E
ffect of pH on the Activity of Platinum Group Metal-Free Catalysts
in Oxygen Reduction Reaction
Santiago Rojas-Carbonell, Kateryna Artyushkova, Alexey Serov, Carlo Santoro, Ivana Matanovic,
and Plamen Atanassov
*
Department of Chemical & Biological Engineering, Center for Micro-Engineered Materials (CMEM), University of New Mexico, Advanced Materials Lab, 1001 University Blvd SE, Albuquerque, New Mexico 87131 United States
*
S Supporting InformationABSTRACT: The impact of the electrolyte’s pH on the catalytic activity of platinum group metal-free (PGM-free) catalysts toward the oxygen reduction reaction (ORR) was studied. The results indicate that the ORR mechanism is determined by the affinity of protons and hydroxyls toward multiple functional groups present on the surface of the PGM-free catalyst. It was shown that the ORR is limited by the proton-coupled electron transfer at pH values below 10.5. At higher pH values (>10.5), the reaction occurs in the outer Helmholtz plane (OHP), favoring hydrogen peroxide produc-tion. Using a novel approach, the changes in the surface chemistry of PGM-free catalyst in a full pH range were studied by X-ray photoelectron spectroscopy (XPS). The variations in the surface concentration of nitrogen and carbon species are
correlated with the electron transfer process and overall kinetics. This study establishes the critical role of the multitude of surface functional groups, presented as moieties or defects in the carbonaceous “backbone” of the catalyst, in mechanism of oxygen reduction reaction. Understanding the pH-dependent mechanism of ORR provides the basis for rational design of PGM-free catalysts for operation across pH ranges or at a specific pH of interest. This investigation also provides the guidelines for developing and selecting ionomers used as “locally-confined electrolytes”, by taking into account affinities and possible interactions of specific functional groups of the PGM-free catalysts with protons or hydroxyls facilitating the overall ORR kinetics.
KEYWORDS: oxygen reduction reaction, PGM-free catalysts, pH-dependent activity, XPS, mechanism
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INTRODUCTIONFuel cells transform the chemical energy contained in the fuel into useful electricity. Commonly, the fuel cells utilize oxygen as an oxidant due to its high redox potential and its availability in the atmosphere at practically no additional cost.1The oxygen reduction reaction (ORR) is one of the most studied electrochemical processes due to its importance in living organisms as well as in electrolytic and power generation technologies. The ORR is of particular importance to fuel cell technology, where it is used across a wide range of temperatures and in extreme pH regions: acidic2,3 and alkaline.2,4 New biological fuel cells and bioelectrochemical reactors are emerging as a technology that utilizes the ORR in a broad range of pH of various enzymes, supramolecular assemblies of complex biocatalysts, and microorganisms and their communities (biofilms).5−7
In acidic media, Pt and other platinum-group-metal (PGM) catalysts exhibit high efficiency in the reaction of oxygen reduction.8−10 In alkaline media, Pt has excellent electro-catalytic activity toward ORR, but it very sensitive to poisoning by anions and consequently, prone to losing its activity.11Also,
Pt cannot be utilized under conditions at which pollutants are present (e.g., microbial fuel cells) due to fast poisoning.12−14 The high cost of PGMs and poisoning issues led to extensive research in the attempt to identify readily available low-cost alternative materials with performances comparable to that of platinum.15
From the alternative PGM-free catalysts, two groups of materials stand out as suitable for applications across the range of pH: (i) metal-free carbon−nitrogen systems N−C16,17
and (ii) the M−N−C type of materials where PGM-free transition metal is incorporated into the carbon−nitrogen matrix (M = Fe, Co, Ni, etc.).2,18−20 Carbonaceous materials have several characteristics that make them promising alternatives to PGM catalysts, such as high surface area, high chemical stability, high electrical conductivity, low-cost, and commercial availability. Moreover, carbonaceous materials are resistant to the influence of pollutants.21,22 The catalytic activity of these materials Received: November 22, 2017
Revised: February 18, 2018 Published: March 1, 2018
pubs.acs.org/acscatalysis
toward ORR, however, is quite low in acidic16,17 and alkaline media,16,17 while in neutral media the output can be comparable to those of PGM-free M−N-C materials.22−24
Therefore, presently carbonaceous materials are mainly used as catalyst supports rather than as catalysts in both acidic and alkaline media.25−28
In PGM-free M−N-C catalyst, earth-abundant transition metals M such as Mn, Fe, Co, or Ni are atomically dispersed within the carbon−nitrogen rich matrix, designated as N− C.29−31These are usually synthesized from metal, nitrogen, and carbon-containing precursors through high-temperature treat-ment (pyrolysis).32−39 M−N−C catalysts have demonstrated good performances in acidic media,40−42 outstanding and unprecedented results in neutral media,30,43−45and comparable and even superior to Pt performance in alkaline media.46−49
As mentioned before, three types of fuel cells based on different pH values are proton exchange membrane fuel cells (PEMFCs; pH∼ 1), microbial fuel cells (MFCs; pH ∼ 7), and alkaline membrane fuel cells (AMFCs; pH ∼ 13). The mechanism of the ORR for different pH media differs.50 For example, in acidic media, protons are combined with oxygen at the cathode in several alternative pathways. The ORR can follow a 2e−transfer mechanism with the production of H2O2, a 4e− transfer mechanism with H2O as a final product (full reaction as 1/2O2+ 2H++ 2e−→ H2O51), and a sequential 2 × 2e−mechanism in which hydrogen peroxide is being reduced by the same or a different catalytic site with the production of water.52H2O2produced during the first step of the ORR can also be converted to water through chemical decomposition.53 The role of different nitrogen and metal surface moieties in these three types of mechanisms in acidic media was reported recently.52In alkaline media, OH− plays a crucial role in the ORR mechanism, being its primary product. The reaction can occur via a 2e− mechanism, forming HO2− and OH−, or a direct 4e− mechanism, with the generation of OH−51 (full reaction as O2+ H2O + 2e−→ 4OH−). A sequential 2× 2e− mechanism can also take place.2,4
Another difference between ORR at different pH levels is based on inner-sphere and outer-sphere reactions. In alkaline media, surface hydroxyl groups promote surface-independent outer-sphere electron transfer,54 while in acidic media immediate involvement of active sites such as metal coordinated to nitrogen in inner-sphere electron transfer is of direct relevance.55,56
The mechanism of ORR in neutral media is far less studied and, probably, the most challenging due to the limited availability of both H+ and OH−. Recently, Malko et al. have conducted a pH study on the Fe−N−C catalyst and it was shown that the reaction follows the acidic-type mechanism until pH of 11, after which the mechanism switches to the alkaline-type.57
It is known that M−N−C catalysts consist of different types of nitrogen (e.g., pyridinic, pyrrolic, graphitic, metal−nitrogen directly coordinated, etc.) as well as different types of metal-containing (atomically dispersed and zerovalent metal-rich phases) moieties that contribute differently to the electron transfer mechanism.52,58Moreover, the type and abundance of surface oxygenated species, such as−OH, −OC, and O, and the degree of protonation of nitrogen may also be affected by changes in proton concentration.59−65 The influence of the changes in the catalysts surface chemistry caused by different pH levels on the ORR performance has not yet been spectroscopically identified. The effects of the chemistry and
morphology of the PGM-free catalysts have been previously studied for extreme pH values, mainly for pH values arround 1, 7, and 14.31,57,66−69 To the best of our knowledge, a comprehensive study that explains the connection between the change in the electrolyte’s pH, its effect on interfacial chemistry spectroscopically measured, and the M−N-C catalytic activity has never been presented.
In this study, Fe−N−C catalyst was synthesized using the sacrificial support method (SSM)70utilizing a charge-transfer organic salt, nicarbazin, as the organic precursor containing both N and C atoms.71 The dependence of the surface chemistry of the catalyst on pH was studied by X-ray photoelectron spectroscopy (XPS). Electrochemical measure-ments using the rotating ring disk electrode (RRDE) technique were performed in 18 different electrolytes with pH values covering the entire range from pH 1.1 to pH 13.5. For thefirst time, an unambiguous relationship between electrolyte pH, surface chemistry, and catalyst ORR performance has been established. The results of this work may impact not only traditional low-temperature fuel cell technologies, such as polymer electrolyte membrane fuel cells (PEMFC) and alkaline membrane fuel cells (AMFC) and corresponding electrolyzer technologies but also a broad spectrum of newly introduced bioelectrochemical devices such as biofuel cells, microbial fuel cells, and microbial bioreactors.
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EXPERIMENTAL SECTION■
CATALYST SYNTHESISThe PGM-free catalyst used in this study is iron-nicarbazin derived and prepared following the sacrificial support method based on the procedure described before72and used in practical applications, including automotive PEMFC and AMFC. The details of the synthesis are included in the Supporting Informationsection.
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ELECTROCHEMICAL MEASUREMENTSThe rotating ring disk electrode (RRDE) technique was used to assess the electrochemical performance of the PGM-free catalyst. An optimized loading (Figure S1) of 175μg cm−2of catalyst was obtained by depositing ink, made from a suspension of 5 mg of catalyst in 850 μL of isopropanol and 150μL of Nafion (0.5 wt % in isopropanol), onto the mirror polished glassy carbon disk. This loading ensured a complete coverage of the glassy carbon.
The electrolytes used were prepared by selecting the appropriate buffer solution that would provide buffering capacity for the required pH range. The pH values were measured using an Orion Star A111 pH meter (Thermo Scientific). The buffer solutions used are shown inTable 1.
The RRDE cyclic voltammetry (CV) was carried out using WEB30-Pine bipotentiostat and a Pine Instruments Rotator (Pine Instruments, Raleigh, NC) as described in detail in the
Supporting Information. The potential for the ring was selected to be within the region where the water is thermodynamically stable and in the region where the peroxide is electrochemically oxidized as shown in the Pourbaix diagram inFigure S2.
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CHEMICAL CHARACTERIZATIONThe surface chemistry of PGM-free catalyst was analyzed with a XPS Kratos Ultra DLD spectrometer as has been described before.52 A subset of the electrolytes was used to make a
suspension of catalysts to study the chemistry of catalyst resulting from interacting with different buffer solutions. This was attained by adding 100μL of the electrolyte to 1 mg of catalyst, which was then left to interact for 2 days. The XPS spectra from the catalyst and dried suspension of catalysts with electrolytes were obtained. Three areas per sample were analyzed. High-resolution N 1s, C 1s, and O 1s spectra were processed using CasaXPS software.
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RESULTS: SURFACE CHEMISTRYHigh-resolution XPS experiments were performed to under-stand the surface chemical composition of the catalyst.Figures S3 and S4show high-resolution N 1s and C 1s spectra obtained and fitted with individual peaks according to the previously established curvefit.45,73,74There are multiple types of nitrogen in this catalyst, such as pyridinic, nitrogen coordinated to iron Nx-Fe, pyrrolic or hydrogenated, graphitic, quaternary, and multiple types of nitrogen coordinated with oxygen. Protonated nitrogens such as protonated pyridine are also detected in pristine catalyst powder. Carbon consists of graphitic carbon, aliphatic carbon, carbon coordinated to nitrogen, and different types of surface carbon−oxygen species such as phenolic (−C− OH), lactone and pyrone (−CO), and carboxylic (−COOH).
There are several important considerations for understanding the influence of the electrolyte pH on the oxygen reduction reaction. Thefirst is the prevalence and concentration of either protons or hydroxyl ions in the electrolyte, which determines the inner- vs outer-sphere electron transfer mechanism of the reaction. A clear distinction between these mechanisms is expected at high and very low pH values while in the middle range,75between 5 and 9, this separation is unclear and both mechanisms can be occurring during ORR. Excess of either hydronium ions or hydroxyls can also result in their specific binding to the moieties at the surface of the catalyst, therefore modifying possible active sites.
The second consideration is protonation/deprotonation of nitrogen and carbon−oxygen groups that are present on the surface of the catalyst layer. Several groups59−65,75−82 have investigated the acid dissociation constants (pKa) displayed by
carbon-based materials that contain nitrogen and oxygen species. The values of pKa associated with different oxygen and nitrogen-containing moieties that can lose or accept proton are summarized inTable 2. At pH values higher than the pKa,
the nitrogen and oxygen species loses its protons and becomes either anionic or neutral. Oppositely, at pH values lower than the pKa, that chemical species will capture protons, becoming neutral or cationic. Based on these pKa values, it can be predicted the pH ranges at which surface carbon and nitrogen moieties will change their state, affecting, therefore, the affinity of oxygen toward binding and the type of possible oxygen reduction steps in which these species participate. Figure 1
graphically represents the expected effect the pH has on the chemistry of functional groups of the PGM-free catalyst.
Within the f irst pH region (up to pH 2.5) all the functional groups on the surface of the catalyst are expected to be protonated, except the graphitic nitrogen (Figure 1(a)). At the same time, excess hydronium ions present in the electrolyte can interact via specific hydrogen bonding with the lone electron pair of nitrogen atoms coordinated to iron in Fe−NXcenters. For clarity,Figure 1(a) shows the hydrogen bonding only for one of the iron centers. However, it is expected to occur for the majority of them.
In the second pH region (Figure 1(b)) between pH 2.5 and 5.3, successive deprotonation of carboxylic acid occurs. Moreover, due to the lower concentration of protons in the electrolyte, the specific adsorption of hydronium ions onto the nitrogen of the Fe−NXcenters is less probable.
The third pH region corresponds to the range between 5.3 and 7.5. This region is characterized by deprotonation of the pyridinic nitrogen moieties. (Figure 1(c)).
The fourth pH region is between pH 7.5 and 9, in which the pyrone-type groups act as a base by capturing protons from the electrolyte (Figure 1(d)).
The f if th pH region of changes in the surface chemistry of the PGM-free catalyst is caused by the deprotonation of the phenolic groups (between pH 9 and 10.5) (Figure 1(e)).
At last, at pH values above 10.5 (sixth pH region), all the functional groups but the hydrogenated ones should be deprotonated, leading to a significant dependence of the ORR on the increased concentration of hydroxyl groups (Figure 1(f)). At high pH, hydroxyl groups can specifically adsorb onto the positively charged atoms, such as nitrogen coordinated to iron centers and quaternary nitrogen, making a double layer effect dominant in the mechanism of oxygen reduction.
To confirm the changes in the surface composition expected due to different pH levels as suggested inFigure 1, the chemical composition of catalysts mixed with a subset of 10 electrolytes and DI water was studied by XPS. The pH of all suspensions containing catalyst and electrolyte measured after 2 days of Table 1. Buffers Used as Electrolytes for Each pH Value
pH Constituents pH Constituents
1.1 Phosphoric acid 6.1 Potassium diphosphate, potassium
dihidrogenphosphate 1.3 Sulfuric acid 7.2 Potassium diphosphate,
potassium dihydrogen phosphate
1.6 Perchloric acid 8.4 Potassium diphosphate, potassium
dihidrogenphosphate 2.4 Phosphoric acid, potassium
dihydrogen phosphate
9.6 Sodium bicarbonate and sodium carbonate 2.8 Citric acid-potassium
dihydrogen phosphate
9.8 Boric acid, potassium hydroxide 3.6 Malic acid, potassium
hydroxide
10.6 Sodium bicarbonate and sodium carbonate 4.6 Malic acid, potassium
hydroxide
11.2 Boric acid, potassium hydroxide 5.2 Acetic acid, potassium
acetate
12.5 Boric acid, potassium hydroxide 5.5 Potassium diphosphate,
potassium dihidrogenphosphate
13.5 Potassium hydroxide
Table 2. Acid Dissociation Constants for the Functional Groups of PGM-free Catalysts
mixing was found to be the same as that of the corresponding buffer. For catalyst suspension in DI water, the pH was 7.65, indicating the slightly basic character of the catalyst due to the strong tendency of surface nitrogen and carbon−oxygen groups toward protonation.
The XPS as an analytical technique has a unique sensitivity to intermolecular interactions such as hydrogen bonding and chemisorption. The ability of ex situ XPS to detect adsorbates such as intermediates of the catalytic reaction on the surface of the material has been earlier demonstrated.83−85Here, much stronger effects than physical adsorption due to the intermolecular interaction of the catalyst with constituents of buffers are expected. Changes in the electron density on the nitrogen atom due to the participation in intermolecular bonds will be reflected in the shift in the binding energy of the 1s electrons of the nitrogen atoms. Electrostatic bonding of opposite charges can shift the position of the nitrogen energy by 1−2 eV, while the protonation of the nitrogen results in a shift toward higher binding energy up to 2.5−3 eV.85−87It was previously demonstrated that photoelectron spectroscopic
analysis is successful in determining the pKa values of the protonated amines and carboxylic groups in biomolecules.74
In suspensions of the catalyst with electrolytes, the atomic percent of nitrogen detected may decrease due to the contribution of elements from the buffer into the total elemental composition. However, there is no decrease or increase in the absolute amount of nitrogen in the catalyst itself upon mixing with the buffer expected. On the other hand, due to protonation and intermolecular bonds formation with electrolyte, a rearrangement in the distribution of different peaks is observed, which originates from the contribution of different types of nitrogen to the total nitrogen spectrum. To evaluate the protonation and intermolecular bonds formation, high-resolution nitrogen spectra of inks made with buffers of specific pH values were curve-fitted using the peak set established for the catalyst itself as shown in Figure S3. The surface concentrations of species identified in the N 1s spectrum are plotted in Figure 2 in relative percentage as a stacked bar totaling to 100% of total nitrogen and as combined plots of relative percentage of N species as a function of pH. Individual plots are shown inFigure S5for clarity.
Figure 1.Schematic of the suggested surface chemistry for the PGM-free catalyst at corresponding pH range, as described in the text: (a) highly acidic (pH < 2.5), (b) slightly acid (2.5 < pH < 5.3), (c) neutral (5.3 < pH < 7.5), (d) slightly alkaline 1 (7.5 < pH < 9), (e) slightly alkaline 2 (9 < pH < 10.5), and (f) highly alkaline (pH > 10.5). The interrupted red lines in the upper right corner of the chemical structure signify that this is a representation of the functional groups that are present in a long-range graphitic-like structure.
Figure 2shows that the variation in the electrolyte pH leads to significant changes in the concentrations of detected nitrogen surface species of the catalysts. It is important to note the existence of pH regions that correspond to the groups identified inFigure 1is associated with specific changes in the chemical species related to their corresponding acid dissocia-tion constants. The relative surface concentradissocia-tion of pyridinic nitrogen steadily increases from the value at lower pH to the plateau around neutral/slightly alkaline pH ∼ 9 (Figure 2.b). Significant protonation of pyridine occurs at pH lower than its pKa, while at pH above the pKa of 6.5, deprotonation of pyridine occurs, explaining the increase in the relative amount of pyridinic nitrogen.59,82 For the highest pH (13.5), the specific adsorption of hydroxyl groups to pyridine nitrogens leads to the significant decrease in their relative surface concentration. The samefigure plots the changes in the surface concentration of hydrogenated nitrogen with the pH. This peak value includes pyrrolic nitrogen and hydrogenated pyridinic nitrogen.88 It should be noticed that pyrrolic nitrogen is hydrogenated within the full pH range due to its high pKa. A decrease in the surface concentration of hydrogenated nitrogen with pH is a result of dehydrogenation of the pyridinic nitrogen moieties and their conversion to unprotonated pyridine. The inverse relationship between trends of pyridine and hydro-genated nitrogen with pH confirms the deprotonation of pyridine at the highly alkaline pH of 12.
The surface concentration of nitrogen coordinated with iron as a function of the pH is shown in Figure 2(c). A gradual decrease in surface concentration of Fe−NX centers with the transition from acidic to alkaline pH is observed mainly due to specific adsorption of hydroxyls ions on Fe−NX centers. The adsorbed hydroxyls may serve as inhibitors of oxygen binding to these moieties. Specific binding of hydroxyl ions onto Fe coordinated to nitrogen was previously reported.89−92
The next peak in Figure 2(c) has a contribution from both graphitic nitrogen and protonated nitrogen Ngr/N+. The intensity of this peak is the largest at the lowest pH mainly
due to the highest amount of protonated nitrogens. The intensity of this peak decreases at higher pH values as nitrogens become deprotonated. Figure 2(d) shows the contribution from other types of cationic nitrogens, such as quaternary and protonated graphitic nitrogen species present in the carbon matrix. This is the least abundant type of nitrogen present, accounting for 2−6% of the total surface nitrogen detected.
Lastly, the species whose contribution increases the most upon transition from acid to alkaline pH are nitrous oxides (Figure 2(d)). As the concentration of OH− ions in the solution increases, there is an increase in their specific adsorption on the surface moieties with affinity to hydroxyls, such as atomically dispersed iron−nitrogen centers and other nitrogen moieties as was discussed above.
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RESULTS: ELECTROCHEMICAL ACTIVITYThe results from the linear sweep voltammetry of the catalyst at 18 different pH values are presented inFigure 3. The peroxide yield and the overall number of electrons transferred in ORR were calculated using the ring and disk current densities (Figure S6). There are substantial changes in the shape of the voltammograms observed with the change of pH of the electrolyte, consistent with the different surface chemistries at pH regions defined inFigure 1and spectroscopically observed inFigure 2.
It was previously reported that the heterogeneous decom-position of the hydrogen peroxide can affect the interpretation of the results from the RRDE experiments.53 In the current report, no evidence of such decomposition was observed, confirmed by the linearity of the Koutecky−Levich plots (Figure S7).
The interpretation of the ORR reaction mechanism at different pH regions has to take into account hydrogen peroxide yield and number of electrons transferred (Figure S6). Low hydrogen peroxide yield may be explained by a complex, combined serial and direct reaction pathway that could possibly Figure 2.Surface concentration expressed as relative percentage of chemical species at different pH levels: (a) overall distribution of nitrogen, (b) pyridinic N and hydrogenated N−H, (c) Nx-M and graphitic N, and (d) protonated N and nitrous oxides NOx.
be supported at a multitude of alternative active sites catalyzing ORR with comparable rates:
1. If no active sites producing hydrogen peroxide exist, 4-electron“direct” reduction of oxygen to water will take place.
2. If a limited amount of species producing hydrogen peroxide are present, a 2 × 2 electron transfer mechanism will be partially displayed, resulting in slightly less than 4 electrons transferred as calculated.
3. In the case of substantial presence of surface species, catalyzing the production of hydrogen peroxide along with a comparable amount of species reducing all/most of the generated peroxide further to water, a typical 2× 2 bifunctional (two separate sites) electron transfer mechanism will be established with less than 4 electron transferred calculated and a strong dependence on catalyst loading.
The higher hydrogen peroxide yield may be a result from 4. A large amount of species producing hydrogen peroxide
with a small concentration of the species reducing H2O2 to water, demonstrated as 2× 2 electron transfer with less than 4 electrons transferred.
At highly acidic pH values (Figure 3(a)), the oxygen reduction produced small amounts of peroxide indicating either a direct 4 electron mechanism or almost complete reduction of all generated hydrogen peroxide to water. The LSV corresponding to pH 1.1 was obtained in a phosphoric acid electrolyte. Other studies have proposed a possible inhibition effect of phosphate onto the ORR catalytic activity of the PGM-free catalysts,93 as well as Pt catalyst,93,94 which is indeed observed in comparison with the voltammograms obtained in other electrolytes within this highly acidic region. As the pH increased to mildly acidic (Figure 3(b)), the peroxide yield increased, reaching a first maximum at pH of 3.6, and then decreased as the pH reached a neutral value. This observation points toward a change in the ORR mechanism with thefirst step of oxygen reduction to hydrogen peroxide being more efficient than the H2O2reduction to water. As discussed later, this variation in ORR activity corresponds to the changes in the surface chemistry of the carboxylic acid groups, as these moieties transition from being protonated to being deproto-nated in this pH range. Around neutral pH (Figure 3(c)), low ring current densities, low peroxide yields, and overall 4 electrons transferred are again observed (Figure S6(c)). Figure 3.Graphic representation of ring and disc current densities vs potential in the ORR at different pH values of the electrolyte. Scan rate of 5 mV s−1, 1600 rpm, oxygen saturated electrolyte. Each graph is associated with a particular pH range, as described in the text: (a) highly acidic (pH < 2.5), (b) slightly acid (2.5 < pH < 5.3), (c) neutral (5.3 < pH < 7.5), (d) slightly alkaline 1 (7.5 < pH < 9), (e) slightly alkaline 2 (9 < pH < 10.5) and (f) highly alkaline (pH > 10.5). The pH was buffered with the electrolytes listed inTable 1.
As the pH increased to the mildly alkaline (Figure 3(d) and (e)), the peroxide generation occurred at much higher overpotentials. In this region, the highest overpotential and lowest peroxide yield were registered at pH 9.8.
As pH further increased to highly alkaline (Figure 3(f)), the peroxide yield increased as the overpotential for its generation decreased, directly indicating a shift to a mechanism that proceeds predominantly via a 2× 2e− transfer, with peroxide generation being a more efficient reaction. Still, the high overall number of electrons transferred (>3.4), as shown in Figure S6(c), clearly indicates that most of the produced peroxide is further reduced to water by the active sites. At pH 13.5, a significant decrease in the limiting current density is accompanied by an increase in the peroxide detected in the ring, pointing toward the specific adsorption of hydroxide ions onto the iron centers, as has been previously proposed by others.31,54,67 This effect leads to a prevalence of an OHP reaction, where to be further reduced to water the intermediate peroxide needs to access other moieties. This second step is successfully occurring at the surface of the catalyst, as the overall number of electrons transferred at pH 13.5 is 3.4 as measured by the relationship of disk and current density. The overall ORR activity at pH 13.5 is higher and has lower onset potential. Nevertheless, when comparing the ring current
density, it is evident that peroxide is detected at much lower overpotentials for pH 13.5 when compared with pH 9.8.
Li et al.68 have proposed a model to establish the pH dependency for nitrogen-doped graphitic carbon materials. Although they proposed that the mechanism in the acid media follows a 2-electron transfer process, we found from the ring current densities ofFigure 3(a) that the peroxide detected is much lower than in the case of the alkaline electrolytes (Figure3(f)), where they claim that the ORR occurs in a 4-electron transfer manner. Ourfindings are in agreement with the results obtained by Liang et al.67 on nitrogen doped carbons.
The potential at a current density of 100μA cm−2falls within the potential window of a kinetically limited region, allowing assessment of the variations in the kinetics of the ORR with pH. The half-wave potential (E1/2) is a commonly used descriptor to compare the catalytic activity toward the ORR.95,96 These two metrics provide different information depending on whether they are calculated versus RHE or Ag/ AgCl reference electrode. The first one eliminates the dependence of potential from pH. Both E@100μA cm−2and E1/2are presented inFigure 4. Even for E1/2calculated versus RHE electrode (Figure 4(a)), a pH dependence is observed, particularly, in the direction toward the alkaline environment. Figure 4.Potential at which the ORR current density reaches 100μA cm−2vs (a) RHE and (c) Ag/AgCl electrode. The half-wave potential for the ORR vs RHE electrode (b) and vs Ag/AgCl electrode (d). Each of the shaded regions corresponds to the six pH regions identified inFigure 1.
For potentials reported vs the Ag/AgCl electrode (Figure 4(b) andFigure 4(d)), there is an expected 59 mV per one unit of pH change that is derived from a different proton concentration. For potentials reported vs RHE (Figure 4(a)), there is no pH dependence of the potential observed in the case of pH values below 2.5. At pH higher than 2.5, the potential dependence on pH indicates that measured potential is not solely dependent on the proton concentration. The half-wave potential also confirms the existence of different pH regions (Figure 4(b)).
These results indicate the differences in the electrocatalytic ORR mechanism when the pH of the electrolyte is varied. The changes in the surface chemistry proposed inFigure 1 are in good agreement with the shifts in the trends observed in the kinetic region of potentials ofFigure 4.
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DISCUSSION: ORR ACTIVITY DESCRIPTORS OFTHE PGM-FREE CATALYSTS
To obtain a better insight into how changes in the catalyst’s surface chemistry impact the ORR mechanism, the electron transfer and kinetic parameters for ORR were calculated and presented in Figure 5. The shaded regions correspond to six different pH regions described inFigure 1, which are associated with changes in the surface chemistry of the catalyst as a function of pH.
Figure 5(a) indicates a logarithmic decrease of the exchange current density as pH changes from acidic toward neutral. The observed decrease is associated with the limitation of the electron transferred to the protons and hydroxyls, which are present in much lower concentrations at neutral pH. A logarithmic decrease in the exchange current density for pH higher than 10.5 points toward the selective adsorption of hydroxyl ions onto the catalyst surface, with the introduction of the additional electron transfer barrier, which is required for the outer Helmholtz plane (OHP) mechanism. The hydroxyl adsorption at the surface was observed by XPS as discussed above.
Figure 5(b) depicts the electron transfer coefficient as a function of the electrolyte pH. This parameter indicates the magnitude of the ORR overpotential required for surface species to obtain the necessary activation energy sufficient to the initiation of the ORR. The charge transfer coefficient is the lowest for acidic pH and increases with higher pH. It has been proposed that the charge transfer coefficient indicates whether the rate-determining step (RDS) is determined by a proton-coupled electron transfer (values around 0.5) or by the availability of reactant species, i.e., protons or hydroxyls, at the electrode (values deviating the 0.5 value).97For highly alkaline and neutral pH values, the RDS is caused by the availability of the H+/OH− species. At neutral pH, protons have to have access to the catalyst surface for the ORR to be carried out, and for highly alkaline pH, the hydroxyl ions need to reach the boundary of the outer Helmholtz plane to initiate the oxygen reduction reaction. Liang et al.67 have suggested the direct participation of surface species toward the ORR at acid pH values and the simultaneous inner and other Helmholtz plane reactions at highly alkaline pH values. Our reaction mechanism analysis is homologous and yielded comparable results (Figure 5) to the ones obtained by these authors. The present paper goes into further detail about what are these chemical species, as it correlates surface chemistry changes measured spectro-scopically with the electrochemical results obtained and includes the iron species of the Fe-MNCs catalysts.
The ORR kinetic parameters were determined from the Koutecky−Levich (K-L) equation as described in the
Supporting Information. The calculated kinetic current densities are displayed inFigure 6. In the highly acidic region,
there is a difference of about an order of magnitude between the kinetic current densities obtained in phosphoric acid versus perchloric and sulfuric acids (Figure 3(a)). Phosphoric acid is known to be an inhibitor of ORR for iron-containing porphyrins and iron-free nitrogen−carbon composites.98,99 Lower transport-limited current density is also observed when phosphoric acid is used. Dihydrogen phosphate ion interacts with the hydrogenated species of the PGM-free catalyst.99On the other hand, the activity of the PGM-free catalyst is decreased to a much lesser extent than in the case of platinum.100
For the second region of pH (from 2.5 to 5.3), where changes in the surface chemistry relate to carboxylic acid species, the kinetic current density decreases (Figure 6). This can be explained by the ability of carboxylic groups to carry out the reduction of oxygen to peroxide, as is evidenced by the increase in the ring current density shown inFigure 3(b), at the pH region close to the pKavalue of these species.
A significant increase in the kinetic current density is observed for the third pH region, which corresponds to pH values where the pyridinic nitrogen is most sensitive to proton concentration due to its pKa value (Figure 6). It was shown before that pyridinic nitrogen plays a critical role in the ORR catalytic activity of the PGM-free catalysts.52 The pyridinic nitrogen has a pKaclose to the neutral pH at which availability of protons is limited. Pyridines that exist in their protonated form become deprotonated at the pH around its pKa value providing protons necessary for ORR. The increase in the kinetic current density caused by the pyridinic nitrogens around neutrality might come from the proximity of their pKa to the neutral value. Since the pKa value of pyridinic nitrogen is slightly lower than 7, it is possible for the pyridinic nitrogen to deprotonate the water molecules via an acid−base coupling, increasing the availability of protons for the ORR. This can be interpreted as a cocatalysis effect of the pyridinic nitrogen species and explain multiple observations of positive correlation of their abundance with higher activity of PGM-free catalysts in ORR.
This effect diminishes with the further pH increase. In the fourth andfifth pH regions, where the changes in the lactone-pyrone and phenolic functional groups are expected, the kinetic current density decreases (Figure 6). The effect of the abovementioned groups on the ORR activity is significantly less pronounced compared to the one caused by the pyridinic nitrogen. The Koutecky−Levich slope, as well as the logarithm Figure 6.Logarithm of the kinetic current density as a function of the change in pH.
of the ratio of the kinetic current density and the electron transfer current density, is analyzed in more detail inFigure S8.
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DISCUSSION: CORRELATION BETWEEN SURFACECHEMISTRY AND ELECTROCHEMICAL ORR ACTIVITY
It was reported in the previous studies that the Fe−NXcenters are the active sites that bind oxygen and reduce it via a direct four electron transfer mechanism.29,52,101,102Therefore, it is of interest to investigate the correlations between the chemistry of this particular type of nitrogen, the chemistry modified by the variations of proton concentration, and obtained electro-chemical activity. The relationship between the relative surface concentration of Fe−NX centers and electrochemical parame-ters at different pH values is shown inFigure 7.
Figure 7(a) shows that the extreme pH values (acidic and alkaline) led to similar surface concentrations of Fe−NXcenters being significantly lower compared to the other pH values. However, electron transfer parameters observed at extreme pH values are drastically different. This can be explained by (i) the specific adsorption of protons onto the nitrogen atoms coordinating the iron centers; (ii) the specific adsorption of hydroxyls on the iron centers that occurs at higher pH affecting the electron transfer process between the catalyst and the reactive species (Figure 1(a)).
Figure 7(a) shows that the increase in the surface concentration of the Fe−NX centers causes a decrease in the electron transfer coefficient. Importantly, the surface concen-tration of these centers is higher at lower pH. At alkaline pH, the hydroxyls adsorbed onto the iron centers introduce an additional electron transfer barrier, increasing the overpotential that is required to overcome this resistance.
The exchange current densities shown inFigure 7.b confirm thefindings described above. In fact, the highest currents are observed when the majority of the iron centers are available to carry out the electron transfer at acidic pH values. With the pH of the electrolyte changed to alkaline, adsorption of hydroxyls onto the iron centers serves as inhibition of the ORR. The increase in the exchange current density is observed for the pH values 10.6 and 12.4pH values at which the phenolic and carboxylic acid species become deprotonated (Figure 7.b). Protons released may neutralize the adsorbed hydroxyls and, therefore, reactivate the iron center.
As the surface concentration of deprotonated nitrogen increases, the electron transfer coefficient increases (Figure
S9.a). This indicates that at acidic pH, pyridinic nitrogen provides the protons required to carry out the proton-coupled electron transfer, which becomes more difficult as they become deprotonated at higher pH. Between 2.8 and 7.2, small changes in the amount of pyridinic nitrogen have a significant impact on the exchange current density (Figure S9.b), which supports the vital role of the protons from the pyridinic nitrogen for carrying out the ORR. Trends observed for hydrogenated pyridines (Figure S10.a) are inversely related to the trends for deprotonated pyridinic nitrogen (Figure S9). At acidic pH, a high amount of hydrogenated pyridine groups facilitate the proton-coupled electron transfer, indicated by the lower electron transfer coefficients. As the amount of hydrogenated nitrogen decreases at higher pH due to deprotonation, a significant increase occurs in the electron transfer coefficient, shown by the data point at pH 7.2. The same trend is observed for the exchange current density in Figure S10.b, which indicates that the electron transfer happens readily when protonated nitrogen is present.
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CONCLUSIONSIn conclusion, the present study shows the effect of the electrolytes’ pH on the ORR electrochemical activity of PGM-free catalysts. The change in the concentration of protons and hydroxyls in the electrolyte leads to changes in the surface chemistry of the catalyst itself, which was spectroscopically confirmed. Based on the pKa values of multiple functional groups that exist in the PGM-free catalysts, it was possible to identify and explain the trends in the electrochemical activity of the PGM-free catalyst toward ORR and correlate them with the changes in the surface chemistry as a function of pH.
We report that the pH has a strong effect on the chemical state and on the accessibility of Fe−NX, centers by the molecular oxygen and thus is critical in general ORR activity not only based on the massflow (availability of H+or OH−as reactants). The increase in the amount of hydroxyls in the electrolyte decreases the amount of accessible Fe−NXcenters due to specific adsorption. At pH values below 10.5, protons available at the surface of the PGM-free catalyst such as protonated nitrogen and carbon functional group neutralize this specifically adsorbed OH− enabling Fe−NX as active centers toward the most beneficial 4 e−ORR to water as product. At pH higher than 10.5, the excess OH−adsorbed at the Fe−NX centers contributes to the outer Hemlotz plane oxygen reduction mechanism. Protonation of pyridinic nitrogen plays Figure 7.Surface concentration expressed as relative percentage of Fe−NXspecies of the PGM-free catalyst after being exposed to the different pH
levels vs ORR electron transfer parameters. The numbers next to the squares indicate the respective buffer pH: (a) electron transfer coefficient and (b) exchange current density.
an extremely important role in the pH effect on the oxygen reduction mechanism, as having a pKavalue close to the neutral pH, it can provide protons for the oxygen reduction reaction itself and the neutralization of the adsorbed OH−on the Fe− NXcenters. By this, we show that the transition between“acid” and “alkaline” mechanism of ORR on PGM-free catalysts of transition metal−nitrogen−carbon type happens at pH 10.5, with multiple pH zones in variation of the “acid” mechanism determined by the pKavalues of surface hydrogenated species.
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ASSOCIATED CONTENT*
S Supporting InformationThe Supporting Information is available free of charge on the
ACS Publications websiteat DOI:10.1021/acscatal.7b03991. The Supporting Information contains (I) The catalyst loading test for one of the pH values, (II) the Purbaix diagram that was created to determine the appropriate oxidative current in the ring of the rotating ring disk electrode to determine the peroxide yield, (III) the curve fitted XPS spectra of the nitrogen species in the PGM-free catalyst, (IV) the curve fitted XPS spectra of the carbon species in the PGM-free catalyst, (V) the surface concentration expressed as relative percentage of chemical species at different pH values, (VI) the hydrogen peroxide yield and overall electron transfer in the ORR at all pH values, (VII) representative Koutecky−Levich plot for the ORR carried out at pH neutral pH, (VIII) Koutecky−Levich slope, (IX) the logarithm of the ratio of kinetic current density and electron transfer current density vs pH, (X) surface concentration expressed as relative percentage of pyridinic nitrogen of the PGM-free catalyst after being exposed to the different pH levels vs ORR kinetic and electron transfer parameters, and (X) surface concen-tration expressed as relative percentage of pyrrolic nitrogen of the PGM-free catalyst after being exposed to the different pH levels vs ORR kinetic and electron transfer parameters (PDF)
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AUTHOR INFORMATION Corresponding Author *E-mail:[email protected]. ORCID Kateryna Artyushkova:0000-0002-2611-0422 Ivana Matanovic:0000-0002-9191-8620 Plamen Atanassov:0000-0003-2996-472X NotesThe authors declare no competingfinancial interest.
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ACKNOWLEDGMENTSThis work was supported by Center for Micro-Engineered Materials.
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