Dee., 1960
COMPARATIVE
ROLESO F OXYGEN AND
observed. As the power is decreased the four weak lines increase in amplitude, while the three strong lines decrease. At about 10 mw. only the four, single-quantum transition lines remain (Fig. 6). These transitions are more easily saturated and as a result are observed with low amplitudes at high power levels ; consequently, a decrease in power effects mainly the two-quantum transitions. At pressures above 0.20 mm. the spectrum consists of only three lines; the two 3P1 oxygen transitions and a strong, single line instead of the four (or seven lines, depending on the microwave power) lines observed a t low pressures (Fig. 7). This dependence of the spectrum on pressure was further investigated. From 0.020 to 0.120 mm. the width of the resolved single-quantum transitions shows a linear increase from 0.05 to 0.25 gauss. These values are in reasonable agreement with the Van Vleck-Weisskopf theory. In the pressure range from 0.20 to 9.0 mm., where the four lines are broadened into a single line, the width of this line increases again linearly from 2.5 to 6.2 gauss. From about 0.120 to 0.20 mm. there is a smooth transition from the four lines to a single line. Thus the appearance of Ihe spectrum a t high pressures is apparently the result of pressure broadening of the closely spaced individual lines. At low pressures a single broad line may also be obtained by over-modulating. Using this tech-
INHIBITORS I?1' I R O N PASSIVATIOX
1877
nique, which makes the detection somewhat easier, the atomic oxygen was observed down to a pressure of 0.005 mm. (total pressure in the flow system). Above 10 mm. total pressure detection becomes difficult due to the width of this line. Atomic oxygen, in contrast with atomic nitrogen is readily detected when air is passed through the discharge, demonstrating the longer lifetime of atomic oxygen under these conditions. Hydr~gen.'~-The spectrum of atomic hydrogen consists of two lines centered a t about g = 2.003, with a separation of approximately 1420 Mc.1 s ~ c .This ~ spectrum was observed over the pressure range of 3.0 to 8.0 mm. Optimum pressure in our case mas about 3 mm. The intensity was low, saturation was noticed even at low microwave power and the over-all reproducibility was poor. It seems likely that these results are diie to a low and poorly reproducible production of hydrogen atoms in the 2450 &IC./sec. discharge. Acknowledgments.-The author wishes to acknowledge the assistance of Mr. C. A. Hauck, Mr. B. M. Shields, and Mr. J. Saia in the various experimental phases of this work. (19) After this work was completed, Hildebrandt, Booth and Barth [ J . Chem. Phws., 81, 273 (195Y)lin a Letter to theEditor, reported their observations on the e.p.r. spectrum of atomic hydrogen. The effect of pressure on the line width for the atomic hydrogen lines ie corered in more detail, with results similar to those reported here.
THE COMPARATIVE ROLES OF OXYGEN AND INHIBITORS I N THE PASSIVATION OF IRON. I. SON-OXIDIZING INHIBITORS BY G. H. CARTLEDGE Chemistry Division,, Oak Ridge National Laboratory, Operated by Union Carbide Corporation for the U . S. Atomic Energy Commission, Oak Ridge, Tennessee Received M a y 7, 1960
Conflicting views are held as to the respective roles of oxidizing agent and inhibitor in the passivation of iron in aerated solutions of certain inhibitors. In the first paper of this series, experiments with inhibitors having no oxidizing properties are presented. It was found that passivation a t a noble potential may be achieved under conditions in which oxygen is necessarily the oxidizing agent, though without the inhibitor it leads to corrosion. It was found also that the passive potential is destroyed by low concentrations of sulfate ions, and thus exhibits the same behavior as that previously observed with oxidizing inhibitors. By means of cathodic polarization measurements, the magnitude of the cathodic current density in the reduction of oxygen on passive iron a t potentials above or somewhat below the Flade potential was determined. It was found, further, that addition of sulfate ions in low concentration sensitizes the system to activation under cathodic polarization. The results are interpreted as further evidence that the inhibitor's essential function is related to its adsorption, in competition with other ions or molecules.
Introduction In most of the earlier work on the passive state of iron, passivation was produced either by anodic polarization or by use of a vigorous oxidizing agent such as nitric acid. Recently, several studies have been concerned with passivation in aerated solutions of certain compounds commonly referred to as inhibitor^.'-^ Stern discussed the polarization curves of iron and stainless steel and superimposed upon them schematic polarizati on curves for an added oxidation-reduction couple to show how its normal electrode potential, exchange curj
(1) G . H. Cartledge, Z. Elektrochem., 62, 684 (1958). Contains references to several prior papers. (2) Milton Stern, J . EEectrochem. Soc., 106,638 (1958) (3) M. Cohen and d 4 . F. Beck, 8.Elektrochem., 61, 696 (1958).
rent, polarization characteristics and concentration enter into its ability to passivate the metal by its own reduction. Uhlig considers the passive film to consist of chemisorbed ~ x y g e n and , ~ Uhlig and King6have suggested that oxidizing agents like chromate and pertechnetate ions passivate iron by being reduced on large, discrete rathodic areas, thereby producing a sufficiently large local-cell current density on small anodic areas for passivation to result. A similar explanation is inherent in the treatment by Stern.2 These interpretations are obviously of an entirely different type from the strictly chemical point of view proposed by (4) H. H. Uhlig, L b d , 6%,626 (1958). (5) H H.Uhlig and P. F. King. J. Electrochen Soc., 106, 1 (1969). (6) T. P Hoar and U. R. Evans, zbzd., 99, 212 (1952).
1878
G. H. CARTLEDGE
Hoar and Evans,* Evans,’ and Cohen and Beck.* It is generally recognized that oxidation of iron by some means is involved in passivation processes, that is, processes in which both a low corrosion rate and a noble electrode potential result. If a reducible inhibitor is used in aerated solutions, experiments show that it may be reduced along with oxygen during passivation under certain conditions, but it remains to be demonstrated whether its function in the maintenance of passivity depends primarily upon such oxidizing and precipitating properties. The alternative view has been expressed that the inhibiting substance, without which passivation in aerated solutions does not occur, is adsorbed as such and that the adsorbed ions in some way alter the kinetics of one or more of the reactions involved in the corrosion and passivation processes. Cartledges*9 proposed the hypothesis that a special kind of interfacial electrostatic field arises by adsorption of XOP- particles when X has a high formal positive charge and the X-0 bonds have considerable polar character. For short distances from the interface within the solid substrate, this field was calculated to be opposite in sign from the image charge due to the negative ion considered as a unit.’ The adsorption of such ions was therefore assumed to alter specifically the kinetics of one or more processes involved in the corrosion reactions in such a way as to permit passivation. A related view was proposed by Kabanov and Leikis.Io It is obvious that an adequate interpretation of the passivation reaction presupposes an understanding of the process to be inhibited. I n a recent paper, Heusler” has given convincing evidence that the dissolution of iron involves an initial step by which a catalytic intermediate is formed electrochemically from iron and a hydroxide ion and remains adsorbed on the surface. The effect of competitive adsorption of inhibiting species on the surface then becomes understandable. This effect has been studied quantitatively for the anodic dissolution of iron in sulfuric acid with carbon monoxide or iodide ions as inhibitor.I2 In whatever way an inhibitor may affect the polarization characteristics of the electrode processes of iron and protons in the active region, however. it is also necessary to know specifically what cathodic processes are effective in producing and maintaining the higher potentials of the passive state, and what the relation of inhibitors to these processes is. Previous experiment^'^ showed that passivation of iron may be effected in an aerated solution of pertechnetate ions with consumption of so little TcO4- by reduction that oxygen is clearly the chief oxidizing agent. Yet without the pertechnetate a t sufficient concentration corrosion proceeds rapidly under otherwise similar conditions. It, is also known that both CrOr-and Tc04- may react with iron to form a film of (7) U. R. Evans, Z . EZeklrochem, 68, 619 (1958). (8) G. H. Cartledge, J . Am. Chem Soe , 77, 2658 (1955). (9) G. H. Cartledge, Corrosion, 11, 3350 (1955). (10) B N. Kabanov and D. 1. Leikia, Z . Elektrochcm., 61,660 (1958) (11) K. E. Heusler, ibtd, 68, 582 (1958). (12) K. E. Heusler and G. H. Cartledge, in procem of pnblicstbm. (13) G . FI Cartledge, Ta18 JOURNAL. S@, 979 (1065).
Vol. 64
insoluble oxides, but even in aerated solution this film does not long retard corrosion in the absence of the unreduced ions. According to Hoar and Evans16this excess of inhibitor is needed to precipitate in situ any iron(I1) ions emerging through faults in the protective film. The data of Brasher and c o - w ~ r k e r s ~ ~do, - ’ ~ indeed, show that the amount of Cr51 on a passive surface increases slowly on long exposure to solutions containing Cr6104--. Yet the f a ~ t that l ~ ~a variety ~ ~ of ions, such as S04--and Re04-, may destroy the inhibiting action of Tc04-, CrOk--, etc., apparently supports the suggestion that adsorption of the inhibitor in competition with other anions is an important, if not dominant, factor in its action. It is therefore an experimental problem to determine whether such reduction of a reducible inhibitor as actually occurs is really the essential part of its action in the passivation process. The experiments to be described constitute the first of a series of studies directed to an attempt to determine definitely the respective roles of oxygen and a variety of inhibitors in producing and maintaining the passivit,y of iron. The first portion of the study deals with the combined action of oxygen and inhibitors which are unable, of themselves, to oxidize iron or ferrous ions to the state found in the passive film. The action of certain potentially oxidizing inhibitors will be considered in subsequent papers. Passivation in the Presence of Non-oxidizing Inhibitors.-As non-oxidizing inhibitors, benzoate, phthalate and phosphate ions were chosen. It has been reported previously that neither benzoate1g nor phosphate*O inhibits in the absence of oxygen. Measurements were therefore made to determine (a) whether these three inhibitors in the presence of oxygen lead to an electrode potential that is more noble than the Flade potential, (b) whether the noble potential, if attained, is sensitive to low concentrations of added foreign ions, as was previously found with the reducible X04n- inhibitors,l8 and (c) the behavior of electrodes passivated in such systems when polarized cathodically. Experimental Phthalate was used a t concentrations of 5 X 10-3 to 5 X 10-2 f and acidities corresponding to pH values between 5.40 and 7.05. (Passivation was not achieved in an f solution a t pH 5.4.) The elecoxygenated 1.0 X trodes were made of electrolytic iron; they were abraded with 2/0 emery and cleaned in acetone and water. A stream of air or oxygen bubbled rapidly through the cell. The electrode was connected to a saturated calomel halfcell (S.C.E.) through a Haber-Luggin capillary and bridge leading to another compartment. A Vibrating Reed Electrometer and Brown Recorder were used in measuring the c.m.f.’s. The potentials recorded were the essentially (14) D. M. Brasher, A. H. Kingsbury and A. D. Mercer, Nature. 180, 27 (1957). (15) D. M. Brasher and C. P. De, ibid., 180, 28 (1957). (16) D. M. Brasher and A. H. Kingsbury, Trans. Faraday SOC.,64, 1214 (1958). (17) G. H. Cartledge, THIS JOURNAL, 60, 28 (1956). (18) R. F. Sympson and G. H. Cartledge, ibid., 60, 1037 (1956). (19) D. M. Brasher, “Chemistry Research” 1954, Her Majesty’s Stationery Office, London, 1955, p. 12 (Personal Communication). Cf. also F. Wormwell and A. D. Mercer, J . A p p Z . Chem., 8 , 150 (1952). (20) M. J. Pryor and M. Cohen, J . Electrochem. SOC.,98, 263 (1851); M. J. Pryor, M. Cohen and F. Brown, ibid.. 99, 542 (1952).
Dec., 1960
COMPARATIVE ROLESOF OXYGENAND INHIBITORS IN IRON PASSIVATION
stable values attained after an exposure to the environment lasting 16 hr , or longer. As shown in Table I , potentials more noble than the Flade potential were obtained; the potentials varied somewhat for different specimens and with the conditions of passivation. The most noble potentials were achieved when the solution was allowed to accumulate the iron ions that passed into solution during the passivation process, as in the experiment with footnote reference e in Table I. If the solution was replaced with fresh electrolyte before complete passivation, the electrode potential became noble to the calculated Flade potential,21 but without the appearance of visible f ilm or of coloration in the solution. Potassium sulfate was subsequontly added to certain of the solutions to determine whether the nobility is sensitive to the presence of foreign ions. It was observed that the potential remained stable in the 5.0 X 10-2 f phthalate solution a t pH 7.05 until the sulfate-ion concentration reached 3.3 X 1 0 - 2 f , when instability developed. I n the similar test with phthalate 5.0 X 10-8 f at a pH of 6.08, addition of sulfate to a concentration of 4 X 10-sf caused immediate debasing. The fall of potential in phthalate solutions was so irregular as to make it uncertain whether a halt sometimes observed in the vicinity of the calculated Flade potential was significant. I n any event, complete passivation was achieved under the conditions indicated.
1879
i
I
Y L
1,2, 3,4,
-400,
.4501
I
P A S S I V A T I O N AND ACTiVnT'ON Ik A E R A T E D BENZOATE SOLUilON,ZP
0.0105f 0.01(6 f
E, = c o ' c u l a t e d Flode p o t e n t i a l
I -!--O -I--"
.r.__~i~.~-_~i-.i..-~!-i_i 4 5 0 7
.3
8
TIME (irs),
Fig. 1.-Potential-time curves for activation of iron passivated in aerated benzoate solutions. Flade otentials are calculated from El = 338 - 59 pH mv. us. {C.E.
tential in inhibited systems to slightly lower pH values.z2J3 TABLE I For the phosphate solutions, phosphoric acid was diluted PASSIVATION IN SOLUTIONS OF NON-OXIDIZINQ INHIBITORS, with triply distilled water and neutralized by sodium hydroxide to give solutions of various concentrations and pH 20-23 ' Inhibitor
pH
Potential niv. (S.C.E.)
Flade potential calcd. mv. (S.C.E.), 20'
Phthalate, 5 . 0 X 10-zf 6.08" +10 - 16 5.80h e 0 Phthalate, 1 . 0 X 10-2,f 6.20" -5 5 d - 22 Phthalate, 1 . 0 X lO-*f Phthalate, 1 . 0 X 1 0 - 2 j G.22b +115 - 23 Phthalate, 2 . 5 X f 5.92* $81' - 6 7.05" -17 - 72 Phthalate, 5.0 X 10+f Benzoate, 1.05 X 1C)-2f 5.60" +35 12 5.60" +40 12 Benzoate, 1.05 X .f 5.3P $71 27 Benzoate, 1.00 X 10-2f Benzoate, 1.16 X 10-Zf 5.13" +lo2 +39 Benzoate, 1.16 X 1 0 - 2 f 5.13" +I16 $39 4.80" Active .. Benzoate, 1.25 X 10-2f Benzoate, 1.00 X 10-'f 5.40* 3 24 Phosphate, 4 . 8 X 10-? f 7.20" - 55 - 81 Phosphate, 1 .0 x 10-2 .f 7.llb 78 - 76 Phosphate, 1 . 0 X 10-l.f 7.02b - 60 -69 - 9 69 Phosphate, 1 . 0 X 10-l.f 7.02b 5 Aerated; Oxygenated. In 90 min. potential rose to -10 mv., then fell t o -70 mv., with formation of a slight turbidity in 3 days. Stable and electrode remained bright, although the potential never reached the Flade potential; the solution was clear after 3 days. "Visible film and solution ha,d perceptible brown coloration; weak test for Fe3+.
+ + +
+
+
-
-
With fully aerated and buffered benzoate solutions, aa with the phthalate solutiona, the range of concentrations and pH values within which complete passivation is possible is rather limited. Thus, a 1.25 X 10-2 f benzoate solution a t pH 4.80 left the metal at an active potential in four tests. With 1.0 X 10-af benzoate solution in the pH range of 5.0 to 7.5 and thoroughly aerated or oxygenated, the potential in vmious experiments rose well above the Flade potential (Table I). The potentials again became unstable when a sufficient concentration of sulfate ions was present, and definite halts very close to the Flade potential were obtained during the activation, aa seen for four experiments in Fig. 1, for pH values of 5.27, 5.60 and 5.70, respectively. These experiments supplement the results previously reported for inorganic inhibitors and extend the relationship between pH value and Flade po(21) U. F. Franck, Z. Natwforaha., 4a, 378 (1040).
values. With rapid oxygenation, passivation was not f solutions at pH's of 5.68 or 6.61. obtained in 5.0 X Between 5 X 10-3 f and 1.0 X 10-1 f, passivation above the calculated Flade potential was usually attained when the pH value was 7 or somewhat higher. Certain data are shown in Table I. In a 4.8 X 10-Sf solution at pH 7.20, sensitivity to added sulfate ions was demonstrated. The passive potential (-55 mv.) was maintained only 12 min. after addition of potassium sulfate to a concentration of 1.4 X 10-2 f. There was no evidence of a halt at the Flade potential during the ensuing activation.
Discussion These experiments demonstrate that the attainment of passive potentials does not require reduction of the inhibitor itself, since neither of the three inhibitors used is reducible under the conditions in effect. Furthermore, the passivity produced was shown to be destroyed by small concentrations of sulfate ions, as was previously shown for oxidizing inhibit~rs.'~J* With benzoate, a t least, there was an indication of a Flade potential, t,hough, unfortunately, the pH range in which passivation may be attained barely extends into the region in which the inhibitor is well buffered, so that the experiments were necessarily limited to a narrow range of acidities. At any rate, it is clear that oxygen even a t 0.2 atm. pressure can produce passivation very effectively without the help of an (22) G. H. Cartledge and R. F. Sympson, THIBJOURNAL,61, 973 (1957). (23) The irregularities in certain of the curves of Fig. 1 would leave some doubt aa to the significance of halts near the theoretical Flade potential were it not for the cumulative evidence of many previous experiments with four other inhibitors, aa given in ref. 22. The experiments were necessarily conducted under marginal conditions between activation and passivation, and the curves clearly demonstrate the d e c t s of the opposed proces8es in operation. At pH valuea of 6 or higher the reactions aseociated with the Flade potential are known to be aluggiah and somewhat erratic, but use of the experimental relation between Bfand pH at the values of pH required for paseivation in phthalate and phosphate solutions seems to be justified by the experiments on molybdate and tungatate" around pH 7, even though no reliable halts were observed with phthalate and phosphate solutions. According to theae measurements, the Flade potential is a property of the iron-film-proton system, and not of the inhibitor, unless it be in a iecondnry effect.
1880
G. H. CARTLEDGE
Vol. 64
oxidizing inhibitor, provided a non-oxidizing in- total cathodic current passing when a passivated hibitor is present under suitable conditions of con- iron electrode is cathodically polarized a t potencentration and acidity. Although the pH value is tials in the neighborhood of the Flade potential, a crucial condition, the failure of the perrhenate under conditions of concentration and acidity ionz4 or most other anions to inhibit in the same that permit passivation before application of currange of acidities shows that the maintenance of rent. It was thought that such measurements a low acidity is not the sole reason for the effective- should give results bearing on the role played by a ness of the anions which do inhibit. The experi- reducible inhibitor in the maintenance of passivity, ments do not of themselves show what the specific -that is, whether it actually adds significantly to effect of the inhibitor is, but definitely eliminate the available cathodic current according to the the requirement that it shall contribute directly proposal of references 2 and 5 . or acts by some other to the cathodic process by which the noble mixed mechanism. as is necessarily true of the inhibitors potential is achieved. The results are consistent devoid of oxidizing properties. As a background with the hypothesis that competing adsorption for these studies, measurements of the cathodic is involved. polarization of iron electrodes passivated in oxyCathodic Processes on Passivated Iron.-Previgenated phthalate solutions were made. Here the ous experiments have shown that passivation in a possible cathodic processes are reduction of either chromate solution in contact with air gives a film oxygen (perhaps tia HOz and HzOs) or the comcontaining both iron and chromium, with the ponents of the passive film itself, with ensuing actiiron compound predominating. 3,26 If the specimen vation. The potentials used were too noble for is first exposed to air, less chromium is taken up.I6 reduction of protons to be a factor, and use of The amount of chromium in the film is also a func- triply distilled water and pre-electrolyzed phthalate tion of the acidity, becoming less as the pH value solutions eliminated significant concentrations of increases.16s26 The chromium taken up was shown other reducible species. by autoradiography to be rather uniformly disExperimental tributed over the surface when an active specimen The galvanostatic procedure was used in making cathodic was passivated, except that higher concentrations under conditions which preceding experiof radioactivity were found a t inclusions, scratches, polarizations ments showed to be requisite for maintenance of passivity or other active sites.3 Similar observations were Polarizing current was drawn from a battery of up to 135 made when Tc04- was used as i n h i b i t ~ r . ~ ~volts ' ~ through high resistances, and the current was dedrop in potential across suitable preIt is unfortunate that, in the experiments with termined from thePotentials were measured by use of a cision resistors. Cr5'O4', no experimental differentiation between Vibrating Reed Electrometer or a Leeds and Eorthrup adsorbed G O 4 = and Cr(OH)3or Cr203 precipitated pH Indicator coupled to a Brown Recorder. Polarizations by reduction has been made, except that Brasher w r e made at the constant room temperature of 24". and Del5 considered that their data indicated ap- The cathodes were strips of electrolytic iron having an of approximately 1 cm.2. Before use, they were proximately a monolayer of CrO4= to be quickly area abraded with 2 / 0 emery, scrubbed in water, and then adsorbed, while a slow reduction continues over degreased m t h acetone. Just before the start of an exlong periods of time. Cohen and Beck3 found a periment, film formed during storage of the electrode in considerably larger uptake from NazCrOl (50 an oven a t 110' was removed by treatment with 1 N until the surface brightened. A Haber-Luggin p.p.m.) within the f i s t few minutes (up to 15 H2S04 capillary and bridge led to a flask containing a solution min.) on surfaces freshly reduced in hydrogen. of the same composition as that in the cell. A saturated Their data refer to the rapid initial action of some calomel electrode dipped into the solutionTin the flask. chromate species upon active iron or an iron(I1) The anode was platinum; to avoid changes in acidity, the were in the same cell. oxide when oxygen is either absent or limited by electrodes The phthalate stock solution was prepared from bufferdiffusion control from exerting its full effect. The grade potassium hydrogen phthalate, which was dissolved situation is different after the fast initial reaction in triply distilled water from a silica duplex still. This is ended and passivity is to be maintained a t the solution was pre-electrolyzed in a stream of helium with platinum electrodes in a divided cell. In the first experimore noble potential. The reported results there- ments, the catholyte from a 24-36 hr. electrolysis was again fore leave it somewhat uncertain to what extent electrolyzed anodically. It was found, however, that if the chromate ion is actually reduced in the mainte- this electrolysis was conducted a t too high a current density nance of passivity, and, whatever the amount may the product contained a trace of some reducible impurity which manifested itself by a distortion of the subsequent be, they give no basis for deciding whether such cathodic polarization curves a t the lowest current densities reduction as actually occurs is the primary source used. Since the anodic pre-electrolysis was found not to of the continuing inhibitory action in the presence change the general results of the polarization measurements, it was subsequently omitted. After cathodic pre-electrolyof oxygm. the solutions required very little sodium hydroxide The experiments of the preceding section demon- sis, to bring them to the desired pH values. During polarizastrated the ability of oxygen at atmospheric pres- tions in oxygen, cylinder gas was passed successively sure or less to produce passivity in the presence of through a long column of Ascarite, a tube packed with inhibitors that are unable of themselves to supple- glass wool, and a water saturator. ment the cathodic current available from reduction Results of oxygen. Further studies were therefore made In preliminary measurements it was demonto determine quantitatively the relative contributions of oxygen and a reducible inhibitor to the strated that cathodic polarization of well-inhibited passive specimens in oxygen can usually be carried (24) G . H. Cartledge, THISJOURNAL, 60, 32 (1956). a t least 150 mv. below the Flade potential without ( 2 5 ) J. E. 0 . M a m e and M. J Pryor, J . Chem. Soc., 1831 (1940). activation during the time required for the experi( 2 6 ) R. A. Powers and N. Hackerman, J . Electrochen. Soc., 100, 314 (1953). ment. It was found also that the open-circuit
Dec., 1960
COMPARATIVE ROLESOF OXYGENAND INHIBITORS IK IROX PASSIVATION
1881
(mv L S S.C.E. b r P l a t i n u m ' potential before polarization varied somewhat with +ZOO 1250 t300 '20 the conditions of passivation, particularly with reI spect to the amount of iron that was allowed to accumulate in the solution before passivation was complete. If the pH value was much below 6 and the phthalate solution too concentrated, brown solutions and visible films on the electrodes resulted. I n the measurements to be presented, the solutions were 0.0100 f, pH was close to 6.0, and the electrolyte in the cell was replaced with fresh solution a t least once, and usually two or three times, before polarization of the passive electrode was started. Under these conditions the electrode remained bright and the solution was colorless. Since passivation in phthalate solutions is achieved only a t pH values as high as approximately 6, deviations from Tafel lines became apparent a t current densities in the neighborhood of amp./cm.2, owing, presumably, to depletion CATHODIC DCLAR ZATION of oxygen or hydrogen ions required for the reducCURVES IY CXYGENATED O OIGO f PHTHALATE tion reaction. Further, the steady-state corrosion S O L U T I O N S , 24' \ current density of passive iron in 0.05 f phthalate \ E at pH 6 was reported by Weil and B ~ n h o e f f e r ~ ~ 10 E ' ''4 -233 -150 -"C 5c c +5c to be approximately 2 X amp./cm.2 and to be E L E C T R O D E PGTEllr'IAL I r n v v s S C E constant over a wide range of potentials. The Fig. 2.-Cathodic polarization of passive iron or platinum present measurements indicated steady-state cor- electrodes in oxygenated phthalate solutions. Curves are rosion rates of the same order of magnitude or some- extrapolated to E, and j,, which are the observed steadywhat less, hence it was possible to utilize data ob- state open-circuit potential and indicated current density, tained a t the lowest current densities by correcting respectively. Current densities are in amperes per cm.2 of apparent surface. "7, points established potentiostatithem, when necessary, for a constant anodic cur- cally; 0 , galvanostatic data for a different electrode; 0 , rent density in the range of lo-' amp./ points determined with decreasing current density. depending upon the conditions. In Fig. 2, the curve marked IV-1 is representative current density and stable potentials were obof the polarization data obtained in numerous series tained when the electrode was first used. After of measurements on passivated electrodes. The standing overnight in the phthalate solution the Flade potential calculated for pH 6.0 (at 20') is electrode had lost some of its activity. The open-11 mv. (S.C.E.), and it is clear that a normal circuit potential and indicated pseudo-exchange polarization curve was obtained to as low as current are most likely associated with the oxi-140 mv. S.C.E. Curve V-A,3 shows the open- dation of some organic material, as mentioned in circuit potential, E,, and three points determined connection with anodic pre-electrolysis. The deleterious effect of sulfate ions on the paspotentiostatically on a different electrode in 0.0100 f phthalate, also a t pH 6.0. When the polari- sivating inhibitors was clearly confirmed by addizations were measured first with increasing cathodic tion of sodium sulfate during a polarization meascurrent, the potential quickly returned to points on urement. Thus, a passivated iron electrode was the curve upon diminishing the current. Such given a test polarization; it showed typical bepoints are indicated on curve IV-1 by squares. havior from its corrosion potential of +60 mv. Different electrodes passivated spontaneously to S.C.E. to -168 mv. at a current density of 5.05 amp./cm.2, Current was then cut off Open-circuit, potentials that ranged from ap- X proximately the Flade potential to $125 mv., and the electrode stabilized overnight to a potenS.C.E., with indicated steady-state corrosion rates tial of +89 mv. S.C.E. Potassium sulfate was f that varied from about to 3 X amp./ added to make the phthalate solution 5 X in sulfate and a polarizing current density of 1.12 cm.2. Tafel slopes averaged 70 It 6 mv./decade in 14 polarizations with 5 electrodes variously X lo+ amp./cm.2 was applied. The potential fell rapidly at first, retarded a t a value which fell treated in successive measurements. At -75 mv. S.C.E., the (average current density for these near the Tafel line previously obtained, and then 14 polarizations was 3.6 f 1.4 X amp./ debased irregularly to an unsteady, oscillating cm.2. This value occasionally was as high as value approximately - 165 mv. S.C.E. The 2 X amp./'cm. when a specimen was first potential remained almost unchanged while the current density was increased to 5.02 to polarized. amp./cm.2, when activation set in. When the For comparison, Fig. 2 shows a polarization current was cut off, the specimen failed to become curve for reduction of oxygen on a smooth plati- repassivated and subsequently corroded heavily. num electrode in 0.011 f phthalate a t pH 5.51. Discussion Before use, the electrode was cleaned in a mixture The rapid re-ennobling of the electrode potential of HC1 and H20:,. Rapid response to changes in on diminution of the applied cathodic current and (27) K. G,Weil and K. F. Bonhoeffer, Z. phzlszk. Chem., X.F., 4, the stability of the observed potentials a t the cur1 7 3 (1955).
\>,
L__ J
_ _
1882
G. H. CARTLEDGE
rent densities used indicate that reduction of the passive film itself was unimportant kinetically by comparison with the reduction of oxygen. Incomplete studies in the benzoate system show that the generd relationships are similar to those in the phthalate system, with some differences in detail. For passivation to be maintained in the absence of an externally applied current it is necessary that the cathodic current density available from the effective passivator exceed the steady-state corrosion current density; that is, the polarization curve for reduction of the passivator must intersect the polarization curve of iron a t a potential more noble than the Flade potential.28 The steadystate corrosion rate varies with the nature and concentration of the electrolyte, as well as with pH, as may be seen by comparing the phthalate data of Weil and BonhoeffeF with Vetter's data for sulfate solutions.29 From the present experiments it is clear that at pH values appreciably below 6 in phthalate solutions of 5 X to 5 X low2j the oxygen current is inadequate for complete passivation, whereas it is ample a t pH 6 or higher. By extrapolation from Fig. 2 it is seen that, in the region of passive iron potentials, reduction of oxygen on the platinum electrode exceeded that on the iron electrode by a factor of about lo4. This is a demonstration of an effect similar to that of platinum on the passivation of stainless steel, as discussed by T o m a ~ c h o wand , ~ ~ by Stern and Wissenberg3I for platinum and titanium. (28) This assunies the validity of Vetter's demonstration t h a t the paeaive film is free of poi=, BO t h a t both the anodic and cathodic prooesses operata over the same total area. Cf.K. J. Vetter, 2. h'lektrochsm., 55, 274 (1951). (29) K.J. Vetter, ibid., b9, 67 (1955). (30) N. D. Tomaschow, rbid., 62, 717 (1958). (31) Milton Stern and Herman Wissenberg, J . Electrochem. Soc., 106, 759 (1959).
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If the curves of Fig. 2 are extrapolated to the corresponding reversible oxygen electrode potentials, it may be estimated that tlhe exchange currents are of the order of 10-14 amp./cm.2 on the platinum surface and lo-'* amp./cm.2 on passive iron. These numbers may be compared with to 10-lo smp./cmq2found by Bockris and Huqa2 for smooth platinum in sulfuric acid of pH 1.25, and lo-*" ampJcm.2 obtained by extrapolation of the data of Wade and Hackermanas on passive iron a t 5' and a t pH 4. The present measurements therefore show that oxygen alone a t 1 atm. or less is reduced rapidly enough on a passive iron electrode to maintain the mixed potential in the passive region in spite of a continuing corrosion current densit'y, which is of the order of - lo-* amp./cm.2 in the phthalate system a t pH 6. The non-oxidizing inhibitor is essential, however, and its effectiveness is disturbed by the addition of foreign ions such as the sulfate ion, which sensitize the system to activation. Since it is difficult to see how such ions enter directly into the electrochemical processes a t the low concentrations involved, it seems most reasonable to assume that they compete with oxygen or inhibitor a t sites that are active in the electrochemical reactions. I n the following paper, the degree of participation of a reducible inhibitor in the total cathodic process will be examined. Acknowledgment.-It is a pleasure to acknowledge helpful discussions with my colleagues, E. J. Kelly, R. E. Meyer and Franz A. Posey in connection with this series of studies. (32) J. O'M. Bockris and A. K. M. S. Huq, Proc. Roy. Soe. (London), A231, 277 (1956,. (33) W. H. Wade and N. Hackerman, Trans. Puradag Soc.. 68, 1 (1957).
THE COMPARATIVE ROLES OF OXYGEN AND INHIBITORS IN THE PASSIVATION OF IRON. 11. THE PERTECHNETATE ION BYG. H. CARTLEDGE Chemistry Division, Oak Ridge National Laboratory, Operated by Union Carbide Corporation for the U. S. Atomic Energg Commisszon,Oak Ridge, Tennessee Received May 7, 1960
Galvanostatic and potentiostatic polarizations of passive iron electrodes have been made in solutions of a phthalate or pertechnetate and mixtures of them. By measuring the polarizations both in oxygen and in essentially oxygen-free helium the relative contributions of oxygen and the reducible inhibitor to the total cathodic current have been determined. An acceleration of the cathodic processes by the reduction product of the pertechnetate ion was demonstrated. It was shown also that, in all cases, reduction of oxygen is the principal cathodic process a t passive potentials.
The preceding paper in this series' demonstrated that passivation of iron can be achieved under the oxidizing action of oxygen alone a t atmospheric pressure, provided a suitable non-oxidizing inhibitor is present. When a reducible inhibitor is available, it may supplement the cathodic current due to reduction of oxygen, and some reduction of such inhibitors is generally observed. The purpose of this study is to determine the extent to (1) G.H.Cartledge, THISJOURNAL, 64,1877 (1960).
which the inhibitor itself is reduced under passivating conditions, in comparison wit,h the reduct,ion of oxygen under a total pressure (including water vapor) approximating 1 atm. For this purpose, the pertechnetate and chromate ions and osmium(VII1) oxide have been used as inhibitors. The pertechnetate ion differs in important ways from the chromate ion with respect to properties that are important for theories of inhibition which assume oxidizing and precipitating power as the
