Cation–Hydroxide–Water Coadsorption Inhibits the Alkaline

Letter pubs.acs.org/JPCL

Cation−Hydroxide−Water Coadsorption Inhibits the Alkaline Hydrogen Oxidation Reaction Hoon Taek Chung,∥,† Ulises Martinez,∥,† Ivana Matanovic,‡,§ and Yu Seung Kim*,† †

MPA-11: Materials Synthesis and Integrated Devices, Los Alamos National Laboratory, Los Alamos, New Mexico 87545, United States ‡ Department of Chemical and Biological Engineering, Center for Micro-Engineered Materials (CMEM), The University of New Mexico, Albuquerque, New Mexico 87231, United States § T-1: Physics and Chemistry of Materials, Los Alamos National Laboratory, Los Alamos, New Mexico 87545, United States S Supporting Information *

ABSTRACT: Rotating disk electrode voltammograms and infrared reflection absorption spectra indicate that the hydrogen oxidation reaction of platinum in 0.1 M tetramethylammonium hydroxide solution is adversely impacted by time-dependent and potential-driven cation−hydroxide−water coadsorption. Impedance analysis suggests that the hydrogen oxidation reaction inhibition is mainly caused by the hydrogen diffusion barrier of the coadsorbed trilayer rather than intuitive catalyst site blocking by the adsorbed cation species. These results give useful insights on how to design ionomeric binders for advanced alkaline membrane fuel cells.

A

electrodes may occur at high pH and low potential. Kohl’s group reported that the tetramethylammonium adsorption reduces Pt surface charge density and HOR current density.8 Recently, Janik’s group reported that the specifically adsorbed potassium drives the adsorption of OH with coadsorbed H2O to more positive potentials, making the surface hydroxide formation less stable.9 The coadsorption study explains the hydrogen adsorption−desorption potential shift in cyclic voltammetry at higher pH; however, the study did not explicitly explain how the cation−hydroxide−water coadsorption affects the HOR activity of Pt. Here we investigate the HOR behavior of Pt in an organic cation solution using a rotating disk electrode. For this study, high-purity tetramethylammonium hydroxide (TMAOH) is used because TMAOH tethered polymers are the most popular polymer electrolytes for AMFCs. Direct evidence of cation− hydroxide−water coadsorption is provided by infrared reflection absorption spectroscopy (IRRAS). Further impedance studies propose a new HOR inhibition mechanism associated with the coadsorbed layer on the Pt surface. Figure 1a compares the HOR voltammograms of a Pt polycrystalline electrode in 0.1 M HClO4 and TMAOH solutions. The initial HOR voltammograms were obtained from the respective anodic scans [−0.1 to 1.2 V vs reversible

lkaline membrane fuel cells (AMFCs) offer significant advantages over traditional proton exchange membrane fuel cells (PEMFCs). The most significant advantage is the increased electrocatalytic activity and stability of nonprecious metal group catalysts for the oxygen reduction reaction.1 However, the performance of AMFCs using nonprecious, or even precious metal catalysts, is yet much inferior to that of PEMFCs.2 One of the most critical shortcomings causing the inferior AMFC performance is the sluggish activity of the hydrogen oxidation reaction (HOR). Previous studies performed by Shao-Horn’s group indicate that HOR kinetics in alkaline electrolyte are several orders of magnitude slower than in acid electrolyte.3 Up until recently, relatively little research has been devoted to the understanding of the poor HOR activity in alkaline electrolytes. Markovic’s group alluded to the importance of an optimal balance between the adsorption− dissociation of H2 and the adsorption of hydroxide (OH), suggesting that fine-tuning the OH adsorption energy through improved oxophilicity of electrocatalysts could increase the HOR activity.4 However, studies by Durst et al. using oxophilic iridium catalysts disproved that OH adsorption is the ratedetermining step; high H-binding energy instead might be responsible for the slow HOR kinetics. 5 Besides the mechanistic explanation for the slow HOR kinetics in alkaline electrolyte, possible cationic group adsorption is suggested for the attribute of slow HOR activities. Density functional theory (DFT) calculations by Mills et al.6 and Matanovic et al.7 suggest that specific adsorption of alkali metal cation onto Pt © XXXX American Chemical Society

Received: September 5, 2016 Accepted: October 23, 2016 Published: October 24, 2016 4464

DOI: 10.1021/acs.jpclett.6b02025 J. Phys. Chem. Lett. 2016, 7, 4464−4469

Letter

The Journal of Physical Chemistry Letters

solutions. The voltammograms in 0.1 M HClO4 feature a peak in the oxidation current at ca. 0.2 V that is likely caused by the oxidation of excess H2 evolved at −0.1 V. The kinetic current density at 0.01 V measured as a gauge for HOR activity of Pt electrode in the HClO4 solution is 0.32 mA cm−2, significantly higher than that of 0.10 mA cm−2 obtained in the TMAOH solution (inset in Figure 1a). This much lower initial HOR activity of Pt in the alkaline electrolyte than in acidic electrolyte is consistent with the previous result by Sheng et al.3 The diffusion-limited HOR current densities in 0.1 M HClO4 and TMAOH are similar (∼0.73 mA cm−2). Note that the limiting current density at 900 rpm is lower than the reported HOR limiting current density (∼2 mA cm−2) because of the low scan rate (5 mV s−1)10 and high altitude of Los Alamos that decreases the limiting current by ca. 12%. In addition to measuring the initial HOR activity of Pt electrode, the HOR activity from the anodic scans was measured after applying 0.1 V vs RHE to the Pt electrode for 15 min. The kinetic current density at 0.01 V in the HClO4 solution remained virtually the same as 0.32 mA cm−2, while the current density measured in the 0.1 M TMAOH solution further decreased from 0.10 to 0.03 mA cm−2 at 0.01 V (inset). Furthermore, the HOR limiting current density measured at 0.6 V in the TMAOH electrolyte decreased from 0.72 to 0.60 mA cm−2. The timedependent HOR current density decay is also observed with carbon-supported Pt in 0.1 M TMAOH solution (unpublished results), although the decay rate with carbon-supported Pt was much lower, for example, ∼33% HOR current density decrease at 0.01 V after 2 h of exposure to 0.1 V vs RHE. Figure 1b compares the cyclic voltammograms (CVs) of Pt polycrystalline electrode in 0.1 M HClO4 and TMAOH solutions. CVs of the Pt polycrystalline electrode in 0.1 M HClO4 and TMAOH solutions show a well-defined hydrogen adsorption−desorption region with corresponding peaks for the Pt(110) and Pt(100) planes as indicated in the figure. These peaks are shifted to more positive potentials when the electrolyte is changed from HClO4 to TMAOH. The peak shifts for the Pt(110) and Pt(100) are 0.156 and 0.092 V, respectively. The positive shift we observe with TMAOH is similar to the shift calculated and observed for potassium.9,11 The hydroxide−water−cation coadsorption phenomenon can help explain the difference in CV between HClO4 and TMAOH. As the pH is increased, cation adsorption becomes more favorable, increasing its

Figure 1. (a) HOR voltammograms of Pt in 0.1 M HClO4 and TMAOH electrolytes. The HOR voltammograms of Pt were obtained before and after applying 0.1 V vs RHE to Pt electrode for 15 min. The HOR voltammograms were performed at 25 °C; rotating speed, 900 rpm; and scan rate, 5 mV s−1. (b) CVs of Pt in 0.1 M HClO4 and TMAOH electrolytes. The CVs were obtained at 25 °C; rotating speed, 900 rpm; and scan rate, 50 mV s−1.

hydrogen electrode (RHE)] in hydrogen-saturated solutions after a few cycles between 0.0 and 1.2 V in nitrogen-saturated

Figure 2. IRRAS spectra change in the (a) fingerprint and (b) O−H stretching regions during chronoamperometry of Pt in 0.1 M TMAOH electrolyte at 0.1 V vs RHE. FTIR spectra of 0.1 M TMAOH electrolytes are also shown for comparison purpose. 4465

DOI: 10.1021/acs.jpclett.6b02025 J. Phys. Chem. Lett. 2016, 7, 4464−4469

Letter

The Journal of Physical Chemistry Letters

Figure 3. (a) HOR voltammograms of Pt in 0.1 M TMAOH electrolytes. The HOR voltammograms were measured before and after applying 0.1, 0.3, or 0.9 V vs RHE for 15 min. The HOR voltammograms were performed at 25 °C; rotating speed, 900 rpm; and scan rate, 5 mV s−1. (b) IRRAS spectra change in the fingerprint region after chronoamperometry of Pt in 0.1 M TMAOH electrolyte for 20 min as a function of electrode potential vs RHE. (c) Integrated methyl peak absorbance after the chronoamperometry experiment.

originates from the OH stretch in water, and the peak at ∼2800 cm−1 originates from the hydroxide anion. This is in stark contrast with the OH stretching peak of the 0.1 M TMAOH electrolytes in which only one broad solvated OH stretching peak is observed at 3200 cm−1. Two separate peaks for the OH and H2O suggest that the OH anions are not complexed with H2O but segregated. The relatively large hydroxide peak compared to H2O suggests more hydroxide molecules are adsorbed than water molecules. For the normalized IRRAS spectra after 15 min exposure in 0.1 M TMAOH solution, the values for water and hydroxide molecules per TMA+ cation were calculated to 0.5 and 3.5, respectively, which is equivalent to approximately 30 N KOH. The coadsorbed trilayer containing relatively high concentration of hydroxide might not be located directly at the interface but at the near surface instead with no effect on the local pH change of the Pt surface, as we observed the peak shift of the CV is negligible. This is consistent with previous observations that the structure and dynamics of water near the electrode surface and the change in the electric field near the electrode surface have a small effect on the adsorption potential of hydrogen.13,14 While we do not have further analytical data on the coadsorbed trilayer structure, the IRRAS experiments again suggest that the coadsorption is not a simple and mobile adsorption, yet it is more in line with an accumulated solid layer formation in which lateral interactions play a significant role. These IRRAS spectra are direct evidence for the existence of TMA + /H 2 O/OH coadsorbed trilayer on the Pt surface, which are in good agreement with the theoretical studies that predict the existence of H2O/OH bilayer15 and the recent DFT calculation in which potassium-specific adsorption occurs with the coadsorption of OH and water.9 The time-dependent organic cation− hydroxide−water adsorption behavior suggests that HOR current density decay may be related with the coadsorbed trilayer. The chronoamperometric studies at different potentials revealed the dependence of the HOR activity on the Pt electrode potential. Figure 3a compares the HOR voltammogram of the Pt electrode in 0.1 M TMAOH electrolyte before and after 0.1, 0.3, or 0.9 V is applied for 15 min. The HOR activity and limiting current obtained after exposure to 0.1 V for 15 min decrease substantially from the initial value, and with

coverage on the surface. On the basis of density functional studies of hydroxide−water−potassium (co)adsorption on different Pt surfaces, Janik and McCrum proposed that the coadsorbed potassium disturbs the solvation of adsorbed hydroxide.9 This in turn destabilizes the surface hydroxide and drives the desorption peaks measured in cation-containing electrolyte to higher potentials. After 0.1 V vs RHE is applied to the Pt electrode for 15 min, the intensity of the (110) peak decreases in TMAOH. This may indicate that the pH change on the Pt surface in TMAOH is not significant and the peak suppression may be explained by the decrease in hydrogen adsorption−desorption. The time-dependent HOR current density decay and corresponding CV change due to the possible trilayer coadsorption are also observed with alkali metal cation solutions. Figure S1 shows the CV change in 0.1 M NaOH solution before and after 0.1 V exposure for 120 min. IRRAS experiments were performed to identify the adsorbate species on Pt surface. In these experiments, a polycrystalline Pt electrode was immersed in water and the concentrated TMAOH solution was added to a concentration of 0.1 M. The infrared spectra of the Pt surface were recorded at 0.1 V as a function of the immersing time. The characteristic peaks of adsorbed species are obtained after subtracting the reference background for water. The IRRAS spectrum in the fingerprint region for the Pt electrode surface shows that TMA + characteristic peaks at 1488 cm−1 υas(CH3), 1419 cm−1 υs(CH3), and 950 cm−1 υas(C−N) are evolved with time (Figure 2a). A slight redshift for the CH3 vibration peaks was observed for adsorbed TMA+: 1477 to 1488 cm−1 for υas(CH3) and 1419 to 1423 cm−1 for υas(CH3), indicating an increase of the interaction between the Pt surface and the adsorbed TMA+ species. The absence of characteristic peaks of possible degraded species from TMAOH12 such as methanol (1021 cm−1, υ(C−O)), dimethyl ether (∼1200 cm−1, υ(C−O)), and trimethyl amine (1375 cm−1, υ(CH3), 1205 cm−1, υ(C−N)) indicates that the adsorbed species on the Pt electrode is not decomposed TMA. 1H NMR study of TMAOH solution also attests no degradation of the TMA+ and no-impurity effect (Figure S2 and the Supporting Information). In addition to TMA+ adsorption peak, the characteristic peaks of H2O and OH adsorptions were also evolved with time at ∼3500 and ∼2800 cm−1, respectively (Figure 2b). The peak at ∼3500 cm−1 4466

DOI: 10.1021/acs.jpclett.6b02025 J. Phys. Chem. Lett. 2016, 7, 4464−4469

Letter

The Journal of Physical Chemistry Letters

Figure 4. Impedance measured at 900 and 2500 rpms at (a) 0.01 V vs RHE, (b) 0.05 V vs RHE, and (c) 0.1 V vs RHE. (d) HOR measured at 900 and 2500 rpm. (e) Impedance spectra comparison before and after applying 0.1 V for 15 min. The frequency range for impedance measurement was 100 kHz to 0.1 Hz.

(400 Hz to 100 kHz). The appearance of two semicircles for the Pt electrode under alkaline conditions is consistent with previous observation by Koper and co-workers.16 As the impedance measuring potential increased to 0.05 and 0.1 V, the semicircle at low frequencies notably increased while the semicircle at high frequencies is almost invariant. The diameter of the low-frequency semicircle decreases with increasing rotating speed (Figure 4b,c). These results indicate that the semicircle at low frequencies does not originate from a kineticcontrolled process but from a diffusion-controlled process. The higher HOR current densities of the electrode measured at the higher rotating speed shown in Figure 4d also indicate that the semicircle at low frequencies is H2 diffusion-related. This refutes the previous belief15 that the origin of both low- and high-frequency semicircles is the two kinetic processes occurring at different time scales. Figure 4e shows the comparison of impedance spectra of Pt electrode before and after exposure of the electrode at 0.1 V vs RHE for 15 min. The size of the semicircle at high frequencies slightly increases (inset), while the size of the semicircle at low frequencies significantly increases. This indicates that the HOR inhibition after the 15 min holding test at 0.1 V as shown in Figures 1 and 3 is predominantly due to H2 diffusion limitation through the coadsorbed trilayer. The significant impact of time-dependent adsorption is also observed with a segmented frequency impedance measurement. In this measurement, each EIS of the Pt electrode was measured from low to high frequency after a holing potential at 1.2 V for 30 s. In that way, we can minimize the cation−water−hydroxide coadsorption during the EIS measurement. When the EIS data measured from high to low frequency are compared, the size of the low-frequency semicircle decreases (Figure S3). This experiment suggests that

increasing applied potentials, higher HOR activity and limiting current are obtained. However, even after the electrode was held at 0.9 V for 15 min, the initial HOR activity was not fully recovered. The effect of electrode potential on the cation adsorption was investigated using IRRAS (Figure 3b). The IRRAS results of the CH3 fingerprint region for the Pt electrode in the TMAOH solution were taken after applying different potentials from 0.05 to 0.90 V for 20 min. As the potential of the Pt electrode increases, the methyl bend at 1488 cm−1 gradually decreases, as shown quantitatively in Figure 3c. A trace of the methyl bend observed at 0.9 V is consistent with the rotating disk electrode result that the TMA+ desorption was incomplete at the high potential. These results indicated that the TMA+ adsorption decreases as the electrode potential increases. The fact that the catalytic activity could be restored only by excursions to very positive potential, ca. >0.9 V, where the Pt surface is being oxidatively cleaned may also suggest that the reported phenomena may not be caused by simple physisorption but rather by complex chemisorption process. These results again attest that HOR activity and limiting current decrease after potential hold tests in alkaline electrolytes as shown in Figure 1 are related to the coadsorbed layer. Electrochemical impedance spectroscopy (EIS) analysis is performed to investigate the HOR inhibition mechanism by organic cation−hydroxide−water coadsorption. Panels a, b, and c of Figure 4 compare the faradic impedance of the Pt electrode at 0.01, 0.05, and 0.1 V vs RHE, respectively, at two different electrode rotating speeds, ca. 900 and 2500 rpm. At 0.01 V vs RHE (Figure 4a), identical impedance spectra were obtained at 900 and 2500 rpm. The impedance spectra are composed of two semicircles: one larger semicircle at low frequencies (0.1− 400 Hz) and the other smaller semicircle at high frequencies 4467

DOI: 10.1021/acs.jpclett.6b02025 J. Phys. Chem. Lett. 2016, 7, 4464−4469

Letter

The Journal of Physical Chemistry Letters

Chulsung Bae and Dr. Angela Mohanty for 1H NMR chracterizations. This work was supported by the U.S. Department of Energy, Energy Efficiency and Renewable Energy, Fuel Cell Technology Office Incubator Program (DE-EE0006962). Los Alamos National Laboratory is operated by Los Alamos National Security LLC under Contract No. DEAC52-06NA25396.

the semicircle at low frequencies even at 0.01 V, i.e., kinetic region, is also related to the diffusion-controlled process. The H2 diffusion-related HOR inhibition mechanism looks counterintuitive because H2 is a small molecule and the diffusion of H2 is fast. A possible explanation is the extremely low solubility and diffusivity of H2 through the coadsorbed layer. Rüetschi reported that the H2 solubility and diffusivity could decrease significantly in a caustic solution. For example, H2 solubility and diffusivity decrease 7.5 and 4.5 times at 6 N KOH, respectively.17 Considering that the coadsorbed layer has a very high hydroxide population and is segregated from water, i.e., nonsolvated, it is expected that the solubility and diffusivity of hydrogen become extremely low (the “salting-out effect”).18 Although the coadsorbed layer thickness and the compositional distribution of the coadsorbed layer are currently unknown, the H2 permeability barrier in the coadsorbed layer seems to be the major attribute for low HOR activities in TMAOH solution from our EIS analysis. The H2 diffusion-related HOR inhibition by the coadsorbed layer gives few useful insights on the electrode design aspect of AMFCs. First, the structure of the cationic group plays a critical role in the coadsorption phenomenon observed. Less HOR inhibition was observed from bulky hydrophobic cationic group such as tetrabutyl ammonium or tetrabutyl phosphonium (unpublished results). Second, flexible polymer electrolytes may help desorb the cationic group when the anode electrode is exposed to relatively high potentials due to the greater mobility of the tethered cationic group. We have demonstrated that the HOR activities of Pt with flexible perfluorinated ionomeric binder19 are much better than that with rigid polyaromatic ionomeric binder20 after exposure to 1.4 V for 30 s.21 In summary, we have demonstrated that the HOR on Pt is substantially inhibited by time-dependent and potential driven cation−hydroxide−water coadsorption under alkaline conditions. IRRAS study indicates that the cation adsorption drives hydroxide and water coadsorbed onto Pt surface with much higher concentration of the hydroxide molecules compared to water. The coadsorption decreases with electrode potential in spite of incomplete removal of coadsorbents at relatively high potential ca. 0.9 V vs RHE. EIS analysis suggests that the inhibited HOR is largely due to the hydrogen permeability barrier through the accumulated coadsorbed layer.

■

■

(1) Chung, H. T.; Won, J. H.; Zelenay, P. Active and Stable Carbon Nanotube/Nanoparticle Composite Electrocatalyst for Oxygen Reduction. Nat. Commun. 2013, 4, 1922. (2) Piana, M.; Boccia, M.; Filpi, A.; Flammia, E.; Miller, H. A.; Orsini, M.; Salusti, F.; Santiccioli, S.; Ciardelli, F.; Pucci, A. H2/air Alkaline Membrane Fuel Cell Performance and Durability, using Novel Ionomer and Non-platinum Group Metal Cathode Catalyst. J. Power Sources 2010, 195, 5875−5881. (3) Sheng, W.; Gasteiger, H. A.; Shao-Horn, Y. Hydrogen Oxidation and Evolution Reaction Kinetics on Platinum: Acid vs Alkaline Electrolytes. J. Electrochem. Soc. 2010, 157, B1529−B1536. (4) Strmcnik, D.; Uchimura, M.; Wang, C.; Subbaraman, R.; Danilovic, N.; van der Vliet, D.; Paulikas, A. P.; Stamenkovic, V. R.; Markovic, N. M. Improving the Hydrogen Oxidation Reaction Rate by Promotion of Hydroxyl Adsorption. Nat. Chem. 2013, 5, 300−306. (5) Durst, J.; Siebel, A.; Simon, C.; Hasche, F.; Herranz, J.; Gasteiger, H. A. New Insights into the Electrochemical Hydrogen Oxidation and Evolution Reaction Mechanism. Energy Environ. Sci. 2014, 7, 2255− 2260. (6) Mills, J. N.; McCrum, I. T.; Janik, M. J. Alkali Cation Specific Adsorption onto fcc(111) Transition Metal Electrodes. Phys. Chem. Chem. Phys. 2014, 16, 13699−13707. (7) Matanovic, I.; Atanassov, P.; Garzon, F. H.; Henson, N. J. Density Functional Theory Study of the Alkali Metal Cation Adsorption on Pt (111), Pt (100), and Pt (110) Surfaces. ECS Trans. 2014, 61, 47−53. (8) Unlu, M.; Abbott, D.; Ramaswamy, N.; Ren, X. M.; Mukerjee, S.; Kohl, P. A. Analysis of Double Layer and Adsorption Effects at the Alkaline Polymer Electrolyte-Electrode Interface. J. Electrochem. Soc. 2011, 158, B1423−B1431. (9) McCrum, I. T.; Janik, M. J. pH and Alkali Cation Effects on the Pt Cyclic Voltammogram Explained using Density Functional Theory. J. Phys. Chem. C 2016, 120, 457−471. (10) Markovića, N. M.; Sarraf, S. T.; Gasteiger, H. A.; Ross, P. N., Jr. Hydrogen Electrochemistry on Platinum Low-Index Single-Crystal Surfaces in Alkaline Solution. J. Chem. Soc., Faraday Trans. 1996, 92, 3719−3725. (11) Sheng, W.; Zhuang, Z.; Gao, M.; Zheng, J.; Chen, J. G.; Yan, Y. Correlating Hydrogen Oxidation and Evolution Activity on Platinum at Different pH with Measured Hydrogen Binding Energy. Nat. Commun. 2015, 6, 5848. (12) Macomber, C. S.; Boncella, J. M.; Pivovar, B. S.; Rau, J. A. Decomposition Pathways of an Alkaline Fuel Cell Membrane Material Component via Evolved Gas Analysis. J. Therm. Anal. Calorim. 2008, 93, 225−229. (13) Karlberg, G. S.; Jaramillo, T. F.; Skulason, E.; Rossmeisl, J.; Bligaard, T.; Nørskov, J. K. Cyclic Voltammograms for H on Pt(111) and Pt(100) from First Principles. Phys. Rev. Lett. 2007, 99, 126101. (14) Sakong, S.; Naderian, M.; Mathew, K.; Hennig, R. G.; Groβ, A. Density Functional Theory Study of the Electrochemical Interface between a Pt Electrode and an Aqueous Electrolyte using an Implicit Solvent Method. J. Chem. Phys. 2015, 142, 234107. (15) Jinnouchi, R.; Nagoya, A.; Kodama, K.; Morimoto, Y. Solvation Effects on OH Adsorbates on Stepped Pt Surfaces. J. Phys. Chem. C 2015, 119, 16743−16753. (16) Schouten, K. J. P.; van der Niet, M. J. T. C.; Koper, M. T. M. Impedance Spectroscopy of H and OH Adsorption on Stepped SingleCrystal Platinum Electrodes in Alkaline and Acidic Media. Phys. Chem. Chem. Phys. 2010, 12, 15217−15224.

ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jpclett.6b02025. Details of the experiments, additional 1H NMR spectrum of TMAOH, and impedance measurements at segmented frequencies (PDF)

■

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Author Contributions ∥

REFERENCES

H.T.C. and U.M. contributed equally to this work.

Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS We thank Mr. Joseph Dumont at Los Alamos National Laboratory for performing FTIR analysis. We also thank Dr. 4468

DOI: 10.1021/acs.jpclett.6b02025 J. Phys. Chem. Lett. 2016, 7, 4464−4469

Letter

The Journal of Physical Chemistry Letters (17) Rüetschi, P. Solubility and Diffusion of Hydrogen in Strong Electrolytes and the Generation and Consumption of Hydrogen in Sealed Primary Batteries. J. Electrochem. Soc. 1966, 113, 301−305. (18) Rüetschi, P.; Amile, R. F. Solubility of Hydrogen in Potassium Hydroxide and Sulfuric Acid. Salting-out and Hydration. J. Phys. Chem. 1966, 70, 718−723. (19) Kim, D. S.; Fujimoto, C. H.; Hibbs, M. R.; Labouriau, A.; Choe, Y. K.; Kim, Y. S. Resonance Stabilized Perfluorinated Ionomers for Alkaline Membrane Fuel Cells. Macromolecules 2013, 46, 7826−7833. (20) Fujimoto, C.; Kim, D. S.; Hibbs, M.; Wrobleski, D.; Kim, Y. S. Backbone Stability of Quaternized Polyaromatics for Alkaline Membrane Fuel Cells. J. Membr. Sci. 2012, 423−424, 438−449. (21) Yim, S. D.; Chung, H. T.; Chlistunoff, J.; Kim, D. S.; Fujimoto, C.; Yang, T. H.; Kim, Y. S. A Microelectrode Study of Interfacial Reactions at the Platinum-Alkaline Polymer Interface. J. Electrochem. Soc. 2015, 162, F499−F506.

4469

DOI: 10.1021/acs.jpclett.6b02025 J. Phys. Chem. Lett. 2016, 7, 4464−4469