I S D VSTRIAL A S D E,VGINEERISG CHEMISTRY
April, 1925
iron type. With carbon contents that are reasonably constant and pouring temperatures that are sufficiently consistent to eliminate differences in structure, this curve explains the role of chemical composition in the corrosion of this particular alloy. It also serves as a convenient check in metallurgical control work, as it will indicate in advance of the time necessary to run a complete chemical analysis, even though the silicon content is up to standard, whether other impurities are present in an amount seriously to affect the resistance of the alloy to the standard sulfuric acid solution. General Deduction
If this paper, and the discussion which it brings out, is useful in establishing the value of accelerated laboratory corrosion tests made under proper conditions, and indicates a method by which it is possible to secure data to interpret the results of the short-time tests into a fairly accurate prediction of the life of apparatus in commercial service, it is then hoped that a standard method will be adopted so that all manufacturers of acid-proof alloys can publish data secured in the same manner which will be satisfactory to the prospective user of equipment for m y particular corrosive conditions. The assistance of is very earnestly solicited in arriving a t a suitable this SOCIETY standard for reporting laboratory corrosion tests.
Corrosion of Iron in Absence of Oxygen By J. W. Shipley, I. R. McHaffie, and N. D. Clare UNIVERSITY OF
MANITOBA,
WINNIPEG,CANADA
The corrosion of iron in the absence of oxygen is proportional to the hydrogen-ion concentration down to a pH of 9.4, when hydrogen evolution and solution of iron ceases. The rate of corrosion in the absence of oxygen is determined by the hydrogen-ion concentration, and the continuance of the corrosion by the total available acidity. If the solution in contact with the iron is not buffered, corrosion will continue, but at a diminishing rate and diminishing hydrogen-ion concentration until a pH of 9.4 is reached. Oxidation of iron requires a potential a little less than the potential required for the evolution of hydrogen. Consequently, in the presence of oxygen, hydrogen will not be evolved at a H-ion concentration corresponding to the solubility of ferric hydroxide or lower. The submerged corrosion of iron imbedded in impervious clay is due (1) to the buffer action of the clay at a Hion concentration sufficiently acidic in the absence of oxygen to produce the evolution of hydrogen, and (2) to the lowering of hydrogen overvoltage by the thin film of solution on the extremely fine subdivision of the clay in contact with the metal. The location of pits is determined by the constituents of the iron. When oxygen is present the migration of electrically charged colloidal particles of ferric hydroxide to cathodic areas on the metal perpetuates and enlarges these cathodic areas. The metal underneath becomes anodic. If the noncorrodible constituents are sufficiently cohesive, as is the case in gray cast iron, graphitic softening rather than pitting is produced, and no holing of the metal results.
UCH of the recent investigation into the mechanism of the corrosion of iron has been directed towards the effect of oxygen upon the course of the corrosion. This has been particularly the case in the laboratories of the National Tube Company and the Massachusetts Insti-
M
381
tute of Technology, where, under the direction of F. N. Speller and W. G. Whitman, many valuable observations have added much to our knowledge of the role played by oxygen in the destruction of submerged iron. Evans,' by direct observation of the behavior of drops of salt solutions on the surface of polished steel, showed the formation of cathodic areas in the peripheries of the drops where ferric hydroxide formed, while within the periphery of adhering hydroxide, corrosion proceeded from an anodic surface. McKay2gives some very interesting observations regarding corrosion by electrolyte concentration cells, in which agitation of the solution determined the setting-up of anodic areas on otherwise uniform pieces of metal. He also describes experiments showing greatly accelerated corrosion of homogeneous metals in acid solution due to differences in concentrations of dissolved materials in the solution. A more recent paper by the same author3 stresses the importance of electrolyte concentration cells, and describes how these may operate in the local cell corrosion of iron. Two of the present authors, in a report of the Advisory Research Council, Ottawa, October, 1923,* found: (a) In the absence of oxygen the corrosion of iron was proportional to the hydrogen-ion concentration. Below a pH of 9.4 iron would not free hydrogen from solution. ( b ) The concentration pH 9.4 corresponded to that produced by the solubility of ferrous hydroxide. Oxygen in solution produced ferric hydroxide, and the in(c) solubility of this compound, as compared with ferrous hydroxide, increased the hydrogen-ion concentration to a pH of 7.0, buffering the solution a t this concentration and thus stimulating corrosion. (d) A layer or film of ferrous hydroxide was first formed in contact with the iron, but this was dissolved by the more acid solution in contact with ferric hydroxide produced by the inward diffusing oxygen.
The usual reactions concerned in the corrosion of iron are the formation of a corrosion cell with solution of iron a t the anode and either the production of hydrogen or (if oxygen be present) the oxidation of the iron and the formation of hydroxyl a t the cathode. The cathode may be any substance in contact with the iron or even a constituent of the metal itself. The operation of a corrosion cell requires iron to be in contact with a more electronegative substance and both to be immersed in an electrolyte. The rate of corrosion is determined by the difference in potential and the electrolytic resistance. Speller, Wilson, and Whitmans have made a detailed study of the rate of corrosion when the electrolyte is kept circulating through a pipe. The corrosion cells are dealt with en masse and the damage to the pipe is calculated from the oxygen absorbed. These observers assume the cathode areas to be the formation of films of the hydroxides of iron in contact with the metal, and conclude that the diffusion of oxygen to the hydroxide film is the determining factor in the rate of corrosion. They have particularly stressed changes in the H-ion concentration produced in the solution in contact with the iron as influenced by the production of hydroxyl ions from diffusing oxygen. The mechanism of the corrosion as explained by Wilson and Whitman corresponds closely with that observed by Evans and is somewhat similar to what the writers have outlined above. The practical value of the rate of corrosion established by the oxygen absorption method of Speller is somewhat dubious, J . SOC.Chem. I n d . , 48, 316 (1924). THIS J O U R N A L , 16, 555 (1923). * I b i d . , 17, 23 (1925). 4 Shipley and McHaffie, Can. Chcm. Met., 8, 121 (1924). 1 Speller and Kendall, THIS JOURNAL, 16, 134 (1923); Speller and Texter, I b i d . , 16,393 (1924); Wilson, I b i d . , 16, 127 (1923); Whitman, Russell, and Altieri, I b i d . , 16, 665 (1924); Whitman and Russell, J. 5'06. Chcm. I n d . , I S , 193 (1924). 1
382
INDUSTRIAL AND ENGINEERING CHEMISTRY
for it does not take into account local pitting. Translating the total corrosion within 150 feet of piping into average specific penetration in inches per year per cubic centimeter of oxygen per liter, will not determine the specific penetration where a pit is forming below an accumulation of ferric hydroxide. The use of the term “protective film” is somewhat questionable, for in many cases such films are exactly the opposite, providing, as they do, a cathodic area producing with the iron a corrosion cell. The mechanism of corrosion, even with the reactions as simplified by Wilson, Whitman, and others, is still in its application to specific conditions extremely complicated. The accumulation on cathodic areas of ferric hydroxide, the changing rate of diffusion of oxygen to the cathodic areas, and the varying H-ion concentration in contact with the metal are but some of the variables omnipresent in ordinary corrosion cell action. Corrosion in the Absence of Oxygen
The removal of oxygen as a factor in corrosion reduces the initial corrosion cell reaction to the solution of iron at the anode and the evolution of hydrogen a t the cathode. There are natural conditions6 under which corrosion proceeds where oxygen cannot be the controlling factor. Such is the case where iron is buried in an impervious clay through which the soil water has little, if any, movement, either vertically or laterally. One of the writers’ has investigated such clays where the ground water level differs by several feet within a few yards, and maintains this difference throughout the season. The direct evidence that oxygen is not the controlling factor in the corrosion of pipes buried in such soils is twofold. On exhumation the soil for several inches from the pipe is found to be impregnated with bluish black hydrated oxides of iron. I n the second place, the products of corrosion remaining in situ, as in the graphitic softening of cast iron, were sometimes found to contain no ferric iron. Interest in the problem of corrosion in the absence of oxygen was particularly stimulated by observing the continuous evolution of hydrogen over a period of several months from a quantity of iron submerged in a wet clay, contained in a glass flask. The clay and the water had been degassed by vigorous boiling before sealing. Hydrogen was evolved continuously from the flask for 4 months, the characteristic bluish green hydroxide of iron appearing in ihe clay. The soil solution above the clay had a pH of 7.6 both before and after treatment, which is much below that a t which Whitman first observed the evolution of hydrogen when oxygen was present. The clay had a high lime content and was impregnated with gypsum. The continuance of the corrosion and the constancy of the pH suggested a buffer action on the part of the clay. Evolution of Hydrogen from Oxygen-Free Buffered Solutions
A preliminary series3 had indicated that the evolution of hydrogen from buffered solutions was a function of the hydrogen-ion concentration down to a pH of 9.4, the hydrogen-ion concentration found by the writers corresponding to a saturated solution with respect to ferrous hydroxide. At a lower concentration of hydrogen ions than this equilibrium value, the evolution of hydrogen apparently ceased. A series of buffered solutions was prepared and placed in flasks, together with 5 grams of 99.8 per cent No. 30 iron wire, exposing a surface of 112 sq. cm. The solution of pH 5 was prepared from potassium acid phthalate and potassium Webb. “The Corrosion of Yeovil Water Mains in Clay,”-paper
read
in 1922 before the Institution of Water Engineers, London, England, Medinger, J . Gasbel , 68, 73 (1918) Smith and Shipky, J. Eng. I n s l . C a n a d a , 40, 642 (1921); Shipley, J . SOC.Chcm. Znd., 41, 311 (1922).
Vol. 17, No. 4
chloride, while those from pH 6 to pH 9 were prepared from sodium phosphate and hydrochloric acid. The solutions were made oxygen-free by boiling under reduced pressure and the flasks were filled in an atmosphere of hydrogen. The evolution of hydrogen began a t once, with the exception of pH 6, which refused to act, and a duplicate was set up.
Original PH
Days observed
5
1s 39 144 116
6
7 8
TabIe I -GAS EVOLVEDTotal Per day Cc. cc.
47 82 20.2 19.0
Final PH
2.6 2.1 0.14 0.16
6.4 6.2 9.4 9.4
The change in the hydrogen-ion concentration indicated that the buffer action had not held, so a second series was prepared in which the concentration of the buffer solutions was much increased. The solutions were also prepared to overlap so that any salt effect might be detected. Table I1
-GAS EVOLVED-
No. 1
2 3
4 5 6
7 8
9 10
Original Days Total BUFFERSOLUTION D H observed Cc. KHPh NaOH 5.0 14 17 NaOH KHPh 6.0 45 20 KHzPOd 6.0 60 55 NaOH 6.0 45 50 NaOH KHzPOI 60 31 NaOH 7.0 KHzPOd 45 36 NaOH 8.0 KHePO4 HC1 HsBOa None NaOH 8.0 60 &BO8 f HCl f 60 None 9,0 NaOH . 4- HCI None 60 10.0 NaOH 60 None KHlPO4 NaOH 10.0
++ ++ + ++ +
+
+
Perday Cc.
Final PH
1.2 0.44 0.90 1.10 0.52 0.80
6.6 7.0 6.4 6.0 7.4 9.4
..
8.0
..
9.0
.. ..
10.0 10.0
The evolution of hydrogen was by no means regularNo. 3, for example, evolving 20 cc. in less than 24 hours and then slowing down until only 5 cc. were evolved in the last 20 days. The wire used in No. 6 had previously been immersed for a time in the buffered solution of pH 5, and then exposed to the air, before finally being placed in the flask with oxygen-free solution. This treatment apparently stimulated the evolution of hydrogen a t first, but finally only 4 cc. were evolved in the last 30 days. After 2 weeks the seal of flask No. 1 broke and admitted air, thus terminating this member of the series. Supplementary to the above series, several experiments were carried out to determine whether iron would evolve hydrogen in a solution of pH 10. Wire treated with hydrochloric acid, wire allowed to rust, and wire immersed in copper sulfate solution, each carefully washed without removing all the adhering rust or copper, gave no evolution of hydrogen after 60 days’ immersion in oxygen-free solutions buffered to a pH of 10. The solutions used were made up as Nos. 9 and 10, Table 11. The writers had several times observed that iron wire immersed in wet clay would continuously evolve hydrogen over a long period but that no hydrogen was evolved when the iron was immersed in the clay solution alone. Two explanations were probable: (1) that the soil buffered the solution at a pH of about 8 and thus kept the hydrogen concentration above the equilibrium value of 9.4, below which hydrogen is not evolved, and (2) that the extremely fine subdivision of the particles of clay in contact with the iron, exposing innumerable thin films of solution to the metal, lowered the potential a t which hydrogen would be evolved. Calcium sulfate and carbonate are characteristic constituents of the clay, so they were included in a series for observing the rate of evolution of hydrogen. Five grams of No. 30 iron wire were used in flasks filled and set up as previously described. The solutions were made oxygen-free in every case. The soil in solution No. 4 was treated for several hours in an autoclave a t a pressure of 1 atmosphere (15
INDUSTRIAL A N D ENGINEERING CHEMISTRY
April, 1925
pounds) and 121' C. to destroy bacteria. The soil water was prepared by frequently shaking distilled water with the clay, over a period of several weeks, and then,jafter settling, decanting without filtering. Table 111 No. SOLUTION 1 Satd. s o h . of Cas04 CaCO: 2 Soil water 3 Clay and soil water 4 Autoclaved clay soil water 5 NazHPO4 HCl Bas04 solid 6 NazHPO4 t HCI Ottawasand
pH 7.8 8.0 8.0 8.0 8.0 8.0
+
+
+
++
Days -GAS EVOLVEDevolving Total Per day gas Cc. Cc. 22 None 22 None 22 35 1.6 120 175 1,s 14 42 3.00 27 17.6 0.7
The barium sulfate was prepared in a very fine state of subdivision. The rate of evolution of gas was much greater where the finely divided barium sulfate was in contact with the iron than where coarse sand was used. Heyn and Bauer8 found that cast iron immersed in clay saturated with distilled water corroded almost twice as fast as when immersed in sand and distilled water. The buffer action of the clay was readily demonstrated by shaking up separate portions with dilute acid and alkali and, after settling, determining the p H colorimetrically. A considerable quantity of either acid or alkali was required to change the p H appreciably. The soil solution out of contact with the clay had no buffer action whatsoever and, it will be observed, did not generate hydrogen in contact with iron. It has been shown by two of the authorsg that the system CaSO,-CaC03-C02 and HzO has a strong buffer action a t a pH of 5.1. This is well within the concentration where hydrogen evolution accompanies the corrosion of iron, so that such a solution should continuously evolve hydrogen in the presence of iron. Three systems were set up:
+
++
A-Cas04 solid CaC03 solid COS saturated solution, and not freed from dissolved 0 2 5.000 grams No. 30 iron wire B-Same as A but freed from dissolved 0 2 C-Water saturated with COS and free from dissolved O2 5.000 grams No. 30 iron wire
+
A continuous supply of COZ, freed from 0 2 and under slight pressure, was provided each solution, and provision was made for collecting any evolved gases. System A did not evolve gas for a week but continuously absorbed C02. This was due to the removal of the hydrogen by the dissolved oxygen. Systems B and C, immediately after setting up, simultaneously absorbed COZ and evolved H2. The COZ evolved with the HPwas absorbed in caustic potash solution. Table I V
-HZ EVOLVED-
SOLIJTIOX Cas04 solid CaCOi solid CO?. mesent ., with 0% ~. = - ~ - ~ at first.. . , . . . . . , . . . ... Sameas.4 but no Oxpresent HzO COz
.A
+
B
+
~
~
+
C
..
Days Total pH evolving C c . 5.1 5.1
4.0
124 115 100
180 208 124
Per day Cc.
Final
1.45 1.80 1.24
6.4 6.4 6.4
pH
After the series was dismantled each solution had a p H of 6.4. This must correspond to the hydrogen-ion concentration of a solution saturated in respect to COz and ferrous carbonate. All three were actively generating hydrogen when the series was discontinued, thus indicating that no polarizing film was being produced. Effect of Aerating One Electrode
EvansJIO and Whitman and Russellx1 describe the effect of aerating one of the electrodes in an otherwise uniform solution. They found that the aerated electrode became
* M i l t . K g l . Mafnialprujungsoml, 18, 62 (1910). Shipleg and McHaffie, J. SOC.Chrm. I n d . , 42, 317 (1923). Ibid., IS, 316 (1924). I * I b i d . , I S , 193 (1924).
383
cathodic. The latter observers explain the phenomenon on the basis of hydroxide film formation, due to the action of the oxygen. Bancroft,12 however, explains the action on the basis of frictional electrification produced by the gas stream on the electrode. I n experiments in this laboratory, two glass tubes each containing 6.1 meters (20 feet) of coiled No. 30 iron wire connected to a potentiometer through a resistance of 100 ohms were filled with buffered oxygen-free solutions of the same pH and connected by a short glass tube filled with the same solution. Neither hydrogen nor coal gas freed from oxygen, when passed through the solution in one of the tubes, had any effect upon the potentiometer reading, but the moment air was passed in, an electric current was generated with the aerated electrode becoming cathodic. The rate of passing in the gases did not affect the voltages. The potentiometer readings were of the order 0.0005 to 0.0025 volt and the current continued to flow after the air current was stopped. Apparently, Bancroft's observations on the effect of hydrogen and coal gas do not cover the conditions under which the present authors worked. They also observed that with solutions having a p H below 9, hydrogen was evolved in the oxygen-free anode tube, showing that local corrosion was proceeding, while on the cathode a voluminous deposit of rust indicated a similar condition of local activity. Factors Controlling Corrosion
The corrosion of iron in oxygen-free solutions buffered a t H-ion concentrations above that corresponding to a p H 9.4 should not be accompanied by any film formation of hydroxides. The evolution of hydrogen is, then, a function of the H-ion concentration and is the characteristic corrosion reaction. Hydrogen will plate out on the more electronegative points or areas on the metal and the rate of evolution will be determined by the e. m. f. developed, the electrolytic conductivity, and the hydrogen overvoltage. The corrosion for the same metal will be much more uniformly distributed over the surface than where film formation occurs, and the penetration into the metal will be largely determined by the constituents of the iron. The appearance of the iron wire corroded in the buffered solutions, as in Tables I and 11, bears out this reasoning, as in every case where hydrogen was evolved the wire was uniformly covered with a iine, loosely adhering powder. The particles represent the local cathodic areas on the surface of the metal. The directive tendencies of the constituents of iron, particularly cast iron, have been studied micrographicaliy by two of the authors,' and the penetration of the corrosion was observed to be according to the e. m. f. existing between the constituents of the metal. The rate of evolution of the hydrogen will drop off for a buffered solution unless stirring is provided, since the liquid in contact with the iron cannot be kept a t the H-ion concentration of the surrounding solution. This was quite apparent in the writers' experiments, but the results were sufficiently quantitative to indicate the effect of H-ion concentration on the rate of the evolution of hydrogen. The electrochemical explanation of the increased evolution of hydrogen with acidity from a pH of 9.4 lies in the concentration effect of the hydrogen ion, the greater this concentration the greater the driving force or e. m. f . E = 0.058 In (Hi) The writers' failure to obtain the evolution of hydrogen from solutions more alkaline than corresponds to a p H of 9.4 is in agreement with the fact that pure iron will not continuously corrode in the presence of water alone. If iron is left in contact with water originally neutral-that is, a t
+
'0
lz
J. Phrr Chcm., 18, 842 (1924).
I S D U S T R I A L A N D ENGINEERING CHEMISTRY
384
Vol. 17, No. 4
points out, one-third of the total deposit may consist of ferrous iron. The oxidation of the Fe++ to Fe+++ is, it is believed, the controlling reaction in corrosion in the presence of ferric hydroxide and oxygen. The oxidation must take place within the cathodic film, since the outer layer is observed to be ferric hydroxide, while ferrous iron is diffusing from the metal surface outwards into the film. Wilson does not consider the oxidation of Fe++ to Fe+++ as exerting an important influence on the rate of corrosion, but considers it to affect only the character of the deposit formed. The writers cannot agree with this conclusion, for if the relative solubilities of the two hydroxides of iron (Fe++) AF = - R T In = -20,35013 are compared it will be seen that ferric hydroxide can exist (H+)2 (Fe++) (OH)*= 3.9 X 10-l6 indicate that iron should cease going into solution (corroding) (Fe+++) (OH)a = 1.1 X 1 0 - 8 8 a t a pH of 9.76. This value is very close to that a t which the writers found iron ceased to evolve any appreciable in a much more acid solution than ferrous hydroxide. Ferquantity of hydrogen. Wilson points out that the free rous hydroxide can exist in solutions having a pH of 9.4 or energy decrease for iron in an alkaline solution is conditioned higher, while ferric hydroxide will persist in solutions having by the ratio (H+)2:(Fe++) and that, after the solubility a pH of 7.0 or higher. As a consequence, the concentration product for ferrous hydroxide is reached, further increase of iron in solutions in contact with ferrous hydroxide may of (OH)- will decrease the H + as rapidly as the Fe++, thus be much higher than when in contact with ferric hydroxide preserving the ratio constant. This is quite true, but the and as a result concentration polarization of Fe+T will be H + a t which theoretically iron ceases to corrode, p H 9.76, higher in the case of ferrous hydroxide than in that of ferric and that a t which ferrous hydroxide first forms, p H 9.4, hydroxide, thus diminishing the rate of solution of the iron. are so close together as to be almost identical. The “natural water region” of corrosion, as described b y Wilson and Whitman as existing between a p H of about Film Formation 4 and 10 a t 22” C . , where the rate of corrosion is apparently Ferrous hydroxide forms, in the absence of oxygen, a t a independent of the H-ion concentration in the main solution, H-ion concentration corresponding to a pH of 9.4 and will may have its explanation in the controlling action of t h e persist a t all concentrations lower than this value. Ferric neglected reaction, the oxidation of Fe++ to Fe+T+. But hydroxide forms, in the presence of oxygen, a t a H-ion con- this reaction again is dependent upon the rate of diffusion centration corresponding to a p H of 7.0 and will persist a t of oxygen into the cathode area. Another very important consideration in respect to film all concentrations lower than this value. Because of the fact that iron going into solution electrolytically produces formation is the local character of its emplacement. Since ferrous ions, the solution next to the corroding metal will colloidal ferric hydroxide is probably the first form of the always contain ferrous iron and if the (OH)- is sufficiently hydroxide, and being positively charged, it will migrate high a precipitate of ferrous hydroxide will form. The towards a cathodic area and the first deposition of this subwriters know practically nothing about the electrochemical stance in completely submerged corrosion will be determined properties of this precipitate. I n the presence of oxygen, by local cathodic areas on the surface of the iron. This localhowever, Fe++ is transformed to Fe+++ and the precipitate ization of deposits of ferric hydroxide enhances the local now formed is ferric hydroxide, a substance electronegative corrosion cell e. m. f., and determines the formation of pits, to iron and much more insoluble than the ferrous hydroxide. for the iron surface below the deposit will be more strongly In the colloidal condition this hydroxide is positively charged anodic than adjacent areas. The writers do not believe and will migrate to the cathode. One of the authors’ has that the film first formed is closely adherent to the iron or observed this migration in a concentration cell on the surface that it is continuous, but they do believe that it must be of a plate of iron, the colloidal ferric hydroxide actually exceedingly tenuous and porous. This is due to the dilute passing through the connecting U-tube, and depositing on solution of Fe++ or Fe+++ and (OH)- ions forming first the colloid, and later, the coagulated hydroxide. the surface of the iron in the other division of the cell. The constitution of the film in the absence of oxygen and Protectiveness of Film Formation insoluble salts ie therefore ferrous hydroxide and, so far as the evidence goes, this compound does not appear to stimulate Speller and Texter, Whitman, Evans, and others stress corrosion but, if anything, inhibits it. When oxygen is the protective nature of deposits formed upon the surface present in solution, the outer layer of the film will consist of the metal. So far as ferric hydroxide is concerned the of ferric hydroxide, the inner, next to the corroding iron, protectiveness is somewhat doubtful, especially when the may be ferrous hydroxide, and the middle layers may consist process of pitting is considered. The writers’ observations of a mixture of ferrous and ferric hydroxides. If, however, have always associated deep pitting with an overlying deposit the H-ion concentration next to the surface of the iron is so of ferric hydroxide. This is particularly observable where high that the solubility product for ferrous hydroxide cannot hot water systems are perforated. They have found steam be attained, then no solid ferrous hydroxide will appear. pipes immersed in a hot water tank, in which aerated water That this condition of acidity ever arises in the presence of continually circulated, completely perforated from the outferric hydroxide and oxygen is doubtful, since, as Wilson side in 18 months’ service. The pits were very heavily enand Whitman point out, (OH)- is being continually formed crusted with ferric hydroxide. The water was lake water in the cathode area of the ferric hydroxide, according to the with a pH of 8.2. Again, when they permitted tap water reaction H 2 0 1/2 O2= 2(OH)- 2 0. Moreover, deposits to circulate slowly through a glass tube in which a length of ferric hydroxide over corroding areas are very frequently of No. 30 iron wire extended, areas covered with ferric found to be bluish black next to the iron and, as Wilson hydroxide soon formed, especially where the wire touched the glass, and in a few days the wire beneath these “pro18 Lewis, “Thermodynamics,” p. 607. a p H of 7.0-the water will, after some little time, have a p H of 9.4,with no apparent corrosion. Iron has gone into solution until the solubility product for the formation of ferrous hydroxide has been satisfied but,not enough hydrogen has been generated to saturate the solution. If a stream of oxygen-free water were permitted to flow over iron, the iron should corrode continuously. Wilson and Whitman have found corrosion to proceed beyond an alkalinity of pH 9.4;a t least, oxygen was consumed but the rate of absorption of oxygen dropped very rapidly after a pH of 9.5 was reached. Calculations based on the free energy of iron, where
-
+
+
INDUSTRIAL A N D ENGINEERING CHEMISTRY
April, 1925
tective layers” was corroded completely through, and broke, while other sections of the wire were as bright as when first immersed. Hydrogen-Ion Concentration and Corrosion of Iron
In the absence of oxygen, the rate of corrosion of iron will be proportional to the acidity or hydrogen-ion concentration of the solution in contact with the cathode. The writers believe this holds not only in the acid region but down t o a pH of 9.4. Their results indicate that the H-ion concentration determines the rate and total acidity the continuance of the reaction. Consequently, above a pH of 9.4 any substances present that will hold the acidity constant or “buffer” the reaction will be the controlling factor in the corrosion. Substances such as gypsum, limestone, and carbon dioxide, when occurring together, give this buffer action a t a pH of 5.1, while carbon dioxide and water give a n acidity of pH 4.0. The continuance of the corrosion of iron embedded in clay is partially explained by the buffer action of the clay. Where the solution is stagnant, migration of H T to the cathode will be the controlling factor. The foregoing cpsiderations hold where conductivity and hydrogen overvoltage remain undisturbed. When oxygen is present the H-ion concentration is determined a t the cathode by the solubility of ferric hydroxide. At the anode it will be dependent upon the migration of
(OH)-.
The anodic reaction is Fe
and the cathodic 3/2 H20
+ 2 @+
+ 3/4+ + +
0 2
+.
Fe++ 3(OH)-
(1)
+3@
Fe++ @ + F e + + + F e + + + 3(OH)- + Fe(0H)a or summing (2), (3), and (4) 3/2 H10 3/4 00 F e + ++. Fe (OH)3
+
+2 @
(2) (3) (4) (5)
If hydrogen is evolved a t the cathode, there will be an excess of (OH)- over that necessary for the production of ferric hydroxide and these (OH)- will eventually migrate towards the anode producing a precipitate of ferrous hydroxide if the solubility product of this hydroxide is reached. The H - at the anode will therefore be that determined by the Fe-’. These considerations show that the H-ion concentration cannot alone determine the rate of the corrosion when ferric hydroxide is present. The rate of diffusion of oxygen to the cathode is, as Wilson and Whitman point out, the determining factor, while the production and precipitation of ferric hydroxide a t the cathode is the characteristic reaction. At an acidity where ferric hydroxide can no longer remain insoluble, the H-ion concentration again functions as the controlling factor and the reactions concerned in the corrosion are independent of the oxygen supply. Hydrogen Evolution in the Corrosion of Iron
In a corrosion cell the cathodic hydrogen overvoltage drops off with the current density; consequently, as the H-ion concentration diminishes and the rate of corrosion drops, as a sort of compensating measure, hydrogen evolution will continue a t a lower e. m. f. The presence of oxygen increases the current density and consequently the cathodic hydrogen overvoltage, so that the evolution of hydrogen will not extend down to so low a H-ion concentration as in its absence. Moreover, the oxidation potential for iron lies a little below that for the evolution of hydrogen, so that in low concentrations of H + a t the cathode, such as the solubility of ferric hydroxide produces, the oxidation of the iron will be the preferential reaction and no hydrogen will be .evolved. When the H-ion concentration increases beyond
385
the solubility of ferric hydroxide the possibility of hydrogen evolution obtains, and a t a concentration corresponding to a p H of 5.4, Whitman found hydrogen gas to be evolved in the presence of oxygen. The writers found the evolution of hydrogen in the absence of oxygen to continue down to a p H of 9.4 or a t a concentration in respect to H + 10,000 times more dilute.
Corrosion of Iron By W. R. Whitney GENERAL ELECTRIC Co., SCHENECTADY, N. Y.
The paper describes some experiments in which the wearing away, or cutting, of iron by water in motion in presence of air was followed. It is an attempt to show that in some cases, where i t has been customary to consider the erosion of metals in contact with moving water as due entirely to a mechanical effect of the water, it is quite probable that the removal of the metal is preceded by a chemical corrosion.
I
T MAY be a question to some whether enough, and much more, has not already been written on this subject; but it will probably prove true that as new viewpoints arise much of the theory of corrosion will be rewritten, and even many of the old experiments will have to be repeated under influence of modern knowledge. For example, possibly many experiments on corrosion which, in the hands of different investigators, failed to give identical results, may be now explained, not only by differences in chemical purity or conditions of local stresses, but by differences in surface atomic orientations of the crystal structure as disclosed by X-ray crystal analysis. A sand-papered surface of pure iron may well act “energetically” in quite a different way from the normally oriented or annealed crystalline material. It is not the purpose of this paper to add to the quantity of theoretical matter which has so generally permeated the subject of corrosion of iron, nor yet, on the other hand, to deprecate or depreciate such theoretical work. Iron is our most generally useful element, and the value of even the crudest guess may prove to be a positive quantity. But a few experiments that have been made in this laboratory and which fit into the general subject may be useful and suggestive, and ought to be recorded. Corrosion of Soft Iron in Presence of Distilled Water
Mr. Fuller, of this laboratory, has shown that ordinary soft iron, when cleaned, as by surface grinding or by the use of emery paper, starts corroding very rapidly in the air where a drop of distilled water is in contact with the metal. Within less than 2 minutes a greenish color appears, and in 4 minutes a suspension of rust becomes visible throughout the drop of water. This indicates an exceedingly rapid attack of iron by water and air, which in general must be greatly slowed down by some cause in a short time; otherwise, no iron mould last as long as it does. This reduction of speed may he the effect of the covering when this is not removed. This coating is not easily and simply removed when dry, though it may possibly be removed before it dries. According to experiments reported by Fuller, as above, successive applications of the same size water drop to the identical spot on iron, resulting rust being removed by filter paper after each water application, showed that the quantity of corrosion rapidly decreased with successive applications. However, since the iron area was rougheningand therefore increasing, it seems as though some rptardirig wrface was produced.
