Controlling kinetic branching in C0₂-to-fuels catalysis

Controlling Kinetic Branching in C02-to-Fuels Catalysis

by

Anna Lydia Nakamura Wuttig

B. A. Chemistry

Princeton University, 2013

SUBMITTED TO THE DEPARTMENT OF CHEMISTRY IN PARTIAL FULFULLMENT OF

THE REQUIREMENTS FOR THE DEGREE OF

DOCTOR OF PHILOSOPHY IN CHEMISTRY

AT THE

MASSACHUESETTS INSTITUTE OF TECHNOLOGY

June 2018

0 2018 Massachusetts of Institute of Technology. All rights reserved.

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Department of Chemistry

May 11, 2018

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Accepted by: ...

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Yogesh Surendranath

Paul M Cook Career Development Assistant Professor of Chemistry

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Thesis Supervisor

MASSACHUSETTS INSTITUTE OF TECHNOLOGY

JUN 2

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Robert W. Field

Haslam and Dewey Professor of Chemistry

Chairman, Committee of Graduate Students

This Doctoral Thesis has been examined by a Committee of the Department of Chemistry as

follows:

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Alexander Radosevich

Associate Professor of Chemistry

Chairman, Thesis Committee

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Yogesh Surendranath

Paul M Cook Career Development Assistant Professor of Chemistry

Thesis Supervisor

Signature redacted______

Daniel Nocera

Patterson Rockwood Professor of Chemistry at Harvard University

Committee Member

Controlling Kinetic Branching in C02-to-Fuels Catalysis

by

Anna Lydia Nakamura Wuttig

Submitted to the Department of Chemistry

on May 11, 2018 in Partial fulfillment of the

requirements for the Degree of

Doctor of Philosophy in Chemistry

Abstract

While electrified transition-metal surfaces mediate the synthesis of carbonaceous

products from C02, these processes suffer efficiency losses due to the multitude of products

accessible over a narrow potential range. Selective product formation requires knowledge and

control over various branch points in the reaction pathway. In this work, we will present two

specific branch points: (1) the requirements for selective activation of C02 over H' to form the

two-electron reduced CO product; and (2) the requirements for the accumulation of

surface-bound CO species that can be reduced to higher order products beyond CO. Using model Au and

Cu electrocatalysts, we uncover mechanistic insights into these branch points.

We identify the differential proton-coupling requirements for CO

versus H' activation

on polycrystalline Au surfaces that establish a mechanistic basis for CO versus H2

product

selectivity. Electrokinetic data are consistent with a mechanism of CO production involving

rate-limiting single electron transfer to CO2 with concomitant adsorption to surface-active sites,

followed by one electron, two proton transfer, and CO liberation from the surface. In contrast,

the data suggest a

evolution mechanism involving rate-limiting single electron transfer

coupled with proton transfer from bicarbonate, hydronium, and/or carbonic acid to form

adsorbed H species, followed by sequential one electron, one proton transfer or H recombination

reactions.

We elucidate the differential CO electrosorption dynamics on polycrystalline Au and Cu

surfaces using temperature-dependent in-situ surface-enhanced infrared absorption spectroscopy,

establishing a mechanistic basis for potential-dependent CO binding. On Au surfaces, we

observe that reversible linearly-bonded CO electrosorption is a water substitution process, where

bound CO species readily dissociate from the surface upon negative potential bias. Conversely,

labile CO species accumulate upon negative potential bias on Cu surfaces via a charge

displacement reaction with carbonate, providing a pool of reactant primed for further reduction

to higher order products. The enthalpy and entropy of electrosorption are also quantified.

Thesis Supervisor: Yogesh Surendranath

Title: Assistant Professor of Chemistry

The Artist will find in Science a more stable foundation,

And the Scientist will draw from Art a more certain intuition.

Contents

I. Defining Two Kinetic Branchpoints in C02-to-Fuels Catalysis --- 16

1. 1. References 19

II. Selective Activation of CO2 to CO over H' to H2 22

11.1. Role of Impurity Ions --- 23

11.1.1. Introduction 23

11.1.2. Group 11 Metal Surfaces Accumulate Impurity Metal Poisons during

Electrocatalysis --- 26 11.1.3. Impurity Ion Coordination Inhibits Impurity Deposition and Enables Sustained

Electrocatalysis --- 30

11.1.4. Conclusions 38

11.1.5. Experimental Detail --- 38

11.2. Proton Coupling Requirements for Selective CO2 Activation --- 45

11.2.1. Introduction___._ 45

11.2.2. CO Evolution Rate is Independent of the Proton Donor Environment --- 49 11.2.3. H2 Production is Dependent on the Proton Donor Environment --- 68

11.2.4. Conclusions and Mechanistically-Informed Improvements to Catalyst Design 73 11.2.5. Experimental Details --- 75

11.2.6. Simulation Details 82

11.2.7. Derivations of Diagnostic Kinetic Criteria --- 87

11.3. References 113

III. Towards Further Activation of CO to Higher Order Products --- 121 III. 1. Surface Adsorbate Population and Dynamics during CO2 Reduction ... 121

111.1.1. Introduction____- 121

111.1.2. Surface Adsorbate Population and Dynamics on Electrified Polycrystalline Au Surfaces during CO2 Reduction --- 125

111.1.3. Synthesis and Characterization of Cu Films for Surface Enhanced Infrared Absorption Spectroscopy --- 139 111.1.4. Surface Adsorbate Population and Dynamics on Electrified Polycrystalline Cu

Surfaces during CO2 Reduction --- 143

111.1.6. Experimental Details --- 158

111.2. Surface Dynamics and Thermodynamics of Reversible CO Electrosorption - 171

111.2.1. Introduction 171

111.2.2. Reversible COatop Electrosorption is an Interfacial Water Substitution Process on

Reconstructed Au Surfaces 173

111.2.3. Reversible CO Electrosorption is a Proton-Coupled Carbonate Charge Displacement Reaction on Polarized Cu Surfaces --- 187 111.2.4. Determination of Enthalpy and Entropy of CO Electrosorption on Au --- 193

111.2.5. Conclusions and Future Directions 207

111.2.6. Experimental Details --- 208

List of Figures

Figure 11.1 Surface chemistry following CO2RR catalysis on group 11 metals ₂₆

determined by narrow scan XPS and voltammetry

Figure 11.2 Survey XPS spectra of group 11 metals surfaces following CO2RR 28

Figure 11.3 Cyclic voltammetry of group 11 metals surfaces following CO2RR at 28

varying durations

Figure 11.4 Cyclic voltammetry of group 11 metals surfaces following CO2RR 28

with addition of ZnSO4, CuSO4, and/or Pb(N03)2

Figure 11.5 Faradaic efficiencies for CO2RR and HER product formation on Au ₃₁

foil in electrolytes of varying purity

Figure 11.6 Faradaic efficiencies for CO2RR and HER product formation on Au ₃₁

foil in pre-electrolyzed electrolyte

Figure 11.7 Partial current densities for CO2RR and HER product formation on ₃₂

Au foil in electrolytes of varying purity

Figure 11.8 Faradaic efficiencies for CO2RR and HER product formation on Ag ₃₃

foil in electrolytes of varying purity

Figure 11.9 Partial current densities for CO2RR and HER product formation on ₃₄

Ag foil in electrolytes of varying purity

Figure 11.10 Faradaic efficiencies for CO2RR and HER product formation on Cu ₃₅

foil in electrolytes of varying purity

Figure 11.11 Surface faceting and CO2RR activity of SEIRAS-active Au film, ₅₄

surface faceting of Au foils and RCE

Figure 11.12 Cu UPD on SEIRAS-active Au film, Au foils, and Au RCE 55

Figure 11.13 Kouteckf-Levich plots of CO partial current density operated in C02- 56

saturated 0.1 M NaHCO3

Figure 11.14 Electrokinetic data for CO evolution catalysis on Au surfaces 57

Figure 11.15 Comparison of CO2RR conducted in bicarbonate electrolytes with ₆₀

and without phosphate

Figure 11.16 Simulated bicarbonate concentration gradients and current densities ₆₀ for CO2RR reaction sequence involving CPET or ET in 0.1 M electrolyte

Figure 11.17 Simulated bicarbonate concentration gradients and current densities 61

for CO2RR reaction sequence involving CPET or ET in 1 M electrolyte

Figure 11.18 Simulated carbonate concentration gradients and current densities for 62

CO2RR reaction sequence involving CPET or ET

Figure 11.19 Simulated CO2 concentration gradients and current densities for ₆₃ CO2RR reaction sequence involving CPET or ET

Figure 11.20 Proposed mechanism for CO evolution and concurrent H2 evolution 67

on Au surfaces

Figure 11.21 Kouteckf-Levich plots of H_{2 partial current density for C02 68}

reduction occurring in various bicarbonate buffer concentrations

Figure 11.22 Tafel plots of H2 partial current density for CO₂ reduction occurring ₆₉ in bicarbonate buffer

Figure 11.23 _{Electrokinetic data for concurrent H}

2 evolution catalysis on Au 70

surfaces

Figure 11.24 Faradaic efficiencies for CO and H2 as a function of bicarbonate 74 concentration

Figure 111.1 Potential-dependent adsorption dynamics of principle electrolyte

species and CO on Au films determined by SEIRAS, CO and C02-saturated 0.1 M 126

NaHCO₃

Figure 111.2 Stark tuning slopes for adsorbed CObridge and COatop on Au in CO and 127

C02-saturated 0.1 M NaHCO3

Figure 111.3 Time-dependent SEIRA spectra of Au and corresponding integrated 129

band intensities of COatop and CObridge in CO2-saturated 0.1 M NaHCO3 Figure 111.4 Potential-dependent adsorption dynamics of principle electrolyte

species and CO on Aumod films determined by SEIRAS CO and C02-saturated 0.1 131

M NaHCO3

Figure 111.5 Cyclic voltammetry of Au rotating cone electrode following CO2 134

reduction catalysis

Figure 111.6 Cyclic voltammetry of Au rotating cone electrode following CO2 135

reduction catalysis with and without surface oxide formation

Figure 111.7 Chronoamperometry traces of oxide-derived Au surfaces in 0.95 atm 135

C02-saturated and 0.05 atm Ar-saturated 0.1 M NaHCO3

Figure 111.8 Cyclic voltammetry of oxide-derived Au surfaces following CO₂ 136

Figure 111.9 Cyclic voltammetry of oxide-derived Au surfaces following CO2 137

reduction catalysis in the presence of CO

Figure 111.10 Atomic force microscopy images of Si substrates along the synthetic 139 sequence used to prepared SEIRAS-active Cu films

Figure 111.11 X-ray photoelectron spectra of Cu thin film on Si substrate 140 Figure 111.12 Atomic force microscopy image and corresponding height profile of 140 Cu thin film on Si substrate

Figure 111.13 Atomic force microscopy image of Cu film after polarization and 141 cyclic voltammetry of Cu film on Si substrate

Figure 111.14 Roughness factor measurements of Cu film on Si substrates 142 Figure 111.15 Cyclic voltammograms collected during SEIRAS acquisition in 1.0 143 atm C02-saturated 0.1 M NaHCO3 on Cu films

Figure 111.16 SEIRA spectra collected in 1.0 atm C02-saturated 0.1 M NaHCO3 44 and during choronoamperometry traces with '3CO addition on Cu films

Figure 111.17 SEIRA spectra collected in 1.0 atm C02-saturated 0.1 M NaHCO_{3 145} with an expanded spectral window on Cu films

Figure 111.18 Final SEIRA spectra collected in 1.0 atm C02-saturated 0.1 M 145 NaHCO3 over sequential runs on Cu films

Figure 111.19 Time-dependence of the evolution of the integrated band intensity ₁₄₇ corresponding to electrogenerated CO on Cu film surfaces

Figure 111.20 CO adsorption profiles at pH 6.9 with varying CO partial pressure 148 on Cu films

Figure 111.21 Cyclic voltammograms recorded on Cu films at pH 6.9 with varying 148 CO partial pressures

Figure 111.22 Faradaic efficiencies for C2H4, CO, H2, and CH4 on Cu films 150 prepared on Si substrates

Figure 111.23 CO adsorption profiles on polarized Cu films at varying pH 151

Figure 111.24 CO adsorption profiles at pH 10.1 at varying CO partial pressures 152 Figure 111.25 Cyclic voltammetry and SEIRA spectra recorded at pH 7.8 and pH 153 10.1 in (bi)carbonate electrolytes with varying CO partial pressure

Figure 111.26 Potential dependence of integrated band intensity of IR peak

spanning 1880 to 1760 cm-1 in 0.1 M (bi)carbonate buffer, pH 10.1, with 0.25 atm 154

CO

Figure 111.27 Calculated dominant C-O stretching vectors for surface-bound ₁₅₄ OCCO and CHO on Cu(100)

Figure 111.28 Cyclic voltammogram and corresponding SEIRA specta in Ar- 174

saturated 0.1 M (bi)carbonate electrolyte, pH 9.2, on Aumod films

Figure 111.29 Potential-dependent adsorption dynamics of principle electrolyte

species and CO on Aumod films determined by SEIRAS, CO-saturated 0.1 M 175

(bi)carbonate buffer, pH 9.2

Figure 111.30 Stark tuning slopes for adsorbed COhollow, CObridge _{and COatop on 176}

Aumod in CO-saturated 0.1 M (bi)carbonate buffer, pH 9.2

Figure 111.31 Cyclic voltammograms and SEIRA spectra on Aumod _{in 0.1 M 180}

(bi)carbonate buffer, varying CO partial pressure, pH 9.2, 24' C

Figure 111.32 Cyclic voltammograms and SEIRA spectra on Aumod in CO and 181

C02-saturated 0.1 M NaHCO3 with Pt counter

Figure 111.33 Integrated band intensity of COatop, _{interfacial water, and adsorbed 182}

(bi)carbonate on Aumod films in CO-saturated 0.1 M (bi)carbonate buffer, pH 9.2

Figure 111.34 Proposed stoichiometry of potential-dependent CO binding on Au 186

and Cu surfaces

Figure 111.35 Cyclic voltammograms and SEIRA spectra on Cu films in Ar- 187

saturated 0.1 M (bi)carbonate electrolyte, pH 9.2

Figure 111.36 Cyclic voltammograms, SEIRA spectra, and integrated band

intensities of bound-CO, interfacial water, and carbonate on Cu films in

CO-saturated 0.1 M (bi)carbonate electrolyte, pH 9.2, varying CO pressure, cyclic 188

voltammograms of polycrystalline Cu foils in CO-saturated 0.1 M (bi)carbonate electrolyte, pH 9.2

Figure

111.37

saturated 0.1 M (bi)carbonate buffer

Figure 111.38 Changes in CO solubility and Ag/AgCl reference electrode potential 195

as a function of temperature

Figure 111.39 Cyclic voltammograms and SEIRA spectra on Aumod in 0.1 M 199

(bi)carbonate buffer, varying CO partial pressure, pH 9.2, 10 C

Figure 111.40 Cyclic voltammograms and SEIRA spectra on Aumod in 0.1 M 200 (bi)carbonate buffer, varying CO partial pressure, pH 9.2, 38" C

Figure 111.41 Cyclic voltammograms and SEIRA spectra on Aumod in 0.1 M 201 (bi)carbonate buffer, varying CO partial pressure, pH 9.2, 50" C

Figure 111.42 Integrated band intensity of COatop _{on Aumoc in 0.1 M (bi)carbonate 202} buffer, varying CO partial pressure, pH 9.2, 10, 38, and 50 C

Figure 111.43 Langmuir adsorption isotherm of COatop on Aumod _{at pH 9.2, 10, 24, 203} 38, and 50' C at the peak potential

Figure 111.44 Integrated COatop and COridge _{band intensities as a function of 203} normalized total CO coverage

Figure 111.45 COatop band on Au films as a function of exogenous CO partial 204 pressure at -0.17 V vs SHE

Figure 111.46 Langmuir adsorption isotherm of COatop on Aumod _{at pH 9.2, 10, 24, 205} 38, and 50' C at varying applied potential

Figure 111.47 Temperature-dependence of standard free energy of electrosorption ₂₀₆ of COatop on Au at varying applied potential

Figure 111.48 Experimental setup for temperature-dependent SEIRA 210 measurements

List of Tables

Table 11.1 Diagnostic criteria of proposed paths for CO production from CO2 with 50 HCO3- as a proton donor

Table 11.2 Diagnostic criteria of proposed paths for CO production from CO2 with 51

H20 as a proton donor

Table 11.3 Diagnostic criteria of proposed paths for CO production from CO2 with 52

H30' as a proton donor

Table 11.4 Input parameters used for simulations 85

Table 111.1 Calculated vibrational modes of surface-bound CHO and OCCO on 155

Cu(l00) surface

Table 111.2 Table of fitting parameters for Langmuir adsorption isotherm of COatop 204

on Au film surfaces at the peak potential

Table 111.3 Table of fitting parameters for Langmuir adsorption isotherm Of COatop 205 on Au at 50 mV positive of the peak potential

Table 111.4 Table of fitting parameters for Langmuir adsorption isotherm of COatop 206

on Au at 50 mV negative of the peak potential

List of Schemes

Scheme 1.1 Possible catalytic cycles depicting sequential C02 and CO activation 18 Scheme 11.1 Possible pathways for involvement of bicarbonate anion for CO2 47

electroreduction on an electrode surface

List of Abbreviations

CO2RR - C02 reduction reaction

HER - hydrogen evolution reaction OER- oxygen evolution reaction ORR - oxygen reduction reaction EDTA -ethylenediaminetetraacetic acid

Ci - unpurified C02-saturated 0.1 M NaHCO3 electrolyte RHE - reversible hydrogen electrode

XPS - X-ray photoelectron spectroscopy

ICP-MS -Inductively Coupled Plasma Mass Spectrometry CV - cyclic voltammogram

UPD - underpotential deposition

PCET - proton-coupled electron transfer RLS - rate-limiting step

PT - proton transfer

CPET - concerted proton electron transfer ET - electron transfer

Ox - surface coverage of species x

PCO2 - partial pressure of CO2 Pco - partial pressure of CO

p - electrochemical reaction order / - overpotential

E - applied potential

jo - exchange-current density

SHE - standard hydrogen electrode RCE - rotating cone electrode

SEIRAS - surface enhanced infrared absorption spectroscopy jco - rate of CO evolution

GC - gas chromatography

ECSA - electrochemically active surface area

8 - symmetry factor Pi - phosphate buffer

jH2 - rate of hydrogen evolution jTOT - total observed current density

KL - Kouteckf-Levich

SERS - surface-enhanced Raman spectroscopy FTIR - Fourier transform infrared

ATR - attenuated total reflectance

COatop - linearly bonded CO

CObridge - 2-fold bridging CO

COg - gaseous CO

COhollow - 3-fold bridging CO

Fx - surface concentration of species x AFM - atomic force microscopy IR - infrared

CObound - surface-bound CO

PZC - potential of zero charge MT - mass transfer

Part I

Defining Two Kinetic Branchpoints in C02-to-Fuels Catalysis

Contents

I.

I. Defining Two Kinetic Branchpoints in C02-to-Fuels Catalysis

Reproduced in partfrom Wuttig, Yaguchi, Motobayashi, Osawa, Surendranath, Proc. Nati. Acad. Sci, USA, 2016, 113 (32), pp E4585-E4593.

Product selectivity is a principal design consideration for the development of practical

catalysts. Catalyst selectivity is dictated by: (1) the relative free-energy barriers for progress along

competing reaction pathways; and (2) the relative rates of reactant delivery to active sites.

Enzymes fine-tune these parameters with exquisite precision to achieve selectivity.

Nature

augments the coordination environment of metallo-cofactor active sites to optimize the binding

strengths of reaction partners and to pre-organize reaction participants towards low-barrier

pathways.

Additionally, many active sites reside at the terminus of molecular channels that gate

the coordinated delivery of substrates

_'

_{required for selective transformations. Efforts to prepare}

artificial catalysts with product selectivities rivaling that of Nature require a detailed understanding

of these factors.

Currently, our understanding of how to systematically modulate selectivity in

heterogeneous catalysts remains poor.6-

Unlike (bio)molecular catalysts, which, ideally consist of

a uniform ensemble of active sites, heterogeneous catalysts consist of a non-uniform distribution

of surface sites,

requiring an understanding of which are active and which are dormant. The

surface site distribution is strongly dependent on the surface nanostructure, oxidation state, and

degree of restructuring.

Superimposed on this distribution are the rate-limiting microscopic steps

that dictate kinetic branching ratios at surface active sites.

For catalysis mediated by electrode

surfaces, these factors are subject to change as a function of the applied potential,"

_{and the}

log-linear relationship between potential and reaction rate often introduces severe transport limitations

that further augment the intrinsic reaction kinetics". Disentangling these correlated processes is

essential to the development of selective electrocatalysts.

In this thesis, we aim to understand these correlated processes in the context of studying

electrocatalytic C02 reduction. Electrocatalytic C02 reduction allows for the storage of

intermittent renewable electricity in energy-dense carbonaceous fuels.

-

C02 reduction reactions

(CO2RRs) are most practically conducted in aqueous electrolytes, in which the undesirable

reduction of protons to H2 presents a principal selectivity challenge. Since the hydrogen evolution

reaction (HER) is thermodynamically accessible over the same potential range as nearly all CO

RR

products,1

-

selective fuel formation relies on the electrode's ability to control the relative rates

of these competing pathways. Despite the ubiquitous nature of competing HER in nearly all

reported electrocatalytic CO

RR schemes, detailed study of the two relative rates remains elusive.

In Section [II], we pinpoint the kinetic requirements to obtain selective C02 activation over H' to

produce CO, Scheme 1.1, blue, on model Au electrode surfaces.

00

H

HO

e-

cc

H+

CO

CO

Activation

C

Activation

\,

C-C.

CO

9e-CO

Scheme 1.1 Possible catalytic cycles depicting sequential CO2 and CO activation. Two themes discussed in this thesis

Mechanistic studies of CO2RR have defined several possible branchpoints for CO2RR catalysis to various carbonaceous products beyond CO. Although the mechanistic proposals that have been put forward thus far differ in the assignment of the rate-limiting step for the formation of each product 21_-32 _{and the surface structures responsible for these rate-limiting steps,}2 _{,29,31,33-43 in all}

cases, these models invoke a common intermediate, surface-bound CO, which precedes the formation of all higher order fuels. Despite its central role in C02-to-fuels catalysis, the energetics and dynamics of CO electrosorption to surfaces under catalytic conditions have yet to be probed experimentally. In Section [III], we will elucidate the surface dynamic processes and quantify the thermodynamics of surface-bound CO on Cu and Au surfaces in-situ, revealing the requirements to further accumulate surface-bound CO.

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Part II

Selective Activation of CO

to CO over H' to

Contents

II. Selective Activation of CO2 to CO over H' to H2

11.1. Role of Impurity Ions --- 23

11.1.1. Introduction 23

11.1.2. Group 11 Metal Surfaces Accumulate Impurity Metal Poisons during

Electrocatalysis --- 26 11.1.3. Impurity Ion Coordination Inhibits Impurity Deposition and Enables Sustained

Electrocatalysis --- 30

11.1.4. Conclusions 38

11.1.5. Experimental Details 38

11.2. Proton Coupling Requirements for Selective C02 Activation

---11.2.1. Introduction 45

11.2.2. CO Evolution Rate is Independent of the Proton Donor Environment 49 11.2.3. H2 Production is Dependent on the Proton Donor Environment 68

11.2.4. Conclusions and Mechanistically-Informed Improvements to Catalyst Design_ 73

11.2.5. Experimental Details 75

11.2.6. Simulation Details 82

11.2.7. Derivations of Diagnostic Kinetic Criteria 87

II. Selective Activation of C02 to CO over H' to H2

1. Role of Impurity Ions

Reproduced in partfrom Wuttig, Surendranath, ACS Catal., 2015, 5 (7), pp 4479-4484.

1.1. Introduction

Catalytic reactions are highly sensitive to the presence of impurities. Trace constituents in the reaction medium can have an outsized effect on reaction efficiency and selectivity by interfering with the active species or mediating turn-over by themselves. For example, metal and halide contaminants promote and poison product activity and selectivity in a myriad array of heterogeneous reactions,- have led to serendipitous catalyst discovery in organic synthesis5 6 and can dramatically alter enzymatic activity 9.

Trace impurities also play a dominant role in electrochemical energy catalysis. For example, trace Co2+ impurities have led to false-positives in the development of molecular oxygen evolution

reaction (OER) catalysts'0 and trace iron,12 and platinum" impurities have been shown to enhance the OER on heterogeneous catalysts. Likewise, Pt surfaces are readily poisoned for the oxygen reduction reaction (ORR) by various impurities4 whereas "metal-free" ORR" and C02 reductioni activity has been ascribed to trace metal impurities. Trace metal impurities have also been shown to convolute the facet-dependence of HER activity on Au surfaces.i7

The electrochemical reduction of CO2 to fuels is of growing interest because it provides an

attractive platform for the storage of intermittent renewable energy in energy dense chemical bonds.18-20 Relative to other energy conversion reactions such as ORR, OER, and HER, CO2RR

2 1-3

products thermodynamically accessible over a narrow potential range." The development of practical CO2RR catalysts requires unparalleled control over product selectivity, which can be

easily compromised by impurities that interact with or irreversibly alter the surface.

Group 11 metal surfaces are regarded as the most promising heterogeneous catalysts for this reaction because they display low to moderate overpotentials for CO production, and, in the case of copper, generate higher order products including methane and ethylene.243 2 However, planar

group 11 metal surfaces are known to lose their catalytic activity and selectivity for CO2 reduction

over the time scale of minutes to hours under steady state electrolysis.33_-39

For example, copper surfaces lose one-half of their catalytic activity for methane production within 20 min of polarization,'3 5 40

and CO production selectivity on Au34 _{and Ag}36 _{decreases within minutes of} electrolysis. Despite posing a clear obstacle to practical implementation of CO2 reduction

technologies, the mechanistic basis for this activity loss remains poorly understood.

Researchers have posited that the deactivation is due the deposition of trace metal ion2 3'40 or organic impurities40 in the electrolyte, and others have suggested that this activity loss is unavoidable, resulting from the accumulation of catalytic intermediates that poison the surface over time- 3,3,-44. Based on these hypotheses, contemporary strategies to prolonging catalyst lifetimes include periodic oxidative pulsing of the electrode to remove adsorbed organics34-36 and long-term (>9 hrs) pre-electrolysis using a sacrificial electrode to scavenge trace metal ion impurities in the electrolyte.40

The former is impractical because it progressively alters the catalyst surface structure,45 _{and pre-electrolysis has been shown, in many cases, to be ineffective}46

'47 and irreproducible35,36 and is both energy and time intensive. We note that high surface electrodes3941,47 should exhibit reduced sensitivity to trace impurities, but these systems are difficult to probe

development of new CO

RR catalysts would benefit from a clear understanding of the principle

CO

RR deactivation mechanism on these surfaces and the development of simple strategies for

sustaining catalyst activity and selectivity over time.

We demonstrate that metal ion deposition is the principle mode of catalyst deactivation for Cu,

Ag, and Au metal surfaces and show, for the first time, that catalyst deactivation can be entirely

eliminated by irreversibly coordinating trace metal ions in situ with ethylenediaminetetraacetic

acid (EDTA) or ex-situ with a solid-supported iminodiacetate resin. The high binding affinity

and rapid complexation kinetics (ki 1A010 M-Is-1)495 1 - of EDTA and solid phase analogues make

this a general strategy for maintaining high CO2RR activity over time, regardless of structure of

the catalyst.

______________________________________________ -I

1.2. Group 11 Metal Surfaces Accumulate Impurity Metal Poisons during Electrocatalysis

a

d

r I I I~ I~p~ II

Z 2p" ZT 2p Zr 2p'

1060 1030 1000 460 430 400 1060 1030 1000 960 930 900 1060 1030 1000 960 930 900

Binding Energy /eV Binding Energy 9 V Binding Energy /eV

d

so

E/VvsRHE E/VvsRHE E/VvsRHE

Figure 11.1. Surface chemistry following CO2RR catalysis on group 11 metal surfaces. Narrow scan X-ray

photoelectron spectra of Cu (a), Ag (b), and Au (c) rotating electrodes before (black) catalysis and after 45 min of CO2RR catalysis in native Ci electrolyte (green) and EDTA-containing electrolyte (red). Cyclic voltammetry (CV)

traces of rotating electrodes at 2500 RPM in native Ci electrolyte (green) or EDTA-containing Ci electrolyte (red) after (d) 120 min of CO2RR catalysis at -1.00 V on Cu; (e) 45 min of CDR catalysis at -0.90 V on Ag; and (f) 45 min of

CO2RR catalysis at -0.70 V on Au. In all cases, CV scans were recorded at 50 mV s-1 scan rate with a positive

direction of scan.

To determine if metal impurities that electrodeposit natively under CO2RR conditions, we

examined group 11 electrocatalysts by following catalysis in unpurified C02-saturated 0.1 M

NaHCO3 (Ci) electrolyte (see Section [II.1.5] for experimental details). Despite 18 MQ. cm water

measurements of copper rotating disk electrodes following

45

min of electrolysis at -1.0 V vs

RHE, (in Section [II.1], all potentials are quoted versus the reversible hydrogen electrode, RHE),

a potential typical for selective C02 reduction, reveals the build-up of zinc and lead impurities

(Figures II.1a and 2). Similarly, prolonged electrolyses of silver (-0.9 V) and gold (-0.7 V)

rotating cone electrodes reveal the accumulation of zinc and copper impurities, (Figures II.1b, c

and 11.2). Previous studies conducted at similar electrolysis potentials were unable to detect trace

metal impurity signals via XPS.

Our use of a rotating electrode serves to accelerate the rate of

diffusion-limited metal deposition, increasing the surface impurity population (see calculation in

section [11.1.6]). These XPS results indicate that all group 11 metal surfaces are subject to

contamination via impurity deposition even in cases where high purity electrolytes are employed.

In line with the XPS results, 1.1L0.1

< 0.05 pM, and < 0.2 pM, of Zn2+, Pb2+, and Cu,2+

respectively were detected by Inductively Coupled Plasma Mass Spectrometry (ICP-MS). The

thermodynamic

redox potentials

for Zn2+ (-0.54 V), Pb2+ (0.06 V), and

(0.54 V) under

these reduced concentrations are still close to or positive of the thermodynamic potentials for

CO2RR catalysis, therefore, we expect that impurity deposition will also occur on high surface

area group 11 catalysts that operate at lower CO2RR overpotentials, albeit at slower rates.

To gain further insight into changes in surface composition during CO

RR, we characterized

electrodes using cyclic voltammetry (CV) following (-1 s time delay) reductive polarization. The

first CV sweeps recorded after short (15 min) and prolonged (>45 min) CO

RR catalysis in native

Ci electrolyte reveal the progressive rise, Figure II.3a, of broad oxidative features at -0.22 V vs.

RHE and 0.26 V for Cu, Figure II.ld (green). When the same experiment was performed on Ag

electrodes, we observed a continuous rise, Figure II.3b, in broad oxidative features at -0.25 V

and 0.50 V, Figure II.le (green). Similarly, the data collected on Au reveals a rise, Figure II.3c,

of features at 0.11 V, 0.50 V, and 0.85 V with the appearance of distinct shoulders upon longer

electrolysis, Figure II.lf (green). These features are not observed prior to CO2RR (Figures

II.3a-c, black) or upon subsequent cycling of the electrode after CO

RR. Together, these observations

Co

1000 800

600

400 200

Binding Energy

/

eV

0

Figure 11.2. Survey X-ray photoelectron spectra of copper (red), silver (blue), and gold (green) electrodes following 45 min electrolysis (-1.00 V for Cu, -0.90 V for Ag and -0.70 V for Au) in untreated Ci electrolyte. Black dotted lines denote peak positions of Pb, Cu, and Zn impurities detected. All other peaks in the spectra are unchanged relative to the initial electrode prior to CO2RR catalysis.

removal

(M

CO

RR catalysis.

To assign these stripping waves to metal ion impurities, we collected CV traces following

CO

RR in Ci electrolyte, containing 50sM of various M

salts of metals detected via XPS. This

intentional introduction of an impurity metal ion allows us to probe the stripping behavior of a

a Cu 4 Zn0 stripping 3 Pb0 stripping 2 0 -0.4 -0.2 0.0 0.2 0.4 EIV vs RHE -0.2 0.0 0.2 0.4 b 0.3 0.2 -0.11 . -0.4 -0.2 0.0 El -0.4 -0.2 0.0 0.2 0.4 0.6 0.8 VvsRHE 0.2 0.4 0.6 0.8 C 0. 0. 0. 0.

0.

Figure 11.4. Cyclic voltammograms (CV) of copper (a), silver (b) and gold (c) electrodes following 12 min electrolysis (-1.00 V for Cu, -0.90 V for Ag and -0.70 V for Au) in untreated Ci electrolyte with 50 pM ZnSO4 (red), CuSO4 (green) and/or Pb(N03)2 (blue). CVs of each electrode following prolonged (black; Cu: 120 min, Ag &

Ag ZnO stripping Cu0stripping _&tj I . . .. -A

SPb 4d5r2

C2p-32

single M2

+ candidate under experimental conditions relevant to CO₂RR catalysis. For Cu

electrodes polarized at CO2RR potentials in Zn2+-containing electrolytes, the first CV scan

fol-lowing CO2RR catalysis reveals a broad Zn stripping feature at -0.30 V, Figure II.4a (red), in

line with one oxidative wave, at -0.22 V, Figure II.ld (green), observed for CO2RR performed

in native Ci electrolyte. An analogous experiment performed with Pb2+-containing electrolyte reveals oxidative features at 0.15 V and 0.26 V (Figure II.4a, blue), attributed to stripping of bulk and underpotential deposited (UPD) lead on Cu, as previously characterized in acidic electrolytes. Consistent with the low concentrations of Pb in native Ci electrolyte (see above), the stripping feature observed in Figure II.1d (green) matches the UPD stripping potential of 0.26 V in Figure II.4a (blue). Together, these results indicate that Cu is susceptible to progressive

fouling by Zn and Pb deposition during CO2RR catalysis.

Similar experiments conducted with Ag and Au working electrodes in Zn2+-containing electrolytes reveal Zn stripping features at -0.35 V and -0.23 V for Ag (Figure II.4b, red), and at -0.10 V and 0.12 V for Au (Figure II.4c, red). In both cases, the more positive peak is attributed to Zn UPD stripping features,55-57 which match the stripping waves at -0.25 V on Ag (Figure II.le, green) and 0.11 V on Au (Figure II.1f, green) observed following CO2_{RR in native Ci}

electrolyte. Analogous experiments performed with Ag and Au electrodes in Cu2+-containing

electrolytes reveal broad stripping features at 0.50 V and 0.60 V for Ag (Figure II.4b, green), and 0.65 V and 0.85V for Au (Figure II.4c, green). These features are close to the stripping waves

observed at 0.50 V on Ag (Figure II.le, green), and 0.56 V and 0.85 V on Au (Figure II.lf, green)

following CO2RR in native Ci electrolyte. Slight differences in potentials may be due to alloying

of co-deposited Cu and Zn from native Ci electrolytes,5 8 _{as opposed to the deposition of Cu alone}

such as Fe and Ni, would be expected to occur at values less positive than that of Cu,

_,

_leading

us to assign the aforementioned peaks to Cu stripping. Together with XPS data, analysis of

stripping voltammograms demonstrates that Au and Ag electrodes are susceptible to fouling by

Pb, Zn, and Cu deposition during CO

RR catalysis.

1.3. Impurity Ion Coordination Inhibits Impurity Deposition and Enables Sustained

Electrocatalysis

Impurity metal deposition leads to dramatic changes in CO2RR selectivity on group 11 catalysts. At -0.70 V, CO production on polycrystalline Au foil electrodes operated in native Ci electrolytes commences with -60% Faradaic efficiency but declines to -10% over the course of 2 hours (Figure II.5a, black squares, error bars shown in Figures 11.5, 7-11 are the standard deviation of three independent measurements) consistent with literature reports.34' 36 _{The large error}

bars observed for electrolyses performed in native Ci electrolyte reveal a high degree of run-to-run variability, consistent with trace metal impurities strongly influencing catalytic efficiency. Catalyst deactivation for CO production is accompanied by a rise, from -23% to -77%, in Faradaic efficiency for HER. The observed CO and H2 account for -90% of the steady state current, and no

other gas phase products were detected, suggesting that solution phase products, such as formate,60

may account for the balance. Notwithstanding, CO2RR selectivity is progressively eroded over

time for electrolysis performed from native Ci electrolyte. In order to suppress the observed

deactivation, we employed established preelectrolysis

hands, this method proved ineffective, leading to similar deactivation in CO production Faradaic

a

b

.

J

T

time min time/ min

Figure 11.5. Faradaic efficiencies for CO_{2RR and HER product formation on gold foil in electrolytes of varying purity.} Activity of Au for CO (a) and H2 (b) formation at -0.70 V in native Ci electrolyte (black squares), Ci electrolyte

containing 3.4 pM EDTA (red circles), and Chelex-treated Ci electrolyte (blue triangles). Data are reported for the average and standard deviation of three independent runs for each condition.

Seeking an alternative to preelectrolysis methods, we envisioned that introducing EDTA to Ci

electrolytes would prevent metal deposition, thereby, enhancing long-term catalysis. CV scans of

a Au electrode recorded immediately following

45

minutes of CO

RR catalysis in Ci electrolyte

containing 3.4 pM EDTA (Figure II.lf, red) do not exhibit the Zn and Cu stripping features

observed for the same experiment conducting in the absence of EDTA (Figure II.lf, green). We

also do not observe Zn and Cu features by XPS after

45

min of CO

RR catalysis in

EDTA-containing electrolyte (Figure II.1c, red), indicating an impurity-free surface. A broad oxidative

0.8-0 8 0.6-0.4 2 0.2 -V 2O 0.2-I.C E aE EE.. 0.0 I.I.I. 0 20 40 60 80 100120140 time

Figure 11.6. Faradaic Efficiency for H2 (red circles) and CO (black squares) production on Au foil in

feature is observed beginning at 0.60 V (Figure II.1f, red), which we attribute to the oxidation of

long-lived surface-adsorbed CO formed during electrolysis (see Section

We speculate

that in untreated electrolytes, this weak feature is masked by the copper stripping wave.

In line with the absence of metal fouling observed by CV, we find that ex-situ and in-situ metal

ion chelation significantly attenuates catalyst deactivation. Au electrodes display sustained CO2RR

activity when operated in Ci electrolyte containing 3.4 gM EDTA (Figure II.5a, red circles). The

concentration of EDTA was chosen based on the M2- impurity concentration measured by

ICP-MS of native Ci electrolyte (see above). The initial Faradaic efficiency for CO production is -80%

with only a slight decay to -70% after two hours of steady state electrolysis. Further enhancement

in catalytic activity is afforded by purification of Ci electrolyte with solid-supported iminodiacetate

resin (Chelex) prior to electrolysis; CO production on Au is sustained at -80% over two hours

(Figure II.5a, blue triangles). In addition, the use of EDTA-containing or Chelex-treated Ci

elec-trolyte enables greater reproducibility in product selectivity and catalyst activity, as evidenced by

tighter errors bars compared to the data obtained from native Ci electrolyte. Furthermore, the

sustained CO

RR selectivities for Au are reflected in sustained partial current densities for CO

a

1.4

b

0.7

1.

I

₀₆

_-0.70

V

C1.20

j

i

0.65-E

o0.8

00.4-E

0.6

0.3

-0.70 V

0.21

01-

20

40

60

80

100 120

20

40

60

80 100 120

time/

min

time/

min

Figure 11.7. Partial current densities for CO2RR and HER product formation on gold foils in electrolytes of varying

purity. Activity of Au for CO (a) and H2 (b) formation at -0.70 V in native Ci electrolyte (black squares), Ci electrolyte

(Figure II.7a) and H

production (Figure II.7b), indicating that impurity chelation provides for

sustained intrinsic rates of product formation. Upon treatment of the electrolyte with either EDTA

or Chelex, the current densities for CO production remains high (1.2 mA cm-

_{) relative to HER}

(0.1 mA cm-

_{). Taken together with the CV and XPS data, these results suggest that impurity metal}

deposition over the course of CO

RR catalysis is the principal source of electrode deactivation on

Au surfaces.

Silver electrodes display similar deactivation profiles when operated in native Ci electrolyte.

At -0.90 V, CO production on Ag in native CI electrolyte commences with -50% Faradaic

efficiency but declines over the course of two hours to -33% (Figure II.8a, black squares). This

deactivation is accompanied by a corresponding rise from -38 to -52% in current going to the

HER over the same period. Similar to the results observed on Au, following 45 minutes of CO

RR

catalysis in Ci electrolyte containing 3.4 pM EDTA, no stripping waves are observed by CV

(Figure II.le, red) and XPS spectra show no Zn or Cu peaks (Figure II.1b, red), indicating a

a

b

0.7

r0.6 -

_0.6

0.5-

tE.5

c-2

0.4

L 0.4

-U-

Ag

0.3

0 0.3 - -0.90 V

-20

40

60

80

100 120

20

40

60

80

100 120

time / min

_time

_{/ min}

Figure 11.8. Faradaic efficiencies for CO2RR and HER product formation on silver foils in electrolytes of varying

purity. Activity of Ag for CO (a) and H2 (b) formation at -0.90 V in native Ci electrolyte (black squares), Ci electrolyte

containing 3.4 pM EDTA (red circles), and Chelex-treated Ci electrolyte (blue triangles).

Ag

-0.90

V

I

I

.

I

.

I

.

I

metal impurity free surface. The trailing cathodic feature ending at -0.15 V is attributed to residual

catalytic current, and this feature is not observed in subsequent scans.

Consistent with the CV data, EDTA enables sustained and improved CO production,

commencing at

-58%

and declining only slightly to 52% after two hours of electrolysis (Figure

II.8a, red circles). As for Au, further enhancements in long-term catalytic activity are observed for

Ci electrolytes purified by treatment with Chelex prior to electrolysis: CO production on Ag is

sustained at -60% over the entire two-hour period (Figure II.8a, blue triangles). Consistent with

the retention in FE, the total currents for product formation are preserved at 0.65 mA cm-

for CO

(Figure II.9a) and at 0.4 mA cm

for H

formation (Figure II.9b). The error bars for the Ag data

are large possibly due to variance in the sulfuric acid etching treatment prior to each run. As for

Au electrodes, these results suggest that impurity metal deposition during CO

RR catalysts is the

principal source of deactivation of Ag electrodes.

a

0.9

0.8

0.6

o 0.5

0.4

n 3

b o.9

0.8

E0.7

0.6

E

0.5

0.4

0.3

-0.90

V

-20

40

60

80

100 120

20

40

60

80

100 120

time

/ min

time

/ min

Figure 11.9. Partial current densities for CO2RR and HER product formation on silver foils in electrolytes of varying

purity. Activity of Ag for CO (a) and H2 (b) formation at -0.90 V in native Ci electrolyte (black squares), Ci

electrolyte containing 3.4

sM

-1

M

Ag

-0.90 V

I ' ' I ' ' * ' ' '

Copper electrodes display complex deactivation profiles when operated at -1.00 V in native

Ci electrolyte. At early times, the principal gas phase CO2RR product is C

H

with -12% Faradaic

efficiency, Figure II.10c. However, after two hours of electrolysis this C2 product is not observed

at all. Similarly, CH

production commences at -7% Faradaic efficiency, but is no longer observed

after two hours of electrolysis, Figure II.10a. The decline in selectivity for higher order CO2RR

products is accompanied by a rise in

production yield from

-65%

to

-85%,

Figure II.10d.

These results are consistent with previous reports of Cu at similar potentials." Over the same time

0.30

0.25

0.20

0.15

0.10

0.05

0.00

0.2C

0.15

0.1

0.05

0.00

-1

1

iii

-

~

20

40

60

80

100 120

time/

min

b

d

.2

Cu

'a

T

0.06

0.04

0.02

0.00

1.0

0.9

0.8

0.7

0.6

0.5

Cu

20

40

60

80

100 120

time

min

Cu

-L

-20

40

60

80

100 120

20

40

60

80

100 120

time

/ min

_{time / min}

Figure II.10. Faradaic efficiencies for CO2RR and HER product formation on copper foils in electrolytes of varying

purity. Activity of Cu for CH4 (a), CO (b), C2H4 (c) and H2 (d) formation at -1.00 V in native Ci electrolyte (black

squares), Ci electrolyte containing 3.4pM EDTA (red circles), and Chelex-treated Ci electrolyte (blue triangles).

a

Cu

period, the minority production of CO,

-5%,

remains constant within error over the two hour

electrolysis period, Figure II.10b. We only examine gas phase products in this study, and note

that copper also mediates the production of liquid products including ethanol, propanol and

formate,

under these conditions.

The data obtained in the case of Cu differ from those collected when using Ag and Au; we do

not observe enhanced long-term CO

RR activity on Cu electrodes operated in EDTA-containing

Ci electrolyte. The deactivation profiles for C

H

and CH

production as well as the rise in H

production are similar to that observed for native Ci electrolyte (Figures II.10a, 10c, and 10d, red

circles). Surprisingly, the introduction of EDTA to the electrolyte promotes a decline in CO

production Faradaic efficiency from

-5%

to -0% over the course of two hours, Figure II.10b.

While these observations seem to imply that EDTA is ineffective at preventing metal deposition

on Cu electrode surfaces, voltammetry scans recorded immediately following 2 hours of

electrolysis in the presence of EDTA do not display any stripping waves (Figure II.ld, red), and

XPS spectra recorded following electrolysis show no Zn or Pb peaks (Figure II.1a, red),

suggesting a metal impurity-free surface. The trailing cathodic feature ending at 0.00 V is

attributed to residual catalytic current, and is not observed in subsequent scans. Based on these

results, we propose that the observed inefficacy of EDTA in preventing CO

RR deactivation on

Cu may result from direct interaction of the chelator with the surface and/or chelator-induced

surface restructuring. Interestingly, the rise in CO production observed for Cu operated in native

Ci electrolyte, which is not seen in EDTA-containing Ci electrolyte, suggests that Pb or Zn metal

deposition, observed by XPS (Figure II.1a) and CV (Figure II.ld), may promote release of CO

intermediates from Cu surfaces.

The observed deactivation of Cu electrodes can, however, be eliminated by pre-treatment of

Ci electrolyte with Chelex. The principal CO

RR product is CH

with -20% Faradaic efficiency

sustained over two hours (Figure II.10a, blue triangles). Similarly, C

H4 and CO production is

sustained at

-5%

_{(Figure II.10c, blue triangles) and -2% (Figure II.10b, blue triangles),}

respectively, over the same time period. It appears that removal of metal ion impurities from the

electrolyte alters product selectivity even at early reaction times, as shown by the substantial

enhancement in methane yield and corresponding decrease in C

H

Faradaic efficiency.

Cu

-

₄

-

4

20

40

60

80

100 120

time/

min

20

40

60

80

time / min

b

_0.08

1

0.06

E

E 0.04

0

'0.02

0.00

d

E

E

100 120

1.8

1.6

1.4

1.2

1.0

0.8

0.6

0.4

Cu

20

40

60

80

100 120

time/

min

- CU

-j

I-

l

i

i

111 11

$-20

40

60

time

80

Smin

100 120

Figure 11.11. Partial current densities for CO2RR and HER product formation on copper foil in electrolytes of varying purity. Activity of Cu for CH4 (a), CO (b), C2H4 (c) and H2 (d) formation at -1.00 V in native Ci electrolyte (black

squares), Ci electrolyte containing 3.4pM EDTA (red circles), and Chelex-treated Ci electrolyte (blue triangles).

a

0.6

0.0

E

E

0.30

0.25

0.20

E 0.15

0.10

0.05

0.00

Cu