CHAPTER VI

THE 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, NH_{4}Cl, on the passage of a series of electric sparks.[1] In ammonium chloride, HCl is combined with NH_{3}, consequently N with H_{3} 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, NH_{4}Cl, and ammonium sulphate, (NH_{4})_{2}SO_{4}. They are saline substances which are formed because ammonia, NH_{3}, combines with all acids, HX, forming ammonia salts, NH_{4}X. Sal-ammoniac, NH_{4}Cl, 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_, NH_{4}Cl, as from any other ammoniacal salt, by heating it with lime. Calcium hydroxide, CaH_{2}O_{2}, as an alkali takes up the acid and sets free the ammonia, forming calcium chloride, according to the equation 2NH_{4}Cl + CaH_{2}O_{2} = 2H_{2}O + CaCl_{2} + 2NH_{3}. In this reaction the ammonia is evolved as a gas.[6]

[1] The ammonia in the air, water, and soil proceeds from the decomposition of the nitrogenous substances of plants and animals, and also probably from the reduction of nitrates. Ammonia is always formed in the rusting of iron. Its formation in this case depends in all probability on the decomposition of water, and on the action of the hydrogen at the moment of its evolution on the nitric acid contained in the air (Cloez), or on the formation of ammonium nitrite, which takes place under many circumstances. The evolution of vapours of ammonia compounds is sometimes observed in the vicinity of volcanoes. At a red heat nitrogen combines directly with B Ca Mg, and with many other metals, and these compounds, when heated with a caustic alkali, or in the presence of water, give ammonia (_see_ Chapter XIV., Note 14, and Chapter XVII., Note 12). These are examples of the indirect combination of nitrogen with hydrogen.

[2] If a silent discharge or a series of electric sparks be passed through ammonia gas, it is decomposed into nitrogen and hydrogen. This is a phenomenon of dissociation; therefore, a series of sparks do not totally decompose the ammonia, but leave a certain portion undecomposed. One volume of nitrogen and three volumes of hydrogen are obtained from two volumes of ammonia decomposed. Ramsay and Young (1884) investigated the decomposition of NH_{3} under the action of heat, and showed that at 500°, 1-1/2 p.c. is decomposed, at 600° about 18 p.c., at 800° 65 p.c., but these results were hardly free from the influence of 'contact.' The _presence_ of free ammonia--that is, ammonia not combined with acids--in a gas or aqueous solution may be recognised by its characteristic smell. But many ammonia salts do not possess this smell. However, on the addition of an alkali (for instance, caustic lime, potash, or soda), they evolve ammonia gas, especially when heated. The presence of ammonia may be made visible by introducing a substance moistened with strong hydrochloric acid into its neighbourhood. A white cloud, or visible white vapour, then makes its appearance. This depends on the fact that both ammonia and hydrochloric acid are volatile, and on coming into contact with each other produce solid sal-ammoniac, NH_{4}Cl, which forms a cloud. This test is usually made by dipping a glass rod into hydrochloric acid, and holding it over the vessel from which the ammonia is evolved. With small amounts of ammonia this test is, however, untrustworthy, as the white vapour is scarcely observable. In this case it is best to take paper moistened with mercurous nitrate, HgNO_{3}. This paper turns black in the presence of ammonia, owing to the formation of a black compound of ammonia with mercurous oxide. The smallest traces of ammonia (for instance, in river water) may be detected by means of the so-called Nessler's reagent, containing a solution of mercuric chloride and potassium iodide, which forms a brown coloration or precipitate with the smallest quantities of ammonia. It will be useful here to give the thermochemical data (in thousands of units of heat, according to Thomsen), or the quantities of heat _evolved_ in the formation of ammonia and its compounds in quantities expressed by their formulæ. Thus, for instance, (N + H_{3}) 26·7 indicates that 14 grams of nitrogen in combining with 3 grams of hydrogen develop sufficient heat to raise the temperature of 26·7 kilograms of water 1°. (NH_{3} + nH_{2}O) 8·4 (heat of solution); (NH_{3},nH_{2}O + HCl,nH_{2}O) 12·3; (N + H_{4} + Cl) 90·6; (NH_{3} + HCl) 41·9.

[3] The same ammonia water is obtained, although in smaller quantities, in the dry distillation of plants and of coal, which consists of the remains of fossil plants. In all these cases the ammonia proceeds from the destruction of the complex nitrogenous substances occurring in plants and animals. The ammonia salts employed in the arts are prepared by this method.

[4] The technical methods for the preparation of ammonia water, and for the extraction of ammonia from it, are to a certain extent explained in the figures accompanying the text.

[5] Usually these crystals are sublimed by heating them in crucibles or pots, when the vapours of sal-ammoniac condense on the cold covers as a crust, in which form the salt comes into the market.

[6] On a small scale ammonia may be prepared in a glass flask by mixing equal parts by weight of slaked lime and finely-powdered sal-ammoniac, the neck of the flask being connected with an arrangement for drying the gas obtained. In this instance neither calcium chloride nor sulphuric acid can be used for drying the gas, since both these substances absorb ammonia, and therefore solid caustic potash, which is capable of retaining the water, is employed. The gas-conducting tube leading from the desiccating apparatus is introduced into a mercury bath, if dry gaseous ammonia be required, because water cannot be employed in collecting ammonia gas. Ammonia was first obtained in this dry state by Priestley, and its composition was investigated by Berthollet at the end of the last century. Oxide of lead mixed with sal-ammoniac (Isambert) evolves ammonia with still greater ease than lime. The cause and process of the decomposition are almost the same, 2PbO + 2NH_{4}Cl = Pb_{2}OCl_{2} + H_{2}O + 2NH_{3}. Lead oxychloride is (probably) formed.

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,3NH_{3} 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.

[7] [Illustration: FIG. 44.--Carré's apparatus. Described in text.]

This is evident from the fact that its absolute boiling point lies at about +130° (Chapter II., Note 29). It may therefore be liquefied by pressure alone at the ordinary, and even at much higher temperatures. The latent heat of evaporation of 17 parts by weight of ammonia equals 4,400 units of heat, and hence liquid ammonia may be employed for the production of cold. Strong aqueous solutions of ammonia, which in parting with their ammonia act in a similar manner, are not unfrequently employed for this purpose. Suppose a saturated solution of ammonia to be contained in a closed vessel furnished with a receiver. If the ammoniacal solution be heated, the ammonia, with a small quantity of water, will pass off from the solution, and in accumulating in the apparatus will produce a considerable pressure, and will therefore liquefy in the cooler portions of the receiver. Hence liquid ammonia will be obtained in the receiver. The heating of the vessel containing the aqueous solution of ammonia is then stopped. After having been heated it contains only water, or a solution poor in ammonia. When once it begins to cool the ammonia vapours commence dissolving in it, the space becomes rarefied, and a rapid vaporisation of the liquefied ammonia left in the receiver takes place. In evaporating in the receiver it will cause the temperature in it to fall considerably, and will itself pass into the aqueous solution. In the end, the same ammoniacal solution as originally taken is re-obtained. Thus, in this case, on heating the vessel the pressure increases by itself, and on cooling it diminishes, so that here heat directly replaces mechanical work. This is the principle of the simplest forms of _Carré's ice-making machines_, shown in fig. 44. C is a vessel made of boiler plates into which the saturated solution of ammonia is poured; m is a tube conducting the ammonia vapour to the receiver A. All parts of the apparatus should be hermetically joined together, and should be able to withstand a pressure reaching ten atmospheres. The apparatus should be freed from air, which would otherwise hinder the liquefaction of the ammonia. The process is carried on as follows:--The apparatus is first so inclined that any liquid remaining in A may flow into C. The vessel C is then placed upon a stove F, and heated until the thermometer _t_ indicates a temperature of 130° C. During this time the ammonia has been expelled from C, and has liquefied in A. In order to facilitate the liquefaction, the receiver A should be immersed in a tank of water R (_see_ the left-hand drawing in fig. 44). After about half an hour, when it may be supposed that the ammonia has been expelled, the fire is removed from under C, and this is now immersed in the tank of water R. The apparatus is represented in this position in the right-hand drawing of fig. 44. The liquefied ammonia then evaporates, and passes over into the water in C. This causes the temperature of A to fall considerably. The substance to be refrigerated is placed in a vessel G, in the cylindrical space inside the receiver A. The refrigeration is also kept on for about half an hour, and with an apparatus of ordinary dimensions (containing about two litres of ammonia solution), five kilograms of ice are produced by the consumption of one kilogram of coal. In industrial works more complicated types of Carré's machines are employed.

[8] Below 15° (according to Isambert), the compound AgCl,3NH_{3} is formed, and above 20° the compound 2AgCl,3NH_{3}. The tension of the ammonia evolved from the latter substance is equal to the atmospheric pressure at 68°, whilst for AgCl,3NH_{3} the pressures are equal at about 20°; consequently, at higher temperatures it is greater than the atmospheric pressure, whilst at lower temperatures the ammonia is absorbed and forms this compound. Consequently, all the phenomena of dissociation are here clearly to be observed. Joannis and Croisier (1894) investigated similar compounds with AgBr, AgI, AgCN and AgNO_{3}, and found that they all give definite compounds with NH_{3}, for instance AgBr,3NH_{3}, 2AgBr,3NH_{3} and AgBr,2NH_{3}; they are all colourless, solid substances which decompose under the atmospheric pressure at +3·5, +34° and +51°.

[9] The liquefaction of ammonia may be accomplished without an increase of pressure, by means of refrigeration alone, in a carefully prepared mixture of ice and calcium chloride (because the absolute boiling point of NH_{3} is high, about +130°). It may even take place in the severe frosts of a Russian winter. The application of liquid ammonia as a motive power for engines forms a problem which has to a certain extent been solved by the French engineer Tellier.

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]

[10] The combustion of ammonia in oxygen may be effected by the aid of platinum. A small quantity of an aqueous solution of ammonia, containing about 20 p.c. of the gas, is poured into a wide-necked beaker of about one litre capacity. A gas-conducting tube about 10 mm. in diameter, and supplying oxygen, is immersed in the aqueous solution of ammonia. But before introducing the gas an incandescent platinum spiral is placed in the beaker; the ammonia in the presence of the platinum is oxidised and burns, whilst the platinum wire becomes still more incandescent. The solution of ammonia is heated, and oxygen passed through the solution. The oxygen, as it bubbles off from the ammonia solution, carries with it a part of the ammonia, and this mixture explodes on coming into contact with the incandescent platinum. This is followed by a certain cooling effect, owing to the combustion ceasing, but after a short interval this is renewed, so that one feeble explosion follows after another. During the period of oxidation without explosion, white vapours of ammonium nitrite and red-brown vapours of oxides of nitrogen make their appearance, while during the explosion there is complete combustion and consequently water and nitrogen are formed.

[11] This may be verified by their densities. Nitrogen is 14 times denser than hydrogen, and ammonia is 8-1/2 times. If 3 volumes of hydrogen with 1 volume of nitrogen gave 4 volumes of ammonia, then these 4 volumes would weigh 17 times as much as 1 volume of hydrogen; consequently 1 volume of ammonia would be 4-1/4 times heavier than the same volume of hydrogen. But if these 4 volumes only give 2 volumes of ammonia, the latter will be 8-1/2 times as dense as hydrogen, which is found to be actually the case.

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, H_{2}SO_{4}, in absorbing ammonia, forms (on evaporating the solution) two salts, according to the relative quantities of ammonia and acid. One salt is formed from NH_{3} + H_{2}SO_{4}, and consequently has the composition NH_{5}SO_{4}, and the other is formed from 2NH_{3} + H_{2}SO_{4}, and its composition is therefore N_{2}H_{8}SO_{4}. 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, NH_{4}Cl, be compared with table salt, NaCl; and the ammonium hydrogen sulphate, NH_{4}HSO_{4}, with the sodium hydrogen sulphate, NaHSO_{4}; or ammonium nitrate, NH_{4}NO_{3}, with sodium nitrate, NaNO_{3}.[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 NH_{4}, which is called _ammonium_. If table salt be called 'sodium chloride,' then sal-ammoniac should be and is called 'ammonium chloride.'

[12] Aqueous solutions of ammonia are lighter than water, and at 15°, taking water at 4° = 10,000, their specific gravity, as dependent on _p_, or the percentage amount (by weight) of ammonia, is given by the expression _s_ = 9,992-42·5_p_ + 0·21_p_^2; for instance, with 10 p.c. _s_ = 9,587. If _t_ represents the temperature between the limits of +10° and +20°, then the expression (15-_t_)(1·5 + 0·14_p_) must be added to the formula for the specific gravity. Solutions containing more than 24 p.c. have not been sufficiently investigated in respect to the variation of their specific gravity. It is, however, easy to obtain more concentrated solutions, and at 0° solutions approaching NH_{3},H_{2}O (48·6 p.c. NH_{3}) in their composition, and of sp. gr. 0·85, may be prepared. But such solutions give up the bulk of their ammonia at the ordinary temperature, so that more than 24 p.c. NH_{3} is rarely contained in solution. Ammoniacal solutions containing a considerable amount of ammonia give ice-like crystals which seem to contain ammonia at temperatures far below 0° (for instance, an 8 p.c. solution at -14°, the strongest solutions at -48°). The whole of the ammonia may be expelled from a solution by heating, even at a comparatively low temperature; hence on heating aqueous solutions containing ammonia a very strong solution of ammonia is obtained in the distillate. Alcohol, ether, and many other liquids are also capable of dissolving ammonia. Solutions of ammonia, when exposed to the atmosphere, give off a part of their ammonia in accordance with the laws of the solution of gases in liquids, which we have already considered. But the ammoniacal solutions at the same time absorb carbonic anhydride from the air, and ammonium carbonate remains in the solution.

Solutions of ammonia are required both for laboratory and factory operations, and have therefore to be frequently prepared. For this purpose the arrangement shown in fig. 46 is employed in the laboratory. In works the same arrangement is used, only on a larger scale (with earthenware or metallic vessels). The gas is prepared in the retort, from whence it is led into the two-necked globe A, and then through a series of Woulfe's bottles, B, C, D, E. The impurities spurting over collect in A, and the gas is dissolved in B, but the solution soon becomes saturated, and a purer (washed) ammonia passes over into the following vessels, in which only a pure solution is obtained. The bent funnel tube in the retort preserves the apparatus from the possibility both of the pressure of the gas evolved in it becoming too great (when the gas escapes through it into the air), and also from the pressure incidentally falling too low (for instance, owing to a cooling effect, or from the reaction stopping). If this takes place, the air passes into the retort, otherwise the liquid from B would be drawn into A. The safety tubes in each Woulfe's bottle, open at both ends, and immersed in the liquid, serve for the same purpose. Without them, in case of an accidental stoppage in the evolution of so soluble a gas as ammonia, the solution would be sucked from one vessel to another--for instance, from E into D, &c. In order to clearly see the necessity for _safety tubes_ in a gas apparatus, it must be remembered that the _gaseous pressure_ in the interior of the arrangement must exceed the atmospheric pressure by the height of the sum of the columns of liquid through which the gas has to pass.

[13] The analogy between the ammonium and sodium salts might seem to be destroyed by the fact that the latter are formed from the alkali or oxide and an acid, with the separation of water, whilst the ammonium salts are directly formed from ammonia and an acid, without the separation of water; but the analogy is restored if we compare soda to ammonia water, and liken caustic soda to a compound of ammonia with water. Then the very preparation of ammonium salts from such a hydrate of ammonia will completely resemble the preparation of sodium salts from soda. We may cite as an example the action of hydrochloric acid on both substances.

NaHO + HCl = H_{2}O + NaCl Sodium hydroxide Hydrochloric acid Water Table salt

NH_{4}HO + HCl = H_{2}0 + NH_{4}Cl Ammonium hydroxide Hydrochloric acid Water Sal-ammoniac

Just as in soda the hydroxyl or aqueous radicle OH is replaced by chlorine, so it is in ammonia hydrate.

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, NH_{4}Cl + HgNa = NaCl + HgNH_{4}. 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, NH_{4} = NH_{3} + 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, NH_{4}OH,[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 NH_{4}OH = NH_{3} + H_{2}O.

[14] Weyl (1864) by subjecting sodium to the action of ammonia at the ordinary temperature and under considerable pressures, obtained a liquid, which was subsequently investigated by Joannis (1889), who confirmed the results obtained by Weyl. At 0° and the atmospheric pressure the composition of this substance is Na + 5·3NH_{3}. The removal (at 0°) of ammonia from the liquid gives a solid copper-red body having the composition NH_{3}Na. The determination of the molecular weight of this substance by the fall of the tension of liquid ammonia gave N_{2}H_{6}Na_{2}. It is, therefore, free ammonium in which one H is replaced by Na. The compound with potassium, obtained under the same conditions, proved to have an analogous composition. By the decomposition of NH_{3}Na at the ordinary temperature, Joannis (1891) obtained hydrogen and sodium-amide NH_{2}Na in small colourless crystals which were soluble in water. The addition of liquid ammonia to metallic sodium and a saturated solution of sodium chloride, gives NH_{2}Na_{2}Cl, and this substance is sal-ammoniac, in which H_{2} is replaced by Na_{2}.

If pure oxygen be passed through a solution of these compounds in ammonia at a temperature of about -50°, it is seen that the gas is rapidly absorbed. The liquid gradually loses its dark red colour and becomes lighter, and when it has become quite colourless a gelatinous precipitate is thrown down. After the removal of the ammonia, this precipitate dissolves easily in water with a considerable evolution of heat, but without giving off any gaseous products. The composition of the sodium compound thus obtained is NH_{2}Na_{2}HO, which shows that it is a hydrate of bisodium-ammonium. Thus, although free ammonium has not been obtained, still a sodium substitution product of it is known which corresponds to it as a salt to a hydrate. Ammonium amalgam was originally obtained in exactly the same way as sodium amalgam (Davy); namely, a piece of sal-ammoniac was taken, and moistened with water (in order to render it a conductor of electricity). A cavity was made in it, into which mercury was poured, and it was laid on a sheet of platinum connected with the positive pole of a galvanic battery, while the negative pole was put into connection with the mercury. On passing a current the mercury increased considerably in volume, and became plastic, whilst preserving its metallic appearance, just as would be the case were the sal-ammoniac replaced by a lump of a sodium salt or of many other metals. In the analogous decomposition of common metallic salts, the metal contained in a given salt separates out at the negative pole, immersed in mercury, by which the metal is dissolved. A similar phenomenon is observed in the case of sal-ammoniac; the elements of ammonium, NH_{4}, in this case are also collected in the mercury, and are retained by it for a certain time.

[15] We may mention, however, that under particular conditions hydrogen is also capable of forming an amalgam resembling the amalgam of ammonium. If an amalgam of zinc be shaken up with an aqueous solution of platinum chloride, without access of air, then a spongy mass is formed which easily decomposes, with the evolution of hydrogen.

[16] We saw above that the solubility of ammonia in water at low temperatures attains to the molecular ratio NH_{3} + H_{2}O, in which these substances are contained in caustic ammonia, and perhaps it may be possible at exceedingly low temperatures to obtain ammonium hydroxide, NH_{4}HO, in a solid form. Regarding solutions as dissociated definite compounds, we should see a confirmation of this view in the property shown by ammonia of being extremely soluble in water, and in so doing of approaching to the limit NH_{4}HO.

[17] In confirmation of the truth of this conclusion we may cite the remarkable fact that there exist, in a free state and as comparatively stable compounds, a series of alkaline hydroxides, NR_{4}HO, which are perfectly analogous to ammonium hydroxide, and present a striking resemblance to it and to sodium hydroxide, with the only difference that the hydrogen in NH_{4}HO is replaced by complex groups, R = CH_{3}, C_{2}H_{5}, &c., for instance N(CH_{3})_{4}HO. Details will be found in organic chemistry.

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]

[18] The fact that ammoniacal salts are decomposed when ignited, and not simply sublimed, may be proved by a direct experiment with sal-ammoniac, NH_{4}Cl, which in a state of vapour is decomposed into ammonia, NH_{3}, and hydrochloric acid, HCl, as will be explained in the following chapter. The readiness with which ammonium salts decompose is seen from the fact that a solution of ammonium oxalate is decomposed with the evolution of ammonia even at -1°. Dilute solutions of ammonium salts, when boiled give aqueous vapour having an alkaline reaction, owing to the presence of free ammonia given off from the salt.

Ammonia is not only capable of combining with acids, but also with many salts, as was seen from its forming definite compounds, AgCl,3NH_{3} and 2AgCl,3NH_{3}, 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.

[19] Isambert studied the dissociation of ammoniacal compounds, as we have seen in Note 8, and showed that at low temperatures many salts are able to combine with a still greater amount of ammonia, which proves an entire analogy with hydrates; and as in this case it is easy to isolate the definite compounds, and as the least possible tension of ammonia is greater than that of water, therefore the ammoniacal compounds present a great and peculiar interest, as a means for explaining the nature of aqueous solutions and as a confirmation of the hypothesis of the formation of definite compounds in them; for these reasons we shall frequently refer to these compounds in the further exposition of this work.

[20] Chapter V., Note 2.

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, CuSO_{4}. As we saw in Chapter I., it gives with water blue crystals, CuSO_{4},5H_{2}O; but it also absorbs ammonia in the same molecular proportion, forming a blue substance, CuSO_{4},5NH_{3}, 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, NH_{2}Na (according to the equation NH_{3} + Na = NH_{2}Na + 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 NH_{2} remains, the resultant substances NH_{2}R are called _amides_, whilst the substitution products, NHR_{2}, in which only NH remains, are called _imides_,[20 bis] and those in which none of the ammoniacal hydrogen remains, NR_{3}, are known as _nitrides_. Free amidogen, N_{2}H_{4}, 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.

[20 bis] Imide, NH, has not been obtained in a free state, but its hydrochloric acid salt, NHHCl, has apparently been obtained (1890) by Maumené by igniting the double bichloride of platinum and ammonium chloride, PtCl_{2}NH_{4}Cl = Pt + 2HCl + NHHCl. It is soluble in water, and crystallises from its solution in hexagonal rhombic prisms. It gives a double salt with FeCl_{3} of the composition FeCl_{3}3NHHCl. The salt NHHCl is similar (isomeric) with the first possible product of the metalepsis of ammonia, NH_{2}Cl, although it does not resemble it in any of its properties.

[21] Free _amidogen_ or _hydrazine_, N_{2}H_{4}, or 2NH_{2}, was prepared by Curtius (1887) by means of ethyl diazoacetate, or triazoacetic acid. Curtius and Jay (1889) showed that triazoacetic acid, CHN_{2}.COOH (the formula should be tripled), when heated with water or a mineral acid, gives (quantitatively) oxalic acid and amidogen (hydrazine), CHN_{2}.COOH + 2H_{2}O = C_{2}O_{2}(OH)_{2} + N_{2}H_{4}--_i.e._ (empirically), the oxygen of the water replaces the nitrogen of the azoacetic acid. The amidogen is thus obtained in the form of a salt. With acids, amidogen forms very stable salts of the two types N_{2}H_{4}HX and N_[2]H_{4}H_{2}X_{2}, as, for example, with HCl, H_{2}SO_{4}, &c. These salts are easily crystallised; in acid solutions they act as powerful reducing agents, evolving nitrogen; when ignited they are decomposed into ammoniacal salts, nitrogen, and hydrogen; with nitrites they evolve nitrogen. The sulphate N_{2}H_{4},H_{2}SO_{4} is sparingly soluble in cold water (3 parts in 100 of water), but is very soluble in hot water; its specific gravity is 1·378, it fuses at 254° with decomposition. The hydrochloride N_{2}H_{4},2HCl crystallises in octahedra, is very soluble in water, but not in alcohol; it fuses at 198°, evolving hydrogen chloride and forming the salt N_{2}H_{4}HCl; when rapidly heated it decomposes with an explosion; with platinic chloride it immediately evolves nitrogen, forming platinous chloride. By the action of alkalis the salts N_{2}H_{4},2HX give _hydrate of amidogen_, N_{2}H_{4},H_{2}O, which is a fuming liquid (specific gravity 1·03), boiling at 119°, almost without odour, and whose aqueous solution corrodes glass and india-rubber, has an alkaline taste and poisonous properties. The reducing capacities of the hydrate are clearly seen from the fact that it reduces the metals platinum and silver from their solutions. With mercuric oxide it explodes. It reacts directly with the aldehydes RO, forming N_{2}R_{2} and water; for example, with benzaldehydes it gives the very stable insoluble _benzalazine_ (C_{6}H_{5}CHN)_{2} of a yellow colour. We may add that hydrazine often forms double salts; for example, MgSO_{4}N_{2}H_{4}H_{2}SO_{4} or KClN_{2}H_{4}HCl, and that it is also formed by the action of nitrous acid upon aldehyde-ammonia. The products of the substitution of the hydrogen in hydrazine by hydrocarbon groups R (R = CH_{3}, C_{2}H_{5}, C_{6}H_{5}, &c.) were obtained before hydrazine itself; for example, NHRNH_{2}, NR_{2}NH_{2}, and (NRH)_{2}.

The heat of solution of the sulphuric acid salt (1 part in 200 and 300 parts of water at 10°·8) is equal to -8·7 C. According to Berthelot and Matigon (1892), the heat of neutralisation of hydrazine by sulphuric acid is +5·5 C and by hydrochloric acid +5·2 C. Thus hydrazine is a very feeble base, for its heat of saturation is not only lower than that of ammonia (+12·4 C. for HCl), but even below that of hydroxylamine (+9·3 C.) The heat of formation from the elements of hydrated hydrazine -9·5 C was deduced from the heat of combustion, determined by burning N_{2}H_{4}H_{2}SO_{4} in a calorimetric bomb, +127·7 C. Thus hydrazine is an endothermal compound; its passage into ammonia by the combination of hydrogen is accompanied by the evolution of 51·5 C. In the presence of an acid these figures were greater by +14·4 C. Hence the direct converse passage from ammonia into hydrazine is impossible. As regards the passage of hydroxylamine into hydrazine, it would be accompanied by the evolution of heat (+21·5 C.) in an aqueous solution.

Amidogen must be regarded as a compound which stands to ammonia in the same relation as hydrogen peroxide stands to water. Water, H(OH), gives, according to the law of substitution, as was clearly to be expected, (OH)(OH)--that is, peroxide of hydrogen is the free radicle of water (hydroxyl). So also ammonia, H(NH_{2}), forms hydrazine, (NH_{2})(NH_{2})--that is, the free radicle of ammonia, NH_{2}, or amidogen. In the case of phosphorus a similar substance, as we shall afterwards see, has long been known under the name of liquid phosphuretted hydrogen, P_{2}H_{4}.

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-H_{2}O = MX. If aqueous ammonia, NH_{4}OH, 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. Aqueous ammonia. Ammonium salt. Basic hydrate. In solution In solution In solution As precipitate

[21 bis] In practice, the applications of ammonia are very varied. The use of ammonia as a stimulant, in the forms of the so-called 'smelling salts' or of spirits of hartshorn, in cases of faintness, &c., is known to everyone. The volatile carbonate of ammonium, or a mixture of an ammonium salt with an alkali, is also employed for this purpose. Ammonia also produces a well-known stimulating effect when rubbed on the skin, for which reason it is sometimes employed for external applications. Thus, for instance, the well-known volatile salve is prepared from any liquid oil shaken up with a solution of ammonia. A portion of the oil is thus transformed into a soapy substance. The solubility of greasy substances in ammonia, which proceeds from the formation both of emulsions and soaps, explains its use in extracting grease spots. It is also employed as an external application for stings from insects, and for bites from poisonous snakes, and in general in medicine. It is also remarkable that in cases of drunkenness a few drops of ammonia in water taken internally rapidly renders a person sober. A large quantity of ammonia is used in dyeing, either for the solution of certain dyes--for example, carmine--or for changing the tints of others, or else for neutralising the action of acids. It is also employed in the manufacture of artificial pearls. For this purpose the small scales of a peculiar small fish are mixed with ammonia, and the liquid so obtained is blown into small hollow glass beads shaped like pearls.

In nature and the arts, however, ammonium salts, and not free ammonia, are most frequently employed. In this form a portion of that _nitrogen_ which is necessary for the formation of albuminous substances is _supplied to plants_. Owing to this, a large quantity of ammonium sulphate is now employed as a fertilising substance. But the same effect may be produced by nitre, or by animal refuse, which in decomposing gives ammonia. For this reason, an ammoniacal (hydrogen) compound may be introduced into the soil in the spring which will be converted into a nitrate (oxygen salt) in the summer.

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]

[22] As certain basic hydrates form peculiar compounds with ammonia, in some cases it happens that the first portions of ammonia added to a solution of a salt produce a precipitate, whilst the addition of a fresh quantity of ammonia dissolves this precipitate if the ammoniacal compound of the base be soluble in water. This, for example, takes place with the copper salts. But alumina does not dissolve under these circumstances.

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, H_{2}O, NH_{3}, 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, H_{2}O, may be divided in two ways, because it contains 3 atoms: into H and (HO) on the one hand, and into H_{2} and O on the other. Consequently, being given RH, its substitution products will be R(HO) according to the first form, and R_{2}O according to the second; being given RH_{2}, its corresponding substitution products will be RH(OH), R(OH)_{2}, RO, (RH)_{2}O, &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 NH_{4}Cl--correspond with hydroxides of the alkalis Na(OH) or NH_{4}(OH). In hydrocarbons--for instance, C_{2}H_{6}--the hydrogen is replaceable by chlorine and by hydroxyl. Thus ordinary alcohol is C_{2}H_{6}, in which one atom of H is replaced by (OH); that is, C_{2}H_{5}(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 H_{2}O_{2}. The other form of substitution--namely, that of O in the place of H_{2}--is also a common chemical phenomenon. Thus alcohol, C_{2}H_{6}O, or C_{2}H_{5}(OH), when oxidising in the air, gives acetic acid, C_{2}H_{4}O_{2}, or C_{2}H_{3}O(OH), in which H_{2} is replaced by O.

[23] When the element chlorine, as we shall afterwards more fully learn, replaces the element hydrogen, the reaction by which such an exchange is accomplished proceeds as a substitution, AH + Cl_{2} = ACl + HCl, so that two substances, AH and chlorine, react on each other, and two substances, ACl and HCl, are formed; and further, two molecules react on each other, and two others are formed. The reaction proceeds very easily, but the substitution of one element, _A_, by another, _X_, does not always proceed with such ease, clearness, or simplicity. The substitution between oxygen and hydrogen is very rarely accomplished by the reaction of the free elements, but the substitution between these elements, one for another, forms the most common case of oxidation and reduction. In speaking of the law of substitution, I have in view the substitution of the elements one by another, and not the direct reaction of substitution. The law of substitution determines the cycle of the combinations of a given element, if a few of its compounds (for instance, the hydrogen compounds) be known. A development of the conceptions of the law of substitution may be found in my lecture given at the Royal Institution in London, 1889.

[24] If hydrogen peroxide be taken as a starting point, then still higher forms of oxidation than those corresponding with water should be looked for. They should possess the properties of hydrogen peroxide, especially that of parting with their oxygen with extreme ease (even by contact). Such compounds are known. Pernitric, persulphuric, and similar acids present these properties, as we shall see in describing them.

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, NH_{3}, or aqueous ammonia, NH_{4}(OH), by substituting their hydrogen by hydroxyl, or H_{2} by oxygen. And such is the case. The two extreme cases of such substitution will be as follows: (1) One atom of H in NH_{3} is substituted by (OH), and NH_{2}(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, NH_{3}(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, NH_{4}(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 NO_{2}(OH), or NHO_{3}, as we find to be really the case, for NHO_{3} 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 NHO_{2}, equal to N(OH)_{3}-H_{2}O, and nitrous anhydride, N_{2}O_{3} = 2N(OH)_{3}-3H_{2}O. 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, HN_{3}. 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.

[25] The compound of hydroxylamine with hydrochloric acid has the composition NH_{2}(OH)HCl = NH_{4}ClO--that is, it is as it were oxidised sal-ammoniac. It was prepared by Lossen in 1865 by the action of tin and hydrochloric acid in the presence of water on a substance called ethyl nitrate, in which case the hydrogen liberated from the hydrochloric acid by the tin acts upon the elements of nitric acid--

C_{2}H_{5}·NO_{3} + 6H + HCl = NH_{4}OCl + H_{2}O + C_{2}H_{5}·OH Ethyl nitrate Hydrogen Hydroxylamine Water Alcohol from + HCl HCl and Sn

Thus in this case the nitric acid is deoxidised, not directly into nitrogen, but into hydroxylamine. Hydroxylamine is also formed by passing nitric oxide, NO, into a mixture of tin and hydrochloric acid--that is, by the action of the hydrogen evolved on the nitric oxide, NO + 3H + HCl = NH_{4}OCl--and in many other cases. According to Lossen's method, a mixture of 30 parts of ethyl nitrate, 120 parts of tin, and 40 parts of a solution of hydrochloric acid of sp. gr. 1·06 are taken. After a certain time the reaction commences spontaneously. When the reaction has ceased the tin is separated by means of hydrogen sulphide, the solution is evaporated, and a large amount of sal-ammoniac is thus obtained (owing to the further action of hydrogen on the hydroxylamine compound, the hydrogen taking up oxygen from it and forming water); a solution ultimately remains containing the hydroxylamine salt; this salt is dissolved in anhydrous alcohol and purified by the addition of platinum chloride, which precipitates any ammonium salt still remaining in the solution. After concentrating the alcoholic solution the hydroxylamine hydrochloride separates in crystals. This substance melts at about 150°, and in so doing decomposes into nitrogen, hydrogen chloride, water, and sal-ammoniac. A sulphuric acid compound of hydroxylamine may be obtained by mixing a solution of the above salt with sulphuric acid. The sulphate is also soluble in water like the hydrochloride; this shows that hydroxylamine, like ammonia itself, forms a series of salts in which one acid may be substituted for another. It might he expected that by mixing a strong solution of a hydroxylamine salt with a solution of a caustic alkali hydroxylamine itself would be liberated, just as an ammonia salt under these circumstances evolves ammonia; but the liberated hydroxylamine is immediately decomposed with the formation of nitrogen and ammonia (and probably nitrous oxide), 3NH_{3}O = NH_{3} + 3H_{2}O + N_{2}. Dilute solutions give the same reaction, although very slowly, but by decomposing a solution of the sulphate with barium hydroxide a certain amount of hydroxylamine is obtained in solution (it is partly decomposed). Hydroxylamine in aqueous solution, like ammonia, precipitates basic hydrates, and it deoxidises the oxides of copper, silver, and other metals. Free hydroxylamine was obtained by Lobry de Bruyn (1891). It is a solid, colourless, crystalline substance, without odour, which does not melt below 27°. It has the property of dissolving metallic salts; for instance, sodium chloride. Hydroxylamine, when rapidly heated with platinum, decomposes with a flash and the formation of a yellow flame. It is almost insoluble in ordinary solvents like chloroform, benzine, acetic ether, and carbon bisulphide. Its aqueous solutions are tolerably stable, contain up to 60 per cent. (sp. gr. 1·15 at 20°), and may be kept for many weeks without undergoing any change. Lobry de Bruyn used the hydrochloric salt to prepare pure hydroxylamine. The salt was first treated with sodium methylate (CH_{3}NaO), and then methyl alcohol was added to the mixture. The precipitated sodium chloride was separated from the solution by filtration. (The methyl alcohol is added to prevent the precipitated chloride of sodium from coating the insoluble hydrochloric salt of hydroxylamine.) The methyl alcohol was driven off under a pressure 150-200 mm., and after extracting a further portion of methyl alcohol by ether and several fractional distillations, a solution was obtained containing 70 per cent. of free hydroxylamine, 8 per cent. water, 9·9 per cent. chloride of sodium, and 12·1 per cent. of the hydrochloric salt of hydroxylamine. Pure free hydroxylamine, NH_{3}O, is obtained by distilling under a pressure of 60 mm.; it then boils at 70°, and solidifies in a condenser cooled to 0° in the form of long needles. It melts at 33°, boils at 58° under a pressure of 22 mm., and has a sp. gr. of about 1·235 (Brühl). Under the action of NaHO it gives NH_{3} and NHO_{2} or N_{2}O, and forms nitric acid (Kolotoff, 1893) under the action of oxidising agents. Hydroxylamine is obtained in a great number of cases, for instance by the action of tin on dilute nitric acid, and also by the action of zinc on ethyl nitrate and dilute hydrochloric acid, &c. The relation between hydroxylamine, NH_{2}(OH), and nitrous acid, NO(OH), which is so clear in the sense of the law of substitutions, becomes a reality in those cases when reducing agents act on salts of nitrous acid. Thus Raschig (1888) proposed the following method for the preparation of the hydroxylamine sulphate. A mixture of strong solutions of potassium nitrite, KNO_{2}, and hydroxide, KHO, in molecular proportions, is prepared and cooled. An excess of sulphurous anhydride is then passed into the mixture, and the solution boiled for a long time. A mixture of the sulphates of potassium and hydroxylamine is thus obtained: KNO_{2} + KHO + 2SO_{2} + 2H_{2}O = NH_{2}(OH),H_{2}SO_{4} + K_{2}SO_{4}. The salts may be separated from each other by crystallisation.

[25 bis] In order to illustrate the application of the law of substitution to a given case, and to show the connection between ammonia and the oxides of nitrogen, let us consider the possible products of an oxygen and hydroxyl substitution in caustic ammonia, NH_{4}(OH). It is evident that the substitution of H by OH can give: (1) NH_{3}(OH)_{2}; (2) NH_{2}(OH)_{3}; (3) NH(OH)_{4}; and (4) N(OH)_{5}. They should all, like caustic ammonia itself, easily part with water and form products (hydroxylic) of the oxidation of ammonia. The first of them is the hydrate of hydroxylamine, NH_{2}(OH) + H_{2}O; the second, NH(OH)_{2} + H_{2}O (and also the substance NH(OH)_{4} or NH_{3}O_{2}), containing, as it does, both hydrogen and oxygen, is able to part with all its hydrogen in the form of water (which could not be done by the first product, since it contained too little oxygen), forming, as the ultimate product, 2NH_{2}(OH)_{3}-5H_{2}O = N_{2}O--that is, it corresponds with nitrous oxide, or the lower degree of the oxidation of nitrogen. So, also, nitrous anhydride corresponds with the third of the above products, 2NH(OH)_{4}-5H_{2}O = N_{2}O_{3}, and nitric anhydride with the fourth, 2N(OH)_{5}-5H_{2}O = N_{2}O_{5}. As, in these three equations, two molecules of the substitution products (-5H_{2}O) are taken, it is also possible to combine two different products in one equation. For instance, the third and fourth products: NH(OH)_{4} + N(OH)_{5}-5H_{2}O corresponds to N_{2}O_{4} or 2NO_{2}, that is, to peroxide of nitrogen. Thus all the five (see later) oxides of nitrogen, N_{2}O, NO, N_{2}O_{3}, NO_{2}, and N_{2}O_{5}, may be deduced from ammonia. The above may be expressed in a general form by the equation (it should be remarked that the composition of all the substitution products of caustic ammonia may be expressed by NH_{3}O{5-_a_}, where _a_ varies between 0 and 4):

NH_{5}O_{5 - _a_} + NH_{5}O{5 - _b_} - 5H_{2}O = N_{2}O_{5 - (_a_ + _b_)},

where _a_ + _b_ can evidently be not greater than 5; when _a_ + _b_ = 5 we have N_{2}--nitrogen, when = 4 we have N_{2}O nitrous oxide; when _a_ + _b_ = 3 we have N_{2}O_{2} or NO--nitric oxide, and so on to N_{2}O_{5}, when _a_ + _b_ = 0. Besides which it is evident that intermediate products may correspond with (and hence also break up into) different starting points; for instance, N_{2}O is obtained when _a_ + _b_ = 2, and this may occur either when _a_ = 0 (nitric acid), and _b_ = 2 (hydroxylamine), or when _a_ = _b_ = 1 (the third of the above substitution products).

[26] Nitric acid corresponds with the anhydride N_{2}O_{5}, which will afterwards be described, but which must be regarded as the highest saline oxide of nitrogen, just as Na_{2}O (and the hydroxide NaHO) in the case of sodium, although sodium forms a peroxide possessing the property of parting with its oxygen with the same ease as hydrogen peroxide, if not on heating, at all events in reactions--for instance, with acids. So also nitric acid has its corresponding peroxide, which may be called pernitric acid. Its composition is not well known--probably NHO_{4}--so that its corresponding anhydride would be N_{2}O_{7}. It is formed by the action of a silent discharge on a mixture of nitrogen and oxygen, so that a portion of its oxygen is in a state similar to that in ozone. The instability of this substance (obtained by Hautefeuille, Chappuis, and Berthelot), which easily splits up with the formation of nitric peroxide, and its resemblance to persulphuric acid, which we shall afterwards describe, will permit our passing over the consideration of the little that is further known concerning it.

[27] Phosphorus (Chapter XIX.) gives the hydride PH_{3}, corresponding with ammonia, NH_{3}, and forms phosphorous acid, PH_{3}O_{3}, which is analogous to nitrous acid, just as phosphoric acid is to nitric acid; but phosphoric (or, better, orthophosphoric) acid, PH_{3}O_{4}, is able to lose water and give pyro-and meta-phosphoric acids. The latter is equal to the ortho-acid minus water = PHO_{3}, and therefore nitric acid, NHO_{3}, is really meta-nitric acid. So also nitrous acid, HNO_{2}, is meta-nitrous (anhydrous) acid, and thus the ortho-acid is NH_{3}O_{3} = N(OH)_{3}. Hence for nitric acid we should expect to find, besides the ordinary or meta-nitric acid, HNO_{3} (= 1/2N_{2}O_{3},H_{2}O), and ortho-nitric acid, H_{3}NO_{4} (= 1/2N_{2}O_{3},3H_{2}O), an intermediate pyro-nitric acid, N_{2}H_{4}O_{7}, corresponding to pyrophosphoric acid, P_{2}H_{4}O_{7}. We shall see (for instance, in Chapter XVI., Note 21) that in nitric acid there is indeed an inclination of the ordinary salts (of the meta-acid), MNO_{3}, to combine with bases M_{2}O, and to approximate to the composition of ortho-compounds which are equal to meta-compound and bases (MNO_{3} + M_{2}O = M_{3}NO_{4}).

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, NHO_{3} + 8H = NH_{3} + 3H_{2}O.

2. In 1890 Curtius in Germany obtained a gaseous substance of the composition HN_{3} (hydrogen trinitride), having the distinctive properties of an acid, and giving, like hydrochloric acid, salts; for example, a sodium salt, NaN_{3}; ammonium salt, NH_{4}N_{3} = N_{4}H_{4}; barium salt, Ba(N_{3})_{2}, &c., which he therefore named hydronitrous acid, _HN__{3}.[28 bis] The extraordinary composition of the compound (ammonia, NH_{3}, contains one N atom and three H atoms; in HN_{3}, 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 HN_{3} 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, HN_{3}, and the long known ammonia, NH_{3}, and nitric acid, HNO_{3}, 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, NH_{4}OH, there should correspond, according to the law of substitution, an ortho-nitric acid (_see_ Note 27), H_{3}NO_{4} = NO(OH)_{3}, which equals NH_{4}(OH) with the substitution in it of (_a_) two atoms of hydrogen by oxygen (O--H_{2}) 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, H_{2}NH_{4}NO_{4}; bi-substituted, H(NH_{4})_{2}NO_{4}; and tri-substituted, (NH_{4})_{3}NO_{4}. 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 H_{2}NH_{4}NO_{4}-4H_{2}O = N_{2}O--nitrous oxide, and from the second H(NH_{4})_{2}NO4-4H_{2}O = HN_{3}--hydronitrous acid, and from the third (NH_{4})_{3}NO-4H_{2}O = N_{4}H_{4}--the ammonium salt of the same acid. The composition of HN_{3} should be thus understood, whilst its acid properties are explained by the fact that the water (4H_{2}O) from H(NH_{4})_{2}N_O{4} 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 HN_{3} may be regarded as HCN where C has been replaced by N_{2}). Moreover, besides explaining the composition and acid properties of HN_{3}, the above method gives the possibility of foretelling the closeness of the bond between hydronitrous acid and nitrous oxide, for N_{2}O + NH_{3} = HN_{3} + H_{2}O. 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, NH_{2}Na (obtained by heating Na in a current of NH_{3}), and acting upon it (when heated) with nitrous oxide, N_{2}O, when 2NH_{2}Na + N_{2}O = NaN_{3} + NaHO + NH_{3}. The resultant salt, NaN_{3}, gives hydronitrous acid when acted upon by sulphuric acid, NaN_{3} + H_{2}SO_{4} = NaHSO_{4} +HN_{3}. The latter gives, with the corresponding solutions of their salts, the insoluble (and easily explosive) salts of silver, AgN_{3} (insoluble, like AgCl or AgCN), and lead, Pb(N_{3})_{2}.

[28] The formation of ammonia is observed in many cases of oxidation by means of nitric acid. This substance is even formed in the action of nitric acid on tin, especially if dilute acid be employed in the cold. A still more considerable amount of ammonia is obtained if, in the action of nitric acid, there are conditions directly tending to the evolution of hydrogen, which then reduces the acid to ammonia; for instance, in the action of zinc on a mixture of nitric and sulphuric acids.

[28 bis] Curtius started with benzoylhydrazine, C_{6}H_{5}CONHNH_{2} (hydrazine, see Note 20 bis). (This substance is obtained by the action of hydrated hydrazine on the compound ether of benzoic acid). Benzoylhydrazine under the action of nitrous acid gives benzoylazoimide and water:

C_{6}H_{5}CONHNH_{2} + NO_{2}H = C_{6}H_{5}CON_{3} + 2H_{2}O.

Benzoylazoimide when treated with sodium alcoholate gives the sodium salt of hydronitrous acid:

C_{6}H_{5}CON_{3} + C_{2}H_{3}ONa = C_{6}H_{5}O_{2}C_{2}H_{3} + NaN_{3}.

The addition of ether to the resultant solution precipitates the NaN_{3}, and this salt when treated with sulphuric acid gives gaseous hydronitrous acid, HN_{3}. It has an acrid smell, and is easily soluble in water. The aqueous solution exhibits a strongly acid reaction. Metals dissolve in this solution and give the corresponding salts. With hydronitrous acid gaseous ammonia forms a white cloud, consisting of the salt of ammonium, NH_{4}N_{3}. This salt separates out from an alcoholic solution in the form of white lustrous scales. The salts of hydronitrous acid are obtained by a reaction of substitution with the sodium or ammonium salts. In this manner Curtius obtained and studied the salts of silver (AgN_{3}), mercury (HgN_{3}), lead (PbN_{6}), barium (BaN_{6}). With hydrazine, N_{2}H_{4}, hydronitrous acid forms saline compounds in the composition of which there are one or two particles of N_{3}H per one particle of hydrazine; thus N_{5}H_{5} and N_{8}H_{6}. The first was obtained in an almost pure form. It crystallises from an aqueous solution in dense, volatile, lustrous prisms (up to 1 in. long), which fuse at 50°, and deliquesce in the air; from a solution in boiling alcohol it separates out in bright crystalline plates. This salt, N_{5}H_{5}, has the same empirical composition, NH, as the ammonium salt of hydronitrous acid, N_{4}H_{4}, and imide; but their molecules and structure are different. Curtius also obtained (1893) hydronitrous acid by passing the vapour of N_{2}O_{5} (evolved by the action of HNO_{3} on As_{2}O_{3}) into a solution of hydrazine, N_{2}H_{4}. Similarly Angeli, by acting upon a saturated solution of silver nitrite with a strong solution of hydrazine, obtained the explosive AgN_{3} in the form of a precipitate, and this reaction, which is based upon the equation N_{2}H_{4} + NHO_{2} = HN_{3} + 2H_{2}O, proceeds so easily that it forms an experiment for the lecture table. A thermal investigation of hydronitrous acid by Berthelot and Matignon gave the following figures for the heat of solution of the ammonium salt N_{3}HNH_{3} (1 grm. in 100 parts of water)-708 C., and for the heat of neutralisation by barium hydrate +10·0 C., and by ammonia +8·2 C. The heat of combustion of N_{4}H_{4} (+163·8 C. at a constant vol.) gives the heat of formation of the salt N_{4}H_{4} (solid) as -25·3 C. and (solution)-32·3 C.; this explains the explosive nature of this compound. In its heat of formation from the elements N_{3}H =-62·6 C., this compound differs from all the hydrogen compounds of nitrogen in having a maximum absorption of heat, which explains its instability.

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,[29] nitrous and nitric oxides, peroxide of nitrogen, and nitric acid, NHO_{3}, 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 NO_{2} form N_{2}O_{3}, _and the intermediate oxides may, in splitting up, give a higher and lower oxide_. So N_{2}O_{4} gives N_{2}O_{3} and N_{2}O_{5}, or, in the presence of water, their hydrates.

[29] According to the thermochemical determinations of Favre, Thomsen, and more especially of Berthelot, it follows that, in the formation of such quantities of the oxides of nitrogen as express their formulæ, if gaseous nitrogen and oxygen be taken as the starting points, and if the compounds formed be also gaseous, the following amounts of heat, expressed in thousands of heat units, are _absorbed_ (hence a minus sign):--

N_{2}O N_{2}O_{2} N_{2}O_{3} N_{2}O_{4} N_{2}O_{5} -21 -43 -22 -5 -1 -22 +21 +17 +4

The difference is given in the lower line. For example, if N_{2}, or 28 grams of nitrogen, combine with O--that is, with 16 grams of oxygen--then 21,000 units of heat are absorbed, that is, sufficient heat to raise 21,000 grams of water through 1°. Naturally, direct observations are impossible in this case; but if charcoal, phosphorus, or similar substances are burnt both in nitrous oxide and in oxygen, and the heat evolved is observed in both cases, then the difference (more heat will be evolved in burning in nitrous oxide) gives the figures required. If N_{2}O_{2}, by combining with O_{2}, gives N_{2}O_{4}, then, as is seen from the table, heat should be developed, namely, 38,000 units of heat, or NO + O = 19,000 units of heat. The differences given in the table show that the maximum absorption of heat corresponds with nitric oxide, and that the higher oxides are formed from it with evolution of heat. If liquid nitric acid, NHO_{3}, were decomposed into N + O_{3} + H, then 41,000 heat units would be required; that is, an evolution of heat takes place in its formation from the gases. It should be observed that the formation of ammonia, NH_{3}, from the gases N + H_{3} evolves 12·2 thousand heat units.

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_, NHO_{3}, 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, KNO_{3}, is common nitre, and the sodium salt, NaNO_{3}, 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, NaHSO_{4}, or a normal salt, Na_{2}SO_{4}, whilst nitric acid is formed from the nitre and is volatilised. The decomposition is expressed by the equations: (1) NaNO_{3} + H_{2}SO_{4} = HNO_{3} + NaHSO_{4}, if the acid salt be formed, and (2) 2NaNO_{3} + H_{2}SO_{4} = Na_{2}SO_{4} + 2HNO_{3}, 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 NaHSO_{4} itself acts like an acid (its hydrogen being replaceable as in acids), according to the equation NaNO_{3} + NaHSO_{4} = Na_{2}SO_{4} + HNO_{3}.

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.

[29 bis] This often gives rise to the supposition that sulphuric acid possesses a considerable degree of affinity or energy compared with nitric acid, but we shall afterwards see that the idea of the relative degree of affinity of acids and bases is, in many cases, exceedingly unbiassed; it need not be accepted so long as it is possible to explain the observed phenomena without admitting any supposition whatever of the degree of the force of affinity, because the latter cannot be measured. The action of sulphuric acid upon nitre may be explained by the fact alone that the resultant nitric acid is volatile. The nitric acid is the only one of all the substances partaking in the reaction which is able to pass into vapour; it alone is volatile, while the remainder are non-volatile, or, more strictly speaking, exceedingly difficultly volatile substances. Let us imagine that the sulphuric acid is only able to set free a small quantity of nitric acid from its salt, and this will suffice to explain the decomposition of the whole of the nitre by the sulphuric acid, because once the nitric acid is separated it passes into vapour when heated, and passes away from the sphere of action of the remaining substances; then the free sulphuric acid will set free a fresh small quantity of nitric acid, and so on until it drives off the entire quantity. It is evident that, in this explanation, it is essential that the sulphuric acid should be in excess (although not greatly) throughout the reaction; according to the equation expressing the reaction, 98 parts of sulphuric acid are required per 85 parts of Chili nitre; but if this proportion be maintained in practice the nitric acid is not all disengaged by the sulphuric acid; an excess of the latter must be taken, and generally 80 parts of Chili nitre are taken per 98 parts of acid, so that a portion of the sulphuric acid remains free to the very end of the reaction.

[30] It must be observed that sulphuric acid, at least when undiluted (60° Baumé), corrodes cast iron with difficulty, so that the acid may be heated in cast-iron retorts. Nevertheless, both sulphuric and nitric acids have a certain action on cast iron, and therefore the acid obtained will contain traces of iron. In practice sodium nitrate (Chili saltpetre) is usually employed because it is cheaper, but in the laboratory it is best to take potassium nitrate, because it is purer and does not froth up so much as sodium nitrate when heated with sulphuric acid. In the action of an excess of sulphuric acid on nitre and nitric acid a portion of the latter is decomposed, forming lower oxides of nitrogen, which are dissolved in the nitric acid. A portion of the sulphuric acid itself is also carried over as spray by the vapours of the nitric acid. Hence sulphuric acid occurs as an impurity in commercial nitric acid. A certain amount of hydrochloric acid will also be found to be present in it, because sodium chloride is generally found as an impurity in nitre, and under the action of sulphuric acid it forms hydrochloric acid. Commercial acid further contains a considerable excess of water above that necessary for the formation of the hydrate, because water is first poured into the earthenware vessels employed for condensing the nitric acid in order to facilitate its cooling and condensation. Further, the acid of composition HNO_{3} decomposes with great ease, with the evolution of oxides of nitrogen. Thus the commercial acid contains a great number of impurities, and is frequently purified in the following manner:--Lead nitrate is first added to the acid because it forms non-volatile and almost insoluble (precipitated) substances with the free sulphuric and hydrochloric acids, and liberates nitric acid in so doing, according to the equations Pb(NO_{3})_{2} + 2HCl = PbCl_{2} + 2NHO_{3} and Pb(NO_{3})_{2} + H_{2}SO_{4} = PbSO_{4} + 2NHO_{3}. Potassium chromate is then added to the impure nitric acid, by which means oxygen is liberated from the chromic acid, and this oxygen, at the moment of its evolution, oxidises the lower oxides of nitrogen and converts them into nitric acid. A pure nitric acid, containing no impurities other than water, may be then obtained by carefully distilling the acid, treated as above described, and particularly if only the middle portions of the distillate are collected. Such acid should give no precipitate, either with a solution of barium chloride (a precipitate shows the presence of sulphuric acid) or with a solution of silver nitrate (a precipitate shows the presence of hydrochloric acid), nor should it, after being diluted with water, give a coloration with starch containing potassium iodide (a coloration shows the admixture of other oxides of nitrogen). The oxides of nitrogen may be most easily removed from impure nitric acid by heating for a certain time with a small quantity of pure charcoal. By the action of nitric acid on the charcoal carbonic anhydride is evolved, which carries off the lower oxides of nitrogen. On redistilling, pure acid is obtained. The oxides of nitrogen occurring in solution may also be removed by passing air through the nitric acid.

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, HNO_{3}, 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, HNO_{3}, which corresponds with the salts, NMO_{3}, 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, HNO_{3} 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] HNO_{3}; 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.

[31] Dalton, Smith, Bineau, and others considered that the hydrate of constant boiling point (see Chapter I., Note 60) for nitric acid was the compound 2HNO_{3},3H_{2}O, but Roscoe showed that its composition changes with a variation of the pressure and temperature under which the distillation proceeds. Thus, at a pressure of 1 atmosphere the solution of constant boiling point contains 68·6 p.c., and at one-tenth of an atmosphere 66·8 p.c. Judging from what has been said concerning solutions of hydrochloric acid, and from the variation of specific gravity, I think that the comparatively large decrease in the tensions of the vapours depends on the formation of a hydrate, NHO_{3},2H_{2}O (= 63·6 p.c.). Such a hydrate may be expressed by N(HO)_{5}, that is, as NH_{4}(HO), in which all the equivalents of hydrogen are replaced by hydroxyl. The constant boiling point will then be the temperature of the decomposition of this hydrate.

The variation of the specific gravity at 15° from water (_p_ = 0) to the hydrate NHO_{3},5H_{2}O (41·2 p.c. HNO_{3}) is expressed by _s_ = 9992 + 57·4_p_ + 0·16_p^2_, if water = 10,000 at 4°. For example, when _p_ = 30 p.c., _s_ = 11,860. For more concentrated solutions, at least, the above-mentioned hydrate, HNO_{3},2H_{2}O, must be taken, up to which the specific gravity _s_ = 9570 + 84·18_p_-0·240_p_^2; but perhaps (since the results of observations of the specific gravity of the solutions are not in sufficient agreement to arrive at a conclusion) the hydrate HNO_{3},3H_{2}O should be recognised, as is indicated by many nitrates (Al, Mg, Co, &c.), which crystallise with this amount of water of crystallisation. From HNO_{3},2H_{2}O to HNO_{3} the specific gravity of the solutions (at 15°) _s_ = 10,652 + 62·08_p_-0·160_p_^2. The hydrate HNO_{3},2H_{2}O is recognised by Berthelot on the basis of the thermochemical data for solutions of nitric acid, because on approaching to this composition there is a rapid change in the amount of heat evolved by mixing nitric acid with water. Pickering (1892) by refrigeration obtained the crystalline hydrates: HNO_{3},H_{2}O, melting at -37° and HNO_{3},3H_{2}O, melting at -18°. A more detailed study of the reactions of hydrated nitric acid would no doubt show the existence of change in the process and rapidity of reaction in approaching these hydrates.

[32] The normal hydrate HNO_{3}, corresponding with the ordinary salts, may be termed the monohydrated acid, because the anhydride N_{2}O_{5} with water forms this normal nitric acid. In this sense the hydrate HNO_{3},2H_{2}O is the pentahydrated acid.

[33] For technical and laboratory purposes recourse is frequently had to _red fuming nitric acid_--that is, the normal nitric acid, HNO_{3}, containing lower oxides of nitrogen in solution. This acid is prepared by decomposing nitre with half its weight of strong sulphuric acid, or by distilling nitric acid with an excess of sulphuric acid. The normal nitric acid is first obtained, but it partially decomposes, and gives the lower oxidation products of nitrogen, which are dissolved by the nitric acid, to which they impart its usual pale-brown or reddish colour. This acid fumes in the air, from which it attracts moisture, forming a less volatile hydrate. If carbonic anhydride be passed through the red-brown fuming nitric acid for a long period of time, especially if, assisted by a moderate heat, it expels all the lower oxides, and leaves a colourless acid free from these oxides. It is necessary, in the preparation of the red acid, that the receivers should be kept quite cool, because it is only when cold that nitric acid is able to dissolve a large proportion of the oxides of nitrogen. The strong red fuming acid has a specific gravity 1·56 at 20°, and has a suffocating smell of the oxides of nitrogen. When the red acid is mixed with water it turns green, and then of a bluish colour, and with an excess of water ultimately becomes colourless. This is owing to the fact that the oxides of nitrogen in the presence of water and nitric acid are changed, and give coloured solutions.

Markleffsky (1892) showed that the green solutions contain (besides HNO_{3}) HNO_{2} and N_{2}O_{4}, whilst the blue solutions only contain HNO_{2} (_see_ Note 48).

The action of red fuming nitric acid (or a mixture with sulphuric acid) is in many cases very powerful and rapid, and it sometimes acts differently from pure nitric acid. Thus iron becomes covered with a coating of oxides, and insoluble in acids; it becomes, as is said, passive. Thus chromic acid (and potassium dichromate) gives oxide of chromium in this red acid--that is, it is deoxidised. This is owing to the presence of the lower oxides of nitrogen, which are capable of being oxidised--that is, of passing into nitric acid like the higher oxides. But, generally, the action of fuming nitric acid, both red and colourless, is powerfully oxidising.

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 + 2HNO_{3} = Ca(NO_{3})_{2} + H_{2}O. 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(NO_{3})_{_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]

[34] Hydrogen is not evolved in the action of nitric acid (especially strong) on metals, even with those metals which evolve hydrogen under the action of other acids. This is because the hydrogen at the moment of its separation reduces the nitric acid, with formation of the lower oxides of nitrogen, as we shall afterwards see.

[35] Certain basic salts of nitric acid, however (for example, the basic salt of bismuth), are insoluble in water; whilst, on the other hand, all the normal salts are soluble, and this forms an exceptional phenomenon among acids, because all the ordinary acids form insoluble salts with one or another base. Thus, for sulphuric acid the salts of barium, lead, &c., for hydrochloric acid the salts of silver, &c., are insoluble in water. However, the normal salts of acetic and certain other acids are all soluble.

[36] _Ammonium nitrate_, NH_{4}NO_{3}, is easily obtained by adding a solution of ammonia or of ammonium carbonate to nitric acid until it becomes neutral. On evaporating this solution, crystals of the salt are formed which contain no water of crystallisation. It crystallises in prisms like those formed by common nitre, and has a refreshing taste; 100 parts of water at _t_° dissolve 54 + 0·61_t_ parts by weight of the salt. It is soluble in alcohol, melts at 160°, and is decomposed at about 180°, forming water and nitrous oxide, NH_{4}NO_{3} = 2H_{2}O + N_{2}O. If ammonium nitrate be mixed with sulphuric acid, and the mixture be heated to about the boiling point of water, then nitric acid is evolved, and ammonium hydrogen sulphate remains in solution; but if the mixture be heated rapidly to 16O°, then nitrous oxide is evolved. In the first case the sulphuric acid takes up ammonia, and in the second place water. Ammonium nitrate is employed in practice for the artificial production of cold, because in dissolving in water it lowers the temperature very considerably. For this purpose it is best to take equal parts by weight of the salt and water. The salt must first be reduced to a powder and then rapidly stirred up in the water, when the temperature will fall from +15° to -10°, so that the water freezes.

Ammonium nitrate absorbs ammonia, with which it forms unstable compounds resembling compounds containing water of crystallisation. (Divers 1872, Raoult 1873.) At -10° NH_{4}NO_{3},2NH_{3} is formed: it is a liquid of sp. gr. 1·15, which loses all its ammonia under the influence of heat. At +28° NH_{4}NO_{3},NH_{3} is formed: it is a solid which easily parts with its ammonia when heated, especially in solution.

Troost (1882) investigated the tension of the dissociation of the compounds formed, and came to the conclusion that a definite compound corresponding to the formula 2NH_{4}NO_{3},3NH_{3} is formed, because the tension of dissociation remains constant in the decomposition of such a compound at 0°. Y. Kouriloff (1893), however, considers that the constancy of the tension of the ammonia evolved is due to the decomposition of a saturated solution, and not of a definite compound. During decomposition the system is composed of a liquid and a solid; the tension only becomes constant from the moment the solid falls down. The composition 2NH_{4}NO_{3},3NH_{3} corresponds to a saturated solution at 0°, and the solubility of NH_{4}NO_{3} in NH_{3} increases with a rise of temperature.

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, C_{6}H_{6}, for instance, acts according to the equation C_{6}H_{6} + HNO_{3} = H_{2}O + C_{6}H_{5}NO_{2}. Nitrobenzene is produced. The substance taken, C_{6}H_{6}, 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 NO_{2} 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, CO_{2}, 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]

[36 bis] This is explained by saying that in true nitro-compounds the residue of nitric acid NO_{2} takes the place of the hydrogen in the hydrocarbon group. For example, if C_{6}H_{5}OH be given, then C_{6}H_{4}(NO_{2})OH will be a true nitro-compound having the radical properties of C_{6}H_{5}OH. If, on the other hand, the NO_{2} replace the hydrogen of the aqueous radicle (C_{6}H_{5}ONO_{2}), then the chemical character varies, as in the passage of KOH into KONO_{2} (nitre) (_see_ Note 37 and Organic Chemistry).

[37] The compound ethers of nitric acid in which the hydrogen of the aqueous radicle (OH) is replaced by the residue of nitric acid (NO_{2}) are frequently called nitro-compounds. But in their chemical character they differ from true nitro-compounds (for details _see_ Organic Chemistry) and do not burn like them.

The action of nitric acid on cellulose, C_{6}H_{10}O_{5}, is an example. This substance, which forms the outer coating of all plant cells, occurs in an almost pure state in cotton, in common writing-paper, and in flax, &c.; under the action of nitric acid it forms water and nitrocellulose (like water and KNO_{3} from KHO), which, although it has the same appearance as the cotton originally taken, differs from it entirely in properties. It explodes when struck, bursts into flame very easily under the action of sparks, and acts like gunpowder, whence its name of pyroxylin, or gun-cotton. The composition of gun-cotton is C_{6}H_{7}N_{3}O_{11} = C_{6}H_{10}O_{5} + 3NHO_{3}-3H_{2}O. The proportion of the group NO_{2} in nitrocellulose may be decreased by limiting the action of the nitric acid and compounds obtained with different properties; for instance, the (impure) well-known _collodion cotton_, containing from 11 to 12 per cent. of nitrogen, and _pyro-collodion_ (Mendeléeff, 1890), containing 12·4 per cent. of nitrogen. Both these products are soluble in a mixture of alcohol and ether (in collodion a portion of the substance is soluble in alcohol), and the solution when evaporated gives a transparent film, which is insoluble in water. A solution of collodion is employed in medicine for covering wounds, and in wet-plate photography for giving on glass an even coating of a substance into which the various reagents employed in the process are introduced. Extremely fine threads (obtained by forcing a gelatinous mixture of collodion, ether, and alcohol through capillary tubes in water) of collodion form artificial silk.

[37 bis] The property possessed by nitroglycerin (occurring in dynamite), nitrocellulose, and the other nitro-compounds, of burning with an explosion, and their employment for smokeless powder and as explosives in general, depends on the reasons in virtue of which a mixture of nitre and charcoal deflagrates and explodes; in both cases the elements of the nitric acid occurring in the compound are decomposed, the oxygen in burning unites with the carbon, and the nitrogen is set free; thus a very large volume of gaseous substances (nitrogen and oxides of carbon) is rapidly formed from the solid substances originally taken. These gases occupy an incomparably larger volume than the original substance, and therefore produce a powerful pressure and explosion. It is evident that in exploding with the development of heat (that is, in decomposing, not with the absorption of energy, as is generally the case, but with the evolution of energy) the nitro-compounds form stores of energy which are easily set free, and that consequently their elements occur in a state of particularly energetic motion, which is especially strong in the group NO_{2}: this group is common to all nitro-compounds, and all the oxygen compounds of nitrogen are unstable, easily decomposable, and (Note 29) absorb heat in their formation. On the other hand, the nitro-compounds are instructive as an example and proof of the fact that the elements and groups forming compounds are united in definite order in the molecules of a compound. A blow, concussion, or rise of temperature is necessary to bring the combustible elements C and H into the most intimate contact with NO_{2}, and to distribute the elements in a new order in new compounds.

As regards the composition of the nitro-compounds, it will be seen that the hydrogen of a given substance is replaced by the complex group NO_{2} of the nitric acid. The same is observed in the passage of alkalis into nitrates, so that the reactions of substitution of nitric acid--that is, the formation of salts and nitro-compounds--may be expressed in the following manner. In these cases the hydrogen is replaced by the so-called _radicle of nitric acid_ NO_{2}, as is evident from the following table:--

{Caustic potash KHO. {Nitre K(NO_{2})O.

{Hydrate of lime CaH_{2}O_{2}. {Calcium nitrate Ca(NO_{2})_{2}O_{2}.

{Glycerin C_{3}H_{5}H_{3}O_{3}. {Nitroglycerin C_{3}H_{5}(NO_{2})_{3}O_{3}.

{Phenol C_{6}H_{5}OH. {Picric acid C_{6}H_{2}(NO_{2})_{3}OH, &c.

The difference between the salts formed by nitric acid and the nitro-compounds consists in the fact that nitric acid is very easily separated from the salts of nitric acid by means of sulphuric acid (that is, by a method of double saline decomposition), whilst nitric acid is not displaced by sulphuric acid from true nitro-compounds; for instance, nitrobenzene, C_{6}H_{5}·NO_{2}. As nitro-compounds are formed exclusively from hydrocarbons, they are described with them in organic chemistry.

The group NO_{2} of nitro-compounds in many cases (like all the oxidised compounds of nitrogen) passes into the ammonia group or into the ammonia radicle NH_{2}. This requires the action of reducing substances evolving hydrogen: RNO_{2} + 6H = RNH_{3} + 2H_{2}O. Thus Zinin converted nitrobenzene, C_{6}H_{5}·NO_{2}, into aniline, C_{6}H_{5}·NH_{2}, by the action of hydrogen sulphide.

Admitting the existence of the group NO_{2}, as replacing hydrogen in various compounds, then nitric acid may be considered as water in which half the hydrogen is replaced by the radical of nitric acid. In this sense nitric acid is nitro-water, NO_{2}OH, and its anhydride dinitro-water, (NO_{2})_{2}O. In nitric acid the radical of nitric acid is combined with hydroxyl, just as in nitrobenzene it is combined with the radical of benzene.

It should here be remarked that the group NO_{3} may be recognised in the salts of nitric acid, because the salts have the composition M(NO_{3})_{n}, just as the metallic chlorides have the composition MCl_{n}. But the group NO_{3} does not form any other compounds beyond the salts, and therefore it should he considered as hydroxyl, HO, in which H is replaced by NO_{2}.

[38] The nitro-compounds play a very important part in mining and artillery. Detailed accounts of them must be looked for in special works, among which the works of A. R. Shuliachenke and T. M. Chelletsoff occupy an important place in the Russian literature on this subject, although historically the scientific works of Abel in England and Berthelot in France stand pre-eminent. The latter elucidated much in connection with explosive compounds by a series of both experimental and theoretical researches. Among explosives a particularly important place from a practical point of view is occupied by ordinary or black gunpowder (Chapter XIII., Note 16), fulminating mercury (Chapter XVI., Note 26), the different forms of gun-cotton (Chapter VI., Note 37), and nitro-glycerine (Chapter VIII., Note 45, and Chapter XII., Note 33). The latter when mixed with solid pulverulent substances, like magnesia, tripoli, &c., forms dynamite, which is so largely used in quarries and mines in driving tunnels, &c. We may add that the simplest true nitro-compound, or marsh gas, CH_{4}, in which all the hydrogens are replaced by NO_{2} groups has been obtained by L. N. Shishkoff, C(NO_{2})_{4}, as well as nitroform, CH(NO_{2})_{3}.

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, 2NHO_{3} = H_{2}O + 2NO_{2} + O. The decomposition is complete at a white heat--that is, nitrogen is formed, 2NHO_{3} = H_{2}O + N_{2} + O_{5}. Hence it is easily understood that nitric acid may part with its oxygen to a number of substances capable of being oxidised.[39] It is consequently an _oxidising agent_. Charcoal, as we have already seen, burns in nitric acid; phosphorus, sulphur, iodine, and the majority of metals also decompose nitric acid, some on heating and others even at the ordinary temperature: the substances taken are oxidised and the nitric acid is deoxidised, yielding compounds containing less oxygen. Only a few metals, such as gold and platinum, do not act on nitric acid, but the majority decompose it; in so doing, an oxide of the metal is formed, which, if it has the character of a base, acts on the remaining nitric acid; hence, with the majority of metals the result of the reaction is usually not an oxide of the metal, but the corresponding salt of nitric acid, and, at the same time, one of the lower oxides of nitrogen. The resulting salts of the metals are soluble, and hence it is said that nitric acid _dissolves_ nearly all metals.[40] This case is termed the solution of metals by acids, although it is not a case of simple solution, but a complex chemical change of the substances taken. When treated with this acid, those metals whose oxides do not combine with nitric acid yield the oxide itself, and not a salt; for example, tin acts in this manner on nitric acid, forming a hydrated oxide, SnH_{2}O_{3}, which is obtained in the form of a white powder, Sn + 4NHO_{3} = H_{2}SnO_{3} + 4NO_{2} + H_{2}O. Silver is able to take up still more oxygen, and to convert a large portion of nitric acid into nitrous anhydride, 4Ag + 6HNO_{3} = 4AgNO_{3} + N_{2}O_{3} + 3H_{2}O. 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 + 10NHO_{3} = 4Zn(NO_{3})_{2} + N_{2}O + 5H_{2}O.[41] Sometimes, and especially with dilute solutions of nitric acid, the deoxidation proceeds as far as the formation of hydroxylamine and ammonia, and sometimes it leads to the formation of nitrogen itself. The formation of one or other nitrogenous substance from nitric acid is determined, not only by the nature of the reacting substances, but also by the relative mass of water and nitric acid, and also by the temperature and pressure, or the sum total of the conditions of reaction; and as in a given mixture even these conditions vary (the temperature and the relative mass vary), it not unfrequently happens that a mixture of different products of the deoxidation of nitric acid is formed.

[39] [Illustration: FIG. 49.--Decomposition of nitrous oxide by sodium.]

Nitric acid may be entirely decomposed by passing its vapour over highly incandescent copper, because the oxides of nitrogen first formed give up their oxygen to the red-hot metallic copper, so that water and nitrogen gas alone are obtained. This forms a means for determining the composition both of nitric acid and of all the other compounds of nitrogen with oxygen, because by collecting the gaseous nitrogen formed it is possible to calculate, from its volume, its weight and consequently its amount in a given quantity of a nitrogenous substance, and by weighing the copper before and after the decomposition it is possible to determine the amount of oxygen by the increase in weight. The complete decomposition of nitric acid is also accomplished by passing a mixture of hydrogen and nitric acid vapours through a red-hot tube. Sodium also decomposes the oxides of nitrogen at a red-heat, taking up all the oxygen. This method is sometimes used for determining the composition of the oxides of nitrogen.

[40] The application of this acid for etching copper or steel in engraving is based on this fact. The copper is covered with a coating of wax, resin, &c. (etching ground), on which nitric acid does not act, and then the ground is removed in certain parts with a needle, and the whole is washed in nitric acid. The parts coated remain untouched, whilst the uncovered portions are eaten into by the acid. Copper plates for etchings, aquatints, &c., are prepared in this manner.

[41] The formation of such complex equations as the above often presents some difficulty to the beginner. It should be observed that if the reacting and resultant substances be known, it is easy to form an equation for the reaction. Thus, if we wish to form an equation expressing the reaction that nitric acid acting on zinc gives nitrous oxide, N_{2}O, and zinc nitrate, Zn(NO_{3})_{2}, we must reason as follows:--Nitric acid contains hydrogen, whilst the salt and nitrous oxide do not; hence water is formed, and therefore it is as though anhydrous nitric acid, N_{2}O_{5}, were acting. For its conversion into nitrous oxide it parts with four equivalents of oxygen, and hence it is able to oxidise four equivalents of zinc and to convert it into zinc oxide, ZnO. These four equivalents of zinc oxide require for their conversion into the salt four more equivalents of nitric anhydride; consequently five equivalents in all of the latter are required, or ten equivalents of nitric acid. Thus ten equivalents of nitric acid are necessary for four equivalents of zinc in order to express the reaction in whole equivalents. It must not be forgotten, however, that there are very few such reactions which can be entirely expressed by simple equations. The majority of equations of reactions only express the chief and ultimate products of reaction, and thus none of the three preceding equations express all that in reality occurs in the action of metals on nitric acid. In no one of them is only one oxide of nitrogen formed, but always several together or consecutively--one after the other, according to the temperature and strength of the acid. And this is easily intelligible. The resulting oxide is itself capable of acting on metals and of being deoxidised, and in the presence of the nitric acid it may change the acid and be itself changed. The equations given must be looked on as a systematic expression of the main features of reactions, or as a limit towards which they tend, but to which they only attain in the absence of disturbing influences.

Thus the action of nitric acid on metals consists in their being oxidised, whilst the acid itself is converted, according to the temperature, concentration in which it is taken, and the nature of the metal, &c., into lower oxides, ammonia, or even into nitrogen.[42] Many compounds are oxidised by nitric acid like metals and other elements; for instance, lower oxides are converted into higher oxides. Thus, arsenious acid is converted into arsenic acid, suboxide of iron into oxide, sulphurous acid into sulphuric acid, the sulphides of the metals, M_{2}S, into sulphates, M_{2}SO_{4}, &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, C_{10}H_{16}, bursts into flame when poured into fuming nitric acid. In virtue of its oxidising property, nitric acid _removes the hydrogen_ from many substances. Thus it decomposes hydriodic acid, separating the iodine and forming water; and if fuming nitric acid be poured into a flask containing gaseous hydriodic acid, then a rapid reaction takes place, accompanied by flame and the separation of violet vapours of iodine and brown fumes of oxides of nitrogen.[43]

[42] Montemartini endeavours to show that the products evolved in the action of nitric acid upon metals (and their amount) is in direct connection with both the concentration of the acid and the capacity of the metals to decompose water. Those metals which only decompose water at a high temperature give, under the action of nitric acid, NO_{2}, N_{2}O_{4}, and NO; whilst those metals which decompose water at a lower temperature give, besides the above products, N_{2}O, N, and NH_{3}; and, lastly, the metals which decompose water at the ordinary temperature also evolve hydrogen. It is observed that concentrated nitric acid oxidises many metals with much greater difficulty than when diluted with water; iron, copper, and tin are very easily oxidised by dilute nitric acid, but remain unaltered under the influence of monohydrated nitric acid or of the pure hydrate NHO_{3}. Nitric acid diluted with a large quantity of water does not oxidise copper, but it oxidises tin; dilute nitric acid also does not oxidise either silver or mercury; but, on the addition of nitrous acid, even dilute acid acts on the above metals. This naturally depends on the smaller stability of nitrous acid, and on the fact that after the commencement of the action the nitric acid is itself converted into nitrous acid, which continues to act on the silver and mercury. Veley (Oxford 1891) made detailed researches on the action of nitric acid upon Cu, Hg, and Bi, and showed that nitric acid of 30 p.c. strength does not act upon these metals at the ordinary temperature if nitrous acid (traces are destroyed by urea) and oxidising agents such as H_{2}O_{2}, KClO_{3}, &c. be entirely absent; but in the presence of even a small amount of nitrous acid the metals form nitrites, which, with HNO_{3}, form nitrates and the oxides of nitrogen, which re-form the nitrous acid necessary for starting the reaction, because the reaction 2NO + HNO_{3} + H_{2}O = 3HNO_{2} is reversible. The above metals are quickly dissolved in a 1 p.c. solution of nitrous acid. Moreover, Veley observed that nitric acid is partially converted into nitrous acid by gaseous hydrogen in the presence of the nitrates of Cu and Pb.

[43] When nitric acid acts on many organic substances it often happens that not only is hydrogen removed, but also oxygen is combined; thus, for example, nitric acid converts toluene, C_{7}H_{8}, into benzoic acid, C_{7}H_{6}O_{2}. In certain cases, also, a portion of the carbon contained in an organic substance burns at the expense of the oxygen of the nitric acid. So, for instance, phthalic acid, C_{8}H_{6}O_{4}, is obtained from naphthalene, _{10}H_{8}. Thus the action of nitric acid on the hydrocarbons is often most complex; not only does nitrification take place, but also separation of carbon, displacement of hydrogen, and combination of oxygen. There are few organic substances which can withstand the action of nitric acid, and it causes fundamental changes in a number of them. It leaves a yellow stain on the skin, and in a large quantity causes a wound and entirely eats away the membranes of the body. The membranes of plants are eaten into with the greatest ease by strong nitric acid in just the same manner. One of the most durable blue vegetable dyes employed in dyeing tissues is _indigo_; yet it is easily _converted into a yellow substance_ by the action of nitric acid, and small traces of free nitric acid may be recognised by this means.

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_, N_{2}O_{5}; 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° (2AgNO_{3} + Cl_{2} = 2AgCl + N_{2}O_{5} + 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, P_{2}O_{5}, 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, H_{2}O,2N_{2}O_{5} = N_{2}O_{5},2HNO_{3}, whilst the greater part of the nitric acid gives the anhydride according to the equation 2NHO_{3} + P_{2}O_{5} = 2PHO_{3} + N_{2}O_{5}. On heating, nitric anhydride decomposes with an explosion, or gradually, into nitric peroxide and oxygen, N_{2}O_{5} = N_{2}O_{4} + O.

_Nitrogen peroxide_, N_{2}O_{4}, and _nitrogen dioxide_, NO_{2}, 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, O_{3} loses heat in passing into O_{2}, whilst N_{2}O_{4} absorbs heat in forming NO_{2}.

Nitric acid in acting on tin and on many organic substances (for example, starch) gives brown vapours, consisting of a mixture of N_{2}O_{3} and NO_{2}. A purer product is obtained by the decomposition of lead nitrate by heat, Pb(NO_{3})_{2} = 2NO_{2} + O + PbO, when non-volatile lead oxide, oxygen gas, and nitrogen peroxide are formed. The latter condenses, in a well-cooled vessel, to a brown liquid, which boils at about 22°. The purest peroxide of nitrogen, solidifying at -9°, is obtained by mixing dry oxygen in a freezing-mixture with twice its volume of dry nitric oxide, NO, when transparent prisms of nitrogen peroxide are formed in the receiver: they melt into a colourless liquid at about -10°. When the temperature of the receiver is above -9°, the crystals melt,[44] and at 0° give a reddish yellow liquid, like that obtained in the decomposition of lead nitrate. The vapours of nitrogen peroxide have a characteristic odour, and at the ordinary temperature are of a dark-brown colour, but at lower temperatures the colour of the vapour is much fainter. When heated, especially above 50°, the colour becomes a very dark brown, so that the vapours almost lose their transparency.

[44] According to certain investigations, if a brown liquid is formed from the melted crystals by beating above -9°, then they no longer solidify at -10°, probably because a certain amount of N_{2}O_{3} (and oxygen) is formed, and this substance remains liquid at -30°, or it may be that the passage from 2NO_{2} into N_{2}O_{4} is not so easily accomplished as the passage from N_{2}O_{4} into 2NO_{2}.

Liquid nitrogen peroxide (that is, a mixture of NO_{2} and N_{2}O_{4}) is employed in admixture with hydrocarbons as an explosive.

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 NO_{2} (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, N_{2}O_{4}, 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 (N_{2}O_{4} gives colourless and transparent vapours, whilst those of NO_{2} 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 N_{2}O_{4} with 21 parts by weight of NO_{2}.[45] It is evident that a decomposition here takes place the peculiarity of which consists in the fact that the product of decomposition, NO_{2}, is polymerised (_i.e._ becomes denser, combines with itself) at a lower temperature; that is, the reaction

N_{2}O_{4} = NO_{2} + NO_{2}

is a reversible reaction, and consequently the whole phenomenon represents a _dissociation_ in a homogeneous gaseous medium, where the original substance, N_{2}O_{4}, and the resultant, NO_{2}, 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 N_{2}O_{4} is not then decomposable. Consequently the limits of dissociation here are -10° and 135° at the atmospheric pressure.[46] Within the limits of these temperatures the vapours of nitrogen peroxide have not a constant density, but, on the other hand, above and below these limits definite substances exist. Thus above 135° N_{2}O_{4} has ceased to exist and NO_{2} alone remains. It is evident that at the ordinary temperature there is a partially dissociated system or mixture of nitrogen peroxide, N_{2}O_{4}, and nitrogen dioxide, NO_{2}. In the brown liquid boiling at 22° probably a portion of the N_{2}O_{4} has already passed into NO_{2}, and it is only the colourless liquid and crystalline substance at -10° that can be considered as pure nitrogen peroxide.[47]

[45] Because if _x_ equal the amount by weight of N_{2}O_{4}, its volume will = _x_/46, and the amount of NO_{2} will = 100-_x_, and consequently its volume will = (100-_x_)/23. But the mixture, having a density 38, will weigh 100; consequently its volume will = 100/38. Hence _x_/46 + (100-_x_)/23 = 100/38, or _x_ = 79·O.

[46] The phenomena and laws of dissociation, which we shall consider only in particular instances, are discussed in detail in works on theoretical chemistry. Nevertheless, in respect to nitrogen peroxide, as an historically important example of dissociation in a homogeneous gaseous medium, we will cite the results of the careful investigations (1885-1880) of E. and L. Natanson, who determined the densities under various conditions of temperature and pressure. The degree of dissociation, expressed as above (it may also he expressed otherwise--for example, by the ratio of the quantity of substance decomposed to that unaltered), proves to increase at all temperatures as the pressure diminishes, which would he expected for a homogeneous gaseous medium, as a decreasing pressure aids the formation of the lightest product of dissociation (that having the least density or largest volume). Thus, in the Natansons' experiments the degree of dissociation at 0° increases from 10 p.c. to 30 p.c., with a decrease of pressure of from 251 to 38 mm.; at 49°·7 it increases from 49 p.c. to 93 p.c., with a fall of pressure of from 498 to 27 mm., and at 100° it increases from 89·2 p.c. to 99·7 p.c., with a fall of pressure of from 732·5 to 11·7 mm. At 130° and 150° the decomposition is complete--that is, only NO_{2} remains at the low pressures (less than the atmospheric) at which the Natansons made their determinations; but it is probable that at higher pressures (of several atmospheres) molecules of N_{2}O_{4} would still be formed, and it would be exceedingly interesting to trace the phenomena under the conditions of both very considerable pressures and of relatively large volumes.

[47] Liquid nitrogen peroxide is said by Geuther to boil at 22°-26°, and to have a sp. gr. at 0° = 1·494 and at 15° = 1·474. It is evident that, in the liquid as in the gaseous state, the variation of density with the temperature depends, not only on physical, but also on chemical changes, as the amount of N_{2}O_{4} decreases and the amount of NO_{2} increases with the temperature, and they (as polymeric substances) should have different densities, as we find, for instance, in the hydrocarbons C_{5}H_{10} and C_{10}H_{20}.

It may not be superfluous to mention here that the measurement of the specific heat of a mixture of the vapours of N_{2}O_{4} and NO_{2} enabled Berthelot to determine that the transformation of 2NO_{2} into N_{2}O_{4} is accompanied by the evolution of about 13,000 units of heat, and as the reaction proceeds with equal facility in either direction, it will be exothermal in the one direction and endothermal in the other; and this clearly demonstrates the possibility of reactions taking place in either direction, although, as a rule, reactions evolving heat proceed with greater ease.

The above explains the action of nitrogen peroxide on water at low temperatures. N_{2}O_{4} then acts on water like a mixture of the anhydrides of nitrous and nitric acids. The first, N_{2}O_{3}, 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 NO_{2}, proper to nitric acid; and in nitrogen peroxide one atom of the hydrogen of water is replaced by NO and the other by NO_{2}, as is seen from the formulæ--

H} NO} NO } NO_{2}} H} O; NO} O; NO_{2}} O; NO_{2}} O;

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, HNO_{3}, and nitrous, HNO_{2}, acids. The latter, as we shall afterwards see, splits up into water and the anhydride, N_{2}O_{3}. If, however, warm water act on nitrogen peroxide, only nitric acid and monoxide of nitrogen are formed: 3NO_{2} + H_{2}O = NO + 2NHO_{3}.

Although NO_{2} 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, 2NO_{2} + Hg = HgNO_{3} + NO, being itself deoxidised into nitric oxide, into which the dioxide in many other instances passes, and from which it is easily formed.[48]

[48] Nitric acid of sp. gr. 1·51 in dissolving nitrogen peroxide becomes brown, whilst nitric acid of sp. gr. 1·32 is coloured greenish blue, and acid of sp. gr. below 1·15 remains colourless after absorbing nitrogen peroxide (Note 33).

_Nitrous anhydride_, N_{2}O_{3}, corresponds[49] to nitrous acid, NHO_{2}, which forms a series of salts, the nitrites--for example, the sodium salt NaNO_{2}, the potassium salt KNO_{2}, the ammonium salt (NH_{4})NO_{2},[50] the silver salt AgNO_{2},[51] &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 NHO_{2} from its salts, it always gives water and the anhydride, whilst the latter, as an intermediate oxide, partially or wholly splits up into NO + NO_{2}. But the salts of nitrous acid are distinguished for their great stability. Potassium nitrate, KNO_{3}, 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, KNO_{3} + Pb = KNO_{2} + PbO.[51 bis] The resultant salt is soluble in water, whilst the oxide of lead is insoluble. With sulphuric and other acids the solution of potassium nitrite[52] immediately evolves a brown gas, nitrous anhydride: 2KNO_{2} + H_{2}SO_{4} = K_{2}SO_{4} + N_{2}O_{3} + H_{2}O. The same gas (N_{2}O_{3}) is obtained by passing nitric oxide at 0° through liquid peroxide of nitrogen,[53] 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°,[54] but then partially decomposing into NO + NO_{2}. 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.[55] Nitrous acid acts directly on ammonia, forming nitrogen and water, HNO_{2} + NH_{3} = N_{2} + 2H_{2}O.[56]

[49] Nitrogen peroxide as a mixed substance has no corresponding independent salts, but Sabatier and Senderens (1892) showed that under certain conditions NO_{2} combines directly with some metals--for instance, copper and cobalt--forming Cu_{2}NO_{2} and CoNO_{2} as dark brown powders, which do not, however, exhibit the reactions of salts. Thus by passing gaseous nitrogen dioxide over freshly reduced (from the oxides by heating with hydrogen) copper at 25°-30°, Cu_{2}NO_{2} is directly formed. With water it partly gives off NO_{2} and partly forms nitrite of copper, leaving metallic copper and its suboxide. The nature of these compounds has not yet been sufficiently investigated.

[50] Ammonium nitrite may be easily obtained in solution by a similar method of double decomposition (for instance, of the barium salt with ammonium sulphate) to the other salts of nitrous acid, but it decomposes with great ease when evaporated, with evolution of gaseous nitrogen, as already mentioned (Chapter V.) If the solution, however, be evaporated at the ordinary temperature under the receiver of an air-pump, a solid saline mass is obtained, which is easily decomposed when heated. The dry salt even decomposes with an explosion when struck, or when heated to about 70°--NH_{4}NO_{2} = 2H_{2}O + N_{2}. It is also formed by the action of aqueous ammonia on a mixture of nitric oxide and oxygen, or by the action of ozone on ammonia, and in many other instances. Zörensen (1894) prepared NH_{4}NO_{2} by the action of a mixture of N_{2}O_{3} and other oxides of nitrogen on lumps of ammonium carbonate, extracting the nitrite of ammonium formed with absolute alcohol, and precipitating it from this solution by ether. This salt is crystalline, dissolves in water with absorption of heat, and attracts moisture from the air. The solid salt and its concentrated solutions decompose with an explosion when heated to 50°-80°, especially in the presence of traces of foreign acids. Decomposition also proceeds at the ordinary temperature, but more slowly; and in order to preserve the salt it should be covered with a layer of pure dry ether.

[51] Silver nitrite, AgNO_{2}, is obtained as a very slightly soluble substance, as a precipitate, on mixing solutions of silver nitrate, AgNO_{3}, and potassium nitrite, KNO_{2}. It is soluble in a large volume of water, and this is taken advantage of to free it from silver oxide, which is also present in the precipitate, owing to the fact that potassium nitrite always contains a certain amount of oxide, which with water gives the hydroxide, forming oxide of silver with silver nitrate. The solution of silver nitrite gives, by double decomposition with metallic chlorides (for instance, barium chloride), insoluble silver chloride and the nitrite of the metal taken (in this case, barium nitrite, Ba(NO_{2})_{2}).

[51 bis] Leroy (1889) obtained KNO_{2} by mixing powdered KNO_{3} with BaS, igniting the mixture in a crucible and washing the fused salts; BaSO_{4} is then left as an insoluble residue, and KNO_{2} passes into solution: 4KNO_{3} + BaS = 4KNO_{2} + BaSO_{4}.

[52] Probably potassium nitrite, KNO_{2}, when strongly heated, especially with metallic oxides, evolves N and O, and gives potassium oxide, K_{2}O, because nitre is liable to such a decomposition; but it has, as yet, been but little investigated.

[53] There are many researches which lead to the conclusion that the reaction N_{2}O_{3} = NO_{2}-NO is reversible, _i.e._ resembles the conversion of N_{2}O_{4} into NO_{2}. The brown colour of the fumes of N_{2}O_{3} is due to the formation of NO_{2}.

If nitrogen peroxide be cooled to -20°, and half its weight of water be added to it drop by drop, then the peroxide is decomposed, as we have already said, into nitrous and nitric acids; the former does not then remain as a hydrate, but straightway passes into the anhydride, and, hence, if the resultant liquid be slightly warmed vapours of nitrous anhydride, N_{2}O_{3}, are evolved, and condense into a blue liquid, as Fritzsche showed. This method of preparing nitrous anhydride apparently gives the purest product, but it easily dissociates, forming NO and NO_{2} (and therefore also nitric acid in the presence of water).

[54] According to Thorpe, N_{2}O_{3} boils at +18°. According to Geuther, at +3°·5, and its sp. gr. at 0° = 1·449.

[55] In its oxidising action nitrous anhydride gives nitric oxide, N_{2}O_{3} = 2NO + O. Thus its analogy to ozone becomes still more marked, because in ozone it is only one-third of the oxygen that acts in oxidising; from O_{3} there is obtained O, which acts as an oxidiser, and common oxygen O_{2}. In a physical aspect the relation between N_{2}O_{3} and O_{3} is revealed in the fact that both substances are of a blue colour when in the liquid state.

[56] This reaction is taken advantage of for converting the amides, NH_{2}R (where R is an element or a complex group) into hydroxides, RHO. In this case NH_{2}R + NHO_{2} forms 2N + H_{2}O + RHO; NH_{2}, is replaced by HO, the radicle of ammonia by the radicle of water. This reaction is employed for transforming many nitrogenous organic substances having the properties of amides into their corresponding hydroxides. Thus aniline, C_{6}H_{5}·NH_{2}, which is obtained from nitrobenzene, C_{6}H_{5}·NO_{2} (Note 37), is converted by nitrous anhydride into phenol, C_{6}H_{5}·OH, which occurs in the creosote extracted from coal tar. Thus the H of the benzene is successively replaced by NO_{2}, NH_{2}, and HO; a method which is suitable for other cases also.

As nitrous anhydride easily splits up into NO_{2} + NO, so, like NO_{2}, with warm water it gives nitric acid and nitric oxide, according to the equation 3N_{2}O_{3} + H_{2}O = 4NO + 2NHO_{3}.

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.[57]

[57] The action of a solution of potassium permanganate, KMnO_{4}, on nitrous acid in the presence of sulphuric acid is determined by the fact that the higher oxide of manganese, Mn_{2}O_{7}, contained in the permanganate is converted into the lower oxide, MnO, which as a base forms manganese sulphate, MnSO_{4}, and the oxygen serves for the oxidation of the N_{2}O_{3} into N_{2}O_{5}, or its hydrate. As the solution of the permanganate is of a red colour, whilst that of manganese sulphate is almost colourless, this reaction is clearly seen, and may be employed for the detection and determination of nitrous acid and its salts.

_Nitric oxide_, NO.--This permanent gas[58] (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 N_{2}O_{3} and NO_{2} are produced) is poured into a flask containing metallic copper.[59] 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

2NHO_{3} + 3Cu = H_{2}O + 3CuO + 2NO.

The second reaction consists in the formation of copper nitrate--

6NHO_{3} + 3CuO = 3H_{2}O + 3Cu(NO_{3})_{2}.

[58] The absolute boiling point = -93° (_see_ Chapter II., Note 29).

[59] Kammerer proposed preparing nitric oxide, NO, by pouring a solution of sodium nitrate over copper shavings, and adding sulphuric acid drop by drop. The oxidation of ferrous salts by nitric acid also gives NO. One part of strong hydrochloric acid is taken and iron is dissolved in it (FeCl_{2}), and then an equal quantity of hydrochloric acid and nitre is added to the solution. On heating, nitric oxide is evolved. In the presence of an excess of sulphuric acid and mercury the conversion of nitric acid into nitric oxide is complete (that is, the reaction proceeds to the end and the nitric oxide is obtained without other products), and upon this is founded one of the methods for determining nitric acid (in nitrometers of various kinds, described in text-books of analytical chemistry), as the amount of NO can be easily and accurately measured volumetrically. The amount of nitrogen in gun-cotton, for instance, is determined by dissolving it in sulphuric acid. Nitrous acid acts in the same manner. Upon this property Emich (1892) founds his method for preparing pure NO. He pours mercury into a flask, and then covers it with sulphuric acid, in which a certain amount of NaNO_{2} or other substance corresponding to HNO_{2} or HNO_{3} has been dissolved. The evolution of NO proceeds at the ordinary temperature, being more rapid as the surface of the mercury is increased (if shaken, the reaction proceeds very rapidly). If the gas be passed over KHO, it is obtained quite pure, because KHO does not act upon NO at the ordinary temperature (if heated, KNO_{2} and N_{2}O or N_{2}, are formed).

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 N_{2} and 2N_{2}O_{3}, but complete decomposition into N_{2} and O_{2} 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 = N_{2}O_{3}, 2NO + O_{2} = 2NO_{2}. 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 + H_{2}O + O_{3} = 2HNO_{3}, 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 HNO_{3}, is absorbed, and the water entirely fills the bell jar previously containing the gas.[60] It is evident that nitric oxide[61] in combining with oxygen has a strong tendency to give the higher types of nitrogen compounds, which we see in nitric acid, HNO_{3} or NO_{2}(OH), in nitric anhydride, N_{2}O_{5} or (NO_{2})_{2}O, and in ammonium chloride, NH_{4}Cl. 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 H_{2}, be indicated by X_{2}, then the three compounds of nitrogen above named should be considered as compounds of the type or form NX_{5}. For example, in nitric acid X_{5} = O_{2} + (OH), where O_{2} = X_{4}, and OH = X; whilst nitric oxide is a compound of the form NX_{2}. 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. NX_{2} passes consecutively into NX_{3}--namely, into N_{2}O_{3} and NHO_{2}, NX_{4} (for instance NO_{2}) and NX_{5}.

[60] This transformation of the permanent gases nitric oxide and oxygen into liquid nitric acid in the presence of water, and with the evolution of heat, presents a most striking instance of liquefaction produced by the action of chemical forces. They perform with ease the work which physical (cooling) and mechanical (pressure) forces effect with difficulty. In this the motion, which is so distinctively the property of the gaseous molecules, is apparently destroyed. In other cases of chemical action it is apparently created, arising, no doubt, from latent energy--that is, from the internal motion of the atoms in the molecules.

[61] Nitric oxide is capable of entering into many characteristic combinations; it is absorbed by the solutions of many acids, for instance, tartaric, acetic, phosphoric, sulphuric, and metallic chlorides (for example, SbCl_{5}, BiCl_{3}, &c., with which it forms definite compounds; Besson 1889), and also by the solutions of many salts, especially those formed by suboxide of iron (for instance, ferrous sulphate). In this case a brown compound is formed which is exceedingly unstable, like all the analogous compounds of nitric oxide. The amount of nitric oxide combined in this manner is in atomic proportion with the amount of the substance taken; thus ferrous sulphate, FeSO_{4}, absorbs it in the proportion of NO to 2FeSO_{4}. Ammonia is obtained by the action of a caustic alkali on the resultant compound, because the oxygen of the nitric oxide and water are transferred to the ferrous oxide, forming ferric oxide, whilst the nitrogen combines with the hydrogen of the water. According to the investigations of Gay (1885), the compound is formed with the evolution of a large quantity of heat, and is easily dissociated, like a solution of ammonia in water. It is evident that oxidising substances (for example, potassium permanganate, KMnO_{4}, Note 57) are able to convert it into nitric acid. If the presence of a radicle NO_{2}, composed like nitrogen peroxide, must be recognised in the compounds of nitric acid, then a radicle NO, having the composition of nitric oxide, may be admitted in the compounds of nitrous acid. The compounds in which the radicle NO is recognised are called _nitroso-compounds_. These substances are described in Prof. Bunge's work (Kief, 1868).

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 oxide[62] in which a feebly incandescent piece of charcoal is extinguished.

[62] A mixture of nitric oxide and hydrogen is inflammable. If a mixture of the two gases be passed over spongy platinum the nitrogen and hydrogen even combine, forming ammonia. A mixture of nitric oxide with many combustible vapours and gases is very inflammable. A very characteristic flame is obtained in burning a mixture of nitric oxide and the vapour of the combustible carbon bisulphide, CS_{2}. The latter substance is very volatile, so that it is sufficient to pass the nitric oxide through a layer of the carbon bisulphide (for instance, in a Woulfe's bottle) in order that the gas escaping should contain a considerable amount of the vapours of this substance. This mixture continues to burn when ignited, and the flame emits a large quantity of the so-called ultra-violet rays, which are capable of inducing chemical combinations and decompositions, and therefore the flame may be employed in photography in the absence of sufficient daylight (magnesium light and electric light have the same property). There are many gases (for instance, ammonia) which when mixed with nitric oxide explode in a eudiometer.

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.[63] 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, SO_{2}, 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, SO_{3}, and if it does form sulphuric acid with water and the oxygen of the atmosphere, SO_{2} + H_{2}O + O = H_{2}SO_{4}, 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--

3SO_{2} + 2NHO_{3} + 2H_{2}O = 2H_{2}SO_{4} + 2NO.

[63] The oxides of nitrogen naturally do not proceed directly from oxygen and nitrogen by contact alone, because their formation is accompanied by the absorption of a large quantity of heat, for (_see_ Note 29) about 21,500 heat units are absorbed when 16 parts of oxygen and 14 parts of nitrogen combine; consequently the decomposition of nitric oxide into oxygen and nitrogen is accompanied by the evolution of this amount of heat; and therefore with nitric oxide, as with all explosive substances and mixtures, the reaction once started is able to proceed by itself. In fact, Berthelot remarked the decomposition of nitric oxide in the explosion of fulminate of mercury. This decomposition does not take place spontaneously; substances even burn with difficulty in nitric oxide, probably because a certain portion of the nitric oxide in decomposing gives oxygen, which combines with another portion of nitric oxide, and forms nitric peroxide, a somewhat more stable compound of nitrogen and oxygen. The further combinations of nitric oxide with oxygen all proceed with the evolution of heat, and take place spontaneously by contact with air alone. It is evident from these examples that the application of thermochemical data is limited.

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