see caption — Fig. 42.—The dry distillation of bones on a large scale. The bones are heated in the vertical cylinders C (about 1½ metre high and 30 centimetres in diameter). The products of distillation pass through the tubes T, into the condenser B, and receiver F. When the distillation is completed the trap H is opened, and the burnt bones are loaded into trucks V. The roof M is then opened, and the cylinders are charged with a fresh quantity of bones. The ammonia water is preserved, and goes to the preparation of ammoniacal salts, as described in the following drawing.

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 crystals⁵ 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

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.

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,²¹ 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;²² 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 chemistry²³ 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.

MX	+	NH 4 (OH)	=	NH 4 X	+	MHO
Salt of the metal. In solution		Aqueous ammonia. In solution		Ammonium salt. In solution		Basic hydrate. As 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.²⁴

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,²⁵ 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.²⁶ 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,²⁷ 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.²⁸ (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.²⁹ 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,³⁰ 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.

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.³¹ 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.

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, HN₃.³² 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 ³⁰), 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:—

N 2 O, nitrous oxide; hydrate NHO.
N 2 O 2 , nitric oxide, NO.
N 2 O 3 , nitrous anhydride; hydrate NHO 2 .
N 2 O 4 , peroxide of nitrogen, NO 2 .
N 2 O 5 , nitric anhydride; hydrate NHO 3 .

Of these compounds,³³ 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.³⁴ 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,³⁵ 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,³⁶ 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,³⁷ 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;³⁸ 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.³⁹ 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.⁴⁰ 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.⁴¹

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.⁴² 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.⁴³ 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 compounds⁴⁴ 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.⁴⁵

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.⁴⁶ 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.⁴⁷ 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.⁴⁸ 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.⁴⁹ 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.⁵⁰

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,⁵¹ 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 NO₂.⁵² 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.⁵³ 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.⁵⁴

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—

	H		O;	NO		O;	NO		O;	NO 2		O;
	H	NO	NO 2	NO 2
or	H 2 O	N 2 O 3	N 2 O 4	N 2 O 5

In fact, nitrogen peroxide at low temperatures gives with water (ice) both nitric, HNO₃, and nitrous, HNO₂, acids. The latter, as we shall afterwards see, splits up into water and the anhydride, N₂O₃. If, however, warm water act on nitrogen peroxide, only nitric acid and monoxide of nitrogen are formed: 3NO₂ + H₂O = NO + 2NHO₃.

Although NO₂ is not decomposed into N and O even at 500°, still in many cases it acts as an oxidising agent. Thus, for instance, it oxidises mercury, converting it into mercurous nitrate, 2NO₂ + Hg = HgNO₃ + NO, being itself deoxidised into nitric oxide, into which the dioxide in many other instances passes, and from which it is easily formed.⁵⁵

Nitrous anhydride, N₂O₃, corresponds⁵⁶ to nitrous acid, NHO₂, which forms a series of salts, the nitrites—for example, the sodium salt NaNO₂, the potassium salt KNO₂, the ammonium salt (NH₄)NO₂,⁵⁷ the silver salt AgNO₂,⁵⁸ &c. Neither the anhydride nor the hydrate of the acid is known in a perfectly pure state. The anhydride has only been obtained as a very unstable substance, and has not yet been fully investigated; and on attempting to obtain the acid NHO₂ from its salts, it always gives water and the anhydride, whilst the latter, as an intermediate oxide, partially or wholly splits up into NO + NO₂. But the salts of nitrous acid are distinguished for their great stability. Potassium nitrate, KNO₃, may be converted into potassium nitrite by depriving it of a portion of its oxygen; for instance, by fusing it (at not too high a temperature) with metals, such as lead, KNO₃ + Pb = KNO₂ + PbO.⁵⁹ The resultant salt is soluble in water, whilst the oxide of lead is insoluble. With sulphuric and other acids the solution of potassium nitrite⁶⁰ immediately evolves a brown gas, nitrous anhydride: 2KNO₂ + H₂SO₄ = K₂SO₄ + N₂O₃ + H₂O. The same gas (N₂O₃) is obtained by passing nitric oxide at 0° through liquid peroxide of nitrogen,⁶¹ or by heating starch with nitric acid of sp. gr. 1·3. At a very low temperature it condenses into a blue liquid boiling below 0°,⁶² but then partially decomposing into NO + NO₂. Nitrous anhydride possesses a remarkable capacity for oxidising. Ignited bodies burn in it, nitric acid absorbs it, and then acquires the property of acting on silver and other metals, even when diluted. Potassium iodide is oxidised by this gas just as it is by ozone (and by peroxide of hydrogen, chromic and other acids, but not by dilute nitric acid nor by sulphuric acid), with the separation of iodine. This iodine may he recognised (see Ozone, Chapter IV.) by its turning starch blue. Very small traces of nitrites may be easily detected by this method. If, for example, starch and potassium iodide are added to a solution of potassium nitrite (at first there will be no change, there being no free nitrous acid), and then sulphuric acid be added, the nitrous acid (or its anhydride) immediately set free liberates iodine, which produces a blue colour with the starch. Nitric acid does not act in this manner, but in the presence of zinc the coloration takes place, which proves the formation of nitrous acid in the deoxidation of nitric acid.⁶³ Nitrous acid acts directly on ammonia, forming nitrogen and water, HNO₂ + NH₃ = N₂ + 2H₂O.⁶⁴

As nitrous anhydride easily splits up into NO₂ + NO, so, like NO₂, with warm water it gives nitric acid and nitric oxide, according to the equation 3N₂O₃ + H₂O = 4NO + 2NHO₃.

Being in a lower degree of oxidation than nitric acid, nitrous acid and its anhydride are oxidised in solutions by many oxidising substances—for example, by potassium permanganate—into nitric acid.⁶⁵

Nitric oxide, NO.—This permanent gas⁶⁶ (that is, unliquefiable by pressure without the aid of cold) may be obtained from all the above-described compounds of nitrogen with oxygen. The deoxidation of nitric acid by metals is the usual method employed for its preparation. Dilute nitric acid (sp. gr. 1·18, but not stronger, as then N₂O₃ and NO₂ are produced) is poured into a flask containing metallic copper.⁶⁷ The reaction commences at the ordinary temperature. Mercury and silver also give nitric oxide with nitric acid. In these reactions with metals one portion of the nitric acid is employed in the oxidation of the metal, whilst the other, and by far the greater, portion combines with the metallic oxide so obtained, with formation of the nitrate corresponding with the metal taken. The first action of the copper on the nitric acid is thus expressed by the equation

2NHO₃ + 3Cu = H₂O + 3CuO + 2NO.

The second reaction consists in the formation of copper nitrate—

6NHO₃ + 3CuO = 3H₂O + 3Cu(NO₃)₂.

Nitric oxide is a colourless gas which is only slightly soluble in water (

1 / 20

of a volume at the ordinary temperature). Reactions of double decomposition in which nitric oxide readily takes part are not known—that is to say, it is an indifferent, not a saline, oxide. Like the other oxides of nitrogen, it is decomposed into its elements at a red heat (starting from 900°, at 1,200° 60 per cent. give N₂ and 2N₂O₃, but complete decomposition into N₂ and O₂ only takes place at the melting point of platinum, Emich 1892). The most characteristic property of nitric oxide is its capacity for directly and easily combining with oxygen (owing to the evolution of heat in the combination). With oxygen it forms nitrous anhydride and nitrogen peroxide, 2NO + O = N₂O₃, 2NO + O₂ = 2NO₂. If nitric oxide is mixed with oxygen and immediately shaken up with caustic potash, it is almost entirely converted into potassium nitrite; whilst after a certain time, when the formation of nitric peroxide has already commenced, a mixture of potassium nitrite and nitrate is obtained. If oxygen is passed into a bell jar filled with nitric oxide, brown fumes of nitrous anhydride and nitric peroxide are formed, even in the absence of moisture; these in the presence of water give, as we already know, nitric acid and nitric oxide, so that in the presence of an excess of water and oxygen the whole of the nitric oxide is easily and directly converted into nitric acid. This reaction of the re-formation of nitric acid from nitric oxide, air, and water, 2NO + H₂O + O₃ = 2HNO₃, is frequently made use of in practice. The experiment showing the conversion of nitric oxide into nitric acid is very striking and instructive. As the intermixture of the oxygen with the oxide of nitrogen proceeds, the nitric acid formed dissolves in water, and if an excess of oxygen has not been added the whole of the gas (nitric oxide), being converted into HNO₃, is absorbed, and the water entirely fills the bell jar previously containing the gas.⁶⁸ It is evident that nitric oxide⁶⁹ in combining with oxygen has a strong tendency to give the higher types of nitrogen compounds, which we see in nitric acid, HNO₃ or NO₂(OH), in nitric anhydride, N₂O₅ or (NO₂)₂O, and in ammonium chloride, NH₄Cl. If X stand for an atom of hydrogen, or its equivalents, chlorine, hydroxyl, &c., and if O, which is, according to the law of substitution, equivalent to H₂, be indicated by X₂, then the three compounds of nitrogen above named should be considered as compounds of the type or form NX₅. For example, in nitric acid X₅ = O₂ + (OH), where O₂ = X₄, and OH = X; whilst nitric oxide is a compound of the form NX₂. Hence this lower form, like lower forms in general, strives by combination to attain to the higher forms proper to the compounds of a given element. NX₂ passes consecutively into NX₃—namely, into N₂O₃ and NHO₂, NX₄ (for instance NO₂) and NX₅.

As the decomposition of nitric oxide begins at temperatures above 900°, many substances burn in it; thus, ignited phosphorus continues to burn in nitric oxide, but sulphur and charcoal are extinguished in it. This is due to the fact that the heat evolved in the combustion of these two substances is insufficient for the decomposition of the nitric oxide, whilst the heat developed by burning phosphorus suffices to produce this decomposition. That nitric oxide really supports combustion, owing to its being decomposed by the action of heat, is proved by the fact that strongly ignited charcoal continues to burn in the same nitric oxide⁷⁰ in which a feebly incandescent piece of charcoal is extinguished.

The compounds of nitrogen with oxygen which we have so far considered may all be prepared from nitric oxide, and may themselves be converted into it. Thus nitric oxide stands in intimate connection with them.⁷¹ The passage of nitric oxide into the higher degrees of oxidation and the converse reaction is employed in practice as a means for transferring the oxygen of the air to substances capable of being oxidised. Starting with nitric oxide, it may easily be converted, with the aid of the oxygen of the atmosphere and water, into nitric acid, nitrous anhydride, and nitric peroxide, and by their means employed to oxidise other substances. In this oxidising action nitric oxide is again formed, and it may again be converted into nitric acid, and so on continuously, if only oxygen and water be present. Hence the fact, which at first appears to be a paradox, that by means of a small quantity of nitric oxide in the presence of oxygen and water it is possible to oxidise an indefinitely large quantity of substances which cannot be directly oxidised either by the action of the atmospheric oxygen or by the action of nitric oxide itself. The sulphurous anhydride, SO₂, which is obtained in the combustion of sulphur and in roasting many metallic sulphides in the air is an example of this kind. In practice this gas is obtained by burning sulphur or iron pyrites, the latter being thereby converted into oxide of iron and sulphurous anhydride. In contact with the oxygen of the atmosphere this gas does not pass into the higher degree of oxidation, sulphuric anhydride, SO₃, and if it does form sulphuric acid with water and the oxygen of the atmosphere, SO₂ + H₂O + O = H₂SO₄, it does so very slowly. With nitric acid (and especially with nitrous acid, but not with nitrogen peroxide) and water, sulphurous anhydride, on the contrary, very easily forms sulphuric acid, and especially so when slightly heated (about 40°), the nitric acid (or, better still, nitrous acid) being converted into nitric oxide—

3SO₂ + 2NHO₃ + 2H₂O = 2H₂SO₄ + 2NO.

The presence of water is absolutely indispensable here, otherwise sulphuric anhydride is formed, which combines with the oxides of nitrogen (nitrous anhydride), forming a crystalline substance containing oxides of nitrogen (chamber crystals, which will be described in Chapter XX.) Water destroys this compound, forming sulphuric acid and separating the oxides of nitrogen. The water must be taken in a greater quantity than that required for the formation of the hydrate H₂SO₄, because the latter absorbs oxides of nitrogen. With an excess of water, however, solution does not take place. If, in the above reaction, only water, sulphurous anhydride, and nitric or nitrous acid be taken in a definite quantity, then a definite quantity of sulphuric acid and nitric oxide will be formed, according to the preceding equation; but there the reaction ends and the excess of sulphurous anhydride, if there be any, will remain unchanged. But if we add air and water, then the nitric oxide will unite with the oxygen to form nitrogen peroxide, and the latter with water to form nitric and nitrous acids, which again give sulphuric acid from a fresh quantity of sulphurous anhydride. Nitric oxide is again formed, which is able to start the oxidation afresh if there be sufficient air. Thus it is possible with a definite quantity of nitric oxide to convert an indefinitely large quantity of sulphurous anhydride into sulphuric acid, water and oxygen only being required.⁷² This may be easily demonstrated by an experiment on a small scale, if a certain quantity of nitric oxide be first introduced into a flask, and sulphurous anhydride, steam, and oxygen be then continually passed in. Thus the above-described reaction may be expressed in the following manner:—

nSO₂ + nO + (n + m)H₂O + NO = nH₂SO₄,mH₂O + NO

if we consider only the original substances and those finally formed. In this way a definite quantity of nitric oxide may serve for the conversion of an indefinite quantity of sulphurous anhydride, oxygen, and water into sulphuric acid. In reality, however, there is a limit to this, because air, and not pure oxygen, is employed for the oxidation, so that it is necessary to remove the nitrogen of the air and to introduce a fresh quantity of air. A certain quantity of nitric oxide will pass away with this nitrogen, and will in this way be lost.⁷³

The preceding series of changes serve as the basis of the manufacture of sulphuric acid or so-called chamber acid. This acid is prepared on a very large scale in chemical works because it is the cheapest acid whose action can be applied in a great number of cases. It is therefore used in immense quantities.

The process is carried on in a series of chambers (or in one divided by partitions as in fig. 50, which shows the beginning and end of a chamber) constructed of sheet lead. These chambers are placed one against the other, and communicate by tubes or special orifices so placed that the inlet tubes are in the upper portion of the chamber, and the outlet in the lower and opposite end. The current of steam and gases necessary for the preparation of the sulphuric acid passes through these chambers and tubes. The acid as it is formed falls to the bottom of the chambers or runs down their walls, and flows from chamber to chamber (from the last towards the first), to permit of which the partitions do not reach to the bottom. The floor and walls of the chambers should therefore be made of a material on which the sulphuric acid will not act. Among the ordinary metals lead is the only one suitable.⁷⁴

For the formation of the sulphuric acid it is necessary to introduce sulphurous anhydride, steam, air, and nitric acid, or some oxide of nitrogen, into the chambers. The sulphurous anhydride is produced by burning sulphur or iron pyrites. This is carried on in the furnace with four hearths to the left of the drawing. Air is led into the chambers and furnace through orifices in the furnace doors. The current of air and oxygen is regulated by opening or closing these orifices to a greater or less extent. The ingoing draught in the chambers is brought about by the fact that heated gases and vapours pass into the chambers, whose temperature is further raised by the reaction itself, and also by the remaining nitrogen being continually withdrawn from the outlet (above the tower K) by a tall chimney situated near the chambers. Nitric acid is prepared from a mixture of sulphuric acid and Chili saltpetre, in the same furnaces in which the sulphurous anhydride is evolved (or in special furnaces). Not more than 8 parts of nitre are taken to 100 parts of sulphur burnt. On leaving the furnace the vapours of nitric acid and oxides of nitrogen mixed with air and sulphurous anhydride first pass along the horizontal tubes T into the receiver B B, which is partially cooled by water flowing in on the right-hand side and running out on the left by o, in order to reduce the temperature of the gases entering the chamber. The gases then pass up a tower filled with coke, and shown to the left of the drawing. In this tower are placed lumps of coke (the residue from the dry distillation of coal), over which sulphuric acid trickles from the reservoir M. This acid has absorbed in the end tower K the oxides of nitrogen escaping from the chamber. This end tower is also filled with coke, over which a stream of strong sulphuric acid trickles from the reservoir M. The acid spreads over the coke, and, owing to the large surface offered by the latter, absorbs the greater part of the oxides of nitrogen escaping from the chambers. The sulphuric acid in passing down the tower becomes saturated with the oxides of nitrogen, and flows out at h into a special receiver (in the drawing situated by the side of the furnaces), from which it is forced up the tubes h′ h′ by steam pressure into the reservoir M, situated above the first tower. The gases passing through this tower (hot) from the furnace on coming into contact with the sulphuric acid take up the oxides of nitrogen contained in it, and these are thus returned to the chamber and again participate in the reaction. The sulphuric acid left after their extraction flows into the chambers. Thus, on leaving the first coke tower the sulphurous anhydride, air, and vapours of nitric acid and of the oxides of nitrogen pass through the upper tube m into the chamber. Here they come into contact with steam introduced by lead tubes into various parts of the chamber. The reaction takes place in the presence of water, the sulphuric acid falls to the bottom of the chamber, and the same process takes place in the following chambers until the whole of the sulphurous anhydride is consumed. A somewhat greater proportion of air than is strictly necessary is passed in, in order that no sulphurous anhydride should be left unaltered for want of sufficient oxygen. The presence of an excess of oxygen is shown by the colour of the gases escaping from the last chamber. If they be of a pale colour it indicates an insufficiency of air (and the presence of sulphurous anhydride), as otherwise peroxide of nitrogen would be formed. A very dark colour shows an excess of air, which is also disadvantageous, because it increases the inevitable loss of nitric oxide by increasing the mass of escaping gases.⁷⁵

Nitrous oxide, N₂O,⁷⁶ is similar to water in its volumetric composition. Two volumes of nitrous oxide are formed from two volumes of nitrogen and one volume of oxygen, which may be shown by the ordinary method for the analysis of the oxides of nitrogen (by passing them over red-hot copper or sodium). In contradistinction to the other oxides of nitrogen, it is not directly oxidised by oxygen, but it may be obtained from the higher oxides of nitrogen by the action of certain deoxidising substances; thus, for example, a mixture of two volumes of nitric oxide and one volume of sulphurous anhydride if left in contact with water and spongy platinum is converted into sulphuric acid and nitrous oxide, 2NO + SO₂ + H₂O = H₂SO₄ + N₂O. Nitric acid, also, under the action of certain metals—for instance, of zinc⁷⁷—gives nitrous oxide, although in this case mixed with nitric oxide. The usual method of preparing nitrous oxide consists in the decomposition of ammonium nitrate by the aid of heat, because in this case only water and nitrous oxide are formed, NH₄NO₃ = 2H₂O + N₂O (a mixture of NH₄Cl and KNO₃ is sometimes taken). The decomposition⁷⁸ proceeds very easily in an apparatus like that used for the preparation of ammonia or oxygen—that is, in a retort or flask with a gas-conducting tube. The decomposition must, however, be carried on carefully, as otherwise nitrogen is formed from the decomposition of the nitrous oxide.⁷⁹

Nitrous oxide is not a permanent gas (absolute boiling point +36°); it is easily liquefied by the action of cold under a high pressure; at 15° it may be liquefied by a pressure of about 40 atmospheres. This gas is usually liquefied by means of the force pump⁸⁰ shown in fig. 51. As it is liquefied with comparative ease, and as the cold produced by its vaporisation is very considerable,⁸¹ it (as also liquid carbonic anhydride) is often employed in investigations requiring a low temperature. Nitrous oxide forms a very mobile, colourless liquid, which acts on the skin, and is incapable in a cold state of oxidising either metallic potassium, phosphorus, or carbon; its specific gravity is slightly less than that of water (0° = 0·910, 10° = 0·856, 35° = 0·60, 39° = 0·45, Villard, 1894). When evaporated under the receiver of an air-pump, the temperature falls to -100°, and the liquid solidifies into a snow-like mass, and partially forms transparent crystals. Both these substances are solid nitrous oxide. Mercury is immediately solidified in contact with evaporating liquid nitrous oxide.⁸²

When introduced into the respiratory organs (and consequently into the blood also) nitrous oxide produces a peculiar kind of intoxication accompanied by spasmodic movements, and hence this gas, discovered by Priestley in 1776, received the name of ‘laughing gas.’ On a prolonged respiration it produces a state of insensibility (it is an anæsthetic like chloroform), and is therefore employed in dental and surgical operations.

Nitrous oxide is easily decomposed into nitrogen and oxygen by the action of heat, or a series of electric sparks; and this explains why a number of substances which cannot burn in nitric oxide do so with great ease in nitrous oxide. In fact, when nitric oxide gives some oxygen on decomposition, this oxygen immediately unites with a fresh portion of the gas to form nitric peroxide, whilst nitrous oxide does not possess this capacity for further combination with oxygen.⁸³ A mixture of nitrous oxide with hydrogen explodes like detonating gas, gaseous nitrogen being formed, N₂O + H₂ = H₂O + N₂. The volume of the remaining nitrogen is equal to the original volume of nitrous oxide, and is equal to the volume of hydrogen entering into combination with the oxygen; hence in this reaction equal volumes of nitrogen and hydrogen replace each other. Nitrous oxide is also very easily decomposed by red-hot metals; and sulphur, phosphorus, and charcoal burn in it, although not so brilliantly as in oxygen. A substance in burning in nitrous oxide evolves more heat than an equal quantity burning in oxygen; which most clearly shows that in the formation of nitrous oxide by the combination of nitrogen with oxygen there was not an evolution but an absorption of heat, there being no other source for the excess of heat in the combustion of substances in nitrous oxide (see Note ³³). If a given volume of nitrous oxide be decomposed by a metal—for instance, sodium—then there remains, after cooling and total decomposition, a volume of nitrogen, exactly equal to that of the nitrous oxide taken; consequently the oxygen is, so to say, distributed between the atoms of nitrogen without producing an increase in the volume of the nitrogen.

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