January, 1925
INDUSTRIAL A N D ENGINEERING CHEMISTRY
might be used. The resin-bearing coal was ground in a ball mill with a brine solution of 1.2 specific gravity, and upon settling the resin was found floating on the surface of the brine, but quantitative separations were not undertaken. COMM~RCIAL ASPECTSOB THE Rmm-Being insoluble in
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alcohol, the resin cannot be made into “spirit varnishes,” but a concentrated petroleum ether solution acts in a similar way, producing a hard, smooth, transparent, yellow coating, which is somewhat brittle. By melting the resin it dissolves in linseed oil and gives a fairly hard, transparent surface.
The Common Occurrence of Corrosion by Electrolyte Concentration Cells’ By Robert J. McKay THEINTERNATIONAL NICKEL Co.,NEWYORK,N . Y
M AN earlier paper2 at-
It should be noted that The results of recent corrosion study in the writer’s this definition of electrolyte tention was directed to laboratory, and in others, indicate widespread corrosive concentration cells takes in the accelerated corroaction by electrolyte concentration cells. Consideration a greater range than the sion of copper-bearing alloys of these cells shows, for instance, how pitting may occur concentration cells comby electrolytic cells whose on the most homogeneous metals, and that, contrary to monly used in the demone. m. f. is due only to differwidely accepted views, the theoretical considerations stration laboratory. In the ences in the concentration would lead one to expect pitting, rather than uniform corlatter the osmotic pressure of the copper ion in the rosion, under a broad range of common, well-defined concorrodingsolution. At that of attainable concentrations ditions. Pits should be expected in all cases where any time no reference could be is large in comparison with solid or colloidal product can remain as a film or mass in found in the literature on the solution pressure of the close proximity to the metal. Their location is detercorrosion to the action of electrode, the electrode is mined by conditions in the solution, rather than in the such cells, and therefore the same as the ion whose metal. For the explanation of the phenomena, the reathe theoretical discussion concentration is varied, and soning of electrochemistry is necessary, but the original was confined to this particthe measurablee. m. f. agrees concept of the “electrolytic theory of corrosion” is insufular case, though qualitawith that calculated by ficient. tive tests indicated widethe Newst formula. These simde t m e s are selected smead distribution. The purpose of the present paper is to demonstrate that for demonstration purposes. In practice; ‘we must ordinacorrosion by such concentration cells is of common occur- rily consider, not these simple cells, but combinations in rence, rather than of the limited application shown in the which are salts of different elements from the electrodes and former paper. I n fact, it is pow believed that a great many, wherein the absolute value of the e. m. f. therefore depends on possibly all, of the so-called paradoxes and anomalies of cor- more complicated reactions. Such cells, of course, include rosion are readily and simply explainable by consideration the important class of “oxidation-reduction” cells. The of these cells. Many of the present inaccuracies in the e. m. f. of these cells is a function of the difference in concensolution of corrosion problems will be avoided only when tration4and disappears when the concentrations are equalized. it is generally understood that such cells are as common and The concept which is considered new is that of corrosion by cells caused by differences in concentration of solution as as important as the galvanic or voltaic cell. Electrolytic corrosion, as generally understood by engi- distinct from differences in the metal. The great changes in the single potential of electrodes of neers, is corrosion due primarily to an e. m. f . set up between two electrodes of different materials in the same electrolyte. the same metal immersed in different solutions were pointed ~ data show that the single potentials in This was the concept of the “electrolytic theory” of corrosion3 out by B a ~ e r . His and there is abundance of proof that corrosion is caused by salt solutions containing oxygen are much more valuable for such galvanic cells. However, the attempt to explain all the predicting corrosion than the potential as usually determined phenomena of various types of corrosion on this basis has against a normal solution of the salt. But, in his interesting data on the localization of corrosion by contact between two been unsuccessful. Electrolyte concentration cells may be defined as cells different metals, he failed to note the relatively great effect of whose e. m. f. is set up by two electrodes of the same material contact of the same metal with two concentrations of solution. Probably the most active. and widespread corroding agent in different electrolytes. Lewis4 defines a concentration cell as “two similar electrodes dipping into solutions of the same is oxygen and no corrosion test should be made, or practical salt, the solutions being at different concentrations of the corrosion problem answered, without the most careful considsalt. The source of e. m. f . is to be found in the tendency of eration and thorough knowledge of the oxygen content of the two solutions to equalize their concentrations.” This the corroding solution and its supply to the corroded surface.61’ latter type of cell is at least as prevalent as the former, and Usually, oxygen in aqueous solution is peculiarly adapted to as powerful. Theoretically, then, it should be as active in the setting up of concentration cells, owing to its low maxiproducing corrosion, and recent experiments in the writer’s mum concentration under atmospheric pressure and the laborattory, and published results by others, have shown this resulting high percentage concentration changes with slight to be true. corrosion. The author has measured the e. m. f. of monel
I
1
a 8 4
Received Auaust 9, 1924. McKay, Trans. A m . Electrochem. SOC.,41,201(1922). Whitney, J . A m . Chcm. SOC.,26, 294 (1903). “A System of Physical Chemistry,” 1916. Longmans, Green & Co.
6 0. Bauer, Mitt. kgl. Matmial-pr&fununssamt,36, 114 (1918); 2. Mdadlkunde, 10, 124 (1919). 0 Thompson and McKay, THIS JOURNAL, 16,1114(1923). 1 Whitman, Russell, Welling, and Cochrane, Ibid., 15,672 (1923).
I N D U S T R I A L A N D ENGINEERING C H E M I S T R Y
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metal electrodes in solutions of varying oxygen concentration.2 Evans has also shown8 qualitatively that such cells are active under a different set of conditions, although he neglected the probably important changes in salt concentration at the water line, and has showng that corrosion of iron by oxygen occurs a t some other point than that of maximum oxygen concentration. The latter phenomenon was quantitatively studied by Astonlo and useful practical conclusionswere reached. The exact mechanism was not discussed, but the action was apparently attributed to some change in the surface of the metal. Action of Oxygen Concentration Cells
The corrosion of iron by oxygen-bearing solutions is of very great economic importance. The prevalence of pitting in the corrosion of even the purest iron has not been satisfactorily explained by the galvanic or voltaic cell, although it has been given much thought on account of its economic importance. This pitting is an example of the phenomena which seem to be completely explained by consideration of the action of oxygen concentration cells. These cells are a type of the general class of concentration cells, and their action in causing pitting is discussed below. When iron comes in contact with oxygen-bearing solution it reacts with the oxygen and water, forming ferric hydroxide. The precipitate tends to flocculate, or may congregate in one point by settling or be distributed by currents in the solution. Where a mass of precipitate or other solid material is held by any chance in contact with the iron, the oxygen in the solution in the pores of the mass near the iron immediately disappears in reacting with the iron. In the region around this mass, the oxygen content is renewed by convection and diffusion, and we have a concentration cell set up with a low concentration of oxygen in the mass of precipitate and a higher concentration in its neighborhood. The metal in contact with the higher concentration becomes the cathode, and that in contact with the low oxygen concentration in the precipitate, the anode. Iron goes into solution under the precipitate and the action is autocatalytic, as the longer it goes on the thicker the precipitated mass becomes and the more efficient as a diaphragm. A pit forms and deepens indefinitely. I n the later stages, precipitation occurs as ferrous hydroxide and there is always a layer of ferrous hydroxide in contact with the metal. Quantitative results, showing the conditions existing in these circumstances, have been given by Whitman, Russell, and Altieri.ll The reaction by which iron dissolves divides itself into three parts, which take place at different points, as follows:
++
--.+
+ --.+ + + 2 ( - )
At cathodic surface:z '6H/ 1 1 3Ha0 a(+) At anodic surface: 2Fe +4(0H) --f 2Fe(OH)n 4(-) In corrosion product: 2Fe(OH)t 2(OH) ZFe(0H)r 3Hs0
+ + ll/aOz + 2Fe ---f ZFe(0H)a
(1) (2) (3) (4)
These reactions assume the reaction of dissolved oxygen directly with the hydrogen ion of water without formation of hydrogen gas. Whether or not hydrogen gas actually forms and the action of the oxygen is a depolarization, does not affect the final conclusions. The mechanism in this case is especially adapted to cause pitting. The precipitate of iron hydroxides forms a physical separation between the two portions of this concentration cell at its liquid-liquid contact. This action is fundamental in localizing the corrosion. The formation of ferric hydroxide also operates chemically, in two distinct ways, to accelerate the action. First, oxygen is removed by its reaction with ferrous hydroxide; and second, iron is removed by precipitaa J . Insi. Metals, 1928, advance proof. .e Chem. Me!. Eng., 30, 949 (1924). 10 Trans. A m . Elscirochem. Soc., 29,449 (19113). 11 Tnrs JOURNAL, 16, 66G (1924).
Vol. 17, No. 1
tion, which, if not removed, would retard the reaction by the mass law and would set up a counter e. m. f. Effect of Corroding Film
This action of iron is a specific case of a quite general phenomenon. Whenever corrosion takes place, the corroding film changes chemically, tending towards exhaustion of the corroding constituent. At the point where corrosion starts any solid or colloidal product of corrosion concentrates and tends to prevent the replenishing of the corroding medium. In neighboring areas, where for any reason the corrosion has been less active, the corroding medium has better access to the metal, and the differences in concentration in the contact film set up an electric cell, with the protected portion the anode and that open to corroding solution the cathode. Any point over which the amount of solid corrosion product begins to concentrate thus becomes anodic and becomes more so as corrosion preceeds. The amount of corrosion depends upon the supply of corroding agent to the cathodic areas, but most of the metal disintegration occurs at the anodic areas, which thus become pits. The action of the current is in the direction described because the corroding reaction must be such as to furnish plus charges to the metal, and the areas over which a greatter concentration of available plus charges are present are cathodic whereas those over which lesser concentrations are present must be anodic. This effect is more important than any differences in potential existing between different parts of a single metal of ordinary purity. Illustrations of Theory
The conditions outlined are easily proved experimentally in the natural water corrosion of iron. It is well known that the formation of iron rust will stimulate corrosion a t the point where rust exists. The writer has experimentally proved the conditions to be analagous in the case of corrosion of nickel alloys by alkaline oxidizing agents, such as strong bleaching agents. Here solid nickelous and nickelic hydrates are formed and all the corrosion is concentrated under films of hydrate. Pitting may be started at predetermined points by properly inoculating them with hydrate. I n addition to this general tendency towards pitting, there is an opposite one in which concentration cells may localize corrosion by a somewhat different mechanism. For instance, in the corrosion of copper by acid solutions containing dissolved oxygen,*the dissolved copper sets up a counter e. m. f., which is found by practical experiment to be greater than the oxidizing or original e. m. f. and the direction of current will usually be determined by the concentration of copper ion. This is in accord with the foregoing reasoning, the copper ion becoming the active corroding agent. In fact, an analysis of the mechanism will show that the difference between the two cases is largely one of the solubility of the corrosion product. The natural water corrosion of iron is the best example of the effect of concentration differences in causing pitting. The considerations involved are of general application, but probably are at their maximum importance in this case. The mistake of deducing positive general conclusions from one specific instance should not be made. Thus, iron may be corroded by water moving at such high velocity that rust films do not form, or by acids which dissolve them. There are many cases of nonferrous corrosion where corrosion products cannot precipitate. I n these cases the foregoing reasoning obviously cannot apply. In such cases, corrosion may be accelerated by other types of concentration cells or by galvanic cells, or the anodic and cathodic parts of the reaction may take place so close to each other as to make consideration of the electric current produced of no practical value.
