Applied Science, Faculty of
Chemical and Biological Engineering, Department of
DSpace
UBCV
Moore, Colin E.
Doctor of Philosophy - PhD
University of British Columbia
Catalyst durability in renewable energy systems (including CO₂ electroreduction and Hydrogen fuel cells) is vital to the overall lifetime of the systems. To protect hydrogen fuel cells during cell voltage reversal, oxygen evolution reaction (OER) catalysts (typically RuO₂ or IrO₂) are added to the anode of proton exchange membrane fuel cells (PEMFCs). In this thesis, the durability of these OER catalysts was investigated by measuring the mass changes of heat-treated IrOx powders (350, 450 and 550 °C) drop cast on an electrochemical quartz crystal microbalance (EQCM). IrOx catalysts showed difference in frequency response due to the differential uptake of water and formation of oxyhydroxide species during cyclic voltammetry (CV) experiments (1.2, 1.4 and 1.6 to 0.05 V vs. RHE) confirmed by XPS. Platinum on carbon catalyst (Pt/C) suffered from carbon corrosion during potential holds at 1.8 and 2.0 V. The addition of IrOx powders to Pt/C protected the layer against carbon corrosion creating a simulated reversal tolerant anode. Additionally, OER catalysts were ranked by OER activity and dissolved Ir³⁺ in the electrolyte after an ex situ accelerated stress test (AST) in a representative PEMFC anode environment. The ex situ results are compared with reversal times obtained in a single cell PEMFC subjected to anode accelerated stress test (AAST). Generally, catalysts with higher OER mass activity and Ir³⁺ dissolution had longer reversal times. Heat treatment of unsupported IrOx increased OER durability in the fuel cell anode environment. Catalyst instability strongly affects the viability of electroreduction of CO₂ (ERC) systems as well. To improve the durability and efficiency of ERC catalysts, five Bimetallic Sn-Pb catalyst compositions were electrodeposited (from fluoroborate or oxide media) on two carbon supports. Sn majority catalysts with 15 to 35% Pb generated stable faradaic efficiencies (FE) up to 95% during constant potential electrolysis whereas pure Sn experienced an extensive (up to 30%) decrease in FE. After electrolysis, XRD analysis showed a SnO₂ phase present on 35% Pb catalysts but not on pure Sn catalysts. It is proposed that the presence of Pb in Sn majority catalysts stabilized SnO₂ and enhanced faradaic efficiency durability during ERC.
ELECTROCATALYSTS FOR VOLTAGE REVERSAL TOLERANT ANODES IN PROTON EXCHANGE MEMBRANE FUEL CELLS AND CO2 REDUCTION TO FORMATE by Colin E. Moore B.Sc., The University of North Carolina at Chapel Hill, 2009 M.Sc., Eötvös Loránd University, Budapest, Hungary, 2014 A THESIS SUBMITTED IN PARTIAL FULFILLMENT OF THE REQUIREMENTS FOR THE DEGREE OF DOCTOR OF PHILOSOPHY in THE FACULTY OF GRADUATE AND POSTDOCTORAL STUDIES (Chemical and Biological Engineering) THE UNIVERSITY OF BRITISH COLUMBIA (Vancouver) November 2019 © Colin E. Moore, 2019 ii The following individuals certify that they have read, and recommend to the Faculty of Graduate and Postdoctoral Studies for acceptance, the dissertation entitled: ELECTROCATALYSTS FOR VOLTAGE REVERSAL TOLERANT ANODES IN PROTON EXCHANGE MEMBRANE FUEL CELLS AND CO2 REDUCTION TO FORMATE submitted by Colin E. Moore in partial fulfillment of the requirements for the degree of Doctor of Philosophy in Chemical and Biological Engineering Examining Committee: Elöd Gyenge Supervisor Natalia Kremliakova Supervisory Committee Member David Wilkinson Supervisory Committee Member Saavas Hatzikiriakos University Examiner Michael Wolf University Examiner Additional Supervisory Committee Members: Kevin Smith Supervisory Committee Member Supervisory Committee Member iii Abstract Catalyst durability in renewable energy systems (including CO2 electroreduction and Hydrogen fuel cells) is vital to the overall lifetime of the systems. To protect hydrogen fuel cells during cell voltage reversal, oxygen evolution reaction (OER) catalysts (typically RuO2 or IrO2) are added to the anode of proton exchange membrane fuel cells (PEMFCs). In this thesis, the durability of these OER catalysts was investigated by measuring the mass changes of heat-treated IrOx powders (350, 450 and 550 °C) drop cast on an electrochemical quartz crystal microbalance (EQCM). IrOx catalysts showed difference in frequency response due to the differential uptake of water and formation of oxyhydroxide species during cyclic voltammetry (CV) experiments (1.2, 1.4 and 1.6 to 0.05 V vs. RHE) confirmed by XPS. Platinum on carbon catalyst (Pt/C) suffered from carbon corrosion during potential holds at 1.8 and 2.0 V. The addition of IrOx powders to Pt/C protected the layer against carbon corrosion creating a simulated reversal tolerant anode. Additionally, OER catalysts were ranked by OER activity and dissolved Ir3+ in the electrolyte after an ex situ accelerated stress test (AST) in a representative PEMFC anode environment. The ex situ results are compared with reversal times obtained in a single cell PEMFC subjected to anode accelerated stress test (AAST). Generally, catalysts with higher OER mass activity and Ir3+ dissolution had longer reversal times. Heat treatment of unsupported IrOx increased OER durability in the fuel cell anode environment. Catalyst instability strongly affects the viability of electroreduction of CO2 (ERC) systems as well. To improve the durability and efficiency of ERC catalysts, five Bimetallic Sn-Pb catalyst compositions were electrodeposited (from fluoroborate or oxide media) on two carbon supports. Sn majority catalysts with 15 to 35% Pb generated stable faradaic efficiencies (FE) up to 95% during constant potential electrolysis whereas pure Sn experienced an extensive (up to 30%) iv decrease in FE. After electrolysis, XRD analysis showed a SnO2 phase present on 35% Pb catalysts but not on pure Sn catalysts. It is proposed that the presence of Pb in Sn majority catalysts stabilized SnO2 and enhanced faradaic efficiency durability during ERC. v Lay Summary Long lasting catalysts are crucial to implement strategies to mitigate climate change. A catalyst is a material that improves the rate of a reaction. Increasing catalyst lifetime is essential for these technologies to compete against their fossil fuel derived counterparts. In this work, efficient techniques are developed to investigate the durability of hydrogen polymer electrolyte fuel cell catalysts essential for protecting against damage during startup and shutdown procedures. It is described how heat treatments and catalyst supports can improve catalyst durability and longevity in harsh environments. To improve catalysts that transform carbon dioxide to value added products, it was determined that combining tin and lead into a single catalyst has better performance than either tin or lead alone. This is because lead stabilizes the catalytically active form of tin. vi Preface The work presented in this thesis including the literature review, industrial collaboration research proposal, experimental design and execution, data analysis and interpretation and thesis preparation was completed by Colin E. Moore under the direct supervision of Professor Előd L. Gyenge at the Department of Chemical and Biological Engineering, the University of British Columbia. Thomas Rowe, assisted in electrode preparation and characterization in Chapter 6 under the direct supervision of the author. Publications that are partly included in this work are listed:  A version of Chapter 5 has been published (Moore CE, Eastcott J, Cimenti M, Kremliakova N, Gyenge EL (2019) Novel methodology for ex situ characterization of iridium oxide catalysts in voltage reversal tolerant proton exchange membrane fuel cell anodes Journal of Power Sources 417:53–60.) The author performed all ex situ experiments and prepared the publication and its content. Dr. Natalia Kremliakova co-authored the publication and is responsible for significant hypothesis formulation and experimental guidance. Dr. Jennie Eastcott provided some ex situ experimental procedures. Dr. Max Cimenti performed in situ fuel cell experiments during an industrial collaboration with Automotive Fuel Cell Cooperation (formerly of Burnaby, BC).  A version of Chapter 4 has been published (Moore CE, Afsahi F, Young AP, Gyenge EL (2019) Vibrating Powders: Electrochemical Quartz Crystal Microbalance Study of IrO2 and Pt/C Catalyst Layers for Voltage Reversal Tolerant Anodes in Fuel Cells Journal of Physical Chemistry C 123: 23361-23373) The author is responsible for the preparation of this publication and its content. This work was carried out as a part of an industrial vii collaboration with Ballard Power Systems (Burnaby, BC) who provided laboratory equipment and fuel cell expertise.  A version of Chapter 6 has been published (Moore CE, Gyenge EL (2017) Tuning the Composition of Electrodeposited Bimetallic Tin-Lead Catalysts for Enhanced Activity and Durability in Carbon Dioxide Electroreduction to Formate ChemSusChem 10:3512–3519.) The author is responsible for the preparation of this publication and its content. This work was carried out as a part of an industrial collaboration with Mantra Energy Alternatives (formerly of Vancouver, BC). The work has also been presented at the following conferences by the author:  Moore, C. E.; Afsahi, F.; Young, A. P.; Knights, S.; Gyenge, E. L. (2019) Vibrating Powders: Eqcm Measurement of Potential Dependent IrO2 Mass Changes for Voltage Reversal Tolerant PEMFC Anodes 235th ECS Meeting 1629  Moore CE, Gyenge EL (2018) Tuning the Composition of Bimetallic Electrodeposited Sn-Pb Catalysts for Enhanced Activity and Durability in CO2 Electroreduction to Formate 233rd ECS Meeting 1829  Moore CE, Eastcott J, Cimenti M, Kremliakova N, Gyenge EL (2018) Novel Methodology for Ex-Situ Characterization of Catalysts in Reversal Tolerant PEM-FCs 233rd ECS Meeting 1771 Chapters 1-3 and 7 include portions of previously published material from the references listed above in the Journal of Power Sources, the Journal of Physical Chemistry C and ChemSusChem. viii Table of Contents Abstract ......................................................................................................................................... iii Lay Summary .................................................................................................................................v Preface ........................................................................................................................................... vi Table of Contents ....................................................................................................................... viii List of Tables .............................................................................................................................. xiv List of Figures ............................................................................................................................. xvi List of Symbols ........................................................................................................................ xxvii List of Abbreviations ............................................................................................................... xxix Acknowledgements ................................................................................................................ xxxiii Dedication ............................................................................................................................... xxxiv Chapter 1: Introduction ................................................................................................................1 1.1 Motivation ....................................................................................................................... 1 1.2 Catalyst durability in fuel cells and carbon electroreduction.......................................... 1 1.3 Proton exchange membrane fuel cell overview .............................................................. 2 1.4 Department of Energy targets ......................................................................................... 5 1.5 PEMFC failure modes..................................................................................................... 6 1.6 Cell reversal mitigation strategies ................................................................................... 8 1.7 Choosing effective oxygen evolution catalysts for reversal tolerant anodes ................ 10 1.8 Electrochemical quartz crystal microbalance as an ex situ oxygen evolution reaction catalyst durability investigation tool ......................................................................................... 11 1.9 Electroreduction of CO2 proposal and addition to thesis .............................................. 12 1.10 Carbon dioxide conversion and storage technologies ................................................... 12 ix 1.11 Durability requirements for electroreduction of CO2 ................................................... 14 1.12 Industrial application requirements for electroreduction of CO2 .................................. 15 Chapter 2: Literature review ......................................................................................................17 2.1 Voltage reversal in PEMFCs ........................................................................................ 17 2.1.1 Reversal tolerant anodes and oxygen evolution reaction catalyst development ....... 17 2.1.2 Oxygen evolution reaction catalyst activity and stability ......................................... 17 2.1.3 Oxygen evolution reaction kinetics .......................................................................... 22 2.1.4 Oxygen evolution reaction catalyst crystallinity and chemical form ........................ 24 2.1.5 Nanoparticle catalysts ............................................................................................... 25 2.1.6 Oxide supports for IrO2 catalysts .............................................................................. 26 2.1.7 Mixed IrO2 RuO2 catalysts........................................................................................ 27 2.1.8 Reversal tolerant anode related patents ..................................................................... 28 2.1.9 EQCM measurement of rigid IrO2 layers ................................................................. 28 2.1.10 EQCM measurement of catalyst powders ............................................................. 29 2.2 CO2 electroreduction ..................................................................................................... 29 2.2.1 CO2 electroreduction thermodynamics ..................................................................... 29 2.2.2 Possible chemical products ....................................................................................... 31 2.2.3 Faradaic efficiency of metallic catalysts ................................................................... 32 2.2.4 Bi and tri-metallic catalysts ...................................................................................... 36 2.2.5 Support effects .......................................................................................................... 38 2.3 Knowledge gap ............................................................................................................. 38 2.4 Thesis objectives and layout ......................................................................................... 39 2.4.1 IrO2 EQCM objectives of the thesis.......................................................................... 39 x 2.4.2 Oxygen evolution reaction catalyst objectives of the thesis ..................................... 40 2.4.3 Electroreduction of CO2 goals of the thesis .............................................................. 40 2.4.4 Thesis layout ............................................................................................................. 40 Chapter 3: Methods and materials .............................................................................................42 3.1 Surface and material analysis........................................................................................ 42 3.1.1 Scanning electron microscopy .................................................................................. 42 3.1.2 Energy dispersive X-ray ........................................................................................... 43 3.1.3 X-ray diffraction ....................................................................................................... 43 3.1.4 Brunauer-Emmett-Teller surface area analysis ......................................................... 44 3.1.5 Scanning transmission electron microscopy ............................................................. 45 3.1.6 X-ray photoelectron spectroscopy ............................................................................ 45 3.1.7 Inductively couple plasma mass spectrometry ......................................................... 46 3.2 OER experimental goals and implementation .............................................................. 47 3.3 Iridium oxide heat treatment and surface area analysis ................................................ 48 3.3.1 Ink preparation .......................................................................................................... 48 3.3.2 Quartz crystal microbalance measurements.............................................................. 49 3.3.3 Characterization and accelerated potential stress testing .......................................... 50 3.3.4 Cu deposition EQCM calibration.............................................................................. 50 3.4 Oxygen evolution ex situ characterization methods ..................................................... 55 3.4.1 Types of oxygen evolution reaction catalysts ........................................................... 55 3.4.2 IrOx heat treatment .................................................................................................... 56 3.4.3 Electrode preparation ................................................................................................ 57 3.4.4 Electrochemical characterization .............................................................................. 58 xi 3.4.5 Ex situ stress testing and ICP-MS ............................................................................. 59 3.4.6 Fuel cell anode accelerated stress test ....................................................................... 59 3.5 Electroreduction of CO2 methodology.......................................................................... 61 3.5.1 Pb-Sn electrodeposition on two types of fibrous carbon substrates ......................... 61 3.5.2 Electrochemical characterization .............................................................................. 64 3.5.3 Cyclic voltammetry and H-cell experiments ............................................................ 65 3.5.4 Determination of faradaic efficiency ........................................................................ 66 Chapter 4: Electrochemical quartz crystal microbalance measurement of IrOx mass changes........................................................................................................................................................68 4.1 Reversal tolerant anode OER PEMFC objectives ........................................................ 68 4.2 Introduction ................................................................................................................... 69 4.3 Improving on ex situ AST methodology ....................................................................... 70 4.4 Frequency to mass relationship ..................................................................................... 71 4.5 Heat-treated IrOx XRD, XPS and SEM Characterization ............................................. 72 4.6 Cyclic voltammetry and relative capacitance (RC) ...................................................... 76 4.7 EQCM response of heat-treated vs. untreated IrO2 ....................................................... 79 4.8 EQCM response to potentiostatic holds ........................................................................ 85 4.9 Pt/C with IrO2 potential cycling and constant potential polarization ........................... 86 4.10 Conclusions ................................................................................................................... 91 Chapter 5: Ex situ testing of reversal tolerant anode catalysts ...............................................93 5.1 Introduction ................................................................................................................... 93 5.2 In situ vs. ex situ protocols for RTA investigations ...................................................... 93 5.3 Temperature controlled accelerated stress testing ........................................................ 94 xii 5.4 Ir dissolution after ex situ AST: unsupported and supported IrOx comparison ............ 95 5.5 Ex-situ oxygen evolution reaction activity ................................................................... 98 5.6 Fuel cell reversal performance vs. ex situ Ir dissolution and OER activity ................ 101 5.7 Metal cation migration during ex situ accelerated stress test ...................................... 103 5.8 In situ Ir migration across the membrane electrode assembly during fuel cell AAST 104 5.9 Ex situ performance of the IrOx/TiO2 OER catalyst combined with hydrogen oxidation reaction Pt catalyst .................................................................................................................. 106 5.10 Conclusions ................................................................................................................. 107 Chapter 6: Bimetallic Sn-Pb catalysts for electroreduction of CO2 to formate ...................110 6.1 Introduction ................................................................................................................. 110 6.2 Pb/Sn catalyst composition ......................................................................................... 110 6.3 Electrochemical properties .......................................................................................... 112 6.4 Formate faradaic efficiency and total formate ............................................................ 116 6.5 Catalyst changes after formate production ................................................................. 121 6.6 Low Pb content efficiency stabilization ...................................................................... 123 6.7 Conclusions ................................................................................................................. 126 Chapter 7: Conclusions and recommendations for future work ...........................................128 7.1 Electroreduction of CO2 and PEMFC to combat climate change ............................... 128 7.2 EQCM measurement of IrOx and Pt/C conclusions .................................................... 129 7.3 Ex situ accelerated stress test for reversal tolerant anodes conclusions...................... 130 7.4 Bimetallic Sn-Pb catalysts for electroreduction of CO2 conclusions.......................... 131 7.5 Contributions to knowledge ........................................................................................ 131 7.6 Recommendations for future work ............................................................................. 132 xiii 7.6.1 EQCM oxygen evolution reaction catalyst future work ......................................... 132 7.6.2 PEMFC oxygen evolution reaction catalyst future work ........................................ 134 7.6.3 Electroreduction of CO2 future work ...................................................................... 135 7.6.4 Gas diffusion electrodes .......................................................................................... 137 7.6.5 Flow cell.................................................................................................................. 138 7.6.6 Nanostructures ........................................................................................................ 138 7.6.7 Bi and trimetallic cathodes ...................................................................................... 139 7.6.8 Support materials .................................................................................................... 139 Bibliography ...............................................................................................................................141 Appendices ..................................................................................................................................158 Appendix A ............................................................................................................................. 158 A.1 Electrochemical details ........................................................................................... 158 Appendix B ............................................................................................................................. 159 B.1 IrOx EQCM supplemental information ................................................................... 159 Appendix C ............................................................................................................................. 170 C.1 IrOx ex situ characterization .................................................................................... 170 C.2 Additional metal supported IrOx OER catalysts ..................................................... 175 Appendix D ............................................................................................................................. 180 D.1 ERC supplemental information............................................................................... 180 xiv List of Tables Table 2.1. Faradaic efficiency for CO2 electroreduction on various metallic electrodes from Hori et al. [100] Current density and potential vs. the SHE included for each metallic electrode catalyst. Methane (CH4), ethane (C2H4), ethanol (EtOH), propanol (PrOH), carbon monoxide (CO), Formate (HCOO-) and hydrogen gas (H2) faradaic efficiency included. Total faradaic efficiency does not add up to 100% due to experimental error in quantification. ......................................... 35 Table 3.1. Sensitivity of analytical techniques to determine Ir3+ concentration (ng mL-1) in 1 M H2SO4. [36] .................................................................................................................................... 47 Table 3.2. BET Surface Area of IrOx Untreated and Heat-treated at 350, 450 and 550 °C. ........ 48 Table 3.3 BET surface areas and IrOx contents of three unsupported and one TiO2 supported catalyst in Chapter 5...................................................................................................................... 56 Table 3.4. Electrodeposition bath compositions using Sn and Pb tetrafluoroborate salts. ........... 63 Table 3.5. Electrodeposition bath compositions using SnO and PbO. ......................................... 64 Table 5.1 Elemental makeup on the cathode determined by EDX. Relocated metal area and spot analysis were averaged from the SEM images in Figure 5.6. ..................................................... 106 Table 6.1. EDX analysis of Pb and Sn electrodeposited carbon cloth and carbon paper using the fluoroborate bath. Other electrodeposition conditions given in Section 3.5.1. ........................... 111 Table 6.2. EDX analysis of Pb and Sn electrodeposited carbon cloth and carbon paper using the oxide bath. Other experimental conditions given in Section 3.5.1. ............................................ 112 Table 6.3. Total formate production rate during a 2 hour electrolysis at -2.0 V vs. Ag/AgCl in 0.5 M KHCO3. ............................................................................................................................ 120 Table 6.4 Standard reduction potential of Sn and Pb in aqueous solutions at 25 °C (V vs. NHE). [132] ............................................................................................................................................ 126 xv Appendix Table A.1.1 Ir-O bond dissociation energy Hf 298 is 352 kJ mol-1 [36]. The basic thermodynamic properties of Ir and IrO2. ................................................................................... 158 Appendix Table B.1.1. XPS quantification of Ir3+/4+ and oxygen environments (IrO2, IrOx(OH)y and H2O) in pristine and used (after potential cycling from 1.2, 1.4 and 1.6 to 0.05 V vs. RHE in 0.09 M H2SO4) heat treated IrOx (untreated, 350, 450 and 550 °C). .......................................... 169 Appendix Table C.1.1 Concentration of metals in the electrolyte after the ex situ AST. Experimental conditions: N2 purged 0.09 M H2SO4, 30,000 cycles from 0.05 V to 1.2 V vs. RHE with a 1 s hold at each potential. ................................................................................................. 170 Appendix Table C.1.2 IrOx/TiO2 elemental atomic % determined by EDX before and after the ex situ AST. ..................................................................................................................................... 172 Appendix Table C.1.3 Cathode effective platinum surface area (EPSA) determined by CO stripping at BOT and EOT (after 300 H2/Air AAST cycles), respectively. Experiments performed in a single cell 20 cm2 MEA. ...................................................................................................... 175 Appendix Table C.2.1. BET surface area and IrOx content of three unsupported IrOx catalysts and six metal oxide supported IrOx catalysts. .................................................................................... 175 xvi List of Figures Figure 1.1. (a) Basic schematic of a hydrogen PEMFC. The anode and cathode are separated by a selectively permeable proton conducting polymer membrane. Water is the only product of the chemical reaction. [32] (b) A theoretical fuel cell (FC) polarization curve showing the kinetic and mass transfer controlled regions modified from Banerjee et al. [33] ............................................. 3 Figure 1.2 Schematic of OER protection in RTAs. (1) Normal PEMFC anode with not-RTA material or modification, (2) mechanism of OER catalyst protection and (3) mechanism of hygroscopic metal oxide (MxOy) protection. .................................................................................. 9 Figure 2.1. (a) Theoretical model of Ru dissolution and O2 evolution on Ru/RuO2 electrodes through a RuO4 intermediate. (b) Model of Ir/IrOx speciation/dissolution and oxygen evolution [124,125]. ...................................................................................................................................... 18 Figure 2.2. Pourbaix diagram of Ir. Possible speciation of Ir at pH -2 to 0 in the potential of near reversal (Red square, 0 V to 1.2 V) and far reversal (green square, 1.2 V to 1.9 V) includes Ir(0), Ir(III), and IrO2[126]. .................................................................................................................... 20 Figure 2.3. Dissolution of Ru with (grey) and without (red) oxide film on a Ru sputtered gold RDE electrode measured by in-line ICP-MS. [118] .............................................................................. 21 Figure 2.4. Tafel plots of an Ir electrode in acid solutions at pH 0.9 (circles), pH 1.9 (crosses) and pH 2.9 (triangles). [131]................................................................................................................ 23 Figure 2.5. First OER scan for bulk and nanoparticle Ru, Ir and Pt at 1600 RPM and 6 mV/s in 0.1 M HClO4. The current is normalized to the number of active sites determined by CO stripping. [60] ................................................................................................................................................ 25 xvii Figure 2.6. Extended cell reversal test for a RuO2 containing RTA. The O2 (squares, dashed) and CO2 (triangles, dotted) current consumed is displayed along with the cell potential (diamonds, solid). [45] ..................................................................................................................................... 26 Figure 2.7. Box diagram of hypothetical ERC process. Products of ERC are dependent on reaction conditions but include formic acid (HCOOH), syngas (CO, H2), methanol (MeOH) and C2 to C4 alcohols and hydrocarbons. The process should be carried out using energy from renewable sources........................................................................................................................................... 31 Figure 3.1 Current (black curve) and frequency (red curve) profiles of Cu deposition and stripping on an Au plated EQCM crystal at a sweep rate of 20 mV s-1. Electrolyte: 0.09 M H2SO4 containing 1 mM CuSO4 at 25 °CRT. ............................................................................................................ 51 Figure 3.2 Sensitivity factor graph correlating the change in frequency (Δf) with the reduction charge (Q). .................................................................................................................................... 53 Figure 3.3 Diagram of the working electrode setup. Teflon based body and screw cap holding the GDL/MPL. Electrical contact was provided by a graphite disk and Pt wire ................................ 58 Figure 4.1 Diagram of the OER durability objectives that are detailed in Chapter 4 and 5. This includes projects at AFCC and Ballard. ........................................................................................ 68 Figure 4.2. (a) Deposition and drying (under a heat lamp) of an ink composed of IrO2 and 2-propanol deposited on the surface of an EQCM crystal. The ink was drop cast in three 5 μL steps (at 0, 12 and 22 minutes) to observe the detachment of particles during multiple depositions. (b) Plot of the deposit frequency change after each ink drop vs. the mass of IrO2 present in the volume of ink. The second y-axis is the calculated mass change from the observed frequency change after each deposition.............................................................................................................................. 71 xviii Figure 4.3. XRD scans of untreated (black), 350°C (red), 450°C (blue) and 550°C (green) heat-treated IrOx. Two crystalline phases were identified (Ir metallic, PDF2 00-006-0598 (ICDD, 2018) and IrO2, PDF2 00-015-0870 (ICDD, 2018)) which are indicated by their Miller indices. ......... 73 Figure 4.4. XPS of the Ir region (a and b) and Oxygen region (c and d) of pristine (a and c) and post CV accelerated stress test (b and d) of IrOx: untreated (black), 350°C (red), 450°C (blue) and 550°C (green) heat-treated. CV cycling performed between 0.05 to 1.2, 1.4 and 1.6 V vs. RHE, 30 times each at 25 °C in 0.09 M H2SO4. .......................................................................................... 75 Figure 4.5. SEM images of pristine (a) untreated IrOx and (b) IrOx 550 °C drop cast on an Au EQCM crystal. Images were acquired at 138 kx. ......................................................................... 76 Figure 4.6. CV response of IrO2 untreated (a) and IrOx heat-treated (350 °C (b), 450 °C (c) and 550 °C (d)) from 1.2 (black), 1.4 (red) and 1.6 V (blue) to 0.05 vs. RHE, 30 times each at 25 °C in 0.09 M H2SO4. .............................................................................................................................. 78 Figure 4.7. EQCM response of untreated (a) and heat-treated (350 °C (b), 450 °C (c) and 550 °C (d)) IrOx during cyclic voltammetry experiments from: 1.2 V (black), 1.4 V (red) and 1.6 V (blue) to 0.05 V vs. RHE. The cycles were repeated 30 times each. “fs” – indicates first potential sweep. Temperature: 25 °C. Electrolyte: 0.09 M H2SO4. ......................................................................... 81 Figure 4.8 Combined frequency and current responses for (a) background Au EQCM crystal, (b) untreated IrO2, (c) 350 °C heat treated IrO2, (d) 450 °C heat treated IrO2 and (e) 550 °C heat treated IrO2 during the final CV scan from 1.6 V to 0.05 V vs. RHE at 25 °C in 0.09 M H2SO4. ........... 84 Figure 4.9. Five minute potential holds at 0.9, 1.2, 1.4, 1.6 and 0.05 V vs. RHE with (a) freshly prepared and (b) CV cycled untreated IrOx. Electrolyte: 0.09 M H2SO4 purged with N2. Scan rate: 100 mV s-1. Temperature: 25 °C. .................................................................................................. 85 xix Figure 4.10. Frequency response of (a) Ketjen black, (b) Pt/C alone and with (c) untreated IrOx and (d) IrOx heat-treated at 550 °C. Catalyst mixtures were 2:1 Pt:Ir metals basis and total solids loading were 2.7, 5.5 and 7.5 μg cm-2 for a, b and both c and d respectively. CV cycled from 1.2 (black), 1.4 (red) and 1.6 V (blue) to 0.05 V vs. RHE. Electrolyte: 0.09 M H2SO4 purged with N2. Scan rate: 100 mV s-1. Temperature: 25 °C. ................................................................................. 87 Figure 4.11. The first CV cycle from 1.6 to 0.05 V vs. RHE of (a) Ketjen black, (b) Pt/C alone and with (c) untreated IrO2 and (d) IrO2 heat-treated at 550 °C. Catalyst mixtures were 2:1 Pt:Ir metals basis. CVs were recorded in 0.09 M H2SO4 purged with N2 at 100 mV s-1 at 25 °C. .................. 89 Figure 4.12. Five minute potential holds at 1.6, 1.8 and 2.0 V vs. RHE with freshly prepared Pt/C (black) and Pt/C with untreated IrO2 (red). Pt/C with untreated IrO2 had to be stopped after 1.8 V vs. RHE because of damage to the crystal. Electrolyte: 0.09 M H2SO4 purged with N2. 25 °C... 91 Figure 5.1 The temperature effect on Ir dissolution of the IrOx/TiO2 catalyst after the 30,000 stepping cycles from 0.05 to 1.2 V vs. RHE in 0.09 M H2SO4 (ex situ AST). ............................. 95 Figure 5.2 Electrolyte Ir concentration after the ex situ accelerated stress test for OER catalysts (three unsupported IrOx and one TiO2 supported). Experimental conditions: N2 purged 0.09 M H2SO4, 40 °C, 30,000 cycles from 0.05 V to 1.2 V vs. RHE with a 1 s hold at each potential. For direct comparison between the unsupported and supported catalysts, the Ir concentration and loss for the supported catalyst was normalized to 100% IrOx content according to Table 3.3. ........... 97 Figure 5.3 Ex situ OER activity characterization: (a) mass activities of three unsupported IrOx catalysts (Low SA, High SA and High SA HT) and IrOx/TiO2 at the BOT, (b) Mass activities at 1.6 V vs. RHE at BOT and at the end of test (EOT, after 30,000 voltage stepping cycles) (mass activity is given in A g-1 Ir metal), (c) The OER overpotentials at 5 mA cm-2geom at BOT and EOT. (d) Specific activities (mA cm-2) at 1.6 V vs. RHE at BOT and EOT. Temperature 40 °C. ...... 100 xx Figure 5.4 (a) Reversal tolerance times for fuel cell anodes containing either unsupported OER catalysts (IrOx Low SA, High SA or High SA HT) or supported OER catalyst (IrOx/TiO2) at the BOT and after an in situ AAST (EOT). In parenthesis the ex situ EOT mass activity values from Figure 5.3a are also provided for easier comparison of the in situ (fuel cell) and ex situ performance. (b) Fuel cell reversal loss (min) vs. Ir concentration after ex situ AST. ............... 103 Figure 5.5 BSE detector images at 500x magnification of (a) freshly deposited and (b) ex situ AST cycled IrOx/TiO2 (81.5 μg cm-2) on a GDL/MPL. The bright dots were confirmed by EDX to contain Ir ..................................................................................................................................... 104 Figure 5.6 SEM images of an MEA after in situ AAST at two magnifications. An area spectrum (a) and point spectrum (b) were taken by EDX to determine the elemental makeup on the cathode...................................................................................................................................................... 105 Figure 5.7 The ex situ interaction effect between Pt HOR and IrOx/TiO2 OER catalysts. (a) Concentration of Ir in the electrolyte after ex situ AST with and without HOR catalyst. (b) Mass activity at 1.6 V vs. RHE (A g-1Ir). Experiments performed in a flooded (N2 purged 0.09 M H2SO4) three electrode cell 30,000 cycles from 0.05 V to 1.2 V vs. RHE with a 1 s hold at each potential. Catalysts were prepared with either IrOx/TiO2 alone or IrOx/TiO2 and carbon supported platinum HOR catalyst. Tests were run three times each at 40°C (red) and 60°C (blue). ......................... 107 Figure 6.1. Cyclic voltammetry of pure Sn (a) and Pb (b) catalysts generated by electrodeposition on carbon cloth using fluoroborate baths . Legend: (1) red N2 purged and (2) black CO2 purged 0.5 M KHCO3 solution. Potential sweep from 0.1 V to -2.0 V vs. Ag/AgCl at 20 mV/s and 293 K...................................................................................................................................................... 114 Figure 6.2. Cyclic voltammetry of bimetallic Pb/Sn catalysts generated by electrodeposition on carbon cloth from fluoroborate baths containing the following molar ratios of salts (Table 1): (a) xxi Sn91%-Pb9%, (b) Sn75%-Pb25% and (c) Pb84%-Sn16% (1) red N2 purged and (2) black CO2 purged 0.5 M KHCO3 solution. Potential sweep from 0.1 V to -2.0 V vs. Ag/AgCl at 20 mV/s and 293 K. .......................................................................................................................................... 115 Figure 6.3. Cumulative formate production faradaic efficiency (FE) (a, c) and total moles of formate synthesized per unit electrode area (b, d) for catalysts produced by electrodeposition using the oxide bath. Chronoamperometry at -2.0 V vs. Ag/AgCl in 0.5 M KHCO3, 293 K, 2 hours. Substrates: (a, b) carbon paper and (c, d) carbon cloth. .............................................................. 117 Figure 6.4. Cumulative formate production faradaic efficiency (FE) (a, c) and total moles of formate synthesized (b, d) for catalysts produced by electrodeposition using the fluoroborate bath. Chronoamperometry at -2.0 V vs Ag/AgCl in 0.5 M KHCO3 at 293 K, 2 hours. Substrates: (a, b) carbon paper and (c, d) carbon cloth. .......................................................................................... 118 Figure 6.5. Fluoroborate deposited Sn100% on carbon cloth: (a) freshly prepared and (b) after CO2 electroreduction at -2.0 V vs. Ag/AgCl in CO2 purged 0.5 M KHCO3 293 K. Catalysts layers were imaged with focused ion beam scanning electron microscopy (FIB-SEM) before and after 2 hours of polarization at -2.0 V vs. Ag/AgCl. ........................................................................................ 121 Figure 6.6 Fluoroborate deposited Pb100% on carbon cloth: (a) freshly prepared and (b) after CO2 electroreduction at -2.0 V vs. Ag/AgCl for 2 hours. ................................................................... 122 Figure 6.7 Fluoroborate deposited Sn75%-Pb25% on carbon cloth: (a) freshly prepared catalyst and (b) after CO2 electroreduction at -2.0 V vs. Ag/AgCl for 2 hours. ...................................... 124 Figure 6.8 XRD spectrum of fluoroborate deposited Sn75%-Pb25% on carbon cloth after polarization for 2 hours at 2.0 V vs. Ag/AgCl in 0.5 M KHCO3 at 293 K. Peak labels are for Miller indices of crystalline Sn (red, S), PDF2 01-072-1012 (ICDD, 2016) and crystalline SnO2 (blue, O), PDF2 00-004-0673 (ICDD, 2016). ....................................................................................... 125 xxii Figure 7.1. Schematic of a continuously operating gas diffusion electrode setup from Bitar et al. [240] ............................................................................................................................................ 137 Appendix Figure B.1.1 STEM images of pristine (a) Pt/C and CV cycled Pt/C with (b) untreated IrOx and (c) IrOx heat-treated at 550 °C. Images were taken at 630 kx magnification. Significant electron beam induced polymerization of the carbon support and other contaminants occurred while capturing the CV cycled images (b and c) which causes the blurry appearance of the images and prevents comparison of the carbon structure. ...................................................................... 159 Appendix Figure B.1.2. SEM images of untreated IrOx (a, b) and IrOx 350 °C (c, d) each catalyst before (a, c) and after (c, d) the stress test (CV cycles from 0.05 to 1.2, 1.4 and 1.6 V vs. RHE, 30 times each). ................................................................................................................................. 160 Appendix Figure B.1.3. SEM images of IrOx 450 °C (a, b) and IrOx 550 °C (c, d) each catalyst before (a, c) and after (c, d) the stress test (CV cycles from 0.05 to 1.2, 1.4 and 1.6 V vs. RHE, 30 times each). ................................................................................................................................. 161 Appendix Figure B.1.4. XPS survey scans of untreated IrOx (a) pristine and (b) after the CV stress test (CV cycles from 0.05 to 1.2, 1.4 and 1.6 V vs. RHE, 30 times each 100 mV s-1) in N2 purged 0.09 M H2SO4 at RT. .................................................................................................................. 162 Appendix Figure B.1.5. Current response of Pt/C (a) alone and with (b) untreated IrOx and (c) IrOx heat-treated at 550 °C. Catalyst mixtures were 2:1 Pt:Ir metals basis. CVs were recorded in 0.09 M H2SO4 purged with N2 at 100 mV s-1 at 25 °C. ........................................................................... 163 Appendix Figure B.1.6. EDS mapping image of Pt/C + untreated IrOx after CV stress test in 0.09 M H2SO4 purged with N2 at 100 mV s-1. Oxygen (blue), Pt (yellow) and Ir (red) color intensity is proportional to their respective EDS signal strengths................................................................. 164 xxiii Appendix Figure B.1.7. The frequency (a) and CV (b) response of an Au EQCM crystal CV cycled from 1.2 (black), 1.4 (red) and 1.6 V (blue) to 0.05 vs. RHE, 30 times each at 100 mV s-1. Electrolyte: N2 purged 0.09 M H2SO4. 25 °C. ............................................................................ 165 Appendix Figure B.1.8. Deconvolution of XPS peaks in the Ir region of pristine IrOx (a) untreated and heat-treated at (b) 350 °C (c) 450 °C (d) 550 °C. ................................................................ 166 Appendix Figure B.1.9. Deconvolution of XPS peaks in the Oxygen region of pristine IrOx (a) untreated and heat-treated at (b) 350 °C (c) 450 °C (d) 550 °C.................................................. 167 Appendix Figure B.1.10. Deconvolution of XPS peaks of IrOx (a-Ir and b-Oxygen) untreated and (c-Ir and d-Oxygen) heat-treated at 550 °C after potential cycling from 1.2, 1.4 and 1.6 to 0.05 V vs. RHE in 0.09 M H2SO4. .......................................................................................................... 168 Appendix Figure C.1.1 XRD IrOx/TiOx showing two distinct phases, TiO2 (green, circles), PDF2 03-065-5714 (ICDD, 2017), IrO2 (blue, triangles), PDF2 01-088-0288 (ICDD, 2017) as well as an amorphous peak at 2 =12.5o. ..................................................................................................... 171 Appendix Figure C.1.2 SEM and EDX of IrOx/TiO2 before (a) and after (b) the ex situ AST (30,000 stepping cycles from 0.05 to 1.2 V vs RHE; 0.09 M H2SO4). .................................................... 171 Appendix Figure C.1.3 SEM images of (a) Low SA, (b) High SA HT and (c) High SA IrOx at 3000x magnification. Images were generated using a backscatter detector after the ex situ AST (30,000 stepping cycles from 0.05 to 1.2 V vs RHE; 0.09 M H2SO4) ........................................ 172 Appendix Figure C.1.4 XRD of freshly prepared IrOx High SA HT powder (red) and as received IrOx High SA (green) and IrOx Low SA (black). One crystalline phase is identified as IrO2 (PDF2 00-015-0870 (ICDD, 2017)) for the IrOx High SA HT (red lines and Miller indices). ............. 173 Appendix Figure C.1.5 Cell polarization curves at BOT, 100, 200 and 300 H2/Air AAST cycles, respectively. Anode Pt/Ir loading was 50/20 μg cm-2 and cathode Pt loading was 250 μg cm-2. OER xxiv catalysts were (a) IrOx/TiO2 (b) IrOx Low SA (c) IrOx High SA HT. Cell polarization was carried out in “differential mode” with a stoichiometry of 6 at the maximum current. .......................... 174 Appendix Figure C.2.1. Electrolyte Ir concentration after accelerated stress test. Performed in a flooded (N2 purged 0.09 M H2SO4) three electrode cell 30,000 cycles from 0.05 V to 1.2 V vs. RHE with a 1 s hold at each potential. ........................................................................................ 176 Appendix Figure C.2.2. OER activity was measured in 0.09 M H2SO4 by LSV from 1.9 V to 1.3 V vs. RHE with a sweep rate of 5 mV s-1. 81.5 μg cm-2 of each catalyst was freshly deposited onto a GDL.MPL support material. (a) The overpotentials at 5 mA for each catalyst is plotted at the BOT and the EOT after 30k voltage stepping cycles (b) five metal oxide supported (IrOx/MxOy #1, IrOx/MOy #2, IrOx/MOy #3 IrOx/MOy, IrOx/MOy, #4 and IrOx/TiO2) were analyzed................ 177 Appendix Figure C.2.3. (a) Reversal tolerance time at 0.2 A under N2 of various IrOx based OER catalysts using a freshly prepared anode at the BOT and after an in situ AAST on a second freshly prepared MEA (EOT). (b) Scatter plot of Ir dissolution after ex situ AST vs. in situ reversal loss (min). Metal oxide supported IrOx catalysts (solid red box) and unsupported IrOx catalysts (dashed blue box) were tested using both methods .................................................................................. 178 Appendix Figure D.1.1 Chronoamperometry profiles for CO2 electroreduction at -2.0 V vs. Ag/AgCl for 2 hours in CO2 purged 0.5 M KHCO3. 293 K. Legend: (a) Catalysts supported on carbon paper obtained from the oxide based bath and (b) Catalysts supported on carbon cloth obtained from the oxide based bath, (c) Catalysts supported on carbon paper obtained from the fluoroborate based bath and (d) Catalysts supported on carbon cloth obtained from the fluoroborate bath. (1) Sn100%, (2) Sn80%-Pb20% (oxide) and Sn91%-Pb9% (fluoroborate), (3) Sn67%-Pb33% (oxide) and Sn75%-Pb25% (fluoroborate), (4) Sn50%-Pb50% (oxide) and Sn16%-Pb84% (fluoroborate), (5) Pb100%. ........................................................................................................ 180 xxv Appendix Figure D.1.2 XRD spectrum of fluoroborate deposited Sn100% on carbon cloth after polarization for 2 hours at 2.0 V vs. Ag/AgCl in 0.5 M KHCO3. Peak labels are for the Miller indices of crystalline Sn (S) (PDF2 00-004-0673 (ICDD, 2016)). ............................................. 181 Appendix Figure D.1.3 XRD spectrum of fluoroborate deposited Sn75%-Pb25% on carbon cloth after polarization for 2 hours at 2.0 V vs. Ag/AgCl in 0.5 M KHCO3 at 293 K. Peak labels are for Miller indices of crystalline Sn (S) PDF2 00-004-0673 (ICDD, 2016) and crystalline SnO2 (O) PDF2 03-065-2873 (ICDD, 2016). ............................................................................................. 182 Appendix Figure D.1.4 XRD spectrum of fluoroborate deposited Pb100% on carbon cloth after polarization for 2 hours at -2.0 V vs. Ag/AgCl. Peak labels are for the Miller indices of crystalline lead (Pb) (P), PDF2 03-065-2873 (ICDD, 2016), lead oxide (PbO) (O), PDF2 01-078-1666 (ICDD, 2016), massicot (PbO) (M), PDF2 01-077-1971 (ICDD, 2016) and plumbonacrite (Pb5(CO3)O(OH)2) (A), PDF2 01-073-8925 (ICDD, 2016). ...................................................... 183 Appendix Figure D.1.5 XRD spectrum of fluoroborate deposited Sn16%-Pb84% on carbon cloth after polarization for 2 hours at -2.0 V vs. Ag/AgCl. Peak labels are for the Miller indices of litharge (OPb) (L), PDF2 00-005-0561 (ICDD, 2016), lead oxide hydrate (3PbO∙H2O) (H), PDF2 00-022-1134 (ICDD, 2016) and potassium lead carbonate hydroxide (KPb2(CO3)2OH) (C), PDF2 00-037-0502 (ICDD, 2016). ....................................................................................................... 184 Appendix Figure D.1.6 XRD spectrum of Sn91%-Pb9% on carbon cloth after polarization for 2 hours at 2.0 V vs. Ag/AgCl. Peak labels are for the miller indices of tin (Sn) (S), PDF2 03-065-0298 (ICDD, 2016) and lead (Pb) (P), PDF2 03-065-2873 (ICDD, 2016). ............................... 185 Appendix Figure D.1.7 FIB-SEM images of Sn16%-Pb84% deposited without Triton (a) freshly prepared and (b) after 2 hours of use as a CO2 reduction catalyst layers at 5000X magnification...................................................................................................................................................... 186 xxvi Appendix Figure D.1.8 Megazyme formate testing kit (a) absorbance spectrum at four formate concentrations in 0.5 M KHCO3 and blank KHCO3 solution. (b) Calibration curve and linear regression used for calculation of formate concentrations in electrolyte solutions. ................... 186 xxvii List of Symbols Symbol Definition Units (S.I. Or Customary) 𝝍𝑪𝑶𝟐 Current efficiency of ERC % 𝝏𝝓(𝒙)𝝏𝒙 Potential gradient V m-1 A Area m2 or cm2 b Tafel slope mV dec-1 Cf Sauerbrey constant Hz μg-1 Cj Convection of species j m2 s-1 Dj Mass diffusivity of species j m2 s-1 E°298K Standard reduction potential V vs. SHE F Faraday’s constant 96485 C mol-1e- f Frequency Hz f0 Quartz fundamental frequency Hz i Current density A m-2 or mA cm-2 i0 Exchange current density A or mA Jj(x) Flux of species j mol m-2 K Crystallite shape factor usually 0.89 to 0.94 m/z Mass to charge ratio kg C-1 Mw Molecular weight Kg mol-1 or g mol-1 n Number of electrons p Partial pressure atm Q Charge C xxviii Symbol Definition Units (S.I. Or Customary) R Gas constant 8.314 J mol-1 K-1 T Temperature Kelvin or °C v(x) Velocity in the x direction m s-1 vrxn Rate of mass transfer mol m-2 s-1 vrxn Rate of reaction mol m-3 s-1 Z Impedance Ω zj Charge of species j ΔG Gibbs free energy J ΔH Enthalpy joules ΔS Entropy joules α Transfer coefficient ca. 0.5 β Line broadening at half max intensity radians η Overpotential V θ Bragg angle radians λ X-ray wavelength m or nm μq Shear modulus of quartz 2.947x1011 g cm-1 s-2 ρq Density of quartz 2.648 g cm-3 τ Crystallite size m or nm ω Frequency Hz xxix List of Abbreviations Abbreviation Definition AAST Fuel cell anode accelerated stress test (in situ) AST Accelerated stress test (ex situ) ATO Antimony tin oxide Bar(a) Bar absolute pressure BET Brunauer Emmett and Teller BF Bright field BOP Balance of plant BOT Beginning of test BSE Backscatter electrons CCS Carbon capture and storage CCU Carbon capture and utilization CD Current density CL Catalyst layer CO Carbon Monoxide CO2 Carbon dioxide CV Cyclic voltammetry DFT Density field theory DOE Department of energy ECSA Electrochemically active surface area EDX Energy dispersive X-rays EE Energy efficiency xxx Abbreviation Definition EOT End of test EQCM Electrochemical quartz crystal microbalance EtOH Ethanol FC Fuel cell FDH Formate dehydrogenase enzyme FE Faradaic efficiency FE-SEM field emission scanning electron microscope GC Gas chromatography GDL Gas diffusion layer H2 Hydrogen HAADF High-angle annular dark field HCOO- Formate HOR Hydrogen oxidation reaction HPLC High performance liquid chromatrography ICE Internal combustion engine ICP-MS Inductively coupled plasma mass spectrometry ILTEM Identical location transmission electron microscopy IrOx High SA Iridium oxide high surface area IrOx High SA HT Iridium oxide high surface area heat-treated IrOx Low SA Iridium oxide low surface area IrOx/TiO2 Iridium oxide supported on titanium dioxide LOD limit of detection xxxi Abbreviation Definition LSV Linear sweep voltammetry MEA Membrane electrode assembly MeOH Methanol MMO Mixed metal oxide MPL Micro-porous layer NADH/NAD+ Nicotinamide adenine dinucleotide O2 Oxygen OCV Open circuit voltage OER Oxygen evolution reaction ORR Oxygen reduction reaction Pb Lead PEMFC Proton exchange membrane fuel cell PFSA Perfluorosulfonic acid PGM Platinum group metals PrOH Propanol QCM Quartz crystal microbalance RHE Reversible hydrogen electrode RTA Reversal tolerant anode SEM Scanning electron microscopy SHE Standard hydrogen electrode SLPM Standard liters per minute SnOx Tin dioxide xxxii Abbreviation Definition STEM Scanning transmission electron microscopy UV Ultraviolet VP-SEM Variable pressure scanning electron microscope XANES X-ray absorption near edge structure XPS X-ray photo-electron spectroscopy XRD X-ray diffraction xxxiii Acknowledgements I would like to thank Prof. Gyenge for his guidance and expertise in preparing this thesis. Thanks also to Dr. Natalia Kremliakova and Dr. Max Cimenti for their extensive PEMFC knowledge and persistent enthusiasm for the project. Thank you to Dr. Foroughazam Afsahi for continued discussion and troubleshooting of EQCM, Dr. Ping He, Dr. Siyu Ye, Alan Young and Shanna Knights for supporting on the project after the closure of AFCC. The important input of my committee members, Prof. Kevin Smith and Prof. David Wilkinson is greatly appreciated. Thank you to labmates, especially Thomas Rowe for generating and testing catalyst materials for support both inside and outside the lab. The financial support from Mathematics of Information Technology and Complex Systems (MITACS) and Natural Sciences and Engineering Research Council (NSERC) of Canada is gratefully acknowledged. I would like to thank Derrick Horne from UBC Bioimaging and Jacob Kabel from UBC Materials Engineering for their EDX and SEM imaging assistance, Anita Lam from UBC X-ray crystallography for XRD crystal phase determination. I would like to thank Dr. Myles Bos for his XPS interpretation. Xin Zhang and Dennis Hsiao are acknowledged for their STEM and XPS analysis respectively. To my sister Katie and my mom Diane for always being there to listen and give advice. Finally, I would like to thank my wife Noémi. This is as much hers as it is mine. She has supported me in countless ways and I would not have continued without her. xxxiv Dedication This thesis is dedicated to my daughter Emilia. You give me so much purpose and inspiration. You are the most important thing to me and I hope you always feel my love in your heart. 1 Chapter 1: Introduction 1.1 Motivation Climate change caused by human activity increasing CO2 concentration in the atmosphere (i.e. global average CO2 concentration of 400 ppm by volume) has far-reaching economic, political and societal consequences. In 2004, 77% of all greenhouse gas emissions came from CO2 and total emissions of CO2 rose by 80% between 1970 and 2004. [1] Reducing atmospheric CO2 emissions should be aggressively pursued to help keep global temperature rise below 1.5 °C using a multi-pronged approach of increasing the efficiency of current systems, implementation of carbon capture technologies and transitioning to a low-carbon economy. [2] Efforts to reduce CO2 emissions include replacing coal and natural gas power plants with wind [3,4], solar [5], hydroelectric [6,7] and nuclear power. [8–10] Using intermittent power sources such as wind and solar requires energy storage methods to capture energy during non-peak demand times. [11] Hydrogen generated by electrolysis is ideal for this purpose because it can be generated, stored and used again in PEMFCs. [12] Similarly, CO2 captured via direct air capture or from flue gas can be electroreduced to formate, CO or other hydrocarbons. [13] These fuels generated from non-peak demand renewable energy sources can then be used power other processes. [14,15] 1.2 Catalyst durability in fuel cells and carbon electroreduction In order for PEMFCs and ERC to be used in industrial or commercial applications, the systems must be robust and long lasting as compared to current technologies. Catalyst breakdown, through dissolution, detachment and deactivation, contributes to the limited lifetimes of these systems. In this thesis, the durability and longevity of PEMFCs and ERC systems will be increased through improving the electrocatalysts that are at the heart of these technologies. It is imperative to generate new catalysts that are resistant to known breakdown mechanisms. Therefore, the focus of this 2 thesis is not only on generating new catalysts but also on uncovering the mechanism of improved durability. 1.3 Proton exchange membrane fuel cell overview PEMFCs use thin polymer membranes which transfer protons and electrically isolate the anodic and cathodic reactions. They were first developed at General Electric and used in the Gemini space missions. [16] PEMFCs have been proposed for use in energy storage/utilization from intermittent sources [8,11], public transportation [17], unmanned aerial vehicles [18], and automotive applications. [19,20] They promise lower point of use pollution and better air quality than current fossil fuel based technologies. There are pilot programs of PEMFC cars including the Toyota Mirai [21], Honda Clarity [22], and the Mercedes GLC F-Cell [23] among others. Fleets of fuel cell busses have also been deployed in the U.S. [24], Canada [17], Scotland [25], and Brazil. [26] However, issues with cost versus diesel-powered buses have prevented the bus fleet programs from continuing or expanding. [27] Automotive mixed drive applications have also been suggested in which the PEMFC system work in concert with batteries, super capacitors or photovoltaic cells. [20] Many comprehensive PEMFC reviews have been published [16,28–30] covering numerous aspects of PEMFC operation and function. The basic schematic of a hydrogen PEMFC is shown in Figure 1.1a. A solid polymer electrolyte membrane electrically isolates anode and cathode reactions while providing proton conductivity. The polymer electrolyte is selectively permeable to protons and minimizes the crossover of H2 and air. Water formed on the cathode and any excess fuel from the anode flow through the cell and are exhausted or recycled. Special bi-polar plates ensure proper flow of fuel and oxidant through the PEMFC stack. The schematic presented below shows only a single cell but in a real application, several cells would be connected in series to 3 produce a higher voltage fuel cell stack. The number of cells is determined based on the power needs of the particular application. [31] \ Figure 1.1. (a) Basic schematic of a hydrogen PEMFC. The anode and cathode are separated by a selectively permeable proton conducting polymer membrane. Water is the only product of the chemical reaction. [32] (b) A theoretical fuel cell (FC) polarization curve showing the kinetic and mass transfer controlled regions modified from Banerjee et al. [33] EOC measuredKinetic lossesOhmic lossesMass transfer limited regionb 4 Figure 1.1b shows the polarization curve of a typical PEMFC. The voltage is at its maximum at minimum current called the open circuit potential (Eoc). As the current density increases, kinetic losses reduce the cell potential. The ohmic region follows where cell resistance is the major factor reducing cell voltage. Finally, the flow of fuel and oxidant limit the maximum current density in the mass transfer limited region. On the right axis, the power density curve is shown. This has a maximum at ca. 2.2 V just before the mass transfer limited region. [33–35] A PEMFC running on hydrogen fuel (H2) and air uses two chemical reactions. [36] H2 → 2H+ + 2e- E°25 °C = 0.00 V vs. SHE Equation 1.1 O2 + 4H+ + 4e- → 2H2O E°25 °C = 1.23 V vs. SHE Equation 1.2 Equation 1.1 is the hydrogen oxidation reaction (HOR) which consumes the hydrogen fuel on the anode side and produces protons that are transported to the cathode. Equation 1.2 is the oxygen reduction reaction (ORR) which occurs on the cathode side of a PEMFC. This reaction consumes the protons produced at the anode to produce two moles of water for each mole of diatomic oxygen. The standard reduction potentials listed above are for standard conditions (1 atm, 298 K). To determine the equilibrium cell potential (Ee, cell) and Gibbs free energy, for a PEMFC, one must use the Nernst equation. This equation is dependent on the partial pressures (p) of the product (H2O) and reactants (H2 and O2). 𝐸𝑒,𝑐𝑒𝑙𝑙 = 𝐸𝑇,𝑐𝑒𝑙𝑙𝑜 +𝑅𝑇2𝐹𝑙𝑛 (𝑝𝐻2𝑝𝑂21/2𝑝𝐻2𝑂) Equation 1.3 𝐸𝑇𝑜 = 𝐸298𝑜 +∆𝑆298𝑜𝑛𝐹(𝑇 − 298) Equation 1.4 5 To determine the equilibrium cell potential at elevated temperature one must use Equation 1.4 using the entropy of the reaction (ΔS). Thermodynamic equilibrium potential can never be reached in practice due to charge and/or mass transfer kinetic limitations (usually expressed as overpotentials), and ohmic potential losses. However, a major goal of PEMFC engineering and design is to minimize these energy losses and maximize the cell potential. In modern PEMFCs, ohmic cell losses are minimized by using thin (e.g. 20 to 50 µm) polymer electrolyte separator to transfer protons from the anode to the cathode. The ohmic resistance depends on membrane hydration, membrane thickness and temperature. [35] The polymer electrolyte is typically made from perfluorosulfonic acid (PFSA) under the commercial name Nafion® and others. While the polymer electrolyte must be proton conducting, it also has to be electrically insulating and gas impermeable to prevent shorting and fuel/oxidant crossover, respectively. The polymer electrolyte is sandwiched between thin anode and cathode layers made from catalyst inks either deposited directly onto the membrane or onto a graphitic micro-porous layer (MPL) backed by a carbon paper gas diffusion layer (GDL). This is collectively known as the membrane electrode assembly (MEA). 1.4 Department of Energy targets The US Department of Energy (DOE) has developed guidelines to help steer the course of research towards achieving common goals. The DOE targets for 2020 are to reduce the cost of MEA production to $14/kW and have total Platinum-group metal (PGM) loading of 0.15 mg cm-2, which include, in order of atomic number: Ru, Ru, Pd, Os, Ir and Pt. Cost targets for the fuel cell stack are $20/kW. The target for the entire fuel cell system including the balance of plant (BOP) and thermal system but excluding hydrogen storage, batteries and electronics is $40/kW. BOP is any 6 supporting component that was not included in the MEA and stack assembly but are necessary for the functioning of the fuel cell system. The fuel cell durability should target to 5,000 h, which is equivalent to 150,000 miles of driving. This is to compete with internal combustion engines. Testing conditions should encompass real conditions such as variable loading, shock absorption and heat cycles and achieve a current density of 300 mA cm-2 at 0.8 V. [19] Pt/C HOR catalysts on the anode are much more active than Pt/C ORR catalysts on the cathode. Therefore, to reach DOE targets for PEMFCs, the anode catalyst can be modified to reduce the PGM loading by using non-PGM catalysts or by engineering the PGMs to have a higher mass-activity. The loading of HOR catalysts on the anode is much lower than the ORR catalysts on the cathode. Even so, it is still advantageous to pursue PGM loading reduction on both electrodes. To improve stack lifetime, engineering safeguards into the fuel cell stacks such as cell potential sensors or constant hydrogen circulation could solve some of the issues with fuel starvation. However, these additions would increase the BOP costs of producing a PEMFC stack. The goal should be to reduce the total cost of the fuel cell system and limit the possibility of failure in the cell while at the same time improving the durability of the cell and resistance to voltage reversal induced damage. 1.5 PEMFC failure modes Several modes of performance degradation exist in PEMFCs [37] including platinum dissolution from the cathode [38,39], water management failure [40], membrane degradation due to metal migration and peroxide attack [41], mechanical breakdown of the membrane [42], mechanical breakdown of seals [43], and anode degradation caused by fuel starvation. [44,45] These failure modes often occur in concert and one failure can lead to cascading failures in other areas of the 7 cell. This thesis will focus on anode degradation due to fuel starvation resulting in cell voltage reversal and carbon corrosion of the GDL/MPL. During normal PEMFC operation, an excess of oxygen and hydrogen are supplied to the cathode and anode respectively. The hydrogen flow can fail to reach the anode either because of water blocking the flow channels [44,46], ice formation [37] or during transient load change such as start-up/shutdown procedures. When one of these situations occurs, current continues to flow through the cell stack even though there is a lack of H2 present at the anode of a particular cell. The other cells of the stack drive the starved cell into reversal. Even after flushing the stack with fuel, localized fuel starvation can occur due to poor distribution of fuel on the anode. [45] In the absence of H2, water, the platinum catalyst, and the carbon-based components of MEA can oxidize to complete the load circuit. When this occurs, the potential on the anode can rise to over 1.9 V and the cell voltage becomes negative. If carbon oxidation is the primary reaction, the carbon-based structures in the anode layer will corrode by the following reaction. [36] E°25 °C = 0.21 V C + 2H2O → CO2 + 4H+ + 4e- Equation 1.5 This could cause loss of electric conductivity due to corrosion of the carbon support and loss of adhesion of the catalyst nanoparticles to the support with catastrophic effects for cell performance. [47] The MPL and GDL are also subject to corrosion, increasing the contact resistance between the catalyst layer (CL) and MPL/GDL. [48] Carbon corrosion at the anode caused by fuel starvation can be observed during operation as cell voltage reversal to negative cell voltage (< -2 V vs SHE). [49] Cell voltage reversal is a transient condition that will subside once the cell is again flooded with fuel but the damage to the cell can cause permanent cell voltage loss of 0.5 V or more. The cell voltage loss is a result of irreversible 8 damage to the carbonaceous components of the MEA. This kind of damage can cause the catastrophic failure of a particular cell resulting in the need for disassembly and servicing. [50] 1.6 Cell reversal mitigation strategies Systems mitigation, by monitoring individual cell voltage and controlling individual cell electrical connections and fuel/oxidant flow rates, can be an effective strategy to limit cell voltage reversal. [51] However, this approach adds complexity and cost to the PEMFC stack. Importantly, these systems may not even be able to change fuel flow quickly enough to prevent a voltage reversal event. In addition, areas within the anode can experience localized cell voltage reversal that may not cause a change in cell voltage. [52] Therefore, it is advantageous to build fuel cells that have cell voltage reversal tolerance incorporated into the anode catalyst layer. To address fuel starvation induced cell failure, extensive research has been performed to produce reversal tolerant anodes (RTAs). [44,45,50,53,54] This can be accomplished by changing the anode structure [55], by using catalysts with corrosion resistant support [45,56] or by imbedding OER catalysts. [57] During a reversal event, the OER catalysts lower the onset potential required to oxidize the water present at the anode and prevent the potential rising high enough to oxidize the carbon support material. [45,58–61]. Different oxides (including SiO2) have been shown to protect carbon components of the cathode from corrosion by preventing the adsorption of water on the carbon surface. [62,63] Additionally, metal oxides have been suggested to cover the surface to the anode with a more corrosion resistant layer. [56] Figure 1.2 shows a schematic of the possible routes of RTA protection. 9 Figure 1.2 Schematic of OER protection in RTAs. (1) Normal PEMFC anode with not-RTA material or modification, (2) mechanism of OER catalyst protection and (3) mechanism of hygroscopic metal oxide (MxOy) protection. Finally, a fast OER reaction will increase the local H+ activity while lowering the local H2O activity on the fuel cell anode surface. According to the Nernst equation, these local activity changes will trigger the increase of the equilibrium potential for the carbon corrosion (Equation 1.5). Thereby, the overpotential for ERC is diminished at a given anode potential, which in turn lowers the carbon corrosion rate. Oxygen evolution is the reverse reaction to oxygen reduction (Equation 1.2), which takes place on the cathode of a hydrogen-oxygen PEMFC. This reaction is thermodynamically unfavorable when compared to carbon corrosion (0.21 V vs 1.23 V) and therefore kinetic promoter catalysts are needed. Even though carbon oxidation is much more energetically favorable, there is large kinetic inhibition and higher activation energy for carbon oxidation as compared to OER. This means that 10 there is a high overpotential for carbon corrosion, which increases the Carbon corrosion onset potential into the region of OER. The sluggish kinetics of carbon corrosion is fortunate in this case and allows us to use OER catalysts as a cell voltage reversal mitigation strategy. It is not the goal of this thesis to show that OER catalysts are capable of protecting Pt/C and other carbon containing components of the anode. This capability has been described extensively. [45,49,53,64–66] Instead, this thesis will focus on the degradation of the OER catalysts themselves. 1.7 Choosing effective oxygen evolution catalysts for reversal tolerant anodes A good OER catalyst candidate should have a few specific properties. First, an effective OER catalyst should limit the anode potential during a reversal event by promoting proton generation coupled with oxygen evolution, thereby, slowing down or even inhibiting the carbon corrosion reaction. OER catalysts should also be stable and durable under normal PEMFC operating conditions (pH ≤ -1 and potential of almost 0 V in the presence of H2). If the OER catalyst dissolves and loses OER active surface area, this will lead to a decrease in the reversal tolerance time of the anode. Therefore, it is expected that a link will exist between the rate of dissolution of OER catalysts and the decrease in reversal tolerance time after an accelerated stress test. When incorporated into the fuel cell anode, OER catalyst should not produce any metal cations that affect the catalyst activity at the cathode. Finally, to protect the membrane from degradation, OER catalysts should not contain Fenton’s metals (Fe, Cr, Co, Cu, Ni, etc.) that could catalyze the formation of peroxyl and hydroxyl radicals. Peroxyl and hydroxyl radicals induce membrane breakdown and lead to cell performance loss due to high proton resistance or short-circuiting. [67,68] Extensive research has been carried out on RTAs using OER catalysts. [44,45,53,54] However, most of these results were obtained using in situ testing that requires time, equipment and expertise 11 to fabricate the MEA and carry out long term experiments (on the order of days to weeks). Ex situ characterization can be faster and less material intensive than running a fuel cell stack. 1.8 Electrochemical quartz crystal microbalance as an ex situ oxygen evolution reaction catalyst durability investigation tool The EQCM uses a piezoelectric sensor (in this case a vibrating quartz crystal) where the vibration frequency is measured concurrently with potential. This vibration frequency is inversely proportional to the mass of the crystal according to Equation 1.6. ∆𝑓 = 𝐶𝑓∆𝑚 Equation 1.6 where Δf is the observed frequency change (Hz), Δm is the mass change per unit area (µg cm-2) and Cf is the quartz crystal sensitivity factor (referred to also as the Sauerbrey constant) (Hz μg-1 cm2). The conversion of observed frequency to mass via Equation 1.6 is only valid for a rigid uniform film on the surface to the crystal. Instead of having a uniform rigid resonator, particle aggregates are being investigated which can exhibit viscous or elastic behavior during CV cycling. [69–72] This complicates the interpretation of frequency changes of powders. Despite these limitations, combining EQCM with other detection and analysis techniques can still yield valuable data points that can be used to develop ex situ stress test conditions. In Chapter 4, the effect of heat treatment on potential dependent mass changes of IrO2 powders is investigated for the first time using EQCM. IrO2 and Pt/C are combined to form a layer that mimics a voltage reversal tolerant PEMFC anode catalyst. The addition of IrO2 to Pt/C is investigated for its ability to protect against carbon corrosion and Pt dissolution at high potentials. With this layer potential dependent frequency changes are investigated and insight is provided into the processes occurring at the electrode surface in order to de-convolute the EQCM response into its constituent 12 parts. EQCM is a powerful method that is rapid and cost effective compared to in situ investigation or ex situ investigation without inline detection of dissolution, corrosion or other chemical changes. 1.9 Electroreduction of CO2 proposal and addition to thesis Before the work on reversal tolerant anodes began, separate thesis proposal was written which described the plan for a project on electroreduction of CO2 to value added products (including formate) in trickle bed and gas diffusion electrolyzers. This was a separate Mitacs project that began in collaboration with Mantra Energy Alternatives Ltd. (formerly of Vancouver, B.C.). While the new OER project is a departure from the original topic, the common theme between the ERC and the OER work is electrocatalyst durability and longevity. This thesis will encompass the OER project at AFCC published in the Journal of Power Sources [73] and continued using EQCM at Ballard as well as the ERC work and the paper published in ChemSusChem. [74] The study of Sn-Pb catalysts is expanded on by investigating the interactions between electrodeposition conditions, substrates (carbon cloth vs. teflonated carbon paper) and electrocatalytic activity. Furthermore, insights are made with regard to degradation modes and bimetallic catalyst durability. 1.10 Carbon dioxide conversion and storage technologies Despite efforts to reduce CO2 emissions, there are still industries that produce CO2 as necessary byproduct and not as a means to produce energy, including breweries where yeast give off CO2 during the fermentation process and concrete factories where carbonates involved in the production of concrete dissociate forming CO2. One possible route would be to bury CO2 underground into spent oil wells in a process called carbon capture and sequestration (CCS). [75,76] The process relies on the formation of carbonate deposits after injection in underground aqueous aquifers. However, potential environmental side effects of carbon sequestration exist [77] 13 and importantly, it produces little financial incentive for the companies to reduce their CO2 emissions other than carbon tax avoidance. Emphasizing the financial incentive for new technologies, generation of value added products provides a more viable and sustainable route to deal with the problem of CO2 emission mitigation on a massive scale. This is a broad field of study and including aspects of chemistry, engineering and economics known as carbon capture and utilization (CCU). Informative reviews of the economic issues involved have been prepared. [76,78,79] CO2 conversion to value added products such as CO, HCOOH, CH4, alcohols or hydrocarbons [80,81] is intensely investigated by all possible catalytic pathways such as: photocatalytic reduction, [82,83] homogeneous chemical processes, [84–86] heterogeneous thermocatalysis [84], biocatalytic reactors [87,88] and ERC. [89] Each process has its own advantages and disadvantages. For instance, photobioreactors can run on sunlight and convert CO2 directly to sugars and other longer chain carbons. Photosynthetic parts of plants and algae are like nature’s photobioreactors. Photosynthetic plants are the pioneers of carbon sequestration and their structure can inform new strategies for how to reduce atmospheric CO2 concentrations. Homogeneous chemical reduction of CO2 with H2 is possible using ruthenium catalysts even without the use of acid or base. [90] Heterogeneous catalysis can be carried out by reverse water gas shift [91,92], direct Fischer Tropsch synthesis [93,94] and CO2 to methanol synthesis. [95] However, these processes involve high heat and pressure conditions which reduce the overall energy efficiency of the processes. CO2 is useful directly as a stable solvent for extractions and is very non-reactive towards most species. [96] It can be used as a gaseous fertilizer [97], as a solvent in polycarbonate plastic manufacture and in the manufacture of carbonate salts and salicylic acid. Photocatalytic reduction uses photocatalysts such as metal oxides, mixed metal 14 oxide semiconductors and mixed supramolecular photocatalysts [83,98] to make value added products from CO2. Hydrocarbon fuels can also be synthesized in photobioreactors utilizing algae [88], or electrolytic bioreactors using methanogenic bacterial films [87]. ERC reactors can be altered depending on the catalyst material to produce a variety of products. This allows the technology to be adapted to the specific situation. [99,100] ERC reactors can have good energy efficiency (EE) (>70%), current density (CD) (>100mA cm2) or faradaic efficiency (FE) (>80%) but obtaining all three in one system has proven difficult. [89] For ERC it is essential for the energy to come from renewable resources. The chemical products could generate an economic incentive for industry to adopt CO2 conversion technologies. [79] 1.11 Durability requirements for electroreduction of CO2 For economic viability, an industrial electrochemical process should operate at current densities typically between 100 and 700 mA cm-2 for 10,000 hours continuously, with faradaic efficiencies above 90% and cell voltage drift of only a few µV h-1. In an effort to develop such a process, Pb and Sn electrodes were directly compared in a continuous flow reactor and key parameters including flow rate, and current density were optimized. It was concluded that Sn was capable of a 25% higher rate of HCOO- production than Pb at 12.2 mA cm-2. [101] This was due to higher flow rates improving the performance of the Sn electrode but not the Pb electrode, which was under electrode kinetic limitation. [102] Li and Oloman achieved much higher current densities using diverse porous Sn cathodes in trickle-bed reactors in the range of 60 to 300 mA cm-2. [103] However, the faradaic efficiencies were only between 40 to 60%. Over time, the competing hydrogen evolution reaction (HER) became the dominant reaction. Promisingly, Sn based gas diffusion electrodes [104] were also tested in a semi-batch mode and current densities of 200 mA cm-2 were achieved with current efficiencies about 85%. [105] 15 1.12 Industrial application requirements for electroreduction of CO2 Li and Oloman stated that it is necessary for a practical industrial application that a system be capable of >70% faradaic efficiency at >100 mA cm-2 current density and specific energy < 500 kWh kmol-1. [106] In the same paper, the best ERC results were found to be a current density of 130 mA cm-2 and a faradaic efficiency of 86% and specific electrochemical energy of 260 kWh kmol-1formate in a trickle bed reactor. The current density was later improved to 310 mA cm-2 with faradaic efficiencies ranging from 90% to as low as 45% after three hours of electrolysis. [101] The temporal stability of the Sn based catalysts was not sufficient for industrial applications. Several studies have shown that the faradaic efficiency of Sn cathodes decreases over the first four to eight hours of electroreduction and the HER becomes the predominant reaction. [101–103] This is because of tin dioxide (SnOx) layer is will be electrochemically reduced to Sn(0) which catalyzes HER instead of ERC. [107] H2 is still a valuable product but it can be produced with lower overpotential and higher energy efficiency using precious metal catalysts. A few companies are developing ERC technology for widespread use. Liquid Light, Monmouth Junction NJ, is focused on generating value added products from CO2 using low energy processes from renewable sources. They perform hydrogenation and purification of the products to form alcohols, olefins, glycols and organic acids. Their technology is designed to use mixed feedstocks to minimize the need for upstream feedstock separation. [108] Dioxide Materials, Boca Raton FL, is a multifaceted company with products including water electrolyzers, anion exchange membranes and ERC electrolyzers. They claim a maximum of 600 mA cm-2 at 3.25 V and 95% faradaic efficiency for CO using their Sustanion® membrane. They claim a 6 μV hr-1 increase in operating voltage with up to 6 months of functioning without any loss of performance. They state that a breakthrough came from the use of an ionic liquid with metallic Ag. [109,110] 16 While commercial improvements are promising, further improvements to faradaic efficiency stability are still needed to achieve an economically viable process for ERC to formate. The stability of ERC catalysts will be improved by changing the catalyst material, support material and surface structure. These experiments are performed with a focus on generating industrially viable processes in terms of activity and lifetime. 17 Chapter 2: Literature review 2.1 Voltage reversal in PEMFCs 2.1.1 Reversal tolerant anodes and oxygen evolution reaction catalyst development OER catalysts are added to Pt/C in the anode catalyst layer to protect carbon components of PEMFC anodes during transient fuel starvation conditions such as start-up/shut-down. [45] OER catalysts such as IrO2 or RuO2 promote the oxidation of water on the anode surface instead of the corrosion of carbon components and possibly Pt dissolution. [45,111] The durability of OER catalysts has been investigated using single cell setups focused mainly on electrolyser conditions. [112–116] However, the conditions in a PEMFC anode during normal operation (ca. 0 V and low pH) can induce instability in these metal oxides. IrO2 in particular can form soluble Ir3+ species [116] that migrate across the membrane electrode assembly and deposit on the cathode causing poisoning of the ORR catalyst (Pt/C). Typically, the suitability of OER catalysts for RTAs is investigated using prototype MEA single cell testing, which is time consuming and expensive to perform. [53,117] Therefore, effective ex situ methods are needed to provide OER catalytic activity and durability related information in a simulated PEMFC environment that reliably translates to fuel cell anode performance with respect to voltage reversal tolerance. Furthermore, in order to validate these new methods, it is important to have a series of catalysts with known properties and well-understood OER activity. 2.1.2 Oxygen evolution reaction catalyst activity and stability Transition metals have been studied for their OER activity in both the metallic [118–120] and oxide forms. [115,118,121,122] Metallic transition metals are generally oxidized at high potentials and this can lead to chemical instability in some cases. Monometallic oxides were investigated by Danilovic et al. who reported OER activity trends at 5 mA cm-2. They determined that the 18 overpotential (η) increases as: Os < Ru < Ir < Pt < Au. They also found that durability to dissolution in an OER environment increases with the nobility of the metals and decreases with the density of surface defects. [121] When choosing an OER catalyst for RTAs, the best compromise between increased activity and durability should be found. Osmium cannot be used because of the formation of OsO4, which is volatile and toxic. [123] This leaves RuO2 and IrO2 as the best possible catalyst candidates for RTAs. RuO2 has higher activity than IrO2 activity but IrO2 catalysts are more stable than RuO2 catalysts under electrolysis conditions in acidic media. One proposed mechanism of Ru/RuO2 dissolution (Figure 2.1a) was presented by Kötz et al. Ru destabilization takes place through volatile RuO4 in acid media. [124] Figure 2.1. (a) Theoretical model of Ru dissolution and O2 evolution on Ru/RuO2 electrodes through a RuO4 intermediate. (b) Model of Ir/IrOx speciation/dissolution and oxygen evolution [124,125]. Taniguchi et al. determined that, during reversal, Ru was eliminated from the cell particularly in the region around the fuel outlet. [44] A 30 wt% Pt-15 wt% Ru alloy catalyst supported on carbon black was used on the anode. There was reduced CO tolerance on the anode after cycling and the fuel starvation caused massive permanent damage to the Pt/Ru catalyst. 19 IrO2 is more stable than RuO2 in the high potential OER region. However, during normal fuel cell operation, IrO2 can form soluble Ir3+ species (Figure 2.1b). Ir3+ ions migrate to the cathode, thereby diminishing the reversal tolerance of the anode or negatively affecting the performance of the cathode side and thereby decreasing the performance and durability of the cell. [73] The Ir Pourbaix diagram shows thermodynamic stability of Ir species as a function of pH and potential (Figure 2.2). The normal operating conditions of fuel cell anodes is 0 V vs. RHE and 0 pH. For automotive applications when the PEMFC is not in use, air can diffuse in the anode raising the potential to 1.2 V vs. RHE. In this region (outlined in red), metallic Ir is the stable species. The green box outlines the cell reversal region (1.2 V to 1.9 V) where it would be expected for most of the carbon corrosion to take place. IrO2 is the stable species in the reversal region. Danilovic et al. confirmed this speciation by characterizing Ir and Ru thin films using X-ray absorption near edge structure (XANES) analysis at 0.05 and 1.45 V vs. RHE. [121] They found that at 0.05 V vs. RHE, both Ir and Ru had XANES spectra indicating the metallic phase. Whereas, at 1.45 V vs. RHE both metals had XANES spectra typical of their oxide forms. 20 Figure 2.2. Pourbaix diagram of Ir. Possible speciation of Ir at pH -2 to 0 in the potential of near reversal (Red square, 0 V to 1.2 V) and far reversal (green square, 1.2 V to 1.9 V) includes Ir(0), Ir(III), and IrO2[126]. Cherevko et al. described the dissolution of Ir and Ru and their oxides in thin films using in-line inductively coupled plasma mass spectrometry (ICP-MS). [118] There is an initial dissolution at 0.05 V followed by dissolution peaks at the beginning of the potential step to 1.2 V and at the OER ramp from 1.2 V to 1.4 V (Figure 2.3). They determined that the dissolution of Ir and the generated oxide were similar to the Ru response shown in the figure. Interestingly, there was no dissolution of a thermally generated Ir oxide layer at 0.05 V. This could be due to the difference between the 21 Ir4+/Ir3+ content of electrochemically generated IrOx vs. thermally generated IrOx. XPS was performed on thermally generated IrOx and did not show any Ir3+ in the film. XPS for the electrochemically generated IrOx was not shown and could contain Ir3+ which would explain this difference in dissolution. OER activity in acidic media was found to decrease as: Ru > Ir ≈ RuO2 > IrO2 and dissolution in acidic media was found to increase as: IrO2 ≪ RuO2 < Ir ≪ Ru. Figure 2.3. Dissolution of Ru with (grey) and without (red) oxide film on a Ru sputtered gold RDE electrode measured by in-line ICP-MS. [118] 22 Neither IrO2 nor RuO2 are ideal catalysts, especially in their bulk forms. Therefore, further modifications, including using nanoparticle catalysts, improved catalyst supports or mixed metal oxides should be investigated. 2.1.3 Oxygen evolution reaction kinetics Polarization curves to ca. 1.9 V vs. RHE are used for OER activity benchmarking and examples of this can be found across literature [121,127–129]. Important information can be determined from this analysis including the OER onset potential and Tafel slope. The onset potential is the potential at a defined current density where appreciable reaction occurs. The Tafel equation is an empirical relationship between surface overpotential (η) and faradaic current (i). It can be derived from the Erdely-Gruz-Volmer equation that describes the current based on the surface overpotential, exchange current density, transfer coefficient and temperature. 𝜂 = 𝑎 + 𝑏 𝑙𝑜𝑔 |𝑖| Equation 2.1 𝑎 = 𝑙𝑜𝑔 (𝑖0), 𝑏 =𝛼2.3𝑅𝑇 Equation 2.2 where 𝑖0 is the exchange current density, b is the Tafel slope and α is the charge transfer coefficient. In the case of a one electron transfer reaction with an α of 0.5 at 25 °C, b is 118 mV dec-1. Bockris showed that Tafel slopes change based on the current density on oxide catalysts. [130] The Tafel slope of IrO2 changed from ca. 30-40 mV dec-1 to 120 mV dec-1. [131] This change is again dependent on current density and there is an inflection point where the change begins at around 10 mA cm-2. This can be seen in Figure 2.4 where there are two distinct Tafel slope regions. The change in Tafel slope is due to the effect of the oxygen bubbles that form and detach from the surface of the oxide electrode. [129] 23 Figure 2.4. Tafel plots of an Ir electrode in acid solutions at pH 0.9 (circles), pH 1.9 (crosses) and pH 2.9 (triangles). [131] Impedance spectroscopy can be used to probe the electrochemical environment of a sample. In impedance spectroscopy, the cell is probed with a small magnitude alternating signal. This gives us a quasi-steady response that can be approximated by a theoretical circuit and averaged over time [132]. From this technique, Bode (log |Z| vs. log ω; Z is impedance and ω is frequency) or Nyquist plots can be generated and analyzed. It is also possible to determine the polarization resistance, which is a measure of the kinetics of the reaction. The double layer capacitance can be determined by fitting a line to the linear region of the Bode plot and extrapolating to the x-intercept. Analyzing the Nyquist plot uncovers interfacial behavior differences in catalysts (i.e. Warburg 24 impedance etc.) which gives information about adsorbed species on the surface of the electrode. [133] 2.1.4 Oxygen evolution reaction catalyst crystallinity and chemical form The activity and durability of IrO2 catalysts is dependent on the crystallinity vs. amorphous nature of the catalysts. IrO2 catalysts with more amorphous character have been shown to be more catalytically active for the OER, whereas more crystalline IrO2 catalysts are more stable to dissolution. [114,115,134] This could be caused by the increased Ir3+ presence in amorphous (non-heat treated) IrOx catalysts. Ir3+ has been suggested to be more active for the OER reaction due to more electrophilic O- and increased formation of O-O pairs. [135] Reier et al. investigated the activity and durability of IrO2 calcined on TiO2 substrates from Ir acetate at different calcination temperatures. [134] They determined that increasing calcination temperature caused a decrease in OER activity partially due to the decrease in electrochemically active surface area. da Silva et al. prepared IrO2 catalysts at different calcination temperatures 100 to 500 °C and found that crystallinity and durability to dissolution increased with increasing calcination temperature whereas OER activity decreased (as seen by lower OER onset potential). [115] The catalysts were investigated by identical location transmission electron microscopy (ILTEM) where IrO2 heat-treated at 500 °C showed no morphological changes after 500 aging cycles (1.1 to 1.6 V vs. RHE, 50 mV s-1). Other catalysts showed a decrease in OER activity during potential cycling that was attributed to loss of Ir through dissolution as well as loss of active surface area through agglomeration and particle detachment. It is interesting that whereas the amorphous IrO2 that did not undergo heat treatment showed higher initial OER activity loss after potential cycling, the final OER activity was still the highest of any IrO2 catalyst at the end of testing. Similar heat treatments 25 have been used to produce more durable but less active IrO2 and metal oxide supported IrO2 OER catalysts. [134,136–138] 2.1.5 Nanoparticle catalysts Ir, Ru and Pt carbon supported nanoparticles (2.0 nm, 4.3 nm and 2.5 nm, respectively) and bulk metals were investigated for their OER properties by Reier et al. Ir nanoparticles lost their metallic CV peaks (hydrogen adsorption peaks near 0 V) during voltage cycling and the OER activity remained similar to the bulk metal. [60] At the Ir surface, IrOx species will be formed during normal polarization above 1.2 V in acidic solutions. The Ru nanoparticles suffered from corrosion at high potentials (>1.2 V) making them unsuitable for long term OER, which is a confirmation of previous studies. Additionally, Pt nanoparticles showed low initial OER activity that decreased to zero over time (Figure 2.5). Figure 2.5. First OER scan for bulk and nanoparticle Ru, Ir and Pt at 1600 RPM and 6 mV/s in 0.1 M HClO4. The current is normalized to the number of active sites determined by CO stripping. [60] 26 Ralph et al. tested in situ extended reversal using a Ru nanoparticle RTA. [45] CO2 and O2 concentrations were monitored at the anode outlet. This analysis gives us a clear picture of the processes occurring at the anode. The OER protects the MEA from corrosion until eight minutes into the test. From this point on, carbon corrosion begins and the cell voltage drops below -1.8 V and the cell is destroyed after 13 minutes of reversal (Figure 2.6). Figure 2.6. Extended cell reversal test for a RuO2 containing RTA. The O2 (squares, dashed) and CO2 (triangles, dotted) current consumed is displayed along with the cell potential (diamonds, solid). [45] 2.1.6 Oxide supports for IrO2 catalysts On the cathode, PEMFCs can lose activity through platinum crystal size growth (Ostwald ripening) or Pt dissolution and migration. The dissolved platinum re-deposits and forms a band at the cathode/membrane interface. [139] Like Pt, IrO2 will dissolve during normal operation and re-deposit throughout the MEA. This process is similar to electro-dialysis systems for water 27 treatment. [140] Re-deposition can cause issues at the cathode including loss of activity and cell voltage. Additionally, membrane stability can suffer due to the re-deposition causing physical polymer breakdown. Mixed metal oxide (MMO) supports have been combined with OER catalysts to increase activity and durability via ion exchange properties of metal oxides with promising results. [138,141–149] The metal oxide supports are able to retain soluble metal species formed by the OER catalyst, preventing migration to the cathode side of the assembly. [150] Oh et al. determined that an antimony tin oxide (ATO) support suppresses Ir3+ dissolution via a metal/metal-oxide support interaction when compared to carbon supported IrO2. [141] They suggest that the ATO prevents the complete oxidation of Ir to IrO2 and Ir is stabilized at an average oxidation state of +3.2. They compared this to a carbon supported catalyst which displayed an IrO2 oxidation state of +4. Abbasi et al. investigated RuO2 supported on Ti/IrO2 + Ta2O5 for activity and lifetime using an accelerated stress test (AST) of 100 mA cm-2 in 1 M Na2SO4. They determined that the addition of RuO2 up to 30 wt% increased durability and the activity of the electrode. [151] There was no effect of the RuO2 loading on the OER onset potential. They concluded that the major mode of deactivation was due to dissolution of the MMO catalyst. Dissolution of MMO catalysts can produce redox couples that generate peroxides and damage the membrane of the MEA. [67] 2.1.7 Mixed IrO2 RuO2 catalysts Kötz et al. improved the stability of RuO2 catalysts by making admixtures of IrO2 and RuO2. The addition of just 20% IrO2 decreased the corrosion of RuO2 by 96% in the high potential, low pH OER conditions. [152] Ralph et al. determined that the most promising RTA was composed of two Ru based catalysts. First, a 40 wt% Pt-20 wt% Ru HOR catalyst and second, a RuO2-IrO2 (90:10 atomic ratio Ru:Ir) OER catalyst both supported on Shawinigan acetylene black. They were 28 able to achieve a cell voltage reversal time of 1630 minutes with a permanent cell voltage drop of only 27 mV. [45] 2.1.8 Reversal tolerant anode related patents Several patents have been filed and granted covering many aspects of RTA design. Ford and Daimler own patents relating to the addition of Ru and Ir species to the anode of PEMFCs. Zhang, and Haas et al. invented an anode with coverage of Ir or IrO2 supported on TiO2 catalyst only on certain areas of the anode to reduce the loading and cost. [153] Beattie et al. invented reversal tolerant anode catalysts with high catalyst loading on oxidation resistant supports such as graphitic carbon or Ti4O7. [154] Zhang et al. were granted a patent related to anode structure where they suggest separating the HOR and OER catalysts by an ion conducting membrane. This would make the stack thicker and also increase proton transfer overpotentials in a single cell. [155] Electrical protection has also been patented to protect the cell from damage by adding diodes to the cell. [51,156] Overall, the patent literature is focused on anode catalyst composition and loading. [55,157–160] It suggests using RuO2 or IrO2 on different types of metal oxide or carbon supports. [56,57] There are mentions of support structure (corrosion resistant carbons) but the details of the OER (crystal size etc.) are not explicitly covered in the patent literature. The missing piece from all of these patents is the stability of the OER catalyst itself in a PEMFC environment. There is space in the patent literature for proprietary OER catalyst treatments that improve stability and reduce dissolution. 2.1.9 EQCM measurement of rigid IrO2 layers Electrochemical changes of an electrodeposited IrO2 layer were investigated using EQCM by Birss et al. [161] They reported a frequency increase during the oxidation sweep at potentials greater 29 than 0.2 V vs. SCE accompanied by a frequency decrease on the reverse reduction scan virtually without any hysteresis of the frequency profile. After ruling out possible viscoelastic effects, the observed frequency changes were mainly attributed to the potential dependent uptake and ejection of water and ions in the hydrous Ir oxide film during the redox shuttle between Ir3+/4+ oxyhydroxides. This is also consistent with Burke and Scannell’s explanation of the hydrous Ir oxide cyclic voltammograms. [162] However, the formation of Ir oxyhydroxide was not confirmed by separate surface analytical techniques; only the presence of counter ions in the hydrous films were confirmed by XPS and ion chromatography. [161] 2.1.10 EQCM measurement of catalyst powders EQCM has been previously used to study electrochemical processes occurring on catalyst powder layers. [163–167] Ofstad et al. prepared a Pt/C ink containing Nafion® which was then sprayed onto EQCM quartz crystals. [166] They were able to identify increased Pt dissolution and lower onset potential of Pt oxide formation for Pt/C layers in the presence of chloride ions. Platinum dissolution or degradation from Pt/C is more significant than dissolution from electrodeposited Pt layers. Rice et al. calculated the mass of Pt oxide on Pt/C using single cycle asymmetric voltammetry coupled to EQCM. They assumed a perfectly rigid uniform layer in their calculations and reported very precise mass calculations. This work focused only on PEMFC cathode side catalysts. [165] 2.2 CO2 electroreduction 2.2.1 CO2 electroreduction thermodynamics The equilibrium cell voltage depends on the reactions taking place at the anode and cathode. 𝐸𝑒,𝑐𝑒𝑙𝑙∘ = 𝐸𝑐𝑎𝑡ℎ𝑜𝑑𝑒 − 𝐸𝑎𝑛𝑜𝑑𝑒 Equation 2.3 30 𝐸𝑒,𝑐𝑒𝑙𝑙∘ = 𝐸𝐶𝑂2∘ − 𝐸𝑂2∘ = −0.43 𝑉 − 1.23 𝑉 = −1.66 𝑉 Equation 2.4 These standard potentials represent an ideal thermodynamic scenario without the interference of HER in the reaction. In practice, there will be substantial overpotential and mixed potential due to HER that adds an extra energetic and faradaic penalty to the reaction. The choice of cathode and anode material as well as the reactor setup and physical dimensions can reduce these overpotentials and improve performance. Cell voltage is determined by the equilibrium kinetic overpotential (𝜂𝑘𝑖𝑛), mass transfer overpotential (𝜂𝑀𝑇) and the ohmic overpotential 𝜂𝑂ℎ𝑚𝑖𝑐. 𝐸𝑐𝑒𝑙𝑙 = 𝐸𝑒,𝑐𝑒𝑙𝑙∘ − (𝜂𝑘𝑖𝑛 + 𝜂𝑀𝑇 + 𝜂𝑂ℎ𝑚𝑖𝑐) Equation 2.5 The kinetic and mass transfer overpotentials will depend on the anode and cathode catalysts while the ohmic overpotential is dependent on the resistivity of the cell. The mass transfer overpotential (𝜂𝑀𝑇) can be broken down further into components: 𝜂𝑀𝑇 = 𝜂𝑀 + 𝜂𝐶 + 𝜂𝐷 Equation 2.6 where 𝜂𝑀 is migration overpotential, 𝜂𝐶 is convection overpotential and 𝜂𝐷 is diffusion overpotential. In ERC, the first reaction step has been found to be the main energy barrier and activation step. This reaction controls the 𝜂𝑘𝑖𝑛 for the cathode reaction and can be seen in to Equation 2.7. [168] E° = -1.9 V CO2 + e- → CO2∙- Equation 2.7 The potential barrier to form the CO2 radical is significantly higher than the standard potentials for any of the overall reactions. Finding catalysts that reduce the activation overpotential is essential for the energy efficiency of the process. 31 2.2.2 Possible chemical products Ideally, a reactor would reduce CO2 in a one-step process to ethanol, methanol or hydrocarbons such as propane, pentane or even octane. [169] These could be used again as solvents, in chemical synthesis or stored and used later as fuels. Long chain hydrocarbons are ideal because of their high energy density, ease of transport and long shelf life. The reduction of CO2 to long chain hydrocarbons is only possible by Fischer-Tropsch chemistry using CO and H2 at high temperature (150-300 °C). [170] Using ERC, only a few products are formed efficiently and at a high rate (Figure 2.7). The products are limited to short one to two carbon chains and these are difficult to produce as the primary product. Nonetheless, these short carbon chains and small reaction products can still be effectively stored and used as fuels in fuel cells, chemical feed stocks or used directly to power other processes. Figure 2.7. Box diagram of hypothetical ERC process. Products of ERC are dependent on reaction conditions but include formic acid (HCOOH), syngas (CO, H2), methanol (MeOH) and C2 to C4 alcohols and hydrocarbons. The process should be carried out using energy from renewable sources. 32 Half-cell reactions and standard reduction potentials in basic media for these reactions are shown below. Potentials are versus the standard hydrogen electrode (SHE) at pH 7 and 298K calculated from thermodynamic data. [168,171] E° = -0.43 V CO2 + H2O + 2e- → HCOO- + OH- Equation 2.8 E° = -0.52 V CO2 + H2O + 2e- → CO + 2OH- Equation 2.9 E° = -0.25 V CO2 + 6H2O + 8e- → CH4 + 8OH- Equation 2.10 E° = -0.34 V 2CO2 + 8H2O + 12e- → C2H6 + 12OH- Equation 2.11 E° = -0.38 V CO2 + 6H+ + 6e- → CH3OH + H2O Equation 2.12 These small molecule products of ERC are particularly useful as fuels. Carbon monoxide (CO) can be combined with H2 to form syngas which is a commonly fuel mixture. Formate (HCOO-) and methane can be used in fuel cells or burned. The overall cell voltage equation for ERC with formation of oxygen (O2) at the anode can be seen in Equation 2.13 and Equation 2.14. Cathode: E° = -0.43 V CO2 + H2O + 2e- → HCOO- + OH- Equation 2.13 Anode: E° = 1.23 V 1/2O2 + H2O + 2e- → 2OH- Equation 2.14 2.2.3 Faradaic efficiency of metallic catalysts Faradaic processes consist of a charge transfer across a metal liquid interface that causes a reduction or oxidation of a chemical species. Faraday’s law states that the amount of chemical reaction caused by a flow of current passed is determined by the moles of electrons passed across a metal solution interface. [132] Reactions that proceed according to Faraday’s law will form reactants as follows: A + e- → B Equation 2.15 33 where A undergoes a one electron reduction to form B. In this case, 1 mole of electrons will react with one mole of reactant (A) to form one mole of product (B). Faraday’s constant is the amount of charge carried by 1 mole of electrons, or 96,485 coulombs mol-1. Non-Faradaic processes are anything that occurs at the electrode surface that consumes charge but does not result in a charge transfer. These processes include the formation of an electric double layer. The rate of ERC is controlled by the current density passing through the electrode surface. It is possible to increase the rate of ERC by simply passing more current through the cell. However, a problem occurs because of the competing nature of ERC and the HER (Equation 2.16). Increasing the current density of an operating reactor to leads to increased rate of the parasitic HER. [103] Therefore it is necessary to balance the increase in the rate of reaction to the loss of faradaic efficiency. E° = 0 V H+ + e- → H2∙- Equation 2.16 It is possible to have more than one ERC reaction taking place. The total cathodic current density can be defined as: 𝑖𝑐 = 𝑖𝑐,1 + 𝑖𝑐,2 + 𝑖𝐻2 Equation 2.17 where 𝑖𝑐,1 is the current density of the desired ERC reaction, 𝑖𝑐,2 is the current density of other ERC reactions and iH2 is the current density of HER. The current or faradaic efficiency for ERC is therefore: 0% ≥ 𝜓𝐶𝑂2 =𝑖𝑐,1𝑖𝑐≤ 100% Equation 2.18 34 where 𝜓𝐶𝑂2 is the faradaic efficiency of ERC and 𝜓𝐶𝑂2 is a function of the catalyst, temperature, the potential, pH and the partial pressure of CO2. Many papers have been published on metal electrode catalysts for ERC. Post transition metals (Pb, Hg, Sn, In, Sn, Cd, Tl, and Bi) in particular have been studied as cathode side catalysts for room temperature ERC due to their high H2 overvoltage and specificity for producing HCOO-. [100–103,168,172–177] Hori et al. investigated various metal cathodes for ERC faradaic efficiency and selected results from this paper can be seen in Table 2.1. 35 Table 2.1. Faradaic efficiency for CO2 electroreduction on various metallic electrodes from Hori et al. [100] Current density and potential vs. the SHE included for each metallic electrode catalyst. Methane (CH4), ethane (C2H4), ethanol (EtOH), propanol (PrOH), carbon monoxide (CO), Formate (HCOO-) and hydrogen gas (H2) faradaic efficiency included. Total faradaic efficiency does not add up to 100% due to experimental error in quantification. Electrode material Potential vs. SHE (V) Current Density (mA cm-2) Faradaic Efficiency (%) CH4 C2H4 EtOH PrOH CO HCOO- H2 Total Pb -1.63 5.0 0.0 0.0 0.0 0.0 0.0 97.4 5.0 102.4 Hg -1.51 0.5 0.0 0.0 0.0 0.0 0.0 99.5 0.0 99.5 In -1.55 5.0 0.0 0.0 0.0 0.0 2.1 94.9 3.3 100.3 Sn -1.48 5.0 0.0 0.0 0.0 0.0 7.1 88.4 4.6 100.1 Bi -1.56 1.2 - - - - - 77 - - Cu -1.44 5.0 33.3 25.5 5.7 3.0 1.3 9.4 20.5 103.5 Au -1.14 5.0 0.0 0.0 0.0 0.0 87.1 0.7 10.2 98.0 Ag -1.37 5.0 0.0 0.0 0.0 0.0 81.5 0.8 12.4 94.6 Zn -1.54 5.0 0.0 0.0 0.0 0.0 79.4 6.1 9.9 95.4 Use of Cu as a cathode material especially interesting because it is the only monometallic catalyst capable of forming EtOH, CH4 and C2H4. [178] Cu electrodes are tunable based on the pH, buffer, and surface characteristics. [179] However, obtaining a single main product is important to reduce product separation related costs. In recent years, extensive research on tin (Sn) metal has been carried out because of its high faradaic efficiency for producing HCOO- from CO2. [101–103,107,172] Sn catalytic activity is 36 dependent on the metastable SnOx layer on the surface of the electrode. Kanan et al. found that in the region where ERC takes place, SnOx is only metastable and will reduce to metallic Sn over time. The newly formed Sn(0) surface does not have ERC activity but instead is more active for HER. [107] The morphology of the Sn catalyst has an effect on catalytic activity as well. [180] Nanostructured SnOx catalysts on graphene supports showed tunable characteristics based on the nanostructure size. [181] Lead (Pb) metal catalysts have also been identified as promising catalysts for ERC to HCOO-. [102,175] Hori et al. compared Sn and Pb and found HCOO- as the primary product with faradaic efficiency of 99% and 88% for Sn and Pb, respectively. These experiments were performed at a current density of 5.0 mA cm-2 in 0.1 M KHCO3 solution. [100,176] The Norskov group identified Pb and Ag as the most promising monometallic catalysts for HCOO- based on theoretical DFT calculations. [182] These results were based purely on the thermodynamics of CO2 adsorption and electroreduction. The theoretical predictions are difficult to extrapolate to industrially relevant electrocatalysts due to a multitude of complex interacting factors that can influence the performance. These factors include electrode polarization related effects, changes of the catalyst surface over time, competitive adsorption and support related effects. On the experimental side, the electroreduction mechanism on Pb was studied by in-situ FTIR at pH 8.6 revealing that the reaction takes place essentially as a surface hydrogenation involving two adsorbed species Pb-Had and Pb-CO2, ad (the latter formed by the reduction of HCO-3). [175] 2.2.4 Bi and tri-metallic catalysts Most CO2 electroreduction to formate studies investigate monometallic catalysts. [183–188] However, bimetallic catalysts have been extensively investigated for CO2 electroreduction to CO and other electrode reactions including borohydride oxidation [189,190], oxygen reduction, 37 [191,192] oxidation of alcohols and HCOOH [193,194] and hydrogen generation. [195] Bimetallic compositions can provide synergistic catalytic effects due to selective adsorption sites, electronic interactions and/or third-body (or ensemble) effects. [185,196] There have been reports of bimetallic alloy electrodes with special interest in Cu –Au electrodes. [184,186–188] Theoretical studies of bimetallic electrode surfaces were carried out by the group of Dr. Rossmeisl. [185] They proposed compounds containing transition metal elements and post transition metal elements in a single cathode to produce CO as a primary product. Each of these studies focused on bimetallic electrodes but there have been few studies on trimetallic electrodes. [197–199] These density field theory (DFT) computational based studies can cut down the work needed to characterize new classes of electrode materials. Decreasing the time and energy needed to find an industrially viable catalyst with synergistic effects for overpotential, current density and catalyst lifetime. Even with all of this research interest, there are still fundamental questions about the role of the alloying metals in the longevity of metal catalysts for ERC. On the experimental side, Choi et al. investigated CO2 electroreduction to HCOO- using Sn/Pb “alloy” materials found that a catalyst with 56.3% Sn and 43.7% Pb displayed they highest faradaic efficiency for HCOO- at 79.8% and the highest current density at 45.7 mA cm-2 of any of the studied alloys. The experiments were carried out over two hours in an H-Cell at -2.0 V vs. Ag/AgCl with 0.5 M KHCO3 as the electrolyte solution. [200] Whereas Sn/Pb compounds are commonly referred to as “alloys”, it has been shown by elemental mapping that Sn/Pb solid mixtures are not true alloys. Sn and Pb EDX signals are in fact separate from each other and form complex grain patterns. [201] Kumawat et al. investigated Sn-Pb-In tri-metallic catalysts for their ERC activity. [202] They found that the eutectic compositions had the best faradaic efficiency and performance. 38 They ranked their eutectic compositions according to their formate production rate: Pb29Bi46Sn25 > Pb25Sn75 > Pb44Bi56. 2.2.5 Support effects While these metal alloys can be used on their own as catalysts, it can be beneficial to make catalysts using support materials. The inert nature of carbon supports as well as their good electronic conductivity makes them strong candidates. Carbon supports have been used previously in direct methanol fuel oxidation [203] and lead/acid batteries. [204] To investigate the effect of carbon support structure on ERC, Pd was deposited on carbon substrates with different morphologies including carbon nanocoils, carbon nanofibers and ordered mesoporous materials with Vulcan as a reference. The rate of ERC was affected by the support conductivity, morphology, and size/shape of the deposited Pd. [205] High surface area carbon supported Sn catalysts have been investigated in aqueous conditions [172] to determine their effect on flow cell efficiency. It was found that gaseous CO and H2 bubbles formed on the surface of the electrode reduced the efficiency of the reaction. The support material could play a role in the catalyst activity beyond simply adding real surface area. The geometry of catalyst supports can change the interaction of the catalyst with the adsorbed chemicals either by influencing the orientation of the CO2 to the catalyst or by forming a cluster of atoms forming active sites in what is called the ensemble effect. [206–208] 2.3 Knowledge gap The knowledge gap in literature pertaining to PEMFC and ERC catalyst durability can be summarized as follows. 39  OER catalysts for RTAs are tested using time consuming and expensive in situ stress tests which limits the number of catalysts that can be investigated  An ex situ test for OERs has not been developed which reliably identifies promising OER candidates for further investigation  The mass changes of IrO2 powders have not been investigated during potential cycling  Heat treatment and its effect on EQCM response of IrO2 powders has not been investigated  Simulated RTAs using EQCM have not been described previously  Sn/Pb catalysts for ERC were reported previously but the mechanism of their improved faradaic efficiency durability was not described  Deposition conditions and catalyst substrate have not been investigated as variables for Sn/Pb catalysts In this thesis, these gaps in literature will be addressed while describing the methods by which the results were determined. 2.4 Thesis objectives and layout 2.4.1 IrO2 EQCM objectives of the thesis The objectives to investigate potential dependent mass changes of IrO2 during CV cycling are as follows:  Generate a novel ex-situ method to determine the suitability of OER catalysts in a PEMFC anode environment  Determine the effect of heat treatment on potential dependent mass changes of IrOx catalysts heat-treated at three temperatures (350, 450 and 550 °C)  Develop a method a better understanding of the EQCM response of IrOx catalyst powders 40  Generate an ex situ reversal tolerant anode capable of protecting Pt/C from carbon corrosion using IrOx  Determine the crystal phases and speciation of the IrOx after heat treatment and CV cycling by XRD and X-ray photo-electron spectroscopy (XPS) respectively  Use XRD and XPS data to explain differences in durability and activity of IrOx catalysts 2.4.2 Oxygen evolution reaction catalyst objectives of the thesis The objectives to investigate ex situ testing of OER catalysts are as follows:  Validate an ex situ accelerated stress test against an in situ anode accelerated stress test for OER evaluation  Investigate the effect of Pt/C catalysts on IrO2 durability in a simulated PEMFC environment 2.4.3 Electroreduction of CO2 goals of the thesis The objectives to improve ERC catalysts lifetime and efficiency are as follows:  Obtain repeatable electrodeposited Sn/Pb bimetallic catalysts at specified ratios using two different deposition bath compositions: tetrafluoroborate salts and metal oxides  Determine which catalyst has the best the faradaic efficiency stability over time  Determine the mechanism of superior performance for some catalyst compositions using XRD and SEM 2.4.4 Thesis layout This thesis will consist of two main parts. First catalysts for reversal tolerant anodes are discussed in Chapters 4 and 5 followed by electroreduction of CO2 to formate in Chapter 6. While Chapters 4 and 5 have the same overall goals, namely finding and describing better RTA catalysts, the goals are achieved in different ways. In Chapter 4, the focus is on fundamental electrochemistry and 41 chemistry to explain potential dependent mass changes occurring to heat treated IrOx catalysts. In Chapter 5, a more applied approach is used to compare in situ derived fuel cell data to ex situ results to compare different IrOx based reversal tolerant anode catalysts. 42 Chapter 3: Methods and materials 3.1 Surface and material analysis Surface analysis techniques were used to determine chemical and physical properties of catalysts before and after CV cycling, potential holds and stress tests. This allows us to make qualitative observations about changes that occur to the surface structure of the catalysts. The capabilities and limitations of these techniques are outlined below. 3.1.1 Scanning electron microscopy SEM is a surface sensitive technique capable of capturing images in the nanometer range. SEM requires that samples are either inherently conductive or non-conductive and coated with a thin layer of gold. All of the catalysts in this thesis were sufficiently conductive for SEM. The advantage of SEM is that it allows us to identify qualitative characteristics of the catalysts before and after applying the AST or electrolysis. Qualitative information such as cracking, surface roughness, catalyst size, and distribution are evidence of catalyst activity and breakdown mechanisms. SEM is limited however by the resolution of the detector and the magnification limiting the scale of the observable features to a few nm. SEM imaging in Chapter 5 was performed using a Philips XL30ESEM with backscatter electron (BSE) detector on freshly deposited GDL/MPL and stress tested GDL/MPL from the ex situ AST and MEA from the in situ AAST. SEM images in Chapter 4 were generated using a TESCAN VEGA3. IrOx layers drop cast on EQCM crystals were imaged both before and after the cycling protocol was applied. In Chapter 6, Sn/Pb catalysts were imaged before and after catalysis using a Hitachi S-4700 field emission scanning electron microscope (FE-SEM). 43 3.1.2 Energy dispersive X-ray Energy dispersive X-ray (EDX) analysis is a technique for elemental identification and quantification. Coupled with SEM, EDX allows for spatial resolution of catalyst coverage before and after AST or electrolysis. EDX does suffer from lower quantification accuracy as compared to XPS and therefore is better utilized as a tool to determine the presence or absence of a metal. In addition, as with many x-ray based techniques, the identification of small atoms such as carbon, oxygen, nitrogen and hydrogen is very limited. Finally, EDX cannot be used to accurately determine the stoichiometry of mineral deposits. In Chapter 5, EDX measurements were carried out using a Philips XL30ESEM with attached EDX detector. In Chapter 6, catalysts were analyzed using a Hitachi S-2600N variable pressure scanning electron microscope (VP-SEM) with an X-ray detector with full spectrum EDX to determine their elemental makeup at three different points on each catalyst. The Sn/Pb ratios are averaged over these three spots and the error presented is the standard error of the mean. 3.1.3 X-ray diffraction XRD produces X-ray diffraction patterns of a material which can be compared to a library of known materials to determine crystal phases present in the sample. A given catalyst may have several separate crystal phases present in a single sample and XRD is powerful enough to deconvoluting complex crystal structures. Once the phases are identified, it is possible to generate a picture of the different crystal phases present. If a sample is not crystalline, it will not generate distinct peaks. These broad peaks that are not useful for identification of crystal phases. Similar to SEM and EDX experiments, XRD can be carried before and after the AST or electrolysis to identify changes happening in the catalyst. 44 XRD analysis is limited to catalysts with high loading (> 10 mg). While this is not a problem with the powder catalysts and high loading Sn/Pb catalysts, the limitation comes from the protocols for the mass of catalyst that is able to be deposited on the MEA or EQCM crystal. It was not feasible to carry out XRD analysis on catalysts after AST using the current protocol. Higher loading would need to be used in order to obtain an interpretable XRD signal from the catalyst layer. This would not be possible on the EQCM because the crystal becomes overloaded and the frequency response is disrupted. This loading limitation was not an issue for the fluoroborate bath ERC catalysts because the loading of Pb and Sn was high enough to make XRD measurements. In Chapter 4, 5 and 6, IrOx crystal structures were analyzed as received and after heat treatment by powder XRD using a Bruker axs-D8 Advance with a Cu radioactive x-ray source along with a LynxEye silicon strip detector. XRD measurements were taken at 40 kV and 40 mA. Tetrafluoroborate Sn/Pb crystal structures were analyzed as before and after ERC by XRD. Crystallite sizes were calculated using the Scherrer equation shown below. 𝜏 =𝐾𝜆𝛽𝑐𝑜𝑠𝜃 Equation 3.1 Where τ is the crystallite size, K is the shape factor (usually close to 0.9), λ is the X-ray wavelength, β is the line broadening at half maximum intensity and θ is the Bragg angle. 3.1.4 Brunauer-Emmett-Teller surface area analysis In this technique, the surface of a catalyst is covered with N2. The pressure of a vessel containing a known mass of catalyst is closely measured and from this, it is possible to determine the volume of adsorbed N2. Using the effective cross-sectional area of N2 (0.162 nm2), the surface area of the catalyst is calculated based on the assumption that a monolayer of N2 was formed on the surface. 45 This technique determines the total gas accessible surface area of a catalyst. The calculated surface area is greater than either the electroactive surface area or the OER active surface area. In Chapters 4 and 5, a NOVA 2000e surface area and pore size analyzer (Quantachrome Instruments) was used to determine the surface area of the catalyst powders. The Brunauer-Emmett-Teller (BET) theory was used to calculate the total surface area of each catalyst prior to potential cycling. 3.1.5 Scanning transmission electron microscopy Scanning transmission electron microscopy (STEM) is an imaging technique where a beam of electrons is passed through a thin sample. The beam of electrons is scanned across the surface of a material to generate an image. The image can be acquired at higher magnification than SEM. STEM can be coupled with EDX obtain spatially resolved elemental information. Additionally, the crystal structure of the sample can be investigated using lattice fringe technique. Using this structural information, the chemical environment of a catalyst can be estimated. To use TEM effectively and obtain clear images, it is necessary to start with a very clean sample. Any organic impurities will blur the obtained images making fine details difficult to identify. This also makes lattice fringe analysis and other small-scale measurements impossible. STEM in Chapter 4 was carried out using an FEI Osiris with a high-angle annular dark field (HAADF) STEM detector and a bright field (BF) detector. 3.1.6 X-ray photoelectron spectroscopy X-ray photoelectron spectroscopy is a technique that identifies the speciation of heavy elements (atomic number > 6) in a catalyst sample. XPS allows us to explore the changes that occur after stress testing in terms of the redox processes occurring on the electrode surface. Unlike XRD, XPS is highly sensitive to trace elements. Ideally, this would be in line with the stress test itself to 46 produce potential resolved speciation information and reduce the possibility of contamination or oxidation after the AST. The XPS spectra is generated in a similar way to SEM-EDX spectra except that in this case the radiation used comes from an X-ray source instead of a beam of electrons. The penetration depth of the X-ray beam decreases with increasing sample density. The decrease in photoelectron yield with increasing depth makes XPS is a surface sensitive technique. Interpretation of XPS spectra can be difficult even for someone experienced in interpreting XPS results. It is not always possible to determine the exact speciation of a sample. This is because the interpretation is based on previous spectra of known pure species (similar to XRD spectra). However, the complication occurs when there are multiple possible species in the same sample. In this case, it is only possible to determine the general chemical environments of the atoms in the sample. For example, it is possible to identify oxygen bound to two hydrogens or one hydrogen but not the ratio of OH to H2O in a given sample. This problem prevents us from making definitive statements about the speciation of IrOx catalysts after the CV stress tests in Chapter 4. XPS in Chapter 4 was performed using a Kratos Analytical Axis Ultra DLD. 3.1.7 Inductively couple plasma mass spectrometry The most sensitive technique for determining Ir metal content (and the content of other metals) in solution is by ICP-MS. An aerosolized sample is injected into an argon plasma. Any material will be ionized and sent onto the MS detector where it is separated based on their mass to charge ratio (m/z). The limit of detection is dependent on the MS detector (quadrupole, time of flight etc.) and the ion being measured. The preparation of the sample is important to obtain accurate concentrations. ICP-MS can identify only dissolved atoms. Additionally, identification by m/z can produce inaccurate results when the 47 material does not dissociate to single atoms in the argon plasma. This can leave dimers of atoms which show can be misidentified as a heavier species. Electrolytes were collected and analyzed for metal content to determine the catalyst dissolution. The Ir detection limits for several analytical methods are presented in Table 3.1. [36] In Chapter 5, ICP-MS analysis was carried out using standards with a commercial ICP-MS service (Exova, Surrey, BC, Canada). Table 3.1. Sensitivity of analytical techniques to determine Ir3+ concentration (ng mL-1) in 1 M H2SO4. [36] Wavelength Flame emission Flame atomic absorption Electrothermal atomic absorption Argon ICP Ir Detection limit / ng mL-1 208.88 212.68 224.27 400 500 0.5 20 20 3.2 OER experimental goals and implementation In situ testing is both time consuming and expensive to perform. There is a need in industry to improve early screening protocols to identify promising catalyst candidates. A new strategy has been developed to benchmark OER catalysts for use on PEMFC anodes. Harsh conditions exist at the fuel cell anode (low pH and reducing conditions) for OER catalysts. These catalysts are normally designed for use only in stable oxidizing conditions on the anode of water electrolyzers where maximum activity is the goal. New classes of OER catalysts will have to be specially developed to remain stable and active at the PEMFC anode. The goal of this work is to develop an ex situ test to determine the stability of different monometallic and MMO catalysts. The tests performed must account for the complex interactions 48 in a PEMFC anode. Various techniques were used to elucidate the relationship between catalyst structure and in situ activity and durability. While the support plays an important role in the activity and durability of an OER catalyst, the oxidation state and crystal structure of the IrO2 catalyst itself could be of equal importance. 3.3 Iridium oxide heat treatment and surface area analysis In Chapter 4, IrO2 (Alfa Aesar, 99%) was heat-treated at three temperatures (350, 450 and 550 °C) in air using a tube furnace (Barnstead Thermolyne 21100). First the catalyst was placed in an alumina crucible. The powder was confirmed not to have any large chunks before the crucible was placed into the tube furnace. Air was introduced to the tube using a custom gas manifold and the temperature was increased to the set-point. The heat was maintained for two hours and then allowed to decrease to 50 °C before removing the crucible. Table 3.2 presents the BET surface area of each heat-treated catalyst. Table 3.2. BET Surface Area of IrOx Untreated and Heat-treated at 350, 450 and 550 °C. IrOx Surface area m2 g-1 Untreated 28.2 350 °C 27.4 450 °C 22.5 550 °C 19.0 3.3.1 Ink preparation The IrOx ink was prepared by sonicating (Branson 3510 ultrasonic) 4 mg of IrOx powder in 1.2 mL 2-propanol (Alfa Aesar, 99.5%) for 1.5 hours until a well dispersed dark black ink was formed. 10 µL of this solution was drop cast onto the EQCM crystal (Stanford Research Systems, 5 MHz 49 AT-cut quartz, 1.37 cm2 electroactive area, 0.4 cm2 piezo-active area, 331 μm crystal thickness, gold coated) and allowed to dry. The target crystal loading was 24 μg cm-2 electroactive area. The goal of this technique was to obtain a uniform and repeatable ink deposition. Platinum on carbon (Pt/C, Tanaka TEC10E50E, Pt 50 wt% on “high surface area carbon”) was also admixed with two types of IrOx, untreated and heat-treated at 550 °C. 3 mg of the Pt/C and 1 mg of the IrOx was sonicated in 1.2 mL of 2-propanol for 1.5 hrs. The ratio of Pt metal to Ir metal was 2:1 in the ink. 3 μL of the resulting ink was drop cast onto the EQCM quartz crystal. This resulted in a total solids loading of 7.3 μg cm-2. Pt/C alone was prepared in the same way but without the addition of IrOx which resulted in a total solids loading of 5.5 μg cm-2. Ketjen Black (EC600JD, 1270 m2 g-1, 34.0 nm primary particle radius) was also prepared into an ink in the same way. In this case, 1.5 mg of the powder was dissolved in 1.2 mL of 2-propanol. This was to maintain the same carbon loading as in the 50 wt% Pt/C whereas the total solids loading was 2.7 μg cm-2. 3.3.2 Quartz crystal microbalance measurements The ink drying QCM measurements and EQCM measurements were made using a QCM 200 system (Stanford Research Systems) connected to the integrated analog to digital converter of a Metrohm Autolab PGSTAT302N and recorded using Metrohm NOVA software. EQCM measurements were carried out in a three electrode cell configuration consisting of a platinum counter electrode, reversible hydrogen electrode (RHE) and EQCM crystal secured in a Teflon holder as the working electrode. All experiments were carried out using N2 purged 0.09 M H2SO4 electrolyte prepared with double deionized water (Thermo Barnstead GenPure, 18.2 MΩ cm-1). The electrolyte composition was chosen to mimic the pH (~ 1) of the anode catalyst layer in a PEMFC. [73] 50 3.3.3 Characterization and accelerated potential stress testing Each ink composition was investigated using the following stress test. Thirty CV cycles were carried out at each of three potential windows (1.2, 1.4 and 1.6 V to 0.05 vs. RHE) in 0.09 M H2SO4 at 25 °C. The concentration of 0.09 M H2SO4 was chosen because at concentration the pH is 1. The potential was swept from high to low potential to measure frequency changes taking place from near open circuit voltage (OCV) (ca. 0.9 V vs. RHE) to 0.05 V vs. RHE. The following potential windows were chosen to approximate three PEMFC anode conditions: a gas switching stress test (0.05 to 1.2 V vs. RHE), “low” voltage reversal conditions (0.05 to 1.4 V vs. RHE) and “high” reversal conditions (0.05 to 1.6 V vs. RHE). Potential holds at 1.8 and 2.0 V vs. RHE were carried out to investigate high reversal conditions. However, the current at these potentials caused damage to the EQCM crystal and are not presented here. 3.3.4 Cu deposition EQCM calibration Copper deposition was carried out using cyclic voltammetry from 0.5 V to 0.0 V and the results are presented in Figure 3.1. The total area under the Cu deposition and dissolution curves in (peaks at ca. 0.24 and 0.33 V vs. RHE, respectively) appear unequal. However, the total charge consumed during deposition and dissolution is very close (Figure 3.2). This is because deposition continues past the switching potential of 0 V vs. RHE and the current continues to be negative until dissolution begins at ca. 0.27 V vs. RHE. 51 Figure 3.1 Current (black curve) and frequency (red curve) profiles of Cu deposition and stripping on an Au plated EQCM crystal at a sweep rate of 20 mV s-1. Electrolyte: 0.09 M H2SO4 containing 1 mM CuSO4 at 25 °CRT. The Sauerbrey equation can be used to estimate mass changes on an EQCM using Equation 3.2. [209] ∆𝑓 = 𝐶𝑓∆𝑚 Equation 3.2 where: Δf - observed frequency change (Hz), Δm - change in mass per unit area (g cm-2), Cf - sensitivity factor (or Sauerbrey constant) for the crystal (the theoretical value is 56.6 Hz μg-1 cm2 for a 5 MHz AT-cut quartz crystal at room temperature). 0.0 0.1 0.2 0.3 0.4 0.5-0.6-0.30.00.30.60.91.2f / HzI / mAE / V vs. RHE-40-30-20-100 52 In order to use this equation, the value of Cf must be calculated. ∆𝑓 =106𝑀𝑤𝐶𝑓𝑄𝑛𝐹𝐴 Equation 3.3 where: Mw - apparent molar mass of the depositing species (63.546 g mol-1), Q - integrated charge during the reduction (C), A - active deposition area of the working (liquid contact) electrode (1.37 cm2) F - Faraday’s constant (9.648 x 104 C mol-1), n - number of electrons transferred (i.e. n = 2 for Cu deposition from CuSO4). (The factor of 106 converts from μg in Cf to g in Mw) This equation can be rearranged to calculate the Cf from the slope of a Δf vs Q graph. 𝐶𝑓 =𝑛𝐹𝐴106𝑀𝑤∆∆𝑓∆𝑄 Equation 3.4 53 Figure 3.2 Sensitivity factor graph correlating the change in frequency (Δf) with the reduction charge (Q). Using the slope from Figure 3.2 and Equation 3.4, the following sensitivity factor is obtained. Cf = 47.8 Hz cm2 μg-1 1/Cf = 20.9 ng Hz-1 cm-2 The theoretical sensitivity factor of a 5 MHz AT cut crystal is given below: [210] Cf = 56.6 Hz cm2 μg−1 1/Cf = 17.7 ng Hz-1 cm-2 ref. [210] To compare these results to Vatankhah et al.[211] and Ignaszak et al.,[212] the results were converted from a 5 MHz crystal to a 9MHz crystal using the Sauerbrey equation where the sensitivity factor is expressed as follows: [209] -0.004 -0.003 -0.002 -0.001 0.000-50-40-30-20-100f / HzQ / CAdj. R-Square 0.99759 Value Standard ErrorHz Intercept 0.26595 0.10522Hz/C Slope 11487.99196 39.5727 54 ∆𝑓 = −2𝑓02𝐴√𝜌𝑞𝜇𝑞∆𝑚 Equation 3.5 where: f0 – fundamental frequency (Hz), ρq – Density of quartz (2.648 g cm-3), μq – Shear modulus of quartz (2.947x1011 g cm-1 s-2), The Sauerbrey equation shown above can be rearranged to produce Equation 3.6. 𝐶𝑓,9 𝑀𝐻𝑧 = 𝐶𝑓,5 𝑀𝐻𝑧𝑓0,9 𝑀𝐻𝑧2𝑓0,5 𝑀𝐻𝑧2 Equation 3.6 Using Equation 3.6, The 5 GHz crystal Cf was converted to a 9 GHz crystal to be 154.8 Hz cm2 μg-1. This result can be compared to what was found in Ignaszak et al. [212] They determined that the Sauerbrey constant is 221.7 Hz cm2 μg−1; which is higher than the theoretically value of 183 Hz cm2 μg−1. [212,213] For this reason Ignaszak et al. concluded that multi-layer Cu deposition occurred. Vatankhah et al. reported that the sensitivity factor for a 9 MHz crystal varied from 166.7 Hz cm2 μg−1 for monolayer and 238.1 Hz cm2 μg−1 for more than seven layers. [211] For calibration, the sensitivity factor was lower than the literature value, suggesting that there was a sub-monolayer of Cu deposited on the surface of the gold electrode. To obtain the true sensitivity factor relevant for the powder catalyst films, multiple deposition and calibration experiments should be carried out for each tested film type on the EQCM. Instead, it is simpler and more 55 effective to use the theoretical sensitivity factor of 56.6 Hz cm2 μg−1 for calculating mass change estimates for our films and inks. It should be understood that these estimates are affected by a multitude of factors that must be accounted for including viscosity changes close to the electrode, non-rigid film formation, adsorption of ions etc. 3.4 Oxygen evolution ex situ characterization methods Several methods were used to investigate the durability and activity of OER catalysts in Chapter 5. Typical voltammetry methods were used along with surface characterization (BET, SEM, etc.) and material characterization (XRD, XPS and EDX). Analytical methods including ICP-MS and EQCM were used to measure mass and material changes to the OER catalysts. 3.4.1 Types of oxygen evolution reaction catalysts Three unsupported and one metal oxide (TiO2) supported IrOx catalysts were investigated in this work. The supported catalyst (IrOx/TiO2) was chosen in order to evaluate the effect of the TiO2 support (i.e., ion exchange capacity and oxygen vacancy formation during potential cycling) on the activity and stability of IrOx. The unsupported catalysts were obtained as either high surface area (IrOx High SA) or low surface area (IrOx Low SA) catalysts (Table 1). Additional metal oxide supported IrOx catalysts were investigated. These results are included in Appendix C.2. Table 3.3 presents the BET surface area and IrOx content of the four tested catalyst materials. Previous studies show a decrease in BET surface area with heat treatment: from 141 m2 g-1 (untreated), to 75 m2 g-1 (treatment at 300 °C in air) and 55 m2 g-1 (treatment at 400 °C in air) [115]. Information about the catalyst formulation and production as well as the identity of the suppliers has been restricted for publication. 56 Table 3.3 BET surface areas and IrOx contents of three unsupported and one TiO2 supported catalyst in Chapter 5. IrOx Low SA IrOx High SA IrOx High SA HT (350 °C) IrOx/ TiO2 BET Surface area (m2 g-1) 47.7 91.6 47.1 28 IrOx content (wt. %) 100 100 100 75.6 3.4.2 IrOx heat treatment Several studies have used heat treatment to improve dissolution stability of IrOx catalysts. [114,115,134] The exact temperature has varied but, in general, at temperatures of ca. 350 °C, most studies began to see differences between the heat-treated and as prepared materials. The temperature in Chapter 5 was chosen to maintain good catalytic activity while at the same time improving durability to dissolution. The choice of Oxygen as the gas used in the tube furnace was to maintain an oxidative environment and remove any contaminants that could be introduced by using ambient air. The IrOx High SA was heat-treated at 350 °C to increase the crystallinity of the catalyst (denoted as IrOx High SA HT). First, 0.5 g of catalyst was weighed out into a 2 cm x 6 cm alumina crucible. Any large catalyst chunks were crushed using a mortar and pestle. The crucible was placed into a ThermoScientific Lindberg Blue M three-zone tube furnace. Oxygen was introduced through a custom made gas mixing unit. A heat treatment program was used to increase the temperature at ramp rate of 5°C min-1 to 350°C and held for two hours. After allowing to cool for at least 2 hours, 57 the crucible was removed from the furnace and the catalyst was weighed again. The prepared catalyst was then used to make the electrodes. 3.4.3 Electrode preparation The catalysts inks in Chapter 5 were prepared by bath sonicating (Branson 3510 ultrasonic) 4 mg of catalyst in 3.6 mL 2-propanol (99.5%, Alfa Aesar) to homogenize and disperse the catalyst. A combined GDL/MPL (5 wt % Teflon coating, Fuel Cell Store, Freudenberg H2315) was cut out with a 2.5 cm diameter metal punch. Then, the homogenized catalyst mixture was deposited dropwise on the MPL side of the GDL/MPL and dried using a heat lamp. A target of 0.4 mg catalyst was deposited on each GDL/MPL for a total of 81.5 μg cm-2 loading. The dried electrode was finally rinsed with deionized water and placed into the electrode holder (Figure 3.3) and the three-electrode cell setup. 58 Figure 3.3 Diagram of the working electrode setup. Teflon based body and screw cap holding the GDL/MPL. Electrical contact was provided by a graphite disk and Pt wire Only the front side of the GDL/MPL was exposed to the electrolyte and the electrical contact was made using a platinum wire on the GDL side of the GDL/MPL. The holder contained a 2 cm diameter O-ring which, along with a screw cap, to hold the 2.5 cm diameter GDL/MPL disks. This cell setup was similar to one used previously by Pinaud et al. [214] but without the flow channels on the GDL side of the layer. 3.4.4 Electrochemical characterization A water jacketed three electrode cell was used consisting of a platinum counter electrode, RHE reference electrode and a GDL/MPL working electrode in a Teflon cell holder in 200 mL N2 purged 0.09 M H2SO4 electrolyte prepared with double deionized water (Thermo Barnstead 59 GenPure, 18.2 MΩ cm-1). The temperature was controlled using a thermostated water bath (Fisher Isotemp 3016HS) filled with distilled water. An inert environment was maintained by bubbling N2 through the solution for the duration of the experiment. Evaporation was minimized by sealing the top of the three-electrode cell with a Teflon lid and lab film (Parafilm, “M”). Electrochemical experiments (CV, linear sweep voltammetry (LSV) and potential stepping) were controlled by a Princeton Applied Research 263A potentiostat controlled by Corrware software (Scribner Associates Inc.). 3.4.5 Ex situ stress testing and ICP-MS Initially, CV cycling (100 cycles from 0.05 V to 0.8 V vs RHE at 500 mV s-1) was performed to condition the electrode. A potential stepping procedure was then employed to stress and dissolve the OER catalysts. Potential stepping was used with one second holds at 0.05 V and 1.2 V for 30,000 step cycles. After every 5,000 step cycles, four CVs from 1.3 V to 1.9 V vs. RHE (two at 20 mV s-1 and two at 5 mV s-1) were taken to determine the OER activity of the electrode. Additionally, CVs were recorded after the conditioning and again after every 5,000 stepping cycles from 0.05 V to 1.4 V vs. RHE at 100 mV s-1 scan rate and the same temperature as the AST. This was to monitor the change in capacitive current at intervals during the AST. The final electrolyte volume was measured using a graduated cylinder and placed into Nalgene containers for dissolved metals determination. Typical electrolyte volume loss due to evaporation was on the order of 5-10%. ICP-MS (Exova, Surrey) was used to determine the dissolved metals concentration of the electrolyte. 3.4.6 Fuel cell anode accelerated stress test The fuel cell voltage reversal experiments were performed using a MEA in a single-cell configuration. A Baltic™ quick connect liquid cooled cell was employed with custom made plates 60 with a straight flow field and an active area of 20 cm2. The anode Pt/Ir(metal basis) loading was 50/20 μg cm-2 with an ionomer to carbon weight ratio of 0.7. Cathode Pt loading was 250 μg cm-2. Cell temperature and humidity were controlled via a test program with the corresponding thermostats and gas humidifiers (Scribner Associates Inc.). Fuel cell performance tests were carried out in ‘differential mode’ with a stoichiometry of 6 at the max current (Appendix Figure C.1.5). Using a fresh MEA, the anode reversal time (minutes) was measured at 30 °C by holding the cell at a constant current of 0.2 A cm-2 with air flowing on the cathode (100% RH, 2 LSTD min-1 (SLPM), 20.9% mole fraction O2, 2.5 bar absolute pressure) and N2 on the anode (100% RH, 1 SLPM, 2.7 bar(a)) until the cell voltage reached -1.85 V. This is referred to as the beginning of test (BOT) reversal time. Due to the damaging nature of the voltage reversal, a new MEA was used for the AAST. Gas switching cycles were carried out consisting of 2 minutes of normal operation at 0.5 V with H2 at the anode (100% RH, 1 SLPM, 2.7 bar(a), H2 stoichiometry: 2.4) and air at the cathode. This was followed by 2 minutes of air on both the cathode and the anode at OCV. A N2 purge of the anode (to cell voltage < 0.1 V) was used between each gas switch to prevent mixing of H2 and air. The AAST was performed at 30 °C to reduce carbon corrosion and induce minimal stress on the cathode. This is necessary in order to attribute all the measured changes to the anode degradation. To confirm this hypothesis, cathode electrochemically active surface area (ECSA) measurements were carried out at the BOT and end of test (EOT) (Table S3, Supplemental Information). The anode reversal time measurement described in the previous section was repeated at the end of 300 gas switching cycles (i.e., after the AAST), and is referred to as the EOT reversal time. The reversal 61 loss is defined as the difference between the BOT and EOT reversal times. The standard error of these experiments is approximately 7% based on previous replicates. 3.5 Electroreduction of CO2 methodology The preparation of Sn-Pb bimetallic catalysts was carried out using electrodeposition techniques at constant current to control the total deposited metal on the carbon substrates. The bath compositions were chosen to investigate a range of Sn-Pb compositions with the understanding that the bath compositions would not necessarily translate to the final deposited metal composition. In order to determine the final composition, it was first attempted to dissolve both the Pb and Sn at the same time and measure the concentrations by ICP-MS. However, the two metals do not readily dissolve in the same acids or acid combinations. Another approach could have been to electro-oxidize the deposited metals under constant potential conditions. In this way, it would be possible either to detect a change in potential that indicated that all metal had been dissolved or to take a sample of the electrolyte for ICP-MS analysis. In this way, the composition of the bulk material could be determined. The chosen method for Sn:Pb ratio determination was EDX which determines composition has an interaction volume on the order of μm [215] and is sufficient for determining Sn:Pb ratio in this study. 3.5.1 Pb-Sn electrodeposition on two types of fibrous carbon substrates Pb-Sn electrodeposition was carried out on two fibrous carbon substrates: un-teflonated carbon cloth (Fuel Cell Earth CCP40, 99% carbon content, thickness: 0.38 mm, thread count: 18 x 18 yarn cm-1, bulk density: 0.75 g cm-3) and teflonated carbon paper (Toray carbon paper 060, 5 %wt. Teflon content, thickness: 0.19 mm, bulk density 0.44 g cm-3), respectively. The substrates were sonicated at 25 °C in ethanol for 20 minutes followed by oven drying at 60 °C in air. The carbon 62 substrates were masked with Kapton polyimide film (Matrix Technologies) to expose a 1 cm2 geometric electrode area for deposition. Two electrodeposition bath compositions were used that differed in the Pb and Sn precursor compounds employed. The first bath composition contained different concentrations of tin(II) tetrafluoroborate (Sn(BF4)2) (Pure Chemistry Scientific, 48 %wt., aq) and lead(II) tetrafluoroborate (Pb(BF4)2) (Sigma, 50 %wt., aq) as shown in Table 3.4. Additionally, the 100 mL deposition baths contained 0.44 M boric acid (HBO3), tetrafluoroboric acid (HBF4), and either bovine gelatin or peptone (Sigma). [216,217] The high Pb concentration baths (100% Pb and Pb(84%)-Sn(16%)) were made using a patented method developed for lead acid battery electrodes. [216] Electroplating high Sn concentration baths (100% Sn, Sn(91%)-Pb(9%), Sn(75%)-Pb(25%)) under the same conditions produced electrodes that were mechanically unstable and another recipe had to be used. [217] In this procedure, the HBF4 concentration was increased to 2.46 mol L-1 and peptone was substituted for gelatin to improve catalyst layer stability as well. After adding the other components to the bath, 5% Triton X-102 by volume (Sigma) was dissolved at room temperature. The critical micelle concentration of Triton X-102 in water is 0.33 mM and the HLB number (hydrophile-lipophile balance) is 14.2. [218] The high HLB number indicates very good water dispersion and surface wetting properties, both characteristics are important for electrodeposition. The electrodeposition was carried out at 25 °C at a constant current density of 50 mA cm-2 for 1 hour with a total of 180 C of charge passing through the cell. A Parstat 2243 potentiostat was used along with Electrochemical Power Suite (Princeton Applied Research) to record and collect the electrochemical data. The electrodeposition cell was set up with an Ag/AgCl, KCl std. reference electrode along with Sn and Pb counter electrodes. The Sn and Pb counter electrodes 63 were used to maintain the metal ion concentration of the deposition bath. The deposition was followed by two five minute sonications in isopropyl alcohol and two washes in double distilled water. Afterwards, the catalysts were electrochemically conditioned by cycling from 0.1 V to -2.0 V vs. Ag/AgCl in N2 purged KHCO3 until a stable cyclic voltammogram was observed. Table 3.4. Electrodeposition bath compositions using Sn and Pb tetrafluoroborate salts. Nomenclature Sn(BF4)2 conc. [mol L-1] Pb(BF4)2 conc. [mol L-1] Peptone [g L-1] Gelatin [g L-1] HBF4 conc. [mol L-1] Sn100% 0.400 N/A 3 N/A 2.46 Sn91%-Pb9% 0.429 0.042 3 N/A 2.46 Sn75%- Pb25% 0.315 0.106 3 N/A 2.46 Sn16%- Pb84% 0.212 1.08 N/A 1 0.328 Pb100% N/A 1.06 N/A 3 0.738 In addition to the tetrafluoroborate salt based electrodeposition bath presented in Table 4, a second set of catalyst depositions were performed using metal oxide based precursors. The plating bath contained 1 M HBF4 with different concentrations of PbO (Fisher, 99.9%) and SnO (Sigma, 97%) as shown by Table 3.4. The total concentration of SnO and PbO was 10 mM in each case. The depositions were carried out galvanostatically, at 16 mA cm-2 until 3.6 C of charge had passed through the cell (225 s). Thus, in this case lower loadings were generated such that the results to be directly comparable with those presented by Choi et al. [200] A conventional three electrode cell was used with a Ag/AgCl reference electrode and a platinum wire counter electrode. The rest of the electrodeposition conditions were the same as presented before. 64 Table 3.5. Electrodeposition bath compositions using SnO and PbO. Nomenclature SnO conc. [mM] PbO conc. [mM] Sn100% 10 N/A Sn80%-Pb20% 8 2 Sn70%-Pb30% 6.7 3.3 Sn50%-Pb50% 5 5 Pb100% N/A 10 3.5.2 Electrochemical characterization CV, chronopotentiometry and chronoamperometry were performed using a conventional three electrode cell with a silver/silver chloride (Ag/AgCl) reference electrode (Fisher) and a platinum wire (99.99%, Sigma) counter electrode. The electrolyte was 0.5 mol L-1 potassium bicarbonate (KHCO3, 95%, Acros Organics) solution purged with either N2 or CO2 (99%, Praxair) for 30 minutes. A catalyst with a geometric surface area of 1 cm2 was immersed into the solution. CV experiments were performed in a jacketed cell at 25 °C. A blanket of gas was maintained over the solution during the CV experiment. Chronoamperometry was carried out using an h-cell setup with a grafted polyolefin separator and a glass frit to isolate the counter and reference electrodes, respectively. CO2 gas was continuously bubbled into solution to ensure saturation and stir the solution while also dislodging H2 bubbles from the electrode surface. Electrolyte samples (0.4 mL) were taken from the cathode compartment every 20 minutes during 2 hours of electroreduction for a total of 7 samples per run. The formate concentration of the samples were determined using a Megazyme enzyme based spectroscopic kit (K-Form). [219,220] The absorbance of each sample was determined at 340 nm using a Shimadzu UV-2450a. A calibration curve of five HCOO- 65 standards was generated using the megazyme kit. The HCOO- concentrations of each chronoamperometry sample was determined based on linear regression using the standard curve (Appendix Figure D.1.8). The faradaic efficiency was determined from the formate concentrations and the measured current passed through the electrode. 3.5.3 Cyclic voltammetry and H-cell experiments CV was used to determine the activity and electrochemical properties of novel catalysts. CVs of Sn/Pb catalysts show cathodic and anodic waves corresponding to reduction and oxidations, respectively, taking place at the electrode surface. Using this technique it is possible to determine the onset potential of electroreduction of CO2 or H2 on the catalyst surface under varying conditions. However, separating HER and ERC reduction waves is not possible because their onset potentials overlap. Using the CV response as a starting point, the potential or current density can be set to maximize the efficiency of ERC. Additionally, CVs were used for electrode cleaning prior to electrolysis to ensure consistent results. As the current density is increased, electroactive species will react at the electrode surface as soon as they meet it. This is called the mass transfer controlled current density. The rate of a reaction is governed by the rate at which the reactants are brought to the surface. 𝑣𝑟𝑥𝑛 = 𝑣𝑚𝑡 = 𝑖/𝑛𝐹𝐴 Equation 3.7 where 𝑣𝑟𝑥𝑛 is the rate of reaction and 𝑣𝑚𝑡 is the rate of mass transfer. The Nernst-Planck equation is used to determine the rate mass transfer and is written for one dimensional mass transfer along the x-axis below: 𝐽𝑗(𝑥) = −𝐷𝑗 (𝜕𝐶𝑗(𝑥)𝜕𝑥) −𝑧𝑗𝐹𝑅𝑇𝐷𝑗𝐶𝑗𝜕𝜙(𝑥)𝜕𝑥+ 𝐶𝑗𝑣(𝑥) Equation 3.8 66 Equation 3.8 describes the flux of species j (𝐽𝑗(𝑥)) at distance x from the electrode surface. The flux of species j depends on the diffusion coefficient (𝐷𝑗) and the concentration gradient of species j (𝜕𝐶𝑗(𝑥)𝜕𝑥), the migration of a charged species j based on the potential gradient (𝜕𝜙(𝑥)𝜕𝑥) and the charge of species j (𝑧𝑗) and the convection (𝐶𝑗) due to velocity (𝑣(𝑥)). This means that it is possible to increase the rate of reaction by increasing the concentration of dissolved CO2 or increasing the convection speed of the electrolyte near the electrode surface. Since CO2 is not a charged species, increasing the electric field will not affect the mass transfer. Constant current and constant voltage experiments will allow us to measure the faradaic efficiency of ERC under different conditions. Constant current experiments will more closely approach the conditions in an industrial reactor. The experiments will be performed in an H-cell setup. The H-cell design separates the cathodic and anodic processes which prevents the consumption of the ERC product (HCOO-) at the cathode. The separation of the two compartments is maintained by a semipermeable polymer membrane at the junction of the H-cell. This membrane should be selectively permeable and not allow the product of electroreduction to pass through; while allowing charged species (i.e. H+) to cross the membrane and complete the electrical circuit. 3.5.4 Determination of faradaic efficiency Faradaic efficiency of HCOO- was determined using a calibration curve using an enzyme based kit. The formate dehydrogenase (FDH) enzyme selectively catalyzes the reduction of nicotinamide adenine dinucleotide (NAD+) by HCOO-. The NADH formed absorbs light at 340 nm and this absorption can be quantified using a UV-Vis spectrophotometer. This is a very sensitive technique capable of determining HCOO- concentrations down to the 10 mM level. [219] HCOO- is also detectable directly by UV spectroscopy at 260 nm but this is only possible at higher concentrations of ~0.1 M or higher. Other methods (HPLC or GC) should be used for multiple product 67 determination. Gas and liquid phase chromatography are powerful tools to measure different products of CO2 reduction. Chromatography coupled with mass spectrometry can provide qualitative identification as well as quantitative detection depending on the mode of scanning. These techniques are necessary in the case that the product of reduction is not well characterized. However, for Sn/Pb electrodes in this work, HCOO- (and H2) is assumed to be the primary product of electroreduction. 68 Chapter 4: Electrochemical quartz crystal microbalance measurement of IrOx mass changes 4.1 Reversal tolerant anode OER PEMFC objectives The objectives of the OER portion of the thesis (Chapters 4 and 5) are outlined in Figure 4.1. The diagram includes specific goals of the OER durability project including how electrochemical techniques will inform the reversal tolerant anode design process. Figure 4.1 Diagram of the OER durability objectives that are detailed in Chapter 4 and 5. This includes projects at AFCC and Ballard. To produce an industrially viable reversal tolerant anode, it is necessary consider the structure of the anode itself. A good anode should strike a balance between retaining enough water to participate in the OER but also prevent flooding of the anode and insufficient fuel reaching the HOR catalyst. In Chapters 4 and 5, the focus is on the connection between physical properties (crystallinity, particle size and support effect) of OER catalysts and the effect on the durability and activity of OER catalysts in a PEMFC environment. Ex situ methods will be used to determine fundamental OER durability and activity data. Dissolution will be determined from the 69 concentration of catalyst metal ions in the electrolyte after AST by ICP-MS. The EQCM technique will give us potential resolved mass changes of the layer itself. These mass changes describe sorption processes, chemical changes (i.e. oxidation/reduction) and dissolution processes that must be de-convoluted using the CV data. In addition, qualitative descriptions of catalyst durability by analysis of SEM-EDX imaging identify cracks, migration or other defects in the deposited catalyst. OER activity will be determined by calculating Tafel slopes and finding the potential at a given current density from LSV and CV plots. Tafel slopes are commonly used to determine activity of OER catalysts. To validate techniques in this thesis, ex situ data will be compared with in situ results in Chapter 5 to build a bridge between fundamental and applied knowledge. The main objective of Chapter 5 is to correlate ex situ tests to real world performance in situ and find criteria that make a catalyst suitable for OER in a PEMFC anode environment. 4.2 Introduction Fundamental understanding of the interaction between Pt and IrO2 catalysts in PEMFC anodes is important to mitigate degradation during cell voltage reversal events by catalyzing the OER instead of carbon corrosion and Pt dissolution. The potential dependent mass changes of untreated and different heat treated (350, 450 and 550 °C) IrO2 and IrO2+Pt/C powders were investigated in a simulated PEMFC environment using the EQCM technique. During successive potential cycling the frequency changes associated with IrO2 were correlated with the potential dependent uptake of water and formation of iridium oxyhydroxide species; confirmed by XPS. Heat treatment of the IrO2 powders increased crystallinity which triggered changes in the frequency response during accelerated cyclic voltammetry stress tests. During anodic polarization, the untreated IrO2 exhibited a steady increase in frequency beginning at ca. 0.6 V vs. RHE; whereas the heat treated 70 IrO2 first revealed a frequency decrease, then levelled off before increasing. These observations were explained based on differences in potential dependent water uptake and Ir oxyhydroxide formation as a function of IrO2 structural variations from amorphous to crystalline and Ir3+/4+ ratio. Pt/C alone when subjected to potential hold at 1.8 V and 2.0 V vs. RHE experienced significant mass decrease due mainly to carbon corrosion. The addition of IrO2 to Pt/C in the catalyst layer did not show any mass decrease during the same potential hold. The OER catalyzed by IrO2 protected Pt/C from degradation (i.e., carbon corrosion and Pt dissolution) at high potentials and created an ex situ voltage reversal tolerant anode. 4.3 Improving on ex situ AST methodology To develop efficient ex situ techniques, in Chapter 5, the dissolution of unsupported and supported IrOx catalysts is investigated using an ex situ accelerated stress test based on potential stepping. Ir3+ can exist in a soluble state under certain conditions of T, V, and pH. The dissolved Ir3+ formed during this test was compared with the in situ voltage reversal tolerance of a membrane electrode assembly fabricated using the same IrOx catalyst. [221] For the unsupported catalysts, iridium dissolution (measured by ICP-MS) in the three electrode ex situ cell was correlated to in situ voltage reversal loss in a single cell MEA. Reversal loss was defined as the reversal time of a freshly prepared MEA minus the reversal time of an MEA after it had undergone an AAST. While Ir3+ dissolution can be effective in determining promising OER catalyst candidates, it has one major limitation. ICP-MS only provides the Ir3+ concentration in the bulk of the electrolyte at the end of a stress test and it does not account for Ir3+ possibly consumed at the counter electrode. To address this issue, here the EQCM technique is employed for both IrOx and IrOx-Pt/C powder catalyst films. 71 4.4 Frequency to mass relationship Figure 4.2 shows the change in frequency during multiple ink deposition and drying steps on an EQCM crystal. This has been performed previously with similar results. [210] Figure 4.2. (a) Deposition and drying (under a heat lamp) of an ink composed of IrO2 and 2-propanol deposited on the surface of an EQCM crystal. The ink was drop cast in three 5 μL steps (at 0, 12 and 22 minutes) to observe the detachment of particles during multiple depositions. (b) Plot of the deposit frequency change after each ink drop vs. the mass of IrO2 present in the volume of ink. The second y-axis is the calculated mass change from the observed frequency change after each deposition. Figure 4.2a shows the frequency response of three consecutive 5 μL deposition steps on an Au EQCM crystal. During the first deposition, the frequency decreases by ca. 930 Hz (6.5 μg) because of the viscous coupling of the ink. This is similar to immersing the electrode into an ink solution where the frequency decrease is dependent on the viscosity of the solution. [222] During subsequent deposition steps, the frequency first increases indicating detachment of particles from the Au surface. The frequency drops again during the drying step. The mass scale in Figure 4.2 (and later EQCM figures) was calculated using a conversion factor of 7.08 ng Hz-1. Which was obtained by multiplying the sensitivity factor (Cf, Equation 1.6) with the piezo-active area 10 20 30 40 50 602000400060008000mcalculated / gf deposition / Hzm Inkdeposition / g102030405060baR2 = 0.99970 5 10 15 20 25 30-8000-6000-4000-20000+5 L+5 L+5 Lmass / gf / Hztime / min2-propanolviscous coupling detachment of particles drying time shortenedDryingcrystal nearly overloaded 6050403020100 72 (0.4 cm2). Vatankhah et al. showed that there is variation in this sensitivity factor based on Cu film thickness. [211] Using the Cu deposition technique (Figure 3.1 and Figure 3.2), the integral sensitivity factor was found to be 47.8 Hz cm2 μg-1 (vs. the theoretical Cf of 56.6 Hz μg-1 cm2 for the 5 MHz AT-cut quartz crystal). However, in order to obtain the true sensitivity factor that is relevant for the film thicknesses obtained with the various IrO2 and Pt/C catalyst powders, multiple Cu deposition calibration procedures would be needed to approximate the actual catalyst film thickness. Even if this would be accomplished, the sensitivity factor for a rigid film (Cu layer) will be different than the one for a less rigid powder film. For these reasons, it was decided to use the integral sensitivity factor of a standard AT-cut 5 MHz quartz crystal for mass calculations using Equation 1.6 (Figure 3.1 and Figure 3.2). The use of the 5 MHz theoretical sensitivity factor was supported by the multiple IrOx ink deposition experiment in which a known mass was deposited on the surface of the EQCM. The integral sensitivity factor calculated mass change was 18.4 ± 0.6 μg while the mass calculated based on the mass of IrOx in the ink was 17.1 μg for one 5 μL deposition; a 7.6% difference (Figure 4.2b). 4.5 Heat-treated IrOx XRD, XPS and SEM Characterization Figure 4.3 shows the XRD spectra of each IrOx catalyst. The untreated IrOx has one identified metallic Ir phase and an amorphous phase (presumed to be IrO2). The heat-treated catalysts all have identifiable Ir metal and IrO2 crystal phases. Metallic Ir peaks are much smaller in the IrOx heat-treated at 550 °C. Similar heat treatment performed by da Silva et al. determined that IrO2 heat-treated up to 300 °C was similar to untreated IrO2 and remained amorphous in nature. However, heat treatment at 400 °C caused IrO2 to become more crystalline and at 500 °C the peaks became sharper with crystal sizes between 3.4 and 13.7 nm respectively. [115] Previous findings 73 were confirmed by identifying that starting at 350 °C, IrO2 becomes more crystalline and less amorphous in character (Figure 4.3). Additionally, increased peak height and a narrowing of the peak width is observed as the heat treatment temperature was increased to 450 and 550 °C. Figure 4.3. XRD scans of untreated (black), 350°C (red), 450°C (blue) and 550°C (green) heat-treated IrOx. Two crystalline phases were identified (Ir metallic, PDF2 00-006-0598 (ICDD, 2018) and IrO2, PDF2 00-015-0870 (ICDD, 2018)) which are indicated by their Miller indices. The surface chemistry of untreated and heat-treated IrO2 powders was investigated by XPS (Figure 4.4). For the pristine samples (i.e., unexposed to the CV accelerated stress test), the Ir 4f regions of the spectra (Figure 4.4a), present a small peak shift to lower binding energies with increasing heat treatment temperature; but, in essence, for all samples two characteristic peaks are observed with 3 eV energy difference at ca. 61.7 eV and 64.6 to 64.7 eV. Metallic Ir, which was identified by XRD (Figure 4.3), is normally seen at 60.8 eV and generates asymmetric peaks typical for metallic conductors. [137,223] The absence of the metallic Ir XPS signal indicates that it is covered 2-Theta - Scale5 10 20 30 40 50 60 70 80 90550°C450°C350°CUntreatedIr (111) Ir (200)Ir (220) Ir (311)Ir (222)IrO2 (101)IrO2 (110)IrO2 (211)IrO2 (200)IrO2 (220)IrO2 (310) 74 by a surface oxide film. The peaks with a doublet separation of 3 eV belong to Ir4+ 4f 7/2 and 5/2, corresponding to anhydrous rutile-type IrO2. [223] The shift toward slightly lower binding energies with increasing heat treatment temperature reflects the increased content of crystalline vs. amorphous IrO2. Accelerated CV cycling of the samples produced somewhat wider peaks in the Ir 4f region as a result of Ir3+ presence with peaks at 62.6 eV and 65.6 eV. [223] Deconvolution and quantification of these peaks is presented in Appendix Figure B.1.8, Appendix Figure B.1.10 and Appendix Table B.1.1. The oxygen regions of the pristine (uncycled) IrOx powders (Figure 4.4c), changed with increasing heat treatment temperature. Thus, the lattice oxide peak at ca. 530 eV became more pronounced and sharper. The peak at 531.5 eV became a smaller due to the oxidation of oxyhydroxide to lattice oxide. It is presumed that the OH character is an indication of the presence of Ir3+ in the IrOx catalyst. This is similar to previous reports [58,115,137] showing the difference in XPS spectra between untreated and heat-treated IrOx. Figure 4.4b and d present the XPS spectra of the IrOx samples after being subjected to CV accelerated stress tests. The most significant changes induced by successive potential cycling are observed in the oxygen region. For the untreated iridium oxide, there is a peak broadening due to OH and presence of a new oxygen peak at ca. 532.4 eV (Figure 4.4d). These changes are due to hydrated iridium oxyhydroxide formation. [58,125] Additionally, with potential cycling, the main oxide peak for IrO2 550 °C shifted toward higher binding energies, from 530 to 532.4 eV (Figure 4.4d). The latter peak (at 532.4 eV) is characteristic for O (1s) in surface adsorbed water. [223] Deconvolution and quantification of the relative percentages of each oxygen environment are presented in Appendix Figure B.1.9, Appendix Figure B.1.10 and Appendix Table B.1.1. .Previously, Augustynski et al. reported the presence of sulfate ions in electrochemically prepared 75 IrO2 films. [224] The XPS survey scans (Appendix Figure B.1.4) did not identify any sulfur present in the IrO2 layers after CV accelerated stress tests. Figure 4.4. XPS of the Ir region (a and b) and Oxygen region (c and d) of pristine (a and c) and post CV accelerated stress test (b and d) of IrOx: untreated (black), 350°C (red), 450°C (blue) and 550°C (green) heat-treated. CV cycling performed between 0.05 to 1.2, 1.4 and 1.6 V vs. RHE, 30 times each at 25 °C in 0.09 M H2SO4. SEM imaging was performed to determine morphological changes from heat treatment of the IrOx catalysts (Figure 4.5). Particles and their aggregates were in general larger after heat treatment at 550 °C (Figure 4.5b) than the untreated IrO2 (Figure 4.5a). Untreated IrOx had more small particles 70 68 66 64 62 60400800120016002000 Untreated 550oCNo metallic Ir on surfaceBinding Energy / eVIntensity / counts s-1Intensity / counts s-1536 534 532 530 528560064007200 Untreated 550oCIntensity / counts s-1Lattice OxideBinding Energy / eVOHH2O528530532534536200040006000529.9 eVIntensity / counts s-1450550350UntreatedLattice Oxide Binding Energy / eVOH606264666870020004000600064.6 eV 61.7 eVIntensity / counts s-1 Binding Energy / eVUntreated350450550No metallic Ir on surfaceIrO2bcad 76 dispersed over the Au substrate surface. In contrast, IrOx heat-treated to 550 °C had nearly all of its particles contained within larger aggregates and very few smaller particles on the Au surface. This is because the heat-treated IrOx formed aggregates that were not broken up by sonication during ink preparation. The incomplete coverage of the Au substrate with IrOx aggregates is a limitation of investigating drop cast powders. Investigation of the deposited IrOx at 350 and 450 °C (Appendix Figure B.1.2 and Appendix Figure B.1.3) showed morphology similar to IrO2 heat-treated at 550 °C. Additionally, SEM images taken after CV accelerated stress tests did not have any apparent differences to those of pristine films (Appendix Figure B.1.2 and Appendix Figure B.1.3). Figure 4.5. SEM images of pristine (a) untreated IrOx and (b) IrOx 550 °C drop cast on an Au EQCM crystal. Images were acquired at 138 kx. 4.6 Cyclic voltammetry and relative capacitance (RC) CVs were obtained concurrently to the frequency response for each IrOx layer (Figure 4.6). The CVs show very little change from the beginning to the end of the stress test. Comparing untreated 77 IrOx (Figure 4.6a) to the heat-treated IrOx (Figure 4.6b-d), heat treatment causes a large decrease in the capacitive current. The area under the CV from 0.5 to 1.0 V vs. RHE was integrated to find the total charge for each catalyst layer (expressed in millicoulombs, mC). The total charge in the above potential region is due to both electric double layer charging and contributions from faradaic charge transfer. Isolating the purely capacitive component is not possible based on the CVs showed in Figure 4.6. The total charge was found to be: 13 mC, 4.0 mC, 1.4 mC and 0.8 mC for IrO2 untreated, 350 °C, 450 °C and 550 °C heat treated respectively. For the untreated IrOx, the total charge was more than 15 times larger than the 550 °C heat-treated IrOx total charge (Figure 4.6). The decrease of total charge (Figure 4.6) with heat treatment is much larger than the BET surface area loss which decreased from 28.2 to 19.0 cm2 g-1 (Table 3.2). The decrease in total charge with increasing heat treatment temperature is caused by a decrease in the surface area of the aggregates, possibly a decrease in the capacitance of the aggregates themselves and a decrease in the coverage of the film on the surface of the gold electrode (as seen in SEM imaging, Figure 4.5). Additionally, heat treatment caused a decrease in observed current at ca. 0.7 V vs. RHE. Previous studies have found that this wave corresponds to the oxidation of Ir3+ to Ir4+. [135,225] It is proposed that the thermal oxidation of Ir3+ to Ir4+ during heat treatment decreases the proportion of Ir3+ in the sample. The peaks at ca. 1.1 and 1.4 V vs. RHE were also present on the CVs of an uncoated Au EQCM crystal. This suggests that they are due to Au reduction and oxidation, respectively (Figure 4.6 and Appendix Figure B.1.6). The size of these peaks could be proportional to the exposed area of Au not covered by IrOx. 78 Figure 4.6. CV response of IrO2 untreated (a) and IrOx heat-treated (350 °C (b), 450 °C (c) and 550 °C (d)) from 1.2 (black), 1.4 (red) and 1.6 V (blue) to 0.05 vs. RHE, 30 times each at 25 °C in 0.09 M H2SO4. The current at 1.6 V vs. RHE decreases with increasing heat treatment temperature. This can be attributed to the loss of OER activity with increasing heat treatment temperature which has been reported previously by several groups. [115,134] Additionally, oxidation of amorphous Ir3+ to Ir4+ by heat treatment decreases the formation of an electrophilic O- during OER. Pfeifer et al. found that the presence of electrophilic O- corresponds to increased OER activity. They proposed that this effect is caused by the increased reactivity of O- which promotes O-O bond formation. [135] 0.0 0.3 0.6 0.9 1.2 1.5 1.8-0.60.00.61.21.8ci / mA cm-2E / V vs. RHE0.0 0.3 0.6 0.9 1.2 1.5 1.8-10123i / mA cm-2E / V vs. RHEa0.0 0.3 0.6 0.9 1.2 1.5 1.8-0.40.00.40.81.21.6di / mA cm-2E / V vs. RHE0.0 0.3 0.6 0.9 1.2 1.5 1.8-0.80.00.81.62.4bi / mA cm-2E / V vs. RHEIrOx 350 C IrOx UntreatedIrOx 450 C IrOx 550 C Au peaksIr4+Ir3+ 79 4.7 EQCM response of heat-treated vs. untreated IrO2 Interpreting the EQCM response of a porous layer deposited from an ink can be challenging. An observed increase in frequency cannot be automatically attributed to potential dependent mass decrease of the deposited film. It could also be caused by viscoelastic effects, detachment of particles from the film, a chemical change of the film or a decrease in the viscosity of the liquid near the electrode surface. [226–229] Figure 4.7 shows the frequency response of the untreated and heat-treated (at 350, 450 and 550 °C) IrOx during successive CV tests. Regardless of the IrOx sample, the first potential sweep starting at 1.15 V vs. RHE and going in the cathodic direction, was associated with a decrease of frequency (i.e., mass increase) up to either 0.7 V vs. RHE (for untreated and 350 °C samples) or 1.0 V vs. RHE (for 450 and 550 °C samples). Further scanning to 0.05 V vs. RHE induced a frequency increase (i.e., mass decrease) for all samples starting at approximately 0.7 V vs. RHE. The untreated IrOx had the smallest overall frequency increase (44 Hz or -310 ng) compared to all the heat-treated samples. To determine the cause of this frequency increase, electrolyte samples were taken from near the electrode surface during the initial sweep and were tested for Ir content by ICP-MS. No dissolved Ir was detected (limit of detection: 1 μg L-1 or 4 ng dissolved Ir in the sample) in the electrolyte. The total loading of Ir on the EQCM crystal was 28.9 μg. To reach the ICP-MS limit of detection (LOD), 0.00014% of the deposited Ir would need to dissolve into the solution during the first scan from 1.15 V vs. RHE. Additionally, the 4 ng dissolution needed to reach the ICP-MS LOD is 75 times smaller than the smallest first scan mass change of ca. 300 ng in the untreated IrO2 (Figure 4.7a). Therefore, the frequency increase is neither due to dissolution of soluble Ir species nor detachment of particles from the surface but can be associated with potential dependent surface reactions and possibly viscoelastic effects of the particle aggregates. 80 As pointed out by Burke and Scannell, the hydrous Ir oxide surface layer can be regarded as a porous gel or polymer. [162] Birss et al. argued that in the case of an electrodeposited IrO2 film viscoelastic effects are negligible. [161] However, here particle aggregates are investigated. It is speculated that during the first sweep after immersion of the EQCM probe in the electrolyte, the aggregation of particles is not yet stabilized and changes in layer viscosity could take place due to hydration. Such viscosity changes could influence the measured frequency response. Lu et al. investigated the viscoelastic effects of binderless carbon blacks with different aggregate sizes. [69] They reported that the interparticle movement (e.g., sliding of the layers relative to each other) contributes to the viscous deformation but overall the elastic character is dominant. Similar effects can be expected for hydrous Ir oxide particle aggregates as well. Kanazawa showed that the resonant frequency increases abruptly with increasing viscosity as the film transitions from elastic to viscous character. [230] Outside this transition region, the frequency is constant and independent of viscosity. In the transition region, for a 6 µm thick viscoelastic film with a density of 1200 kg m-3 and shear modulus of 2x107 N m-2 the expected frequency increase is approximately 0.13 MHz. [230] The latter value is orders of magnitude higher than the frequency changes measured in the present work even when accounting for possibly thinner film (Figure 4.7). Therefore, in accordance with previous studies, the elastic character of the film dominates and the frequency changes are related to film mass changes. Regarding surface reactions during the first cycle, the reduction of mainly anhydrous IrO2 at potentials lower than 0.7 V vs. RHE causing a mass decrease on the electrode according to: 2𝐼𝑟𝑂2,(𝑠) + 2𝐻+ + 2𝑒− → 𝐼𝑟2𝑂3,(𝑠) + 𝐻2𝑂 Equation 4.1 This could explain the observed first cycle frequency increases below 0.7 V vs. RHE (Figure 4.7). 81 As shown by the XRD spectra (Figure 4.3), the Ir oxide species have different compositions for the heat-treated, crystalline samples as compared to the untreated, amorphous species. The latter has the lowest oxide content (Figure 4.3) and, therefore, exhibits the smallest frequency increase (mass decrease) according to Equation 4.1 during the first cycle (Figure 4.7). Figure 4.7. EQCM response of untreated (a) and heat-treated (350 °C (b), 450 °C (c) and 550 °C (d)) IrOx during cyclic voltammetry experiments from: 1.2 V (black), 1.4 V (red) and 1.6 V (blue) to 0.05 V vs. RHE. The cycles were repeated 30 times each. “fs” – indicates first potential sweep. Temperature: 25 °C. Electrolyte: 0.09 M H2SO4. Following the first sweep, successive potential cycling coupled with frequency measurements were performed using three start potentials: 1.2, 1.4 and 1.6 V vs. RHE (Figure 4.7). The corresponding cyclic voltammograms are shown in Figure 4.6. Two observations are made. First, 0.0 0.3 0.6 0.9 1.2 1.5 1.8-150-100-50050100150200f / HzE / V vs. RHE8004000-400-800-1200cmass / ng0.0 0.3 0.6 0.9 1.2 1.5 1.8-120-80-4004080120160mass / ngf / HzE / V vs. RHE8004000-400-800-1200d0.0 0.3 0.6 0.9 1.2 1.5 1.8-300306090mass / ngf / HzE / V vs. RHEa1500-150-300-450-6000.0 0.3 0.6 0.9 1.2 1.5 1.8-60060120180240300b5000-500-1000-1500-2000 f / HzE / V vs. RHEmass / ngIrOx 350 C IrOx UntreatedIrOx 450 C IrOx 550 C bcdafsfsfsfs 82 the frequency decreases with each successive cycle for all samples. Second, the potential dependent frequency profiles of the 450 and 550 °C heat-treated IrO2 samples are similar from the first cycle through subsequent cycles, whereas in case of the 350 °C heat-treated and especially the untreated IrOx, the frequency profiles changed between the first cycle and the subsequent cycles. The frequency response for the Au EQCM crystal alone was investigated using the same CV cycling protocol to filter out and possible interference from the uncovered Au surface. There was only a small decrease in frequency (ca.-10 Hz or 71 ng) during the first 30 cycles from 1.2 to 0.05 V vs. RHE. Afterwards, the frequency response was stable (i.e., frequency change lower than 5 Hz) during successive cycling from 1.4 or 1.6 V to 0.05 V vs. RHE (Appendix Figure B.1.7). Hence, the observed frequency profile changes during successive cycling (Figure 4.7) are intrinsic features of the differently heat treated IrO2 films. The gradual frequency decreases (i.e., mass increase) with each successive cycle is due to chemical process occurring that increase the mass of the adhered film. The XPS response of these films after the CV stress tests (Figure 4.4d) shows larger peaks for OH and surface adsorbed H2O suggesting that compositional modifications occurred over time in the film due water uptake and formation of: IrO2xH2O, Ir2O3xH2O and various Ir oxyhydroxide species (e.g., IrO(OH)y). For the untreated sample (Figure 4.7a), all successive scans show a frequency increase peak in the potential region 0.6 V to 1.1 V vs. RHE. In contrast, successive scans of the heat-treated IrOx in the same potential region display either a fairly constant frequency (for the sample heat-treated at 350 °C, Figure 4.7b) or a frequency decrease valley for the 450 and 550 °C treated samples (Figure 4.7c and d). In other words, Figure 4.7 reveals that increasing the heat treatment temperature of IrOx above 350 °C increases the apparent mass between ca. 0.6 V to 1.1 V vs. RHE, whereas the untreated IrOx displays a low apparent mass in the same potential region. This behavior is 83 independent of the start potential and is an intrinsic feature of IrOx as a function of heat treatment. Furthermore, there is little hysteresis between the anodic and cathodic scanning direction of each cycle (Figure 4.8). In between each set of thirty cycles, there is a frequency gap before the beginning of the next set of thirty cycles. The gap is caused by a five second hold at the starting potential (max potential minus 0.05 V vs. RHE). The final frequency increases at 1.4 and 1.6 V correspond to the Ir4+ to Ir5+/6+ transition and the beginning of the OER (Figure 4.7). This causes a further frequency increase from oxygen evolution from the surface of the electrode. [125] In Figure 4.8 we overlay the final CV scan (1.6 V to 0.05 V vs. RHE) and the corresponding frequency response of Figure 4.7. During polarization from 0.05 V to 1.6 V, the frequency of the untreated IrOx increased by ca. 60 Hz (-425 ng) (Figure 4.7a). Birss et al. reported the formation of an iridium oxyhydroxide at low potentials which is oxidized to less hydrated IrO2 and causes a mass decrease when polarizing from 0 to 1.6 V vs. RHE. [161] They suggested that the oxyhydroxide Ir3+ species has a higher molar mass than IrOx. The differences in the frequency profiles between untreated and heat-treated (especially above 350 °C) IrO2 (Figure 4.8d and e) can be explained in relation to the XPS spectra shown by Figure 4.4d. As the heat treatment temperature is increased, the increasingly crystalline IrOx has a higher water coverage on the crystal surface (as revealed by XPS), a process which in terms of frequency response, competes with the mass changes induced by the redox reaction between Ir3+ and Ir4+ oxyhydroxide species. This surface adsorbed water to oxide and hydroxide ratio could be significantly less below the catalyst surface, but the XPS results clearly show a higher peak area of H2O in the IrOx 550 °C vs. untreated IrOx (Appendix Table B.1.1). 84 Figure 4.8 Combined frequency and current responses for (a) background Au EQCM crystal, (b) untreated IrO2, (c) 350 °C heat treated IrO2, (d) 450 °C heat treated IrO2 and (e) 550 °C heat treated IrO2 during the final CV scan from 1.6 V to 0.05 V vs. RHE at 25 °C in 0.09 M H2SO4. 0.0 0.3 0.6 0.9 1.2 1.5 1.8-1.0-0.50.00.51.01.52.0-250-200-150-100-50050f / Hzi / mA cm-2E / V vs. RHE0.0 0.3 0.6 0.9 1.2 1.5 1.8-1.0-0.50.00.51.01.52.02.53.0-300-250-200-150-100-50050f / Hzi / mA cm-2E / V vs. RHE0.0 0.3 0.6 0.9 1.2 1.5 1.8-2-10123-60-50-40-30-20-10010f / Hzi / mA cm-2E / V vs. RHE0.0 0.3 0.6 0.9 1.2 1.5 1.8-0.50.00.51.01.52.0-250-200-150-100-50050f / Hzi / mA cm-2E / V vs. RHEIrO2 350 C IrO2 UntreatedIrO2 450 C IrO2 550 C ccdbAu a0.0 0.3 0.6 0.9 1.2 1.5 1.8-0.4-0.20.00.2-20-1001020f / HzE / V vs. RHEi / mA cm-2 85 4.8 EQCM response to potentiostatic holds The frequency response of untreated IrOx was measured under potentiostatic mode (Figure 4.9). The frequency change during 5-minute potential holds at 0.9, 1.2, 1.4 and 1.6 V vs. RHE (but not at 0.05 V vs. RHE) was similar for the pristine and CV stress tested untreated IrOx. When polarized at positive potentials, the frequency response has an initial frequency increase (25-75 Hz or -170-530 ng) which then decreases over five minutes. This is caused by the oxidation of oxyhydroxide species to lighter IrO2. At 1.6 V vs. RHE, the frequency is lower than at 1.4 V vs. RHE indicating the formation of a heavier layer. The main difference occurs when polarizing to 0.05 V vs. RHE where the pristine film has a large frequency increase of ca. 900 Hz (-7 μg) while the CV cycled film has a frequency decrease of ca. 200 Hz (1.4 μg). The frequency increase of the pristine film is similar to the first CV cycle frequency increase shown in Figure 4.7 and discussed previously. This effect is not observed in the CV cycled film because of different Ir speciation and water uptake characteristics. Figure 4.9. Five minute potential holds at 0.9, 1.2, 1.4, 1.6 and 0.05 V vs. RHE with (a) freshly prepared and (b) CV cycled untreated IrOx. Electrolyte: 0.09 M H2SO4 purged with N2. Scan rate: 100 mV s-1. Temperature: 25 °C. 0 400 800 1200 1600 2000-300-250-200-150-100-500Time / s1.4 Vf / Hzmass / ng0.9 V1.2 V1.6 V0.05 VUntreated IrO2After CV Cycling20001500100050000 400 800 1200 1600-2000200400600800Time / s1.4 Vf / Hzmass / ng0.9 V1.2 V1.6 V0.05 VUntreated IrO2Pristine15000-1500-3000-4500-6000baxx 86 4.9 Pt/C with IrO2 potential cycling and constant potential polarization The frequency response of carbon and Pt/C containing layers was measured with and without IrOx (Figure 4.10). The frequency increases during successive CV cycling, which indicates a mass decrease of the deposited film. Pt/C alone had the largest mass decrease followed by Pt/C + untreated IrO2, Ketjen black and Pt/C + IrO2 550 °C. Wickman et al. determined that most of the mass loss can be attributed to the dissolution of Pt from a Pt/C catalyst. [167] A smaller mass loss contribution from the carbon support was also observed. da Silva et al. investigated the degradation of Pt/IrOx bi-functional catalysts by ICP-MS. Dissolution of both Pt and Ir was observed by ICP-MS during potential cycling (and square wave potential steps) from 0.6 to 1.6 V vs. RHE. [231,232] In the present case however, Pt dissolution combined with carbon corrosion cannot explain the total observed mass loss of the film because the Pt/C theoretical mass decrease (ca. -18 μg) was larger than the total deposited mass (10.3 μg). Additionally, the layers were visually inspected after each experiment and there was no change in the appearance of the film that would indicate that all of the Pt and carbon had been corroded. The Pt/C + untreated IrOx and Pt/C + IrOx 550 °C experiments were repeated three times and the calculated decrease in mass was 11.1 ± 1.5 μg and 2.6 ± 1.2 μg respectively. This can be compared to the calculated total solids loading of 10.3 μg. Therefore, in the presence of the carbon support, viscoelastic effects must be considered as well to explain the observed frequency increase. Lu et al. investigated the viscoelastic properties and their relationship to aggregate size of carbon black. [69] They found that the viscous character increases and the elastic character decreases with increasing aggregate size. The effect of these changes on the EQCM response has been to some extent investigated using conducting and non-conducting polymers. [70–72,230] In the case of the Pt/C film, there seems to be enhanced viscous character induced by extensive potential cycling. The decreasing 87 rigidity of the film could lead to significant phase shift and possibly dissipation before reaching the outer film surface. In this case, the Sauerbrey equation no longer applies. [71] Figure 4.10. Frequency response of (a) Ketjen black, (b) Pt/C alone and with (c) untreated IrOx and (d) IrOx heat-treated at 550 °C. Catalyst mixtures were 2:1 Pt:Ir metals basis and total solids loading were 2.7, 5.5 and 7.5 μg cm-2 for a, b and both c and d respectively. CV cycled from 1.2 (black), 1.4 (red) and 1.6 V (blue) to 0.05 V vs. RHE. Electrolyte: 0.09 M H2SO4 purged with N2. Scan rate: 100 mV s-1. Temperature: 25 °C. Adding IrOx to Pt/C caused lower frequency increase during successive CV sweeps compared to Pt/C alone. Pt/C + IrOx 550 °C had the smallest frequency increase of all the tested catalysts (Figure 4.10). Since the frequency increase cannot be caused solely by Pt dissolution and carbon corrosion, the addition of IrO2 must also have some impact on the physical properties of the film. 0.0 0.3 0.6 0.9 1.2 1.5 1.8-500050010001500200025003000mass / ngf / HzE / V vs. RHE30000-3000-6000-9000-12000-15000-18000-210000.0 0.3 0.6 0.9 1.2 1.5 1.80100200300400500600mass / ngf / HzE / V vs. RHE0-700-1400-2100-2800-3500-42000.0 0.3 0.6 0.9 1.2 1.5 1.803006009001200mass / ngf / HzE / V vs. RHE0-2000-4000-6000-80000.0 0.3 0.6 0.9 1.2 1.5 1.8040080012001600mass / ngf / HzE / V vs. RHE0-2200-4400-6600-8800-11000-13200bcdPt/CPt/C+IrOx (unt.)Pt/C+IrOx (550 C)a Ketjen 88 The IrOx particles were well incorporated into the Pt/C particles (Appendix Figure B.1.6) and likely increased the rigidity of the film mitigating the viscous behavior. Figure 4.11 shows the frequency responses accompanying the first cycle voltammetry sweeps for Pt/C and Pt/C + IrOx. The complete set of thirty CV cycles with three potential ranges are presented in Appendix Figure B.1.5. In the case of Ketjen black, the frequency increases by ca. 230 Hz during the cathodic sweep and then, after a fairly constant frequency domain up to 1.2 V vs. RHE, decreases by ca. 170 Hz up to 1.6 V vs. RHE (Figure 4.11a). The latter frequency drop is due to the formation of oxidized functional groups on carbon coupled with oxide formation on the uncovered Au substrate. The oxidation and reduction peaks at 1.4 and 1.1 v vs. RHE respectively, can be attributed to exposed Au substrate oxidation and reduction, as discussed previously (Figure 4.8a). Furthermore, for Ketjen black the peak at ca. 0.6 V vs. RHE on the cathodic sweep can be attributed to the two-electron O2 reduction (Figure 4.11a). For each of the Pt/C containing layers (Figure 4.11b, c and d) the CVs and frequency responses have the same basic shape. During the cathodic sweep there is oxide reduction (both Pt and Au) with the associated frequency increase followed by levelling off in the hydrogen underpotential deposition region. The anodic sweep begins with a steep frequency decrease suggested previously to be caused by ingress of water into the support. [165] It is noted that this frequency decrease was absent in case of the carbon support alone (compare Figure 4.11a with Figure 4.11b, c or d). Due to the presence of Pt the catalyst layer becomes more hydrophilic with enhanced water uptake and strong water adsorption on Pt. 89 Figure 4.11. The first CV cycle from 1.6 to 0.05 V vs. RHE of (a) Ketjen black, (b) Pt/C alone and with (c) untreated IrO2 and (d) IrO2 heat-treated at 550 °C. Catalyst mixtures were 2:1 Pt:Ir metals basis. CVs were recorded in 0.09 M H2SO4 purged with N2 at 100 mV s-1 at 25 °C. Thereafter, the frequency levels off from ca. 0.6 to 0.9 V vs. RHE and increases slightly during Pt oxidation between 0.9 and 1.1 V vs. RHE. This increase is due to the conversion of Pt hydroxides to an oxide monolayer. da Silva et al. found that Pt dissolution in Pt/IrOx bi-functional catalysts was at a maximum near 0.6 V vs. RHE. [231] Further scanning to 1.6 V produces a frequency decrease as indicated previously due to more extensive surface oxidation of both Pt and the exposed Au substrate. 0.0 0.3 0.6 0.9 1.2 1.5 1.8-0.9-0.6-0.30.00.30.60.9200250300350400450500f / HzE / V vs. RHEi / mA cm-20.0 0.3 0.6 0.9 1.2 1.5 1.8-0.4-0.20.00.20.40.6f / HzE / V vs. RHEi / mA cm-2700800900100011000.0 0.3 0.6 0.9 1.2 1.5 1.8-0.6-0.30.00.30.60.9030060090012001500E / V vs. RHEi / mA cm-2f / Hz0.0 0.3 0.6 0.9 1.2 1.5 1.8-1.5-1.0-0.50.00.51.01.5f / HzE / V vs. RHEi / mA cm-202004006008001000bc dPt/CPt oxidation /DissolutionPt reductiona KetjenPt/C+IrOx (unt.) Pt/C+IrOx (550 C) 90 Constant potential polarization for 300 s each at 1.6, 1.8 and 2.0 V were also applied to Pt/C and Pt/C + IrOx catalyst (Figure 4.12). These experiments were carried out to simulate a typical fuel cell voltage reversal event where the cell is starved of fuel and held at a fixed current until a potential limit is reached. The potential holds at 1.6 V vs. RHE were similar for Pt/C alone and Pt/C with untreated IrOx. Starting at 1.8 V vs. RHE, however, Pt/C alone experienced a sharp frequency increase (mass decrease) of 150 Hz or -1300 ng, respectively. The slope of the frequency change increased at 2.0 V vs. RHE (Figure 4.12) as a result of carbon corrosion and Pt dissolution. Pt dissolution should only occur at the beginning of polarization during the formation of Pt oxides. In contrast, the Pt/C + IrOx displayed a slight frequency decrease (mass increase) at 1.8 V vs RHE. This difference in the frequency response indicates that instead of carbon corrosion and Pt dissolution, OER is favored in the presence of IrOx, creating a durable catalyst able to withstand voltage spike events. The slight frequency decrease with IrOx could be due to oxidation of the metallic Ir under the IrO2 surface. The evolved gas bubbles from the Pt/C + untreated IrOx layer were small and fine (indicating O2), whereas, the evolved gas bubbles from the Pt/C alone were fewer, larger and adhered to the surface (indicating CO2). After the test, the electrode was agitated to remove any bubbles that were attached to the surface. No significant change in the end frequency was observed after removal of the bubbles indicating that these frequency increases of the Pt/C alone were persistent and not caused by adhered gas bubbles changes the density and viscosity of the electrolyte near the surface. Furthermore, during constant potential hold experiments, the viscous behavior of the Pt/C film does not appear to have an effect in contrast to the extensive potential sweep cycles (Figure 4.12 vs. Figure 4.10b). Appendix Figure B.1.6 shows the EDS mapping of the Pt/C + IrOx catalyst. The incorporation of the IrOx likely improved the carbon corrosion resistance of the layer. 91 Figure 4.12. Five minute potential holds at 1.6, 1.8 and 2.0 V vs. RHE with freshly prepared Pt/C (black) and Pt/C with untreated IrO2 (red). Pt/C with untreated IrO2 had to be stopped after 1.8 V vs. RHE because of damage to the crystal. Electrolyte: 0.09 M H2SO4 purged with N2. 25 °C. 4.10 Conclusions EQCM was used to investigate the frequency changes of untreated and heat-treated IrOx catalyst powders without or with Pt/C in a simulated PEMFC. The combination of Pt/C with IrO2 catalysts is relevant for simulating a PEMFC anode designed for voltage reversal tolerance. The goal is to develop methodologies to better understand the voltage reversal tolerance of fuel cell anode catalysts and more efficiently screen possible catalyst formulations. The untreated and heat-treated (at 350, 450 and 550 °C) crystalline IrOx catalysts behaved differently in terms of frequency change during thirty successive potential cycles. This was caused by potential dependent uptake of water in all IrOx samples and formation of iridium oxyhydroxide layers in untreated IrOx with different compositions, confirmed by XPS. The exact chemical makeup of the layer was difficult 0 300 600 9000200400600Pt/C + IrO2 Unt2.0 V1.8 VTime / s1.6 Vmass / ngf / HzPt/C0-1000-2000-3000-4000x 92 to determine because of the complex nature of the iridium chemistry; but there was no observed sulfur in the XPS scan suggesting that presence of SO4-2 ions in the electrolyte did not affect the frequency responses. The IrO2 layers remain relatively stable under these CV cycling conditions and there is no significant iridium dissolution (confirmed by ICP-MS). Moreover, the frequency responses of Pt/C + IrOx catalyst layers were investigated during both potential cycling and constant potential polarization and compared to Pt/C and Ketjen black alone. During potential holds at 1.8 and 2.0 V vs. RHE, a large mass decrease was observed in case of Pt/C caused by carbon corrosion. The addition of IrOx prevented the mass loss as shown by the virtually constant frequency, thereby, it protected Pt/C from corrosion (i.e., simulating ex situ the voltage reversal tolerant anode). The major challenge in the use of EQCM for powder catalyst investigation is producing layers with consistent and repeatable frequency behavior and separating (or minimizing) viscous layer related effects from intrinsic potential dependent mass changes experienced by the catalyst layers. In this work, the IrOx and Pt/C deposition conditions were tightly controlled and each experiment was repeated to ensure the consistency of the results. It is demonstrated that EQCM could be a rapid and powerful tool to probe the potential dependent mass changes associated with a myriad of powder catalyst materials. . 93 Chapter 5: Ex situ testing of reversal tolerant anode catalysts 5.1 Introduction Cell failure due to fuel starvation is a major issue affecting fuel cell durability and lifetime. During transient conditions, e.g., start-up/shut-down, freeze startup etc., the anode or localized regions of the anode can become starved of fuel putting the cell into reversal conditions with cell voltage reaching as low as -2 V. This can induce corrosion of the carbonaceous MEA components and massive failure of cell stack. To counteract this failure mode, it has been proposed to use OER catalysts to oxidize water, forming O2 instead of oxidizing carbon. Currently, the testing of these OERs is done in situ in fuel cells which is time consuming and expensive. [44,53,64] A quick and efficient ex situ testing protocol is needed to screen OER catalysts for use in RTAs. This combined with an understanding of the relationship between OER catalyst structure, ex situ characteristics and in situ activity can lead to better RTAs. In this chapter, OER catalysts will be investigated in a simulated PEMFC anode for their ability to mitigate cell voltage reversal induced carbon corrosion. Ex situ methods including CV, LSV, ICP-MS, XRD, EDX-SEM will be used to characterize OER catalysts. In situ voltage reversal data will be used to validate ex situ dissolution results. It will be shown that this ex situ AST is able to identify promising OER catalyst candidates for use in RTAs for PEMFCs. 5.2 In situ vs. ex situ protocols for RTA investigations The in situ tests are similar in principle to ex situ three electrode cell test. Much like the ex situ case, potential or current controlled techniques are used to probe a test cell. In situ test cells use a full membrane electrode assembly in single or multi-cell configurations. Cell temperature and humidity are controlled via a test program and the corresponding thermostats and gas humidifiers. Pt/C is capable of retaining water near the catalyst surface. This creates quasi-flooded conditions 94 which are similar to the three electrode cell used in the ex situ experiments. Initial reversal tolerance (minutes) is measured with a reversal event continued until the cell voltage reaches ~-2.0 V. This test can do significant damage to the MEA. Therefore, a new MEA is then put through an AAST to age the MEA in a fashion similar to what would occur during normal operation. The stress test cycles air and fuel or air and N2 mixtures to simulate repeated cell startups and shutdowns. Additionally, the cell performance is determined before and after the AAST by generating a cell polarization curve to ensure the health of the cell. The polarization curve provides information on the poisoning of the cathode catalyst due to OER metal ion migration from the anode across the membrane. After the AAST, the same process as before (reversal event to -2.0 V) is used to measure the reversal tolerance time. The reversal time retention is the difference between the initial reversal tolerance time and the reversal tolerance time after AAST. This difference will be compared to the dissolved metals content of the ex situ AST electrolyte. For both the polarization curves and the extended reversal, a ranking of the in situ RTA catalysts will be generated from most durable to least durable and compared to the ranked results of the ex situ tests. 5.3 Temperature controlled accelerated stress testing Low temperature PEMFCs have a normal operating range between 50-100 °C. [30] In order to more accurately mimic the in situ fuel cell conditions in an ex situ environment, a temperature range between 20 and 80 °C was employed to investigate the TiO2 supported IrOx catalyst. In Figure 5.1, the left y-axis shows the total dissolved Ir present in the electrolyte and the right y-axis shows the Ir loss as a percentage of the total deposited Ir in the electrode. The Ir concentration in the electrolyte after the ex situ AST increased exponentially with temperature (Figure 5.1). The error bars are from the standard error of 3-5 repeated tests. The TiO2 support is capable of reducing 95 Ir4+ to Ir3+. Increasing temperature would increase the rate of this reaction and thereby increase the rate of dissolution of Ir3+ to the bulk electrolyte. Because of the absence of this dissolution mechanism in unsupported catalysts, it is expected that unsupported catalysts would have a smaller increase in dissolved Ir3+ with increasing temperature. Figure 5.1 The temperature effect on Ir dissolution of the IrOx/TiO2 catalyst after the 30,000 stepping cycles from 0.05 to 1.2 V vs. RHE in 0.09 M H2SO4 (ex situ AST). Appendix Table C.1.1 shows the elemental composition of the electrolyte solutions. Titanium is only present in case of IrOx/TiO2, however, trace amounts of various metals (Sn, Al, Zn, Pb, Ni, Ag, Fe and Cu) were detected in all of the samples except for the blank; which suggests that the source of these could be from impurities in the OER catalyst precursors. 5.4 Ir dissolution after ex situ AST: unsupported and supported IrOx comparison Previously, McCrory et al. carried out studies of Ir and Ru based OER catalysts in both metallic and oxide forms under H2O electrolyzer conditions. It was determined that every tested oxide 96 catalyst was unstable under acidic conditions with the exception of IrOx. [59] In the current study, four IrOx based OER catalysts were investigated using the ex situ accelerated stress testing protocol. This test was designed to mimic conditions that would be experienced by the anode during normal operation (at 0.05 V vs. RHE) and during start-up/shut-down (at 1.2 V) in a hydrogen starved environment. As shown by Figure 5.2, IrOx/TiO2 and the IrOx High SA produced the highest Ir concentration in the electrolyte after ex situ AST. This is most likely due to the amorphous state of the oxide and high BET surface area (Table 3.3), which causes higher dissolution. In contrast, the IrOx High SA HT sample produced the lowest dissolved Ir3+ concentration present in the electrolyte at the end of the AST. 97 Figure 5.2 Electrolyte Ir concentration after the ex situ accelerated stress test for OER catalysts (three unsupported IrOx and one TiO2 supported). Experimental conditions: N2 purged 0.09 M H2SO4, 40 °C, 30,000 cycles from 0.05 V to 1.2 V vs. RHE with a 1 s hold at each potential. For direct comparison between the unsupported and supported catalysts, the Ir concentration and loss for the supported catalyst was normalized to 100% IrOx content according to Table 3.3. Heat treatment of the IrOx (i.e., High SA HT) caused a significant decrease in Ir dissolution compared with pristine IrOx High SA. Additionally, despite having almost identical surface area (47.7 and 47.1 m2 g-1 respectively), the IrOx Low SA had higher Ir3+ dissolution than the IrOx High SA HT (Figure 5.2). Previous studies showed that IrOx undergoes a morphological change to a more crystalline phase with heat treatment at 400 °C in air [115], which improves the stability. It is difficult to directly compare IrOx/TiO2 to unsupported IrOx in terms of dissolution. The IrOx/TiO2 catalyst had a BET surface area of 28 m2 g-1 which is lower than any of the unsupported catalysts (Table 3.3). Thus, the high dissolution rate of the supported catalyst compared to any of IrOx High SA IrOx Low SA IrOx High SA HT IrOx/TiO20102030405060 Ir Loss / %Ir Concentration / µg L-10.00.51.01.52.02.53.03.54.0 98 the unsupported catalysts cannot be attributed to the surface area effect (Figure 5.2). It is proposed that it is due to a catalyst/oxide support interaction. At 0.05 V vs. RHE TiO2 undergoes partial reduction to Ti3+ which is a strong reducing agent [126]. This promotes the reduction of Ir4+ to Ir3+ which increases the dissolution of the IrOx. Furthermore, TiO2 is a pH-dependent ion exchanger, which is known to absorb Ir3+. Once the ion exchange capacity of the TiO2 was reached, the Ir3+ concentration in the electrolyte would begin to rise. In order to investigate the structure of the supported catalyst, XRD analysis was carried out. It showed distinct IrO2 and TiO2 crystalline phases as well as an amorphous bump at low angle (Appendix Figure C.1.1). The crystallite size was estimated to be 3.7 nm based on the application of the Scherrer equation to the peak at 2 = 35o. The amorphous IrOx component also contributes to the high dissolution rate observed for the supported catalyst (Figure 5.3). IrOx High SA HT was investigated by XRD (Appendix Figure C.1.4). The heat-treated catalyst had one identified phase (IrO2) with an estimated crystallite size of 4.3 nm (peak at 2 = 35o). Un-heat-treated IrO2 powders did not have identifiable crystalline phases. 5.5 Ex-situ oxygen evolution reaction activity Figure 5.3 presents the OER activities of all the investigated IrOx catalysts at the beginning and at the end of ex situ tests (BOT and EOT, respectively). The activity of IrOx catalysts has been suggested to increase as a function of Ir3+ content. [135] It is expected from XPS results in Chapter 4 that with increasing heat treatment temperature there will be decreased Ir3+ content and increased crystallinity. For the unsupported catalysts, the relationship is robust between the OER activity and Ir dissolution as follows from highest to lowest: IrOx High SA > IrOx Low SA> IrOx High SA HT. Thus, the IrOx High SA has the highest OER mass activity (60.5 A g-1 at 1.6 V vs. RHE) and lowest overpotential at 5 mA cm-2 (0.37 V vs. RHE) while the IrOx High SA HT has the lowest 99 mass activity (19.2 A g-1 at 1.6 V vs. RHE) and highest overpotential (0.50 V vs. RHE). Furthermore, the IrOx Low SA, which is more amorphous and likely has higher Ir3+ content, has higher activity (29.2 A g-1 at 1.6 V vs. RHE) than the IrOx High SA HT in spite of very similar surface areas 47.7 and 47.1 m2 g-1 respectively. Due to increased crystallinity and lower Ir3+ content it is expected based on literature that IrOx High SA HT would have lower activity than the amorphous IrOx Low SA [114,115,137] (Figure 5.3). 100 Figure 5.3 Ex situ OER activity characterization: (a) mass activities of three unsupported IrOx catalysts (Low SA, High SA and High SA HT) and IrOx/TiO2 at the BOT, (b) Mass activities at 1.6 V vs. RHE at BOT and at the end of test (EOT, after 30,000 voltage stepping cycles) (mass activity is given in A g-1 Ir metal), (c) The OER overpotentials at 5 mA cm-2geom at BOT and EOT. (d) Specific activities (mA cm-2) at 1.6 V vs. RHE at BOT and EOT. Temperature 40 °C. The OER activity was also determined after 30k cycles of ex situ AST (i.e., referred EOT activity). All unsupported catalysts experienced small decreases in activity (± < 3 A g-1), whereas IrOx/TiO2 had increased OER activity after 30k cycles (33.8 to 39.0 A g-1, Figure 5.3a). TiO2 is capable of reducing Ir4+ to Ir3+, which increases the OER activity of IrOx catalysts. This could contribute to the 15% increase in activity. Additionally, there could be an increase in electrochemically active 00.20.40.60.811.21.41.6IrOx High SA IrOx Low SA IrOx High SAHTIrOx/TiO2Specific activity at 1.6 V / mA cm-21.3 1.4 1.5 1.6 1.7 1.8 1.904080120160200240 IrOx High SA HT IrOx High SA IrOx Low SA IrOx/TiO2Mass activity / A g-1E / V vs. RHE00.10.20.30.40.50.6IrOx High SA IrOx Low SA IrOx High SA HT IrOx/TiO2ηat 5mA cm-2 / V0.010.020.030.040.050.060.070.0IrOx High SA IrOx Low SA IrOx High SAHTIrOx/TiO2Mass activity at 1.6 V / A g-1BOTEOTcba(BET Surface area / m2 g-1)(Ir3+ dissolution / μg L-1)(91.6)(47.7)(47.1)(28)(33 12)(20.0 0.8)(8.8 0.4)(52 9)d 101 surface area for the supported catalyst due to Ir dissolution and redeposition onto the support. Finally a strong catalyst/support interaction could lead to increased OER activity of the TiO2 supported catalyst. [134,141,233,234] Certain oxide supports have been reported to improve performance durability of OER catalysts and heterogeneous catalysts by changing the crystal structure and oxidation state of the catalyst surface. Having higher mass activity at the end of the test means that anode designs could incorporate less Ir metal therefore have lower cost associated with generating a given anode layer design. In Figure 5.3b, the mass activity is normalized to the mass of the precious metal (Ir) present in each catalyst. Furthermore, comparing the ex situ results of unsupported IrOx in Figure 5.2 and Figure 5.3, it is observed that generally the higher the electrolyte Ir concentration, the higher the OER BOT and EOT activities. In other words, generally the higher the ex situ OER activity, the lower the stability of the IrOx catalyst. The OER mechanism for IrO2 involves the formation of soluble Ir3+ species which are transported away from the surface [125]. 5.6 Fuel cell reversal performance vs. ex situ Ir dissolution and OER activity The IrOx catalysts were also investigated in a single MEA fuel cell according to a protocol described in Section 2.5, devised to determine the anode performance under voltage reversal conditions. The reversal times at the BOT and the end of the AAST test (EOT) are plotted in Figure 5.4a. The reversal loss, defined as the difference between the BOT and EOT reversal times, is a reflection of catalytic durability and is plotted against the Ir concentration of the electrolyte measured after the ex situ AST (Figure 5.4b). The BOT reversal times of the unsupported catalysts (IrOx Low SA, IrOx High SA and IrOx High SA HT) and TiO2 supported IrOx (Figure 5.4a) correlate well with the respective BET surface areas (Table 3.3), i.e., the higher the surface area, the longer the BOT reversal time. Furthermore, the ex situ OER tests revealed the highest BOT 102 mass activity for IrOx High SA (Figure 5.3a), which also translated into the longest fuel cell (in situ) BOT reversal time (Figure 5.4a). For the rest of the catalysts, the correlation between in situ BOT reversal time and ex situ BOT OER mass activity is good with the exception of IrOx/TiO2. The supported IrOx catalyst had a higher ex situ BOT mass activity than either IrOx Low SA or IrOx High SA HT (Figure 5.3a) but had a shorter BOT reversal time. Regarding the EOT performance, there was virtually no correlation between the ex situ mass activity and in situ reversal time (Figure 5.4a). There are a number of differences between the two types of EOT experiments that are responsible for the observed discrepancy. First, it is important to note that the fuel cell AAST protocol is not the same with the ex situ AST. Second, the presence of Pt in the anode catalyst layer can also affect the performance of the OER catalyst, as discussed in section 3.7. Third, the Ir dissolution and redeposition mechanisms are different in the fuel cell and ex situ. In the fuel cell, the dissolved Ir migrates through the membrane to the cathode side assisted by the PFSA acting as a strong ligand (see section 3.5), hence, it is unavailable for redeposition on the anode. These combined effects render the ex situ mass activity EOT data uncertain for predicting the in situ EOT reversal time trends among the various catalysts. With regard to reversal loss (Figure 5.4b), IrOx/TiO2 had the best durability, namely, the smallest reversal loss, which corresponds also to the highest ex situ AST Ir concentration and improved ex situ EOT mass activity after AST (Figure 5.3b). Generally, the higher the ex situ Ir concentration after AST, the higher the reversal loss, with the exception of IrOx/TiO2 (Figure 5.4b). Thus, for unsupported IrOx catalysts, the ex situ Ir concentration after AST is a good predictor of in situ durability as expressed by reversal loss. The poor fuel cell durability of IrOx High SA (91.6 m2 g-1) is not predicted by the ex situ EOT OER activity but is predicted by the ex situ Ir concentration 103 level. Finally, it is also noted that heat treatment of the unsupported IrOx is a viable method to improve the OER durability in the fuel cell (Figure 5.4a and b). Figure 5.4 (a) Reversal tolerance times for fuel cell anodes containing either unsupported OER catalysts (IrOx Low SA, High SA or High SA HT) or supported OER catalyst (IrOx/TiO2) at the BOT and after an in situ AAST (EOT). In parenthesis the ex situ EOT mass activity values from Figure 5.3a are also provided for easier comparison of the in situ (fuel cell) and ex situ performance. (b) Fuel cell reversal loss (min) vs. Ir concentration after ex situ AST. 5.7 Metal cation migration during ex situ accelerated stress test BSE images were taken of an IrOx/TiO2 catalyst freshly deposited on the MPL side of a GDL/MPL and after the ex situ AST (Figure 5.5a and b). The BSE detector increases the contrast between heavy (Ir, Ti, etc.) and light (C, O, etc.) atomic weight elements. The bright spots in the image were confirmed by EDX to contain both Ir and Ti atoms (Appendix Figure C.1.2 and Appendix Table C.1.2). After ex situ AST the catalyst is not well dispersed, which is caused by dissolution and redeposition of the Ir and Ti onto the crack of the MPL. The catalyst relocalization could be due to the cracks having surface sites that are favorable to the anchoring of metal ions and nucleation of the particles and possibly to variations in potential between the cracks and the surface b0200400600800IrOx High SA IrOx Low SA IrOx High SAHTIrOx / TiO2Reversal time / minBOT EOTa0 10 20 30 40 50 60 700200400600800Reversal Loss / minIr Electrolyte Conc. / g L-1Unsupported IrO2IrOx/TiO2(EOT OER mass activityat 1.6 V/ A g-1)(59.6)(26.9)(18.1)(39.0) 104 of the MPL. Cracks of this nature were seen in all of the tested MPL layers. Additionally, the unsupported IrOx catalysts had similar relocalization to the cracks of the MPL. However, it was not as severe as in the case of the IrOx/TiO2. BSE images were also taken of the unsupported IrOx catalysts (Appendix Figure C.1.3). These figures showed a difference in crystallite size (High SA was lower than either Low SA or High SA HT) as well as catalyst migration to the cracks of the GDL/MPL. Figure 5.5 BSE detector images at 500x magnification of (a) freshly deposited and (b) ex situ AST cycled IrOx/TiO2 (81.5 μg cm-2) on a GDL/MPL. The bright dots were confirmed by EDX to contain Ir 5.8 In situ Ir migration across the membrane electrode assembly during fuel cell AAST EDX elemental quantification was carried out for spot and area analysis of in situ crossover. In Figure 5.6, the area of spectrum 1 is between the cathode layer and the MPL. The bright thin region near the top of the image is the anode layer where the Ir based catalyst was originally admixed. EDX quantification can be seen in Table 5.1. The Ir atomic percentage in the area (0.83%) and spot (1.92%) analysis is small compared to the proportion of C from the support and F from PFSA ionomer. The localization of Ir on the cathode side of the MEA can be described by the following mechanism. The Ir3+ originally 105 present in the IrOx based OER catalyst (80.2% Ir4+, 19.8% Ir3+ for IrOx High SA HT; confirmed by XPS) or formed by Ir4+ reduction to Ir3+ during normal operation of the fuel cell anode under H2, migrates in the ionomer environment across the PEM. Ir3+ migration is driven by concentration gradient (chemical potential). The localization of Ir3+ and re-deposition is then determined by the local chemical and potential environment at the cathode. In order to have deposition in a specific location, it is necessary to have both the presence of a liquid phase with Ir3+ and the right local potential conditions to drive nucleation and precipitation of Ir either as metallic Ir or Ir3+/4+ oxide or hydroxide. Figure 5.6 SEM images of an MEA after in situ AAST at two magnifications. An area spectrum (a) and point spectrum (b) were taken by EDX to determine the elemental makeup on the cathode. 106 Table 5.1 Elemental makeup on the cathode determined by EDX. Relocated metal area and spot analysis were averaged from the SEM images in Figure 5.6. Element (shell) Atomic % C (K) 81±1 O (K) 7.6±0.9 F (K) 7±1 S (K) 0.5±0.1 Cl (K) 0.27±0.04 Ir (M) 1.4±0.5 Pt (M) 1.8±0.3 5.9 Ex situ performance of the IrOx/TiO2 OER catalyst combined with hydrogen oxidation reaction Pt catalyst IrOx/TiO2 was combined with carbon supported platinum (Pt/C, 47.1% wt) HOR catalyst and deposited on GDL/MPL for ex situ studies. The mass of the OER and HOR catalyst were 81.5 μg cm-2 and 326 μg cm-2, respectively. As confirmed by XPS the freshly prepared IrOx contains both Ir3+ and Ir4+ species. The presence of Pt in the anode catalyst formulation decreased the Ir dissolution (Figure 5.7a). It is proposed that Pt could catalyze the disproportionation of dissolved Ir3+ to Ir0 and Ir4+ thereby, decreasing the Ir3+ concentration at the surface and reducing the transport of dissolved Ir3+ into the bulk of the electrolyte. Regarding ex situ OER performance (Figure 5.7b), the presence of Pt in the catalyst layer lowered the Ir dissolution (Figure 5.7a) and increased the BOT mass activities. While Pt does contribute to 107 the overall OER activity (only at high potentials above 1.8 V vs. RHE), it is not as effective an OER catalyst as IrOx. Nonetheless, the presence of Pt/C in the layer has the effect of increasing the OER activity of the layer when normalized to Ir metal loading. Furthermore, the ex situ AST caused an increase of the OER mass activities (i.e., EOT activity) for all catalysts. This increase was temperature dependent and was larger in case of IrOx/TiO2 at 60 °C. Comparing the EOT activities at 60 °C suggests that Pt reduces the temperature sensitivity of mass activity increase induced by ex situ AST. Figure 5.7 The ex situ interaction effect between Pt HOR and IrOx/TiO2 OER catalysts. (a) Concentration of Ir in the electrolyte after ex situ AST with and without HOR catalyst. (b) Mass activity at 1.6 V vs. RHE (A g-1Ir). Experiments performed in a flooded (N2 purged 0.09 M H2SO4) three electrode cell 30,000 cycles from 0.05 V to 1.2 V vs. RHE with a 1 s hold at each potential. Catalysts were prepared with either IrOx/TiO2 alone or IrOx/TiO2 and carbon supported platinum HOR catalyst. Tests were run three times each at 40°C (red) and 60°C (blue). 5.10 Conclusions Unsupported and TiO2 supported IrOx OER catalysts were investigated for voltage reversal mitigation performance in a PEM fuel cell environment. An ex situ AST protocol was developed 0.015.030.045.060.075.090.0105.0IrOx/TiO2 w/HOR IrOx/TiO2 w/HORMass activity at 1.6 V / A g-1Ir40 °C 60 °CBOTEOT01234502040608040 40 +HOR 60 60 +HORIr loss %Ir Concentration / µg L-1Temperature / °Ca b 108 to screen and predict the fuel cell (in situ) performance with regard to voltage reversal time and reversal loss (i.e., durability) as determined in the fuel cell AAST. Generally, the ex situ OER activity could be directly related to the resistance of IrOx to dissolution, i.e., more active catalysts released more Ir3+ into the electrolyte solution during AST. Furthermore, the catalyst with the highest ex situ OER mass activity, such as the high surface area unsupported IrOx, produced the longest BOT reversal time in the fuel cell. The BOT reversal time also showed a strong direct correlation with the BET surface areas of IrOx catalysts. Regarding the fuel cell voltage reversal time loss (a way to measure OER durability in the fuel cell environment), generally unsupported catalysts with low ex situ Ir concentration after AST showed better durability, i.e., lower reversal loss. In this respect, the heat-treated high surface area IrOx had the best performance with a reversal loss of 221 minutes and a low ex-situ Ir loss of only 0.5%. The only exception to this trend is the TiO2 supported IrOx, which displayed the lowest reversal loss in the fuel cell during the AAST (approximately 20 minutes), in spite of the highest ex situ Ir3+ dissolution of approximately 3.5% by weight. This could be explained by the changes in the TiO2 support during AAST forming a nonstoichiometric oxide described as: Ti(IV)1−xTi(III)xO(2−x/2)□x/2, where □ is oxygen vacancy. The high Ir dissolution is due to formation of Ti3+ sites on the support acting as reducing agent for IrOx generating dissolved Ir3+. The low reversal loss (i.e., high OER durability of the anode) on the other hand, could be explained by the formation of oxygen vacancies in the TiO2 support during the AAST cycles contributing to the enhanced OER activity for IrOx/TiO2. In separate ex situ experiments it was found that the presence of Pt (the HOR catalysts) decreases the Ir3+ dissolution and increases the OER mass activity. Heat treatment of high surface area unsupported IrOx catalysts was an effective strategy to produce more durable voltage reversal tolerant anodes. This heat treatment could decrease the portion of Ir3+ in 109 the catalyst leading to lower activity by higher durability. Overall, the high surface area heat-treated IrOx catalyst showed the best compromise between BOT reversal time and OER stability in the fuel cell environment. 110 Chapter 6: Bimetallic Sn-Pb catalysts for electroreduction of CO2 to formate 6.1 Introduction ERC for industrial applications requires catalysts with long lifetimes (up to 10,000 hours) and high faradaic efficiency (above 90%). The current standard to perform ERC are Sn catalysts. These catalysts rely on naturally occurring SnO2 layers to maintain high efficiency for ERC to formate. However, SnO2 is only metastable at ERC potentials and it quickly reduces to form metallic Sn. [107] Whereas SnO2 is an excellent catalyst for ERC to formate, metallic Sn is not and results in the electrolysis of water and formation of H2 instead of formate. Therefore, much research effort has been expended to find efficient and long lifetime ERC catalysts. Choi et al. investigated several Sn/Pb bimetallic cathodes (which they refer to as alloys) and found that a near eutectic composition of 56.3% Sn and 43.7% Pb had the best performance in terms of FE and current density for ERC to formate. In Chapter 6, the Choi work is significantly expanded on in this chapter by studying the effect of a new catalyst electrodeposition bath composition (tetrafluoroborate salt) and a new carbon cloth substrate on FE and FE stability. Overall, the work in Chapter 6 adds three new conditions to the preparation of Sn/Pb bimetallic catalysts. Additionally, the mechanism of improved FE stability, which was not covered in the Choi paper, is investigated. 6.2 Pb/Sn catalyst composition Electrodeposited Pb/Sn catalysts supported on carbon cloth and teflonated (5 %wt. PTFE) carbon fiber paper were prepared with five different compositions each. Two different electrodeposition bath compositions were employed: a fluoroborate and an oxide bath (see: Experimental Methods). Table 6.1 and Table 6.2 show the Pb/Sn compositions (as determined by EDX) and calculated mass loadings for all the samples prepared with either the fluoroborate or oxide bath. The loading 111 was calculated assuming 100% current efficiency for depositing the metals. The conditions and compositions explored in Table 6.2 followed the recommendations by Choi et al. [200] The fluoroborate electrodeposition bath is more advantageous due to the higher solubility of Sn and Pb fluoroborate salts compared to oxides. Hence, higher catalyst loadings can be achieved and the post-deposition recycling of the bath is more convenient. Significant variation in the atomic ratio was observed in some samples (±4 to 12%). This could be indicative of non-uniform elemental distribution across the electrode surface. Table 6.1. EDX analysis of Pb and Sn electrodeposited carbon cloth and carbon paper using the fluoroborate bath. Other electrodeposition conditions given in Section 3.5.1. Fluoroborate bath molar composition Atomic ratio Sn:Pb Loading [mg cm-2] Carbon Cloth Carbon Paper Sn100% 100:0±0 100:0±0 111 Sn91%-Pb9% 97.7:2.3±1.5 88.0:12.0±4.9 118 Sn75%-Pb25% 65.2:34.8±4.2 85.5:14.5±2.3 131 Sn16%-Pb84% 1.9:98.1±1.9 3.9:96.1±1.5 176 Pb100% 0:100±0 0:100±0 193 112 Table 6.2. EDX analysis of Pb and Sn electrodeposited carbon cloth and carbon paper using the oxide bath. Other experimental conditions given in Section 3.5.1. Oxide bath molar composition Atomic ratio Sn:Pb Loading [mg cm-2] Carbon Cloth Carbon Paper Sn100% 100:0±0 100:0±0 2 Sn80%-Pb20% 62.0:38.0±0.50 45:55±6.7 3 Sn67%-Pb33% 40.3:59.7±0.19 48:52±12 3 Sn50%-Pb50% 42.3:57.7±1.1 31.0:69.0±1.4 3 Pb100% 0:100±0 0:100±0 4 6.3 Electrochemical properties CV was employed to investigate the electrochemical properties of Sn/Pb catalyst layers. Figure 6.1 displays the CV response of Sn100% and Pb100% catalysts on carbon cloth prepared using the fluoroborate bath (Table 6.1). These CVs are similar to those obtained in previous investigations of Pb and Sn metal. [175,235] In Figure 6.1a, the Sn100% CV in N2 purged KHCO3 shows characteristic peaks for monolayer SnO or Sn(OH)2 formation followed by a further oxidation to SnO2 or Sn(OH)4 corresponding to peaks i and ii respectively. Peaks iii and iv are the corresponding reductions. [235] In Figure 6.1b, the Pb100% CV in N2 has one clear peak for oxidation (i) and two poorly resolved peaks for reduction (ii and iii). Peak i is identified as the oxidation of Pb to PbO or Pb(OH)2 while the two corresponding reduction peaks (ii and iii) are identified as the reduction of PbO back to metallic Pb. [126,236] The Pb CV shows more extensive oxidation and reduction under N2 than the Sn electrode as evidenced by the higher peak current densities. 113 In the presence of CO2, the CV of Sn100% shows a sharp increase in cathodic current at potentials lower than -1.25 V which is due to HER combined with CO2 electroreduction. (Figure 6.1a). After purging with CO2 gas, the pH of the solution decreased from 8.4 to 7.5. The drop in pH causes the Sn100% CV to shift towards positive potentials. A shift of ca. 50 mV is expected according to the Nernst equation for proton concentration dependent waves. [132] Peaks i and ii are shifted accordingly, while oxidative peaks iii and iv are shifted by more than 50 mV suggesting an additional non-Nernstian effect related to electrode kinetics. The HER overpotential on Sn100% is reduced due to the decreased pH and combined with CO2 electroreduction this manifests in the CV as a higher current at -2.0 V. With respect to Pb100% (Figure 6.1b), in the presence of CO2 the current at -2.0 V is smaller than in N2 purged electrolyte, which is contrary to the situation on pure Sn. It is proposed that in the presence of CO2 the surface chemistry of the Pb electrode is changing, with possible surface carbonation, reducing the active sites available for HER and/or CO2 reduction. This aspect will be further investigated. 114 Figure 6.1. Cyclic voltammetry of pure Sn (a) and Pb (b) catalysts generated by electrodeposition on carbon cloth using fluoroborate baths . Legend: (1) red N2 purged and (2) black CO2 purged 0.5 M KHCO3 solution. Potential sweep from 0.1 V to -2.0 V vs. Ag/AgCl at 20 mV/s and 293 K. Figure 6.2 shows the cyclic voltammetry behavior of three bimetallic Pb/Sn catalyst compositions electrodeposited on carbon cloth with the fluoroborate bath. At potentials more negative than -1.6 V the reduction currents in the presence of CO2 for all three bimetallic compositions are higher (in absolute value) than in the case of N2 purged electrolyte. The most significant enhancement of the reduction currents between CO2 and N2 purged electrolytes was observed for Pb84%-Sn16%, where the reduction current at -2.0 V increased by about 70%, followed closely by Sn91%-Pb9% (Fig. 2c and a, respectively). The total reduction currents at -2.0 V (given by both ERC and HER), however, were smaller (in absolute value) than in case of either pure Sn or pure Pb. These results point toward an interaction effect between Pb and Sn increasing the faradaic (or current) efficiency for CO2 electroreduction. Furthermore, the reversal observed in case of pure Pb of lower reduction current with CO2 present at potentials below -1.8 V (Figure 6.1b), was not observed for any bimetallic catalyst (Figure 6.2) indicating the possibility of Sn surface segregation. 115 Next, chronoamperometry experiments at -2.0 V were performed for 2 hours. While CV measures the potential sweep rate dependent instantaneous current response, the 2 hour constant potential experiments are indicative of the longer-term performance. Figure 6.2. Cyclic voltammetry of bimetallic Pb/Sn catalysts generated by electrodeposition on carbon cloth from fluoroborate baths containing the following molar ratios of salts (Table 1): (a) Sn91%-Pb9%, (b) Sn75%-Pb25% and (c) Pb84%-Sn16% (1) red N2 purged and (2) black CO2 purged 0.5 M KHCO3 solution. Potential sweep from 0.1 V to -2.0 V vs. Ag/AgCl at 20 mV/s and 293 K. 116 6.4 Formate faradaic efficiency and total formate Figure 6.3 and Figure 6.4 show the cumulative formate faradaic efficiency and the total moles of formate produced per unit electrode area, for the two types of electrodeposition bath compositions, oxide and fluoroborate, respectively. For both electrodeposition conditions two substrates were employed, carbon cloth and teflonated carbon paper (Table 6.1 and Table 6.2). The experiments were performed in an H-cell with a cation exchange membrane separator to ensure minimal crossover of formate to the anode compartment. The current profiles as a function of time accompanying the faradaic efficiency and formate production data (Figure 6.3 and Figure 6.4), are shown in the Appendix Figure D.1.1. 117 Figure 6.3. Cumulative formate production faradaic efficiency (FE) (a, c) and total moles of formate synthesized per unit electrode area (b, d) for catalysts produced by electrodeposition using the oxide bath. Chronoamperometry at -2.0 V vs. Ag/AgCl in 0.5 M KHCO3, 293 K, 2 hours. Substrates: (a, b) carbon paper and (c, d) carbon cloth. 118 Figure 6.4. Cumulative formate production faradaic efficiency (FE) (a, c) and total moles of formate synthesized (b, d) for catalysts produced by electrodeposition using the fluoroborate bath. Chronoamperometry at -2.0 V vs Ag/AgCl in 0.5 M KHCO3 at 293 K, 2 hours. Substrates: (a, b) carbon paper and (c, d) carbon cloth. For oxide based electrodeposition, the bimetallic composition Sn67%-Pb33% and Sn100% gave the highest faradaic efficiency (about 80%) and formate production on teflonated carbon and carbon cloth, respectively (Figure 6.3). Higher efficiency of 90 to 95% was obtained with the fluoroborate deposited Sn75%-Pb25% catalyst on teflonated carbon paper (Figure 6.4a). The latter faradaic efficiency exceeds the values reported by Choi et al. for bimetallic catalysts. [200] 119 Generally, the teflonated carbon paper substrate promotes increased faradaic efficiency compared to carbon cloth (compare Figure 6.3a and c, Figure 6.4a and c). This could be due to the hydrophobic nature of the Teflon enhancing CO2 adsorption to the surface of the electrode and decreasing flooding of the electrode. This would increase the available CO2 concentration near the catalyst. This is well-represented by the case of Pb100%. On carbon cloth the faradaic efficiency with Pb100% was 20% or below regardless of the deposition method, whereas on teflonated carbon for Pb100%, the faradaic efficiency ranged between 40 to 80% (Figure 6.3c and a, respectively). Regarding the formate production rate (Table 6.3), the catalysts that produced formate at the highest rate were fluoroborate Sn100% on carbon cloth (6.1 ± 0.3 μmol cm-2 min-1) followed by oxide Sn100% on carbon cloth and fluoroborate Sn100% on carbon paper (5.1 ± 0.4 μmol cm-2 min-1). However, the formate faradaic efficiency for these catalysts was lower than for catalysts containing approx. 35 %at. Pb. Also, the longer term stability of Sn100% is poorer compared to selected bimetallics. On teflonated carbon paper, the faradaic efficiency of Sn100%, produced by either the oxide or fluoroborate deposition bath, decreased by as much as 30% over 2 hours of polarization (Figure 6.3a and Figure 6.4a). The stability of Sn100% fluoroborate on carbon cloth was better with a maximum of 10% decrease (Figure 6.4c). In this study the majority of faradaic efficiency loss is attributed to the parasitic HER as demonstrated previously. [200] Hence, the decreased efficiency is accompanied by increased rate of H2 gas evolution on the electrode surface. In contrast to pure Sn, the Pb containing Sn rich bimetallic catalysts such as fluoroborate deposited Sn75%-Pb25% and oxide deposited Sn67%-Pb33% on carbon paper (Table 6.1 and Table 6.2) start at high faradaic efficiencies (95% and 80% in the first 20 min, respectively) and show superior stability over time (Figure 6.3a and Figure 6.4a). The beneficial interaction effect between Sn 120 majority and Pb minority components will be further investigated by surface characterization techniques (see: section 6.6). For all substrates and catalyst deposition methods, pure Pb has the lowest formate generation rate and faradaic efficiency. Table 6.3. Total formate production rate during a 2 hour electrolysis at -2.0 V vs. Ag/AgCl in 0.5 M KHCO3. Electrodeposition Bath Composition Carbon paper Carbon cloth a) Fluoroborate Bath Formate production rate / [μmol cm-2 min-1] Sn100% 5.1±0.4 6.1±0.3 Sn91%-Pb9% 4.7±0.4 4.2±0.1 Sn75%- Pb25% 4.70±0.09 5.0±0.3 Sn16%- Pb84% 3.9±0.3 3.3±0.1 Pb100% 3.63±0.10 0.81±0.09 b) Oxide Bath Sn100% 1.5±0.2 5.6±0.2 Sn80%-Pb20% 2.3±0.1 2.4±0.1 Sn67%-Pb33% 4.3±0.3 1.7±0.1 Sn50%-Pb50% 1.06±0.6 3.1±0.1 Pb100% 0.94±0.09 1.1±0.3 121 6.5 Catalyst changes after formate production Figure 6.5. Fluoroborate deposited Sn100% on carbon cloth: (a) freshly prepared and (b) after CO2 electroreduction at -2.0 V vs. Ag/AgCl in CO2 purged 0.5 M KHCO3 293 K. Catalysts layers were imaged with focused ion beam scanning electron microscopy (FIB-SEM) before and after 2 hours of polarization at -2.0 V vs. Ag/AgCl. Figure 6.5a shows fresh Sn100% on carbon cloth prepared using the fluoroborate bath. The deposit completely covers the fiber surface as a fairly uniform coating which some surface cracking. In Figure 6.5b, the same catalyst was imaged after 2 hours of ERC. The surface shows more cracking with wider gaps between the platelets of Sn, which is due to mechanical stresses induced by hydrogen blistering and embrittlement due to parasitic HER. During the 2 hour experiment no 122 large scale catalyst detachment is visible which would expose the carbon cloth substrate. Polycrystalline Sn with Miller indices (2,1,1), (2,2,0), (2,0,0), (1,0,1) were identified by XRD as the main components of the catalyst after 2 hours CO2 electroreduction (Appendix Figure D.1.2). Figure 6.6 Fluoroborate deposited Pb100% on carbon cloth: (a) freshly prepared and (b) after CO2 electroreduction at -2.0 V vs. Ag/AgCl for 2 hours. The fluoroborate Pb100% catalyst on carbon cloth (Figure 6.6) also shows significant morphology changes with polarization. The freshly prepared catalyst surface contains disk-like Pb deposits, which change into thin blade-like structures with polarization. Large crystals that were converted to blade like structures after electrolysis were seen in Sn16%-Pb84% catalysts deposited without triton (Appendix Figure D.1.7). The blade-like morphology is related to the complex chemical 123 composition of the catalyst with different lead carbonate and oxide species identified by XRD (Appendix Figure D.1.4). Appendix Figure D.1.4 shows XRD peaks due to crystalline Pb with Miller indices (1,1,1), (2,0,0), (2,2,0) and (3,1,1) together with lead oxide (PbO) and massicot (i.e. PbO with orthorhombic structure) as major components. In addition, hydrated lead carbonate such as plumbonacrite (Pb5(CO3)3O(OH)2) was also identified. Similarly, on the Pb84%-Sn16% catalyst surface as well, oxides and carbonates are present (Appendix Figure D.1.5). The formation of these carbonate and oxide species are responsible for the poor performance of the Pb100% catalyst by blocking the active sites (Figure 6.4c). 6.6 Low Pb content efficiency stabilization Figure 6.8 and Figure 6.7 show the XRD spectrum and the SEM images, respectively, for Sn75%-Pb25% on carbon cloth. The freshly deposited catalyst is multifaceted with a petal-like geometry (Figure 6.7a). This catalyst shows extreme changes after CO2 reduction; it becomes sponge-like and deeply pitted with large cracks on the surface (Figure 6.7b). 124 Figure 6.7 Fluoroborate deposited Sn75%-Pb25% on carbon cloth: (a) freshly prepared catalyst and (b) after CO2 electroreduction at -2.0 V vs. Ag/AgCl for 2 hours. However, the catalyst morphology changes did not have a negative effect on performance, the faradaic efficiency was stable at approximately 70% after 40 min. (Figure 6.4c). The apparent increase of the surface area could have compensated some of the intrinsic catalytic activity loss over time. The XRD spectrum of the used catalyst reveals polycrystalline Sn and SnO2 as the major components with Pb in solid solution Sn phase (Figure 6.8). The presence of SnO2 and its proposed electrocatalytic role in stabilizing the radical anion CO2- [100] could explain the higher faradaic efficiency of this catalyst compared to either pure Sn or pure Pb (Figure 6.4a and c). The naturally occurring SnO2 film should be reduced over time to Sn(0) at highly negative potentials similar to 125 the Sn100% or Sn91%-Pb9% cases (Appendix Figure D.1.6). However, as shown by Figure 6.8, the presence of SnO2 after 2 hours of ERC demonstrate that it is stabilized by approximately 35 %at. Pb content and SnO2 reduction is minimized. Figure 6.8 XRD spectrum of fluoroborate deposited Sn75%-Pb25% on carbon cloth after polarization for 2 hours at 2.0 V vs. Ag/AgCl in 0.5 M KHCO3 at 293 K. Peak labels are for Miller indices of crystalline Sn (red, S), PDF2 01-072-1012 (ICDD, 2016) and crystalline SnO2 (blue, O), PDF2 00-004-0673 (ICDD, 2016). The mechanism of this stabilization is likely due to the oxidation of Sn0 by PbO2. The standard reduction potential of the reduction of PbO2 to PbSO4 is higher than the standard reduction potential of Sn4+ to Sn2+ and Sn2+ to Sn0 (Table 6.4).This makes it thermodynamically possible for the PbO2 seen in the XRD of Pb catalysts to oxidize Sn0 to the catalytically active SnO2. 60 40T01-072-1012 (*) - Tin Oxide - SnO - Y: 31.98 % - Tetragonal - a 3.80290 - b 3.80290 - c 4.83820 - Primitive - 4/nmm (129) - 2 - 69.9703 - I/Ic PDF 15.3 - S-Q 8.1 % - F30=1000(0.0000,33) - alpha 90.0000-004-0673 (*) - Tin, syn - Sn - Y: 46.57 % - Tetragonal - a 5.83100 - b 5.83100 - c 3.18200 - Body-centered - I41/amd (141) - 4 - 108.190 - I/Ic PDF 2. - S-Q 91.9 % - F29= 72(0.0129,33) - alpha 90.000File: 60 40T.raw - Start: 4.885 ° - End: 89.923 ° - Step: 0.030 ° - Step time: 145.6 s - Anode: Cu - WL1: 1.5406 - WL2: 1.54439 - kA2 Ratio: 0.5 - Generator kV: 40 kV - Generator mA: 40 mA - Creation: Lin (Cps)010203040502-Theta - Scale5 10 20 30 40 50 60 70 80 90O (1,1,0)O (1,1,2)0 (2,2,1)O (1,0,1)S (2,0,0)S (1,0,1)S (2,2,0)S (2,1,1) 126 Table 6.4 Standard reduction potential of Sn and Pb in aqueous solutions at 25 °C (V vs. NHE). [132] Reaction Potential / V Sn4+ → Sn2+ + 2e- 0.15 Sn2+ → Sn0 + 2e- -0.375 PbO2 + SO42- + 4H+ → PbSO4 + 2H2O 1.698 6.7 Conclusions The electrocatalytic performance of five different Sn-Pb catalysts compositions was investigated toward CO2 electroreduction to formate. The catalysts were prepared by electrodeposition on either teflonated (5 %wt.) carbon paper or un-teflonated carbon cloth using both fluoroborate and oxide based deposition media, respectively. Thus, interaction effects involving the catalyst composition, loading and carbon substrate could be studied. The electrocatalytic activity was assessed by cyclic voltammetry combined with chronoamperometry at -2 V vs. Ag/AgCl for 2 hours. While typically, Sn100% catalysts deposited on either carbon cloth or teflonated carbon paper performed the best in terms of HCOO- production rate, the faradaic efficiency of these catalysts dropped by as much as 30% over a 2 hour period. Both low and high loading Sn100% catalysts, prepared by oxide and fluoroborate based deposition bath, respectively, experienced rather similar degradation rates. The Sn100% deposited on carbon cloth had somewhat better durability than on carbon paper, with faradaic efficiency losses of about 10 to 15% in 2 hours. The faradaic efficiency degradation for pure Sn is due to reduction of SnO2 to metallic Sn during ERC either by direct electron transfer or by adsorbed hydrogen. In contrast, bimetallic Sn rich catalysts containing between approximately 15 and 35 %at. Pb generated higher and more stable faradaic efficiencies. The highest faradaic efficiency, between 127 90 to 95% over the 2 hour period was obtained with the fluoroborate bath deposited catalyst supported on teflonated carbon paper and containing approximately 35 %at. Pb (Table 1, Fig. 4a, Sn75%-Pb25% salt molar composition). During electroreduction, the catalyst morphology changed to a spongy pitted surface containing also crystalline SnO2 (1,0,1), (1,1,0), (1,1,2) and (2,1,1) as identified by XRD. The beneficial catalytic role of SnO2 in CO2 electroreduction has been described previously in the literature[107] and it could explain the superior performance of the bimetallic formulation. These findings invite further research into better understanding the stabilization mechanism of SnO2 for highly active and durable CO2 reduction electrocatalyst synthesis. 128 Chapter 7: Conclusions and recommendations for future work 7.1 Electroreduction of CO2 and PEMFC to combat climate change Climate change is the most important issue facing the world today. To address and correct this troubling trend, the global community should take a multipronged approach using several technologies to find an effective overall solution. It was the goal of this thesis further the fight against climate change through electrochemistry. Specifically, the main goal in this work was to improve the efficiency and durability of ERC and OER catalysts in renewable energy systems. Improving catalyst lifetime and efficiency makes these technologies more affordable to potential consumers. To improve efficiency, this thesis has focused on using catalysts to reduce their chemical reaction energy barriers. To improve durability, it was investigated how different catalyst formulations and supports affect the dissolution of metals and the faradaic efficiency stability over time. All of this work is focused on simple methodology to produce impactful and insightful results as quickly and efficiently as possible. PEMFC manufacturers in particular have expressed a need for better methodology to identify OER catalyst candidate materials. OER catalysts can improve the overall lifetime of fuel cell stacks. This lifetime has to be better than their competition in order to make a meaningful impact in the market. Of course, this is not just controlled by the cost and lifetime of automotive fuel cells. The availability of hydrogen and the overall hydrogen infrastructure is another major limitation for widespread use of this technology for consumers. However, stack durability is still an essential part of the affordability of these systems and should be at least as good as the internal combustion engines. OER catalyst durability in PEMFCs is particularly important for the stability of the carbonaceous components during start-up/shut-down cycles. The start-up/shut-down stability is especially important in automotive applications where the stack will be unused to long periods of 129 time. This thesis contributed to improving OER catalyst testing methodology by investigating changes to Ir dissolution in a simulated PEMFC environment. Ir dissolution was correlated to voltage reversal data from a single cell test setup. Additionally, PEMFC environment approximation was improved by adding a Pt/C catalyst. This system was used along with the EQCM to gain insight into the potential dependent mass changes occurring at the anode catalyst. The ERC work in this thesis focused on increasing catalyst durability. The situation with Sn catalysts is very poor in terms of faradaic efficiency stability over time and the approach of this work included adding a second metal, changing the loading and using two different carbon support materials. Whereas the combination of Sn and Pb for use in ERC had been investigated once previously to this thesis, [200] an important mechanistic explanation was uncovered to explain the observed faradaic efficiency stability. 7.2 EQCM measurement of IrOx and Pt/C conclusions The main conclusions of the EQCM measurement of potential depended IrOx and Pt/C mass changes investigation are outlined below: The potential dependent mass changes of IrOx powders were investigated for the first time using EQCM in a simulated PEMFC anode environment. Heat-treated (at 350, 450 and 550 °C) and untreated IrOx catalysts exhibited different behavior in terms of frequency changes during potential cycling. This behavior was identified for the first time. The explanation for this difference in behavior is that the frequency response was caused by the potential dependent uptake of water and formation of oxyhydroxides with different compositions. XPS analysis confirmed that heat treatment reduced the oxyhydroxide composition of the catalyst after CV cycling as compared to untreated IrOx. The first EQCM based reversal tolerant anode system was also developed. During potential holds at 1.8 and 2.0 V vs. RHE, Pt/C exhibited large frequency increases caused by carbon 130 corrosion. This carbon corrosion was protected during the same potential holds by the addition of OER catalyst to the layer. 7.3 Ex situ accelerated stress test for reversal tolerant anodes conclusions Main conclusions of the ex situ AST for OER catalyst investigation are outlined below: For the first time, an ex situ accelerated stress test was compared to in situ reversal time data generated from a single cell test station. Unsupported catalysts with low ex situ Ir concentration after AST showed better durability and lower in situ reversal time loss. Heat-treated high surface area IrOx had the best performance with a reversal loss of 221 minutes and a low ex-situ Ir loss of only 0.5%. The TiO2 supported IrOx, had the lowest reversal time loss in the fuel cell during the AAST (ca. 20 minutes), and the highest ex situ Ir3+ dissolution of approximately 3.5% by Ir mass. The difference between unsupported and metal oxide supported IrOx catalysts could be explained by one of two mechanisms. First, the nonstoichiometric oxide described as: Ti(IV)1−xTi(III)xO(2−x/2)□x/2, where □ is oxygen vacancy could be formed. The low reversal time loss could be explained by the formation of oxygen vacancies in the TiO2 which would improve the stability of OER activity for IrOx/TiO2 even though there is significant Ir dissolution. Second, the high Ir dissolution could be caused by formation of Ti3+ sites on the support acting as reducing agent for IrOx generating dissolved Ir3+. This would cause high Ir3+ which both increases the dissolution of Ir and also improves OER activity (due to more electrophilic oxygen in the Ir3+ catalyst) In separate ex situ experiments, the presence of Pt (the HOR catalysts) in the IrOx catalyst layer caused a decrease both the Ir3+ dissolution and the OER mass activity. Heat treatment of high surface area unsupported IrOx catalysts was an effective strategy to produce more durable voltage 131 reversal tolerant anodes. The heat treatment lowered the dissolution of Ir3+ ex situ and decreased the voltage reversal time loss in situ by oxidizing Ir3+ to Ir4+ and increasing IrOx crystallinity. 7.4 Bimetallic Sn-Pb catalysts for electroreduction of CO2 conclusions Main conclusions of the bimetallic Sn-Pb investigation are outlined below: Similar to previous studies, Sn100% catalysts began with high faradaic efficiency which dropped during 2 hour potentiostatic experiments Pb performed the best in terms of faradaic efficiency stability but had the lowest absolute faradaic efficiency. The catalyst with the highest faradaic efficiency, between 90 to 95% over the 2 hour period, was with the fluoroborate bath deposited catalyst supported on teflonated carbon paper and containing approximately 35 %at. Pb. In this work, I was able to find a possible mechanism for the improved faradaic efficiency stability of Sn/Pb catalysts. Sn catalysts containing 35 %at. Pb contained crystalline SnO2 as identified by XRD. SnO2 was not identified in any other catalyst composition even before catalysis. Pb oxides and oxide carbonates were identified in Pb catalysts and it is suggested that Pb oxides and oxide carbonates are able to stabilize the metastable SnO2. 7.5 Contributions to knowledge This thesis was able to make several significant contributions to scientific knowledge. These contributions were focused on the durability of electrocatalysts in PEMFC anodes and ERC cathodes. The main contribution of the PEMFC anode work was the confirmation that Ir dissolution can be used as a metric to rank OER catalysts prior to in situ testing. This finding creates the opportunity to select OER catalysts quickly before using them in costly MEA fuel cell experiments. The second major finding arising from this work is that the addition of Pt/C causes less dissolution of Ir into the electrolyte after the ex situ AST. 132 In the EQCM work, it was shown that the addition of IrOx to a Pt/C layer can protect the layer from carbon corrosion at high potentials. This method can be used to test OER catalysts that become unstable at high potentials such as RuO2 based catalysts. In order to make this a viable testing method, the EQCM crystal would need withstand higher current in the high potential region. The ERC work in this thesis was a replication of a previous report that Sn/Pb catalysts have higher faradaic efficiency stability than Sn or Pb alone. While, replication studies are not given much priority in modern science, many publications have shown that replication studies are important to show that an observed effect is not caused by some outside uncontrolled variable. [237–239] This work strengthens the link between Sn/Pb bimetallic catalysts and improved faradaic efficiency for ERC. The major contribution to knowledge that arose from this work is that SnO2 remains stable after ERC on Sn majority catalysts with 15 to 35 %at. Pb. This was not previously identified as the mechanism for improved faradaic efficiency stability. 7.6 Recommendations for future work The time and material limitations of this thesis meant that much work was left undone. This was especially true in the ERC work, where extensive scale up testing was initially planned but not carried out to fruition. This ERC work was initially outlined in a research proposal and is presented here along with suggested future work for the PEMFC OER catalyst and EQCM sections. 7.6.1 EQCM oxygen evolution reaction catalyst future work Pt/C was added to the catalyst layers and performed the same CV stress test and there was a large frequency increase observed. In fact, the increase in frequency corresponded to a theoretical mass decrease larger than the total deposited mass of Pt/C. This mass decrease could be caused by viscoelastic changes in the carbon particle aggregates because a similar mass decrease was seen 133 when CV stress testing Ketjen black under the same conditions. The presence of Pt did increase the magnitude of mass changes whereas the addition of IrOx decreased the magnitude. Further investigation, including physical modelling of the viscoelastic interactions of the Pt/C and IrOx particles would be necessary to better understand the observed frequency response. During potential holds at 1.8 and 2.0 V vs. RHE, a large mass decrease was observed caused by the corrosion of Pt/C. The addition of IrOx to the film protected the Pt/C from corrosion effectively creating an ex situ voltage reversal tolerant anode. Potential holds at ca. 1.8 to 2.0 V vs. RHE could be used to determine the stability of OER catalysts in PEMFC anode conditions. The limitations to this technique are that the EQCM crystal itself should be robust enough to carry current to the face to the electrode without delaminating from the crystal surface. This could be achieved by making special EQCM crystals with thicker conductors (typically made of titanium). This type of testing would be especially useful for Ru based OER catalysts. Ru forms volatile RuO4 at high potentials, which is removed from the surface of the electrode. Therefore, the durability of Ru based catalysts mixed with Pt catalysts could be investigated at high potentials. The same technique could not be used to effectively investigate Ir based OER catalysts because of their high stability at potentials of ca. 2.0 V vs. RHE. Further studies of heat-treated IrOx catalysts mixed with Pt/C should be carried out. This would investigate the mechanism of how heat-treated catalysts improve the voltage reversal protection of Pt/C. This should be performed at high potentials >1.8 V vs. RHE and under constant current conditions. Additionally, simulated drive cycles could be performed with “voltage reversal” events at certain intervals to further investigate the durability of OER catalysts ex situ. The mechanism of OER catalyst protection of Pt/C should also be investigated. This would require a mass balance of the catalyst layer before and after CV cycling and analysis of ECSA changes due to Pt dissolution, 134 detachment, Oswald ripening, and coalescence. Finally, an investigation of pH effects on OER catalyst mass changes and dissolution during CV cycling could produce interesting findings relevant to the conditions at the PEMFC anode. Modifying the heat treatment procedure to use different atmospheres such as N2 or H2 could generate IrOx catalysts with different properties. Reducing environments would likely generate IrOx with higher proportions of Ir3+ and metallic Ir in the catalyst. Subsequent investigation of these powders using EQCM would further connect the preparation of OER catalysts to their frequency response during CV cycling. 7.6.2 PEMFC oxygen evolution reaction catalyst future work Further investigation of OER dissolution migration in real fuel cell stacks would greatly compliment this work. This would determine if the simulated PEMFC environment had the necessary components to approximate the real PEMFC anode accurately. The Pt/C addition and subsequent decrease in Ir dissolution should be studied in more detail as well. This is the most interesting research direction for this work in the future because the real PEMFC anode contains Pt/C. Therefore, understanding the interaction between Pt/C and different supported and unsupported OER catalysts would be important and interesting not only from an academic point of view but from an industrial perspective as well. This work could be carried out using the method described in Chapter 5. Additionally, sensitive surface techniques could be used to investigate the chemical state of the catalysts before and after the ex situ AST. Specifically, XPS would be a useful addition to determine the oxidation state of Ir compounds. Modifications to the method used in Chapter 5 to account for the difference in chemical environment during normal operation and cell shutdown should be investigated as well. This could be done by purging solutions alternatively with H2 and O2 during potential holds at 0.05 and 1.2 V 135 vs. RHE, respectively. The reducing and oxidizing chemical environments could increase the dissolution rate of Ir3+ into the electrolyte. Increasing cell temperature should be investigated as well. While the 40 °C temperature was useful to test conditions during the start-up of a cell where cell potentials are relatively low, it would be useful to investigate Ir3+ dissolution at elevated temperature, ca. 80 °C. This would more closely approximate the temperatures expected during normal PEMFC operation. The fabrication of a specialized cell would be needed to prevent electrolyte evaporation during experimentation and to reducevariation in Ir3+ electrolyte concentration. Additionally, dynamic vapor sorption could be used to probe the hydration of catalysts in different humidified conditions (and elevated temperatures) and compare to the hydration occurring in the flooded three electrode cell. Finally, investigating other supports including MMO and carbon supports would be a valuable addition to the literature. This would further strengthen the link between the support effect and OER activity/dissolution. 7.6.3 Electroreduction of CO2 future work Global warming is the most pressing problem in the world today. It is essential that the global industrial carbon footprint be eliminated to prevent a global climate catastrophe. Removal of atmospheric CO2 and sequestration to underground wells is part of the solution to this problem. However, this approach lacks financial incentive for companies to reduce their carbon emissions decreasing the probability of practical implementation [75,76]. A viable alternative is needed to curb CO2 emissions. ERC is a controllable process that can produce various products based on the electroreduction system used. The tunability of ERC to fit a particular industrial application is one of the strengths of the process which could make it widely applicable. Future implementation of 136 ERC will hinge on the ability to generate longer lasting robust and efficient catalysts capable of ERC at industrially viable current densities. Future work should further investigate synergistic effects of bi- and tri-metallic electrodes. While much work has been done on single metal electrodes, there is much more that can be discovered by investigating multi-metal electrodes. Changes to the activity of a given cathode should be described and the underlying mechanisms of ERC should be elucidated. The causes of catalyst efficiency breakdown will also be a crucial point of future work. Using this data it would be possible to make better and longer lasting catalyst materials. Future work should also include the effect of different catalyst support materials including carbon, and metal mesh on activity. Hydrophobic polymer coatings could be used to concentrate CO2 at the electrode surface. This would increase the maximum current density achievable at the electrode surface. Additionally, it could increase the longevity of the electrode surface by preventing formation of non-conducting deposits at the active sites. These new catalysts could be investigated in a gas diffusion cell setup. Gas diffusion cells hold promise because of the possibility of higher current densities and ERC rates. A flow cell apparatus could be assembled with the goal of current density near 100 mA cm-2 and faradaic efficiencies of up to 90%. Flow rate, pressure and cell design should be found to determine the optimal working conditions for the flow cell. Importantly, it should be demonstrated that the novel catalysts have increased longevity in terms of maintaining high faradaic efficiency. This would ultimately bring this technology to the point of industrial application and would be achieved through electrochemisty and chemical engineering. 137 7.6.4 Gas diffusion electrodes In order to overcome mass transfer limitations, gas diffusion cells can be used to achieve high current densities [105]. A possible schematic of a gas diffusion electrode can be seen Figure 7.1. Figure 7.1. Schematic of a continuously operating gas diffusion electrode setup from Bitar et al. [240] A gas diffusion half-cell can test the efficacy of electroreducing pure gas phase CO2 to HCOO-. The gas diffusion electrode can be fabricated utilizing commercially available gas diffusion electrode materials. On one side of the catalyst would be a stream of pressurized CO2 and on the other an aqueous electrolyte. The gas diffusion electrode can be customized both from the commercial manufacturer and in the laboratory by adding additional coatings or metallic catalysts. The main advantage of the GDE would be the ability to reduce CO2 at an electrode surface directly rather than in the electrolyte solution. Using a gas diffusion configuration could lead to a higher current density process because the process would not limited by the solubility and diffusion of 138 CO2 to the surface of the cathode. The discussion of mass transfer can be seen in the section describing the H-cell configuration (section 3.5.3). 7.6.5 Flow cell Flow cell experiments are needed to demonstrate the scale up of CO2 electroreduction to industrially scale. Current densities of 100 mA cm-2 and faradaic efficiencies of 90% need to be targeted for an industrial process. The essential factor that must be improved is the lifetime of the catalyst material which should be at least 100 hours. Using Sn-Pb bimetallics could be a big step towards the industrial goal of 10,000 hours of catalyst lifetime without significant (~10%) loss of efficiency. It was demonstrated in Chapter 6 that Faradaic efficiency improvements were achievable using Sn/Pb electrodes. However, further long lifetime investigations are needed. Additionally, to optimize the system, several variables would be tested including, pressure in the system (from 1 atm to 30 atm) and flow rate, catalyst thickness and metal loading. Increasing the flow of electrolyte to the catalyst surface will allow us to reduce the mass transfer overpotential. Increasing the pressure would increase the CO2 concentration in the electrolyte. Reducing the thickness and resistance of the components would decrease IR drop in the cell and allow us to obtain higher thermodynamic efficiently. 7.6.6 Nanostructures Nanostructured catalysts have come into intense research interest in recent years with the advent of sensitive surface imaging techniques and nanostructure generating procedures. The efficiency of these nanostructured catalysts has been wide ranging with the efficiency reported varying depending on how the catalyst was generated. This suggests that there is a strong relationship between the surface morphology of the catalyst material and the rate and efficiency of catalysis. Zhang et al. generated catalysts utilizing high surface area (carbon black and grapheme) and SnO2 139 nanoparticles. It was determined that the efficiency was best (>93%) using 5 nm particles in an aqueous NaHCO3 solution. Importantly, they found that the catalyst was stable for as long as 18 hours. [181] These findings could be used to optimize Sn-Pb electrodes for ERC. Varying the electrodeposition conditions including the bath composition and addition of surfactants will change the electrode surface characteristics. Additionally, performing pulsed deposition could control the size of the particles formed by controlling the number of nucleation sites on the substrate. This along with oxidizing pretreatment with nitric acid (HNO3) for example could change the oxidation state of the electrode surface. This will need to be applied carefully as surface oxidation could form a passivation layer reducing the conductivity of the electrode. 7.6.7 Bi and trimetallic cathodes ERC can be further improved by the development of cathode materials to reduce overpotential and increase faradaic efficiency. This can be achieved through bi- and tri-metallic catalyst electrodes while controlling catalyst surface roughness and nano-scale surface features. Metals that could be considered for this catalyst generation include the post transition metals: Sn, Pb, In, Bi. One tri-metallic catalyst has been investigated for ERC by Kumawat et al. [202] Experiments on bi- and tri-metallic cathode materials should out using also factorial based experimental design to determine the best conditions under which to perform ERC. 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E° = 0.223 V IrO2 + 4H+ + e- → Ir3+ + 2H2O IrO2 + 4H+ + 4e- → Ir + 2H2O Ir3+ + 3e- → Ir Appendix Equation A.1 E° = 0.935V Appendix Equation A.2 E° = 1.156V Appendix Equation A.3 159 Appendix B B.1 IrOx EQCM supplemental information Appendix Figure B.1.1 STEM images of pristine (a) Pt/C and CV cycled Pt/C with (b) untreated IrOx and (c) IrOx heat-treated at 550 °C. Images were taken at 630 kx magnification. Significant electron beam induced polymerization of the carbon support and other contaminants occurred while capturing the CV cycled images (b and c) which causes the blurry appearance of the images and prevents comparison of the carbon structure. 160 Appendix Figure B.1.2. SEM images of untreated IrOx (a, b) and IrOx 350 °C (c, d) each catalyst before (a, c) and after (c, d) the stress test (CV cycles from 0.05 to 1.2, 1.4 and 1.6 V vs. RHE, 30 times each). 161 Appendix Figure B.1.3. SEM images of IrOx 450 °C (a, b) and IrOx 550 °C (c, d) each catalyst before (a, c) and after (c, d) the stress test (CV cycles from 0.05 to 1.2, 1.4 and 1.6 V vs. RHE, 30 times each). 162 Appendix Figure B.1.4. XPS survey scans of untreated IrOx (a) pristine and (b) after the CV stress test (CV cycles from 0.05 to 1.2, 1.4 and 1.6 V vs. RHE, 30 times each 100 mV s-1) in N2 purged 0.09 M H2SO4 at RT. baIrO2 Untreated IrO2 untreatedAfter CV stress test02004006008001000O 2sIr 4f 5/2Ir 5sIr 5p 3/2Ir 4d 5/2Ir 4d 3/2Ir 4p 3/2Ir 4p 1/2Ir 4sC 1sO 1sIntensity / counts s-1 Binding Energy / eVO KL11000 800 600 400 200 0F KL1Au KL1Au 4p 1/2Au 4f 5/2Au 5sIr 4f 5/2Au 4d 5/2Au 4d 3/2Au 4p 3/2F 1sC 1sO 1sO KL1Sulfur regionIntensity / counts s-1Binding Energy / eVIr 4d 3/2Ir 4d 5/2 163 Appendix Figure B.1.5. Current response of Pt/C (a) alone and with (b) untreated IrOx and (c) IrOx heat-treated at 550 °C. Catalyst mixtures were 2:1 Pt:Ir metals basis. CVs were recorded in 0.09 M H2SO4 purged with N2 at 100 mV s-1 at 25 °C. 0.0 0.3 0.6 0.9 1.2 1.5 1.8-1.5-1.0-0.50.00.51.0E / V vs. RHEi / mA cm-20.0 0.3 0.6 0.9 1.2 1.5 1.8-0.8-0.40.00.40.8E / V vs. RHEi / mA cm-20.0 0.3 0.6 0.9 1.2 1.5 1.8-1.5-1.0-0.50.00.51.01.5E / V vs. RHEi / mA cm-2ab cPt/CPt/C Unt Pt/C 550 164 Appendix Figure B.1.6. EDS mapping image of Pt/C + untreated IrOx after CV stress test in 0.09 M H2SO4 purged with N2 at 100 mV s-1. Oxygen (blue), Pt (yellow) and Ir (red) color intensity is proportional to their respective EDS signal strengths. 165 Appendix Figure B.1.7. The frequency (a) and CV (b) response of an Au EQCM crystal CV cycled from 1.2 (black), 1.4 (red) and 1.6 V (blue) to 0.05 vs. RHE, 30 times each at 100 mV s-1. Electrolyte: N2 purged 0.09 M H2SO4. 25 °C. 0.0 0.3 0.6 0.9 1.2 1.5 1.8-0.4-0.3-0.2-0.10.00.10.20.3E / V vs. RHEi / mA cm-20.0 0.3 0.6 0.9 1.2 1.5 1.8-18-15-12-9-6-303mass / ngf / HzE / V vs. RHE120100806040200-20ba 166 Appendix Figure B.1.8. Deconvolution of XPS peaks in the Ir region of pristine IrOx (a) untreated and heat-treated at (b) 350 °C (c) 450 °C (d) 550 °C. Ir 4f/5Residual STD = 7.51985Ir 4fx 102102030405060CPS72 69 66 63 60Binding Energy (eV)Ir 4f/8Residual STD = 8.01889x 102102030405060CPS72 69 66 63 60Binding Energy (eV)Ir 4f/8Ir 4fResidual STD = 7.78288x 10210203040506070CPS72 69 66 63 60Binding Energy (eV)Ir 4f/8Residual STD = 5.9539x 102102030405060CPS72 69 66 63 60Binding Energy (eV)Untreated 350 C550 C450 CadcbIr4+Ir3+ Ir3+Ir4+ 167 Appendix Figure B.1.9. Deconvolution of XPS peaks in the Oxygen region of pristine IrOx (a) untreated and heat-treated at (b) 350 °C (c) 450 °C (d) 550 °C. O 1s/6O 1sResidual STD = 0.814638x 102253035404550556065CPS536 534 532 530 528 526Binding Energy (eV)O 1s/3Residual STD = 0.868114x 10225303540455055606570CPS536 534 532 530 528 526Binding Energy (eV)O 1s/6O 1sResidual STD = 1.1095x 1022530354045505560CPS536 534 532 530 528 526Binding Energy (eV)O 1s/6O 1sResidual STD = 1.27867x 10225303540455055606570CPS536 534 532 530 528 526Binding Energy (eV)Untreated 350 C550 C450 CadcbH2OOHO2 168 Appendix Figure B.1.10. Deconvolution of XPS peaks of IrOx (a-Ir and b-Oxygen) untreated and (c-Ir and d-Oxygen) heat-treated at 550 °C after potential cycling from 1.2, 1.4 and 1.6 to 0.05 V vs. RHE in 0.09 M H2SO4. Ir 4f/7Residual STD = 2.40177x 102101214161820CPS70 68 66 64 62 60Binding Energy (eV)O 1s/5O 1sResidual STD = 0.787842x 10260626466687072CPS536 534 532 530 528Binding Energy (eV)Ir 4f/9Residual STD = 1.75801x 10210121416182022CPS69 68 67 66 65 64 63 62 61 60Binding Energy (eV)O 1s/7Residual STD = 0.959222x 10264666870727476788082CPS536 534 532 530 528Binding Energy (eV)UntreatedIrUntreated Oxygen550 COxygen550 CIradcbH2OOHO2Ir4+Ir3+ Ir3+Ir4+ 169 Appendix Table B.1.1. XPS quantification of Ir3+/4+ and oxygen environments (IrO2, IrOx(OH)y and H2O) in pristine and used (after potential cycling from 1.2, 1.4 and 1.6 to 0.05 V vs. RHE in 0.09 M H2SO4) heat treated IrOx (untreated, 350, 450 and 550 °C). Atomic % of Ir or O in each form Pristine Ir3+ Ir4+ O2 OH H2O Untreated 44 56 22 39 39 350 °C 43 57 18 48 34 450 °C 42 58 22 43 35 550 °C 41 59 23 40 37 Used Untreated 52 48 4 42 54 550 °C 48 53 8 3 88 170 Appendix C C.1 IrOx ex situ characterization Appendix Table C.1.1 Concentration of metals in the electrolyte after the ex situ AST. Experimental conditions: N2 purged 0.09 M H2SO4, 30,000 cycles from 0.05 V to 1.2 V vs. RHE with a 1 s hold at each potential. Electrolyte concentration / mg L-1 Limit of Detection Blank 0.09 M H2SO4 IrOx Low SA IrOx High SA IrOx High SA HT IrOx/TiO2 Titanium 0.002 <0.002 <0.002 <0.002 <0.002 0.003 Aluminum 0.001 0.001 0.053 0.023 0.021 0.015 Boron 0.002 <0.002 <0.002 <0.002 0.008 0.028 Chromium 0.00005 <0.000050 <0.000050 0.000500 0.000220 0.028255 Cobalt 0.00002 <0.000020 <0.000020 0.000400 0.000090 0.000229 Copper 0.0005 <0.0005 <0.0005 0.0210 0.0032 0.0066 Iron 0.002 <0.002 <0.002 0.050 0.020 0.123 Lead 0.00001 <0.00001 0.00053 0.00110 0.00034 0.00038 Manganese 0.001 <0.001 <0.001 <0.001 <0.001 0.015 Nickel 0.0002 <0.0002 0.0057 0.1359 0.0069 0.0174 Silver 0.00001 <0.00001 0.00066 0.00187 0.00023 0.00244 Tin 0.0001 <0.0001 0.0053 0.0057 0.0017 0.0024 Zinc 0.0005 <0.0005 0.0380 0.0473 0.0230 0.0131 Calcium 0.01 <0.01 <0.01 0.20 0.13 0.13 Magnesium 0.02 <0.02 <0.02 <0.02 <0.02 <0.02 Potassium 0.04 <0.04 <0.04 <0.04 0.10 0.09 Sodium 0.005 0.006 0.085 0.157 0.123 0.210 171 Appendix Figure C.1.1 XRD IrOx/TiOx showing two distinct phases, TiO2 (green, circles), PDF2 03-065-5714 (ICDD, 2017), IrO2 (blue, triangles), PDF2 01-088-0288 (ICDD, 2017) as well as an amorphous peak at 2 =12.5o. Appendix Figure C.1.2 SEM and EDX of IrOx/TiO2 before (a) and after (b) the ex situ AST (30,000 stepping cycles from 0.05 to 1.2 V vs RHE; 0.09 M H2SO4). 172 Appendix Table C.1.2 IrOx/TiO2 elemental atomic % determined by EDX before and after the ex situ AST. Element (shell) Atomic % Before After Oxygen (K) 69.16 67.06 Titanium (K) 5.70 6.33 Iridium (M) 25.14 26.60 Appendix Figure C.1.3 SEM images of (a) Low SA, (b) High SA HT and (c) High SA IrOx at 3000x magnification. Images were generated using a backscatter detector after the ex situ AST (30,000 stepping cycles from 0.05 to 1.2 V vs RHE; 0.09 M H2SO4) 173 Appendix Figure C.1.4 XRD of freshly prepared IrOx High SA HT powder (red) and as received IrOx High SA (green) and IrOx Low SA (black). One crystalline phase is identified as IrO2 (PDF2 00-015-0870 (ICDD, 2017)) for the IrOx High SA HT (red lines and Miller indices). Lin (Cps)01020304050607080902-Theta - Scale6 10 20 30 40 50 60 70 80High SA HT101110211220310High SALow SA 174 Appendix Figure C.1.5 Cell polarization curves at BOT, 100, 200 and 300 H2/Air AAST cycles, respectively. Anode Pt/Ir loading was 50/20 μg cm-2 and cathode Pt loading was 250 μg cm-2. OER catalysts were (a) IrOx/TiO2 (b) IrOx Low SA (c) IrOx High SA HT. Cell polarization was carried out in “differential mode” with a stoichiometry of 6 at the maximum current. 175 Appendix Table C.1.3 Cathode effective platinum surface area (EPSA) determined by CO stripping at BOT and EOT (after 300 H2/Air AAST cycles), respectively. Experiments performed in a single cell 20 cm2 MEA. Cathode EPSA / cm2pt cm-2geom BOT EOT IrOx/TiO2 115.5 124.5 IrOx High SA 111.6 109.4 IrOx High SA HT 111.2 106.9 C.2 Additional metal supported IrOx OER catalysts Three unsupported IrOx and five metal oxide supported catalysts were investigated as a part of this thesis. Due to intellectual property constraints with the suppliers of the catalysts, the identity of the metals used in the four of the five support materials is not available for publication. The supported metals are labeled as follows: IrOx/MOy #1, IrOx/MOy #2, IrOx/MOy #3 IrOx/MOy, IrOx/MOy #4 and IrOx/TiO2. The unsupported catalysts were obtained as either high surface area (IrOx High SA) or low surface area (IrOx Low SA) catalysts. Appendix Table C.2.1. BET surface area and IrOx content of three unsupported IrOx catalysts and six metal oxide supported IrOx catalysts. IrOx/MOy #1 IrOx/MOy #2 IrOx/MOy #3 IrOx/MOy #4 IrOx/TiO2 BET Surface area (m2 g-1) 9.5 ± 0.5 18.8 28 IrO2 content (wt. %) 37.27 57.2 55.76 48 75.6 176 Appendix Figure C.2.1. Electrolyte Ir concentration after accelerated stress test. Performed in a flooded (N2 purged 0.09 M H2SO4) three electrode cell 30,000 cycles from 0.05 V to 1.2 V vs. RHE with a 1 s hold at each potential. The IrOx/MOy #2 and #4 catalysts produced the lowest dissolved Ir present in the electrolyte at the end of AST indicating promising durability. This shows that, with the exception of IrOx/TiO2, non-carbon metal oxide supports increase the durability of IrOx catalysts towards dissolution in the ex situ test. IrO x/MxO y #1IrO x/MxO y #2IrO x/MxO y #3IrO x/MxO y #4IrO x/TiO 20102030405060Ir Concentration / µg L-1 177 Appendix Figure C.2.2. OER activity was measured in 0.09 M H2SO4 by LSV from 1.9 V to 1.3 V vs. RHE with a sweep rate of 5 mV s-1. 81.5 μg cm-2 of each catalyst was freshly deposited onto a GDL.MPL support material. (a) The overpotentials at 5 mA for each catalyst is plotted at the BOT and the EOT after 30k voltage stepping cycles (b) five metal oxide supported (IrOx/MxOy #1, IrOx/MOy #2, IrOx/MOy #3 IrOx/MOy, IrOx/MOy, #4 and IrOx/TiO2) were analyzed. The activity of metal oxide supported catalysts does not match up perfectly with the dissolution data. The IrOx/TiO2 catalyst had the highest dissolution followed by IrOx/MOy #1; but these two are switched with respect to activity. IrOx/MOy #1 has the highest activity (0.372 V) of the metal oxide supported catalysts. IrOx/MOy #4 also has higher activity than IrOx/MOy #2 and IrOx/MOy #3; 0.372, 0.570 and 0.562 V respectively. It would be expected from the better resistance to corrosion that IrOx/MOy #4 would have lower activity than either #2 or #3. The activity of IrOx/MOy #2 and #3 are nearly identical with #3 being slightly more active. However, this does not match up with the dissolution results; #3 has higher dissolution than #2. What is clear is that activity alone, while being a useful point of reference, is not enough on its own to predict the durability of OER catalysts in RTAs. Two catalysts also had increased activity after 30k cycles: 1.3 1.4 1.5 1.6 1.7 1.8 1.905101520 IrOx/MxOy #1 IrOx/MxOy #2 IrOx/MxOy #3 IrOx/MxOy #4 IrOx/TiO2i / mA cm-2E / V vs. RHE(b)00.10.20.30.40.50.6ηat 5mA cm-2BOTEOT(a) 178 IrOx/MOy #1 (0.372 V to 0.336 V) and IrOx/TiO2 (0.416 V to 0.398 V). IrOx/MOy #3 and #4 had small changes in activity (± <0.02 V) and these also had the best corrosion resistance in the ex situ AST. This suggests if there is little or no change to OER activity before and after the ex situ AST, then that catalyst is likely to have good resistance to corrosion. Appendix Figure C.2.3. (a) Reversal tolerance time at 0.2 A under N2 of various IrOx based OER catalysts using a freshly prepared anode at the BOT and after an in situ AAST on a second freshly prepared MEA (EOT). (b) Scatter plot of Ir dissolution after ex situ AST vs. in situ reversal loss (min). Metal oxide supported IrOx catalysts (solid red box) and unsupported IrOx catalysts (dashed blue box) were tested using both methods Catalysts were tested in a single MEA fuel cell with a protocol that mimics the in situ environment of a full PEMFC stack. The reversal time at the BOT and the EOT are plotted in Appendix Figure C.2.3. The in situ reversal loss time was plotted against the Ir concentration after the ex situ AST (Appendix Figure C.2.3b). Of the metal oxide supported catalysts, the IrOx/MOy #1 catalyst has the highest loss of reversal time (155 min). The ex situ results correlate well with the in situ results with the exception of the IrOx/TiO2 catalyst which had the highest dissolution of all the tested b0 10 20 30 40 50 600200400600800Reversal Loss / minIr Electrolyte Conc. / g L-1Unsupported IrO2Supported IrO2050100150200Reversal time / minBOTEOTa 179 catalysts ex situ and was fourth highest in terms of reversal loss (23 min). Additionally, the reversal losses of IrOx/MOy #2 (-20 min) and #3 (-16 min) were switched compared with the ex situ dissolution test. However, the difference between these two catalysts is small (4 min) and is within the margin of error. When compared to the unsupported OERs, the reversal loss of the metal oxide supported catalysts is less than would be expected from the ex situ dissolution results. In fact, all of the supported catalysts except for the IrOx/MOy #1 and IrOx/TiO2 had higher reversal retention times after the AAST. The increase in reversal time for metal oxide supported IrOx could be explained by increased IrOx surface area after cycling induced dissolution/re-deposition of IrOx. This shows that reversal time is a better indicator of the long term reversal tolerance of the cell while metal dissolution can be used to explain and predict both reversal tolerance as well as Ir3+ migration to the membrane and cathode leading to cell performance loss. 180 Appendix D D.1 ERC supplemental information Appendix Figure D.1.1 Chronoamperometry profiles for CO2 electroreduction at -2.0 V vs. Ag/AgCl for 2 hours in CO2 purged 0.5 M KHCO3. 293 K. Legend: (a) Catalysts supported on carbon paper obtained from the oxide based bath and (b) Catalysts supported on carbon cloth obtained from the oxide based bath, (c) Catalysts supported on carbon paper obtained from the fluoroborate based bath and (d) Catalysts supported on carbon cloth obtained from the fluoroborate bath. (1) Sn100%, (2) Sn80%-Pb20% (oxide) and Sn91%-Pb9% (fluoroborate), (3) Sn67%-Pb33% (oxide) and Sn75%-Pb25% (fluoroborate), (4) Sn50%-Pb50% (oxide) and Sn16%-Pb84% (fluoroborate), (5) Pb100%. 181 Appendix Figure D.1.2 XRD spectrum of fluoroborate deposited Sn100% on carbon cloth after polarization for 2 hours at 2.0 V vs. Ag/AgCl in 0.5 M KHCO3. Peak labels are for the Miller indices of crystalline Sn (S) (PDF2 00-004-0673 (ICDD, 2016)). Sn CCT00-004-0673 (*) - Tin, syn - Sn - Y: 70.20 % - d x by: 1. - WL: 1.5406 - Tetragonal - a 5.83100 - b 5.83100 - c 3.18200 - alpha 90.000 - beta 90.000 - gamma 90.000 - Body-centered - I41/amd (141) - 4 - File: Sn CCT.raw - Start: 4.868 ° - End: 89.912 ° - Step: 0.030 ° - Step time: 145.6 s - Anode: Cu - WL1: 1.5406 - WL2: 1.54439 - kA2 Ratio: 0.5 - Generator kV: 40 kV - Generator mA: 40 mA - Creation: Lin (Cps)010203040502-Theta - Scale5 10 20 30 40 50 60 70 80 90S (2,0,0)S (1,0,1)S (2,2,0)S (2,1,1) 182 Appendix Figure D.1.3 XRD spectrum of fluoroborate deposited Sn75%-Pb25% on carbon cloth after polarization for 2 hours at 2.0 V vs. Ag/AgCl in 0.5 M KHCO3 at 293 K. Peak labels are for Miller indices of crystalline Sn (S) PDF2 00-004-0673 (ICDD, 2016) and crystalline SnO2 (O) PDF2 03-065-2873 (ICDD, 2016). 60 40T01-072-1012 (*) - Tin Oxide - SnO - Y: 31.98 % - Tetragonal - a 3.80290 - b 3.80290 - c 4.83820 - Primitive - P4/nmm (129) - 2 - 69.9703 - I/Ic PDF 15.3 - S-Q 8.1 % - F30=1000(0.0000,33) - alpha 90.0000-004-0673 (*) - Tin, syn - Sn - Y: 46.57 % - Tetragonal - a 5.83100 - b 5.83100 - c 3.18200 - Body-centered - I41/amd (141) - 4 - 108.190 - I/Ic PDF 2. - S-Q 91.9 % - F29= 72(0.0129,33) - alpha 90.000File: 60 40T.raw - Start: 4.885 ° - End: 89.923 ° - Step: 0.030 ° - Step time: 145.6 s - Anode: Cu - WL1: 1.5406 - WL2: 1.54439 - kA2 Ratio: 0.5 - Generator kV: 40 kV - Generator mA: 40 mA - Creation: Lin (Cps)010203040502-Theta - Scale5 10 20 30 40 50 60 70 80 90O (1,1,0)O (1,1,2)0 (2,2,1)O (1,0,1)S (2,0,0)S (1,0,1)S (2,2,0)S (2,1,1) 183 Appendix Figure D.1.4 XRD spectrum of fluoroborate deposited Pb100% on carbon cloth after polarization for 2 hours at -2.0 V vs. Ag/AgCl. Peak labels are for the Miller indices of crystalline lead (Pb) (P), PDF2 03-065-2873 (ICDD, 2016), lead oxide (PbO) (O), PDF2 01-078-1666 (ICDD, 2016), massicot (PbO) (M), PDF2 01-077-1971 (ICDD, 2016) and plumbonacrite (Pb5(CO3)O(OH)2) (A), PDF2 01-073-8925 (ICDD, 2016). Pb-CCT01-073-8925 (N) - Plumbonacrite, syn - Pb5(CO3)3O(OH)2 - Y: 10.45 % - d x by: 1. - WL: 1.5406 - Hexagonal - a 9.09210 - b 9.09210 - c 24.92300 - alpha 90.000 beta 90.000 - gamma 120.000 - Primi01-077-1971 (*) - Massicot, syn - PbO - Y: 13.71 % - d x by: 1. - WL: 1.5406 - Orthorhombic - a 5.89310 - b 5.49040 - c 4.75280 - alpha 90.000 - beta 90.000 - gamma 90.000 - Primitive - Pbcm (57) - 4 -01-078-1666 (*) - Lead Oxide - PbO - Y: 20.09 % - d x by: 1. - WL: 1.5406 - Orthorhombic - a 5.62530 - b 5.62530 - c 5.02590 - alpha 90.000 - beta 90.000 - gamma 90.000 - Base-centered - Cmma (67)03-065-2873 (I) - Lead - Pb - Y: 87.39 % - d x by: 1. - WL: 1.5406 - Cubic - a 4.94930 - b 4.94930 - c 4.94930 - alpha 90.000 - beta 90.000 - gamma 90.000 - Face-centered - Fm-3m (225) - 4 - 121.236 -File: Pb-CCT.raw - Start: 5.000 ° - End: 89.975 ° - Step: 0.030 ° - Step time: 145.6 s - Anode: Cu - WL1: 1.5406 - WL2: 1.54439 - kA2 Ratio: 0.5 - Generator kV: 40 kV - Generator mA: 40 mA - Creation: Lin (Cps)01020304050607080902-Theta - Scale5 10 20 30 40 50 60 70 80P (1,1,1)P (2,0,0)P (2,2,0)P (3,1,1)O (1,1,1)O (0,2,0)O (2,0,2)O (1,3,1)A (1,1,5)A (3,0,0)M (1,1,1)M (1,0,10) 184 Appendix Figure D.1.5 XRD spectrum of fluoroborate deposited Sn16%-Pb84% on carbon cloth after polarization for 2 hours at -2.0 V vs. Ag/AgCl. Peak labels are for the Miller indices of litharge (OPb) (L), PDF2 00-005-0561 (ICDD, 2016), lead oxide hydrate (3PbO∙H2O) (H), PDF2 00-022-1134 (ICDD, 2016) and potassium lead carbonate hydroxide (KPb2(CO3)2OH) (C), PDF2 00-037-0502 (ICDD, 2016). 10 90T00-037-0502 (I) - Potassium Lead Carbonate Hydroxide - KPb2(CO3)2OH - WL: 1.5406 - Hexagonal - a 5.34800 - b 5.34800 - c 13.99000 - alpha 90.000 - beta 90.000 - gamma 120.000 - 2 - 346.522 - F00-022-1134 (N) - Lead Oxide Hydrate - 3PbO·H2O - WL: 1.5406 - Tetragonal - a 8.01600 - b 8.01600 - c 9.38900 - alpha 90.000 - beta 90.000 - gamma 90.000 - Primitive - P4/mnc (128) - 4 - 603.302 -03-065-2809 (I) - Litharge, syn - OPb - WL: 1.5406 - Tetragonal - a 3.97200 - b 3.97200 - c 5.01800 - alpha 90.000 - beta 90.000 - gamma 90.000 - Primitive - P4/nmm (129) - 2 - 79.1679 - I/Ic PDF 23.3 File: 10 90T.raw - Start: 5.000 ° - End: 90.005 ° - Step: 0.030 ° - Step time: 145.6 s - Anode: Cu - WL1: 1.5406 - WL2: 1.54439 - kA2 Ratio: 0.5 - Generator kV: 40 kV - Generator mA: 40 mA - Creation: Lin (Cps)0102030402-Theta - Scale5 10 20 30 40 50 60 70 80 90L (1,0,1)L (1,1,0)L (1,1,2)L (2,1,1)C (1,0,3)H (2,0,1)H (3,0,1)H (2,1,4)H (2,0,2)C (1,1,0)C (2,0,3) 185 Appendix Figure D.1.6 XRD spectrum of Sn91%-Pb9% on carbon cloth after polarization for 2 hours at 2.0 V vs. Ag/AgCl. Peak labels are for the miller indices of tin (Sn) (S), PDF2 03-065-0298 (ICDD, 2016) and lead (Pb) (P), PDF2 03-065-2873 (ICDD, 2016). 90 10T03-065-2873 (I) - Lead - Pb - Y: 32.36 % - d x by: 1. - WL: 1.5406 - Cubic - a 4.94930 - b 4.94930 - c 4.94930 - alpha 90.000 - beta 90.000 - gamma 90.000 - Face-centered - Fm-3m (225) - 4 - 121.236 -00-004-0673 (*) - Tin, syn - Sn - Y: 57.32 % - d x by: 1. - WL: 1.5406 - Tetragonal - a 5.83100 - b 5.83100 - c 3.18200 - alpha 90.000 - beta 90.000 - gamma 90.000 - Body-centered - I41/amd (141) - 4 - File: 90 10T.raw - Start: 4.868 ° - End: 89.912 ° - Step: 0.030 ° - Step time: 145.6 s - Anode: Cu - WL1: 1.5406 - WL2: 1.54439 - kA2 Ratio: 0.5 - Generator kV: 40 kV - Generator mA: 40 mA - Creation: Lin (Cps)01020304050607080901002-Theta - Scale5 10 20 30 40 50 60 70 80 90S (2,0,0)S (1,0,1)P (1,1,1) S (2,2,0)S (2,1,1)P (2,0,0)P (2,2,0) 186 Appendix Figure D.1.7 FIB-SEM images of Sn16%-Pb84% deposited without Triton (a) freshly prepared and (b) after 2 hours of use as a CO2 reduction catalyst layers at 5000X magnification. Appendix Figure D.1.8 Megazyme formate testing kit (a) absorbance spectrum at four formate concentrations in 0.5 M KHCO3 and blank KHCO3 solution. (b) Calibration curve and linear regression used for calculation of formate concentrations in electrolyte solutions. Appendix Figure D.1.8 shows the results of a calibration curve generated using a Megazyme enzyme based spectroscopic kit (K-Form). [219,220] The principles and methodology for this kit are presented below. 280 300 320 340 360 380 4000123Absorbance / nmBlank1x10-2M1x10-3M5x10-3M5x10-4M1x10-4M0.000 0.002 0.004 0.006 0.008 0.0100.00.40.81.21.6Absorbance at 340 nmFormate Concentration / mol L-1y = 169x +0.0117R2 = 0.9986a b 187 The K-Form kit uses the enzyme formate dehydrogenase (FDH) to catalyze the oxidization of HCOOH to CO2 by nicotinamide adenine dinucleotide (NAD+). HCOOH + NAD+ → CO2 + NADH + H+ Appendix Equation D.1 This reaction results in the production of NADH which is stoichiometrically equivalent to initial moles of formic acid. The K-Form kit contained: Bottle 1: Buffer (pH 7.6) Bottle 2: NAD+ Bottle 3: formate dehydrogenase (stable at -20 °C) Bottle 4: Sodium formate (~ 2 g) Each of these bottles were prepared according to the kit instructions. A series of formate concentration standards was prepared in 0.5 M KHCO3. These standards are later referred to as “sample.” The following was pipetted into 1.6 mL plastic Eppendorf® tubes: Distilled water (at 25 °C) 200 μL Sample 10 μL Buffer solution 20 μL NAD+ solution 20 μL This solution was mixed manually (by flicking) and allowed to rest for 5 minutes. After which, 5 μL of formate dehydrogenase solution was added. This final solution was then analyzed after the absorbance signal stabilized (approximately 10 minutes) as described in section 0.
Thesis/Dissertation
eng
Chemical and Biological Engineering
Vancouver : University of British Columbia Library
Attribution-NonCommercial-NoDerivatives 4.0 International
Graduate
Electrocatalysts for voltage reversal tolerant anodes in proton exchange membrane fuel cells and CO2 reduction to formate
Text