ACTION O F METALS ON NITRIC ACID BY WILDER D. BANCROFT
For nearly thirty years I have carried the reduction of nitric acid by metals in my notebooks as a major problem to be taken up as soon as possible, and a t differenttimes during that period I have suggested it to graduate students as a desirable thesis subject-all in vain. Three years ago I found in Mr. L. H. Milligan a student who was both competent to tackle the problem and interested in it. Rather to my dismay it turned out that the methods of analysis were not satisfactory and consequently Mr. Milligan’s thesis, which will be published soon, had to be on analytical methods and not on the real problem at all. Fortunately he was able to do enough additional work to make it seem certain that the underlying theory was sound and that the problem now is primarily one of collecting data. Since it may be many years before we get the whole question cleaned up and since we would like very much to get other laboratories interested in the subject, it has seemed worth while to present a general outline of the situation as we now see it and to follow this up with the experimental papers as fast as we can get them finished. Some day it is to be hoped that the general theory of the reduction of nitric acid will be worked out as clearly and definitely as is the theory of the reduction of nitrobenzene. There are two quite distinct problems in the reduction of nitric acid: the question of the intermediate stages; and the part played by any given metal. The solving of the second problem carries with it the solving of the first. Acworth and Armstrong’ believe that the direct reduction products are nitrous acid, hyponitrous acid, hydroxylamine, and ammonia, with nitrogen peroxide, nitric oxide, nitrous oxide, and nitrogen as secondary products. Ostwald2 believes that the reduction of nitric acid starts from a dissociation into N 0 2 ‘ + 0 H ’ and that consequently the direct reduction products are nitrogen peroxide, nitrous acid, nitric oxide, hyponitrous acid, and either nitrogen or hydroxylamine and ammonia; with nitrous oxide presumably a dehydration product of hyponitrous acid. The same view is taken by the unknown author (presumably Partington) in Thorpe’s Dictionary of Applied Chemistry, 4, 56 I (1922).
Mellor3 is distinctly non-committal; but rather inclines to Ostwald’s view. “Some consider that the first product of the reaction is a nitrate of the metal and nascent hydrogen: Cu+zHNO3= Cu(N03)2+2H. The nascent hydrogen is then supposed to reduce the nitric acid to nitrous acid. With some metals, the reduction of the nitric acid proceeds much further, say through the stages : NO~+HNO~+NO+N20+-+N2--tNH3. Free hydrogen is seldom evolved because it is so rapidly oxidized by the nitric acid. However,. I
J. Chem. Soc. 32, 56 (1877). Grundriss allg. Chemie, 440 (1899). “Modern Inorganic Chemistry”, 514 (1912).
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WILDER D. BANCROFT
free hydrogen is said to have been obtained by the action of nitric acid on manganese and on magnesium. The reducing actions indicated above can be represented’ symbolically : Z H N O ~ + Z H = Z H ~ O + Z N O Z zHN03+811 = 5HzO+Nz0 Z H N O ~ + ~ H = Z H Z HNOz O+~ 2HN03+ 1oH=6 HzO+Nz zHN03+6H = 4HzO+zN0 2”03+ 16H= OH20+zNH3 Some believe that the acid first oxidizes the metal2 to the oxide, e.g. 3 Cu+ 2HN03=3CuO+HzO+zNO; and that the oxide then dissolves in the acid to form the nitrate: CuO+zHN03=Cu(N03)z+HzO. W h e n digerences of opinion can reasonab1.y be entertained, it 5 0 1 1 0 ~that ~ our knowledge of the facts is incomplete, and that more experimental work i s needed. The metals copper, silver, mercury, and bismuth have no perceptible action on cold dilute sulphuric and hydrochloric acids, and accordingly it is not likely that they will reduce nitric acid by the action of nascent hydrogen. V. H. Veley (1890) proved that these metals have no action on cold dilute nitric acid unless a trace of nitrous acid or a lower nitrogen oxide is present. Nitrous acid may be present in the nitric acid as an impurity; it may be formed by the incipient decomposition of nitric acid when it is warmed; or it may be formed in the acid by electrolysis produced by local currents of electricity set up by impurities in the metal (‘local action’). Once the action has started, the evolution of nitric oxide, and the formation of nitrate proceeds quickly. According to Veley, therefore, the dissolution of this metal in nitric acid proceeds: Cu+ 3HN03+HN02+H20, is the resultant of a series of consecutive reactions: Cu+4HNO2 = Cu(N02)z+2Hz0+zNO; followed by Cu(N02)2+ zHN03 = C U ( N O ~ ) ~ + ~ H Nthe O ~small ; trace of nitrous acid thus acts as a catalytic agent; nitrous acid is continuously produced and continuously decomposed : 3 H N 0 2=HN03+2NO+H20.” Roscoe and Schorlemmer3 dodge the whole question. “AS already mentioned, nitric acid dissolves a large number of metals with formation of nitrates. Hydrogen is not evolved at the same time, as is the case with sulphuric and hydrochloric acids, but in its place lower oxides of nitrogen and even nitrogen itself and ammonia are formed. The explanation frequently given of this change is that hydrogen is first produced, but that it at once acts on the excess of nitric acid present, forming water and the lower oxides of nitrogen. Thus, for example, the formation of nitric oxide by the action of copper on nitric acid is supposed to take place in the two following stages:CU+ z H N O = ~ C U ( N O B ) P 2H + 6H 2HN03 = ZNO +4H20 “According to Veley,4 however, this explanation is not correct, inasmuch as pure copper, mercury, and bismuth do not dissolve in pure dilute nitric acid,
+
Cf. Montemartini: Gam. chim. ita]. 22, I, 339 (1892).
* This is the view held at one time by Higley: Am. Chem. J. 17, 26 (1895). It is still
taught at Harvard. “ A Treatise on Chemistry”, 1, 544 (1920). Proc Roy. SOC.46, 216 (1890); 52, 27 (1893); Phil Trans. 182 A, 312 (1891); J. SOC. Chem. Ind. 10, 204 (1891).
ACTION OF METALS ON N I T R I C ACID
477
but dissolve readily when nitrous acid is present the change at any moment being directly proportional to the mass of the nitrous acid in the solution, and the more rapid the greater the proportion of the former to the latter. He therefore believes that the reaction is started either by traces of nitrous acid already present or by impurities in the metal inducing a local electric current; the first product of the reduction of the metal is nitrous acid, and the production of lower oxides of nitrogen he regards as due to the subsequent changes occurring between nitrous acid and the metallic nitrate or nitrite in presence of an excess of nitric acid, the nitrous acid being decomposed as fast as it is formed. “The relative proportion of the products formed varies not only with the concentration of the acid and the temperature, but also with the nature of the metal employed. Thus silver when treated with acids of sp. gr. 1.05-1.4 yields nitric oxide and nitrogen peroxide, but no nitrous oxide, whereas lead with an acid of sp. gr. I .n yields a gas of which forty percent is nitrous oxide.” Norris‘ states explicitly that the first action of nitric acid on metals is an oxidizing one. The square brackets indicate substances which are not actually set free as such. “Nitric acid oxidizes all the metals except the so-called noble metals, gold and platinum, for example. The nitrates formed are soluble in water and, consequently, nitric acid dissolves these metals. . . . In writing equations for the reactions involved, it is best to separate them into steps and combine the partial equations in the way already explained. For exaniple, the oxidation of copper by nitric acid can be represented as follows: 2”03= ~NO+Hz0+[30] 3 CU+[3OJ = 13CUOl [3CUOI 6HNO3 = 3Cu(N03)2 3 H z 0 ~CU+~HNO~=~CU(NO~)~+~H~O+ZNO.” Howe2 takes an intermediate viewpoint. He says that “the reaction between zinc and nitric acid may be looked upon as taking place in two stages: 3Zn+ 6HN03= 3Zn(N03)2 + 6H 6H+ 2 H N 0 3 = 4 HZO+ZNO, or, summing up these reactions: 3Zn+8 NH03=3Zn(NOs)~+4H20+z NO. It is perhaps better to look at this reaction from the standpoint of the oxidation of the metal, in which case we have the two equations: 3Zn+zHNO3 = 3ZnO+HzO+z NO 3ZnO 6HN03= jZn(N03)2 + 3 H 2 0 which, when summed up, become as above: 3Zn+8HN03 = 3Zn(N03)2+4H20+ nNO.” This is playing safe withTa vengeance. Unfortunately, as wc shall see later, neither set of equationszrepresents the facts.
+
+
+
“A Textbook of Inorganic Chemistry for Colleges,” 327 (1921). “Inorganic Chemistry for Schools and Colleges,” 270 (1920).
4i8
WILDEH. D. BANCROFT
A great deal of work has been done on the reduction of nitric acid by metals with very little in the way of theoretical results. When this sort of thing happens over a period of years, it is usually safe to assume that the problem has been attacked from the wrong point of view, On looking over the literature it is easy enough to see where people have gone wrong. There has been the tacit assumption that the action of the metal depends on a single factor, whereas I shall show that we must consider not less than three factors, which do not necessarily run parallel as we change from one metal to another. Acworth and Armstrong’ say that “we regard the action of the metal as consisting simply in the displacement of the hydrogen of the acid and the formation of the corresponding nitrate in the manner expressed by the following equation, in which R denotes the amount of metal which is equivalent in combining or displacing power to two atoms of hydrogen: R 2HNO 3 = R (Nos)2+ 2H. Under no circumstances, however, is the hydrogen thus displaced evolved as such; it at once acCs on the free acid present (and even, in some cases, as we shall show latJer on, also on the metallic nitrate which has been formed), reducing it more or less completely to nitrous acid, nitrosidic acid [hyponitrous acid], hydroxylamine, or ammonia.” This is, of course, a special statement of the electrolytic theory of corrosion, long before it was developed as a general theory by Whitney2in 1903. A few years later I pointed out3 that “the most striking characteristic of an electrolytic action is that it occurs in two places-at the anode and at the cathode. This peculiarity can be made less marked by bringing the electrodes nearer and nearer together. When the distance between them vanishes, we have a chemical reaction in the ordinary sense of the word and not an electrochemical reaction. Any chemical reaction, therefore, which can be made to take place electrolytically, must consist of an anode and a cathode p r o c e ~ s . ~ Considering the matter in this light, we see that there is a possibility of the anode and cathode processes interfering and of one perhaps masking the other. “In some cases it is easy enough to tell what the anode and the cathode processes are. If we dissolve zinc in sulphuric acid, the formation of zinc SUIphate is the anode process and the evolution of hydrogen is the cathode process. Now we know that pure zinc does not dissolve readily in sulphuric acid. Consequently we should expect to find a difficulty of some sort if we electrolyze sulphuric acid between the electrodes. We find this in the form of the SQcalled ‘excess voltage’ at the cathode; and in the electrolytic process we can obtain a mare or less quantitative measurement of the phenomenon though we are still far from knowing the cause of it. “A less simple case is that of copper in dilute nitric acid. Copper reacts chemically with dilute nitric acid, setting free nitric oxide. The formation of
+
?
J. Chem. SOC.32, 56 (1877). J. Am. Chem. SOC.25,394 (1903). Trans. Am. Electrochem. SOC.9, 13 (1906). Cf. Trauhe: Ber. 26, 1473 (1893); Haber: Z. physik. Chem. 34, 514 (1900).
ACTION O F METALS ON NITRIC ACID
479
copper nitrate must be the anode process and the reduction of the nitric acid the cathode process. When we start to test this we find difficulties. Everybody knows that we get ammonia instead of nitric oxide if we electrolyze dilute nitric acid, using a copper cathode. We have here an apparent contradiction, the chemical reduction yielding nitric oxide and the electrochemical one ammonia. Mr. Turrentine was good enough to solve the mystery for me. When copper reacts chemically with nitric acid, the anode product, copper nitrate, is formed at the same spot that the reduction takes place. In the electrolytic reduction of nitric acid with a copper cathode, the reduction takes place in a solution practically free from copper salt. The conditions are therefore not the same in the two cases. Mr. Turrentine therefore electrolyzed a solution of nitric acid and copper nitrate using a copper cathode. A gas was evolved at the cathode which proved on analysis to be chiefly nitric oxide. This experiment can be done in another form which is more striking. If dilute nitric acid be electrolyzed between copper electrodes, there will at first be no evolution of gas at the cathode. Gas will begin to appear as soon as the blue solution formed at the anode comes in contact with the cathode. A corollary to this is that ammonia would be formed in the chemical reaction between copper and nitric acid if the concentration of the copper salt could be kept sufficiently low. There did not semi to br any salt which one could add to the solution without introducing more complications than were eliminated. The difficulty was overcome by Mr. Turrentine in a distinctly ingenious manner. Strips of copper were hung vertically in a tall vessel. The copper nitrate flowed to the bottom of the vessel and the copper was removed by electrolytic precipitation in the form of cupric hydroxide. No current flowed through the copper strips and there was no copper cathode; but ammonia was formed. “These experiments were performed to prove that the difference between the electrochemical and the chemical corrosion of copper by nitric acid was an apparent one only and due to an unsuspected difference in the conditions. In addition they illustrate the superior flexibility of the electrochemical method over the chemical method. In the electrochemical method there is no difficulty in varying the concentration of the copper salt at the cathode between any desired limits, whi’le this is very difficult to do in the case of the chemical method. This is in addition to the advantage, which the electrochemical method always has, of permitting a wide variation in the rate of reaction for constant temperature and constant concentration. If we are ever to have a thorough knowledge of the chemical reactions between nitric acid and the metals we must study the problem electrochemically.” On the assumption that all corrosion is electrolytic in nature, it is evident that at least three independent factors must be taken into account in the reduction of nitric acid by a metal: the specific reducing power of the metal which may be measured by the hydrogen over-voltage in a corresponding sulphuric acid solution; the catalytic action of the metal on the various reduction products; and the catalytic action of the metallic nitrate on the various reduction products. This analysis of the problem is of no value unless
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WILDER D. BANCROFl
we can devise methods of distinguishing the effects due to these three causes. Fortunately that is a very simple matter. We can determine the reduction products on adding different, dissolved, reducing agents having no catalytic action, which will give us the reduction products due to different electromotive forces. Of course it is possible that any given reducing agent may exert a catalytic action itself or through one of its oxidation products, and a t least one case of this sort has already been found by Mr. Milligan. The catalytic action can be recognized and eliminated by using other reducing agents of about the same reducing power or by adding the same reducing agent to solutions of the different reduction products of nitric acid. This last method will be effective only in case the catalytic action is on one of the first reduction products of nitric acid, such as nitrous acid. There is no reason a t present t o suppose that there will be any special difficulty in distinguishing between reducing action and catalytic action. Electrolysis of a nitric acid solution with the metal under consideration as cathode will give the combined effect of the specific reducing action of the cathode plus the possible catalytic action of the cathode metal. If the electrolytic reduction agrees with those obtained by the corresponding chemical reduction, the cathode metal has no catalytic action. If there is a difference between the sets of results, the difference gives us valuable information in regard to the catalytic action of the cathode metal. It will probably not be difficult then to formulate the catalytic action of the metal clearly and definitely. The catalytic action of the metallic nitrate can be detected by adding the salt to nitric acid and then rcpeatinp the experiments with the dissolved reducing agents. These results will be supplemented by repeating the electrolytic reductions in presence of the metallic nitrate, after which there should be no difficulty in accounting for all the products obtained when the metal reacts direct with the nitric acid under the conditions of ordinary corrosion. In some experiments, on which Mr. Milligan will report in detail later, it was found that when acidified ferrous sulphate is added to an excess of nitric acid, the main reduction product is nitric oxide. Since acidified ferrous sulphate is a weak reducing agent, this is about what one would haire predicted. With a stronger reducing agent, acidified stannous chloride, the reaction runs practically completely (over ninety percent) to hydroxylamine if the temperature is kept low. If the temperature is allowed to rise, the reduction product is practically completely nitrous oxide. This has nothing t o do with the stannous chloride because the same result is obtained if we heat the hydroxylamine solution and nitric acid solution. The nitrous oxide is formed by the reaction between nitrous acid and hydroxylamine HNOZ+NH,OH =NzO+HzO. Since there is not much nitrous acid in the solution at any one time, we may have to postulate a preliminary reduction of nitric acid by hydroxylamine, Z H N O ~ + Z N H ~= O ~HNOz+Nz0+3H20. H
ACTION O F METALS ON N I T R I C ACID
481
Acidified titanous chloride is a still more powerful reducing agent and will reduce hydroxylamine quantitatively to ammonia. It was therefore clear that it should reduce nitric acid t o ammonia; but Mr. Milligan found that the main reduction product was nitric oxide, the same compound which is the chief product obtained with ferrous sulphate. I n other words, titanous chloride does not reduce nitric acid so far as does the weaker reducing agent, stannous chloride. One can see what havoc this would play with any theory which made the reducing action depend solely upon the chemical potential. The abnormal result with tit,anons chloride is not due to the catalytic action of the final oxidation product, because hydroxylamine is obtained on adding stmannouschloride to a nitric acid soiution, which has h e n reduced in part by titanous chloride. The next step-which has not yet been taken-is to add titanous chloride to nitrous acid, t o hyponitrous acid, and to hydroxylamine,~ as~to determine what side reaction is catalyzed. In this particular case, it is probably the decomposition of nitrous acid which is catalyzed, though it may be that the reduction proceeds to nitrohydroxylaminel which then loses water, H2N~03= 2NO+H20. The result with titanous chloride make it possible that the data with ferrous sulphate are due in part to catalytic action. This can be checked by finding whether ferrous sulphate does or does not reduce a hyponitrous acid solution. It is evident from what has already been done that Armstrong was right in saying that the direct reduction products are H N 0 2 ,H2N202,NH20H, and "3, while NO2, NO, NzO and Nz are formed by secondary reactions, very probably as follows:H N 0 3 + HNOz = 2NO2fHzO 3HNO2= 2NO+HNO3+Hz0 HzNz03= 2NO HzO HzNzOz = NzO +HzO HNOzfHNzOH =NzO+2H20 H2Nz02+2NHzOH= 2N2+4HzO HNOz NH3 = Nz 2 1 1 2 0 . Thc reaction between nitric and nitrous acids to form nitrogen peroxide will increase relatively with increasing concentration of nitric, acid. It is simply the reverse of one of the reactions in the synthesis of nitric acid by the arc process. It is sometimes stated that the formation of nitrogen peroxide is due to the oxidation of nitric oxide by concentrated nitric acid; but this is clearEy a mistaken way of looking a t things. There is no reason to postulate an intermediate formation of nitric oxide. We get nitric oxide by the decomposition of nitrous acid and this reaction will predominate as the concentration of nitric acid decreases and that of nitrous acid increases. This equilibrium was studied carefully by Noyes and Edgar2 over a dozen years ago.
+
+
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Angeli: Gaze. chim. ital. 26 11, 17 (1896).
* J. Am. Chem. SOC.33, 292 (1911).
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WILDER D. BANCROFT
We get nitrous oxide either through the dehydration of hyponitrous acid or through the interaction of nitrous acid and hydroxy1amine.l If this second reaction passes through the hyponitrous acid stage as seems probable, these are merely different ways of saying the same thing. It is usually assumed that nitrogen comes froni the interaction of nitrous acid and ammonia, apparently because ammonium nitrite will decompose into nitrogen and water under certain conditions. This seems to me inadequate because we get nitrogen evolved in cases where there is no reason to suppose that ammonia is formed at all. We must therefore postulate a reaction between hyponitrous acid and hydroxylamine. HzNzOz+z NHzOH = 2N2+4 HzO. Oesterheld2 was forced to the same conclusion when he found that the electrolytic oxidation of an alkaline hydroxylamine solution a t a platinum anode gave a mixture of gases approximating 5 1YoNzand 49% NzO. “The formation of nitrous oxide is undoubtedly due to the oxidation of hydroxylamine to hyponitrous acid which then breaks down into nitrous oxide and water: NHzOH+O =NOH+H20 2NOH+HzN202+NzO H2 0 . It is impossible, however, that the nitrogen formed from the hydroxylamine can be a decomposition product of ammonium nitrite as was assumed to be the case when studying the oxidation of ammonia. In the first place there is no ammonia present and, in the second place, there can be no oxidation of hydroxylamine to nitrous acid a t the low anode potential which was maintained, The simplest explanation is that the nitrogen is due to the reaction between hydroxylamine and nitroxyl which is the first oxidation product, O:NH+NH20H=Nz+2H2 0. A. Angeli has observed the same reaction when nitroxyl splits off from benzene sulphohydroxamic acid in presence of hydroxylamine. While two molecules of nitroxyl may combine to form hyponitrous acid or nitrous oxide and water, yet there is a quantitative evolution of nitrogen so long as there is a sufficient excess of hydroxylamine.” It is also possible that nitrogen may be formed simultaneously with nitrous oxide by a different decomposition of hyponitrous acidJ3perhaps sHzNzOz= 4HzO+HN03+4Nz. While it is perfectly certain that nitrogen peroxide, nitric oxide, nitrous oxide, and nitrogen are due to secondary reactions and are not direct reduction products of nitric acid, this does not mean that hydroxylamine and ammonia may not be formed by reduction of the oxides of nitrogen. In fact, we know that under certain conditions this may happen. The simplest supposition to make would be that the stages were NOZ+NO-+NZO(+N~)+NH~OH +NH3; but there seems to be good reason to believe that matters are not
+
V. Meyer: Ann. 175, 141 (1875). 2. anorg. Chem. 86, 129 (1914). SBerthelot and Ogier: Compt. rend. 96, 30 84 (1883); Hantnsch and Kaufman : Ann. 232, 317 (1896); Ray and Ganguli: J. Chem. SOC.91, 1866 (1907). 1 2
ACTION O F METALS ON N I T R I C ACID
’
483
so simple as this. It seems probable that hydroxylamine is not always a preliminary stage to the formation of ammonia, that nitrous oxide is not always a preliminary stage in the formation of hydroxylamine or ammonia from nitrogen peroxide or nitric oxide, and that the reduction of nitrogen peroxide does not necessarily give rise to nitric oxide as an intermediate stage. Divers and Shimidzu‘ have found that “zinc dissolves in dilute nitric acid, forming nitric oxide, nitrous oxide, ammonia and zinc nitrates, and sometimes a trace of hydroxyammonium [hydroxylamine] nitrate. Should it dissolve so far as to neutralise or nearly neutralise the acid, zinc nitrite is also formed. By mixing much sulphuric acid with the nitric acid, the dissolution of the zinc produces sulphates instead of nitrates and nitrites, hydroxyammonium and ammonium sulphates, nitrous oxide and a little nitrogen, little if any nitric oxide, and, lastly, hydrogen.” I n this paper I shall not take up the question why sulphuric acid increases the yield of hydroxylamine, but rather why it does not appear when nitric acid alone is used. Since Divers did not check his temperatures, it is possible that the temperature rose in the experiments in which there was no sulphuric acid and that the hydroxylamine was destroyed by reaction with nitrous or nitric acid. This is not very probable and would lead to the principal product being nitrous oxide and not ammonia. Also, in presence of sulphuric acid the yield of hydroxylamine increases with rising temperature.2 When once formed, hydroxylamine is not reduced readily to ammonia (p. 613). “In one experiment, I O cc of a solution of hydroxyammonium chloride containing 0.033 gram of hydroxyamine, were made up to IOO cc with water, containing a little sulphuric acid. The mixture was poured upon 45 grams of granulated zinc, and thus exposed to a relatively very large surface of zinc. It was left in contact with the zinc for two hours, dilute sulphuric acid being occasionally added, so as to keep up effervescence. The solution, still effervescing, was poured off, and titrated for hydroxyamine, and the whole of this was found unchanged. Indeed, the anomaly was observed, of a little more iodine being required than equalled the whole hydroxyamine. But we find that in presence of a concentrated solution of zinc salt, a clear excess of iodine solution is decolorised over that calculated for the hydroxyamine. “We then tried the action of zinc alone upon a dilute solution of hydroxyammonium chloride. Here there was a marked destruction of the hydroxyamine in two hours, one-eighth to one-fourth disappearing. The nitrate appears to be somewhat more unstable. Acidified with a little sulphuric acid, however, the nitrate gave in ten minutes on zinc, a third more hydroxyamine than it contained in consequence of the conversion of part of its nitric acid to that base. (The nitrate consists by weight of one part base t o t w o parts acid, so that about a third of its acid became hydroxyamine.) The presence of zinc sulphate or ammonium chloride or of free ammonia, appears to be without marked effect. 2
J. Chem. SOC.47, 597 (1885). Divers and Shimidau: J. Chem. SOC. 47, 619 (1885).
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WILDER D. BANCROFT
“Hydroxyamine is ’thus seen to be only very slowly decomposed by the contact of zinc with its cold solution, even in absence of acid. It was otherwise, however, in the case of hot solutions, for in these the hydroxyamine is destroyed by zinc with relative rapidity, although still wonderfully resistant to decomposition. Boiled for a quarter of an hour in I O O cc, and less, of solution, upon 35 grams of zinc, 0.033 gram was reduced to 0.02 gram. Heated to 7oo-8o0 for one hour upon 35 grams of zinc, 0.033 gram hydroxyamine in IOO cc was reduced to 0.008 gram.” Divers and Shimidzul state specifically that ammonia and hydroxylamine appear to be not consecutive products but products of parallel and independent reductions. “As pointed out in the previous paper ammonia is abundant from the first, during the reduction of nitric acid by sulphuric acid and zinc. Its non-production from hydroxyamine, in the presence, that is, of free acid, gives its appearance during this reduction a new interest. Nitrous oxide is also produced. We have made no attempt to estimate the very small quantities of this gas formed in our experiments. Deville’s determinations may be referred to on this point, but we ourselves in Japan have no means of doing so, for purposes of illustration. “A few determinations only of ammonia have been made by us, and further work in this direction is desirable. Having poured off the acid solution from the zinc at those times when it was calculated that the hydroxyamine would have just reached its maximum, one part was titrated for hydroxyamine, and another titrated for ammonia, by adding first sodium carbonate until some zinc carbonate was precipitated, then iodine to des troy hydroxyamine, then precipitating nearly all the zinc by sodium sulphide, and distilling the motherliquor with potassium hydroxide. The ammonia was received in acid and titrated. I n this way we obtained ammonia equivalent to much of the nitric acid not convertible into hydroxyamine. (‘By using a very great excess of sulphuric acid nearly, if not all, the nitric acid unconverted to hydroxyamine appears to be converted into ammonia. This is in accordance with Kinnear’s results,2 and we believe it to be correct,. but there is difficulty in determining closely a very small quantity of ammonia in presence of very much zinc salt, and, without more trials than we have made, we cannot assert that it is so. “Under favourable conditions, we have got as much as, and even more than seven-tenths of the nitric acid converted to hydroxyamine, but, with unsuitable zinc, we have sometimes failed to get more than two-tenths converted. In our former paper, we mentioned a yield of eight-tenths of the nitric acid as hydroxyamine, but we were not confident as to this yield having been really obtained. “In a preceding section, we have attributed the production of hydroxyamine entirely to the specific action of the zinc and the nitric acid upon SUIphuric acid. We can now almost assert that ammonia is not a product of that 1
J. Chem. SOC.47, 615 (1885). Kinnear: Chem. News, 46, 63 (1882).
ACTION OF METALS ON NITRIC ACID
485
specific action, and that the ammonia freely produced from the first during the dissolution of the zinc in the mixed acids, is formed solely by the direct action of the nitric acid upon the zinc. Uncertainty thus seems cleared away a s t o the specific hydrogenising action of the sulphuric acid upon the nitric acid. For, although nitric acid in contact with zinc and water can readily hydrogenise itself without the aid of any other acid, the product is then ammonia, not hydroxyamine, whereas when nitric acid is hydrogenised by sulphuric acid, the product is hydroxyamine and not ammonia. “In the paper on the production of hydroxyamine from nitric acid, we have said that a perfect metal for making hydroxyamine would be one active towards dilute nitric acid, inactive towards sulphuric or hydrochloric acid, and inactive towards hydroxyamine. With our present knowledge we would alter this, and say, that such a metal should be inactive to nitric acid alone, and active only to a mixture of nitric acid with either hydrochloric or sulphuric acid. The above considerations at once explain why it is that a metal may freely produce ammonia from nitric acid and yet produce hydroxyamine from it only with difficulty, or not a t all. For, in such cases, there may be strong action of the metal upon nitric acid, or independently upon both acids, whereby ammonia will be formed, and yet be no specific action upon the two acids combined, which would give rise to hydroxyamine. Or again, the effect of t.his specific action may be as nothing compared with that of the superior activity of the metal towards the independent acids. “Granulated zinc may differ widely in its power of yielding hydroxyamine, but a given preparation of it, or mixtures of similar preparations, will show under the same circumstances the same power at one time that it has shown at another. The differences observed in this relation between different preparations of granulated zinc, depend in part upon the nature of the zinc melted down, of course, but they also depend upon differences in the process of granulating. Partly, no doubt, in consequence of the fact that apparently both the condition of the zinc before it has been granulated, and the character of the operation of granulating it, may exert marked influence upon its power to produce hydroxyamine from nitric acid, but at any rate we have failed in the attempts we have made to secure at will granulated zinc of a uniform and high power in this respect, or to ascertain any sufficient knowledge of the causes of its varying power. “On one occasion, there seemed to us to be a marked difference between the same zinc when granulated a t diffreent temperatures. By running the zinc when almost boiling hot, and then again, from the same melting, after the zinc had cooled almost to the solidifying point, a much more productive lot of granules was obtained from the second running than from the first. But on trying the effect of this difference in operating again and again, we got no confirmation of our first result. We also varied the process by pouring from a height, and pouring near the water, and of granulating in boiling water and granulating in cold water, but from all our trials we learnt nothing. At one spell of work we could repeatedly granulate zinc turning out to be of
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WILDER D. BANCROFT
nearly equal power, but on another occasion we could not count on producing more zinc of the same power.” There is independent confirmation of the view that the formation of ammonia does not always pass through the hydroxylamine stage. Tafell says that “nitric acid in sulphuric solution is reduced electrolytically at, an amalgamated cathode almost exclusively to hydroxylaniine, although such a cathode has a distinct electrolytic reducing action on hydroxylamine sulphate. On the other hand nitric acid is reduced almost completely to ammonia at a copper cathode which is effective only at low current densities and which is quite inadequate to reduce hydroxylamine sulphate any further. “From these facts the only safe conclusion to draw is that the electrolytic reduction of nitric acid to ammonia a t a copper cathode does not pass through the hydroxylamine stage. . . . I assume that certain cathode materials, such as copper for instance, by a chemical action divert the reduction process somewhere between nitric acid and hydroxylamine from the specific electrolytic course, so that the reaction goes direct t o ammonia, passing round the hydroxylamine stage. I cannot tell certainly at what stage this switch takes place; but it seems not improbable that it occurs a t dihydroxylamine, ”(OH) 2.’’ There is another way in which one can get ammonia without passing through the hydroxylamine stage. Hantzsch and Kaufmann2 found that hyponitrous acid may break down either to give nitrous oxide or to give ammonia, HzNzOz +N20+HzO 3HzNz03 +2NH3+N203. Any substance, which catalyzed the second reaction without catalyzing either the first one or the reduction to hydroxylamine, would give ammonia. We do not know yet what actually happens; but it seems certain that hydroxylamine is not necessarily an intermediate stage in the formation of ammonia. Hantzsch3 suggests that the syn form of hyponitrous acid would be expected to split off water very easily; but he leaves it rather vague as to what the characteristics of the anti form would be. We find other difficulties with nitric oxide and nitrogen peroxide. Divers and Haga4 say that “every chemist is supposed to know that when nitric oxide is passed into a hydrochloric acid solution engaged in dissolving tin, ammonia and hydroxylamine are produced. Strictly speaking, however, this is by no means the case. We assert that, in the change that goes on, not a trace of ammonia is formed. Experience to the contrary must have been all vitiated by neglecting to exclude air while the gas was in contact with the mixture, and to use pure nitric oxide, Hydroxylamine is formed, but if the nitric oxide never comes in contact with oxygen, not a trace, as we have said, Z. anorg. Chem. 31, 289
(1902).
Ann. 293, 332 (1896). *Ann. 292, 340 (1896). J. Chem. SOC.47, 623 (1885).
ACTION OF METALS ON NITRIC ACID
487
of ammonia is formed. V. Dumreicher has already shown that acid stannous chloride will of itself convert nitric oxide to hydroxyamine; whether he then found ammonia or not we cannot learn. This will have depended upon the care with which he excluded air. But his observation was of great interest a t the time, as serving to show what little foundation there was for the hypothesis of the action of nascent hydrogen in explaining the formation of hydroxyamine from nitric oxide by treating it with tin and hydrochloric acid. So far, indeed, as we can ascertain by experiments which we have made, the presence of the metal does not directly help to form hydroxyamine, and is of no service but to generate stannous chloride. It is objectionable on two grounds; it leaves the solution charged with unused stannous chloride whenever the preparation of hydroxyamine is stopped; and, secondly, the hydrogen it produces carries off the nitric oxide that would otherwise lie on the surface of the solution and be absorbed. “Tin and hot concentrated hydrochloric acid do not sensibly destroy hydroxyammonium chloride. This is our experience, and it confirms that of Mr. Shimidzu and one of us, as to the want of action of zinc and sulphuric action upon the hydroxyamine. Further stannous chloride in acid solution has no action upon the hydroxyammonium chloride, even when boiled with it; that is, for a reasonable time; an aqueous solution could itself not be boiled very long without suffering some change. Were either tin and hydrochloric acid, or stannous chloride, capable of converting hydroxyamine into ammonia, what we assert concerning the non-production of ammonia from nitric oxide by tin and hydrochloric acid could not, of course, be true. According to V. Dumreicher, however, a t a temperature of IOO’, hydroxyamine is quickly and completely reduced to ammonia by excess of stannous chloride. “Coming now to the use of stannous chloride, our experience is, that with exclusion of air, not a trace of ammonia is formed by its action upon an acid solution of hydroxyamine. In working, we prefer to use a mixture of ferrous sulphate, nitric acid, and sulphuric acid as the source of the nitric oxide. It yields the gas in a pure state, and the generator can be charged and worked so as t,o furnish easily, except for a short time a t first, a gentle stream of the gas for 24 hours at a stretch. To destroy any nitric peroxide, formed through entrance of traces of air into the generator, the gas is sent through a washbottle of potassium hydroxide solution. This bottle is scarcely necessary, however, as reaction between ferrous sulphate and the acids does not begin in the cold; and before starting the disengagement, of nitric oxide, a current of hydrogen is sent through the whole series of flasks and bottles until all air has been driven out. Only then is the evolution of nitric oxide commenced and passed into the apparatus, the current of hydrogen being at the same time discontinued. The remote end of the arrangement is provided with a washbottle as a guard against any reflux of air. When the process is finished, or when it has to be interrupted, the current of nitric oxide is stopped and its place taken by one of hydrogen again, until all nitric oxide has been chased away. The flasks of tin chloride may then be opened with impunity, air hav-
488
WILDER D. BANCROFT
ing no effect upon the hydroxyamine already formed, not even when stannous chloride is still present, and being so, goes on changing into the stannic salt by contact with the air. “The reaction between acid stannous chloride and nitric oxide is not quite simple, however. The greater part of the nitric oxide becomes hydroxyammonium chloride, but a part of it is always converted into nitrogen. No nitrous oxide is formed. Since much of the nitric oxide passes unabsorbed, we could not well take the proportions of it to the hydroxyamine formed a t the same time. But we many times collected it in quantity, the nitric oxide we were using being tested and found pure, as from its source it was sure to he. . . . “There is a remarkable fact about the reaction between nitric oxide and stannous chloride, we have yet to mention. At a temperature of fully IOO’, there is no action between nitric oxide and acid stannous chloride solution. The stannous chloride remains unchanged1 for hours together, and neither nitrogen nor hydroxyamine is formed. The only thing we have noticed is always a trace of ammonia in experiments at IOO’, and this, we are inclined to believe, is due to some other cause than this reaction alone. At goo, the action is still exceedingly small, but as the temperature descends from about So0, it becomes rapidly greater with the descent. “We have tried cooling the tin-solution by a freezing mixture, but have not succeeded in getting better results than a t the ordinary temperature. Further experiments are needed, to be sure that a low temperature does not favourthereaction. . . . “When nitric acid is in presence of enough water to prevent any reaction taking place between it and hydrochloric acid, it is not perceptibly acted upon by stannous chloride. It may be left for a day or two mixed with these reagents, and remain, as well as the stannous chloride, unchanged. We have not yet established this fact by quantitativp estimation, air having been allowed access to the mixture in our experiments hitherto, but we have no reason to believe that it is incorrect. “If such a mixture is sufficiently dilute, it may even be boiled without change, but if not too dilute, reaction is thus set up. Another way of bringing about reaction in such a dilute solution, is to add sulphuric acid, already slightly diluted so that the heating effect of mixing it with water may be avoided, since that alone would afford an explanation of its action. It is thus seen that the addition of sulphuric acid, by diminishing the weakness of the solution in acid, serves to induce the reaction between the tin salt and the nitric and hydrochloric acids. “When dilute nitric acid added to the solution of stannous chloride in hydrochloric acid is not too dilute, reaction between them proceeds and continues until either the stannous chloride has all become stannic chloride, or 1 [Mr. Milligan suggests that this may be due chiefly to the sparing solubility of nitric oxide a t the higher temperature.]
ACTION OF METALS ON NITRIC ACID
489
until the solution has become too dilute in hydrochloric acid. In the latter case, we believe that we can cause the reaction to go on again, by passing hydrochloric acid gas into the solution, but we have not tried that yet. We have only a t present added sulphiiric acid as mentioned above with very favourable results. We have not succeeded, however, in destroying all nitric acid. “The reaction which takes place results in the conversion of so much of the nitric acid as is decomposed into hydroxyammonium chloride and nothing else-no ammonia and no gaseous body, provided the stannous chloride remains in excess. This statement is in direct opposition to that made by von Dumreicher, which is, that “acid stannous chloride acting on nitric acid yields hydroxylamine, but the reaction proceeds further, and the hydroxylamine is nearly simultaneously reduced to ammonia.” When the stannous chloride is in insufficient quantity, the conversion of the last portion of it is followed, first, by the solution becoming yellow and smelling of nitrosyl chloride, and then by an effervescence of nitrous oxide. If, on the disappearance of this yellow colour, more stannous chloride is at once added, the solution is at once bleached; and if this additional quantity of stannous chloride is insufficient, the liquid again becomes yellow after a time. The phenomena observed in this case, in fact, strongly remind one of those noticed by Vernon Harcourtl, in a solution of hydrogen peroxide and hydrogen iodide to which sodium thiosulphate is added each time coloration by iodine reappears. When we proceeded in this way, the solution soon became charged with nitrous oxide, which escaped with effervescence on stirring.” If Divers is correct in saying that nitric oxide is not reduced to ammonia by acidified stannous chloride and that ammonia is formed to some extent in presence of nitrogen peroxide, this means that the reduction of nitrogen peroxide by acidified stannous chloride does not pass through the nitric oxide stage. As there are no data to show whether the amount of ammonia is or is not equivalent to the amount of nitrogen peroxide, we do not know whether the ammonia is formed from Nz03or from the hypothetical first reduction product of nitrogen peroxide, N02H2. This reaction should be studied experimentally as soon as possible. Mr. Milligan does not feel certain that Divers is right and von Dumreicher wrong in the points at issue between them. If nitrous oxide is formed solely as a dehydration product of hyponitrous acid, we ought to get it on reduction of nitric oxide and under favourable conditions. If it is formed from the reaction between hyponitrous acid and hydroxylamine this would not be the case, because there is no hyponitrous acid when we start with nitric oxide. As a matter of fact we do not know a t all what is the first reduction product of nitric oxide. The most probable product is hyponitrous acid; but in that case there is no apparent reason why Divers should not have obtained some nitrous oxide. The whole matter of the reduction of nitric oxide calls for systematic study. J. Chem. SOC.20,460 (1861).
JVILDER D. BASCROFT
490
Guye and Schneider‘ have shown that at 280’ nitrous oside is reduced very slowly relatively to nitric oxide by nickel and hydrogen and that no nitrous oxide is formed during the reduction of nitric oxide. Nitrous oxide is evidently not ordinarily an intermediate stage in the formation of hydroxylamine, ammoiiia, or nitrogen from nitric oxide. The reduction to nitrous oside can be carried out by means of sulphurous acid.2 “Two molecules of nitric oxide react with one of sulphurous acid over water in several hours to forni aqueous sulphuric acid and one volume of nitrous oxide; but, according to R. Weber, the change is not complete even after fourteen days and it i. not rapid at the teniperature of the lead chamber, In presence of platinum black, the reduction to nitrous oxide and even to nitrogen talies place readily, the rate of reaction increasing with rising temperature. According to Lunge the dry gaqes do not react; but, in presence of water, there is a powerful reaction which causes a complete conversion to nitrous oxide when there is an excess of sulplnurour acid present, no nitrogen A mixture of nitric oxide, sulphurous acid and oxygen being formed always gives some nitrous oxide in presence of water even though there be a large excess of oxygen; in presence of sulphuric acid of spccific gravity 1.32, there is no appreciable trace either of nitrous oxide or of nitrogen. “When nitric oxide is iiiixed with a n equal volume of hydrogen sulphide, there is a decomposition in the course of a few hours into ammonium sulphide with a small amount of nityous oxide.” It seems probable that nitrous oxide is not in thiq case an intermediate pi oduct in the ammonia formation. The method of forination of nitrogen from nitric oxide is not sclttled definitely for all cases. Felgate3 s h o n ~ ithat nitric oxide and colloidal nickel react to forni nitrogen and nickel oxide. In this case there is no proof of any intermediate hydrogenation, though this may occur. Since acidified stannous chloride reduces nitric oxidc to hydroxylamine, we must postulate hydrogenation. If we me going to have hytlrogeii add on two 1))- two with the splitting off of water, as seems to occur with nitric acid, we are rather forced to postulate some such reactionq ns the follon7ing: 2NO+zHz --t (HN0H)z --t Nz+2H20 (HNOH)z+H2 +z N H ~ O H (Hl”rJOH)2+2H?--t N ~ H ~ + Z H ~ O (HNOH)z+sH? ---j Z N H ~ + Z H Z O . Since ammonia is probably not formed through reduction of hydrazine, it ,vould simplify matters if we could w i t e the last reaction: z;SO+jHz +2HYOH+;jH2 +2NH3$2H?O. The difficulty with this is that it i q purely paper woik and that we have no proof of any of these stages. I hope that this formulation of the problem may prove to be the fiist step in the solution. Sabatier and Sendereiis4 report that palladium sponge and hydrogen reduce nitric oxide completely to ammonia. Helvetica chim. Acta, 1, 133 (1918). Stickstoff,” 89 (1903) Chem. News. 108, 178 (1913). Compt. rend., 114, 1429 (18921.
* Spiegel: “Der
ACTION O F METALS ON N I T R I C ACID
491
The last paragraph has brought up the question of hydrazine which is usually not assigned any definite place among the reduction products of nitric acid. Hantzschl considers that the salts of dinitrososulphonic acid have the formula
/O\
and Duden2showed that potassium
RO -N - N.SOaR dinitrososulphonate gives hydrazine when reduced a t oo with sodium amalgam or with zinc dust and ammonia. KO .NzO.S03K+4H2 +NzH4+KzS03+2HzO. According to Spiegel,3 careful reduction with sodium converts hydrazine into ammonia. It was found by von Bracke14that hydrazine is obtained when a solution of sodium bisulphite, silver hyponitrite, and hydrochloric acid (the acid being equivalent to the silver) is reduced with zinc dust and glacial acetic acid in the cold. The hydrazine was not isolated as such, but a yellow precipitate of benzalazine was obtained on adding benzaldehyde. The presence of sulphite is apparently essential, just as the presence of sulphuric acid is desirable in the reduction of nitric acid to hydroxylamine. I should like to see the experiment tried of running nitric acid with and without nitrous acid slowly into a cold, concentrated solution of sulphurous acid. According to Spiege1,s “sulphurous acid reduces nitric acid very slowly unless sulphuric acid is present. If one boils one volume of nitric acid (sp. gr. 1.4)with five volumes of aqueous sulphurous acid, nitric and nitrous oxides are formed; in presence of less water there is a sudden, violent evolution of nitric oxide before boiling begins. A mixture of nitric and sulphuric acids reacts with sulphurous acid in different ways, depcnding on the concentration. If sulphur dioxide is passed into a mixture of oil of vitriol with ten percent of strong nitric acid, lead chamber crystals are formed if the saturated mixture is left t o itself for twenty-four hours. If t’hese are redissolved and more sulphur dioxide is passed in, the liquid becomes dark violet if left standing in a tightly closed bottle, and contains a mixture of sulphurous acid, which passes off in the air or in a vacuum, and nitrous acid.6 A mixture of nitric acid and dilute sulphuric acid (at least sp.g.1.34) is readily reduced by sulphurous acid to nitric oxide. If one drops liquid sulphurous acid into pure nitric acid monohydrate, there are formed red fumes and lead chamber crystals, which disappear again when an excess of the reducing agent is added, so that finally there is only a slightly colored sulphuric acid. On heating a mixture of the two acids, Fr6my observed a formation of nitrous oxide. . . . “Hydrogen sulphide is not attacked a t ordinary temperature by nitric acid containing no oxides of nitrogen, and there is no reaction when hydrogen Ber. 27, 3264 (1894). Ber. 27, 3498 (1894). “Der Stickstoff,” 571 (1903). Ber. 33, 2115 (1900). “Der Stickstoff,” 223, 102 (1903). 6 Girard and Pabst: Bull. (2) 30, 531 (1878).
492
WILDER D. BANCROFT
sulphide is passed into a pure acid of 1.18 sp. g. If the nitric acid contains only so much of the oxides of nitrogen as forms on standing exposed to air a t 2 5 ’ , the whole of the acid will be decomposed with formation of sulphur, sulphuric acid, ammonia, nitric oxide, and nitrogen.” “Nitrous acid in a good deal of water loses two thirds of its oxygen in presence of sulphurous acid and is reduced to nitrous oxide. When less water is present or when sulphur dioxide is passed into a solution of nitric acid in sulphuric acid of 1.4spg., nitric oxide is formed. Nitrous acid dissolved in oil of vitriol does not react at all with sulphurous acid.” Mr. Milligan believes that there is some reason for postulating the intermediate existence of the hypothetical acid, H2N205. He suggests writing the main chain of reductions:
/ “03
+HzNzO~+HNOz
.1
+HzN203
.1
*H2N20
”3
+NHzOH
+NHa.
- c l d
NO NO2 NzO Nz When this work was reported at the intersectional meeting of the American Chemical Society held in Syracuse in February, Professor R. A. Baker suggested that another factor t o be taken into account was the direct oxidizing action of nitric acid a t any rate in concentrated solutions. This is the point of view taken by Freer and Higleyl and we know that nitrogen peroxide will oxidize iron2 to passive iron. While not denying the possibility of direct oxidation, which one must assume if one puts carbon in fused potassium nitrate, I see no reason a t present for postulating anything but electrolytic corrosion under all ordinary conditions. It will be soon enough to modify this hypothesis when it has proved inadequate. Postulating a second hypothesis unnecessarily gives rise to indecision and delay. The general results of this paper are:I . The action of nitric acid on metals can be considered profitably as a case of electrolytic corrosion, the anode reaction being the formation of the nitrate. This is the view advocated by Armstrong in 1877. 2. The nature of the metal shows itself in at least three ways: in the reducing power as measured by over-voltage; in the catalytic action of the metal; and in the catalytic action of the salt of the metal. These three factors are to be studied in detail. 3 . Acidified ferrous sulphate reduces nitric acid practically completely to nitric oxide; acidified stannous chloride reduces it practically completely t o hydroxylamine; acidified titanous chloride, though a stronger reducing agent than stannous chloride, reduces nitric acid mainly t o nitric oxide. 4. The direct reduction products of nitric acid are nitrous acid, hyponitrous acid, hydroxylamine, and ammonia, while nitrogen peroxide, nitric Am Chem. J. 15, 71 (1893); 17, 18 (1895); 18, 587 (1896); 21, 377 (1899). Young and Hogg: J. Phys. Chem. 19, 617 (1915).
ACTION OF METALS ON NITRIC ACID
493
oxide, nitrous oxide, and nitrogen are due to secondary reactions. This is the view advocated by Armstrong in 1877. 5 . Hydroxylamine is not necessarily an intermediate stage in the formation of ammonia and nitric oxide is apparently not a necessary stage in the reduction of nitrogen peroxide, though the evidence in regard to the latter point needs confirmation. 6. I n the ordinary reduction of nitric acid, nitrogen probably occurs as the result of a reaction between hydroxylamine and hyponitrous acid more often than as the result of a reaction between nitrous acid and ammonia. In the reduction of nitric oxide, it seems possible that nitrogen may result from the dehydration of a hypothetical substance, NHOH or (NH0H)z. 7. Nitrous oxide is not ordinarily a direct reduction product of nitric oxide and it is probably never an intermediate stage in the reduction of nitric oxide to hydroxylamine, ammonia, or nitrogen. 8. Hydrazine is a direct reduction product of hyponitrous acid; but is not ordinarily an intermediate stage in the formation of ammonia. The presence of sulphite is apparently essent.ia1to the formation of hydrazine. 9. The presence of sulphuric acid increases the tendency to form hydroxylamine during the reduction of nitric acid ; but no satisfactory explanation of this has yet been offered. IO. While it is probable that nitrous oxide results, ordinarily from the dehydration of the syn form of hyponitrous acid, it is not known what the relation is between the syn and anti forms on the one hand and the production of hydroxylamine, ammonia, nitrogen, and hydrazine on the other hand. Cornell University.
