ACTIOK O F XITRIC ACID O X METALS BY N. R. DHAR
In three previous papers1 the action of nitric acid on metals has been studied. I n the first paper it has been proved that, contrary to the prevailing views, both ferrous and ferric salts accelerate largely the action of copper on nitric acid and an explanation of this behaviour has also been advanced. In the same paper it was also observed that out of fifty-six substances (electrolytes as well as non-electrolytes) investigated for their catalytic behaviour on the action of copper and nitric acid, twenty-two acted as accelerators and an equal number exerted a negative action in all concentrations, whilst eight of them are slight accelerators in small concentrations and are retarders in concentrated solutions. Only four of these substances have been found to be without effect in small concentrations, whilst in concentrated solution even these four substances are negative catalysts. In the second paper we have found that other metals like silver, mercury, lead, and an alloy of copper and nickel, brass, and an alloy of iron and nickel, etc., all dissolve more readily in dilute nitric acid in presence of ferrous salts. Moreover with the exception of mercury all the above metals and alloys dissoive more readily in nitric acid in presence of ferric salts as well. It has been also observed in the same paper that there is a periodicity in the action of nitric acid on the alloy of nickel and iron. I n the third paper interesting results have been obtained for the action of mercury and nitric acid. We shall discuss this question further later on. Recently Prof. Bancroft2 in an importapt paper has drawn the attention of chemists t o this important problem again. He has thrown out some important suggestions which are worth while investigating. For the last twelve years I have been interested in oxidation and reduction processes, their kinetics, their temperatures coefficients and the influence of catalysts, etc. I have always considered that the action of nitric acid on metals is a complicated oxidation-reduction process. Nitric acid is very rich in oxygen and contains 76% by weight of oxygen and is a very convenient oxidizing agent. Nitric acid is a fairly unstable oxidising agent and it can be converted into all sorts of reduction products and hence it is such a suitable and useful reagent in the chemical laboratory especially because most of the reduction products are more or less soluble in water. We know that substances like starch, sugar, camphor, pentamethylbenzene, chloronitrotoluene, glycerol, metabutyltoluene, etc., can be oxidised conveniently by nitric acid aided by catalysts like vanadium pentoxide and so on. Similarly most common Dhar: Proc. Akad. Wet. Amsterdam, 28, 545 (1919); Banerji and Dhar: Z. anorg Chem. 122, 73 (1922): Palit and Dhar: 124, 91 (1924). J. Phys. Chem. 28, 475 (1924).
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metals with the exception of gold and platinum are oxidised readily by nitric acid. Personally, I do not find any fundamental difference in the nature of the oxidation of metals and of other reducing substances like starch, sugar, camphor by nitric acid. We have made certain experiments on the oxidation of sugar and starch alone and in presence of catalysts like vanadium pentoxide, silver oxide, molybdic acid, manganous sulphate, sodium-nitrate and the following experimental results were obtained : With vanadium pentoxide as catalyst the oxidation of sucrose to oxalic acid takes place a t about 70' more rapidly than in the absence of the catalyst. Consequently the yield of oxalic acid from a known weight of sugar is much greater in presence of vanadium pentoxide provided the solution is not boiled. If the solution is boiled for concentrating it, the oxalic acid is further oxidised to carbonic acid and water, and consequently, in all cases when the solutions were boiled no yield of oxalic acid was obtained. Similar results were obtained by Kaumann, Moeser, and Lindenbauml. Consequently the partial oxidation of sugar to oxalic acid and the oxidation of oxalic acid to carbon dioxide and water by nitric acid are increased by vandium pentoxide. Similarly manganous salts have been found to be accelerators of both the above reactions and hence the yield of oxalic acid when sugar is oxidised by nitric acid in its presence is much smaller than in its absence when the solutions are concentrated by boiling. The yield of oxalic acid when sodium nitrate is used as a catalyst is a little less than in its absence because nitrous acid is also a feeble accelerator in the Oxidation of oxalic acid by nitric acid. On the other hand, with molybdic acid and silver oxide the yield of oxalic acid even when the solutions were boiled for concentration is greater than in the absence of these catalysts. Consequently the rate of oxidation of oxalic acid by nitric acid is decreased in presence of the above catalysts. I n a foregoing paper2 it has been proved that the rate of oxidation of oxalic acid by chromic acid is retarded by molybdic acid. On the other hand, Kempf3 has shown that the rate of oxidation of oxalates by potassium persulphate is accelerated by silver oxide. I have already shown that metallic copper, zinc, etc., dissolve copiously in a solution of ammonium nitrite4. Instead of ammonium nitrite a solution of any soluble nitrite and an ammonium salt can be used. If copper foil or copper wire is placed in a beaker containing a solution of a nitrite and an ammonium salt, the solution becomes blue in a few minutes. It has also been proved that other unstable nitrite solutions like ethyl ammonium nitrite can also dissolve copper a t the ordinary temperatures, whilst solutions of sodium and potassium nitrites which are fairly stable do not dissolve metals like copper. It is likely that the solvent power of nitrite solutions depends on their unstability. Ammonium nitrite will be partially hydrolysed into J. prakt. Chem. ( 2 ) 275, 143 (1907). Dhar: J. Chem. SOC. 111, 707 (1917). Ber. 38,3975 (1905). Dhar: Z. anorg. Chem. 119, 174 (1921)
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",OH and HK02 and both these substances are good solvents of copper. This is probably the reason for the dissolution of copper in ammonium nitrite solutions. Moreover it is well known that persulphate solutions slowly dissolve metals and the explanation is not far t o seek. Persulphate solutions are known to decompose in aqueous solution into acid sulphate and oxygen Kz S208+H20= 2KHSO4+0. The oxygen that is set free acts on the metal and forms oxide, and immediately the oxide dissolves in the free sulphuric acid. Thus the persulphate oxidises the metal and the acid dissolves the oxide. I n this way, the metals dissolve in persulphates by the joint action of oxygen and acid. S o other salt solutions can readily dissolve metals like Cu, h g , etc. If strips of copper are placed in test tubes and covered with solutions of persulphates, the solutions soon become blue, showing that the metal is attacked. If a fairly concentrated solution is employed the liquid becomes quite warm. There is no evolution of gas if the temperature is low, the two sulphates being the only products (K&08+Cu =K2SO4+C~iSod). The interaction of persulphate and metals in other cases, viz: h g , Au, Mg, Zn, Cd, Hg, Al, Ti, Sn, Pb, As, Sb, Bi, Cr, Sc, Te, U, &In, Fe, Si, Co, Pd, Pt, etc. has been investigated. All the metals except h u . and Pt, react with persulphate solutions either passing directly into solutions or remaining undissolved in the form of oxides or basic salts. The results show that those metals go into solution as anions which in their general chemical behaviour exhibit a marked non-metallic character e.g., Cr, Mn, Se, As, Mo, etc. Some metals of this type, however, such as Sb, are transformed into insoluble oxides. Elements which are distinctly metallic in character pass into solution as cations, the persulphate being decomposed sometimes with evolution of gas. With Zn, Hg, Cd, Al, Si, Co, etc., no gas is evolved and as for example in the case of Zn, the reaction may be represented by Zn+KzS&= Zii S04+K2S04. When gas is evolved it is found to be hydrogen due to the action of hydrogen ion set free by hydrolysis on metals like Zn, Mg, etc. Oxygen gas niay also be given off due the decomposition of persulphates. = 2KHS04+O) is accelerated by Au, and Pt, This action (K2S20s+R20 resulting in the decomposition of the persulphate. It is evident therefore that the chemical changes involved in the action of metals on persulphate solutions are closely allied t o those happening in the action of nitric acid on metals. Similarly metals like Cu, Ag, etc., can dissolve slowly in oxidising agents like H20z,chromic acid, potassium permanganate, etc., aided by acids like sulphuric acid, Just as nitric acid can oxidise several substances, numerous oxidations of organic substances have been effected by persulphates. I am inclined to the view that the action of nitric acid on metals is more or less similar to that of metals in persulphate solutions. If you allow nitric acid to come in contact with metals, they mill try to reduce the nitric acid to the nitrous state and will be converted into oxide and the oxide thus formed will immediately dissolve in the excess of the nitric acid. I n other words, this action is mainly an oxidation of the metals to the oxide state and its consequent solution in the nitric acid. It is needless t o mention that there
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are several side reactions along with the main chemical change. I shall discuss the special function of nitrous acid later on. I do not see any advantage in the electrolytic theory of corrosion. The assumption that all corrosion is electrolytic in nature does not lead to any important fact with regard to the solution of the problem in question. To my mind the hypothesis of Acworth and hrmstrongl who regard the first change as consisting of the replacement of the hydrogen of the acid by metals and the formation cf the metallic nitrate is also not suitable. It is well known that metals like Cu, Ag, Hg, etc., hardly replace hydrogen in acids like HC1, HzSO4. On the other hand, the above metals readily dissolve in ordinary dilute nitric acid. Theoretically, we can assume that an infinitely small amount of hydrogen is set free when metals like copper is put in contact with acids. I n order to explain the rapid action of metals on nitric acid we have to imagine that the hydrogen set free reacts instantaneously on the unchanged nitric acid; but we know from our experience in chemical dynamics that gases react rather slowly with liquids. Consequently there is considerable difficulty based on kinetics in the hypothesis of Acworth and Armstrong with regard to the action of nitric acid on metals Moreover, on the hypothesis of hcworth and Armstrong the action of metals on nitric acid is put in an entirely different category from that of nitric acid on substances like sugar, starch, camphor, etc., because in the latter cases we cannot make the assumption of the replacement of the hydrogen by the reducing agents. Similarly if we assume that all corrosion is electrolytic in nature me are faced with the same difficulty, because we have to assume that the action of nitric acid on metals is an entirely different type of reaction from that of the action of nitric acid on substances like sugar, starch, oxalic acid, charcoal, sulphur, camphcr, etc., on the other hand, if me imagine that the action of nitric acid on metals is essentially the formation of the oxide of the metal and its subsequent solution in the nitric acid or nitrous acid, then this action is immediately brought in line not only with the oxidation of other substances by nitric acid but all oxidation reactions in general. It is well known that when sugar is added to hot concentrated nitric acid, partial oxidation of the sugar molecules into oxalic acid takes place and nitrous fumes are given out, but if manganous sulphate i s present complete oxidation takes place and the sugar is converted into carbon dioxide and water by nitric acid simply because the oxidation of oxalic acid by nitric acid is largely accelerated by manganous salts. Harcourt and Esson2 have shown that the action of potassium permanganate on oxalic acid is accelerated by manganous salts. I n a previous paper3, I have shown that the action of chromic acid and oxalic acid is largely accelerated by managanous salts, similarly the oxidation of the oxalates by persulphates is accelerated by manganous salts. Consequently the oxidation of oxalic acid by nitric acid is more or less allied to its oxidation by other oxidising agents. J. Chem. SOC. 32, 56 (1877). ZPhil. Trans. 156, 198 (1866). 3 Dhar: J. Chem. SOC.111, 707 (1917).
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It is well known that manganese can exhibit different valences. Skraball has assumed the existence of various intermediate stages in the oxidation of oxalic acid by potassium permanganate. Similarly, intermediate stages have been observed in other oxidation-reduction processes. Thus we have proved that in the oxidation of arsenious acid by iodic acid, hypoiodous acid and probably iodous acid are intermediate stages. This is possible because compounds in the intermediate stages can exist in the free state. Thus manganates, manganites, manganic and manganous salts which are lower stages of oxidation than permanganate are known to exist. Different stages of the oxidation of chromium have also been supported by Jablczynski2. Similarly nitric acid can readily pass into nitrous acid, hyponitrous acid, etc. Considering these facts I do not find any special reason for assuming either that the action of nitric acid on metals is essentially electrolytic or that it is really a displacement of hydrogen in acids. On the other hand, the simpler assumption that the action of nitric acid on metals is a formation of its oxide and subsequent solution is more satisfactory and brings the action in line along with other oxidation reactions. More than eighty years ago Millon3 discovered the marked accelerating effect of nitrous acid on the action of nitric acid on metals like Cu, Ag, Hg, Bi, etc., and suggested that nitrous acid behaved more or less like an inorganic ferment, though this important observation was ridiculed by Gay-Lussac4. Following up the idea of Millon, Veley5 has shown that nitrous acid plays a very important part in the oxidation of metals and that metals like Cu, Ag, Hg, Bi, etc., have no action on cold dilute nitric acid unless a trace of nitrous acid is present. The nitrous acid may be present in the nitric acid as an impurity, or it may be formed by the incipient decomposition of nitric acid when it is warmed. According t o Veley, therefore, the dissolving of copper in nitric acid occurs according to the following equations: Cu+3HN03= C U ( N O ~ ) ~ + H N O ~ + H ~This O . reaction is a resultant of a series of consecu+ ~ K O ,by Cu tive reactions: C U + ~ H N O ~ = C U ( N O ~ ) ~ + ~ H ~ Ofollowed (N02)2+zHN03=CU(KO~)~+ZHT\TO~. The small trace of nitrous acid thus acts as a catalytic agent. It is well known that nitrous acid is continuously produced and decomposed according to the equation: 3HN02 HN03+2KO+H20. From my experiments in this line I am also convinced that nitrous acid is an important factor not only in the action of nitric acid on metals but also in its action on substances like sugar, starch, etc. It seems certain that nitric acid is at first reduced to nitrous acid by the action of metals like Bj, Hg, Cu, Ag, etc, and FeS04, sugar, etc. There is a lot of experimental support of the above view. Thus R&yeobtained mercurous nitrite Hgz(NOz)z(notmercurous nitrate) in the solid state by the action of
a
Z. anorg. Chem. 42,
I
(1904).
8. anorg. Chem. 60, 38 (1900). Compt. rend., 14, 104 (1842); Ann. Chim. Phys. (3) 6, 95 (1842).
Ann. Chim. Phys. (3) 6, 385 (1842). Phil. Trans. 182A, 279 (1891). Z. anorg. Chem. 1 2 , 365 (1896).
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mercury and dilute nitric acid. We have studied the action of mercury and nitric acid in detail and we always find that the first product of the reaction of mercury and dilute nitric acid is yellow crystals of HgZ(N02)Z. Strong nitric acid decomposes Hg2(N02)2 and hence with a strong nitric acid we get white crystals of Hg2(N03)2. If we start with dilute nitric acid and mercury at the ordinary temperature we always get yellow crystals of mercurous nitrite after some hours, but if the reaction is allowed to proceed in excess of nitric acid, say for a few days in open air in a beaker, the crystals of mercurous nitrite disappear gradually and white crystals of mercurous nitrate appear, because on evaporation of water, the nitric acid becomes inore concentrated and hence can decompose mercurous nitrite. It has been observed that the greater the surface of mercury the smaller is the yield of mercurous nitrite formed by the action of mercury and dilute nitric acid, because the active substance nitrous acid which is formed by the action of mercury on nitric acid is spread over a larger area. It has been observed that by keeping the amount of mercury the same and by changing the amount of nitric acid of definite concentration, the quantity of mercurous nitrite formed is greater a t the beginning in the beaker containing smaller quantities of nitric acid, whilst after the reaction has proceeded for twenty-four hours it has been observed that the yield of mercurous nitrite in the beakers containing an excess of acid is much greater. The explanation is not far t o seek; in the beakers containing smaller quantities of nitric acid the solution becomes saturated with respect to mercurous nitrite more readily than in beakers containing larger quantities of nitric acid and hence crystals of mercurous nitrite separate out more readily in beakers containing smaller quantities of nitric acid. On the other hand, in beakers containing larger quantities of nitric acid the yield of mercurous nitrite will be greater in the long run beCause of the larger amount of one of the reacting substances. In order to find out under what conditions the maximum yield of mercurous nitrite is obtained by the action of nitric acid on mercury, we made experiments with nitric acid of different concentrations and at different temperatures with and without catalysts. Acids of concentration 15.1%) 19.1%~ 22.870, 26% and 29.2% were used. It has been established by us that acid of 26% concentration produces the maximum yield of mercurous nitrite a t the temperature of about 30’. RAyl observes: “I have sometimes been struck with the remarkable fact that under analogous conditions the yield of mercurous nitrite has been very poor. This abnormal behaviour of nitric acid in isolated cases led me to undertake a close and systematic investigation of the disturbing causes. I t was soon discovered that the retarding effect was due to the presence of minute quantities of iron in the acid, for whenever the acid was redistilled in glass retorts no such anomaly was noticed. The residue after distillation was invariably ferric iron.” RAY proved that in presence of ferric iron less mercury dissolves in dilute nitric acid than in its absence. The interval of reaction was J. Chem. SOC. 99, 1012 (1911).
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four hours and in this time no crystals of mercurous nitrite separated. We repeated some of these experiments carefully and we confirmed the experiments of R$y and we also found that when the interval of the reaction is four and a half hours, less mercury dissolves in nitric acid in presence of ferric salts. In other words, ferric nitrate is a retarder in the action of nitric acid and mercury, though we have observed that ferrous sulphate accelerates the action of mercury and nitric acid. We have also observed an interesting fact that though the amount of mercury dissolved by dilute nitric acid in presence of ferric nitrate is less than in its absence, the yield of mercurous nitrite is much greater in the presence of ferric nitrate than in its absence. We have, in all our experiments with nitric acid of 2 6 % concentration, got much larger quantities of mercurous nitrite formed in presence of ferric nitrite than in its absence. This observation contradicts the statement of R$y already quoted. We have also observed that in presence of mercurous nitrate the yield of mercurous nitrite is much greater than in its absence. We have found out that sodium nitrite is a great accelerator in the action of mercury and dilute nitric acid. It is an interesting fact that in presence of sodium nitrite mercury dissolves copiously in nitric acid but there is hardly any yield of mercurous nitrite. This observation can be explained very readily from the following facts: It is well known that mercuric nitrite, silver nitrite, etc., are much less ionised than ordinary salts in aqueous solutions and these nitrites readily form complex nitrites with alkali or alkaline earth nitrites'. From the researches of R$yz it is clear that mercurous nitrite decomposes in presence of water according to the following equation: Hgz(N0z) J Hg+Hg(KOa)z If to this mixture sodium nitrite be added the mercuric nitrite will immediately combine with sodium nitrite forming a complex nitrite of the type ?rTazHg(N02)4. Hence the equilibrium will be displaced from left to right. In other words, the whole of the mercurous nitrite will be converted into sodium mercurinitrite. Moreover, there is the possibility of the formation of sodium mercuronitrite formed by the interaction of sodium nitrite and mercurous nitrite and this complex nitrite is likely to be more soluble than mercurous nitrite. Consequently the action of nitric acid on mercury in presence of sodium nitrite produces hardly any mercurous nitrite. The greater yield of mercurous nitrite by the action of mercury and dilute nitric acid in presence of mercurous nitrate is due to the decreased solubility of mercurous nitrite due to the presence of a common ion. It is difficult to explain at this stage the increased yield of mercurous nitrite in presence of ferric nitrite. It has been observed that in presence of urea the action of nitric acid on mercury does not produce any mercurous nitrite because the urea present completely destroys the nitrous acid originally existing in the nitric acid as well as that produced by the reduction of Abegg and Pick: Z. anorg. Chem. 51, 965 (1912). J. Chem. Soc. 87, 777 (1905).
I
(1906): R$y and Dhar: J. Chem. SOC. 101,
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nitric acid by mercury or by nitric oxides; hence very little mercury dissolves in nitric acid in presence of urea and so there is practically no yield of mercurous nitrite in this case. Stansbiel has suggested that with copper and silver in the presence of dilute nitric acid, approximately half the nitric acid is reduced t o nitrous acid. The experimental results show that with silver the following expresses the reaction : 2Ag+ 2 H N 0 3 = AgK03+&NO2 +HzO In the case of copper the reaction is most probably as follows: 3 C u S 6 I l N O s = CU(N03, KO2)+3H20. Stansbie observes that no gas is given off when dilute nitric acid is used and the amount of reaction taking place is increased by the addition of free nitrous acid. Consequently it is clear that nitrous acid plays a very important part in the action of nitric acid on metals like copper, silver, mercury, bismuth, etc. The formation of nitrous acid by the action of metals on nitric acid can be readily demonstrated in the following way: If metallic copper is covered with 2 0 % nitric acid a t the ordinary temperature, a t the beginning there is very little chemical change. After a few minutes, blue colour appears at the bottom of the tube, due to the dissolving of copper and there is hardly any evolution of gas. S o w if urea is added to the reacting mixture, copious evolution of nitrogen and carbon dioxide takes place, showing the existence of nitrous acid in the mixture. Curiously enough the action of nitric acid on lead is not markedly accelerated by nitrous acid which is really produced in this case also as a first step in the reaction*. This behaviour can be partly explained from the following facts : Lead nitrate is a sparingly soluble substance in comparison with silver nitrate, copper nitrate, mercurous nitrate, etc., especially in presence of nitric acid. It is quite likely that when lead is brought in presence of fairly dilute nitric acid there is the possibility of the formation of lead nitrite and lead nitrate just as in the previous cases we always get a mixture of nitrite and nitrate. In the case of lead, it seems certain that lead nitrite is much more soluble than lead nitrate and consequently lead nitrate comes out of the solution very readily and the amount of lead ions from the dissociation of lead nitrite is very small. Moreover the lead nitrite formed in the solution is gradually decomposed by the excess of nitric acid with the formation of sparingly soluble lead nitrate. Consequently the chemical change between lea,d and nitric acid results mainly in the formation of lead nitrate. Hence the catalytic action of nitrous acid cannot be prominent in this reaction. On the other hand with mercury and silver the case is otherwise, because with those met)als, the nitrites are much less soluble than the nitrates and these nitrites tend to come out as precipitates. Hence in the action of nitric acid on mercury, silver, copper, etc., the action of nitrous acid is more prominent J. Sac. Chem. Ind. 32, 311 (1913). Compare Veley: J. Sac. Chem. Ind. 10, 206 (1891).
than in the case of lead. Looking at the problem from another point of view, we can explain the difference in the behaviour of lead and mercury towards nitric acid from the following facts: It is well known that nitrous acid decomposes according to the following equation : 3HNO2* "Os+ 2NO+H20. With mercury, silver, etc., the (.nitrites being sparingly soluble come out readily as crystals in their action with nitric acid. Consequently nitrous acid is being continually removed from the solution due to the precipitation of the nitrites. Hence, in order to establish the above equilibrium, more and more nitric acid will be converted into nitrous acid. In other words, in these cases, it is t o be expected on theoretical grounds that nitric acid will gradually pass into the nitrous state and hence the action of nitric acid will be far less prominent than that of nitrous acid which really appears to be the chief solvent of these metals and these conclusions are actually corroborated by the experiments of Millon, Veley, RAY, Stansbie and others. On the other hand, with lead the case is different because lead nitrate which is a very sparingly soluble substance comes out as a precipitate more readily than the nitrite. In other words, nitric acid is removed from the sphere of the chemical change due to the separation of lead nitrate. In order to establish the foregoing equilibrium nitrous acid will be converted into nitric acid and in this chemical change the nitric acid will play more prominent part than the nitrous acid. Now Millon and T'eley have pointed out that the presence of ferrous sulphate "which removes the nitrous acid as fast as it might be formed" serves to prevent the chemical change between nitric acid and the metals. But we have found that ferrous sulphate markedly accelerates the action of nitric acid on copper, silver, mercury, nickel, silver coin, nickel-copper alloy (Ni soyG,Cu 50%), brass and iron-nickel alloy (Fe 50%) Ni 50%). Other ferrous salts have also been found to be accelerators. In explaining the accelerating action of ferrous salts, I have advanced the argument "that a part of the ferrous ion reduces the nitric acid to nitric oxide and passes into the ferric state. The nitric oxide dissolves in the ferrous salt solution forming the unstable bivalent ion, FeNO' .' The dissolved nitric oxide then reduces a part of the nitric acid according to the following equation: HKOs+2NO+H20+ 3HN02. It is quite possible that some nitrous acid is produced by the direct reduction of nitric acid by ferrous ions. The formation of nitrous acid either by the direct reduction of nitric acid by ferrous salts or by the indirect reduction through the intervention of nitric oxide is proved by the following experiment. If nitric acid of 20% strength be taken in a test tube and a crystal of ferrous ammonium sulphate or ferrous sulphate be added to it, almost immediately the crystal is covered with the deep-brown FeXO" ion and a little nitric oxide also escapes. If urea crystals are now added, they are immediately oxidised with the evolution of carbon dioxide and nitrogen, indicating the presence of nitrous acid. So in the presence of ferrous salts, nitrous acid, which is the active substapce in the action of nitric acid on cop-
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per, is formed when we have an excess of nitric acid. This’explains the accelerating influence of ferrous salts in the complete dissolving of copper in zoyo nitric acid. As a matter of fact, the accelerating influence of ferrous salts is slightly greater than the accelerating effect of sodium nitrite on the dissolution of copper in nitric acid. The greater the concentration of the ferrous salt, the greater is the acceleration. Ferric sulphate, ferric nitrate and ferric chloride exert a marked accelerating effect though their activity is slightly less than that of sodium nitrite and the accelerating effect is proportional to the concentration of the ferric salt, It would appear that the acid radicle in this case also, plays no part. The explanation of this activation seems to be in the reduction of ferric salts by the nitric oxide which is a product of the chemical change between nitric acid and copper. The ferrous salt, may thus be formed, will reduce a part of the nitric acid to nitrous acid, which activates the action of nitric acid on copper. It seems plausible that a part of the ferric salt would be reduced to the ferrous state by the metallic copper. It is well known that when a solution of a ferric salt is shaken with metallic copper, the ferric salt is partly reduced to the ferrous state and the copper is oxidised to the cupric salt, an equilibrium being set up: aFeCls+Cu~-zFeC12+CuC12. The ferrous salt thus formed reduces the nitric acid to nitrous acid, which accelerates the action of nitric acid on copper. I n a similar way the accelerating effect of arsenious oxide, strychnine sulphate, phthalic anhydride, etc., may be explained on the basis of the formation of nitrous acid by the action of these reducing agents on the nitric acid’.” “In this connection, it is interesting to observe, that several reactions in which nitric acid is the oxidising agent, are autocatalytic. As for example, the action of nitric acid on metals like Cu, Ag, Bi, Hg, etc., on starch, on sugar, on arsenious oxide, on hydrogen iodide2, on nitric oxide3, etc., become more pronounced as the chemical changes proceed. The explanation is not far t o seek. The nitrous acid is the active substance and its concentration and hence the reaction velocity increase with the progress of the chemical change. I n all these cases I have found that the chemical change becomes more rapid when a nitrite is added at the commencement of the reaction. It has been observed that the chemical change between nitric acid and copper may be practically stopped by agitating vigorously the tube containing copper and nitric acid, because the nitrous acid cannot accumulate round the copper.” Just as the action of nitric acid on metals is accelerated by nitrous acid, I have repeatedly observed that the action of the same substance on starch, sugar, oxalic acid, arsenious oxide, etc., is also accelerated by nitrous acid. Dhar: Proc. Akad. Wet. Amsterdam 28, 545 (1919).
* Eckstadt: Z. anorg. Chem. 29, 57 (1901).
Lewis and Edgar: J. Am. Chem. Soc. 33, 292 (191I ) .
Here again we find there is no essential difference in the nature of the action between nitric and the metals on the one hand and nitric acid and reducing agents like sugar, starch, e t a , on the other hand, it is interesting to note that the oxidation of iodine to iodic acid by concentrated nitric acid is not accelerated by nitrous acid. Ostwald has observed that nitric acid in which copper has already been dissolved readily acts on further quantities of copper and he has compared this phenomenon with memory and habit. Apparently, the reason for this peculiarity is that the nitric acid in which copper is already dissolved, contains free nitrous acid, which markedly accelerates the change. Hence, from the foregoing experimental evidence we find that nitrous acid is likely to be always produced in the first stage. The nitrous acid can be further reduced to hyponitrous acid, or hydroxylamine, or, hydrazine or to ammonia. As I have already said nitrous acid is an unstable substance and it readily decomposes into nitric aoid and nitric oxide, 3 H N 0 2 S - H n ’ O s +zNO+HzO). Hence the production of nitric oxide by the action of nitric acid on metals like copper, silver, mercury, bismuth, etc., and reducing agents like ferrous sulphate, titanous chloride, sugar, starch, etc., is really due to the decomposition of nitrous acid which is the first product in the reduction of nitric acid. Hence in most reactionsof nitric acid and reducing agents we get the evolution of nitric oxide due to the decomposition of nitrous acid. It seems likely that nitrous oxide which is sometimes a product of the reaction between nitric acid and a reducing agent is mainly due to the decomposition of hyponitrous acid, HzNzOz= XzO+HzO. At the same time one cannot ignore the possibility of the formation of 1320 by the action of hydroxylamine and nitrous acid or by the decomposition of ammonium nitrate. But to my mind the formation of nitrous oxide in the action of nitric acid on metals is mainly due to its production by the decomposition of hyponitrous acid. It is well known that more nitrous oxide is formed by the action of lead on nitric acid than that produced by the action of copper and nitric acid. Lead reacts much more readily than copper producing with equal concentration of nitric acid relatively more nitrous oxide and less of the higher oxides of nitrogen1. This is certainly due to the fact that with lead more hyponitrous acid is formed than with copper because lead is a better reducing agent than copper (compare electrode potentials of copper and lead). Consequently the production of nitric oxide of nitrous oxide will depend on the reducing power of the metal in question. Bancroft has noted that the direct reduction products of nitric acid are nitrous acid, hyponitrous acid, hydroxylamine and ammonia while nitrogen peroxide, nitric oxide, nitrous oxide, and nitrogen are due to secondary reactions. Similarly Stansbie2 has observed that the liberation of nitrogen Compare Freer and Higley: Am. Chem. J. 21, 377 (1899). ,J. SOC.Chem. Ind. 32,311 (1913).
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peroxide, which is the result of the interaction of nitrous and nitric acids, takes place as the concentration of the nitric acid increases, thus HN02+ H N 0 , ~ H z 0 + ~ N 0 2But . it is very difficult to make a definitely hard and fast rule about the direct and indirect products of the reaction between nitric acid and metals, simply because all the reduction products of nitric acid are fairly unstable and are interconvertible one into the other under suitable conditions. Thus Reynolds and Taylor1have shown that in presence of light strong nitric acid decomposes according to the following equation: 4 H N O 3 ~ 2 H 2 O + 2 N z O 4 + O 2According . to this reaction K204 can be also a direct product of the decomposition of nitric acid. SpiegeP has observed that nitrous acid in a good deal of water loses two thirds of its oxygen in presence of H2S03and is reduced to nitrous oxide when less water is present or when SOz is passed into a solution of nitric acid in HzS04, nitric oxide is formed. This is due t o the fact that the reduction of nitric acid or nitrous acid to hyponitrous acid is more favourable in dilute solutions than in concentrated ones. It seems probable that very dilute solutions of nitric acid are more apt to pass into the hyponitrous stage by the action of suitable reducing agents than concentrated solutions of the acid. The hyponitrous acid formed from dilute solutions will decompose into nitrous oxide and water. Commenting on the work of Divers on the formation of hydroxylamine, Bancroft has observed that “the presence of sulphuric acid increases the tendency to form hydroxylamine during the reduction of nitric acid by zinc;” but that no satisfactory explanation of this has yet been offered. I suggest the following explanation of the increased yield of hydroxylamine in presence of sulphuric acid. Ammonium hydroxide is certainly more basic than hydroxylamine hydroxide. Consequently the formation of ammonia is undoubtedly more difficult than the formation of hydroxylamine in strongly acid solutions. Noreover metallic zinc alone is a better reducing agent than zinc and sulphuric acid because by the action of zinc and sulphuric acid hydrogen is set free, which is far less reducing in its properties than metallic zinc. (Compare electrode potentials of hydrogen and zinc). Consequently with the better reducing agent zinc alone, nitric acid can pass into ammonia whilst with zinc sulphuric acid which is not as good a reducing agent as zinc alone we get mainly hydroxylamine. Divers and Shimidzu, have observed: “In one experiment I O cc of a solution of -hydroxyammonium chloride containing 0.033 gram of hydroxyamine were made up to I O O cc with water containing a little H z S 0 4 . 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 HzS04 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. TTe then tried the action of zinc 1
J. Chem. SOC.101, 131 (1912). Stickstoff,” 571 (1903). J. Chem. SOC.47, 587 (1885).
* “Der
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N. R. DHBR
alone upon a dilute solution of hydroxyammonium chloride. Here there was a marked destruction of the hydroxyamine in two hours, one-eigth to onefourth disappearing.” It appears from the above observations that hydroxylamine is reduced by metallic zinc alone to ammonia and not by zinc and acid. As has been already explained this happens because metallic zinc alone is a better reducing agent than zinc and sulphuric acid. Divers observes that hydroxylamine is destroyed, but I am sure it is converted into ammonium salts. Divers has divided metals into two classes as regards their behaviour with nitric acid, namely those which yield nitrite and nitrate with water but exert no further action forming neither ammonia nor hydroxylamine. In this class are placed Ag, Hg, Cu, and Bi. Secondly those which form ammonia and generally also hydroxylamine but do not yield nitrite or nitrous acid with free nitric acid. On the other hand they readily form nitrite. To this class belong Zn, Cd, Mg, AI, Pb, Fe, and the alkali metals. I am of the opinion that this division is a sound one. With the metals hg, Hg, Cu, Bi, and HNO, we always get nitrous acid as the first reduction product and this unstable substance decomposes into nitric oxide and nitric acid. So the main gaseous product obtained by the action of dilute nitric acid on Ag, Hg, Cu, Bi, etc., at the ordinary temperature is nitric oxide. Small quantities by hyponitrous acid may also be formed by the reduction of nitric or notrous acid by the above metals. The hyponitrous acid being unstable will mainly decompose into nitrous oxide and water. This is certainly the origin of nitrous oxide obtained under certain conditions in the oxidation of metals like Cu, Hg, Bi, etc., by nitric acid. It is well known that in the action of copper and nitric acid more and more nitrous oxide is set free instead of nitric oxide as the cupric nitrate goes on increasing in the system. It seems likely that in presence of cupric nitrate and in dilute solutions of nitric acid, nitrous acid can be reduced to the hyponitrous state which readily breaks up into nitrous oxide and water. Traces of nitrogen are sometimes set free by the action of metals like Cu, Ag, Hg, etc., and an explanation is difficult to find. Kitrogen can be formed either by the action of ammonia and nitrous acid or by the action of hydroylamine and hyponitrous acid. It is not sound to assume the formation of either ammonia or hydroxylamine by the action of nitric acid on metals like Cu, Ag, etc. One has to take recourse to the suggestion of Hantzsch and Kaufmannl who state that hyponitrous acid may also break down according to the following equation: 3H2X\;2O::= 2NH,+zN203 It is likely that ammonia will immediately react with N203 with the formation of nitrogen and water. This is the possible explanation of the formation of nitrogen in the action of nitric acid on metals like Ag, Cu, etc. Just as nitrous acid is known to decompose according to the equation 3HN02+ HN03+2K0+H20, similarly it is very likely that hyponitrous acid ‘Ann. 232, 3 1 7 (1896).
ACTION O F S I T R I C A C I D ON METAIX
155
ill be in equilibrium with nitrous acid according to the equation3H2KzOZe 2HN02+2Sz+2Hz0 and the nitrous acid produced will be in equilibrium with nitric acid. Hence in a solution we should get HN03*HNOz* H2K202.Moreover, Berthelot and Ogierl and Rhy and Ganguliz,have shown that hyponitrous acid can also decompcse according to the following equation : 5RzP\Tz02= 4HzO 2 HxO 3 4 x 2 Metals like Zn, Mg, Sn, Fe, etc., which are strong reducing agents and which set free hydrogen by their action on dilute acids can certainly reduce nitric acid to hydroxylamine, hydrazine, and ammonia; sometimes free hydrogen also escapes. It is very likely that the observations of von Dumreicher3 claiming the possibility of the reduction of nitric oxide to ammonia by acidified stannous chloride are correct. It appears to me that Divers is incorrect in saying that nitric oxide is not reduced to ammonia by acidified stannous chloride and that ammonia is only formed to some extent in presence of nitrogen peroxide. Experiments on these lines are in progress jn this laboratory. When nitric acid is reduced theoretically the following compounds may be produced : (1) (2) ( 31 (4) (5) H,N~O,--tH,~~O,--tH~~~O~--tH,N,O~~H~T\’~Oz~H~~~O (6) (7) (8) +XH20H--+T\T2H,+KHa.
+
+
Of these Xos. ( I ) , (3) and ( 5 ) are extremely unstable and are of doubtful existence. It is most interesting that in the above chain of compounds, those with an even number of oxygen atoms are much stabler than the ones mitl1 an odd number of oxygen atoms. It may be that in compounds with an even number of oxygen atoms the valency of nitrogen is saturated and is possibly five whilst in the compounds with an odd number of oxygen atoms, the valency of nitrogen is unsaturated and the compounds are in a strained condition. It will be very interesting if one could isolate Compounds Kos. ( I ) and (5). If we take metals like silver, copper, mercury bismuth, etc., which do not set free hydrogen by their action on dilute acids (HC1, HzS04,etc.), theoretically we should get compounds Sos. ( I ) , ( 2 ) , ( 3 ) , (4),(s), but not compounds ( 6 ) , (7) and (8). The above metals can take up oxygen from nitric acid but as they cannot produce hydrogen by their action on acids or water, they should not hydrogenate nitric acid and form hydroxylamine, hydrazine, 01 ammonia. So with the above metals the formation of hydroxylamine and hydrazine and ammonia seems practically impossible. They can form compounds ( I ) , ( 2 ) , (3) (4) directly and by the decomposition of these compounds we can get NOz, NO, SZO, and Sz. With metals like Zn, Mg, Al, Sn, etc.. which can readily give off hydrogen from acids, along with nitrates and nitrites we can expect the formation of Compt. rend. 96, 30, 84 (1883). J. Chem. SOC.91, 1866 (1907). 3 J. Chem. SOC.20, 460 (1861).
I 56
Ir;.
R. DHAR
hydroxylamine, hydrazine and ammonia and the actual hydrogenation of nitric acid is possible because the above metals are very good reducing agents and can actually add hydrogen to nitric acid whilst at the same time they take away oxygen from nitric acid. On the other hand, metals like Cu, Hg, Ag, Bi, etc., can only take away oxygen from nitric acid and cannot add hydrogen to it and that is why with the above metals and nitric acid we get HN02, (HN0)2 and their decomposition products and not NHZOH, NzH4 and “,OH. It is well known that hyponitrous acid can be reduced by hydrazine under favourable conditions. Consequently it seems very likely that hydrazine can also be a direct reduction product of nitric acid. Bancroft has observed: “In 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.” This statement seems very sound in all cases where reduction of nitric acid is effected by metals like Zn, R l g , Sn, etc., where there is the possibility of the formation of hydroxylamine. As hydrcxylamine hydroxide is less basic than ammonium hydroxide, it is natural that the former compound would be formed inore readily in an acid solution than the latter. Moreover, ammonia being certainly in a higher state of reduction than hydroxylamine, is formed with greater difficulty in the reduction of nitric acid by reducing agents than hydroxylamine. Hence, on the whole, there is far greater possibility of the formation of hydroxylamine than that of ammonia and the former compound will react on hyponitrous acid forming nitrogen. But in the case of metals like Ag, Cu, Iig, Bi, etc., the formation of hydroxylamine or hydrazine or ammonia is extremely doubtful as has already been emphasised. Consequently, the formation of traces of nitrogen sometimes obtained in the action of metals like Ag, Cu, Hg, etc., on nitric acid cannot be satisfactorily explained by the reaction between hydroxylamine and hyponitrous acid as suggested by Bancroft. As we have already observed, we have to explain the formation cf nitrogen in these cases by the decomposition of hyponitrous acid according to following equations : ;HzT\r,O, = 4H20f 2HN03+4K2 3H2XzO2 = 2Hlr’02+2S2+2H20 Veleyl suggested that nitrous acid decomposes according to the following equation : 3HT\J02+HN03+2NO+Hz0 Lewis and Edgar2have studied the equilibrium 3HS02q=tHN03+ 2XO+ HzO, and have shown that the equilibrium constant changes with the concentration of nitric acid. It seems likely that secondary reactions interfere with the main reaction and affect the equilibrium constants of the main reaction. ‘Proc. Roy. SOC. 5 2 , 27 (1893); compare Veley and Manley: 62, (1901).
J. Am Chem. SOC.33,
292
(1911).
223
(1897); 68, 128
ACTION O F KITRIC ACID ON METALS
I57
Following the analogy of nitrous acid, I suggest that hyponitrous acid will undergo change according to the following equilibrium relation. ~H~N~O~=ZHT\TO~+N~+H~O (1) Although R$y and Gangulil have shown that hyponitrous acid decomposes into nitric acid and not to nitrous acid.
+
5HzKzOz = 4HzO 2HN03+ 4Nz (2) I am of the opinion that hyponitrous acid decomposes into nitrous acid which in its turn passes into nitric acid. We are trying to establish equilibrium experimentally. De Girard and de Saporta2, Dey and Sen3 and Sommer4have studied the action of nitrous acid on hydrazine. Solutions of metallic nitrites and hydrazine disulphate react vigorously even a t oo, the gases evolved containing two volumes of nitrous oxide to one volume of nitrogen. Sommer has shown that hydrazine nitrite XzHdHNOz decomposes according to the equation NzH4HNO2= "3+N20 +HzO and this decomposition is very greatly accelerated by nitrous acid. The decomposition of ammonium nitrite has also been shown to be accelerated by the presence of nitrous acid5. If hydrazine disulphate and barium nitrite react together, the products are ammonium nitrite and nitrous oxide. Consequently under similar conditions hydrazine nitrite is less stable than ammonium nitirite. From the foregoing facts it is clear that hydraxine is more readily acted upon by nitrous acid and hence the possibility of its existence as a reduction product of nitric acid is much less than that of ammonia. Moreover Tanatar6, Purgotti and Zanichelli' and Gutbier and XeundlingeP have shown that hydrazine decomposes into KHBand Nzin presence of platinum. It is quite possible that other catalysts might accelerate the decompcsition of hydrazine according to the equation 3r\TpH4=N2+qNH3and that in the reduction of nitric acid by reducing agents, the hydrazine formed will partly decompose according t o the above equation. Consequently me seldom get hydrazine in the reduction of nitric acid by metals like Zn, AI, Fe, Mg, etc. I am of the opinion that hydrazine can be a direct product in the reduction of nitric acid by metals like Zn, Mg, etc., although hydrazine is not usually found in such reaction and we get ammonia instead, because hydrazine is more readily oxidized by nitrous acid than ammonia. It has been already stated that the following substances are likely to be produced in the reduction of nitric acid by metals or other reducing agents: 1 J. Chem. SOC. 91, 1866 (1907). *Bull. 31 I I I , 905 (1904). 3 2 . anorg. Chem. 71, 236 (1911). anorg. Chem. 83, 119 (1913). 5Veley: J. Chem. Soc. 83, 736 (1903); Blanchard: Z. physik. Chem. 41, 681 (1902). 6 Z. physik. Chem. 40, 475; 41, 37 (1902). 7 Gaze. 34 I, 57 (1904). SZ. physik. Chem. 84, 203 (1913).
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N. R . DHAR
H Z N Z OHzN204, ~, H~N203,HzNzOZ,H2N20, SH20H, N z H ~and "3. Now some of these compounds are extremely unstable and decompose readily into SOZ, NO, K2O and N2, whilst some others will react on each other also forming more or less the same substances-that is why all the direct reduction products of nitric acid cannot be obtained a t the same time. In this connection it will be interesting to note that according to Sommer, hydrazoic acid is formed as a product of the secondary reaction between hydrazine and nitrous acid and is due to the action of nitrous acid on the ion S 2 H 6derived from the dissociation of X2HJYO2 (K2H4HNO2).
Summary I. It appears that there is no fundamental difference in the nature of the oxidation of metals and of other substances like starch, sugar, etc., by nitric acid. 2. The action of nitric acid on metals is similar to that of solutions of persulphates on metals. 3. The hypothesis that the action of nitric acid on metals is electrolytic in nature and that of Acworth and Armstrong who regard that hydrogen of the acid is first replaced by the metals, seem inadequate. 4. The partial oxidation of sugar by nitric acid to oxalic acid is accelerated by vanadium pentoxide and manganous sulphate. 5 . The complete oxidation of oxalic acid by nitric acid is accelerated by vanadium pentoxide, manganous sulphate and nitrous acid but retarded by molybdic acid, silver oxide, etc. 6 . The oxidations of metals, starch, sugar, oxalic acid, arsenious oxide, etc., by nitric acid are all autocatalytic and are accelerated by nitrous acid. 7. The action of nitric acid on metals like copper, silver, etc., is accelerated by ferrous and ferric salts, because these salts help in the formation of the active substance, nitrous acid. 8. Metallic copper disolves in ammonium nitrite solution (or in a mixture of any ammonium salt and a nitrite) because the products of hydrolysis S H 4 0 H and H N 0 2 derived from ammonium nitrite (NH4N02+Hz0 S N H 4 0 H + H N 0 2 ) exert a marked solvent action on copper. 9. It seems certain that nitric acid is a t first reduced to nitrous acid, by the action of metals like Bi, Hg, Cu, Xg, etc., FeS04, sugar and similar reducing agents. Sitric oxide (NO) is produced in these reactions mainly by the decomposition of nitrous acid. IO. Mercury dissolves copiously in nitric acid t o which sodium nitrite is added, but no mercurous nitrite is produced in this way. Nitric acid of ~ 6 concentration 7 ~ by its action on mercury produces the maximum yield of mercurous nitrite. 11. By the action of nitric acid on metals like Cu, Ag, Hg, and Bi, we which can only can get H N 0 2 or HzNzOzbut not NHzOH, NzH, and "3, be formed by the action of metals like Zn, Mg, Al? Fe, Cd, etc. , which can
A C T I O S O F N I T R I C A C I D OK METALS
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actually hydrogenate nitric acid, whilst Cu, hg, Bi, Hg, etc., can only take away oxygen from nitric acid and cannot add hydrogen to it, as these metals do not set free hydrogen from acids or from water. I 2. It seems probable that hyponitrous acid will be decomposed according to the following reaction : 3HzNzOz~2H~Oz+n’z+~H,O, Hence in a solution containing hyponitrous acid we get H,N202=HNO2 aHN03. 13. Hydrazine can be a direct product in the reduction of nitric acid by metals, although ive seldom get it in this way because it is much more readily oxidized by nitrous acid than ammonia. 14. The formation of traces of nitrogen obtained in the oxidation of metals like Cu, Ag, etc., by nitric acid is due to the decomposition of hyponitrous acid. The formation of nitrogen in these cases cannot be due to the interaction of hydroxylamine and hyponitrous acid or that of ammonia and nitrous acid, because hydroxylamine and ammonia are not formed by the action of nitric acid on copper, silver, etc. Chemical Laboratory, Allahabad University, Allahabad, India, August 7, lQ2.4.
