Footnotes:[1]The ammonia in the air, water, and soil proceeds from the decomposition of the nitrogenous substances of plants and animals, and also probably from the reduction of nitrates. Ammonia is always formed in the rusting of iron. Its formation in this case depends in all probability on the decomposition of water, and on the action of the hydrogen at the moment of its evolution on the nitric acid contained in the air (Cloez), or on the formation of ammonium nitrite, which takes place under many circumstances. The evolution of vapours of ammonia compounds is sometimes observed in the vicinity of volcanoes. At a red heat nitrogen combines directly with B Ca Mg, and with many other metals, and these compounds, when heated with a caustic alkali, or in the presence of water, give ammonia (seeChapter XIV., Note14, and Chapter XVII., Note 12). These are examples of the indirect combination of nitrogen with hydrogen.[2]If a silent discharge or a series of electric sparks be passed through ammonia gas, it is decomposed into nitrogen and hydrogen. This is a phenomenon of dissociation; therefore, a series of sparks do not totally decompose the ammonia, but leave a certain portion undecomposed. One volume of nitrogen and three volumes of hydrogen are obtained from two volumes of ammonia decomposed. Ramsay and Young (1884) investigated the decomposition of NH3under the action of heat, and showed that at 500°, 1½ p.c. is decomposed, at 600° about 18 p.c., at 800° 65 p.c., but these results were hardly free from the influence of ‘contact.’ Thepresenceof free ammonia—that is, ammonia not combined with acids—in a gas or aqueous solution may be recognised by its characteristic smell. But many ammonia salts do not possess this smell. However, on the addition of an alkali (for instance, caustic lime, potash, or soda), they evolve ammonia gas, especially when heated. The presence of ammonia may be made visible by introducing a substance moistened with strong hydrochloric acid into its neighbourhood. A white cloud, or visible white vapour, then makes its appearance. This depends on the fact that both ammonia and hydrochloric acid are volatile, and on coming into contact with each other produce solid sal-ammoniac, NH4Cl, which forms a cloud. This test is usually made by dipping a glass rod into hydrochloric acid, and holding it over the vessel from which the ammonia is evolved. With small amounts of ammonia this test is, however, untrustworthy, as the white vapour is scarcely observable. In this case it is best to take paper moistened with mercurous nitrate, HgNO3. This paper turns black in the presence of ammonia, owing to the formation of a black compound of ammonia with mercurous oxide. The smallest traces of ammonia (for instance, in river water) may be detected by means of the so-called Nessler's reagent, containing a solution of mercuric chloride and potassium iodide, which forms a brown coloration or precipitate with the smallest quantities of ammonia. It will be useful here to give the thermochemical data (in thousands of units of heat, according to Thomsen), or the quantities of heatevolvedin the formation of ammonia and its compounds in quantities expressed by their formulæ. Thus, for instance, (N + H3) 26·7 indicates that 14 grams of nitrogen in combining with 3 grams of hydrogen develop sufficient heat to raise the temperature of 26·7 kilograms of water 1°. (NH3+ nH2O) 8·4 (heat of solution); (NH3,nH2O + HCl,nH2O) 12·3; (N + H4+ Cl) 90·6; (NH3+ HCl) 41·9.[3]The same ammonia water is obtained, although in smaller quantities, in the dry distillation of plants and of coal, which consists of the remains of fossil plants. In all these cases the ammonia proceeds from the destruction of the complex nitrogenous substances occurring in plants and animals. The ammonia salts employed in the arts are prepared by this method.[4]The technical methods for the preparation of ammonia water, and for the extraction of ammonia from it, are to a certain extent explained in the figures accompanying the text.[5]Usually these crystals are sublimed by heating them in crucibles or pots, when the vapours of sal-ammoniac condense on the cold covers as a crust, in which form the salt comes into the market.[6]On a small scale ammonia may be prepared in a glass flask by mixing equal parts by weight of slaked lime and finely-powdered sal-ammoniac, the neck of the flask being connected with an arrangement for drying the gas obtained. In this instance neither calcium chloride nor sulphuric acid can be used for drying the gas, since both these substances absorb ammonia, and therefore solid caustic potash, which is capable of retaining the water, is employed. The gas-conducting tube leading from the desiccating apparatus is introduced into a mercury bath, if dry gaseous ammonia be required, because water cannot be employed in collecting ammonia gas. Ammonia was first obtained in this dry state by Priestley, and its composition was investigated by Berthollet at the end of the last century. Oxide of lead mixed with sal-ammoniac (Isambert) evolves ammonia with still greater ease than lime. The cause and process of the decomposition are almost the same, 2PbO + 2NH4Cl = Pb2OCl2+ H2O + 2NH3. Lead oxychloride is (probably) formed.[7]see captionFig.44.—Carré's apparatus. Described in text.This is evident from the fact that its absolute boiling point lies at about +130° (Chapter II., Note29). It may therefore be liquefied by pressure alone at the ordinary, and even at much higher temperatures. The latent heat of evaporation of 17 parts by weight of ammonia equals 4,400 units of heat, and hence liquid ammonia may be employed for the production of cold. Strong aqueous solutions of ammonia, which in parting with their ammonia act in a similar manner, are not unfrequently employed for this purpose. Suppose a saturated solution of ammonia to be contained in a closed vessel furnished with a receiver. If the ammoniacal solution be heated, the ammonia, with a small quantity of water, will pass off from the solution, and in accumulating in the apparatus will produce a considerable pressure, and will therefore liquefy in the cooler portions of the receiver. Hence liquid ammonia will be obtained in the receiver. The heating of the vessel containing the aqueous solution of ammonia is then stopped. After having been heated it contains only water, or a solution poor in ammonia. When once it begins to cool the ammonia vapours commence dissolving in it, the space becomes rarefied, and a rapid vaporisation of the liquefied ammonia left in the receiver takes place. In evaporating in the receiver it will cause the temperature in it to fall considerably, and will itself pass into the aqueous solution. In the end, the same ammoniacal solution as originally taken is re-obtained. Thus, in this case, on heating the vessel the pressure increases by itself, and on cooling it diminishes, so that here heat directly replaces mechanical work. This is the principle of the simplest forms ofCarré's ice-making machines, shown in fig.44. C is a vessel made of boiler plates into which the saturated solution of ammonia is poured; m is a tube conducting the ammonia vapour to the receiver A. All parts of the apparatus should be hermetically joined together, and should be able to withstand a pressure reaching ten atmospheres. The apparatus should be freed from air, which would otherwise hinder the liquefaction of the ammonia. The process is carried on as follows:—The apparatus is first so inclined that any liquid remaining in A may flow into C. The vessel C is then placed upon a stove F, and heated until the thermometertindicates a temperature of 130° C. During this time the ammonia has been expelled from C, and has liquefied in A. In order to facilitate the liquefaction, the receiver A should be immersed in a tank of water R (seethe left-hand drawing in fig.44). After about half an hour, when it may be supposed that the ammonia has been expelled, the fire is removed from under C, and this is now immersed in the tank of water R. The apparatus is represented in this position in the right-hand drawing of fig.44. The liquefied ammonia then evaporates, and passes over into the water in C. This causes the temperature of A to fall considerably. The substance to be refrigerated is placed in a vessel G, in the cylindrical space inside the receiver A. The refrigeration is also kept on for about half an hour, and with an apparatus of ordinary dimensions (containing about two litres of ammonia solution), five kilograms of ice are produced by the consumption of one kilogram of coal. In industrial works more complicated types of Carré's machines are employed.[8]Below 15° (according to Isambert), the compound AgCl,3NH3is formed, and above 20° the compound 2AgCl,3NH3. The tension of the ammonia evolved from the latter substance is equal to the atmospheric pressure at 68°, whilst for AgCl,3NH3the pressures are equal at about 20°; consequently, at higher temperatures it is greater than the atmospheric pressure, whilst at lower temperatures the ammonia is absorbed and forms this compound. Consequently, all the phenomena of dissociation are here clearly to be observed. Joannis and Croisier (1894) investigated similar compounds with AgBr, AgI, AgCN and AgNO3, and found that they all give definite compounds with NH3, for instance AgBr,3NH3, 2AgBr,3NH3and AgBr,2NH3; they are all colourless, solid substances which decompose under the atmospheric pressure at +3·5, +34° and +51°.[9]The liquefaction of ammonia may be accomplished without an increase of pressure, by means of refrigeration alone, in a carefully prepared mixture of ice and calcium chloride (because the absolute boiling point of NH3is high, about +130°). It may even take place in the severe frosts of a Russian winter. The application of liquid ammonia as a motive power for engines forms a problem which has to a certain extent been solved by the French engineer Tellier.[10]The combustion of ammonia in oxygen may be effected by the aid of platinum. A small quantity of an aqueous solution of ammonia, containing about 20 p.c. of the gas, is poured into a wide-necked beaker of about one litre capacity. A gas-conducting tube about 10 mm. in diameter, and supplying oxygen, is immersed in the aqueous solution of ammonia. But before introducing the gas an incandescent platinum spiral is placed in the beaker; the ammonia in the presence of the platinum is oxidised and burns, whilst the platinum wire becomes still more incandescent. The solution of ammonia is heated, and oxygen passed through the solution. The oxygen, as it bubbles off from the ammonia solution, carries with it a part of the ammonia, and this mixture explodes on coming into contact with the incandescent platinum. This is followed by a certain cooling effect, owing to the combustion ceasing, but after a short interval this is renewed, so that one feeble explosion follows after another. During the period of oxidation without explosion, white vapours of ammonium nitrite and red-brown vapours of oxides of nitrogen make their appearance, while during the explosion there is complete combustion and consequently water and nitrogen are formed.[11]This may be verified by their densities. Nitrogen is 14 times denser than hydrogen, and ammonia is 8½ times. If 3 volumes of hydrogen with 1 volume of nitrogen gave 4 volumes of ammonia, then these 4 volumes would weigh 17 times as much as 1 volume of hydrogen; consequently 1 volume of ammonia would be 4¼ times heavier than the same volume of hydrogen. But if these 4 volumes only give 2 volumes of ammonia, the latter will be 8½ times as dense as hydrogen, which is found to be actually the case.[12]Aqueous solutions of ammonia are lighter than water, and at 15°, taking water at 4° = 10,000, their specific gravity, as dependent onp, or the percentage amount (by weight) of ammonia, is given by the expressions= 9,992 - 42·5p+ 0·21p2; for instance, with 10 p.c.s= 9,587. Iftrepresents the temperature between the limits of +10° and +20°, then the expression (15 -t)(1·5 + 0·14p) must be added to the formula for the specific gravity. Solutions containing more than 24 p.c. have not been sufficiently investigated in respect to the variation of their specific gravity. It is, however, easy to obtain more concentrated solutions, and at 0° solutions approaching NH3,H2O (48·6 p.c. NH3) in their composition, and of sp. gr. 0·85, may be prepared. But such solutions give up the bulk of their ammonia at the ordinary temperature, so that more than 24 p.c. NH3is rarely contained in solution. Ammoniacal solutions containing a considerable amount of ammonia give ice-like crystals which seem to contain ammonia at temperatures far below 0° (for instance, an 8 p.c. solution at -14°, the strongest solutions at -48°). The whole of the ammonia may be expelled from a solution by heating, even at a comparatively low temperature; hence on heating aqueous solutions containing ammonia a very strong solution of ammonia is obtained in the distillate. Alcohol, ether, and many other liquids are also capable of dissolving ammonia. Solutions of ammonia, when exposed to the atmosphere, give off a part of their ammonia in accordance with the laws of the solution of gases in liquids, which we have already considered. But the ammoniacal solutions at the same time absorb carbonic anhydride from the air, and ammonium carbonate remains in the solution.see captionFig.46.—Apparatus for preparing solutions of ammonia.Solutions of ammonia are required both for laboratory and factory operations, and have therefore to be frequently prepared. For this purpose the arrangement shown in fig.46is employed in the laboratory. In works the same arrangement is used, only on a larger scale (with earthenware or metallic vessels). The gas is prepared in the retort, from whence it is led into the two-necked globe A, and then through a series of Woulfe's bottles, B, C, D, E. The impurities spurting over collect in A, and the gas is dissolved in B, but the solution soon becomes saturated, and a purer (washed) ammonia passes over into the following vessels, in which only a pure solution is obtained. The bent funnel tube in the retort preserves the apparatus from the possibility both of the pressure of the gas evolved in it becoming too great (when the gas escapes through it into the air), and also from the pressure incidentally falling too low (for instance, owing to a cooling effect, or from the reaction stopping). If this takes place, the air passes into the retort, otherwise the liquid from B would be drawn into A. The safety tubes in each Woulfe's bottle, open at both ends, and immersed in the liquid, serve for the same purpose. Without them, in case of an accidental stoppage in the evolution of so soluble a gas as ammonia, the solution would be sucked from one vessel to another—for instance, from E into D, &c. In order to clearly see the necessity forsafety tubesin a gas apparatus, it must be remembered that thegaseous pressurein the interior of the arrangement must exceed the atmospheric pressure by the height of the sum of the columns of liquid through which the gas has to pass.[13]The analogy between the ammonium and sodium salts might seem to be destroyed by the fact that the latter are formed from the alkali or oxide and an acid, with the separation of water, whilst the ammonium salts are directly formed from ammonia and an acid, without the separation of water; but the analogy is restored if we compare soda to ammonia water, and liken caustic soda to a compound of ammonia with water. Then the very preparation of ammonium salts from such a hydrate of ammonia will completely resemble the preparation of sodium salts from soda. We may cite as an example the action of hydrochloric acid on both substances.NaHO+HCl=H2O+NaClSodium hydroxideHydrochloric acidWaterTable saltNH4HO+HCl=H2O+NH4ClAmmonium hydroxideHydrochloric acidWaterSal-ammoniacJust as in soda the hydroxyl or aqueous radicle OH is replaced by chlorine, so it is in ammonia hydrate.[14]Weyl (1864) by subjecting sodium to the action of ammonia at the ordinary temperature and under considerable pressures, obtained a liquid, which was subsequently investigated by Joannis (1889), who confirmed the results obtained by Weyl. At 0° and the atmospheric pressure the composition of this substance is Na + 5·3NH3. The removal (at 0°) of ammonia from the liquid gives a solid copper-red body having the composition NH3Na. The determination of the molecular weight of this substance by the fall of the tension of liquid ammonia gave N2H6Na2. It is, therefore, free ammonium in which one H is replaced by Na. The compound with potassium, obtained under the same conditions, proved to have an analogous composition. By the decomposition of NH3Na at the ordinary temperature, Joannis (1891) obtained hydrogen and sodium-amide NH2Na in small colourless crystals which were soluble in water. The addition of liquid ammonia to metallic sodium and a saturated solution of sodium chloride, gives NH2Na2Cl, and this substance is sal-ammoniac, in which H2is replaced by Na2.If pure oxygen be passed through a solution of these compounds in ammonia at a temperature of about -50°, it is seen that the gas is rapidly absorbed. The liquid gradually loses its dark red colour and becomes lighter, and when it has become quite colourless a gelatinous precipitate is thrown down. After the removal of the ammonia, this precipitate dissolves easily in water with a considerable evolution of heat, but without giving off any gaseous products. The composition of the sodium compound thus obtained is NH2Na2HO, which shows that it is a hydrate of bisodium-ammonium. Thus, although free ammonium has not been obtained, still a sodium substitution product of it is known which corresponds to it as a salt to a hydrate. Ammonium amalgam was originally obtained in exactly the same way as sodium amalgam (Davy); namely, a piece of sal-ammoniac was taken, and moistened with water (in order to render it a conductor of electricity). A cavity was made in it, into which mercury was poured, and it was laid on a sheet of platinum connected with the positive pole of a galvanic battery, while the negative pole was put into connection with the mercury. On passing a current the mercury increased considerably in volume, and became plastic, whilst preserving its metallic appearance, just as would be the case were the sal-ammoniac replaced by a lump of a sodium salt or of many other metals. In the analogous decomposition of common metallic salts, the metal contained in a given salt separates out at the negative pole, immersed in mercury, by which the metal is dissolved. A similar phenomenon is observed in the case of sal-ammoniac; the elements of ammonium, NH4, in this case are also collected in the mercury, and are retained by it for a certain time.[15]We may mention, however, that under particular conditions hydrogen is also capable of forming an amalgam resembling the amalgam of ammonium. If an amalgam of zinc be shaken up with an aqueous solution of platinum chloride, without access of air, then a spongy mass is formed which easily decomposes, with the evolution of hydrogen.[16]We saw above that the solubility of ammonia in water at low temperatures attains to the molecular ratio NH3+ H2O, in which these substances are contained in caustic ammonia, and perhaps it may be possible at exceedingly low temperatures to obtain ammonium hydroxide, NH4HO, in a solid form. Regarding solutions as dissociated definite compounds, we should see a confirmation of this view in the property shown by ammonia of being extremely soluble in water, and in so doing of approaching to the limit NH4HO.[17]In confirmation of the truth of this conclusion we may cite the remarkable fact that there exist, in a free state and as comparatively stable compounds, a series of alkaline hydroxides, NR4HO, which are perfectly analogous to ammonium hydroxide, and present a striking resemblance to it and to sodium hydroxide, with the only difference that the hydrogen in NH4HO is replaced by complex groups, R = CH3, C2H5, &c., for instance N(CH3)4HO. Details will be found in organic chemistry.[18]The fact that ammoniacal salts are decomposed when ignited, and not simply sublimed, may be proved by a direct experiment with sal-ammoniac, NH4Cl, which in a state of vapour is decomposed into ammonia, NH3, and hydrochloric acid, HCl, as will be explained in thefollowing chapter. The readiness with which ammonium salts decompose is seen from the fact that a solution of ammonium oxalate is decomposed with the evolution of ammonia even at -1°. Dilute solutions of ammonium salts, when boiled give aqueous vapour having an alkaline reaction, owing to the presence of free ammonia given off from the salt.[19]Isambert studied the dissociation of ammoniacal compounds, as we have seen in Note 8, and showed that at low temperatures many salts are able to combine with a still greater amount of ammonia, which proves an entire analogy with hydrates; and as in this case it is easy to isolate the definite compounds, and as the least possible tension of ammonia is greater than that of water, therefore the ammoniacal compounds present a great and peculiar interest, as a means for explaining the nature of aqueous solutions and as a confirmation of the hypothesis of the formation of definite compounds in them; for these reasons we shall frequently refer to these compounds in the further exposition of this work.[20]Chapter V., Note2.[20 bis]Imide, NH, has not been obtained in a free state, but its hydrochloric acid salt, NHHCl, has apparently been obtained (1890) by Maumené by igniting the double bichloride of platinum and ammonium chloride, PtCl2NH4Cl = Pt + 2HCl + NHHCl. It is soluble in water, and crystallises from its solution in hexagonal rhombic prisms. It gives a double salt with FeCl3of the composition FeCl33NHHCl. The salt NHHCl is similar (isomeric) with the first possible product of the metalepsis of ammonia, NH2Cl, although it does not resemble it in any of its properties.[21]Freeamidogenorhydrazine, N2H4, or 2NH2, was prepared by Curtius (1887) by means of ethyl diazoacetate, or triazoacetic acid. Curtius and Jay (1889) showed that triazoacetic acid, CHN2.COOH (the formula should be tripled), when heated with water or a mineral acid, gives (quantitatively) oxalic acid and amidogen (hydrazine), CHN2.COOH + 2H2O = C2O2(OH)2+ N2H4—i.e.(empirically), the oxygen of the water replaces the nitrogen of the azoacetic acid. The amidogen is thus obtained in the form of a salt. With acids, amidogen forms very stable salts of the two types N2H4HX and N_[2]H4H2X2, as, for example, with HCl, H2SO4, &c. These salts are easily crystallised; in acid solutions they act as powerful reducing agents, evolving nitrogen; when ignited they are decomposed into ammoniacal salts, nitrogen, and hydrogen; with nitrites they evolve nitrogen. The sulphate N2H4,H2SO4is sparingly soluble in cold water (3 parts in 100 of water), but is very soluble in hot water; its specific gravity is 1·378, it fuses at 254° with decomposition. The hydrochloride N2H4,2HCl crystallises in octahedra, is very soluble in water, but not in alcohol; it fuses at 198°, evolving hydrogen chloride and forming the salt N2H4HCl; when rapidly heated it decomposes with an explosion; with platinic chloride it immediately evolves nitrogen, forming platinous chloride. By the action of alkalis the salts N2H4,2HX givehydrate of amidogen, N2H4,H2O, which is a fuming liquid (specific gravity 1·03), boiling at 119°, almost without odour, and whose aqueous solution corrodes glass and india-rubber, has an alkaline taste and poisonous properties. The reducing capacities of the hydrate are clearly seen from the fact that it reduces the metals platinum and silver from their solutions. With mercuric oxide it explodes. It reacts directly with the aldehydes RO, forming N2R2and water; for example, with benzaldehydes it gives the very stable insolublebenzalazine(C6H5CHN)2of a yellow colour. We may add that hydrazine often forms double salts; for example, MgSO4N2H4H2SO4or KClN2H4HCl, and that it is also formed by the action of nitrous acid upon aldehyde-ammonia. The products of the substitution of the hydrogen in hydrazine by hydrocarbon groups R (R = CH3, C2H5, C6H5, &c.) were obtained before hydrazine itself; for example, NHRNH2, NR2NH2, and (NRH)2.The heat of solution of the sulphuric acid salt (1 part in 200 and 300 parts of water at 10°·8) is equal to -8·7 C. According to Berthelot and Matigon (1892), the heat of neutralisation of hydrazine by sulphuric acid is +5·5 C and by hydrochloric acid +5·2 C. Thus hydrazine is a very feeble base, for its heat of saturation is not only lower than that of ammonia (+12·4 C. for HCl), but even below that of hydroxylamine (+9·3 C.) The heat of formation from the elements of hydrated hydrazine -9·5 C was deduced from the heat of combustion, determined by burning N2H4H2SO4in a calorimetric bomb, +127·7 C. Thus hydrazine is an endothermal compound; its passage into ammonia by the combination of hydrogen is accompanied by the evolution of 51·5 C. In the presence of an acid these figures were greater by +14·4 C. Hence the direct converse passage from ammonia into hydrazine is impossible. As regards the passage of hydroxylamine into hydrazine, it would be accompanied by the evolution of heat (+21·5 C.) in an aqueous solution.Amidogen must be regarded as a compound which stands to ammonia in the same relation as hydrogen peroxide stands to water. Water, H(OH), gives, according to the law of substitution, as was clearly to be expected, (OH)(OH)—that is, peroxide of hydrogen is the free radicle of water (hydroxyl). So also ammonia, H(NH2), forms hydrazine, (NH2)(NH2)—that is, the free radicle of ammonia, NH2, or amidogen. In the case of phosphorus a similar substance, as we shall afterwards see, has long been known under the name of liquid phosphuretted hydrogen, P2H4.[21 bis]In practice, the applications of ammonia are very varied. The use of ammonia as a stimulant, in the forms of the so-called ‘smelling salts’ or of spirits of hartshorn, in cases of faintness, &c., is known to everyone. The volatile carbonate of ammonium, or a mixture of an ammonium salt with an alkali, is also employed for this purpose. Ammonia also produces a well-known stimulating effect when rubbed on the skin, for which reason it is sometimes employed for external applications. Thus, for instance, the well-known volatile salve is prepared from any liquid oil shaken up with a solution of ammonia. A portion of the oil is thus transformed into a soapy substance. The solubility of greasy substances in ammonia, which proceeds from the formation both of emulsions and soaps, explains its use in extracting grease spots. It is also employed as an external application for stings from insects, and for bites from poisonous snakes, and in general in medicine. It is also remarkable that in cases of drunkenness a few drops of ammonia in water taken internally rapidly renders a person sober. A large quantity of ammonia is used in dyeing, either for the solution of certain dyes—for example, carmine—or for changing the tints of others, or else for neutralising the action of acids. It is also employed in the manufacture of artificial pearls. For this purpose the small scales of a peculiar small fish are mixed with ammonia, and the liquid so obtained is blown into small hollow glass beads shaped like pearls.In nature and the arts, however, ammonium salts, and not free ammonia, are most frequently employed. In this form a portion of thatnitrogenwhich is necessary for the formation of albuminous substances issupplied to plants. Owing to this, a large quantity of ammonium sulphate is now employed as a fertilising substance. But the same effect may be produced by nitre, or by animal refuse, which in decomposing gives ammonia. For this reason, an ammoniacal (hydrogen) compound may be introduced into the soil in the spring which will be converted into a nitrate (oxygen salt) in the summer.[22]As certain basic hydrates form peculiar compounds with ammonia, in some cases it happens that the first portions of ammonia added to a solution of a salt produce a precipitate, whilst the addition of a fresh quantity of ammonia dissolves this precipitate if the ammoniacal compound of the base be soluble in water. This, for example, takes place with the copper salts. But alumina does not dissolve under these circumstances.[23]When the element chlorine, as we shall afterwards more fully learn, replaces the element hydrogen, the reaction by which such an exchange is accomplished proceeds as a substitution, AH + Cl2= ACl + HCl, so that two substances, AH and chlorine, react on each other, and two substances, ACl and HCl, are formed; and further, two molecules react on each other, and two others are formed. The reaction proceeds very easily, but the substitution of one element,A, by another,X, does not always proceed with such ease, clearness, or simplicity. The substitution between oxygen and hydrogen is very rarely accomplished by the reaction of the free elements, but the substitution between these elements, one for another, forms the most common case of oxidation and reduction. In speaking of the law of substitution, I have in view the substitution of the elements one by another, and not the direct reaction of substitution. The law of substitution determines the cycle of the combinations of a given element, if a few of its compounds (for instance, the hydrogen compounds) be known. A development of the conceptions of the law of substitution may be found in my lecture given at the Royal Institution in London, 1889.[24]If hydrogen peroxide be taken as a starting point, then still higher forms of oxidation than those corresponding with water should be looked for. They should possess the properties of hydrogen peroxide, especially that of parting with their oxygen with extreme ease (even by contact). Such compounds are known. Pernitric, persulphuric, and similar acids present these properties, as we shall see in describing them.[25]The compound of hydroxylamine with hydrochloric acid has the composition NH2(OH)HCl = NH4ClO—that is, it is as it were oxidised sal-ammoniac. It was prepared by Lossen in 1865 by the action of tin and hydrochloric acid in the presence of water on a substance called ethyl nitrate, in which case the hydrogen liberated from the hydrochloric acid by the tin acts upon the elements of nitric acid—C2H5·NO3+6H+HCl=NH4OCl+H2O+C2H5·OHEthyl nitrateHydrogen fromHCl and SnHydroxylamine + HClWaterAlcoholThus in this case the nitric acid is deoxidised, not directly into nitrogen, but into hydroxylamine. Hydroxylamine is also formed by passing nitric oxide, NO, into a mixture of tin and hydrochloric acid—that is, by the action of the hydrogen evolved on the nitric oxide, NO + 3H + HCl = NH4OCl—and in many other cases. According to Lossen's method, a mixture of 30 parts of ethyl nitrate, 120 parts of tin, and 40 parts of a solution of hydrochloric acid of sp. gr. 1·06 are taken. After a certain time the reaction commences spontaneously. When the reaction has ceased the tin is separated by means of hydrogen sulphide, the solution is evaporated, and a large amount of sal-ammoniac is thus obtained (owing to the further action of hydrogen on the hydroxylamine compound, the hydrogen taking up oxygen from it and forming water); a solution ultimately remains containing the hydroxylamine salt; this salt is dissolved in anhydrous alcohol and purified by the addition of platinum chloride, which precipitates any ammonium salt still remaining in the solution. After concentrating the alcoholic solution the hydroxylamine hydrochloride separates in crystals. This substance melts at about 150°, and in so doing decomposes into nitrogen, hydrogen chloride, water, and sal-ammoniac. A sulphuric acid compound of hydroxylamine may be obtained by mixing a solution of the above salt with sulphuric acid. The sulphate is also soluble in water like the hydrochloride; this shows that hydroxylamine, like ammonia itself, forms a series of salts in which one acid may be substituted for another. It might he expected that by mixing a strong solution of a hydroxylamine salt with a solution of a caustic alkali hydroxylamine itself would be liberated, just as an ammonia salt under these circumstances evolves ammonia; but the liberated hydroxylamine is immediately decomposed with the formation of nitrogen and ammonia (and probably nitrous oxide), 3NH3O = NH3+ 3H2O + N2. Dilute solutions give the same reaction, although very slowly, but by decomposing a solution of the sulphate with barium hydroxide a certain amount of hydroxylamine is obtained in solution (it is partly decomposed). Hydroxylamine in aqueous solution, like ammonia, precipitates basic hydrates, and it deoxidises the oxides of copper, silver, and other metals. Free hydroxylamine was obtained by Lobry de Bruyn (1891). It is a solid, colourless, crystalline substance, without odour, which does not melt below 27°. It has the property of dissolving metallic salts; for instance, sodium chloride. Hydroxylamine, when rapidly heated with platinum, decomposes with a flash and the formation of a yellow flame. It is almost insoluble in ordinary solvents like chloroform, benzine, acetic ether, and carbon bisulphide. Its aqueous solutions are tolerably stable, contain up to 60 per cent. (sp. gr. 1·15 at 20°), and may be kept for many weeks without undergoing any change. Lobry de Bruyn used the hydrochloric salt to prepare pure hydroxylamine. The salt was first treated with sodium methylate (CH3NaO), and then methyl alcohol was added to the mixture. The precipitated sodium chloride was separated from the solution by filtration. (The methyl alcohol is added to prevent the precipitated chloride of sodium from coating the insoluble hydrochloric salt of hydroxylamine.) The methyl alcohol was driven off under a pressure 150–200 mm., and after extracting a further portion of methyl alcohol by ether and several fractional distillations, a solution was obtained containing 70 per cent. of free hydroxylamine, 8 per cent. water, 9·9 per cent. chloride of sodium, and 12·1 per cent. of the hydrochloric salt of hydroxylamine. Pure free hydroxylamine, NH3O, is obtained by distilling under a pressure of 60 mm.; it then boils at 70°, and solidifies in a condenser cooled to 0° in the form of long needles. It melts at 33°, boils at 58° under a pressure of 22 mm., and has a sp. gr. of about 1·235 (Brühl). Under the action of NaHO it gives NH3and NHO2or N2O, and forms nitric acid (Kolotoff, 1893) under the action of oxidising agents. Hydroxylamine is obtained in a great number of cases, for instance by the action of tin on dilute nitric acid, and also by the action of zinc on ethyl nitrate and dilute hydrochloric acid, &c. The relation between hydroxylamine, NH2(OH), and nitrous acid, NO(OH), which is so clear in the sense of the law of substitutions, becomes a reality in those cases when reducing agents act on salts of nitrous acid. Thus Raschig (1888) proposed the following method for the preparation of the hydroxylamine sulphate. A mixture of strong solutions of potassium nitrite, KNO2, and hydroxide, KHO, in molecular proportions, is prepared and cooled. An excess of sulphurous anhydride is then passed into the mixture, and the solution boiled for a long time. A mixture of the sulphates of potassium and hydroxylamine is thus obtained: KNO2+ KHO + 2SO2+ 2H2O = NH2(OH),H2SO4+ K2SO4. The salts may be separated from each other by crystallisation.[25 bis]In order to illustrate the application of the law of substitution to a given case, and to show the connection between ammonia and the oxides of nitrogen, let us consider the possible products of an oxygen and hydroxyl substitution in caustic ammonia, NH4(OH). It is evident that the substitution of H by OH can give: (1) NH3(OH)2; (2) NH2(OH)3; (3) NH(OH)4; and (4) N(OH)5. They should all, like caustic ammonia itself, easily part with water and form products (hydroxylic) of the oxidation of ammonia. The first of them is the hydrate of hydroxylamine, NH2(OH) + H2O; the second, NH(OH)2+ H2O (and also the substance NH(OH)4or NH3O2), containing, as it does, both hydrogen and oxygen, is able to part with all its hydrogen in the form of water (which could not be done by the first product, since it contained too little oxygen), forming, as the ultimate product, 2NH2(OH)3- 5H2O = N2O—that is, it corresponds with nitrous oxide, or the lower degree of the oxidation of nitrogen. So, also, nitrous anhydride corresponds with the third of the above products, 2NH(OH)4- 5H2O = N2O3, and nitric anhydride with the fourth, 2N(OH)5- 5H2O = N2O5. As, in these three equations, two molecules of the substitution products (-5H2O) are taken, it is also possible to combine two different products in one equation. For instance, the third and fourth products: NH(OH)4+ N(OH)5- 5H2O corresponds to N2O4or 2NO2, that is, to peroxide of nitrogen. Thus all the five (see later) oxides of nitrogen, N2O, NO, N2O3, NO2, and N2O5, may be deduced from ammonia. The above may be expressed in a general form by the equation (it should be remarked that the composition of all the substitution products of caustic ammonia may be expressed by NH3O5 -a, whereavaries between 0 and 4):NH5O5-a+ NH5O5-b- 5H2O = N2O5-(a+b),wherea+bcan evidently be not greater than 5; whena+b= 5 we have N2—nitrogen, when = 4 we have N2O nitrous oxide; whena+b= 3 we have N2O2or NO—nitric oxide, and so on to N2O5, whena+b= 0. Besides which it is evident that intermediate products may correspond with (and hence also break up into) different starting points; for instance, N2O is obtained whena+b= 2, and this may occur either whena= 0 (nitric acid), andb= 2 (hydroxylamine), or whena=b= 1 (the third of the above substitution products).[26]Nitric acid corresponds with the anhydride N2O5, which will afterwards be described, but which must be regarded as the highest saline oxide of nitrogen, just as Na2O (and the hydroxide NaHO) in the case of sodium, although sodium forms a peroxide possessing the property of parting with its oxygen with the same ease as hydrogen peroxide, if not on heating, at all events in reactions—for instance, with acids. So also nitric acid has its corresponding peroxide, which may be called pernitric acid. Its composition is not well known—probably NHO4—so that its corresponding anhydride would be N2O7. It is formed by the action of a silent discharge on a mixture of nitrogen and oxygen, so that a portion of its oxygen is in a state similar to that in ozone. The instability of this substance (obtained by Hautefeuille, Chappuis, and Berthelot), which easily splits up with the formation of nitric peroxide, and its resemblance to persulphuric acid, which we shall afterwards describe, will permit our passing over the consideration of the little that is further known concerning it.
Footnotes:
[1]The ammonia in the air, water, and soil proceeds from the decomposition of the nitrogenous substances of plants and animals, and also probably from the reduction of nitrates. Ammonia is always formed in the rusting of iron. Its formation in this case depends in all probability on the decomposition of water, and on the action of the hydrogen at the moment of its evolution on the nitric acid contained in the air (Cloez), or on the formation of ammonium nitrite, which takes place under many circumstances. The evolution of vapours of ammonia compounds is sometimes observed in the vicinity of volcanoes. At a red heat nitrogen combines directly with B Ca Mg, and with many other metals, and these compounds, when heated with a caustic alkali, or in the presence of water, give ammonia (seeChapter XIV., Note14, and Chapter XVII., Note 12). These are examples of the indirect combination of nitrogen with hydrogen.
[1]The ammonia in the air, water, and soil proceeds from the decomposition of the nitrogenous substances of plants and animals, and also probably from the reduction of nitrates. Ammonia is always formed in the rusting of iron. Its formation in this case depends in all probability on the decomposition of water, and on the action of the hydrogen at the moment of its evolution on the nitric acid contained in the air (Cloez), or on the formation of ammonium nitrite, which takes place under many circumstances. The evolution of vapours of ammonia compounds is sometimes observed in the vicinity of volcanoes. At a red heat nitrogen combines directly with B Ca Mg, and with many other metals, and these compounds, when heated with a caustic alkali, or in the presence of water, give ammonia (seeChapter XIV., Note14, and Chapter XVII., Note 12). These are examples of the indirect combination of nitrogen with hydrogen.
[2]If a silent discharge or a series of electric sparks be passed through ammonia gas, it is decomposed into nitrogen and hydrogen. This is a phenomenon of dissociation; therefore, a series of sparks do not totally decompose the ammonia, but leave a certain portion undecomposed. One volume of nitrogen and three volumes of hydrogen are obtained from two volumes of ammonia decomposed. Ramsay and Young (1884) investigated the decomposition of NH3under the action of heat, and showed that at 500°, 1½ p.c. is decomposed, at 600° about 18 p.c., at 800° 65 p.c., but these results were hardly free from the influence of ‘contact.’ Thepresenceof free ammonia—that is, ammonia not combined with acids—in a gas or aqueous solution may be recognised by its characteristic smell. But many ammonia salts do not possess this smell. However, on the addition of an alkali (for instance, caustic lime, potash, or soda), they evolve ammonia gas, especially when heated. The presence of ammonia may be made visible by introducing a substance moistened with strong hydrochloric acid into its neighbourhood. A white cloud, or visible white vapour, then makes its appearance. This depends on the fact that both ammonia and hydrochloric acid are volatile, and on coming into contact with each other produce solid sal-ammoniac, NH4Cl, which forms a cloud. This test is usually made by dipping a glass rod into hydrochloric acid, and holding it over the vessel from which the ammonia is evolved. With small amounts of ammonia this test is, however, untrustworthy, as the white vapour is scarcely observable. In this case it is best to take paper moistened with mercurous nitrate, HgNO3. This paper turns black in the presence of ammonia, owing to the formation of a black compound of ammonia with mercurous oxide. The smallest traces of ammonia (for instance, in river water) may be detected by means of the so-called Nessler's reagent, containing a solution of mercuric chloride and potassium iodide, which forms a brown coloration or precipitate with the smallest quantities of ammonia. It will be useful here to give the thermochemical data (in thousands of units of heat, according to Thomsen), or the quantities of heatevolvedin the formation of ammonia and its compounds in quantities expressed by their formulæ. Thus, for instance, (N + H3) 26·7 indicates that 14 grams of nitrogen in combining with 3 grams of hydrogen develop sufficient heat to raise the temperature of 26·7 kilograms of water 1°. (NH3+ nH2O) 8·4 (heat of solution); (NH3,nH2O + HCl,nH2O) 12·3; (N + H4+ Cl) 90·6; (NH3+ HCl) 41·9.
[2]If a silent discharge or a series of electric sparks be passed through ammonia gas, it is decomposed into nitrogen and hydrogen. This is a phenomenon of dissociation; therefore, a series of sparks do not totally decompose the ammonia, but leave a certain portion undecomposed. One volume of nitrogen and three volumes of hydrogen are obtained from two volumes of ammonia decomposed. Ramsay and Young (1884) investigated the decomposition of NH3under the action of heat, and showed that at 500°, 1½ p.c. is decomposed, at 600° about 18 p.c., at 800° 65 p.c., but these results were hardly free from the influence of ‘contact.’ Thepresenceof free ammonia—that is, ammonia not combined with acids—in a gas or aqueous solution may be recognised by its characteristic smell. But many ammonia salts do not possess this smell. However, on the addition of an alkali (for instance, caustic lime, potash, or soda), they evolve ammonia gas, especially when heated. The presence of ammonia may be made visible by introducing a substance moistened with strong hydrochloric acid into its neighbourhood. A white cloud, or visible white vapour, then makes its appearance. This depends on the fact that both ammonia and hydrochloric acid are volatile, and on coming into contact with each other produce solid sal-ammoniac, NH4Cl, which forms a cloud. This test is usually made by dipping a glass rod into hydrochloric acid, and holding it over the vessel from which the ammonia is evolved. With small amounts of ammonia this test is, however, untrustworthy, as the white vapour is scarcely observable. In this case it is best to take paper moistened with mercurous nitrate, HgNO3. This paper turns black in the presence of ammonia, owing to the formation of a black compound of ammonia with mercurous oxide. The smallest traces of ammonia (for instance, in river water) may be detected by means of the so-called Nessler's reagent, containing a solution of mercuric chloride and potassium iodide, which forms a brown coloration or precipitate with the smallest quantities of ammonia. It will be useful here to give the thermochemical data (in thousands of units of heat, according to Thomsen), or the quantities of heatevolvedin the formation of ammonia and its compounds in quantities expressed by their formulæ. Thus, for instance, (N + H3) 26·7 indicates that 14 grams of nitrogen in combining with 3 grams of hydrogen develop sufficient heat to raise the temperature of 26·7 kilograms of water 1°. (NH3+ nH2O) 8·4 (heat of solution); (NH3,nH2O + HCl,nH2O) 12·3; (N + H4+ Cl) 90·6; (NH3+ HCl) 41·9.
[3]The same ammonia water is obtained, although in smaller quantities, in the dry distillation of plants and of coal, which consists of the remains of fossil plants. In all these cases the ammonia proceeds from the destruction of the complex nitrogenous substances occurring in plants and animals. The ammonia salts employed in the arts are prepared by this method.
[3]The same ammonia water is obtained, although in smaller quantities, in the dry distillation of plants and of coal, which consists of the remains of fossil plants. In all these cases the ammonia proceeds from the destruction of the complex nitrogenous substances occurring in plants and animals. The ammonia salts employed in the arts are prepared by this method.
[4]The technical methods for the preparation of ammonia water, and for the extraction of ammonia from it, are to a certain extent explained in the figures accompanying the text.
[4]The technical methods for the preparation of ammonia water, and for the extraction of ammonia from it, are to a certain extent explained in the figures accompanying the text.
[5]Usually these crystals are sublimed by heating them in crucibles or pots, when the vapours of sal-ammoniac condense on the cold covers as a crust, in which form the salt comes into the market.
[5]Usually these crystals are sublimed by heating them in crucibles or pots, when the vapours of sal-ammoniac condense on the cold covers as a crust, in which form the salt comes into the market.
[6]On a small scale ammonia may be prepared in a glass flask by mixing equal parts by weight of slaked lime and finely-powdered sal-ammoniac, the neck of the flask being connected with an arrangement for drying the gas obtained. In this instance neither calcium chloride nor sulphuric acid can be used for drying the gas, since both these substances absorb ammonia, and therefore solid caustic potash, which is capable of retaining the water, is employed. The gas-conducting tube leading from the desiccating apparatus is introduced into a mercury bath, if dry gaseous ammonia be required, because water cannot be employed in collecting ammonia gas. Ammonia was first obtained in this dry state by Priestley, and its composition was investigated by Berthollet at the end of the last century. Oxide of lead mixed with sal-ammoniac (Isambert) evolves ammonia with still greater ease than lime. The cause and process of the decomposition are almost the same, 2PbO + 2NH4Cl = Pb2OCl2+ H2O + 2NH3. Lead oxychloride is (probably) formed.
[6]On a small scale ammonia may be prepared in a glass flask by mixing equal parts by weight of slaked lime and finely-powdered sal-ammoniac, the neck of the flask being connected with an arrangement for drying the gas obtained. In this instance neither calcium chloride nor sulphuric acid can be used for drying the gas, since both these substances absorb ammonia, and therefore solid caustic potash, which is capable of retaining the water, is employed. The gas-conducting tube leading from the desiccating apparatus is introduced into a mercury bath, if dry gaseous ammonia be required, because water cannot be employed in collecting ammonia gas. Ammonia was first obtained in this dry state by Priestley, and its composition was investigated by Berthollet at the end of the last century. Oxide of lead mixed with sal-ammoniac (Isambert) evolves ammonia with still greater ease than lime. The cause and process of the decomposition are almost the same, 2PbO + 2NH4Cl = Pb2OCl2+ H2O + 2NH3. Lead oxychloride is (probably) formed.
[7]see captionFig.44.—Carré's apparatus. Described in text.This is evident from the fact that its absolute boiling point lies at about +130° (Chapter II., Note29). It may therefore be liquefied by pressure alone at the ordinary, and even at much higher temperatures. The latent heat of evaporation of 17 parts by weight of ammonia equals 4,400 units of heat, and hence liquid ammonia may be employed for the production of cold. Strong aqueous solutions of ammonia, which in parting with their ammonia act in a similar manner, are not unfrequently employed for this purpose. Suppose a saturated solution of ammonia to be contained in a closed vessel furnished with a receiver. If the ammoniacal solution be heated, the ammonia, with a small quantity of water, will pass off from the solution, and in accumulating in the apparatus will produce a considerable pressure, and will therefore liquefy in the cooler portions of the receiver. Hence liquid ammonia will be obtained in the receiver. The heating of the vessel containing the aqueous solution of ammonia is then stopped. After having been heated it contains only water, or a solution poor in ammonia. When once it begins to cool the ammonia vapours commence dissolving in it, the space becomes rarefied, and a rapid vaporisation of the liquefied ammonia left in the receiver takes place. In evaporating in the receiver it will cause the temperature in it to fall considerably, and will itself pass into the aqueous solution. In the end, the same ammoniacal solution as originally taken is re-obtained. Thus, in this case, on heating the vessel the pressure increases by itself, and on cooling it diminishes, so that here heat directly replaces mechanical work. This is the principle of the simplest forms ofCarré's ice-making machines, shown in fig.44. C is a vessel made of boiler plates into which the saturated solution of ammonia is poured; m is a tube conducting the ammonia vapour to the receiver A. All parts of the apparatus should be hermetically joined together, and should be able to withstand a pressure reaching ten atmospheres. The apparatus should be freed from air, which would otherwise hinder the liquefaction of the ammonia. The process is carried on as follows:—The apparatus is first so inclined that any liquid remaining in A may flow into C. The vessel C is then placed upon a stove F, and heated until the thermometertindicates a temperature of 130° C. During this time the ammonia has been expelled from C, and has liquefied in A. In order to facilitate the liquefaction, the receiver A should be immersed in a tank of water R (seethe left-hand drawing in fig.44). After about half an hour, when it may be supposed that the ammonia has been expelled, the fire is removed from under C, and this is now immersed in the tank of water R. The apparatus is represented in this position in the right-hand drawing of fig.44. The liquefied ammonia then evaporates, and passes over into the water in C. This causes the temperature of A to fall considerably. The substance to be refrigerated is placed in a vessel G, in the cylindrical space inside the receiver A. The refrigeration is also kept on for about half an hour, and with an apparatus of ordinary dimensions (containing about two litres of ammonia solution), five kilograms of ice are produced by the consumption of one kilogram of coal. In industrial works more complicated types of Carré's machines are employed.
[7]
see captionFig.44.—Carré's apparatus. Described in text.
Fig.44.—Carré's apparatus. Described in text.
This is evident from the fact that its absolute boiling point lies at about +130° (Chapter II., Note29). It may therefore be liquefied by pressure alone at the ordinary, and even at much higher temperatures. The latent heat of evaporation of 17 parts by weight of ammonia equals 4,400 units of heat, and hence liquid ammonia may be employed for the production of cold. Strong aqueous solutions of ammonia, which in parting with their ammonia act in a similar manner, are not unfrequently employed for this purpose. Suppose a saturated solution of ammonia to be contained in a closed vessel furnished with a receiver. If the ammoniacal solution be heated, the ammonia, with a small quantity of water, will pass off from the solution, and in accumulating in the apparatus will produce a considerable pressure, and will therefore liquefy in the cooler portions of the receiver. Hence liquid ammonia will be obtained in the receiver. The heating of the vessel containing the aqueous solution of ammonia is then stopped. After having been heated it contains only water, or a solution poor in ammonia. When once it begins to cool the ammonia vapours commence dissolving in it, the space becomes rarefied, and a rapid vaporisation of the liquefied ammonia left in the receiver takes place. In evaporating in the receiver it will cause the temperature in it to fall considerably, and will itself pass into the aqueous solution. In the end, the same ammoniacal solution as originally taken is re-obtained. Thus, in this case, on heating the vessel the pressure increases by itself, and on cooling it diminishes, so that here heat directly replaces mechanical work. This is the principle of the simplest forms ofCarré's ice-making machines, shown in fig.44. C is a vessel made of boiler plates into which the saturated solution of ammonia is poured; m is a tube conducting the ammonia vapour to the receiver A. All parts of the apparatus should be hermetically joined together, and should be able to withstand a pressure reaching ten atmospheres. The apparatus should be freed from air, which would otherwise hinder the liquefaction of the ammonia. The process is carried on as follows:—The apparatus is first so inclined that any liquid remaining in A may flow into C. The vessel C is then placed upon a stove F, and heated until the thermometertindicates a temperature of 130° C. During this time the ammonia has been expelled from C, and has liquefied in A. In order to facilitate the liquefaction, the receiver A should be immersed in a tank of water R (seethe left-hand drawing in fig.44). After about half an hour, when it may be supposed that the ammonia has been expelled, the fire is removed from under C, and this is now immersed in the tank of water R. The apparatus is represented in this position in the right-hand drawing of fig.44. The liquefied ammonia then evaporates, and passes over into the water in C. This causes the temperature of A to fall considerably. The substance to be refrigerated is placed in a vessel G, in the cylindrical space inside the receiver A. The refrigeration is also kept on for about half an hour, and with an apparatus of ordinary dimensions (containing about two litres of ammonia solution), five kilograms of ice are produced by the consumption of one kilogram of coal. In industrial works more complicated types of Carré's machines are employed.
[8]Below 15° (according to Isambert), the compound AgCl,3NH3is formed, and above 20° the compound 2AgCl,3NH3. The tension of the ammonia evolved from the latter substance is equal to the atmospheric pressure at 68°, whilst for AgCl,3NH3the pressures are equal at about 20°; consequently, at higher temperatures it is greater than the atmospheric pressure, whilst at lower temperatures the ammonia is absorbed and forms this compound. Consequently, all the phenomena of dissociation are here clearly to be observed. Joannis and Croisier (1894) investigated similar compounds with AgBr, AgI, AgCN and AgNO3, and found that they all give definite compounds with NH3, for instance AgBr,3NH3, 2AgBr,3NH3and AgBr,2NH3; they are all colourless, solid substances which decompose under the atmospheric pressure at +3·5, +34° and +51°.
[8]Below 15° (according to Isambert), the compound AgCl,3NH3is formed, and above 20° the compound 2AgCl,3NH3. The tension of the ammonia evolved from the latter substance is equal to the atmospheric pressure at 68°, whilst for AgCl,3NH3the pressures are equal at about 20°; consequently, at higher temperatures it is greater than the atmospheric pressure, whilst at lower temperatures the ammonia is absorbed and forms this compound. Consequently, all the phenomena of dissociation are here clearly to be observed. Joannis and Croisier (1894) investigated similar compounds with AgBr, AgI, AgCN and AgNO3, and found that they all give definite compounds with NH3, for instance AgBr,3NH3, 2AgBr,3NH3and AgBr,2NH3; they are all colourless, solid substances which decompose under the atmospheric pressure at +3·5, +34° and +51°.
[9]The liquefaction of ammonia may be accomplished without an increase of pressure, by means of refrigeration alone, in a carefully prepared mixture of ice and calcium chloride (because the absolute boiling point of NH3is high, about +130°). It may even take place in the severe frosts of a Russian winter. The application of liquid ammonia as a motive power for engines forms a problem which has to a certain extent been solved by the French engineer Tellier.
[9]The liquefaction of ammonia may be accomplished without an increase of pressure, by means of refrigeration alone, in a carefully prepared mixture of ice and calcium chloride (because the absolute boiling point of NH3is high, about +130°). It may even take place in the severe frosts of a Russian winter. The application of liquid ammonia as a motive power for engines forms a problem which has to a certain extent been solved by the French engineer Tellier.
[10]The combustion of ammonia in oxygen may be effected by the aid of platinum. A small quantity of an aqueous solution of ammonia, containing about 20 p.c. of the gas, is poured into a wide-necked beaker of about one litre capacity. A gas-conducting tube about 10 mm. in diameter, and supplying oxygen, is immersed in the aqueous solution of ammonia. But before introducing the gas an incandescent platinum spiral is placed in the beaker; the ammonia in the presence of the platinum is oxidised and burns, whilst the platinum wire becomes still more incandescent. The solution of ammonia is heated, and oxygen passed through the solution. The oxygen, as it bubbles off from the ammonia solution, carries with it a part of the ammonia, and this mixture explodes on coming into contact with the incandescent platinum. This is followed by a certain cooling effect, owing to the combustion ceasing, but after a short interval this is renewed, so that one feeble explosion follows after another. During the period of oxidation without explosion, white vapours of ammonium nitrite and red-brown vapours of oxides of nitrogen make their appearance, while during the explosion there is complete combustion and consequently water and nitrogen are formed.
[10]The combustion of ammonia in oxygen may be effected by the aid of platinum. A small quantity of an aqueous solution of ammonia, containing about 20 p.c. of the gas, is poured into a wide-necked beaker of about one litre capacity. A gas-conducting tube about 10 mm. in diameter, and supplying oxygen, is immersed in the aqueous solution of ammonia. But before introducing the gas an incandescent platinum spiral is placed in the beaker; the ammonia in the presence of the platinum is oxidised and burns, whilst the platinum wire becomes still more incandescent. The solution of ammonia is heated, and oxygen passed through the solution. The oxygen, as it bubbles off from the ammonia solution, carries with it a part of the ammonia, and this mixture explodes on coming into contact with the incandescent platinum. This is followed by a certain cooling effect, owing to the combustion ceasing, but after a short interval this is renewed, so that one feeble explosion follows after another. During the period of oxidation without explosion, white vapours of ammonium nitrite and red-brown vapours of oxides of nitrogen make their appearance, while during the explosion there is complete combustion and consequently water and nitrogen are formed.
[11]This may be verified by their densities. Nitrogen is 14 times denser than hydrogen, and ammonia is 8½ times. If 3 volumes of hydrogen with 1 volume of nitrogen gave 4 volumes of ammonia, then these 4 volumes would weigh 17 times as much as 1 volume of hydrogen; consequently 1 volume of ammonia would be 4¼ times heavier than the same volume of hydrogen. But if these 4 volumes only give 2 volumes of ammonia, the latter will be 8½ times as dense as hydrogen, which is found to be actually the case.
[11]This may be verified by their densities. Nitrogen is 14 times denser than hydrogen, and ammonia is 8½ times. If 3 volumes of hydrogen with 1 volume of nitrogen gave 4 volumes of ammonia, then these 4 volumes would weigh 17 times as much as 1 volume of hydrogen; consequently 1 volume of ammonia would be 4¼ times heavier than the same volume of hydrogen. But if these 4 volumes only give 2 volumes of ammonia, the latter will be 8½ times as dense as hydrogen, which is found to be actually the case.
[12]Aqueous solutions of ammonia are lighter than water, and at 15°, taking water at 4° = 10,000, their specific gravity, as dependent onp, or the percentage amount (by weight) of ammonia, is given by the expressions= 9,992 - 42·5p+ 0·21p2; for instance, with 10 p.c.s= 9,587. Iftrepresents the temperature between the limits of +10° and +20°, then the expression (15 -t)(1·5 + 0·14p) must be added to the formula for the specific gravity. Solutions containing more than 24 p.c. have not been sufficiently investigated in respect to the variation of their specific gravity. It is, however, easy to obtain more concentrated solutions, and at 0° solutions approaching NH3,H2O (48·6 p.c. NH3) in their composition, and of sp. gr. 0·85, may be prepared. But such solutions give up the bulk of their ammonia at the ordinary temperature, so that more than 24 p.c. NH3is rarely contained in solution. Ammoniacal solutions containing a considerable amount of ammonia give ice-like crystals which seem to contain ammonia at temperatures far below 0° (for instance, an 8 p.c. solution at -14°, the strongest solutions at -48°). The whole of the ammonia may be expelled from a solution by heating, even at a comparatively low temperature; hence on heating aqueous solutions containing ammonia a very strong solution of ammonia is obtained in the distillate. Alcohol, ether, and many other liquids are also capable of dissolving ammonia. Solutions of ammonia, when exposed to the atmosphere, give off a part of their ammonia in accordance with the laws of the solution of gases in liquids, which we have already considered. But the ammoniacal solutions at the same time absorb carbonic anhydride from the air, and ammonium carbonate remains in the solution.see captionFig.46.—Apparatus for preparing solutions of ammonia.Solutions of ammonia are required both for laboratory and factory operations, and have therefore to be frequently prepared. For this purpose the arrangement shown in fig.46is employed in the laboratory. In works the same arrangement is used, only on a larger scale (with earthenware or metallic vessels). The gas is prepared in the retort, from whence it is led into the two-necked globe A, and then through a series of Woulfe's bottles, B, C, D, E. The impurities spurting over collect in A, and the gas is dissolved in B, but the solution soon becomes saturated, and a purer (washed) ammonia passes over into the following vessels, in which only a pure solution is obtained. The bent funnel tube in the retort preserves the apparatus from the possibility both of the pressure of the gas evolved in it becoming too great (when the gas escapes through it into the air), and also from the pressure incidentally falling too low (for instance, owing to a cooling effect, or from the reaction stopping). If this takes place, the air passes into the retort, otherwise the liquid from B would be drawn into A. The safety tubes in each Woulfe's bottle, open at both ends, and immersed in the liquid, serve for the same purpose. Without them, in case of an accidental stoppage in the evolution of so soluble a gas as ammonia, the solution would be sucked from one vessel to another—for instance, from E into D, &c. In order to clearly see the necessity forsafety tubesin a gas apparatus, it must be remembered that thegaseous pressurein the interior of the arrangement must exceed the atmospheric pressure by the height of the sum of the columns of liquid through which the gas has to pass.
[12]Aqueous solutions of ammonia are lighter than water, and at 15°, taking water at 4° = 10,000, their specific gravity, as dependent onp, or the percentage amount (by weight) of ammonia, is given by the expressions= 9,992 - 42·5p+ 0·21p2; for instance, with 10 p.c.s= 9,587. Iftrepresents the temperature between the limits of +10° and +20°, then the expression (15 -t)(1·5 + 0·14p) must be added to the formula for the specific gravity. Solutions containing more than 24 p.c. have not been sufficiently investigated in respect to the variation of their specific gravity. It is, however, easy to obtain more concentrated solutions, and at 0° solutions approaching NH3,H2O (48·6 p.c. NH3) in their composition, and of sp. gr. 0·85, may be prepared. But such solutions give up the bulk of their ammonia at the ordinary temperature, so that more than 24 p.c. NH3is rarely contained in solution. Ammoniacal solutions containing a considerable amount of ammonia give ice-like crystals which seem to contain ammonia at temperatures far below 0° (for instance, an 8 p.c. solution at -14°, the strongest solutions at -48°). The whole of the ammonia may be expelled from a solution by heating, even at a comparatively low temperature; hence on heating aqueous solutions containing ammonia a very strong solution of ammonia is obtained in the distillate. Alcohol, ether, and many other liquids are also capable of dissolving ammonia. Solutions of ammonia, when exposed to the atmosphere, give off a part of their ammonia in accordance with the laws of the solution of gases in liquids, which we have already considered. But the ammoniacal solutions at the same time absorb carbonic anhydride from the air, and ammonium carbonate remains in the solution.
see captionFig.46.—Apparatus for preparing solutions of ammonia.
Fig.46.—Apparatus for preparing solutions of ammonia.
Solutions of ammonia are required both for laboratory and factory operations, and have therefore to be frequently prepared. For this purpose the arrangement shown in fig.46is employed in the laboratory. In works the same arrangement is used, only on a larger scale (with earthenware or metallic vessels). The gas is prepared in the retort, from whence it is led into the two-necked globe A, and then through a series of Woulfe's bottles, B, C, D, E. The impurities spurting over collect in A, and the gas is dissolved in B, but the solution soon becomes saturated, and a purer (washed) ammonia passes over into the following vessels, in which only a pure solution is obtained. The bent funnel tube in the retort preserves the apparatus from the possibility both of the pressure of the gas evolved in it becoming too great (when the gas escapes through it into the air), and also from the pressure incidentally falling too low (for instance, owing to a cooling effect, or from the reaction stopping). If this takes place, the air passes into the retort, otherwise the liquid from B would be drawn into A. The safety tubes in each Woulfe's bottle, open at both ends, and immersed in the liquid, serve for the same purpose. Without them, in case of an accidental stoppage in the evolution of so soluble a gas as ammonia, the solution would be sucked from one vessel to another—for instance, from E into D, &c. In order to clearly see the necessity forsafety tubesin a gas apparatus, it must be remembered that thegaseous pressurein the interior of the arrangement must exceed the atmospheric pressure by the height of the sum of the columns of liquid through which the gas has to pass.
[13]The analogy between the ammonium and sodium salts might seem to be destroyed by the fact that the latter are formed from the alkali or oxide and an acid, with the separation of water, whilst the ammonium salts are directly formed from ammonia and an acid, without the separation of water; but the analogy is restored if we compare soda to ammonia water, and liken caustic soda to a compound of ammonia with water. Then the very preparation of ammonium salts from such a hydrate of ammonia will completely resemble the preparation of sodium salts from soda. We may cite as an example the action of hydrochloric acid on both substances.NaHO+HCl=H2O+NaClSodium hydroxideHydrochloric acidWaterTable saltNH4HO+HCl=H2O+NH4ClAmmonium hydroxideHydrochloric acidWaterSal-ammoniacJust as in soda the hydroxyl or aqueous radicle OH is replaced by chlorine, so it is in ammonia hydrate.
[13]The analogy between the ammonium and sodium salts might seem to be destroyed by the fact that the latter are formed from the alkali or oxide and an acid, with the separation of water, whilst the ammonium salts are directly formed from ammonia and an acid, without the separation of water; but the analogy is restored if we compare soda to ammonia water, and liken caustic soda to a compound of ammonia with water. Then the very preparation of ammonium salts from such a hydrate of ammonia will completely resemble the preparation of sodium salts from soda. We may cite as an example the action of hydrochloric acid on both substances.
Just as in soda the hydroxyl or aqueous radicle OH is replaced by chlorine, so it is in ammonia hydrate.
[14]Weyl (1864) by subjecting sodium to the action of ammonia at the ordinary temperature and under considerable pressures, obtained a liquid, which was subsequently investigated by Joannis (1889), who confirmed the results obtained by Weyl. At 0° and the atmospheric pressure the composition of this substance is Na + 5·3NH3. The removal (at 0°) of ammonia from the liquid gives a solid copper-red body having the composition NH3Na. The determination of the molecular weight of this substance by the fall of the tension of liquid ammonia gave N2H6Na2. It is, therefore, free ammonium in which one H is replaced by Na. The compound with potassium, obtained under the same conditions, proved to have an analogous composition. By the decomposition of NH3Na at the ordinary temperature, Joannis (1891) obtained hydrogen and sodium-amide NH2Na in small colourless crystals which were soluble in water. The addition of liquid ammonia to metallic sodium and a saturated solution of sodium chloride, gives NH2Na2Cl, and this substance is sal-ammoniac, in which H2is replaced by Na2.If pure oxygen be passed through a solution of these compounds in ammonia at a temperature of about -50°, it is seen that the gas is rapidly absorbed. The liquid gradually loses its dark red colour and becomes lighter, and when it has become quite colourless a gelatinous precipitate is thrown down. After the removal of the ammonia, this precipitate dissolves easily in water with a considerable evolution of heat, but without giving off any gaseous products. The composition of the sodium compound thus obtained is NH2Na2HO, which shows that it is a hydrate of bisodium-ammonium. Thus, although free ammonium has not been obtained, still a sodium substitution product of it is known which corresponds to it as a salt to a hydrate. Ammonium amalgam was originally obtained in exactly the same way as sodium amalgam (Davy); namely, a piece of sal-ammoniac was taken, and moistened with water (in order to render it a conductor of electricity). A cavity was made in it, into which mercury was poured, and it was laid on a sheet of platinum connected with the positive pole of a galvanic battery, while the negative pole was put into connection with the mercury. On passing a current the mercury increased considerably in volume, and became plastic, whilst preserving its metallic appearance, just as would be the case were the sal-ammoniac replaced by a lump of a sodium salt or of many other metals. In the analogous decomposition of common metallic salts, the metal contained in a given salt separates out at the negative pole, immersed in mercury, by which the metal is dissolved. A similar phenomenon is observed in the case of sal-ammoniac; the elements of ammonium, NH4, in this case are also collected in the mercury, and are retained by it for a certain time.
[14]Weyl (1864) by subjecting sodium to the action of ammonia at the ordinary temperature and under considerable pressures, obtained a liquid, which was subsequently investigated by Joannis (1889), who confirmed the results obtained by Weyl. At 0° and the atmospheric pressure the composition of this substance is Na + 5·3NH3. The removal (at 0°) of ammonia from the liquid gives a solid copper-red body having the composition NH3Na. The determination of the molecular weight of this substance by the fall of the tension of liquid ammonia gave N2H6Na2. It is, therefore, free ammonium in which one H is replaced by Na. The compound with potassium, obtained under the same conditions, proved to have an analogous composition. By the decomposition of NH3Na at the ordinary temperature, Joannis (1891) obtained hydrogen and sodium-amide NH2Na in small colourless crystals which were soluble in water. The addition of liquid ammonia to metallic sodium and a saturated solution of sodium chloride, gives NH2Na2Cl, and this substance is sal-ammoniac, in which H2is replaced by Na2.
If pure oxygen be passed through a solution of these compounds in ammonia at a temperature of about -50°, it is seen that the gas is rapidly absorbed. The liquid gradually loses its dark red colour and becomes lighter, and when it has become quite colourless a gelatinous precipitate is thrown down. After the removal of the ammonia, this precipitate dissolves easily in water with a considerable evolution of heat, but without giving off any gaseous products. The composition of the sodium compound thus obtained is NH2Na2HO, which shows that it is a hydrate of bisodium-ammonium. Thus, although free ammonium has not been obtained, still a sodium substitution product of it is known which corresponds to it as a salt to a hydrate. Ammonium amalgam was originally obtained in exactly the same way as sodium amalgam (Davy); namely, a piece of sal-ammoniac was taken, and moistened with water (in order to render it a conductor of electricity). A cavity was made in it, into which mercury was poured, and it was laid on a sheet of platinum connected with the positive pole of a galvanic battery, while the negative pole was put into connection with the mercury. On passing a current the mercury increased considerably in volume, and became plastic, whilst preserving its metallic appearance, just as would be the case were the sal-ammoniac replaced by a lump of a sodium salt or of many other metals. In the analogous decomposition of common metallic salts, the metal contained in a given salt separates out at the negative pole, immersed in mercury, by which the metal is dissolved. A similar phenomenon is observed in the case of sal-ammoniac; the elements of ammonium, NH4, in this case are also collected in the mercury, and are retained by it for a certain time.
[15]We may mention, however, that under particular conditions hydrogen is also capable of forming an amalgam resembling the amalgam of ammonium. If an amalgam of zinc be shaken up with an aqueous solution of platinum chloride, without access of air, then a spongy mass is formed which easily decomposes, with the evolution of hydrogen.
[15]We may mention, however, that under particular conditions hydrogen is also capable of forming an amalgam resembling the amalgam of ammonium. If an amalgam of zinc be shaken up with an aqueous solution of platinum chloride, without access of air, then a spongy mass is formed which easily decomposes, with the evolution of hydrogen.
[16]We saw above that the solubility of ammonia in water at low temperatures attains to the molecular ratio NH3+ H2O, in which these substances are contained in caustic ammonia, and perhaps it may be possible at exceedingly low temperatures to obtain ammonium hydroxide, NH4HO, in a solid form. Regarding solutions as dissociated definite compounds, we should see a confirmation of this view in the property shown by ammonia of being extremely soluble in water, and in so doing of approaching to the limit NH4HO.
[16]We saw above that the solubility of ammonia in water at low temperatures attains to the molecular ratio NH3+ H2O, in which these substances are contained in caustic ammonia, and perhaps it may be possible at exceedingly low temperatures to obtain ammonium hydroxide, NH4HO, in a solid form. Regarding solutions as dissociated definite compounds, we should see a confirmation of this view in the property shown by ammonia of being extremely soluble in water, and in so doing of approaching to the limit NH4HO.
[17]In confirmation of the truth of this conclusion we may cite the remarkable fact that there exist, in a free state and as comparatively stable compounds, a series of alkaline hydroxides, NR4HO, which are perfectly analogous to ammonium hydroxide, and present a striking resemblance to it and to sodium hydroxide, with the only difference that the hydrogen in NH4HO is replaced by complex groups, R = CH3, C2H5, &c., for instance N(CH3)4HO. Details will be found in organic chemistry.
[17]In confirmation of the truth of this conclusion we may cite the remarkable fact that there exist, in a free state and as comparatively stable compounds, a series of alkaline hydroxides, NR4HO, which are perfectly analogous to ammonium hydroxide, and present a striking resemblance to it and to sodium hydroxide, with the only difference that the hydrogen in NH4HO is replaced by complex groups, R = CH3, C2H5, &c., for instance N(CH3)4HO. Details will be found in organic chemistry.
[18]The fact that ammoniacal salts are decomposed when ignited, and not simply sublimed, may be proved by a direct experiment with sal-ammoniac, NH4Cl, which in a state of vapour is decomposed into ammonia, NH3, and hydrochloric acid, HCl, as will be explained in thefollowing chapter. The readiness with which ammonium salts decompose is seen from the fact that a solution of ammonium oxalate is decomposed with the evolution of ammonia even at -1°. Dilute solutions of ammonium salts, when boiled give aqueous vapour having an alkaline reaction, owing to the presence of free ammonia given off from the salt.
[18]The fact that ammoniacal salts are decomposed when ignited, and not simply sublimed, may be proved by a direct experiment with sal-ammoniac, NH4Cl, which in a state of vapour is decomposed into ammonia, NH3, and hydrochloric acid, HCl, as will be explained in thefollowing chapter. The readiness with which ammonium salts decompose is seen from the fact that a solution of ammonium oxalate is decomposed with the evolution of ammonia even at -1°. Dilute solutions of ammonium salts, when boiled give aqueous vapour having an alkaline reaction, owing to the presence of free ammonia given off from the salt.
[19]Isambert studied the dissociation of ammoniacal compounds, as we have seen in Note 8, and showed that at low temperatures many salts are able to combine with a still greater amount of ammonia, which proves an entire analogy with hydrates; and as in this case it is easy to isolate the definite compounds, and as the least possible tension of ammonia is greater than that of water, therefore the ammoniacal compounds present a great and peculiar interest, as a means for explaining the nature of aqueous solutions and as a confirmation of the hypothesis of the formation of definite compounds in them; for these reasons we shall frequently refer to these compounds in the further exposition of this work.
[19]Isambert studied the dissociation of ammoniacal compounds, as we have seen in Note 8, and showed that at low temperatures many salts are able to combine with a still greater amount of ammonia, which proves an entire analogy with hydrates; and as in this case it is easy to isolate the definite compounds, and as the least possible tension of ammonia is greater than that of water, therefore the ammoniacal compounds present a great and peculiar interest, as a means for explaining the nature of aqueous solutions and as a confirmation of the hypothesis of the formation of definite compounds in them; for these reasons we shall frequently refer to these compounds in the further exposition of this work.
[20]Chapter V., Note2.
[20]Chapter V., Note2.
[20 bis]Imide, NH, has not been obtained in a free state, but its hydrochloric acid salt, NHHCl, has apparently been obtained (1890) by Maumené by igniting the double bichloride of platinum and ammonium chloride, PtCl2NH4Cl = Pt + 2HCl + NHHCl. It is soluble in water, and crystallises from its solution in hexagonal rhombic prisms. It gives a double salt with FeCl3of the composition FeCl33NHHCl. The salt NHHCl is similar (isomeric) with the first possible product of the metalepsis of ammonia, NH2Cl, although it does not resemble it in any of its properties.
[20 bis]Imide, NH, has not been obtained in a free state, but its hydrochloric acid salt, NHHCl, has apparently been obtained (1890) by Maumené by igniting the double bichloride of platinum and ammonium chloride, PtCl2NH4Cl = Pt + 2HCl + NHHCl. It is soluble in water, and crystallises from its solution in hexagonal rhombic prisms. It gives a double salt with FeCl3of the composition FeCl33NHHCl. The salt NHHCl is similar (isomeric) with the first possible product of the metalepsis of ammonia, NH2Cl, although it does not resemble it in any of its properties.
[21]Freeamidogenorhydrazine, N2H4, or 2NH2, was prepared by Curtius (1887) by means of ethyl diazoacetate, or triazoacetic acid. Curtius and Jay (1889) showed that triazoacetic acid, CHN2.COOH (the formula should be tripled), when heated with water or a mineral acid, gives (quantitatively) oxalic acid and amidogen (hydrazine), CHN2.COOH + 2H2O = C2O2(OH)2+ N2H4—i.e.(empirically), the oxygen of the water replaces the nitrogen of the azoacetic acid. The amidogen is thus obtained in the form of a salt. With acids, amidogen forms very stable salts of the two types N2H4HX and N_[2]H4H2X2, as, for example, with HCl, H2SO4, &c. These salts are easily crystallised; in acid solutions they act as powerful reducing agents, evolving nitrogen; when ignited they are decomposed into ammoniacal salts, nitrogen, and hydrogen; with nitrites they evolve nitrogen. The sulphate N2H4,H2SO4is sparingly soluble in cold water (3 parts in 100 of water), but is very soluble in hot water; its specific gravity is 1·378, it fuses at 254° with decomposition. The hydrochloride N2H4,2HCl crystallises in octahedra, is very soluble in water, but not in alcohol; it fuses at 198°, evolving hydrogen chloride and forming the salt N2H4HCl; when rapidly heated it decomposes with an explosion; with platinic chloride it immediately evolves nitrogen, forming platinous chloride. By the action of alkalis the salts N2H4,2HX givehydrate of amidogen, N2H4,H2O, which is a fuming liquid (specific gravity 1·03), boiling at 119°, almost without odour, and whose aqueous solution corrodes glass and india-rubber, has an alkaline taste and poisonous properties. The reducing capacities of the hydrate are clearly seen from the fact that it reduces the metals platinum and silver from their solutions. With mercuric oxide it explodes. It reacts directly with the aldehydes RO, forming N2R2and water; for example, with benzaldehydes it gives the very stable insolublebenzalazine(C6H5CHN)2of a yellow colour. We may add that hydrazine often forms double salts; for example, MgSO4N2H4H2SO4or KClN2H4HCl, and that it is also formed by the action of nitrous acid upon aldehyde-ammonia. The products of the substitution of the hydrogen in hydrazine by hydrocarbon groups R (R = CH3, C2H5, C6H5, &c.) were obtained before hydrazine itself; for example, NHRNH2, NR2NH2, and (NRH)2.The heat of solution of the sulphuric acid salt (1 part in 200 and 300 parts of water at 10°·8) is equal to -8·7 C. According to Berthelot and Matigon (1892), the heat of neutralisation of hydrazine by sulphuric acid is +5·5 C and by hydrochloric acid +5·2 C. Thus hydrazine is a very feeble base, for its heat of saturation is not only lower than that of ammonia (+12·4 C. for HCl), but even below that of hydroxylamine (+9·3 C.) The heat of formation from the elements of hydrated hydrazine -9·5 C was deduced from the heat of combustion, determined by burning N2H4H2SO4in a calorimetric bomb, +127·7 C. Thus hydrazine is an endothermal compound; its passage into ammonia by the combination of hydrogen is accompanied by the evolution of 51·5 C. In the presence of an acid these figures were greater by +14·4 C. Hence the direct converse passage from ammonia into hydrazine is impossible. As regards the passage of hydroxylamine into hydrazine, it would be accompanied by the evolution of heat (+21·5 C.) in an aqueous solution.Amidogen must be regarded as a compound which stands to ammonia in the same relation as hydrogen peroxide stands to water. Water, H(OH), gives, according to the law of substitution, as was clearly to be expected, (OH)(OH)—that is, peroxide of hydrogen is the free radicle of water (hydroxyl). So also ammonia, H(NH2), forms hydrazine, (NH2)(NH2)—that is, the free radicle of ammonia, NH2, or amidogen. In the case of phosphorus a similar substance, as we shall afterwards see, has long been known under the name of liquid phosphuretted hydrogen, P2H4.
[21]Freeamidogenorhydrazine, N2H4, or 2NH2, was prepared by Curtius (1887) by means of ethyl diazoacetate, or triazoacetic acid. Curtius and Jay (1889) showed that triazoacetic acid, CHN2.COOH (the formula should be tripled), when heated with water or a mineral acid, gives (quantitatively) oxalic acid and amidogen (hydrazine), CHN2.COOH + 2H2O = C2O2(OH)2+ N2H4—i.e.(empirically), the oxygen of the water replaces the nitrogen of the azoacetic acid. The amidogen is thus obtained in the form of a salt. With acids, amidogen forms very stable salts of the two types N2H4HX and N_[2]H4H2X2, as, for example, with HCl, H2SO4, &c. These salts are easily crystallised; in acid solutions they act as powerful reducing agents, evolving nitrogen; when ignited they are decomposed into ammoniacal salts, nitrogen, and hydrogen; with nitrites they evolve nitrogen. The sulphate N2H4,H2SO4is sparingly soluble in cold water (3 parts in 100 of water), but is very soluble in hot water; its specific gravity is 1·378, it fuses at 254° with decomposition. The hydrochloride N2H4,2HCl crystallises in octahedra, is very soluble in water, but not in alcohol; it fuses at 198°, evolving hydrogen chloride and forming the salt N2H4HCl; when rapidly heated it decomposes with an explosion; with platinic chloride it immediately evolves nitrogen, forming platinous chloride. By the action of alkalis the salts N2H4,2HX givehydrate of amidogen, N2H4,H2O, which is a fuming liquid (specific gravity 1·03), boiling at 119°, almost without odour, and whose aqueous solution corrodes glass and india-rubber, has an alkaline taste and poisonous properties. The reducing capacities of the hydrate are clearly seen from the fact that it reduces the metals platinum and silver from their solutions. With mercuric oxide it explodes. It reacts directly with the aldehydes RO, forming N2R2and water; for example, with benzaldehydes it gives the very stable insolublebenzalazine(C6H5CHN)2of a yellow colour. We may add that hydrazine often forms double salts; for example, MgSO4N2H4H2SO4or KClN2H4HCl, and that it is also formed by the action of nitrous acid upon aldehyde-ammonia. The products of the substitution of the hydrogen in hydrazine by hydrocarbon groups R (R = CH3, C2H5, C6H5, &c.) were obtained before hydrazine itself; for example, NHRNH2, NR2NH2, and (NRH)2.
The heat of solution of the sulphuric acid salt (1 part in 200 and 300 parts of water at 10°·8) is equal to -8·7 C. According to Berthelot and Matigon (1892), the heat of neutralisation of hydrazine by sulphuric acid is +5·5 C and by hydrochloric acid +5·2 C. Thus hydrazine is a very feeble base, for its heat of saturation is not only lower than that of ammonia (+12·4 C. for HCl), but even below that of hydroxylamine (+9·3 C.) The heat of formation from the elements of hydrated hydrazine -9·5 C was deduced from the heat of combustion, determined by burning N2H4H2SO4in a calorimetric bomb, +127·7 C. Thus hydrazine is an endothermal compound; its passage into ammonia by the combination of hydrogen is accompanied by the evolution of 51·5 C. In the presence of an acid these figures were greater by +14·4 C. Hence the direct converse passage from ammonia into hydrazine is impossible. As regards the passage of hydroxylamine into hydrazine, it would be accompanied by the evolution of heat (+21·5 C.) in an aqueous solution.
Amidogen must be regarded as a compound which stands to ammonia in the same relation as hydrogen peroxide stands to water. Water, H(OH), gives, according to the law of substitution, as was clearly to be expected, (OH)(OH)—that is, peroxide of hydrogen is the free radicle of water (hydroxyl). So also ammonia, H(NH2), forms hydrazine, (NH2)(NH2)—that is, the free radicle of ammonia, NH2, or amidogen. In the case of phosphorus a similar substance, as we shall afterwards see, has long been known under the name of liquid phosphuretted hydrogen, P2H4.
[21 bis]In practice, the applications of ammonia are very varied. The use of ammonia as a stimulant, in the forms of the so-called ‘smelling salts’ or of spirits of hartshorn, in cases of faintness, &c., is known to everyone. The volatile carbonate of ammonium, or a mixture of an ammonium salt with an alkali, is also employed for this purpose. Ammonia also produces a well-known stimulating effect when rubbed on the skin, for which reason it is sometimes employed for external applications. Thus, for instance, the well-known volatile salve is prepared from any liquid oil shaken up with a solution of ammonia. A portion of the oil is thus transformed into a soapy substance. The solubility of greasy substances in ammonia, which proceeds from the formation both of emulsions and soaps, explains its use in extracting grease spots. It is also employed as an external application for stings from insects, and for bites from poisonous snakes, and in general in medicine. It is also remarkable that in cases of drunkenness a few drops of ammonia in water taken internally rapidly renders a person sober. A large quantity of ammonia is used in dyeing, either for the solution of certain dyes—for example, carmine—or for changing the tints of others, or else for neutralising the action of acids. It is also employed in the manufacture of artificial pearls. For this purpose the small scales of a peculiar small fish are mixed with ammonia, and the liquid so obtained is blown into small hollow glass beads shaped like pearls.In nature and the arts, however, ammonium salts, and not free ammonia, are most frequently employed. In this form a portion of thatnitrogenwhich is necessary for the formation of albuminous substances issupplied to plants. Owing to this, a large quantity of ammonium sulphate is now employed as a fertilising substance. But the same effect may be produced by nitre, or by animal refuse, which in decomposing gives ammonia. For this reason, an ammoniacal (hydrogen) compound may be introduced into the soil in the spring which will be converted into a nitrate (oxygen salt) in the summer.
[21 bis]In practice, the applications of ammonia are very varied. The use of ammonia as a stimulant, in the forms of the so-called ‘smelling salts’ or of spirits of hartshorn, in cases of faintness, &c., is known to everyone. The volatile carbonate of ammonium, or a mixture of an ammonium salt with an alkali, is also employed for this purpose. Ammonia also produces a well-known stimulating effect when rubbed on the skin, for which reason it is sometimes employed for external applications. Thus, for instance, the well-known volatile salve is prepared from any liquid oil shaken up with a solution of ammonia. A portion of the oil is thus transformed into a soapy substance. The solubility of greasy substances in ammonia, which proceeds from the formation both of emulsions and soaps, explains its use in extracting grease spots. It is also employed as an external application for stings from insects, and for bites from poisonous snakes, and in general in medicine. It is also remarkable that in cases of drunkenness a few drops of ammonia in water taken internally rapidly renders a person sober. A large quantity of ammonia is used in dyeing, either for the solution of certain dyes—for example, carmine—or for changing the tints of others, or else for neutralising the action of acids. It is also employed in the manufacture of artificial pearls. For this purpose the small scales of a peculiar small fish are mixed with ammonia, and the liquid so obtained is blown into small hollow glass beads shaped like pearls.
In nature and the arts, however, ammonium salts, and not free ammonia, are most frequently employed. In this form a portion of thatnitrogenwhich is necessary for the formation of albuminous substances issupplied to plants. Owing to this, a large quantity of ammonium sulphate is now employed as a fertilising substance. But the same effect may be produced by nitre, or by animal refuse, which in decomposing gives ammonia. For this reason, an ammoniacal (hydrogen) compound may be introduced into the soil in the spring which will be converted into a nitrate (oxygen salt) in the summer.
[22]As certain basic hydrates form peculiar compounds with ammonia, in some cases it happens that the first portions of ammonia added to a solution of a salt produce a precipitate, whilst the addition of a fresh quantity of ammonia dissolves this precipitate if the ammoniacal compound of the base be soluble in water. This, for example, takes place with the copper salts. But alumina does not dissolve under these circumstances.
[22]As certain basic hydrates form peculiar compounds with ammonia, in some cases it happens that the first portions of ammonia added to a solution of a salt produce a precipitate, whilst the addition of a fresh quantity of ammonia dissolves this precipitate if the ammoniacal compound of the base be soluble in water. This, for example, takes place with the copper salts. But alumina does not dissolve under these circumstances.
[23]When the element chlorine, as we shall afterwards more fully learn, replaces the element hydrogen, the reaction by which such an exchange is accomplished proceeds as a substitution, AH + Cl2= ACl + HCl, so that two substances, AH and chlorine, react on each other, and two substances, ACl and HCl, are formed; and further, two molecules react on each other, and two others are formed. The reaction proceeds very easily, but the substitution of one element,A, by another,X, does not always proceed with such ease, clearness, or simplicity. The substitution between oxygen and hydrogen is very rarely accomplished by the reaction of the free elements, but the substitution between these elements, one for another, forms the most common case of oxidation and reduction. In speaking of the law of substitution, I have in view the substitution of the elements one by another, and not the direct reaction of substitution. The law of substitution determines the cycle of the combinations of a given element, if a few of its compounds (for instance, the hydrogen compounds) be known. A development of the conceptions of the law of substitution may be found in my lecture given at the Royal Institution in London, 1889.
[23]When the element chlorine, as we shall afterwards more fully learn, replaces the element hydrogen, the reaction by which such an exchange is accomplished proceeds as a substitution, AH + Cl2= ACl + HCl, so that two substances, AH and chlorine, react on each other, and two substances, ACl and HCl, are formed; and further, two molecules react on each other, and two others are formed. The reaction proceeds very easily, but the substitution of one element,A, by another,X, does not always proceed with such ease, clearness, or simplicity. The substitution between oxygen and hydrogen is very rarely accomplished by the reaction of the free elements, but the substitution between these elements, one for another, forms the most common case of oxidation and reduction. In speaking of the law of substitution, I have in view the substitution of the elements one by another, and not the direct reaction of substitution. The law of substitution determines the cycle of the combinations of a given element, if a few of its compounds (for instance, the hydrogen compounds) be known. A development of the conceptions of the law of substitution may be found in my lecture given at the Royal Institution in London, 1889.
[24]If hydrogen peroxide be taken as a starting point, then still higher forms of oxidation than those corresponding with water should be looked for. They should possess the properties of hydrogen peroxide, especially that of parting with their oxygen with extreme ease (even by contact). Such compounds are known. Pernitric, persulphuric, and similar acids present these properties, as we shall see in describing them.
[24]If hydrogen peroxide be taken as a starting point, then still higher forms of oxidation than those corresponding with water should be looked for. They should possess the properties of hydrogen peroxide, especially that of parting with their oxygen with extreme ease (even by contact). Such compounds are known. Pernitric, persulphuric, and similar acids present these properties, as we shall see in describing them.
[25]The compound of hydroxylamine with hydrochloric acid has the composition NH2(OH)HCl = NH4ClO—that is, it is as it were oxidised sal-ammoniac. It was prepared by Lossen in 1865 by the action of tin and hydrochloric acid in the presence of water on a substance called ethyl nitrate, in which case the hydrogen liberated from the hydrochloric acid by the tin acts upon the elements of nitric acid—C2H5·NO3+6H+HCl=NH4OCl+H2O+C2H5·OHEthyl nitrateHydrogen fromHCl and SnHydroxylamine + HClWaterAlcoholThus in this case the nitric acid is deoxidised, not directly into nitrogen, but into hydroxylamine. Hydroxylamine is also formed by passing nitric oxide, NO, into a mixture of tin and hydrochloric acid—that is, by the action of the hydrogen evolved on the nitric oxide, NO + 3H + HCl = NH4OCl—and in many other cases. According to Lossen's method, a mixture of 30 parts of ethyl nitrate, 120 parts of tin, and 40 parts of a solution of hydrochloric acid of sp. gr. 1·06 are taken. After a certain time the reaction commences spontaneously. When the reaction has ceased the tin is separated by means of hydrogen sulphide, the solution is evaporated, and a large amount of sal-ammoniac is thus obtained (owing to the further action of hydrogen on the hydroxylamine compound, the hydrogen taking up oxygen from it and forming water); a solution ultimately remains containing the hydroxylamine salt; this salt is dissolved in anhydrous alcohol and purified by the addition of platinum chloride, which precipitates any ammonium salt still remaining in the solution. After concentrating the alcoholic solution the hydroxylamine hydrochloride separates in crystals. This substance melts at about 150°, and in so doing decomposes into nitrogen, hydrogen chloride, water, and sal-ammoniac. A sulphuric acid compound of hydroxylamine may be obtained by mixing a solution of the above salt with sulphuric acid. The sulphate is also soluble in water like the hydrochloride; this shows that hydroxylamine, like ammonia itself, forms a series of salts in which one acid may be substituted for another. It might he expected that by mixing a strong solution of a hydroxylamine salt with a solution of a caustic alkali hydroxylamine itself would be liberated, just as an ammonia salt under these circumstances evolves ammonia; but the liberated hydroxylamine is immediately decomposed with the formation of nitrogen and ammonia (and probably nitrous oxide), 3NH3O = NH3+ 3H2O + N2. Dilute solutions give the same reaction, although very slowly, but by decomposing a solution of the sulphate with barium hydroxide a certain amount of hydroxylamine is obtained in solution (it is partly decomposed). Hydroxylamine in aqueous solution, like ammonia, precipitates basic hydrates, and it deoxidises the oxides of copper, silver, and other metals. Free hydroxylamine was obtained by Lobry de Bruyn (1891). It is a solid, colourless, crystalline substance, without odour, which does not melt below 27°. It has the property of dissolving metallic salts; for instance, sodium chloride. Hydroxylamine, when rapidly heated with platinum, decomposes with a flash and the formation of a yellow flame. It is almost insoluble in ordinary solvents like chloroform, benzine, acetic ether, and carbon bisulphide. Its aqueous solutions are tolerably stable, contain up to 60 per cent. (sp. gr. 1·15 at 20°), and may be kept for many weeks without undergoing any change. Lobry de Bruyn used the hydrochloric salt to prepare pure hydroxylamine. The salt was first treated with sodium methylate (CH3NaO), and then methyl alcohol was added to the mixture. The precipitated sodium chloride was separated from the solution by filtration. (The methyl alcohol is added to prevent the precipitated chloride of sodium from coating the insoluble hydrochloric salt of hydroxylamine.) The methyl alcohol was driven off under a pressure 150–200 mm., and after extracting a further portion of methyl alcohol by ether and several fractional distillations, a solution was obtained containing 70 per cent. of free hydroxylamine, 8 per cent. water, 9·9 per cent. chloride of sodium, and 12·1 per cent. of the hydrochloric salt of hydroxylamine. Pure free hydroxylamine, NH3O, is obtained by distilling under a pressure of 60 mm.; it then boils at 70°, and solidifies in a condenser cooled to 0° in the form of long needles. It melts at 33°, boils at 58° under a pressure of 22 mm., and has a sp. gr. of about 1·235 (Brühl). Under the action of NaHO it gives NH3and NHO2or N2O, and forms nitric acid (Kolotoff, 1893) under the action of oxidising agents. Hydroxylamine is obtained in a great number of cases, for instance by the action of tin on dilute nitric acid, and also by the action of zinc on ethyl nitrate and dilute hydrochloric acid, &c. The relation between hydroxylamine, NH2(OH), and nitrous acid, NO(OH), which is so clear in the sense of the law of substitutions, becomes a reality in those cases when reducing agents act on salts of nitrous acid. Thus Raschig (1888) proposed the following method for the preparation of the hydroxylamine sulphate. A mixture of strong solutions of potassium nitrite, KNO2, and hydroxide, KHO, in molecular proportions, is prepared and cooled. An excess of sulphurous anhydride is then passed into the mixture, and the solution boiled for a long time. A mixture of the sulphates of potassium and hydroxylamine is thus obtained: KNO2+ KHO + 2SO2+ 2H2O = NH2(OH),H2SO4+ K2SO4. The salts may be separated from each other by crystallisation.
[25]The compound of hydroxylamine with hydrochloric acid has the composition NH2(OH)HCl = NH4ClO—that is, it is as it were oxidised sal-ammoniac. It was prepared by Lossen in 1865 by the action of tin and hydrochloric acid in the presence of water on a substance called ethyl nitrate, in which case the hydrogen liberated from the hydrochloric acid by the tin acts upon the elements of nitric acid—
Thus in this case the nitric acid is deoxidised, not directly into nitrogen, but into hydroxylamine. Hydroxylamine is also formed by passing nitric oxide, NO, into a mixture of tin and hydrochloric acid—that is, by the action of the hydrogen evolved on the nitric oxide, NO + 3H + HCl = NH4OCl—and in many other cases. According to Lossen's method, a mixture of 30 parts of ethyl nitrate, 120 parts of tin, and 40 parts of a solution of hydrochloric acid of sp. gr. 1·06 are taken. After a certain time the reaction commences spontaneously. When the reaction has ceased the tin is separated by means of hydrogen sulphide, the solution is evaporated, and a large amount of sal-ammoniac is thus obtained (owing to the further action of hydrogen on the hydroxylamine compound, the hydrogen taking up oxygen from it and forming water); a solution ultimately remains containing the hydroxylamine salt; this salt is dissolved in anhydrous alcohol and purified by the addition of platinum chloride, which precipitates any ammonium salt still remaining in the solution. After concentrating the alcoholic solution the hydroxylamine hydrochloride separates in crystals. This substance melts at about 150°, and in so doing decomposes into nitrogen, hydrogen chloride, water, and sal-ammoniac. A sulphuric acid compound of hydroxylamine may be obtained by mixing a solution of the above salt with sulphuric acid. The sulphate is also soluble in water like the hydrochloride; this shows that hydroxylamine, like ammonia itself, forms a series of salts in which one acid may be substituted for another. It might he expected that by mixing a strong solution of a hydroxylamine salt with a solution of a caustic alkali hydroxylamine itself would be liberated, just as an ammonia salt under these circumstances evolves ammonia; but the liberated hydroxylamine is immediately decomposed with the formation of nitrogen and ammonia (and probably nitrous oxide), 3NH3O = NH3+ 3H2O + N2. Dilute solutions give the same reaction, although very slowly, but by decomposing a solution of the sulphate with barium hydroxide a certain amount of hydroxylamine is obtained in solution (it is partly decomposed). Hydroxylamine in aqueous solution, like ammonia, precipitates basic hydrates, and it deoxidises the oxides of copper, silver, and other metals. Free hydroxylamine was obtained by Lobry de Bruyn (1891). It is a solid, colourless, crystalline substance, without odour, which does not melt below 27°. It has the property of dissolving metallic salts; for instance, sodium chloride. Hydroxylamine, when rapidly heated with platinum, decomposes with a flash and the formation of a yellow flame. It is almost insoluble in ordinary solvents like chloroform, benzine, acetic ether, and carbon bisulphide. Its aqueous solutions are tolerably stable, contain up to 60 per cent. (sp. gr. 1·15 at 20°), and may be kept for many weeks without undergoing any change. Lobry de Bruyn used the hydrochloric salt to prepare pure hydroxylamine. The salt was first treated with sodium methylate (CH3NaO), and then methyl alcohol was added to the mixture. The precipitated sodium chloride was separated from the solution by filtration. (The methyl alcohol is added to prevent the precipitated chloride of sodium from coating the insoluble hydrochloric salt of hydroxylamine.) The methyl alcohol was driven off under a pressure 150–200 mm., and after extracting a further portion of methyl alcohol by ether and several fractional distillations, a solution was obtained containing 70 per cent. of free hydroxylamine, 8 per cent. water, 9·9 per cent. chloride of sodium, and 12·1 per cent. of the hydrochloric salt of hydroxylamine. Pure free hydroxylamine, NH3O, is obtained by distilling under a pressure of 60 mm.; it then boils at 70°, and solidifies in a condenser cooled to 0° in the form of long needles. It melts at 33°, boils at 58° under a pressure of 22 mm., and has a sp. gr. of about 1·235 (Brühl). Under the action of NaHO it gives NH3and NHO2or N2O, and forms nitric acid (Kolotoff, 1893) under the action of oxidising agents. Hydroxylamine is obtained in a great number of cases, for instance by the action of tin on dilute nitric acid, and also by the action of zinc on ethyl nitrate and dilute hydrochloric acid, &c. The relation between hydroxylamine, NH2(OH), and nitrous acid, NO(OH), which is so clear in the sense of the law of substitutions, becomes a reality in those cases when reducing agents act on salts of nitrous acid. Thus Raschig (1888) proposed the following method for the preparation of the hydroxylamine sulphate. A mixture of strong solutions of potassium nitrite, KNO2, and hydroxide, KHO, in molecular proportions, is prepared and cooled. An excess of sulphurous anhydride is then passed into the mixture, and the solution boiled for a long time. A mixture of the sulphates of potassium and hydroxylamine is thus obtained: KNO2+ KHO + 2SO2+ 2H2O = NH2(OH),H2SO4+ K2SO4. The salts may be separated from each other by crystallisation.
[25 bis]In order to illustrate the application of the law of substitution to a given case, and to show the connection between ammonia and the oxides of nitrogen, let us consider the possible products of an oxygen and hydroxyl substitution in caustic ammonia, NH4(OH). It is evident that the substitution of H by OH can give: (1) NH3(OH)2; (2) NH2(OH)3; (3) NH(OH)4; and (4) N(OH)5. They should all, like caustic ammonia itself, easily part with water and form products (hydroxylic) of the oxidation of ammonia. The first of them is the hydrate of hydroxylamine, NH2(OH) + H2O; the second, NH(OH)2+ H2O (and also the substance NH(OH)4or NH3O2), containing, as it does, both hydrogen and oxygen, is able to part with all its hydrogen in the form of water (which could not be done by the first product, since it contained too little oxygen), forming, as the ultimate product, 2NH2(OH)3- 5H2O = N2O—that is, it corresponds with nitrous oxide, or the lower degree of the oxidation of nitrogen. So, also, nitrous anhydride corresponds with the third of the above products, 2NH(OH)4- 5H2O = N2O3, and nitric anhydride with the fourth, 2N(OH)5- 5H2O = N2O5. As, in these three equations, two molecules of the substitution products (-5H2O) are taken, it is also possible to combine two different products in one equation. For instance, the third and fourth products: NH(OH)4+ N(OH)5- 5H2O corresponds to N2O4or 2NO2, that is, to peroxide of nitrogen. Thus all the five (see later) oxides of nitrogen, N2O, NO, N2O3, NO2, and N2O5, may be deduced from ammonia. The above may be expressed in a general form by the equation (it should be remarked that the composition of all the substitution products of caustic ammonia may be expressed by NH3O5 -a, whereavaries between 0 and 4):NH5O5-a+ NH5O5-b- 5H2O = N2O5-(a+b),wherea+bcan evidently be not greater than 5; whena+b= 5 we have N2—nitrogen, when = 4 we have N2O nitrous oxide; whena+b= 3 we have N2O2or NO—nitric oxide, and so on to N2O5, whena+b= 0. Besides which it is evident that intermediate products may correspond with (and hence also break up into) different starting points; for instance, N2O is obtained whena+b= 2, and this may occur either whena= 0 (nitric acid), andb= 2 (hydroxylamine), or whena=b= 1 (the third of the above substitution products).
[25 bis]In order to illustrate the application of the law of substitution to a given case, and to show the connection between ammonia and the oxides of nitrogen, let us consider the possible products of an oxygen and hydroxyl substitution in caustic ammonia, NH4(OH). It is evident that the substitution of H by OH can give: (1) NH3(OH)2; (2) NH2(OH)3; (3) NH(OH)4; and (4) N(OH)5. They should all, like caustic ammonia itself, easily part with water and form products (hydroxylic) of the oxidation of ammonia. The first of them is the hydrate of hydroxylamine, NH2(OH) + H2O; the second, NH(OH)2+ H2O (and also the substance NH(OH)4or NH3O2), containing, as it does, both hydrogen and oxygen, is able to part with all its hydrogen in the form of water (which could not be done by the first product, since it contained too little oxygen), forming, as the ultimate product, 2NH2(OH)3- 5H2O = N2O—that is, it corresponds with nitrous oxide, or the lower degree of the oxidation of nitrogen. So, also, nitrous anhydride corresponds with the third of the above products, 2NH(OH)4- 5H2O = N2O3, and nitric anhydride with the fourth, 2N(OH)5- 5H2O = N2O5. As, in these three equations, two molecules of the substitution products (-5H2O) are taken, it is also possible to combine two different products in one equation. For instance, the third and fourth products: NH(OH)4+ N(OH)5- 5H2O corresponds to N2O4or 2NO2, that is, to peroxide of nitrogen. Thus all the five (see later) oxides of nitrogen, N2O, NO, N2O3, NO2, and N2O5, may be deduced from ammonia. The above may be expressed in a general form by the equation (it should be remarked that the composition of all the substitution products of caustic ammonia may be expressed by NH3O5 -a, whereavaries between 0 and 4):
NH5O5-a+ NH5O5-b- 5H2O = N2O5-(a+b),
wherea+bcan evidently be not greater than 5; whena+b= 5 we have N2—nitrogen, when = 4 we have N2O nitrous oxide; whena+b= 3 we have N2O2or NO—nitric oxide, and so on to N2O5, whena+b= 0. Besides which it is evident that intermediate products may correspond with (and hence also break up into) different starting points; for instance, N2O is obtained whena+b= 2, and this may occur either whena= 0 (nitric acid), andb= 2 (hydroxylamine), or whena=b= 1 (the third of the above substitution products).
[26]Nitric acid corresponds with the anhydride N2O5, which will afterwards be described, but which must be regarded as the highest saline oxide of nitrogen, just as Na2O (and the hydroxide NaHO) in the case of sodium, although sodium forms a peroxide possessing the property of parting with its oxygen with the same ease as hydrogen peroxide, if not on heating, at all events in reactions—for instance, with acids. So also nitric acid has its corresponding peroxide, which may be called pernitric acid. Its composition is not well known—probably NHO4—so that its corresponding anhydride would be N2O7. It is formed by the action of a silent discharge on a mixture of nitrogen and oxygen, so that a portion of its oxygen is in a state similar to that in ozone. The instability of this substance (obtained by Hautefeuille, Chappuis, and Berthelot), which easily splits up with the formation of nitric peroxide, and its resemblance to persulphuric acid, which we shall afterwards describe, will permit our passing over the consideration of the little that is further known concerning it.
[26]Nitric acid corresponds with the anhydride N2O5, which will afterwards be described, but which must be regarded as the highest saline oxide of nitrogen, just as Na2O (and the hydroxide NaHO) in the case of sodium, although sodium forms a peroxide possessing the property of parting with its oxygen with the same ease as hydrogen peroxide, if not on heating, at all events in reactions—for instance, with acids. So also nitric acid has its corresponding peroxide, which may be called pernitric acid. Its composition is not well known—probably NHO4—so that its corresponding anhydride would be N2O7. It is formed by the action of a silent discharge on a mixture of nitrogen and oxygen, so that a portion of its oxygen is in a state similar to that in ozone. The instability of this substance (obtained by Hautefeuille, Chappuis, and Berthelot), which easily splits up with the formation of nitric peroxide, and its resemblance to persulphuric acid, which we shall afterwards describe, will permit our passing over the consideration of the little that is further known concerning it.