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Design and modification of electrocatalysts for use in fuel cells and CO₂ reduction
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Design and modification of electrocatalysts for use in fuel cells and CO₂ reduction
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Design and Modification of Electrocatalysts for Use in Fuel Cells and CO2 Reduction by Vicente Galvan A dissertation Presented to the FACULTY OF THE USC GRADUATE SCHOOL UNIVERSITY OF SOUTHERN CALIFORNIA In Partial Fulfillment of the Requirements for the Degree DOCTOR OF PHILOSOPHY (CHEMISTRY) August 2022 Copyright 2022 Vicente Galvan ii Dedicated to My family, friends and loved ones iii Acknowledgements Looking back at when I first started my Ph.D. journey this day appeared so far, it almost seemed impossible. Yet I have been fortunate to be surrounded and work alongside many wonderful and exemplary people who have made this once farfetched idea, possible. First, I would like to express my sincerest gratitude to my Ph.D. advisor Prof. G. K. Surya Prakash for his mentorship, guidance, and support not only in my lab work but also in life. I have really enjoyed working under his tutelage; his “hands-off” approach has provided me the freedom of exploring my scientific interests and giving me a more profound appreciation for chemistry. I would also like to thank Prof. Sri Narayan who was always available to help me in questions with regards to electrochemistry and allowing me to join his group meetings to further broaden my knowledge of electrochemistry. Every discussion we had was always a learning experience and fostered even more excitement towards electrochemistry. I would also like to thank my committee members Prof. Nicos Petasis, Prof. Ralf Haiges, Prof. Barry Thompson, Prof. Kathryn Shing and Prof. Jayakanth Ravichandran, for their dedication and generosity with their time in helping, advising, and guiding me during the screening, qualifying and dissertation defense process. I have been privileged to have been a part of a research group full of fantastic, intelligent, talented, and friendly people who have made my graduate experience enjoyable. I would first like to thank the people in my immediate group in the fuel cell/electrochemistry lab. Dr. Marc Iulliucci and Dr. Dean E. Glass for being my mentors when I first joined the group and providing a conducive, learning and positive atmosphere. Dr. Bo Yang, who was always a constant in the lab iv throughout my time, providing insight and helping in my research and issues with instrumentation. Dr. Eugene Kong, a collaborator from chemical engineering, who helped in improving my knowledge of fuel cells and electrochemistry and was always happy to talk about soccer with me. Dr. Amanda Baxter, thank you for asking fundamental electrochemistry and surface analysis questions and helping me learn more. Thank you for allowing me to mentor you during the latter years of your time at USC. Mr. Adam Ung of the Hogen-Esch group, thank you for your help in fundamental questions about all chemistry and for making me think deeper about my results. I would also like to thank Mr. Juan Pablo de los Rios for all the help and your cooperation these past years as the future of the fuel cell/electrochemistry lab. It was a great pleasure working alongside you and thank you for allowing me to accompany you to the gym on a regular basis. It has been a pleasure watching you grow and become an expert in electrochemistry. Thank you all for your friendships, which have made my time in the lab much more enjoyable and thank you for helping me grow as a scientist and person. I will cherish our moments, in and outside of the lab, together; I know I have made lifelong friends. I would like to thank the rest of the past and present members of the Prakash group for providing me a friendly and collaborative atmosphere. Special thanks to Archith, Alex, Aisha, Antonio, CJ, Colby, Daniel, Jothi, Kavita, Matt, Nazanin, Raktim, Sahar, Sayan, Sankar, Vinayik, Xanath, and Ziyue who made my time in the lab memorable and joyful, thank you for your companionship over the years. I would also like to thank the senior scientists of the Prakash group, late Dr. Golam Rasul, Dr. Alain Goeppert, Dr. Thomas Matthew, and Dr. Patrice Batamack for their willingness to help and guide me through my projects. It was a pleasure to learn from such brilliant and enthusiastic scientists. Along with the Prakash group I would like to thank the members of the Narayan lab who opened their lab to me and made it feel like I was also a part of v the group. I would like to thank Dr. Dan Fang, Dr. Buddhine Jayathilake, Dr. Lena Hoober- Burkhardt, Dr. Debanjan Mitra, Dr. Irshad Ahamed, Dr. Sundar, Advaith, Bilal, Rodrigo, and Sarah for your help with electrochemical instrumentation, discussion about research and your friendship. I would also like to thank Prof. Chulsung Bae from Rensselaer Polytechnic Institute for providing anion exchange membranes and binders for several projects. Furthermore, I am grateful to the staff of The Loker Hydrocarbon Institute and Chemistry Department; Dr. Robert Aniszfeld, Jessie May, David Hunter, Carole Phillips, Michele Dea and Magnolia Benitez for happily helping me through all the paperwork and making sure things are taken care of so I could focus on my research. Special thanks to Jessie May and Gloria Canada; Jessie who is always happy to talk about anything and whose joyful and bubbly personality brings fun and joy to the Institute and Gloria who was always a joy to talk to even on those very late nights and who worked tirelessly to keep the Institute presentable and clean. I would also like to thank people from my time as an undergraduate for providing the building blocks and the opportunity to explore research in electrochemistry. I would especially like to thank my undergraduate research advisor Dr. Frank A. Gomez for allowing me to do research in his group and exposing me to the realm of laboratory research, electrochemistry, and fuel cells. He always believed in me and pushed me to do the best that I could, I am grateful for all the help you have provided. I would also like to thank Dr. John Haan from California State University, Fullerton for his collaborations, discussions about electrochemistry and fuel cells during my time as an undergraduate and time during graduate school. On a more personal note, I would like to express my gratitude to my partner Shanna McCue for being an endless source of support and inspiration throughout my time at USC. Coming home to your presence always made even the toughest days in lab brighter. I would also like to thank my vi childhood best friends, particularly Francisco Alvarez, Diego Avalos, Cristian Carrillo, Ivan Garcia, Luis Garcia, Diego Godinez, Jose Gamboa, Jesus Laguna, and Javier Roa. Even though I see you all less than desired, you are all always supportive and quick to offer a helping hand. I would also like to thank my family who have been there with me for the entirety of my life, encouraging, supporting, and believing in me throughout my education. Vicente Galvan, Angela Galvan, Claudia Galvan, Esmeralda Galvan, Ana Galvan, Veronica Macedo and Imelda Macedo; without your love, encouragement, and support this would not have been possible. To my parents, Vicente and Angela Galvan, a formal education was something not readily available to both. Nonetheless, they ingrained the importance of a formal education into me and my sisters, they always support us, and remind us to focus on school and push the boundaries. This dissertation is especially dedicated to the both of you. Thank you! Electrochemistry rocks! vii TABLE OF CONTENTS Acknowledgements ........................................................................................................................ iii List of Tables ................................................................................................................................. ix List of Figures ................................................................................................................................. x Abstract ....................................................................................................................................... xvii Chapter 1: Introduction ................................................................................................................... 1 1.1. The Climate Crisis and Clean Energy .................................................................................. 1 1.2. Fuel Cells.............................................................................................................................. 5 1.3. Workings of a Fuel Cell ...................................................................................................... 6 1.4. Fuel Cell Types .................................................................................................................... 8 1.4.1. Proton Exchange Membrane Fuel Cell .......................................................................... 9 1.4.2. Alkaline Fuel Cells ...................................................................................................... 11 1.4.3. High Temperature Fuel Cells ...................................................................................... 12 1.5. Fuel Cell Applications ........................................................................................................ 14 1.6. CO2 Reduction.................................................................................................................... 16 1.7. Scope of Dissertation ......................................................................................................... 18 1.8. References .......................................................................................................................... 19 Chapter 2: Partially Fluorinated Carbon Supported Catalysts for Improved Proton and Anion Exchange Membrane Fuel Cells ................................................................................................... 29 2.1. Fluorinated Carbon Supports ............................................................................................ 29 2.2. Platinum Supported on Partially Fluorinated Carbons for Improved Performance and Stability in H2 Proton Exchange Membrane Fuel Cells ............................................................ 31 2.2.1. Experimental ................................................................................................................ 31 2.2.2. Results and Discussion ................................................................................................ 35 2.3. Development of Partially Fluorinated Carbon Supported Palladium Nanoparticles for Alkaline Direct Liquid Fuel Cells ............................................................................................. 58 2.3.1. Experimental ................................................................................................................ 58 2.3.2. Results and Discussion ................................................................................................ 61 2.4. Conclusion .......................................................................................................................... 78 2.5. References ......................................................................................................................... 79 Chapter 3: Development and Improvements in Direct Formate Fuel Cells .................................. 90 3.1. Direct Formate Fuel Cells .................................................................................................. 90 3.2. Reduced Graphene Oxide Supported Palladium Nanoparticles for Enhanced Electrocatalytic Activity Towards Formate Electrooxidation in an Alkaline Medium ............ 93 3.2.1. Experimental ................................................................................................................ 93 3.2.2. Results and Discussion ................................................................................................ 96 viii 3.3. Improving Alkaline Direct Formate Fuel Cell Performance with PdIrO 2/MWCNT Electrocatalysts........................................................................................................................ 107 3.3.1. Experimental .............................................................................................................. 107 3.3.2. Results and Discussion .............................................................................................. 110 3.4. Studies on the Simultaneous Production of Alkali Hydroxide and Electricity with a Direct Formate Fuel Cell .................................................................................................................... 126 3.4.1. Experimental .............................................................................................................. 126 3.4.2. Results and Discussion .............................................................................................. 127 3.5. Conclusion ........................................................................................................................ 136 3.6. References ........................................................................................................................ 137 Chapter 4: Improvements Towards Alkaline Direct Methanol Fuel Cells from Polymeric Materials to Catalyst ................................................................................................................................... 152 4.1. Alkaline Direct Methanol Fuel Cells ............................................................................... 152 4.2. Ionomer Significance in Alkaline Direct Methanol Fuel Cell to Achieve High Power with a Quarternized poly(terphenylene) Membrane ....................................................................... 156 4.2.1. Experimental .............................................................................................................. 156 4.2.2. Results and Discussion .............................................................................................. 159 4.3. Development of Nickel Gallium Electrocatalysts for Methanol Electrooxidation and Platinum Free Alkaline Direct Methanol Fuel Cells ............................................................... 174 4.3.1. Experimental .............................................................................................................. 174 4.3.2. Results and Discussion .............................................................................................. 178 4.4. Improving Noble Metal Free Alkaline Direct Methanol Fuel Cells with CeO2/C Supported Ni electrocatalyst ..................................................................................................................... 190 4.4.1. Experimental .............................................................................................................. 190 4.4.2. Results and Discussion .............................................................................................. 193 4.5. Conclusion ........................................................................................................................ 202 4.6. References ........................................................................................................................ 203 Chapter 5: Electrochemical CO2 Reduction ............................................................................... 223 5.1. CO2 Reduction .................................................................................................................. 223 5.2. Selective Electrochemical CO2 Reduction on a Copper Single Atom Catalyst ............... 225 5.2.1. Experimental .............................................................................................................. 225 5.2.2. Results and Discussion .............................................................................................. 229 5.3. Conclusion ........................................................................................................................ 243 5.4. References ........................................................................................................................ 243 ix List of Tables Table 2.1. Summary of BET and Pore Volume for the carbon supports ..................................... 35 Table 2.2. Reactivation, charge transfer and diffusion coefficients for the prepared catalysts with various fuels. ................................................................................................................................. 70 Table 2.3. Obtained activation energies, onset potentials and Tafel slopes for the various fuels with the prepared catalysts. ........................................................................................................... 74 Table 3.1. Summary of the catalysts performance ..................................................................... 105 Table 3.2. Summary of the physical characterization for the prepared catalyst. ....................... 117 Table 3.3. Summary of half-cell results for the prepared catalysts. ........................................... 124 Table 3.4. Summary of half-cell results ..................................................................................... 133 Table 4.1. The ion exchange capacity (IEC) and ion conductivity of the TPN-TMA membrane and the AEIs. ..................................................................................................................................... 160 Table 4.2. Summary of AEI half-cell performance, the uncertainty values were determined from the standard deviation of the repeated experiments. ................................................................... 167 Table 4.3. Comparison of maximum power densities for ADMFC ........................................... 173 Table 4.4. Elemental composition of the catalysts from EDS and XPS. ................................... 184 Table 4.5. Summary of half-cell experiments ............................................................................ 188 Table 4.6. Elemental composition of the catalysts from EDS and XPS. ................................... 197 Table 4.7. Summary of half-cell experiments ............................................................................ 200 Table 5.1. Crystallite and particle sizes obtained from powder XRD and TEM. ...................... 230 Table 5.2. Element weight percentage obtained from EDS and XPS. ....................................... 236 Table 5.3. Products obtained for the prepared catalysts. ............................................................ 240 x List of Figures Figure 1.1. A) Worldwide energy consumption by regions. Adapted from [1] B) Worldwide energy consumption by source. Adapted from [2].......................................................................... 2 Figure 1.2. CO2 levels for the past 800,000 years. Adapted from [3]............................................ 3 Figure 1.3. Electricity generation from various sources. Adapted from [10] ................................ 4 Figure 1.4. Ragone plot of various power sources. Adapted from [11] ......................................... 5 Figure 1.5. Schematic of a fuel cell single stack ............................................................................ 7 Figure 1.6. Polarization and power curve for a fuel cell. Adapted from [14] ................................ 8 Figure 1.7. Schematic of different types of fuel cells and their operating conditions. .................. 9 Figure 1.8. Typical power ranges for various fuel cells. Adapted from [56] ............................... 15 Figure 1.9. Typical products for ECO2R based on electrocatalyst. Adapted from [67] .............. 18 Figure 2.1. A) Mesoporous size distribution curves for XC72R and CFx and B) corresponding N2 adsorption-desorption isotherms of XC72R and CFx. C) TGA curves for the prepared catalysts under an air atmosphere. ............................................................................................................... 35 Figure 2.2. TEM images of the catalysts: A) Pt-1F; B) Pt-1C; C) Pt-2F; D) Pt-2C .................... 37 Figure 2.3. A) XRD pattern of the catalysts on CFx and XC72R supports. B) CV scans of the catalysts in 0.5 M H2SO4 solution at 20 mV/s under Ar. The solid lines indicate the CFx support; the dashed lines indicate the XC-72R support. ............................................................................. 38 Figure 2.4. LSV scans of the catalysts in 0.5 M H2SO4 solution at 5 mV/s: A) of the Pt-2F; B) Pt- 2C. ................................................................................................................................................. 39 Figure 2.5. A) LSV scans before and after stability experiments at scan rate of 5 mV/s at 1600 RPM. The solid lines indicate the CFx support; the dashed lines indicate the XC72R support. B) Change in ECSA after start stop and 3000 cycles. ....................................................................... 40 Figure 2.6. TEM micrographs of A) Pt-2C and B) Pt-2F after start stop stability and C) Pt-2C and D) Pt-2F after 3000 cycles. ........................................................................................................... 44 Figure 2.7. LSV scans of the catalysts in 0.5 M H2SO4 solution at 5 mV/s: A) of the Pt-1F; B) Pt- 1C; C) Pt-2F; and D) Pt-2C. ......................................................................................................... 45 xi Figure 2.8. LSV scans of A) Pt-1and B) Pt-2 in 0.5M H2SO4 under H2 gas before and after stability experiments at scan rate of 5 mV/s at 1600 RPM. The solid lines indicate before stability, the dotted lines indicate after start-stop experiments and the dashed lines indicate after 3000 CV cycles. c) Change in ECSA after start-stop and 3000 cycles. ....................................................... 46 Figure 2.9. SEM and EDS images for A) Pt-2C; B) Pt-2C after 3000 cycles; C) Pt-2C after start- stop; D) Pt-2F; E) Pt-2F after 3000 cycles and F) Pt-2F after start-stop in HOR potential region. For elemental mapping red corresponds to C, green corresponds to oxygen, blue corresponds to fluorine, magenta corresponds to Pt and yellow corresponds to sulfur. ....................................... 49 Figure 2.10. SEM and EDS images for A) Pt-1C; B) Pt-1C after 3000 cycles; C) Pt-1C after start- stop; D) Pt-1F; E) Pt-1F after 3000 cycles and F) Pt-1F after start-stop in HOR potential region. For elemental mapping red corresponds to C, green corresponds to oxygen, blue corresponds to fluorine, magenta corresponds to Pt and yellow corresponds to sulfur. ....................................... 51 Figure 2.11. TEM micrographs of A) Pt-1C and B) Pt-1F after 3000 cycles and C) Pt-1C and D) Pt-1F after start stop experiments. TEM micrographs of E) Pt-2C and F) Pt-2F after 3000 cycles and G) Pt-2C and H) Pt-2F after start-stop experiments............................................................... 54 Figure 2.12. Polarization curves for the Pt-1, Pt-2 anode catalyst MEAs at: A) 25 ℃; B) 60 ℃. C) EIS curves for the Pt-1 and Pt-2 catalyst MEAs at 0.4 V at 60 ℃ and D) the equivalent circuit used to fit the spectra. ................................................................................................................... 55 Figure 2.13. Contact angle measurement image for a) Pt-1C and b) Pt-1F ................................. 56 Figure 2.14. Contact angle measurement image for a) Pt-2C and b) Pt-2F ................................. 56 Figure 2.15. A) Polarization curves for the MEAs with Pt-1 anode and Pt-2 cathode electrodes at 25 ℃ and 60 ℃. Solid lines indicate CFx supported and dashed lines indicate XC72R supported catalysts. ........................................................................................................................................ 57 Figure 2.16. A) MEAs held at constant currents with Pt-1 anode and Pt-2 cathode electrodes at 60 ℃. B) Polarization curves with Pt-1 anode and Pt-2 cathode electrodes after stability experiments with a forward and reverse scan. Solid lines indicate CFx supported and dashed lines indicate XC72R supported catalyst. ........................................................................................................... 58 Figure 2.17. A) TGA plot and B) XRD pattern of the prepared catalysts. .................................. 62 Figure 2.18. TEM micrographs of A) Pd/CFx and B) Pd/C ........................................................ 63 Figure 2.19. A) XPS PD 3d spectra of the Pd/CFx catalysts. B) CV scans of the prepared catalysts in 1M KOH at a scan rate of 20 mV/s under flow of N2. ............................................................. 64 Figure 2.20. CV scans of the prepared catalysts at scan rate of 20 mV/s under N2 with A) 1M formate and 1M KOH; B) 1M methanol and 1M KOH; C) 1M glycerol and 1M KOH and D) 1M ethylene glycol and 1M KOH ....................................................................................................... 66 xii Figure 2.21. CV scans of A) Pd/C and B) Pd/CFx in 1M formate and 1M KOH; C) Pd/C and D) Pd/CFx in 1M methanol and 1M KOH; E) Pd/C and F) Pd/CFx in 1M glycerol and 1M KOH; G) Pd/C and H) Pd/CFx in 1M ethylene glycol and 1M KOH at various scan rates. ........................ 68 Figure 2.22. Plot of Forward peak current vs square root of scan rate and plot of peak potential vs natural log of scan rate in A and B) 1M formate and 1M KOH; C and D) 1M methanol and 1M KOH; E and F) 1M glycerol and 1M KOH and G and H) 1M ethylene glycol and 1M KOH. .... 70 Figure 2.23. Tafel plots for the prepared catalysts under A) 1M formate and 1M KOH; B) 1M methanol and 1M KOH; C) 1M glycerol and 1M KOH; and D) 1M ethylene glycol and 1M KOH. ....................................................................................................................................................... 72 Figure 2.24. CV scans of A) Pd/C and B) Pd/CFx in 1M formate and 1M KOH; C) Pd/C and D) Pd/CFx in 1M methanol and 1M KOH; E) Pd/C and F) Pd/CFx in 1M glycerol and 1M KOH; G) Pd/C and H) Pd/CFx in 1M ethylene glycol and 1M KOH at various temperatures. ................... 73 Figure 2.25. Arrhenius plots for Pd/CFx and Pd/C in A) 1M formate and 1M KOH; B) 1M methanol and 1M KOH; C) 1M glycerol and 1M KOH and D) 1M ethylene glycol and 1M KOH. ....................................................................................................................................................... 74 Figure 2.26. Fuel cell results for alkaline direct glycerol fuel cells with Pd/CFx and Pd/C anodes and Pt/C cathode electrodes and cell temperature of 80 ℃. With A) 1M glycerol; B) 2M glycerol and C) 3M glycerol with 6M KOH. .............................................................................................. 76 Figure 2.27. Fuel cell results for alkaline direct ethylene glycol fuel cells with Pd/CFx and Pd/C anode and Pt/C cathode electrodes and cell temperature of 80 ℃. With A) 1M ethylene glycol; B) 2M ethylene glycol and C) 3M ethylene glycol with 6M KOH. .................................................. 76 Figure 2.28. Fuel cell results for alkaline direct methanol fuel cells with Pd/CFx and Pd/C anode and Pt/C cathode electrodes and cell temperature of 80 ℃. With A) 1M methanol; B) 2M methanol and C) 3M methanol with 6M KOH. ............................................................................................ 77 Figure 2.29. Fuel cell results for alkaline direct formate fuel cells with Pd/CFx and Pd/C anode and Pt/C cathode electrodes and cell temperature of 80 ℃. With A) 1M formate; B) 2M formate and C) 3M formate with 6M KOH. .............................................................................................. 78 Figure 3.1. Characterizations of the catalyst: A) The TGA profile for Pd/rGO catalyst. The TGA experiments were carried out in an air atmosphere with a heating rate of 10 ℃/min. B) Raman Spectra and; C) XRD patterns. ...................................................................................................... 97 Figure 3.2. SEM images of: A) GO; B) Pd/rGO; TEM image of Pd/C C) and D) and Pd/rGO E) and F). ........................................................................................................................................... 99 Figure 3.3. HRTEM images for A) Pd/rGO and B) Pd/C ............................................................ 99 Figure 3.4. XPS spectra of: A) Survey scan; B) C1s of Pd/rGO; C) C1s of GO and; D) Pd 3d of Pd/rGO. ....................................................................................................................................... 100 xiii Figure 3.5. CV scan of the catalyst in: A) 1.0 M NaOH and; B) 1.0 M HCOONa and 1.0 M NaOH at 20 mV/s. C) Chronoamperommetry scans of the catalyst at -0.65 V vs Ag/AgCl at 1000 RPM in 1.0 M HCOONa and 1.0 M NaOH. ........................................................................................ 102 Figure 3.6. CV scans at different scan rates in 1.0 M HCOONa and 1.0 M NaOH for: A) Pd/rGO and B) Pd/C. C) Plot of Forward peak current vs square root of scan rate and; D) plot of corresponding potential vs ln (v)................................................................................................. 103 Figure 3.7. A) Tafel Plots for the catalyst in 1.0 M HCOONa and 1.0 M NaOH at 0.5 mV/s. CV scans of the catalysts at different temperatures in 1.0 M HCOONa and 1.0 M NaOH for B) Pd/rGO and C) Pd/C at 20 mV/s and D) Arrhenius plots for the catalysts. ............................................. 105 Figure 3.8: A) EIS curves of the catalyst. B) Equivalent circuit of the EIS spectra. C) Single cell DFFC polarization curves for the catalysts at 30 and 60 ℃. ...................................................... 106 Figure 3.9. A) TGA profile and B) powder XRD pattern of the prepared catalysts. ................. 111 Figure 3.10. A) N2 adsorption-desorption isotherms of XC72R and MWCNT and B) corresponding mesoporous size distribution curves for XC72R and MWCNT ......................... 112 Figure 3.11. SEM images and EDS mapping for A) Pd/C and B) Pd/MWCNT where red corresponds to C and green corresponds to Pd. SEM images and EDS mapping for C) IrO2C; D) PdIrO2C (15:5); E) PdIrO2C (10:10); F) PdIrO2C (5:15) and G) PdIrO2MWCNT where red corresponds to carbon, green corresponds to O, blue corresponds to Pd and yellow corresponds to Ir. ................................................................................................................................................. 115 Figure 3.12. TEM micrographs of A) PdC; B) PdIrO2C (15:5); C) PdIrO2C (10:10); D) PdIrO2C (5:15); E) PdMWCNT and F) PdIrO2MWCNT.......................................................................... 117 Figure 3.13. CV scans of the prepared catalysts in A) 1M KOH and B) 1M HCOOK and 1M KOH at a scan rate of 20 mV/s under a flow of N2. ............................................................................. 118 Figure 3.14. CV scans in 1M HCOOK and 1M KOH at various scan rate for A) PdC; B) PdIrO2C (15:5); C) PdIrO2C (10:10); D) PdIrO2C (5:15); E) PdMWCNT and F) PdIrO2MWCNT ........ 120 Figure 3.15. A) Plot of Forward peak current vs square root of scan rate and B) plot of peak potential vs natural log of scan rate. ........................................................................................... 121 Figure 3.16. A) Tafel slopes obtained from LSVs at 0.5 mV/s and B) change in current over time holding the potential at 0.4 V vs. RHE for the prepared catalysts. ............................................. 122 Figure 3.17. EIS spectra at various potentials vs. MMO for A) PdC; B) PdIrO2C (15:5); C) PdIrO2C (10:10); D) PdIrO2C (5:15); E) PdMWCNT and F) PdIrO2MWCNT ......................... 124 Figure 3.18. A) Polarization and power curves for the DFFCs with the prepared catalysts and B) the DFFC with PdIrO2MWCNT anode electrode held at constant current of 200 mA/cm 2 ....... 125 xiv Figure 3.19. A) CV scans of the formate salts at a scan rate of 20 mV/s under flow of N 2. B) CV scans of the hydroxide base at a scan rate of 5 mV/s under flow of N2...................................... 128 Figure 3.20. CV scans at various scan rates for A) HCOOLi; B) HCCONa; C) HCOOK; D) HCOORb and E) HCOOCs with Pd/C catalysts under flow of N2............................................. 130 Figure 3.21. A) Plot of Forward peak current vs square root of scan rate and B) plot of peak potential vs natural log of scan rate. ........................................................................................... 131 Figure 3.22. A) Tafel plots for the various formate salts and B) EIS at 0.1 V vs. Ag/AgCl for the various formate salts. .................................................................................................................. 132 Figure 3.23. Constant potential experiments for the A) Li, Na, K and B) Rb and CS formate salts held at 0.1 V vs. Ag/AgCl. .......................................................................................................... 133 Figure 3.24. A) Polarization and power curves for 1M HCOONa with varying membranes at 80 ℃. B) The potential of the DFFCs held at constant current. ...................................................... 134 Figure 3.25. A) Polarization and power curves for 1M HCOOM with varying cations at 80 ℃. B) The potential of the DFFCs held at constant current. ................................................................. 135 Figure 4.1. A) Chemical Structures of the polymer electrolytes used in this study a) m-TPN1- TMA, b) m-TPN1-Pip and c) m-TPN1-DMIm and B) 1 H NMR of a) m-TPBr, b) m-TPN1-TMA, c) m-TPN1-Pip and d) m-TPN1-DMIm ...................................................................................... 159 Figure 4.2. CV scan utilizing different AEI in the electrode in A) 1.0 M KOH and B 1.0 M MeOH and 1.0 M KOH at 20 mV/s. ....................................................................................................... 160 Figure 4.2: CV in 1M KOH under N2 (black curve) and CO (Red curve) with A) DMIm AEI electrode, B) Pip AEI electrode, C) TMA AEI electrode. .......................................................... 161 Figure 4.3: CV in 1 M MeOH +1M KOH at varying scan rates with A) DMIm AEI electrode B) Pip AEI electrode and C) TMA AEI electrode. .......................................................................... 163 Figure 4.4. A) Plot of Forward peak current vs square root of scan rate; B) plot of peak potential vs natural log of scan rate and; C) Tafel plot for the electrodes containing the three different AEI. ..................................................................................................................................................... 165 Figure 4.5. A) Chronoamperometry scans of the three AEI electrodes at 0.69 V vs RHE in 1.0 M MeOH and 1.0 M KOH. B) Nyquist plot for the three AEI electrodes at 0.74 V vs RHE and the equivalent circuit. ........................................................................................................................ 166 Figure 4.6: EIS under 1M MeOH + 1M KOH at varying potentials with A) DMIm AEI electrode, B) Pip AEI electrode and C) TMA AEI electrode. ..................................................................... 167 Figure 4.7. Single cell ADMFC polarization curve for the anode electrode with varying AEIs A) at 60 ℃ and B) at 80 ℃. .............................................................................................................. 168 xv Figure 4.8. Single cell ADMFC polarization curve for the anode electrode with varying AEIs at 60 ℃ A) with Sustainion membrane, B) with Tokuyama A201 membrane. 2M MeOH without added KOH using C) Sustainion membrane, D) Tokuyama A201 membrane and E) TPN membrane. ................................................................................................................................... 169 Figure 4.9: A) Power density vs time plot for single cell ADMFC constant potential experiments held at 0.4 V. B) 1 H NMR and C) 13 C NMR of the ADMFC fuel before and after constant potential experiment with TMA anode binder. To prepare the NMR samples, 500 µl of the fuel was mixed with 500 µl of D2O and 100 µl of isopropyl alcohol (IPA) as an internal standard. .................. 171 Figure 4.10. Single cell polarization curve for A) ADMFC without and with MPL cathode GDE and B) change in catalyst to AEI in anode GDE. C) Power density vs time plot for single cell ADMFC constant potential experiments held at 0.64 V ............................................................. 172 Figure 4.11. A) TGA curve for prepared catalyst B) XRD pattern of the prepared catalyst where • corresponds to Ni and ° corresponds to Ni(OH)2. SEM images of B) Ni/C C) GaNi3/C D) Ga2Ni/C E) Ga3Ni/C F) Ga4Ni/C and G) Ga/C. ........................................................................................ 179 Figure 4.12. SEM and EDS images for A) Ni/C B) GaNi3/C C) Ga2Ni/C D) Ga3Ni/C E) Ga4Ni/C and F) Ga/C. Where blue corresponds to Ni, yellow corresponds to Ga, green corresponds to O, and red corresponds to C............................................................................................................. 182 Figure 4.13. A) Survey scan B) Ni 2p C) Ga 2p D) C 1s and E) O 1S XPS spectra of the prepared catalysts. ...................................................................................................................................... 183 Figure 4.14. CV experiments of the prepared catalysts in A)1M KOH. CVs at varying scan rates in 1M KOH for B) Ni/C C) Ga4Ni/C D) Ga3Ni/C E) Ga2Ni/C F) GaNi3/C under N2 flow. ...... 185 Figure 4.15. A) CVs of the prepared catalyst in1M KOH with 1M methanol at a scan rate of 20 mV/s B) Nyquist plot for the prepared catalysts at a constant potential of 0.7 V vs. MMO and C) Stability experiments of the prepared catalysts holding a constant potential of 0.7 V vs. MMO. ..................................................................................................................................................... 187 Figure 4.16. Polarization curve for fuel cells with A) varying anode electrode, B) varying methanol concentration, C) varying KOH concentration and D) α-MnO2/C cathode electrode. 189 Figure 4.17. A) TGA plot and B) XRD spectra for the prepared catalyst. ................................ 194 Figure 4.18. SEM images of A) Ni/C and B) NiCeO2/C. SEM and EDS images of C) Ni/C, where blue corresponds to Ni, red corresponds to C and green corresponds to O and D) NiCeO2/C, where blue corresponds to Ce, yellow corresponds to Ni, red corresponds to C and green corresponds to O. ................................................................................................................................................. 195 Figure 4.19. XPS A) survey scan, high resolution scans of the B) Ce 3d, C) Ni 2p, D) C 1s and E) O 1s region for the prepared catalysts......................................................................................... 197 Figure 4.20. CV scans of the prepared catalyst in 1M KOH A) at scan rate of 20 mV/s, B) Ni/C at various scan rates and C) NiCeO2/C at various scan rates under flow of N2. ............................. 198 xvi Figure 4.21. A) CV at a scan rate of 20 mV/s, B) potentiostatic hold at 0.7 V vs. Ag/AgCl and C) EIS at 0.5V vs. Ag/AgCl of the prepared catalyst in 1 M KOH + 1M MeOH under flow of N 2. ..................................................................................................................................................... 199 Figure 4.22. Polarization curves with A) NiCeO2/C anode and B) Ni/C electrodes. C) NiCeO2/C anode electrode with 1.5M MeOH + 6M KOH at varying temperatures. D) Comparison of ADMFCs with Pt/C and FegCN cathode electrodes................................................................... 201 Figure 4.23. A) ADMFCs with NiCeO2/C anode and FegCN cathode electrodes held at a constant potential of 0.6 V. B) 1 H NMR of the fuel reservoir before and after the experiment, 100 µL of n- butanol was added as an internal standard. ................................................................................. 202 Figure 5.1. A) TGA profile for the prepared catalyst under air atmosphere with a ramp rate of 10°/min. B) Powder XRD pattern for the prepared catalyst. ...................................................... 230 Figure 5.2. SEM and EDS mapping of A) CuXC72-np where carbon corresponds to red, oxygen corresponds to green and copper corresponds to blue; B) CugCN-np where carbon corresponds to red, oxygen corresponds to blue, nitrogen corresponds to green and copper corresponds to yellow; C) CugCN-SAC where carbon corresponds to red, nitrogen corresponds to green and copper corresponds to blue; and D) gCN where carbon corresponds to red, copper corresponds to yellow and nitrogen corresponds to green. ............................................................................................. 233 Figure 5.3. TEM micrographs of A) CuXC72-np; B) CugCN-np; C) CugCN-SAC; and D) gCN. ..................................................................................................................................................... 234 Figure 5.4. XPS A) survey scan; and high-resolution scans for B) Cu 2p; C) N 1s; and D) C 1s for the prepared catalyst. .................................................................................................................. 236 Figure 5.6. LSV scans under Ar and CO2 flow at scan rates of 20 mV/s in 0.1M NaHCO3 for A) CuXC72-np; B) CugCN-np; C) CugCN-SAC; and D) gCN. ..................................................... 239 Figure 5.7. Current values obtained from constant potential electrolysis experiments in 0.1M NaHCO3 under flow of CO2 for A) CuXC72-np; B) CugCN-np; C) CugCN-SAC; and D) gCN. ..................................................................................................................................................... 240 Figure 5.8. A) Faradaic efficiencies of ethylene glycol for CugCN catalyst; B) liquid sample and C) gas sample GC spectra for CugCN-SACs at various potentials; and D) current and FE of CugCN-SAC at -1.0 V vs. Ag/AgCl over 10 h. .......................................................................... 242 xvii Abstract This dissertation comprises of five chapters discussing developments in polymeric materials to catalysts for use in fuel cells and electrochemical CO2 reduction. In Chapter 1, an introduction to fuel cells and electrochemical CO2 reduction is provided. Chapter 2 discusses further studies of partially fluorinated carbon (CFx) supported platinum nanoparticles for hydrogen oxidation reaction (HOR) and oxygen reduction reaction (ORR). The Pt/CFx stability is compared to Pt supported on Vulcan XC72R for both HOR and ORR. The Pt/CFx displayed improved stability due to the improved interaction between the Pt and CFx support. Moreover, the Pt/CFx displayed improved electrocatalytic activity for HOR compared to Pt/XC72R. The improvement is due to improved mass transport and water management. Furthermore, the Pt/CFx was used as both he anode and cathode electrocatalyst in a H2-proton exchange membrane fuel cell (PEMFC) and displayed a 20% improvement in terms of peak power compared to a H2-PEMFC with Pt/XC72R. Furthermore, the CFx support was also utilized to prepared Pd/CFx catalysts for alkaline direct liquid fuel cells. The Pd/CFx catalysts displayed improvements over Pd/C arising from increased electrochemically active surface area and improved mass transport. The Pd/CFx catalyst also displayed improved performance in alkaline direct liquid fuel cells for various liquid fuels (glycerol, methanol, ethylene glycol and formate). The use of CFx as a catalyst support is robust and aids in improving both acidic and alkaline fuel cells, aiding mass transport and stability. In chapter 3, developments in formate fuel cells are discussed from catalyst design to cell structure. First, we have explored the activity of palladium supported on reduced graphene oxide (Pd/rGO) towards the formate oxidation reaction (FOR) in an alkaline medium. The reduction of GO to rGO and synthesis of Pd nanoparticles were confirmed using X-Ray Diffraction (XRD), xviii Raman and X-Ray Photoelectron Spectroscopy (XPS). The surface morphology was evaluated by Scanning Electron Microscopy (SEM) and Transmission Electron Microscopy (TEM). Half-cell studies demonstrated superior electrocatalytic activity and stability towards formate electrooxidation for Pd/rGO than commercial Pd/C catalysts. A low metal loading DFFC, fabricated with a Pd/rGO anode catalyst displayed a 15% increase in maximum power density at 60 ℃ compared to the commercial Pd/C catalyst. Improvements were further made by combining IrO2 with Pd and MWCNTs. The addition of IrO2 and MWCNTs both improved the ECSA of the catalyst and improved kinetics of FOR. Moreover, the improvements were demonstrated in a DFFC providing one of the highest reported peak power densities of a DFFC, 299 mW/cm 2 , while utilizing less than 1 mgpd/cm 2 . Finally, the role of cations on FOR was studied and found to play a significant role in current observed. HCOONa displayed the highest performance due to improved mass transport and more facile removal of reactive intermediates from the catalyst surface. Moreover, the formate salts were utilized with a cation exchange membrane to coproduce electricity and alkali hydroxide. The first section of this chapter is reprinted with permission from ACS Appl. Energy Mater. 2019, 2, 10, 7104–7111. Copyright 2019 American Chemical Society. Chapter 4 discusses improvements in alkaline direct methanol fuel cells (ADMFCs), specifically changes in the ionomer and membrane materials and catalysts. The first portion of the chapter reports the activity of the methanol oxidation reaction (MOR) in half-cell experiments with varying ionomers and the use of a poly(terphenylene) (TPN) membrane in an alkaline direct methanol fuel cell (ADMFC). The results demonstrate that small changes in the cation structures of the ionomer have a significant role in MOR on the PtRu/C catalyst. Moreover, with the use of a TPN membrane and the prepared anode containing ionomers, high power densities are achieved with less than 1 mgPtRu/cm 2 in the catalyst layer. Next, nickel-based catalysts were prepared and xix studied for methanol oxidation. Firstly, GaNi based catalysts were prepared at various ratios and it was found that a 3 to 1 Ga to Ni provided the highest performance in half-cell studies. The GaNi catalysts were also demonstrated to be effective in ADMFCs, displaying a similar trend in catalytic activity as the half-cell experiments. Lastly, a NiCeO2/C catalyst was also prepared and displayed improved catalytic activity compared to Ni/C due to increased ECSA, and kinetics. Moreover, both GaNi/C and NiCeO2/C catalysts were utilized in precious metal free ADMFCs by using MnO 2/C and FegCN catalysts for the cathode electrode. The first section of this chapter is reprinted with permission from ACS Appl. Energy Mater. 2021, 4, 6, 5858–5867. Copyright 2021 American Chemical Society In chapter 5, copper single atom catalysts (SACs) were prepared with a facile and scalable procedure and utilized for electrochemical CO2 reduction (ECO2R). The SAC formation was confirmed with XPS, SEM, TEM and XRD analysis. The CugCN-SACs were then utilized in an H-cell for ECO2R, the catalyst was able to reduce CO2 to ethylene glycol with high selectivity and maintain Faradaic Efficiencies near 80% for production of ethylene glycol for 10 h at a constant potential of -1.0 V vs Ag/AgCl. Thus, demonstrating one of the most effective catalysts for ECO2R to ethylene glycol. 1 Chapter 1: Introduction 1.1. The Climate Crisis and Clean Energy The consumption of all forms of energy has been increasing with the rapid industrial developments and increasing populations across the globe. Even though a 1% decrease was observed in 2020 due to the global COVID-19 pandemic, the U.S. Energy Information Administration (EIA) and International Energy Agency (IEA) project that the increasing rate of energy use will continue over the next few decades as shown in Figure 1.1. The increase is expected to come primarily from the continued development of countries in the Asia Pacific region. 2 Figure 1.1. A) Worldwide energy consumption by regions. Adapted from [1] B) Worldwide energy consumption by source. Adapted from [2] Fossil fuels are expected to be the primary sources of energy, especially with the resurgence of economies from the COVID-19 pandemic (Figure 1.1B). Even though fossil fuels have a relatively simple combustion process and have been used for centuries, they are fraught with flaws. One of the most evident drawbacks of fossil fuel use is the increase in CO 2 emissions which has been increasing every year and has reached levels over 420 ppm, significantly higher than any other point in the last 800,000 years of our earth (Figure 1.2).[3,4] This significant increase in CO2 levels has led to unnatural temperature rises, climate change and ocean acidification.[5] The use of fossil fuels also emits other polluting gases such as methane, NOx, SOx, and particulate matter, which are all detrimental to the well-being of living beings and the environment.[6] 3 Figure 1.2. CO2 levels for the past 800,000 years. Adapted from [3] Moreover, these emissions have harmful effects on human health as well, between 4-7 million premature deaths are due to air pollutions. Another major drawback is the limited availability of fossil fuels and their location in the unstable regions of the world.[7] Even though new methods of extraction and fuel sources have surfaced, such as fracking for extraction of natural gas, it has also been shown to be damaging to the environment and human health.[8] Moreover, leaks, spills and transportation accidents can have severe consequences. Such as in the ruptured gas line in the Gulf of Mexico in 2021, which resulted in a fire in the ocean with unknown ramifications to ocean life and the environment.[9] In recent times, the development and advancements in clean and renewable energy has been significant, and energy generation from these sources has improved as shown in Figure 1.3. 4 Figure 1.3. Electricity generation from various sources. Adapted from [10] Renewable energy sources that have aided in alleviating the energy demands are solar, wind, hydroelectric power and electrochemical devices (batteries, capacitors and fuel cells) with most of the current energy needs being met by solar, wind and hydroelectric. However, these technologies all have major drawbacks: the lack of control of energy production, their intermittency and depend on regional characteristics. Solar energy will not produce energy in the absence of sunlight and wind turbines will fail to produce energy without wind. Electrochemical energy, which consists of several different technologies as shown in Figure 1.4, can be utilized regardless of regional characteristics and produce or store energy regardless of the environment. Moreover, the different technologies span a large area of power and energy needs. One of the many types of electrochemical energy devices which has gained increased attention are fuel cell technologies. 5 Figure 1.4. Ragone plot of various power sources. Adapted from [11] 1.2. Fuel Cells Over the past few decades fuel cell technologies have gained an increased popularity in industrial and personal use. However, fuel cells have been around for much longer than that, with the earliest developments dating back to 1838 and 1839, where Schoenbein first introduced the concept of a fuel cell and Grove was the first to bring this concept into development, although, under a different name “the gas battery”.[12–14] It wasn’t until 1889 that Ludwig Mond and Charles Langer devised the first practical device and termed it “fuel cell”. However, it wasn’t until 1932 that significant improvements were made to fuel cell technologies, Francis Bacon developed the first practical H2/O2 fuel cell under alkaline conditions. Bacon developed fuel cells that were utilized in submarines by the Royal Navy during World War II and his work was also utilized in the 1960s and 1970s for NASAs spacecraft missions, notably the Apollo mission.[13] In the 1960s Dupont introduced a fluorocarbon, cation exchange membrane, Nafion. With the introduction of Nafion the focus shifted from alkaline fuel cells to proton exchange membrane fuel cells.[15] 6 Moreover, several types of high temperature fuel cells were developed in the 1960s such as the phosphoric acid fuel cell, molten carbonate fuel cell, and solid oxide fuel cell. In the 1990s, the direct methanol fuel cell (DMFC) was developed by a collaborative work between the University of Southern California and the Jet Propulsion Laboratory of NASA.[16] Since then, several types of fuel cells have been introduced and major improvements have been made. 1.3. Workings of a Fuel Cell Fuel cells convert the chemical energy in a fuel into electrical energy via catalysts and uses a potential difference between the fuel and an oxidant to drive current. A fuel cell is composed of three active materials: a fuel electrode (anode), an oxidant electrode (cathode) and an electrolyte sandwiched in between the anode and cathode. Other typical components to a fuel cell are gas diffusion layers (GDLs), bipolar flow fields, and current collectors as shown in Figure 1.5. Fuel cells and batteries have some similarities in their design, however, differ in their functionality, batteries are utilized as energy storage and conversion devices whereas fuel cells are only energy conversion devices. 7 Figure 1.5. Schematic of a fuel cell single stack Moreover, batteries use the chemical energy stored in the electrode material to produce electricity at specified potentials, thus can only function so long as the material is not depleted. Once depleted the battery must be recharged or replaced, depending on whether it is a rechargeable battery or primary battery. However, in a fuel cell, the fuel and oxidant supplied are used to produce electricity where the internal components are not consumed, therefore a fuel cell can theoretically continue to work as long as reactants are supplied. Fuel cell performance is typically assessed by a polarization curve, shown in Figure 1.6, where a higher potential, current density and power density are desired. 8 Figure 1.6. Polarization and power curve for a fuel cell. Adapted from [14] The cell voltage is often lower than the theoretical voltage due to parasitic losses like fuel crossover (from anode to cathode), inefficient catalysts, poisoning of electrolyte and dehydration of electrolyte membrane. There are three regions seen in an ideal polarization curve displaying losses in the fuel cell potential. The first region in the low current range is termed activation or kinetic losses which arises from the rate of charge transfer occurring at the electrode surface. The second region is the ohmic losses and is due to the resistance of the fuel cell external components, electrode, and membrane. The third region is the mass transport or diffusion losses and occurs from the change in reactant concentration at the surface of the electrode.[17] In commercial applications, multiple fuel cells are connected in series and parallel to compose a fuel cell stack to obtain the voltage and current for the desired application. 1.4. Fuel Cell Types One of the major contributors to the rapid progression and advancements in fuel cell development is the myriad of environments in which fuel cells have been shown to operate. The 9 different types of fuel cells, shown in Figure 1.7, vary based on the membrane/electrolyte, operating temperature, reactants used. The wide range of efficiencies and power generations of fuel cells allow for a plethora of applications. Figure 1.7. Schematic of different types of fuel cells and their operating conditions. 1.4.1. Proton Exchange Membrane Fuel Cell The most common low temperature fuel cell that uses hydrogen as a fuel, is the proton exchange membrane fuel cell (PEMFC). PEMFCs use a polymeric material, most commonly a perfluorosulfonic acid-based membrane like Nafion, to conduct protons from the anode to the cathode. PEMFCs have a wide range of applications from transportation to portable applications due to their high-power density, fast start-up time, high efficiency, relatively low operating 10 temperature and facile fabrication. The reaction within a hydrogen based PEMFC is as follows [18]: Anode: 2H2 → 4H + + 4e - E0 = 0.00 V (1) Cathode: O2 + 4H + + 4e - → 2H2O E0 = 1.23 V (2) Overall: 2H2 + O2 → 2H2O E0 = 1.23 V (3) As shown one of the major advantages of hydrogen based PEMFCs is the only product released is water. However, PEMFCs are not without their drawbacks such as the need for noble metal catalyst for optimal fuel cell performance and stability under acidic environment. However, there have been recent studies showing nonnoble metals as single atom catalysts that display promising results under acidic conditions. Other drawbacks are the low tolerance of CO and impurities in the hydrogen fuel that can result in catalyst or membrane poisoning, and the storage and transportation of hydrogen gas under very high pressure.[19,20] PEMFCs, however, are not limited to gaseous reactants, direct liquid fuel cells (DLFCs) have also been developed.[21,22] The use of liquids as a fuel source makes the storage and transportation of the fuel more facile than gaseous reactants. Moreover, integrating liquids to existing infrastructure is a lot more facile than their gaseous counterparts. Several liquids have been studied in PEMFCs such as ethanol, methanol, formic acid, ethylene glycol and glycerol. The most studied being the direct methanol fuel cell (DMFC); since its introduction in the 1990s, which sparked the methanol economy concept, there have been several improvements to its performance and has been used in several applications.[16,23–26] One of the major advantages of using methanol is the significantly larger volumetric energy density when compared to hydrogen, 4.8 11 kWh/L and 0.4 kWh/L respectively. The reactions within a DMFC are similar to a PEMFC as shown below: Anode: CH3OH + H2O → CO2 + 6H + + 6e - E0 = 0.02 V (4) Cathode: 3/2 O2 + 6H + + 6e - → 3H2O E0 = 1.23 V (5) Overall: 2H2 + O2 → CO2 + 2H2O E0 = 1.21 V (6) DMFCs also have disadvantages that have prevented major commercialization such as methanol crossover and CO poisoning on the platinum catalyst resulting in Pt-Ru catalyst showing the highest power density in DMFCs. A higher catalyst loading than in typical hydrogen based PEMFCs is also required driving up the cost.[27] 1.4.2. Alkaline Fuel Cells Alkaline fuel cells (AFCs) utilize hydroxide-based electrolytes and unlike PEMFCs, they conduct OH - from the cathode to the anode. The reactions are shown below: Anode: 2H2 + 4OH - → 4H2O + 4e - E0 = -0.83 V (7) Cathode: O2 + 2H2O + 4e - → 4OH - E0 = 0.40 V (8) Overall: 2H2 + O2 → 2H2O E0 = 1.23 V (9) Even though AFCs were the first developed fuel cells, the lack of a stable and effective anion exchange membrane (AEM) resulted in a larger focus on studying PEMFCs. Interest in AFCs remained due to advantages posed by working under alkaline conditions such as faster oxygen reduction kinetics, the stability and catalytic activity of nonnoble metals, they do not have a corrosion problem and can operate at low temperatures.[28,29] Work resurged with developments 12 in AEMs, especially in recent years there has been a significant effort into developing AEMs, which can bring forth all the advantages of AFCs. Several AFCs have shown high-power, long- term stability and functionality with nonnoble metal catalysts.[30–34] However, some drawbacks remain for AFCs such as the electrolyte intolerance to CO2, development of proper ionomeric binders and stability are still not on par with PEMFCs Similar to PEMFCs, AFCs have also been shown to work with a myriad of liquid fuels, alkaline direct liquid fuel cells (ADLFCs), and have displayed higher performance in many cases compared to PEM counterpart.[22,35] Ethanol, ethylene glycol, glycerol, formate and methanol have been shown to work under alkaline conditions, moreover several other liquid fuels have been shown to work such as urea, glucose, mannitol, proponal and ascorbate.[36–42] Many of these fuels function without platinum and often display higher performance with non-noble metal electrocatalysts. There is also less fuel crossover, especially crossover due to electroosmotic drag. Of all the ADLFCs, the one that has shown the most promise has been formate. Unlike many of the other alkaline liquid fuels, formate oxidation is independent of pH, allowing the formate oxidation to occur without any supporting electrolyte.[43] Further discussion of formate fuel cells will be discussed in Chapter 3. ADLFCs also have additional drawbacks than the disadvantages of AFCs, such as requiring additional electrolyte added to the fuel, lower power output and requiring high loading of catalyst. 1.4.3. High Temperature Fuel Cells There are also fuel cells which operate at higher temperatures, relative to PEMFCs and AFCs, ranging from 200-1000 ℃. The two most developed being the molten carbonate fuel cell (MCFC) and solid oxide fuel cell (SOFC). The MCFC was first introduced in 1960s by Broers and 13 Ketelaar, where they demonstrated a fuel cell with an electrolyte composed of lithium, sodium and potassium carbonate mixture impregnated on a porous agglomerate disc of magnesium oxide. [44,45] MCFCs now function utilizing a lithium and potassium carbonate mixture electrolyte which shuttle carbonate ions from the cathode to the anode. Moreover, unlike other fuel cells MCFCs use CO2 as a reactant which is typically recirculated from the anode as shown below.[18] Anode: 2H2 + 2CO3 2- → 2H2O + 2CO2 + 4e - (7) Cathode: O2 + 2CO2 + 4e - → 2CO3 2- (8) Overall: 2H2 + O2 → 2H2O (9) The typical operating temperature for MCFC is between 500-700 ℃, this allows for use of nonnoble metal electrocatalysts, nickel-based catalysts are typically used. Moreover, they are highly efficient and allow for internal reforming of hydrocarbon fuels and thus can utilize various types of fuels.[46] However, some of the issues with MCFCs are electrolyte contamination and evaporation, catalyst dissolution, low power densities and necessity of CO2 in the cathode stream.[47,48] The most developed high temperature fuel cell is the solid oxide fuel cell (SOFC), which also runs at the highest temperature of any type of fuel cell between 700-1000 ℃. SOFCs were first introduced in the 1930s by Baur and Preis who utilized solid oxide electrolytes and zirconium, lanthanum, or yttrium.[49] Since then, several developments have been made in SOFCs. Some of the major advantages of SOFCs are they contain no corrosive components, fuel flexible, long life durability and work without precious-metal catalysts at its high temperature of operation.[50,51] The cell functions by utilizing a ceramic electrolyte that conducts oxide ions from the cathode to the anode and the typical reactions are shown below [18,52]: 14 Anode: 2H2 + 2O2 - → 2H2O + 4e - E0 = -0.25 V (10) Cathode: O2 + 4e - → 2O2 - E0 = -0.25 V (11) Overall: 2H2 + O2 → 2H2O E0 = -0.25 V (12) SOFCs have found use as auxiliary power systems and other stationary applications. However, some drawbacks of utilizing SOFCs are the long start up times due to the high operating temperature, the need for such high temperatures and strict material requirements.[53] The final high temperature fuel cell is the phosphoric acid fuel cell (PAFC), which is the most mature and first commercially used fuel cell. PAFCs utilize phosphoric acid as the electrolyte, which is contained in a silicon carbide matrix, platinum catalysts contained in a porous carbon electrode. The appeal of PAFCs is the high tolerance of CO2, simple water management, high efficiency when cogenerating heat and electricity. However, some of the drawbacks of PAFCs are the higher catalysts loadings required, they are typically large, slow start up times and lower power output than other fuel cells.[54,55] 1.5. Fuel Cell Applications The various operating conditions of fuel cells, differing in efficiencies and power generations have allowed for them to find commercial use in several different applications (Figure 1.8). Low temperature fuel cells have mostly found use in portable devices due to the lower power, not needing many auxiliary components and low operating temperatures. Whereas high temperature fuel cells have found use in stationary applications and auxiliary power systems. 15 Figure 1.8. Typical power ranges for various fuel cells. Adapted from [56] As the focus of this dissertation will be on low temperature fuel cells, the remainder of this discussion will be centered on applications of PEMFCs and AFCs. One of the major appeals in utilizing fuel cells is the energy density of the system, as shown in Figure 1.5, which allows for longer discharge times before needing to be refueled or recharged when compared to other power systems. One of the past major applications of AFCs has been in aerospace technologies, dating back to the Apollo mission in 1963. However, due to the historically lack of a stable and effective AEM, AFCs have not been utilized in many commercial applications. But with the recent advancements in AEMs several proofs of concepts have been shown like a scooter being powered by an alkaline direct formate fuel cell stack.[57] PEMFCs have found more success in commercial products and applications. In the early 2000s Toshiba demonstrated DMFC based portable chargers and cell phones.[24,58] In recent years several car manufactures have released fuel cell vehicles to the public such as the Toyota Mirai, Hyundai Nexo and Honda Clarity and other vehicles are in development such as Hyperion XP1 and Roland Gumpert’s Nathalie. Moreover, 16 fuel cells have also found considerable attention for larger commercial vehicles and modes of transportations such as trains, buses, and commercial trucks due to the ample space to accommodate various components.[59,60] Furthermore, there has been a surge in utilizing fuel cells for aviation purposes, such as ZeroAvia’s work on fuel cell powered aircrafts and several airports around the world having constructed hydrogen refueling stations for fuel cells. 1.6. CO2 Reduction With the ever-increasing levels of CO2 and greenhouses gases, renewable energy sources are not the only avenue for mitigating environmental catastrophe. Renewable energy sources allow us to decrease emissions, however another method to improve the climate crisis is to convert the greenhouse gases into chemically useful fuels and products. There are several methods to convert CO2 into value added products based on hydrothermal and electrochemical conversion.[61] In the aforementioned cases CO2 is reduced to hydrocarbons with proton and electron sources transferred by catalysts. The catalysts are crucial due to the relative stability of CO2 and intermediates with high energy barriers, in electrochemical reduction these intermediates are often avoided by proton coupled electron transfer (PCET).[62] One of the main attractions of electrochemical CO2 reduction (ECO2R) is coupling the system with renewable energy sources to further curb CO2 emissions. The most attractive method of ECO2R is to use aqueous solutions containing inorganic salts, but issues arise such as low solubility of CO2 and high concentrations of H + . Moreover, the low energy barrier required for hydrogen evolution reaction (HER) often results in lower efficiency and selectivity for the desired products.[63,64] This has resulted in studies exploring nonaqueous and ionic liquid electrolytes, where CO2 solubility is improved and HER is suppressed. However, these systems have issues of 17 their own, such as, higher cost, toxicity and safety hazards and lower conductivities compared to aqueous solutions limiting their practical applications. Another approach to mitigating HER is the catalyst utilized.[65] A plethora of systems have been studied and as a result several products have been observed with some shown below [66]: CO2 + 2H + + 2e - → HCOOH E0 = -0.25 V (13) CO2 + 2H + + 2e - → CO + H2O E0 = -0.11 V (14) CO2 + 4H + + 4e - → CH2O + H2O E0 = -0.07 V (15) CO2 + 6H + + 6e - → CH3OH + H2O E0 = 0.02 V (16) CO2 + 8H + + 8e - → CH4 + H2O E0 = 0.17 V (17) 2CO2 + 10H + + 10e - → C2H6O2 + H2O E0 = 0.06 V (18) 2CO2 + 12H + + 12e - → C2H5OH + H2O E0 = 0.09 V (19) 2CO2 + 12H + + 12e - → C2H4 + 4H2O E0 = 0.08 V (20) 3CO2 + 16H + + 16e - → C3H6O + 5H2O E0 = 0.09 V (21) Even though the catalyst used is a major factor in product formation other aspects also play a role such as electrolyte, pH and applied potential. Furthermore, some general knowledge of products that can be obtained with certain metals has been gained over the years as shown in Figure 1.9, however new design of catalyst, changes in morphology, catalyst supports could yield different products and improve selectivity and efficiencies. 18 Figure 1.9. Typical products for ECO2R based on electrocatalyst. Adapted from [67] 1.7. Scope of Dissertation The main goals of this dissertation have been to explore new and commercial catalysts for their use in electrochemical devices; particularly PEMFCs, AFCs and ECO2R. Chapter 2 focuses on utilizing a partially fluorinated carbon (CFx) as a catalyst support for PEMFCs and various liquid AFCs and compare fuel cell performance to commercial (XC-72R) carbon support. The stability of CFx supported platinum catalysts for oxygen reduction reaction, stability and catalytic activity for hydrogen oxidation reaction were studied in half-cell and full fuel cell experiments. Moreover, the catalytic activity of palladium nanoparticles supported on CFx for oxidation of various liquid fuels was also studied in half-cell and fuel cell experiments. Chapter 3 covers the investigation of several palladium-based catalysts for use in alkaline direct formate fuel cells 19 (ADFFC). Furthermore, a formate fuel cell with a cation membrane was utilized in which electricity and a hydroxide base were coproduced. Chapter 4 focuses on developments in alkaline direct methanol fuel cells (ADMFC), first several ionomers for the anode electrode were studied in both half-cell and full fuel cell configurations. Then work on nickel-based catalysts for the electrooxidation of methanol was pursued. Finally, chapter 5 covers the electrochemical reduction of CO2 to ethylene glycol utilizing a single atom copper catalyst. 1.8. References (1) Kahan, A. 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Power Sources 2021, 506 (June), 230215. https://doi.org/10.1016/j.jpowsour.2021.230215. 29 Chapter 2: Partially Fluorinated Carbon Supported Catalysts for Improved Proton and Anion Exchange Membrane Fuel Cells 2.1. Fluorinated Carbon Supports Developing the electrocatalysts plays a significant role in the practical implementation of fuel cells in everyday use. Currently, the most effective electrocatalysts are platinum-based catalysts supported on high surface area carbon. However, Pt is not without its drawbacks such as its low tolerance levels of CO and its high cost and low availability, which is a significant issue in terms of mass production.[1–3] Thus, significant efforts have been made to mitigate these issues. One such effort has been incorporation of more earth abundant metals to reduce the amount of incorporated platinum including core-shell catalyst design and platinum alloyed catalysts.[4–7] Altering the catalyst support has also been another avenue of improving the catalysts by utilizing other carbon allotropes such as carbon nanotubes, carbon aerogels, graphene and high or low surface area carbons and by doping the carbon with metals or heteroatoms.[8–14] Another major challenge of fuel cells is the degradation of the catalysts due to Pt oxidation, dissolution, agglomeration of nanoparticles and carbon corrosion.[15–17] The carbon supports undergo corrosion which result in catalyst particle detachment and agglomeration leading to loss in fuel cell performance or failure.[18–22] Several strategies have been employed to mitigate the corrosion of carbon from surface modifications, changing architecture design and altering the fuel cell testing procedures.[23–27] The functionalization of carbon supports has also been utilized to improve the stability and electronic properties of the carbon supports. Even though N doping of carbon has been the most 30 extensively studied heteroatom doped carbon support, recently partial fluorination of carbon has gained significant interest.[28–31] The high electronegativity of fluorine has been demonstrated to result in strong bond formations with transition metals. Moreover, the high covalent bond energy of the C-F bond results in high thermodynamic stability.[32–35] Furthermore, the incorporation of fluorine changes the electronic structure of the carbons and has been beneficial for several applications such as electrochemical CO2 reduction, cathode materials in lithium-ion batteries, supercapacitors, and sensors.[36–42] Sun et al. were the first to demonstrate fluorinated carbon materials could provide viable catalysts for ORR in both acidic and alkaline media.[43,44] Since then, several groups have studied heteroatom doped carbon nanostructures with fluorine.[45–48] Berthon-Fabry et al. demonstrated the use of fluorinated carbon aerogels for ORR, in which the fluorination process was carried out either prior to or after the deposition of platinum. They found that the fluorination must be done prior to deposition of platinum or else amorphous platinum is formed resulting in decreased ORR catalytic activity.[49] Our group has employed commercially available partially fluorinated carbon (CFx, where x is 0.09-0.11), which has electrical conductivity similar to non- fluorinated XC72R, as a platinum support for the cathode catalyst with a ~40wt% platinum loading.[50] This led to enhanced performances in both proton exchange membrane (PEM) based direct methanol fuel cells (DMFCs) and hydrogen fuel cells. Furthermore, Chatenets group has studied the role of partially fluorinated carbon supported platinum towards ORR, again with 40wt% of platinum. They found no difference in stability under simulated “load cycling”, however under a start-stop procedure they found the fluorinated carbon supported catalysts to retain higher ECSA and attributed this to mitigation in carbon corrosion.[51,52] Li et al. provided further insight into the improvements in stability of fluorinated carbon supports and determined that the 31 improvement arises from improved charge transfer between the platinum and the fluorinated carbon and, similar to the studies of Chatenet et al., improved resistance to carbon corrosion from introduction of C-F bonds as opposed to only C-O and C-C bonds in nonfluorinated carbons. [51,53] Furthermore, fluorination of graphene with Ni nanoparticles were recently shown to improve the oxidation of methanol under alkaline conditions.[54] It was shown that the fluorination aided in OH - accessibility and decreased the energy barrier for methanol oxidation. Furthermore, the work by Glass studied the role of various platinum loadings on CFx for ORR.[55] Glass demonstrated that 30wt% Pt on CFx provided the greatest improvement over non fluorinated carbon supported Pt in PEMFCs due to improved electrochemically active surface area, water management and mass transport of reactants to the Pt. In this chapter the work of Glass, a previous graduate student in our group, is expanded on by extensively studying the stability of the prepared catalysts to observe the benefits provided by the CFx support. Furthermore, given the improvements in ORR with CFx, the catalysts were also studied for hydrogen oxidation reaction (HOR). Extensive studies of the catalysts in H2-PEMFCs were also completed to demonstrate the practical application of CFx. Moreover, Pd based nanoparticles were synthesized on CFx and utilized as electrocatalysts for a myriad of liquid fuels under alkaline conditions. Thus, demonstrating the versatility of partially fluorinated carbon supported catalysts in both proton and anion exchange membrane fuel cells. 2.2. Platinum Supported on Partially Fluorinated Carbons for Improved Performance and Stability in H2 Proton Exchange Membrane Fuel Cells 2.2.1. Experimental Catalyst Synthesis 32 The prepared catalysts were obtained by impregnation and reduction with NaBH 4.[55] First, Vulcan Carbon (XC72R, FuelCellStore) and CFx (9-11 wt% fluorine loading, Advanced Research Chemicals Inc., resistance less than 10 Ω) were vigorously stirred and sonicated in separate 100 mL Millipore water (Direct-Q UV, 18.2 M ) until well dispersed. Then an appropriate amount of 1g/mL aqueous solution of H2PtCl6 (Sigma Aldrich) was added to obtain loadings of 10 and 30 wt% platinum, referred to as Pt-1 and Pt-2, respectively. The pH of the solution was adjusted to neutral by adding 1.0 M NaOH (Macron). The solution was then heated to 80 ℃, reduced with an aqueous solution of NaBH4, stirred at 80 ℃ for 2 h, and then stirred at room temperature overnight. The obtained catalyst powder was then washed and centrifuged with Millipore water until the supernatant had reached a neutral pH, the powders were then dried in an oven at 65 ℃. Catalyst Characterization The N2 adsorption and desorption measurements were carried out on a Quantachrome Nova 2200E surface area and pore volume analyzer; the specific surface area and pore size distribution were determined by the Brunauer-Emmett-Teller (BET) and Barrett-Joyner-Halenda (BJH) analysis, respectively. SEM and EDS measurements were performed on a JEOL JSM-7001F electron microscope with an acceleration voltage of 20 kV. TEM measurements were performed on a JEOL JEM-2100F electron microscope with an acceleration voltage of 200 kV. Powder X- ray Diffraction (XRD) measurements were performed on a Rigaku X-Ray diffractometer using Cu-K (0.15458 nm) radiation with a scan rate of 6°/min. Thermogravemetric analysis (TGA) was performed on a TGA-50 Thermogravimetric Analyzer (Shimadzu). The catalysts were heated up at a rate of 10 ℃/min under air. Contact angle measurements were obtained by a Rame-hart Contact Angle Goniometer. 33 Half-Cell Measurements The half-cell tests were conducted with a standard three electrode setup, with a glassy carbon rotating disk electrode (GCE-RDE, 0.195 cm 2 ), Pt wire and Hg/HgSO4 (0.5 M H2SO4 filling solution, 0.68 V vs. RHE) as the working, counter and reference electrodes, respectively. The measurements were obtained on a Solartron SI 1287 Potentiostat. The catalyst inks were prepared by adding 1 mg of catalyst to a solution consisting of 100 µL of isopropyl alcohol, 100 µL of ethanol and 800 µL Millipore water, and 10 mg of 5 wt% Nafion ionomer solution (Ion Power). The solution was sonicated and 20 L was pipetted onto the working electrode and dried. The cell was purged with ultra-pure Ar (99.999%) for 15 min before the experiments, CV scans under N2 were performed from 0 and 1.28 V vs. RHE. For the ORR measurements, the cell was purged with ultra-pure O2 (99.994%) for 15 min. Linear sweep voltammetry (LSV) measurements were performed by scanning the potential from 1.18 V to 0.18 V vs. RHE at 5 mV/s at rotation rates of 400, 800, 1200, 1600, 2000, and 2400 RPM. Start-stop stability measurements were performed under Ar by scanning the potential from 1.0 to 1.5 V vs. RHE at 0.5 mV step increments and holding each potential for 3 seconds. Further durability measurements were conducted under O2 flow and cycling 3000 CV scans from 0.45 to 0.95 V vs. RHE. For the HOR measurements, the cell was purged with ultra-pure H2 (99.999%) for 30 minutes. LSV measurements were performed by scanning the potential from -0.08 V to 0.115 V vs. RHE at 5 mV/s at rotations of 400, 800, 1200, 1600, 2000 and 2400 RPM. Start-stop stability measurements were performed under Ar by scanning the potential from 0.17 to -0.27 V vs. RHE at 0.5 mV step increments and holding each potential for 3 seconds. Further durability measurements were conducted under H2 flow and cycling 3000 CV scans from -0.03 to 0.22 V vs. RHE. 34 Membrane Electrode Assembly (MEA) and Fuel Cell Measurements Catalyst inks were made from a 1:3:5 weight ratio of catalyst to Millipore water to 5% Nafion ionomer solution, respectively, for both the cathode and anode. The inks were then sonicated and painted onto a 5 cm 2 Toray Carbon Paper electrodes (E-TEK, TGH-060, 10% wet proofing). For HOR, Pt-1 or Pt-2 catalysts were painted onto the anode electrode until a platinum loading of 0.5 mg/cm 2 was obtained. For the cathode 40% Pt/C was painted onto the electrode until a platinum loading of 1.6 mg/cm 2 was obtained. Furthermore, an MEA with Pt-1 and Pt-2 as the anode and cathode electrodes, respectively, was also prepared with a platinum loading of 0.5 mg/cm 2 of platinum on both anode and cathode electrodes. The electrodes were then dried in an oven at 110 ℃. The proton form of Nafion-211 was prepared and sandwiched between the cathode and anode electrodes and hot pressed at 140 ℃ for 5 min at 500 psi. The MEA was then hydrated overnight at room temperature. The fuel cell polarization measurements were performed using a Fuel Cell Test System 890B (Schribner Associated). The MEAs were tested at an ambient cell temperature (~23 ℃) and 60 ℃. Humidified hydrogen was passed through the anode compartment at 100 mL/min at 85 ℃, while oxygen was passed through the cathode compartment at 50 mL/min at 85 ℃. The relative humidity for each of the gases was 100%. Electrochemical impedance spectroscopy (EIS) measurements were performed with a Solartron SI 1287 and SI 1260 Impedance-Phase Analyzer. EIS measurements were taken a 0.4 V with an amplitude of 10 mV at a frequency range of 100 kHz to 100 mHz. 35 2.2.2. Results and Discussion Figure 2.1. A) Mesoporous size distribution curves for XC72R and CFx and B) corresponding N2 adsorption-desorption isotherms of XC72R and CFx. C) TGA curves for the prepared catalysts under an air atmosphere. Table 2.1. Summary of BET and Pore Volume for the carbon supports SA BET (m 2 /g) Pore V olume (cm 3 /g) XC72R 194 0.232 CFx 174 0.574 The specific surface area of the CFx and XC72R supports were determined via BET, shown in Figure 2.1A and B with the values reported in Table 2.1, there is a slight decrease in surface area for the CFx compared to the XC72R. Furthermore, the pore volume for the CFx was larger than XC72R due to the disruption of the carbon structure by introduction of fluorine, as shown previously.[49] TGA tests were performed on the catalysts to confirm the platinum 36 loading on each of the catalysts shown in Figure 2.1C and calculated similar to a previously reported procedure.[55] TEM images of the catalysts were obtained to observe any differences in metal morphology (Figure 2.2). The average particle size of the Pt was found to be 3.712±1.832, 3.58±1.322, 4.640±1.626, and 3.654±1.222 for Pt-1F, Pt-2F, Pt-1C and Pt-2C, respectively. The trend of the nanoparticle size is similar to that obtained previously, where the Pt size for the CFx supported catalyst is smaller than for the XC72R supports. The difference in size could be due to the hydrophobic nature of the CFx aiding in the particle dispersion during the impregnation process. 37 Figure 2.2. TEM images of the catalysts: A) Pt-1F; B) Pt-1C; C) Pt-2F; D) Pt-2C XRD patterns were obtained and shown in Figure 2.3A. There are three characteristic peaks at 2 values of: 40.0°, 46.6°, and 67.9° signifying the face centered cubic (fcc) structure of platinum with crystallographic indices of (111), (200), and (220), respectively. The broad peaks at a 2 value of 25.2° signifies the (002) graphitic phase of both the XC72R and CFx supports. The patterns for XC72R and CFx display very similar features showing the fluorination of the high surface area carbon does not affect the overall crystallographic morphology of the platinum or carbon support. Previous reports on fluorination of carbon also show no significant differences in the XRD patterns for the fluorinated and nonfluorinated carbons. This could arise from the minor amount of fluorine present on the carbon which does not result in a drastic change in the crystallographic morphology. However, previous reports have shown that the catalyst structure does change if the fluorination treatment is done after synthesis of the Pt, but a loss in catalytic activity is observed.[49,53] 38 Figure 2.3. A) XRD pattern of the catalysts on CFx and XC72R supports. B) CV scans of the catalysts in 0.5 M H2SO4 solution at 20 mV/s under Ar. The solid lines indicate the CFx support; the dashed lines indicate the XC-72R support. The electrochemical performance of the prepared catalysts was first studied in H2SO4 under Ar gas. CV scans were obtained to determine the electrochemical surface area (ECSA) of the catalyst from hydrogen adsorption region (Figure 2.3B). The ECSA was calculated from the area above the curve in the potential region between 0 and 0.4 V and subtracted from the double layer capacitance portion of the CV near 0.4 V. The ECSA was calculated from this total charge and based on the charge required for hydrogen desorption of 210 C/cm 2 .[56] The calculated ECSA values were 74.6, 51.9, 44.3, and 33.5 m 2 /g for Pt-1F, Pt-1C, Pt-2F and Pt-2C, respectively. The ECSA value of the CFx supported catalysts was larger than that of the XC72R supported catalysts. Moreover, the ECSA of the Pt-1 catalysts was greater than the Pt-2, this could arise from the smaller nanoparticle size and better Pt dispersion shown in the TEM micrographs. Another aspect that could aid in the ECSA improvement is the interaction between the tetrafluoroethylene 39 backbone of the Nafion binder and CFx allowing for the Pt to be more accessible. Figure 2.4. LSV scans of the catalysts in 0.5 M H2SO4 solution at 5 mV/s: A) of the Pt-2F; B) Pt- 2C. LSV scans for the Pt-2 catalysts were obtained to assess the catalytic activity towards ORR (Figure 2.4). The onset potential for ORR is similar for both catalysts, however there is a difference in the current density. The Pt-2F displayed a higher current density than the Pt-2C, likely arising from the improved ECSA and mass transport of reactants to the catalyst, similar to the previously reported results.[55] Furthermore, the stability of the catalysts was studied with two different stability experiments, a start-stop procedure, and an accelerated stress experiment. Start-stop experiments were conducted in the potential window of 1.0 V to 1.5 V vs. RHE and the ECSA as well as the LSV experiments were obtained before and after the start-stop experiments. The current density of the LSVs after the start-stop for the nonfluorinated carbon support were similar to the initial values, whereas the CFx supported catalyst displayed a slight increase in the current density (Figure 2.5). Further, the ECSA of the XC72R supported catalyst had a slight decrease whereas the CFx supported catalyst displayed an increase. The stability was further observed by cycling the catalyst in the potential window between 0.46 and 0.95 V vs. 40 RHE under a flow of O2 for 3000 cycles. A negative potential shift in the onset potential is seen in both catalysts, likely due to catalyst agglomeration, however the CFx supported catalyst did not decrease in current density, whereas the XC72R catalyst decreased significantly. Furthermore, the ECSA in both catalysts decreased, however the CFx catalyst had a loss of less than 20% of the original ECSA whereas the XC72R catalyst decreased by nearly 80% of the initial ECSA. The improved stability of fluorinated carbons arises from improved interaction between the metal and the carbon support and mitigation of carbon corrosion due to the presence of C-F bonds, as opposed to only C-C and C-O bonds in the XC72R catalyst.[52,53] The free dangling groups of carbon supports are often the most prone to corrosion and it has been shown that the C-F bonds formed on these free dangling groups result in a more robust support. Moreover, the metal to support interaction has been shown to increase with heteroatom doped carbons which further improves the stability of the catalyst.[13,51,53] Figure 2.5. A) LSV scans before and after stability experiments at scan rate of 5 mV/s at 1600 RPM. The solid lines indicate the CFx support; the dashed lines indicate the XC72R support. B) Change in ECSA after start stop and 3000 cycles. EDS mapping of the catalysts were obtained prior to and after the stability experiments (Figure 2.6). From the figure we can see that there is no change in the elemental distribution for 41 the CFx supported catalyst, except for sulfur present from the sulfuric acid electrolyte. For the XC72R supported catalyst there is a larger presence of sulfur and oxygen, the larger presence of oxygen indicates corrosion of the carbon. Moreover, there is some fluorine present with the XC72R supported catalyst due to the Nafion added while making the catalyst inks. To further assess the differences, TEM micrographs of the catalysts after the stability experiments were obtained (Figure 2.6). From the figure we can see there is a change in the particle size of the Pt for both catalysts. However, the agglomeration and increase in particle size is more prominent in the XC72R supported Pt. The CFx supported Pt has more particles that remained similar in size to those before the stability experiments and this arises from the improved interaction between the Pt and support as shown previously by a charge transfer between Pt and the fluorinated support resulting in improved stability.[53] 42 43 Figure 2.5. SEM and EDS images for A) Pt-2C; B) Pt-2C after 3000 cycles; C) Pt-2C after start- stop; D) Pt-2F; E) Pt-2F after 3000 cycles and F) Pt-2F after start-stop in ORR potential region. For elemental mapping red corresponds to C, green corresponds to oxygen, blue corresponds to fluorine, magenta corresponds to Pt and yellow corresponds to sulfur. 44 Figure 2.6. TEM micrographs of A) Pt-2C and B) Pt-2F after start stop stability and C) Pt-2C and D) Pt-2F after 3000 cycles. Given the improved activity and stability of the CFx supported Pt, the catalysts were also screened for HOR activity. The LSVs of the catalysts were obtained under a flow of H2, shown in Figure 2.7. The CFx supported Pt show slightly higher current densities than the XC72R supported catalysts, while the Pt-1F catalyst displayed the highest current density of the screened catalysts. 45 The improvement in HOR for the CFx supported catalyst arises from the improved ECSA and could also be due to improved H2 diffusion. Figure 2.7. LSV scans of the catalysts in 0.5 M H2SO4 solution at 5 mV/s: A) of the Pt-1F; B) Pt- 1C; C) Pt-2F; and D) Pt-2C. Furthermore, the stability of the catalysts was examined following similar procedures to the ORR stability experiments, shown in Figure 2.8. Start-stop measurements were completed from the potential window of 0.17 V to -0.27 V vs. RHE. After the start-stop measurement, there was a decrease in current density for all catalysts, however the decrease was much more significant for the XC72R supported Pt. Moreover, when observing the ECSA the CFx supported Pt had a loss between 5-15%, whereas the XC72R catalysts had a decrease in ECSA between 20-60%. The stability was also examined after 3000 cycles between -0.03 and 0.22 V vs. RHE, again a decrease in the current density in the LSVs was seen across all the catalysts, where a more drastic drop was 46 observed in the XC72R catalysts. Furthermore, the ECSA decreased 50-70% for the XC72R and 15-30% for the CFx catalysts. Figure 2.8. LSV scans of A) Pt-1and B) Pt-2 in 0.5M H2SO4 under H2 gas before and after stability experiments at scan rate of 5 mV/s at 1600 RPM. The solid lines indicate before stability, the dotted lines indicate after start-stop experiments and the dashed lines indicate after 3000 CV cycles. c) Change in ECSA after start-stop and 3000 cycles. Moreover, EDS mapping of the catalysts before and after the stability experiments were obtained (Figure 2.9 and 2.10). Pt-2C and Pt-1C both show fluorine present from the Nafion added to the catalyst inks. Furthermore, after stability experiments, all catalysts display sulfur from the electrolyte. The Pt-2 catalysts both show similar distribution of elements before and after the stability experiments. However, the Pt-2C does display a slightly higher amount of sulfur and oxygen present when compared to the Pt-2F. Similarly, the Pt-1C displays a larger presence of 47 sulfur and oxygen than the Pt-1F. Furthermore, the Pt-1F catalysts again show minimal differences in the elemental distribution between the initial catalyst and catalyst after stability experiments. 48 49 Figure 2.9. SEM and EDS images for A) Pt-2C; B) Pt-2C after 3000 cycles; C) Pt-2C after start- stop; D) Pt-2F; E) Pt-2F after 3000 cycles and F) Pt-2F after start-stop in HOR potential region. For elemental mapping red corresponds to C, green corresponds to oxygen, blue corresponds to fluorine, magenta corresponds to Pt and yellow corresponds to sulfur. 50 51 Figure 2.10. SEM and EDS images for A) Pt-1C; B) Pt-1C after 3000 cycles; C) Pt-1C after start- stop; D) Pt-1F; E) Pt-1F after 3000 cycles and F) Pt-1F after start-stop in HOR potential region. For elemental mapping red corresponds to C, green corresponds to oxygen, blue corresponds to fluorine, magenta corresponds to Pt and yellow corresponds to sulfur. TEM micrographs of the catalysts after the stability experiments were also obtained, shown in Figure 2.11. The Pt-1 catalysts displayed minimal changes in the nanoparticle sizes, although there was more Pt agglomeration seen in the Pt-1C. Moreover, the Pt-2 catalysts displayed significant changes. Pt-2C had a significant increase in Pt agglomeration and size in nanoparticles, whereas the Pt-2F had minor increase in average particle size. The larger difference between the Pt-1 and Pt-2 is due to the amount of Pt present on the carbon supports. With higher Pt loading 52 there is higher propensity for Pt agglomeration and thus the improvement of the Pt stability from CFx is more prominent at higher loadings. Nonetheless, the partial fluorination of carbon also aids in improving stability of the Pt under acidic conditions in the hydrogen oxidation region. 53 54 Figure 2.11. TEM micrographs of A) Pt-1C and B) Pt-1F after 3000 cycles and C) Pt-1C and D) Pt-1F after start stop experiments. TEM micrographs of E) Pt-2C and F) Pt-2F after 3000 cycles and G) Pt-2C and H) Pt-2F after start-stop experiments. To fully assess the efficacy of the prepared catalysts, the Pt-1 and Pt-2 catalysts were utilized as the anode electrode catalysts (Figure 2.12) in H2-PEMFCs. The synthesized catalysts were used at the anode and a commercial Pt/C catalyst was used at the cathode electrode. With the Pt-1F and Pt-2F catalysts, an increase in peak power densities of 23% and 20%, respectively, was observed at 60 ℃ when compared to the XC72R supported catalysts. This increase in power densities could be attributed to the higher ECSA of the CFx support compared to XC72R as well as better diffusivity of H2 (improved mass transport). Furthermore, the Pt-1F catalysts displayed a higher power density than the Pt-2F, potentially arising from the higher ECSA and improved Pt utilization. 55 Figure 2.12. Polarization curves for the Pt-1, Pt-2 anode catalyst MEAs at: A) 25 ℃; B) 60 ℃. C) EIS curves for the Pt-1 and Pt-2 catalyst MEAs at 0.4 V at 60 ℃ and D) the equivalent circuit used to fit the spectra. To further understand the differences in catalytic activity, EIS experiments were conducted. Depressed semi-circles were obtained in the Nyquist plots, shown in Figure 2.12C and the spectra was fit to the equivalent circuit shown in Figure 2.12D. From the circuit, L1 and L2 represent the inductance of the system arising from the wires and equipment, Rmem and R ct represent the membrane and charge transfer resistances, respectively, CPE represents the pseudo- capacitive nature of the catalyst layers, and Ws represents the Warburg short element for the finite diffusion of reactants and products through the porous catalyst layers. There is a 30% decrease in the Rct of the CFx supported catalysts compared to the XC72R supported catalyst for both the Pt- 1 and Pt-2 at 60 ℃. Furthermore, there is a decrease in the Rmem for CFx supported catalysts compared to the XC72R supported catalysts, which further indicates an improved triple phase boundary and improved water management leading to higher conductivity. To confirm that the 56 CFx support aids in water management due to differences in hydrophobicity, contact angle measurements of the electrodes were measured, which are shown in Figure 2.13 and 14, with values of 135.35 ± 3.3, 136.98 ± 5.5, 140.69 ± 3.6 and 144.59 ± 4.2 for Pt-1C, Pt-2C, Pt1F and Pt- 2F, respectively. The CFx containing gas diffusion electrodes have a higher average contact angle measurement, confirming the greater hydrophobicity of the electrodes resulting in improved water management. Figure 2.13. Contact angle measurement image for a) Pt-1C and b) Pt-1F Figure 2.14. Contact angle measurement image for a) Pt-2C and b) Pt-2F Furthermore, a H2-PEMFC utilizing CFx supported Pt as both the anode and cathode electrodes was constructed and compared to a PEMFC with XC72R support for both electrodes in 57 Figure 2.15A and Figure 2.15B at 25 °C and 60 °C, respectively. The use of the catalysts in both the anode and cathode were carried out to ensure that the increased hydrophobicity would not result in too much water rejection and result in reduced fuel cell performance. However, an increase in peak power density of 20% was observed while using the CFx support compared to the XC72R supported PEMFC. Therefore, the hydrophobicity of the CFx aids in mitigating catalyst flooding but still allows for proper hydration of the membrane to maintain optimal ionic conductivity. Figure 2.15. A) Polarization curves for the MEAs with Pt-1 anode and Pt-2 cathode electrodes at 25 ℃ and 60 ℃. Solid lines indicate CFx supported and dashed lines indicate XC72R supported catalysts. Moreover, stability experiments were conducted holding the PEMFCs at constant currents of 200 mA/cm 2 and 1A/cm 2 shown in Figure 2.16. Both MEAs displayed no change in the potential while holding the current at 200 mA/cm 2 , however at 1 A/cm 2 there was a slight decrease in the potential and the XC72R supported catalysts showed large fluctuations throughout the allotted time. Moreover, polarization curves of the PEMFCs after the stability experiments were obtained and a reverse scan was performed in the mass transport loss region (Figure 2.16B). The MEA with CFx supported catalysts displayed minimal change in performance in the reverse scan, whereas 58 the XC72R supported catalyst displayed a larger decrease in performance. The minimal difference in performance for the MEA with CFx likely arises from improved water management from the hydrophobicity of the CFx supported catalyst, thus showing the benefits of utilizing CFx for both anode and cathode electrodes. Figure 2.16. A) MEAs held at constant currents with Pt-1 anode and Pt-2 cathode electrodes at 60 ℃. B) Polarization curves with Pt-1 anode and Pt-2 cathode electrodes after stability experiments with a forward and reverse scan. Solid lines indicate CFx supported and dashed lines indicate XC72R supported catalyst. 2.3. Development of Partially Fluorinated Carbon Supported Palladium Nanoparticles for Alkaline Direct Liquid Fuel Cells 2.3.1. Experimental Catalyst Synthesis The palladium on partially fluorinated carbon (Pd/CFx) catalyst was synthesized by impregnation and sodium borohydride reduction. 0.30 g of CFx from section 2.2 was dispersed in 100 mL of Millipore water under vigorous stirring and sonication for 1 hour. Then, 75 mg of PdCl2 (Sigma-Aldrich) was added to the solution and then stirred vigorously and sonicated for 1 hour. 59 The pH of the solution was adjusted to 10 by the addition of 3.0 M NaOH. The solution was then placed in an oil bath and heated to 80 ℃, followed by slowly adding 50 mL of 0.7 M NaBH 4 solution and the reaction mixture was stirred for 1 hour at 80 ℃. The resulting solution was then vacuum filtered and the obtained solid was washed with Millipore water before drying in the oven overnight. Palladium on non-partially fluorinated carbon was also prepared in a similar procedure with the difference of XC72R utilized instead of CFx, the catalyst will be referred to as Pd/C. Catalyst Characterization Powder X-Ray Diffraction (XRD) measurements were performed on a Rigaku X-Ray diffractometer with a Cu-Kα (0.154056 nm) radiation source and a scan rate of 6°/min from a 2θ value of 10° to 90°. Transmission Electron Microscopy (TEM) images were taken on a JEOL JEM 2100F with an acceleration voltage of 200 keV. Thermogravimetric analysis (TGA) was performed on a TGA-50 Thermogravimetric analyzer (Shimadzu) under air. X-ray Photoelectron spectroscopy (XPS) data was collected on a Kratos Axis Ultra DLD using a mono Al anode with a pass energy of 160 keV for the survey scan and 20 keV for the high-resolution scans. Half-Cell Measurements A standard rotating disk electrode (RDE) setup was utilized for the half-cell measurements. All measurements were recorded on a Solatron SI 1287 potentiostat. Catalyst ink solutions were prepared by adding 1.0 mg of catalyst to 1.0 mL solution composed of 10% isopropyl alcohol, 90% Millipore water and 30 mg of an alkaline anion binder (Tokuyama, AS-4). After sonication, 20 μL of the solution was pipetted onto a glassy carbon electrode (Pine Instruments) with a surface area of 0.195 cm 2 , in order to obtain 20 μg of catalyst. After drying, the electrode was placed in a three-cell testing system with the RDE as the working electrode, a platinum wire counter electrode 60 and a Hg/HgO reference electrode. Prior to electrochemical testing, the cell was purged with ultra- pure argon (99.999%) for 15 minutes. Membrane Electrode Assembly (MEA) and Fuel Cell Measurements For the catalyst layers, the prepared catalysts and Pt/C (Alfa Aesar, 40 wt % metal) were used for the anode and cathode, respectively. Either Pd/C or Pd/CFx and a quaternized poly(terphenylene) (TPN) binder were mixed at a catalyst to binder solution weight ratio of 1:3 in 200 µl of Millipore water and 200 µl of isopropyl alcohol followed by bath sonication for 8 minutes and probe sonication for 2 minutes. The catalyst inks were painted onto a 4 cm 2 teflonized carbon cloth gas diffusion layer (GDL) until a mass loading of Pd of 0.7 mg/cm 2 was acheived for the anode gas diffusion electrode (GDE). The cathode GDE was obtained by mixing Pt/C with Millipore water and Nafion (5 wt%) ionomer at a ratio of 1:10:3, followed by bath sonication for 8 minutes and probe sonication for 2 minutes. The catalyst ink was painted on a teflonized carbon paper with a microporous layer (22BB, SGL Carbon) until a Pt mass loading of 1.3 mg/cm 2 was achieved. The prepared anode GDEs were placed in 1 M NaOH solution for 20 h and then washed with Millipore water. A TPN membrane was sandwiched in between the anode and cathode GDEs and placed into a fuel cell hardware. The fuel cell polarization measurements were performed using a Fuel Cell Test System 890B (Schribner Associates). Prior to testing, a solution of 0.5 M KOH was passed through the cell for 10 minutes followed by Millipore water for 10 minutes, at room temperature. The cell temperature was raised to 60 ℃ and a 1M solution of the fuel with 1 M KOH was delivered through the anode compartment at 5 mL/min and recirculated to the fuel reservoir and humidified oxygen was delivered to the cathode compartment at 100 mL/min. The open circuit voltage (OCV) was monitored for 20 minutes, then the cell was held at 0.5 V for 20 minutes and then held at 0.2 V for 61 10 minutes. The cell temperature was increased to 80 ℃ and the fuel reservoir was replaced with fresh fuel and fuel cell measurements were obtained. 2.3.2. Results and Discussion The metal loading on the prepared catalysts was first determined via TGA and shown in Figure 2.17A, the loading of the metal for both catalysts was near the target metal loading of 20 wt%. The XRD patterns of the prepared catalysts were then obtained and shown in Figure 2.17B. The patterns for the two catalysts are similar and both show diffraction peaks near 25°, 40°, 46°, 68°, 82° and 86°, which correspond to graphite (002) and palladium (111), (200), (220), (311), and (222) Miller indices, respectively. There is a minor negative shift in the peak positions for the CFx supported Pd, which could indicate a difference in the catalyst structure. Moreover, the (111) peak was utilized to determine the crystallite size of the catalysts using the Debye-Scherrer equation and found to be 11 ± 1 nm and 5.8 ± 0.6 nm for Pd/C and Pd/CFx, respectively. Thus, the partial fluorination of the carbon results in a smaller crystallite size, which could arise from improved metal to support interaction during the particle synthesis, resulting in less agglomeration. Furthermore, Braggs law was utilized to determine the lattice spacing of the (111) peak and was calculated to be 0.23 nm for both catalysts. 62 Figure 2.17. A) TGA plot and B) XRD pattern of the prepared catalysts. TEM micrographs were obtained to further analyze any differences in the surface morphology of the catalysts, shown in Figure 2.18. The average particle size of the catalysts was found to be 4.26 ± 1.2 and 8.72 ± 2.4 nm for Pd/CFx and Pd/C, respectively. Similar to what was observed in section 2.2 with Pt/CFx catalysts, the CFx supported catalysts are smaller, however the difference is much more prominent with the Pd/CFx. Moreover, the morphology of the catalysts is different with the Pd/C displaying more agglomeration and less spherical shape, whereas the Pd/CFx catalysts are more spherical. Furthermore, the trend for the size of the nanoparticles is similar to the crystallite size determined from the XRD, where the Pd/C is nearly twice the size of the Pd/CFx. 63 Figure 2.18. TEM micrographs of A) Pd/CFx and B) Pd/C To further analyze the catalysts XPS analysis was completed for the Pd/CFx and the Pd 3d core spectrum is shown in Figure 2.19A. The binding energies of Pd 3d can be resolved into two doublets corresponding to the 3d5/2 and 3d3/2 spin-orbital coupling. The curves are deconvoluted into two pairs of binding energies corresponding to the Pd 0 and Pd-oxide. Moreover, the larger intensity of the Pd 0 peaks indicates a larger presence of metallic Pd phase and very minimal amount of Pd-oxide. 64 Figure 2.19. A) XPS PD 3d spectra of the Pd/CFx catalysts. B) CV scans of the prepared catalysts in 1M KOH at a scan rate of 20 mV/s under flow of N2. To study the electrochemical activity of the prepared catalysts, CV scans in 1 M KOH electrolyte were first obtained, shown in Figure 2.19B. Both catalysts show similar features, in the forward scan Peak I is ascribed to the oxidation of hydrogen and removal of adsorbed hydrogen, peak II is due to oxygen evolution reaction and formation of Pd-oxides, peak III is due to the reduction of any formed oxides on the surface and peak IV is due to hydrogen evolution and hydrogen adsorption. The reduction peak (III) can be utilized to obtain the electrochemically active surface area (ECSA) of the catalysts, by integrating the area of the reduction peak and then utilizing equation 1: [57] ECSA=Q/Sl (1) Where Q is the charge for the reduction of the oxidized Pd, S is the proportionality constant (405 µC/cm 2 ), and l is the catalyst loading. The calculated ECSA values are 26.42 m 2 /g and 39.8 m 2 /g for Pd/C and Pd/CFx respectively; the CFx supported catalyst displays a significantly larger ECSA than the Pd/C. This likely arises from the smaller nanoparticle size and improved dispersion of the nanoparticles on the CFx. 65 The catalytic activity for oxidation of the liquid fuels (formate, methanol, glycerol, and ethylene glycol) for the prepared catalysts was then studied and CVs are shown in Figure 2.20. All CVs present two oxidation peaks corresponding to the oxidation of the fuels occurring between - 0.7 and 0.1 V vs. MMO. The peak in the forward scan is the primary oxidation of the fuel which increases in current until the surface is covered in oxide species or intermediates, deactivating the surface. Then in the reverse scan these oxide and intermediate species are reduced and removed from the surface resulting in reactivation of catalysts and an increase in the current.[58,59] Moreover, from the peak current in the forward scan (If) and reverse scan (Ib), the reactivation coefficient (Ib/If) could be determined, shown in Table 2.2. From the obtained values it is observed that, with the exception of methanol, the CFx support aids in reactivation of the catalysts, thus the CFx aids in removal of oxide and intermediate species. Moreover, the glycerol and ethylene glycol CVs display a third peak at the higher potentials (over 0.3 V vs. MMO), which likely arises from oxygen evolution or oxidation of intermediates on the surface. Irrespective of the fuel, the Pd/CFx displayed an increased peak current density compared to the Pd/C, indicating improved catalytic activity. This improvement likely arises from the increased ECSA. 66 Figure 2.20. CV scans of the prepared catalysts at scan rate of 20 mV/s under N2 with A) 1M formate and 1M KOH; B) 1M methanol and 1M KOH; C) 1M glycerol and 1M KOH and D) 1M ethylene glycol and 1M KOH To further understand the differences in catalytic activity, CV scans were obtained at various scan rates (v), shown in Figure 2.21. As expected, an increase in the peak current is observed as the scan rate increases. Furthermore, from the CV scans two plots are obtained: one plotting the potential of the peak current (Ep) at the various scan rates against the natural log of the scan rate (ln(v)) and the second plot from the peak current (Ip) against the square root of the scan rate (v 1/2 ) (Figure 2.22). Both plots for all the fuels display a linear correlation, moreover the linear correlation of the Ep vs ln(v) confirms that all the fuel oxidations are diffusion limited processes. 67 68 Figure 2.21. CV scans of A) Pd/C and B) Pd/CFx in 1M formate and 1M KOH; C) Pd/C and D) Pd/CFx in 1M methanol and 1M KOH; E) Pd/C and F) Pd/CFx in 1M glycerol and 1M KOH; G) Pd/C and H) Pd/CFx in 1M ethylene glycol and 1M KOH at various scan rates. Furthermore, this linear relationship can be utilized to calculate the value of αn’ (Table 2.2) using equation 2: E p ln v = RT 2F(αn’) (2) Where in equation 2, R is the gas constant, T is temperature, F is Faraday constant and α is the charge transfer coefficient and n’ is the electrons in the rate determining step (RDS). The αn’ values can then be utilized to obtain the diffusion coefficient from equation 3 and the slope from the Ip vs v 1/2 plots shown in figure 2.22. I p v 1/2 = 2.99 X 10 5 n(αn’) 1/2 CD 1/2 (3) From equation 3, C is the concentration of fuel in solution, D is the diffusion coefficient and n is the number of electrons in the entire electrochemical process. Previous studies demonstrate the diffusion of methanol to the catalyst varied by altering the carbon nanotubes supporting the catalysts, thus partial fluorination of the carbon would likely affect the diffusion as well.[60] From 69 the values obtained and shown in table 2.2, there is an increase in the diffusion coefficient from all the fuels when comparing the Pd/CFx to the Pd/C catalysts. Thus, indicating the partially fluorinated catalyst support better aids in utilization of liquid fuels improving mass transport resulting in improved oxidation currents. 70 Figure 2.22. Plot of Forward peak current vs square root of scan rate and plot of peak potential vs natural log of scan rate in A and B) 1M formate and 1M KOH; C and D) 1M methanol and 1M KOH; E and F) 1M glycerol and 1M KOH and G and H) 1M ethylene glycol and 1M KOH. Table 2.2. Reactivation, charge transfer and diffusion coefficients for the prepared catalysts with various fuels. Reactivation Coefficient αn' Diffusion Coefficient (x10 -8 ) Pd/CFx Pd/C Pd/CFx Pd/C Pd/CFx Pd/C Formate 1.28 0.97 0.98 0.86 240 120 Methanol 0.19 0.23 0.97 0.38 2.6 2.5 71 Glycerol 0.56 0.23 0.94 0.90 1.8 0.21 Ethylene Glycol 0.46 0.25 0.96 0.94 7.2 1.6 Tafel plots were obtained for the catalysts from linear sweep voltammetry scans at a scan rate of 0.5 mV/s and are shown in Figure 2.23. The Tafel slopes were obtained and are provided in Table 2.3, except for the Tafel slope of glycerol. This is due to glycerol displaying two Tafel slopes; for Pd/C Tafel-a was between -0.05 V to 0.00 V vs. MMO and Tafel-b occurred between 0.00 to 0.05 V vs. MMO and for Pd/CFx Tafel-a was between -0.014 V to 0.05 V vs. MMO and Tafel-b was between 0.05 V to 0.07 V vs. MMO. Tafel-a slope was 252.9 mV/dec and 216.8 mV/dec for Pd/C and Pd/CFx, respectively. Tafel-b slope was 144.1 mV/dec and 91.9 mV/dec for Pd/C and Pd/CFx, respectively. The Tafel slope values for the CFx supported catalysts were all lower than the XC72R supported catalysts for all fuels indicating faster charge transfer kinetics due to the higher accessibility to the fuel (mass transport) and the increased ECSA from the CFx. 72 Figure 2.23. Tafel plots for the prepared catalysts under A) 1M formate and 1M KOH; B) 1M methanol and 1M KOH; C) 1M glycerol and 1M KOH; and D) 1M ethylene glycol and 1M KOH. Furthermore, the role of temperature was also studied in the half-cell experiments. CVs were completed from 5 to 35 ℃ for all the fuels, shown in Figure 2.24. The peak current density increases with an increase in temperature for both catalysts across all fuels due to improved kinetics. However, the Pd/CFx maintains a higher peak current density at all temperatures, when compared to Pd/C. 73 Figure 2.24. CV scans of A) Pd/C and B) Pd/CFx in 1M formate and 1M KOH; C) Pd/C and D) Pd/CFx in 1M methanol and 1M KOH; E) Pd/C and F) Pd/CFx in 1M glycerol and 1M KOH; G) Pd/C and H) Pd/CFx in 1M ethylene glycol and 1M KOH at various temperatures. The current density for each temperature at -0.03 V vs. MMO was used to formulate Arrhenius plots shown in Figure 2.25 in order to obtain activation parameters for both catalysts (Table 2.3) utilizing equation 4: Ln (Ip) = Ln (A) - 𝐸 𝑎𝑝𝑝 𝑅𝑇 (4) In equation 4, the apparent activation energy is given by Eapp, A is the pre-exponential constant, R is gas constant and T is the operating temperature.[61] The Pd/CFx displayed a decrease in activation energy across all fuels compared to Pd/C owing to improved removal of unreactive species on the surface and diffusion of the fuels to the catalysts. 74 Figure 2.25. Arrhenius plots for Pd/CFx and Pd/C in A) 1M formate and 1M KOH; B) 1M methanol and 1M KOH; C) 1M glycerol and 1M KOH and D) 1M ethylene glycol and 1M KOH. Table 2.3. Obtained activation energies, onset potentials and Tafel slopes for the various fuels with the prepared catalysts. Eapp (kj/mol) Onset Potential (V vs. MMO) Tafel Slope (mV/dec) Pd/CFx Pd/C Pd/CFx Pd/C Pd/CFx Pd/C Formate 9.0 11.5 -0.74 -0.64 155.8 211.5 Methanol 24.1 24.2 -0.31 -0.30 56.3 166.5 Glycerol 17.4 21.1 -0.31 -0.31 - - Ethylene Glycol 22.1 24.7 -0.37 -0.40 72.3 129.1 75 In order to fully assess the efficacy of the catalysts, the catalysts were utilized in the anode electrode for alkaline direct liquid fuel cells. The various fuel cells were studied at three different concentrations of fuel (1-3 M) while keeping the flow rates, concentration of base (6M KOH) and cell temperature constant (80 ℃). The polarization and power curves for the alkaline direct glycerol fuel cell (ADGFC) are shown in Figure 2.26. At 1M glycerol the Pd/CFx displayed a significantly higher peak power density when compared to Pd/C, however as the concentration increased to 2 M the power density of the Pd/CFx ADGFC increased sparingly while the Pd/C increased to a comparable peak power as the Pd/CFx containing ADGFC. Furthermore, both ADGFCs had a decrease in peak power when the concentration increased to 3 M. This decrease likely arises from an increase in fuel crossover and catalyst surface poisoning, which is also observed in the decreased OCV. 76 Figure 2.26. Fuel cell results for alkaline direct glycerol fuel cells with Pd/CFx and Pd/C anodes and Pt/C cathode electrodes and cell temperature of 80 ℃. With A) 1M glycerol; B) 2M glycerol and C) 3M glycerol with 6M KOH. The polarization and power curves for the alkaline direct ethylene glycol fuel cells (ADEGFC) are shown in Figure 2.27. At the three concentrations studied the ADEGFC with Pd/CFx produced a higher peak power compared to the Pd/C ADEGFC. Moreover, similar to the ADGFC, as the concentration of fuel increased to 3M a drop in peak power was observed for both ADEGFCs, this decrease could arise from mass transport issues, catalyst surface poisoning or crossover. Figure 2.27. Fuel cell results for alkaline direct ethylene glycol fuel cells with Pd/CFx and Pd/C anode and Pt/C cathode electrodes and cell temperature of 80 ℃. With A) 1M ethylene glycol; B) 2M ethylene glycol and C) 3M ethylene glycol with 6M KOH. The polarization and power curves for the alkaline direct methanol fuel cells (ADMFC) are displayed in Figure 2.28. At the three concentrations, the Pd/CFx ADMFC produced higher power 77 than the Pd/C ADMFC. Moreover, unlike the ADEGFC and ADGFC, the power increased as the concentration of methanol increased. At 3M methanol, the peak power produced was amongst the highest reported power densities for ADMFCs with Pd catalysts.[62,63] Figure 2.28. Fuel cell results for alkaline direct methanol fuel cells with Pd/CFx and Pd/C anode and Pt/C cathode electrodes and cell temperature of 80 ℃. With A) 1M methanol; B) 2M methanol and C) 3M methanol with 6M KOH. The polarization and power curves for the alkaline direct formate fuel cells (ADFFC) are shown in Figure 2.29. The peak power increased as the concentration of formate increased, similar to the ADMFC. Furthermore, the Pd/CFx ADFFC displayed a higher power compared to the Pd/C ADFFC. 78 Figure 2.29. Fuel cell results for alkaline direct formate fuel cells with Pd/CFx and Pd/C anode and Pt/C cathode electrodes and cell temperature of 80 ℃. With A) 1M formate; B) 2M formate and C) 3M formate with 6M KOH. The benefit of utilizing a partially fluorinated carbon support for Pd catalysts for direct liquid alkaline fuel cells can be observed from the fuel cell results. All the studied fuels displayed an increase in peak power when utilizing Pd/CFx compared to Pd/C anode electrode catalysts. This increase arises from the improved ECSA, kinetics and diffusion of reactants to the catalysts. Furthermore, the highest power of the fuels was obtained from ADFFCs due to the faster oxidation kinetics as has previously been shown.[64] 2.4. Conclusion Herein, the advantages of utilizing a partially fluorinated carbon support for both anode and cathode electrodes in fuel cells was demonstrated. The stability of Pt/CFx and Pt/C catalysts under acidic conditions was studied in both the ORR and HOR regions. The CFx supported catalysts 79 demonstrated improved stability in terms of ECSA retention and current density after two different stability experiments. The improved stability was due to the C-F bonds present in the CFx which improve the metal to support interaction as well as less carbon corrosion and lower Pt agglomeration overtime. Furthermore, the Pt/CFx was found to improve the HOR in of H2- PEMFCs and improve the water management of the fuel cell. Moreover, the Pt/CFx could be utilized as both the anode and cathode electrode catalysts and improve H2-PEMFCs performance by 20% when compared to Pt supported on nonfluorinated carbon. The use of CFx was also explored as a support for Pd catalysts in alkaline direct liquid fuel cells. Pd/CFx was observed to improve the catalytic activity of glycerol, ethylene glycol, formate and methanol oxidation compared to the Pd/C catalyst in half cell experiments. This improvement was found to be due to an increase in ECSA, smaller nanoparticles and increased accessibility of catalysts to the reactants (improved mass transport). The improvement of Pd/CFx was also observed in alkaline fuel cell studies, where the Pd/CFx demonstrated higher peak power densities compared to fuel cells constructed with Pd/C. 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Direct Formate Fuel Cells Among the liquid fuels, methanol, ethanol, ethylene glycol, urea and glycerol have gained the most attention.[1–11] These liquid fuels have also been the most commonly used fuels in alkaline direct fuel cells (ADFC), where faster kinetics for the oxidation of particular fuels and oxygen reduction reactions allow for the use of non-noble metals as catalysts.[7,12–18] Formic acid has also been used as a fuel in DLFCs and has advantages over other various alcohols such as higher theoretical cell potential, lower fuel crossover and higher power densities.[19–23] However, formic acid is corrosive, which leads to an unstable catalyst structure. Furthermore, high catalyst loadings are required for acceptable performance, and the system suffers from inherent mass transport limitations. Nevertheless, most of the challenges of formic acid fuel cells can be overcome by using aqueous formate salts as the liquid fuel in ADFCs. Direct formate fuel cells (DFFC) have recently been shown to be one of the most efficient and effective fuels under alkaline conditions.[1,24–28] Formate salts have also been approved for environmental use and can be produced by the reduction of carbon dioxide with high selectivity and efficiency.[29–31] Moreover, the formate oxidation reaction (FOR) has sufficiently fast kinetics such that the oxidation can occur in the absence of an additional electrolyte, unlike many ADFCs that require the addition of supporting electrolyte to obtain adequate power densities. Haan et al. have demonstrated that the FOR is independent of pH between 9 and 14 and were able to achieve 105 mW/cm 2 power density with a DFFC operating with 1 M HCOOK. Similarly, Zhao et al. were able to produce 130 mW/cm 2 power density with a DFFC running with 5 M HCOOK.[25,32] 91 Formates ability to operate in the absence of supporting electrolyte provides ample room for its integration in commercial use. The use of DFFCs in the absence of added hydroxide has also been studied in a cation exchange membrane fuel cell, where the cation from the formate salt is the charge carrier across the membrane.[33–35] Moreover, Li et al. demonstrated the use of the cation exchange membrane DFFC (CEM-DFFC) to simultaneously produce electricity and alkali hydroxide base in the cathode.[34] However, they did not quantify the produced base and assumed production of alkali hydroxide would occur at the cathode end. Thus, our group expanded on this work and demonstrated the integration of CO2 capture with alkali hydroxides and conversion to formate which was then utilized in a CEM-DFFC to produce electricity and regenerate the alkali hydroxide at the cathode. The hydroxide was quantified and utilized to recapture CO2.[31] Furthermore, Sun et al. utilized a CEM-DFFC to coproduce electricity and H2 on the cathode via water reduction.[35] Thus, demonstrating the practicality of FOR to generate electricity and other useful products. The use of formate in CEM-DFFC and the absence of added electrolyte adds importance into the choice of cation due to the interaction of the cation with the catalyst surface and transport across the membrane. Several studies have shown that alkali metal cations in electrolyte solution have an impact on electrocatalytic activity towards oxidation of carbon monoxide, urea, glycerol, ethylene glycol, methanol, benzaldehyde and the reduction of carbon dioxide and oxygen.[36–46] Previdillo et al. studied the role of cations for FOR on Pt surfaces and found the catalytic activity increased in the sequence of Li<Na<K.[47] However, studies have shown difference in electrochemical response to cations depending on the metal used.[39] Moreover, given the wide acceptance of Pd based catalysts being the most effective towards FOR and lack of mechanistic 92 insight on FOR on Pd based catalysts, it is important to understand the role of cations on Pd surface, especially in DFFCs where the cation is the charge carrier across the membrane. Palladium has been shown to be a more effective catalyst for the FOR in both half-cell and fuel cell experiments, when compared to platinum-based catalysts.[1] Moreover, the FOR on palladium catalysts has been shown to proceed predominantly through dehydrogenation to form CO2, whereas on platinum catalysts the reaction proceeds through dehydration to form CO which is known to poison the platinum catalyst surface, although to a much lesser extent in an alkaline than in an acidic environment.[48–51] Much effort has gone into improving the activity and durability of the Pd catalyst, mainly by designing Pd-based alloys such as; PdCu, PdAg, PdRh and PdCo.[52–56] The morphology of the Pd can also be tuned to improve the catalytic activity, for example Ding et al. recently prepared mesoporous free-standing Pd nanotube arrays which displayed a 6.5 time gravimetric current improvement when compared to commercial Pt/C.[57] Catalyst supports have been used to increase the catalytic activity of metal nanoparticles in fuel cells via effectively dispersing the metal nanoparticles as well as altering the electronic structure of the catalyst.[20,58] Some catalyst supports that are currently being investigated are carbon blacks, carbon nanotubes and metal oxide supports, such as cerium oxide.[48,59–61] Graphene is one of the most promising carbon supports due to its many characteristics such as high charge carrier mobility, capacitance and functionalization ability.[62–65] Thus, graphene has been integrated into a plethora of applications from energy devices to sensors.[13,16,66–68] Recently, graphene was shown to improve the catalytic activity of platinum for formate electooxidation by reducing the adsorption strength of CO onto the platinum surface.[69] However, there were no previous reports utilizing palladium on a graphene support for formate electrooxidation under alkaline conditions. Herein, we report the use of palladium on 93 reduced graphene oxide (Pd/rGO) support for the FOR in DFFCs in order to improve the catalytic activity of the Pd for more effective formate electrooxidation. The electrochemical behavior, kinetic parameters and fuel cell performance of the catalysts were investigated. Furthermore, Pd bimetallic catalysts with IrO2 were also prepared on XC72R and multiwalled carbon nanotubes (MWCNT) to observe differences in catalytic activity. The addition of the metal oxide and MWCNTs resulted in improved performance of FOR observed in both half-cell and DFFC studies. Finally, various formate salts were studied to gain insight into the role of the cation for FOR on PdC catalyst. Furthermore, the formate salts with various cations were used in a CEM-DFFC to produce electricity and alkali hydroxide. 3.2. Reduced Graphene Oxide Supported Palladium Nanoparticles for Enhanced Electrocatalytic Activity Towards Formate Electrooxidation in an Alkaline Medium 3.2.1. Experimental Catalyst synthesis Graphene Oxide (GO) was synthesized from graphite flakes (325 mesh, 99.8%, Alfa Aesar) using a modified Hummer's method;[70] 2.0 g of graphite powder was dispersed in 50 mL concentrated H2SO4 (Macron) under magnetic stirring for 1 hour. 5.0 g of KMnO4 (Macron) was then added while continuously stirring for 2 hours in an ice bath. Afterwards, the temperature was raised to 40 ℃ in an oil bath and the reaction mixture was stirred for 1 hour. Then, 100 mL of Millipore water (Direct-Q UV, 18.2 MΩ) was added and finally 20 mL of 30% H2O2 (Macron) was added to terminate the reaction. The resulting suspension was vacuum filtered and washed with 10% HCl and then Millipore water before being dried in an oven at 80 ℃, to obtain GO. 94 The palladium on reduced graphene oxide (Pd/rGO) catalyst was synthesized via impregnation via sodium borohydride reduction. 0.30 g of the previously synthesized GO was dispersed in 100 mL of Millipore water under vigorous stirring and sonication for 1 hour. Then, 75 mg of PdCl2 (Sigma-Aldrich) was added to the solution and then stirred vigorously and sonicated for 1 hour. The pH of the solution was adjusted to 10 by the addition of 3.0 M NaOH. The solution was then placed in an oil bath and heated to 80 ℃, followed by slowly adding 50 mL of 0.7 M NaBH4 solution and the reaction mixture was stirred for 1 hour. The resulting solution was then vacuum filtered and the obtained solid was washed with Millipore water before being freeze-dried. Catalyst Characterization Powder X-Ray Diffraction (XRD) measurements were performed on a Rigaku X-Ray diffractometer with a Cu-Kα (0.154056 nm) radiation source and a scan rate of 6°/min from a 2θ value of 10° to 90°. Raman measurements were performed on a Horiba XploRA ONE Raman microscope with a 532 nm wavelength laser excitation. Scanning Electron Microscopy (SEM) images were obtained from a JEOL JSM-7001F electron microscope with an acceleration voltage of 15 keV. Transmission Electron Microscopy (TEM) images were taken on a JEOL JEM 2100F with an acceleration voltage of 200 keV. Thermogravimetric analysis (TGA) was performed on a TGA-50 Thermogravimetric analyzer (Shimadzu). X-ray Photoelectron spectroscopy (XPS) data was collected on a Kratos Axis Ultra DLD using a mono Al anode with a pass energy of 160 keV for the survey scan and 20 keV for the high-resolution scans. Electrochemical Measurements 95 A standard rotating disk electrode (RDE) setup was utilized for the half-cell measurements. All measurements were obtained on a Solatron SI 1287 potentiostat. Catalyst ink solutions were prepared by adding 1.0 mg of catalyst to 1.0 mL solution composed of 10% isopropyl alcohol, 90% Millipore water and 30 mg of an alkaline anion binder (Tokuyama, AS-4). After sonication, 20 μL of the solution was pipetted onto a glassy carbon electrode (Pine Instruments) with a surface area of 0.195 cm 2 , in order to obtain 20 μg of catalyst. After drying, the electrode was placed in a three-cell testing system with the RDE as the working electrode, a platinum wire counter electrode and an Ag/AgCl reference electrode. Prior to electrochemical testing, the cell was purged with ultra-pure argon (99.999%) for 15 minutes. Electrochemical impedance Spectroscopy (EIS) measurements were performed on a Solatron SI1287 potentiostat and SI1260 impedance Phase Analyzer at a potential of -0.3 V vs. reference with an amplitude of 5 mV. Membrane Electrode Assembly (MEA) and Fuel Cell Measurements Single fuel cell assemblies were used to further assess the performance of the catalyst. Either commercially available Pd/C (Alfa Aesar, 20 wt% palladium) or the synthesized Pd/rGO were used in the MEA as the anode catalyst, while commercially available Pt/C (Alfa Aesar, 40 wt% platinum) was used as the cathode catalyst. Catalyst inks were formulated by mixing the catalyst with appropriate amounts of water and an alkaline anion binder (Tokuyama, 5% AS-4), 22 mg, 400 mg and 73 mg, respectively, and sonicated for 10 minutes. The resulting catalyst inks were then hand-brush painted onto 4 cm 2 non-teflonized carbon cloth until 0.8 mg/cm 2 of Pd was deposited for the anode and 1.6 mg/cm 2 of Pt was painted onto a 10% teflonized carbon cloth for the cathode. The MEA was prepared by sandwiching the anode and cathode catalyst layers to the anion exchange membrane (Tokuyama, A201) and hot pressing at 110 ℃ for 5 minutes at 500 kilograms of force. 96 The fuel cell polarization measurements were performed using a Fuel Cell Test System 890B (Schribner associated). The MEAs were tested at cell temperature of 30 and 60 ℃. A solution of 1 M HCOONa in 1 M NaOH at ambient temperature was delivered through the anode compartment at a flow rate of 2.5 mL/min while humidified oxygen was heated to 85 ℃ and passed through the cathode compartment at 100 mL/min. 3.2.2. Results and Discussion TGA tests were performed on the catalyst to determine the palladium loading. The target loading was 20% metal, while the actual loading was found to be approximately 19% (Figure 3.1). Figure 3.1B depicts the Raman spectra of Pd/rGO and GO. There are three main features for graphene: the D band at approximately 1350 cm -1 , the G band at approximately 1590 cm -1 and the 2D band at 2700 cm -1 . The D band measures defects and impurities at the edges of the graphene lattice, while the G band arises from the sp 2 carbon bond stretching in the graphene lattice and the 2D band can be associated with the number of layers of graphene.[71] The shift in the 2D band from 2707 cm -1 to 2695 cm -1 implies the presence of few layers of graphene in Pd/rGO. The intensity ratios of the D to G bands (ID/IG) can be utilized to compare the quantity of defects in the graphitic lattice relative to the sp 2 hybridized carbon domains. The ID/IG of the GO and Pd/rGO are 0.96 and 1.07, respectively; this increase confirms the partial reduction of GO to rGO.[71] The XRD patterns corresponding to GO and Pd/rGO are shown in Figure 3.1C. The XRD of the GO exhibits a peak at approximately 11°, which corresponds to the diffraction from (100) Miller index of GO. The presence of this peak confirms the presence of several kinds of oxygen- containing functional groups in the graphitic layers. For the Pd/rGO, there are peaks at 40.1°, 46.3°, 68.5°, 82.2°, which correspond to the (111), (200), (220), and (311) Miller indices of face- 97 centered cubic (fcc) Pd, respectively.[16] The average crystallite size of the catalysts was calculated using the Debye-Scherrer equation and based on the broadening of the (111) peak, the Pd/rGO shows an average size of 4.9 nm and Pd/C shows a size of approximately 4.3 nm. Using Bragg’s law, the lattice spacing for the (111) was calculated to be 0.23 nm for both Pd/rGO and Pd/C. Also, the peak corresponding to GO is diminished and a new peak is present at 25°, corresponding to the (002) Miller Index of rGO, further confirming the reduction of GO. Figure 3.1. Characterizations of the catalyst: A) The TGA profile for Pd/rGO catalyst. The TGA experiments were carried out in an air atmosphere with a heating rate of 10 ℃/min. B) Raman Spectra and; C) XRD patterns. The structure of the Pd/rGO catalyst was analyzed by SEM and TEM. Figures 3.2A and B show the SEM images of the GO and Pd/rGO, respectively. Figure 3.2A exhibits the consistency of few layer graphitic carbon with typical wrinkle behavior observed previously.[72] After reduction, the surface morphology had changed to translucent graphene layers decorated with Pd 98 nanoparticles. The Pd nanoparticles can further be analyzed in the TEM images shown in Figure 3.2C-F where the average diameter of the Pd nanoparticles was found to be 5.8 ± 1.8 nm for both carbon and rGO supported catalyst, which is comparable in size to Pd supported on rGO reported elsewhere.[20,73] However, the average particle size from the XRD data (4.3 nm) is smaller than this value but is within the standard deviation. The lattice fringes can be seen in Figures 3.3A and B corresponding to the fcc facets of Pd with a d-spacing of approximately 0.23 nm which corroborates with the obtained XRD data. Moreover, the particles are more effectively dispersed throughout the rGO support as opposed to the carbon supported particles. A) C) B) D) 99 Figure 3.2. SEM images of: A) GO; B) Pd/rGO; TEM image of Pd/C C) and D) and Pd/rGO E) and F). Figure 3.3. HRTEM images for A) Pd/rGO and B) Pd/C XPS experiments were carried out to further characterize the Pd/rGO catalyst. The survey scan (Figure 3.4A) shows peaks due to O, Pd, C and Si from the Si wafer used as a substrate. Figure 3.4B depicts the C1s XPS spectrum of the graphene lattice. A major peak is observed, which deconvolutes into the C=C, C-C, C-O and C=O binding species. The Pd 3d core level XPS spectrum of Pd/rGO is shown in Figure 3.4D. The binding energies of Pd 3d can be resolved into E) F) B) A) 100 two doublets corresponding to the 3d5/2 and 3d3/2 spin-orbital coupling. The curves are deconvoluted into two pairs of binding energies corresponding to the Pd 0 having peaks at 335.6 and 340.9 eV and Pd-oxide having peaks at 337.3 and 341.7 eV. Moreover, the larger intensity of the Pd 0 peaks indicates a larger presence of metallic Pd phase. Figure 3.4. XPS spectra of: A) Survey scan; B) C1s of Pd/rGO; C) C1s of GO and; D) Pd 3d of Pd/rGO. Cyclic voltammetry (CV) experiments were carried out for each of the catalysts in 1 M NaOH as shown in Figure 3.5A. The forward scan contains two processes: the hydrogen desorption region, labeled I and the oxidation of the Pd surface shown in peak II and oxygen evolution after 0.4 V. The reverse scan shows two processes as well: the reduction of oxidized Pd surface to Pd, labeled III and the hydrogen adsorption region, peak IV. The Pd/rGO catalyst was more prone to surface oxidation as seen by the forward scan and confirmed by the surface reduction in the reverse 101 scan, peak III. However, due to the complexity of the surface oxidation process, it can only be confirmed that the surface is oxidized by a mixture of oxygen containing adsorbates. Moreover, the cathodic peak (III) can be utilized to estimate the electrochemically active surface area (ECSA) of the Pd catalyst.[74] The values of the ECSA were computed using the equation 1: ECSA=Q/Sl (1) Where Q is the columbic charge for the reduction of oxidized Pd which is found by integrating the area of peak III, S is the proportionality constant (405 µC/cm 2 ) relating charge with area and l is the catalyst loading.[75] The Pd-utilization was also calculated assuming the surface area for 100% utilization of 1 g Pd to be 448 m 2 , from the charge required to reduce 1 g of Pd 2+ .[76] The Pd/rGO has a larger ECSA as well as better Pd utilization, when compared to the Pd/C catalyst, shown in Table 3.1. This could be due to better dispersion of the catalyst on the graphene support. CV experiments were completed for each catalyst in a solution of 1 M HCOONa and 1 M NaOH and evaluated using the two oxidation peaks. The first peak in the forward scan is the primary oxidation of formate and increases until approximately -0.2 V vs. Ag/AgCl in which the surface is covered in oxide species resulting in deactivation of the catalyst. In the reverse scan, these oxide species are reduced, resulting in reactivation of the surface promoting a spike in oxidation current at -0.20 V and -0.25 V for Pd/rGO and Pd/C, respectively. The peak current densities were obtained prior to the deactivation in the forward scan. Figure 3.5B shows a higher peak current density for both the forward and the reverse scan for Pd/rGO compared to Pd/C as well as a more favorable onset potential, the onset potential was set to when 1 mA/cm 2 of current was seen in the CV, as reported in Table 3.1. The improvement in both peak current density and onset potential imply an important role of the graphene support. This could potentially be due to an increase in accessibility of the incoming reactants to the higher surface area of the Pd.[21] 102 Moreover, the defect sites of rGO can facilitate electron transport and charge transfer accessibility from bulk medium to the electrode and electrolyte interface. Figure 3.5. CV scan of the catalyst in: A) 1.0 M NaOH and; B) 1.0 M HCOONa and 1.0 M NaOH at 20 mV/s. C) Chronoamperommetry scans of the catalyst at -0.65 V vs Ag/AgCl at 1000 RPM in 1.0 M HCOONa and 1.0 M NaOH. Chronoamperometric technique was used to observe the stability of the catalyst under a constant potential of -0.3 V for 5 hours shown in Figure 3.5C. Both catalysts show initial decay in current with time due to catalyst deactivation from adsorbed species onto the Pd surface, however, the current remained higher for Pd/rGO compared to Pd/C throughout the duration of the experiment, which could be due to higher catalytic activity and lower extent of catalyst deactivation. Moreover, the rGO support provides better dispersion of the Pd nanoparticles, which could allow for easier access to the catalytic sites. 103 CVs were obtained at different scan rates to observe the behavior of formate oxidation, shown in Figure 3.6A and B. It can be seen that the peak current density increases with an increase in scan rate, whereas the potential for the peak current shifts in the positive direction that can be attributed to the iR drop.[77] The linear relationship between peak current density and square root of scan rate, shown in Figure 3.6C suggests the FOR is mass transport controlled and involves a diffusion-controlled rate determining process. Figure 3.6. CV scans at different scan rates in 1.0 M HCOONa and 1.0 M NaOH for: A) Pd/rGO and B) Pd/C. C) Plot of Forward peak current vs square root of scan rate and; D) plot of corresponding potential vs ln (v). The slope of the lines also provides insight into the kinetics of the reaction, with the larger slope on Pd/rGO as compared to Pd/C suggesting improved electrooxidation kinetics. The linear relationship between peak potential and natural log of scan rate displayed in Figure 3.6D can be utilized to calculate the charge transfer coefficient (α), shown in equation 2: 104 Ep = A + 𝑅𝑇 (1−∝)𝑛𝐹 ln(v) (2) Where A is a constant related to the standard rate constant and formal potential, R is the gas constant, T is the operating temperature, α is the charge transfer coefficient, n is the number of electrons involved in FOR (n=2) and F is Faraday’s constant.[1,78] Thus, the charge transfer coefficients were found to be 0.648 and 0.731 for Pd/C and Pd/rGO, respectively. Moreover, the linear relationship indicates the oxidation of formate is an irreversible process.[77] To further study the kinetics of formate electrooxidation, Tafel plots shown in Figure 3.7A were obtained from LSV scans at 0.5 mV/s, with the Tafel slopes reported in Table 3.1. There is a decrease in the Tafel slope going from Pd/C to Pd/rGO, which indicates faster charge-transfer kinetics for the Pd/rGO catalyst.[79] Moreover, the lower onset potential to obtain higher currents with the Pd/rGO signifies a more efficient catalyst. This improvement could arise from the defect sites of the rGO that allows for faster electron transport.[16] The temperature effect on the oxidation of formate was also investigated (Figure 3.7B and C) to obtain the activation parameters (Table 3.1). The peak current density increases with an increase in temperature for both catalysts due to improved kinetics. However, the Pd/rGO maintains a higher peak current density at all temperatures, when compared to Pd/C. The current density for each temperature at low overpotentials was used to formulate an Arrhenius plot shown in Figure 3.7D in order to calculate the activation energy for both catalysts. The temperature can be related to the natural log of current in equation 3: Ln (Ip) = Ln (A) - 𝐸 𝑎𝑝𝑝 𝑅𝑇 (3) 105 Where the apparent activation energy is given by Eapp, A is the pre-exponential constant, R is gas constant and T is the operating temperature.[77] The Pd/rGO displayed a decrease in activation energy compared to Pd/C owing to improved kinetics for the electrooxidation of formate. Figure 3.7. A) Tafel Plots for the catalyst in 1.0 M HCOONa and 1.0 M NaOH at 0.5 mV/s. CV scans of the catalysts at different temperatures in 1.0 M HCOONa and 1.0 M NaOH for B) Pd/rGO and C) Pd/C at 20 mV/s and D) Arrhenius plots for the catalysts. Table 3.1. Summary of the catalysts performance Pd/rGO Pd/C ECSA (m 2 /g) 79.22 54.24 Pd-Utilization (%) 18.80 12.11 Eo (V vs Ag/AgCl) -0.868 -0.842 Ip (mA/cm 2 ) 50.60 29.71 Tafel Slope (mV/dec) 144.36 161.18 Eapp (kJ/mol) 8.35 11.57 Rsol (ohms) 14.24 15.32 Rct (ohms) 42.49 386.1 OCV (V) 60 ℃ 0.858 0.837 Power Density (mW/cm 2 ) 60 ℃ 93.8 81.2 106 EIS was used to evaluate the mobile charges in the electrode-electrolyte interface and gain further insight of the electrocatalytic activity. The obtained curves, shown in Figure 3.8A all resemble a depressed semi-circle, and an equivalent circuit model was used to explain the behavior of the electrode, Figure 3.8B. From the circuit Rsol is the solution resistance, Rct is the charge transfer resistance, CPE1 is the constant phase element, which is due to the pseudo-capacitive nature of the catalyst ink and Ws1 is the Warburg element, which is due to the finite diffusion of the formate in solution to the electrode surface. The impedance results are provided in Table 3.1, the charge transfer resistance is significantly enhanced for the Pd/rGO confirming that the rGO support aids in the mobility of charges and improves the electrocatalytic activity of the Pd. Figure 3.8: A) EIS curves of the catalyst. B) Equivalent circuit of the EIS spectra. C) Single cell DFFC polarization curves for the catalysts at 30 and 60 ℃. The formate oxidation activity was further assessed using a DFFC single cell stack fuel cell. The polarization and power density curves for the DFFC with the Pd/rGO and Pd/C anode electrodes are presented in Figure 3.8C. The performance of the single cell fuel cell was observed at two different temperatures, 30 and 60 ℃. At both temperatures, the DFFC with Pd/rGO as the anode catalyst performed better than the DFFC utilizing Pd/C as the anode catalyst in terms of maximum power density and OCV (Table 3.1). There is a near 100% increase in power density 107 when the temperature is increased from 30 ℃ to 60 ℃, due to the improved kinetics of oxidation.[24,26] Furthermore, the OCV increases by 100 mV for both DFFCs, when the temperature increases from 30 ℃ to 60 ℃, however the Pd/rGO DFFC maintains a greater OCV than the Pd/C DFFC irrespective of temperature. The increase in performance from the Pd/rGO DFFC when compared to the Pd/C DFFC, irrespective of the temperature, could be due to the increased ECSA of the Pd nanoparticles as well as the improved kinetics due the rGO support. These results corroborate with those found in the half cell studies, where the Pd/rGO had a higher peak current density and a more negative onset potential compared to the Pd/C electrodes. This section of the chapter is reprinted with permission from ACS Appl. Energy Mater. 2019, 2, 10, 7104–7111. Copyright 2019 American Chemical Society.[80] 3.3. Improving Alkaline Direct Formate Fuel Cell Performance with PdIrO2/MWCNT Electrocatalysts 3.3.1. Experimental Catalyst Synthesis The Pd based catalysts were prepared by a green reduction utilizing methanol as the solvent and a reducing agent.[21] First, the carbon support, either XC72R (C) or multiwalled carbon nanotubes (MWCNT), was dispersed in methanol (0.5 mg/mL). The mixture was then sonicated for 2 h at room temperature. Appropriate amounts of Pd acetate and IrO2 powder were then added to the mixture to obtain 15:5, 10:10, 5:15 weight ratios of Pd to IrO2. The mixtures were then sonicated for 30 min, followed by stirring and heating up to 50 ℃. The mixtures were then stirred 108 at 50 ℃ for 1 h. The mixture was then cooled and separated by centrifugation. The remaining solid was washed by centrifugation with water four times. The resulting solid was then dried in an oven overnight. PdC and PdMWCNT were also prepared in a similar procedure apart from the addition of IrO2. Catalyst Characterization Powder X-Ray Diffraction (XRD) measurements were performed on a Rigaku X-Ray diffractometer with a Cu-Kα (0.154056 nm) radiation source and a scan rate of 4°/min from a 2θ value of 10° to 80°. Scanning Electron Microscopy (SEM) images and Energy Dispersive X-ray Spectroscopy (EDS) mapping were obtained from a JEOL JSM-7001F electron microscope with an acceleration voltage of 19 keV. Transmission Electron Microscopy (TEM) images were taken on a JEOL JEM 2100F with an acceleration voltage of 200 keV. Thermogravimetric analysis (TGA) was performed on a TGA-50 Thermogravimetric analyzer (Shimadzu) under air at a ramp rate of 10 ℃/min. Half-Cell Measurements A standard rotating disk electrode (RDE) setup was utilized for the half-cell measurements. All measurements were obtained on a Solatron SI 1287 potentiostat. Catalyst ink solutions were prepared by adding 2.0 mg of catalyst to 1.0 mL solution composed of 200 µL isopropyl alcohol, 800 µL Millipore water and 7 mg of a quarternized poly(terphenylene) (TPN). After sonication, 20 μL of the solution was pipetted onto a glassy carbon electrode (Pine Instruments) with a surface area of 0.195 cm 2 . After drying, the electrode was placed in a three-cell testing system with the RDE as the working electrode, a platinum wire counter electrode and an Hg/HgO reference 109 electrode (4.24 M KOH filling solution) was used as a reference electrode and all values were converted to Reversible Hydrogen Electrode (RHE) scale using the following equation: ERHE = EHg/HgO + 0.019 + 0.059 x pH Prior to electrochemical testing, the cell was purged with ultra-pure argon (99.999%) for 15 minutes. Electrochemical impedance Spectroscopy (EIS) measurements were performed on a Solatron SI1287 potentiostat and SI1260 impedance Phase Analyzer at several potentials with an amplitude of 5 mV. Membrane Electrode Assembly (MEA) and Fuel Cell Testing For the catalyst layers, the prepared catalysts and Pt/C (Alfa Aesar, 40 wt % metal) were used for the anode and cathode, respectively. The Pd based catalysts and TPN binder were mixed at a catalyst:binder ratio of 1:3 in 200 µL of Millipore water and 200 µL of isopropyl alcohol followed by bath sonication for 8 minutes and probe sonication for 2 minutes. The catalyst inks were painted onto a 4 cm 2 teflonized carbon cloth gas diffusion layer (GDL) until a mass loading of metal of 0.7 mg/cm 2 was obtained for the anode gas diffusion electrode (GDE). The cathode GDE was obtained by mixing Pt/C with Millipore water and Nafion (5 wt%) ionomer at a ratio of 1:10:3, followed by bath sonication for 8 minutes and probe sonication for 2 minutes. The catalyst ink was painted on a teflonized carbon paper with a microporous layer (22BB, SGL Carbon) until a Pt mass loading of 1.3 mg/cm 2 was achieved. The prepared anode GDEs were placed in 1 M NaOH solution for 20 h and then washed with Millipore water. A TPN membrane was sandwiched in between the anode and cathode GDEs and placed into a fuel cell hardware. The fuel cell polarization measurements were performed using a Fuel Cell Test System 890B (Schribner Associates). Prior to testing, a solution of 0.5 M KOH was passed through the 110 cell for 10 minutes followed by Millipore water for 10 minutes, at room temperature. The cell temperature was raised to 60 ℃ and a solution of 4 M HCOOK and 4 M KOH was delivered through the anode compartment at 5 mL/min and recirculated to the fuel reservoir and humidified oxygen was delivered to the cathode compartment at 100 mL/min. The open circuit voltage (OCV) was monitored for 20 minutes, then the cell was held at 0.5 V for 20 minutes and then held at 0.2 V for 10 minutes. The fuel was replaced by fresh fuel (4M HCOOK and 4M KOH) and fuel cell measurements were obtained. 3.3.2. Results and Discussion The metal loadings of the prepared catalysts were then obtained from TGA, shown in Figure 3.9A. The TGA profiles for the catalysts display similar metal loadings between 16 and 23 wt%, moreover all loadings are near the target value of 20 wt%. The XRD patterns of the prepared catalysts is shown in Figure 3.9B, all catalysts present a peak near 25° corresponding to carbon (002) miller index. The catalyst with only Pd display peaks near 40°, 46°, and 68°, which correspond to the (111), (200), (220), and (311) Miller indices of fcc Pd, respectively. The IrO2C pattern displayed peaks near 28°, 35°, 40°, 54°, 58°, 66°, 69°, and 73° corresponding to (110), (101), (200), (211), (220), (112), (301), and (311) Miller indices, respectively. The catalyst with both Pd and IrO2 display all the aforementioned peaks, however the signals arising from IrO2 are slightly shifted to a more positive 2θ. Moreover, there is a difference in the peaks near 68° and 70°, indicating a difference in the structure of the catalysts from interaction between Pd and IrO 2. The average crystallite size was determined utilizing the Debye-Scherrer equation and the broadening of the signal near 40° and shown in Table 3.2. The mixture of the Pd and IrO2 decreases the crystallite size for the catalysts prepared on XC72R, however the catalysts supported on 111 MWCNTs displayed similar crystallite sizes. The difference observed in the XC72R supported catalyst likely arises from the differences in surface functionality and pore size. Figure 3.9. A) TGA profile and B) powder XRD pattern of the prepared catalysts. BET measurements of the bare XC72R and MWCNTs were obtained and are shown in Figure 3.10. The pore size and surface area of the supports were obtained and found to be 0.232 mL/g and 194 m 2 /g, and 1.655mL/g and 204 m 2 /g for XC72R and MWCNT, respectively. Thus, further providing evidence that the difference in pore size could result in differences in crystallite size observed from XRD. 112 Figure 3.10. A) N2 adsorption-desorption isotherms of XC72R and MWCNT and B) corresponding mesoporous size distribution curves for XC72R and MWCNT To further study the differences of the surface morphology, SEM images and EDS mapping were obtained, and are shown in Figure 3.11. From the SEM images the XC72R supported catalysts show similar morphologies where the nanoparticles are too small to be discernable. However, the IrO2 containing catalysts did show a higher propensity of larger agglomerations, which resulted in mainly being IrO2. The catalysts prepared on MWCNTs, were also similar in morphologies and also had larger IrO2 particles. There are differences in morphologies between the XC72R and MWCNT catalysts, the MWCNTs present a more agglomerate like morphology with larger pores whereas the XC72R shows smaller pores and a more level surface. Moreover, EDS was utilized to determine the elemental composition of the catalysts, as shown in Table 3.2. The elemental composition determined from EDS are similar to the metal weight loading determined by TGA. Moreover, the EDS confirms that the larger particles observed are IrO2 and for the most part there is a uniform distribution of the elements throughout the catalysts surface. 113 114 115 Figure 3.11. SEM images and EDS mapping for A) Pd/C and B) Pd/MWCNT where red corresponds to C and green corresponds to Pd. SEM images and EDS mapping for C) IrO 2C; D) PdIrO2C (15:5); E) PdIrO2C (10:10); F) PdIrO2C (5:15) and G) PdIrO2MWCNT where red corresponds to carbon, green corresponds to O, blue corresponds to Pd and yellow corresponds to Ir. To gain further insight into the morphology, TEM micrographs of the catalysts were obtained and are shown in Figure 3.12. The catalysts prepared on XC72R display similar morphology with the metal nanoparticles dispersed on the carbon particles. The catalysts supported on the MWCNT also have similar morphologies with the carbon nanotubes entangled containing the metal nanoparticles on the surface. There is minimal difference in the morphology of the Pd and PdIrO2 particles across the catalysts. However, the size of the particles varies as shown in Table 3.2, presenting a similar trend to the crystallite sizes obtained via XRD. Thus, the incorporation of IrO2 aids in preventing catalyst agglomeration during the synthesis of the 116 nanoparticles. Moreover, the particles are larger than previous reports and this is likely due to the increased fabrication temperature resulting in faster growth of the particles.[21] 117 Figure 3.12. TEM micrographs of A) PdC; B) PdIrO2C (15:5); C) PdIrO2C (10:10); D) PdIrO2C (5:15); E) PdMWCNT and F) PdIrO2MWCNT Table 3.2. Summary of the physical characterization for the prepared catalyst. Catalyst SEM-EDS XRD TEM Pd (wt%) IrO2 (wt %) Crystallite Size (nm) Particle Size (nm) PdC 20 - 11 15.8 ± 6.1 PdIrO2C (15:5) 16 7 7 9.6 ± 4.4 PdIrO2C (10:10) 13 6 9 6.9 ± 1.9 PdIrO2C (5:15) 7 9 9 9.7 ± 5.7 PdMWCNT 19 - 9 7.2 ± 2.9 PdIrO2MWCNT 16 8 9 7.2 ± 4.2 IrO2C - 16 14 - The electrochemical activities of the catalysts were first studied in 1M KOH and are shown in Figure 3.13A. Similar to the CVs obtained in section 3.2, all the Pd catalyst contained four 118 distinct electrochemical signals. The hydrogen desorption/oxidation region near 0.3 V vs. RHE, the Pd oxide formation and oxygen evolution region after 1 V vs. RHE, the reduction of Pd oxides at 0.7 V vs. RHE and the hydrogen adsorption and hydrogen evolution region after 0.2 V vs. RHE. The catalyst only containing IrO2 did not present any apparent electrochemical signals. Moreover, the current for the catalysts prepared with MWCNTs presented a significantly higher current value than the XC72R counterparts. The Pd oxide reduction peak was utilized to obtain the ECSA from equation 1 in section 3.2 and the calculated values are shown in Table 3.3. The addition of minor amounts of IrO2 improved the ECSA significantly, however when the Pd and IrO2 content were similar or more IrO2 than Pd there was a decrease in the ECSA. Changing the catalyst support to MWCNT also resulted an increase in the ECSA and a significant increase in the PdIrO2MWCNT. The increase likely arises from the higher surface area of the CNTs and smaller catalyst particle size. Figure 3.13. CV scans of the prepared catalysts in A) 1M KOH and B) 1M HCOOK and 1M KOH at a scan rate of 20 mV/s under a flow of N2. Furthermore, CV scans of the catalyst in an electrolyte solution of HCOOK and KOH were obtained and are shown in Figure 3.13B. Similar CVs to those shown in section 3.2 were obtained; 119 in the forward scan there is an increase in the current after 0.2 V vs. RHE and peak current occurs between 0.7 V 0.9 V vs. RHE. Furthermore, the peak in current in the forward scan is ascribed to surface oxidation or coverage by intermediates impeding the accessibility of catalytic sites. Then in the reverse scan the surface oxides and intermediates are removed and an increase in the current is observed related to oxidation of formate.[81,82] Moreover, the peak current in the forward (If) and the reverse (Ib) scan can be utilized to obtain the reactivation coefficient (Ib/If). The reactivation coefficient was found to be near 1 for all the catalysts, indicating the presence of IrO2 and the MWCNTs does not play a role in oxide removal from the surface. To gain further insight into the differences in catalytic activity of the catalysts, the CVs were completed at several scan rates and shown in Figure 3.14. As expected, there is an increase in the peak current (Ip) as the scan rate (v) increases, further the peak current can be plotted against the square root of scan rate (Figure 3.15A) giving a linear relationship. The linear relationship from Figure 3.15A confirms the FOR is a diffusion limited process. 120 Figure 3.14. CV scans in 1M HCOOK and 1M KOH at various scan rate for A) PdC; B) PdIrO2C (15:5); C) PdIrO2C (10:10); D) PdIrO2C (5:15); E) PdMWCNT and F) PdIrO2MWCNT Furthermore, plotting the peak current potential (Ep) and the natural log of scan rate (Figure 3.15B) can be utilized to obtain the value of αn’ using equation 4. E p ln v = RT 2F(αn’) (4) From equation 4, R is the gas constant, T is temperature, F is Faraday constant and α is the charge transfer coefficient and n’ is the electrons in the rate determining step (RDS). The αn’ values for the three electrodes were calculated and utilized to obtain the diffusion coefficient of the catalysts. The diffusion coefficient was found from equation 5 and the slope of the plots in Figure 3.15A. I p v 1/2 = 2.99 X 10 5 n(αn’) 1/2 CD 1/2 (5) 121 In equation 5, C is the concentration of formate in solution, D is the diffusion coefficient and n is the number of electrons in the entire electrochemical process. The values for formate diffusivity are shown in Table 3.3. Previous study has demonstrated differences in methanol diffusion to catalysts by altering the catalyst support.[83] Similarly, an increase in the diffusion of formate is observed when the catalyst support is changed from XC72R to MWCNTs. However, there is a decrease in the diffusion of formate when IrO2 is introduced into the catalyst. Figure 3.15. A) Plot of Forward peak current vs square root of scan rate and B) plot of peak potential vs natural log of scan rate. Tafel plots were obtained from LSV scans from an overpotential of 400 mV to OCV at a scan rate of 0.5 mV/s to obtain any differences in kinetics towards FOR from the catalysts (Figure 3.16A). The onset potentials for the catalysts were obtained from the Tafel plots and are shown in Table 3.3, the addition of IrO2 improves the onset potential, except at excessive IrO2 loading. Moreover, the use of MWCNTs further improves the onset potential and this arises from the improved surface area and ECSA. The Tafel slopes were also obtained and shown in Table 3.3. A similar trend is observed in the Tafel slopes, where a decrease in the Tafel slope is observed when the carbon support is changed to MWCNTs and a further decrease in the slope with the minor 122 addition of IrO2. Thus, an improvement in the kinetics is obtained from utilizing MWCNTs as the support as well as the addition of IrO2. The stability of the catalysts was also studied in half-cell experiments by potentiostatic experiments, holding the potential at 0.4 V vs. RHE (Figure 3.16B). Overtime, there is a decrease in current observed for many of the catalysts. Thus, the poisoning rate was determined from equation 6. σ = 100 I o ×( dI dt ) t>500s (6) In equation 6, dI dt is the slope of the linear portion of the current decay and Io is the initial current. The poisoning rates are shown in Table 3.3 and are near 1% decay apart from the PdIrO2C (5:15) and PdIrO2MWCNT, which have significantly lower decay rates. Figure 3.16. A) Tafel slopes obtained from LSVs at 0.5 mV/s and B) change in current over time holding the potential at 0.4 V vs. RHE for the prepared catalysts. To obtain further insight into the differences in the catalytic activity of the catalysts EIS experiments were completed at several potentials, as shown in Figure 3.17. The Nyquist plots obtained were fit using a similar equivalent circuit as shown in Figure 3.8B. For all the catalysts 123 and potentials, a similar solution resistance was obtained with a value near 30 Ω. Thus, the addition of IrO2 and use of MWCNTs as support do not change the resistance of the electrolyte. However, a significant change in the Rct is observed when increasing the potential, a decrease in the Rct is observed due to faster charge transfer. Moreover, there is a difference in the Rct at similar potentials across the prepared catalysts. The Rct values at -0.4 V vs. MMO are shown in Table 3.3 and a similar trend to the calculated ECSA was observed. An improvement in charge transfer is observed with minor addition of IrO2 and further improvement is observed when MWCNTs are utilized as the catalyst support. This likely arises from an improved interaction between the catalysts and the MWCNT and the greater surface area of the MWCNTs compared to XC72R. 124 Figure 3.17. EIS spectra at various potentials vs. MMO for A) PdC; B) PdIrO2C (15:5); C) PdIrO2C (10:10); D) PdIrO2C (5:15); E) PdMWCNT and F) PdIrO2MWCNT Table 3.3. Summary of half-cell results for the prepared catalysts. ECSA (m 2 /g) Diffusion Coefficient (x10 -7 ) Onset Potential (V vs. RHE) Tafel Slope (mV/dec) Rct at -0.4 V vs. MMO (Ω) Poisoning Rate (%) PdC 30.0 9.09 0.077 193.43 234.9 1.1 PdIrO2C (15:5) 62.8 6.86 0.037 185.79 77.12 1.3 PdIrO2C (10:10) 19.5 1.14 0.058 193.16 272.3 1.2 PdIrO2C (5:15) 10.0 1.68 0.092 213.25 919.9 0.56 PdMWCNT 50.2 154.92 0.029 179.44 83.4 1.4 PdIrO2MWCNT 102.2 28.84 0.028 166.14 71.8 0.29 Select catalysts were utilized as the anode electrocatalysts in DFFCs to demonstrate the practical efficacy of the catalysts. The PdC, PdIrO2C (15:5), PdMWCNT and PdIrO2MWCNT were utilized and the results of the DFFCs are shown in Figure 3.18A. Regardless of the catalysts the OCV for the DFFCs were all near 0.91 V. However, the peak power density for the catalysts varied and followed a similar trend to the observed half-cell results. An increase in peak power 125 density was observed with the addition of IrO2 due to improved charge transfer and increased ECSA. Moreover, the power further increased with the use of MWCNTs as the catalyst support due to further improvements in the charge transfer and improved surface area. Moreover, the stability of the PdIrO2MWCNT containing DFFC was studied by holding the cell at a constant current of 200 mA/cm 2 (Figure 3.18B). There is minimal decay in the potential over the allotted time and the decrease was due to consumption of the formate and formation of carbonates resulting in a decrease in ionic conductivity. Furthermore, the DFFC reported here provides one of the highest reported power densities for a DFFCs of 299.2 mW/cm 2 . Hong et al. recently reported a DFFC with higher peak power, however they utilized a higher catalyst loading and higher flow rates of reactants, the DFFC with similar formate flow rate provided lower peak power (250 mW/cm 2 ) than the power obtained by the PdIrO2MWCNT.[56] Figure 3.18. A) Polarization and power curves for the DFFCs with the prepared catalysts and B) the DFFC with PdIrO2MWCNT anode electrode held at constant current of 200 mA/cm 2 126 3.4. Studies on the Simultaneous Production of Alkali Hydroxide and Electricity with a Direct Formate Fuel Cell 3.4.1. Experimental Half-Cell Measurements All electrochemical analyses were carried out in a three-electrode system on a Solatron SI 1287 potentiostat. A Pt wire and Ag/AgCl in 1M KCl were used as counter and reference electrode, respectively. The catalyst ink for the working electrode was composed of 5.0 mg of Pd/C (20 wt% Pd) and 1.0 mL of a 100 mL stock solution composed of 10% isopropyl alcohol, 90% Millipore water and 10 mg of Nafion (5 wt%) binder and the mixture was sonicated for 8 min. After sonication, 20 µL of the catalyst suspension was pipetted onto a rotating glassy carbon disk electrode (RDE, Pine Instruments) with a surface area of 0.195 cm 2 . Prior to electrochemical experiments, the cell was purged with N2 for 30 minutes. Electrochemical impedance spectroscopy (EIS) measurements were performed on a Solatron SI 1287 potentiostat and SI 1287 impedance Phase analyzer at varying potentials with an amplitude of 5 mV. Membrane electrode assembly and Fuel Cell Characterization Nafion 117, 211 and Fumasep F1850 membranes were prepared by modifying a reported procedure, the membranes were first pretreated by boiling them in 3% hydrogen peroxide, Millipore water, 1M H2SO4 and Millipore water for 1 hour each; then the pretreated membranes were placed in 2.5 M MOH (M=Li, Na, K, Rb, Cs) at room temperature for 2 hours and stored in Millipore water.[33,34] 127 The anode catalyst ink was formulated by mixing 20 wt% Pd/C, ethanol and a neutral binder (Fluorinated ethylene Propylene 55 wt% in water) in a ratio of 6:17:1, respectively, and sonicating the mixture for 10 minutes. The resulting catalyst ink was then hand-brush painted onto a teflonized carbon paper (4 cm 2 ) until a loading of 2.0 mg of Pd was obtained. The cathode coated electrode was obtained by mixing 40 wt% Pt/C with water and Nafion binder solution (5 wt%) in a ratio of 1:5:3, respectively. The catalyst ink was then hand-brush painted onto a teflonized carbon paper (4 cm 2 ) until a loading of 2.0 mg of catalyst was obtained. The MEA was obtained by placing the pretreated membrane in between the anode and cathode electrodes and pressed at 130 ℃ for 5 minutes under 500 psi. The fuel cell measurements were performed using a Fuel Cell Test System 980B (Schribner associated). The MEAs were tested at 80 ℃, with non-heated 1M HCOOM delivered through the anode compartment at a flow rate of 2.5 ml/min and humidified oxygen was heated to 85 ℃ and passed through the cathode compartment at 100 ml/min. The alkali hydroxide was quantified by 13 C NMR following our previous report.[31] 3.4.2. Results and Discussion Figure 3.19A shows the CV of the different HCOOM salts. The forward scan shows an increase in current until approximately 0.2 V vs. Ag/AgCl for all the formate salts, in which there is a decrease due to surface coverage of oxide species. The peak current was highest for Na + and then decreased as there was an increase in cation size. Moreover, the oxidation peak from the reverse scan (Ib) remained the same as peak current in the forward scan (If) for Li + and Na + and slightly decreased for K + and significantly decreased for Rb + and Cs + . The reverse oxidation peak is believed to originate from reactivation of catalytic sites by removal of oxygen containing species 128 on the surface of the catalysts, thus the decreased reverse oxidation peak could be indicative of more strongly bound oxygen containing species on the catalysts. Previous studies have shown an increase in Pd-cation interaction in order of increasing cation size.[84,85] As it is believed the cations adsorb onto the metal surfaces in the form of OHadsM + (H2O)x, thus increasing cation atomic size resulting in more strongly bound oxygen containing adsorbates. Even though previous cation effects on FOR show a different trend (increase in activity as cation size increases) these studies were done on Pt catalysts and several studies have shown variations of cation effects based on metals used. Moreover Hiltrop et al. have showed that the cation effect for glycerol oxidation on Pd/CNTs also varied depending on the functional groups on the carbon supports.[39] Figure 3.19. A) CV scans of the formate salts at a scan rate of 20 mV/s under flow of N 2. B) CV scans of the hydroxide base at a scan rate of 5 mV/s under flow of N2. CV experiments were also completed in MOH solution, Figure 3.19B. There are four distinct peaks in the CV scans which correspond to hydrogen adsorption and hydrogen evolution (I), desorption of hydrogen (II), oxidation of Pd surface and oxygen evolution (III) and reduction of oxidized Pd surface to Pd (IV). In the forward scan there appears to be no apparent difference in the oxidation of the Pd surface, however there is significant change in the reduction charge. This change has been observed for Pt and was attributed to higher interfacial concentration of 129 OHadsM + (H2O)x however the charge increased with increased ionic radius.[47] Whereas the difference is not linear here, this could be from different interactions between the cations and Pd, as well as the carbon support. Moreover, this agrees with the results obtained by the reverse scan in FOR, where the oxidation peak in the reverse scan was lower in current as the ionic radius of the cation increased. Recently, Choun et al. showed that HCOO - electrooxidation on Pd/C surface is limited by equation 8 and the Hads arises from both dissociation of H2O and electrooxidation of HCOO - . [48,86] There is a difference in the peak current for the desorption of hydrogen, shown in Table 3.4, in which a similar trend is seen in previous results. Thus, the improvement of FOR could be due to a more facile removal of Hads in the presence of Na + . This improvement could be due to the OH from the OHadsM + (H2O)x adsorbed onto the surface. HCOO - → Hads + COO - ads (7) Hads + OH - → H2O + e - (8) COO ads − → CO2 + e - (9) 130 Figure 3.20. CV scans at various scan rates for A) HCOOLi; B) HCCONa; C) HCOOK; D) HCOORb and E) HCOOCs with Pd/C catalysts under flow of N2. To further try and understand the role of the cations, CV experiments were completed at different scan rates, as shown in Figure 3.20. Moreover, the peak current (Ip) and square root of scan rate resulted in a linear relationship, as shown in Figure 3.21A, thus the FOR is determined to be a diffusion limited process. Moreover, by obtaining the slope of E p vs ln(v), from Figure 3.1B, αn’ can be determined utilizing equation 4 and is approximately 0.5 for all HCOOM solutions; since the rate determining step must be an 1e - process, we can assume α≈ 0.5. 131 Furthermore, the diffusion coefficients can be obtained from Figure 3.21A and equation 5, similar values were obtained to previously reported values.[87–89] However, the diffusion coefficients follow a similar trend as the CV measurements, where the highest diffusion coefficient was obtained for HCOONa. Even though surface coverage of cations should increase in the order of Li>Na>K>Rb>Cs, the values of diffusion coefficients show the opposite trend. Higher activity may be due to the added presence of hydroxide from the cation-hydroxide adsorption. Variations to this trend are not uncommon, one possible reason for this difference is the use of the carbon support which has oxygen functional groups that may bind to the cations. Cations have been shown to bind to carbon substrates via cation-pi interactions and the binding energy is higher in order of Li>Na>K, therefore this could lead to more Li + and Na + binding to the carbon surface allowing for more metal active sites whereas K + , Rb + and Cs + more favorably bind to the metal surface.[84,90] Figure 3.21. A) Plot of Forward peak current vs square root of scan rate and B) plot of peak potential vs natural log of scan rate. To further gain insight into the kinetics of the reaction, Tafel plots were obtained by running LSVs at a scan rate of 0.5 mV/s, as shown in Figure 3.22A. From the Tafel plots, the Tafel slopes were obtained and are shown in Table 3.4. The obtained values corroborate with diffusion coefficients, indicating faster charge transfer kinetics for HCOONa. 132 Figure 3.22. A) Tafel plots for the various formate salts and B) EIS at 0.1 V vs. Ag/AgCl for the various formate salts. EIS experiments were completed at 0.1 V vs. Ag/AgCl, the Nyquist plots were fitted to an equivalent circuit shown in Figure 3.8B. From the equivalent circuit Rsol is the solution resistance, Rct is the charge transfer resistance, CPE1 is the constant phase element from the pseudo-capacitive nature of the catalyst, and Ws1 is the Warburg element due to the finite diffusion of reactants. The Rsol for the salts were similar between 20-25 Ω, apart from HCOOLi which had an Rsol of 42 Ω. This likely arises form the difference in hydration of Li + . Moreover, the Rct of the salts varied significantly and HCOONa displayed the lowest Rct, which likely arises from the higher accessibility of the formate to the catalyst surface. The stability of the formate oxidation was studied by constant potential experiments, shown in Figure 3.23. An initial drop in current was observed for all the salts, however they all stabilized after 500 s. Thus, the poisoning rate was calculated from equation 6. The poisoning rates were found to be similar for all the cations, shown in Table 3.4. The similar poisoning rates suggests that any long-term instability would be independent of the cations. 133 Figure 3.23. Constant potential experiments for the A) Li, Na, K and B) Rb and CS formate salts held at 0.1 V vs. Ag/AgCl. Table 3.4. Summary of half-cell results Cation Tafel (mV/dec) αn' Diffusion coefficient (x10-7) Rct at 0.1V vs. Ag/AgCl (Ω) Hdes charge (mC) Poisoning Rate (%) Li 177 0.52 3.73 625.7 0.89 1.4 Na 166 0.53 8.29 158.2 1.86 1.2 K 176 0.56 5.94 776.6 0.98 1.6 Rb 185 0.54 3.54 1179 0.11 1.4 Cs 186 0.51 3.51 1810 0.25 1.3 Furthermore, to view practical systems, the HCOOM fuels were studied in fuel cell experiments, utilizing a CEM which had been converted to the respective alkali metal. Even though power produced is lower than the AEM-DFFC, CEM-DFFC has the added benefit of producing hydroxide base in the cathode. First, CEM-DFFC experiments were conducted with HCOONa and studied with three commercial membranes (Figure 3.24). From the three membranes studied Nafion 211 produced the highest peak power, even though it had a lower OCV. The 134 decrease in OCV is due to crossover of water decreasing accessibility of O2 at the cathode electrode and crossover of formate. However, the decreased thickness is also the reason for the increased power, as it decreases the ohmic resistance and the Nafion 211 membrane has a higher ionic conductivity compared to the other two membranes. Furthermore, to confirm the production of hydroxide base, the DFFCs were held at a constant current over 20 h (Figure 3.24B). There is an initial drop in the potential but then begins to stabilize after 10000 s. The product at the cathode was collected and initially the pH was obtained to confirm that an alkali solution was obtained, the amount of hydroxide was then confirmed by 13 C NMR after bubbling the cathode product solution with CO2. Figure 3.24. A) Polarization and power curves for 1M HCOONa with varying membranes at 80 ℃. B) The potential of the DFFCs held at constant current. Based on the obtained results from the various membranes, formate salts with varying cations were then studied in CEM-DFFCs, shown in Figure 3.25. There is an increase in power generated in sequence of Na>K>Li>Rb>Cs, which is different than previously reported studies of DFFC utilizing an AEM, which reported that HCOOK outperforms HCOONa.[26] This difference could arise from the use of non-carbon supported Pd, anion exchange membrane and addition of supporting electrolyte in previous studies. 135 Constant current experiments were completed to observe the stability of the DFFC as well as verify the production of hydroxide (Figure 3.25B). The potential was relatively stable for all DFFCs except for HCOOLi in which a drop was observed after 10 hours. This drop is attributed to the lower solubility of LiOH causing degradation of the membrane, which was verified when the DFFC assembly was taken apart. White precipitates formed on the surface of the membrane and cathode electrode leading to physical degradation of the membrane. Furthermore, the constant potential half-cell experiments did not demonstrate any differences in rate of degradation for the varying formate salts. Even though, the half-cell potentiostatic times were not the same duration as the fuel cell experiments the rates of degradations are all similar, thus confirming the loss in fuel cell performance for HCOOLi is not due to issues with the FOR but rather with the LiOH formation. Moreover, all the formate salts were able to produce hydroxide base of various cations. The LiOH was not quantifiable due to the formation of the precipitate, however the obtained precipitate and aqueous products were confirmed by pH meter to be caustic. Figure 3.25. A) Polarization and power curves for 1M HCOOM with varying cations at 80 ℃. B) The potential of the DFFCs held at constant current. 136 3.5. Conclusion Formate has been shown to be one of the most promising fuels for alkaline direct liquid fuel cells. The role of the catalysts and their supports are vital in fuel cell performance, thus the use of rGO as a support material for Pd nanoparticles synthesized via coreduction of the metal-salt and GO for the FOR in DFFCs was demonstrated. XPS studies revealed that the catalyst particles of the Pd/rGO were mostly in the Pd 0 state with minimal PdO present. The Pd/rGO catalyst exhibited a greater catalytic activity and stability towards the FOR compared to the commercial Pd/C catalyst. This enhancement could be attributed to the better dispersion of the metal catalyst, which leads to a greater electrochemically active surface area and better catalyst utilization. Furthermore, it was found that the activation energy for the FOR was lowered with Pd/rGO. The rGO supported catalyst also had a 15% greater power density than Pd/C in a DFFC with Pd metal loading lower than 1 mg/cm 2 . To further improve upon the catalysts, Pd nanoparticles were combined with IrO2 to observe changes in catalytic activity. Moreover, to further illustrate the role of the catalyst support MWCNTs were utilized as the support for the PdIrO2 catalysts. Both the addition of IrO2 and use of MWCNTs demonstrated greater catalytic activity compared to Pd/C. The improvements were due to improved ECSA and improved kinetics arising from the greater surface area of the MWCNTs and synergistic effects between Pd and IrO2. The improvements were further exemplified in DFFC studies where a 50% improvement in peak power density was observed from the PdIrO2MWCNT and the PdC DFFC. The peak power of 299 mW/cm 2 being amongst the highest reported for DFFCs and with low metal loading. 137 Lastly, the role of cations on FOR in the absence of added electrolyte was also studied. 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Alkaline Direct Methanol Fuel Cells Several of the issues observed in PEMFCs can be mitigated by operating DLFCs under alkaline conditions.[1–7] However, the development of highly conductive and stable anion exchange membranes (AEM) and anion exchange ionomers (AEI) has been a challenging feat to achieve.[8] Nonetheless, large strides in their developments have been made in recent years with the use of different polymer backbones in which AEM hydrogen fuel cells have been able to achieve power densities comparable to their PEM fuel cell counterparts.[6,9–13] Direct methanol fuel cells (DMFC) have a longstanding history as one of the most promising liquid fuel cells, already having some commercial systems developed.[14–22] The use of methanol has a plethora of advantages as a fuel such as potential generation from renewable energies, high power and energy densities (6100 Wh/kg), oxidation which does not involve C-C bond breaking and the potential of using platinum group metal (PGM)-free catalysts, such as Ni based catalyst, for methanol oxidation reaction (MOR) under alkaline conditions.[23–26] Even though operating DMFCs in alkaline conditions has several advantages, the performance for alkaline direct methanol fuel cells (ADMFC) has been problematic because of lack of a suitable alkaline hydroxide-ion conducting membrane and an ionomer binder. Our group previously demonstrated a record-breaking power density of 160 mW/cm 2 for ADMFCs, however the performance was obtained utilizing high catalyst loadings.[27] Several developments were made thereafter in attempts to improve the performance; catalyst modification and changes in the AEM being the two most prominent areas of interest.[28–39] However, it wasn’t until recently with the use of a polybenzimidazole- (PBI) based membrane containing graphene oxide (GO) fillers that 153 significant improvements to ADMFC performance were achieved.[40,41] Moreover, the researchers were able to decrease the loading of Pt-Ru by about 75% compared to the ADMFC studied by our group, which is an important step towards reducing manufacturing costs. The improvements of cell performance were ascribed to properties of the membrane such as its high ionic conductivity and decrease in methanol permeability from the addition of the GO fillers. To further attempt to decrease the manufacturing costs of FCs, significant efforts have gone into utilizing transition metal catalyst in alkaline fuel cells. Most efforts have gone into developing electrocatalyst for the ORR and developing bimetallic catalyst with noble metal catalyst for the anode reaction. However, non-noble metal-based catalyst have been shown to have activity for the MOR, formate oxidation reaction, glycerol oxidation reaction, ascorbate electrooxidation reaction and urea electrooxidation reaction.[42–48] Fleischmann was the first to demonstrate that Ni electrodes can effectively oxidize methanol under alkaline conditions.[42] Since then, several researchers have developed various Ni based catalysts to improve the catalytic activity via preparing different Ni-oxides, altering the catalyst support, and incorporating other metals or heteroatoms into the catalyst.[49–66] Many reports provide improved current densities and stability in half-cell experiments. However, majority of the developed catalyst have only been shown to be active in half-cell studies and do not get incorporated into ADMFCs and those that have been shown in fuel cell studies have yielded low power outputs. [67,68] Gallium nickel-based catalysts have been shown to be one of the most effective catalysts for the conversion of CO2 to MeOH in hydrothermal reactions.[69,70] Moreover, Ga and other P block metals have been utilized in photochemical and electrochemical reactions to improve catalytic activity.[71–75] Recently, Hsu et al. demonstrated the incorporation of Au-GaOOH nanorods into commercial Pt/C for MOR under photochemical conditions.[76] The prepared 154 catalysts displayed improved CO tolerance as well as increased currents when compared to the commercial Pt/C. Another metal-oxide which has been heavily studied is CeO2 due to its rapid conversion between Ce 3+ and Ce 4+ which allows for storage and release of oxygen and well dispersion of metal nanoparticles amongst other factors.[77–82] CeO2 has been shown to improve MOR with several catalysts, amongst the earliest reports demonstrating the CeO2 aiding in oxidation of CO to CO2 from Pt surfaces. It has also been synthesized with a nanocrystaline zeolite and demonstrated to maintain current towards MOR after 1000 cycles.[78] Furthermore, it has recently been utilized with an iron-based catalyst and demonstrated to improve catalytic activity from the oxygen vacancies present and improved catalyst stability.[82] Membranes and electrocatalysts are two important components in fuel cells. However, a third component that is significant in developing an effective membrane electrode assembly (MEA) and in turn a high performing fuel cell is the ionomeric binder. The ionomeric binder has two major functions: 1) bind the catalyst nanoparticles to the substrate and form a porous catalyst layer and 2) create the triple phase boundary layer for ion transport. PEM fuel cells have the advantage of utilizing the well-developed Nafion based binder, however, there is no standard alkaline ionomer binder in AEM-based technology at this moment. Jervis et al. showed that hydrogen oxidation reaction (HOR) under alkaline conditions with Nafion binder underestimated the catalytic activity due to the acidic ionomer binder acting as an insulator of hydroxide transport.[83] Nonetheless, several groups have developed anion ionomer-free binders such as polytetrafluoroethylene (PTFE) and Nafion as a cation ionomer to transfer Na + in alkaline-acid fuel cells.[84–86] Even though these DLFCs have high power densities, the use of two liquids and often use of harmful oxidants makes commercialization of these systems difficult. Thus, 155 development of an effective AEI is imperative to further the advancement of conventional AEM fuel cells. Recently, there have been several advancements in developing AEIs in order to improve their stability under normal operating conditions. However, unlike PEMFCs, the use of a single AEI for both anode and cathode has been difficult to achieve due to the different properties needed in order to obtain optimal performance. Thus, two different AEIs are often employed for the anode and cathode of AEM fuel cells to achieve increased performance.[6,7,87,88] Further, the interaction between the AEIs and the metal catalyst has only recently started to gain interest from researchers.[10,83,87,89–92] Several of these studies have provided insight into cation-hydroxide interaction between the cations of AEIs and the catalysts, which can result in changes in performance and impact stability. Kohl et al. were the first to show the use of cation-hydroxide adsorption of quaternary ammonium cations with varying alkyl chains and poly(diallyldimethylammonium) (PTMA) cation onto a Pt surface resulting in decreased rate of methanol oxidation.[92] Recently, Santori et al. showed that changes in AEI structure did not heavily impact ORR for Pt/C and HOR for PtRu/C, however when utilizing PGM-free catalysts for ORR the changes in AEI had significant impact and changes in AEI also had a large impact when Pd-CeO2/C was used as an electrode catalyst for HOR.[91] However, apart from the work by Kohl et al., which mainly focused on cations in solution and not the AEI binder itself, not much work has been done looking into changes in AEI structure for DLFCs, where operating conditions differ to those of H2/O2 fuel cells. In this chapter, we provide insight into the role of AEIs in DLFCs and how minor changes in the chemical structure effect operating performance in ADMFCs. Moreover, we show stable short-term durability, highly conductive AEM for ADMFCs, which provides improvement to current state of the art ADMFCs by operating under low noble metal loadings while still achieving 156 high power densities. We also synthesized gallium nickel oxide-based catalysts supported on carbon for MOR. By altering the elemental composition of the catalysts, we demonstrate improvements in the half-cell performance when incorporating Ga into the Ni catalysts. Furthermore, we prepared Ni nanoparticles on a CeO2/C support, improving the catalytic activity towards MOR. Finally, ADMFCs with high-power densities for non-platinum group metal based ADMFCs were demonstrated. Further, we incorporated noble metal free ORR catalysts to demonstrate a completely noble metal free ADMFC with competitive power densities to reported ADMFCs. Thus, demonstrating the potential of utilizing nonnoble metals in ADMFCs and aiding in decreasing the cost of the system. 4.2. Ionomer Significance in Alkaline Direct Methanol Fuel Cell to Achieve High Power with a Quarternized poly(terphenylene) Membrane 4.2.1. Experimental Polymer Electrolytes Quaternized poly(terphenylene) (TPN) mebranes and binders were received from Professor Bae’s group.[9] The AEI binders were composed of TPN with three different cations – alkyltrimethylammonium (TMA), 1,2-dimethylimidazolium (DMIm) and N-methylpiperidinium (Pip)– and they were dissolved in a 7:3 mixture of isopropyl alcohol and water (3 wt%). The chemical structures of the polymer electrolytes are shown in Figure 4.1A. Half-Cell Testing Catalyst inks were prepared by mixing 2 mg of PtRu/C (Fuel Cell Store, 20 wt % metal), 5 mg of binder solution, 100 µl of Millipore water (Direct-Q UV, 18.2 MΩ) and 100 µl of isopropyl 157 alcohol. The mixture was bath sonicated for 8 minutes and 20 µl was drop cast on a polished glassy carbon rotating disk electrode (RDE) with a geometric area of 0.195 cm 2 . Half-cell measurements were performed using a Solatron SI 1287 potentiostat with a standard three electrode cell at room temperature. A platinum wire served as the counter electrode, a Hg/HgO electrode (4.24 M KOH filling solution) was used as a reference electrode and all values were converted to Reversible Hydrogen Electrode (RHE) scale using the following equation. ERHE = EHg/HgO + 0.019 + 0.059 x pH Prior to electrochemical testing, the cell was purged with nitrogen (99.99% purity) for 15 minutes and the electrodes were stabilized by performing 100 cyclic voltammograms (CV) at 100 mV/s between 0.0 and 1.1 V vs RHE. All half-cell measurements were normalized to geometric area of the glassy carbon electrode. Electrochemical impedance Spectroscopy (EIS) measurements were performed on a Solatron SI1287 potentiostat and SI1260 Impedance Phase Analyzer at varying potentials with an amplitude of 5 mV. Experiments were completed in triplicates to ensure reproducibility; the standard deviations are provided in Table 4.2. Membrane Electrode Assembly (MEA) and Fuel Cell Measurements For the catalyst layers, PtRu/C (Fuel Cell store, 20 wt% metal) and Pt/C (Alfa Aesar, 40 wt % metal) were used for the anode and cathode, respectively. The PtRu/C and AEI binders were mixed at a catalyst:AEI ratio of 1:3, unless stated otherwise, in 200 µl of Millipore water and 200 µl of isopropyl alcohol followed by bath sonication for 8 minutes and probe sonication for 2 minutes. The catalyst inks were painted onto a 4 cm 2 teflonized carbon cloth gas diffusion layer (GDL) until a mass loading of Pt-Ru of 0.7 mg/cm 2 was obtained for the anode gas diffusion electrode (GDE). The cathode GDE was obtained by mixing Pt/C with Millipore water and Nafion 158 (5 wt%) ionomer at a ratio of 1:10:3, followed by bath sonication for 8 minutes and probe sonication for 2 minutes. The catalyst ink was painted on either a teflonized carbon paper (060, Toray) or a teflonized carbon paper with a microporous layer (22BB, SGL Carbon) until a Pt mass loading of 1.3 mg/cm 2 was achieved. The prepared anode GDEs were placed in 1 M NaOH solution for 20 h and then washed with Millipore water. The AEM was sandwiched in between the anode and cathode GDEs and placed into a fuel cell hardware. The fuel cell polarization measurements were performed using a Fuel Cell Test System 890B (Schribner Associates). Prior to testing, a solution of 0.5 M KOH was passed through the cell for 10 minutes followed by Millipore water for 10 minutes, at room temperature. The cell temperature was raised to 60 ℃ and a solution of 2 M methanol (MeOH) and 4 M KOH was delivered through the anode compartment at 5 mL/min and recirculated to the fuel reservoir and humidified oxygen was delivered to the cathode compartment at 100 mL/min. The open circuit voltage (OCV) was monitored for 20 minutes, then the cell was held at 0.5 V for 20 minutes and then held at 0.2 V for 10 minutes. The fuel was replaced by fresh fuel and fuel cell measurements were obtained. 159 Figure 4.1. A) Chemical Structures of the polymer electrolytes used in this study a) m-TPN1- TMA, b) m-TPN1-Pip and c) m-TPN1-DMIm and B) 1 H NMR of a) m-TPBr, b) m-TPN1-TMA, c) m-TPN1-Pip and d) m-TPN1-DMIm 4.2.2. Results and Discussion Polymer electrolytes were received from Prof. Bae’s group, following a previously reported procedure.[9] Three tertiary amines, trimethylamine, N-methylpiperidine and 1,2- dimethylimidazole, were reacted with the precursor polymer m-TPBr to form ionic polymers which were named as m-TPN1-TMA, m-TPN1-Pip and m-TPN1-DMIm, respectively. m-TPN1- TMA was synthesized for use both as an ionomeric binder and membrane electrolyte. The precursor and ionic polymers were characterized by 1 H NMR spectroscopy (Figure 4.1B). The molecular weights of precursor polymer were Mn = 70,000 g/mol (Ð = 2.3) for membrane purpose and Mn = 20,000 g/mol (Ð = 2.2) for ionomeric binder purpose. The ion exchange capacity and the ionic conductivity of the membrane electrolyte and AEIs were determined and shown in table 4.1. Solution ionic conductivity measurements was utilized to determine the ionic conductivity of the AEIs due to their inability to form freestanding films when utilized in fuel cells. The ionic conductivity values are comparable to those previously reported for similar AEIs.[87] Moreover, 160 there is a slight increase in IEC and ionic conductivity of the TMA AEI compared to Pip and DMIm. Table 4.1. The ion exchange capacity (IEC) and ion conductivity of the TPN-TMA membrane and the AEIs. IEC (meq/g) Ion conductivity (mS/cm) Ion conductivity (mS/cm) addition of 4M KOH NMR Titration TPN-TMA 2.11 2.10 127 - TMA 2.11 2.11 1.33 137 PiP 1.95 1.95 1.28 126 DMIm 1.86 1.86 1.21 117 Figure 4.2. CV scan utilizing different AEI in the electrode in A) 1.0 M KOH and B 1.0 M MeOH and 1.0 M KOH at 20 mV/s. CV experiments were completed in 1 M KOH for the catalysts with all three binders (Figure 4.2A). They all present five distinct peaks; in the anodic scan peak I is ascribed to hydrogen desorption, the following peaks (II, III, IV) are ascribed to adsorption of OH - and formation of different metal oxide species. The cathodic scan presents two peaks, peak V due to reduction of the oxide species and peak VI is adsorption of hydrogen.[16,93,94] There is significant change in current for the catalyst with the m-TPN1-TMA AEI. To further assess the catalyst, the electrochemically active surface area (ECSA) was determined via CO stripping (Figure 4.3), using equation 1. ECSA = Q Cl (1) 161 In equation 1, Q is the CO stripping charge, C is the proportionality constant of 420 µC/cm 2 and l is the mass loading of catalyst on the electrode. The obtained ECSA for the three catalysts are shown in Table 4.2 and are all within the range of previously reported ECSA for PtRu/C.[95] Figure 4.2: CV in 1M KOH under N2 (black curve) and CO (Red curve) with A) DMIm AEI electrode, B) Pip AEI electrode, C) TMA AEI electrode. Moreover, the catalyst with TMA binder resulted in a larger ECSA of c.a. 30% compared to the other two catalysts, suggesting that the metal catalyst particles are more exposed to reactants when m-TPN1-TMA is used as a binder. The difference in ECSA could arise from the surface coverage of the AEIs, where the cation-hydroxide adsorption covers that catalyst surface. Even though it has been well documented that TMA + cations tend to have higher surface coverage than larger cations due to higher mobility and thus should result in lower catalytic activity, deviations from this have been reported.[10,87,90–92] Kohl et al. observed higher Pt coverage by TMA + cations compared to larger PTMA + , however, the observed current densities for MOR were higher 162 in the presence of TMA + than PTMA + . This was ascribed to a more negative potential at the outer Helmholtz plane for PTMA + which inhibits hydroxide transport to the surface.[92] Since the anodic peak for CO oxidation is dependent on adsorbed hydroxyl species, the lower ECSA for the electrodes with m-TPN1-DMIm and -Pip AEIs arises from a decrease in accessibility to hydroxide, which is evident by the lower CO oxidation peak (Figure 4.3).[96] CV experiments were completed in 1 M MeOH and 1 M KOH (Figure 4.2B). There are two oxidation peaks present which are both due to methanol electrooxidation. The onset potential in the anodic scan, defined here as when the current reaches 1 mA/cm 2 geo, for the three binders were all c.a. 0.4 V vs RHE. However, the peak current in the forward scan (If) is higher for TMA and lowest for DMIm containing electrode indicating the activity for methanol electrooxidation is dependent on the binder. This increase in catalytic activity could be due to higher ECSA for TMA containing electrode. The presence of the oxidation peak in the cathodic scan (Ib) is believed to arise from the reduction of oxide species and reactivation of catalyst active sites.[94,97] The reactivation coefficient (Ib/If) for the three binders is approximately 0.3, thus the removal of oxide species is not affected by the AEI composition. 163 Figure 4.3: CV in 1 M MeOH +1M KOH at varying scan rates with A) DMIm AEI electrode B) Pip AEI electrode and C) TMA AEI electrode. To gain further insight, CV experiments were conducted at varying scan rates (Figure 4.3). As expected, the peak current (Ip) increases with increased scan rate (v) and the peak current shows a linear relationship with the square root of scan rate (Figure 4.4A). This linear relationship confirms the methanol electrooxidation is a diffusion limited process. Furthermore, a linear relationship between the peak current potential (Ep) and the natural log of scan rate (Figure 4.4B) can be utilized to obtain the value of αn’ (Table 4.2) using equation 2. E p ln v = RT 2F(αn’) (2) In equation 2, R is the gas constant, T is temperature, F is Faraday constant and α is the charge transfer coefficient and n’ is the electrons in the rate determining step (RDS). The αn’ values for the three electrodes (Table 4.2) are similar, with DMIm having the lowest value. Assuming the RDS is a one electron process, the charge transfer coefficients can be obtained for each AEI 164 containing catalyst. More importantly, the value of αn’ can be used to obtain the diffusion coefficient with equation 3 and the slope of the plots in Figure 4.4A. I p v 1/2 = 2.99 X 10 5 n(αn’) 1/2 CD 1/2 (3) In equation 3, C is the concentration of MeOH in solution, D is the diffusion coefficient and n is the number of electrons in the entire electrochemical process. The values for MeOH diffusivity are similar to those found in a previous study, where an increase in MeOH diffusivity was observed by utilizing partially exfoliated carbon nanotubes as metal support in comparison to non-exfoliated carbon nanotubes as the metal support.[98] Here, the electrode with TMA AEI has the highest diffusion coefficient value signifying the catalysts are more accessible to MeOH. The formation of a thin film by AEIs covering the active catalysts has been reported to result in diffusion barriers for O2 and H2 accessibility for several different metal catalysts. Kim et al. observed improvements in H2 diffusion by modifying the polymer backbone of the studied AEIs and further minor improvements by replacing TMA + with a triethylammonium cation.[87] Moreover, Santori et al. showed that synthesized PPO based AEIs result in a diffusion barrier from the thin film covering the Pt/C catalysts for ORR and PtRu/C for HOR.[91] Therefore, it is likely with the larger size of MeOH compared to H2 and O2 that the thin film formation of the AEI results in a diffusion barrier. 165 Figure 4.4. A) Plot of Forward peak current vs square root of scan rate; B) plot of peak potential vs natural log of scan rate and; C) Tafel plot for the electrodes containing the three different AEI. Tafel plots were obtained from linear sweep voltammetry scans at a scan rate of 0.5 mV/s (Figure 4.4C) to obtain further insight into the role of the AEIs in the kinetics of methanol electrooxidation. The Tafel slope for MOR has been reported to have two regions; the first (Tafel- a) occurring at lower potentials corresponds to the dehydrogenation of MeOH and the second (Tafel-b) occurring at higher potentials arises from CO oxidation.[31] Here, Tafel-a is between the potential window 0.4 and 0.5 V vs RHE and Tafel-b occurs between 0.5 to 0.65 V vs RHE. The value for Tafel-a slope was smaller for TMA indicating faster charge transfer kinetics for MOR due to the higher accessibility to MeOH and the larger ECSA. The highest slope value was from DMIm AEI and is due to the lower ECSA and lower accessibility to MeOH. However, the slope for Tafel-b is similar between the three electrodes, with TMA having a slightly higher slope, even though CO stripping experiments displayed higher oxidation current for the TMA AEI electrode. However, this could be explained by the slightly higher onset potential for CO oxidation (Figure 4.2) for the TMA AEI electrode. Chronoamperometry (CA) was completed over a span of 1 h at a potential of 0.69 V vs RHE to observe the short-term stability of the AEIs (Figure 4.5A). There is an initial significant drop in current but eventually the current begins to stabilize for the three electrodes. Thus, the 166 poisoning rate was determined for each of the electrodes (Table 4.2) calculated by the following equation[99] σ = 100 I o ×( dI dt ) t>1000s (4) In equation 4, dI dt is the slope of the linear portion of the current decay and Io is the initial current. The poisoning rates were found to be similar for the three AEIs, where the poisoning rate for DMIm AEI containing electrode was the highest. This suggests that any long-term instability would be independent of the AEI. Figure 4.5. A) Chronoamperometry scans of the three AEI electrodes at 0.69 V vs RHE in 1.0 M MeOH and 1.0 M KOH. B) Nyquist plot for the three AEI electrodes at 0.74 V vs RHE and the equivalent circuit. EIS experiments were completed at varying potentials (Figure 4.6), the Nyquist plots were fitted to an equivalent circuit (Figure 4.5B) where Rsol is the solution resistance, Rct is the charge transfer resistance, CPE1 is a constant phase element due to the pseudo-capacitive nature of the catalyst inks and Ws1 is Warburg element which is associated with the finite diffusion of MeOH to the electrode. The Rct increased for all three electrodes when the potential decreased, which is due to a decrease in oxidation of MeOH. However, the electrode with DMIm had the largest Rct across all potentials which could be due to the lower diffusion of MeOH and lack of catalytic sites available. Furthermore, the TMA containing electrode had the lowest Rct due to higher accessibility 167 to reactants. The change in diffusion barrier is further confirmed by observing the diffusion resistance, obtained from the Warburg element, where the electrode with TMA AEI has the lowest resistance and the electrode containing DMIm has the highest resistance. Figure 4.6: EIS under 1M MeOH + 1M KOH at varying potentials with A) DMIm AEI electrode, B) Pip AEI electrode and C) TMA AEI electrode. Table 4.2. Summary of AEI half-cell performance, the uncertainty values were determined from the standard deviation of the repeated experiments. ECSA (m 2 /g) Onset potential (V vs RHE) αn’ D (x10 -8 cm 2 /s) Tafel-a (mV/dec) Tafel-b (mV/dec) R ct (Ω) (0.74 V vs RHE) Ws1-R (Ω) (0.74 V vs RHE) Poison Rate (%/min) DMIm 22.6±0.6 0.42±0.03 0.78±0.03 3.36± 0.04 149±2 205±1 29.0±0.3 61.0±16.5 0.72±0.05 Pip 23.4±0.8 0.40±0.03 0.83±0.03 8.45± 0.03 132±3 205±1 20.3±0.2 10.1±1.0 0.61±0.03 TMA 30.3±0.9 0.39±0.03 0.85±0.02 24.8± 0.03 123±1 211±2 18.3±0.2 6.9±0.9 0.65±0.02 In order to assess the effects of the AEI in a more practical environment, the AEIs were utilized in the anode electrodes of a single cell stack ADMFC. Fuel cell experiments were obtained 168 by scanning the potential from OCV to short circuit potential; the peak power density for the fuel cell with m-TPN1-TMA AEI was highest at 151 mW/cm 2 . Further, a commercial alkaline ionomer binder (Sustainion® from Dioxide Material) was also utilized to compare at 60 ℃ and the results are shown in Figure 4.7A. Figure 4.7. Single cell ADMFC polarization curve for the anode electrode with varying AEIs A) at 60 ℃ and B) at 80 ℃. To demonstrate that the improvement of the ADMFC is not simply due to improved AEI- membrane triple phase boundary interactions, two commercial membranes were also utilized in ADMFCs (Figure 4.8). For both ADMFCs with commercial membranes the power density is lower than the TPN-AEM ADMFC, likely due to improved ionic conductivity of our TPN AEM. Moreover, a similar trend is seen in which the ADMFC with TMA AEI displays the highest power density. Furthermore, the ADMFCs were tested in the absence of added hydroxide (Figure 4.8C- E), there is a significant decrease in the overall power output across all three membranes, likely due to the increase in resistivity and need for stoichiometric amounts of hydroxide for MOR. Nonetheless, the ADMFC utilizing the TPN membrane displays significantly higher power output compared to the commercial membranes. Further, the ADMFC with the TMA AEI maintains highest power output across all AEIs irrespective of the membrane. 169 Figure 4.8. Single cell ADMFC polarization curve for the anode electrode with varying AEIs at 60 ℃ A) with Sustainion membrane, B) with Tokuyama A201 membrane. 2M MeOH without added KOH using C) Sustainion membrane, D) Tokuyama A201 membrane and E) TPN membrane. One of the longstanding issues with previous AEMs has been instability at high temperatures under highly alkaline environments, however the prepared TPN membrane has been shown to be stable at elevated temperatures under alkaline conditions in previous studies.[8–10,87] Therefore, further experiments were completed at 80 ℃ and an increase in power density for the three ADMFCs is observed compared to the ADMFCs at 60 ℃ due to improved kinetics (Figure 4.7B). Nonetheless, the ADMFC with m-TPN1-TMA AEI remains the highest power output, 170 increasing to 212.7 mW/cm 2 . These results are in agreement with the obtained results from the half-cell studies, indicating the change in diffusion coefficient and ECSA result in improved fuel cell performance. In order to observe the short-term stability of the ADMFC, constant potential experiments were completed at 0.4 V for 1 h, and the results are shown in Figure 4.9. There was a gradual decrease in power output, over the course of 1 h, which is attributed to consumption of fuel. This was confirmed via 1 H NMR (Figure 4.9B), where after the course of 1 h the concentration of methanol was found to be c.a. 0.15 M compared to the initial 2.0 M. There is also a significant decrease of the methanol peak in 13 C NMR (Figure 4.9C) and an increase in carbonate peak, which indicates conversion of hydroxide to carbonate from the CO2 as the final product of MOR. Furthermore, there are minor peaks corresponding to HCOO - seen in 1 H NMR and 13 C NMR. This could arise from insufficient hydroxide to complete the oxidation of MeOH. Matsuoka et al. also identified HCOO - as the only product of MOR on Pt in half cell experiments with increased concentration of HCOO - as the time increased.[100] Moreover, Haisch et al. showed that the MOR on Pt surface in RDE experiments are limited by availability of OH - and performance is influenced by a pH shift at the electrode interface from the formation of carbonates.[101] Furthermore, when the fuel was replaced with fresh MeOH-KOH solution (time 60 minutes) the power was regained for all three ADMFCs, confirming the drop in ADMFC performance was due to depletion of fuel and not due to degradation of the membrane, AEIs or catalyst. These results corroborate with the half-cell experiments in which reactivation of the catalyst and poisoning rates were similar for all three binders. 171 Figure 4.9: A) Power density vs time plot for single cell ADMFC constant potential experiments held at 0.4 V. B) 1 H NMR and C) 13 C NMR of the ADMFC fuel before and after constant potential experiment with TMA anode binder. To prepare the NMR samples, 500 µl of the fuel was mixed with 500 µl of D2O and 100 µl of isopropyl alcohol (IPA) as an internal standard. Further optimization of the fuel cell was done utilizing m-TPN1-TMA as the anode AEI binder, due to its higher performance. Water management is a prominent issue in AEM fuel cells, and more water is produced in the anode and back diffusion occurs from the anode to the cathode in DLFCs. Recent studies for ADMFCs displayed high power output by incorporating reduced graphene oxide into the membrane matrix to reduce the crossover of methanol from anode to cathode and aiding in the prevention of water transport.[41] Furthermore, with the need to have O2 stream humidified the probability of cathode flooding in DLFCs is exacerbated. Recent results from a direct formate fuel cell (DFFC) have shown that cell performance was improved when the O2 humidification was decreased from 100% to 70% due to decrease in cathode flooding, however further decreasing the humidification to 50% resulted in a loss in performance confirming the need for O2 humidification.[102] A common strategy in PEM fuel cells to improve water management 172 is the use of microporous layer (MPL), however their use in AEM fuel cells have been conflicting.[103] Previous results by Kasper et al. showed decrease in performance when a MPL was added to the anode GDL and no significant change in cell performance when utilizing a MPL cathode GDL.[104] However, these results were obtained with a H2/O2 AEM fuel cells where less water is present when compared to DLFCs. Chen et al. observed that in an alkaline direct ethanol fuel cell the back diffusion of water results in cathode flooding, especially in lower current density region and slow cathode flow rates.[105] Thus, in order to improve water management in the cell, a GDL with a carbon-PTFE MPL was used as the cathode substrate. It can be seen in Figure 4.10A that an increase in power from 212.7 to 233.3 mW/cm 2 was obtained with the addition of a MPL on the cathode GDL. This increase in performance could be due to rerouting excess water to the anode exhaust. Figure 4.10. Single cell polarization curve for A) ADMFC without and with MPL cathode GDE and B) change in catalyst to AEI in anode GDE. C) Power density vs time plot for single cell ADMFC constant potential experiments held at 0.64 V 173 The ratio of catalyst to AEI was studied for the anode GDE (Figure 4.10B) as this has been shown to be a significant factor in fuel cell performance. With an increase in the binder content a major drop in performance is observed, which may arise from the catalyst being less accessible to the fuel due to a much thicker diffusion layer. When decreasing the binder content to a weight ratio of 1:1 an increase in performance is obtained of nearly 10%, and further decreasing the binder content results in drop in performance. The drop in performance with less AEI could be due to insufficient ion conductivity. Nonetheless, the obtained power density with the 1:1 ratio is amongst the highest for ADMFCs (Table 4.3). Furthermore, the stability of the ADMFC was observed by constant potential experiment at a potential of 0.64 V for 19 h, shown in Figure 4.10C. The power density had slight fluctuations for the first 8 h and then stabilized for the remainder of the time. Even though an ADMFC with higher power performance has been reported,[41] the ADMFC presented here utilizes much less precious metal in the catalyst layers while achieving high power, lowering the cost of the ADMFC. When ADMFC performance is normalized to mgPtRu the ADMFC presented here displays 359.6 mW/mgPtRu, while the ADMFC reported by Chang et al. displays 155.2 mW/mgPtRu. Further, the authors fail to mention any stability experiments to show the viability of implementing the ADMFC into practical systems. The short-term stability experiments shown here are enough to show the prospect of utilizing this ADMFC system in practical applications. Table 4.3. Comparison of maximum power densities for ADMFC Membrane Anode catalyst loading (mgPtRu/cm 2 ) Fuel ([MeOH]/[KOH]) Maximum power density (mW/cm 2 ) Ref. Tokuyama A201 8.0 1.0 M /2.0 M 160 (90℃) [28] Tokuyama A201 1.8 2.0 M /2.0 M NaOH 140 (80℃) [27] 174 KOH doped PVA 0.45 3.0 M /6.0 7.10 (80℃) [106] QPPO/PSF/TiO2 0.4 mgPt 2.0 M /2.0 M 118 (60 ℃) [107] AQPVBH-C 1.0 1.0 M /1.0 M NaOH 53.2 (60 ℃) [32] Polymer fiber membrane 1.0 1.0 M /4.0 M 93 (60 ℃) [35] QCS/QSiO2@PVDF 4.0 2.0 M /5.0 M 98.7 (80℃) [37] QCS/PVA-1%- LDH@CNT 4.0 2.0 M /5.0 M 107 (80℃) [36] PVA-b-PVBTAC - 1.5 M MeOH 99.6 (80℃) [34] PBI 2.0 2.0 M /6.0 M 253.9 (80℃) [41] PBI/NGO 2.0 2.0 M /6.0 M 310.3 (80℃) [41] TPN 0.7 2.0 M /4.0 M 251.7 (80℃) This Work This section of the chapter is reprinted with permission from ACS Appl. Energy Mater. 2021, 4, 6, 5858–5867. Copyright 2021 American Chemical Society.[108] 4.3. Development of Nickel Gallium Electrocatalysts for Methanol Electrooxidation and Platinum Free Alkaline Direct Methanol Fuel Cells 4.3.1. Experimental Catalyst synthesis GaNi/C The gallium nickel supported on Vulcan XC72-R carbon (GaNi/C) was synthesized via impregnation utilizing hydrazine hydrate as the reducing agent. A 0.05 g of XC72-R was dispersed 175 in 30 mL of ethylene glycol and vigorously stirred and sonicated for 1 h. Followed by the addition of Ga2(SO4)3 and NiCl2 at appropriate amounts in order to obtain weight ratios of Ga:Ni of 4:1, 3:1, 2:1 and 1:3. The aqueous mixture was stirred and sonicated for 1 h. The aqueous mixture was then placed in an oil bath and heated to 100 ℃, followed by the slow addition of 5 mL of hydrazine hydrate. The reaction mixture was stirred for 40 minutes. The resulting solution was vacuum filtered, and the obtained powder was washed and centrifuged 4 times with Millipore water (Direct-Q UV, 18.2 MΩ). Ni/C and Ga/C were also prepared following a similar procedure, however without the addition of Ga2(SO4)3 and NiCl2, respectively. MnO2/C for ORR The α-MnO2 catalyst was prepared by adding 1.014 g MnSO4·H2O to 20 mL Millipore water and stirred and sonicated. 0.95 g KMnO4 was then dissolved in 20 mL Millipore water and slowly added to the solution followed by 30 min of stirring and sonication. The solution was then transferred to an oil bath and kept under reflux at 100 ℃ for 20 h. Afterwards, the solution was washed and centrifuged with Millipore water until the pH of the supernatant was neutral and dried in an oven. The α-MnO2 was then added to an aqueous solution containing XC72-R, the reaction mixture was stirred overnight. The α-MnO2 loading was 20 wt%. Catalyst Characterization Powder X-Ray Diffraction (XRD) was performed on a Rigaku Ultima Diffractometer with a Cu Kα (0.154 nm) radiation source and a scan rate of 4°/min from a 2θ value of 20° to 90°. Scanning electron microscopy (SEM) and energy dispersive X-ray spectroscopy (EDS) micrographs were obtained from a JEOL-JSM-7001F electron microscope with an accelerating voltage of keV. Thermogravimetric analysis (TGA) was performed on a Shimadzu TGA-50. X- 176 ray photoelectron spectroscopy (XPS) spectra were obtained from a Kratos Axis Ultra DLD using a mono Al anode with a pass energy of 160 keV for the survey scan and 20 keV for the high- resolution scan. A Shirley background was utilized for all datasets and all datasets were referenced to C 1s binding energy of 284.8 eV. Electrochemical Measurements Half-cell experiments were completed utilizing a standard three electrode cell, with a glassy carbon rotating disk electrode (RDE), platinum wire and Hg/HgO (4.24 M KOH filling solution) as the working, counter, and reference electrodes, respectively. All measurements were obtained on a Solatron SI 1287 potentiostat and Solatron SI1260 impedance phase analyzer for electrochemical impedance spectroscopy (EIS). Catalyst ink solutions were prepared by adding 2.5 mg of catalysts to 900 µl of Millipore water, 100 µl of isopropyl alcohol and 8 mg of TPN- TMA alkaline binder solution. The ink was sonicated for 10 minutes and then 20 µl of the ink was drop cast onto the glassy carbon electrode (Surface area = 0.195 cm 2 ) and dried. Prior to all experiments the electrolyte solution was purged with ultrapure N2 for 15 min. The electrochemically active surface area (ECSA) of the catalyst was obtained using the capacitance method. The CV scans were completed between 0.05 and 0.15 V vs. MMO at various scan rates. The peak current was then plotted against the scan rate to obtain the double layer capacitance value (Cdl). The ECSA was then calculated using equation 5, where c is the specific charge density 40 µF/cm 2 and m is the mass of the catalyst (0.04 mg).[109,110] ECSA = C dl c × m (5) Membrane electrode assembly (MEA) and Fuel Cell Testing 177 Single fuel cell assemblies were utilized to further assess the performance of the catalysts. The as synthesized GaNi/C was used as the anode catalyst and either Pt/C (40 wt % Platinum) or MnO2/C was used as the cathode catalyst in the MEA. The anode catalysts were hand brush painted onto carbon cloth (4 cm 2 ) and the catalysts ink were prepared by adding 20 mg of catalysts, 200 µl of Millipore water, 200 mg of isopropyl alcohol and 40 mg of TMA binder. The inks were bath sonicated for 10 min followed by probe sonication for 2 min, with a 10 s on and 10 s off cycle. The catalysts were applied until a loading of 4 mg/cm 2 was obtained. The cathode gas diffusion electrode (GDE) was prepared by hand painting the catalysts onto a teflonized carbon paper with a microporous layer (22BB, SGL Carbon). The Pt/C inks were prepared by adding 20 mg of Pt/C to 200 mg of Millipore water and 60 mg of Nafion dispersion (5 wt %). The MnO 2/C ink was prepared by adding 30 mg of catalyst to 200 mg of Millipore water and 90 mg of Nafion dispersion (5 wt %). Both were then bath sonicated for 10 minutes. The MEA was prepared by sandwiching the TPN-TMA membrane in between the anode and cathode electrodes and cold pressing for 5 minutes at 500 PSI. The fuel cell polarization measurements were performed using a Scribner fuel cell test system 890B. The MEAs were conditioned by passing 0.5 M KOH through the cell for 10 min followed by Millipore water for 10 min at room temperature. The cell temperature was then raised to 60 ℃ and a solution of 1 M methanol (MeOH) and 1 M KOH was delivered through the anode compartment at a flow rate of 10 mL/min and recirculated to the fuel reservoir. Humidified oxygen was delivered to the cathode compartment at a flow rate of 100 mL/min. The open circuit voltage (OCV) was monitored for 20 min; then the cell was held at 0.2 V for 30 min. The fuel was then replenished with fresh fuel and the fuel cell measurements were obtained. 178 4.3.2. Results and Discussion Catalyst Characterization In order to determine the metal content of the catalysts TGA was completed for all catalysts. The target loading was 50% and as shown in Figure 4.11A, the actual loadings vary between 45-55%. Figure 4.11B shows the XRD pattern for the prepared catalysts. The catalyst containing only Ni presents two distinct crystal phases, those corresponding to Ni(OH)2 and Ni, while the catalyst with only Ga displays peaks corresponding to GaOOH. Further, the catalysts synthesized with both Ni and Ga display a significantly different pattern. The signal for Ni (111) at 45° diminishes significantly and broadens for all catalysts except for the GaNi3/C catalyst. The decrease in Ni (111) indicates minimal mixture of intermetallic GaNi/C and Ni/C. Moreover, there are two broad peaks observed which correspond to a Ga-Ni oxide and indicate a decrease in particle size or a loss of crystallinity.[111] Further, there is a broad signal at c.a. 25° corresponds to the carbon support. 179 Figure 4.11. A) TGA curve for prepared catalyst B) XRD pattern of the prepared catalyst where • corresponds to Ni and ° corresponds to Ni(OH)2. SEM images of B) Ni/C C) GaNi3/C D) Ga2Ni/C E) Ga3Ni/C F) Ga4Ni/C and G) Ga/C. To gain insight into the morphology of the catalyst SEM images were obtained. The Ga/C catalyst display rod like structures, confirming the preparation of GaOOH rods.[73,76] However, 180 the rod like structures are not present in the GaNi catalysts, rather spherical clusters are observed more akin to the morphology of the Ni/C catalyst. The elemental composition was obtained by EDS and composite mappings are shown in Figure 4.12. There is uniform dispersion of Ni and Ga across the carbon substrate. Furthermore, the weight percentage of the elements are provided in Table 4.4 and the metal content for the catalyst is similar to that determined by TGA, between 40- 60 wt %. 181 182 Figure 4.12. SEM and EDS images for A) Ni/C B) GaNi3/C C) Ga2Ni/C D) Ga3Ni/C E) Ga4Ni/C and F) Ga/C. Where blue corresponds to Ni, yellow corresponds to Ga, green corresponds to O, and red corresponds to C. To gain further insight into the elemental composition XPS analysis was completed on the catalysts (Figure 4.13). Based on the survey scan, only Ni, Ga, C and O were present in the catalyst. However, the observed amount of metal content differs from that quantified by EDS which is due to the differences in penetration depth between XPS (<5 nm) and EDS (1-3µm). Thus, the surface of the catalysts differs from the bulk, where the surface of the catalysts contains a higher oxygen content indicating a higher degree of oxidation of the surface catalysts. Figure 4.13B shows the high-resolution scans of the Ni 2p region displaying two peaks at approximately 854 and 874 eV, 183 corresponding to Ni 2p3/2 and Ni 2p1/2 as well as their two satellite signals near 862 and 880 eV, respectively. Furthermore, the high-resolution scans of the Ga 2p region display one signal at 1118 eV, which displays a shift in the binding energy when comparing the Ga/C to Ni containing catalysts. This shift in binding energy could arise from the difference in Ga structure and lack of GaOOH present. Figure 4.13. A) Survey scan B) Ni 2p C) Ga 2p D) C 1s and E) O 1S XPS spectra of the prepared catalysts. 184 Table 4.4. Elemental composition of the catalysts from EDS and XPS. C (wt%) O (wt%) Ga (wt%) Ni (wt%) EDS XPS EDS XPS EDS XPS EDS XPS Ni/C 40.75 35.12 6.36 37.46 - - 52.58 27.42 Ga/C 49.47 61.89 11.30 22.07 39.02 16.04 - - Ga4Ni/C 37.01 53.22 13.27 23.94 39.55 13.62 10.17 9.22 Ga3Ni/C 39.69 36.77 17.06 39.64 31.03 10.54 12.22 13.05 Ga2Ni/C 46.14 45.58 11.38 25.88 25.89 9.02 16.59 19.52 GaNi3/C 45.09 41.96 6.61 45.18 12.10 4.91 36.19 7.95 Electrochemical Characterization Prior to studying the electrochemical catalytic activity of the catalysts for MOR, CV experiments of the catalysts in only the supporting electrolyte were first completed (Figure 4.14A). The Ni/C catalyst displays four major peaks, the first and second oxidation peaks occurring at 0.5 V and 0.7 V vs. MMO are due to the conversion of Ni(OH)2 to NiOOH, where the first is the more ordered state of NiOOH and the second due to a disordered NiOOH structure. The oxidation peak at 1.0 V vs MMO is due to oxygen evolution reaction. In the reverse scan, there is a reduction peak at 0.4 V vs. MMO which is due to the reduction of NiOOH to Ni(OH)2. As seen in Figure 4.14A the addition of Ga to the catalysts shifts the major oxidation to ca 0.7 V vs. MMO, indicating the formation of a disordered Ni 3+ , which is due to a disordered structure from the Ga incorporation as compared to the Ni(OH)2. The ECSA of the catalysts was obtained utilizing the capacitance method, with the CVs at the varying scan rates shown in figure 4.14 and the calculated ECSA 185 values shown in Table 4.5. Here we can see that the Ni/C has the largest ECSA, likely due to the higher Ni content in the catalysts. Figure 4.14. CV experiments of the prepared catalysts in A)1M KOH. CVs at varying scan rates in 1M KOH for B) Ni/C C) Ga4Ni/C D) Ga3Ni/C E) Ga2Ni/C F) GaNi3/C under N2 flow. The MOR was then assessed in half-cell experiments, as shown in Figure 4.15A there is a significant increase in the current in the forward scan after 0.5V vs. MMO which is attributed to the oxidation of MeOH. Furthermore, as the oxidation of Ni 2+ to Ni 3+ occurs at 0.5V vs. MMO 186 this is further evidence of Ni 3+ being the active species for MOR.[112] Moreover, the addition of Ga into the catalysts has a prominent role in the electrooxidation of MeOH. There is an enhancement in the peak current density when the Ga:Ni is 3:1 and a decrease when more Ga is present. Furthermore, the current decreases when Ga:Ni is 1:3, this could be a due to higher mixture of GaNi and Ni. Moreover, based on the XPS elemental analysis the surface of the GaNi3/C contains minimal amount of metal content, which could also result in the decreased performance. While the catalyst with only Ga displays oxidation for methanol, it is orders of magnitude lower than Ni containing catalysts. Moreover, the CV for the Ga/C catalyst displays two oxidation peaks, one in the forward scan and one in the reverse scan. The oxidation in the reverse scan has been attributed to removal of adsorbed species onto the surface of the catalysts. The increased catalytic activity for MOR with Ga incorporation into Ni could arise from the change in the structure of the catalysts exposing more of the Ni sites and the Ga aids in removal of intermediates. Furthermore, since the ECSA of Ga3Ni/C is lower than the Ni/C catalysts, this signifies an intrinsic increase in the catalytic activity for MOR. 187 Figure 4.15. A) CVs of the prepared catalyst in1M KOH with 1M methanol at a scan rate of 20 mV/s B) Nyquist plot for the prepared catalysts at a constant potential of 0.7 V vs. MMO and C) Stability experiments of the prepared catalysts holding a constant potential of 0.7 V vs. MMO. To gain further insight into the differences in catalytic activity EIS experiments were completed. As shown in Figure 4.15B, all catalysts display a depressed semi-circle shape in the Nyquist plot, with similar x-intercept values in the high frequency region indicating similar solution resistance. The Rct, however, displayed significant differences across all the catalysts, with Ga4Ni/C having the largest resistance and Ga3Ni/C having the lowest resistance; similar to the trend observed with the CVs. Moreover, the improvement is also attributed to the charge transfer resistance of the catalyst. To further asses the efficacy of the catalysts short term stability experiments were completed by holding a constant potential of 0.7 V vs. MMO for 1 h. All catalysts displayed minimal decay in current density throughout the 1 h, indicating stability of the catalysts. Moreover, the decay rate decreased with the addition of Ga into the structure indicating the Ga aids in removal of any adsorbed intermediates improving the stability of the catalyst. 188 Table 4.5. Summary of half-cell experiments ECSA (cm 2 /g) Ip (mA/cm 2 ) Rct (Ω) Poison Rate (%) Ni/C 49.5 104.1 15.38 0.13 Ga/C - 1.1 - 0.06 Ga4Ni/C 25.2 27.6 43.98 0.03 Ga3Ni/C 36.8 156.2 9.99 0.02 Ga2Ni/C 26.9 53.5 16.37 0.05 GaNi3/C 35.4 61.0 17.44 0.10 In order to fully display the efficacy of the catalysts for ADMFCs, the catalysts were utilized in single cell ADMFCs. As seen in Figure 4.16A all the prepared catalyst are able to oxidize MeOH in an ADMFC. The OCV for all the catalysts was similar between 0.5 and 0.6 V, with the lowest OCV from the Ga4Ni/C. Furthermore, a similar trend to the half-cell performance is seen in the fuel cells, where the highest power density is obtained from Ga3Ni/C followed by the Ni/C and lowest performance from the Ga4Ni/C. The improved performance of the Ga3Ni/C is likely due to the improved charge transfer resistance from addition of Ga to the catalyst. In order to further improve the fuel cell performance, different parameters were altered while utilizing the Ga3Ni/C catalyst as the anode electrode. The flow rate of the MeOH was varied, and increasing the flow rate increases the peak power density however, the power begins to decrease at 200 ml/min and significantly decreases at 300 mL/min. This initial increase could arise from improved accessibility to MeOH to the catalyst and then decreases once there is excess liquid causing catalyst flooding and increased crossover of fuel stream to the cathode. The concentration of the MeOH was then changed to observe its role in the fuel cell performance. An increase in power was observed when the concentration was increased to 1.5 M but then decreases with 2 M and further decreases when the concentration is below 1 M. The power remained similar when the supporting 189 electrolyte concentration was increased to 8 M but decreased when it was less than 4M. This could be due to the maximum conductivity of KOH being achieved at 4M. The highest performing ADMFC was obtained with a peak power density of 25 mW/cm 2 , which is on par with several previously reported ADMFC and DMFCs with precious metals and amongst the highest reported for nonnoble metal DMFCs.[56,67,68,113,114] Figure 4.16. Polarization curve for fuel cells with A) varying anode electrode, B) varying methanol concentration, C) varying KOH concentration and D) α-MnO2/C cathode electrode. Furthermore, a completely precious metal free ADMFC was also prepared utilizing α- MnO2/C catalyst for the cathode electrode with either Ni/C or Ga3Ni/C as the anode electrode. The peak power for the GaNi/C containing ADMFC remained higher than that of the Ni/C. However, utilizing the noble metal free ADMFC did display a decrease in the peak power density of the ADMFC compared to the Pt/C containing ADMFC. This decrease is due to the higher catalytic 190 activity of Pt/C compared to MnO2/C. Nonetheless, the results display one of the first noble metal free DMFCs, with peak power density not far from reported DMFCs with noble metals. 4.4. Improving Noble Metal Free Alkaline Direct Methanol Fuel Cells with CeO2/C Supported Ni electrocatalyst 4.4.1. Experimental Catalyst synthesis CeO2/C The CeO2/C was prepared via mixing 2 g of Vulcan XC72-R and 5 g of Ce(NO3) in 125 mL of Millipore water and sonicating for 30 minutes. 1 mL of 2M KOH was then added to the solution mixture and stirred at room temperature for 2 h. The solution mixture was then centrifuged, and the supernatant was discarded. The powder was washed and centrifuged with water until the pH was neutral. The obtained powder was dried at 55 ℃ overnight. 4 g of the obtained powder was then placed in a tube furnace and heated to 400 ℃ with a ramp rate of 5 ℃/min and held at 400 ℃ for 4 hours under a flow of Ar (15mL/min). NiCeO2/C The NiCeO2/C was prepared via hydrazine reduction, 0.205 g of NiCl2 was mixed with 0.05 g of the prepared CeO2/C in 30 mL of ethylene glycol. The solution mixture was stirred and heated to 100 ℃, followed by the addition of 5 mL of hydrazine hydrate. The solution mixture was kept stirring at 100 ℃ for 2 h. The powder was obtained by centrifuging and washing with water. A catalyst of Ni/C was also prepared following a similar procedure however Vulcan XC72R was utilized in place of the CeO2/C. 191 FegCN for ORR The FegCN catalyst was obtained by direct pyrolysis of the mixture. 0.102 g of FeCl 3 and 3.8 g of urea were mixed in a quartz tube and placed in a tube furnace. The mixture was then heated to 1000 ℃ at a rate of 10 ℃/min and held at 1000 ℃ for 3 under a flow of Ar (100 mL/min). Catalyst Characterization Powder X-Ray Diffraction (XRD) was performed on a Rigaku Ultima Diffractometer with a Cu Kα (0.154 nm) radiation source and a scan rate of 4°/min from a 2θ value of 20° to 90°. Scanning electron microscopy (SEM) and energy dispersive X-ray spectroscopy (EDS) micrographs were obtained from a JEOL-JSM-7001F electron microscope with an accelerating voltage of 15 keV. Thermogravimetric analysis (TGA) was performed on a Shimadzu TGA-50 under air at a flow rate of 40 mL/min. X-ray photoelectron spectroscopy (XPS) spectra were obtained from a Kratos Axis Ultra DLD using a mono Al anode with a pass energy of 160 keV for the survey scan and 20 keV for the high-resolution scan. A Shirley background was utilized and the C 1s (284.8 eV) was used as a reference. Electrochemical Measurements Half-cell experiments were completed utilizing a standard three electrode cell, with a glassy carbon rotating disk electrode (RDE), platinum wire and silver-silver chloride (saturated KCl filling solution) as the working, counter, and reference electrodes, respectively. All measurements were obtained on a Solatron SI 1287 potentiostat and Solatron SI1260 impedance phase analyzer for electrochemical impedance spectroscopy (EIS). Catalyst ink solutions were prepared by adding 2.5 mg of catalysts to 900 µl of Millipore water, 100 µl of isopropyl alcohol and 8 mg of a TMA alkaline binder solution. The ink was sonicated for 10 minutes and then 20 µl 192 of the ink was drop cast onto the glassy carbon electrode (surface area = 0.195 cm 2 ) and dried. Prior to all experiments the electrolyte solution was purged with ultrapure N2 for 15 min. The electrochemically active surface area (ECSA) of the catalyst was obtained using the capacitance method. The CV scans were completed between 0.05 and 0.15 V vs. Ag/AgCl at various scan rates. The peak current was then plotted against the scan rate to obtain the double layer capacitance value (Cdl). The ECSA was then calculated using equation 5 from section 4.3.1. Membrane electrode assembly (MEA) and Fuel Cell Measurements Single fuel cell assemblies were utilized to further assess the performance of the catalysts. The as synthesized catalysts were used as the anode catalyst and either Pt/C (Alfa Aesar, 40 wt % Platinum) or FegCN was used as the cathode catalyst in the MEA. The anode catalysts were hand brush painted onto carbon cloth (4 cm 2 ) and the catalysts ink were prepared by adding 20 mg of catalysts, 200 µl of Millipore water, 200 mg of isopropyl alcohol and 22 mg of TMA binder solution. The inks were bath sonicated for 10 min followed by probe sonication for 2 min, with a 10 s on and 10 s off cycle. The catalysts were applied until a loading of 4 mg/cm 2 was obtained. The cathode gas diffusion electrode (GDE) was prepared by hand painting the catalysts onto a teflonized carbon paper with a microporous layer (22BB, SGL Carbon). The Pt/C inks were prepared by adding 20 mg of Pt/C to 200 mg of Millipore water and 60 mg of Nafion dispersion (5 wt %). The FegCN ink was prepared by adding 20 mg of catalyst to 200 mg of Millipore water and 60 mg of Nafion dispersion (5 wt %). Both were then bath sonicated for 10 minutes and painted onto the carbon paper until a loading of 4 mg/cm 2 was obtained. The MEA was prepared by sandwiching the TPN-TMA membrane in between the anode and cathode electrodes and cold pressing for 5 minutes at 500 psi. 193 The fuel cell polarization measurements were performed using a Scribner Fuel Cell Test System 890B. The MEAs were conditioned by passing 0.5 M KOH through the cell for 10 min followed by Millipore water for 10 min at room temperature. The cell temperature was then raised to 60 ℃ and a solution of 1 M MeOH and 1 M KOH was delivered through the anode compartment at a flow rate of 10 mL/min and recirculated to the fuel reservoir. Humidified oxygen was delivered to the cathode compartment at a flow rate of 100 mL/min. The OCV was monitored for 20 min; then the cell was held at 0.2 V for 30 min. The fuel was then replenished with fresh fuel and the fuel cell measurements were obtained. 4.4.2. Results and Discussion The prepared catalyst was characterized via TGA in order to determine the mass of the metal. The TGA plots are shown in Figure 4.17A, here it can be seen that the amount of Ce on the carbon was roughly 60 wt%, similarly the Ni content in the Ni/C was approximately 60 wt%. For the NiCeO2/C catalyst there was a mass loss of nearly 15 wt%, indicating majority of the catalyst was Ni and Ce. Powder XRD analysis were completed on the prepared materials (Figure 4.17B). Here, we can see that the Ni/C catalyst contains peaks for two phases of Ni; Ni(OH)2 and metallic Ni. Thus, as shown in section 4.3.2 this synthetic procedure can result in the formation of metal hydroxides. The CeO2/C only displayed peaks corresponding to CeO2 and the NiCeO2/C contained peaks corresponding to CeO2 and Ni, indicating proper mixture of the two metals. Moreover, there were no peaks corresponding to Ni(OH)2 or any other Ni-oxide, indicating the CeO2 aids in formation of metallic Ni. 194 Figure 4.17. A) TGA plot and B) XRD spectra for the prepared catalyst. To gain further insight into the elemental composition and morphology of the catalyst, SEM images were obtained (Figure 4.18). Both catalysts have similar morphology of large clusters, however the CeO2 containing catalyst does contain larger agglomerates which are the CeO2. To further confirm the elemental distribution, EDS images were obtained, and we can see that the Ni is well dispersed throughout the substrate (Figure 4.18C and D). Moreover, the Ni is well dispersed on the CeO and C in the NiCeO2/C catalysts. The mass percentage of the elements are provided in table 4.6 and they are similar to the target values and to the values obtained by TGA. 195 Figure 4.18. SEM images of A) Ni/C and B) NiCeO2/C. SEM and EDS images of C) Ni/C, where blue corresponds to Ni, red corresponds to C and green corresponds to O and D) NiCeO2/C, where blue corresponds to Ce, yellow corresponds to Ni, red corresponds to C and green corresponds to O. To gain further insight into the elemental composition of the catalyst surface, XPS analysis was completed on the samples (Figure 4.19A-E). From the survey scan we can see that the Ni/C catalyst contains only Ni, C and O; while the NiCeO2/C contains Ni, Ce, C, and O. Furthermore, the survey scan was utilized to obtain the mass percentage of the elements, shown in table 4.6. Here we can see that the element composition between EDS and XPS differ, and this is due to the difference in penetration depth. The values obtained from EDS provide the bulk composition, whereas the values obtained from XPS provide the composition of the surface (<5 nm). There is a larger presence of oxygen on the surface, which indicates oxidation of the Ni catalysts, moreover 196 there is less of the Ce on the surface. However, both catalysts have similar Ni content on the surface, thus any difference in catalytic activity would be due to the addition of Ce and not from surface available Ni. High resolution scans of Ce 3d contains six distinct peaks denoting the Ce 3d3/2 and Ce 3d5/2 at 881.68 eV, 889.42 eV, 897.84 eV, 900.55 eV, 905.51 eV and 916.38 eV which are denoted as v, v, v, u, u’ and u, respectively.[82,115] The presence of u” is the primary characteristic peak for Ce 4+ and peak v is the fingerprint peak for CeO2. Moreover, all the peaks are characteristic of Ce 4+ , even though there is broadening of peak u’, there are no indications of Ce 3+ peaks.[116] The high-resolution scans of the Ni 2p displays four peaks near 855, 872, 862 and 882 eV which correspond to the Ni 2p3/2, Ni 2p1/2 and two satellite peaks, respectively. All these peaks are characteristic of Ni 2+ , however, there is a shift in the binding energies of the Ni 2p3/2 when comparing the two catalysts (Figure 4.19C).[117,118] This shift in binding energy to a lower binding energy could indicate a partial electron transfer from the CeO2 to the Ni in the NiCeO2/C catalyst.[79] Furthermore, the O 1s and C 1s analysis of both catalyst demonstrated similar spectra. 197 Figure 4.19. XPS A) survey scan, high resolution scans of the B) Ce 3d, C) Ni 2p, D) C 1s and E) O 1s region for the prepared catalysts. Table 4.6. Elemental composition of the catalysts from EDS and XPS. EDS XPS Ni (wt%) Ce (wt%) C (wt%) O (wt%) Ni (wt%) Ce (wt%) C (wt%) O (wt%) Ni/C 64.27 - 31.81 3.92 19.14 - 44.36 36.50 NiCeO2/C 46.86 25.01 19.2 9.01 19.60 5.84 36.79 37.77 The electrochemical activity of the catalyst was first observed in 1M KOH (Figure 4.20A), here we can see that the Ni containing catalyst have four distinct regions. The oxidation peak labeled I occurs from the oxidation of Ni 2+ to Ni 3+ , similarly with peak II and peak II is attributed to a disordered state of Ni 3+ . Peak labeled III is due to the oxygen evolution reaction, then in the reverse scan peak IV is the reduction of Ni 3+ to Ni 2+ . While the CeO2/C catalysts doesn’t show any 198 apparent oxidation or reduction in the tested potential window apart from oxygen evolution after 0.8 V vs. Ag/AgCl. Moreover, the current for peak I and II are significantly higher for the NiCeO2/C catalysts than Ni/C, even though the Ni content on the surface is similar for both catalyst and less in the bulk for NiCeO2/C. Thus, the CeO2 likely aids in Ni accessibility and conversion from Ni 2+ to Ni 3+ . To further asses the electrochemical activity of the catalyst, the ECSA was obtained utilizing the capacitance method (Figure 4.20B and C) and shown in table 4.7.[110] Here we can see that the CeO2 addition improves the ECSA even though the Ni content in both electrocatalysts is similar. Therefore, this increase could arise from the lowered binding energy of Ni (Figure 4.19C) and from higher accessibility of Ni. Figure 4.20. CV scans of the prepared catalyst in 1M KOH A) at scan rate of 20 mV/s, B) Ni/C at various scan rates and C) NiCeO2/C at various scan rates under flow of N2. The catalysts were then studied in an electrolyte solution containing 1M MeOH and KOH in order to observe the electrocatalytic activity towards MOR. CVs were obtained in the same 199 potential window as the experiments in the absence of MeOH shown in figure 4.21A. A significant increase in current is observed after 0.5 V vs. Ag/AgCl, which arises from oxidation of methanol. Furthermore, this confirms the activation of methanol occurs with the formation of Ni 3+ as mentioned in section 4.3.2, moreover CeO2/C displayed negligible currents towards MOR. We can see that the peak current for NiCeO2/C is significantly higher than that of Ni/C and thus acts as a better electrocatalysts. This improvement arises from the improved ECSA and accessibility of Ni. Figure 4.21. A) CV at a scan rate of 20 mV/s, B) potentiostatic hold at 0.7 V vs. Ag/AgCl and C) EIS at 0.5V vs. Ag/AgCl of the prepared catalyst in 1 M KOH + 1M MeOH under flow of N2. To further asses the difference in catalytic activity towards MOR, potentiostatic experiments were completed on the Ni/C and NiCeO2/C catalysts (Figure 4.21B). Both catalysts displayed similar trends with a slight initial increase followed by a slight decrease with an improved poisoning rate for the NiCeO2/C catalyst, thus the addition of CeO2 provides improvements to the catalyst. EIS was conducted at a constant potential of 0.5V vs. Ag/AgCl 200 (Figure 4.21C), here we can see that both display a depressed semicircle in the Nyquist plot and have similar high frequency resistance values, indicating the electrolytic resistance is the same. Furthermore, the Rct, shown in table 4.7, are significantly different for the two catalysts, where the addition of CeO2 significantly decreases the Rct. This drastic decrease in the Rct likely arises from the change in binding energies observed in the XPS. Figure 4.21D displays the Tafel plots that were extrapolated from linear scan voltammetry experiments at 0.5 mV/s for both electrocatalysts with the Tafel slopes shown in table 4.7. There is a decrease in the Tafel slope for the NiCeO2/C catalysts when compared to Ni/C, indicating improved charge transfer kinetics. Table 4.7. Summary of half-cell experiments ECSA (cm 2 /g) Ip (mA/cm 2 ) Rct (Ω) at 0.5 V vs. Ag/AgCl Tafel Slope (mV/dec) Poison Rate (%) Ni/C 54.2 103.5 236 176.4 0.53 NiCeO2/C 65.5 155.2 82.1 253.7 0.39 The prepared catalysts were then utilized in ADMFCs, and their performance was assessed at various temperatures (Figure 4.22A and B). Both ADMFCs display in increase in power as the temperature is increased and this is due to improved kinetics. However, NiCeO 2/C displayed higher power densities at all temperatures when compared to Ni/C, confirming the catalyst is more effective at oxidizing MeOH. Furthermore, the OCV of the NiCeO2/C ADMFC was 50 mV higher than the Ni/C further demonstrating improved catalytic activity. This improvement arises from the improved ECSA and R ct from the addition of CeO2. The ADMFC was further improved by adjusting the fuel concentration increasing the peak power density to 28 mW/cm 2 (Figure 4.22C), which to the best of our knowledge is amongst the highest reported performing noble metal free ADMFCs and comparable to various noble metal containing DMFCs.[56,67,68,113,114] The 201 cathode electrocatalyst was also replaced with a noble metal free catalyst, FegCN to further demonstrate a platinum group metal free ADMFC. As shown in figure 4.22D the peak power density of the ADMFC without platinum was able to produce higher power than the ADMFC with Ni/C and Pt/C electrocatalyst. Figure 4.22. Polarization curves with A) NiCeO2/C anode and B) Ni/C electrodes. C) NiCeO2/C anode electrode with 1.5M MeOH + 6M KOH at varying temperatures. D) Comparison of ADMFCs with Pt/C and FegCN cathode electrodes. The stability and products of the platinum group metal free ADMFC were then studied via holding the fuel cell at a constant potential for 3 hours (Figure 4.23A). It can be seen that the current decreases initial but then stabilizes after about 1 hour. The products were collected after the 3 hours and analyzed via 1 H NMR (Figure 4.23B). The only observable product is formate which has been shown to be the product of MOR for Ni based catalysts. Moreover, we can see that there is trace amounts of MeOH after the experiment, and this could be due to consumption of methanol from the reaction as well as boiling of MeOH from holding the cell at 90 ℃. 202 Figure 4.23. A) ADMFCs with NiCeO2/C anode and FegCN cathode electrodes held at a constant potential of 0.6 V. B) 1 H NMR of the fuel reservoir before and after the experiment, 100 µL of n- butanol was added as an internal standard. 4.5. Conclusion Herein, we show one of the highest performing ADMFCs with great stability at high temperatures. Moreover, we provide insight into the importance of AEIs in DLFCs and how small changes can have large impacts on catalytic activity. Half-cell experiments showed that changes in the cation of the AEIs leads to differences in catalyst accessibility, in which we observed TMA cation give the largest ECSA and have highest accessibility to MeOH. Utilizing the prepared AEIs in single cell ADMFCs a similar trend to that of the half-cell results was observed, where the TMA containing AEI ADMFC produced the highest power density output of 251.7 mW/cm 2 . Moreover, these results were achieved utilizing less than 1 mgPtRu/cm 2 in the catalyst layer, decreasing the overall cost of the ADMFC. Further, improvements may be achieved by developing a more appropriate AEI for the cathode GDE which could bring closer the commercialization of ADMFCs. We also demonstrated GaNi based catalysts which were effective in oxidizing MeOH in alkaline conditions. The incorporation of Ga into a Ni catalyst improved the catalytic activity and stability of the catalysts, where a 3:1 ratio of Ga:Ni gave the highest performance in half cell 203 experiments. This improvement was attributed to improvements in R ct for the prepared catalyst. Furthermore, the catalysts were utilized in an ADMFC in which similar trends observed in half cell experiments were observed in terms of power density. The catalytic activity of Ni was further improved by preparing the nanoparticles on a CeO2/C support. The improvement was observed in half-cell experiments and fuel cell performance. The improved activity is attributed to a decrease in binding energy observed in Ni XPS. This decrease results in an improved ECSA, Rct and kinetics which were observed in half- cell experiments. Furthermore, both the GaNi and NiCeO2/C catalysts were utilized as anode GDE in ADMFCs with MnO2/C and FegCN cathode electrodes to demonstrate platinum group metal free ADMFCs. 4.6. 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The Role of the Auger Parameter in XPS Studies of Nickel Metal, Halides and Oxides. Phys. Chem. Chem. Phys. 2012, 14 (7), 2434–2442. https://doi.org/10.1039/c2cp22419d. 223 Chapter 5: Electrochemical CO2 Reduction 5.1. CO2 Reduction Electrochemical CO2 reduction (ECO2R) utilizing renewable energy sources is considered a promising approach to a more sustainable carbon economy. The catalysts utilized in ECO2R are one of the key factors in the performance and products of ECO2R. Of the many metals used as catalysts, copper metal has remained one of the few that can not only reduce CO2 but produce a plethora of products ranging from C1 up to C3 products.[1–8] The reason for the large variety of products obtained by Cu based catalysts is due to the moderate binding energy of Cu to carbon containing intermediates.[9,10] Cu also has the added benefit of being a low cost and highly abundant transition metal, which would allow for more facile commercialization. However, the use of Cu catalysts often results in poor selectivity and a large mixture of products. Several strategies have been employed to improve the product selectivity such as modifying the Cu surface, the catalysts porosity and incorporating other elements to act as cocatalysts or to modify the catalysts’ electronic structure.[1,11–13] Of the many products obtained by ECO2R, production of ethylene glycol is of great interest due to its wide use such as in CO2 capture technology, polyester production, manufacturing technologies, automotive and transportation and as a coolant.[14,15] Currently, ethylene glycol is produced by an oil-based method utilizing ethylene oxide, however the current method is costly. [16] An alternative to the oil-based method is coal to ethylene glycol, which converts coal to syngas which is then processed to hydrogenate dimethyl oxalate to ethylene glycol. This process, even though being widely adopted in China, requires multiple steps and complicated reaction setup impeding its widespread implementation.[17,18] Thus, ECO2R would allow for a simpler and cost- 224 effective way to produce ethylene glycol. Multiple studies utilizing Cu catalysts have been able to obtain ethylene glycol, however typically with a mixture of other products and at low current densities. However, there have been reports of ECO2R to ethylene glycol with high efficiencies and selectivity with Au, Ru and Fe based electrocatalysts.[19–21] Tamura et al. demonstrated the use of an imidazolium surface assembled monolayer on Au shifted the ECO2R product from CO to ethylene glycol compared to a bare Au electrode. Moreover, the Faradaic efficiencies achieved were up to 80%, however this was only achieved after 5 h of electrolysis.[20] Moreover, Calvinho et al. demonstrated the introduction of phosphide to Fe catalysts shifts the product formation to mainly ethylene glycol, with Faradaic efficiencies of over 50%.[21] Moreover, they demonstrated that the reaction pathway differs from that of Cu based catalysts. It has been shown that the formation of C1+ products with Cu catalysts are obtained through a CO intermediate and a proton coupled electron transfer step.[4,22] However, Calvinho et al. showed that the ethylene glycol formation on Fe2P occurs through hydride transfer chemistry and is poisoned by CO intermediates, further demonstrating the significance of catalysts structure on reactivity with CO2. In recent years, atomic catalysts have been gaining a lot of attention in electrochemistry due to their similarities to homogeneous catalysts.[23–26] Their use in ECO2R has resulted in higher atom utilization, improvements in selectivity and efficiencies. Xu et al. were able to prepare Cu single atom catalysts (SACs) via an amalgamated Cu-Li method.[27] The obtained Cu-SACs were able to achieve 90% Faradaic efficiencies for CO 2 to ethanol. Moreover, Yang et al. demonstrated the use of a zeolitic imidazolate framework assisted synthesis of Cu-SACs to obtain near 50% Faradaic efficiency of methanol and 100% FE for C1 products.[28] Thus, demonstrating that differences in SAC synthesis and structure could be utilized to selectively obtain C1+ products. Nonetheless, the synthetic procedures, like most of the procedures to obtain SACs, involve 225 multiple steps, are not simple to scale up and involve carbon precursors that also are not easily attainable.[29–31] However, recently several researchers have demonstrated the use of urea and various carbon precursors to obtain graphitic carbon nitrogen (gCN) with graphene like structure containing electron lone pairs that can bind to metal ions and form SACs.[32–36] Zheng et al. recently utilized this method to obtain Ni-SACs for ECO2R to obtain near 100% selectivity for CO at high current densities.[33] Cometto et al. utilized a similar method to prepare CugCN-SACs, however they started with a melamine precursor and reduced Cu particles with NaBH4.[37] Moreover, the major product obtained from the prepared catalysts was formate in bicarbonate and hydrogen in phosphate electrolyte. This chapter discusses the preparation of CugCN-SACs from the pyrolysis of urea and commercial carbon and characterized to ensure the obtained catalysts are SACs. Moreover, the prepared catalysts are utilized in ECO2R to demonstrate the robustness and versatility of Cu-SACs. Similar to previous work, the SAC obtained demonstrate high selectivity for a single product, herein ethylene glycol is the product obtained. Moreover, the stability of the catalysts is observed over several hours in which no decay in performance is observed. Thus, demonstrating one of the most efficient catalysts for ethylene glycol production from ECO2R. 5.2. Selective Electrochemical CO2 Reduction on a Copper Single Atom Catalyst 5.2.1. Experimental Catalyst Synthesis 226 The Cu single atom catalyst (CugCN-SAC) was prepared by modifying a previously reported procedure.[33] 2 g of Vulcan carbon (XC72R) was first dispersed in 100 mL of 9 M nitric acid, stirred and kept under reflux at 90 ℃ for 3 h. The XC72R was then washed with water and centrifuged until the supernatant had a neutral pH and dried at 100 ℃ overnight. Then 500 mg of HNO3 treated XC72R was added to 100 mL of Millipore water and stirred and sonicated until a homogenous dispersion was obtained. To the dispersion, 230 mg of Cu(NO3)2 was then added slowly and kept stirring overnight in order to obtain Cu II -XC72R. The mixture was then centrifuged to obtain the powder and dried in an oven at 50 ℃ overnight. Then 0.15g Cu II -XC72R was mixed with 1.5 g of solid urea in a quartz tube and placed in a tube furnace. The solid mixture was then heated to 800 ℃ at a ramp rate of 4.4°/min under Ar and held at 800 ℃ for 1 h. The resulting powder obtained will be referred to as CugCN-SAC. A control without Cu was prepared following the same procedure with the exception of the addition of Cu(NO3)2 and the obtained sample is denoted as gCN. Furthermore, two separate Cu nanoparticle containing catalysts were also prepared utilizing ascorbic acid as a reducing agent. Firstly, CuXC72-np were prepared by adding 0.2 g of HNO3 treated XC72, 0.067 g of Cu(NO3)2 and 0.1 g of polyvinylpyrrolidone (PVP) to 40 mL of Millipore water. The mixture was sonicated and stirred for 40 min, then the mixture was kept stirring and heated to 60 ℃ and held for 1h before the slow addition of 2 g of L-ascorbic acid. The mixture was stirred at 60 ℃ for 1 h followed by centrifugation and washing the catalyst with acetone. The resulting powder was then dried overnight in an oven. Secondly, Cu nanoparticles on gCN were also prepared (CugCN-np) following a similar procedure. 50 mg of aforementioned gCN, 16.5 mg of Cu(NO3)2, and 25 mg of PVP were added to 10 mL of Millipore water. The mixture was stirred and sonicated for 40 min followed by heating to 60 ℃ and stirring for 1 h. 227 Then 0.5 g of L-ascorbic acid was slowly added, and the mixture was stirred at 60 ℃ for 1 h and then centrifuged and washed with acetone. The resulting powder was dried in an oven overnight. Catalyst Characterization Powder X-Ray Diffraction (XRD) measurements were performed on a Rigaku X-Ray diffractometer with a Cu-Kα (0.154056 nm) radiation source and a scan rate of 4°/min from a 2θ value of 10° to 80°. Scanning Electron Microscopy (SEM) images and Energy Dispersive X-ray Spectroscopy (EDS) mapping were obtained from a JEOL JSM-7001F electron microscope with an acceleration voltage of 19 and 20 keV. Transmission Electron Microscopy (TEM) images were taken on a JEOL JEM 2100F with an acceleration voltage of 200 keV. Thermogravimetric analysis (TGA) was performed on a TGA-50 Thermogravimetric analyzer (Shimadzu) under air at a ramp rate of 10 ℃/min. X-ray Photoelectron spectroscopy (XPS) data was collected on a Kratos Axis Ultra DLD using a mono Al anode with a pass energy of 160 keV for the survey scan and 20 keV for the high-resolution scans. Half-Cell Measurements A standard three electrode cell was used for half-cell experiments, where a glassy carbon rotating disk electrode (GC-RDE) with surface area 0.195 cm 2 was used as the working electrode, Pt wire and Ag/AgCl were used as the counter and reference electrode, respectively. The catalysts were drop cast onto the GC-RDE. The catalyst ink composed of 7 mg of catalyst, 900 µL of ethanol and 100 µL of Nafion solution (5 wt%). The ink was sonicated for 8 min and then 15 µL were drop cast onto the working electrode. The cell was purged with Ar for 30 min and then with CO2 for 30 min and with continuous flow of CO2 throughout the experiment. 228 The H-Cell measurements were obtained utilizing a custom-built H-type glassware with Nafion 117 as the membrane separating the anode and cathode chambers containing 0.1M NaHCO3. In the cathode compartment, the working electrode composed of teflonized carbon paper with a microporous layer (22BB, SGL Carbon) with size of 1 x 4 cm, Ag/AgCl reference electrode and the anode compartment was made of a Pt wire. The catalysts were hand brush painted onto the carbon paper with an ink consisting of 8 mg catalyst, 100 µL Nafion solution (5 wt%), and 200 µL Millipore water; the catalyst active layer was 1 x 1.5 cm. Prior to experiments, the cell was purged with Ar for 20 min while monitoring the OCV, then the gas was switched to CO2 and the solution was purged for 30 min while monitoring the OCV and each experiment was then completed. CO2 Reduction Product Analysis Each potential was held for 1 h after which 1 mL of liquid product was obtained from the cathode compartment and gas sample was also obtained. For gas product analysis, a Varian CP- 3800 equipped with a Carboxen fused silica capillary column and a thermal conductive detector was used. For liquid product analysis a Thermo Trace GC 2000 equipped with an Agilent Technologies DB-wax column and with a flame ionization detector was utilized. For all liquid samples, butanol was added as a standard. To quantify the liquid product, a calibration curve was obtained from stock ethylene glycol and butanol. The Faradaic Efficiency (FE) was calculated from equation 1 FE % = n ×F ×mol product Q × 100% (1) From equation 1, n is the number of electrons involved in the reaction, F is Faraday’s constant (96485 C/mol) and Q is the charge input for the electrolysis time. Moreover, the liquid products were confirmed with 1 H NMR. 229 5.2.2. Results and Discussion The prepared catalysts were first analyzed via TGA in order to obtain the temperature profile of the catalysts and the metal loading (Figure 5.1A). From the TGA profile the Cu loading for both the nanoparticle catalysts display roughly 10 wt%, whereas the CugCN-SAC does not have any difference in plateau weight compared to the gCN. Thus, indicating that the CugCN- SAC does not contain metallic Cu particles. Furthermore, the temperature at which the mass loss begins is much earlier for the CugCN-SAC compared to gCN, indicating the incorporation of Cu decreases the thermal stability of the carbon network. Moreover, the gCN containing catalysts are more thermally stable than the XC72 catalysts. The XRD pattern of the catalysts were obtained to observe any differences in the crystal structure of the catalyst and shown in Figure 5.1B. The gCN and CugCN-SAC XRD patterns are similar displaying a broad peak corresponding to (002) and (100) Miller Indices of graphitized carbon. Whereas the CuXC72 and CugCN-np XRD patterns display the (002) carbon Miller index as well as three peaks corresponding to metallic copper. The absence of diffraction peaks corresponding to Cu on the CugCN-SAC catalyst further indicates that any presence of Cu in the catalysts are mostly SACs and not as crystallite particle. Furthermore, the broadening of the (111) peak was utilized to obtain the crystallite size from the Debye-Scherrer equation and shown in Table 5.1. The CugCN-np and CuXC72-np catalysts displayed similar crystallite sizes, thus indicating that the use of gCN does not alter the catalyst size significantly. This could be due to the use of PVP having a more significant role in controlling the size of the nanoparticles. 230 Figure 5.1. A) TGA profile for the prepared catalyst under air atmosphere with a ramp rate of 10°/min. B) Powder XRD pattern for the prepared catalyst. Table 5.1. Crystallite and particle sizes obtained from powder XRD and TEM. Crystallite Size from XRD (nm) Particle Size from TEM (nm) CuXC72-np 16.8 ± 2 4.1 ± 1.1 CugCN-np 17.7 ± 2 8.1 ± 4.2 CugCN-SAC - - gCN - - To observe any differences in morphology SEM images were obtained and are shown in Figure 5.2. Similar to the XRD patters the gCN and CugCN-SAC catalysts display similar morphologies, porous flat surface, however the gCN does display some large aggregates. The CuXC72-np and CugCN-np catalysts, however, show different structures. CugCN-np also has a similar porous structure observed in the other gCN catalysts, however there are also several large particles present. The CuXC72-np catalysts also has large particles present and a different morphology, where the surface looks more flat and less porous. To observe the elemental distribution and the profile of the large particles observed, EDS mapping was performed for all the 231 images. From the EDS, the gCN only presented carbon, oxygen, and nitrogen present and the large agglomerates observed were agglomerates of the gCN. Moreover, the CugCN-SAC catalyst displayed presence of Cu, carbon, oxygen and nitrogen, and the Cu are seen dispersed throughout the gCN structure. The CuXC72-np and the CugCN-np show Cu, carbon, oxygen and nitrogen. However, the Cu observed for these two catalysts are much larger aggregates confirming the formation of Cu particles, whereas in the CugCN-SAC the Cu the particles are too small to be observed in SEM, indicating the potential formation of SACs or nanoparticles. The weight percent of the elements were also obtained and are shown in Table 5.2. The amount of copper observed in the CugCN-np and CuXC72-np agrees with the weight percentage obtained from TGA. Moreover, there is a 2 wt% of copper in the CugCN-SAC, which likely indicates the amount of Cu are SAC and not nanoparticles. 232 233 Figure 5.2. SEM and EDS mapping of A) CuXC72-np where carbon corresponds to red, oxygen corresponds to green and copper corresponds to blue; B) CugCN-np where carbon corresponds to red, oxygen corresponds to blue, nitrogen corresponds to green and copper corresponds to yellow; C) CugCN-SAC where carbon corresponds to red, nitrogen corresponds to green and copper corresponds to blue; and D) gCN where carbon corresponds to red, copper corresponds to yellow and nitrogen corresponds to green. To obtain further insight into the differences in the morphology of the catalysts, TEM micrographs were obtained and are shown in Figure 5.3. The gCN and CugCN-SAC again show similar morphologies, where all that is observable are the gCN particles with size near 30 nm. The CuXC72-np display a different morphology from the gCN, where the structure is a more sheet like surface, corroborating with the morphology observed in the SEM images. The CugCN-np presents a similar structure to the other gCN containing catalysts with large particles roughly 30 nm in size and smaller particles near 8 nm. Moreover, the CuXC72-np also present small nanoparticles with the average size shown in Table 5.1. The average particle size observed in the TEM are smaller than the crystallite size observed in XRD. This difference arises from the particle size observed in TEM only being a small representation of the particle size, this is confirmed by the SEM/EDS showing the presence of much larger Cu particles. Moreover, the particle distribution between the two vary significantly where the Cu is more dispersed throughout the entire carbon support for the CugCN-np and more dispersed in a smaller region for the CuXC72-np. Furthermore, no Cu nanoparticles were observed on the CugCN-SAC further indicating the formation of SACs. 234 Figure 5.3. TEM micrographs of A) CuXC72-np; B) CugCN-np; C) CugCN-SAC; and D) gCN. XPS analysis was completed to gain further insight into the elemental composition of the catalysts surface and the results shown in Figure 5.4. The survey scan corroborates with the elemental composition observed from EDS; where gCN only contains carbon, nitrogen and oxygen and CugCN-SAC and CugCN-np contain copper, carbon, nitrogen, and oxygen. The weight percentages of the elements are provided in Table 5.2 and align with that observed in EDS. Apart 235 from copper in the CugCN-np. This likely arises from the difference in penetration depth of EDS and XPS, where XPS is only a few nanometers through the surface. Thus, there is a difference between the bulk and surface of the CugCN-np where 2 wt% of copper is on the surface. However, the CugCN-SACs have similar copper content on the surface and the bulk, indicating a more uniform distribution of the copper. High resolution scans were also obtained and confirm the presence of copper in the copper containing catalysts, where the two Cu 2p peaks are observed. Interestingly, the copper high resolution scans in CugCN-np and CugCN-SAC differ; CugCN-np contains two peaks near 945 and 965 eV indicating the presence of oxidized copper. Moreover, the Cu 2p3/2 and Cu 2p1/2 peaks for the CugCN-SAC are shifted to lower binding energies, which could arise from the charge transfer of the gCN support and the Cu. From the high-resolution spectra of the N 1s the catalysts contain two major types of nitrogen species, pyrrolic and pyridinic. The CugCN-np contains a significant amount of pyrrolic nitrogen, and this is from the PVP. The gCN and CugCN-SAC contain more pyridinic nitrogen than pyrrolic; moreover, there is a slight increase in the binding energy for CugCN-SAC in the pyridinic region and this arises from the Cu- N bonds present. Furthermore, the presence of copper in both EDS and XPS for the CugCN-SACs, the shift in N 1s and absence of nanoparticles in TEM indicate the formation of SACs. 236 Figure 5.4. XPS A) survey scan; and high-resolution scans for B) Cu 2p; C) N 1s; and D) C 1s for the prepared catalyst. Table 5.2. Element weight percentage obtained from EDS and XPS. EDS XPS C (wt%) N (wt%) O (wt%) Cu (wt%) C (wt%) N (wt%) O (wt%) Cu (wt%) CuXC72 -np 75.59 2.14 11.56 10.71 - - - - CugCN- np 75.66 2.28 10 12.06 81.43 4.10 11.75 2.73 CugCN- SAC 80.09 12.25 5.3 2.36 79.64 2.88 15.99 1.49 gCN 78.55 13.68 7.76 - 76.07 3.12 20.82 - The electrochemical performance was first studied in a single chamber half-cell, the electrochemical double layer capacitance was obtained by completing cyclic voltammetry (CV) experiments at various scan rates in a non-Faradaic potential region and plotting the peak current against the scan rates (Figure 5.5). The capacitance of the catalysts was obtained by the slopes of the current vs scan rate plots, the values are shown in Figure 5.5E-H. There is a decay in capacitance with the introduction of the copper in both the nanoparticles and the SACs. However, 237 the decay is much more prominent in the nanoparticles. Moreover, the use of XC72 has a much lower capacitance value, indicating the advantage of utilizing the gCN. 238 Figure 5.5. CV scans at various scan rates in 0.1M NaHCO3 under a flow of Ar for A) CuXC72- np; B) CugCN-np; C) CugCN-SAC; and D) gCN. The corresponding peak current at each scan rate to obtain the capacitance of the catalysts for E) CuXC72-np; F) CugCN-np; G) CugCN-SAC; and H) gCN. LSVs of the catalysts under Ar and CO2 were obtained to quickly assess catalytic activity towards ECO2R, as shown in Figure 5.6. From the LSVs the catalysts containing gCN demonstrate a lower onset potential, which is denoted as the potential when the current reaches 1 mA/cm 2 , under CO2 compared to Ar. The decreased onset potential could arise from activation and reduction of CO2; therefore, these catalysts could be utilized for ECO2R. 239 Figure 5.6. LSV scans under Ar and CO2 flow at scan rates of 20 mV/s in 0.1M NaHCO3 for A) CuXC72-np; B) CugCN-np; C) CugCN-SAC; and D) gCN. The catalysts were then applied to a carbon paper electrode and placed in a custom H-cell to detect electrolysis products. Constant potential electrolysis experiments were completed in a potential window of -0.5 V to -1.8 V vs. Ag/AgCl, where each potential was held for 1 h. The currents obtained during the electrolysis were plotted against the applied potential and shown in Figure 5.7 for each of the catalysts. The obtained products at each potential are shown in Table 5.3, the catalysts all provide different products. The gCN produces a mixture of CO and H2, as has been previously shown, moreover the addition of Ni improves the FE and product yield of CO.[33] CuXC72-np interestingly only produced H2, even though Cu based catalysts are known to produce several C1+ products. Furthermore, the CugCN-SAC and CugCN-np produced a mixture of ethylene glycol and H2, where the CugCN-SAC was more selective in ethylene glycol produced. The products obtained differ from those of previously reported CugCN-SACs. This likely arises from the differences in synthesizing the catalysts, resulting in different catalyst structure and exposed active sites.[27,28,37] 240 Figure 5.7. Current values obtained from constant potential electrolysis experiments in 0.1M NaHCO3 under flow of CO2 for A) CuXC72-np; B) CugCN-np; C) CugCN-SAC; and D) gCN. Table 5.3. Products obtained for the prepared catalysts. Products Potential (V vs. Ag/AgCl) CuXC72-np CugCN-np CugCN-SAC gCN -0.5 - - - - -1.0 H2 H2/Ethylene glycol Ethylene glycol H2 -1.2 H2 H2/Ethylene glycol Ethylene glycol H2 -1.4 H2 H2 H2/Ethylene glycol H2/CO 241 -1.6 H2 H2 H2/Ethylene glycol H2/CO -1.8 H2 H2 H2/Ethylene glycol H2/CO CugCN-np only produced minor amounts of ethylene glycol at -1.0 and -1.2 V vs. Ag/AgCl. The FE for the ethylene glycol produced for the two catalysts are shown in Figure 5.8. The CugCN-np reached a maximum FE for ethylene glycol of approximately 5% and then at higher potentials H2 was observed as the only product. The FE for ethylene glycol for the CugCN-SAC was near 80% and is amongst the highest reported values.[20] Moreover, no other liquid products were observed and at -1.0 and -1.2 V vs. Ag/AgCl there was no H2 detected via GC. There was a decrease in the FE as the potential increased, however at potentials more positive than -1.4 V vs. Ag/AgCl the FE remained above 40% and ethylene glycol remained the only detectable liquid product across all potentials. Thus, demonstrating the advantage of utilizing SACs to improve product selectivity. The stability of the CugCN-SAC was then studied over the course of several hours at a constant potential of -1.0 V vs. Ag/AgCl, shown in Figure 5.8D. 242 Figure 5.8. A) Faradaic efficiencies of ethylene glycol for CugCN catalyst; B) liquid sample and C) gas sample GC spectra for CugCN-SACs at various potentials; and D) current and FE of CugCN-SAC at -1.0 V vs. Ag/AgCl over 10 h. Previous reports demonstrated an increase in the FE of the ethylene glycol over the span of 8 h, however as seen in the Figure 5.8D, the FE of the CugCN-SAC remains relatively stable throughout 10 h of electrolysis.[20] Therefore, the catalyst is stable under the observed conditions and can efficiently produce ethylene glycol. Insight into the mechanism can also be obtained, even though there is uncertainty in the mechanism for C2+ products on Cu catalysts, majority of work demonstrates a pathway involving CO intermediate.[6,22] This process is further evidenced herein based on the production of CO on the bare gCN and absence of CO in both CugCN catalysts. Thus, indicating the CO that would be formed by the gCN further reacts with the Cu surface and forms the desired products in the case of CugCN-SAC. While the H-cell work is impressive and an indicator of the advantages of CugCN-SAC, further work is needed to implement the catalyst in an electrolyzer cell to fully demonstrate the efficacy of the catalyst for production of ethylene glycol. 243 5.3. Conclusion A simple and scalable procedure to obtain Cu-SACs is demonstrated. The obtained catalysts were characterized via XPS, SEM, TEM and XRD to confirm the formation of SACs. Moreover, the SAC’s electrocatalytic activity towards CO2 reduction was compared to Cu nanoparticle catalysts. The CugCN-SAC were able to reduce CO2 with high selectivity to ethylene glycol and produce ethylene glycol for 10 h at a low potential maintaining FE above 80%. Thus, demonstrating one of the most efficient and selective catalysts for ethylene glycol production. 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Abstract (if available)
Abstract
This dissertation comprises of five chapters discussing developments in polymeric materials to catalysts for use in fuel cells and electrochemical CO2 reduction. In Chapter 1, an introduction to fuel cells and electrochemical CO2 reduction is provided.
Chapter 2 discusses further studies of partially fluorinated carbon (CFx) supported platinum nanoparticles for hydrogen oxidation reaction (HOR) and oxygen reduction reaction (ORR). The Pt/CFx stability is compared to Pt supported on Vulcan XC72R for both HOR and ORR. The Pt/CFx displayed improved stability due to the improved interaction between the Pt and CFx support. Moreover, the Pt/CFx displayed improved electrocatalytic activity for HOR compared to Pt/XC72R. The improvement is due to improved mass transport and water management. Furthermore, the Pt/CFx was used as both he anode and cathode electrocatalyst in a H2-proton exchange membrane fuel cell (PEMFC) and displayed a 20% improvement in terms of peak power compared to a H2-PEMFC with Pt/XC72R. Furthermore, the CFx support was also utilized to prepared Pd/CFx catalysts for alkaline direct liquid fuel cells. The Pd/CFx catalysts displayed improvements over Pd/C arising from increased electrochemically active surface area and improved mass transport. The Pd/CFx catalyst also displayed improved performance in alkaline direct liquid fuel cells for various liquid fuels (glycerol, methanol, ethylene glycol and formate). The use of CFx as a catalyst support is robust and aids in improving both acidic and alkaline fuel cells, aiding mass transport and stability.
In chapter 3, developments in formate fuel cells are discussed from catalyst design to cell structure. First, we have explored the activity of palladium supported on reduced graphene oxide (Pd/rGO) towards the formate oxidation reaction (FOR) in an alkaline medium. The reduction of GO to rGO and synthesis of Pd nanoparticles were confirmed using X-Ray Diffraction (XRD), Raman and X-Ray Photoelectron Spectroscopy (XPS). The surface morphology was evaluated by Scanning Electron Microscopy (SEM) and Transmission Electron Microscopy (TEM). Half-cell studies demonstrated superior electrocatalytic activity and stability towards formate electrooxidation for Pd/rGO than commercial Pd/C catalysts. A low metal loading DFFC, fabricated with a Pd/rGO anode catalyst displayed a 15% increase in maximum power density at 60 ℃ compared to the commercial Pd/C catalyst. Improvements were further made by combining IrO2 with Pd and MWCNTs. The addition of IrO2 and MWCNTs both improved the ECSA of the catalyst and improved kinetics of FOR. Moreover, the improvements were demonstrated in a DFFC providing one of the highest reported peak power densities of a DFFC, 299 mW/cm2, while utilizing less than 1 mgpd/cm2. Finally, the role of cations on FOR was studied and found to play a significant role in current observed. HCOONa displayed the highest performance due to improved mass transport and more facile removal of reactive intermediates from the catalyst surface. Moreover, the formate salts were utilized with a cation exchange membrane to coproduce electricity and alkali hydroxide. The first section of this chapter is reprinted with permission from ACS Appl. Energy Mater. 2019, 2, 10, 7104–7111. Copyright 2019 American Chemical Society.
Chapter 4 discusses improvements in alkaline direct methanol fuel cells (ADMFCs), specifically changes in the ionomer and membrane materials and catalysts. The first portion of the chapter reports the activity of the methanol oxidation reaction (MOR) in half-cell experiments with varying ionomers and the use of a poly(terphenylene) (TPN) membrane in an alkaline direct methanol fuel cell (ADMFC). The results demonstrate that small changes in the cation structures of the ionomer have a significant role in MOR on the PtRu/C catalyst. Moreover, with the use of a TPN membrane and the prepared anode containing ionomers, high power densities are achieved with less than 1 mgPtRu/cm2 in the catalyst layer. Next, nickel-based catalysts were prepared and studied for methanol oxidation. Firstly, GaNi based catalysts were prepared at various ratios and it was found that a 3 to 1 Ga to Ni provided the highest performance in half-cell studies. The GaNi catalysts were also demonstrated to be effective in ADMFCs, displaying a similar trend in catalytic activity as the half-cell experiments. Lastly, a NiCeO2/C catalyst was also prepared and displayed improved catalytic activity compared to Ni/C due to increased ECSA, and kinetics. Moreover, both GaNi/C and NiCeO2/C catalysts were utilized in precious metal free ADMFCs by using MnO2/C and FegCN catalysts for the cathode electrode. The first section of this chapter is reprinted with permission from ACS Appl. Energy Mater. 2021, 4, 6, 5858–5867. Copyright 2021 American Chemical Society
In chapter 5, copper single atom catalysts (SACs) were prepared with a facile and scalable procedure and utilized for electrochemical CO2 reduction (ECO2R). The SAC formation was confirmed with XPS, SEM, TEM and XRD analysis. The CugCN-SACs were then utilized in an H-cell for ECO2R, the catalyst was able to reduce CO2 to ethylene glycol with high selectivity and maintain Faradaic Efficiencies near 80% for production of ethylene glycol for 10 h at a constant potential of -1.0 V vs Ag/AgCl. Thus, demonstrating one of the most effective catalysts for ECO2R to ethylene glycol.
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Galvan, Vicente
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Core Title
Design and modification of electrocatalysts for use in fuel cells and CO₂ reduction
School
College of Letters, Arts and Sciences
Degree
Doctor of Philosophy
Degree Program
Chemistry
Degree Conferral Date
2022-08
Publication Date
07/09/2022
Defense Date
06/16/2022
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University of Southern California
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University of Southern California. Libraries
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alkaline fuel cell,carbon dioxide,carbon nanotubes,catalysts,ceria,CFx,CO₂ reduction,copper,electrochemistry,ethylene glycol,formate fuel cell,formate oxidation,fuel cells,gallium,graphene,graphene oxide,hydrogen oxidation,iridium oxide,methanol fuel cell,methanol oxidation,nickel,OAI-PMH Harvest,oxygen reduction,palladium,platinum,reduced graphene oxide,single atom catalyst,Vulcan carbon,XC-72
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Prakash, Surya (
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), Ravichandran , Jayakanth (
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galvanv@usc.edu,vicentegalvan5@gmail.com
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Tags
alkaline fuel cell
carbon dioxide
carbon nanotubes
catalysts
ceria
CFx
CO₂ reduction
copper
electrochemistry
ethylene glycol
formate fuel cell
formate oxidation
fuel cells
gallium
graphene
graphene oxide
hydrogen oxidation
iridium oxide
methanol fuel cell
methanol oxidation
nickel
oxygen reduction
palladium
platinum
reduced graphene oxide
single atom catalyst
Vulcan carbon
XC-72