the predominant cyclohexyl CH2 groups. Compound VI1 was also examined with lithium fluoride optics to obtain higher resolution and its spectrum showed a weak absorption band a t 2960 em.-' in addition to that a t 3030 em.-' The compounds VI11 through X, which contain only tertiary cyclopropyl ring hydrogens, showed a single absorption near 3005 em.-' This would be anticipated since two absorptions in the 3000-cm.-' region would not appear if the ring unit C H R or CR2were present rather than the CH2 ring group (20). We noted that the parent compound of compounds IX and X, i.e., norcarane, has the characteristic cyclopropyl CH2 stretching frequencies at 3077 and 3003 em.-' For obvious reasons, samples XI and XII, which show absorptions in this region because of unsaturation, will not be discussed. The data presented in Table I show that significant absorptions appear consistently in the 1040- to 1008-cm.-l range with the majority centered near 1020-cm.-1 A deviation noted in this series is in the spectrum of compound VIII, which contains the cyclopropyl ring in a substituted form as part of a spiro and bicyclic structure. Since similar compounds of this type are not available for comparison, the significance of this deviation is not known. As in other forms of substituted rings, Le., the nortricyclenes, compound VI11 also shows absorption near 850 cm.-l From spectra reported by von Doering (8),compounds IX and X, with heavy ring substitution, exhibit absorptions a t
frequencies lower than 860 em.-' Bands appear a t 840 and 830 em.-', respectively. Likewise, cornpound XI absorbs near 850 em.-' and compound XI1 shows medium absorption a t 835 em.-' Generally, for compounds in this series, the band near 860 em.-' proved to be less useful than bands in the 3050- and 1020-cm.-' regions for the identification of the cyclopropyl ring. From these studies, the usefulness of the assignments for the cyclopropyl ring system in the 3050-cm.-' C-H stretching and 1020-cm.-l ring deformation regions has been confirmed and is in accord with the works of Wiberley et al. (19, 20) and Slabey (18). Although these assignments in some instances have limited value as noted, the benefit derived from their use in certain structural studies involving polycyclic hydrocarbons and their derivatives is well established. ACKNOWLEDGMENT
The authors thank W. R.Smith, J. C. Alm, R.H. Campbell, J. I;.Driscoll, and A. Bekebrede for their contributions to and helpful discussions of this work LITERATURE CITED
(1) Allen, C. F. H., Davis, T. J., Humphlett, W. J., Stewart, D. W., J . Org. Chem. 22, 1291 (1957). (2) Bartleson, J. D., Burk, R. E., Lankelma, H. P., J . Am. Chem. SOC.68,
2513 (1946). (3) ~, Barton. D. H. R.. J . Chem. SOC. 1951, 1444. (4) Bridson-Jones, F. S., Buckley, G. D.,
Cross, L. H., Driver, A. Y . , I b i d . , 1951, 2999. ( 5 ) Cleveland, F. F., Murray, &I. J., Gallaway, R. S., J . Chenz. Phys. 15, 742 (1947). (6) Cole, 9. R. H., IbLd., 1954, 3807. (7) c ~ x p c ~ i ments tltwribcti bclon. supl)oi’t the contcmtion that this \mvc is not tlric to the rrtluction of n rceit1u:il p1:itinuni oxide but to blie prcscwccl of frcdily forincd, finely divided, pl!itiiiiuu nwtal that rcwilts from :dternately oxidizing and reducing :L platinum cl(~ctrodc~.Iii ot1it.r ~ o r d s tlic , cxlectrotle surfarc bccomes platinized. Conclusive evidence that the small wave a t the foot, of t’he chronopotentiogram is not due to platinum oxide was obtained by chemicnlly st’ripping all oxidized p1:itinuni from an oxidized and a partiallj. rccluctd platinum electrode and analyzing the stripping solutions for platinum sprctrophotometrically. -4solution that is 0.05F in HZSO4 and 0.1F in IC1 vcry rapidly dissolves the platinum oxides from an oxidized platinum electrode in the form of iodoplatinum compleses ( 2 ) . Because of overlapping spectra it is not possible to analyze iodide stripping solutions quantitatively for Pt(I1) and Pt(1T’). However, the prescnce of m a l l amounts of PtJB-’ and Pt14-2 can be readily detected by measuring the absorbance of acidic iodide solutions a t 498 mp. (Any Is- present in the solution does not absorb a t this wave length.) The following experiment was performed : A large, platinum gauze electrode having a n area of approximately 100 sq. cm. was oxidized potentiostatically at 1.4 volts us. S.C.E. in oxygen-free 1F H2S04. The oxidized electrode was washed with distilled water and then immersed in 100 ml. of the iodidestripping solution that was being continuously swept with nitrogen. This process was repeated twice more, resulting in an easily perceptible red-orange color in the iodide-stripping solution. The absorbance of an aliquot of this solution in a 10-em. cell was measured at 498 m l with a portion of the original iodide solution as the blank. The absorbance was 0.345. The experiment was then repeated with a fresh stripping solution, the only difference being that after each oxidation of the electrode it was reduced potentiostatically a t 0 volt us. S.C.E. before being immersed in the iodidestripping solution. As can be seen from Figure 1, curves 1 and 3, a t 0 volt os. S.C.E. the major portion of the platinum oside on the electrode will have been reduced, but any oxide that contributes to the small wave a t the foot of the chronopotentiogram will not have been reduced. The absorbance of the stripping solution obtained in this way was less than 0.003. Thus the amount of oxidized platinum present on an electrode that displays the small wave is a t most 1% of that on a fully oxidized electrode. As is apparent from Figure 1, curve 1, the number of coulombs corresponding to
936
ANALYTICAL CHEMISTRY
I
0.8,
TIME
Figure 2. Cathodic chronopotentiograms in air-free 1 F HzS04 1. Reduction of unoxidized a g e d electrode 2. Repeat of 1, ofter oxidation of electrode and recording of one previous cathodic chronopotentiogram 3. Reduction of unoxidized aged electrode following platinization in H,PtCI, Current density 600 pa. per sq. cm. throughout
the transition time for the small wave is about one quarter of that corresponding to the complete oxide-reduction wa\ e. Thus, the amount of oxide stripped by the iodide solution from the reduced electrode is not nearly adequate to account for the size of the small wave. If the small wave is not due to platinum oxide reduction, what does cause it? Lingane observed that a reduced platinum electrode that has stood a day or more in sulfuric or perchloric acid gives no small wave at the foot of chronopotentiograms and that no osygen reduetion wave is observed with such electrodes in oxygen-saturated solutions. Both waves can be restored by osidation of the electrode. The author has confirmed these observations, but has found t h a t it is also possible to cause the two waves to reappear if the aged electrode is very lightly platinized in a n acidic chloroplatinate solution. Passing cathodic currents of 20 to 30 ma. per sq. cm. through the electrode for 5 to 10 seconds in a 0.5F solution of H2PtC16 is ample to restore the waves. Furthermore, an aged electrode that does not display the small nave at the foot of the cathodic chronopotentiogram in oxygen-free 1F sulfuric acid does display such a wave immediately following light platinization. Curve 1 in Figure 2 is a chronopotentiogram obtained with a reduced and aged (24 hours) electrode in oxygenfree 1F H2S04. Note the absence of the small wave. The electrode was strongly oxidized, a cathodic chronopotentiogram was recorded that showed the oxide being reduced, and then curve 2 was recorded. Kote the presence of the small wave at the foot of the chronopotentiogram. The electrode was allowed to age until the small wave in curve 2 disappeared and curve 1 was reproduced. The electrode was then
made the cathodc in a 0.5F HzPtCl, solution and 10 ma. per sq. em. were passed for 10 seconds. After thorough washing. the electrode was replaced in air-free 1F H2S04>and one cathodic chronopotentiogram was recorded to encuie that any adsorbed Pt(1V) was reduced. Curve 3 resulted nhen a second chronogotentiogram was recorded. S o t e the presence of the small wave at the foot of the chronopotentiogram and its similanty to the wave in curve 2 . This wave gradually disappears as the freshly deposited platinum ages and loses its catalytic properties. The reason that the presence of finely deposited platinum on the electrode results in a small prewave just before hydrogen evolution commences is probably associated n i t h the production of adsorbed hydrogen on the freshly deposited platinum. The small potential inflection corresponds to the point at which the platinized surface is just covered ~ i t ha layer of adsorbed hydrogen. The same conclusions based on essentially the same kind of evidence were expressed by Butler and coworkers many years ago ( 5 ) . I n any case the small wave observed by Lingane and in this study clearly results from the presence of freshly formed, finely divided platinum metal on the surface of the platinum electrode. Such a layer results from the reduction of an oxidized electrode as well as from the deposition of metallic platinum on the electrode in solutions of HQPtC16. As one would espect on the basis of the classical use of platinized platinum electrodes to increase the reversibility of the standard hydrogen electrode, the presence of finely divided platinum on electrodes increases the reversibility of reactions carried out at the electrodes. As long ago as 1924 Hammett (7) reported that oxidation and reduction of platinum electrodes increased their “activity” for the subsequent oddation of molecular hydrogen and the reduction of hydrogen ion. Lingane obs e n e d that the reduction of oxygen is much more reversible a t a recently oxidized (and hence presently platinized) electrode. Oxygen reduction failed to occur before reduction of the background a t aged electrodes. Analogous results are obtained in the reduction of Fe(II1) in sulfuric and perchloric acids. The reduction of Fe(II1) occurs at potentials several hundred millivolts less oxidizing at aged electrodes. I n addition, if the electrode is subjected to any treatment that will chemically dissolve the oxide film before it is reduced, the resulting electrode will not have a platinized surface and the reduction of Fe(II1) immediately proceeds less reversibly, no aging being required. Curves 2 and 4 in Figure 1 show the difference between the chronopotentiograms obtained for the reduc-
tion of Fe(1II) a t a n electrode from nhich the oxide has bccn renloved by rcdurtion and with the same electrode from w hic,h the oxide has been removed by c.hcmica1 dissolution. Curve 2 in Figure 1 is typical of cathodic chronopotentiogranis obtained with an electrode t h a t has been alternately oxidized and reduced 2 or 3 times. Curve 4 in Figure 1 rcsulted after the electrode was oxidized, rclmoved from the solution and immcrred in a 12F HC1 solution at 95°C. for 60 seconds (hot, concentrated HC1 dissolves the oxide from the electrode) ; thoroughly washed; replaced in the Fe(II1) solution in 1F H2S04;and a cathodic chronopotentiopram recorded. The decrease in reversibility is so marked t h a t scarcely a n y reduction wave for the Fe(II1) is observed. If a recently oxidized and reduced electrode that has a platinized surface is similarly treated with 12F HC1, essentially no decrease in reversibility for the subsequent reduction of Fe(II1) results. The fact t h a t the behavior of platinum electrodes is affected markedly by the presence of finely divided platinum metal on its surface makes i t necessary t o distinguish among three varieties of electrodes rather than between t h e two common designations of “oxidized” and “reduced.” I n addition t o “oxidized electrodes” one must consider two varieties of “reduced electrodes”: those t h a t are free of platinum oxide and do not have platinized surfaces and those t h a t are free of platinum oxide but do have a layer of freshly r e d u c d platinum metal on their surfaces. Throughout this paper, unless the contrary is specified, a “reduced electrode” is one t h a t does have a platinized surface Platinization of t h e electrode surely increases its true area. However, such a n increase Mould be observable in chronopotentiometric experiments only if the “hills and valleys” created by the platinizatiori of the electrode surface had depths of the order of t h e diffusion layer thickness. No such effect was detected n hen lightly platinized electrodes were used in cathodic chronopotentiometric experiments with Fe(II1) with transition times of a few seconds or longer, so that the effective ehronopotentiometric area is not significantly altered by light platinization. Oxidation of Fe(1I) in Sulfuric Acid. Curve 5 in Figure 1 shon-s a n anodic chronopotentiograni for a 2mM solution of Fe(I1) in H2SO4with a reduced electrode. If t h e electrode is oxidized a n d replaced in t h e ferrous solution, t h e resulting anodic chronopotentiogram is identical to curve 5. If t h e electrode is again strongly oxidized, placed in t h e Fe(I1) solution for one minute, removed from the ferrous solution, washed free of iron, and used t o record a cRthodic chronopoteiitiogram in
TIME
Figure 3.
Cathodic chronopotentiograrns for Fe(lll) in 1 F
HClOi 1. Reduction of oxide film in iron-free 1 F HCIOI 2. Reduction of 2 m M solution of Fe(lll) at reduced electrode 3. Repeat of 2, at oxidized electrode 4. Repeat of 3, while stirring solution Current density 600 pa. per sq. cm. throughout
a n oxygen-free 1F HzS04 solution, no waive corresponding t o the reduction of the oxide film is observed. Thus no effect due t o electrode oxidation is observed in chronopotentiograms for the oxidatioii of Fe(I1) in HzS04 because t h e platinum oxide film is chemically reduced rapidly by the Fe(I1). The rate of the reduction of the platinum oxide decreases with the concentration of Fe(II), but even in 0 . l m M solutions t h e reduction is essentially complete in 2 to 3 minutes. Therefore, with ordinary voltammetric experiments that require a few minutes to perform, it is not possible to observe the effect, if any, of electrode oxidation on subsequent reactions at the electrode in sulfuric arid solutions containing Fe(I1). Essentially the same conclusions were reached b y Baker and hfacxevin (4) with much more concentrated solutions of Fe(I1) in sulfuric acid. There is evidence t o indicate t h a t the oxidation of Fe(1I) in sulfuric acid solutions is rendered less reversible by oxidation of t h e electrode. I n order to observe the effect, however, the experiments must be conducted in such a way t h a t the chemical removal of the platinum oxide b y t h e Fe(I1) is eliminated or slowed down sufficiently for measurements t o be made while it is still present. Baker and hIacNevin (4) performed potentiostatic experiments involving the oxidation of sulfuric acid solutioiis of Fe(I1) a t platinum electrodes covered with a n oxide film b y presetting the potentiostat t o +0.7 volt us. S.C.E. before introducing the oxidized electrode into the solution of Fe(I1). At this potential the coneentration of Fe(I1) at the electrode surface is maintained so near zero that no significant chemical reduction of the film can occur. These experiments showed that oxidation of Fe(I1) a t oxidized elcctrodes proceeds less reversibly than a t reduced electrodes. Curve 6 in Figure 1 shows the initial portion of a chronopotentiogram of the oxidation of Fe(I1) in 1F H2S01 that
was obtained by oxidizing the electrode strongly, stirring the solution for 2 to 3 seconds, turning off the stirrer, and immediately recording the anodic chronopotentiogram. The decrease in rcversibility indicated by the larger jump in the potential at the very start of the chronopotentiogram results from the platinum oxide on the electrode that had not yet been rcduccd b y the Fe(I1). Kolthoff and Sightingale (9) decided on the basis of current-voltage measurements that’ the platinum oxide produced by oxidation of platinum electrodes renders the Fe(I1)-Fe(II1) couple more reversible in 1F sulfuric acid. The present cl-idence sugyrssts that the platinum oxide present on the electrodes used by Kolthoff and Kightingale in 1 F HzS04 solutions of Fc(I1) could not have been sufficiently longlived to contribute to their esprrinicntal results and that the increase in revrrsibility they observed vas due to the finply divided platinum that wvas formed on the surface of t’heir elcctrode when it was oxidized and then clirmicnlly reduced by the Fe(I1). A report by D a h (6) [rontrary to earlier work of Baker and 11acNevin ( 4 ) ]that. oxidation of a platinurn electrode produces an increase in rcvcrsibility for subsequent oxidation of As(II1) in 1F H2S04is probably erroneous for a similar reason. As(II1) in 1F H2S04should rcduce any platinum oxide to give a platinizctl but reduced electrode at which the oxidation of ils(II1) proceeds more reversibly. The reason for the decreases in rcversibility observed by Davis when the electrode was reduced at potcntials of 0.55 volt us. S.C.E. :md bclow may have arisen as follows: The potential of 0.55 volt z’s. S.C.E. is suspiciously close t o the potential of the .4g-a4g+ couple. Since Davis carried out his oxidation wit’h Ag(I1) solutions .rather than electrolytically, it is likely that some Ag(1) was adsorbed on the surface of his electrode. The reduction of this VOL. 33, NO. 7,JUNE 1961
937
Ag(1) to the metal could have been rcsponsiblc for the decrease in rerersibility. Reduction of Fe(II1) in Perchloric Acid. Cathodic chronopotentiograms for a 2 m M solution of Fe(II1) perchlorate in 1F HClOd, as well a s for a 1F HCIOI solution alone, are shown in Figure 3. Curve 2 in Figure 3 is a chronopotentiogram for t h e reduction of Fe(II1) a t a reduced platinum electrode. Curve 3 resulted when t h e electrode was oxidized prior t o t h e recording of the cathodic chronopotentiogram. The t n o waves that appear are due, first, to the reduction of most of the oxide film on the electrode (compare with curve 1, Figure 3), followed by the Fe(II1) reduction that now occurs a t a potential about 250 mv. more reducing than a t the uiiosidizcd electrode. That this is the correct assignment of each electrode reaction to the naves was verified by showing that the transition time for the first wave is nearly independent of the concentration of Fe(1II) and of whether or not the solution is stirred. This is the first reported instance in which the reduction of a species t h a t proceeds relatively reversibly at reduced platinum electrodes becomes irreversible a t oxidized electrodes. It might be expected t h a t once the oxide film is reduced the reduction of the Fe(II1) would be occurring a t a reduced electrode, so t h a t the electrode potential should reach a minimum and then increase t o the potential, Nhhere Fe(II1) is reduced a t a n initially reduced electrode. This is just what happens if the solution is stirred during the recording of t h e chronopotentiogram, as can be seen in curve 4 of Figure 3. This behavior is not observed in unstirred solutions, presumably because the small quantity of oxide remaining on the electrode a t the transition time for the first wave in curve 3 (+0.30 volt us. S.C.E.) requires a longer time to be reduced than the transition time for the Fe(II1). Oxidation of Fe(I1) in Perchloric Acid. Curve 2 in Figure 4 is a chronopotentiogram for t h e oxidation of Fe(I1) at a reduced electrode in a 1F HC104 solution t h a t was 2 m M in both Fe(I1) and Fe(II1). The lack of sharpness in t h e potential inflection results from oxidation of t h e electrode commencing just before the transition time for t h e oxidation of t h e Fe(I1). Curve 1in Figure 4,which is a chronopotentiogram for the oxidation of a platinum electrode in pure I F H2C103, shows that oxidation of the platinum electrode commences a t about +0.7 volt us. S.C.E. The rate of potential increase following the transition time is smaller in curve 2 than in curve 1 because the current available to oxidize the clectrode beyond the transition time 938
0
ANALYTICAL CHEMISTRY
c8nc.-
I
1
TIME
Figure
4.
Anodic chronopotentiograms for Fe(ll) in 1 F HCIOd
1, Oxidation of electrode in iron-free HCIOe 2. Oxidation of Fe(ll) in solution containing 2mM Fe(ll) and 2mM Fellll) at reduced electrode 3. Repeat of 2, after electrode oxidized patentiostatically at 1.4 volts VS. S.C.E. 4. Repeat of 3, after electrode oxidized at 1.6 volts vs. S.C.E. 5 . Repeat o f 4, after electrode oxidized at 3.0 volts vs. S.C.E. Current density 600 p a . per sq. cm. throughout
is much smaller in a solution of Fe(II), where ferrous ions continue to diffuse to the electrode and consume current. Surprisingly, in perchloric acid, the chemical reduction of the oxide film by Fe(I1) proceeds much more slowly than in sulfuric acid, and it is a simple matter to obtain chronopotentiograms with oxidized electrodes in solutions containing Fe(I1) perchlorate. Curves 3, 4, and 5 in Figure 4 are chronopotentiograms for the oxidation of Fe(I1) in 1F HClO, at a n electrode that had previously been oxidized potentiostatically a t 1.4, 1.6, and 3.0 volts us. S.C.E., respectively, and then stirred free of oxygen. The electrode oxidation produces a pronounced decrease in reversibility. The oxidation of Fe(I1) now occurs a t a potential u p to 500 mv. more oxidizing than with a reduced electrode. The much sharper potential inflections are t o be expected, since in this case no current is expended beyond the transition time in oxidation of the electrode. This decrease in reversibility for the oxidation of ferrous perchlorate at oxidized electrodes is analogous with the results obtained by Baker and MacKevin (4) and in the present study (curve 6, Figure 1) for the oxidation of Fe(I1) in sulfuric acid. Furthermore, many other oxidations that show similar oxide film effects are known (1, 3, 6, 8, 12). I n fact, except for the papers of Kolthoff and Nightingale (9) and Davis ( 6 ) ,all of the recently reported instances in which oxidation of the electrode produces any effect on subsequent oxidations at the electrode have involved a decrease in reversibility. This appears to be a general rule. It could be that the oxide film produces an apparent irreversibility in subsequent oxidations simply by providing a n ohmic resistance across which the overvoltage appears. A purel) ohmic resistance can be ruled out, however, because the magnitude of the overvoltage is not linearly dependent on the current density. To rule out the possibility t h a t the oxide might behave as a nonohmic resistance in series with
the electrode, experiments were performed in which platinum electrodes were oxidized in sulfuric acid solutions and then removed, washed free of electrolyte and dried in air, and the current-voltage characteristic of the resulting electrode determined with a very fine gold wire (ea. 0.2-mm. diameter) resting lightly on the surface of the electrode. I n all such experiments the electrode behaved like a pure metallic conductor, and no difference was observed between oxidized and reduced electrodes. Another possible explanation for the decreased reversibility observed for oxidations a t oxidized electrodes is that the layer of finely divided platinum metal has been oxidized and is not present. This is in line with the observation made earlier that unosidized electrodes which have been freed of oxide in ways that do not platinize the electrode surface display larger overvoltages for the reduction of Fe(II1) than do reduced electndes. (The same phenomenon is observed in the reduction of oxygen.) The absence of the platinized platinum layer on oxidized electrodes is doubtless partially responsible for the higher overvoltage, but it cannot account for all the experimental obwrvations t h a t have been made. For example, the oxidation of oxalic acid in I F HzS04 a t platinum electrodes n a s shown by Lingane (12) t o occur before oxidation of the sulfuric acid supporting electrolyte only with reduced electrodes. Experiments in these laboratories have shown that oxalic acid oxidation also occurs before oxygen evolution a t oxidized electrodes that have been chemically stripped of their oxide. Thus t h e pronounced decrease in reversibility for oxalic acid oxidation a t oxidized electrodes is not due solely to the absence of the platinized platinum film. A detailed interpretation which can account more fully for the variety of experimental results that have been obtained remains the object of continued investigation.
ACKNOWLEDGMENT
‘I’lic s u l y o r t of thc Office of Ordnnncc Rcscarch, U. S. ,4rmy, under Contract DAk-04-495-ORD-149 is gratefully acknowledged. Extensive corrcsgondence between t h e author and James ,J. Ling a m contributed t o the interpretation of results, but responsibility for the conclusions rcwlied rests with the author. LITERATURE CITED
(1) Anson, E’. C., J . Am. Chem. SOC.81, 1554 (1959).
(2) Anson, F. C., Ph.D. thesis, IIarvarci
University, 195’7. (3) Anson, F. C., Lingane, J. J., J . Am. Chem. Soc. 79, 1015 (195.7). (4) Baker. B., MacNevin. W.,Ibid.. ~, 75, 1476 (1953). (.5\ Butler. J. A. V.. “Electrical Phenomena‘atInterfac&,’) Chap. IX, Macmillan, New York, 1951. (6) Davis, D. G., Talanta 3,335 (1960). (. 7,) Hammett. L. P., J . Am. Chem. SOC. 4 6 , 7 (i924j. (8) Hickling, A., Wilson, W.,J . Electrochem. SOC.98,429 (1951). (9) Kolthoff, I. hf., Nightingale, E. R., Anal. Chim. Acta 17,329 (1957). \-I
(10) L:iitincn, H. A., Enke, C. G., J . Electrochem. Soc. 107, 773 i 1 9 f i O l (11) Lingane, J. J., “L1ect;oa;alytical Chemistry,” 2nd ed., Chap. XXII, Interscience. S e w York, 1958. (12) Lingane, ‘J. J., J . Electround. Chem. 1. 379 f 196O’i. (13j I b i d , in press. (14) Whiteker, R. A., Davidson, N., J . A m . Chem. SOC.75,308 (1953). RECEIVEDfor review January 10, 1961. Accepted April 11, 1961. Contribution 2641 from the Gates and Crellin Laboratories of Chemistry.
Exchange Current De nsities f o r Fe (I I)- Fe(III) SoI utio ns in Sulfuric Acid and Perchloric Acid FRED C. ANSON California institute of Technology, Pasadena, Calif.
b Exchange current densities for the Fe(ll)-Fe(1ll) couple have been evaluated galvanostatically in sohtions of sulfuric and perchloric acids a t a platinum electrode. The results indicate that the heterogeneous rate constant for this couple is independent of p H between p H 0 and p H 1 and is decreased b y complexation with SUIfate. Perchloric acid solutions of Fe(ll) and Fe(l1l) display an anomalous concentration dependence of the exchange current density that may b e due to contributions to the measured exchange currents from Fe(l1l) or Fe(ll) adsorbed on the electrode.
E
current clensitics a t platinum electrodes for solutions of Fe(I1) and Fe(II1) in a few supporting electrolytes have been measured by a number of experimenters (5,6,8,1b, 14). The agreement among the experiments has not been satisfying, and the discrcpancies have usually been accepted as a n inevitable result of the use of solid electrodcs n i t h their difficult-toreproduce surfaces. hloreorer, several authors were concerned primarily with thc development of the technique employed and were interested only incidentally in the electrode kinetics of the Fe(I1)-Fe(II1) couple. I n short, the data of various workers are not readily comparable, and the electrode kinetic parameters for the Fe(I1)Fe(II1) couple cannot be regarded as well establishcd. Experiments in these laboratories have indicated t h a t exchange current densities a t platinum electrodes for the Fe(I1)-Fe(II1) couple can be reproduced to 1 5 to 10% if sufficient care is paid to solution purity and, especially, to the state of the electrode surface XCHANGE
with respect to the presence or absence of platinum oxides and of freshly formed, finely divided platinum metal (1). In this report heterogeneous rate constants for Fe(I1)-Fe(II1) solutions in sulfuric and perchloric acids have been evaluated with a galvanostatic technique similar to that of Berzins and Delahag (3). EXPERIMENTAL
X constant current in the form of a step function was passed betxeen platinum working arid auxiliary electrodes in Fe(I1)-Fe(II1) solutions. The potential of the working electrode with respect to a third, nonpolarized platinum .reference electrode was observed as a function of time on a Tektronix Model 536 cathode ray oscilloscope (CRO) equipped with a Type D differential preamplifier having a maximum sensitivity of 1 mv. per em. and a Type T time base generator. The CRO was triggered by the voltage being observed. The point at which the trace crossed the central, graduated voltage axis on the CRO screen was observed visually and recorded. This voltage consisted of the overpotential, the ohmic potential drop, and the potential difference due t o concentration polarization corresponding to the particular current density and time for each experiment. The recorded voltage was corrected for ohmic and concentration polarization contributions, and the exchange current density corresponding to the prevailing current density and Fe(I1) and Fe(II1) concentrations mas calculated from the corrected voltage. Apparatus a n d Procedure. T h e constant current source consisted of 300 to 400 volts dropped across a bank of resistors in series with t h e cell. T h e voltage drop across t h e cell never exceeded 2 volts. T h e voltage was supplied b y a regulated
variable power supply (General Radio Co., Cambridge, Mass.). T h e current passed either through a d u m m y resistance t h a t was matched t o the cell resistance or through the cell, depending on the position of a mercury reliiy (Western Electric Co., Type 275B). T o eliminate switching transients associated with toggle switches, the relay was activated by charging a capacitor through a resistor selected so that several hundred milliseconds elapsed between the time the toggle switch was thrown and the mercury relay was activated. The s~$itching transients resulting from the mercury relay decayed in less than 3 psec. T o reduce transients associated with the power supply, a procedure used by Mattsson and Bockris (10) was employed: A resistance of 25,000 to 50,000 ohms was placed parallel to the cell and the dummy resistance in such a way t h a t the current could pass through this resistor even during the instant when the moving relay pole was not in contact n i t h either fived contact. The presence of the resistor caused the current to be somewhat less constant as the voltage drop across the cell varied; but all measurements were completed in a few milliseconds before any significant change in the current had occurred. The voltage and time scales of the CRO were calibrated and checked regularly. The calibrations remained constant within =t3% throughout the course of these experiments (12 months). Chronopotentiometric transition times of 0.2- to 1-second duration were measured with the CRO for use in calculating concentration polarization contributions to the measured overvoltages. The accuracy of the measurements of 7 with the CRO was only about +57,, but this was adequate, inasmuch as concentration polarizaVOL. 33, NO. 7, JUNE 1961
939
