take, for example, copper sulphate, CuSO₄. As we saw in Chapter I., it gives with water blue crystals, CuSO₄,5H₂O; but it also absorbs ammonia in the same molecular proportion, forming a blue substance, CuSO₄,5NH₃, and therefore the ammonia combining with salts may be termed ammonia of crystallisation.

Such are the reactions of combination proper to ammonia. Let us now turn our attention to the reactions of substitution proper to this substance. If ammonia be passed through a heated tube containing metallic sodium, hydrogen is evolved, and a compound is obtained containing ammonia in which one atom of hydrogen is replaced by an atom of sodium, NH₂Na (according to the equation NH₃ + Na = NH₂Na + H). This body is termed sodium amide. We shall afterwards see that iodine and chlorine are also capable of directly displacing hydrogen from ammonia, and of replacing it. In fact, the hydrogen of ammonia may be replaced in many ways by different elements. If in this replacement NH₂ remains, the resultant substances NH₂R are called amides, whilst the substitution products, NHR₂, in which only NH remains, are called imides,[20 bis] and those in which none of the ammoniacal hydrogen remains, NR₃, are known as nitrides. Free amidogen, N₂H₄, is now known in a state of hydration under the name of hydrazine;[21] it combines with acids and resembles ammonia in this respect. In the action of different substances on ammonia it is the hydrogen that is substituted, whilst the nitrogen remains in the resultant compound, so to say, untouched. The same phenomenon is to be observed in the action of various substances on water. In the majority of cases the reactions of water consist in the hydrogen being evolved, and in its being replaced by different elements. This also takes place, as we have seen, in acids in which the hydrogen is easily displaced by metals. This chemical mobility of hydrogen is perhaps connected with the great lightness of the atoms of this element.

In practical chemistry[21 bis] ammonia is often employed, not only for saturating acids, but also for effecting reactions of double decomposition with salts, and especially for separating insoluble basic hydroxides from soluble salts. Let MHO stand for an insoluble basic hydroxide and HX for an acid. The salt formed by them will have a composition MHO + XH - H₂O = MX. If aqueous ammonia, NH₄OH, be added to a solution of this salt, the ammonia will change places with the metal M, and thus form the insoluble basic hydroxide, or, as it is said, give a precipitate.

Thus, for instance, if aqueous ammonia is added to a solution of a salt of aluminium, then alumina hydrate is separated out as a colourless gelatinous precipitate.[22]

In order to grasp the relation between ammonia and the oxygen compounds of nitrogen it is necessary to recognise the general law of substitution, applicable to all cases of substitution between elements,[23] and therefore showing what may be the cases of substitution between oxygen and hydrogen as component parts of water. The law of substitution may be deduced from mechanical principles if the molecule be conceived as a system of elementary atoms occurring in a certain chemical and mechanical equilibrium. By likening the molecule to a system of bodies in a state of motion—for instance, to the sum total of the sun, planets, and satellites, existing in conditions of mobile equilibrium—then we should expect the action of one part, in this system, to be equal and opposite to the other, according to Newton's third law of mechanics. Hence, given a molecule of a compound, for instance, H₂O, NH₃, NaCl, HCl, &c., its every two parts must in a chemical sense represent two things somewhat alike in force and properties, and therefore every two parts into which a molecule of a compound may be divided are capable of replacing each other. In order that the application of the law should become clear it is evident that among compounds the most stable should be chosen. We will therefore take hydrochloric acid and water as the most stable compounds of hydrogen.[24] According to the above law of substitution, if the elements H and Cl are able to form a molecule, HCl, and a stable one, they are able to replace each other. And, indeed, we shall afterwards see (Chapter XI.) that in a number of instances a substitution between hydrogen and chlorine can take place. Given RH, then RCl is possible, because HCl exists and is stable. The molecule of water, H₂O, may be divided in two ways, because it contains 3 atoms: into H and (HO) on the one hand, and into H₂ and O on the other. Consequently, being given RH, its substitution products will be R(HO) according to the first form, and R₂O according to the second; being given RH₂, its corresponding substitution products will be RH(OH), R(OH)₂, RO, (RH)₂O, &c. The group (OH) is the same hydroxyl or aqueous radicle which we have already mentioned in the third chapter as a component part of hydroxides and alkalis—for instance, Na(OH), Ca(OH)₂, &c. It is evident, judging from H(HO) and HCl, that (OH) can be substituted by Cl, because both are replaceable by H; and this is of common occurrence in chemistry, because metallic chlorides—for example, NaCl and NH₄Cl—correspond with hydroxides of the alkalis Na(OH) or NH₄(OH). In hydrocarbons—for instance, C₂H₆—the hydrogen is replaceable by chlorine and by hydroxyl. Thus ordinary alcohol is C₂H₆, in which one atom of H is replaced by (OH); that is, C₂H₅(OH). It is evident that the replacement of hydrogen by hydroxyl essentially forms the phenomenon of oxidation, because RH gives R(OH), or RHO. Hydrogen peroxide may in this sense be regarded as water in which the hydrogen is replaced by hydroxyl; H(OH) gives (OH)₂ or H₂O₂. The other form of substitution—namely, that of O in the place of H₂—is also a common chemical phenomenon. Thus alcohol, C₂H₆O, or C₂H₅(OH), when oxidising in the air, gives acetic acid, C₂H₄O₂, or C₂H₃O(OH), in which H₂ is replaced by O.

In the further course of this work we shall have occasion to refer to the law of substitution for explaining many chemical phenomena and relations.

We will now apply these conceptions to ammonia in order to see its relation to the oxygen compounds of nitrogen. It is evident that many substances should be obtainable from ammonia, NH₃, or aqueous ammonia, NH₄(OH), by substituting their hydrogen by hydroxyl, or H₂ by oxygen. And such is the case. The two extreme cases of such substitution will be as follows: (1) One atom of H in NH₃ is substituted by (OH), and NH₂(OH) is produced. Such a substance, still containing much hydrogen, should have many of the properties of ammonia. It is known under the name of hydroxylamine,[25] and, in fact, is capable, like ammonia, of giving salts with acids; for example, with hydrochloric acid, NH₃(OH)Cl—which is a substance corresponding to sal-ammoniac, in which one atom of hydrogen is replaced by hydroxyl.[25 bis] (2) The other extreme case of substitution is that given by ammonium hydroxide, NH₄(OH), when the whole of the hydrogen of the ammonium is replaced by oxygen; and, as ammonium contains 4 atoms of hydrogen, the highest oxygen compound should be NO₂(OH), or NHO₃, as we find to be really the case, for NHO₃ is nitric acid, exhibiting the highest degree of oxidation of nitrogen.[26] If instead of the two extreme aspects of substitution we take an intermediate one, we obtain the intermediate oxygen compounds of nitrogen. For instance, N(OH)₃ is orthonitrous acid,[27] to which corresponds nitrous acid, NO(OH), or NHO₂, equal to N(OH)₃ - H₂O, and nitrous anhydride, N₂O₃ = 2N(OH)₃ - 3H₂O. Thus nitrogen gives a series of oxygen compounds, which we will proceed to describe. We will, however, first show by two examples that in the first place the passage of ammonia into the oxygen compounds of nitrogen up to nitric acid, as well as the converse preparation of ammonia (and consequently of the intermediate compounds also) from nitric acid, are reactions which proceed directly and easily under many circumstances, and in the second place that the above general principle of substitution gives the possibility of understanding many, at first sight unexpected and complex, relations and transformations, such as the preparation of hydronitrous acid, HN₃. In nature the matter is complicated by a number of influences and circumstances, but in the law the relations are presented in their simplest aspect.

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:—

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.

Nitric acid so obtained always contains water. It is extremely difficult to deprive it of all the admixed water without destroying a portion of the acid itself and partially converting it into lower oxides, because without the presence of an excess of water it is very unstable. When rapidly distilled a portion is decomposed, and there are obtained free oxygen and lower oxides of nitrogen, which, together with the water, remain in solution with the nitric acid. Therefore it is necessary to work with great care in order to obtain a pure hydrate of nitric acid, HNO₃, and especially to mix the nitric acid obtained from nitre, as above described, with sulphuric acid, which takes up the water, and to distil it at the lowest possible temperature—that is, by placing the retort holding the mixture in a water or oil bath and carefully heating it. The first portion of the nitric acid thus distilled boils at 86°, has a specific gravity at 15° of 1·526, and solidifies at -50°; it is very unstable at higher temperatures. This is the normal hydrate, HNO₃, which corresponds with the salts, NMO₃, of nitric acid. When diluted with water nitric acid presents a higher boiling point, not only as compared with that of the nitric acid itself, but also with that of water; so that, if very dilute nitric acid be distilled, the first portions passing over will consist of almost pure water, until the boiling point in the vapours reaches 121°. At this temperature a compound of nitric acid with water, containing about 70 p.c. of nitric acid,[31] distils over; its specific gravity at 15° = 1·421. If the solution contain less than 25 p.c. of water, then, the specific gravity of the solution being above 1·44, HNO₃ evaporates off and fumes in the air, forming the above hydrate, whose vapour tension is less than that of water. Such solutions form fuming nitric acid. On distilling it gives monohydrated acid,[32] HNO₃; it is a hydrate boiling at 121°, so that it is obtained from both weak and strong solutions. Fuming nitric acid, under the action not only of organic substances, but even of heat, loses a portion of its oxygen, forming lower oxides of nitrogen, which impart a red-brown colour to it;[33] the pure acid is colourless.

Nitric acid, as an acid hydrate, enters into reactions of double decomposition with bases, basic hydrates (alkalis), and with salts. In all these cases a salt of nitric acid is obtained. An alkali and nitric acid give water and a salt; so, also, a basic oxide with nitric acid gives a salt and water; for instance, lime, CaO + 2HNO₃ = Ca(NO₃)₂ + H₂O. Many of these salts are termed nitres.[34] The composition of the ordinary salts of nitric acid may be expressed by the general formula M(NO₃)_n, where M indicates a metal replacing the hydrogen in one or several (n) equivalents of nitric acid. We shall find afterwards that the atoms M of metals are equivalent to one (K, Na, Ag) atom of hydrogen, or two (Ca, Mg, Ba), or three (Al, In), or, in general, n atoms of hydrogen. The salts of nitric acid are especially characterised by being all soluble in water.[35] From the property common to all these salts of entering into double decompositions, and owing to the volatility of nitric acid, they evolve nitric acid when heated with sulphuric acid. They all, like the acid itself, are capable of evolving oxygen when heated, and consequently of acting as oxidising substances; they therefore, for instance, deflagrate with ignited carbon, the carbon burning at the expense of the oxygen of the salt and forming gaseous products of combustion.[36]

Nitric acid also enters into double decompositions with a number of hydrocarbons not in any way possessing alkaline characters and not reacting with other acids. Under these circumstances the nitric acid gives water and a new substance termed a nitro-compound. The chemical character of the nitro-compound is the same as that of the original substance; for example, if an indifferent substance be taken, then the nitro compound obtained from it will also be indifferent; if an acid be taken, then an acid is obtained also.[36 bis] Benzene, C₆H₆, for instance, acts according to the equation C₆H₆ + HNO₃ = H₂O + C₆H₅NO₂. Nitrobenzene is produced. The substance taken, C₆H₆, is a liquid hydrocarbon having a faint tarry smell, boiling at 80°, and lighter than water; by the action of nitric acid nitrobenzene is obtained, which is a substance boiling at about 210°, heavier than water, and having an almond-like odour: it is employed in large quantities for the preparation of aniline and aniline dyes.[37] As the nitro-compounds contain both combustible elements (hydrogen and carbon), as well as oxygen in unstable combination with nitrogen, in the form of the radicle NO₂ of nitric acid, they decompose with an explosion when ignited or even struck, owing to the pressure of the vapours and gases formed—free nitrogen, carbonic anhydride, CO₂, carbonic oxide, CO, and aqueous vapour. In the explosion of nitro compounds[37 bis] much heat is evolved, as in the combustion of gunpowder or detonating gas, and in this case the force of explosion in a closed space is great, because from a solid or liquid nitro-compound occupying a small space there proceed vapours and gases whose elasticity is great not only from the small space in which they are formed, but owing to the high temperature corresponding to the combustion of the nitro-compound.[38]