Polyaniline

ARTICLE pubs.acs.org/JPCC

RRDE and Voltammetric Study of ORR on Pyrolyzed Fe/Polyaniline Catalyst. On the Origins of Variable Tafel Slopes Jerzy Chlistunoff* Materials Physics and Applications Division, Los Alamos National Laboratory, Los Alamos, New Mexico 87545, United States ABSTRACT: Catalytic activity of heat-treated iron and polyaniline-based oxygen reduction catalysts was studied in aqueous acidic media using the rotating ring disk (RRDE) technique and linear potential scan voltammetry employing stationary electrodes. The stationary voltammograms of the catalyst exhibit the presence of a reversible surface red
ox reaction at 0.647 V vs RHE. It is shown that molecular oxygen reversibly adsorbs on the catalyst surface at potentials more positive than the formal potential of the surface red
ox couple and that the adsorption occurs through either the oxidized form of this couple or an atom in its close proximity. The Tafel plots for oxygen reduction reaction (ORR) exhibit variable slopes ranging from 60 mV dec
1 at the lowest overpotentials to more than 240 mV dec
1 at high overpotentials. The kinetic data obtained from the RRDE experiments for various catalyst loadings and from the linear potential scan voltammetry of adsorbed oxygen demonstrate that the high Tafel slopes originate from intrinsic features of the reduction mechanism rather than incomplete catalyst utilization. It is postulated that the surface red
ox couple is FeIII/FeII and that it takes an active part in ORR in the whole range of overpotentials. The proposed ORR mechanism involves a simple mediation by the FeIII/FeII couple at low overpotentials and a concerted process of charge transfer and oxygen
oxygen bond splitting at high overpotentials.

1. INTRODUCTION The demonstration of electrocatalysis of oxygen reduction reaction (ORR) by cobalt phthalocyanine more than 40 years ago by Jasinski1,2 spurred numerous studies of macrocyclic N4complexes of transition metals as ORR catalysts. The catalytic activity for ORR has been reported for the complexes of transition metals (particularly Fe and Co) with porphyrins, phthalocyanines, tetraaza-annulenes, and related macrocycles. References to the extensive literature on the subject can be found in review articles.3
9 A few years after the Jasinski’s pioneering study,1,2 it was discovered that pyrolysis of macrocyclic complexes of transition metals led to materials with improved activity and stability compared to that of the parent compounds.3 The latter is known to be rather low for the isolated complexes, especially in acidic environments. In 1986, Johansson and Larsson demonstrated that the planar N4-macrocyclic coordination was not necessary in the transition metal and nitrogen precursor to produce effective oxygen reduction catalysts upon pyrolysis.10 Later, Gupta and co-workers11 synthesized Co- and Fe-based ORR catalysts using individual transition metal and nitrogen precursors. The discoveries by Johansson and Larsson10 and by Gupta et al.,11 opened the way to much less expensive catalysts, and some of the most active materials created to date have been synthesized using rather inexpensive nitrogen precursors such as ammonia and phenantroline12 or polyaniline13 instead of costly N4-macrocycles. While it is generally accepted that both nitrogen and transition metal precursors are necessary to achieve high catalytic activity of the pyrolyzed material, the nature of the active site(s) remains a subject of an ongoing debate. The majority of researchers seem to favor the hypothesis of the transition metal being an inherent r 2011 American Chemical Society

part of the active site,5 but according to others, the active site may involve no metal.14
23 If ORR catalysis involves transition metal centers, the coordination of the surface metal must play an important role in both catalyst activity and stability. Consequently, the question of metal coordination has been addressed many times. The N4-coordination, such as that in N4-macrocycles, was suggested for the catalysts produced at lower temperatures (T e 550
C).24,25 Simultaneous presence of two active sites, whose abundances depend on pyrolysis temperature, was suggested by Lef
evre and co-workers based on time-of-flight secondary ion mass spectrometry (ToF SIMS).26 According to these authors, the most abundant active site for the catalysts pyrolyzed at high temperatures (700
C g T g 900
C) is most likely of an N2-type, where two nitrogen atoms are incorporated in a graphene sheet as in phenantroline molecule.12,26 On the other hand, a linear correlation between the catalyst activity and the population of in-plane Fe
N4 centers embedded in a graphene-type matrix was reported by Koslowski et al.27 The assumption of N-coordinated transition metal being an active center leads to another question, namely, whether a single metal center is sufficient to break the strong oxygen
oxygen bond and thus catalyze four-electron reduction of oxygen to water or two such centers arranged, e.g., as in cofacial porphyrins,28 are required. The latter notion originates from a quite common point of view that the splitting of dioxygen molecule requires a simultaneous interaction with two active sites.29 As low peroxide generation rates Received: September 2, 2010 Revised: January 28, 2011 Published: March 17, 2011 6496

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The Journal of Physical Chemistry C are relatively common for ORR on heat-treated catalysts17,30
34 and no closely spaced metal centers have been detected on their surface, the assistance of other surface or subsurface atoms in splitting the O
O bond has been proposed.35,36 The ambiguity in the determination of the active site of the pyrolyzed ORR catalysts has numerous origins. First, the heattreated catalysts are highly dispersed materials, inherently difficult to study, whose structure and composition likely vary with the distance from the surface, e.g., due to the surface segregation38 of catalyst components during pyrolysis or because of the acid leaching13,17,36 after pyrolysis. Second, a vast majority of contemporary surface characterization techniques have nonzero penetration depths, i.e., probe a finite surface layer of a material rather than its surface. Third, the active sites may be very few, which can make them difficult to detect. Finally, different active sites may contribute to the catalytic activity of various materials or even to that of a single material. On the other hand, oxygen reduction on many heat-treated ORR catalysts exhibits strikingly similar properties. These are, for example, low hydrogen peroxide generation rates17,30
34 and curved Tafel plots with Nernstian slopes (∼60 mV dec
1 at ambient temperatures) at low overpotentials12,15,16,25,32,35,38
40 and very high (>120 mV dec
1) slopes at high overpotentials.12,16,36,39
43 The Nernstian Tafel slopes have been linked to a red
ox catalysis involving the transition metal38,40 or one-electron transfer followed by a chemical step,15 but the high Tafel slopes have not been much discussed. As suggested by some reports,12 they were most likely attributed to an incomplete catalyst utilization resulting from mass transport losses at high current densities. However, Gojkovic et al.16 indicated that the existing theoretical models of porous electrodes44 could not satisfactorily explain the high Tafel slopes they observed for heat-treated iron(III) tetramethoxyphenyl porphyrin catalysts. More specifically, a reaction order for oxygen of 0.5 could be expected from the theory44 for the potential range characterized by the extremely high (∼500 mV dec
1) Tafel slopes, whereas the order was one in the whole potential range.16 While the relatively high hydrogen peroxide generation yields observed for those catalysts31 may have contributed to the apparent deviation of the experimental results16 from the theory,44 other origins of the high Tafel slopes cannot be excluded. In this paper, we present a detailed electrochemical study of ORR on pyrolyzed Fe and polyaniline-based oxygen reduction catalysts developed at Los Alamos National Laboratory.13 We report kinetic data for ORR and demonstrate that the curvature of the measured Tafel plots reflects intrinsic properties of the ORR mechanism. Due to the similarity of our results to those reported for other heat-treated nonprecious metal ORR catalysts, the conclusions from this study may have important implications for the future direction of the related research.

2. EXPERIMENTAL SECTION The electrochemical measurements were performed using either a CH Instruments bipotentiostat model CHI760D (RRDE and selected voltammetric measurements) or a Pine model AFCBP1 bipotentiostat (RRDE) or a PAR model 283 potentiostat (voltammetry). A positive feedback (85% compensation level, 10% undershoot) was used for iR drop compensation in all voltammetric experiments with the exception of those employing RRDE. The working electrode in the majority of RRDE experiments and selected voltammetric experiments was a Pine model AFE7R9GCPT electrode with a 5.61 mm diameter glassy carbon

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Figure 1. Cyclic voltammograms of oxygen-free (1) and oxygensaturated (2) solution of sulfuric acid (0.5 mol dm
3) and their difference (3). Stationary (no rotation) GC disk (5.61 mm diameter) of RRDE 1 used as a working electrode. Catalyst loading 0.7 mg cm
2 (batch 1). Scan rate 100 mV s
1. Equilibration time at an open circuit potential was ∼20 min.

disk and a Pt ring with an inner diameter of 6.25 mm and an outer diameter of 7.92 mm. This electrode is called hereafter RRDE 1. The geometric surface area of the disk of RRDE 1 was 0.2472 cm2, and the nominal collection efficiency of this RRDE was 37%. In a limited number of RRDE experiments, a Pine model AFDT21GCPT electrode with a GC disk (7.64 mm diameter) and a Pt ring (8.44 mm outer diameter and 7.98 mm inner diameter) was used. This electrode is called RRDE 2. The collection efficiency of both RRDEs was calibrated using a 1 mmol dm
3 potassium ferricyanide solution in 1 mol dm
3 KCl. The working electrode in majority of voltammetric experiments was a GC disk (3 mm diameter) from Bioanalytical Systems Inc. The counter electrode in all RRDE experiments was a graphite rod, whereas a Pt mesh, separated from the main cell compartment by a glass frit, was used in voltammetric experiments. The reference electrode was either a Vycor-separated hydrogen electrode employing 6% H2 (in Ar carrier gas) in the identical electrolyte solution as that used in the cell or an Ag/AgCl (3 mol dm
3 NaCl) electrode (Bioanalytical Systems). The latter electrode was calibrated against the former. Unless otherwise stated, all potentials in this work are referred to the reversible hydrogen electrode at zero elevation, i.e., 100% H2 at 101.325 kPa. The solutions were deaerated using ultrahigh purity (UHP) Ar. The concentration of oxygen in solution for oxygen adsorption measurements was adjusted by simultaneous purging of high-purity oxygen and UHP argon through separate fine PTFE tubing (0.8 mm). After the purging, the oxygen and argon tubing were lifted above the solution level, and the gases allowed to flow above the solution to prevent its contact with air. The actual oxygen concentration in solution was measured using a limiting current of oxygen reduction on a Pt microdisk (50 μm diameter) from Bioanalytical Systems. If the oxygen concentration before and after the adsorption measurements differed by more than 5%, the measurement was repeated. The stock solution of sulfuric acid (0.5 mol dm
3) was obtained by diluting a 96% H2SO4 (Veritas double distilled from Vycor, GFS Chemicals) with deionized water (18.2 MΩ cm, Millipore) to a volume. The respective HClO4 solution (0.5 mol dm
3) was obtained from 70% HClO4 (Veritas redistilled, GFS Chemicals). The catalyst was prepared using the procedure described previously,13 and its specific surface area from BET measurements was ∼400 m2 g
1. The catalyst ink was prepared by ultrasonically dispersing 20 mg of the catalyst in 1 cm3 of a mixture of 6497

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Table 1. Surface Characterization of Catalyst Samples batch specific surface area (m2 g
1) Qrev/Aa (μC cm
2)

Figure 2. Oxygen reduction peak current density in 0.5 mol dm
3 H2SO4 after background correction (curve 3 in Figure 1) plotted vs the square root of the scan rate. Experimental details as in Figure 1 with the exception of scan rate. The straight solid line and the dashed line represent the hypothetical contributions from diffusing and adsorbed oxygen, respectively.

1

∼400

4.61 ( 0.27

2

∼400

2.67 ( 0.25

3







Qrev/Qtotb 0.16 ( 0.01 0.075 ( 0.009 0.075

a Qrev/A is baseline-corrected charge of the reversible system normalized versus the actual surface area of the catalyst sample, as calculated from BET data and catalyst loading. b Qtot is total cathodic/anodic charge (including Qrev) measured between þ0.8 and þ0.2 V vs RHE in oxygenfree solution. Assuming that Qtot is proportional to the specific surface areas of similar catalysts, the Qrev/Qtot ratios may be used to compare relative surface concentrations of the reversible system without knowing the corresponding specific (BET) surface areas. The ratio of the respective surface concentrations determined in this way for batches 1 and 2 is close to 2.1 and thus exceeds the corresponding ratio (∼1.7) determined using the BET and catalyst loading data.

water (30% vol), isopropanol (70% vol), and 5% Nafion solution (Solution Technology, 5% vol). Three catalyst batches with slightly different properties were used in the experiments. The catalyst loading in all experiments was between 0.03 and 0.78 mg cm
2.

3. RESULTS 3.1. Adsorption of Oxygen. A typical stationary cyclic voltammogram of oxygen in sulfuric acid solution is presented in Figure 1 together with the corresponding background voltammogram recorded in the absence of oxygen. As shown in Figure 1, the reduction of oxygen overlaps with a reversible surface red
ox process with a formal potential of 0.647 V vs RHE, visible in the voltammogram of the deoxygenated solution (curve 1 in Figure 1). The contribution of ORR in the total current measured in oxygensaturated solution is relatively low, and the shape of the backgroundcorrected oxygen reduction peak resembles that for an electrochemical reaction of a surface-confined species rather than for a mass transport controlled process. In accordance with this finding, the observed parabolic dependence of the oxygen reduction peak current on the square root of the scan rate (Figure 2) revealed that oxygen adsorbed on the catalyst surface significantly contributes to the measured reduction current. The corrected voltammograms of oxygen, which were used to create the plots in Figure 2, were further utilized to determine the reduction charge of adsorbed oxygen. The oxygen reduction currents measured at different scan rates were integrated, and the essentially linear plot of the measured oxygen reduction charge vs the inverse square root of scan rate (not shown) was extrapolated to zero, i.e., to the infinite scan rate to yield the desired reduction charge of adsorbed oxygen (Qads). The ratio of Qads to the charge of the reversible surface red
ox system (Qrev, determined from the baseline-corrected peaks of the reversible system) was found to be 1.82. The low value of the ratio suggested that the adsorption site for oxygen may be structurally related to the surface atoms participating in the one-electron reversible red
ox reaction. The conclusion was supported by a similar experiment performed for a different catalyst sample (batch 2). In this case, the surface charge density corresponding to the reversible system was ∼1.7 times lower than that observed for batch 1 (curve 1 in Figure 1), as calculated from the catalyst loadings and the experimental BET surface areas (Table 1). This finding indicated that the accessible surface concentration of the species responsible for the reversible surface red
ox reaction was ∼1.7 times lower

Figure 3. Oxygen reduction charge in 0.5 mol dm
3 H2SO4 from cyclic voltammetry (see text) plotted vs the equilibration time at 0.91 V vs RHE. Stationary GC disk (3 mm diameter) used as a working electrode. Catalyst loading 0.78 mg cm
2 (batch 2). Scan rate 324 mV s
1.

for batch 2 than that for batch 1 (see also footnote b to Table 1). In spite of the difference, the Qads/Qrev ratio determined for batch 2 was 1.88, i.e., very similar to that determined for batch 1. More detailed information on the adsorption of oxygen was obtained from the studies of its dependencies on equilibration time (adsorption kinetics) and bulk concentration (adsorption thermodynamics). These experiments were carried out exclusively for a single loading (0.78 mg cm
2) of the catalyst batch 2. The measurements aimed at the determination of the adsorption kinetics were performed as follows. First, the electrode with the catalyst was equilibrated at a constant potential, where virtually no reduction within the reversible red
ox system could occur. Even though any potential more positive than ∼0.9 V fulfilled this condition, we selected 0.91 V for all oxygen adsorption experiments to avoid excessive oxidation of the catalyst surface. However, it has to be noted that oxygen adsorption was not affected by the equilibration potential between 0.9 and 1.04 V vs RHE. After the equilibration, two voltammetric cycles at a scan rate of 324 mV s
1 were recorded. The end of the second cycle was the beginning of another equilibration period. This procedure guaranteed negligible surface concentration of oxygen at the beginning of every equilibration period, whereas the high scan rate minimized the contribution of oxygen diffusing from the bulk of solution to the measured reduction charge. The voltammograms were corrected for the background current recorded under identical experimental conditions in the absence of oxygen. Figure 3 6498

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shows the charge of oxygen reduction measured in the first cycle plotted against the equilibration time at þ0.91 V vs RHE. The plot in Figure 3 was fitted using the following equation, derived from the Langmuir adsorption model45 Q ORR ¼ Q dyf þ Q ads ¼ Q dyf þ nFAΓO2 ¼ Q dyf þ

nFAkads CiO2 Γ¥ kads CiO2 þ kd Γ¥

! !! kads CiO2 þ kd Γ¥ 1
exp
t Γ¥

ð1Þ

where QORR is the total oxygen reduction charge, Qdyf stands for the (approximately constant) reduction charge of oxygen diffusing from the bulk of solution during the cathodic scan, Qads is the reduction charge of adsorbed O2, kads, and kd are respectively the adsorption and desorption rate constants, CiO2 is the oxygen concentration at the solution/catalyst interface, t is the equilibration time at 0.91 V, F is the Faraday constant, A is the actual surface area of the catalyst deposited on the electrode, ΓO2 is the surface concentration of oxygen, Γ¥ is the surface concentration of oxygen at the maximum coverage, and n is the number of electrons transferred in ORR. We assumed that n is 4 (see section 3.5) and the surface area (A) is that from the BET measurements, i.e., 220 cm2 for the data in Figure 3. We also approximated CiO2 in eq 1 by the bulk concentration of oxygen CO2. The highest surface concentration of adsorbed oxygen (Γ¥) was assumed to be that of the surface red
ox species (Γredox) and equal to 2.77  10
11 mol cm
2, as calculated from its measured reduction/ oxidation charge. The last assumption implies that oxygen molecule is adsorbed on the surface exclusively through the single electroactive surface species or a neighboring surface atom (atoms) remaining in a structural relationship with this species. While such a hypothesis may seem unjustified, its validity is strongly supported by the measured equilibrium surface concentrations of oxygen, presented below in this section. When the adsorption rate constant kads and the sum kadsCO2 þ kdΓ¥ were treated as independent parameters, the fitting resulted in a negative value for the desorption rate constant kd. When kads and kd were forced to have the same sign, the following parameters were obtained: kads = 1.07  10
7 cm s
1 and kd/kads = 6.35  104 cm
1. Using the above kinetic parameters, we calculated the time necessary to attain specific coverage of adsorption sites by oxygen for its different bulk concentrations. The calculations revealed that the equilibration of the electrode for 12 min is sufficient to attain at least 99% of the oxygen equilibrium coverage in 0.5 mol dm
3 H2SO4 with oxygen concentrations from 15% to 100% of the saturated oxygen concentration (8.65  10
4 mol dm
3) at Los Alamos elevation (2100 m above sea level). The experiments aimed at the determination of equilibrium surface concentrations of oxygen and the adsorption isotherm were carried out using the same voltammetric protocol as that used in the kinetic experiments (vide supra). In accordance with the kinetic data, the equilibration time selected for the measurements of equilibrium surface concentration of oxygen was 12 min in the majority of experiments. Longer equilibration times were found unfeasible because of the very long duration of such experiments. Oxygen reduction charges were determined from the first cycle at 50, 100, 144, 200, and 324 mV s
1 and plotted against the inverse square root of the scan rate (Figure 4). The

Figure 4. Oxygen reduction charge in 0.5 mol dm
3 H2SO4 from cyclic voltammetry after 12 min equilibration at 0.91 V vs RHE (see text) plotted vs the inverse square root of scan rate. Stationary GC disk (3 mm diameter) used as a working electrode. Catalyst loading 0.78 mg cm
2 (batch 2). Oxygen concentration from top (mmol dm
3): 0.86, 0.76, 0.66, 0.53, 0.44, 0.35, 0.29, 0.17.

reduction charge of adsorbed oxygen (Qads) was determined from the extrapolation of the plots to zero (Figure 4). The dependences of the reduction charge of adsorbed oxygen (Qads) on the bulk oxygen concentration were fitted using the expression derived from the Langmuir adsorption isotherm Q ads ¼ ¼

nFAkads CO2 Γ¥ Q max kads CO2 ¥ ¼ kads CO2 þ kd Γ kads CO2 þ kd Γ¥ Q max K ads CO2 K ads CO2 þ 1

ð2Þ

where Kads is the adsorption equilibrium constant and Qmax is the reduction charge of adsorbed oxygen corresponding to Γ¥. While the reduction charges of adsorbed oxygen obtained from different experiments for similar bulk concentrations of oxygen and identical catalyst loadings were generally comparable, the variation of calculated Qmax was more significant. This was because the measured ORR charges were determined in a relatively narrow range of bulk oxygen concentrations that corresponded to low surface coverages by adsorbed oxygen, whereas Qmax was indirectly obtained from a long extrapolation of the nonlinear function (eq 2), the latter being rather sensitive to the initial slope of the Qads vs CO2 relationship. However, the common feature of all the isotherms obtained was that the ratio of the maximum reduction charge of adsorbed oxygen to the charge of the reversible surface system (Qmax/Qrev) was always a small number. It was typically higher than 4 but never exceeded 8. This fact has important implications. First, the maximum surface concentration of adsorbed oxygen is at least 2 orders of magnitude less than what can be expected from the specific surface area of the catalyst assuming close packing of oxygen molecules. Consequently, if the adsorbed molecules are uniformly distributed over the catalyst surface, there are virtually no interactions between them. This finding supports the choice of Langmuir adsorption isotherm. Second, as already suggested by the Qads/ Qrev values measured for the two different batches of the catalyst in oxygen-saturated solutions (vide supra), there likely exists a simple stoichiometric and structural relationship between the adsorption sites and the surface atoms responsible for the oneelectron reversible surface red
ox process, the possibility of direct oxygen adsorption on these atoms not being excluded. 6499

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Figure 5. Reduction charge of oxygen adsorbed at 0.91 V vs RHE in 0.5 mol dm
3 H2SO4 determined using the procedure presented in Figure 4 (see also text). The nonlinear weighted regression fits were obtained using adsorption models based on Langmuir adsorption isotherm with the following constraints (see text): unconstrained (solid line); Qmax = 2  Qrev (dotted line); Qmax = 4  Qrev (dashed line). An extrapolation of the unconstrained fit to higher oxygen concentrations is shown in the inset. Stationary GC disk (3 mm diameter) used as a working electrode. Catalyst loading 0.78 mg cm
2 (batch 2).

Figure 5 shows the reduction charge of adsorbed oxygen plotted vs bulk oxygen concentration. For these data, the overall best fit was obtained. The maximum reduction charge (Qmax) for this fit (3380 μC) is approximately 6.5 times higher than the reduction/oxidation charge (518 μC) of the reversible surface red
ox system (Qrev) measured in the absence of oxygen. The adsorption equilibrium constant (Kads) retrieved from the fit is ∼570 dm3 mol
1 and the corresponding free energy of adsorption, ΔG0ads =
15.7 kJ mol
1. From the fitting parameters, we also obtained a value of 8.38  104 cm
1 for the ratio of the desorption and adsorption rate constants (kd/kads), in good agreement with the value of 6.35  104 cm
1 obtained from the apparent adsorption kinetics (vide supra). In Figure 5, together with the best unconstrained fit of the data, there are also shown two fits obtained using partially constrained models, where the maximum oxygen reduction charge (Qmax) was assumed to be either four or two times as high as the reduction/oxidation charge of the reversible red
ox system. These models would apply if oxygen molecules in the adsorbed state were forming a bond with a single-surface red
ox center (or an atom in its immediate vicinity) or two such neighboring centers, respectively. As seen from Figure 5, the model of oxygen molecule strongly interacting with two surface red
ox centers is rather unfeasible, whereas the one-center model results in quite a satisfactory fit of the experimental points. The adsorption equilibrium constant (Kads) for the latter model equals 860 dm3 mol
1, and the corresponding free energy of adsorption, ΔG0ads =
16.7 kJ mol
1. These numbers remain in a reasonable agreement with those obtained from the unconstrained adsorption model. 3.2. Kinetics of Reduction of Adsorbed Oxygen. While the simultaneous reduction of oxygen diffusing from the bulk of solution and adsorbed on the catalyst surface was helpful for the determination of the equilibrium surface concentration of adsorbed oxygen, it is rather unfavorable for the determination of the reduction kinetics of adsorbed oxygen. The contribution from the diffusing species becomes quite significant at low scan rates (Figures 2 and 4) and makes such kinetic measurements virtually impossible under these conditions. In order to eliminate this contribution, the following procedure was applied. As in the

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Figure 6. Differential voltammograms for oxygen adsorbed at 0.91 V vs RHE in 0.5 mol dm
3 H2SO4 saturated with oxygen. Stationary GC disk (3 mm diameter) used as a working electrode. Catalyst loading 0.78 mg cm
2 (batch 2). Scan rates are listed in the figure.

experiments aimed at the determination of the adsorption kinetics and equilibrium (vide supra), two-cycle voltammograms were recorded. After the background correction, the current measured in the second cycle was subtracted from the current in the first cycle. While the contributions from the diffusing and adsorbed oxygen in the total reduction current are somewhat interrelated, one can demonstrate that the former is always significantly smaller in such differential voltammograms than the latter. However, it has to be noted that the major (adsorption) component in the differential voltammograms always corresponds to the reduction of only a fraction of surface concentration of oxygen attained during the equilibration period due to a contribution from the adsorption in the second cycle. The effect is more pronounced at slow scan rates, which result in longer times between the termination of oxygen reduction in the anodic scan of the first cycle and its resumption in the cathodic scan of the second cycle. Consequently, the use of the differential voltammograms to determine the equilibrium surface coverage by oxygen was not attempted. The experiments were performed exclusively for the catalyst batch 2. A number of differential voltammograms obtained using the above-described procedure for a single oxygen concentration is shown in Figure 6. As seen from Figure 6, the reduction peaks are quite symmetrical at lower scan rates, but their descending sections become somewhat elongated upon the scan rate increase. The effect is not due to an unaccounted for contribution from the solution (diffusing) species, as the latter decreases with the scan rate (Figures 2 and 4). Consequently, it most likely reflects the changes in the intrinsic kinetics of the reduction of adsorbed oxygen. This conclusion remains in agreement with the fact that the observed asymmetry does not result barely from the appearance of a “tail” in the descending section of the reduction peak but is accompanied by a slower current increase in its ascending section as well and, consequently, by an increase in the peak width (Figure 6). The latter cannot be caused by an undercompensated iR drop, as the position and shape of the reduction peaks of adsorbed oxygen at a constant scan rate were virtually independent of oxygen concentration. The dependence of the peak width for adsorbed oxygen on the square root of scan rate is shown in Figure 7 for several oxygen concentrations. From the linear plots in Figure 7, the average peak width at half height in the limit of zero scan rate (and the lowest overvoltage) was determined to be 40 ( 9 mV. According to the theory of irreversible processes of surface confined species,46 this number corresponds to the cathodic transfer coefficient RnR = 1.56 (1.28
2.02), which demonstrates 6500

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Figure 7. Peak widths at half height for the differential voltammograms of adsorbed oxygen (see text) in 0.5 mol dm
3 H2SO4. Equilibration performed for 2 min at 0.91 V vs RHE. Stationary GC disk (3 mm diameter) used as a working electrode. Catalyst loading 0.78 mg cm
2 (batch 2). For clarity, the plots for subsequent oxygen concentration shifted by multiples of 50 mV along the potential axis. Oxygen concentration in solution (mmol dm
3) and the respective shifts in mV (in parentheses) from the bottom: 0.83 (0), 0.74 (50), 0.63 (100), 0.59 (150), 0.51 (200), 0.39 (250), 0.33 (300), 0.22 (350).

Figure 8. Apparent cathodic transfer coefficient for the reduction of adsorbed oxygen from differential voltammograms (see text) in 0.5 mol dm
3 H2SO4 calculated from the peak width at half height.46 Experimental conditions as in Figure 7. Oxygen concentration in solution (mmol dm
3): 0.83 (black), 0.74 (red), 0.63 (navy blue), 0.59 (green), 0.51 (orange), 0.39 (purple), 0.33 (yellow), 0.22 (cyan).

that the second electron transfer is the rate-determining step of the reduction of adsorbed oxygen at low overvoltages. On the other hand, the significant increase in peak width with both the scan rate and cathodic overpotential (Figures 6 and 7) and the corresponding decrease in the apparent transfer coefficient (Figure 8) indicate the presence of potential-driven changes in the reaction mechanism. Given the likely link between ORR and the surface red
ox system, we also explored the possibility of the electron transfer reaction of the latter being responsible for the low transfer coefficients in ORR (Figure 8). However, voltammograms of deoxygenated solutions recorded at higher scan rates than those used in the experiments involving oxygen demonstrated no meaningful departure of the cathodic transfer coefficient of the reversible red
ox system from 0.5. This finding demonstrated that the electron transfer kinetics of this system was not responsible for the low cathodic transfer coefficient of ORR. 3.3. Effects of Rotation Rate on Steady State Voltammetry of Oxygen. Neither linear potential scan (5 mV s
1) nor staircase voltammograms of oxygen for the catalyst loadings on the order of a few hundred micrograms per 1 cm2, typically used in

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Figure 9. Altitude-corrected Koutecky
Levich plot for the limiting oxygen reduction current (forward scan) in 0.5 mol dm
3 H2SO4 saturated with oxygen (full symbols) recorded using RRDE 1. For comparison, the theoretical plot (open symbols) for four-electron ORR is also shown. Enhanced oxygen transport at higher rotation rates results in a nonzero intercept (not shown) with the sign opposite to that of the measured current. Scan rate 5 mV s
1.

Figure 10. Staircase voltammograms for ORR in 0.5 mol dm
3 H2SO4 recorded for two different catalyst loadings (batch 1) using RRDE 1. Potential step height 10 mV. Step duration 30 s. Current measurement time 1 s.

the present work, strictly obeyed Levich and Koutecky
Levich equations.47 The limiting currents were higher than those expected from the rotation rate, geometric surface area of the electrode, and the transport characteristics of oxygen in the studied solutions. The effect increased with the rotation rate (Figure 9) and the size of the disk. The latter finding indicated that the linear disk velocity was responsible for the enhanced transport of oxygen toward the electrode surface. A closer examination of the seemingly smooth electrode surface under microscope revealed that the catalyst layer was significantly rough. Such roughness most likely produced local flow turbulences that increased the apparent rate of transport. Due to this shortcoming of our catalyst ink formulation, all kinetic data reported in this paper were determined from the data recorded under conditions assuring the lowest error possible, i.e., using the smaller RRDE (see the Experimental section) at a relatively low rotation rate of 400 rpm. 3.4. Catalyst Loading Effect on Steady State Oxygen Reduction Kinetics. In Figure 10, there are shown backgroundcorrected staircase voltammograms of oxygen for two significantly different catalyst (batch 1) loadings. As expected, the higher loading of the catalyst results in a faster apparent kinetics of oxygen reduction. However, the voltammogram recorded for the smaller loading (0.03 mg cm
2) is more drawn out, as if there were significant catalyst utilization problems in the thinner catalyst layer. 6501

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Figure 11. Altitude-corrected Tafel plots for ORR in 0.5 mol dm
3 H2SO4 obtained from the data in Figure 10.

The explanation of the apparent mismatch between the slopes of the voltammograms recorded for the high (0.7 mg cm
2) and the low (0.03 mg cm
2) catalyst loadings is that already stated in section 3.2, i.e., the potential-dependent transfer coefficient of ORR. The correctness of the above explanation is confirmed after the conversion of the measured oxygen reduction currents in Figure 10 to the corresponding kinetic currents using standard mass transfer correction48 and the calculations of the respective volumetric current densities (Figure 11). The resulting Tafel plots for volumetric current densities are presented in Figure 11 together with the current densities expressed against the electrode geometric surface area (in the inset). The vertical shifts of the plots in the inset by the numbers equal to the logarithms of the respective film thicknesses make the Tafel plots overlap in the whole range of overpotentials (Figure 11). As manifested by the small oscillations of the kinetic current at the highest overvoltages for the high catalyst loading (0.7 mg cm
2), the accuracy of the kinetic current determination slightly suffers in this potential range due to the closeness of the measured currents to the limiting current. This is not the case for the lower loading, where the measured currents are quite substantially lower than the limiting current even at the highest practical overvoltages (Figure 10). In spite of the said inaccuracies, the agreement between the volumetric kinetic current densities determined for both catalyst loadings is excellent. This finding leaves little doubt that the curvature of the Tafel plots does not originate from incomplete catalyst utilization44 or uncompensated resistance in the catalyst layer but reflects changes in the intrinsic ORR kinetics. The Tafel plots exhibit a 60 mV dec
1 slope at low overpotentials and significantly higher slopes, exceeding 240 mV per current decade at high overpotentials. Assuming that the reversible surface red
ox center is the catalytically active site, one obtains 0.02 e site
1 s
1 for the turnover frequency at 20
C and 0.8 V vs RHE. This number has been calculated assuming that all surface red
ox centers are involved in the catalysis. Unlike the reduction of adsorbed oxygen (sections 3.1 and 3.2), which involves oxygen molecules being adsorbed on ∼15% of the active sites, the ORR under RRDE (and fuel cell) conditions involves all active centers with specific probabilities determined by the rates of all processes affecting the measured kinetics. 3.5. Hydrogen Peroxide Generation and Catalyst Activity in Peroxide Reduction/Oxidation. In Figure 12, there are shown linear scan RRDE voltammograms of a 0.5 mol dm
3 H2SO4 solution containing either 8.65  10
4 mol dm
3 oxygen or 6.4  10
3 mol dm
3 H2O2, recorded using the same electrode and identical experimental conditions. In spite of almost eight times higher concentration of peroxide than that of oxygen, the currents due to oxygen reduction are very comparable to those resulting from hydrogen peroxide reduction or oxidation (Figure 12). Moreover,

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Figure 12. Hydrodynamic voltammograms of oxygen and hydrogen peroxide in 0.5 mol dm
3 H2SO4 recorded using RRDE 2. Catalyst loading 0.7 mg cm
2 (batch 1). Scan rate 5 mv s
1. Rotation rate 4900 rpm.

Figure 13. Molar proportion of hydrogen peroxide generated by ORR under staircase voltammetric conditions in 0.5 mol dm
3 H2SO4 recorded using RRDE 1. Potential step height 10 mV. Step duration 30 s. Current measurement time 1 s. Catalyst (batch 1) loading listed in the figure.

the increase in hydrogen peroxide reduction current with cathodic potential is essentially linear in the whole studied potential range. Similar, rather unusual electrochemical characteristics of hydrogen peroxide have been reported by Jaouen and Dodelet30 for a number of heat-treated oxygen reduction catalysts. The behavior was rationalized by assuming oxo-ferryl-like cation radical to be a common rate-limiting reaction intermediate for both oxygen and hydrogen peroxide reduction reactions.30 The local turbulences interfering with oxygen transport described in section 3.3 are at least partially responsible for the effect of catalyst loading on peroxide generation (Figure 13). Another possible cause of the effect can be linked to the occurrence of the sequential mechanism of ORR, where oxygen is first reduced to hydrogen peroxide, which further undergoes various chemical and electrochemical reactions to eventually produce water.49 Based on the measured reduction currents of hydrogen peroxide (Figure 12), one can estimate that no more than 25% of oxygen molecules reaching the catalyst surface can undergo the reduction to peroxide. The respective number would be even lower if the observed hydrogen peroxide reduction currents (Figure 12) were in part due to the disproportionation of peroxide on the catalyst surface and subsequent reduction of oxygen produced in the reaction. These numbers indicate that the peroxide route is not dominant for the studied catalysts. 3.6. Reaction Orders of ORR. The determination of the order of an electrochemical reaction for one of its reactants has to be 6502

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The Journal of Physical Chemistry C

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Figure 14. Kinetic ORR currents in H2SO4 solutions of various concentrations under staircase voltammetric conditions recorded using RRDE 1. Potential step height 10 mV. Step duration 30 s. Current measurement time 1 s. Catalyst (batch 3) loading 0.5 mg cm
2. Acid concentration (mol dm
3): 0.5 (black); 0.4 (red); 0.3 (light green); 0.2 (blue); 0.1 (purple).

performed at a constant electrode potential. Such determination for hydronium cation is complicated by a nonzero and acid concentration dependent liquid junction potential between the reference electrode and the acidic solution.50,51 The only potential scale free of the junction potential in strongly acidic electrolytes is the reversible hydrogen electrode scale. However, the use of this scale in the determination of the reaction order for Hþ is fundamentally wrong, as it violates the constant potential condition required for the reaction order determination. However, it can be easily demonstrated that the reaction order for Hþ (ROHþ) may be expressed by the following approximate equation: ! !   DlogðjÞ DlogðjÞ RnR F  þ ROHþ ¼ DlogðaHþ Þ DlogðCÞ 2:3RT E

! dERHE d logðCÞ

ERHE

ð3Þ

In the above equation, j stands for current density, aHþ is molar activity of Hþ, E is a constant electrode potential, ERHE is the potential measured vs RHE, RnR is the transfer coefficient, and C is acid concentration. In Figure 14, there are shown Tafel plots for oxygen reduction in sulfuric acid solutions with five different concentrations ranging from 0.1 to 0.5 mol dm
3. The plots in Figure 14 completely overlap in the entire potential range studied. Consequently, the first term in eq 3 equals zero. As the dissociation of HSO4
does not occur substantially for the studied concentrations of sulfuric acid, the derivative (dERHE)/(d log(C)) equals approximately 60 mV. For the linear section of the plots at low overpotentials, the slope is 60 mV dec
1, and, therefore, the corresponding reaction order for Hþ is one. Correspondingly, the reaction order for Hþ gradually decreases with the cathode overpotential, following the changes in the Tafel slope (RnRF)/(2.3RT). No meaningful effect of acid concentration on hydrogen peroxide generation was detected in the whole range of potentials. The reaction order for oxygen was determined using staircase voltammetry for a single catalyst loading (batch 1, loading 0.03 mg cm
2) in a 0.5 mol dm
3 H2SO4 solution saturated with either oxygen or air. The ratio of the measured oxygen reduction currents for the oxygen- and air-saturated solutions was equal to 4.94 ( 0.12 at all potentials studied (0.04
0.74 V vs RHE), i.e., very close to the ratio of pure oxygen pressure and its partial

pressure in air (4.78) in the whole studied potential range. As the solubility of oxygen in moderately concentrated sulfuric acid solutions in water strictly obeys Henry’s law,52 the result implies that the reaction order for oxygen is one in the whole range of cathodic overpotentials. The result remains in agreement with the previous finding that the changes in ORR Tafel slope with overpotential reflect intrinsic properties of the oxygen reduction mechanism rather than catalyst utilization problems. If the Tafel slopes at the highest overpotentials, which more than quadruple the slopes at the lowest overpotentials, were due to the incomplete catalyst utilization, the reaction order for oxygen would be less than one according to the existing theories of porous electrodes.44,53,54 Oxygen concentration was found to have an effect on hydrogen peroxide generation rates. The molar proportions of peroxide determined from the ring current were approximately two times higher for the air-saturated solution than those for oxygensaturated solutions. The result indicates the presence of peroxide-scavenging reactions governed by the second or higher order kinetics. We also explored the effects of anions on the rate of oxygen reduction. In the experiments, mixtures of sulfuric acid and perchloric acid solutions with a total analytical concentration of 0.5 mol dm
3 were studied. The individual concentrations were varied from 0 to 0.5 mol dm
3 in 0.1 mol dm
3 steps. The equilibrium concentration of Hþ in such solution is virtually independent of the solution composition and approximately equal to 0.5 mol dm
3, and the equilibrium concentrations of the anions are exactly equal (ClO4
) to or well approximated (HSO4
) by the analytical concentrations of the respective acids. No effect of the solution composition on the ORR kinetics in the potential range corresponding to the Nernstian Tafel slope was detected from staircase voltammetry. At higher overpotentials, a slight increase in the ORR rate upon the transition from pure H2SO4 to pure HClO4 was observed. Due to the small magnitude of the effect, it was impossible to determine if the effect was real or caused by an experimental error. The latter is more significant at the high overpotentials, where the measured currents approach their transport-controlled limit.

4. DISCUSSION AND CONCLUSIONS As demonstrated by the independence of the volumetric current density on the catalyst loading and by the kinetic data for oxygen adsorbed on the catalyst surface, the observed curvature of ORR Tafel plots is an intrinsic property of the reaction mechanism and does not result from an incomplete catalyst utilization caused by transport and conductivity problems in the catalyst layer.44,53,54 Moreover, the reaction order for oxygen is equal to one in the whole range of overpotentials, whereas incomplete catalyst utilization should result in the reaction orders less than unity in the range of the extremely high Tafel slopes, based on the theoretical models.44,53,54 However, there exist differences between the kinetic data obtained from the steady state measurements employing RRDE and the data from the stationary voltammetry. In particular, the Tafel slope of 60 mV per current decade at low overpotentials in RRDE experiments implies that the apparent cathodic transfer coefficient RnR equals one, whereas a value of 1.56 has been obtained for the similar potential range from the voltammetry of adsorbed oxygen. As will be demonstrated below, the difference is only superficial and only reflects the presence and the absence of 6503

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The Journal of Physical Chemistry C oxygen transport step in the RRDE and the voltammetric experiments, respectively. During the earlier stage of this work, a hypothesis was put forward that ORR is mediated in the whole potential range by the reversible surface red
ox system. The curvature of the experimental Tafel plots (Figure 11) indicated that the ORR kinetics was determined by more than just two parameters, i.e., a single apparent electrochemical rate constant and its associated transfer coefficient, as would be the case if the Tafel plots were linear. Consequently, additional details of the ORR mechanism and the relevant kinetic parameters could potentially be extracted from the steady state (kinetic) RRDE currents using a quasi-steadystate approximation55 to derive kinetic equations for various hypothetical ORR mechanisms and using them to fit the experimental kinetic data. Before the finding of the potential-dependent cathodic transfer coefficient (section 3.2) made the application of the quasi-steady=state approximation55 impractical due to the unknown functional dependence of RnR on E, a number of attempts to fit or model the ORR Tafel plots were made. While these calculations led to incorrect conclusions on the origin of the Tafel slopes at high overpotentials, they also revealed that there was always a range of low to intermediate overpotentials, where a Nernstian slope was observed, whenever oxygen reduction was mediated by a single red
ox center. It is intuitively understood and can also be easily demonstrated within the quasi-steady-state approximation that the Nernstian Tafel slope of a mediated reduction reaction can only be observed when the electrochemical oxidation reaction of the mediator is the fastest process, i.e., at potentials more positive than the standard potential of the mediator and when the other chemical and electrochemical processes are relatively slow. Under such conditions, the overall kinetics is governed by the (low) concentration of the reduced form of the mediator, which is then described to a good approximation by the Nernst equation. The situation is different under non-steady-state conditions, such as in voltammetry of adsorbed oxygen. As no transport step is involved in the reduction of adsorbed oxygen, the final oxidation state of the mediator and the net rate of its electrochemical (and chemical too) oxidation have no effect on the reduction process. Every single red
ox center participates only once in the reduction and transfers as many electrons as required to reduce a single oxygen molecule, i.e., four for the majority of active centers (see sections 3.3 and 3.5). After the reduction of adsorbed oxygen molecule, the mediator participates in the reduction of oxygen diffusing from the bulk of solution, but this contribution was virtually eliminated from the measured voltammograms (section 3.2). In principle, the decrease of the apparent transfer coefficient (RnR) for adsorbed oxygen from 1.56 to ∼0.2 with overpotential (Figure 8) may indicate a change in the reduction mechanism, where the second electron transfer controls the overall rate at low overpotentials, whereas the first electron transfer becomes rate determining at higher overpotentials. However, the extremely low cathodic transfer coefficient makes it difficult to rationalize such an option. Such low transfer coefficients are frequently indicative of dissociative electron transfer processes56 and are rather uncommon for uncomplicated outer sphere56
58 electron transfers or electron transfers followed by a bond formation, e.g., through the protonation of the resulting anion radical. As the first ET is not likely to be associated with any bond-breaking processes, its activation barrier should be rather symmetric, which would result in the transfer coefficient close to 0.5. This is not the case for the second ET. Given the low rates of hydrogen peroxide reduction

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(Figure 12), the interaction of hydrogen peroxide with the active center is quite weak. At the same time, hydrogen peroxide is not a major product of ORR, which indicates that the O
O bond is broken before the third and fourth ET processes take place. Consequently, the O
O bond splitting is most likely associated with the second ET. As the O
O bond is very strong,59 its splitting is the most likely rate-determining step in the whole range of cathodic overpotentials. This conclusion is identical with that reached on similar grounds by Jaouen et al.30 However, our finding of the very low and potential-dependent transfer coefficient for ORR reveals additional details of the reduction mechanism. Consequently, the changes in the apparent transfer coefficient (Figure 8) must reflect corresponding changes in the mechanism of O
O bond splitting. At low overpotentials, the transfers of the first and the second electron are independent processes, as indicated by the value of RnR = 1.56, and the second ET is accompanied by the splitting of the O
O bond. However, RnR significantly drops at high overpotentials. Significant changes in the cathodic transfer coefficient of dissociative electron transfer reactions may occur, when the initially sequential mechanism, i.e., an ET followed by the bond splitting, is gradually replaced by the simultaneous ET and bond-breaking mechanism.56,60 The latter may result in a transfer coefficient significantly lower than 0.5.56,60 If such a mechanism change applied exclusively to the transfer of the second electron and the breaking of the O
O bond in ORR on the present catalyst, the apparent transfer coefficient could decrease with overpotential, but should never drop below one if the first electron was transferred in an independent step. However, the value of RnR drops well below 0.5 at high overpotentials. Consequently, the potential-driven change in the mechanism of reduction of adsorbed oxygen may likely be from the sequential transfer of two electrons and the following dissociation of the O
O bond to the concerted mechanism, where the transfer of two electrons and the O
O bond dissociation occur essentially simultaneously. This hypothesis will be discussed in more detail below, in conjunction with the reversible surface red
ox system. The present data are insufficient to make a definite conclusion regarding the nature of the reversible surface red
ox system, which we identify as the active site. However, based on the results of numerous research studies utilizing high-end material characterization techniques, we postulate that the reversible red
ox couple is FeIII/FeII. Among those studies, some of the most notable employed time of flight secondary ion mass spectrometry (ToF SIMS)26 and M€ossbauer spectroscopy.27,35 The ToF SIMS data26 allowed to link the catalytic activity to two different active sites comprised of N-coordinated Fe,26 whose population strongly depends on the temperature of pyrolysis. In accordance with the hypothesis of N-coordinated iron, Jaouen and Dodelet30 proposed the following mechanism of ORR, based on the known mechanism of ORR on copper/heme oxidases:61
64 e þ Hþ

O2

FeII f FeIII
OO f FeIII Hþ


OOH f • þ FeIV dO e

e þ Hþ

e þ Hþ

þ H2 O f FeIV dO f FeIII OH f FeII þ H2 O ð4Þ Given the significant similarity of the electrochemical behavior of our polyniline-based catalyst to those developed by Dodelet group,30 we presume that the identical or similar oxygen reduction mechanism should be operative in both cases. Consequently, we 6504

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The Journal of Physical Chemistry C Scheme 1. Proposed ORR Mechanism

will try to answer the question if the mechanism presented by eq 4 is compatible with our results. Our kinetic data for the adsorbed oxygen suggest that two electrons must be transferred to the O2 molecule to split the O
O bond, whereas according to eq 4, the bond breakage seems to require only a single electron transfer step. The discrepancy is illusory. In fact, also in Jaouen and Dodelet’s mechanism (eq 4) two electrons are transferred to the molecule before the O
O bond is split. The transfer of the first electron is associated with the reaction between FeII and O2, whereas the second electron transfer is from the electrode to the resulting complex of FeIII with hydroperoxide anion (FeIII
OOH). According to Jaouen and Dodelet,30 the complex of FeIII with hydroperoxide anion may either release hydrogen peroxide molecule upon the protonation (not shown in eq 4) or form the oxo-ferryl cation radical •þFeIVdO (eq 4). The lack of effect of the acid concentration on peroxide generation in the present work remains in agreement with the above proposal. Also, the low activity of the catalyst in hydrogen peroxide reduction remains in agreement with the mechanism outlined by eq 4. However, there also exist discrepancies between our data and the ORR mechanism suggested by Jaouen and Dodelet.30 The latter authors assumed no interaction between dioxygen and N-coordinated FeIII based on the DFT calculations by Anderson and Sidik.65 Our results demonstrate that oxygen adsorption occurs at potentials positive of the formal potential of the surface FeIII/FeII system, i.e., indicate that a relatively strong interaction (ΔG0ads =
15.7 kJ mol
1, see above) between O2 and FeIII may be possible. In principle, the iron center does not have to be directly involved in the adsorption of oxygen, and the adsorption may occur through other surface atoms, e.g., nitrogen, which bond with the iron center. However, a thermodynamically favorable interaction between FeIII and O2 cannot be excluded either. In Anderson and Sidik’s model, ammonia molecules were used as ligands,65 whereas a π-electron rich coordination environment exists on the surface of pyrolyzed oxygen reduction catalysts, where the ligating nitrogen atoms belong to a graphene plane.12,26,35,36,43 Perhaps such an environment may promote oxygen bonding to FeIII. As indicated by the kinetic data for adsorbed oxygen, all reaction steps following the splitting of O
O bond are fast and therefore should have no influence on the overall kinetics of oxygen reduction. Consequently, the correctness of the subsequent steps in the mechanism described by eq 4 cannot be verified. Given the experimental results and the above discussion, we postulate the mechanism of ORR on the pyrolyzed Fe/PANI catalysts as shown in Scheme 1. With a few exceptions, the postulated mechanism is identical to the one proposed by Jaouen and Dodelet30 (eq 4). In the

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mechanism, we incorporated FeIII, whose presence was not explicitly included in the equation proposed by Jaouen and Dodelet30 (eq 4). This change allowed us to take into account the adsorption of oxygen on FeIII. We also omitted the transitional species, oxoferryl cation radical. While we do not question its existence, the independence of the reaction steps associated with the transfer of the second electron in our reaction scheme (formally first in eq 4) is rather questionable at high reduction overvoltages. According to our model, the Nernstian Tafel slope at the potentials positive of the formal potential of the FeIII/FeII red
ox couple (E0Fe(III/II)) under RRDE conditions can be explained by a simple red
ox-mediated mechanism of ORR. The rate of mediated ORR is relatively low, which explains why there is no noticeable decrease of the FeII oxidation peak in the stationary oxygen voltammograms (Figure 1). At higher overvoltages, the mediated mechanism does not apply anymore. Here, FeII is the stable iron species, and oxygen interacts directly with it rather than with FeIII, and the mechanism proposed by Jaouen and Dodelet30 (eq 4) is most likely operative. However, the individual steps contributing to the transition from FeIII—OOH to FeIVdO in eq 4 become progressively less distinguishable with increase in cathodic overpotential. Consequently, a single electrochemical rate constant (kf3) with a potential-dependent apparent transfer coefficient was assigned to that transition. The apparent transfer coefficient (RnR) for this rate constant decreases with the overpotential, and, at the potentials positive of E0Fe(III/II), it is well approximated by the respective plot obtained for adsorbed oxygen (Figure 8). The decrease of RnR indicates the transition from the mechanism with well-separated reaction steps to the mechanism, where the splitting of the O
O bond occurs simultaneously with the charge transfer. The same transition results in the potential-dependent reaction order for Hþ (see section 3.6). It has to be noted that the major steps of the outlined ORR mechanism (Scheme 1), while proposed under the reasonable assumption of surface iron being an active center, are not iron specific. Both the red
ox-mediated catalysis and the concerted electron transfer and O
O bond splitting could be postulated based on the experimental results if the surface red
ox system was not FeIII/FeII. Also, while the proposed mechanism offers a good explanation of the experimental findings and is consistent with other experimental and theoretical studies, it does not exclude alternative explanations of the data presented in this paper. While the adsorption of oxygen on the oxidized active site does not play a very important role in ORR catalysis, the high intrinsic Tafel slopes of ORR are a drawback. In principle, the problem can be mitigated by increasing the surface density of the active sites and also the specific surface area of the catalyst. The most active catalyst (batch 1) studied in this work had a specific surface area of 400 m2 g
1, and the surface iron atoms occupied around 1% of this surface. Simple geometric considerations based on the proposed structure of the active site12,26,35,36,43 reveal that the surface density of the active sites can be increased by no more than 1 order of magnitude, which is equivalent to a 60 mV reduction in ORR overpotential at high cathode potentials. Even less can possibly be achieved by an increase in the specific surface area of the catalyst. As demonstrated by Wu and co-workers for the Fe/polyaniline-based catalysts utilizing Black Pearls 2000 as a high surface area (1440 m2 g
1) carbon support, the dominant microporous structure of the support results in serious water flooding, which leads to subsequent irreversible damage of the cathode catalyst layer.66 6505

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The Journal of Physical Chemistry C A more meaningful improvement of the catalyst performance would be possible if the ORR Tafel slopes at higher overpotentials could be decreased. Within the framework of the proposed mechanism, lower Tafel slopes could be expected if the simultaneous electron transfer and splitting of O
O bond were faster. Are such kinetic changes possible? According to Tributsch and co-workers,35,36 single iron (transition metal) center cannot effectively transfer four electrons to oxygen molecule. Consequently, they postulate that a neighboring atom (group of atoms) facilitates ORR on the pyrolyzed ORR catalysts. Various favorable effects of surface atoms and groups of atoms have been observed for ORR catalyzed by nonpyrolyzed transition metal porphyrins.67
69 Recently, Elbaz and co-workers demonstrated that the transition metal center in CoIII porphyrin can be coordinated by benzoquinones adsorbed on glassy carbon and by quinone functionalities on the surface of aerogel carbon.70 The subsurface bond of the metal center resulted in an effective catalysis of ORR by the porphyrin with virtually no hydrogen peroxide production.71,72 Based on the sixfold coordination of iron, likely present in some pyrolyzed Fe/C/N catalysts,35 a subsurface coordination of the metal center (“interaction with an adjacent graphene layer”) was also suggested for those materials.35 The hypothetical subsurface bonding of the metal center in the nonprecious-metal-based heat-treated catalysts seems to agree with their higher stability in strongly acidic media compared to related macrocyclic complexes of transition metals.73
75 Viewed from the above perspective, studies of the immediate environment of transition metal centers in these catalysts, including possible subsurface coordination bonds, may help understand better the respective mechanisms of ORR and possibly further improve the catalytic activity of the pyrolyzed nonprecious ORR catalysts.

’ AUTHOR INFORMATION Corresponding Author

E-mail [email protected].

’ ACKNOWLEDGMENT Thanks are due to Dr. Gang Wu for providing the catalyst samples and numerous valuable discussions. Financial support from the U.S. DOE Fuel Cell Technologies Program and the Advanced Cathode Catalysts project is gratefully acknowledged. ’ REFERENCES (1) Jasinski, R. Nature 1964, 201, 1212–1213. (2) Jasinski, R. J. Electrochem. Soc. 1965, 112, 526–528. (3) Jahnke, H.; Sch€onborn, M.; Zimmermann, G. Organic dyestuffs as catalysts for fuel cells. In Physical and Chemical Applications of Dyestuffs. Top. Curr. Chem. 1976, 61, 133–181. (4) Zagal, J. H. In Handbook of Fuel Cells: Fundamentals, Technology and Applications; Vielstich, W., Lamm, A., Gasteiger, H., Eds.; John Wiley & Sons, Ltd.: Chichester, U.K., 2003; Vol. 2: Electrocatalysis, pp 544
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