see captionFig.48.—The method of decomposition of nitrous anhydride, also applicable to the other oxides of nitrogen, and to their analysis. NO2is generated from nitrate of lead in the retort A. Nitric acid and other less volatile products are condensed in B. The tube C C contains copper, and is heated from below. Undecomposed volatile products (if any are formed) are condensed in D, which is cooled. If the decomposition be incomplete, brown fumes make their appearance in this receiver. The gaseous nitrogen is collected in the cylinder E.
Fig.48.—The method of decomposition of nitrous anhydride, also applicable to the other oxides of nitrogen, and to their analysis. NO2is generated from nitrate of lead in the retort A. Nitric acid and other less volatile products are condensed in B. The tube C C contains copper, and is heated from below. Undecomposed volatile products (if any are formed) are condensed in D, which is cooled. If the decomposition be incomplete, brown fumes make their appearance in this receiver. The gaseous nitrogen is collected in the cylinder E.
If the vapour of nitric acid is passed through an even moderately heated glass tube, the formation of dark-brown fumes of the lower oxides of nitrogen and the separation of free oxygen may be observed, 2NHO3= H2O + 2NO2+ O. The decomposition is complete at a white heat—that is, nitrogen is formed, 2NHO3= H2O + N2+ O5. Hence it is easily understood that nitric acid may part with its oxygen to a number of substances capable of being oxidised.[39]It is consequentlyanoxidising 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 saltof 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 aciddissolvesnearly 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, SnH2O3, which is obtained in the form of a white powder, Sn + 4NHO3= H2SnO3+ 4NO2+ H2O. Silver is able to take up still more oxygen, and to convert a large portion of nitric acid into nitrous anhydride, 4Ag + 6HNO3= 4AgNO3+ N2O3+ 3H2O. Copper takes up still more oxygen from nitric acid, converting it into nitric oxide, and, by the action of zinc, nitric acid is able to give up a still further quantity of nitrogen, forming nitrous oxide, 4Zn + 10NHO3= 4Zn(NO3)2+ N2O + 5H2O.[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 isdetermined, not only by the nature of the reacting substances, but also by the relative mass of water and nitric acid, and also by the temperature and pressure, or the sum total of the conditions of reaction; and as in a given mixture even these conditions vary (the temperature and the relative mass vary), it not unfrequently happens that a mixture of different products of the deoxidation of nitric acid is formed.
Thus the action of nitric acid on metals consists in their being oxidised, whilst the acid itself is converted, according to the temperature, concentration in which it is taken, and the nature of the metal, &c., into lower oxides, ammonia, or even into nitrogen.[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, M2S, into sulphates, M2SO4, &c.; in a word, nitric acid brings about oxidation, its oxygen is taken up and transferred to many other substances. Certain substances are oxidised by strong nitric acid so rapidly and with so great an evolution of heat that they deflagrate and burst into flame. Thus turpentine, C10H16, bursts into flame when poured into fuming nitric acid. In virtue of its oxidising property, nitric acidremoves the hydrogenfrom 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 rapidreaction takes place, accompanied by flame and the separation of violet vapours of iodine and brown fumes of oxides of nitrogen.[43]
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 correspondingnitric anhydride, N2O5; but Deville first and subsequently Weber and others, discovered the methods of its formation. Deville obtained nitric anhydride by decomposing silver nitrate by chlorine under the influence of a moderate heat. Chlorine acts on the above salt at a temperature of 95° (2AgNO3+ Cl2= 2AgCl + N2O5+ O), and when once the reaction is started, it continues by itself without further heating. Brown fumes are given off, which are condensed in a tube surrounded by a freezing-mixture. A portion condenses in this tube and a portion remains in a gaseous state. The latter contains free oxygen. A crystalline mass and a liquid substance are obtained in the tube; the liquid is poured off, and a current of dry carbonic acid gas is passed through the apparatus in order to remove all traces of volatile substances (liquid oxides of nitrogen) adhering to the crystals of nitric anhydride. These form a voluminous mass of rhombic crystals (density 1·64), which sometimes are of rather large size; they melt at about 30° and distil at about 47°. In distilling, a portion of the substance is decomposed. With water these crystals give nitric acid. Nitric anhydride is also obtained by the action of phosphoric anhydride, P2O5, on cold pure nitric acid (below 0°). During the very careful distillation of equal parts by weight of these two substances a portion of the acid decomposes, giving a liquid compound, H2O,2N2O5= N2O5,2HNO3, whilst the greater part of the nitric acid gives the anhydride according to the equation 2NHO3+ P2O5= 2PHO3+ N2O5. On heating, nitric anhydride decomposes with an explosion, or gradually, into nitric peroxide and oxygen, N2O5= N2O4+ O.
Nitrogen peroxide, N2O4, andnitrogen dioxide, NO2, express oneand 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, O3loses heat in passing into O2, whilst N2O4absorbs heat in forming NO2.
Nitric acid in acting on tin and on many organic substances (for example, starch) gives brown vapours, consisting of a mixture of N2O3and NO2. A purer product is obtained by the decomposition of lead nitrate by heat, Pb(NO3)2= 2NO2+ O + PbO, when non-volatile lead oxide, oxygen gas, and nitrogen peroxide are formed. The latter condenses, in a well-cooled vessel, to a brown liquid, which boils at about 22°. The purest peroxide of nitrogen, solidifying at -9°, is obtained by mixing dry oxygen in a freezing-mixture with twice its volume of dry nitric oxide, NO, when transparent prisms of nitrogen peroxide are formed in the receiver: they melt into a colourless liquid at about -10°. When the temperature of the receiver is above -9°, the crystals melt,[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.
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 thefollowing chapter, the density 23 corresponds with the compound NO2(because the weight corresponding with this molecular formula = 46, and the density referred to hydrogen as unity is equal to half the molecular weight); therefore at temperatures above 135° the existence of nitrogen dioxide only must be recognised. It is this gas which is of a brown colour. At a lower temperature it formsnitrogen peroxide, N2O4, whose molecular weight, and therefore density, is twice that of the dioxide. This substance, which is isomeric with nitrogen dioxide, as ozone is isomeric with oxygen, and has twice as great a vapour density (46 referred to hydrogen), is formed in greater quantity the lower the temperature, and crystallises at -10°. The reasons both of the variation of the colour of the gas (N2O4gives colourless and transparent vapours, whilst those of NO2are 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 N2O4with 21 parts by weight of NO2.[45]It is evident that a decomposition here takes place the peculiarity of which consists in the fact that the product of decomposition, NO2, is polymerised (i.e.becomes denser, combines with itself) at a lower temperature; that is, the reaction
N2O4= NO2+ NO2
is a reversible reaction, and consequently the whole phenomenon represents adissociationin a homogeneous gaseous medium, where the original substance, N2O4, and the resultant, NO2, are both gases. Themeasure of dissociationwill 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 N2O4is not then decomposable. Consequently the limits of dissociation here are -10° and 135° at the atmospheric pressure.[46]Within the limits of these temperaturesthe 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° N2O4has ceased to exist and NO2alone remains. It is evident that at the ordinary temperature there is a partially dissociated system or mixture of nitrogen peroxide, N2O4, and nitrogen dioxide, NO2. In the brown liquid boiling at 22° probably a portion of the N2O4has already passed into NO2, and it is only the colourless liquid and crystalline substance at -10° that can be considered as pure nitrogen peroxide.[47]
The above explains the action of nitrogen peroxide on water at low temperatures. N2O4then acts on water like a mixture of the anhydrides of nitrous and nitric acids. The first, N2O3, may be looked on as water in which each of the two atoms of hydrogen is replaced by the radicle NO, while in the second each hydrogen is replaced by the radicle NO2, proper to nitric acid; and in nitrogen peroxide one atom of the hydrogen of water is replaced by NO and the other by NO2, as is seen from the formulæ—
In fact, nitrogen peroxide at low temperatures gives with water (ice) both nitric, HNO3, and nitrous, HNO2, acids. The latter, as we shall afterwards see, splits up into water and the anhydride, N2O3. If, however, warm water act on nitrogen peroxide, only nitric acid and monoxide of nitrogen are formed: 3NO2+ H2O = NO + 2NHO3.
Although NO2is not decomposed into N and O even at 500°, still in many cases it acts as an oxidising agent. Thus, for instance, it oxidises mercury, converting it into mercurous nitrate, 2NO2+ Hg= HgNO3+ NO, being itself deoxidised into nitric oxide, into which the dioxide in many other instances passes, and from which it is easily formed.[48]
Nitrous anhydride, N2O3, corresponds[49]to nitrous acid, NHO2, which forms a series of salts, the nitrites—for example, the sodium salt NaNO2, the potassium salt KNO2, the ammonium salt (NH4)NO2,[50]the silver salt AgNO2,[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 NHO2from its salts, it always gives water and the anhydride, whilst the latter, as an intermediate oxide, partially or wholly splits up into NO + NO2. But the salts of nitrous acid are distinguished for their great stability. Potassium nitrate, KNO3, may be converted into potassium nitrite bydepriving it of a portion of its oxygen; for instance, by fusing it (at not too high a temperature) with metals, such as lead, KNO3+ Pb = KNO2+ PbO.[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: 2KNO2+ H2SO4= K2SO4+ N2O3+ H2O. The same gas (N2O3) 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 + NO2. Nitrous anhydride possesses a remarkable capacity for oxidising. Ignited bodies burn in it, nitric acid absorbs it, and then acquires the property of acting on silver and other metals, even when diluted.Potassium iodideis 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 theseparation of iodine. This iodine may he recognised (seeOzone, ChapterIV.) 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 actsdirectly on ammonia, forming nitrogen and water, HNO2+ NH3= N2+ 2H2O.[56]
As nitrous anhydride easily splits up into NO2+ NO, so, like NO2, with warm water it gives nitric acid and nitric oxide, according to the equation 3N2O3+ H2O = 4NO + 2NHO3.
Being in a lower degree of oxidation than nitric acid, nitrous acid and its anhydride are oxidised in solutions by many oxidising substances—for example, by potassium permanganate—into nitric acid.[57]
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 N2O3and NO2are 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
2NHO3+ 3Cu = H2O + 3CuO + 2NO.
The second reaction consists in the formation of copper nitrate—
6NHO3+ 3CuO = 3H2O + 3Cu(NO3)2.
Nitric oxide is a colourless gas which is only slightly soluble in water (1/20of 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 N2and 2N2O3, but complete decomposition into N2and O2only takes place at the melting point of platinum, Emich 1892). The most characteristic property of nitric oxide is its capacity for directly and easily combining with oxygen (owing to the evolution of heat in the combination). With oxygen it forms nitrous anhydride and nitrogen peroxide, 2NO + O = N2O3, 2NO + O2= 2NO2. If nitric oxide is mixed with oxygen and immediately shaken up with caustic potash, it is almost entirely converted into potassium nitrite; whilst after a certain time, when the formation of nitric peroxide has already commenced, a mixture of potassium nitrite and nitrate is obtained. If oxygen is passed into a bell jar filled with nitric oxide, brown fumes of nitrous anhydride and nitric peroxide are formed, even in the absence of moisture; these in the presence of water give, as we already know, nitric acid and nitric oxide, so that in the presence of an excess of water and oxygen the whole of the nitric oxide is easily and directly converted into nitric acid. This reaction of the re-formation of nitric acid from nitric oxide, air, and water, 2NO + H2O + O3= 2HNO3, is frequently made use of in practice. The experiment showing the conversion of nitric oxide into nitric acid is very striking and instructive. As the intermixture of the oxygen with the oxide of nitrogen proceeds, the nitric acid formed dissolves in water, and if an excess of oxygen has not been added the whole of the gas (nitric oxide), being convertedinto HNO3, 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, HNO3or NO2(OH), in nitric anhydride, N2O5or (NO2)2O, and in ammonium chloride, NH4Cl. If X stand for an atom of hydrogen, or its equivalents, chlorine, hydroxyl, &c., and if O, which is, according to the law of substitution, equivalent to H2, be indicated by X2, then the three compounds of nitrogen above named should be considered as compounds of the type or form NX5. For example, in nitric acid X5= O2+ (OH), where O2= X4, and OH = X; whilst nitric oxide is a compound of the form NX2. Hence this lower form, like lower forms in general, strives by combination to attain to the higher forms proper to the compounds of a given element. NX2passes consecutively into NX3—namely, into N2O3and NHO2, NX4(for instance NO2) and NX5.
As the decomposition of nitric oxide begins at temperatures above 900°, many substances burn in it; thus, ignited phosphorus continues to burn in nitric oxide, but sulphur and charcoal are extinguished in it. This is due to the fact that the heat evolved in the combustion of these two substances is insufficient for the decomposition of the nitricoxide, 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.
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 fortransferringthe 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 oxidisean indefinitely large quantity of substances which cannot be directly oxidised either by the action of the atmospheric oxygen or by the action of nitric oxide itself. The sulphurous anhydride, SO2, which is obtained in the combustion of sulphur and in roasting many metallic sulphides in the air is an example of this kind. In practice this gas is obtained by burning sulphur or iron pyrites, the latter being thereby converted into oxide of iron and sulphurous anhydride. In contact with the oxygen of the atmosphere this gas does not pass into the higher degree of oxidation, sulphuric anhydride, SO3, and if it does form sulphuric acid with water and the oxygen of the atmosphere, SO2+ H2O + O = H2SO4, it does so very slowly. With nitric acid (and especially with nitrous acid, but not with nitrogen peroxide) and water, sulphurous anhydride, on the contrary, very easily forms sulphuric acid, and especially so when slightly heated (about 40°), the nitric acid (or, better still, nitrous acid) being converted into nitric oxide—
3SO2+ 2NHO3+ 2H2O = 2H2SO4+ 2NO.
The presence of water is absolutely indispensable here, otherwise sulphuric anhydride is formed, which combines with the oxides of nitrogen (nitrous anhydride), forming a crystalline substance containing oxides of nitrogen (chamber crystals, which will be described in Chapter XX.) Water destroys this compound, forming sulphuric acid and separating the oxides of nitrogen. The water must be taken in a greater quantity than that required for the formation of the hydrate H2SO4, because the latter absorbs oxides of nitrogen. With an excess of water, however, solution does not take place. If, in the above reaction, only water, sulphurous anhydride, and nitric or nitrous acid be taken in a definite quantity, then a definite quantity of sulphuric acid and nitric oxide will be formed, according to the preceding equation; but there the reaction ends and the excess of sulphurous anhydride, if there be any, will remain unchanged. But if we add air and water, then the nitric oxide will unite with the oxygen to form nitrogen peroxide, and the latter with water to form nitric and nitrous acids, which again give sulphuric acid from a fresh quantity of sulphurous anhydride. Nitric oxide is again formed, which is able to start the oxidation afresh if there be sufficient air. Thus it is possible with a definite quantity of nitric oxide to convert an indefinitely large quantity of sulphurous anhydride into sulphuric acid, water and oxygen only being required.[64]This may be easily demonstrated by an experiment on a small scale, if a certain quantity of nitric oxide be first introduced into a flask, and sulphurous anhydride, steam, and oxygen be then continually passed in. Thus the above-described reaction may be expressed in the following manner:—
nSO2+nO + (n+m)H2O + NO =nH2SO4,mH2O + NO
if we consider only the original substances and those finally formed. In this way a definite quantity of nitric oxide may serve for the conversion of an indefinite quantity of sulphurous anhydride, oxygen, and water into sulphuric acid. In reality, however, there is a limit to this, because air, and not pure oxygen, is employed for the oxidation, so that it is necessary to remove the nitrogen of the air and to introduce a fresh quantity of air. A certain quantity of nitric oxide will pass away with this nitrogen, and will in this way be lost.[65]
The preceding series of changes serve as the basis of themanufacture of sulphuric acidor so-calledchamber acid. This acid is prepared on a very large scale in chemical works because it is the cheapest acid whose action can be applied in a great number of cases. It is therefore used in immense quantities.
The process is carried on in a series of chambers (or in one divided by partitions as in fig.50, which shows the beginning and end of a chamber) constructed of sheet lead. These chambers are placed one against the other, and communicate by tubes or special orifices so placed that the inlet tubes are in the upper portion of the chamber, and the outlet in the lower and opposite end. The current of steam and gases necessary for the preparation of the sulphuric acid passes through these chambers and tubes. The acid as it is formed falls to the bottom of the chambers or runs down their walls, and flows from chamber to chamber (from the last towards the first), to permit of which the partitions do not reach to the bottom. The floor and walls of the chambers should therefore be made of a material on which the sulphuric acid will not act. Among the ordinary metals lead is the only one suitable.[65 bis]
see captionFig.50.—Section of sulphuric acid chambers, the first and last chambers only being represented. The tower to the left is called the Glover's tower, and that on the right the Gay-Lussac's tower. Less than1/10th of the natural size.
Fig.50.—Section of sulphuric acid chambers, the first and last chambers only being represented. The tower to the left is called the Glover's tower, and that on the right the Gay-Lussac's tower. Less than1/10th of the natural size.
For the formation of the sulphuric acid it is necessary to introducesulphurous anhydride, steam, air, and nitric acid, or some oxide of nitrogen, into the chambers. The sulphurous anhydride is produced by burning sulphur or iron pyrites. This is carried on in the furnace with four hearths to the left of the drawing. Air is led into the chambers and furnace through orifices in the furnace doors. The current of air and oxygen is regulated by opening or closing these orifices to a greater or less extent. The ingoing draught in the chambers is brought about by the fact that heated gases and vapours pass into the chambers, whose temperature is further raised by the reaction itself, and also by the remaining nitrogen being continually withdrawn from the outlet (above the towerK) by a tall chimney situated near the chambers. Nitric acid is prepared from a mixture of sulphuric acid and Chili saltpetre, in the same furnaces in which the sulphurous anhydride is evolved (or in special furnaces). Not more than 8 parts of nitre are taken to 100 parts of sulphur burnt. On leaving the furnace the vapours of nitric acid and oxides of nitrogen mixed with air and sulphurous anhydride first pass along the horizontal tubesTinto the receiverB B, which is partially cooled by water flowing in on the right-hand side and running out on the left byo, in order to reduce the temperature of the gases entering the chamber. The gases then pass up a tower filled with coke, and shown to the left of the drawing. In this tower are placed lumps of coke (the residue from the dry distillation of coal), over which sulphuric acid trickles from the reservoirM. This acid has absorbed in the end towerKthe oxides of nitrogen escaping from the chamber. This end tower is also filled with coke, over which a stream of strong sulphuric acid trickles from the reservoirM. The acid spreads over the coke, and, owing to the large surface offered by the latter, absorbs the greater part of the oxides of nitrogen escaping from the chambers. The sulphuric acid in passing down the tower becomes saturated with the oxides of nitrogen, and flows out athinto a special receiver (in the drawing situated by the side of the furnaces), from which it is forced up the tubesh′ h′by steam pressure into the reservoirM, situated above the first tower. The gases passing through this tower (hot) from the furnace on coming into contact with the sulphuric acid take up the oxides of nitrogen contained in it, and these are thus returned to the chamber and again participate in the reaction. The sulphuric acid left after their extraction flows into the chambers. Thus, on leaving the first coke tower the sulphurous anhydride, air, and vapours of nitric acid and of the oxides of nitrogen pass through the upper tubeminto the chamber. Here they come into contact with steam introduced by lead tubes into various parts of the chamber. The reaction takes place in the presence of water, the sulphuric acid falls to the bottom of the chamber, and the same processtakes place in the following chambers until the whole of the sulphurous anhydride is consumed. A somewhat greater proportion of air than is strictly necessary is passed in, in order that no sulphurous anhydride should be left unaltered for want of sufficient oxygen. The presence of an excess of oxygen is shown by the colour of the gases escaping from the last chamber. If they be of a pale colour it indicates an insufficiency of air (and the presence of sulphurous anhydride), as otherwise peroxide of nitrogen would be formed. A very dark colour shows an excess of air, which is also disadvantageous, because it increases the inevitable loss of nitric oxide by increasing the mass of escaping gases.[66]
Nitrous oxide, N2O,[67]is similar to water in its volumetric composition. Two volumes of nitrous oxide are formed from two volumes ofnitrogen and one volume of oxygen, which may be shown by the ordinary method for the analysis of the oxides of nitrogen (by passing them over red-hot copper or sodium). In contradistinction to the other oxides of nitrogen, it is not directly oxidised by oxygen, but it may be obtained from the higher oxides of nitrogen by the action of certain deoxidising substances; thus, for example, a mixture of two volumes of nitric oxide and one volume of sulphurous anhydride if left in contact with water and spongy platinum is converted into sulphuric acid and nitrous oxide, 2NO + SO2+ H2O = H2SO4+ N2O. Nitric acid, also, under the action of certain metals—for instance, of zinc[68]—gives nitrous oxide, although in this case mixed with nitric oxide. The usual method of preparing nitrous oxide consists in the decomposition of ammonium nitrate by the aid of heat, because in this case only water and nitrous oxide are formed, NH4NO3= 2H2O + N2O (a mixture of NH4Cl and KNO3is sometimes taken). The decomposition[69]proceeds very easily in an apparatus like that used for the preparation of ammonia or oxygen—that is, in a retort or flask with a gas-conducting tube. The decomposition must, however, be carried on carefully, as otherwise nitrogen is formed from the decomposition of the nitrous oxide.[70]
see captionFig.51.—Natterer's apparatus for the preparation of liquid nitrous oxide and carbonic anhydride. The gas first passes though the vessel V, for drying, and then into the pump (a section of the upper part of the apparatus is given on the left). The pistontof the force pump is moved by the crank E and fly-wheel turned by hand. The gas is pumped into the iron chamber A, where it is liquefied. The valve S allows the gas to enter A, but not to escape from it. The chamber and pump are cooled by the jacket B, filled with ice. When the gas is liquefied the vessel A is unscrewed from the pump, and the liquid may be poured from it by inverting it and unscrewing the valvev, when the liquid runs out of the tubex.
Fig.51.—Natterer's apparatus for the preparation of liquid nitrous oxide and carbonic anhydride. The gas first passes though the vessel V, for drying, and then into the pump (a section of the upper part of the apparatus is given on the left). The pistontof the force pump is moved by the crank E and fly-wheel turned by hand. The gas is pumped into the iron chamber A, where it is liquefied. The valve S allows the gas to enter A, but not to escape from it. The chamber and pump are cooled by the jacket B, filled with ice. When the gas is liquefied the vessel A is unscrewed from the pump, and the liquid may be poured from it by inverting it and unscrewing the valvev, when the liquid runs out of the tubex.
Nitrous oxide is not a permanent gas (absolute boiling point +36°); it is easily liquefied by the action of cold under a high pressure; at 15° it may be liquefied by a pressure of about 40 atmospheres. This gas is usually liquefied by means of the force pump[71]shown in fig.51. As it is liquefied with comparative ease, and as the cold produced by its vaporisation is very considerable,[72]it (as also liquid carbonic anhydride) is often employed in investigations requiring a low temperature. Nitrous oxide forms a very mobile, colourless liquid, which acts on the skin, and is incapable in a cold state of oxidising either metallic potassium, phosphorus, or carbon; its specific gravity is slightly less than that of water (0° = 0·910, 10° = 0·856, 35° = 0·60, 39° = 0·45, Villard, 1894). When evaporated under the receiver of an air-pump, the temperature falls to -100°, and the liquid solidifies into a snow-like mass, and partially forms transparent crystals. Both these substances are solid nitrous oxide. Mercury is immediately solidified in contact with evaporating liquid nitrous oxide.[73]
When introduced into the respiratory organs (and consequently into the blood also) nitrous oxide produces a peculiar kind of intoxication accompanied by spasmodic movements, and hence this gas, discovered by Priestley in 1776, received the name of ‘laughing gas.’ On a prolonged respiration it produces a state of insensibility (it is an anæsthetic like chloroform), and is therefore employed in dental and surgical operations.
Nitrous oxide is easily decomposed into nitrogen and oxygen by the action of heat, or a series of electric sparks; and this explains why a number of substances which cannot burn in nitric oxide do so with great ease in nitrous oxide. In fact, when nitric oxide gives some oxygen on decomposition, this oxygen immediately unites with a fresh portion of the gas to form nitric peroxide, whilst nitrous oxide does not possess this capacity for further combination with oxygen.[74]A mixture of nitrous oxide with hydrogen explodes like detonatinggas, gaseous nitrogen being formed, N2O + H2= H2O + N2. The volume of the remaining nitrogen is equal to the original volume of nitrous oxide, and is equal to the volume of hydrogen entering into combination with the oxygen; hence in this reaction equal volumes of nitrogen and hydrogen replace each other. Nitrous oxide is also very easily decomposed by red-hot metals; and sulphur, phosphorus, and charcoal burn in it, although not so brilliantly as in oxygen. A substance in burning in nitrous oxide evolves more heat than an equal quantity burning in oxygen; which most clearly shows that in the formation of nitrous oxide by the combination of nitrogen with oxygen there was not an evolution but an absorption of heat, there being no other source for the excess of heat in the combustion of substances in nitrous oxide (seeNote29). If a given volume of nitrous oxide be decomposed by a metal—for instance, sodium—then there remains, after cooling and total decomposition, a volume of nitrogen, exactly equal to that of the nitrous oxide taken; consequently the oxygen is, so to say, distributed between the atoms of nitrogen without producing an increase in the volume of the nitrogen.