NITROGEN is a chemical element which, on account of its abundance in nature and its relations to life, is of great importance. About three-fourths of the mass of the atmosphere consists of elementary nitrogen; and, as an essential component of all albumenoids, the element per-vades the whole of the animal and vegetable kingdom. Nitrogen minerals are scarce (almost the only ones are Chili saltpetre and native nitre), but traces of the two nitrogen compounds, ammonia and nitric acid, are diffused through-out all soils, besides existing in the atmosphere.

Elementary nitrogen exists only in the one form of nitrogen gas (N2 = 1 molecule), which is easily extracted from the atmosphere. Though resembling air in its general properties, it is easily distinguished from it by its not supporting combustion. According to Regnault, its specific gravity is 0-9703 times that of pure, dry air; andi one litre, measured at t° C. and P millimetres pressure! (strictly speaking the pressure exerted by a column of mercury of 0° C. and the height P, at latitude 45° and sea-level) is W grammes, where

For sea-level and lat. 55° 54' (Edinburgh or Glasgow) the constant is 0-45090; for lat. 51° 32' (London), 0-45072.

According to Dittmar (Reports on the Challenger Expe-dition) 1000 volumes of pure water, when shaken with excess of gas at t° C. and "one.atmosphere's pressure," absorb /3 volumes of the gas measured at 0° and the same pressure,—/3 having for the temperatures given the fol-lowing values:—
1= 0° 15° 25° 30°
(8 = 24-40 17-65 14-95 13-90

Nitrogen is a permanent gas in this sense that no amount of pressure will liquefy it at any temperature lying above the "critical point" of - 123°"8 C. At or a little below this temperature 42 atmospheres reduce it to a liquid (Sarrau).

Chemically, nitrogen gas is characterized by perfect inertness towards all ordinary reagents under ordinary conditions. But at certain higher temperatures boron, magnesium, vanadium, and titanium combine with it directly into nitrides. Nitrogen is capable even of uniting with ordinary oxygen. A mixture of the two gases, it is true, remains unchanged when exposed en masse to any tem-perature, but when it is subjected to a succession of electric sparks a small proportion of the two gases, no doubt through local dissociation into isolated atoms N and O, does unite into nitric oxide, NO, which then combines with more oxygen into red fumes of peroxide, N204.

The part which the nitrogen gas in the atmosphere plays in the economy of nature is as yet a mystery. It certainly is not susceptible of being taken up directly by the plants and utilized in their synthesis of nitrogenous compounds. It plays no active part in the processes of combustion and of animal respiration ; in either it appears to act only as an inert diluent of the oxygen.

In the case of respiration, however, this particular diluent seems to be essential; no animal could live healthily for any considerable period of time in pure oxygen, and we know of no other diluent which could be substituted for the nitrogen without producing poisonous effects. Besides, there can be no doubt that atmospheric nitrogen, in an indirect way, contributes towards the building up of nitrogenous organic matter. Every process of ordinary combustion probably, and every electric dis-charge in the atmosphere certainly, induces the formation of some nitric acid, which by combining with the atmo-spheric ammonia becomes nitrate of ammonia, and from certain experiments of Schonbein's it would appear that nitrogen gas and water are capable of uniting directly into nitrite of ammonia (N2 + 2Hj,0 = N02 . NH4), which, supposing it to be produced in the atmosphere, would promptly be oxidized into nitrate. The nitrate produced by either process is carried down by the rain and conveyed to the roots of the .plants, which assimilate it as part of their nitrogenous organic ma tter. However small the scale may appear on which these processes of atmospheric nitrifica-tion go on when measured by the mass of nitrogen which remains unchanged, as this mass is immense, their absolute effect must be very considerable, and may form an important item in the economy of nature.

The compounds of nitrogen may be arranged under the heads, of ammonia, nitrates, nitro-compounds, organic nitrogen compounds, and cyanides. As all the several classes and their most important members are treated of under CHEMISTRY (vol. v. p. 509-514), we confine our-selves here, in the main, to supplementing that article by such details as are of practical or general scientific interest.

Ammonia

This, the only known compound of hydrogen with nitro-gen, is a gas of the molecular formula NH3. The most convenient process for the preparation of the pure gas is to mix powdered sal-ammoniac with powdered quicklime in a flask and to heat the mixture in a sand-bath. Torrents of ammonia come off, which must be dried by passing it through a closely packed column of solid caustic potash or soda (chloride of calcium absorbs the gas chemically) and collected over mercury, as the gas dissolves most abundantly in water. In addition to what has been said under CHEMISTRY of the properties of ammonia, it may here be mentioned that, though uninflammable in air, it burns brilliantly in oxygen, and that it is liable to the following peculiar kind of oxidation. Pour some strong liquor ammonise into a large flask, so as to produce a moist mixture of the gas and of air, and suspend in this atmo-sphere a recently ignited spiral of thin platinum wire. The wire continues glowing and the flask soon fills with dense white fumes of nitrite and nitrate of am-monia, formed according to equation: 2NH3 + (40 or 30) = H20 + N03H.NH3 or N02H.NH3. The platinum suffers no permanent change; its mode of action probably consists in this that it alternately absorbs (i.e., combines with) oxygen and hands it over to the ammonia. In any case the reaction is interesting as throwing some light upon the process of nitrification (vide infra).

Aqueous ammonia (liquor ammonise), being in constant requisition as a reagent, and also used in medicine and in the arts, is being manufactured industrially. Fresenius recommends the following process. A cast-iron pot, fitted up as a retort, is charged with alternate layers of slaked lime (10 kilos of quicklime plus 4 kilos of water) and a powdered and sifted mixture of 6'5 kilos of chloride and 3'5 kilos of sulphate of ammonium. Eight litres of water are then added and well incorporated with the solids. The retort is now closed, and the outlet-tube joined on to the lower end of an inverted condenser, the upper end of which communicates with a set of Woulfe's bottles charged with water. The ammonia is driven off by judicious application of heat, the inverted condenser serving to make the gas relatively dry before being absorbed. In this operation the tubes conveying the gas to the water must go to the bottom of the bottles, as the solution produced is lighter than pure water; and, of course, the bottles must be kept cool by immersion in a cold water-bath. The strength of aqueous ammonia, for commercial purposes, is readily ascer-tained by means of an hydrometer ranging from 0'85 to TO specific gravity. The relation between specific gravity S at 14° C. (water of 14° = 1) and percentage of ammonia NH3 (p), according to experiments by Carius, is as follows:—

p s s V S
36 35 34 33 32 31 30 _8844 •8856 •8868 •8907 •8929 •8953 •8976 25 20 15 14 13 12 11 •9106 •9251 •9414 _9449 •9484 •9520 _9556 10
9 8 7 6 5 _9593 •9631 •9670 •9709 •9749 •9790

Aqueous ammonia is generally sent out as "liquor fortissimus " of 30 to 35 per cent. For ordinary labora-tory purposes it is usually diluted down to 9 to 10 per cent. This explains the discontinuities in the table.

Two natural sources of ammonia are at present in industrial use. (1) The gaseous exhalations of volcanoes always include ammonia, hydrochloric acid, and sulphurous acid, of which the first two are formed no doubt by the action of steam on deposits of nitrides and chlorides in the interior of the earth. This explains the existence in the fumaroles (smoke-holes), and in the clefts of the lava of Vesuvius, Hecla, and other volcanoes, of deposits of (chiefly) sal-ammoniac. This volcanic ammonia salt is highly valued as a material for the preparation of pure liquor ammoniee; but its supply hardly comes up to the demands of even this small industry. (2) More important are the masses of ammonia formed in the processes of putrefaction which are going on constantly in nature, and of which a mere fraction would satisfy all the demands of industry, if the recovery of such ammonia were not, as a rule, beset with insuperable difficulties. Thus, for instance, all the immense mass of the ammonia of the sewage of our large cities must be allowed to go to waste because we have no economical method for its extraction. Urine, when undiluted, is an easily handled raw material, and in former times actually formed the principal source of ammonia. Human urine contains from 2 to 3 per cent, of urea, or carbamide as it is called in systematic chemistry, because it is the anhydride of carbonate of ammonia. When urine putrefies, this carbamide takes up the ele-ments of water and becomes carbonate of ammonia, CO . (N2H4) +-2H20 = C03(NH4)2. A prompter mode of conversion is to evaporate the urine with a small propor-tion of vitriol, and heat the residue to near the boiling-point of the acid, when the nitrogen of the urea passes at once into the form of sulphate of ammonia. This latter process would apply also to the urine of horses and cattle, which, instead of urea, contains hippuric acid, a compound which, when taken conjointly with water, contains the elements of ammonia and benzoic and acetic acids. At the present time urine plays hardly any part in the ammonia industry; but it may be mentioned that the produce of the urinals of Glasgow is, or lately was, wrought for car-bonate of ammonia.

Preparation of Ammonia from Nitrogenous Organic Matter.— All such matter when subjected to dry distillation gives up part of-its nitrogen as ammonia, of which the greater part condenses with the vapour of water produced, and is thus obtained as part of the aqueous portion of the "tar." Large quantities of such tar-water are being produced incidentally in the manufacture of coal-gas, and it is this material which at present forms the principal source for the industrial production of ammonia and ammonia salts.

It may also be mentioned that the tar-water obtained as a bye-product in the distillation of shale for the production of paraffin oil is rich in ammonia, and has long come to be worked up for sulphate like gas-liquors.

Crude tar-water contains about 1 per cent, of ammonia (more or less according to the quality of coal used, and the way it has been manipulated), mostly in the form of carbonate, part as cyanide, sulphocyanate, and sulphide of ammonium; and this ammonia, of course, is associated with traces of hydrocarbons and other organic matter dissolved, or suspended, in the liquor. In some establish-ments the ammonia is extracted directly in the form of liquor ammonise. The liquor is run into a large iron boiler, and after . addition of some ferrous and ferric salt (for fixing the sulphur of the sulphide as FeS, and the cyanogen of the cyanide as prussiate, (NC)6Fe. (NH4)4), mixed with slaked lime and distilled. The vapours if condensed as they come off, would yield a very dilute liquor contaminated largely with volatile carbon compounds. To obtain a relatively pure gas, the vapour is subjected to a succession of partial condensations by making it pass through the several com-partments of an iron apparatus, similar in its action to the "Coffey's still" which is used for the strengthening and refining of alcoholic liquors (see DISTILLATION, vol. vii. p. 265). The almost pure gas which leaves the last condenser is passed into a mass of water contained in a refrigerated close lead vessel, and thus converted into liquor ammonias of the requisite strength. In the majority of large establishments, however, the ammonia is converted into sulphate, for which purpose it need not be so elaborately purified. The ammoniacal vapour obtained from the crude liquor by simple distillation, or by distillation with lime from out of a steam-boiler, is passed into a quantity of chamber-acid con-tained in a leaden tank, until the acid is almost but not quite neutralized. An excess of alkali would induce the formation of sulphide of iron from the sulphide of ammonium in the vapour and the traces of iron-salt adventitiously present, and lead to a discolora-tion of the salt. Matters are generally arranged so that the bulk of the sulphate formed crystallizes out, on cooling, in the form of a granular magma, which is allowed to drain and to dry, and either sold as it is, or first purified by recrystallization. The pure salt, (NH4)2S04, forms anhydrous, colourless, transparent crystals, iso-morphous with those of the corresponding potash salt, and, like it, belonging to the rhombic system ; even when produced on a large scale they are generally of small dimensions ; when allowed to grow, they assume forms which strikingly remind one of sugar-candy, although the latter is clino-rhombic. The salt is insoluble in alcohol, like most sulphates,—100 parts of water at 0°, 20°, 100° C. dissolve 71'0, 76'3, 97'5 parts of salt. The solution is neutral to litmus. The salt readily unites with another equivalent of sul-phuric acid into crystallizable acid sulphate, S04. (N H4). H, soluble in alcohol. The neutral salt melts at 140° C. Above 280° it emits ammonia and leaves acid salt, which latter then breaks up with formation of acid sulphite S03(NH4)H1, nitrogen, and water. At a red heat it breaks up into sulphur, nitrogen, and water: (NH4)„S04 = S + N2+4H20.

Of other ammonia salts only the hydrochloride and the carbon-ates are industrially important.

The hydrochloride, HC1. NH3 = NH4C1, better known as sal-ammoniac (see vol. i. p. 741-2), is made sometimes by sublimation of a mixture of the sulphate with common salt, but it is more con-veniently produced direct from gas-liquor ammonia by passing it into muriatic acid until the latter is almost neutralized. The liquor, when sufficiently concentrated by evaporation, deposits, on cooling, part of its salt in feathery crystals, which are customarily purified by sublimation. The subliming apparatus consists of two parts, —(1) a hemispherical stoueware basin placed within a close fitting iron one, or an enamelled iron basin, and (2) a hemispherical lead or stoneware lid, or dome, placed on the top of the basin and cemented on to prevent leakage. The dome has a small aperture in the top which remains open to preclude accumulation of pressure. The carefully dried crystallized salt is pressed into the basin, and, after the lid has been fitted on, is exposed to a long-lasting moderate heat. The salt volatilizes (mostly in the form of a mixed vapour of the two components, which reunite on cooling), and condenses in the dome in the form of a characteristically fibrous and tough crust, The salt readily dissolves in water, with consider-able absorption of heat; 30 parts of salt with 100 parts of water at 13°'3 give a mixture of the temperature of - 5°T C. One hundred parts of water at 0°, 10°, 110° dissolve 28'4, 32'8, 77'2 parts of the salt. From its hot saturated solution it crystallizes on cooling in feathery groups of colourless needles. By slow evaporation of the solution it is possible to produce well-developed crystals which belong to the regular system, but look irregular on account of the predominance of the (hemihedric) faces of the trapezohedron.
CO
C02 + 2NH3





Of the carbonates of ammonia there are a large number, and their chemistry still lacks definiteness. The normal salt C03(NH4)2 is so unstable that it can hardly be said to exist. The acid salt C03(NH4)H is easily produced by passing carbonic acid into a saturated solution of the commercial salt, when it comes down as a crystalline precipitate. The commercial salt (important as a medicinal agent and as a chemical reagent) is obtained by subliming a mixture of sal-ammoniac and chalk from an iron retort pro-vided with a lead dome and receiver. It forms hard fibrous crusts or cakes, smelling strongly of ammonia. The salt has a variable composition. The greater part, as a rule, consists of " sesqui-carbonate," 2(NH4).,0 . 3C02 + H20 = (NH4)2C03 + 2(NH4) . HC03. But it also contains carbamate of ammonia,
NH3
0-(NH4)-
as obtainable by the direct union of carbonic anhydride and ammonia.

Of the several ammonia compounds which we have referred to, the sulphate is by far the most important in an industrial sense. Immense quantities of the crude salt are being used as a manure— the German sugar-beet growers alone consume a considerable fraction of the British produce—while to the technical chemist generally it serves as the most convenient starting-point for the manufacture of ammonia, or of other ammonia salts.

In the mode of distilling coal customarily carried on in gas works, only about one-third of the nitrogen is obtained as ammonia in the tar-water, the remaining two-thirds being lost by evaporation into the air, or remaining in the coke in the carbide form. What used to go into the gas is now mostly recovered by efficient scrubbers. But the more efficient condensation of the ammonia actually formed is a matter of chemical engineering which cannot be more than touched on here. According to Bilbey the nitrogen of the coke can be recovered, partly at least, by distilling it at a very high temperature in a current of steam. Bilbey's process, however, has hitherto failed practically to give satisfaction, because the intense heat required means a great expense for fuel, and destroys the retorts at an alarming rate. The analytical chemist has no difficulty in extracting the whole of the nitrogen in a given sample of coal as ammonia by mixing it with soda-lime and heating the mixture in a combustion-tube to redness, and possibly the technical chemist will one day bring this process into a remunera-tive form. What, however, is meanwhile more easy of attainment is the recovery of the large quantities of ammonia which are being produced in the manufacture of coke and in iron smelting (as far as carried out with coal), and which hitherto have been allowed to go to waste. Quite a number of chemists and engineers have tried their hands at this problem. The apparatus proposed, generally speaking, all come to this, that the coal-smoke produced in the furnace, instead of being allowed to have its own way, is sucked out by exhausters, made to pass through refrigerators to deposit at least part of its tar and ammonia water, and the uncondensed combustible gases are led away to be used as fuel for steam-boilers, or, what in the case of coking is far better, led back to the coke oven and consumed there to increase the temperature, and thus improve the qualities of both coke and tar. The very high tem-perature of the oven or furnace smoke throws great difficulties in the way of a perfect condensation of the ammonia. These, in a Scottish iron-work, have been turned most ingeniously by mixing the smoke with the sulphureous vapour formed by the roasting of pyritic shale or coal, whereby the ammonia is converted into sulphite and sulphate, which can easily be condensed in even hot water.

Should it not be possible to produce ammonia synthetically from atmospheric nitrogen ? This question is still waiting for an industrial solution ; scientifically it may be answered in more than one way. Magnesium, boron, and a number of other solid and non-volatile elementary substances, when kept in nitrogen gas at the proper temperature, unite with the nitrogen into solid non-volatile nitrides ; and these, when heated in steam, yield ammonia and the corresponding oxide. Thus we have 3Mg + N>=N„Mg3 and N2Mg3-r3H20 = 2NH3 + 3MgO. Unfortunately the" reagents are all expensive, and there is no economical method for their regeneration. The following method is not subject to this one objection. A mixture of baryta or carbonate of baryta with char-coal, when heated intensely in nitrogen gas, yields cyanide of barium, Ba(NC)2, and this salt when heated in steam gives off ammonia while carbonate of baryta is left, which latter can be used for starting de novo, BaN2C2 + 4H20 = 2NH3 + BaC03 + CO + H2. Where this process fails we are unable to say; what we do know is that nobody has as yet succeeded in working it profitably even as a means for obtaining cyanides, whose value, per unit of nitrogen, is higher than that of ammonia.

Nitrates.

Nitrates (the generic term for nitric acid, HN03, and its salts) are produced naturally by the electric discharges in the atmo-sphere, and in the processes of "nitrification," a fermentative oxidation which always sets in when moist nitrogenous animal or vegetable matter is left to itself in the presence of air and some basic substance (see FERMENTATION, vol. ix. p. 98). This process in former times used to bo carried out for the production of saltpetre, but as an industrial operation is now obsolete. The deposits of native nitre in India and elsewhere which nature has produced for us by the same method are, of course, still being utilized as far as they go. But they amount to very little compared with the immense masses of native nitrate of soda which exist in South America, and which, at present, constitute by far the most im-portant raw material for the nitrate industry.

This native nitrate of soda forms part of a salty earth known to the natives as caliche or terra salitrosa, which abounds especially in the district of Atacama and the Peruvian province of Tarapaca. The caliche there lies from '25 to 1 '5 metres deep, and stretches over a distance of forty leagues ; it is covered by a layer, from a half to two metres thick, of a hard conglomerate of sand, felspar, phosphates, and other mineral matters, which is designated "iostra." In other places the caliche forms part of a sandy deposit which some-times comes to the surface and never goes down to a depth beyond 2'6 metres. The caliche contains from 48 to 75 per cent, of nitrate of soda and from 20 to 40 per cent, of common salt, which are associated with various minor saline components, including iodate of soda, and more or less of insoluble mineral, and also some organic matter, guano amongst other things, which suggests the idea that the nitrate was formed by the nitrification of this kind of excre-mental matter. The caliche is worked up in loco for crude nitrate of soda ; by extracting the salts with hot water, allowing the suspended earth to settle, and then transferring the clarified liquor, first to a cistern where it deposits part of its chloride of sodium at a high temperature, and then to another where, on cooling, it yields a crop of crystals of purified nitrate, The nitre thus refined is imported chiefly from Valparaiso, whence the name of "Chili saltpetre." The mother-liquors used until a few years ago to be thrown away, but are now being utilized for the extraction of their iodine, which, although little in a relative sense, on account of the large masses of raw material wrought, amounts to a good deal absolutely, as is illustrated by the fact that Peruvian iodine has put an end to the kelp industry in Scotland. Chemically pure nitrate of soda can be obtained by repeated recrystallization of Chili saltpetre or by synthesis from pure nitric acid and pure car-bonate of soda. It forms colourless transparent rhombohedra, like those of Iceland spar, only the angles are more nearly equal to right angles, so that the crystals look hke cubes. Hence the name of cubic saltpetre, which is sometimes given to the salt. One hundred parts of water at 0°, 20°, 50°, 100°, 110° C. dissolve 72 9, 87-5, 112, 180, 200 parts of the salt; at 120°, the boiling-point of the saturated solution, 216 parts. It fuses at 330° C. (Carnelley); at higher temperatures it loses oxygen (more readily than the corresponding potash salt) with formation of nitrite N02Na, which, at very high temperatures, is reduced ultimately to a mixture of peroxide, Na202, and oxide, Na20. Industrially the salt is important as being the raw material for the manufacture of nitric acid and of nitrate of potash.

Nitrate of potash (saltpetre), which forms the predominating component of gunpowder, occurs native in India and other parts of the world, and such native nitre has only to be purified by crystallization to become fit for the market. But the bulk of what occurs in commerce is made by double decomposition of Chili saltpetre with (a) caustic potash, If) carbonate of potash, or (c) chloride of potassium, which processes yield (ÍÍ) caustic soda, (b) carbonate of soda, (c) common salt as bye-products. The third form (c) of the method is most largely wrought, the necessary supplies of cheap chloride of potassium being furnished by the works at the Stassfurt deposits. The two raw materials are analysed, and quantities corresponding to NaN03 = 85 of nitrate of soda and KC1 = 74 -5 parts of chloride of potassium respectively are dissolved together in an iron basin in the least quantity of hot water. The solution is boiled down to 1'5 specific gravity (hot), when the common salt formed gradually crystallizes out. It is fished out, allowed to drain, and the runnings are returned to the basin. The highly concentrated mother-liquor is allowed to cool with frequent agitation, so that the saltpetre, which crystallizes and assumes the form of a "meal," is more readily freed from mother-liquor, by judicious washing with cold water. The crude saltpetre which is thus obtained is recrystallized until it is almost chemically pure, because an even slightly contaminated salt is unfit for gunpowder making. Nitrate of potash is isomorphous with the soda salt in this sense that it is possible to obtain it in the form of rhombohedra ; but these rhombohedra have no great stability, and under ordinary conditions the salt always assumes the form of long six-sided prisms of the right-rhombic system. It is far less soluble in water than nitrate of soda, and, unlike it, is absolutely non-hygroscopic. One hundred parts of water at 0°, 10°, 20°, 50°, 80°, 100° C. dissolve 13-3, 211, 31-2, 86, 172, 247 parts of salt. The boiling fully saturated solution is at 114°, and contains 327 parts of salt per 100 of water (Mulder). It fuses at 352° C. (Carnelley) into a colourless liquid, which freezes into a hard compact mass exhibiting a coarse, fibrous fracture if the salt is pure. In an impure salt this structure be-comes the less distinct the greater the proportion of impurities; with a saltpetre which contains 3 '5 per cent, of common salt, the fibres appear only at the edges of the surface of fracture. This is the basis of a now almost forgotten test for purity, and explains that to the present day the term '' refraction " is used to designate the sum total of impurities contained in 100 parts of nitre analysed. When heated sufficiently beyond its fusing point it decomposes similarly to nitrate of soda, only it gives up its oxygen far less readily. A mixture of nitre with charcoal, sulphur, or other com-bustible matter, when kindled, burns off with explosive violence.

Hence its application for the manufacture of gunpowder and in pyrotechnics, and its use in the laboratory as a powerful oxidizing agent in operations of the dry way.

Nitric Acid, HN03 (see CHEMISTRY, p. 5-11 sq.), is prepared from nitrate of potash or soda by distillation with sulphuric acid. The scientific chemist prefers the potash salt because it is more easily purified ; the manufacturer uses nitrate of soda because it is cheaper and a lower temperature and a less excess of oil of vitriol (over and above 1H2S04 for 2RN03) suffice for its successful con-version into acid. For manufacturing purposes the distillation is effected from out of horizontal cylindrical or deep hemispherical retorts made of cast iron (a material which is far less attacked by the acids than one might be inclined to think). The retort com-municates with a series of Woulfe's bottles made of stoneware, which are sometimes provided with taps for letting off the distillate. The retort, after having been charged and connected with the receivers, is heated over a naked fire until all the acid is driven off, and nothing but a (more or less acid) sulphate of soda is left. The details of the manufacture vary according to the kind of acid which is intended to be produced. If the manufacture of ordinary aqua-fortis (of 1-3 to 1 '4 specific gravity) is aimed at the receivers are charged with the proper-proportion of water, and less than 1H2S04 may be used for the decomposition of lN03Na, because the per-oxide of nitrogen from the dissociated part of the acid, by the action of the water and the air in the receivers, is to a great extent recon-verted into nitric acid. The preparation of acid of highest strength demands the full equivalent of vitriol. But even then the product, apart from the contents of the first few condenser bottles nearest to the retort, is strongly contaminated with dissolved peroxide, which imparts to the acid a deep red colour. This impurity can be brought down to about 2 per cent, by blowing air through the gently heated acid, which carries away the peroxide as a vapour. To utilize the latter, the mixture is made to ascend through a tower filled with coke kept moist by a constant rain of water running downwards against the stream of acid fumes. Water alone would convert two-thirds of the nitrogen into nitric acid, thus:—3N204 (+ aqua) = 2NO + (2N206 + aqua); but by the cooperation of the oxygen of the air a larger proportion of acid is being regenerated as aquafortis of about 1'3 specific gravity. Two kinds of acid are chiefly being produced, viz., (1) full-strength fuming acid of l-5 to

1 '52 specific gravity at 15° C., which is largely used (in the tar-colour industry for the making of nitro-bodies, and also for the manu-facture of gun-cotton and nitroglycerin), and (2) aquafortis of 1 '35 to 1'42 specific gravity, for the charging of batteries, and as a re-agent generally. Either acid is sold in two qualities,—as colour-less acid relatively free of peroxide of nitrogen, and as red acid charged with it, which admixture is not always an impurity in a technical sense, but desired for certain purposes to be present. Most commercial acid is contaminated with chlorine and sulphuric acid, which can both be removed by fractional distillation—most easily from the strongest acid. The chlorine accumulates in the first runnings ; the sulphuric acid remains in the retort, if the dis-tillation be stopped in time.

For most purposes the strength of an acid can be determined with sufficient accuracy by means of a hydrometer.

The following table gives the relation between the percentages p of HN03 and the specific gravity S according to Kolb:—

== TABLE ==

Regarding the reactions of nitric acid with elementary substances and with inorganic compounds generally, see CHEMISTRY (ul sup. pp. 511-514). The general effect of the treatment of an organic substance with aquafortis is that the acid is reduced to lower oxides of nitrogen and water, while the organic substance undergoes some kind of oxidation, which sometimes assumes the character of a com-bustion. With regard to the more immediately formed oxidation products it is difficult to speak in general terms; suffice it to say that, according to Carius, all organic substances when heated with nitric acid in sealed-up tubes to 150° to 300° C. are burned com-pletely into carbonic acid, water, and nitrogen or nitrogen oxides. Any other element that may be present passes into its highest state of oxidation: sulphur, for instance, becomes sulphuric acid; phos-phorus, phosphoric, &c. Very remarkable is the action of real nitric acid, when employed in the cold, and under conditions pre-cluding the accumulation of water or low-er oxides of nitrogen. The acid then, if it acts at all, unites with the respective substances, with elimination of water and formation from the body B of a pro-duct B + nHN03-nH20 or B + nN"02-nH. These nitro-bodies, as they are sometimes called, divide themselves into two classes:— (1) nitric esters, i.e., real nitrates formed from alcohols, which term must be understood here to include glycerin, sugar, mannite, cellu-lose, and many other OH-compounds not usually called alcohols (in these the group N02. 0 stands where the OH was in the original substance); and (2) nitro-bodies proper, in which the N02-group has taken the place of an H which was combined directly with carbon. An example is the nitrobenzol, C6H5(N02), produced from benzol, C6H5. H. When such a (true) nitro-body is treated with nascent hydrogen, the N02 as a rule goes out and is replaced by an amidogen NH2; thus C6H5(N02) becomes C6H5(NH2), amido-benzol or aniline. The nitric esters, in the circumstances, yield up their N02 as ammonia.
For other metallic nitrates, see under the respective metals. For nitric esters, see METHYL, GUN-COTTON, and NITROGLYCERIN. For nitric anhydride (N205), peroxide of nitrogen (N204), nitrous acid (Na03), nitric oxide (NO), nitrous oxide (N20, laughing-gas), also chloride of nitrosyle (aqua regia), see CHEMISTRY, p. 511 sq.

Nitrogenous Carbon Compounds.

Of this very numerous family we can only name a few of the more important groups and explain their genetic relations. Cyanides, compounds of the radical (NC), are important to us here chiefly as a link between the two elements on the one hand and nitrogenous organic bodies proper on the other. Hydrocyanic acid, NCH, can be produced synthetically in a number of ways ; we may, for instance, synthesize cyanide of barium, BaN2C2, as shown above, and decompose it by sulphuric acid. A more direct method is to first combine carbon with hydrogen into acetylene and then to subject a mixture of this and nitrogen gas to a spark-current, when hydrocyanic acid is produced, thus:—C2H2 + N2 = 2HNC.

Alkylamines.—When hydrogen is generated within an acidified solution of hydrocyanic acid by means of zinc, the cyanide, by taking over four atoms of hydrogen, passes into (a salt of) methyl-amine, NCH + 4H = NH2 .CH3. It is explained under METHYL how this base may be utilized for the production' of methyl-alcohol, and thus, indirectly, of iodide of methyl and of acetonitrile (pseudo-cyanide of methyl, NC. CH3). This shows the possibility of producing from hydrocyanic acid, in the first instance di- and tri-methylamine, and, more indirectly, the whole series of alkyla-mines (NH2R, NHR2, NR3; where R = CH3, CH3 + CH2 = C2H5, C2H5 + CH2 = C3H7, &c). Closely allied to these alkylamines are the diamines, derived from the bromides of the defines, C„H2„, as the former are from those, or the iodides, of the alcohol-radicals CH3, C2H5, &c. To ethylamine, for instance, corresponds ethy-lene-diamine, C2H4(NH2)2.

Acid-amides are bodies related to acids, as alkylamines are to alcohols. Thus, for instance,
C2H5NH2 corresponds to [CH3.CO]NH2.
Ethylamine. Acetamide.
p ri /COOH , rlT/C
(-2RL4SsC0NH„ L2H4<VC
Polybasic acids form each a series of amides. Thus succinic acid, C2H4(COOH)2, forms two amides, viz.,
'CONH2
^CONH,
Succinamic acid. Succinamide.

All acid-amides are the anhydrides of corresponding ammonia salts, and can be produced from these by actual dehydration ; there is no need of explaining by an equation what is so clearly seen in the formula;. The two succinamides just named, and all analogous "amides," are susceptible of further iehydration: thus
succinamic acid, C.
JX)NH[H]' by losing the bracketed OH and H becomes
C|OjN|jy C|OjN|H2|'
succinimid, C2H4^pQ^>(NH) ;
and succinamide, C2H4
PTT /CN
similarly becomes
suceino-nitrile.
Acid-amides must not be mixed up with amido-acids.
Amido-acids have their (NH2) within the specific, not in the generic, radical. Thus, acetic acid, CH3COOH, by chlorination, becomes chloracetic acid, CH2C1. COOH, and this chloro-acid again gives up its chlorine to ammonia and to caustic potash, and receives back NH2 in the one case and OH in the other, becoming
CH2(NH2). COOH and CH2(OH). COOH
Amido-acetic acid. Oxy-acetic or glycollic acid.
respectively. The latter forms an amide, CH2(OH)CO. NH2, as acetic acid does, and this amide, as is easily seen, is isomeric with amido-acetic acid, either being oxy-acetic minus OH plus NH2; but
P TT COOH 2 4COOH
Succinic acid.
the two differ widely from each other. The amide, when treated with aqueous potash, yields ammonia and glycollate of potash, the amido-acid water and amido-acetate of potash. The substance asparagine, which is contained in so many vegetable juices, and often appears as a decomposition product of albumenoids, is an amide and an amido-acid in one, being amido-succinamic acid.
C2H3(NH2) jCO;NH2_
Asparagine.
With the so-called aromatic bodies a very general method for producing amido-bodies is first to prepare a nitro-body and then reduce it by nascent hydrogen. Thus, for instance, we can convert nitrobenzol, C6H5. N02, above referred to, into amido-benzol, C6H5. NH2. We shall use this body as an example to explain the general mode of formation of a most interesting group of nitro-genous compounds which at present play a great part in the colour industry.





Diazo-bodics.—All NH2 compounds, when treated with water and nitrous acid, virtually NO. OH = |(N203 + H20), yield the corresponding hydroxyl (OH) compounds, with joint elimination of the nitrogen of substance and reagent as nitrogen gas. Thus we have—
C6H,NH2 + OH. NO =_ C,HB. OH + H20 + N2.
Aniline. Phenol.
-CfiH6-N=N-N03.
Nitrate of diazo-benzol.
This is quite a general reaction ; but in the case of aromatic amines the reaction takes a different turn, if, instead of the free amine, we use one of its salts, and, in a relative sense at least, exclude water as a medium. In the case of nitrate of aniline, for instance, the two atoms of oxygen in the nitrous acid combine with all the hydrogen of the generic radicals, while all the rest unites into diazo-benzol:—
HNO,. CBH,. NH„ + NO, OH = 2H„0
And similarly in similar cases. Diazo-benzol has a great tendency to give off the nitrogen gas N2 which is visible in its formula. We need only treat it with water and it yields what nitrous acid water would have given at once:—
C6H5—N=N—N03 + H. OH =_ HN03 + N2 + C6H6. OH.
There is something between the nitrates of amido- and of diazo-benzol, which, it is true, has only a theoretical existence in this case, but can be realized in other analogous cases.
N. H. (C2H5)2 + NO. OH

Nitroso-bodies.—Ethylamine, under the action of water and nitrous acid, yields alcohol, C2H5. OH, as aniline yields phenol, and other monamines yield their alcohols. With diethylamine this reaction cannot take place; what does take place is that the NO of the nitrous acid expels and takes the place of the one loose H in the amine, and forms nitroso-diethyline:—
_ H. OH + N. (NO) . (C2H,)a Nitroso-diethyline.

Hydrazines.—This nitroso-diethyline is a liquid boiling at 177°, devoid of alkaline properties. When treated with zinc-dust and acetic acid it yields a basic substance diethyl-hydrazine, which contains NH2 in lieu of the NO of the nitroso-body.
N. (NO)(C2H5)2 N. (NH2)(C2H5)2.
Nitroso-diethyline. Diethyl-hydrazine.
This hydrazine is only one of a large genus, representatives of which are also obtainable from diazo-bodies by nascent hydrogen. Example—
CBH5—N=N—(HS04) + 4H - H2S04 + C„H6—NH. (NH2).
Sulphate of diazo-benzol. Phenyl-hydrazine.
(NHS)'
NH, NH:!
TJreids are a class of bodies which are related to urea as amido-bodies are to ammonia. Thus
(2NH2)"
CO -
CO
correspond to
(NH)'1
co<Œ
NH,
NH2
Urea.
(NH)' NH2
Urea radicals.
NH,
Oxamide.

We satisfy ourselves with quoting two ureids derived from oxalic acid,C2O.X^Q^. Just as we have two ammonia derivatives,
and C202<^-2
Oxamic acid.
A<[NIIH-
Parabanic acid.
so there exist two urea-derivatives
<
OH [NH. NH2CO] and °:
Oxaluric acid.
The urea radicals are enclosed in square brackets. These are two examples of a large family of bodies interesting on account of their close relationship to uric acid, a constant component of urine. Creatine is a crystalline base which was discovered by Liebig in the juice of flesh. Volhard prepared it synthetically by the successive realization of the following two reactions :—
(1) CHaCl. C00H + NH2CH3=HC1 + CH2. NH(CH3). COOH ;
Chloracetic acid. Methylamine. Methyl-amido-acetic acid.
(sarkosine).
(2) CH2. NH(CH3). COOH + (CN)KH2 =
Sarkosine. Cyanamide.
WH_0 J NH2
JNH—O j _ÑRÑ~,). [CH2. ____].
Creatine.
Betaine and Neurine, two bases derived from trimethylamine, are of great physiological importance. Neurine (CH3)3-N-(OH)
CH. CH2
was discovered by Liebreich as one of the congeners of a complex substance contained in the brain. By oxidation it is converted into betaine
(CH3)3-N
CH2. CO. 0
which Scheibler discovered in the juice of the sugar-beet. Both bases can be produced synthetically,—the betaine by first unit-ing trimethylamine with chloracetic acid into a chloride, C1CH2.C00H + N(CH3)3, and then replacing the CI by OH by means of oxide of silver and water. The hydroxide
N-[OH]<|^bo[H]
formed spontaneously loses the bracketed [OH] and [H] as water, and becomes trimethyl-glycocoll, which is betaine.
Native Alkaloids.—These may be divided into (1) bodies consisting of carbon, nitrogen, and hydrogen only, generally volatile liquids (of which nicotine may be quoted as an example); and (2) such as contain oxygen in addition to the three elements named (quinine, morphia, strychnine, &c). The more important of these bases are noticed in separate articles.
The Albumenoids. —This class comprises those substances known as albumin, fibrin, casein, &c, which, conjointly with carbo-hydrates and fats respectively may be said to form the basis of vegetable and animal life. They consequently are the most important of all nitrogenous carbon compounds; but unfortunately we know little of their chemical constitution. They are all solids containing 53 to 547 per cent, of carbon, 7-l to 7'2 of hydrogen, 15'6 to 15'8 of nitrogen, and 1'7 to 1'8 of sulphur. Glutin and chondrin (glue) are closely allied to them. See the special articles.
Analysis.

In regard to general methods for the determination (or detection) of the element, see CHEMISTRY, vol. v. p. 546-7. But it may be added that the method there given as that of Dumas, for combustible carbon compounds, can be made to apply to metallic nitrates by simply substituting finely divided metallic copper for the oxide of copper as a burning agent, and that Varrentrapp and Will's method applies\ as it stands to metallic nitrides and amides, to metallic cyanides, and to all metal-amines. We confine ourselves here to the more important methods for the detection of certain classes of nitrogen compounds.

1. Ammonia-salts and most acid-amides and amin-acids, when treated with aqueous potash, give off their nitrogen as ammonia gas, which, if sufficient in quantity, is easily detected by (a) its pungent smell, (b) its action on turmeric or red-litmus paper, (c) its forming dense clouds of sal-ammoniac when brought into contact with a glass rod moistened with muriatic acid. Should the vapours of ammonia be too highly attenuated to be thus identified, condense them in dilute muriatic acid, and to the distillate obtained (after neutralization of the free acid by potash or soda) apply '' Nessler's reagent" (a solution of the salt HgI2.2KI in aqueous caustic potash or soda; see MERCURY, vol. xvi. p. 34). The least trace of ammonia, if present, assumes the form of iodide of mercur-ammonium, which, if too little to come down as a precipitate, will announce itself by imparting to the liquid an intense brown or yellow colour. One-hundredth of a milligram of ammonia dif-fused throughout 50 c.c. of liquid (i.e., five million times its weight) can thus be detected with unerring certainty.

2. Nitrous and Nitric Acid. —We assume the acids to be given as solutions of their alkaline salts, and may well do so because other metallic salts of either acid can easily be brought into this form by suitable operations. A solution of alkaline nitrite, when mixed with aqueous sulphuric acid, gives off brown fumes (of N203 and N204) if it is sufficiently concentrated. In more dilute solutions the liberated nitrous acid breaks up into nitric acid and nitric oxide gas, which latter forms brown fumes (of N204) as soon as it comes into contact with the air. In still more dilute solutions the liberated nitrous acid remains undecomposed, but is readily detected by a sufficiency of a strong solution of ferrous sulphate or chloride, which reagents liberate nitric oxide from it and dissolve it with formation of an inky-black solution. This is the most character-istic and a highly delicate test. But in the latter respect it is far surpassed by " Gries's reagent,"—a solution of sulphate of meta-diamido-benzol in dilute sulphuric acid. When this solution is added to even the most dilute solution of nitrous acid, a yellow coloration is developed which attains its maximum of intensity after about fifteen minutes standing. This test is fully as delicate as Nessler's for ammonia.

A nitrate solution, when mixed with aqueous sulphuric acid or that acid and ferrous salt, exhibits no visible change. But when the solution is mixed with its own volume of (concentrated) oil of vitriol, so as to not only liberate but also dehydrate the nitric acid, and a strong solution of ferrous sulphate is cautiously poured on the top of the mixture, nitric oxide is eliminated which dissolves in, and strikes an inky-black colour with, the ferrous salt layer. This is the test for nitrates.

When a solution containing nitrates or nitrites is distilled with concentrated caustic potash or soda and aluminium foil or zinc-foil and iron filings (or any metal or combination which, with the pure alkali-solution, would give off hydrogen), the nitrogen is gradually eliminated as ammonia, which can be condensed, and, however little it may amount to, detected by Nessler's reagent, as above explained. (W. D.)


Footnotes

For the ammonia process of soda-ash manufacture, which, after having been asleep industrially for some thirty years, was brought into a remunerative form bv Solvav and Mond, and is now gradually, but surely, taking the place of the old Leblanc process, see SODIUM.

See CHEMISTRY. Some of the more important cyanides will he discussed under PRUSSIC ACID.