[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 (ChapterV.) 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°—NH4NO2= 2H2O + N2. 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 NH4NO2by the action of a mixture of N2O3and 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, AgNO2, is obtained as a very slightly soluble substance, as a precipitate, on mixing solutions of silver nitrate, AgNO3, and potassium nitrite, KNO2. 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(NO2)2).[51 bis]Leroy (1889) obtained KNO2by mixing powdered KNO3with BaS, igniting the mixture in a crucible and washing the fused salts; BaSO4is then left as an insoluble residue, and KNO2passes into solution: 4KNO3+ BaS = 4KNO2+ BaSO4.[52]Probably potassium nitrite, KNO2, when strongly heated, especially with metallic oxides, evolves N and O, and gives potassium oxide, K2O, 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 N2O3= NO2- NO is reversible,i.e.resembles the conversion of N2O4into NO2. The brown colour of the fumes of N2O3is due to the formation of NO2.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, N2O3, 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 NO2(and therefore also nitric acid in the presence of water).[54]According to Thorpe, N2O3boils 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, N2O3= 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 O3there is obtained O, which acts as an oxidiser, and common oxygen O2. In a physical aspect the relation between N2O3and O3is 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, NH2R (where R is an element or a complex group) into hydroxides, RHO. In this case NH2R + NHO2forms 2N + H2O + RHO; NH2, 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, C6H5·NH2, which is obtained from nitrobenzene, C6H5·NO2(Note37), is converted by nitrous anhydride into phenol, C6H5·OH, which occurs in the creosote extracted from coal tar. Thus the H of the benzene is successively replaced by NO2, NH2, and HO; a method which is suitable for other cases also.[57]The action of a solution of potassium permanganate, KMnO4, on nitrous acid in the presence of sulphuric acid is determined by the fact that the higher oxide of manganese, Mn2O7, contained in the permanganate is converted into the lower oxide, MnO, which as a base forms manganese sulphate, MnSO4, and the oxygen serves for the oxidation of the N2O3into N2O5, 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.[58]The absolute boiling point = -93° (seeChapter II., Note29).[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 (FeCl2), 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 NaNO2or other substance corresponding to HNO2or HNO3has 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, KNO2and N2O or N2, are formed).[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, SbCl5, BiCl3, &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, FeSO4, absorbs it in the proportion of NO to 2FeSO4. 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, KMnO4, Note57) are able to convert it into nitric acid. If the presence of a radicle NO2, 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 callednitroso-compounds. These substances are described in Prof. Bunge's work (Kief, 1868).[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, CS2. 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.[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 (seeNote29) 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.[64]The instance of the action of a small quantity of NO in inducing a definite chemical reaction between large masses (SO2+ O + H2O = H2SO4) is very instructive, because the particulars relating to it have been studied, and show that intermediate forms of reaction may be discovered in the so-called contact or catalytic phenomena. The essence of the matter here is that A (= SO2) reacts upon B (= O and H2O) in the presence of C, because it gives BC, a substance which forms AB with A, and again liberates C. Consequently C is a medium, a transferring substance, without which the reaction does not proceed. Many similar phenomena may be found in other departments of life. Thus the merchant is an indispensable medium between the producer and the consumer; experiment is a medium between the phenomena of nature and the cognisant faculties, and language, customs, and laws are media which are as necessary for the exchanges of social intercourse as nitric oxide for those between sulphurous anhydride and oxygen and water.[65]If the sulphurous anhydride be prepared by roasting iron pyrites, FeS2, then each equivalent of pyrites (equivalent of iron, 56, of sulphur 32, of pyrites 120) requires six equivalents of oxygen (that is 96 parts) for the conversion of its sulphur into sulphuric acid (for forming 2H2SO4with water), besides 1½ equivalents (24 parts) for converting the iron into oxide, Fe2O3; hence the combustion of the pyrites for the formation of sulphuric acid and ferric oxide requires the introduction of an equal weight of oxygen (120 parts of oxygen to 120 parts of pyrites), or five times its weight of air, whilst four parts by weight of nitrogen will remain inactive, and in the removal of the exhausted air will carry off the remaining nitric oxide. If not all, at least a large portion of the nitric oxide may be collected by passing the escaping air, still containing some oxygen, through substances which absorb oxides of nitrogen. Sulphuric acid itself may be employed for this purpose if it be used in the form of the hydrate H2SO4, or containing only a small amount of water, because such sulphuric acid dissolves the oxides of nitrogen. They may be easily expelled from this solution by heating or by dilution with water, as they are only slightly soluble in aqueous sulphuric acid. Besides which, sulphurous anhydride acts on such sulphuric acid, being oxidised at the expense of the nitrous anhydride, and forming nitric oxide from it, which again enters into the cycle of action. For this reason the sulphuric acid which has absorbed the oxides of nitrogen escaping from the chambers in the towerK(seefig.50) is led back into the first chamber, where it comes into contact with sulphurous anhydride, by which means the oxides of nitrogen are reintroduced into the reaction which proceeds in the chambers. This is the use of the towers (Gay-Lussac's and Glover's) which are erected at either end of the chambers.[65 bis]Other metals, iron, copper, zinc, are corroded by it; glass and china are not acted upon, but they crack from the variations of temperature taking place in the chambers, and besides they are more difficult to join properly than lead; wood, &c., becomes charred.[66]By this means as much as 2,500,000 kilograms of chamber acid, containing about 60 per cent. of the hydrate H2SO4and about 40 per cent. of water, may be manufactured per year in one plant of 5,000 cubic metres capacity (without stoppages). This process has been brought to such a degree of perfection that as much as 300 parts of the hydrate H2SO4are obtained from 100 parts of sulphur, whilst the theoretical amount is not greater than 306 parts. The acid parts with its excess of water on heating. For this purpose it is heated in lead vessels. However, the acid containing about 75 per cent. of the hydrate (60° Baumé) already begins to act on the lead when heated, and therefore the further removal of water is conducted by evaporating in glass or platinum vessels, as will he described in Chapter XX. The aqueous acid (50° Baumé) obtained in the chambers is termed chamber acid. The acid concentrated to 60° Baumé is more generally employed, and sometimes the hydrate (66° Baumé) termed vitriol acid is also used. In England alone more than 1,000 million kilograms of chamber acid are produced by this method. The formation of sulphuric acid by the action of nitric acid was discovered by Drebbel, and the first lead chamber was erected by Roebuck, in Scotland, in the middle of the last century. The essence of the process was only brought to light at the beginning of this century, when many improvements were introduced into practice.[67]If the hydrate HNO3corresponds to N2O4, the hydrate HNO,hyponitrous acid, corresponds to N2O, and in this sense N2O ishyponitrous anhydride. Hyponitrous acid, corresponding with nitrous oxide (as its anhydride), is not known in a pure state, but its salts (Divers) are known. They are prepared by the reduction of nitrous (and consequently of nitric) salts by sodium amalgam. If this amalgam he added to a cold solution of an alkaline nitrite until the evolution of gas ceases, and the excess of alkali saturated with acetic acid, an insoluble yellow precipitate of silver hyponitrite, NAgO, will he obtained on adding a solution of silver nitrate. This hyponitrite is insoluble in cold acetic acid, and decomposes when heated, with the evolution of nitrous oxide. If rapidly heated it decomposes with an explosion. It is dissolved unchanged by weak mineral acids, whilst the stronger acids (for example, sulphuric and hydrochloric acids) decompose it, with the evolution of nitrogen, nitric and nitrous acids remaining in solution. Among the other salts of hyponitrous acid, HNO, the salts of lead, copper, and mercury are insoluble in water. Judging by the bond between hyponitrous acid and the other compounds of nitrogen, there is reason for thinking that its formula should he doubled, N2H2O2. For instance, Thoune (1893) on gradually oxidising hydroxylamine, NH2(OH), into nitrous acid, NO(OH) (Note25), by means of an alkaline solution of KMnO4, first obtained hyponitrous acid, N2H2O2, and then a peculiar intermediate acid, N2H2O3, which, by further oxidation, gave nitrous acid. On the other hand, Wislicenus (1893) showed that in the action of the sulphuric acid salt of hydroxylamine upon nitrite of sodium, there is formed, besides, nitrous oxide (according to V. Meyer, NH3O,H2SO4+ NaNO2= NaHSO4+ 2H2O + N2O), a small amount of hyponitrous acid which may be precipitated in the form of the silver salt; and this reaction is most simply expressed by taking the doubled formula of hyponitrous acid, NH2(OH) + NO(OH) = H2O + N2H2O2. The best argument in favour of the doubled formula is the property possessed by hyponitrous acid of forming acid salts, HNaN2O2(Zorn).According to Thoune, the following are the properties of hyponitrous acid. When liberated from the dry silver salt by the action of dry sulphuretted hydrogen, hyponitrous acid is unstable, and easily explodes even at low temperatures. But when dissolved in water (having been formed by the action of hydrochloric acid upon the silver salt), it is stable even when boiled with dilute acids and alkalis. The solution is colourless and has a strongly acid reaction. In the course of time, however, the aqueous solution also decomposes into nitrous oxide and water. The complete oxidation by permanganate of potash proceeds according to the following equation: 5H2N2O2+ 8KMnO4+ 12H2SO4= 10HNO3+ 4K2SO4+ 8MnSO4+ 12H2O. In an alkaline solution, KMnO4only oxidises hyponitrous acid into nitrous and not into nitric acid. Nitrous acid has a decomposing action upon hyponitrous acid, and if the aqueous solutions of the two acids be mixed together they immediately give off oxides of nitrogen. Hyponitrous acid does not liberate CO2from its salts, but on the other hand it is not displaced by CO2.[68]It is remarkable that electro-deposited copper powder gives nitrous oxide with a 10 p.c. solution of nitric acid, whilst ordinary copper gives nitric oxide. It is here evident that the physical and mechanical structure of the substance affects the course of the reaction—that is to say, it is a case of contact-action.[69]This decomposition is accompanied by the evolution of about 25,000 calories per molecular quantity, NH4NO3, and therefore takes place with ease, and sometimes with an explosion.[70]In order to remove any nitric oxide that might be present, the gas obtained is passed through a solution of ferrous sulphate. As nitrous oxide is very soluble in cold water (at 0°, 100 volumes of water dissolve 130 volumes of N2O; at 20°, 67 volumes), it must be collected over warm water. The nitrous oxide is much more soluble than nitric oxide, which is in agreement with the fact that nitrous oxide is much more easily liquefied than nitric oxide. Villard obtained a crystallohydrate, N2O,6H2O, which was tolerably stable at 0°.[71]Faraday obtained liquid nitrous oxide by the same method as liquid ammonia, by beating dry ammonium nitrate in a closed bent tube, one arm of which was immersed in a freezing mixture. In this case two layers of liquid are obtained at the cooled end, a lower layer of water and an upper layer of nitrous oxide. This experiment should be conducted with great care, as the pressure of the nitrous oxide in a liquid state is considerable, namely (according to Regnault), at +10° = 45 atmospheres, at 0° = 36 atmospheres, at -10° = 29 atmospheres, and at -20° = 23 atmospheres. It boils at -92°, and the pressure is then therefore = 1 atmosphere (seeChapter II., Note27).[72]Liquid nitrous oxide, in vaporising at the same pressure as liquid carbonic anhydride, gives rise to almost equal or even slightly lower temperatures. Thus at a pressure of 25 mm. carbonic anhydride gives a temperature as low as -115°, and nitrous oxide of -125° (Dewar). The similarity of these properties and even of the absolute boiling point (CO2+ 32°, N2O +36°) is all the more remarkable because these gases have the same molecular weight = 44 (ChapterVII.)[73]A very characteristic experiment of simultaneous combustion and intense cold may be performed by means of liquid nitrous oxide; if liquid nitrous oxide be poured into a test tube containing some mercury the mercury will solidify, and if a piece of red-hot charcoal be thrown upon the surface of the nitrous oxide it will continue to burn very brilliantly, giving rise to a high temperature.[74]In thefollowing chapterwe shall consider the volumetric composition of the oxides of nitrogen. It explains the difference between nitric and nitrous oxide. Nitrous oxide is formed with a diminution of volumes (contraction), nitric oxide without contraction, its volume being equal to the sum of the volumes of the nitrogen and oxygen of which it is composed. By oxidation, if it could be directly accomplished, two volumes of nitrous oxide and one volume of oxygen would not give three but four volumes of nitric oxide. These facts must be taken into consideration in comparing the calorific equivalents of formation, the capacity for supporting combustion, and other properties of nitrous and nitric oxides, N2O and NO.
[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 (ChapterV.) 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°—NH4NO2= 2H2O + N2. 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 NH4NO2by the action of a mixture of N2O3and 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.
[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 (ChapterV.) 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°—NH4NO2= 2H2O + N2. 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 NH4NO2by the action of a mixture of N2O3and 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, AgNO2, is obtained as a very slightly soluble substance, as a precipitate, on mixing solutions of silver nitrate, AgNO3, and potassium nitrite, KNO2. 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(NO2)2).
[51]Silver nitrite, AgNO2, is obtained as a very slightly soluble substance, as a precipitate, on mixing solutions of silver nitrate, AgNO3, and potassium nitrite, KNO2. 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(NO2)2).
[51 bis]Leroy (1889) obtained KNO2by mixing powdered KNO3with BaS, igniting the mixture in a crucible and washing the fused salts; BaSO4is then left as an insoluble residue, and KNO2passes into solution: 4KNO3+ BaS = 4KNO2+ BaSO4.
[51 bis]Leroy (1889) obtained KNO2by mixing powdered KNO3with BaS, igniting the mixture in a crucible and washing the fused salts; BaSO4is then left as an insoluble residue, and KNO2passes into solution: 4KNO3+ BaS = 4KNO2+ BaSO4.
[52]Probably potassium nitrite, KNO2, when strongly heated, especially with metallic oxides, evolves N and O, and gives potassium oxide, K2O, because nitre is liable to such a decomposition; but it has, as yet, been but little investigated.
[52]Probably potassium nitrite, KNO2, when strongly heated, especially with metallic oxides, evolves N and O, and gives potassium oxide, K2O, 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 N2O3= NO2- NO is reversible,i.e.resembles the conversion of N2O4into NO2. The brown colour of the fumes of N2O3is due to the formation of NO2.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, N2O3, 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 NO2(and therefore also nitric acid in the presence of water).
[53]There are many researches which lead to the conclusion that the reaction N2O3= NO2- NO is reversible,i.e.resembles the conversion of N2O4into NO2. The brown colour of the fumes of N2O3is due to the formation of NO2.
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, N2O3, 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 NO2(and therefore also nitric acid in the presence of water).
[54]According to Thorpe, N2O3boils at +18°. According to Geuther, at +3°·5, and its sp. gr. at 0° = 1·449.
[54]According to Thorpe, N2O3boils 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, N2O3= 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 O3there is obtained O, which acts as an oxidiser, and common oxygen O2. In a physical aspect the relation between N2O3and O3is revealed in the fact that both substances are of a blue colour when in the liquid state.
[55]In its oxidising action nitrous anhydride gives nitric oxide, N2O3= 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 O3there is obtained O, which acts as an oxidiser, and common oxygen O2. In a physical aspect the relation between N2O3and O3is 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, NH2R (where R is an element or a complex group) into hydroxides, RHO. In this case NH2R + NHO2forms 2N + H2O + RHO; NH2, 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, C6H5·NH2, which is obtained from nitrobenzene, C6H5·NO2(Note37), is converted by nitrous anhydride into phenol, C6H5·OH, which occurs in the creosote extracted from coal tar. Thus the H of the benzene is successively replaced by NO2, NH2, and HO; a method which is suitable for other cases also.
[56]This reaction is taken advantage of for converting the amides, NH2R (where R is an element or a complex group) into hydroxides, RHO. In this case NH2R + NHO2forms 2N + H2O + RHO; NH2, 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, C6H5·NH2, which is obtained from nitrobenzene, C6H5·NO2(Note37), is converted by nitrous anhydride into phenol, C6H5·OH, which occurs in the creosote extracted from coal tar. Thus the H of the benzene is successively replaced by NO2, NH2, and HO; a method which is suitable for other cases also.
[57]The action of a solution of potassium permanganate, KMnO4, on nitrous acid in the presence of sulphuric acid is determined by the fact that the higher oxide of manganese, Mn2O7, contained in the permanganate is converted into the lower oxide, MnO, which as a base forms manganese sulphate, MnSO4, and the oxygen serves for the oxidation of the N2O3into N2O5, 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.
[57]The action of a solution of potassium permanganate, KMnO4, on nitrous acid in the presence of sulphuric acid is determined by the fact that the higher oxide of manganese, Mn2O7, contained in the permanganate is converted into the lower oxide, MnO, which as a base forms manganese sulphate, MnSO4, and the oxygen serves for the oxidation of the N2O3into N2O5, 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.
[58]The absolute boiling point = -93° (seeChapter II., Note29).
[58]The absolute boiling point = -93° (seeChapter II., Note29).
[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 (FeCl2), 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 NaNO2or other substance corresponding to HNO2or HNO3has 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, KNO2and N2O or N2, are formed).
[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 (FeCl2), 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 NaNO2or other substance corresponding to HNO2or HNO3has 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, KNO2and N2O or N2, are formed).
[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.
[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, SbCl5, BiCl3, &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, FeSO4, absorbs it in the proportion of NO to 2FeSO4. 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, KMnO4, Note57) are able to convert it into nitric acid. If the presence of a radicle NO2, 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 callednitroso-compounds. These substances are described in Prof. Bunge's work (Kief, 1868).
[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, SbCl5, BiCl3, &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, FeSO4, absorbs it in the proportion of NO to 2FeSO4. 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, KMnO4, Note57) are able to convert it into nitric acid. If the presence of a radicle NO2, 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 callednitroso-compounds. These substances are described in Prof. Bunge's work (Kief, 1868).
[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, CS2. 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.
[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, CS2. 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.
[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 (seeNote29) 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.
[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 (seeNote29) 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.
[64]The instance of the action of a small quantity of NO in inducing a definite chemical reaction between large masses (SO2+ O + H2O = H2SO4) is very instructive, because the particulars relating to it have been studied, and show that intermediate forms of reaction may be discovered in the so-called contact or catalytic phenomena. The essence of the matter here is that A (= SO2) reacts upon B (= O and H2O) in the presence of C, because it gives BC, a substance which forms AB with A, and again liberates C. Consequently C is a medium, a transferring substance, without which the reaction does not proceed. Many similar phenomena may be found in other departments of life. Thus the merchant is an indispensable medium between the producer and the consumer; experiment is a medium between the phenomena of nature and the cognisant faculties, and language, customs, and laws are media which are as necessary for the exchanges of social intercourse as nitric oxide for those between sulphurous anhydride and oxygen and water.
[64]The instance of the action of a small quantity of NO in inducing a definite chemical reaction between large masses (SO2+ O + H2O = H2SO4) is very instructive, because the particulars relating to it have been studied, and show that intermediate forms of reaction may be discovered in the so-called contact or catalytic phenomena. The essence of the matter here is that A (= SO2) reacts upon B (= O and H2O) in the presence of C, because it gives BC, a substance which forms AB with A, and again liberates C. Consequently C is a medium, a transferring substance, without which the reaction does not proceed. Many similar phenomena may be found in other departments of life. Thus the merchant is an indispensable medium between the producer and the consumer; experiment is a medium between the phenomena of nature and the cognisant faculties, and language, customs, and laws are media which are as necessary for the exchanges of social intercourse as nitric oxide for those between sulphurous anhydride and oxygen and water.
[65]If the sulphurous anhydride be prepared by roasting iron pyrites, FeS2, then each equivalent of pyrites (equivalent of iron, 56, of sulphur 32, of pyrites 120) requires six equivalents of oxygen (that is 96 parts) for the conversion of its sulphur into sulphuric acid (for forming 2H2SO4with water), besides 1½ equivalents (24 parts) for converting the iron into oxide, Fe2O3; hence the combustion of the pyrites for the formation of sulphuric acid and ferric oxide requires the introduction of an equal weight of oxygen (120 parts of oxygen to 120 parts of pyrites), or five times its weight of air, whilst four parts by weight of nitrogen will remain inactive, and in the removal of the exhausted air will carry off the remaining nitric oxide. If not all, at least a large portion of the nitric oxide may be collected by passing the escaping air, still containing some oxygen, through substances which absorb oxides of nitrogen. Sulphuric acid itself may be employed for this purpose if it be used in the form of the hydrate H2SO4, or containing only a small amount of water, because such sulphuric acid dissolves the oxides of nitrogen. They may be easily expelled from this solution by heating or by dilution with water, as they are only slightly soluble in aqueous sulphuric acid. Besides which, sulphurous anhydride acts on such sulphuric acid, being oxidised at the expense of the nitrous anhydride, and forming nitric oxide from it, which again enters into the cycle of action. For this reason the sulphuric acid which has absorbed the oxides of nitrogen escaping from the chambers in the towerK(seefig.50) is led back into the first chamber, where it comes into contact with sulphurous anhydride, by which means the oxides of nitrogen are reintroduced into the reaction which proceeds in the chambers. This is the use of the towers (Gay-Lussac's and Glover's) which are erected at either end of the chambers.
[65]If the sulphurous anhydride be prepared by roasting iron pyrites, FeS2, then each equivalent of pyrites (equivalent of iron, 56, of sulphur 32, of pyrites 120) requires six equivalents of oxygen (that is 96 parts) for the conversion of its sulphur into sulphuric acid (for forming 2H2SO4with water), besides 1½ equivalents (24 parts) for converting the iron into oxide, Fe2O3; hence the combustion of the pyrites for the formation of sulphuric acid and ferric oxide requires the introduction of an equal weight of oxygen (120 parts of oxygen to 120 parts of pyrites), or five times its weight of air, whilst four parts by weight of nitrogen will remain inactive, and in the removal of the exhausted air will carry off the remaining nitric oxide. If not all, at least a large portion of the nitric oxide may be collected by passing the escaping air, still containing some oxygen, through substances which absorb oxides of nitrogen. Sulphuric acid itself may be employed for this purpose if it be used in the form of the hydrate H2SO4, or containing only a small amount of water, because such sulphuric acid dissolves the oxides of nitrogen. They may be easily expelled from this solution by heating or by dilution with water, as they are only slightly soluble in aqueous sulphuric acid. Besides which, sulphurous anhydride acts on such sulphuric acid, being oxidised at the expense of the nitrous anhydride, and forming nitric oxide from it, which again enters into the cycle of action. For this reason the sulphuric acid which has absorbed the oxides of nitrogen escaping from the chambers in the towerK(seefig.50) is led back into the first chamber, where it comes into contact with sulphurous anhydride, by which means the oxides of nitrogen are reintroduced into the reaction which proceeds in the chambers. This is the use of the towers (Gay-Lussac's and Glover's) which are erected at either end of the chambers.
[65 bis]Other metals, iron, copper, zinc, are corroded by it; glass and china are not acted upon, but they crack from the variations of temperature taking place in the chambers, and besides they are more difficult to join properly than lead; wood, &c., becomes charred.
[65 bis]Other metals, iron, copper, zinc, are corroded by it; glass and china are not acted upon, but they crack from the variations of temperature taking place in the chambers, and besides they are more difficult to join properly than lead; wood, &c., becomes charred.
[66]By this means as much as 2,500,000 kilograms of chamber acid, containing about 60 per cent. of the hydrate H2SO4and about 40 per cent. of water, may be manufactured per year in one plant of 5,000 cubic metres capacity (without stoppages). This process has been brought to such a degree of perfection that as much as 300 parts of the hydrate H2SO4are obtained from 100 parts of sulphur, whilst the theoretical amount is not greater than 306 parts. The acid parts with its excess of water on heating. For this purpose it is heated in lead vessels. However, the acid containing about 75 per cent. of the hydrate (60° Baumé) already begins to act on the lead when heated, and therefore the further removal of water is conducted by evaporating in glass or platinum vessels, as will he described in Chapter XX. The aqueous acid (50° Baumé) obtained in the chambers is termed chamber acid. The acid concentrated to 60° Baumé is more generally employed, and sometimes the hydrate (66° Baumé) termed vitriol acid is also used. In England alone more than 1,000 million kilograms of chamber acid are produced by this method. The formation of sulphuric acid by the action of nitric acid was discovered by Drebbel, and the first lead chamber was erected by Roebuck, in Scotland, in the middle of the last century. The essence of the process was only brought to light at the beginning of this century, when many improvements were introduced into practice.
[66]By this means as much as 2,500,000 kilograms of chamber acid, containing about 60 per cent. of the hydrate H2SO4and about 40 per cent. of water, may be manufactured per year in one plant of 5,000 cubic metres capacity (without stoppages). This process has been brought to such a degree of perfection that as much as 300 parts of the hydrate H2SO4are obtained from 100 parts of sulphur, whilst the theoretical amount is not greater than 306 parts. The acid parts with its excess of water on heating. For this purpose it is heated in lead vessels. However, the acid containing about 75 per cent. of the hydrate (60° Baumé) already begins to act on the lead when heated, and therefore the further removal of water is conducted by evaporating in glass or platinum vessels, as will he described in Chapter XX. The aqueous acid (50° Baumé) obtained in the chambers is termed chamber acid. The acid concentrated to 60° Baumé is more generally employed, and sometimes the hydrate (66° Baumé) termed vitriol acid is also used. In England alone more than 1,000 million kilograms of chamber acid are produced by this method. The formation of sulphuric acid by the action of nitric acid was discovered by Drebbel, and the first lead chamber was erected by Roebuck, in Scotland, in the middle of the last century. The essence of the process was only brought to light at the beginning of this century, when many improvements were introduced into practice.
[67]If the hydrate HNO3corresponds to N2O4, the hydrate HNO,hyponitrous acid, corresponds to N2O, and in this sense N2O ishyponitrous anhydride. Hyponitrous acid, corresponding with nitrous oxide (as its anhydride), is not known in a pure state, but its salts (Divers) are known. They are prepared by the reduction of nitrous (and consequently of nitric) salts by sodium amalgam. If this amalgam he added to a cold solution of an alkaline nitrite until the evolution of gas ceases, and the excess of alkali saturated with acetic acid, an insoluble yellow precipitate of silver hyponitrite, NAgO, will he obtained on adding a solution of silver nitrate. This hyponitrite is insoluble in cold acetic acid, and decomposes when heated, with the evolution of nitrous oxide. If rapidly heated it decomposes with an explosion. It is dissolved unchanged by weak mineral acids, whilst the stronger acids (for example, sulphuric and hydrochloric acids) decompose it, with the evolution of nitrogen, nitric and nitrous acids remaining in solution. Among the other salts of hyponitrous acid, HNO, the salts of lead, copper, and mercury are insoluble in water. Judging by the bond between hyponitrous acid and the other compounds of nitrogen, there is reason for thinking that its formula should he doubled, N2H2O2. For instance, Thoune (1893) on gradually oxidising hydroxylamine, NH2(OH), into nitrous acid, NO(OH) (Note25), by means of an alkaline solution of KMnO4, first obtained hyponitrous acid, N2H2O2, and then a peculiar intermediate acid, N2H2O3, which, by further oxidation, gave nitrous acid. On the other hand, Wislicenus (1893) showed that in the action of the sulphuric acid salt of hydroxylamine upon nitrite of sodium, there is formed, besides, nitrous oxide (according to V. Meyer, NH3O,H2SO4+ NaNO2= NaHSO4+ 2H2O + N2O), a small amount of hyponitrous acid which may be precipitated in the form of the silver salt; and this reaction is most simply expressed by taking the doubled formula of hyponitrous acid, NH2(OH) + NO(OH) = H2O + N2H2O2. The best argument in favour of the doubled formula is the property possessed by hyponitrous acid of forming acid salts, HNaN2O2(Zorn).According to Thoune, the following are the properties of hyponitrous acid. When liberated from the dry silver salt by the action of dry sulphuretted hydrogen, hyponitrous acid is unstable, and easily explodes even at low temperatures. But when dissolved in water (having been formed by the action of hydrochloric acid upon the silver salt), it is stable even when boiled with dilute acids and alkalis. The solution is colourless and has a strongly acid reaction. In the course of time, however, the aqueous solution also decomposes into nitrous oxide and water. The complete oxidation by permanganate of potash proceeds according to the following equation: 5H2N2O2+ 8KMnO4+ 12H2SO4= 10HNO3+ 4K2SO4+ 8MnSO4+ 12H2O. In an alkaline solution, KMnO4only oxidises hyponitrous acid into nitrous and not into nitric acid. Nitrous acid has a decomposing action upon hyponitrous acid, and if the aqueous solutions of the two acids be mixed together they immediately give off oxides of nitrogen. Hyponitrous acid does not liberate CO2from its salts, but on the other hand it is not displaced by CO2.
[67]If the hydrate HNO3corresponds to N2O4, the hydrate HNO,hyponitrous acid, corresponds to N2O, and in this sense N2O ishyponitrous anhydride. Hyponitrous acid, corresponding with nitrous oxide (as its anhydride), is not known in a pure state, but its salts (Divers) are known. They are prepared by the reduction of nitrous (and consequently of nitric) salts by sodium amalgam. If this amalgam he added to a cold solution of an alkaline nitrite until the evolution of gas ceases, and the excess of alkali saturated with acetic acid, an insoluble yellow precipitate of silver hyponitrite, NAgO, will he obtained on adding a solution of silver nitrate. This hyponitrite is insoluble in cold acetic acid, and decomposes when heated, with the evolution of nitrous oxide. If rapidly heated it decomposes with an explosion. It is dissolved unchanged by weak mineral acids, whilst the stronger acids (for example, sulphuric and hydrochloric acids) decompose it, with the evolution of nitrogen, nitric and nitrous acids remaining in solution. Among the other salts of hyponitrous acid, HNO, the salts of lead, copper, and mercury are insoluble in water. Judging by the bond between hyponitrous acid and the other compounds of nitrogen, there is reason for thinking that its formula should he doubled, N2H2O2. For instance, Thoune (1893) on gradually oxidising hydroxylamine, NH2(OH), into nitrous acid, NO(OH) (Note25), by means of an alkaline solution of KMnO4, first obtained hyponitrous acid, N2H2O2, and then a peculiar intermediate acid, N2H2O3, which, by further oxidation, gave nitrous acid. On the other hand, Wislicenus (1893) showed that in the action of the sulphuric acid salt of hydroxylamine upon nitrite of sodium, there is formed, besides, nitrous oxide (according to V. Meyer, NH3O,H2SO4+ NaNO2= NaHSO4+ 2H2O + N2O), a small amount of hyponitrous acid which may be precipitated in the form of the silver salt; and this reaction is most simply expressed by taking the doubled formula of hyponitrous acid, NH2(OH) + NO(OH) = H2O + N2H2O2. The best argument in favour of the doubled formula is the property possessed by hyponitrous acid of forming acid salts, HNaN2O2(Zorn).
According to Thoune, the following are the properties of hyponitrous acid. When liberated from the dry silver salt by the action of dry sulphuretted hydrogen, hyponitrous acid is unstable, and easily explodes even at low temperatures. But when dissolved in water (having been formed by the action of hydrochloric acid upon the silver salt), it is stable even when boiled with dilute acids and alkalis. The solution is colourless and has a strongly acid reaction. In the course of time, however, the aqueous solution also decomposes into nitrous oxide and water. The complete oxidation by permanganate of potash proceeds according to the following equation: 5H2N2O2+ 8KMnO4+ 12H2SO4= 10HNO3+ 4K2SO4+ 8MnSO4+ 12H2O. In an alkaline solution, KMnO4only oxidises hyponitrous acid into nitrous and not into nitric acid. Nitrous acid has a decomposing action upon hyponitrous acid, and if the aqueous solutions of the two acids be mixed together they immediately give off oxides of nitrogen. Hyponitrous acid does not liberate CO2from its salts, but on the other hand it is not displaced by CO2.
[68]It is remarkable that electro-deposited copper powder gives nitrous oxide with a 10 p.c. solution of nitric acid, whilst ordinary copper gives nitric oxide. It is here evident that the physical and mechanical structure of the substance affects the course of the reaction—that is to say, it is a case of contact-action.
[68]It is remarkable that electro-deposited copper powder gives nitrous oxide with a 10 p.c. solution of nitric acid, whilst ordinary copper gives nitric oxide. It is here evident that the physical and mechanical structure of the substance affects the course of the reaction—that is to say, it is a case of contact-action.
[69]This decomposition is accompanied by the evolution of about 25,000 calories per molecular quantity, NH4NO3, and therefore takes place with ease, and sometimes with an explosion.
[69]This decomposition is accompanied by the evolution of about 25,000 calories per molecular quantity, NH4NO3, and therefore takes place with ease, and sometimes with an explosion.
[70]In order to remove any nitric oxide that might be present, the gas obtained is passed through a solution of ferrous sulphate. As nitrous oxide is very soluble in cold water (at 0°, 100 volumes of water dissolve 130 volumes of N2O; at 20°, 67 volumes), it must be collected over warm water. The nitrous oxide is much more soluble than nitric oxide, which is in agreement with the fact that nitrous oxide is much more easily liquefied than nitric oxide. Villard obtained a crystallohydrate, N2O,6H2O, which was tolerably stable at 0°.
[70]In order to remove any nitric oxide that might be present, the gas obtained is passed through a solution of ferrous sulphate. As nitrous oxide is very soluble in cold water (at 0°, 100 volumes of water dissolve 130 volumes of N2O; at 20°, 67 volumes), it must be collected over warm water. The nitrous oxide is much more soluble than nitric oxide, which is in agreement with the fact that nitrous oxide is much more easily liquefied than nitric oxide. Villard obtained a crystallohydrate, N2O,6H2O, which was tolerably stable at 0°.
[71]Faraday obtained liquid nitrous oxide by the same method as liquid ammonia, by beating dry ammonium nitrate in a closed bent tube, one arm of which was immersed in a freezing mixture. In this case two layers of liquid are obtained at the cooled end, a lower layer of water and an upper layer of nitrous oxide. This experiment should be conducted with great care, as the pressure of the nitrous oxide in a liquid state is considerable, namely (according to Regnault), at +10° = 45 atmospheres, at 0° = 36 atmospheres, at -10° = 29 atmospheres, and at -20° = 23 atmospheres. It boils at -92°, and the pressure is then therefore = 1 atmosphere (seeChapter II., Note27).
[71]Faraday obtained liquid nitrous oxide by the same method as liquid ammonia, by beating dry ammonium nitrate in a closed bent tube, one arm of which was immersed in a freezing mixture. In this case two layers of liquid are obtained at the cooled end, a lower layer of water and an upper layer of nitrous oxide. This experiment should be conducted with great care, as the pressure of the nitrous oxide in a liquid state is considerable, namely (according to Regnault), at +10° = 45 atmospheres, at 0° = 36 atmospheres, at -10° = 29 atmospheres, and at -20° = 23 atmospheres. It boils at -92°, and the pressure is then therefore = 1 atmosphere (seeChapter II., Note27).
[72]Liquid nitrous oxide, in vaporising at the same pressure as liquid carbonic anhydride, gives rise to almost equal or even slightly lower temperatures. Thus at a pressure of 25 mm. carbonic anhydride gives a temperature as low as -115°, and nitrous oxide of -125° (Dewar). The similarity of these properties and even of the absolute boiling point (CO2+ 32°, N2O +36°) is all the more remarkable because these gases have the same molecular weight = 44 (ChapterVII.)
[72]Liquid nitrous oxide, in vaporising at the same pressure as liquid carbonic anhydride, gives rise to almost equal or even slightly lower temperatures. Thus at a pressure of 25 mm. carbonic anhydride gives a temperature as low as -115°, and nitrous oxide of -125° (Dewar). The similarity of these properties and even of the absolute boiling point (CO2+ 32°, N2O +36°) is all the more remarkable because these gases have the same molecular weight = 44 (ChapterVII.)
[73]A very characteristic experiment of simultaneous combustion and intense cold may be performed by means of liquid nitrous oxide; if liquid nitrous oxide be poured into a test tube containing some mercury the mercury will solidify, and if a piece of red-hot charcoal be thrown upon the surface of the nitrous oxide it will continue to burn very brilliantly, giving rise to a high temperature.
[73]A very characteristic experiment of simultaneous combustion and intense cold may be performed by means of liquid nitrous oxide; if liquid nitrous oxide be poured into a test tube containing some mercury the mercury will solidify, and if a piece of red-hot charcoal be thrown upon the surface of the nitrous oxide it will continue to burn very brilliantly, giving rise to a high temperature.
[74]In thefollowing chapterwe shall consider the volumetric composition of the oxides of nitrogen. It explains the difference between nitric and nitrous oxide. Nitrous oxide is formed with a diminution of volumes (contraction), nitric oxide without contraction, its volume being equal to the sum of the volumes of the nitrogen and oxygen of which it is composed. By oxidation, if it could be directly accomplished, two volumes of nitrous oxide and one volume of oxygen would not give three but four volumes of nitric oxide. These facts must be taken into consideration in comparing the calorific equivalents of formation, the capacity for supporting combustion, and other properties of nitrous and nitric oxides, N2O and NO.
[74]In thefollowing chapterwe shall consider the volumetric composition of the oxides of nitrogen. It explains the difference between nitric and nitrous oxide. Nitrous oxide is formed with a diminution of volumes (contraction), nitric oxide without contraction, its volume being equal to the sum of the volumes of the nitrogen and oxygen of which it is composed. By oxidation, if it could be directly accomplished, two volumes of nitrous oxide and one volume of oxygen would not give three but four volumes of nitric oxide. These facts must be taken into consideration in comparing the calorific equivalents of formation, the capacity for supporting combustion, and other properties of nitrous and nitric oxides, N2O and NO.