Designing interfacial structures for selective electrocatalysis

Designing Interfacial Structures for Selective

Electrocatalysis

by

Bing Yan

B.A. Chemistry, Peking University (2014)

Submitted to the Department of Chemistry

in partial fulfillment of the requirements for the degree of

Doctor of Philosophy in Chemistry

at the

MASSACHUSETTS INSTITUTE OF TECHNOLOGY

June 2019

Massachusetts Institute of Technology 2019. All rights reserved.

Signature redacted

A u th o r ...

Department of Chemistry

May 10, 2019

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C ertified b y ...

Yogesh Surendranath

Paul M. Cook Career Development Associate Professor

Thesis Supervisor

Signature redacted

Accepted by ...

..

Robert W. Field

Haslam and Dewey Professor of Chemistry

Chairman, Department Committee on Graduate Students

MASSACHUSETTS INSTITUTE OF TECHNOLOGY

This doctoral thesis has been examined by a Committee of the

Department of Chemistry as follows:

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Professor Mircea Dinca..

airm

fhhesis Committee

sso< e Professor of Chemistry

Signature redacted

Professor Yogesh Surendranath ...

Thesis Supervisor

Paul M Cook Career Development Associate Professor

Signature redacted

Professor Fikile R. Brushett...

..

Member, Thesis Committee

Associate Professor of Chemical Engineering

Designing Interfacial Structures for Selective Electrocatalysis

by

Bing Yan

Submitted to the Department of Chemistry on May 10, 2019, in partial fulfillment of the

requirements for the degree of Doctor of Philosophy in Chemistry

Abstract

Selectivity of electrocatalysis in fuel cells stresses on the tolerance of electrode catalysts to cross-over species. A typical polymer electrolyte membrane fuel cell (PEMFC) employs a Nafion-based membrane to separate the Pt-based cathode and anode. Nevertheless, the intrinsic porosity of Nafion membranes allows reactants and products to cross over from one electrode chamber to the other. As a result, the poor selectivity of Pt leads to mixed reactivity at both electrodes and thus, decreased output voltage and power density. In principle, the membrane can be eliminated if cathode and anode catalysts are selective to the desired half reaction. Herein, the thesis aims to develop selective cathode and anode catalysts. For the cathode, transition metal chalcogenides are selective catalysts for the oxygen reduction reaction. We investigate the structure, activity, surface dynamics, and mechanism of a nickel sulfide catalyst during the oxygen reduction catalysis. By employing the selective nickel sulfide catalyst as the cathode, we construct a proof-of-concept membrane-free fuel cell which significantly outperforms an unselective Pt-cathode congener.

For the anode, there is still a paucity of intrinsically selective fuel oxidation catalysts. To achieve high 02 tolerance, we design a new configuration of elec-trocatalysis by employing a solid oxide-based mixed electron-proton conductor (MEPC) as a condensed membrane to segregate the catalyst and electrolyte while only transporting H-atom equivalents, thus blocking 02 and impurities dissolved in the electrolyte from reaching the catalyst surfaces. We investigate the activity, selectivity, and mechanism of the catalyst/MEPC membrane composite electrode. Thesis Supervisor: Yogesh Surendranath

Contents

List of Figures List of Tables

List of Abbreviations 1 Introduction

1.1 Transition metal chalcogenicis are highly selective catalysts for the

oxygen reduction reaction . . . . 1.2 Mixed electron-proton conductor membranes enable selective

elec-trocatalysis . . . . 2 Selective oxygen reduction electrocatalysis

2.1 Electrocatalysis and surface reconstruction of first-row transition metal chalcogenides . . . . 2.1.1 Introduction . . . . 2.1.2 Phase-pure nickel sulfide undergoes surface reconstruction

under ORR conditions . . . . 2.1.3 Mechanistic insights into self-limiting nickel sulfide surface

reconstruction . . . . 2.1.4 Local environment dictates the ORR activity of nickel sulfides 2.1.5 Conclusions . . . . 2.1.6 Experimental details . . . . 10 22 25 29 32 36 51 51 52 54 60 65 74 74

2.2 Selective metal chalcogenides electrocatalysts for membrane-free fuel l1ls 2.2.1 2.2.2 2.2.3 2.2.4 2.2.5 2.2.6 .Introductio... Introduction . . . .

Evaluation of intrinsic selectivity and activity of Ni3S2 and Pd Optimization of electrolyte concentrations for the membrane-free fuel cell . . . . Construction of a proof-of-concept membrane-free fuel cell . . C onclusions . . . . Experimental details . . . . 82 83 85 88 89 92 93

3 Selective fuel oxidation electrocatalysis 103

3.1 Mixed electron-proton conductors enable spatial separation of bond activation and charge transfer in electrocatalysis . . . 104 3.1.1 Introduction . . . 105 3.1.2 Preparation and characterization of Pt/tungsten oxide

com-posite membranes . . . 107 3.1.3 The Pt/tungsten oxide MEPC membrane electrode for

hy-drogen oxidation electrocatalysis . . . 112 3.1.4 Mechanistic studies of MEPC membrane-mediated

hydro-gen oxidation electrocatalysis . . . 116 3.1.5 HOR current densities are enhanced by roughening the

tung-sten oxide surfaces . . . 123 3.1.6 Conclusions and outlook . . . 127 3.1.7 Experimental details . . . 128 3.2 Quantitative model of the Pt/WOx composite electrode for

hydro-gen oxidation catalysis . . . 133 3.2.1 Introduction . . . 134 3.2.2 Construction of the EH-XH correlation . . . 136 3.2.3 Determination of EH and XH of the Pt/WOx composite

3.2.4 Derivation of the rate law of H-spillover from XH-t correlation 142

3.2.5 Simulation of HOR current using the derived rate law of H-spillover . . . 152 3.2.6 Conclusions and future directions . . . 154 3.2.7 Experimental details . . . 157

List of Figures

1-1 (a) Schematic representation of a typical PEM fuel cell. The anode and cathode are separated by a proton-conducting Nafion mem-brane. (b) Potential diagram of the H2-02 PEM fuel cell.

Overpo-tential is required to drive both H2 oxidation and 02 reduction re-actions. The black arrows denote the desired fuel oxidation and 02 reduction reactions. The red arrows denote H2 and 02 cross-over and the undesired reactions. . . . 30 1-2 Schematic representation of catalyst/MEPC membrane composite

electrode and the comparison of H2 oxidation current for the

cat-alyst/MEPC membrane configuration and traditional Pt I solution configuration. . . . 36 1-3 Proposed MEPC membranes and matching reactions. . . . 40 2-1 (a) Crystal structure, (b) PXRD pattern, (c) SEM images, and (d)

TEM images of Ni3S2 . . . . 55

2-2 High resolution XPS of (a) Ni 2_{p3/2 and (b) S 2p regions. As}

pre-pared Ni3S2, post-oxidation surface, and bulk Ni3S2 are shown in

black, red, and blue respectively. (c) XPS-derived Ni and S atomic ratios and percentages as a function of the number of Ar* sputter-ing cycles. (d) Initial CV scans of Ni3S2 recorded in N2-saturated

1 M sodium phosphate electrolyte, pH 7, at a scan rate of 5 mV

2-3 TEM images of (a) as-prepared and (b) post-electrolysis Ni3S2 nanopar-ticles. Fast-Fourier transforms are shown in the insets of (a) to con-firm the crystalline nature of the bulk material. Elemental compo-sition determined from energy-dispersive X-ray spectroscopy (EDS) at a series of spots along a line from the crystallite edge to the bulk

(c) reveals the variation in Ni:S composition across the crystallite (d). 58 2-4 Polarization of the Ni3S2 nanoparticles and histograms of the

amor-phous surface layer thickness. Polarization of Ni3S2 nanoparticles drop cast onto TEM grids by a) cyclic voltammetry (CV) and c) chronoamperometry (CA). In CA, the polarization time was increased from 23 s (b, black), 2 min (b, red), and 5 min (b, blue). The elec-trochemistry was performed in 02-saturated 1 M NaPi, pH 7, elec-trolyte. The histograms (b, d-f) were constructed by measuring the

amorphous layer thickness across 10-20 particles for each sample. . . 59

2-5 CVs of as-prepared Ni3S2 in (a) pH 4.7, (b) 7.0, (c) 8.4 and (d) 9.9

sodium phosphate electrolyte. The insets of (c) and (d) are zoom-in CVs with negative-going scans ending at 0.0 and 0.3 V. . . . 62 2-6 Calculated Ni-O-S Pourbaix diagram assuming an equilibrium Ni2

+

concentration of 0.001 M. Ni3S2 is oxidized to Ni2l , NiS and/or Ni(OH)2 depending on the acidity of the solution. The letters corre-spond to the equilibrium outline above. . . . 63 2-7 GIXD of (a) as-prepared NiS and (b) annealed NiS. The as-prepared

sample does not display crystalline peaks, while the annealed NiS displays XRD peaks consistent with the millerite NiS phase. (c) and (d) XPS of electrodeposited NiS. The Ni 2p3/2 peaks are split into two chemical environments: NiO at 856.1 eV and NiS at 853.3 eV. NiO is known to be catalytically inert for ORR. . . . 64 2-8 CV on a freshly-prepared Ni3S2 electrode in the presence of 02. The

Ni3S2 electrode undergoes oxidative surface reconstruction while si-multaneously catalyzing ORR. . . . 65

2-9 (a) Linear sweep (5 mV s-1 scan rate) voltammograms (LSVs) and (b) CAs of Ni3S2 (black), as-prepared NiS (blue), and annealed NiS (red). LSVs (a, solid lines) and steady state measurements (b) were recorded in 02-saturated 1 M sodium phosphate electrolyte, pH 7, at a rotation rate of 2000 rpm. LSVs recorded N2-saturated electrolyte are plotted in dashed lines . . . 66 2-10 The overlay of the LSV and steady state ORR data of Ni3S2 and NiS. 66

2-11 Double-layer capacitance traces (a and c) and linear fits to the ag-gregate capacitive current at the open circuit potential (b and d) for electrodeposited NiS (a and b) and surface reconstructed Ni3S2 (c

andd). ... 67

2-12 The Tafel plots of reconstructed Ni3S92 and electrodeposited NiS with linear fitting. . . . 67 2-13 Free energy diagram plotted at the limiting potential on the NiS2

phase. The step involving the activation of 02 as OOH* is the po-tential determining step and the free energy levels of these two in-termediates line up at the limiting potential of 0.88 V. . . . 69 2-14 Free energy diagram plotted at the limiting potential on the Ni3S4

phase. The step involving the reduction of OH* to H20 is the

poten-tial determining step and the free energy levels of these two inter-mediates line up at the limiting potential of 0.75 V. . . . 69 2-15 Free energy diagram plotted at the limiting potential on the NiS

phase. The step involving the reduction of OH* to H20 is the po-tential determining step and the free energy levels of these two in-termediates line up at the limiting potential of 0.86 V. . . . 70 2-16 Free energy diagram plotted at the limiting potential on the Ni9S8

phase. The step involving the reduction of OH* to H20 is the

poten-tial determining step and the free energy levels of these two inter-mediates line up at the limiting potential of 0.7 V. . . . 70

2-17 Free energy diagram plotted at the limiting potential on the Ni3S2

phase. The step involving the activation of 02 as OOH* is the

po-tential determining step and the free energy levels of these two in-termediates line up at the limiting potential of 0.82 V . . . 71 2-18 (a) Scaling between adsorption free energies of OH* and OOH* on

the various stable Ni-S phases. The black dots represent the sorption energies and the green dots represents the ensemble of ad-sorption free energies obtained from the BEEF-vdW functional that enables error estimation. The best-fit line is given by AGOOH ~ AGOH + 3.04, and the intercept has a standard deviation of 0.17 eV. The black dotted line represents the Ic- line. b) ORR activity vol-cano of nickel sulfide phases showing the expected limiting poten-tial (bold line) and the liming potenpoten-tial (dashed line) obtained from a thermodynamic analysis, as a function of the DFT-calculated ad-sorption free energy of the intermediate OH*. c) a-NiS structure constructed from ab initio simulated annealing. . . . 71 2-19 (a) Correlation between DFT-calculated free energy of adsorption

of OH* and that obtained from the structure descriptor, AGOH*

-0.28(0.08CNN(Ni)

+

cor-relation between the structure descriptor and the DFT-computed AGOH*. (b) Structure-activity contour plot of the expected limiting potential, UEL, based on the structure descriptor. The markers are plotted based on the local coordination around the active site for the various Ni-S phases. . . . 73 2-20 (a) Schematic of a membrane-free carbonaceous fuel cell depicting

the desirable fuel combustion (black) and undesirable parasitic short-circuit reactions (red). (b) Energy diagram of the fuel cell depicting the driving forces for the desirable fuel combustion reactions (black arrows) and the undesirable parasitic reactions (red arrows). . . . 84

2-21 The LSV of (a) Ni3S2/C and b) Pt/C recorded in 02-saturated 0.1 M KPj, pH 7, electrolyte with (red) and without (black) 1.0 M HCO2K. (c) LSV of Pd/C in N₂- (black) and 02-saturated (red) 0.1 M KPi,

pH 7, electrolyte containing 1.0 M HCO2K. (d) Overlay of LSV curves recorded in 02-saturated (red) 0.1 M KPi, pH 7, electrolyte containing 1.0 M HCO2K, with the absolute value of the current

den-sity for Ni3S2/C (red) and Pd/C (black). All data were collected at

5 mV s-I scan rate and 2000 rpm rotation rate. . . . 85

2-22 (a) LSV of Ni3S2/C recorded in 02-saturated 0.1 M KPi (black) and

1.0 M KPj (red), pH 7 electrolyte. (b) CV of Pt/C recorded in

02-saturated 0.1 M KP₁, pH 7 electrolyte, containing 1.0 M HCO2K. All data were collected at 5 mV s I scan rate and 2000 rpm rotation rate. 86

2-23 LSV traces of (a) Pd/C and (b) Ni3S2 recorded in 02-saturated 0.1 M KPj, pH 7, electrolyte containing 1.0 M HCO2K (black), 2.0 M HCO2K (red), and 0.2 M HCO2K (blue). Data in (a) and (b) were recorded at

5 mV s-- scan rate and 2000 rpm rotation rate. ... 88

2-24 Calculated power density vs voltage curves based on 3-electrode LSV data recorded in 02-saturated 0.1 M KPj, pH 7, electrolyte con-taining 1.0 M HCO2K (black), 2.0 M HCO2K (red), and 0.2 M HCO2K (b lu e). . . . 88

2-25 Schematic diagram of a flow-through, two-port membrane-free for-mate fuel cell. Inside the cell: graphite tube current collectors, car-bon felt electrodes, carcar-bon paper and separator. The 02-saturated

0.1 M KPI, 1.0 M HCO2K electrolyte is flowed from the cathode to the anode. ... 90

2-26 Cell voltage (left axis) and power density (right axis) vs current den-sity plots for a Ni3S2-Pd membrane-free fuel cell with electrolyte

flowed at (a) 0.1 mL min-1 _{and (b) 1.0 mL min-}1 _{from the cathode}

(Ni3S2/C) to the anode (Pd/C) (black) and from the anode to the

cathode (red). (c) Cell voltage (left axis) and power density (right axis) for a Ni3S2-Pd (black) and Pt-Pd (red) membrane-free fuel cells with electrolyte flowed at 1.0 mL min-1 _{from the cathode to the} an-ode. The error bars represent the standard error of five indepen-dently assembled fuel cells. . . . 91

3-1 Schematic representations of hydrogen oxidation reaction (HOR) cat-alyzed by (a) a traditional singular solid I liquid electrochemical in-terface and (b) WOx MEPC membrane-mediated inin-terfaces. The color scheme of (b): red-oxygen, orange-W(VI), blue-W(V), and green-H . .. .. . .... .. . .. . .. ... . . . . . .. 106

3-2 (a) The schematic representation of the fabrication of the Pt/WOx composite electrode. (b) The cross-section SEM image of magnetron sputter deposited Pt/WOx film, supported on porous polycarbonate membrane. (c) The top-down SEM image of the Pt/WOx film. (d) The TEM image of Pt islands supported on approximately 40 nm thick WOx films which were deposited on a Cu/lacey carbon TEM grid. (e) XPS spectrum and peak-fitting of W 4f region of the as-prepared Pt/WOx films. The grey dots are the experimental data. The red and blue lines represent the fitted W 4f region for W(VI) and W(V), respectively. The black lines depict the fitting envelope and the background. (f) XPS spectrum of Pt 4f region of the as-prepared Pt/W O x film s. . . . 107

3-4 PXRD pattern of a magnetron sputtered WOx film supported on porous polycarbonate substrates. The observed features are all at-tributed to the polycarbonate. (c) Survey XPS of the as-prepared Pt/WOx composite film. Peaks corresponding to Pt, W, 0, and ad-ventitious C are noted. . . . 108

3-5 Survey XPS of the as-prepared Pt/WOx composite film. Peaks cor-responding to Pt, W, 0, and adventitious C are noted. . . . 109

3-6 (a) Initial cyclic voltammograms of the Pt/WOx MEPC membrane electrode catalyzing HOR. H2 (1 atm) was supplied to the gas-facing side of the Pt/WOx working electrode. An increase in the

magni-tude of current density wras observed across initial cycles. (b) Cyclic

voltammogram of the Pt/WOx composite electrode with 1 atm N2

(red) or H2 (black) supplied to the working electrode. The CV of the singular Pt I solution interface for HOR catalysis is plotted as a comparison (blue dash). The data of (a) and (b) were collected in Ar-saturated 0.1 M HClO4, pH 1 electrolyte . . . 111

3-7 Steady-state HOR current densities were collected in Ar-saturated

0.1 M HClO4, with fH2 supplied to the working electrode. The

po-tential was stepped from 0 to 0.48 V in 40 mV increments. (a) The plot of current densities versus time indicate balanced rate of H2 activation and H-spillover by the rate of electrochemical electron-proton charge separation. (b) The plot of current densities versus potential. The error bars represent the standard error of the mean calculated from three independent film preparations. . . . .111

3-8 The HOR activity of a direct Pt I solution interface was examined in the presence of dissolved 02 or CO in 0.1 M HClO4 electrolyte. (a) CVs recorded in the presence of 1 atm H2 (black), 1 atm 02 (red), and a mixture of 0.5 atm H2 and 0.5 atm 02 (blue). (b) CVs recorded in the presence of 1 atm H2 (black), 1 atm CO (red), and a mixture of

0.5 atm H2 and 0.5 atm CO (blue). The horizontal grey dashed line

denotes zero current density. . . . 114 3-9 The steady-state HOR current density of the Pt/WOx com-posite

electrode in the presence of gaseous impurities. The data were col-lected in Ar- (black square), 02- (red circle), and CO-saturated (blue

triangle) 0.1 M HClO4. . . . 114

3-10 Chronoamperograms of (a) the singular Pt I solution interface and (b) the Pt/WOx composite electrode polarized at 0.50 V in 0.1 M HClO4 electrolyte. Arrows denote the time at which 10 mM Cu(C10₄)2

was added to the electrolyte . . . 115 3-11 Cu 2p region of the XPS spectra of the Pt side of the Pt/WOx

com-posite electrode. The spectrum of the as-prepared electrode (black) is compared to the spectrum of the electrode following 10 min of

electrolysis in the presence of 10 mM Cu(C104)2. . . . . 116

3-12 Cross-section SEM images for Pt/WOx composite films supported on porous polycarbonate substrates. WOx films were deposited for (a) 2.5 and (b) 4 h, respectively. Pt was deposited for 30 s for both sam ples. . . . 117

3-13 Top-down SEM images of as-prepared WOx films of varying thick-ness. The WOx films were deposited for (a) 2.5, (b) 3, and (c) 4 h to fulfill the thicknesses of 0.8, 1.0, and 1.4 ym. The WOx films of different thicknesses display the same morphology. . . . 117

3-14 (a) Steady-state current densities vs potential plots for Pt/WOx com-posite electrodes with a time duration of WOx magnetron sputter of

2.5 (black), 3 (red) and 4 h (blue). (b) The steady-state HOR current density at 0.48 V was plotted versus WOx thickness. The time du-ration of Pt sputtering was 30 s for all samples. The error bars in (b) represent the standard error of the mean calculated from three independent film preparations. . . . 118 3-15 TEM images of the Pt/WO, films deposited on TEM grids. The time

durations of Pt magnetron sputter are (a) 20, (b) 25, (c) 30, (d) 40, and (e) 45 s. The thickness of the WO,, film is approximately 40 nm for all sam ples . . . 119 3-16 The steady-state HOR current density was plotted versus the time

duration of Pt sputtering at 0.48 V. The time duration of WOx sput-tering was 3 h for all samples. The error bars represent the standard

_ro f the- ____ -al-1,a-4 fr-- theeinepndn film rpaa

tion s. . . . 120 3-17 Steady-state current density vs potential data for Pt/WOx films

pre-pared with Pt magnetron sputtering time durations of (a) 20, (b) 25, (c) 30, (d) 40, and (e) 45 s. The time duration of WOx magnetron sputtering was 3 h for all samples. . . . 122 3-18 The steady-state HOR current density vs potential plots for Pt/WOx

composite electrodes recorded in electrolytes of varying H' concen-trations. The data were collected in 0.1 (black square), 1.0 (red cir-cle), 1.5 (blue triangle), and 2.0 M (cyan triangle) HClO4. NaClO4

was added to maintain the ionic strength at 2.0 M. . . . 123 3-19 A putative mechanism for HOR catalysis at Pt/WOx compo-site

electrodes. More positive potentials (b) decrease the H concentra-tion within WO, relative to low applied potentials (a). We invoke rate limiting H-spillover which is accelerated by a lower H concen-tration in W O x. . . . 124

3-20 Representative AFM images for (a-c) smooth Pt/WOx composite films prepared by magnetron sputtering and (d-f) surface-roughened Pt/WOx composite films prepared by thermal evaporation and post-annealing. Both samples were supported on AAO substrates. . . . . 125 3-21 SEM images of surface-roughened WOx films prepared by thermal

evaporation (a) before and (b) after 500 'C annealing for 3 h. . . . 125 3-22 Comparison of (a) CVs and (b) CAs between roughened and smooth

Pt/WOx composite electrodes. The data were collected in 0.1 M HC1O4. The composite films were supported on porous AAO sub-strates. ... 126 3-23 A qualitative model for HOR catalysis at the Pt/WOx composite

electrode. We invoke the rate-limiting H-spillover is balanced by irreversible electron-proton charge separation. . . . 134

3-24 Plots of H chemical potential (EH) vs H concentration (XH) for WOx,

membranes prepared by magnetron sputtering. The data are fit to Equation 3.3. . . . 137 3-25 (a) Representative cross-sectional SEM image of WO,, films prepared

by magnetron sputtering. The thickness of the WOx films is 800 nm. (b) Representative top-down SEM image of WOx films reveal smooth pinhole-free condensed surfaces. (c) Representative AFM image of WOx films, revealing a roughness factor of 1.02 0.01. . . . 140

3-26 Steady-state HOR current at HOR potential from 0.12 to 0.52 V . . . . 140 3-27 The open circuit potential (EOcp) was recorded immediately after

re-moving HOR potential (EHOR) of various values from 0.12 to 0.52 V.

The inset figures denote the change of Eocp at the initial 0.3 s. The electrode is 35 s deposition Pt supported on 2 h deposition WO,. . . 141

3-28 Plots of XH-t at varying EHOR from 0.12 to 0.52 V. The data are fitted

to Equation 3.11. The electrode is 35 s deposition Pt supported on 2 h deposition W O. . . . 145

3-29 The fit KH-spillover at varying HOR potentials for Pt/WOx compos-ite electrodes of varying Pt sputtering time durations: 35 s (black square), 20 s (red circle), and 45 s (blue triangle). . . . 146 3-30 Calculated in situ H concentration, x0 at varying HOR potentials for

Pt/WOx composite electrodes of varying Pt sputtering time dura-tions: 35 s (black square), 20 s (red circle), and 45 s (blue triangle). . .146 3-31 The open circuit potential (Eocp) was recorded immediately after

re-moving HOR potential (EHOR) of various values from 0.12 to 0.52 V. The inset figures denote the change of Eocp at the initial 0.3 s. The

electrode is 20 s deposition Pt supported on 2 h deposition WO,. . . 147 3-32 The open circuit potential (Eocp) was recorded immediately after

re-moving HOR potential (EHOR) of various values from 0.12 to 0.52 V. The inset figures denote the change of Eocp at the initial 0.3 s. The

electrode is 45 s deposition Pt supported on 2 h deposition WOx. . . 148

3-33 Plots of xH-t at varying EHOR from 0.12 to 0.52 V. The data are fitted

to Equation 3.11. The electrode is 20 s deposition Pt supported on 2 h deposition W O.. . . . 150 3-34 Plots of XH-t at varying EHOR from 0.12 to 0.52 V. The data are fitted

to Equation 3.11. The electrode is 45 s deposition Pt supported on 2 h deposition W O .. . . . 151 3-35 Comparison of measured (black square) and calculated HOR

cur-rent (red circle) versus varying EHOR .. . . . . . . . . . . .. 152

3-36 (a) The Ohmic resistance dependence on applied potential. (b) Com-parison of measured (black square) and calculated HOR current (red

circle) versus varying EHOR. . . . . . . . .. . . . 154 3-37 A schematic fuel cell with two catalyst/mixed electron-proton

con-ductor composite electrodes. The two electrodes are separated by an ultra thin solid oxide proton conductor to avoid short-circuiting. . 155

List of Tables

2.1 The free energies of adsorption of the oxygen intermediates based

on the respective thermodynamically favorable adsorption site. . . . 68

2.2 Adsorption free energies calculated from DFT, and through the structure-energy descriptor, on the three stable surface sites on the simulated amorphous structure with a composition of Ni:S 1.0:1.0. . . . 72 2.3 Thermodynamic data for Ni compounds. . . . 77 3.1 ICP-MS results of Cu content in Pt/WOx films. . . . 132 3.2 Fit results of XH-t data for Pt/WOx electrodes of 35 s Pt deposition

and 2 h W Ox deposition. . . . 144 3.3 Fit results of XH-t data for Pt/WOx electrodes of 20 s Pt deposition

and 2 h W Ox deposition. . . . 149 3.4 Fit results of XH-t data for Pt/WO electrodes of 45 s Pt deposition

List of Abbreviations

anodic aluminum oxide atomic force microscopy

Bayesian Error Estimation Functional chronoamperometry

coordination number of nearest neighboring Ni atoms coordination number of nearest neighboring S atoms concentration

cyclic voltammetry density functional theory applied potential

energy dispersive spectroscopy

local-density approximations correlation energies non-local correlation energy

Perdew, Burke and Ernzerhof functional correlation energies hydrogen chemical potential of WOx

applied potential for hydrogen oxidation catalysis generalised gradient approximation exchange energy open circuit potential

anodic peak potential Faraday's constant

generalised gradient approximation exchange enhancement factor generalised gradient approximation

HER HOR HRTEM H-spillover H-transport I (i) ICP-MS IHOR IHOR,cal 'HOR,cal iR i k Keq kH-spillover kH-spillover KH-spillover 1 LDA LSV LTMC MAE MEPC Mwox MX nH NHE nwox OCP OCV

hydrogen evolution reaction hydrogen oxidation reaction

high resolution transmission electron microscopy hydrogen spillover

hydrogen atom equivalent transport current

inductively-coupled plasma mass spectrometry current of the hydrogen oxidation reaction

calculated current of the hydrogen oxidation reaction calculated HOR current considering resistance loss uncompensated ohmic potential loss

current density rate constant

equilibrium constant

rate constant of hydrogen spillover

rate constant of reverse hydrogen spillover

rate constant of H-spillover including the constant terms thickness of WOx films

local density approximation linear sweep voltammetry

late transition metal chalcogenides mean absolute error

mixed electron-proton conductor molar mass of WOx

degree of the polynomial

molarity of hydrogen inserted into WOx normal hydrogen electrode

total molarity of WOx open circuit potential open circuit voltage

OD outer diameter

OER oxygen evolution reaction ORR oxygen reduction reaction

P power density

PBE Perdew, Burke and Ernzerhof functional P1 phosphate buffer

PXRD powder X-ray diffraction

QH charge passed to intercalate hydrogen into WO,

R gas constant

R 2 coefficient of determination RDE rotating disk electrode RF roughness factor

RHE reversible hydrogen electrode rls rate-limiting step

R2 ,Ohmc- rcistance

SEM scanning electron microscopy

STEM scanning transmission electron microscopy

t time

T temperature

TEM transmission electron microscopy

U potential

UEL expected limiting potential

UL limiting potenital

UPD underpotential deposition

V voltage

wt% percent by weight

xo in situ hydrogen concentration in WOx

XH hydrogen concentration inserted into WOx

XH,eq hydrogen concentration at equilibrium under 1 atm H2

AE potential difference

AGOH* adsorption free energy of OH*

(AGOH*) mean value of adsorption free energy of OH*

AG c adsorption free energy of OH* calculated through descriptor

AGOOH* adsorption free energy of OOH*

AG DFT adsorption free energy of OH* calculated from DFT

AH enthalpy

AS entropy

/ overpotential

6 surface coverage

6_{H-WOx surface coverage of hydrogen intercalated into WO,,} 6_{Pt-H surface coverage of hydrogen adsorbed Pt sites}

opt surface coverage of hydrogen-free Pt sites

6_{Pt I wOx surface coverage of Pt I WOx boundary} 6_{WOx surface coverage of hydrogen-free WOx sites}

YOH* mean value of adsorption free energy of OH*

YOOH* mean value of adsorption free energy of OOH*

Pwox density of WO, films

L OH* standard deviation of adsorption free energy of OH* OOOH* standard deviation of adsorption free energy of OOH*

Chapter 1

Introduction

Fuel cells directly convert chemical energy stored in fuels and 02 into electricity. 1-3

The theoretical voltage of a fuel cell is the potential difference between the cathode and anode where 02 reduction and fuel oxidation reactions take place respectively. Taking an H2-02 fuel cell as an example, H2 is oxidized to H+ or H20 at the anode

and 02 is reduced to H20 or OH- at the cathode. Therefore, the upper limit

volt-age of a H2-02 fuel cell is 1.23 V. However, both fuel oxidation and 02 reduction reactions require the transfer of multiple electrons and protons, making these reac-tions kinetically sluggish.4 5 As a result, a few hundred millivolts of overpotential are typically applied to drive these reactions. The large overpotential leads to a cell voltage significantly smaller than the theoretical value (Figure 1-1, black). Ad-ditionally, the sluggish kinetics of fuel cell reactions give rise to small current and power. Therefore, it is critical to develop highly active electrocatalysts that facil-itate the conversion of small molecule substrates. An ideal electrocatalyst should minimize the overpotential required to activate the fuels or 02, and meanwhile support rapid turnover of the substrates, thus supporting a fuel cell of high volt-age, current, and power.

Among the wide variety of electrocatalysts developed for fuel cells, Pt-based materials display high activity and durability for both the fuel oxidation reaction and the oxygen reduction reaction (ORR).6-1 1 _{Despite their high cost, Pt-based}

a

b

e-

H

-*

2H+

H

H H

cell voltage

H+u

H

r

oRR

1.

V

anode

ORR

4e-

4H++

02

-

2H

catalyst

catalyst

Figure 1-1: (a) Schematic representation of a typical PEM fuel cell. The anode and cathode are separated by a proton-conducting Nafion membrane. (b) Potential diagram of the H2-02 PEM fuel cell. Overpotential is required to drive both H2 oxidation and 02 reduction reactions. The black arrows denote the desired fuel oxidation and 02 reduction reactions. The red arrows denote H2 and 02 cross-over and the undesired reactions.

However, the poor selectivity of Pt catalysts to fuel oxidation and ORR causes mixed reactivity in the presence of both fuel and 02 substrates.13 To avoid mixed reactivity at each Pt electrode, a Nafion-based polymer-electrolyte membrane (PEM) is often employed to separate the cathode and anode chambers but only conduct protons.14,15 _{Nevertheless, the Nafion membrane is not only expensive, but its} sul-fonic groups are strongly acidic, which restricts the catalyst choices to precious metals. 16 Furthermore, the Nafion membrane is porous in nature and is hydrated

to form water channels through which protons are transported. 17-19 As a result, soluble substrates readily cross over through the membrane from one electrode chamber to the other. The cross-over substrates, either 02 to the anode chamber or fuels to the cathode chamber, compete with the desired substrates for active sites of the Pt catalysts. Therefore, cross-over fuels or 02 leads to severe mixed reactivity at Pt electrodes and thus, minimal output voltage and power of the fuel cells.

To address the challenges of fuel/02 cross-over, the high cost of Nafion mem-branes, and the requirement of acid resistance for catalysts, one promising

ap-proach is to develop selective electrocatalysts which enable the elimination of ion-conducting membranes. Here the selectivity stresses catalysts' tolerance to the re-actants of the other electrode chamber. Specifically, a selective cathode catalyst only catalyzes the ORR even in the presence of fuels and a selective anode catalyst only catalyzes the fuel oxidation reaction even in the presence of 02. With highly selective catalysts, there is no need to separate the cathode and anode chambers of a fuel cell. Instead, the fuels and 02 can be introduced together into one reaction chamber. The resulting mixed-reagent fuel cell is less expensive and displays more flexibility in device design. 2021 However, it is challenging to achieve high

selectiv-ity for electrocatalysts because the driving forces for the desired fuel cell reactions are indeed much smaller than the driving forces for the parasitic undesired reac-tions (Figure 1-1, red). To address the selectivity challenge, this thesis investigated several approaches to developing selective cathode and anode catalysts for fuel cells.

1.1

Transition metal chalcogenides are highly selective

catalysts for the oxygen reduction reaction

The family of transition metal chalcogenides (MxSy, MxSey, and MxTey) has been widely investigated as ORR catalysts.22-34 Remarkably, the ORR activity of these transition metal chalcogenides is not affected by cross-over fuels such as methanol, formic acid, and H2. The high selectivity of these chalcogenides to ORR catalysis is in stark contrast to Pt-based materials which catalyze both ORR and fuel oxida-tion reacoxida-tions unselectively. Addioxida-tionally, transioxida-tion metal chalcogenides display high tolerance to electrolyte ions such as sulfate and phosphate, and impurities dissolved in the electrolyte such as CO. In contrast, the surfaces of Pt catalysts are strongly adsorbed and poisoned by electrolyte ions and impurities.35-37 Therefore, the high selectivity to ORR and high tolerance to electrolyte impurities make tran-sition metal chalcogenides promising candidates of cathode catalysts for fuel cells. A number of studies try to investigate the origin of the high selectivity to ORR of transition metal chalcogenides in the presence of fuels.38 _{For example,} experi-mental results suggest that ruthenium sulfides and selenides are intrinsically poor catalysts for methanol oxidation.39 _{Indeed, methanol oxidation reaction does not}

onset until 1.1 V ( all potentials are reported versus the reversible hydrogen elec-trode, RHE, unless otherwise noted.). This is in stark contrast to Pt catalysts where methanol oxidation onsets at 0.6 V, right in the ORR potential window. The au-thors also noticed that the methanol oxidation activity of Ru-based catalysts are quickly competed by the oxygen evolution reaction (OER) when the potential in-creases, while the OER activity of Pt is quite low with an onset potential >1.5 V. Therefore, the poor activity of Ru-based materials for methanol oxidation is at-tributed to their strong interaction with water which disfavors the adsorption and oxidation of methanol. Computationally, previous studies reveal that the onset potential of methanol oxidation is positive of the onset potential of ORR for Co-and Ru-based sulfides Co-and selenides, suggesting that methanol is not activated in the ORR potential range.4 0 _{In contrast, for Pt-based catalysts, the onset potential of}

methanol oxidation is negative of the onset potential of ORR. As a result, methanol is readily oxidized at the Pt surfaces under ORR potential window. These previ-ous studies suggest that transition metal chalcogenides display intrinsically poor activity of methanol oxidation whose onset potential is positive of the ORR po-tential window. The poor activity of transition metal chalcogenides for methanol oxidation has been attributed to high d-state density. However, it remains chal-lenging to develop a concise correlation between methanol oxidation activity and chalcogenides structure.

The well-studied chalcogenides are both based on precious metals such as Ru and Rh, and earth-abundant metals such as Co and Ni. The Ru- and Rh-based chalcogenides display comparable ORR activity to Pt in strong acid solution. It has been found that the ORR activity of RuSe nanoparticles are dramatically affected by the structure and subtle changes in the synthesis.41 _{Therefore, it is critical to} precisely control the synthetic conditions in order to achieve high activity for ORR

caLa1ys1. 11h L U and R-based \can cha cogCenids1 e prpad by thL eL1A sOLU-sta

synthesis in which a mixture of high-purity elemental metals and chalcogens are sealed in quartz tubes and sintered at elevated temperatures (around 1000 OC).27,4 2

To prepare metal chalcogenides under milder conditions, solution-based strategies have been developed. One route involves the decomposition of carbonyl precur-sors in organic solvents.4344 _{However, this is not ideal because the carbonyl} com-pounds are expensive and the yield of the products is only 40-60%. Alternatively, ruthenium(III) salts and SeO2 are used as the precursors which are reduced by vari-ous reducing agents such as NaBH4 and Li(Et)3N.45,46 These solution-based strate-gies typically require organic solvents and strong reducing agents, which makes it difficult to produce precious metal chalcogenides on an industrial scale.

Comparing to Ru- and Rh-based chalcogenides, earth-abundant metal genides are much cheaper and easy to handle with. The Co- and Ni-based chalco-genides can be prepared by solvothermal/hydrothermal methods using metal chlo-ride salts and thiourea as precursors which are typically heated at 200 'C. These solution-based preparation methods are ideal to fabricate nanoparticulate metal

chalcogenides with well-controlled shapes, sizes, and compositions.47,48 Addition-ally, solid-state synthesis is another common technique to prepare first-row tran-sition metal chalcogenides. For a typical fabrication process, elemental metal and sulfur precursors are mixed in stoichiometry and heated at elevated temperatures (800-1000 C) to prepare the desired phases.49_-53 _{Despite the high temperature} re-quired, solid-state synthesis is a powerful technique to prepare phase-pure chalco-genides. Electrodeposition and plating are also widely used to prepare thin-film first-row transition metal chalcogenides. For these techniques, a conductive sub-strate is used as the working electrode which is polarized in an aqueous solution of metal salts and thiourea precursors. 54-59 A special technique of electrodeposition is electrochemical atomic layer deposition (E-ALD) which enables the deposition of metal chalcogenides in a layer-by-layer fashion.60-62 E-ALD exploits the strong interaction between metal and chalcogen which leads to an under potential depo-sition (UPD) phenomenon that the metal (chalcogen) grows into a monolayer at the surfaces of chalcogen (metal) at a potential positive of metal (chalcogen) bulk electrodeposition. With the technique of E-ALD, a structure of alternating metal-chalcogen layers is generated. However, most of the well-established E-ALD pro-cedures are to synthesize II-VI and IV-VI semiconductor thin films, and there are only very few examples for late transition metal chalcogenides.33,63,64

While the precious metal (Ru and Rh)-based catalysts are chemically stable across a wide pH range, earth-abundant metal-based catalysts are more facile. 65-67 These first-row transition metal chalcogenides can dissolve in strong acid or can be converted into hydroxides/oxides in strong base. 57 They are also prone to oxi-dation and reduction under electrochemical bias. Even under mild pH, these ma-terials can undergo surface reconstruction under catalytic conditions. Therefore, stability and durability tests are necessary when employing earth-abundant metal chalcogenides as electrocatalysts, or the structural instability will interfere with the structural-activity analysis. For example, a number of studies employ Co, Ni, and Fe-based sulfides to catalyze the oxygen evolution reaction (OER) under alkaline media. However, most of these first-row transition metal sulfides are converted

into the corresponding metal oxides under the catalytic conditions. As a result, the measured OER activity is actually for metal oxides coated metal sulfides. While the structural conversions of metal sulfides under OER conditions have been well documented,68 70 there are only a few records for the catalyst dynamics under ORR conditions. Since earth-abundant metal chalcogenides are selective catalysts for ORR and highly tolerant to cross-over fuels and electrolyte impurities, it is nec-essary to investigate their structural-activity correlations in order to improve the catalytic activity and guide the design of new catalysts.

Here we present an example of Ni3S2 surface reconstruction in neutral pH

un-der ORR conditions in Chapter 2. We also demonstrate that the surface dynamics of first-row transition metal chalcogenides can be well predicted from thermody-namic properties of the material. Furthermore, our results indicate that the ORR activity is governed by the local coordination environments of the active center (metal sites) instead of the long-range ordering. These findings will facilitate the design of first-row transition metal chalcogenides catalysts and the prediction of oxygen reduction activity of a new material.

1.2 Mixed electron-proton conductor membranes

en-able selective electrocatalysis

While transition metal chalcogenides represent a small group of electrocatalysts which exhibit intrinsic selectivity to ORR or fuel oxidation reaction, most catalysts display poor selectivity. Particularly, the majority of catalysts for fuel oxidation reaction are Pt- and Pd-based materials which are also active ORR catalysts. 71-76 Previous attempts to suppress the ORR activity via surface modifications are not sufficient to construct a selective anode catalyst. 77-80 This is because the onset po-tential of fuel oxidation is slightly positive of 0 V where the overpopo-tential for the ORR is almost 1.23 V. As a result, despite the suppression of the ORR activity by surface modifications, the catalyst still display mixed reactivity at fuel oxidation potentials in the presence of

02-Mixed Electron-Proton Conductor Electrocatalysis

H

(gas)

e-

H

oxidation catalysis

H-* H.

25 Pt/WO electrode

20

Rapid

H

Charge

E

diffusio 15

-ifusion separation

_E

₁₅

_Ptlsolution

H++ e~-

interface

RLS

0

Bond

H-WO

Electrolyte

0.0

0.2

0.4

activation

-

EIV vs RHE

Figure 1-2: Schematic representation of catalyst/MEPC membrane composite elec-trode and the comparison of H2 oxidation current for the catalyst/MEPC mem-brane configuration and traditional Pt I solution configuration.

To develop a selective fuel oxidation catalyst with high tolerance to 02, we establish a new configuration of electrocatalysis which employs a mixed electron-proton conductor (MEPC) membrane to separate the catalyst and electrolyte. In this configuration, the soluble substrates and impurities in the electrolyte are

re-jected by the MEPC membrane from contacting the catalyst. Only electron-proton pairs, or H-atom equivalents, which are the key intermediate in any fuel cell re-actions, are transported through the membrane (Figure 1-2). As a result, the cata-lyst/MEPC membrane configuration displays great tolerance to soluble substrates and impurities in the electrolyte.

Mixed electron-proton conductors (MEPCs) are a family of metal oxides which, as their name indicates, display good conductivity for both electrons and pro-tons. 81-87 MEPCs are classified into three categories based on their nature of

con-ductivity.88 (1) Some proton-conducting metal oxides exhibit reducible metal cen-ters. These materials are partially reduced under reductive environments (e.g., H2) so that the conduction band is partially filled to give rise to electron conductivity. A particularly interesting sub-category of these reducible metal oxides is that which can intercalate a significant number of H-atom equivalents, with electrons inserted into the conduction band of the material and protons located adjacent to oxygen

AtomsN*" c. Maxry ME1Pfc b-cllyncgn tl is categooryrir sc1lh as MnO2, 3 W ,M03, etc. (2)

"L1AJkL~7 LO M "LI LL-J L - L'.5'. J j 2'_ LALX

Some intrinsic metallic conductors can display proton conductivity to some extent. However, the population of protons being transported through a metallic conduc-tor is typically small. A prospective material is LaMnO3, the conductivity nature of which is still under investigation. (3) Some materials with crystallographic pro-tons may display mixed conductivity when they are either oxidized to give proton vacancies and electron holes or reduced to give interstitial protons and electrons. A simple example of this category is the Ni(OH)2-NiOOH system.

Among the three categories of MEPCs, we are particularly interested in the H-intercalating metal oxides due to their high electron-proton conductivity and a large number of H-atom equivalents intercalated into the materials. These MEPC materials provide a powerful platform to spatially separate the catalyst and elec-trolyte while only transporting H-atom equivalents. In MEPC electrocatalysis, re-gardless of the type of substrates, the bonds of the substrates are activated at cat-alyst I MEPC membrane interface and electron-proton charge transfer take place at MEPC membrane I electrolyte interface. The high H-diffusivity of MEPC

mem-branes effectively translates the electrochemical driving force for electron-proton charge transfer to the chemical driving force for substrate bond activation. Since the two critical functions of electrocatalysis -bond activation and charge transfer

- are spatially segregated in MEPC electrocatalysis, each of the two functions can be optimized independently. Additionally, the catalyst/electrolyte incompatibility is no longer an issue in MEPC membrane-mediated electrocatalysis. The MEPC membranes provide more flexibility for catalyst design and electrolyte choice, thus enabling greater durability of the catalytic system. The new catalyst/MEPC mem-brane configuration also overcomes the solubility limit of gaseous substrates in electrolyte, leading to a higher rate of electrocatalysis (Figure 1-2). In Chapter 3, we present an example of a Pt/W03 MEPC membrane to catalyze the hydrogen oxidation reaction (HOR). We show that the HOR current of the Pt/W03 compos-ite electrode is not affected by 02, CO, and Cu2' dissolved in the electrolyte.

High electron-proton conductivity is crucial for efficient MEPC electrocatalysis. Taking one of the best known MEPC, W03, as an example, the H-atom equivalent diffusivity in W03 is as high as 10-6 cm2 s-1 at room temperature, 89-91 making W03 a promising MEPC membrane candidate. Without electron insertion, the He-only diffusivity of W03 is only 10-9-10-12 cm2 s--1,92-95 orders of magnitude smaller than H-atom equivalent diffusivity. The great enhancement of proton con-ductivity when both proton and electron are transported is attributed to the corre-lated movement of protons and electrons. A possible explanation for the enhance-ment is that the free moveenhance-ment of correlated electrons compensates for the local electrostatic field generated by the movement of ions. Indeed, the enhancement effect has been observed for a number of solid-state materials with mixed electron-ion conductivity.96-98 The enhanced proton conductivity by correlated electrons makes MEPC electrocatalysis feasible at low temperatures.

The electron-proton charge transfer is driven by the electrochemical field at the MEPC membrane I electrolyte interface. Still taking W03 as an example, the electron-proton pairs, or H-atom equivalents, are intercalated into W03 under a potential negative of the H chemical potential of H-W03 (EH), and are

deinter-calated under a potential positive of EH. The thermodynamics and kinetics of H intercalation and deintercalation are complicated by the dependence of EH on the concentration of H in W03 and the space-charge region generated during dein-tercalation. It has been noticed that EH is not a constant, but is dictated by the H concentration intercalated in W03 which can be defined as the atomic ratio of H:W (XH).9 9 Previous studies have found that EH decreases monotonically with increasing XH following the equation:

R T xH

EH = a+ - xH - .

RTb(X _{F I - XH}

As a result, as the electrochemical H intercalation proceeds, XH increases, lead-ing to a decrease of EH, which further causes a decrease in the driving force for H

intercalation. Indeed, the decrease in EH and electrochemical driving force with increasing XH is often denoted as a "back potential" which controls the rate of electrochemical H intercalation.92_,9 9

,10 0 Therefore, the rate-controlling step of elec-trochemical H intercalation is electron-proton combination at the W03 I electrolyte boundary: W03(s) + H+(aq) + e- H-W0 3(s).

For electrochemical H deintercalation, the rate-controlling step is proton trans-port in a space-charge region at the W03 I electrolyte boundary. 91,92 At a potential

positive of EH, H-atom equivalents are electrochemically extracted out of W03,

with electrons moving to the external circuit and protons to the electrolyte. Since electrons cannot transport in electrolyte, electrons and protons are separated at the

W03 I electrolyte boundary, leading to a space-charge region at the interface where

protons are the only charge carrier. We know that the rapid H-diffusivity within

W03 is due to the enhanced proton conductivity by the correlated movement of electrons. In the absence of electrons, proton-only diffusivity in the space-charge region is orders of magnitude lower than H-diffusivity in the bulk W03. As a re-sult, the slow movement of protons in the space-charge region at the W03 I electrolyte interface controls the overall rate of electrochemical H deintercalation. The com-plicated thermodynamics and kinetics of electrochemical H intercalation and

dein-tercalation should be taken into consideration when we try to develop qualitative and quantitative models of MEPC membrane mediated electrocatalysis.

On the other hand, substrate bond activation takes place at the catalyst I W03

interface. Taking the HOR catalysis at Pt surfaces as an example, the H-H bond breaks readily at Pt surfaces to generate Pt-H. When the Pt catalyst is contacted by an MEPC membrane such as W03, H-atoms adsorbed at Pt surfaces migrate across the Pt I W03 interface to intercalate into W03. The surface migration of H-atom equivalents is denoted as H-spillover. Typically, the rate of H-spillover is significantly lower than the rate of H2 activation at Pt surfaces. 101-104 Therefore, it is critical to investigate the rate-controlling factors of H-spillover in order to de-velop a quantitative model for the overall electrocatalysis. The rate of H-spillover depends strongly on the H concentration of W03 in that a small XH leads to a high rate of H-spillover, and vice versa. Since the rate of electron-proton charge transfer also depends on XH, the two functions of substrate bond activation and electron-proton charge transfer are correlated by XH. Remarkably, the value of XH during electrocatalysis can be conveniently determined from the H chemical potential of

W03.105 In Section 3.2, we integrate these factors to develop a quantitative model

for the HOR catalysis.

E/VvsRHE CoO NiO, _{1.23 V OH} 20 CeO. ORR WO HOR " 00.00 V H+/H2

MOOK HER: CO2products

TiO CO2RR - C o

Figure 1-3: Proposed MEPC membranes and matching reactions.

A large variety of solid-state metal oxides exhibit mixed electron-proton con-ductivity, enabling MEPC electrocatalysis over a wide range of reactions. One principle of choosing MEPC membranes is that the H chemical potential should

match the redox potential of the reaction. For example, the H chemical potential of W03 spans between -0.2 and 0.5 V, perfectly matching the redox potential of the H+/H2 pair.99 Therefore, W03 is an ideal candidate to mediate the hydrogen oxidation catalysis. On the other hand, an MEPC membrane with more negative H chemical potentials such as TiO2 is promising to mediate the hydrogen evolution reaction (HER). 10_6-108 _{Some proposed examples of matching MEPC materials and}

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