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I n order to make real progress and use this awakened popular interest, certain fundamental things must be accomplished. I n the past the theory and mechanism of corrosion were so indefinite and apparently complex that the scientist almost gave up in despair. This led the practical man to attempt tests of an engineering nature, without much regard to theory, in an effortto get some results quickly. As might be expected under such conditions, much time and money were spent in order to obtain a relatively few facts. Happily, during this period, progress was being made on the scientific side of the question, both in this country and abroad. Such men as Bancroft, Bengough, Evans, Aston, McKay, Fink, Speller, and Whitman-all added their contributions to the theory, while valuable data have been obtained and correlated by Calcott. Hatfield, Turner and Hamlin, Whittaker and Van Patten, and, although a few years ago many widely divergent views were held, it is a pleasure to report agreement in almost all questions a t the present time. The tests by the practical men mentioned above were brought about partially by the fact that no central agency in which chemists and engineers had representation was in existence. The AMERICAN CHEMICAL SOCIETYnow has a committee on the subject, whose aim is to study the theory and mechanism of corrosion in collaboration with a similar committee of the American Electrochemical Society. No other societies are a t present studying this phase of the subject, and it seems eminently proper that it should remain under these auspices. We have, then, a definite plan to follow in order to get the fundamental data on which to build our theory structure, and has been assigned to sponsor the work. this SOCIETY Let us look a t some of the outstanding things that have been accomplished on the theoretical side of this question. Speller in his new book on the subject, will list the following theories in addition to the electrochemical theory: (1) the acid theory, (2) the colloidal theory, (3) the theory of direct chemical attack by oxygen, (4) the biological theory, (5) the film protection theory, (6) the peroxide theory; and will point out that none of these have given even an approximately general explanation of observed facts. Evans includes even the earlier forms of the electrochemical theory in this category. It is well known that until recently some of the most common and striking phenomena, such as the formation of protective films in some cases and autocatalytic films in others and the prevalence of pitting of even pure metals under certain conditions, were unexplainable by any known theories. The older electrochemical theory did not explain the cause of pitting. In 1916, J. Aston showed that an electrolytic cell could be set up by bubbling air around one of a pair of submerged iron electrodes and not around the other. The aerated electrode was the cathode, and this showed that the reason wet rust promoted further rusting was not that it acted as a cathodic contact material, but that it acted as a diaphragm, screening the underlying metal from the direct access of oxygen. I n 1922, McKay iirst applied the well-known theory of concentration cells to corrosion, and pointed-out that “corrosion may be produced by cells due to differences in the concentration of acids, concentration of dissolved oxygen or hydrogen, concentration of dissolved oxidizing or reducing salts,” in addition to the copper-ion concentration cells which were the subject of his study. This probably furnishes the theoretical key explaining Aston’s results. Closely following this, Evans elaborated on the effect of “oxygen-variation” cells in the corrosion of iron and zinc, and showed their striking properties in many interesting laboratory experiments. He has pointed out how these cells explain “the local and apparently capricious nature of corrosion.”
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In 1923, Pilling and Bedworth published a paper on the oxidation of metals in air a t high temperatures, which was supplemented by the work of Tammann upon the action of halogens on metals a t low temperatures. Both of these papers helped to explain the formation of films and described the different varieties and the effect of each on corrosion. In this work the microscope has played an important part, and such work as that by Desch and Whyte on the microchemistry of the corrosion of brass is very valuable. In 1924, Bancroft prepared a report for the Corrosion Committee of the National Research Council, which describes the electrolytic theory and comments on the present status. He quotes liberally from all the workers on the subject so that this work is really a reference book. Bengough and Stuart, in seven long reports to the Corrosion Committee of the British Institute of Metals, have presented an exhaustive study, principally of the corrosion of condenser tubes, and have commented on the theories involved. They have emphasized the value of the ‘‘newer electrochemical theory.” Whitman and Russell have shown important quantitative relation between the supply of oxygen to steel surfaces and the rate of corrosion, and McKay and Thompson have demonstrated the same relations on nonferrous alloys. The concensus of opinion seems to be in favor of the electrolytic theory of corrosion as amplified by the later workers, and there is sufficient agreement now so that a program to carry the work further should be a distinct possibility. As was stated earlier, the AMERICAN CHEMICAL SOCIETY has an important part to play, and it is hoped that its committee will further this plan and act to play this part in a manner most creditable to the SOCIET?.
Corrosion in Aqueous Solutions By Wilder D. Bancroft CORNELL UNIVERSITY, ITHACA, N. Y.
The important points brought out in corrosion study during approximately the last twenty-five years are discussed in their relation to the present status of corrosion theory. It is shown that all corrosion reactions are electrochemical. The agreement of these theoretical points with experiments in the author’s laboratory is discussed. The importance of further research along such linese. g., the rate of oxygen supply to the metal, the development of a useful accelerated test, the formation and effect of films, and the other matters affecting electrolytic corrosion-is pointed out.
T
H E most striking characteristics of an electrolytic reaction1 is that it occurs a t two places-at the anode and the cathode. This peculiarity can be made less marked by bringing the electrodes nearer and nearer together. When the distance between them vanishes, we usually speak of the reaction as a chemical one and not as an electrochemical one. This is not a sound distinction. Any chemical reaction which can be split into an anode part and a cathode part2 may be an electrochemical reaction. When zinc dissolves in hydrochloric acid, the formation of zinc chloride is the anode reaction and the evolution of hydrogen the cathode reaction. When copper dissolves in dilute nitric acid, the formation of copper nitrate is the anode reaction and the reduction of nitric acid a t a copper cathode is the cathode reaction. 1 2
Bancroft, T r a m . A m . Electrochem. Soc., 9, 13 (1906). Traube, Ber,, 26, 1473 (1893); Haber. Z.fihrsik. Chem., 34,514 (1900).
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When a metal corrodes in an aqueous solution in presence of oxygen, there are two theoretical possibilities. The metal may react direct with oxygen, in which case it is not an electrochemical reaction, or it may react direct with the water or with some electrolyte, in which case it is an electrochemical reaction. It has been shown that iron does not rust a t ordinary temperatures in presence of air containing water vapor,3 provided no liquid water is formed, and that it does rust if there is a condensation of water on the iron either permanently or intermittently. It has also been shown‘ that dry oxygen reacts very slowly with iron even a t 200” C. The products of corrosion in aqueous solution may be quite different from those obtained on direct oxidation, as when sodium goes to sodium peroxide. From all these facts we are justified in believing that a metal does not ordinarily react direct with oxygen in presence of water5 or an aqueous solution and that the corrosion under these conditions is always electrolytic in nature. This was first pointed out by Whitney6 for the case of iron, and it is to him that we owe what is known as the electrolytic theory of corrosion.
tion of the two methods, and has thereby reduced the corrosion in such systems to a minimum. There is some hydrogen involved and consequently some corrosion, but nothing to what it would be in case oxygen were present. Since the first reduction product of oxygen a t a cathode is hydrogen peroxide,IO this substance should be formed a t least temporarily during the corrosion of metals in contact with water and air. Schonbeinll showed that, when lead amalgam is shaken with dilute sulfuric acid and air, hydrogen peroxide is formed in an amount equivalent to the corroded lead. Traube obtained similar results with zinc and water, while Dunstan, Jowett, and Goulding showed the formation of hydrogen peroxide-though not quantitatively-during the corrosion of silver, mercury, copper, bismuth, lead, tin, and zinc. KOhydrogen peroxide was detected during the corrosion of iron, but they point out that this is because it is decomposed as fast as formed and that hydrogen peroxide is probably formed as an intermediate product in the rusting of iron.
Effect of Potential Difference
The depolarizing action of oxgyen increases with the concentration of the gas and, consequently, reducing its concentration to nothing will prevent corrosion so far as this is dependent on oxygen as a depolarizer. This principle has been made use of by Speller,Qwho has removed practically all the oxygen and carbon dioxide from a hot-water supply system, either by a vacuum process, by chemical reaction, or by a combina-
If we have a two-phase alloy this mill tend to act as a voltaic cell and to accelerate corrosion. Cushman12 seems to have been the first to emphasize this point. Of course, the fact was known for special cases long before. Faraday observed that an alloy of steel with one per cent of platinum dissolved with effervescence in dilute sulfuric acid so weak that it hardly acted on common ~ t e e 1 . I ~It is common knowledge that platinum must be removed from the sulfuric acid used with storage battery plates because any precipitation of platinum on the lead plate causes self-discharge of the latter. While the theoretical tendency to corrosion is greater in a two-phase alloy than in a homogeneous one, the actual corrosion may not be as great because of the formation of surface films. Certain of the copper-tin alloys corrode much less readily in sodium sulfate solutions14 than does pure copper or pure tin. Experiments by Vincent15 show that the addition of platinum to copper does not cause the latter to corrode appreciably in a saturated solution of potassium bichromate. Cushman’s generalization is therefore not necessarily true when one is dealing with finite amounts of corrosion, though it does hold in a great many cases. Since differences of homogeneity will tend to cause local voltaic cells and will therefore tend to cause corrosion, the natural corollary is that the most homogeneous metal will corrode the least rapidly if other things are equal. This is absolutely sound; but some people then assumed, perhaps unconsciously, that other things are equal and that therefore the most homogeneous metal will corrode the least rapidly. This assumption is quite unwarranted by the facts, and yet people have gone farther than this, Bengough and Stuart16 declaring that, according to the electrplytic theory of corrosion, an ideally pure and homogeneous metal is assumed to be incorrodible. While some reckless upholders of the electrolytic theory of corrosion have said things like this, no chemist would claim that a pure and theoretically uniform piece pf sodium would not be attacked by water. In like manner every chemist would admit that copper would be corroded by concentrated nitric acid containing nitrous acid, regardless of how pure or how uniform the copper was.
* Dunstan, Jowett, and Goulding, J. Chcm. Soc. (London), 81, 1554 (1905). Pilling and Bedworth, J.Insf Mefals, 29, 529 (1923). Dry oxidation is differentiated sharply from wet oxidation by Haber, Z . Eleklrochcm., I , 445 (1900). J . A m . Chem. SOC.,26, 394 (1903). 7 Ibsd., 26, 399 (1903). 8 Evans, J . Inst. Mefals, SO, 254 (1923). Trans A m . Elecfrochem. Soc., 99, 141 (19211, J. Franklrn I n s f . , 193, 515 (1922).
Traube, Ber., 16, 2434 (1882). Mellor, “Treatise on Inorganic Chemistry,” 1’01. I, 1922, p. 926. I* U.S. De#l. Agr., Farmers’ Bull. 259 (1905); Trans. A m . Electrochem. SOC.,12, 403 (1907); Walker, Cederholm, and Bent, J. A m . Chem. SOL.,29, 1251 (1907). 18 Silliman, “Elements of Chemistry,” Vol. 2, 1831, p. 128. Bancroft, Trans. A m . Elecfrochem.Soc., 9, 19 (1906); Curry, J. Phyr. Chcm., 10, 484 (1906). 11 Unpublished. 1 6 J. I n s f . Metals, as, 54 (1922).
If a metal is more noble than hydrogen under the conditions of the experiment, there is no reason why it should corrode with the evolution of hydrogen. In order to get corrosion we must have present some depolarizer, usually oxygen. Metallic copper does not corrode appreciably in dilute sulfuric acid in the absence of air but does when air is present, the oxygen acting as a depolarizer for the nascent hydrogen which tends to form. Copper does corrode in nitric acid in the absence of air, because nitric acid (or nitric acid in presence of nitrous acid) acts as a depolarizer. If a metal is less noble than hydrogen under the conditions of the experiment, it will tend to corrode, with evolution of hydrogen from the solution. Such a corrosion and such an evolution of hydrogen actually occur when sodium is brought in contact with water or a salt solution in the absence of air, but the corrosion is, or may be, almost negligible with pure zinc or pure iron in water. This is due to what we call overvoltage, the hydrogen not being set free reversibly a t all cathodes. It is consequently necessary to use depolarizers even in the case of metals which are theoretically less noble than hydrogen. Since the potential difference necessary to set free hydrogen decreases with increasing concentation of hydrogen ion, metals should tend to corrode more readily as the concentration of acid i n ~ r e a s e sand , ~ this actually happens. With cadmium in hydrochloric acid the effect of the overvoltage is sufficient to prevent the corrosion* in normal hydrochloric acid when air is absent. Depolarizing Action of Oxygen
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Corrosion of Alloys
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INDUSTRIAL A N D ENGINEERING CHEMISTRY Factors Affecting Production of Oxide Film
Alkaline solutions inhibit the rusting of iron and this is usually attributed to the low concentration of hydrogen ion; but the real reason is that an oxide film forms on the iron, which prevents further rusting. Copper, zinc, and iron do not corrode in a dilute (M/40) potassium bichromate solution, because they become passive owing to the formation of an anode film. Iron remains passive with higher current densities but the films on copper and zinc break down when the voltage is raised too much. It has been suggested that potassium bichromate is an oxidizing agent and should therefore accelerate corrosion, but Vincent has found that there is 100 per cent evolution of hydrogen from a saturated bichromate solution a t a platinum cathode. With a mercury cathode the overvoltage is enough to cause reduction of the bichromate with no evolution of hydrogen. Magnesium is usually cited against the electrolytic theory of corrosion because it corrodes more rapidly in a magnesium chloride solution than in distilled water; but Miss Souders’7 has shown that this is because a coherent and somewhat protecting film forms in distilled water, while this film is peptized to some extent by magnesium chloride. Walker18 attributes the resistance of copper-bearing steel to atmospheric corrosion to the adherent nature of this film of oxide produced when corrosion starts. The lead peroxide in a storage battery grid must protect the lead of the grid from the sulfuric acid, because otherwise the plate would become a t once a short-circuited cell, which is not the case. Mill scale is another case in which we must have a pretty complete covering of the surface because it protects the iron from corrosion, although, like the lead peroxide, it accelerates corrosion when the film is broken.lg Attention seems to have been focused on the disadvantages of a break in a protecting coating of this type, and people have rather overlooked the possibility of a self-healing coating such as we actually have on aluminium and on nickel. The new stainless steel owes its properties to a coating of this type. Highly polished surfaces show a surprising resistance to corrosion. Whether one considers the polished surface to,consist of very fine crystals, or of an amorphous layer, as Beilby does, it should be anode against a normal surface and should therefore corrode more readily if everything else were equal. I n this case one must assume, as Bengough and Stuart20do in another case, that a more coherent film forms on a very smooth surface. While oxygen usually acts as a depolarizing agent, Bengough and Stuart21find that, when a jet of water or of salt solution, carrying entangled air bubbles, impinges on copper immersed in water or in a solution, the copper corrodes most on the side where the air bubbles strike, though this side should apparently be the cathode. Vieweg showed that a similar result could be‘ obtained by bubbling hydrogen or nitrogen against the piece of copper. Since the phenomenon is independent of the nature of the gas it is clear that we are dealing with frictional electrification.22 M c K has ~ ~shown ~ ~that with flowing and stagnant liquids the flowing liquid may be the anode because of an ordinary concentration cell, t h e concentration of the ions of the metal being low in the flowing liquid. I n the case of straight depolarization, the nonaerated portions become the an0des.~4 This is the explanation for the Bancroft, J . Phys. Chem., 28, 834 (1924). Trans. A m . Electrochem. SOC., 29, 438 (1916). 18 Walker, I b i d . , 14, 181 (1908). Po J . Inst. Metals, 28, 95,107 (1922). 2 1 I b i d . , 28, 66, 125 (1922). 22 Bancroft, J. Phys. Chem., 28, 843 (1924). $8 Trans. A m . Electrochem. Soc., 41, 201 (1922). 24 Aston, I b i d . , 29, 458 (1916); Evans, “The Corrosion of Metals,” 1924, p. 80. 17 18
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experiments by Bengough and in which they find that “any selected portions of a metal can be caused to suffer heavy local corrosion if the conditions external to the metal are suitably controlled. A simple way of showing this is to tie a piece of ordinary string round a piece of copper or brass and immerse the whole in sea-water. Active local corrosion will take place beneath the string in spite of the fact that access of oxygen to the corroded area is apparently greatly lessened. Local corrosion a t any selected spot can also be produced beneath cotton-wool, coke, glass fragments (if not in too fine a state of division), paraffin wax (whenever liquid can penetrate beneath the wax), and many other bodies. No such action takes place, however, beneath a deposit of red lead, which is an excellent depolarizer.” Wilson26points out that the rate of corrosion will be proportional to the rate of diffusion of oxygen to the metal in all cases where depolarization is the important factor, and this is the case within most of the ordinary range of corrosion in natural waters. Since we know that the rate of reaction between copper and dissolved ferric salts is practically instant a n e o u ~it, ~is~evident that the rate of diffusion or the rate of stirring will be the important one. Investigations which ignore this point are necessarily faulty. Problems for Further Study
Since an ordinary corrosion test may easily take months, it is evident that an accelerated test is very desirable. The first thing to do is to determine the voltage-current curve for the metal in the solution to be studied, measuring the anode potential, so as to see whether, and under what conditions, the metal becomes passive. We can then probably make our accelerated test an electrolytic one, though we shall have first to determine how high a current density is permissible, where the cathode should be placed, whether and when stirring is permissible, how often the solution should be changed, etc. An accelerated test, electrolytic or otherwise, must of course give correct results in every case to which it is properly applicable. In addition to the development of a satisfactory accelerated test, we have the still more important task of studying the properties and conditions of formation of surface films, because the surface film is the key to the whole problem of corrosion. 26 26
27
J . Inst. Metals, 28, 6 6 , 90,98 (1922). THISJOURNAL, 16, 127 (1923). Schluederberg, J . Phys. Chem., 12, 574 (1908).
The Passing of Sir Edward Thorpe The death of Sir Edward Thorpe on February 23 came as a distinct shock to his many admirers in this country, and removes another great personality from the field of chemistry. He was born in Manchester, England, seventy-nine years ago. H e studied a t Heidelberg, under Bunsen, and a t Bonn, and served for a time as demonstrator in chemistry under Sir Henry Roscoe. In 1870 he became professor of chemistry a t the Andersonian Institution, Glasgow, and four years later, when the Yorkshire College was established a t Leeds, accepted a similar post there. For ten years he was professor of chemistry in the Royal College of Science, South Kensington, London. He then became principal of the Government Laboratories, which position he held for sixteen years. After this he returned to South Kensington as professor of general chemistry and director of the Chemical Laboratories of the Imperial College of Science. He held many offices of honor and distinction, among them being president of the Society of Chemical Industry in 1895, president of the Chemical Society 1899-1901, and president of the British Association a t Edinburgh in 1921. Aside from his original investigations of chemical problems, he was an entertaining lecturer and contributed much to our chemical literature. Perhaps his most outstanding work is his Dictionary of Applied Chemistry. We are fortunate to have as a lasting memorial to him a new edition of this book which is now on the press.
