920
ANALYTICAL CHEMISTRY, VOL. 50, NO. 7, JUNE 1978
(4) R. BrdiEka, Collect. Czech. Chem. Commun.. 12, 522-540 (1947). (5) G. A. Tedoradze et al., "Progress in Electrochemistry of Organic Compounds", A. N. Frumkin and A. B. Ershler, Ed.. Plenum Press, New York, N.Y., 1971, pp 171 ff. (6) E. Laviron, J . Electroanal. Chem. InterfacialEkctrochem., 52, 355, 1974. (7) D. G. Davis and F. R. Martin, J. Am. Chem. Soc., 88, 1365-1371 (1966). (8) T. M. Bednarskiand J. Jordan, J , Am. Chem. Sw., 89, 1552-1558 (1967). (9) R. BrdiEka and K. Wiesner, Vestn. Kral. Cesk, SpolNauk, 1-25, (1943). (10) E. Ponder, "Hemolysis and Related Phenomenon", 2nd ed., Grune and Stratton, New York, N.Y., 1971.
(11) R. BrdiEka and K. Wiesner, Collect Czech. Chem. Commun., 12, 39-63 (1947). (12) P. Delahay and C. Fike, J . Am. Chem. SOC.. 80, 2628-2630 (1958). (13) L. Meites, "Polarographic Techniques", 2nd ed., Interscience Publishers, New York, N.Y., 1965.
for review August
197'i. Accepted February
33
1978.
Heme Catalyzed Reduction of Oxygen and Hydrogen Peroxide at a Mercury Electrode Surface C. F. Kolpin and H. S. Swofford, Jr." Department of Chemistry, University of Minnesota, Minneapolis, Minnesota 55455
A model has been proposed which describes the electrocatalytic reduction of 0, and H,O, in heme-containing aqueous ethanol solutions. This model, in contrast to the "reaction layer" concept originally proposed by BrdiEka, has been based solely on the interaction of the adsorbed heme with diffusing O2 and H,O,. The rate limiting steps in the overall reduction processes are found to be the formation of the adsorbed Fe( 11) heme-0, and Fe( 11) heme-H,O, complexes, respectively. The rate constants for these processes are evaluated using chronopotentiometric and linear scan voltammetric techniques. Our proposed model together with these rate constants (determined from the above mentioned stationary electrode investigations using an HMDE) are shown to precisely describe all aspects of this electrocatalytic reduction system which have been observed in classical DME polarographic investigations.
Iron compounds have long been known to function as catalysts for hydrogen peroxide decomposition in aqueous solution. With free iron, this reaction proceeds via a free radical mechanism which results in the formation of oxygen and water from hydrogen peroxide ( 1 ) . Iron porphyrins in solution display the same net catalytic behavior toward hydrogen peroxide, and the heme group is known to be the active site for peroxide decomposition by the enzyme catalase (2). I t has been pointed out that heme rapidly undergoes both destructive and nondestructive oxidation in the presence of hydrogen peroxide ( 3 ) . Biliverdin, a nonmetal-containing tetrapyrrole chain, is reported to be the first degradative oxidation product ( 4 ) . Although destructive oxidation does occur, we have been unable to provide any experimental evidence to support the formation of the dark green biliverdin compound. Several investigations of the electrocatalytic behavior of iron porphyrin compounds in the reduction of hydrogen peroxide were conducted in the 1930's and 1940's by BrdiEka and co-workers (5-7). BrdiEka and Wiesner coauthored a classic paper concerning this process in 1947 (7). This paper introduced the "reaction layer concept" in which the catalytic process is treated as occurring in a thin layer of solution near the electrode surface. BrdiEka's theory was thought to describe 0003-2700/78/0350-0920$0 1.OO/O
both the catalytic current and reaction kinetics of this system. A reinvestigation of this catalytic process was conducted by V. Hanus in the 1950's (8). A review article by BrdiEka states that Hanus continued to work with the concept of a reaction layer of fixed solution volume (9). Phthalocyanine complexes of iron have also been used as catalysts for oxygen reduction ( 1 0 , I I ) . Kozowa, using graphite electrodes, reported that the iron atom, rather than a site on the macrocyclic ring, served as the active electrocatalytic site (10). Alt and co-workers published a study of oxygen activation by these catalysts based on molecular orbital considerations (12). They found that a conjugated T electron system seems to be a prerequisite for oxidation of the oxygen molecule. In 1957 Brezina reported that gelatin, when added to a 1 FM solution of Fe(II1) (hydroxocomplex) in 0.5 M KOH, resulted in a decrease in the wave for catalytic H202reduction (13). Brezina has studied the electrocatalyzed reduction of oxygen as well as hydrogen peroxide. Using graphite and carbon paste electrodes, he showed that the addition of heme also has a catalytic effect on the reduction of oxygen (14). In 1973, he published the results of his comprehensive study for electrocatalytic oxygen and hydrogen peroxide reduction using carbon paste electrodes (15). The remainder of this paper is devoted to the formulation of a model for the electrocatalytic reduction of O2 and H 2 0 2 a t a mercury electrode in solutions containing heme. The model for this system will be described and the results of supportive experimental work will be presented. These results will be shown to be entirely consistent with our proposed model. The rate constants for the two catalytic reduction processes are calculated on the basis of the formulated model. Finally, the results obtained, using this model system, will be shown to describe precisely and completely the data obtained from BrdiEka's original polarograpic investigation of the electrocatalytic reduction of H 2 0 2in heme-containing solutions (7).
EXPERIMENTAL Equipment. Electrochemical experiments were performed with the aid of a PAR (model 174) Polarographic Analyzer and a
multipurpose electrochemical instrument constructed in house (16). The multipurpose electrochemical instrument was utilized for all chronopotentiometricexperiments and voltammetric work where scan rates greater than 500 mV/s were required. C 1978 American Chemical Society
ANALYTICAL CHEMISTRY, VOL. 50, NO. 7, JUNE 1978
A Hewlett-Packard X-Y recorder (model 2D-2M) and a Tektronix storage oscilloscope (type 564) were used for the collection of experimental data. The storage oscilloscope, equipped with a time base (type 3B3) and a dual trace voltage amplifier (type 3A72),was employed for chronopotentiometric investigations in which the transition times were short. Cyclic voltammetry conducted a t scan rates greater than 200 mV/s was monitored using the storage oscilloscope equipped with the 3A72 voltage amplifier and a dual trace voltage amplifier (type 3A1). Voltages and currents applied to the electrochemical cell were measured with a Dana Digital Multimeter (model 3800A). All electrochemical investigations were conducted using a three-electrode system. All electrochemical potentials reported in this work have been corrected for IR drop. A Barnstead Conductance Bridge (model PM-70CB) was used for all cell resistance measurements. A saturated calomel electrode (SCE) was used as the reference electrode for d experimental work. “Asbestos wick” bridges were used throughout these investigations. A large platinum gauze cylindrical electrode was used as the auxiliary electrode in all cases. Both a dropping mercury electrode (DME) and a hanging mercury drop electrode (HMDE) were employed as indicating electrodes in these investigations (for capillary characteristics, see ref. 17). Both the hanging drop and dropping capillaries were of the “blunt-end” type. The reader is directed to ref. 17 for the basic cell designs used and the provision made for their temperature control. Spectra of Fe(I1) and Fe(II1) heme solutions were recorded with a Beckman (model 24) double-beam scanning spectrophotometer. Glass equipment was used in all experimental work. Reagents. Soluents. The reader is directed to ref. 17 for the specifics concerning the water and ethanol used in this investigation. Acetone (solvent grade) was also obtained from the University of Minnesota Chemical Storehouse and was used without further purification. B u f f e r Solutions and Supporting Electrolyte. All buffer solutions were prepared from reagent grade chemicals. The analytical concentration of the buffering components was 0.5 M in all stock solutions. Chemicals which were used in the preparation of buffer and supporting electrolyte solutions are listed by manufacturer below. NaH2P04: MW 136.09 (Merck Chemical Co.) NaHC03: MW 84.01 (Mallinckrodt Chemical Works) Na2C03: MW 105.99 (J. T. Baker Chemical Co.) Na2C204:MW 134.00 (Matheson, Coleman and Bell Manufacturing Chemists) Na2S04: MW 142.04 (J. T. Baker Chemical Co.) HN03: 14.7 M (Dupont Co. Industrial and Biochemicals Dept.) HCzH3O2:Glacial (Dupont Co. Industrial and Biochemicals Dept.) NaOH: MW 40.00 (Matheson, Coleman and Bell Manufacturing Chemists) KHC8H404: MW 204.23 (Mallinckrodt Chemical Works) KNOB: MW 101.11 (Baker and Adamson) All solutions subjected to electrochemical investigation were 0.1 M in KNOBand 0.05 M in buffer with the exception of those solutions in which the pH was controlled directly by the addition of strong acid or strong base. The molar concentration of sulfate buffer used in the aqueous ethanol solutions was also necessarily less than 0.05 M due to the decreased solubility of sulfate salts in these solvent mixtures. Other Chemicals. Mercury: (Bethlehem Triple Distilled Instrument Mercury) Na2S204:Technical grade (Fisher Scientific Co.). This material was used only as a reducing agent to convert Fe(II1) heme to Fe(I1) heme for the purpose of obtaining spectra of Fe(I1) heme in various solvent mixtures. 30% H202: C.P. with 0.05% Na4P207.10H20added as a preservative (J. T. Baker Chemical Co.) Heme: Equine Type 111; MW 652 (Sigma Chemical Co.). This material was stored in a refrigerator and was used as received. Crystalline heme is somewhat hygroscopic. However, analyses consistently yielded values of 96 to 98% heme by weight.
RESULTS A N D D I S C U S S I O N T h e salient features of the proposed model for the
921
heme-catalyzed reduction of O2 and H202to H 2 0 are as follows: (1)the high degree of “activity” of the electrocatalytic center (Le., the heme iron) is a result of the adsorption of the heme molecule at the mercury electrode surface; (2) the heme iron, in its divalent oxidation state, is the electrocatalytically active site; (3) the coordination of oxygen (or Hz02)to the adsorbed Fe(I1) heme is the rate limiting step in the overall catalytic reduction process (Le., the rate of electron transfer for the catalytic reduction of O2 and H202is very fast relative to the rate of ligand coordination to the heme iron): (4) oxygen and hydrogen peroxide are each reduced directly to water by a single four-electron, and a single two-electron step, respectively. A scheme illustrating the individual steps in the overall process may be depicted as follows: (diffusion)
Fe(II1)-OH heme (bulk) -----
-
Fe(II1)-OH heme (ads) 0, (bulk)
(diffusion)
(a)
0, (surface)
Fe(II1)-OH heme (ads) + H’
+ le- f%
Fe(I1)-H,O heme (ads)
(c)
Fe(I1)-H,O heme (ads) + 0, (sur)
5
Fe(I1) heme-0, (ads) + H,O
+
+
Fe(I1) heme-O,(ads) 4H’ 4eFe(I1) heme-H,O ads + HzO
(d )
2 (e)
This model predicts that the rates for the electrocatalytic reductions should be describable by the following expressions:
Rate = ho, (Fe(I1) heme) adsorbed (0,)surface (1) Rate = h H Z O (Fe(I1) , heme) adsorbed ( H 2 0 2 )surface
(2)
Catalytic Reduction of Oxygen. Chronopotentiornetric Znuestigations o f Catalytic O2 Reduction. The catalytic reduction of oxygen was investigated by chronopotentiometry. The HMDE covered with a complete monolayer of adsorbed heme (7.6 X lo-” mol/cm2) was used as the working electrode in these investigations (17). In this electrochemical configuration, d i f f u s i o n of O2 from the bulk of the solution to the electrode surface is the sole means of oxygen supply. The chronopotentiometric transition time is therefore controlled by this diffusing species in accordance with the Sand equation. The rate a t which the electrocatalytic reduction process occurs (determined by the magnitude of the applied current density) determines the amount of adsorbed heme which must exist in the Fe(I1) state, and hence the potential a t which the electrochemical reaction proceeds during the course of the reduction process. The ability t o constrain the rate of the electrocatalytic reduction process and to simultaneously monitor the value of the potential a t which the catalytic process occurs provides extensive information concerning the nature of this electrocatalytic reduction. The Sand equation for dissolved/diffusing oxygen, and the Nernst expression for adsorbed heme
( E = E ’ F e ( I I ) H / F e ( I I I ) H adsorbed -!- (RT / n F ) In [(Fe(III) heme) adsorbed/(Fe(I1) heme) adsorbed]) may be combined with the second-order rate law for the
922
ANALYTICAL CHEMISTRY, VOL. 50, NO. 7 , JUNE 1978
Table I. Observed and Predicted Potentials for a Typical Chronopotentiogram Resulting from the Electrocatalytic Reduction of 0, at a Heme-Covered Mercury Electrode Surface (see Figure 1)“
-0.40
co*
surface/ t
w
CO, bulk
tl/Z
6-030
0
0
1
03
0.24 0.44 0.64 0.744 0.84 1.04 1.24 1.44 1.64 1.84 2.04 2.24 2.44 2.64
0.49 0.66 0.80 0.86 0.91 1.02
0.72 0.62 0.54 0.50 0.47 0.41 0.36 0.30 0.26 0.22 0.17 0.13
ul 9 h
2-020 0
>
v
w -0J 0
1
0.0 0 0.0
I
1.0
I
I
2.0
3.0
‘
T
E - E,,, , mV Predicted Observed +18
+16
+9 +5 +2
+11
+6 +2
0
-2 -5 1.11 -8 1.20 -13 1.28 - 17 1.36 - 22 1.43 - 27 1.50 - 34 1.56 0.10 - 43 1.62 0.060 -55 = 2.98 s; E ’ F ~ ( I I ) , F ~ (=I I-0.440; I) E,,,
0
-2 -6 -10 -14 - 19 - 25 - 30 - 39 - 49 - 62 = -0.253 V.
TIME (seconds)
Figure 1. Chronopotentiogram resulting from the electrocatalytic reduction of O2to H,O at a HMDE, which is fully covered with heme, under the following experimental conditions: 60‘YO ethanol/H,O, i = 6 FA, i T ” * = 10.2, E’,,, = -0.440 V, electrode area = 0.041 cm‘, t = 25 O C electrocatalytic reduction process (Equation 1) to yield the following expressions a t 25 “C.
Table 11. Chronopotentiometric Quarter-Wave Potential for the Electrocatalytic Reduction of Oxygen as a Function of Applied Current (PA)at Various Oxygen Concentrations (i.e., i T 1 ~ z ) a i, P A
-Ell4 + E’s,,fa,, = 0.059 log i + 0.059 log i ~ l / +* constant AE1/4
=
log
1
-- -
7112
E1/qa - E 1 / d b
ibTb112)
= -0.059 (log i , ~ , ” ~-
- 0.59
k o 2 ( 7 . 6X
(3a)
log ‘“lib
2.0 3.0 4.5 7.5 9.0 11.5 18.0
1 log -p= constant -
E‘,,,)/0.059]
(3d)
The general shape of the chronopotentiograms resulting from the electrocatalytic reduction of oxygen (see Figure 1) may be precisely predicted on the basis of the restrictions imposed on this system by the second-order reaction process which determines the surface concentrations of diffusing oxygen and adsorbed Fe(I1) heme as a function of time (see Table I). From Table I, it is apparent that the experimentally determined potentials do indeed compare closely to the values predicted from the second-order rate law. The fraction of the bulk concentration of O2 which remains a t the electrode
2.28 2.23 2.36 2.20 2.21 2.38 2.55
/
~
0.30
0.48 0.65 0.88 0.96 1.06 1.26
~
~
0.27 4 0.285 0.300 0.310 0.320 0.328 0.345
( b ) ir”2av = 8.2 (nominally air saturated) 5 7.5 10.0 14.0 22.0 34.0 48.0 92.0 125.0 200.0
3.15 1.2e 0.72 0.28 0.1 26 0.050 0.024 0.0082 0.0052 0.0022
25 37.5 50 110 240 628
2.5 1.12 0.60 0.116 0.022 0.0033 0.0014
8.85 8.50 8.50 7.40 7.80 7.60 7.40 8.30 8.20 9.40
0.70 0.87 1.15 1.34 1.53 1.68 1.96 2.10 2.30
0.258 0.266 0.281 0.296 0.310 0.324 0.340 0.367 0.379 0.400
1.40 1.57 1.70 2.04 2.38 2.80 3.00
0.265 0.281 0.293 0.319 0.349 0.376 0.402
1.00
( c ) irl”, = 37.8
1000 [(E114-
-E,,,
log i
~ = 2.3~
1.3 0.55 0.275 0.086 0.060 0.043 0.020
(3b)
lo-”)
ir”2
s
(a) i
(3) From Equation 3 above, Equations 3a-d follow directly when 10(E114-E‘8~)/0.059 is much larger than 1.
7,
39.5 39.7 38.6 37.5 35.6 36.0 37.4
Fixed electrode area; E’,,,
=
-0.448 V; 60% ethanol/
H,O; 25 “C.
surface as a function of time may be expressed as the quantity (1- [ t 1 i z / ~ 1 / 2When ]). the difference between ( t / ~ ) ’and / ~ one becomes small, the relative error in the value of the potential increases substantially and some deviation is observed. The experimental results obtained from our chronopotentiometric investigation of this electrocatalytic reduction process are tabulated and presented in Table 11. The data shown in Table I1 clearly illustrate the predicted linear relationship of as a function of log i (see Equation
ANALYTICAL CHEMISTRY, VOL. 50, NO. 7, JUNE 1978
Table IV. Results of the Chronopotentiometric Investigation of the Electrocatalytic Reduction of 0, in the Presence of Hemen i7 1 1 2 irl/2 = E’sw = E’s,, = 2.3 log -0.448 8.2 log -0.448
Table 111. Observed and Predicted Values for the Separation of -E,,, vs. log i Plots as a Function of Changes in Solution Oxygen Concentration‘ log h1l2
Separation, mV Observed Predicted
2.3 41
33
46
39
8.2 37.8 Fixed electrode area; E’,,, H,O; 25 “C.
= -0.448
923
V; 60% ethanol/
El,,
(l / T 1 ’ * )
El,,
-0.056 0.130 0.280 0.533 0.611 0.683 0.850
-0.274 -0.285 -0.30 -0.310 -0.320 -0.328 -0.345
-0.25 - 0.054 0.071 0.276 0.450 0.651 0.810 1.04 1.14 1.33
-0.258 -0.266 -0.281 -0.296 -0.310 -0.324 -0.340 -0.367 -0.379 -0.400
iTl/237.8 log
E’s,, = -
iT1’2= 13.1 log
E’sw = -0.432
0.448
(1/T1’2)
El,,
-0.199 -0.024 0.110 0.468 0.829 1.24 1.43
-0.265 -0.381 -0.293 -0.319 - 0.349 - 0.376 -0.402
(l/T’”)
(l/T1”)
El,,
-0.08 0.08
-0.255 -0.264 -0.284 -0.313 -0.327
0.42
0.73 0.91
a Values of E , , , are tabulated as a function of log ( 1 / ~ ”at~ fixed ) values of iT1’2and E’,,, . 60% ethanol/ H,O; 25 “C.
05
0.0
0.5
I.o
1 5
Log I& Figure 2. Values of E,,, chrono vs. log l / ~ ’ ” obtained using an HMDE which is fully covered with adsorbed heme (see Table 11). (A)E’,,, = -0.448;is’/*= 2.3.(0) €Isurlaoe = -0.448;iT1” = 8.2.(0) E’slrface
= -0.432;iT”2 = 13.1.(0) €Iwrtaca = -0.448;is’/2= 37.8.Electrode
area = 0.041 cm2 3a). The separation of these lines (generated at three separate oxygen concentrations) in millivolts vs. the values predicted by Equation 3b are as follows (see Table 111). The data from Table I1 may be reduced to a single straight line (expressed by Equation 3d). These results, as well as the equal to results of an additional investigation (with -0.432 V), are presented in Table IV and Figure 2. These results (as illustrated in Figure 2) show that a linear relationship between the experimental values for EIl4and log ( l / ~ ’ ’does ~ ) indeed exist. The slope of this plot is determined to be 65 mV. Furthermore, the plot of experimentally determined values of E114 (obtained with a value of E’,&,, equal to -0.432 V) is observed to be approximately 16 mV positive of the data obtained a t an value of -0.448 V. Some upward curvature of the plot (Figure 2) is observed as the values for the transition time decrease (the reason for this is unknown). I t is believed that this curvature may occur when the thickness of the established diffusion layer becomes comparable to the average distance between the active catalytic centers on the mercury electrode surface. However, this observed deviation is not germane to the interpretation of the data. The rate constant for this electrocatalytic reduction process is evaluated a t relatively long values of the transition ~ ) between -0.3 and time (Le., where values of log ( l / ~ l /are 0.3) where the data in Figure 2 displays linear behavior. The rate constant (Izo,) can be calculated directly from Equation 3c, once the value of the diffusion coefficient for dissolved 02, in the 60% ethanol/H20 solvent, is determined. The diffusion coefficient for O2 in ethanol/H20 solvent mixtures was not determined directly because of the difficulty
of independently determining the concentration of O2 dissolved in these solvent mixtures. Instead, the changes in the experimentally determined values of the diffusion coefficients for Cd(I1) and Pb(I1) were measured in 60% ethanol/H20 (relative to their values in pure H 2 0 ) ; changes in the value of the diffusion coefficient for O2 were assumed to parallel this behavior. Solutions containing 5 x lo4 M of either Cd(I1) or Pb(I1) in 60% ethanol/H20 displayed polarographic diffusion-limited currents equal to 60% of the values observed for these species, at the same molar concentration, in aqueous media. This implies that the value of the diffusion coefficient for each of these ions is larger, by a factor of 2.77, in an aqueous media when compared to 60% ethanol/H20 as a solvent. The ratio of the absolute viscosity for the 60% ethanol/H20 solvent to that for pure H 2 0 a t 25 “C is equal to 2.78 (18). As a result of the correlation of the changes in the experimentally determined values for the diffusion coefficients of Cd(I1) and Pb(I1) in the ethanol/H20 solvent with the known change in the viscosity of the solvent relative to H 2 0 ,the assumed change for the diffusion coefficient for O2 in ethanol/H20 mixtures is justified. These experimental data are in full agreement with the results predicted by the Stokes-Ilkovic Equation (19). The diffusion coefficient for dissolved O2 in H 2 0 is equal to 2.6 x cm2/s a t 25 “C (20). The value of this diffusion coefficient in the 60% (by volume) ethanol/H20 solvent, used in this investigation, is therefore calculated to be 9.4 X lo+ cm2/s (2.6 X 10-’/2.77). The rate constant for the electrocatalytic reduction of oxygen is then determined by direct substitution of known quantities into Equation 3 c , and the rate constant (koJ is calculated to be 5.9 X lo7 L/mol s. This rate constant is the fundamental second-order rate constant for the formation of the adsorbed Fe(I1) heme-oxygen complex from dissolved O2 and adsorbed Fe(I1) heme. Other Elecrochemical Investigations of t h e H e m e Catalyzed Reduction of Oxygen. The model system which has been proposed to describe the electrocatalytic reduction of oxygen was further tested using linear scan voltammetry. The equation which describes the value of the peak current for
924
ANALYTICAL CHEMISTRY, VOL. 50, NO. 7, JUNE 1978
of the peak potential is predicted using Equation 1 and Equation 4 to obtain an expression analogous to Equation 3d (see Equation 5 ) .
Table V. Experimentally Determined Values of the Peak Current (Linear Scan) for the Electrocatalytic Reduction of 0, as a Function of the Potential Sweep Ratea
a
v , VIS
ip P A
0.2 0.1 0.05 0.02
13.0 9.46 6.76 4.44
i,,,,
-0.0295 log u =
iplu112
peak
11.0 8.04 5.7 5 3.77
29.0 29.9 30.2 31.4
HMDE (fixed electrode area); nominally air saturated
60% ethanol/H,O; pH 11.0; EfsU= -0.440.
the diffusion controlled electrocatalytic reduction of oxygen as a function of the voltage sweep rate ( u ) is as follows:
ip = 2.72 at 25
x io5 noznFe3'zAD1/2Co, BdkV1" "C
(4)
The two values of n (the number of electrons transferred) in Equation 4 are a result of the fact that oxygen reduction (a four-electron process) governs the magnitude of the peak current while the dependence of heme iron reduction on potential (a one-electron process) controls the number of catalytically active sites which are present on the electrode surface. The rate of change in the number of active catalytic sites (a one-electron process) controls the shape of the rising portion of this wave in accordance with the argument originally used by BrdiEka (7) in his derivations concerning the classical polarographical behavior of the heme/HzO2 electrocatalytic reduction process. Because of the above discussed effect of adsorbed heme iron, linear scan voltammograms (for which the values of the half peak potential and the peak potential are separated by 56 mV a t 25 "C) should be observed for this catalytic oxygen reduction process. The ultimate diffusion control of the peak current (dissolved bulk solution oxygen) should yield a linear relationship between the value of ipeak and u1J2 for any fixed value of the dissolved oxygen concentration. For this diffusion-controlled process, the surface concentration of dissolved oxygen is equal to one half of the bulk concentration for this species a t the point on the rising portion of the linear scan voltammogram where the current is equal to 85% of its peak value (analogous to the positions where t = (7/4)in chronopotentiometry and i = id/2 in classical polarography). Since the value of ipeak (hence io.85 peak) is proportional t o u 1 / 2 , the instantaneous rate of the electrocatalytic reduction process a t the 85% point is also proportional to v1f2. Since the surface concentration of oxygen, which corresponds to the point on the linear scan voltammogram where the current is equal to 85% of its peak value, is fixed and is equal to one half the concentration of oxygen in the bulk solution, the concentration of adsorbed Fe(I1) heme must change by a value proportional to v1/2 in accordance with the second-order rate law (Equation 1). This change is manifested by a 29.5-mV change in the value of Eo.85 ped(as for each corresponding decade well as El/z peak and change in the rate of the potential scan. This shift in the value
peak
- E',,,
+ constant
(5)
Equation 5 predicts that a plot of log u vs. Eo,85ped should be linear with a slope of 29.5 mV a t 25 OC. The experimental results obtained from the investigation of this electrocatalytic reduction process by linear scan voltammetry are tabulated in Tables V and VI. The results presented in Table V show that the quantity ip/ul/* is indeed reasonably constant. The results which are presented in Table VI clearly illustrate many of the theoretically predicted properties for this electrocatalytic reduction of oxygen. These include the following: (1)the measured value of the slope of a plot of the experimentally determined values for El/z peak, E0.85 peak, or Epeak vs. log u is determined to be between 26 and 28 mV per decade change in u which compares favorably to the value of 29.5 mV which is predicted by theory; (2) the separation between the experimentally determined values for E p e a k and E l l z peak is in all cases between 56 and 58 mV which is in excellent agreement with the theoretical value predicted for a wave whose shape is governed by a one-electron process (i.e., the reduction of the surface adsorbed heme iron to form the active catalyst); and (3) the values for Eo,&pedpredicted from the value for El/*chno (Figure l),the constant applied current for the chronopotentiometric process, and the experimentally measured value for the peak current from linear scan voltammetry are found to be in excellent agreement with the experimentally measured values of this quantity. The value of Eo35 peak is predicted from Equation 6. lchrono --
0.85 ip r
\l
1
The rate constant for the electrocatalytic reduction process may also be calculated from linear scan voltammetric data using Equation 7 which is analogous to Equation 3c presented in the section devoted to the chronopotentiometric investigation of this same process.
0.85 ipeak n 0 p
Table VI. Experimentally Determined Values for the Electrocatalytic Reduction Potentials for 0, as a Function of the Potential Sweep &teasb v
log v
0.2
-0.70 -1.00 -1.30 -1.70
.1 0.05 0.02
p
Eo.,,p
-0.240 -0.231 -0.222 -0.214
- 0.269
Ell,
-0.259 - 0.249 -0.242
-0.298 -0.288 -0.278 -0.270
0.058 0.057 0.056 0.056
- 0.269
-0.260 -0.252 -0.240
a HMDE (fixed electrode area); nominally air saturated 60% ethanol/H,O; pH 11.0; E',,, = -0.440. NOTE: All experimental data which are present in Tables V and VI were collected using the same mercury drop. This same mercury drop was also used to yield the chronopotentiogram shown in Figure 1.
ANALYTICAL CHEMISTRY, VOL. 50, NO. 7, JUNE 1978
Table VII. Calculated Values of the Second-Order Rate Constant ( h o , )from Linear Scan Voltammetric Dataa V
0.2 0.1 0.05 0.02
- 0.269
6.8
I
x lo7
4.0
7.1 X l o 7 7 . 4 x 107
-0.259 -0.249 -0.242 AV
a
b
‘to,
peak
2D
6.2 x i o 7 6.9 x 107
h
v)
a
E
E’mr= -0.440 V ; 60% ethanol/H,O.
0.0
0 ,
The results for this determination of the rate constant (ha) are presented in Table VII. Equation 7 also predicts that, for any fixed potential scan peak will not be a function of the concentration of rate, oxygen in the bulk of the solution. This predicted independence of the values of Eo.ss peak (as well as Ell2peak and Epeak),as the oxygen concentration was varied, was observed in all linear scan voltammetric investigations of this system. I t should be pointed out that the difference between the value of ko, determined by linear scan and that found by chronopotentiometry (5.9 x IO7) represents a maximum deviation of only 4 mV in the potential for the electrocatalytic process. These values for ko, are therefore considered to be within the limits of experimental error. These data, obtained by linear scan voltammetry, lend further support to the proposed model for the electrocatalytic reduction of oxygen. Further investigations of this electrocatalytic reduction process using a stationary electrode (the HMDE) were conducted in stirred solutions a t slow potential sweep rates. Current-voltage curves which displayed stable, diffusionlimited currents (Le., resembling “classical” polarographic waves) were obtained. Analysis of these experimental data was conducted using a Nernst diffusion layer treatment. The experimentally observed positions of E l l 2 Cat for these waves (as a function of id) are in complete agreement with the results predicted by chronopotentiometry and linear scan voltammetry. Another recalculation of the numerical value of ko2 was not carried out using these data. The reported values for all electrocatalytic reduction potentials (using a stationary electrode) were obtained when the electrode was not allowed to “rest” for periods of time longer than 1 min a t potentials positive to the reduction potential for the surface-adsorbed heme iron. Allowing the electrode to rest for longer than several seconds results in an observable negative shift of the potential for the electrocatalytic reduction wave. The magnitude of this “shift” reaches a maximum value of 10 mV after several minutes (zero current flow) in the 60% EtOH/H20 mixtures used in this work. In the absence of cathodic current flow for extended periods of time, the first cathodic half-cycle of a cyclic voltammogram (obtained from a deaerated solution of heme) displays a nonsymmetrical peak for the reduction of adsorbed heme iron. This peak potential is also shifted negative of its “normal value” by an amount similar to that observed for the corresponding electrocatalytic oxygen reduction peak. After this first reduction half-cycle, the surface peak regains its reversible appearance (i.e., the nearly zero mV peak to peak separation and a symmetrical shape) and returns to its “normal reversible potential” for all successive cycles. I t is believed that, during these periods, the orientation of the adsorbed layer of heme slowly changes in the absence of current flow, and that upon passage of a current, the stable, reversible nature of the adsorbed layer is regained. Experimental data for the electrocatalytic reduction process, taken immediately after long “rest” periods with a stationary electrode, are reproducible, but they are no longer representative of the electrochemical system incurred in classical
925
3 - -2.0 - 4.0
-&O
-eo
t
d
I
I
I
- 0.5 0
- 0.30
-0.10
E ( v o l t s ) vs.
-0.70
-0.90
S.C.E
Figure 3. Cyclic voltammogram resulting from the presence of heme in a deaerated 5 X M heme solution under the following experimental conditions: 60% ethanol/H,O, pH 11, 25 ‘C, v = 1.5 VIS, electrode area = 0.041 cm2. a = E, SU,faCw, b = E, c = Ea diffusionr
=
€e, surlacev
e =
E’surfacwv
=
polarography
DME polarography (Le., a totally reversible reduction process for the adsorbed heme iron). Because of the observed loss of symmetry for the current peaks resulting from the reduction of adsorbed heme iron, these potentials were not used for the evaluations of ko, conducted using the stationary HMDE. The electrocatalytic reduction of oxygen, a t an electrode covered with a monolayer of adsorbed heme, was further investigated by classical polarography using a DME. In order to achieve maximum surface coverage throughout virtually the entire lifetime of the drop, higher solution concentrations M to 5 X M). of heme were required (2 X BrdiEka (7) previously published a theory which he thought correctly described the observed heme catalyzed reduction of H202a t a DME. He was, however, totally unaware of the direct, four-electron, electrocatalytic reduction process for diffusing oxygen itself. His expression describing the polarographic current-voltage curves generated by the electrocatalytic reduction of H 2 0 2is nevertheless valid, after appropriate modifications are made to describe the current-voltage curves obtained for both the electrocatalytic reduction of O2 and H 2 0 2using the DME. BrdiEka’s original expression (7), which he believed fully described the electrocatalytic reduction of H202,is given by Equation 8.
0.059
1
log
x + nF
X Adh,20,(heme)PJulk
(8)
However, before BrdiEka’s expression can be properly applied to the true electrocatalytic reduction process, the following modifications must be made: (1) the half-wave potential for the diffusing heme Fe(II)h/Fe(III)h) must be replaced by the formal potential for the surface adsorbed heme (i.e., the value of E \ e a k surface shown in Figure 3); (2) the total solution heme concentration must be replaced by the total surface concentration of adsorbed heme in mol/cm2; (3) the thickness of the reaction layer ( d ) must be eliminated from the expression. Changing the value for the heme concentration from a bulk solution concentration to a surface quantity (i.e., a concen-
926
ANALYTICAL CHEMISTRY, VOL. 50, NO. 7, JUNE 1978
Table VIII. Experimentally Observed and Theoretically Predicted Values of E,,, cat for the Heme Catalyzed Reduction of Oxygen at a Fully Covered DMEa E ‘sur
Predicted
-0.412 -0.440 - 0.448
-0.205 -0.233 -0.241
‘2 x
Table IX. Experimentally Determined Values of E , , , cat as a Function of the Concentration of Dissolved 0, (proportional t o
E,,, cat 0bseriid -0.200 -0.227 -0.235
,,,
0.059 1 1% - icat 0.059 icatrnax -kat 1 log lSuT
=
(
~
X0,+H20
no2FAho,(heme) sur max
)
lo3
@A
Experimental
Predicted
0.78 2.9 5.3 24.5
-0.230 - 0.231 -0.231 - 0.235
-0.232 -0.232 -0.232 -0.232
‘Mercury column height = 58 cm; mechanically controlled drop time = 2 s; E’,,, = - 0.432.
tration with a definite, readily measurable upper limit which occurs far below the concentration level for solution saturation) then predicts that the half-wave potential for the electrocatalytic reduction process should continue to shift in the positive direction only until the electrode surface is fully covered with heme. This prediction is in direct conflict with BrdiEka’s reaction layer theory which predicts that no limit, other than that imposed by the solubility of the heme itself, should exist for the magnitude of the positive shift for this electrocatalytic wave. If the results obtained from the chronopotentiometric and linear scan voltammetric investigations of the electrocatalytic reduction of oxygen (at a fully covered HMDE) are valid, the rate constant (ko,) and the maximum electrode surface coverage by adsorbed heme obtained by these methods should reasonably predict the limiting positive shift of the half wave potential (El12 - E’,,,,); this affect is polarographically observable for this reduction process. The predicted value for the positive shift is obtained from Equation 9 for catalytic currents (after our modifications reflecting the surface controlled nature of the process are made) as follows:
Ecat- E
Ei / 2 cat
ld O,+H,O,
M heme; other conditions as stated in the
text.
id )“
(9)
-
A t the potential where E,,, = E I l 2cat, i,,, = i,,, max/2. xo2 ~~0 = 16.0 for 60% E t O H / H 2 0 and a corrected mercury column height of 7 2 cm. Making these substitutions, E l l 2cat -E’,,, = 207 mV a t 25 O C .
The validity of this predicted shift of the values for the potential was verified from our experimental results (see Table VIII). Plots of the log (i,,, mlu - icat/icat)vs. Ecatfor these waves display the predicted 59-mV slope at 25 “C. We have observed that no further positive shift in the value of Ell2cat occurs when the solution concentration of heme is increased beyond 2 x M (where full electrode coverage is observed). Actually a t heme concentrations above about 6 X lo-* M, the value of E l l z Cat begins to shift in the negatiue direction. This is probably a result of an interaction between the available surface catalytic sites and the increased quantity of diffusing heme. Equation 9 predicts that the value of ElI2Cat should not be a function of the concentration of dissolved oxygen in the bulk of the solution. The results of an investigation, conducted to confirm this prediction, are presented in Table IX. Under these experimental conditions, the collection of nonconcentration dependent terms in the Ilkovic Equation (x)is found to equal 12.4, and E l l 2 cat is therefore predicted
Table X. Half-Wave Potentials for the Electrocatalytic Reduction of Oxygen as a Function of pH PH
E,,, cat
E’sur
0.5 1.5 2.7 7.8
+0.189 +0.189
- 0.024’
9.0 11.0 12.0
+0.180b -0.062 -0.128 -0.237 -0.295
-0.024’ -O.024’jb -0.262 -0.330 -0.442 -0.497
diffusing heme
-0.127 + 0.01 -0.127 _c 0.01 -0.127 ?: 0.01 -0.415 - 0.482 -0.594 -0.651
a Implied from the reduction potentials f o r oxygen a n d Some precipitation of the diffusing Fe(II1) heme. Fe(II1) heme occurred in this solution.
to be 200 mV positive of E’s,r instead of the previous value (207 mV when x = 16). The potential for the electrocatalytic reduction of oxygen to water was investigated over the p H range of 0.5 to 12.0. Solutions of Fe(I1) heme (at all pH values in this range where it is soluble in the solvents used, Le., ethanol/H,O, acetone/H20, and H20),yielded the same absorption spectrum in the visible region of the electromagnetic spectrum. It is therefore believed that the divalent iron of the bulk solution heme is coordinated to H 2 0 at both axial positions in these solvent mixtures over the entire pH range under investigation. Furthermore, the adsorbed heme in an acetone/H20 solvent displayed the same value for the rate constant (ko,) which was observed in the ethanol/H20 solutions. The nature of the axial ligand of the adsorbed heme should certainly have an effect on koz since it is believed that O2 must displace this coordinated ligand before the electrocatalytic reduction process can occur. This suggests the coordination of H20 in the axial position of the adsorbed heme iron. The experimentally determined values of the polarographic half wave potentials for catalyzed O2 reduction (using the DME) as a function of pH are presented in Table X. T h e Heme-Catalyzed Electrochemical Reduction of H 2 0 2 We propose that the electrocatalytic reduction of H202, a t a mercury electrode surface in 60% ethanol/water solution, may be described by the same general model which was developed in the previous section to describe the electrocatalytic reduction of oxygen. The heme-catalyzed reduction of H 2 0 2differs from the case for oxygen reduction in the following ways: (1) The reduction of H 2 0 2is a two electron process. (2) The adsorbed heme-catalyst is destroyed by electrochemically produced H 2 0 2via a slow “side reaction”. (3) The value for the rate constant (kHzot)differs from that for the direct, four-electron, oxygen reduction process (ko,). The value of the rate constant for the H202electrocatalytic reduction process has been evaluated from linear scan voltammetric data, relative to k0,. Equation 7 (for 02) and an analogous expression for H 2 0 2(Equation 10)
ANALYTICAL CHEMISTRY, VOL 50, NO 7, JUNE 1978
1
1+
1 0 ( E ~ . 6 speak H , 0 , - E ' s ~ r ) ~ 0 ~ 0 5 9
1 ((2)(2.72 X 105)n3'2Fe01'2H,0,v1i210-3 are combined to yield Equation 11. +
1 1 0 ( E o . 6 5H , 0 , - E ' s u r ) / 0 . 0 5 9
1 H,o,
1+
0,-E'~U~)/0.059
If one approximates the ratio of D1J20z/D1/2Hzo2 as being equal to one, when the exponential terms are much greater than one Equation 11 may be reduced to the following:
The value of the rate constant for the hydrogen peroxide reduction process was calculated from experimental data obtained after the system was poised for long periods of time a t potentials where no observable current flowed through the mercury electrode. The potential for the electrocatalytic reduction of oxygen, under these conditions, shifts in the negative direction in the manner previously discussed. The values of for H202reduction are found to be 50 f 3 mV negative of the value of Epeak for the direct electrocatalytic reduction of oxygen under identical experimental conditions (at 25 "C). H 2 0 2spontaneously decomposes to O2 and H 2 0 in the presence of heme, and the subsequent reduction of oxygen formed by this side reaction (in these stationary electrode investigations) causes the rising portion of the H202 reduction peak to broaden. Therefore, the difference in the values of the peak potentials for these processes was used to evaluate k ~ as~a function 0 ~ of ho,. The rate constant for the hydrogen peroxide reduction process (kH102)is then calculated by Equation l l a and is found to be equal to 9 x lo6 L/mol s. This rate constant, calculated from linear scan voltammetric data, may be used to predict the value of Ell2for the electrocatalytic reduction of H 2 0 2in DME polarographic experiments. An expression (analogous to Equation 9 for oxygen reduction) describing the shape and the position of the H 2 0 2 catalytic reduction wave on the voltage axis (obtained using a DME which is covered to varying extents with adsorbed heme a t the end of drop life) may be written as follows:
XO, +H20,
j*O1+H,oz
+ nH,O,
F A ( t o t a 1 heme)sur 103kH20,
The experimentally observed values of the polarographic half-wave potentials for the electrocatalytic reduction of H202 at a DME which is partially covered with adsorbed heme and the predicted values for these half-wave potentials (based on Equation 12 and the value for k H 2 O 2 determined by linear scan voltammetry a t a fully covered mercury drop) are presented in Table XI. The values of the half-wave potential for the electrocatalytic reduction of H202(predicted by classical DME polarography with a partially covered electrode and by linear
927
scan voltammetry) a t a fully covered DME are presented in Table XII. These values of the half-wave potential (predicted for an electrode which is fully covered with a monolayer of heme) represent the potentials at which the positive shift of this catalytic H 2 0 2reduction wave will stop. This concept of a maximum attainable shift in the positive direction is again in direct contrast to the concept of a limitless (except for the solubility of the heme) positive shift for the H 2 0 zcatalytic reduction wave which is predicted using the reaction layer theory proposed by BrdiEka ( 7 ) . The results presented in Table XI clearly show the expected changes for the values of the half-wave potential although all experimental values of E , are approximately 10 mV more positive than the linear scan voltammetric data predicts. The fraction of the DME which is covered a t the end of drop life was evaluated from data collected from linear scan voltammetric experiments employing a dropping mercury electrode in deaerated solutions. The absolute position of El,* for the H 2 0 2 reduction is expected to be affected slightly by a buildup in the H202at the electrode surface early in drop life (hence, some H z 0 2would be expected to diffuse away from the electrode surface). As the catalytic efficiency of the growing drop surface increases with droplife, an H,02 concentration gradient favorable to back-diffusion arises. However, we believe that these gradients are sharpest very early in droplife (Le., when the rate of change of catalytic efficiency is greatest) and should manifest only a minor contribution to the instantaneous current a t the end of droplife. Although the El j 2 values obtained using a DME are not exactly the same as those predicted from the stationary electrode investigations, they are certainly close enough to lend strong experimental support to our proposed model system for these electrocatalytic reduction processes. The potentials in Table XI1 obtained from classical polarographic data and measured heme coverages at the DME again correlate (to within 15 mV) with those predicted from the linear scan voltammetric investigations using a fully covered HMDE. We have shown that the presence of H z 0 2at the electrode surface (generated by normal oxygen reduction at a mercury electrode surface which is free of adsorbed heme) prevents the adsorption of intact heme to the electrode from dilute aqueous ethanol solutions of heme (Le., 1 X lo-' M). The degradation of heme, already adsorbed a t the mercury electrode surface is illustrated by the non-reproducibility of successive current-voltage curves obtained with an HMDE from dilute heme containing solutions to which H202has been added. Because of the observed heme degradation problem (which is more severe with the HMDE than with the DME due to the extended lifetime for the hanging drop), the value of kHZo2 obtained from DME polarography at a partially covered drop is thought to be more reliable than the value calculated from linear scan voltammetric experiments a t a fully covered mercury drop. The values of hHzo,calculated from Equation 12 using the data obtained from the DME polarographic investigations on this system are presented in Table XIII. The value for h ~ ~ oobtained ,, from DME polarographic data, is found to be (1.4 f 0.2) X loi Limo1 s. It is noted that this value is somewhat larger than the value obtained from linear scan voltammetric data which is consistent with the proposed additional degradation of the adsorbed heme in the case of linear scan voltammetry using the HMDE. In solutions containing dissolved 02,diffusing oxygen also affects the observed overall electrocatalytic reduction process a t a partially covered drop in the DME polarographic investigations. When oxygen reaches a mercury electrode surface which is partially covered with adsorbed heme, one
928
ANALYTICAL CHEMISTRY, VOL. 50, NO. 7, JUNE 1978
Table XI. Predicted and Experimentally Observed Values of E,,, for the Electrocatalytic Reduction of H,O, Using a DME” Mercury column ht, cm
Ell,
XO,-+H,O,
Slope, mVb
Fr.
Experimental
COV.~
( a ) Bulk solution heme concentration = 5.4 x
44.8 57.8 74.0 93.0
6.3 7.1 8.0 9.0
44.8 57.8 74.0 93.0
6.3 7.1 8.0 9.0
58 62 62 63
0.085 0.072 0.060 0.052 =
1.7 x
-0.338 -0.345 -0.352 -0.359
M
-0.296 -0.305 -0.312 -0.319
0.27 0.23 0.188 0.175
54 54 57
M -0.331 -0.341 -0.347 -0.351
( b ) Bulk solution heme concentration 51
Predicted
0.308 -0.315 -0.324 -0.328 -
a E’,,, = -0.440 V; 60% ethanol/H,O; 25 ”C; nominally air saturated solutions. Slope (mV) is the value of the slope for Fr. cov. designates the fraction of the a plot of log (icat max - icat/icat) vs. Eapplied. The theoretical value is 59 mV. electrode surface which is covered with adsorbed heme at the end of drop life.
Table XII. Values of for the Catalytic Reduction of H,O, at a Fully Covered Dropping Mercury Electrode as a Function of the Mercury Column Height as Predicted from DME and HMDE Investigations‘
Mercury column ht, cm 44.8 57.8 74.0 93.0
[Hemelbulk = 5.4 x using a DME predicted
b
-0.268 -0.274 -0.275 -0.275
[from k H , O , evaluated by lin. scan volt.] using a HMDE
[Hemelbulk = 1.7 X using a DME
-
predicted
b
predicted
-0.262 -0.267 - 0.269 -0.274
-0.275 -0.278 -0.281 -0.284
These values are predicted from polarographic data by measuring E , , , H,O, and the fraction a 60% ethanol/H,O; 25 “C. of the electrode surface covered. The values of E , , , predicted are then calculated from the additional anodic shift of E,,, H - 0 . that is anticipated when the DME reaches maximum surface coverage with adsorbed heme. Table XIII. Experimentally Determined Values of from DME Polarographic Data
kH,O,
Mercury column ht, cm 44.8 57.8 74.0 93.0
Bulk solution heme concentration 5.4 x 1 0 - 6 ~ 1.7 x 10-5 M k H , O , Limo1 S ’H,O, 1.3 x 1 0 7 1.6 X l o 7 1.15 x 107 1.5 x 107 1.2 x 1 0 7 1.5 x 107 1.35 x 107 1.4 x 10’ AV 1.25 x 107 1.5 x 10’
of the following conditions exists: (1)When the electrode is poised a t a potential which is more positive than the reduction potentials for both O2 to H 2 0 2a t a “clean” mercury surface and for O2 to H 2 0 in the presence of adsorbed heme, no reduction occurs. (2) When the electrode is poised at a potential which is more positive than the potential where 0 2 may be electrocatalytically reduced to H20 but more negative than the reduction potential for O2 on “clean” mercury, virtually all the diffusing oxygen is reduced to H 2 0 za t the clean mercury surface. (3) When the electrode is poised a t a potential more negative than the value where both the normal and the catalytic processes could exhibit diffusion limited currents, these processes are experimentally shown t o co-reduce oxygen (via 2e- and 4e- pathways) in amounts which are proportional to the fraction of the electrode surface area which is “clean” and which is covered with a monolayer of adsorbed heme, respectively. Recall that the half-wave potential for the electrocatalytic reduction of oxygen is approximately 50 mV more positive than the half-wave potential for the H202reduction process. The presence of dissolved O2 in solution, therefore leads to
0.0
I
0.0
-0.10
-0.20
E
-030
-0.40
-050
-0.60
( v o l t s ) vs. S.C.E.
Figure 4. The assignment of the relative contributions to the diffusion limited polarographic currents resulting from the reduction of O2 at a DME in 5.4 X M heme solution as predicted from the stationary electrode investigations; 60% EtOH/H,O, pH 11.O (25 OC); mercury column height of 57.8 cm
the existence of two unresolved waves on the lower portion of the polarographic “catalytic H20zreduction wave”. This phenomenon becomes increasingly significant in the correct interpretation of these polarographic data as the fraction of the electrode surface covered by adsorbed heme increases (see Figure 4).
ANALYTICAL CHEMISTRY, VOL 50, NO 7, JUNE 1978
In conclusion, the adsorbed heme at the mercury electrode surface has been clearly shown to be essentially the sole source of the electrocatalytic activity in hemeoxygen and heme-H202 catalytic reduction systems. The rate determining steps for these reductions have been shown t o be
Fe(ll)Heme (ads) + 0 2 (Diffusing) (Fe(I1)Heme-0,) a d s
quarter-wave potential in chronopotentiometry half-wave potential in classical polarography E, E'sur peak potential (in cyclic kooltammetry) for the oxidation and reduction of surface adsorbed heme
EDeakpeak potential for the electrocatalytic reduction
processes by linear scan voltammetry any fixed potential a t which one of the catalytic reduction processes is occurring and at which icat represents the current flowing as a result of the electrochemical reduction process T h e Faraday, 96 500 C/equiv Molar concentration Temperature Voltage sweep rate, V/s Transition time in chronopotentiometry (in sec, generally the lifetime for any single electrochemical process) Combination of the nonconcentration dependent terms in the Ilkovic Equation
kO, +
and
Fe(ll)Heme (ads) + H202 (Diffusing) ( Fe(II)Heme-H,02) ads
kH,O,,
The values for ho, and hH202have been evaluated to be (6.4 A 0.5) X lo7 and (1.4 & 0.2) X 10- L/mol s, respectively. A P P E N D I X I. T A B L E OF VARIABLES BrdiEka "reaction layer" thickness, cm Mercury column height, cm i Current, PA Specific Exuerimentallv Measured Quantities observed catalytic current under the existing exIcat perimental conditions as a function of applied potential maximum observable value for the catalytic current zca, under existing experimental conditions max diffusion-limited current for the electrochemical Id process peak current observed by linear scan or cyclic lpeak voltammetry k Constant rate constant for cat H202reduction according to kB BrdiEka "reaction layer" theory m Mass of mercury flowing with a DME in m g / s n Number of electrons transferred per mol of reacting electroactive substance t Time, s A Electrode area, cm2 Concentration in mmol/L or m o l / L as is approC priate to the expression (For the electrochemical equations, C is expressed in mmol/L when i is expressed in FA) D Diffusion coefficient, cm'/s E Potential. V vs. the SCE Specific Experimentally Measured Quantities:
d h
929
LITERATURE C I T E D F. Haber and J. Weiss, Proc. R . SOC. London, Ser. A . 147, 332-351 (1934). K. Zeile and H. Heilstrom, Z . Physioi Chem., 192, 171-192 (1930). S. B. Brown, P. Jones and A. Suggett, Trans. Faraday SOC.,64, 986-993 (1968). S. B. Brown and P. Jones, Trans. Faraday Soc., 64, 994-998 (1968). R . BrdiEka and C. Tropp, Biochem. Z . , 289, 301 (1937). R. BrdiEka and K. Wiesner, Naturwissenschaften. 31, 247 (1943). R. BrdiEka and K. Wiesner, Collect. Czech. Chem. Commun., 12, 39-83 (1947). V. Hanus, Dissertation, Polarographic Institute of the Czechoslovak Academv o i Sciences. Praaue. 1955. R. BrdiEka, Adv. Polarogr., h o c . Int. Congr., Znd, Cambridge, England, 7959,2 , 655-673 (1960). A . Kozowa, V. E. Zilionis and R. J. Brodd, J . Electrochem. Soc.. 116, 1705- 1709 1197 1). J. Manassen, j.datal.. 33, 133-137 (1974) H. Alt, H. Binder, and G. Sandstede. d . Catal., 28, 8-19 (1973). M. Brezina. Collect. Czech. Commun., 22, 339-348 (1957). M. Brezina, Fresenius' 2. Anal. Chem., 224, 74-84 (1966). M. Brezina and A. Hofmanov6-Matejkov6, Collect. Czech. Chem. Commun., 38, 3024-3031 (1973). R. 6.Fulton, Ph.D. Thesis, University of Minnesota, Minneapolis, Minn., 1968. C. F. Kolpin and H. S. Swofford, Jr., Anal. Chem., 50, preceding paper in this issue. R. C. Weast Ed., "Handbook of Chemistry and Physics", 52nd ed., The Chemical Rubber Co.. Cleveland. Ohio. 1971. K. Schwabe, in "Progress in Poiarograpy". Vol. 1, P. Zuman and I. M. Kolthoff, Ed.. Interscience, New York, N.Y., 1962. I.M. Kolthoff and J. J. Lingane. "Polarography", 2nd ed., Interscience, New York-London, 1952.
RECEIVED for review August 1, 1977. Accepted February 3, 1978.
