ELECTROCATALYSTS AND POLYMER ELECTROLYTES FOR ANION EXCHANGE MEMBRANE FUEL CELLS A Dissertation Presented to the Faculty of the Graduate School of Cornell University In Partial Fulfillment of the Requirements for the Degree of Doctor of Philosophy by Xinyao Lu August 2021 © 2021 Xinyao Lu ELECTROCATALYSTS AND POLYMER ELECTROLYTES FOR ANION EXCHANGE MEMBRANE FUEL CELLS Xinyao Lu, Ph. D. Cornell University 2021 The development of fuel cells is indispensable to enable the hydrogen society. Anion exchange membrane fuel cells (AEMFCs) have triggered great interest over the past few years. In this dissertation, electrocatalysis in alkaline media and polymer electrolytes for AEMFCs have been explored from fundamental aspects to practical applications. For the oxygen reduction reaction (ORR) in alkaline media, electrochemically dealloyed Pd-M (M = Ni, Mn) nanoparticle catalysts were developed to enhance electrocatalytic activity. The electrochemical dealloying process was demonstrated to be effective in selectively leaching out the less noble metal and metal oxides on the surface. The higher atomic concentration of electrochemically active Pd exposed on the surface of the nanoparticles was found to be the reason for the enhanced ORR activity in alkaline media. These findings provide insights for the rational design of the composition and structure of electrocatalysts with enhanced electrocatalytic activity, based on post-synthesis modification methods. From a fundamental perspective, to explain and predict the electrocatalytic activity of heterogenous reactions, the energy of intermediates is usually used as an activity descriptor. However, this was recently called into questions. We developed a novel method, based on fast scan rate cyclic voltammetry, to directly measure the kinetics of the electro-adsorption processes. The Had adsorption reaction, the elementary step of the hydrogen evolution reaction (HER), on Pt(111) in acid was found to be >100x faster than in alkaline media, although the Had binding energy was the same. Together, with cation effects and isotope effects found in alkaline media, we demonstrated that the slow kinetics of the HER at high pH are not due to an unfavorable Had binding energy but to the high barrier of interfacial water reorganization. Polymer electrolytes in AEMFCs, namely anion exchange membranes and ionomers, play important roles in the transport of anions and water molecules. In situ characterization of these materials in their electrochemical environment is critical for understanding the anion transport mechanism and improving the design of them. The anion exchange and water dynamics in a perspective phosphonium-based AEM during the methanol oxidation process were studied with the electrochemical quartz crystal microbalance (EQCM). The results provide insights of the anion exchange process in the membranes during the reaction and emphasize the importance of characterizing the membranes in a hydrated electrochemical environment. The influence of ion exchange capacity (IEC) on the solubility, the ionomer viscoelasticity in water and the transport of charged and uncharged species, of a promising polyethylene piperidinium methyl (PEPM) ionomer were also investigated. The design of ionomers and membranes, with suitable IEC for their different functions in AEMFCs from the aspect of solubility, mechanical properties and mass transport, can be guided via this work. BIOGRAPHICAL SKETCH Xinyao Lu was born and raised in Shanghai, China. She pursued her Bachelor of Science degree in Chemistry at Nanjing University in China from 2012 to 2016. During her junior year, she went to the University of Wisconsin-Madison as an exchange student for one semester, where she had the opportunity to experience an entirely different culture and customs in a foreign country. She spent the summer of 2015 in a research internship at Shanghai Institute of Microsystem and Information Technology synthesizing and characterizing an atomically thin graphite films in an ultra-high vacuum scanning tunneling microscopy (UHV-STM) chamber. She first got exposed to scientific research in the field of electrochemistry when she was an undergraduate student in Prof. Zhong Jin’s lab. Her research experience on the preparation and characterization of non-noble metal catalysts for the oxygen evolution reaction (OER) triggered her interest in research related to energy conversion and storage. In 2016, she was fortunate enough to be accepted to Cornell graduate school and join Prof. Héctor D. Abruña’s research group. Her research was focused on the development and characterization of electrocatalysts and polymer electrolyte for alkaline exchange membrane fuel cells (AEMFCs). In her third year at the Abruña group, she spent one month in Prof. Juan M. Feliu’s lab at University of Alicante in Spain learning the Clavilier method for the fabrication of single crystal electrodes, which was indispensable to her following research at Cornell. iii ACKNOWLEDGMENTS I would like to first express my sincere gratefulness to my advisor Prof. Héctor D. Abruña who provided me with the opportunity to work on the projects of my interest. He has been very supportive to all of his students not only in research but also in enjoying their time outside the lab. I really appreciate the friendly environment he has created in the lab, especially to woman students. His continuous encouragement helped me build my confidence in doing research. His professional suggestions guided me through my research projects. His great passion for everything he likes, such as research, coffee, wines, cycling, influenced my attitudes towards life. I feel blessed to have spent five years learning and growing up in his group. Secondly, I am thankful to have Prof. Francis J. DiSalvo and Prof. Jin Suntivich as my committee members. Prof. DiSalvo has given me many constructive suggestions during the starting time of my first project. I can still remember the time when he guided me through the synthesis process step by step in his office. Besides, I really enjoyed the time chatting with him and listening to him sharing his experiences. I am also grateful for the guidance and instruction from Prof. Suntivich in my research, especially the collaborative work with the excellent people in his group. His enthusiasm for research stimulated me to dive further into the project. I would also like to thank Prof. Juan M. Feliu and people in his group for teaching me the fabrication and characterization technique of single crystal electrodes at University of Alicante in Spain, where I enjoyed the warmth from the sunshine and the Spanish people. iv My wonderful labmates, Dr. Mahdi Ahmadi, Dr. Hongsen Wang, Dr. Luxi Shen, Dr. Yin Xiong, Dr. Na Zhang, Dr. Shuangyan Lang, Xinran Feng, Yao Yang, Jeesoo Seok, Rui Zeng, Huiqi Li, Cara Gannett, Weixuan Xu, Andrés Molina Villarino, Monica Jo Theibault and Mihail Krumov have also been backing me up both mentally and on research during my graduate study. Their company gave me joy in the lab and brought me courage when I was frustrated. I could always learn something new from them every day. Their perseverance and energy for work and life influenced me a lot. I also appreciate the assistance from my excellent collaborators from other groups, Dr. Kristina M. Hugar, Dr. Wei You, Cheyenne Peltier, Dr. Dingyuan Kuo, Bintao Hu and Rishi Agarwal. My work could not be done without them. In addition, I am grateful to my family and friends for being the persons that I could talk to anytime although the geographical distance might be far. My parents have always been my source of power for making progress and pursuing perfection. The support from my friends and my current and previous roommates, Peize Li, Qinwen Wang, Yixiao Lin, Wenyao Zhang, Chengzhang Wan, especially during the COVID period, was indispensable to me. Lastly, I would like to acknowledge the founding sources. This work was supported as part of the Center for Alkaline Based Energy Solutions (CABES), an Energy Frontier Research Center funded by the U.S. Department of Energy, Office of Science, Basic Energy Sciences under Award DE-SC0019445, and the Multidisciplinary Research program of the University Research Initiative (MURI) funded by the U.S. Air Force, Office of Scientific Research, under Award No. v N00014-17-S-F006. This work made use of the Cornell Center for Materials Research Shared Facilities which are supported through the NSF MRSEC program (DMR- 1719875). vi TABLE OF CONTENTS BIOGRAPHICAL SKETCH ......................................................................................... iii ACKNOWLEDGMENTS ............................................................................................. iv CHAPTER 1 INTRODUCTION .......................................................................................................... 1 1.1 Motivation ............................................................................................................ 1 1.2 Alkaline fuel cells fundamentals .......................................................................... 5 1.2.1 Introduction ................................................................................................... 5 1.2.2 Electrocatalysis in alkaline media ................................................................. 7 1.2.3 Polymer electrolyte materials ...................................................................... 11 1.3 Research overview .............................................................................................. 13 1.4 References .......................................................................................................... 14 CHAPTER 2 ELECTROCHEMICALLY DEALLOYED Pd-BASED NANOPARTICLE CATALYSTS FOR THE OXYGEN REDUCTION REACTION (ORR) IN ALKALINE MEDIA .................................................................................................... 27 2.1 Introduction ........................................................................................................ 27 2.2 Experimental section .......................................................................................... 29 2.2.1 Synthesis of PdxNi/C nanoparticles ............................................................. 29 2.2.2 Synthesis of PdMn/C nanoparticles ............................................................. 30 2.2.3 X-ray characterization ................................................................................. 30 2.2.4 Electrochemical dealloying and testing of the nanoparticles ...................... 30 2.2.5 Electron microscopy .................................................................................... 32 2.2.6 X-ray photoelectron spectroscopy (XPS) characterization ......................... 32 2.3 Results and discussion ........................................................................................ 33 2.4 Conclusions ........................................................................................................ 48 2.5 References .......................................................................................................... 48 CHAPTER 3 RATE AND MECHANISM OF ELECTROCHEMICAL FORMATION OF SURFACE BOUND HYDROGEN ............................................................................. 56 3.1 Introduction ........................................................................................................ 56 3.2 Methodology ....................................................................................................... 58 3.2.1 Accessing kinetics of electro-adsorption without an isotherm assumption . 58 3.2.2 Derivation of the equations ......................................................................... 59 3.2.2.1 Equilibrium potential at a given coverage ........................................................ 59 3.2.2.2 Derivation of surface redox kinetics ................................................................. 60 3.2.2.3 Kinetics in acidic media ................................................................................... 63 3.2.2.4 Kinetics in alkaline media ................................................................................ 66 3.3 Experimental section .......................................................................................... 68 vii 3.3.1 Preparation of Pt(111) working electrode ................................................... 68 3.3.2 Preparation of Pt(111)/Ni(OH)2 electrode ................................................... 68 3.3.3 Chemicals .................................................................................................... 68 3.3.4 Electrochemical measurements ................................................................... 69 3.3.5 Curve fitting ................................................................................................. 70 3.4 Results and discussion ........................................................................................ 70 3.4.1 Measurement of rate constant of Had formation on Pt(111) ....................... 70 3.4.2 Mechanism of Had formation in acid vs in base .......................................... 76 3.4.3 Effect of surface modifiers on Had formation rate ....................................... 81 3.5 Conclusion .......................................................................................................... 83 3.6 References .......................................................................................................... 84 CHAPTER 4 CATION AND ISOTOPE EFFECTS ON HYDROGEN ADSORPTION KINETICS ON Pt(111) IN ALKALINE MEDIA ........................................................................... 90 4.1 Introduction ........................................................................................................ 90 4.2 Experimental section .......................................................................................... 93 4.2.1 Preparation of Pt(111) electrode .................................................................. 93 4.2.2 Chemicals. ................................................................................................... 93 4.2.3 Purification of D2O. ..................................................................................... 93 4.2.4 Electrochemical measurements. .................................................................. 93 4.2.5 Curve fitting. ................................................................................................ 95 4.3 Results and discussion ........................................................................................ 95 4.3.1 Cation effects ............................................................................................... 95 4.3.2 Kinetic isotope effects ................................................................................. 99 4.4 Conclusions ...................................................................................................... 103 4.5 References ........................................................................................................ 104 CHAPTER 5 ............................................................................................................... 106 AN ELECTROCHEMICAL QUARTZ CRYSTAL MICROBALANCE STUDY ON ANION EXCHANGE AND WATER DYNAMICS IN A PHOSPHONIUM-BASED ALKALINE ANION EXCHANGE MEMBRANE MATERIAL FOR FUEL CELLS .................................................................................................................................... 106 5.1 Introduction ...................................................................................................... 106 5.2 Experimental section ........................................................................................ 108 5.2.1 Spin-coating of ionomer films for EQCM ................................................. 108 5.2.2 Film Thickness Measurement .................................................................... 109 5.2.3 Acoustic Impedance Spectroscopy ............................................................ 109 5.2.4 EQCM ........................................................................................................ 109 5.2.5 Open-circuit QCM study ........................................................................... 111 5.3 Results and discussion ...................................................................................... 112 5.3.1 Acoustic impedance analysis ..................................................................... 113 5.3.2 EQCM study of methanol oxidation at phosphonium-based AAEM coated electrodes ............................................................................................................ 118 viii 5.3.2.1 Cyclic voltammetry ........................................................................................ 120 5.3.2.2 Methanol oxidation ......................................................................................... 122 5.4 Conclusions ...................................................................................................... 130 5.5 References ........................................................................................................ 131 CHAPTER 6 INVESTIGATION OF THE INFLUENCE OF ION-EXCHANGE CAPACITY ON THE PROPERTIES OF ANION EXCHANGE IONOMERS FOR FUEL CELLS .. 137 6.1 Introduction ...................................................................................................... 137 6.2 Experimental Section ........................................................................................ 140 6.2.1 Preparation of ionomer solutions ............................................................... 140 6.2.3 Acoustic Impedance Spectroscopy ............................................................ 141 6.2.4 Electrochemical measurements ................................................................. 141 6.2.5 MEA testing ............................................................................................... 142 6.3 Results and Discussion ..................................................................................... 143 6.3.1 PEPM Ionomers and their solubility—tunable solubility with different IEC of the ionomers ................................................................................................... 143 6.3.2 Mechanical properties of PEPM ionomers in water—less rigidity due to higher amount of water uptake with increased IEC of the ionomers ................. 144 6.3.3 Mass transport in PEPM ionomers—electrostatic effects and permeability influenced by IEC of the ionomers ..................................................................... 148 6.3.4 MEA performance with ionomers with different IEC—higher peak power density obtained with ionomers with higher IEC. .............................................. 154 6.4 Conclusions ...................................................................................................... 156 6.5 References ........................................................................................................ 157 ix CHAPTER 1 INTRODUCTION 1.1 Motivation Since the 1900s, the world has seen a rapid growth of energy demands with fossil fuels as the major source of energy. The global consumption of fossil fuels in 2019 was more than 20 times of that in 1900.1 Because of the concerns on the depletion of fossil fuels and the pollution and greenhouse gases produced by their combustion, energy conversion with a clean and efficient pathway from renewable energies has drawn great attention from both academia and industry.2,3 Fuel cells are regarded as a promising candidate to solve these issues.4,5 They directly convert chemical energy, from the fuel, into electrical energy. If electrochemically generated hydrogen is used as the fuel, a green hydrogen cycle can be built, so that the renewable energy can be Figure 1.1. Simplified Ragone plot of the energy storage domains for the various electrochemical energy conversion systems compared to an internal combustion engine. Figure reproduced from ref. 5. 1 stored in the form of molecular hydrogen and converted into electrical energy with water as the only by-product. Since the electrochemical energy conversion is not restricted to Carnot cycle, the efficiency of fuel cells (> 60%) is generally much higher than that of combustion engines. When compared with other electrochemical systems for energy storage and conversion, e.g., batteries and supercapacitors (Figure 1.1), fuel cells stand out because of the higher specific energy they can achieve.6,7 Fuel cells are mainly composed of three parts: anode, cathode and electrolyte. The different kinds of fuel cells are presented in Figure 1.2.7 At the anode, the fuels are oxidized, and at the cathode, oxygen is reduced. The ions are transport between the two electrodes by the electrolyte. Categorized by fuels, there are hydrogen fuel cells and direct alcohol fuel cells (DAFC). Categorized by different electrolytes and Figure 1.2. Scheme of the reactions and processes that occur in the various fuel cell systems. Figure reproduced from ref. 6. 2 operation conditions, they can also be divided into alkaline fuel cells (AFC), proton exchange membrane fuel cells (PEMFC), phosphoric acid fuel cells (PAFC), molten carbonate fuel cells (MCFC) and solid oxide fuel cells (SOFC). Among them, AFC and PEMFC can reach high power densities at low temperatures, making them attractive in transportation. However, to compete with the internal combustion engine, there is still a long way to go for them, which requires further research and development. The high cost of the fuel cell stacks prevents their wide application in vehicles. Figure 1.3 shows the assumed advances achievable until 2025.8 Although the estimated cost of an 80-kWnet PEMFC system based on projections to 500k units/year has been brought down greatly to $45/kWnet in 2017, the ultimate DOE target is $30/kWnet, or $2,400/vehicle (~$3,000/ vehicle for internal combustion engine Figure 1.3. Modeled cost of an 80-kWnet fuel cell system based on projection to high-volume manufacturing (2016 U.S. dollars, left axis) and on-road demonstrated fleet average durability (durability record) with DOE cost and durability targets. Yellow star indicates the $75/kW status for durability-adjusted cost in 2018. Figure reproduced from ref. 8. 3 system9). According to the cost breakdown of fuel cell stacks (Figure 1.4)10, with the increase of production volume, the component costs that depend on processing would Figure 1.4. Breakdown of the 2017 projected fuel cell stack cost at 1,000, 100,000, and 500,000 systems per year. Figure reproduced from ref. 10. largely decrease, whereas the cost of precious-metal catalysts and bipolar plates, that are dominated by material costs, remains virtually unchanged. As a result, the largest fraction of the total cost (41%) would have come from catalysts if 500k systems/year were manufactured. However, to ensure the performance and the durability of the fuel cells, Pt-based nanoparticle catalysts are used in both cathode and anode. Lowering and even eliminating the use of the precious metals in the catalysts, while maintaining the high power density, is of great significance to improve the cost efficiency of the fuel cells. In addition, the durability of fuel cells is also important to the commercialization. 5000 hours (~10-year lifetime) under realistic operating conditions is the 2025 DOE durability goal for transportation fuel cells.10 4 1.2 Alkaline fuel cells fundamentals 1.2.1 Introduction Recently, anion exchange membrane fuel cells (AEMFC) have elicited increasing interest.11–17 The alkaline environment is less corrosive for the catalysts, compared to the standard PEMFCs, which enables the use of non-precious metal catalysts and a longer lifetime than PEMFCs. The development of alkaline anion exchange membrane (AAEM) materials has enabled the further development of AFCs and brought their application from space missions to automobiles.7 Hydroxyl ions and water molecules can be conveyed in those polymer electrolytes, while carbonate precipitation, a serious problem in alkaline fuel cells with liquid electrolytes, could be attenuated. The half reactions taking place at the cathode and anode can be written as follows: Oxygen reduction reaction (ORR) in alkaline media at the cathode: O2 + 2H O + 4e − 2 → 4OH − (1.1) Hydrogen oxidation reaction (HOR) in alkaline media at the anode: H2 + 2OH − → 2H2O + 2e − (1.2) The structure of the AEMFCs, including cathode, anode and anion exchange membranes is shown in Figure 1.5. It is notable that water is generated at the anode and consumed as a reactant for the ORR at the cathode in AEMFCs. It back diffuses from anode to cathode, which is the opposite from the case in PEMFCs. The change in the operating pH brings both new opportunities and challenges for the development of fuel cell technologies. Catalysts with high activity, long-term durability and low cost 5 Figure 1.5. Scheme of the AEMFC structure and the processes occurs in it. (Modified from the figure made by Jeesoo Seok) need to be developed for the reactions at the electrodes. To enhance the catalytic activity, understanding the reaction mechanisms in alkaline media is necessary, which are also more complicated than in acidic media. For anion exchange membranes, different kinds of polymer materials have been developed to improve their conductivity and stability, while the anion transport mechanism is still less understood. Since the catalysts and membrane materials need to work together and water distribution in the membrane electrode assembly (MEA) significantly influences the performance, their performance in an MEA also needs to be carefully investigated with the consideration of water management issues. 6 1.2.2 Electrocatalysis in alkaline media The HOR and ORR, together with their reverse reactions to produce hydrogen: hydrogen evolution reaction (HER) and oxygen evolution reaction (OER) complete the electrochemical hydrogen cycle. For these reactions, catalysts are necessary to lower the overpotential. Specifically, the ORR is one of the bottlenecks due to its sluggish kinetics.18 The reaction involves the transfer of four electrons and water as the reactant in alkaline media. The competing reaction is the two-electron transfer reaction that produces peroxide: O2 + H2O + 2e − → OH− + HO −2 (1.3) If peroxide is formed, the efficiency is significantly lower, and the membrane materials can be degraded by the peroxide generated radicals. To characterize the activity of catalysts, rotating disk electrode (RDE) voltammetry is usually conducted. Figure 1.6 depicts the three-electrode configuration used in RDE measurements and a typical ORR polarization curve profile expected. According to the Levich equation (Eqn. 1.4), the diffusion limited current (𝑖 ) should reach 5.5 mA/cm2𝑙 in O2-saturated Figure 1.6. Scheme of typical RDE setup in three-electrode configuration (reproduced from ref. 18) and typical ORR polarization curve of Pd/C with rotating rate of 1600 rpm in 0.1 M NaOH. CE: counter electrode; RE: reference electrode. 7 0.1 M NaOH solution if the reaction proceeds through the four-electron pathway and is tested at a rotation rate of 1600 rpm. With the Koutecky-Levich equation (Eqn. 1.5), the number of electrons transferred at a certain potential within the kinetic region can be estimated from the current at different rotation rates. Using the rotating ring disk electrode (RRDE) technique, the amount of the peroxide side product generated can be detected at the Pt ring, which also provides a method to determine the number of electrons transferred over the range of potentials scanned. Levich equation: 2/3 il = 0.62nFAC ∗D υ−1/6ω1/20 (1.4) 0 Koutecky-Levich equation: 1 1 1 = + (1.5) i i 2/3 k 0.62nFAC∗D υ−1/6ω1/20 0 F is Faraday constant (96,500 C/mol); A is the area of the electrode (0.196 cm2); D0 is the diffusion coefficient of O -5 22 in 0.1 M NaOH solution (1.93×10 cm /s); 𝜐 is the kinematic viscosity of the electrolyte (1.01×10-2 cm2/s); 𝜔 is the angular frequency of rotation, 𝐶∗0 is the concentration of molecular oxygen in 0.1 M NaOH solution (1.26×10-6 mol/cm3), 𝑖𝑘 is the kinetic current. 19 The thermodynamic equilibrium potential for the ORR is 1.23 V vs. RHE. Consequently, the higher the onset potential, the faster the kinetics of the reaction. The half-wave potential (E1/2) at half the limiting current is usually used for comparing the ORR activities of different catalysts. 8 Pd-based catalysts have been extensively studied for the ORR in alkaline media.20,21 Different from the case in acid, Pd/C shows a remarkably high activity that is comparable to that of Pt/C in base.22 According to the Sabatier principle (volcano relation), optimal catalysts adsorb the intermediates at medium strength (not too high, nor too low).23 Thus, tuning the oxygen binding energy by controlling the nanocrystal structure has been utilized for promoting the ORR activity of Pd. Cubic Pd nanoparticles enclosed by Pd(100) facets were found to be more active than spherical Pd nanoparticles.24 Alloying Pd with first row transition metal, such as Fe, Cu, Co and Ni, also modifies the electronic effects as well as the lattice strain on the surface of the nanoparticles, so that higher activity and lower costs were achieved.25–30 Non- platinum-group-metal (non-PGM) catalysts, such as transition metal oxides31–34, transition metal nitrides35,36 and metal-organic-framework(MOF)-derived carbon materials37,38, have also been widely studied to yield activities that are comparable to that of Pt/C and further cut down the cost. The HOR/HER in alkaline media, while faster when compared to the ORR, are around two orders of magnitude slower than in acidic media on the same catalyst surface.39 The reactions can involve two elementary steps: Tafel step/Heyrovsky step, and Volmer step. Tafel step: H2 ↔ 2Had (1.5) Heyrovsky step: 9 H + OH− ↔ H + e−2 ad + H2O (1.6) Volmer step: Had + OH − ↔ H2O + e − (1.7) The decrease in the HOR/HER activity with increasing pH on Pt was correlated with the hydrogen binding energy by Yan et al.40 However, there are still debates in describing the HOR/HER kinetics in alkaline media solely by the hydrogen binding energy. Markovic et al. proposed a bifunctional mechanism in which OHad is an important reactant and the HOR/HER kinetics increase with the oxophilicity of the catalysts.41 This explains the high HOR activity of PtRu/C, the best commercialized catalyst for alkaline HOR in AEMFCs studies, in which Ru provides the site for OHad. For the HER, they found that the activity on Ni(OH)2 modified Pt(111) surface was ~7 times as high as on a bare Pt(111) surface, which can also be explained by the bifunctional mechanism.42 Nevertheless, a study by Zhuang et al. on the HOR activity and the oxophilicity of a series of metal hydroxide/oxide modified Pt surfaces showed that the electronic effect, which weakens the Pt-Had interaction, is of more significance than the oxophilic effect in the enhancement of the HOR activity.43–45 More recently, it has been proposed that the interfacial water structure, which was characterized by the potential of zero total charge, has a great influence on the hydrogen adsorption kinetics, and thus the HOR/HER kinetics on Pt(111) surfaces.46–49 10 1.2.3 Polymer electrolyte materials Anion exchange membranes and anion exchange ionomers are polymer electrolytes in AEMFCs that transport hydroxide ions from the cathode to the anode and also deliver water between the two electrodes. Membranes separate the cathode and anode, while ionomers are mixed with the catalysts in the catalyst layers to provide better conductivity and ionic transport (Figure 1.5). They are usually composed of two parts: a structural part and a functional part. The structural part contains the polymer backbone which maintains the polymer structure integrity and mechanical strength. In the functional part, there are cations attached to the backbone which serve as the sites for anion exchange and transport. A variety of polymer backbones and cations has been developed with the aim of providing high conductivity and high stability to AEMFCs.50,51 Polyphenylenes, polyfluorenes, polystyrene, polyethylene, and polynorbornene are promising backbone materials. However, it has been found recently that polymers with aromatic backbones tend to adsorb on the catalysts, which has a negative influence on the performance of AEMFCs.52,53 Consequently, polyethylene is the polymer backbone of our interest in this dissertation as it has the advantages of an all-aliphatic hydrocarbon structure, ease of synthesis and high thermal stability. Cation categories that have been synthesized include ammonium- based cations, phosphonium-based cations and N-based resonance-stabilized cations.54 Based on the quantitative assessment on the chemical stability of various cation model compounds in KOH/CD3OH solution 55,56, penta-substituted imidazolium and tetrakis(dialkylamino)-phosphonium cations showed the highest stability. The 11 piperidinium and tetraalkylammonium cations are also significantly more stable than other cations. Although polymer electrolyte materials (with fixed/anchored cationic sites) are designed to mitigate carbonate precipitation in the alkaline aqueous electrolytes, carbonation is still an issue in AEMFCs. When the anion exchange membranes are exposed to air, they quickly react with CO2 and the conductivity becomes much lower due to the lower mobility of carbonate ions. Dekel et al. developed a method to measure the true value of the OH- conductivity.57 They applied a current through the membrane to produce hydroxide at cathode and release CO2 at the anode, so that the conductivity could be measured with fully OH- exchanged membranes. The performance of AEMFCs is also lowered by the presence of CO2. 58–60 However, this incorporation of carbonate has been found to be reversible both with in situ electrochemical quartz crystal microbalance (EQCM) study61 and membrane electrode assembly (MEA) studies62,63. It was also shown by simulations that the influence of CO2 on MEA performance can be diminished when operating at high current density.64 Water plays an important role in the stability of polymer electrolyte materials. Faster degradation rates have been observed under less hydrated conditions.65–67 Higher relative humidity helps maintain the conductivity and ion exchange capacity (IEC) of the membranes.68 In situ methods have been employed to investigate the water distribution at the interfaces under different relative humidity (RH) conditions and the mechanisms of water transport in AEMs.69–73 These include neutron 12 reflectometry73, phase-contrast tapping mode and conductive-probe atomic force microscopy71. Water management also has a great impact on MEA performance and durability, which can be optimized by controlling the water diffusion in the membrane.74–77 As a result, the study of polymer electrolytes needs to be combined with more in-situ/operando methods, so that the water and anion transport in them, with electrochemical environment can be better understood, which would guide the better design of these materials. 1.3 Research overview This dissertation covers the study of electrochemical processes and materials related to AEMFCs from fundamental mechanisms to practical aspects. The work presented can be divided into two parts: the study of electrocatalysts in alkaline media and the characterization of anion exchange membrane materials. In the field of electrocatalysis, the enhancement of the ORR activity in alkaline media by electrochemical dealloying process on Pd-based nanoparticle catalysts is described in Chapter 2. The structures of the nanoparticles, before and after the dealloying process, were characterized by electron energy loss spectroscopy (EELS) elemental mapping, CO stripping and X-ray photoelectron spectroscopy (XPS), to rationalize the significant improvement in catalytic activity. Chapter 3 introduces a novel scan-rate- dependent cyclic voltammetric method that we have developed, to directly measure the kinetics of electro-adsorption processes that are vital in electrocatalysis. Hydrogen adsorption on single crystal Pt(111) surfaces was used as a model system, as it 13 provides insight to the kinetics of the Volmer step of the HER. Its kinetics in alkaline media were found to be over 2 orders of magnitude lower than in acidic media. This method was further applied to various alkaline systems (with different cations or deuterated water) to understand the mechanism of the Volmer step in alkaline media in Chapter 4. The role of interfacial water structure is also discussed. In terms of the development and characterization of polymer electrolytes, an electrochemical quartz crystal microbalance (EQCM) study on a phosphonium-based alkaline anion exchange membrane material is described in Chapter 5. Anion exchange and water dynamics accompanying the methanol oxidation reaction in the membranes are discussed. Chapter 6 presents the influence of IEC on the properties of anion exchange ionomers in aqueous environments. QCM and electroanalytical techniques were employed to characterize the mechanical properties of the ionomers in water and the transport of charged and neutral redox species in the ionomers, respectively. Those properties were also related to the MEA performance with ionomers with different IEC. 1.4 References (1) Vaclav Smil (2017). Energy Transitions: Global and National Perspectives. & BP Statistical Review of World Energy. https://ourworldindata.org/grapher/global-fossil-fuel- consumption?country=~OWID_WRL. (2) Satyapal, S. 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Soc. 2020, 167, 054501. 26 CHAPTER 2 ELECTROCHEMICALLY DEALLOYED Pd-BASED NANOPARTICLE CATALYSTS FOR THE OXYGEN REDUCTION REACTION (ORR) IN ALKALINE MEDIA* 2.1 Introduction In AEMFCs, the design of effective cathode catalysts for the ORR remains one of the major challenges that result from the high overpotentials required to drive the four- electron reduction reaction of oxygen. Finding an efficient, low-cost and durable electrocatalyst for the ORR has been the subject of intense study.1 Among catalysts developed for the ORR in alkaline media, precious-metal-based catalysts, specifically Pd-based nanoparticles2,3, still possess the highest efficiency and stability. Thus, optimizing their atomic utilization to lower the cost and enhance their activity becomes both important and attractive.4,5 Nevertheless, catalyst development for improving activity and cost efficiency have been much more widely examined in acidic media than in alkaline media. Alloying a precious metal with a first row transition metal, with a smaller lattice constant, such as Fe6–8, Co9–11 or Ni17–20, is a powerful method for controlling the surface strain of the catalysts for the ORR, which, in turn, tunes the activity.5,16,17 Nanoparticle dealloying―the process through which the less noble metal is partially removed from a nanoparticle alloy―is another approach for further optimizing the structure18–20, lattice strain21,22 and catalytic * Adapted and reproduced in part with permission from Lu, X., Ahmadi, M., DiSalvo, F. J., & Abruña, H. D. (2020). Enhancing the Electrocatalytic Activity of Pd/M (M= Ni, Mn) Nanoparticles for the Oxygen Reduction Reaction in Alkaline Media through Electrochemical Dealloying. ACS Catalysis, 10(10), 5891-5898. Copyright 2020 American Chemical Society. 27 activity23–26 of a catalyst, which has been mostly studied on Pt-based catalysts for the ORR in acidic media27–29. In a previous systematic study, we showed the power of the electrochemical dealloying process for tuning the morphology and ORR activity of Cu3Pt/C nanoparticles in acidic media. 30 Reasoning by analogy, these strategies should also be effective in promoting the ORR activity of catalysts at high pH. There have been some studies in alkaline media investigating the effects of alloying7,31, core-shell structure32 and electronic structure33 of Pd alloy nanoparticles on ORR activity. However, due in part to the different reaction mechanisms at high pH34, further improvement on Pd alloy catalysts with other methods still needs to be explored. In this chapter, the study of an electrochemical dealloying strategy for Pd-base nanoparticles and its effect on activity for the ORR in alkaline media will be discussed. Since Ni and Mn are 3d transition metals that have smaller lattice constant than Pd and also form alloy phases with Pd, they were selected as representatives for the Pd-M nanoparticle catalysts to study the effects of dealloying. The dealloyed PdNi/C and PdMn/C catalysts achieved 36 and 60 mV enhancement in activity (E1/2 value), respectively, by selectively leaching out the less noble element (Ni or Mn) through cycling the potential 10 times between 0.05 V and 1.1 V vs. the reversible hydrogen electrode (RHE) at a sweep rate of 50 mV/s in 0.1 M HClO4. The specific activities of dealloyed PdNi/C and PdMn/C at 0.85 V were 3.3 and 3.7 times higher than that of Pd/C, respectively. The nanoparticle core-shell structure and reduced surface oxide, observed after dealloying, were found to play key roles in the enhanced activity. Furthermore, after the degradation following extensive potential cycling, the activity could be restored by an additional electrochemical dealloying treatment. The 28 half-wave potential of the dealloyed PdNi/C nanoparticles after 16,000 potential cycles with 6 cycles of cycling-dealloying protocol, was only 14 mV lower than for the initial catalyst. This facile post-synthesis strategy might also be applied for optimizing the atomic utilization of other electrocatalysts. 2.2 Experimental section 2.2.1 Synthesis of PdxNi/C nanoparticles Carbon-supported PdxNi (x=3 or 1) nanoparticles with 20 wt% metal loading were synthesized by a two-step method.15 First, calculated volumes of 10 mM NiCl2·6H2O and 10 mM Na2PdCl4 stock solutions were added to deionized water to make a diluted solution with 1 mM fixed concentration of Pd2+. The Na2PdCl4 stock solution was prepared by dissolving PdCl2 and NaCl in a 1:2 molar ratio in gently warmed water. A given amount of Vulcan XC-72 carbon, leading to a 20 wt% metal loading, was dispersed by sonication for around 1 h. A 0.5 M NaBH4 solution, serving as the reducing agent, at a molar ratio of 10:1 to the metal was quickly added under vigorous stirring at room temperature, and the mixture was stirred continuously overnight. After being separated and washed 3 times with water and ethanol through centrifuging at 7,000 rpm, the catalysts were dried in an oven at 70℃. The resulting powders were annealed in a tube furnace under flowing forming gas (5% H2 balance N2) for 2 h at 700℃ to maximize the extent of alloying. For comparison, Pd/C with 20% Pd loading was synthesized with the same NaBH4 reduction method, but without annealing in the forming gas to prevent aggregation. 29 2.2.2 Synthesis of PdMn/C nanoparticles PdMn nanoparticles were synthesized using a Schlenk line. In a typical synthesis, 10.1 mg of MnCl2 and 60 μL of oleic acid were added to 10 mL of diphenyl ether in a three-neck flask. In another flask, with a rubber septum, 31.5 mg of Pd(acac)2 and 80 μL of oleic acid were added to 10 mL of diphenyl ether. The MnCl2 solution was heated to 220℃ and 0.6 mL of 2.5 M n-BuLi were injected into the solution. After stirring for 20 min at 220℃, the Pd(acac)2 solution was injected and stirred for another 20 min before removing from the heat. The solution was then stirred for 10 min, and ultrasonicated for 10 min with suspended carbon powder (Vulcan XC-72) to achieve a 20 wt% metal loading on the carbon support. The mixture was centrifugated and washed 3 times with methanol and hexanes to remove the surfactants. The resulting product was collected by centrifugation. The prepared Pd-Mn nanoparticles were subsequently annealed under forming gas for 30 min at 750 ℃ to achieve the ordered intermetallic structure. 2.2.3 X-ray characterization The obtained catalysts were characterized by powder X-ray diffraction (XRD) using a Rigaku Ultima VI diffractometer with Cu Kα radiation (λ = 0.154178 Å) at a scan rate of 5 deg/min. 2.2.4 Electrochemical dealloying and testing of the nanoparticles All electrochemical experiments were performed using a Princeton Applied Research potentiostat (Model: 273A) in a standard three-electrode cell at room temperature. 10 30 μL of 4 mg/mL catalyst ink, prepared by sonicating a mixture of catalyst and Nafion solution (0.05 wt% Nafion dissolved in isopropanol), were applied to a glassy carbon rotating disk electrode (RDE) as a working electrode and dried under infrared light. Graphite was used as the counter electrode, and Ag/AgCl in 1 M KCl (0.235 V vs. SHE) was used as the reference electrode. The potentials in this work are reported vs. RHE using the following equation: ERHE = Eapplied + 0.235 + 0.0592 pH Before electrochemical testing of the as-prepared catalysts, an electrochemical dealloying protocol was carried out in Ar-saturated 0.1 M HClO4 solution. The potential was scanned from 0.05 V to 1.1 V vs. RHE at 50 mV/s for 10 cycles. Subsequently, the electrode was electrochemically cleaned in Ar-saturated 0.1 M KOH solution by cycling the potential from 0.05 V to 1.2 V vs. RHE at 50 mV/s for 10 cycles. The ORR evaluation was conducted in O2-saturated 0.1 M KOH solution at a rotation rate of 1600 rpm at 5 mV/s. CO stripping voltammograms were carried out to measure the electrochemical surface area (ECSA) of the catalysts. A dosing potential for CO adsorption of 0.2 V vs. RHE was applied to the electrode while bubbling CO for 20 min, followed by Ar purging for 20 min to remove any CO remaining in the electrolyte. The CO stripping scan was carried out immediately after the purging, starting from 0.2 V to 1.2 V vs. RHE at a scan rate of 10 mV/s in 0.1 M KOH. 31 Durability tests were carried out by scanning the potential from 0.6 V to 1.0 V vs. RHE in an Ar-saturated 0.1 M KOH at 100 mV/s. The ORR activities were assessed after every 2000 or 4000 cycles, and the dealloying protocol was repeated 6 times. 2.2.5 Electron microscopy An FEI Tecnai 12 BioTwin transmission electron microscope (TEM) was used to image the PdNi/C and PdMn/C nanoparticles as prepared and after stability testing. The catalysts were dispersed in ethanol by ultrasonication and drop-cast on lacey carbon copper TEM girds. The accelerating voltage was 120 kV for bright field TEM imaging. Size distributions were obtained from TEM images using ImageJ software. The atomic-resolution, annular dark field, scanning transmission electron microscopy (ADF-STEM) images and electron energy loss spectroscopy (EELS) images were acquired on a fifth-order aberration-corrected scanning transmission electron microscope (FEI Titan Themis) operated at 300 keV. The microscope was operated with a ∼30-mrad probe forming semi angle. Pd-M4,5 and Ni-L2 edges were used to extract Pd and Ni EELS maps with an acquisition time of 0.05 and 0.1 s/pixel, correspondingly. 2.2.6 X-ray photoelectron spectroscopy (XPS) characterization XPS was conducted directly on the catalysts with Nafion after different treatments without removing them from the glassy carbon electrode. The measurements were acquired using a Scienta Omicron ESCA2SR in an ultra-high vacuum (UHV) system with operating pressure of ~2 x 10-10 Torr. A monochromatic X-ray source (Al Kα, 1486.6 eV) was used with an analysis area of ~ 4 × 4 mm. A hemispherical analyzer 32 determined electron kinetic energies, using a pass energy of 200 V for wide/survey scans, and 50 V for high resolution scans. Photoelectrons were collected at a 0-degree emission angle. 2.3 Results and discussion Pd:M (M = Ni, Mn) with a ratio of 1:1 and Pd:Ni with a ratio of 3:1 were synthesized by reducing the metal precursors with strong reducing agents, NaBH4 for Pd-Ni/C nanoparticles, and n-BuLi for Pd-Mn/C nanoparticles, followed by high temperature annealing. As shown in the X-ray powder diffraction pattern (Figure 2.1a), both PdNi/C and Pd3Ni/C exhibited a single-phase (fcc) alloy structure. Shifts in the diffraction peaks to higher 2θ angles, observed upon increasing the Ni content, indicate that the smaller Ni atoms are incorporated into the Pd lattice, inducing a lattice contraction. The lattice constant of PdNi/C was determined to be 3.774 Å from the peak diffraction angles and the Bragg Equation. In the case of PdMn/C, the peaks in the XRD match those for intermetallic PdMn with a body-centered tetragonal crystal structure and lattice constants a = 2.877 Å, c = 3.585 Å. The lattice contraction in PdNi/C was further confirmed by atomic-scale STEM imaging of PdNi/C particles (Figure 2.1b). The particle was viewed along the [001] zone axis. A d-spacing of 0.19 nm was determined from the Fourier transform, and is, within error, the same as the value calculated from the (200) peak in the XRD. The bright field TEM (BF-TEM) 33 Figure 2.1. (a) XRD patterns of Pd3Ni/C, PdNi/C and PdMn/C nanoparticles. (b) Atomic- resolution ADF-STEM image of a PdNi/C nanoparticle. The inset shows the Fourier transform of the particle with the diffraction spots corresponding to the lattice d-spacing of (200) facets calculated from the (200) peak in the XRD for PdNi/C. BF-TEM image of (c) PdNi/C, (d) PdMn/C and (e) Pd/C nanoparticles. The insets in (c) (d) and (e) show the size distributions of PdNi/C, PdMn/C and Pd/C measured from the TEM images, respectively. images (Figures 2.1c and d) show that both PdNi and PdMn particles were well- dispersed on the Vulcan carbon support. However, aggregation of the nanoparticles was induced by the high temperature annealing process. The average particle sizes, determined from the TEM images (inset of Figures 2.1c and d) for PdNi/C and PdMn/C, were 8.0 ± 3.9 nm and 13.0 ± 9.6 nm, respectively, which are larger than for the as-synthesized Pd/C nanoparticles (2.8 ± 1.0 nm, Figure 2.1e). Electrochemical dealloying treatments in 0.1 M HClO4 with potential cycling from 0.05 V to 1.1 V (vs. RHE) were employed on the catalysts before ORR performance evaluation. With only 10 cycles of electrochemical dealloying, the surface leaching of Ni on the PdNi/C 34 Figure 2.2. STEM images of PdNi/C (a) as synthesized, (b) after electrochemical dealloying and (c) after stability testing. Corresponding EELS elemental composition maps of Pd (red) and Ni (green) for PdNi/C (d) as synthesized, (e) after electrochemical dealloying, and (f) after stability testing. (g) EELS elemental line profile of Pd and Ni extracted from the white dashed box in (e) from the EELS mapping. nanoparticles led to a core-shell structure with a 1~2-nm Pd shell covering a PdNi core, as evident from the absence of the Ni signal from 1 ~ 3 nm and 13 ~ 15 nm on the relative distance scale in the line profile (Figure 2.2g) in the EELS composite elemental mapping (Figure 2.2e). In contrast, Ni was homogeneously distributed throughout the nanoparticles in the case of the as-synthesized particles (Figure 2.2d). The ORR activity of the Pd-based nanoparticles in alkaline media was significantly improved after the electrochemical dealloying process. Relative to the response prior to dealloying, the cyclic voltammogram of dealloyed PdNi/C, shown in Figure 2.3a, exhibited more prominent Pd features in the H-adsorption region, and the Pd-O reduction peak (~0.7 V) was shifted positively, closer to the peak position in the 35 Figure 2.3. Cyclic voltammograms in 0.1 M NaOH purged with Ar, at a sweep rate of 50 mV/s, at room temperature for (a) PdNi/C, (b) PdMn/C and (c) Pd3Ni/C, before and after dealloying. ORR polarization curves in O2-saturated 0.1 M NaOH of (d) PdNi/C, (e) PdMn/C and (f) Pd3Ni/C, before and after dealloying. Rotation rate: 1600 rpm, sweep rate: 5 mV/s. CV of Pd/C (Figure 2.4). The shift in the Pd-O reduction peak for PdNi/C after dealloying indicates a more Pd-like surface for the dealloyed PdNi/C nanoparticles, which agrees well with the STEM results. Similar phenomena were also found for Figure 2.4. CV of Pd/C in 0.1 M NaOH purged with Ar, sweep rate 50 mV/s, room temperature. 36 PdMn/C (Figure 2.3b), in which the peak at around 0.9 V, likely due to the oxidation of the surface Mn,35 disappeared after the dealloying process. The half-wave potential for the ORR in 0.1 M NaOH of the dealloyed Pd-based nanoparticles shifted positively by 36 mV for PdNi/C and 60 mV for PdMn/C compared to the as- synthesized samples, as shown in Figures 2.3d and e. This can be attributed to the higher Pd atomic concentration on the surface. However, a change in electrochemical properties was not evident for the dealloyed Pd3Ni/C nanoparticles obtained by the same method (Figures 2.3c and f), which suggests that the electrochemically dealloyed PdNi/C catalyst is different from the Pd-rich (Pd3Ni) nanoparticle catalyst. The ORR activity of the various catalysts is compared in Figure 2.5. Among all the Pd-based catalysts tested, PdNi/C showed the highest half-wave potential (E1/2) with a value of 0.87 V, comparable to Pd/C (0.88 V). The relatively low activity of PdMn/C is likely due to the large particle size that resulted from the high temperature annealing step required in its synthesis. Figure 2.5. Comparison of the ORR polarization curves of the dealloyed Pd-based nanoparticle catalysts and Pd/C in O2-saturated 0.1 M NaOH. Rotation rate: 1600 rpm, Sweep rate: 5 mV/s. The inset is the enlarged region near half-wave potential. 37 CO stripping voltammetry was conducted on the dealloyed Pd-based catalysts to determine their electrochemical surface areas (ECSA) for comparison with the as- synthesized samples (Figure 2.6). The very small peak on the forward scan for the as- synthesized PdxNi/C nanoparticles indicates that only a small amount of Pd is exposed on the surface, which explains the low ORR activity of the as-synthesized samples. After electrochemical dealloying, the intensity of the CO stripping peak increased significantly, which confirms the results from the EELS elemental mapping that the Pd atomic concentration on the surface is higher on the dealloyed nanoparticles. We also noted that the CO stripping behavior of PdMn/C catalysts, before and after dealloying, was different from that of PdNi/C catalysts. A broad CO stripping peak appeared for PdMn/C before dealloying, while the peak shifted to lower potential and became narrower after dealloying. This suggests that there are different sites that bind CO to a different extent on PdMn/C before/after dealloying. The sites that bound CO most strongly, likely due to surface Mn species, were removed after electrochemical dealloying, which is different from the case of PdNi/C. However, compared to Pd/C, the peak position for both dealloyed PdNi/C and PdMn/C shifted negatively by around 80 ~ 100 mV, which indicates a weaker CO binding energy for both dealloyed Pd- based nanoparticles, relative to Pd/C. 38 Figure 2.6. CO stripping voltammetry of (a) Pd/C, Pd3Ni/C and PdNi/C before and after dealloying, (b) PdMn/C before and after dealloying, at a scan rate of 10 mV/s in an Ar-saturated 0.1 M NaOH solution. Dosing potential of CO adsorption was 0.2 V vs. RHE. Solid lines are the first cycle; dashed lines are the forward scan of the second cycle. Table 2.1. Average particle size, electrochemical surface area (ECSA) and activities of different catalysts studied in this work Sample Avg size ECSA E1/2(V) Mass activity at Specific activity (nm) (m2/g ) 0.85V (mA/μgPd) at 0.85V Pd (mA/cm2Pd) Pd/C 2.8 ± 1.0 78 0.88 0.28 0.55 PdNi/C 8.0 ± 3.9 37 0.87 0.37 1.79 PdMn/C 13.0 ± 9.6 16 0.87 0.30 2.06 Since E1/2 is not the most appropriate descriptor of activity when catalysts with different size and structure are compared, the ORR activities of different catalysts were normalized to the electrochemical surface area determined by CO stripping and Pd mass, respectively. The electrochemical surface areas were calculated using the following equation from the CO stripping peak in Figure 2.6: Q S(cm2) = 420μC/cm2 × 0.75 39 Area Q = scan rate where Q is the charge passed, obtained by dividing the integrated area of the CO stripping peak by the scan rate. 0.75 is the fraction number for the saturated CO coverage on Pd reported in the literature.36 The average particle size, the ECSA, and the activities of Pd/C, PdNi/C and PdMn/C are summarized in Table 2.1. Although the E1/2 of Pd/C is higher than that of the dealloyed nanoparticles, the mass activities of Figure 2.7. (a) Mass activity of different catalysts. (b) Specific activity of different catalysts. PdNi/C and PdMn/C (Figure 2.7a) reached 0.37 and 0.30 mA/μgPd respectively at 0.85 V, both of which are higher than the value for Pd/C (0.29 mA/μgPd). The specific activities of both the dealloyed Pd-based nanoparticles surpassed that of the Pd/C catalyst by over a factor of 2 (Figure 2.7b), indicating a higher intrinsic electrocatalytic activity. This suggests that mass activity and specific activity are more appropriate for comparing the intrinsic activity of different catalysts. The mass activity, specific activity and E1/2 of the state-of-art Pd-based catalysts are listed and compared with our catalysts in Table 2.2. Although some Pd-based catalysts exhibited similar E1/2 values, their mass and specific activities were quite different from each other. Both mass and specific activity of our dealloyed catalysts are above most of the 40 other state-of-art Pd-based catalysts in the table, which illustrates the advantage of dealloying for enhancing the activity of PdM catalysts. Interestingly, the mass activity of PdMn/C was lower than that of PdNi/C, even though it had the highest intrinsic specific activity. This might be due to the larger particle size after the high temperature annealing. This also suggests that the activity of the dealloyed Pd-based catalysts could be further improved if smaller nanoparticle sizes could be obtained by optimizing the synthesis method. Table 2.2. Comparison of the mass activity, specific activity and E1/2 of the state-of-art Pd-based catalysts and our catalysts Catalyst Mass activity @ 0.85 Specific activity @ 0.85 E1/2 Reference V (mA/μgPd) V (mA/cm2) @1600rpm (V) PdMn/C-BAE 0.093 - 0.87 35 PdFe/C-BAE 0.081 - 0.85 (cit)PdFe/C - 0.17 0.85 37 (cit)PdCo/C - 0.19 0.87 (micro)PdFe/C - 0.25 0.85 (micro)PdCo/C - 0.1 0.85 Pd3Fe NPs/CB 0.049 - 0.80 7 Pd3Co NPs/CB 0.051 - 0.81 Pd3Ni NPs/CB 0.075 - 0.81 Ni@Pd3 0.095 0.13* 0.85 32 PdCu/Vulcan 0.58 1.3 0.91 38 PdCuCo 0.12* - 0.87 39 NPs/C-375℃ PdAuCu 1.781*⸷ 1.08*⸷ 0.93⸷ 40 Pt/PdCo10/C 0.65 - 0.85 41 Dealloyed 0.37 1.79 0.87 This PdNi/C work Dealloyed 0.30 2.06 0.87 This PdMn/C work *The number was obtained at 0.9 V vs. RHE. ⸷The experiment was conducted in 1 M KOH. 41 The PdNi/C and PdMn/C catalysts were very stable in alkaline media. After 4,000 cycles from 0.6 V to 1.0 V in Ar-saturated 0.1 M KOH at 100 mV/s, there was almost no degradation in the ORR half-wave potential. Additional electrochemical dealloying, after degradation during extended stability testing, could restore the ORR activity. This cycling-dealloying process was repeated 6 times, and the changes in the half-wave potential for dealloyed PdNi/C at different stages are presented in Figure 2.8a. The point at 0 cycle is the initial E1/2 of the dealloyed PdNi/C catalyst. The black points are the E1/2 values after the corresponding potential cycles, and each red point was obtained after additional electrochemical dealloying following the corresponding potential cycles. The data clearly show that repeated dealloying results in the recovery of electrocatalytic activity following stability testing. Even after 16,000 cycles, a dealloying treatment yielded a half-wave potential that was only 14 mV lower than the initial value for the dealloyed catalyst, prior to any stability testing. As we can see from the EELS elemental map after stability testing (Figure 2.2f), the PdNi/C particle Figure 2.8. (a) Half-wave potential change of PdNi/C during the stability test before and after the dealloying process in 0.1 M NaOH, (b) CVs of PdNi/C before stability testing and after 16,000 cycles in 0.1 M NaOH purged with Ar, sweep rate 50 mV/s, room temperature, (c) ECSA estimated from Pd-O reduction peak in CV of PdNi/C during stability testing in 0.1 M NaOH before and after dealloying, and the corresponding specific activity at 0.85 V after dealloying. (Note: the calculation of ECSA and specific activity here is different from that in Table 2.1 and Figure 2.7b.) 42 exhibited a homogenous distribution of Ni instead of a core-shell structure, which suggests that the less noble element moved to the surface of the nanoparticles during the potential cycling. This change in elemental distribution, which is further discussed below in the XPS results, could be the reason for the decrease in E1/2 during potential cycling. Thus, the removal of the surface Ni in the repeated dealloying process would lead to the recovery of the E1/2. Also, the degradation could be due to the aggregation of nanoparticles, as can be seen from the larger average particle size from the TEM after stability testing (Figure 2.9). To better understand the effects of aggregation and Figure 2.9. TEM images for PdNi/C after stability testing and the extracted size distribution. See Figure 2.1c for the comparison of PdNi/C before stability testing. dealloying on activity during potential cycling, the ECSA was tracked, from the Pd-O reduction peak in CVs for convenience, for the PdNi/C catalyst before and after the dealloying process during stability testing (Figure 2.8c). The following equation was used for the calculation42: Q S(cm2) = 430μC/cm2 43 Area Q = scan rate where Q is the charge passed, obtained by dividing the integrated area of the Pd-O reduction peak in CV by the scan rate. It needs to be pointed out that the ECSA, estimated from the chemisorption of a monolayer of oxygen, is much smaller than that calculated from the CO stripping method. It is also not as accurate as the CO stripping method because the oxygen adsorption charge on Pd depends on the potential range, making it difficult to precisely determine the upper potential limit at which the formation of the first monolayer of oxygen is completed.43 Nevertheless, it is still meaningful and more convenient to compare the ECSA within the same sample during stability testing. The ECSA exhibited a peaked response with increasing potential cycles and dealloying process in between 2,000 or 4,000 cycles in alkaline media which we ascribe to the presence of and competition between two factors controlling the ECSA. One is the potential cycling in alkaline media, which would lead to the aggregation of the nanoparticles or the diffusion of Ni species to the surface of the nanoparticles. The other is the dealloying process that would remove the surface Ni species that block the active sites of the nanoparticles. The former would cause a decrease in ECSA while the latter would result in a higher ECSA. By calculating the specific activity at 0.85 V from the estimated ECSA, we found that the specific activity was relatively stable within 4,000 cycles. It started decreasing after 4,000 cycles and was stabilized again at around 10,000 cycles. These results indicate the effectiveness of electrochemical dealloying in restoring the ECSA until 8,000 cycles. The small increase of ECSA after re-dealloying at 4,000 cycles suggests that the 44 diffusion of Ni to the surface was not significant until after 4,000 cycles. However, after 10,000 cycles, the aggregation effect started to dominate, because there was less and less Ni in the particles that could be removed, so that the ECSA decreased and could not be restored to the previous value after 10,000 potential cycles. As a result, the intrinsic activity decreased to a limit. A similar recovery of the activity after stability testing, followed by electrochemical dealloying, was also observed for PdMn/C. The ORR polarization curves for PdMn/C and PdNi/C before stability testing, after 12,000 cycles, and after 12000 cycles following a 10-cycle re-dealloying, are shown in Figure 2.10, as an example, to show that the electrochemical dealloying protocol can be an easy and convenient way to recover/restore the electrocatalytic activity of the Pd-based nanoparticle catalysts after long-term use. Figure 2.10. ORR polarization curves for (a) PdNi/C and (b) PdMn/C, before stability testing, after 12,000 cycles and after 12000 cycles following a 10-cycle re-dealloying, in O2-saturated 0.1 M NaOH. Rotation rate: 1600 rpm, Sweep rate: 5 mV/s. 45 To further study the reason(s) behind the electrocatalytic activity change before and after dealloying and the durability tests, the elemental composition and chemical state of the nanoparticles were examined by high-resolution XPS on an as-synthesized PdNi/C sample, after electrochemically dealloying and after the stability testing. The XPS spectra from the Ni-2p and Pd-3d core level regions of PdNi/C could be deconvolved into 3 duplets respectively, shown in Figures 2.11a and b, corresponding to: Ni (2p3/2, 852.5 eV), Ni 2+(2p3/2, 855.4 eV) and a satellite (2p3/2, 861.1 eV); Pd (3d5/2, 335.4 eV), Pd 2+(3d , 337.3 eV) and Pd4+5/2 (3d5/2, 338.3 eV). The relative content of Pd and Ni species after indicated treatments (Figure 2.11c) was extracted from the XPS spectra by integrating their respective peak areas and normalizing them to the Figure 2.11. XPS spectra for the (a) Ni-2p and (b) Pd-3d core level regions of PdNi/C after the indicated treatments. (c) Pd to Ni atomic ratio extracted from XPS measurements. Content of (d) Ni and (e) Pd species extracted from XPS spectra. 46 initial ratio of the two elements. Since the mean free path of electrons is ~1.4 nm for Pd and ~1 nm for Ni, 63% of the XPS signal would come from photoelectrons that are emitted from atomic layers in the top ~1.4 nm (~4-5 layers) from the surface.44 The significant rise in the Pd:Ni ratio after the dealloying treatment agreed with the results from EELS mapping and the electrochemical studies. In essence, a Pd-rich shell with a PdNi alloy core formed after the electrochemical dealloying process. The drop in the Pd:Ni ratio after the stability testing suggests that the degradation in the ORR activity can be related to the lower Pd content on the surface of the nanoparticles, resulting from Ni diffusing to the surface of the nanoparticles. The observation of a homogenous distribution of Pd and Ni in the EELS mapping of a particle, after stability testing (Figure 2.2f), instead of a core-shell structure, as observed after dealloying and before cycling in alkaline media, also provides evidence of the movement of Ni to the surface. This effect could be triggered by the adsorption of OH in alkaline media. DFT calculations by Nørskov et al.16 showed that the OH adsorption energy of Ni (ΔEOH = 0.13 eV) is higher than that of Pd (ΔEOH = 0.92 eV). As a result, the Ni atoms would segregate to the surface during the potential cycling due to their affinity to OH. A similar behavior was reported for the case of PtNi by C. Cui et al.45, and was explained by the more stable Ni-OH surface spontaneously formed at high pH. From the deconvolved peaks, the relative Ni and Pd species content were extracted, and are shown in Figures 2.11d and e. The sample, after dealloying, showed the lowest Ni2+ and Pd4+ content, indicating the removal of surface oxides after the dealloying process. The Ni2+ content rose again after the stability testing, suggesting that Ni atoms moved to the surface of the nanoparticles and were oxidized during the 47 cycling, due to the higher affinity of Ni atoms for hydroxides in the electrolyte45. The XPS results also support the fact that 10-cycle electrochemical dealloying is effective in removing the surface nickel oxide species, which helps expose more (Pd) active sites for ORR. 2.4 Conclusions We have applied and studied the effect of an electrochemical dealloying treatment on Pd-M (M = Ni and Mn) nanoparticles as catalysts for the ORR in alkaline media. The electrocatalytic activity was enhanced by the dealloying process which partially leached out the inactive transition metal species (Mn, Ni) on the surface, and generated Pd-rich surfaces on the nanoparticles. 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(44) Shinotsuka, H.; Tanuma, S.; Powell, C. J.; Penn, D. R. Calculations of Electron Inelastic Mean Free Paths. X. Data for 41 Elemental Solids over the 50eV to 200keV Range with the Relativistic Full Penn Algorithm. Surf. Interface Anal. 2015, 2015, 871–888. (45) Cui, C.; Ahmadi, M.; Behafarid, F.; Gan, L.; Neumann, M.; Heggen, M.; Cuenya, R.; Strasser, P. Shape-Selected Bimetallic Nanoparticle Electrocatalysts : Evolution of Their Atomic-Scale Structure , Chemical Composition , and Electrochemical Reactivity under Various Chemical Environments. 2013, 91–112. 55 CHAPTER 3 RATE AND MECHANISM OF ELECTROCHEMICAL FORMATION OF SURFACE BOUND HYDROGEN 3.1 Introduction A large fraction of energy-critical reactions occurs via surface hydrogen atoms (Had). For example, the hydrogen evolution reaction (HER) forms Had as an intermediate as the reaction transfers two electrons to protons in acid (2H+ + 2e– → H2) or water in alkaline media (2H2O + 2e – → H2 + 2OH –)1–3. According to the Sabatier principle, one can use the Had binding strength to estimate the effectiveness of an HER electrocatalyst4–7. Specifically, a good HER electrocatalyst should bind to Had not too strongly, such that Had becomes a surface-blocking species, nor too weakly, such that forming Had requires excessive input energy. The challenge is that this picture forgoes the critical step of the Had formation rate, and recently, this assumption was called into question7–9. As pointed out by a theoretical analysis by Zeradjanin et al., the rate of the Had electro-adsorption played just as important a role as the Had binding strength in the HER10. In this chapter, we identify the impact of the Had formation rate on the HER by directly measuring the intrinsic rate constant of the Had formation (i.e., under thermoneutral condition). Our best knowledge of the Had formation (“the Volmer step”) kinetics comes from impedance studies. Conway et al.11 and Sibert et al.12 reported on the Had electro- adsorption rate in acid on platinum single crystals. Unfortunately, they could estimate only the Had formation rate, not the intrinsic rate constant, which requires mapping 56 how the applied potential impacts the electro-adsorption current (vide infra). Oelgeklaus et al.13 and Schouten et al.14 got around this issue in alkaline media by assuming that the interaction between Had could be described by the Frumkin isotherm. However, their approach was only applicable to alkaline environment, not acid, where the experimental Had electro-adsorption rate constant is still elusive. Furthermore, it is not yet possible to extract the rate constant of the electro-adsorption without assuming an equilibrium isotherm model, where the validity of the assumption is often unknown. We present a method for quantifying the electro-adsorption rate as a function of driving force and coverage without assuming any isotherm. Applying our methodology to single-crystal Pt(111), we report, for the first time, the rate constant of the Volmer step in acid. We further examine the nature of the Volmer step by investigating its pH dependence, from which we reveal the mechanistic differences between the Volmer step in acidic and alkaline media. We also quantify the Had formation rate on Pt(111) decorated with the surface modifier and confirm the role of surface modifier in accelerating the Volmer step. Our method enables a direct investigation into the effect of the Had formation rate on the HER and paves an avenue for exploring descriptors beyond the Had binding strength. 57 3.2 Methodology 3.2.1 Accessing kinetics of electro-adsorption without an isotherm assumption Our methodology relies on rate-dependent cyclic voltammetry (CV). This method has been used to measure the electron transfer rate constant of surface-bound species15–17. As shown by Laviron, the electron transfer kinetics can be extracted by analyzing how the peak potentials vary with scan rate in the voltametric profiles18. However, this technique applies only to situations in which the surface adsorbates are either non- interactive (i.e., Langmuir isotherm) or interactive in a way that is describable by the mean-field interaction (i.e., Frumkin isotherm)19. As a result, the approach has limited applicability to heterogeneous electrocatalysts, whose surface behaviors often do not behave as either. Motivated by the need for a rate measurement strategy, we have developed a scan rate-dependent CV approach that enables the measurement of the electro-adsorption kinetics without assuming a specific isotherm. The presented methodology can thus be used to characterize the rate and the rate constant of any electro-adsorption process. Herein, we demonstrate the power of this methodology for studying the Volmer step to quantify the rate constant of the elementary step of the HER on Pt(111). The key aspect is that the driving force for electro-adsorption, i.e., the adsorption overpotential ( = E – Eeq, where E is the applied potential and Eeq is the equilibrium potential at a given coverage), is a function of surface coverage (θ). That is, electro- adsorption rates measured at the same electrochemical potential may not have the same driving force if the surface coverage is not the same, since Eeq would be 58 different. To determine the driving force and its impact on the electro-adsorption rate, we must first establish the Eeq vs θ relationship. This information is the equilibrium surface adsorption, which is readily obtainable on well-defined Pt(111) surfaces. Thus, we can use CVs at different scan rates on Pt(111) to access the electro-adsorption rate at different  and . Monitoring how the electro-adsorption rate maps to  and  at different scan rates, we can determine the electro-adsorption rate as a function of coverage and overpotential and use the Butler-Volmer equation to obtain the rate constant or the exchange current density. To our knowledge, the presented method, which approaches the overpotential by considering the deviation between the applied and equilibrium potential as a function of coverage, is the first that can yield the electro-adsorption rate constant without requiring an assumption of an isotherm model. 3.2.2 Derivation of the equations 3.2.2.1 Equilibrium potential at a given coverage For a surface redox reaction (Ox –ad + e ↔ Redad) that follows the Langmuir isotherm, the equilibrium surface coverage of the reduced species (θ) at a given applied potential can be represented as: θ eη = exp (− ) (3.1) 1 − θ kbT where η = E − E0′, E is the applied potential and E0′ is the formal potential. To include the interaction between adsorbates, θ can be represented in a generalized form as: 59 θ eη = exp [− − f(θ)] (3.2) 1 − θ kbT where f(θ) is a function that describes the influences of adsorbate coverage on reaction free energy (e.g. adsorbate interactions). In other words, at a given θ, we can find a corresponding equilibrium potential (Eeq) which satisfies θ eηθ = exp [− − f(θ)] (3.3) 1 − θ kbT ′ ηθ = E 0 eq − E (3.4) kbT θ Eeq = E 0′ − [ln( ) + f(θ)] (3.5) e 1 − θ indicating Eeq is a function of coverage (θ). 3.2.2.2 Derivation of surface redox kinetics If this reaction obeys the first-order rate law, the apparent rate of this reaction can be represented as: kapp(η, θ) = θ kox(η, θ) − (1 − θ)kred(η, θ) (3.6) where kapp , kox , and kred are apparent rate of the redox reaction, the rate of the oxidation reaction, the rate of the reduction reaction, respectively. At equilibrium, kapp(ηθ, θ) = 0 (3.7) 60 kred(ηθ, θ) θ = (3.8) kox(ηθ, θ) 1 − θ The dependencies of kox and kred on η and θ are expressed by a form of the Butler- Volmer equation as (Figure 3.1): eηθ kox( ap ηθ, θ) = k0 exp {α[ + f(θ)]} (3.9) kbT eηθ kred( ap ηθ, θ) = k0 exp {(1 − α)[− − f(θ)]} (3.10) kbT ap where k0 is the apparent rate constant of the redox reaction and α represents the transfer coefficient in the Butler-Volmer model. Figure 3.1 Free energy diagram of a surface adsorbed redox reaction at equilibrium. The free ap energy diagram at equilibrium potential and 𝜃 = 0.5 (left) where 𝑘ox = 𝑘red = 𝑘0 . The free energy diagram at equilibrium (right) where 𝜃𝑘ox = (1 − 𝜃)𝑘red . 𝑘ox and 𝑘red can be ap expressed as a function of 𝑘0 in a form of the Butler-Volmer equation. Under non-equilibrium conditions, we introduce a parameter ξ to describe the driving force at a given adsorbate coverage (Figure 3.2): ξ = E − E = (E − E0′eq )−(E 0′ eq − E ) = η − ηθ (3.11) 61 Figure 3.2 Free energy diagram of a surface adsorbed redox reaction at a given 𝜃. The free energy diagram at equilibrium where the potential is 𝐸eq (left). The free energy diagram at non- equilibrium condition (right) where the potential is 𝐸eq − 𝜉 . 𝑘ox and 𝑘red can be ap expressed as a function of 𝑘0 in a form of the Butler-Volmer equation. Combining equation (3.6) and (3.11), we can obtain kapp = θkox(ηθ + ξ, θ) − (1 − θ)kred(ηθ + ξ, θ) (3.12) The dependencies of kox and kred on η and θ are expressed by a form of the Butler- Volmer equation as: eη kox(η, θ) ap = k0 exp {α[ + f(θ)]} (3.13) kbT eη kred(η, θ) ap = k0 exp {(1 − α)[− − f(θ)]} (3.14) kbT Combining equation (3.11), (3.12), (3.13) and (3.14), we can obtain ap e(ηθ + ξ) e(ηθ + ξ) kapp = θk0 exp {α[ + f(θ) ap ]} − (1 − θ)k0 exp {(1 − α)[− − f(θ)]}(3.15) kbT kbT From equation (3.3), we can get 62 θ −α eηθ ( ) = exp [α [ + f(θ)]] (3.16) 1 − θ kbT θ (1−α) eηθ ( ) = exp [(1 − α) [− − f(θ)]] (3.17) 1 − θ kbT Plugging equation (3.16) and (3.17) into (3.15), θ −α αeξ θ (1−α)ap eξ kapp = k0 {θ ( ) exp ( ) − (1 − θ) ( ) exp [(1 − α) (− )]}(3.17) 1 − θ kbT 1 − θ kbT Rearranging equation (3.17), we can obtain: αeξ (1−α)eξ − kapp = Γ(θ) (ekbT − e kbT ) (3.18) ap Γ(θ) = k θ1−α0 (1 − θ) α (3.19) If  = 0.5, ap eξ eξ kapp = k0 √θ(1 − θ) [exp ( ) − exp (− )] (3.20) 2kbT 2kbT 3.2.2.3 Kinetics in acidic media Considering the Volmer step in acidic media (Pt + H3O + + e– ↔ Pt-Had + H2O), the apparent rate can be represented as: kapp = θ kox − (1 − θ)kred (3.21) where θ, kox, and kred are the coverage of adsorbed hydrogen, the rate of the oxidation reaction, the rate of the reduction reaction, respectively. A schematic of the free energy 63 of this reaction is shown in Figure 3.3. At pH = 0, E = Eeq vs SHE, and θ = 0.5, kox and kred are equal and can be represented as: k = k0ox desaH2O (3.22) k 0red = kadsaH O+ (3.23) 3 where k0 0ads, kdes, aH O+, and aH O are the adsorption and desorption rate constants of 3 2 the Volmer step in acidic media, activities of hydronium and water, respectively. The adsorption and desorption rate constants are defined at the potentials relative to the standard hydrogen electrode (SHE). Since a = a = 1, k0 = k0 = k0H2O H3O+ des ads acid. While at pH > 0, the equilibrium potential is equal to Eeq (V vs RHE) to achieve θ = 0.5 (Figure 3.3). The RHE scale can be converted into the SHE scale by using the following equation: kbT Eeq (V vs RHE) = Eeq + ln (aH O+) (V vs SHE) (3.24) e 3 64 Figure 3.3 Free energy diagram of the Volmer step in acidic media. The free energy diagram at pH = 0 where the potential is 𝐸eq (V vs SHE) (left). 𝑘ox = 𝑘red = 𝑘 0 𝑎𝑐𝑖𝑑 because both of the activity of water and proton are 1. The free energy diagram at pH > 0 where the potential ap ap ap ap is 𝐸eq (V vs RHE) (right). 𝑘ox = 𝑘des𝑎H O = 𝑘red = 𝑘ads𝑎H O+ where 𝑘2 3 des and 𝑘ads can be represented as a function of 𝑘0acid and pH. Si nce k 0 acid is defined at a potential relative to the SHE, the overpotential is equal to Ee q(V vs RHE) − Eeq(V vs SHE). As a result, the kox and kred can be represented as: k = k0 eη ox acidaH O exp {α[ ]} = k 0 a exp [α ln (a +)] (3.25) 2 kbT acid H2O H3O eη kred = k 0 acidaH O+ exp [(1 − α)(− )] = k 0 acidaH O+ exp [−(1 − α) ln (a3 k T 3 H O +)](3.26) 3 b k0acid exp [−(1 − α) ln (a 0 H O+)] and kacid exp [α ln (aH O+)] can be viewed as the 3 3 ap ap apparent adsorption (kads) and desorption rate constant (kdes), respectively. Under non-equilibrium conditions, the dependencies of kox and kred on η (i.e. E − E0′(V vs SHE)) and θ can be expressed by a form of Butler-Volmer equation as: eη k 0ox = kacidaH O exp {α[ + f(θ)]} (3.27)2 kbT 65 eη kred = k 0 acidaH O+ exp {(1 − α)[− − f(θ)]} (3.28)3 kbT θ aH O eη2 θ= exp [− − f(θ)] (3.29) 1 − θ aH O+ k3 bT If  = 0.5, the apparent rate can be expressed as: 0 eξ eξkapp = kacid√aH O+aH O√θ(1 − θ) [exp ( ) − exp (− )] (3.30) 3 2 2kbT 2kbT Therefore, the apparent rate constant in acid can be expressed as: ap k 00 = kacid√aH O+aH O (3.31) 3 2 3.2.2.4 Kinetics in alkaline media Similarly, the free energy diagram of the Volmer step in alkaline media (Pt+ H2O + e– ↔ Pt-H –ad + OH ) is shown in Figure 3.4. At pOH = 0, E = Eeq vs SHE, and θ = 0.5, kox and kred are equal and can be represented as: kox = k 0 desaOH− (3.32) k 0red = kadsaH (3.33) 2O where k0 0ads , kdes , and aOH− are the adsorption and desorption rate constants of the Volmer step in alkaline media, and activity of hydroxide, respectively. The adsorption and desorption rate constants are defined at potentials relative to the SHE. Since aH O =2 a 0 0 0OH− = 1, kdes = kads = kbase. 66 At non-equilibrium conditions, the dependencies of kox and kred on η and θ are expressed as: eη kox = k 0 baseaOH− exp {α[ + f(θ)]} (3.34) kbT 0 eηkred = kbaseaH O exp {(1 − α)[− − f(θ)]} (3.35)2 kbT Figure 3.4 Free energy diagram of the Volmer step in alkaline media. The free energy diagram at pOH = 0 where the potential is 𝐸eq (V vs SHE) (left). 𝑘ox = 𝑘 0 red = 𝑘base because both of the activity of water and hydroxide are 1. The free energy diagram at pOH > ap ap ap ap 0 at equilibrium (right). 𝑘ox = 𝑘 −des𝑎OH = 𝑘red = 𝑘ads𝑎H O where 𝑘des and 𝑘ads can be 2 represented as a function of 𝑘0base and pOH. If  = 0.5, the apparent rate and apparent rate constant can be expressed as: eξ eξ k 0app = kbase√aH OaOH−√θ(1 − θ) [exp ( ) − exp (− )] (3.36) 2 2kbT 2kbT ap k0 = k 0 base√aH OaOH− (3.37) 2 67 3.3 Experimental section 3.3.1 Preparation of Pt(111) working electrode The working electrode was fabricated by the Clavilier method20 from a 99.99% Pt wire (Goodfellow). The geometric surface area of the electrode was 4.582 mm2. Prior to each experiment, the electrode was flame annealed in a propane flame for 15 seconds, followed by cooling down in a 5% hydrogen, 95% argon gas atmosphere and immediately transferred to the electrochemical cell by the protection of a droplet of ultra-pure water, which was saturated by the gas mixture, on top of the electrode surface. 3.3.2 Preparation of Pt(111)/Ni(OH)2 electrode After being flame-annealed and cooled down in the gas mixture, the Pt(111) electrode was immersed in an Ar-saturated 5 mM Ni(NO3)2 solution for 10 s. The resulting electrode was then rinsed by ultrapure water to remove the extra Ni2+ ions, before transferring into the electrochemical cell. The modified layer of Ni(OH)2 could be removed easily by flame annealing of the electrode after the experiment. 3.3.3 Chemicals Perchloric acid (Sigma-Aldrich ACS reagent 70%); Sodium hydroxide (Sigma- Aldrich 99.99% trace metals basis); Sodium perchlorate monohydrate (VWR EMSURE® for analysis, Supelco®); Ni(NO3)2 ·6H2O (Sigma-Aldrich 99.999% trace metals basis). All the solutions were prepared with ultrapure water with a resistivity of 18.2 MΩ cm. 68 3.3.4 Electrochemical measurements The experiments were performed using a BioLogic potentiostat (SP-300) in a three- electrode electrochemical cell at room temperature. A large area coiled Pt wire was used as the counter electrode, and Ag/AgCl in 1 M KCl (0.235 V vs. SHE) was used as the reference electrode. The Pt(111) working electrode was in contact with the electrolyte in the hanging meniscus configuration. The electrolyte was purged with ultra-high-purity Ar (5.0 grade) for at least 5 min before each experiment. An Ar blanket was maintained above the electrolyte during the experiments. The cyclic voltammetry (CV) was conducted with active-feedback compensation at high scan rates. Before each set of scan rate dependent measurements, a CV under quasi- equilibrium conditions at low scan rate (50 mVs-1) was first obtained, and the resistance was subsequently estimated by electrochemical impedance spectroscopy (EIS) measurements. To avoid overcompensation, trials of CV were performed at high scan rates starting from a lower resistance compensation. The compensation was adjusted until the number of current oscillation cycles became two at the beginning of each CV in the double layer region (0.4 ~ 0.45 V vs. RHE). The number of current oscillation cycles was controlled, so that the extent of compensation could be consistent throughout all the measurements. The reference potential was calibrated and converted to RHE scale by using the following equation. ERHE (V) = Eapplied + 0.059 × pH + 0.235 (3.38) 69 3.3.5 Curve fitting The experimental data were fitted by using the following equation to extract the apparent rate constant. ap 1 eξ eξ Log kapp = Log k0 + Log[θ(1 − θ)] + Log [exp ( ) − exp (− )] (3.39) 2 2kbT 2kbT where e, kb, and T are elementary charge, Boltzmann constant, and temperature, ap respectively. kapp, k0 , θ, and ξ represent the apparent rate, apparent rate constant, the coverage of hydrogen, and the adsorption overpotential at a given hydrogen coverage. The parameters of kapp, θ, and ξ were obtained from experiments and T = 298 K was ap used in all curve fittings. The apparent rate constant (k0 ) was adjusted to optimize the fitted results to achieve the highest R2 value. 3.4 Results and discussion 3.4.1 Measurement of rate constant of Had formation on Pt(111) Our experiment starts with the measurement of the Had coverage as a function of applied potential. Figure 3.5a shows the CV of Pt(111) at a scan rate of 0.05 Vs-1 in perchloric acid (pH ~ 1). This voltammogram can be divided into three regions: the Had underpotential deposition (UPD, i.e., the Volmer step), the electrochemical double layer, and the hydroxide electro-adsorption. We focused our attention on the Had UPD region from 0.05 to 0.4 V vs the reversible hydrogen electrode (RHE). Figure 3.5b shows a series of CVs of the Had UPD region at scan rates from 0.8 Vs -1 to 1000 Vs-1. 70 Figure 3.5. Measurements of the electro-adsorption rate. (a) Cyclic voltammogram of Pt(111) in 0.1 M HClO4 + 0.1 M NaClO4 at a scan rate of 0.05 Vs-1. Shaded area shows the hydrogen adsorption region. (b) Cyclic voltammograms of hydrogen adsorption region on Pt(111) at a series of scan rates from 0.8 to 1000 Vs-1 in HClO4. The y axis is the current normalized by the scan rate. (c) The hydrogen coverage (θ) measured as a function of potential and scan rate obtained by integrating the hydrogen adsorption area in HClO4. 240 μC cm-2 was used to represent the charge density of a fully covered Pt(111) surface. (d) Cyclic voltammogram of Pt(111) in 0.1 M NaOH + 0.1 M NaClO4 at a scan rate of 0.05 Vs-1. (e) Cyclic voltammograms of hydrogen adsorption region on Pt(111) at a series of scan rates from 0.1 to 100 Vs-1 in NaOH. The y axis is the current normalized by the scan rate. (f) The hydrogen coverage (θ) measured as a function of potential and scan rate obtained by integrating the hydrogen adsorption area in NaOH. In all measured CVs, we corrected for the Ohmic (iR) drop using active feedback compensation, which has been successfully used in the past to enable ultrafast CV m easurements21. After correcting for the double-layer background current (0.4 to 0.45 V vs RHE), we integrated the area underneath the CV scan in the cathodic direction to obtain the Had surface charge as a function of electrochemical potential. We divided the Had surface charge by the theoretical surface charge density of Pt(111) (240 71 μC/cm2) to obtain the Had coverage (θ). Figure 3.5c shows the Had coverage as a function of applied potential at various scan rates. At higher sweep rates, the curve shifted to more negative potentials, indicating that the kinetics of the Volmer step could not keep up with the scan rate. We applied the same approach to studying the Had coverage in alkaline media (pH ~13, Figures 3.5d-f). Our results confirm negligible differences between the Had binding energies in acid and base (Figure 3.6), consistent with a previous study22. However, we find significant differences between the Had Figure 3.6. Hydrogen binding strength on Pt(111) in acidic and alkaline media. The CVs of Pt(111) in Ar-saturated 0.1 M HClO4 + 0.1 M NaClO4 (blue) and 0.1 M NaOH + 0.1 M NaClO4 (red) at a scan rate of 0.05 Vs-1. electroadsorption kinetics in the two environments. In particular, the Had coverage in al kaline media shows clear signs of irreversibility, i.e., deviation from the equilibrium co verage, at sweep rates as slow as 1 Vs-1 (Figures 3.7a and b), much earlier than the measurements in acidic media. This observation suggests that the Volmer kinetics in base is slower than in acid. 72 Monitoring the Had electro-adsorption rate variations with applied potential and scan rate, we converted the relationship between the applied potential and scan rate to  and θ. To define the equilibrium coverage, Eeq, we used the H isotherm at 0.8 Vs -1 ad in acid and 0.1 Vs-1 in alkaline. The reversibility of the Had electro-adsorption was confirmed by overlaying the isotherms in the adsorption and desorption directions to ensure that they were identical, i.e., reversible (Figures 3.7c and d). In this analysis, the electro- adsorption rate is the electro-adsorption current measured using CV at specific  and  values, normalized to the theoretical saturated surface charge of Pt(111). This Figure 3.7. The hydrogen coverage-potential curves obtained from a CV at a scan rate of 100 Vs-1 (solid line) and the quasi-equilibrium adsorption isotherm (dashed line) in (a) Ar- saturated 0.1 M HClO4 + 0.1 M NaClO4 and (b) 0.1 M NaOH + 0.1 M NaClO4. The hydrogen coverage-potential curves obtained from the forward direction (solid line) and the backward direction (dashed line) in (c) CV at a scan rate of 0.8 Vs-1 in 0.1 M HClO4 + 0.1 M NaClO4 and (d) CV at a scan rate of 0.1 Vs-1 in 0.1 M NaOH + 0.1 M NaClO4. 73 quantity, which we call the apparent electro-adsorption rate, kapp, is used to represent the apparent electro-adsorption rate per site, i.e., the turnover frequency of the Had electro-adsorption. The quantification of the electro-adsorption rate as a function of driving force and coverage, allows us to quantify the intrinsic rate constant of the Had electro-adsorption. In the analysis, we assume a linearized energy landscape, i.e., the Butler-Volmer equation with a symmetric transfer coefficient,  = 0.5. Fitting our experimental data ap to equation (3.20), allows us to obtain the apparent rate constant, k0 , of the Had electro-adsorption kinetics. Fitting the apparent electro-adsorption rate to Equation (3.20) shows that the apparent rate constant of the Had electro-adsorption in acid (HClO4) is two orders of magnitude larger than in base (NaOH, Figure 3.8). In NaOH, the apparent rate constant of Had electro-adsorption extracted from  ~ 0.1 is 9.0 ± 1.1 (s-1). This value is close to the apparent rate constant of the Had electro-adsorption in alkaline media obtained by Schouten et al. using electrochemical impedance spectroscopy (6.4 s-1). Strikingly, in HClO4, the rate constant of the Volmer step extracted from the same coverage is 1057 ± 314 (s-1). Converting this turnover frequency to an exchange current density, i0, we find that the Volmer step has an i0 value of 253± 75 mA/cm 2 in acid and 2.2 ± 0.3 mA/cm2 in alkaline, respectively. Interestingly, the exchange current densities for the HER and its reverse reaction, hydrogen oxidation reaction (HOR) on the Pt/C catalyst at 313 K have been reported to be 216 ± 50 mA/cm2 in Nafion® and 1.0 ± 0.1 mA/cm2 in NaOH23 (0.69 ± 0.03 mA/cm2 in 0.1 M KOH at 294 K)24. The finding that the rate 74 Figure 3.8. Apparent rate constant of hydrogen electro-adsorption in acidic vs in alkaline media at various hydrogen coverages (θ). (a), (c), & (e) Dependency of the apparent rate on the driving force in acidic medium (blue, 0.1 M HClO4 + 0.1 M NaClO4) and in alkaline medium (red, 0.1 M NaOH + 0.1 M NaClO4) at θ = 0.05, 0.10, and 0.20, respectively. Hollow symbols represent three independent measurements; the solid lines are fitted curves by using Equation (3.20). (b), (d), & (f) Apparent rate constant of hydrogen electro-adsorption in acidic (blue) and in alkaline (red) media at θ = 0.05, 0.10, and 0.20, respectively. co nstants for the HER/HOR are comparable points to the critical role of the Volmer ki netics in the HER/HOR processes and that one should not assume that the Volmer 75 step is much faster than the HER/HOR. We further note that the Butler-Volmer equation fits better to the experiment at low overpotentials due to its nature as a linearized energy landscape approximation, which works most effectively at low overpotentials. Our data in the higher overpotential region is likely useful for future works in the creation of a quantum-mechanical model to capture the curvature of the potential energy surfaces, which shows up more prominently in the high overpotential region. 3.4.2 Mechanism of Had formation in acid vs in base To understand why the rate constant of the Volmer step is larger in acid than in alkaline, we examined the effects of pH (reaction order analysis). Following the same procedures as before, we measured the rate constants of the Volmer step in 0.02, 0.1, and 0.2 M HClO4. Each electrolyte was supplemented with NaClO4 to maintain the total ionic strength at 0.2 M (Figure 3.9). Figure 3.11a shows the effects of pH on the apparent rate constant of the Had electro-adsorption in acid. At a Had surface coverage of 0.1, the apparent rate constants were 507 ± 155, 1057 ± 314, and 1947 ± 494 (s-1), respectively. Regardless of the coverage, the dependency of the apparent rate constant on [H+] was ~0.5. This observation agrees well with the first-order PCET kinetics with protons or hydronium ions as a reactant (Pt + H O+3 + e – ↔ Pt-Had + H2O). The derivation of the first-order kinetics yields the following apparent rate constant for the Volmer step (see Chapter 3.2.2.3 for derivation): ap k0 = k 0 acid√aH O+aH (3.31) 3 2O 76 Figure 3.9. Dependency of the adsorption rate on the driving force in acidic solutions. The keff- curves obtained from CVs at various scan rates in 0.02 M, 0.1 M, and 0.2 M HClO4 at hydrogen coverages equal to 0.05 (left panel), 0.1 (middle panel), and 0.2 (right panel). Hollow symbols represent three independent measurements; the solid lines are fitted curves by using Butler Volmer equations. NaClO4 was added to maintain a 0.2 M ionic strength. w here k 0 acid, aH O+ , and aH O are the adsorption/desorption rate constant of the 3 2 V olmer step in acid, and hydronium and water activities. The observed 0.5 reaction or der agrees well with this prediction. A similar observation has been reported in a number of PCET systems16,25,26 when hydronium/water acts as the proton do nor/acceptor. This observation thus suggests that the Volmer step, in acid, follows fir st-order PCET kinetics that involves the transfer of one proton per electron. 77 Figure 3.10. Dependency of adsorption rate on the driving force in alkaline solutions. The keff- curves obtained from CVs at various scan rates in 0.02 M, 0.1 M, and 0.2 M NaOH at hydrogen coverages equal to 0.05 (left panel), 0.1 (middle panel), and 0.2 (right panel). Hollow symbols represent three independent measurements; the solid lines are fitted curves by using Butler Volmer equations. NaClO4 was added to maintain a 0.2 M ionic strength. Interestingly, unlike in HClO4, the apparent rate constant was independent of pH in N aOH solutions (zero reaction order). We measured the apparent rate constant in 0.02, 0.1, and 0.2 M NaOH, all supplemented with NaClO4 to maintain a constant ionic strength of 0.2 M (Figure 3.10). At all coverages studied, the apparent rate constant w as pH-independent (Figure 3.11b). This observation is striking given that the Had electro-adsorption (Pt+ H2O + e – ↔ Pt-Had + OH –) in alkaline should be pH dependent and have an apparent rate constant of the form: 78 ap k0 = k 0 base√aH a − (3.37) 2O OH where k0base and aOH− are the adsorption/desorption rate constants of the Volmer step in alkaline and the hydroxide activity. The observed zero reaction order implies that the adsorption/desorption rate constant (k0base) is pH-dependent and cancels the aOH− factor in alkaline media. Figure 3.11. pH effect on the apparent rate constant of hydrogen electro-adsorption at various hydrogen coverages (θ). (a) Apparent rate constant in acidic media correlates with the acid concentration. (b) The apparent rate constant in alkaline media is independent of the base concentration. Gray lines are linear regressions of the data. Solid square: θ = 0.05; solid circle: θ = 0.10; solid triangle: θ = 0.20. 79 Rossmeisl et al. proposed that pH can impact the activation energy barrier due to the configurational entropy27. While this analysis explains the deceleration of the Volmer step in alkaline compared to acid, the observation that pH affects the reaction order in acid but not in alkaline is inconsistent with the entropy model, which predicts the adsorption/desorption rate to depend on pH in the same way. We thus believe that there are other factors governing the Volmer step beyond configurational entropy. One possible explanation is the rigid interfacial water layer in alkaline, which has been suggested by Ledezma-Yanez et al. based on their laser-induced temperature-jump study28. In alkaline, the electrochemical potential where the HOR/HER occurs (near - 0.8 V vs. standard hydrogen electrode, “SHE”) is far from the point of zero free charge of Pt(111) (~0.3V vs. SHE)29. The presence of this potential gap polarizes the interfacial water molecules, making their re-organization the rate-limiting step in the Volmer step. We can estimate how pH impacts the HOR through the interfacial water by using the model of Lamoureux et al.30, whose analysis estimated that a 25 mV shift in the interfacial electric field could impact the turnover frequency by a factor of 1.4. Since a 10x change in the OH– concentration would correspond to -59 mV shift in the interfacial electric field, the model by Lamoureux et al. suggests that going from 0.02 M to 0.2 M OH– (-59 mV) should decrease in the turnover frequency by a factor of 3. If this 3-fold turnover frequency change translates linearly to k0base, it would offset the ap aOH− increase in Equation (3.37), effectively causing k0 to be pH-independent as observed. Taken together, in proton-rich acid, the factor controlling the Had electro-adsorption rate is the intrinsic PCET kinetics, which depends on molecular details such as the 80 reactant and product states, electronic coupling, overlap integral, and reorganizational energy31. In proton-poor alkaline media, the Had electro-adsorption relies on interfacial water as the proton donor and the reorganization of the strongly polarized interfacial water layer is the rate-limiting step. 3.4.3 Effect of surface modifiers on Had formation rate The mechanistic difference suggests a more flexible interfacial water structure may assist the Had formation in alkaline media. One way of modifying the interfacial water layer is to tune the potential of zero free charge by introducing a surface modifier, e.g., nickel in form of Ni(OH) 282 . The CV at quasi-equilibrium adsorption state on Pt(111) with Ni(OH)2 modifiers is compared with Pt(111) bare electrode in Figure 3.12. The new peaks at around 0.55 V and 0.7 V are associated to the adsorption and desorption of OH on Pt sites near the nickel clusters.32 The coverage of Ni(OH)2, θNi, was estimated to be 0.09, which was calculated by the normalized difference of the charge Figure 3.12. Cyclic Voltammogram of Pt(111)/Ni(OH)2 (green) and Pt(111) (orange) in Ar-saturated 0.1 M NaOH at a scan rate of 0.05 Vs-1. 81 under Hupd region between a bare Pt(111) electrode and a Pt(111)/Ni(OH)2 electrode.28,32–34 Applying the developed CV methodology to Pt(111)/Ni(OH)2 in alkaline media, we found that the inclusion of nickel on Pt(111) indeed promotes the Figure 3.13. Apparent rate constant of hydrogen electro-adsorption of Pt(111)/Ni(OH)2 vs Pt(111) in alkaline medium at various hydrogen coverages (θ). (a), (c), & (e) Relationship between driving force and the apparent rate of hydrogen electro-adsorption on Pt(111)/Ni(OH)2 (green) and Pt(111) (orange) in alkaline medium (0.1 M NaOH) at θ = 0.05, 0.10, and 0.20, respectively. Hollow symbols represent three independent measurements; the solid lines are fitted curves by using Equation (3.20). (b), (d), & (f) The apparent rate constant of hydrogen electro-adsorption on Pt(111)/Ni(OH)2 (green) and Pt(111) (orange) at θ = 0.05, 0.10, and 0.20, respectively. 82 Had electro-adsorption kinetics to 25.4 ± 2.8 (s -1) at  ~ 0.1 (~3 fold increase vs. 8.0 ± 1.7 (s-1) on unmodified Pt(111), Figure 3.13). Our observation thus confirms a long- held hypothesis that the interfacial water layer plays a crucial role in the Had electro- adsorption kinetics and demonstrates how the inclusion of a surface adsorbate modifier can modify this kinetics parameter beyond the Had binding strength. 3.5 Conclusion Our measurement of the Had formation kinetics and the impact of the interfacial water layer at different pH carries an important message that one needs to consider the molecular dynamics of the local environment in electrocatalysis, in addition to the intermediate’s binding strength (e.g., of Had). Within the larger picture of heterogeneous catalysis, this work demonstrates a limitation on the use of time- independent thermodynamic variables, such as Had binding strength, to approximate time-dependent, kinetics phenomena. We expect that the methodology presented in this work will provide a pathway to extract the rate constants of the elementary steps in multi-electron electrochemical processes and transformations. The developed methodology is applicable to other electro-adsorption reactions and to non-metal surfaces, such as oxides. We believe that this development will open new doors in how we can approach electrocatalysis by engineering the rate constants, aside from the binding energy, in an effort to uncover new catalyst-improvement strategies for future and emerging technologies. 83 3.6 References (1) Turner, J. A. Sustainable Hydrogen Production. Science (80-. ). 2004, 305, 972– 975. (2) Seh, Z. W.; Kibsgaard, J.; Dickens, C. F.; Chorkendorff, I.; Nørskov, J. K.; Jaramillo, T. F. Combining Theory and Experiment in Electrocatalysis: Insights into Materials Design. Science (80-. ). 2017, 355, 6321. (3) Stamenkovic, V. R.; Strmcnik, D.; Lopes, P. P.; Marković, N. M. 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Effect of the Interfacial Water Structure on the Hydrogen Evolution Reaction on Pt (111) Modified with Different Nickel Hydroxide Coverages in Alkaline Media. ACS Appl. Mater. Interfaces 2019, 11, 613–623. 89 CHAPTER 4 CATION AND ISOTOPE EFFECTS ON HYDROGEN ADSORPTION KINETICS ON Pt(111) IN ALKALINE MEDIA 4.1 Introduction The kinetics of electro-adsorption of hydrogen is a topic of both fundamental and practical importance. As it is an elementary step of the hydrogen evolution reaction, the study of its kinetics not only helps us understand how the electrocatalytic process takes place, but also guides us to enhance the catalytic activity. We have shown in Chapter 3 that the approximation of reaction kinetics using intermediate energies is not adequate. The apparent rate constant of hydrogen adsorption on Pt(111) was found to be two orders of magnitude lower in alkaline media than in acidic media although the H binding energy is the same in both media. The reason behind the significantly lower H adsorption kinetics on Pt(111) in alkaline media is a follow up question that has given rise to extensive discussions. One of the most widely accepted explanations is that the electric field in the double layer can influence the orientation of the interfacial water as reactive species, and thus the reaction rate.1,2 The interfacial electric field depends on the potential of zero free charge (pzfc). If the H adsorption potential is close to the pzfc, the electric field would be weak and the interfacial water would be easy to reorient, so that the charge transfer through the double layer would be rapid, and vice versa. For Pt in alkaline media, the pzfc is 1.054 V vs. RHE (pH=13.1), measured by CO displacement, which is much more positive to the H adsorption potential at 0~0.4 V vs. RHE, relative to the pzfc in acidic media (0.343 V vs. RHE pH=1.2).3 As a result, the interfacial water would be rigid at the H adsorption 90 potential in alkaline media, which leads to sluggish kinetics. This hypothesis emphasizes the role of the interfacial water structure in H adsorption process in alkaline media. To further study the mechanism in alkaline media, in this chapter, I tried to perturb the interfacial water structure by using different cations in the electrolyte and isotopically substituted water (D2O), so that we could see how the reaction kinetics change with the change of interfacial water structure in alkaline media. The non-covalent interactions of cations with OH adsorbates and their influence on double layer structure and the rates of the electrocatalytic reactions, such as ORR, HOR and methanol oxidation, have been studied by Marković et al.45 They found that cations with higher hydration energies tend to stabilize the OHad by non-covalent interactions and thus decelerate the reactions with OHad as an intermediate. In another study in acidic media by Kumeda et al.6, similar trends in specific activity, with respect to the hydration energy of the added tetra-alkylammonium cations, were observed. They were able to enhance the ORR activity of single crystal Pt electrodes by adding hydrophobic cations to the electrolyte which destabilized the adsorbed OH and water at the electrode/electrolyte interface. On the other hand, H/D isotope effects can also help unravel the reaction mechanism. The kinetic isotope effect (KIE), i.e., the ratio of the rate with hydrogen to the rate with deuterium, describes the differences in activation energy between reactions with different isotopes. Since the vibrational frequencies of H-bond and D- bond are very different and the mass ratio of H/D is 1:2, differences are expected in 91 their zero-point energies, which generally results in a KIE >1, for the reactions involving Had as an intermediate. In addition, if quantum proton tunneling is involved in the process, an even larger KIE (>13) could be observed because heavier isotopes can only tunnel through lower barriers.7 Previously, Sakaushi studied the quantum proton tunneling by KIE in HER and ORR processes on polycrystalline Pt and polycrystalline Au in fully deuterated electrolytes.7,8 However, challenges still remain for experiments with more well-defined surfaces because single crystal electrodes are very sensitive to impurities in the electrolyte. Reliability of measured KIE values depends on the purity of D2O. More recently, the KIE of the HOR on Pt(111) in base was reported to be 3.4 at an overpotential of 350 mV by Rebollar et al.9 They proposed that the large KIE is due to the influence of isotopes outside of the reaction center, D2O, on the vibrational frequencies of the H bonds, which supports that interfacial water plays an important role in HOR kinetics in alkaline media. In this chapter, I investigated the cation effect and the isotope effect on the H adsorption kinetics (Volmer step of HOR) on Pt(111) electrodes. Fast scan rate CV (same method as in Chapter 3) was used to directly measure the rate constant of the electro-adsorption process in alkaline solutions with cations of different size, Li+, Na+, K+ and TMA+, or in NaOH with ultrapure D2O as the solvent. We found that the rate constant of H adsorption on Pt(111) is highest in 0.1 M LiOH solution, among the four solutions studied, which agrees with the mechanism previously proposed that involves cation-water clusters in the reaction process10. A KIE of 3.4 was obtained in 0.1 M NaOH/D2O. The possible contributions of interfacial water to the KIE are discussed. 92 4.2 Experimental section 4.2.1 Preparation of Pt(111) electrode The Pt(111) working electrode was the same as the one used in Chapter 3, which was fabricated by the Clavilier method and annealed in a propane flame before each set of measurement. (Details can be found in Chapter 3.3.1) 4.2.2 Chemicals. Sodium hydroxide 99.99% Suprapur® (Supelco); lithium hydroxide 99.995% trace metals basis (Sigma-Aldrich); potassium hydroxide 99.995% Suprapur® (Supelco); Tetramethylammonium hydroxide (TMAOH) solution 25 wt. % in H2O (Sigma- Aldrich); Sodium perchlorate monohydrate (VWR EMSURE® for analysis, Supelco®); deuterium oxide 99.9 atom % D (Sigma-Aldrich). All the solutions in H2O were prepared with ultrapure water with a resistivity of 18.2 MΩ cm. 4.2.3 Purification of D2O. The D2O (Sigma, 99.9 atom % D) was first filtered through an ion-exchange resin and then distilled with alkaline KMnO4 to remove ionic and organic impurities. The resulting D2O was then distilled twice to get rid of the KMnO4 residues, before being used in the isotope experiments. The final product was ~95 atom % D, determined via mass spectroscopy and conductivity grade (~18 MOhm cm). 4.2.4 Electrochemical measurements. The experiments were performed using the same method as in Chapter 3. In a typical measurement in a three-electrode electrochemical cell, a large area coiled Pt wire was 93 used as the counter electrode, and Ag/AgCl in 1 M KCl (0.235 V vs. SHE) was used as the reference electrode. The Pt(111) working electrode was in contact with the electrolyte in the hanging meniscus configuration. The electrolyte (~20 mL) was purged with ultra-high-purity Ar (5.0 grade) for at least 5 min before each experiment. An Ar blanket was maintained above the electrolyte during the experiments. Cyclic voltammetry (CV) was conducted with active-feedback compensation at high scan rates with a BioLogic potentiostat (SP-300). Before each set of scan rate dependent measurements, a CV under quasi-equilibrium conditions at low scan rate (50 mVs-1) was first obtained, and the resistance was subsequently estimated by electrochemical impedance spectroscopy (EIS) measurements. To avoid overcompensation, trials of CV were performed at high scan rates starting from a lower resistance compensation. The compensation was adjusted until the number of current oscillation cycles became two at the beginning of each CV in the double layer region (0.4 ~ 0.45 V vs. RHE). The number of current oscillation cycles was controlled, so that the extent of compensation could be consistent throughout all the measurements. The reference potential for the experiments in H2O was calibrated and converted to RHE scale by using the following equation: ERHE (V) = Eapplied + 0.059 × pH + 0.235 (4.1) The reference potential for the experiments in D2O was calibrated by RDeE (reversible deuteron electrode) made by electrolysis of D2O at corresponding pH conditions. In acidic media, Ag/AgCl vs. RHE is the same as Ag/AgCl vs. RDeE. However, in alkaline media, there was a measured difference of ~50 mV because the 94 dissociation equilibrium constant (K +diss = [D ] [OD -]) for D2O is 10 -14.87, instead of 10- 14 (Kw). As a result, pD = 13.87 in 0.1 M NaOD, theoretically. 4.2.5 Curve fitting. The experimental data were fitted by using the following equation (derived in Chapter 3.2) to extract the apparent rate constant. ap 1 eξ eξ Log kapp = Log k0 + Log[θ(1 − θ)] + Log [exp ( ) − exp (− )] (4.2) 2 2kbT 2kbT where e, kb, and T are elementary charge, Boltzmann constant, and temperature, ap respectively. kapp, k0 , θ, and ξ represent the apparent rate, apparent rate constant, the coverage of hydrogen, and the adsorption overpotential at a given hydrogen coverage. The parameters of kapp, θ, and ξ were obtained from experiments and T = 298 K was ap used in all curve fittings. The apparent rate constant (k0 ) was adjusted to optimize the fitted results to achieve the highest R2 value. 4.3 Results and discussion 4.3.1 Cation effects Alkaline solutions with four different cations, Li+, Na+, K+ and TMA+, were used. The cyclic voltammograms of Pt(111) at quasi-equilibrium state in 0.1 M alkaline solutions are shown in Figure 4.1. The H adsorption regions are almost the same for the four different alkaline electrolytes, indicating that the H adsorption reaction is thermodynamically similar with the different cations. Interestingly, the OH adsorption 95 Figure 4.1. Cyclic voltammograms of Pt(111) in 0.1 M LiOH, NaOH, KOH and TMAOH at a scan rate of 50 mV/s. regions are quite different. The OH adsorption on Pt(111) in LiOH split into two peaks. The one at lower potential is larger and narrower, and the other is smaller and broader. The total charge density under the two peaks in LiOH (156 µC/cm2) is the same as the charge under the OH adsorption peak in NaOH. This indicates that there is the same amount, but two different kinds of OHad on the Pt(111) surface. The lower peak potential of the higher peak (0.72 V), compared to the OH adsorption peaks with other cations (0.78 V), suggests that a large portion of the OHad on Pt(111) in LiOH is stabilized by Li+ and the interfacial water structure might be changed because of the non-covalent interactions between Li+ and OHad. In addition, there is a small peak at ~0.56 V in the double layer region in the CVs in KOH and TMAOH. We ascribe it to the impurities in the solution that were introduced by the chemicals used. The NaOH solution was the purest among the four, since no impurity peak was found in the CV. 96 Figure 4.2. (a) Cyclic voltammograms of the hydrogen adsorption region on Pt(111) at a series of scan rates from 50 mV/s to 100 V/s in 0.1 M LiOH. The y axis is the current normalized by the scan rate. (b) The hydrogen coverage (θ) measured as a function of potential at different scan rates obtained by integrating the hydrogen adsorption area in (a). 240 μC cm-2 was used to represent the charge density of a fully covered Pt(111) surface. Scan rate dependent CVs from 100 V/s to 50 mV/s were employed focusing on the H adsorption region on Pt(111) in the four different electrolytes. Figure 4.2 shows an example of the CVs obtained in 0.1 M LiOH and the integrated coverage vs. potential plot. The CV at 50 mV/s was used as the quasi-equilibrium state reference (ξ = 0). The adsorption overpotential, ξ, at θ = 0.1 was obtained from Figure 4.2b. The corresponding apparent rates, kapp, at different adsorption overpotentials in four different electrolytes are plotted in Figure 4.3a. The fitting of the points by equation 4.2 gives the apparent rate constants. For each rate constant, three replicate experiments were conducted. The average values and the error bars for the four different cation containing electrolytes are presented in Figure 4.3c. The apparent rate constant of H adsorption in LiOH is 5.3 s-1, much higher (~2x) than in the other three electrolytes (2.7 s-1 in 0.1 M NaOH), which can also be qualitatively observed from 97 Figure 4.3. (a) Dependency of the apparent rate on the driving force in 0.1 M LiOH, NaOH, KOH and TMAOH at θ = 0.10. Lines are the fits to the data points. (b) Dependency of the apparent rate on the driving force in 0.1 M NaOH and 0.1 M NaOH + 0.1 M NaClO4 at θ = 0.10. Lines are the fits to the data points. (c) Apparent rate constant of hydrogen adsorption in 0.1 M LiOH, NaOH, KOH and TMAOH at θ = 0.10. Error bars were estimated from the standard deviation of three replicate measurements. (d) Apparent rate constant of hydrogen adsorption in 0.1 M NaOH and 0.1 M NaOH + 0.1 M NaClO4 at θ = 0.10. Figure 4.3a. The observation of faster H adsorption kinetics in LiOH seems to be opposite to the results from Marković et al. who found that the trend of the activity of ORR, HOR and methanol oxidation followed Li+ <