Article pubs.acs.org/JPCC
Surface Platinum Electrooxidation in the Presence of Oxygen Anusorn Kongkanand* and Joseph M. Ziegelbauer Electrochemical Energy Research Lab, General Motors Global Research and Development, Honeoye Falls, New York 14472, United States ABSTRACT: Understanding the influences of Pt surface oxides is crucial for elucidating the oxygen reduction reaction (ORR) in fuel cells, particularly with an eye toward the design of next-generation Pt-based electrocatalysts with improved activity. Although gaseous O2 is always present during the ORR, the majority of previous Pt surface oxide studies were conducted in the absence of O2 due to experimental limitations. In this study, multiple in situ techniques were applied to study the ORR on platinum in the presence of O2. The thin channel flow Pt electrode and the electrochemical quartz microbalance (EQCM) techniques on Pt polycrystalline electrodes suggest that the influence of O2 in the electrolyte on mass change and charge transfer is negligible. The oxide formation followed a logarithm-growth behavior initiating as early as 0.2 s after potential steps. This suggests that no slow conversion of oxide species (e.g., between OHads and Oads) takes place. Despite the negligible effect of O2 on the measured oxide coverage, steady-state X-ray absorption spectroscopy (XAS) measurements conducted on dispersed Pt nanoparticles suggested that, only when under O2-sparging, place exchange between adsorbed oxide(s) and the Pt surface layer(s) initiated at potentials as low as 0.75 V. This is significantly lower than that observed in an O2-free electrolyte (>1.1 V). The effects of O2 on adsorbed oxide species illustrate the complexities of applying model systems to the real world in order to arrive at a comprehensive description of both the oxygen reduction and the Pt dissolution processes.
■
electroadsorption
δ− − Pt + H2O ⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯→ Ptδ+−Oads + 2H+ aq + 2e
INTRODUCTION One of the challenges that confront proton exchange membrane fuel cells (PEMFCs) designed for transportation applications is to achieve high energy conversion efficiency with the minimum amount of precious platinum catalysts. During highway cruising conditions, where the cathode operates over a range of 0.7−1.0 V, the kinetics of the oxygen reduction reaction (ORR) are sluggish. This issue arises in part from the partial oxidation of the Pt surfaces during the ORR. The surface oxides can be formed by chemical oxidation by gas-phase O2 without fixed potential,1,2 by electrochemical oxidation using water as the oxygen source,3−17 or by a combination of both processes. These oxides can adversely affect the ORR rate as poisons18 or kinetically trapped intermediates.19,20 Recent density function theory calculations suggested that the chemisorption energies of oxygen species influenced the ORR rate.21,22 This conclusion from theory increased the importance of improved experiment-based understanding of the state of surface oxides. Surface oxide growth by electrochemical polarization has been extensively discussed (cf. the foundational works by Conway et al.).3−17 It was established that the oxidation of water molecules to form adsorbed oxygencontaining species (eq 1) is followed by an interfacial place exchange between O(H)ads and Pt atoms which leads to formation of a quasi-3D layer (eq 2).15,16 The growth rate follows logarithm kinetics when the process is limited by the interfacial place exchange at low oxide coverage but follows inverse-logarithm kinetics when the process is limited by the escape of Pt from the inner layer at high oxide coverage. © 2012 American Chemical Society
(1)
quasi‐3D layer
δ− Ptδ+−Oads ⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯→ Pt2 +−O2 −
(2) 3−17
All of the electrochemical experiments mentioned above were performed in the absence of gaseous O2 in order to eliminate interference by ORR currents. Unfortunately, oxides can also be formed in a reaction with gaseous O2, and the possible influence of this effect was largely neglected. In addition, many ORR kinetics studies19,20,23,24 have been performed by correlating the oxide measurement obtained in the absence of O2 with the reactivity measured in oxygen’s presence. Despite the wealth of information and insight generated by these works, they could not address whether gaseous O2 affects the rate of formation or the steady-state coverages of surface oxides. Some coulometric experiments25,26 have shown that exposure to O2 could have a long-lasting impact on the surface oxide coverage of a PEMFC cathode that were still present even after purging with N2 for 2.5 h. However, careful measurements later performed by Liu et al.27 did not reproduce this result. X-ray photoelectron spectroscopic measurements carried out in an electrochemical/ ultrahigh-vacuum transfer system using “explosive emersion” have also shown increased oxide coverages on Pt after exposure of the electrode to O2-sparged electrolyte.28 Taken together, these results are important because they intimate that in order Received: November 29, 2011 Revised: January 18, 2012 Published: January 18, 2012 3684
dx.doi.org/10.1021/jp211490a | J. Phys. Chem. C 2012, 116, 3684−3693
The Journal of Physical Chemistry C
Article
Figure 1. Schematic illustration of the thin channel flow electrode apparatus (a) and flow cell components (b).
the deoxygenated electrolyte. EQCM was employed owing to its ability to monitor exceedingly small changes in mass correlated with charge transfer during oxide formation on Pt. It is noteworthy that although EQCM has been used for quantification of surface oxides in several studies,9,10,14 the effect of O2 has not been appropriately addressed.41 Finally, element-specific in situ XAS conducted on a state-of-the-art nanoscale Pt/C electrocatalyst, via application of the ΔμXANES technique,42,43 offered direct spectroscopic insight into not only the coverages of the adsorbates but also their geometric binding sites on the Pt surfaces. The combination of these techniques provides insight into not only the influence of dissolved O2 on the oxide formation processes but also the conditions resulting in the detrimental dissolution of Pt.
to arrive at an explicit understanding of ORR kinetics, it is necessary to account for both Pt oxide kinetics and electrode history.29 All of these experiments were done after deoxygenation for a few minutes27 to hours25,26 or after entirely removing the electrode from its electrochemical environment.28 Thus, in order to attain a full understanding of the processes described above, true in situ techniques will be required. Further, it is important to note that the ORR activity per Pt surface area (i.e., the specific activity) on a continuous Pt surface is a factor of 5−10 times higher than that observed with Pt nanoparticles.30 Because an overwhelming majority of the aforementioned research was conducted on continuous Pt surfaces (either polycrystalline or single facets), it is not entirely clear if the conclusions drawn from studies involving continuous Pt surfaces fully approximate Pt nanoparticles. Ideally, these types of studies would be conducted strictly on real-world Pt nanoparticles. At the very least, the oxide coverages on both continuous surfaces and nanoparticles need to be compared with the utmost care. The implications of the states and coverages of oxides not only are important for ORR at steady state but also have significance for transient operation over the 0.7−1.0 V range where automotive PEMFCs are operated for the majority of the time. Indeed, ORR activity during transients can be significantly different from operation at a steady state.31−34 On the other hand, Pt dissolution is a process closely related to surface oxide formation and is considered to be one of the major contributors to fuel cell degradation. The dissolution rate varies with many factors such as potential, potential manipulation, particle size, support type, temperature, and contaminations. For details about Pt dissolution we refer to a review article by Sasaki et al.35 and references therein. It was recently found that the presence of oxygen36,37 and hydrogen peroxide38,39 can significantly enhance the rate of Pt dissolution, especially at low potentials (0.2−1.0 V). Therefore, it is of great importance to understand how the formation rates, coverages, and chemistries of Pt oxides influence the ORR activity. In this work, the formation rate, steady-state coverage, and species type of surface oxide were studied via coulometric, gravimetric, and X-ray absorption spectroscopic methods. The coulometric measurements were performed on a thin channel flow electrode40 (TCFE). A Pt electrode was polarized at a given potential in 0.1 M HClO4 for a given amount of time to form surface oxide in presence and absence of O2, and then the electrolyte was replaced with a deoxygenated electrolyte within seconds prior to the quantification of formed oxide using voltammetry. This minimized any change in the oxide properties that could have occurred during equilibration with
■
EXPERIMENTAL SECTION Thin Channel Flow Electrode. The TCFE cell was made of Kel-F, and the setup was similar to that reported by Wakabayashi et al.40 The working electrode was a polycrystalline Pt sheet with a geometric active area of 0.07 cm2geom and a roughness of 1.4 cm2Pt cm−2geom (assuming a H electrosorption charge of 210 μC cm−2Pt). A 0.1 mm thick Teflon sheet with a 15 × 4 mm cut-out window was placed between the working electrode and a Pt-sheet counter electrode; it served as both a seal and a flow channel. The setup is illustrated in Figure 1. The electrolyte reservoirs were filled with 0.1 M HClO4 (GFS Chemicals) and were purged with N2 or O2 for more than 1 h prior to the measurements. The electrolyte flow rates were adjusted through a combination of the gas pressures and needle valves to maintain a typical flow rate of 10 μL s−1 (2.5 cm s−1 across the electrode surface). The cell was equipped with a Ag/ AgCl reference electrode, but all potentials are reported against the reversible hydrogen electrode (RHE) assuming RHE = 0.281 + V vs Ag/AgCl. The TCFE was initially polarized at an oxide-free potential (0.43 V) followed by polarization to a predetermined Ep in the oxide formation region (0.7−1.2 V) for a given duration (tp) in either N2- or O2-saturated electrolyte. Five seconds prior to the end of tp, the four-way valve was switched to fully flood the cell with a N2-saturated electrolyte. Five seconds is sufficient to flow the electrolyte between the valve and the outlet of the cell. At the end of tp, a linear sweep voltammogram (Ep to 0.02 V) and then a cyclic voltammogram (between 0.02 and 0.6 V) were performed to measure both the oxide reduction and hydrogen adsorption/desorption charges. The measurement procedure was similar to one that had been reported earlier,25−27 except only 3 s was required to fully flush the cell with deoxygenated electrolyte (Figure 2). 3685
dx.doi.org/10.1021/jp211490a | J. Phys. Chem. C 2012, 116, 3684−3693
The Journal of Physical Chemistry C
Article
polarization potentials over a range of 0.6−1.2 V. Further, to maximize the stability of the EQCM, the polarization times were limited to 2 min in these tests. A typical 95% confidence interval of the mean was in the range of 1−2.5 ng cm−2geom. Akin to the TCFE experiments, separate potential hold sweeps were performed on the QC in order to quantify the charge required to form oxide(s) at each polarization time (tp) and potential (Ep). X-ray Absorption Spectroscopy. A 46.6 wt % Pt/C (TKK, supported on Vulcan carbon, 2.9 nm particle diameter by X-ray diffraction) catalyst was used for all of the XAS experiments. Electrodes were constructed by hand-painting a mixture of Pt/V and Nafion solution onto 0.5 × 4 cm strips of Zoltek carbon cloth. The final electrode loadings were 96 wt % Pt/V and 4 wt % Nafion, with a geometric Pt loading of 2.4 mgPt cm−2. This Pt loading was chosen to allow the room temperature XAS experiments to be conducted in transmission mode at the Pt L3 edge (11 564 eV). The spectroelectrochemical cell utilized during the XAS experiments was of a previously described design.45 A sealed Ag/AgCl electrode (BAS, measured 0.271 V vs RHE) served as the reference electrode and was shielded from the working electrode by a Vycor tip and a fine glass frit. The counter electrode was a high purity gold wire (Alfa Aesar) of multiple windings with sufficient area for the working electrode. A reservoir containing ∼500 mL of 0.1 M HClO4 electrolyte was sparged continuously with either high purity N2 or O2 while being pumped through the cell via a small peristaltic pump. The flow rate of the electrolyte was kept as low as possible to minimize the leaching of chloride from the sealed reference electrode. Pt L3 edge (11564 eV) XAS data were collected at beamline XOR-9BM at the Advanced Photon Source (Argonne National Laboratories). Three ionization chambers (I0, It, and Ir) were utilized for transmission data collection. The chambers were filled with N2, Ar, or a mixture of both gases to provide for Xray absorptions of 10, 70, and 100% at the Pt L3 edge. The cell was positioned between I0 and It, and a Pt foil was placed downstream of the cell between It and Ir. This allowed the experimental data to be collected simultaneously with the Pt foil reference to aid in energy alignment and normalization. Prior to data collection, the cell was cycled between 0.05 and 1.1 V vs RHE at 10 mV s−1 until the CV reached a steady state (4−5 cycles). Data were collected during anodic potential holds. Between each hold potential the electrode was cycled between 0.05 and 1.05 V, with a final anodic sweep up to the potential of interest. Except for at the double layer (0.54 V) and at 1.05 V, three successive X-ray energy scans were collected at 0.70/0.75/0.80/0.85/0.90/0.95 V. The first scan collected the full EXAFS (−200 eV to 16K, relative to the Pt L3 edge), while the second and third scans were XANES sweeps (−200 to +200 eV). Energy differences in the XANES region could be determined to 0.2 eV. From the start of the anodic potential hold, the XANES regions were sampled beginning at 2, 14, and 20 min. The scan durations were 12 and 7 min for EXAFS and XANES, respectively.
Figure 2. Current density and potential as a function of time during oxide formation measurements with the TCFE. The O2-diffusion limiting current of approximately −4 mA cm−2 during the first 10 s in the O2-saturated electrolyte was excluded from the figure for better clarity. See text for details.
Quartz-Crystal Microbalance. Pt-coated, AT-cut, 9 MHz, 1 in. diameter quartz crystals (QCs), with a temperature coefficient turnaround point of 25 °C, were purchased from Inficon Co. The QC analyzer was a RQCM from Maxtek with a data acquisition time of 0.05 s point−1 and a mass resolution of ∼0.4 ng cm−2. The QCs were placed in a crystal holder made of Teflon with a Kalrez O-ring which exposed only the front face of the crystals. The electrode assembly was then placed into a conventional three-electrode electrochemical cell with a water jacket. The Ag/AgCl (sat KCl) reference electrode (measured 0.278 V vs RHE) was equipped with a Luggin capillary. All reported potentials are normalized to RHE. Coolant water was circulated through the water jacket to maintain the temperature at 25 ± 0.02 °C. The electrochemically active geometric area was 1.4 cm2geom, and the QCM mass-measuring area was 0.37 cm2. The electrochemical surface area of the Pt working electrode was determined to be 4.9 cm2Pt cm−2geom by measuring the adsorption/desorption of hydrogen. Although not all of the changes in measured resonance frequencies of the QC can be attributed to mass changes resulting from adsorbing species, for ease of presentation the frequency responses were converted into mass changes using the Sauerbrey equation44
Δm =
(fq − f )ρq1/2μq1/2 2nfq 2
(3)
where n = 1, fq = 9 MHz, ρq = 2.648 g cm−3, and μq = 2.947 × 1011 g cm−1 s−2, representing the harmonic order, the fundamental frequency of the crystal oscillation, the density, and shear modulus of the quartz crystal, respectively. Even with significant care, a long-term instability of the EQCM caused by small thermal drifts, mechanical stresses, or vibrations has been a major burden in achieving high mass resolution. This has prevented its use at low oxide coverage region (
