Tuning Metal Structures and Surfaces for Selected Electrochemical Reactions

by Shraboni Ghoshal

B.S. in Chemistry, Lady Brabourne College, University of Calcutta M.S. in Chemistry, Indian Institute of Technology, Kanpur

M.S. in Chemistry, Rutgers, The State University of New Jersey

A dissertation submitted to

The Faculty of the College of Science of Northeastern University

in partial fulfillment of the requirements for the degree of Doctor of Philosophy

April 12, 2017

Dissertation directed by

Sanjeev Mukerjee Distinguished Professor of Chemistry and Chemical Biology

ii

Dedication

Dedicated to my wonderful family, who have always stood by me and have never stopped me

from following my dreams.

iii

Acknowledgements

I would like to begin by thanking my advisor Prof. Sanjeev Mukerjee, without whose

encouragement, support and guidance none of this research would have been possible. I feel

honored and privileged to be a part of his team at Northeastern University Center for Renewable

Energy Technology (NUCRET). The world class facility that NUCRET provided for research is

praiseworthy and I am fortunate to have experienced the professional environment at NUCRET.

Also, Prof. Mukerjee has shown extreme patience and kindness which has helped me to

overcome the challenges of my graduate life and has prepared me for facing more challenges in

the world of scientific research.

Next, I would like to thank my Thesis Committee members Prof. Mary Ondrechen, Prof.

Geoffrey Davies and Prof. David Budil for their valuable support. It was Prof. Ondrechen who

had suggested to arrange annual committee meetings which proved to be extremely helpful

toward preparing myself for my thesis completion. Additionally, valuable suggestions from Prof.

K.M Abraham have enriched my research to a significant extent. The outstanding administrative

staff of the Department of Chemistry and Chemical Biology, especially Cara Shockley and

Vernon Bean are immensely appreciated for their help and friendly behavior over the years. My

heartfelt thanks to the Northeastern University College of Science for accepting me into the

graduate program and giving me the opportunity to conduct my PhD research.

The work environment at NUCRET has been truly inspiring. Our executive director Dr.

Serge Pann has been very supportive and helpful, and life has been much easier at NUCRET

since he joined. My special thanks go to Dr. Qingying Jia who has taught me the crux of XAS

with amazing patience and demeanor. I consider myself lucky to have such a knowledgeable and

nice person in my group. I express my gratitude to him for his analysis of the XAS data and

iv

guiding me in writing the research papers. Special thanks to my batch mate and good friend

Jingkun Li with whom I have worked on several projects. Entering graduate school at the same

time, surviving candidacy exams and facing challenges of graduate research- we have travelled

together and her companionship has made this journey easier. The atmosphere of NUCRET has

been amazing and this credit goes to the awesome graduate students Ryan Pavlicek,

AmellAlsudiari and Huong Doan who have been very friendly and helpful. Additionally, the

presence of enthusiastic group of undergraduate students has been charming. I would also like to

thank Dr. Nagappan Ramaswamy and Dr. Michael Bates with whom I had the opportunity to

work and they have taught me many interesting aspects of electrochemistry. I would especially

like to mention Bob Allen whose profound knowledge on catalyst synthesis has taught me how

to think smartly when designing new materials. Also, special thanks to Wentao Liang and Bill

Fowler from Department of Biology for their help with SEM and TEM.

This dissertation work would not have been possible without numerous industry

collaborators and their financial supports. I would like to thank ARPA-E and SAFCell Inc. for

providing funds to carry out my research. Calum R.I Chisholm of SAFCell Inc. has been a great

collaborator and I would like to thank him for all his support throughout the project. In addition

to this, DeNora Tech is highly appreciated for their collaboration with the Chlor- Alkali project.

Chris Allen from DeNora Tech (also a former graduate student of NUCRET) has been of great

help in cell testing.

I have been blessed with some amazing friends who had crucial role in supporting me

when times were difficult. My best friends SurmaTalapatra and Adita Das who belong to

academia as well have helped me go through graduate life in a foreign country far away from

home. Very close circle of friends (Somenath, Ayan, Samragnee, Seemin and Samya) around

v

Boston were strong support throughout my stay at Boston. My career path would have been very

different have I hadn’t met my high school Chemistry teacher Mr. Arun Nag, under whose

tutelage I learnt to appreciate the beauty and mysteries of Chemistry. My best friend and my

husband, Aritra has been a blessing in my life who has endured me during both my best and

worst times with equal affection, and never stopped encouraging me. My parent in-laws have

been extremely loving, caring and supportive which gave me constant motivation to move

forward. Lastly, my dear parents, my sister and Onki - I can never thank you enough for being

always there for me. Without your invaluable support and blessings, I would not have made it

this far.

vi

Abstract of Dissertation

To mitigate serious issues like climate change and depletion of fossil fuels, adapting to

renewable energy technology as primary energy source is an expedient solution. Among all the

available renewable energy technologies, electrochemical energy devices are the most viable

option since they have no geographical restriction and have potential to meet the global energy

demand. To establish modular and scalable electrochemical energy devices, it is important to

have a fundamental understanding of electrochemical interfaces and the factors that influence the

cost, durability and fuel flexibility in these devices. In this dissertation, various metal surfaces

and structures have been investigated to study selected electrochemical reactions that are relevant

to electrochemical energy devices such as fuel cells and electrolyzers. This work mainly focusses

on two vital electrochemical reactions: 1) Oxygen Reduction Reaction (ORR), which is pertinent

to fuel cells and alkaline electrolyzers for the application of oxygen depolarized cathode (ODC)

and 2) Hydrogen Oxidation Reaction (HOR) which is important for anion exchange membrane

fuel cells (AEMFCs) where overpotential losses from hydrogen side is the major cause of overall

performance loss in alkaline systems. To minimize the requirement of platinum in these energy

devices, several categories of catalysts such as non- platinum catalysts, non- precious catalysts

and catalysts with ultra- low platinum loadings have been explored. Development of such low

cost, high performance catalysts are especially required to achieve a sustainable global green

economy.

In Chapter 1, a general overview of effects of fossil fuels and global warming is

presented, which is then followed by discussions on renewable energy production methods and

viable options. Additionally, brief explanation of different types of fuel cells and electrolyzer

plants (such as chlorine generation plant) has been provided. In Chapter 2, development of a

vii

non- platinum catalyst based on Ru chalcogenides are reported for application in solid acid fuel

cells (SAFC). SAFCs are a recent discovery and have several advantages such as simple built

and cheap installation, however they have a major drawback due to their dependence on

unsupported platinum at cathode terminal. Hence, the exorbitant cost associated with the catalyst

is a major hindrance in commercialization of these fuel cells. Ru being a considerably cheaper

metal than Pt, helps to reduce to cost significantly, albeit Ru undergoes gradual oxidation under

operating condition at the cathode side. Strategic modification of Ru surface while preserving the

catalytically active Ru core has been carried out using Se and Mo as dopants that protect Ru

from getting oxidized and at the same time enhance the overall catalytic performance. A

thorough investigation of the Ru based chalcogenide catalyst has been carried out using

electrochemical and in situ spectroscopic experiments that provide mechanistic detail on the

activity of the Ru based chalcogenide catalyst.

In Chapter 3, a platinum- niobium based alloy catalyst PtNb/NbOx-C has been reported.

This is the first reported work on electrocatalyst where niobium is shown to successfully form an

alloy with Pt. Niobium has strong oxophilicity and high negative reduction potential which

prevents it from existing as zero valent state in bimetallic composites. In this work, a strategic

synthetic route causes formation of platinum- niobium composite with coexistence of Nb

oxide(s) and metallic niobium. This unique material has been analyzed extensively with structure

sensitive experiments such as X-ray diffraction, X-ray photoelectron spectroscopy and XAS.

Following this, the catalytic profile of the PtNb/NbOx-C catalyst has been assessed in various

electrochemical reactions, such as ORR and HOR.

In Chapter 4, HOR activities of different Pt based bimetallic systems have been studied in

alkaline medium. Carbon supported PtRu catalysts, denoted by PtRu/C is the state of the art

viii

catalyst for alkaline HOR. PtRu/C exhibits faster HOR kinetics with more than two times higher

exchange current density compared to Pt/C. The mechanistic origin of such superior performance

is unknown and two contrasting school of theories exist, namely binding energy theory and

reactive OHads theory. As per binding energy theory, Ru modifies the electronic structure of Pt

and lowers the binding energy of hydrogen to Pt surface, thereby facilitating the HOR. Per

reactive OHads theory, Ru moieties present in the PtRu/C catalyst provide reactive OHads moieties

(due to the oxophilicity of Ru) that helps in the proton abstraction from the Pt surface, and makes

HOR much more facile. Different types of PtRu/C catalysts prepared by electrodeposition,

solvothermal and chemical vapor deposition have been tested and analyzed. The alkaline HOR

profile has been studied in all these PtRu/C catalysts and the mechanistic origin of the higher

activity of PtRu/C toward alkaline HOR has been investigated based on the different types of

interactions between Pt and Ru in these catalysts. As an extension of this work, carbon supported

PtMo has been studied and its catalytic activity toward alkaline HOR has been analyzed.

In Chapter 5, application of non-precious catalysts composed of Fe-N-C have been

demonstrated as ODC in chlorine generating electrolyzer cells. Mainly two types of catalysts-

FeMOF and Fe-PEI have been tested as ODC in the electrolyzer cells. The active site structures

of these two catalysts have been elucidated using X-ray Absorption Spectroscopy (XAS). Two

different types of electrolyzer cells have been discussed- a half cell which was tested at

NUCRET and consisted of oxygen evolution reaction (OER) and ORR as the anode and cathode

processes. This study was done to determine the performance of the Fe-N-C materials as oxygen

reducing electrocatalyst while acting as ODC. Once tested at NUCRET, these catalysts were

tested at DeNora Tech in a full cell where the chlorine was generated at the anode side using a

dimensionally stable anode (DSA); while the Fe-N-C catalysts served as the ODC. Combining

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preliminary electrochemical measurements, half-cell and full- cell performances, the Fe-N-C

catalysts were critically analyzed and compared to the precious metal state of the art catalysts-

Pt/C and unsupported Ag (developed by DeNora Tech).

A summary and conclusion of the dissertation work has been provided in Chapter 6 along

with future developments of bimetallic catalysts for HOR in alkaline electrolyte.

x

Table of Contents

Dedication…………………………………………………………………………………………ii

Acknowledgements ........................................................................................................................ iii

Abstract of Dissertation ................................................................................................................. vi

Table of Contents ............................................................................................................................ x

List of Figures ............................................................................................................................... 15

List of Tables ................................................................................................................................ 19

List of Abbreviations and Symbols............................................................................................... 20

Chapter 1: Introduction ................................................................................................................. 26

1.1Energy Beyond Fossil Fuels .................................................................................................... 26

1.2Electrochemical Energy Technology ....................................................................................... 28

1.3Electrochemical Energy Conversion Devices .......................................................................... 29

1.3.1 Fuel Cells ......................................................................................................................... 29

1.3.2 Electrolyzer Cells for Generation of Chlorine ................................................................. 35

1.4Electrochemical Reactions of Interest ..................................................................................... 35

1.4.1 Electrocatalysis of Oxygen Reduction Reaction .............................................................. 35

1.4.2 Electrocatalysis of Hydrogen oxidation Reaction............................................................ 36

1.5RDE and R(R)DE Technique for kinetic Study of Electrocatalysts ........................................ 38

xi

1.6X-ray Absorption Spectroscopy ............................................................................................... 43

1.7Scope of Work………………………………………………………………………………..46

1.8 References……………………………………………………………………………………46

Chapter 2: Electrochemical and In Situ Spectroscopic Evidences toward Empowering Ruthenium Based Chalcogenides as Solid Acid Fuel Cell Cathodes .............................................................. 49 2.1 Introduction ............................................................................................................................. 50

2.2 Experimental Section ............................................................................................................. 53

2.2.1 Synthesis of Electrocatalysts ............................................................................................ 53

2.2.2 Physicochemical Characterization ................................................................................... 53

2.2.3 Electrochemical Characterization .................................................................................... 54

2.2.4 X-Ray Absorption Spectroscopy Measurements ............................................................. 54

2.2.5 MEA Testing in Solid Acid Fuel Cell.............................................................................. 56

2.3 Results and Discussions .......................................................................................................... 57

2.3.1 Physicochemical Characterizations ................................................................................. 57

2.3.2 ORR on Ru based Chalcogenide Catalysts ...................................................................... 58

2.3.3 In Situ XAS Experiment .................................................................................................. 66

2.3.4 Solid Acid Fuel Cell Testing............................................................................................ 73

2.4 Conclusion .............................................................................................................................. 77

2.5 Acknowledgements ................................................................................................................. 78

2.6 References ............................................................................................................................... 78

Chapter 3: Mulfunctional Pt-Nb Electrocatalysts Facilitated by Multiphase Niobium ................ 81

3.1 Introduction ............................................................................................................................. 82

xii

3.2 Experimental Section .............................................................................................................. 85

3.2.1 Synthesis of PtNb/NbOx-C Catalyst ................................................................................ 85

3.2.2 PhysicochemicalCharacterizations .................................................................................. 86

3.2.3 Electrochemical Characterizations................................................................................... 86

3.2.4 Durability Measurement .................................................................................................. 87

3.2.5 X-Ray Absorption Spectroscopy Measurements………………………………………..87

3.3 Results and Discussions .......................................................................................................... 89

3.3.1 Catalyst Synthesis ............................................................................................................ 89

3.3.2 Physicochemical Characterization ................................................................................... 89

3.3.3 X-Ray Absorption Spectroscopy ..................................................................................... 93

3.3.4 Oxygen Reduction Reaction in Acid on PtNb/NbOx-C Catalyst……………………….98

3.3.5 CO Stripping Experiment on PtNb/NbOx-C Catalyst………………………………….101

3.3.6 Alkaline Hydrogen Oxidation on PtNb/NbOx-C Catalyst……………………………..104

3.4 Conclusion ............................................................................................................................ 107

3.5 Acknowledgements ............................................................................................................... 108

3.6 References ............................................................................................................................. 108

Chapter 4: Investigation into the Alkaline Hydrogen Oxidation Reaction Mechanism for Bimetallic Platinum Composites................................................................................................. 111 4.1 Introduction ........................................................................................................................... 112

4.2 Experimental Section ............................................................................................................ 116

4.2.1 Synthesis of Electrocatalysts .......................................................................................... 117

4.2.2 Structural Characterizations ........................................................................................... 118

xiii

4.2.3 Electrochemical Characterizations................................................................................. 118

4.3 Results and Discussions ........................................................................................................ 119

4.3.1 Pt/C-Ru Catalysts Obtained by Electrodeposition ......................................................... 119

4.3.2 Ru/C@ Pt Core Shell Electrocatalysts ........................................................................... 124

4.3.3 PtMo(Ox)/C Catalyst ...................................................................................................... 127

4.4 Conclusions ........................................................................................................................... 131

4.5 Acknowledgements ............................................................................................................... 133

4.6 References ............................................................................................................................. 133

Chapter 5: Engendering Non-Precious Catalysts for Oxygen Depolarized Cathodes in Chlor Alkali Electrolyzer Cells ............................................................................................................. 135 5.1 Introduction ........................................................................................................................... 136

5.2 Experimental Section ............................................................................................................ 138

5.2.1 Synthesis of Catalysts .................................................................................................... 138

5.2.2 X-Ray Absorption Spectroscopy ................................................................................... 140

5.2.3 Fabrication of Gas Diffusion Electrode ......................................................................... 140

5.2.4 Rotating Disk Electrode Experiment ............................................................................. 141

5.2.5 Electrolyzer Cell Testing…………………………………………………………………141

5.3 Results and Discussions ........................................................................................................ 142

5.3.1 Structure of Iron Active Centers .................................................................................... 142

5.3.2 Oxygen Reduction Reaction at high pH Electrolyte ...................................................... 143

5.3.3 Electrolyzer cell Data- Half Cell & Full Cell ................................................................ 146

xiv

5.4 Conclusion ............................................................................................................................ 148

5.5 Acknowledgements ............................................................................................................... 149

5.6 References ............................................................................................................................. 149

Chapter 6: Summary and Future Directions ............................................................................... 152

15

List of Figures

Figure 1.1 Primary contributors to Climate Change 26

Figure 1.2 Schematic representation of PEMFC 31

Figure 1.3 Structure of Nafion and representation of 32

Grotthuss mechanism for proton transport

Figure 1.4 Proton transport in CDP 34

Figure 1.5 Cyclic Voltammogram of Pt/C in 0.1M KOH 40

Figure 1.6 Schematic representation of RRDE 42

Figure 1.7 ORR polarization curve for Pt/C in 0.1M HClO4 43

Figure 1.8 Schematic representation of XAS experiment 44

Figure 1.9 Normalized XAS spectrum of Pt/C 44

Figure 2.1 XRD pattern of Ru based catalysts 58

Figure 2.2 CV of Ru based chalcogenide catalysts 59

Figure 2.3 CV of Mo modified Ru catalysts 59

Figure 2.4 ORR polarization curves on Ru catalysts 62

Figure 2.5 Ring current response for Ru catalysts 62

Figure 2.6 ORR on Ru catalysts with different Ru loading 63

Figure 2.7 Phosphate anion poisoning studies 64

Figure 2.8 CDP poisoning studies 66

16

Figure 2.9 In situ XANES and FT-EXAFS of Ru/C 67

Figure 2.10 In situ XANES and FT-EXAFS of RuSeMo/C 68

At Ru edge

Figure 2.11 Δµ signatures at Ru K-edge for Ru catalysts 70

Figure 2.12 In situ XANES and FT-EXAFS of RuSeMo/C 71

At Se K-edge

Figure 2.13 FT EXAFS for Se reference foil 72

Figure 2.14 SAFC polarization curves of Ru catalysts 75

Figure 2.15 Durability test of RuSeMo/C in SAFC 76

Figure 2.16 TGA profiles of Ru catalysts 76

Figure 3.1 Physicochemical characterization of 91

PtNb/NbOx-C catalyst

Figure 3.2 XPS high resolution spectra of 92

PtNb/NbOx-C catalyst

Figure 3.3 XPS high resolution spectra of NbOx supported 93

Pt/C catalysts

Figure 3.4 In situ XANES of PtNb/NbOx-C at Nb K-edge 93

Figure 3.5 In situ XANES and FT-EXAFS of PtNb/NbOx-C 94

catalyst at Nb K-edge

Figure 3.6 In situ XANES and FT-EXAFS of PtNb/NbOx-C 95

catalyst at Pt L3-edge

Figure 3.7 In situ XANES and Δµ of Pt/C and 96

PtNb/NbOx-C catalysts

Figure 3.8 ORR polarization curves and durability measurements 99

For PtNb/NbOx-C catalyst

17

Figure 3.9 ORR polarization curves of Pt/C in presence of 100

H3PO4

Figure 3.10 CO stripping data for Pt/C and PtNb/NbOx-C 102

Figure 3.11 CV and HOR polarization curves of 104

PtNb/NbOx-C in 0.1M KOH

Figure 3.12 Specific activities of Pt catalysts 105

Figure 3.13 HOR polarization curves of Pt/C supported on 106

NbOx

Figure 4.1 CuUPD stripping experiment on Pt/C 120

Figure 4.2 CV and HOR activities of electrodeposited 121

Pt/C-Ru catalysts in 0.1M KOH

Figure 4.3 Plot of kinetic current density for Pt/C-Ru 121

Figure 4.4 CO stripping data of Pt and Pt/C-Ru catalysts 123

Figure 4.5 Physicochemical characterization of Ru/C@ Pt 124

Figure 4.6 HOR of Ru/C@ Pt in 0.1M KOH 125

Figure 4.7 CO stripping profiles of Ru/C@ Pt catalysts 127

Figure 4.8 Physicochemical characterization of PtMo(Ox)/C 129

Figure 4.9 CV and HOR of PtMo(Ox)/C in 0.1M KOH 130

Figure 4.10 CO stripping data of Pt/C and PtMo(Ox)/C 131

Figure 5.1 FT-EXAFS of FeMOF and Fe-PEI catalysts 142

Figure 5.2 ORR polarization curve of Pt/C, FeMOF and Fe-PEI 144

18

Figure 5.3 Steady state polarization curves of Pt/C, FeMOF and 146

Fe- PEI catalysts in half cell and full cell brine electrolyzer

Scheme 3.1 Synthesis of PtNb/NbOx-C catalyst 90

Scheme 4.1 Illustration of alkaline HOR mechanism 113

Scheme 4.2 Electrodeposition process to obtain Pt/C-Ru catalysts 115

Scheme 4.3 Electrode interface of PtMo(Ox)/C catalyst 116

Scheme 4.4 MOCVD process to obtain Ru/C@ Pt catalysts 117

Scheme 5.1 Synthesis of FeMOF and Fe- PEI catalysts 139

Scheme 5.2 Half- cell chlor alkali electrolyzer 142

19

List of Tables

Table 1.1 Oxygen Reduction in Acidic and Alkaline Electrolytes 37

Table 2.1 Physicochemical characterization of Ru catalysts 57

Table 2.2 RDE and RRDE results of Ru catalysts 61

Table 2.3 Semi-quantitative Tafel analysis of Ru and Pt catalysts 65

toward ORR in presence of H3PO4

Table 2.4 In situ EXAFS fit results for Ru catalysts 72

Table 2.5 Comparison of RDE and MEA performance of 75

Ru based catalysts

Table 3.1 EXAFS fit results of Pt/C and PtNb/NbOx-C catalysts 96

Table 3.2 Quantitative assessment of Pt/C and PtNbOx-C catalysts 105

Toward HOR in 0.1M KOH

20

List of Abbreviations and Symbols

ads Adsorbate

AEM Anion Exchange Membrane

AEMFC Alkaline Electrolyte Membrane Fuel Cell

BE Binding Energy

BEV Battery Electric Vehicle

BET Brunauer- Emmett- Teller

BNL Brookhaven National Laboratory

CDP Cesium Dihydrogen Phosphate

CE Counter Electrode

CHP Combined Heat and Power

CV Cyclic Voltammetry

DOE Department of Energy

DSA Dimensionally Stable Anode

ECSA Electrochemical Surface Area

EDS Energy Dispersive X-ray Spectroscopy

eV Electron volt

EXAFS Extended X-ray Absorption Fine Structure

21

fcc Face Centered Cubic

FCV Fuel Cell Vehicle

FT Fourier Transform

FT- EXAFS Fourier- Transformed Extended X-ray Absorption Fine Structure

GCE Glassy Carbon Electrode

GDE Gas Diffusion Electrode

GDL Gas Diffusion Layer

hcp Hexagonal Close Pack

HUPD Under-Potentially Deposited Hydrogen

HER Hydrogen Evolution Reaction

HOR Hydrogen Oxidation Reaction

HT Heat Treated

K-L Koutecky- Levich

MEA Membrane Electrode Assembly

M-N-C Metal- Nitrogen -Carbon

MOCVD Metal Organic Chemical Vapor Deposition

MOF Metal Organic Framework

NSA Near Surface Alloy

22

NSLS National Synchrotron Light Source

NUCRET Northeastern University Center of Renewable Energy Technology

OCV Open Circuit Voltage

OER Oxygen Evolution Reaction

ODC Oxygen Depolarized Cathode

OHP Outer Helmholtz Plane

ORR Oxygen Reduction Reaction

PAFC Phosphoric Acid Fuel Cell

PEI Polyethyleneimine

PEM Proton Exchange Membrane

PEMFC Proton Exchange Membrane Fuel Cell

PGM Precious Group Metal

PTFE Polytetrafluoroethylene

Pt/C Carbon supported Platinum Nanoparticles

PVC Polyvinyl Chloride

PZC Potential of Zero Charge

RDE Rotating Disk Electrode

RE Reference Electrode

23

RDS Rate Determining Step

RHE Reversible Hydrogen Electrode

rpm Rotations Per Minute

RRDE Rotating (Ring) Disk Electrode

RT Room Temperature

SAED Selected Area Electron Diffraction

SAFC Solid Acid Fuel Cell

SEM Scanning Electron Microscope

SLAC Stanford Linear Accelerator Center

SMSI Strong Metal Support Interaction

SSRL Stanford Synchrotron Radiation Light source

TEM Transmission Electron Microscope

TGA Thermogravimetric Analyzer

TKK Tanaka Kikinzoku International KK

TM Transition Metal

UPD Under Potential Deposition

WE Working Electrode

XAS X-ray Absorption Spectroscopy

24

XANES X-ray Absorption Near Edge Spectra

XPS X-ray Photoelectron Spectroscopy

XRD X-ray Diffraction

A Relative Resonance Areas/Geometric Area of Electrode

α Charge Transfer Coefficient

Å Angstroms (10-10 meters)

b Tafel Slope

η Overpotential

ηa Anodic Overpotential

ηc Cathodic Overpotential

H2O2% Peroxide Yield

ω Rotation Rate

υ Kinematic Viscosity of the Electrolyte

μ Absorption Coefficient

μ(E) Absorption Coefficient at energy E

μm Micrometer (10-6 meters)

σ2 Debye-Waller Factor

C Gravimetric Double Layer Capacitance

CN Coordination number

Co* Concentration of Molecular Oxygen

D0 Diffusion Coefficient

25

e- Charge of an Electron (1.602×10-19 C)

E/V Electrode Potential

E0 Standard Potential/Edge Energy/On-Set Potential

E1/2 Half-Wave Potential

F Faraday’s Constant

i/j/jD Current Density

ik/jk/J/JORR(E) Kinetic Current Density

il/jl Limiting Current Density

il,a/jl,a Anodic Limiting Current Density

il,c/jl,c Cathodic Limiting Current Density

i0/j0 Exchange Current Density

I0 Intensities of the Incident

n Number of Electrons Transferred

N Ring Current Collection Efficiency/Number of Neighboring Atoms

psi Pounds per Square Inch

R Bond Length/Universal Gas Constant

T Temperature

f(k) Effective Scattering Amplitude

δ(k) Phase Shift of the Photoelectron

λ(k) Mean Free Path of the Photoelectron

26

Chapter 1

Introduction

1.1 Energy Beyond Fossil Fuels

Fossil fuels have been the primary source of energy serving humans since ancient history.

Gradually, the concerns associated with the usage of fossil fuels were soon identified. Though

fossil fuels have exceptionally high energy density, their reserves are limited. Crude oil resources

are being used at a rate of 4 billion tonnes per year.1 At this rate, all the crude oil resources are

expected to run out by the year 2050.2 Additionally, these non-renewable fuels release harmful

byproducts that pose serious threat to our climate.3 The most common byproduct discharged

from fossil fuels is carbon dioxide (CO2), which is a greenhouse gas. In general, greenhouse

gases like CO2, methane (CH4) and water vapor are essential in order to maintain the temperature

of earth which makes life sustainable on earth.4 However, an imbalance in the concentration

levels of greenhouse gases can severely damage the integrity of the earth’s atmosphere resulting

in undesirable climatic changes. Figure 1.1 illustrates different man-made factors that have

contributed to climate change in the recent years.

Figure 1.1 Primary contributors to climate change.5

An example is the drastic rise in the sea level resulting from the melting of polar ice.

Between 1870 and 2004, there has been a net rise of 7.7 inches in the sea water level, and the

27

rate of increase has accelerated during recent years.6 Such phenomenon can be directly related to

the elevated concentration of CO2 in the atmosphere which has increased from 280 ppm to 400

ppm in 50 years. This is the highest recorded CO2 level in 650000 years, and has expedited the

global warming at a faster rate than reported hitherto.7,8 In view of the anticipated demand of

global energy and elevated levels of pollutants and greenhouse gases in the atmosphere, we are

in dire need of alternative energy technologies that will provide for renewable energy, energy

management, conservation, storage, pollution control and greenhouse gas regulation. The major

sources of renewable sources of energy are- solar, wind, hydropower, geothermal and biofuels.

While hydropower and geothermal sources have geographical restrictions, biofuels emit CO2 as

byproduct thereby defeating the purpose of having an eco-friendly energy source. Wind and solar

energy sources have intermittency issues, questioning their reliability to sustain global

requirements. Besides, in order to fulfill the need of load leveling, interconversion of electrical

energy into chemical energy and vice versa presents a viable option. Examples of such energy

devices are fuel cells, electrolyzers and redox flow batteries. Fuel cells are reliable and clean

energy devices that have slowly formed a niche in the electrical power market. Fuel cells have

several advantages such as fuel flexibility, on site distributed generation improving power

reliability and energy security. The combined heat and power efficiency (a value that generally

varies with the type of fuel cell) further drives the energy economy.9 More detail on fuel cells

will be discussed in the next section. Electrolyzers, on the other hand, are based on the

technology encompassing conversion of electrical energy into chemical energy. Big industries

such as electroplating, extraction of minerals from ores and polymer companies are heavily

dependent on electrolyzer technology where exorbitant power consumption is a serious concern.

Redox flow batteries offer low cost energy storage; however, their low energy density raises

28

concern over their establishment in the commercial market.10 Nonetheless, all these technologies

hold the key to fulfil the global energy requirement and are associated with the advantage of

being applicable in small (portable) to large (grid) systems. To ensure proper establishment of

such devices in the power market, it is necessary to fundamentally understand the factors that

affect the cost, durability and fuel flexibility in such systems.

1.2 Electrochemical Energy Technology

Fuel cells are energy devices that can generate electric power throughoccurrence

electrochemical reactions inside the device. Renowned car companies such as Toyota have

introduced fuel cell vehicles (FCV)s and Tesla have long introduced battery electric vehicles

(BEV)s which are appreciable steps toward a green economy by lowering our dependence on

foreign oil as well as lowering harmful emissions that induce climate change.11 However, there

are numerous challenges that are to be addressed before FCVs and BEVs can entirely fulfill

global demands. Keeping in mind the fact that the following work is based on fuel cell systems

rather than batteries, primary focus will be on the challenges pertaining to fuel cell vehicles. The

foremost concern is that FCVs are significantly more expensive than conventional vehicles and

hybrids. The factors that mainly contribute to the high cost of FCVs are expensive catalysts that

are required to carry out electrochemical reactions to generate electrical power. This issue can be

mitigated by designing non-precious catalysts and incorporating fuel cells that operate at higher

temperature. Secondly, the availability of hydrogen fueling stations is not abundant and hence

puts restriction in the commercialization of FCVs to full extent. Part of ongoing research

focusses on enabling a more viable hydrogen economy and reducing the cost associated with

using hydrogen as a fuel for commercial vehicles. The work discussed in this dissertation will

focus on minimizing the cost associated with the fuel cell stacks inside the vehicle- which

29

comprises incorporation of less expensive catalysts for electrochemical energy conversion and

adoption of different types of fuel cell systems to reduce overall cost of the fuel cell stacks inside

the vehicle.

The other electrochemical technology includes utilization of electrical energy to carry out

chemical reactions and produce materials that are used at large scales. One such example is

chlorine generation plants that produce chlorine using electrolyzer cells.12 Chlorine is one of the

primary constituent of the polymer industry and is consumed in large scale globally. Chlorine can

be generated by HCl electrolysis, where HCl is obtained mostly from polymer industries where

HCl is a common byproduct. Apart from this, a more abundant source of chlorine is found in sea

water in the form of dissolved NaCl. Since sea water is abundant on earth’s surface, this makes

NaCl electrolysis (or chlor alkali process) a more lucrative option as the process in not limited by

the availability of starting material. However, chlor alkali industries are one of major energy

intensive plants which consume high amounts of electrical power to generate chlorine from sea

water. In the US, 2% of the total electric power, which amounts to 87600 GWh/year is consumed

by chlor alkali industry.13 In order to lower the amount of electric consumption by these plants,

fundamental understanding and designing of catalysts and electrodes are required which are

discussed in detail in chapter 5.

1.3 Electrochemical Energy Conversion Devices

1.3.1 Fuel Cells

As a simple definition, fuel cells are devices that can convert chemical energy into

electrical energy. Fuel cells consist of two electrodes- a positive terminal and a negative terminal

where electrochemical reactions occur. These two electrodes are internally connected by an

30

electrolyte (known as membrane) which essentially transports ions across the two terminals. The

electric power generated as a result of the chemical reaction can be utilized by connecting the

two electrodes externally to a power output. The electric current returns to the cell after

completing the electrical circuit. The basic operation of a fuel cell involves hydrogen which acts

as the fuel, and an oxidizer gas which is either oxygen or air (in case of commercial

applications). Hydrogen, which is fed to the anode terminal, undergoes electrochemical oxidation

while oxygen gets electrochemically reduced at the cathode terminal. Therefore, the two

reactions occurring in a fuel cell are hydrogen oxidation reaction (HOR) and oxygen reduction

reaction (ORR). Depending on electrolyte systems, different types of ionic membranes are used.

For example, if protons (generated at the anode) are transported from anode to cathode, then

polymer electrolyte membranessuch as Nafion® are used. On the other hand, if hydroxide ions

(generated at the cathode) are transported from cathode to anode, anion exchange membranes are

used. In this discussion, we will focus on three types of fuel cells relevant to the work namely

proton exchange membrane fuel cells (PEMFCs), anion exchangemembrane fuel cells

(AEMFCs) and solid acid fuel cells (SAFCs). In PEMFCs and SAFCs, the following

electrochemical equations express the reactions occurring at the two terminals:

Anode: H2 ↔ 2H+ + 2e- E°=0.00 V vs SHE (1.1)

Cathode: ½O2 + 2H+ + 2e- ↔ H2O E°= 1.23 V vs SHE (1.2)

31

Figure 1.2 Schematic representation of a PEMFC

Combination of the two electrochemical half reactions gives rise to the overall cell reaction (ECell

= ECathode – EAnode) in which the combustion of hydrogen and oxygen produce water, heat and

electricity.

ECell = ECathode – EAnode (1.3)

Overall: H2 + ½O2 ↔ H2O ECell= 1.23 V (1.4)

In alkaline systems such as AEMFCs, the equations are modified as per:

Anode: H2+2OH- ↔ 2H2O + 2e- E°= -0.828V vs SHE (1.5)

Cathode: ½O2 + +H2O + 2e- ↔ 2OH- E°= 0.401 V vs SHE (1.6)

Using equation 1.3, Ecellfor AEMFC comes out to be 1.23V, same as for PEMFCs and SAFCs

(a)Proton exchange membrane fuel cells: Proton exchange membrane fuel cells or PEMFCs have

a water based, acidic membrane made of polymer as electrolyte. Hence these fuel cells are also

known as polymer exchange membrane fuel cells. Due to the temperature- sensitive nature of the

32

membranes, these fuel cells are typically operated at temperatures lower than 100oC, and hence

belong to the category of low temperature fuel cells. The commercially used membrane in these

fuel cells are known as Nafion®. These membranes are made of perfluorinated ammonium based

polymer that aid in transporting protons as per Grotthuss hopping mechanism.14,15 The mobility

of protons in the in Nafion® is known to be highly dependent upon the water concentration in the

membranes. Hence in PEMFCs, the gases which are fed at the two electrode terminals must be

humidified before it reaches the electrodes. PEMFCs are currently incorporated in light vehicles

and materials handling vehicles. While PEMFCs are potential answers to the green economy

goal, they are also blended with several challenges such as dependence on precious metal

catalysts, lifetime of the membranes and fueling issues.

Figure 1.3 Chemical structure of Nafion® and (inset) schematic representation of Grotthuss hopping mechanism for proton transport

(b) Anion exchange membrane fuel cells:Anion exchange membrane fuel cells or AEMFCs are

viable alternatives to PEMFCs largely due to several factors. (1) Oxygen reduction reaction

occurring at the cathode terminal is much more facile in alkaline electrolyte. This engenders the

utilization of non-precious catalysts instead of precious catalysts, thereby significantly reducing

the cost associated with the fuel cell stack.16 (2) the electro- oxidation kinetics of many non-

33

conventional liquid fuels are enhanced in alkaline electrolyte.17 (3) the electro-osmotic drag

associated with transport of ions across the terminals restricts the crossover of liquid fuels,

therefore allowing more concentrated liquid fuels in AEMFCs.18 These list of advantages make

AEMFCs profitable choice in the electrical power market.

In AEMFCs, hydroxide anions are generated at the cathode during electrochemical

oxygen reduction. The hydroxide anions are then transported by the electronically insulating-

anion conducting membrane (commercially available a AS4 membrane) to the anode terminal,

where hydrogen undergoes oxidation to form water.19 Hence unlike PEMFCs, the byproduct

water is generated at anode instead of the cathode side in case of PEMFCs. A typical AEMFC

operates in the temperature range of 50oC—60oC. While in the case of PEMFCs, the major

performance loss arises from the sluggish kinetics of ORR, in AEMFCs the anodic reaction

involving alkaline HOR is the major cause of performance loss. This results in an increased

loading of precious catalyst required for the catalysis of HOR occurring at the anode side.

Ongoing efforts include understanding the origin of the slow kinetics of alkaline HOR as well as

designing catalysts with minimal precious metal loadings. More detail on HOR and ORR will be

discussed in the next section.

(c)Solid acid fuel cells: Solid acid fuel cells, or SAFCs are a new type of fuel cell that were

introduced in early 2000.20 The electrolyte used in these fuel cells consist of ceramic proton

conductors that transport protons at elevated temperature. Unlike Nafion®, proton transport in

solid acids does not follow the Grotthuss mechanism and hence no water is required to channel

the protons from one terminal to other.21In this fuel cell, electrolytes are made of solid acid ,

which comprise a large metal cation and a tetrahedral anion. These special type of proton

conductors obtain their ability to transport protons at high temperatures (known as superprotonic

34

temperature) when the solid acid molecules undergo random reorientation, facilitating proton

transfer through the vacant sites.20,22 This phenomenon is also referred to as superprotonic

transformation.22 The most common electrolyte used in solid acid fuel cells is cesium dihydrogen

phosphate (CDP) that undergoes superprotonic transformation at 230oC.23 The SAFCs typically

operate at 250oC and thereby belong to the group of intermediate temperature fuel cells. Being an

intermediate temperature fuel cell has several advantages such as flexibility on the purity of fuel

since catalyst poisoning does not occur at high temperature. In addition to this, SAFCs have

other several advantages such as simple built, small size (which reduces the size of the radiator

in a car engine) and total avoidance of fancy humidification provisions as in PEMFCs.

Nonetheless, the major concern associated with SAFC technology is its dependence on a large

amount of platinum to function. Currently, SAFC operation depends mostly on unsupported

platinum catalyst with a loading of approximately 3 mg/cm2 at the cathode terminal.24 This is 20

times higher than the loading requirements in a PEMFCs. Hence, commercialization of SAFCs is

possible only if there is a significant reduction in the loading of precious metals required to

sustain the fuel cell operation. This issue is discussed in detail in Chapter 2 and some alternative

catalysts idea and fundamental study has been provided which can eventually reduce the cost

associated with SAFC.

Figure 1.4 Mechanism of proton transport in CDP

35

1.3.2 Electrolyzer cells for generation of Chlorine

As discussed above, abundance of sea water makes chlor alkali electrolysis a convenient

technology, but massive power consumption is a major concern. During the process of

electrolysis of brine solution, chlorine is generated at the anode while hydrogen is evolved

because of the hydrogen evolution reaction (HER) at the cathode terminal. The electrochemical

reactions can be expressed as:

Anode: 2Cl- ↔ Cl2 + 2e- E°= 1.358V vs SHE (1.7)

Cathode (HER): 2H2O + 2e-↔ H2 + 2OH- E°= -0.828V vs SHE (1.8)

Cathode (ORR): 2H2O + O2 + 2e-↔ 2OH- E°= 0.401V vs SHE (1.9)

Per the standard reduction potential chart, this whole process involving HER and Cl2

generation translates to a voltage requirement of 2.18V (using equation 1.3) at zero current

output. Replacing the cathodic reaction from HER to ORR thermodynamically reduces the

voltage requirement to 0.95V. This technology of using oxygen electrodes in chlorine

electrolyzer plants have been widely adopted. Though an alkaline electrolyte allows relaxation

on the choice of cathode catalysts that reduce oxygen, the main challenge lies in the durability of

the catalyst in an electrolyzer cell.25 In chapter 5, further discussions on chlor alkali electrolyzer

cells, fabrication of stable, robust electrodes, and non-precious catalysts for application in chlor

alkali electrolyzers have been made.

1.4 Electrochemical Reactions of Interest

In electrochemical devices of our interest, the fundamental reactions that decide the fate

of the device are the ORR and the HOR. For convenience, current discussions will be done on

ORR and HOR, in both acid and alkaline electrolytes.

1.4.1 Electrocatalysis of Oxygen Reduction Reaction

36

The rate of ORR mainly dictates the overvoltage losses and the overall current output of

the device. ORR is a slow, irreversible process and suffers from huge kinetic losses. ORR

proceeds through two different pathways with several reaction intermediates which are

illustrated in Table 1.1. It can be understood from the values of Eo in Table 1.1 that 4-electron

reduction pathway is most desirable in order to derive the maximum output from the device. An

overall 4 electron reduction can also be achieved when the ORR follows a series pathway

involving 2 electrons in each step involving peroxide ion as the intermediate. In this dissertation,

ORR in both acid and alkali will be discussed using Platinum, non-platinum and non-precious

catalysts.

1.4.2 Electrocatalysis of Hydrogen Oxidation Reaction

Hydrogen oxidation is an interesting reaction that has very different mechanisms in acid

and alkaline electrolyte. The primary challenge associated with HOR is in alkaline electrolyte

where the reaction kinetics are very slow.26 In acid medium, electrochemical hydrogen oxidation

proceeds at a very fast rate and hence a minimal amount of platinum catalyst is enough to drive

the reaction.27 However, in alkaline electrolyte, HOR suffers fromsluggish kinetics amounting to

two orders of magnitude decrease in exchange current densityas compared to acid medium. The

origin of this loss is highly debated among scientists. Gasteiger et. al believe that that such slow

HOR kinetics in alkali is a result of stronger binding between hydrogen and Pt surface.28 This is

also termed as “binding energy theory”which suggests that hydrogen binds strongly to Pt surface

at higher pH, thereby slowing down the HOR that involves abstraction of hydrides from the Pt

surface. There is another theory that explains the slow kinetics of HOR, namely the bifunctional

theory. It states that the slow HOR kinetics in alkaline electrolyte arises from the non-availability

of reactive hydroxide ions at the electrode interface.29 Interestingly, when Pt is alloyed with Ru

37

or Ni, an enhancement in the HOR performance occurs. This observation, again, is largely

debated on the basis of binding energy theory andbifunctional theory. Gasteiger et al. believe

that such enhancement in the HOR catalysis in PtRu catalysts results from a weaker binding

energy between hydrogen and Pt, facilitated by Ru.30 On the other hand, Markovicet al.30 believe

that such increase in the alkaline HOR catalysis results from the availability of reactive hydroxyl

species furnished by the “oxophillic” Ru and Ni at the electrode- electrolyte interface. Chapters 3

and 4 will discuss more on the fundamental aspects of thealkaline HOR mechanism as well as

some alternative catalyst materials that have superior activity to platinum towardalkaline HOR.

Table 1.1 Oxygen Reduction in Acidic and Alkaline Electrolytes31

38

1.5 RDE and R(R)DE Techniques for Kinetic Study of Electrocatalysts

In order to design catalysts for electrochemical reactions, it is very important to be able to

analyze fundamentally an electrochemical process using the proper analytical tools.

Electrochemistry is basically a branch of chemistry that studies the interface of an electrode and

electrolyte, where reactions and transfer of electrons occur. A typical electrochemical reaction

can be represented as:

On+ + ne- →R (1.10)

Electrolyte Pathway Reactions

Thermodynamic

electrode potential

at standard

conditions

Acidic Media

Direct 4 electrons O2 + 4H+ + 4e- → 2H2O Eo = 1.229 V

Peroxide

O2 + 2H+ + 2e- → H2O2 Eo = 0.67 V

H2O2+ 2H+ + 2e- →

2H2O Eo = 1.77 V

H2O2→ 2H2O + O2

Alkaline

Media

Direct 4 electrons O2 + H2O + 4e- → 4OH- Eo = 0.401 V

Peroxide

O2 + H2O + 2e- → HO2-

+ OH- Eo = -0.065 V

HO2- + H2O + 2e- →

3OH- Eo = 0.867 V

2HO2- → 2OH + O2

-

39

Where n number of electrons are accepted by species O to get reduced to species R. This is a

representation of a reduction reaction. Similarly, if electrons are lost from a species, the reaction

is called an oxidation reaction. The potential of the half-cell can be represented as:

E = Eo’ + [(RT/nF) ln(CO*/CR*)] (1.11)

where Eo’is the formal potential, R is the universal gas constant (8.314 J/K mol) , T is the Kelvin

temperature, n is the number of electrons transferred in the reaction, F is the Faraday constant

(96,485 C/mol) and (CO*/CR*) represents the ratio of the bulk concentrations of O and R. The

formal potential (Eo’) precludes the use of activity coefficients and can be related to the standard

reduction potential with the following equation:

Eo’ = Eo + [(RT/nF) ln([O]/[R]) (1.12)

where [O]/[R] denotes the ratio of activities of O and R.

the overall electrode reaction (1.4) will be governed by several factors such as:

a) Charge transfer between O and electrode surface

b) Mass transfer (the rate at which O reaches the electrode surface from bulk electrolyte)

c) Other chemical reactions occurring along with the electrochemical reaction

d) Other surface phenomenon such as adsorption and desorption

The relationship between current and applied potential is expressed by the Butler Volmer

equation

i = i0 [exp (-αƒη) – exp((1-α)ƒη)] (1.13)

where i is the current density, i0 is the exchange current density, α is the charge transfer

coefficient, η is the overpotential and ƒ = F/RT. In the high overpotential region (η>>50 mV) the

second half of equation (1.9) transforms into theTafel equation:

40

η = [(RT/αF) ln(i0) – (RT/αF) ln(i)] (1.14)

A plot of η vs. ln(i), known as a Tafel plot, which gives information on critical kinetic

parameters such as i0, α and the Tafel slope (-2.3RT/αF). These values are directly related to the

fundamental traits of the catalyst operation.

Fundamental electrochemical studies involve cyclic voltammetry (CV), which can be

considered to be the baseline of all electrochemical characterization of a catalyst in electrolyte. A

CV can be collected by ramping the potential linearly versus time on a stationary electrode

which is generally set at a particular potential. Then the direction of the ramp is reversed and the

electrode is brought back to the original potential. All this while, the current response is collected

which reflects the surface phenomenon that occurs at the electrode. Hence, a CV gives important

information on the properties of any electrochemical process occurring on the electrode and

helps to identify electroactive species and adsorption process.

Figure 1.5 Cyclic Voltammetry of a Pt/C catalyst in 0.1M KOH. Data were collected at room temperature at 10 mV/sec scan rate

Rotating Disk Electrode (RDE) is a technique that is used to investigate the kinetics of an

electrochemical reaction. In other words, it is a tool that can directly analyze the catalytic profile

of the electrocatalyst that facilitates the electrochemical reaction. The Rotating Ring Disk

41

Electrode (RRDE) is a more specialized technique that allows investigation into electrochemical

reactions that involve an intermediate. A common example is ORR, where the intermediate

peroxide can be detected at the ring which is set at a such a potential that is suitable for the

oxidation of the peroxide intermediate. Therefore, in RRDE, two working electrodes are

employed- the first one being the disk electrode which is usually made of glassy carbon. The

Figure 1.6 Schematic representation of RRDE experiment along with the current responses from the disk and ring electrodes for Ru/C catalyst in 0.1M HClO4 collected under room temperature at the scan rate of 20 mV/sec.

second working electrode (usually made of gold or platinum) surrounds the first working

electrode but separated by an insulating layer of Teflon (which allows the independent control of

the potentials of the two electrodes). The working electrodes are rotated to facilitate a convective

transport of oxygen (which is the reactant, and dissolved in the electrolyte solution) to the

electrode surface. This ensures that the electrochemical reduction of oxygen is limited by mass

transport by increasing the transport of oxygen to the electrode surface. As a laminar flow is

created into the electrolyte due to the rotation, this allows the peroxide intermediate (formed at

42

the disk electrode) to flow towards the ring electrode, get oxidized and finally get detected by

measuring the current response at the ring electrode.

A typical RDE profile has three regions: 1) kinetic region 2) Mixed diffusion- kinetic

region and 3) Diffusion controlled region. In the kinetic region, reaction is kinetically controlled,

which means that the electrochemical reaction is limited by the reaction kinetics. In the diffusion

controlled region, the reaction is limited by the transport of reactant to the electrode surface from

the bulk electrolyte. The Levich equation is another important equation that relates the diffusion

Figure 1.7 ORR polarization curve obtained by RDE measurement for Pt/C catalyst in 0.1M HClO4. Data collected under room temperature at 20 mV/sec scan rate.

limiting current with the rotation rate:

ilim=0.62nFD2/3ν-1/6Coω½ (1.15)

where n is the number of electrons transferred, F is the Faraday constant, D is the diffusion

coefficient of O2 in the electrolyte, v is the kinematic viscosity,Co is the concentration of O2 in

the electrolyte and ω is the rotation rate (radian/s). Here ilimis the limiting current that is

generated when the reaction is limited by the rate at which reactant is supplied to the working

electrode (disk electrode) , which in turn, is facilitated by the rotation of the working electrode.

43

The total current can be expressed as a reciprocal sum of limiting current and kinetic current:

1/i = 1/ilim + 1/ik (1.16)

The kinetic current can be obtained from the total current density using the following equation:

ik = (ilim * i) / (ilim – i) (1.17)

1.6 X-ray Absorption Spectroscopy

X-ray Absorption Spectroscopy (XAS) in a type of spectroscopy that uses X-rays to

excite core level electrons into higher energy electron orbitals. This technique is an element

specific spectroscopy and requires access to a synchrotron radiation facility that provides tunable

X-ray. At the synchrotron facility, as the electrons revolve in a circular path, they emit photons

as a result of change in their angular momentum. This emission, known as polychromatic “white

light” is then sent through various optical devices to finally obtain a fixed wavelength at the

beamline station, known as the hutch. XAS is a very strong tool that provides detailed

information on the structure of the material, giving detail on atomic environment, coordination

numbers and oxidation states of the element of interest. XAS is widely used for ex situ

characterization where information on the local geometry and oxidation state of the element of

interest.However, some researchers use XAS under in situ conditions to determine the changes in

structure, oxidation states and surface adsorbates that happen under experimental conditions.

44

Figure 1.8 Schematic representation of the XAS experiment and data collection

A typical X-ray absorption spectrum can be broken down into two regions: a) X-ray

absorption near edge structure (XANES) and b) Extended X-ray absorption fine structure

(EXAFS). While XANES provides information on the valence state and local geometry of the

central atom (element of interest), EXAFS provides details on the local molecular structure that

involves coordination number and identity of the atomic neighbors.

Figure 1.9 Normalized XAS spectrum of Pt/C catalyst in 0.1M HClO4 held at 0.54V

For our experiments, a specially designed spectro-electrochemical flow cell is used to

study the electrode material under ex situ as well as in situ conditions. As shown in schematic,

three types of gas ionization detectors are used to measure incident beam (Io), transmitted signal

45

(It) and the reference signal (Ir). The sample of interest is placed between Io and It, and a

reference foil is placed between Io and Ir. The spectrum is collected typically from 200 eV below

the absorption

edge till 100 eV above the edge. The absorption of X-ray can be expressed by the Beer- Lambert

Law:

It = I0e-µt (1.18)

where It and I0 are the intensities of the transmitted and incident x-rays, µ is the absorption

coefficient and t is the sample thickness. The absorption coefficient is a material specific

property that depends on the constituent atoms and density of the material. When the energy of

the incident X-ray photons matches with the energy of the core level electrons, electrons are

transferred into empty states near the Fermi level causing a sharp increase in the absorption

coefficient. This is the absorption edge, and the sharp increase in µ is called the white line. As

the energy of the x-ray is further increased above the absorption edge the excess energy causes

the photoelectron to interact with the surrounding atoms, and is the reason behind the origin of

EXAFS. As the photoelectron wave is backscattered by the surrounding atoms both constructive

and destructive interferences occur that give rise to a particular EXAFS. Now it can be

understood that why EXAFS is dependent on the nature of the surrounding atoms, as well as the

general atomic environment of the material. The EXAFS spectrum can be described as the sum

of all the above contributing factors by the EXAFS equation:

χ(k) =∑��������

���

� � sin[2�R� + �k�]� (1.19)

where f(k) and δ(k) are scattering properties of the neighboring atoms, N is the number of

neighboring atoms, R is the distance to the neighboring atom, and σ2 is the disorder in the

46

neighboring atoms. By knowing the values of scattering amplitude f(k) and phase shift δ(k) Eq.

1.15 we can determine N, R and σ2. The scattering factors are dependent on Z of the neighboring

atom, which gives us identity of the neighboring atomic species.

XANES region is very sensitive to external perturbations such as presence of adsorbates on the

surface of adsorbing atom. This property is exploited in the specially designed Δµ technique

where XANES spectrum of a clean surface is subtracted from XANES spectrum of the same

material with adsorbates on the surface. Using FEFF8 code, theoretical Δµ signatures can be

generated and compared to the experimental Δµ spectrum. This technique, originally developed

by Ramakeret al. is highly useful to elucidate the surface- adsorption interactions at the

electrode- electrolyte interface of the electrocatalyst.32

1.7 Scope of Work

The topics covered in this dissertation primarily focus on the development of smart

electrocatalysts that have applications in a wide range of energy conversion devices. For

example, a non-platinum catalyst based on chalcogenides has been developed as cathodes in

SAFCs, which is a very recent technology and is dependent on huge loadings of platinum for

application. This dissertation work also focusses on non-precious catalysts that have been used in

chlorine generator plants as cathodes in chlor alkali technology which has been discussed in

Chapter 5. Fabrication of interesting binary platinum- metal catalysts have been developed to

understand the fundamentals of ORR and HOR mechanism and have been covered in Chapters 3

and 4. These binary platinum- metal catalysts comprise alloys, nano-composites as well as

electrodeposited platinum- metal systems which provide vital information on how the catalytic

profile of a platinum based catalyst gets modified based on the type of interactions between

platinum and the binary metal.

47

This work encompasses a wide variety of nanomaterial fabrication that have been studied to

elucidate fundamental reaction mechanisms as well as applied in various energy conversion

devices. As discussed in the previous sections, electrochemical energy conversion devices are

potential answers to accomplish a green economy and minimize the issues of climate change. In

this dissertation work, several systems have been studied comprehensively with viable

electrocatalyst options for both oxygen and hydrogen electrodes in acid and alkaline

environments. This work allows the universality of cheap catalysts with minimal or no platinum

requirement for application in electrochemical energy conversation devices.

1.8 References

(1) Shafiee, S.; Topal, E. Energy policy2009, 37, 181. (2) Sims, R. E.; Mabee, W.; Saddler, J. N.; Taylor, M. Bioresource technology2010, 101,

1570. (3) Hoffert, M. I.; Caldeira, K.; Benford, G.; Criswell, D. R.; Green, C.; Herzog, H.; Jain, A.

K.; Kheshgi, H. S.; Lackner, K. S.; Lewis, J. S. science2002, 298, 981. (4) Rodhe, H. Science1990, 248, 1217. (5) Boykoff, M. T. Area2007, 39, 470. (6) Church, J. A.; White, N. J. Geophysical research letters2006, 33. (7) Marchetti, C. Climatic change1977, 1, 59. (8) Manabe, S.; Wetherald, R. T. Journal of the Atmospheric Sciences1980, 37, 99. (9) Lokurlu, A.; Grube, T.; Höhlein, B.; Stolten, D. International journal of hydrogen

energy2003, 28, 703. (10) Wang, W.; Luo, Q.; Li, B.; Wei, X.; Li, L.; Yang, Z. Advanced Functional

Materials2013, 23, 970. (11) Hirose, K. Faraday discussions2011, 151, 11. (12) Richards, J. A.; Google Patents: 1966. (13) Subbaraman, R.; Tripkovic, D.; Strmcnik, D.; Chang, K.-C.; Uchimura, M.; Paulikas, A.

P.; Stamenkovic, V.; Markovic, N. M. Science2011, 334, 1256. (14) Choi, P.; Jalani, N. H.; Datta, R. Journal of the electrochemical society2005, 152, E123. (15) Zawodzinski, T. A.; Derouin, C.; Radzinski, S.; Sherman, R. J.; Smith, V. T.; Springer,

T. E.; Gottesfeld, S. Journal of the electrochemical society1993, 140, 1041. (16) Ramaswamy, N.; Mukerjee, S. Advances in Physical Chemistry2012, 2012. (17) Varcoe, J. R.; Slade, R. C. Fuel cells2005, 5, 187. (18) Arges, C. G.; Ramani, V. K.; Pintauro, P. N. The Electrochemical Society Interface2010,

19, 31. (19) Li, N.; Guiver, M. D.; Binder, W. H. ChemSusChem2013, 6, 1376. (20) Boysen, D. A.; Uda, T.; Chisholm, C. R.; Haile, S. M. Science2004, 303, 68. (21) Haile, S. M.; Boysen, D. A.; Chisholm, C. R.; Merle, R. B. Nature2001, 410, 910.

48

(22) Yamada, K.; Sagara, T.; Yamane, Y.; Ohki, H.; Okuda, T. Solid State Ionics2004, 175, 557.

(23) Otomo, J.; Tamaki, T.; Nishida, S.; Wang, S.; Ogura, M.; Kobayashi, T.; Wen, C.-j.; Nagamoto, H.; Takahashi, H. Journal of applied electrochemistry2005, 35, 865.

(24) Ghoshal, S.; Jia, Q.; Li, J.; Campos, F.; Chisholm, C. R.; Mukerjee, S. ACS Catalysis. (25) Morimoto, T.; Suzuki, K.; Matsubara, T.; Yoshida, N. Electrochimica acta2000, 45,

4257. (26) Sheng, W.; Gasteiger, H. A.; Shao-Horn, Y. Journal of The Electrochemical Society2010,

157, B1529. (27) Schmidt, T.; Ross, P.; Markovic, N. Journal of Electroanalytical Chemistry2002, 524,

252. (28) Durst, J.; Siebel, A.; Simon, C.; Hasche, F.; Herranz, J.; Gasteiger, H. Energy &

Environmental Science2014, 7, 2255. (29) Strmcnik, D.; Uchimura, M.; Wang, C.; Subbaraman, R.; Danilovic, N.; Van Der Vliet,

D.; Paulikas, A. P.; Stamenkovic, V. R.; Markovic, N. M. Nature chemistry2013, 5, 300. (30) Danilovic, N.; Subbaraman, R.; Strmcnik, D.; Paulikas, A.; Myers, D.; Stamenkovic, V.;

Markovic, N. Electrocatalysis2012, 3, 221. (31) Marie Strickland, K., PhD Thesis, Northeastern University, 2015.

(32) Arruda, T.M,; Shyam, B.; Ziegelbauer, J.M.; Mukerjee, S.; Ramaker, D.E. The Journal of Physical Chemistry2008, 112, 18087

49

Chapter 2

Electrochemical and In Situ Spectroscopic Evidences towards Empowering

Ruthenium Based Chalcogenides as Solid Acid Fuel Cell Cathodes

This Chapter is based on a published paper with the same title

ACS Catal., 2017, 7 (1), pp 581-591

Shraboni Ghoshal1, Qingying Jia1, Jingkun Li1, Fernando Campos2,Calum R.I. Chisholm2 and Sanjeev Mukerjee1

1Department of Chemistry, Northeastern University, Boston, MA 02115, United States

2SAF Cell Inc., 36 S. Chester Ave, Pasadena, CA 91106, United States

50

Chapter 2

Electrochemical and In Situ Spectroscopic Evidences towards Empowering

Ruthenium Based Chalcogenides as Solid Acid Fuel Cell Cathodes

2.1 Introduction

Establishment of fuel cell technology in electrical power market requires addressing

certain concerns; such as high manufacturing cost, demand of ultra-pure fuels costs and cost of

precious metal catalysts.1,2 Majority of these issues can be addressed by employing intermediate

temperature fuel cellswhich work in the temperature range between 180 degrees to 300

degrees.3One of the earliest examples, phosphoric acid fuel cell (PAFC) with typical operating

temperature between 150-2000C clearly demonstrates the potential advantage of combined heat

and power (CHP) with total CHP efficiency in the range of 70% and a concomitant lowering of

the balance of plant involved in reforming fuels such as methane, with demonstrated CO

tolerances in the range of 1-2%.4-6 Although this reduces the cost of fuel, phosphoric acid fuel

cells have limitations such as prolonged membrane electrode assembly (MEA) preparation,

dependence on platinum as anode and cathode catalyst and inadequate heat cultivation.7In fuel

cells, the use of higher temperature ceramic proton conductors based on hydrogen bonded

tetrahedral oxyanions in conjunction with a metal cation such as Cs have been introduced in

2004 by Haile et al.,8,9as exemplified by cesium dihydrogen phosphate, abbreviated as

CDP.10Such solid acids like CDP become viable proton conductors after undergoing super-

protonic phase transition involving dramatic structural reorientation within the hydrogen bonded

network of the oxyanion moieties, resulting in enhancement of proton conductivity by as much

as 4 orders of magnitude.11 For CDP, this transition initiates at 231°C, which makes the

operating temperature of SAFC to be around 250°C.12Such solid acid fuel cells (SAFC) offers

51

several advantages– (i) SAFC allows relaxation on the purity of the feedstock (ii) requires simple

build up and thus minimizes the radiator size, (iii) Unlike Nafion, CDP exhibits anhydrous

proton transport (super-protonic transfer); hence fancy humidification system is not required and

(iv) enables inexpensive stack construction.13,14 However, the primary concern in enabling SAFC

to reach commercial success is the cost associated with the membrane electrode assembly

(MEA). Also, CDP reacts with a vast number of metal oxides and forms non-conductive phases,

thus severely restricting the choice of catalytic materials.Currently, SAFC operation relies

mostly on unsupported platinum catalyst, with an average loading ~ 3 mg/cm2. This translates to

11- 25 g of platinum/KW, which is approximately twenty times higher than the loading

requirement in PEMFC. Therefore, the immediate goal is to significantly reduce the MEA cost

by employing cheaper catalysts.

Ongoing efforts towards engendering non-precious metal electrocatalysts for oxygen

reduction reaction (ORR) have resulted in various catalyst configurations; such as (a) metal-

nitrogen- carbon network (Zelenayet al.,and Dodeletet al.,)15-17, (b) transition metal

nitrides/oxynitrides (Ota et al.,)18-21 and (c) transition metal-chalcogen cluster compounds.22-24

Catalysts of type (a) comprise carbon as the major constituent, with trace amount of active

metals (< 5% by weight). These catalysts have appreciable ORR activity comparable to that of

Pt/C in acid.25,26 However, high carbon content restrains the use of these catalysts in SAFC

owing to both poor mass transport considerations and carbon corrosion. The transition metal

nitrides (Ota et al.,18,20,21,27) mostly based on nitrides and oxy-nitrides of Ta, Nb and Zr,also

suffer from very low density of active sites and relatively lower activity compared to their

contemporaries resulting in poorer ORR activity, even at elevated temperature.Hence, this work

is based on examining the choice of semiconducting transition metal chalcogenides.

52

Semiconducting, crystalline Chevrel phases (such as RuxMo6-xX8) first reported by Alonso Vante

et al.,23,28 have been shown to reduce oxygen both in acid,23 and alkaline environments29 at low

temperatures (ambient). In such catalysts, the transition metal clusters (Ru) act as reservoirs of

electronic charge and form the core of the particle, whereas the chalcogen (X) resides at the

periphery coordinated to the Ru core.30 Ru and chalcogen coexist on the surface of the particle,

where Ru serves as the active center for molecular oxygen binding, while X provides protection

to Ru against undergoing oxidation.30It is believed that the contribution of Mo is mostly to

increase the oxygen binding ability of the catalyst, though no fundamental details on this

phenomenon are present in literature. In addition to their electrocatalytic activity towards ORR,

these catalysts show exceptional endurancewhen exposed to strongly adsorbing anions such as

Cl- and H2PO4-. Such anion resistance has been extensively reported by Ziegelbaueret al.,31 in the

context of HCl recovery process. A report on the application of various chalcogen modified

catalysts (using variable constituents and different synthetic routes) in HCl recovery cell has

been published and these catalysts have shown to outperform the state of the art catalysts e.g.

Pt/C and Rh/C.31 A plenary study on Ru/C with sequential addition of Se and Mo, and their

effect on ORR catalysis in alkaline electrolyte has been reported by Ramaswamy et al.29,32where,

a detailed X-ray absorption spectroscopy (XAS) experiment under in situ conditions have

provided information on structure and mechanism of the resulting catalysts in alkaline media.

In this report, we have critically investigated the prospects of Ru-based chalcogenides as

potential oxygen electrode in SAFCs. Electrochemical characterizations on Ru/C, RuSe/C and

RuSeMo/C have been performed in acid electrolyte to understand the effect of chalcogen

modification on Ru surface. Surface characterization, ORR catalysis profile and anion poisoning

effect have been combined to delineate the complete profile of the electrocatalysis. In situ XAS

53

experiments at Ru, Se, and Mo edges provide fundamental evidence which helpin elucidating the

structural and electronic features of the catalysts in acid electrolyte. Through thermogravimetric,

electrochemical, and in situ XAS experiments, we have successfully explained the remarkable

stability of the chalcogenide catalyst over a wide temperature range. To our knowledge, this is

the first time SAFC operation is demonstrated using a cathode catalyst devoid of Pt or Pd. This

report also constitutes the highest temperature at which a chalcogen has been shown to function

as an ORR electrocatalyst.

2.2Experimental Section

2.2.1. Synthesis of electrocatalysts

The Ru/C, RuSe/C and RuMoSe/C catalysts were synthesized via a wet chemical

reduction method in aqueous medium as described in a prior publication.29 To synthesize Ru/C,

Vulcan XC-72-R carbon and ruthenium chloride were dispersed in water and heated at 80°C.

The slurry was cooled down and alkaline sodium borohydride solution was added as the reducing

agent. It was followed by heating the slurry at 80°C, in order to complete the reduction process.

The slurry was then filtered, dried and pyrolyzed at 500ºC for 2 hours under Argon at a slow

heating rate of 4ºC min-1. To synthesize RuSe/C and RuSeMo/C, calculated amount of selenium

oxide and phosphomolybdic acid were added along with ruthenium chloride. A 46.1% Pt/C

procured commercially (Tanaka, Japan) was used for comparative studies.

2.2.2 Physicochemical Characterization

X- ray diffraction characterization was conducted using a RigakuUltima IV XRD with Cu

Kα source (lambda= 1.541 Å) operated at 40 kV and 44 mA. 2θ/θ scans were conducted using a

0.05 step size and 5 sec hold per step. Scanning electron microscopy (SEM) characterizations

54

were performed using a Hitachi S- 4800 FE- SEM. For validating sample elemental composition,

EDAX Genesis on the same SEM instrument was used. Thermogravimetric analysis (TGA) were

done using TA instruments model Q 600. To collect the data, sample was heated till 250ºC under

air at a heating ramp rate of 10ºC min-1. The sample was isothermally held at 250ºC for 30 mins

under air atmosphere, followed by a cool down procedure to room temperature.

2.2.3 Electrochemical Characterization

All electrochemical measurements were done at room temperature using a rotating ring

disc electrode (RRDE) from Pine Instruments connected to an Autolab (Ecochemie Inc., model

PGSTAT 30) potentiostat/ galvanostat. Tests were conducted in a 50 ml jacketed three electrode

cell in 0.1 M HClO4. For each test, freshly made, reversible hydrogen electrode (RHE), was used

as the reference electrode with a Pt mesh wire serving as a counter electrode. A glassy carbon

disc (geometrical area of 0.247 cm2)with a gold ring was used as working and ring electrode

respectively. All the results shown here were collected after conditioning the electrodes at a scan

rate of 50 mV/sec for 50 scans in a voltage range between (0.05V – 0.8V), or until stable features

were achieved. For Ru based catalysts, the catalyst loading was maintained at 20 ug/cm2geo. Gold

ring used for ring collection data was held at 1.1V vs. RHE while carrying out the ORR

measurements.

2.2.4 X Ray Absorption Spectroscopy (XAS) Measurements

The in situsynchrotron XAS studies were conducted at the Ru (K-edge, 22117 eV), Se

(K-edge, 12658 eV) and Mo (K-edge, 20000 eV) binding energies at SSRL (Stanford Linear

Accelerator Center, CA) beamline 2-2. Detailed description of the spectro-electrochemical cell is

given in a previous publication from our group.33 All the data at the Ru K-edge were collected in

55

transmissionwhile all the Se and Mo K-edge data were collected under fluorescence mode. For

transmission data, typically three gas ionization detector (Io, It/ Ifand I ref) setup was used with

10% photon absorption in I0 and 50%-70% in It and Iref.The fluorescence detector used was a 13

element Ge detector.Typical loading of the electrode used in these studies were based on the

transmitted x-ray absorption cross section designed to provide a step height of unity. Nitrogen

saturated 0.1 M HClO4 was used as the electrolyte along with an RHE made out of the same

electrolyte. Complete details of EXAFS analysis have been described in detail previously.29,34

Briefly, IFEFFIT suite 1.2.9 was used for background subtraction using AUTOBK algorithm and

normalization. Typical K-range window used was 3-14.0 Å-1. The data were processed and fitted

using the Athena35 and Artemis36programs.The χ(R) transforms were modeled using single

scattering paths calculated by the FEFF6 code.37 In addition to the bulk averaged EXAFS

analysis for obtaining short range atomic order around Ru (bond distance, coordination number,

Debye Waller factor etc.) surface specific information was obtained using a previously

developed38,39 subtractive method referred to as delta mu (Δμ) involving the near edge part of the

spectrum (X-ray absorption near edge structure, XANES), where the effect of the invariant bulk

signal was removed from the surface by subtracting XANES measured at a reference potential

(0.1 V vs. RHE) from other potentials of interest. Data analysis for delta mu (Δμ) studies at Ru

K-edge X- ray absorption near-edge structure (XANES) spectra involved specific normalization

procedures detailed elsewhere.29 This involved careful calibration of edge energy (Ru K-edge,

22117 eV) and alignment to the standard reference scan to account for any drift in the beam

energy. A post edge normalization procedure was then applied to the aligned scans via a cubic

spline function, which normalizes the oscillations over a specific energy range (typically 25-200

eV with respect to E0) thus enabling normalized data on a per-atom basis. Difference spectra

56

were obtained using the equation Δμ = μ(V) - μ(0.1 V), where μ(V) is the XANES at various

potentials and μ(0.1 V) is the reference signal at 0.1 V, where Ru surface can be considered to be

clean (i.e., free of surface adsorbates). These experimental difference signature profiles were

then compared with theoretically generated profiles constructed using cluster models and the

FEFF 8.0 code. Detailed modeling of the clusters in the context of surface adsorbed species was

not necessary in this study as the primary focus was to examine the extent of susceptibility

towards anion adsorption. Prior studies of such nature have been reported.40,41

2.2.5 MEA Testing in Solid Acid Fuel Cell

(a) Membrane electrode assembly (MEA) preparation: For a typical SAFC ¾” button cell,

MEAs were made using a standard carbon supported platinum for anode and RuMoSe/C catalyst

as cathode. For the anode side, a nickel mesh was used as a current collector while a stainless

steel mesh served as the current collector on the cathode side. The proton conductor was a layer

of CDP pressed against the catalyst layers. A nickel mesh, carbon paper, anode catalyst and the

CDP layer were compiled together using a Teflon sheet along the circumference to assemble the

half MEA. The cathode catalyst was then mixed with CDP in a ratio of 1:3. 50 mg of this

catalyst CDP mixture was taken and spread on the CDP layer. Metallic shims were used to

spread the catalyst layer evenly. The full MEA was then pressed under a pressure of 3 metric

tons for 3 seconds. A firm and even spreading of the catalyst layer could be seen after taking it

out of the press station. A Teflon tape was used to wrap the MEA along the circumference. This

MEA was used for testing in a solid acid fuel cell.

(b) MEA testing protocol: MEA was placed inside the test fixture and sealed properly to avoid

potential gas leakage. The test fixture was heated in a step wise manner with different types of

gases flows at different stages. In the first step, dry nitrogen gas was purged till 150°C. Once at

57

150°C, humidified nitrogen gas was fed till 250°C, using a very slow ramp rate. When a stable

250°C temperature was reached, humidified H2 and air as feedstocks were introduced to the

anode and cathode chambers respectively. After recording the open circuit voltage (OCV) for 15

minutes, I-V polarization measurements data were collected. After completion of data collection,

cell temperature was gradually reduced, under flowing nitrogen gas, in order to shut down the

test station.

2.3Results and Discussions

2.3.1 Physicochemical Characterizations

Table 2.1 Physicochemical characterization of heat treated Ru/C, RuSe/C and RuSeMo/C catalysts

using EDS

Elemental analysis of Ru/C, RuSe/C, and RuMoSe/C were performed using the energy

dispersive spectra (EDS), and the results are listed in Table 2.1. Representative XRD profiles of

Ru/C and RuMoSe/C collected from 2θ values of 20º- 80ºare shown in Figure 2.1. The pattern

for Ru/C contains distinctive Ru phase with a hexagonal close packed structure ((hcp), space

group P63/mmc) with a lattice constant of a=b= 2.7106 and c= 4.2911, in good agreement with

JCPDS powder diffraction pattern.The XRD pattern for RuSe/C shows the presence of the

primary peaks of Ru hcp structure, along with peaks corresponding to RuSe2. RuSe2 exists as a

pyrite structure with space group Pa3 and lattice constants a=b=c= 5.9297. Based on Ru (101)

crystallite line broadening, the ruthenium crystallite size was estimated to be 8.6 nm for Ru/C

and 5.4 nm for RuSe/C, which corroborates that Se atoms reduce the sintering of underlying Ru

particles during heat treatment. The comparison of the XRD patterns between Ru/C and RuSe/C

Catalyst Ru (wt%) Ru:Se:Mo Ru/C 40% ---

RuSe/C 30% 1: 0.5 RuSeMo/C 30% 1: 0.5: 0.2

58

reveals that the underlying Ru lattice did not undergo significant change upon addition of

selenium. XRD patternof RuSeMo/C was similar to that for RuSe/C, (not shown here), indicating

that Mo does not affect the crystallographic profile of the catalyst.

Figure 2.1 X ray diffraction pattern of Ru/C and RuSe/C with representative crystal structures

2.3.2. ORR on Ru based chalcogenide catalysts

(a) Cyclic Voltammetry (CV): Electrochemical surface characterization: Figure 2.2(a) compares

the CV of Ru/C and RuSe/C catalysts in Ar- saturated 0.1M HClO4 electrolyte. Ru/C shows a

distinctive hydrogen adsorption and desorption peak, which disappearsafter modifying Ru/C

with Se. Such behavior has been reported in other publications as well.42 Double layer

capacitance decreases following Se addition to Ru/C, which indicates an increase in charge

transfer resistance for both hydride and oxide formation originating from Ru and Se

interactions.43 In Ru/C, an anodic current corresponding to Ru-OH formation and a cathodic

current corresponding to the reduction of Ru-OH can be detected. Comparing the CVs of RuSe/C

and Ru/C, it is clear that Se modification significantly suppresses Ru-OH formation. Figure 2.2

(b) shows the effect of heat treatment on RuSe/C. In the untreated sample, the irreversible

oxidation peak at 0.83V in the cyclic voltammogram of figure 2.2(b) indicates oxidation of Se0

59

to Se4+. The broad anodic peak from 0.2 V till 1.0V is due to the combined formation of RuO2

and SeO2. However, after heat treatment, there is a single anodic peak at 0.90V vs. RHE

corresponding to Se oxidation, thus indicating stabilization of the Ru/C moiety. This observed

increase in stability of heat treated RuSe/C is due to an improved alloy formation between the

metals, thereby preventing the formation of Se and Ru oxides.43,44The CV of RuSeMo/C had

similar features as that for RuSe/C as shown in Figure 2.3.

Figure 2.2 Cyclic voltammetry in deoxygenated 0.1M HClO4 electrolyte measured at 20 mV/s scan rate of (a) Ru/C(HT) and RuSe/C (HT) catalysts (b) untreated and heat treated RuSe/C catalysts

Figure 2.3Cyclic voltammogram of Ru/C, RuSe/C and RuSeMo/C (all HT) in Ar purged 0.1M HClO4 at 20 mV/sec sweep rate

60

(b) ORR measurement on chalcogen modified Ru/C catalysts: Figure 2.4(a) compares the ORR

activity profiles of Ru/C, RuSe/C, and RuSeMo/C catalysts in O2 saturated 0.1M HClO4

electrolyte recorded at 1600 rpm. The onset of oxygen reduction reaction in Ru/C occurs at

0.81V vs. RHE and ORR is kinetically controlled even at high overpotential, with no defined

limiting current region. Addition of Se undeniably improves the ORR activity, with the onset

now shifting to 0.85 V vs. RHE. The mixed kinetic-diffusion region between 0.8V and 0.5V is

followed by a well-defined limiting current. Chalcogen modification causes a net 86 mV anodic

shift in the half wave potential compared to Ru/C. The effect of the ternary metal molybdenum is

self-explanatory in the ORR profile of RuSeMo/C. The onset of ORR for RuSeMo/C happens at

0.86V, with a positive shift in half wave potential by 96 mV compared to Ru/C. RuSeMo/C

exhibits a mixed kinetic- diffusion current region between 0.86V and 0.5V, and a prominent

limiting current region can be identified. This is in accordancewith our previously reported data

where similar trends were obtained in alkaline electrolyte.29The representative Tafel plots and

Koutecky- Levich (K-L) plots are presented in figure 2.4(b). Tafel slopes of Ru/C, RuSe/C and

RuSeMo/C exhibit similar features with a typical two slopes region. This seems to indicate that

the mechanism for ORR most likely does not change on chalcogen modification or ternary Mo

addition. Detailed Tafel slope analysis however was not possible as they are not amenable on

supported catalysts. Assessment of the K-L plots provided further insight into the effect of

chalcogen modification on Ru/C. The theoretical Levich slope was calculated using the

following parameters: Diffusivity (DO2) = 1.9x 10-5 cm2/s, solubility (CO2

) = 1.18x 10-6mol/cm3

and kinematic viscosity (ν) = 8.93x10-3 cm2/s. The number of electrons transferred on Ru/C was

determined to be 2.95, which increased to 3.17 on addition of Se. In RuSeMo/C, K-L plot

analysis showed the number of transferred electrons to be 3.5. This supports the theory of

61

peroxide spillover by Mo, which aids in shifting the ORR pathway towards the more efficient 4

electron route. Similar conclusion on the role of Mo was reported by Alonso Vante et al. in their

studies on the effect of Mo on the ORR profiles of Ru/Se systems.45 The respective ring current

response are shown in Figure 2.5 to get an illustrative idea on the H2O2 generation, and the

corresponding results obtained from RRDE experiments as reported in Table 2.2.

It is intriguing to note how the ORR pathway shifts from a 2 electron route for

unmodified Ru/C towards a 4 electron route in Se and Mo modified Ru/C. Vante et al.46 have

explained this in terms of the binding mode of O2 on Ru surface in the different catalysts. While

unmodified Ru/C gets passivated easily, the number of Ru active sites are greatly reduced. This

prompts the oxygen molecule to bind in a terminal mode, which directs ORR to a 2 electron

reduction pathway to H2O2.

Table 2.2 Summarization of results obtained from RDE and RRDE measurements and

comparison between Ru/C, RuSe/C and RuSeMo/C (all HT) catalysts

On the other hand, Se modification helps to preserve the Ru active sites by preventing

them from getting oxidized. This leads to a twofold binding of oxygen on Ru surface, thus

ensuring dissociative adsorption of O2 leading to 4 electrons reduction. All the above results

unequivocally imply thatmodification of Ru/C catalyst with Se and Mo is crucial when

implementing Ru based system as ORR electrocatalyst in a fuel cell. However, considering the

phosphate–rich electrolyte in SAFC, it is important to ascertain that RuSeMo/C is resistant to

phosphate anion poisoning. Therefore, calculated amount of phosphoric acid was doped into

Catalysts ne from K-L analysis ne from RRDE @ 0.2V vs. RHE

Selectivity H2O @ 0.2V vs. RHE

Ru/C 2.95 2.47 80% RuSe/C 3.17 3.23 92.5%

RuSeMo/C 3.5 3.6 96.8%

62

0.1M HClO4 electrolyte and corresponding ORR polarization curves for RuSeMo/C were

collected.

Figure 2.4Oxygen reduction reaction on Ru/C, RuSe/C and RuSeMo/C (all heat treated) in O2- saturated 0.1M HClO4 electrolyte (a) ORR polarization curves measured at 1600 rpm and 20 mV/s scan rate (b) Mass transport corrected Tafel plots and (inset) K-L plots

Figure 2.5 Ring current response for Ru/C, RuSe/C and RuSeMo/C in O2 purged 0.1M HClO4 collected at 1600 rpm at a scan rate of 20 mV/sec.

(c) Effect of H3PO4 addition on ORR performance for RuMoSe/C:The presence of phosphates in

CDP can potentially restrain catalytic activity of the oxygen electrode in a SAFC. Phosphate

anions have strong affinity towards platinum, thereby resulting in significant poisoning of active

sites at concentrations as low as 0.1 mM. 47In order toexamine the tolerance of the chalcogenide

63

catalyst in presence of phosphate ions, oxygen saturated- 0.1 M HClO4 electrolyte was doped

with calculated amount of phosphoric acid. Figures 2.7(a) and 2.7(b) show the ORR profile of

both Pt/Cand RuMoSe/C in the presence of phosphate ions, where Pt/C catalyst is used as

controlexperiment.

Figure 2.6Oxygen reduction reaction on 30% RuSeMo/C and 50% RuSeMo/C (all heat treated) in O2- saturated 0.1M HClO4 electrolyte, collected at 1600 rpm and 20 mV/sec scan rate.

Both the catalysts exhibit three characteristic potential regions: a diffusion controlled region (<

0.4V for RuMoSe/C and < 0.7V for Pt/C), mixed diffusion–kinetic limitation region (0.45V to

0.75V for RuMoSe/C and 0.65V to 0.85V for Pt/C) and a Tafel region (> 0.8V for RuMoSe/C

and > 0.9 V for Pt/C). For both the catalysts, the half wave potential (E1/2) shifts cathodically

with increasing concentration of H3PO4. The E1/2 values at different concentrations of H3PO4 are

shown in table 2. Interestingly, for RuMoSe/C, the overall shift in E1/2 value amounted to 18 mV

in the concentration range of 0 mM to 100 mM H3PO4, whereas for the same concentration range

of H3PO4 a net shift of 100 mV in E1/2 value occurred for Pt/C. Assuming the concentration of O2

does not change while adding the aliquots of H3PO4, the number of exchanged electrons were

calculated using the K-L equation. In Pt/C, n decreases from 4.00 (at 0 mM H3PO4) to 3.87 at

100 mM H3PO4 concentration whereas for RuSeMo/C, n drops from 3.50 (at 0 mM H3PO4) to

64

3.43 at 100 mM H3PO4 concentration. These results clearly show that compared to Pt/C,

phosphate adsorption is significantly suppressed on the chalcogenide catalyst.

Figure 2.7 Phosphate poisoning study in O2- saturated 0.1M HClO4measured at 1600 rpm and 20 mV/sec scan rateon (a) RuSeMo/C and (b) Pt/C

Park et al.48 studied the effect of phosphoric acid on Se modified Ru/ systems, and

reported enhancement of ORR activity. They ascribed adsorption of phosphate anion on Se

atoms and lower Oads on Ru atoms as a possible reason functioning via a combination of steric

and electrostatic repulsions. Our results however did not show any enhancement in ORR activity,

but the effect on onset potential, limiting current or half wave potential is minimal as compared

to Pt/C.48Analysis of the mass transport corrected Tafel plots can give an estimation of the

coverage of phosphate anions on an electrocatalyst surface. Assuming that ORR reaction is first

order with respect to O2 concentration, the Tafel curves are derived as:

� � �� + � log#

#�

65

Table 2.3 Effect of phosphoric acid doping on the ORR performances of Pt/C and RuSeMo/C

catalysts in O2 saturated 0.1M HClO4 electrolyte measured at 1600 rpm and 20mV/sec scan rate

Catalyst Conc. Of

H3PO4

(mM)

Number of

transferred

electrons (n)

E ½

(V)

Tafel slopes

(mV/dec)

Δ

(mV/dec)

Φ

Pt/C

0 4.00 0.893 159/60 99 1 1 4.00 0.864 144/61 83 1.16

10 3.92 0.818 160/76 84 1.16 100 3.87 0.792 154/75 79 1.26

RuSeMo/C

0 3.5 0.725 143/81 62 1 1 3.5 0.721 139/75 64 0.97

10 3.45 0.718 137/63 63 0.98 100 3.43 0.707 134/73 61 1.01

Where E0= 1.23V vs. RHE under standard conditions, b is Tafel slope, j0 is exchange current

density and jk is the kinetic current density. Typically, two different Tafel slopes are ascribed to

difference in the extent of oxide coverage on the catalyst surface in low and high current

densities. He et al.,47,49have utilized this trend of Tafel slopes to correlate effects of oxide

formation vs. those arising from effects of oxide formation along with H2PO4- adsorption. Values

of Δ and Φ can determine the extent of H2PO4- adsorption, in a semi qualitative manner, where Δ

and Φ can be defined as:

$ � %& ' %(,

* � ∆,�/∆,

where th and tl are the Tafel slopes in high and low current density regions, ΔM0 and ΔM are the Δ

values in 0.1M HClO4, with and without H3PO4 doping respectively. A lower ΔM results in

higher Φ, which is indicative of H2PO4- adsorption interfering with oxide formation. As seen in

Table2.3, Φ increases rapidly for Pt/C, whereas it remains mostly uninterrupted for RuMoSe/C.

66

These calculations unanimously support the RDE experimental results and validate the higher

tolerance of the chalcogenide catalyst against phosphate poisoning. Doping experiments were

performed using cesium dihydrogen phosphate and similar results were obtained as shown in

Figure 2.8.

Figure 2.8Oxygen reduction reaction on (a)Pt/C and (b)30% RuSeMo/C (all heat treated) in O2- saturated 0.1M HClO4 electrolyte in presence of cesium dihydrogen phosphate collected at 1600 rpm and 20 mV/sec scan rate

2.3.3In SituXAS Experiment

XAS measurements were carried out under in situ electrochemical conditions in 0.1 M

HClO4 at the K-edges of Ru, Se, and Mo in the RuSeMo/C catalyst, as well as the Ru/C

nanoparticles (NPs) for comparisons.It should be noted that the XAS was conducted at room

temperature whereas the SAFC cell works at ~250 °C. Despite this, it is reasonable to assume the

structure of the catalysts does not change drastically within this relatively low temperature range,

especially the structural difference between the Ru/C and RuSeMo/C. This is manifested in the

stable thermal behavior (up to 250°C) observed using DSC/TGA. Therefore, the in situ XAS, in

combination with ex situ characterization results given above, can provide valuable information

of the structural and electronic properties of the catalysts as shown below.

67

Figure 2.9(a) shows the XANES spectra obtained at the Ru K-edge on the Ru/C–HT

catalyst in Ar-saturated 0.1 M HClO4 electrolyte at various potentials; the corresponding Fourier

transformed (FT) EXAFS spectra are displayed in Figure 2.9(b). The highlighted region in

Figure 2.9(a) shows the anodic shift of the XANES with increasing potentials, indicating the

increase in Ru oxidation state with increasing potentials. Concomitantly, the FT-EXAFS peak at

2.5 Å gradually decreases with increasing potential; due to gradual insertion of oxygen atoms

into the Ru lattice blocking the Ru-Ru scattering. In sharp contrast to Ru/C, the XANES

spectrum for RuSeMo/C at the Ru K-edge does not change with increasing potentials going from

0.1 to 0.9 V (Figure 2.10(a)), thereby indicating that the oxidation state of Ru practically remains

unchanged within the investigated potential window.

Figure 2.9 In situ XANES (a) and FT- EXAFS (b) of Ru/C (HT) catalyst in Ar saturated 0.1 M HClO4 electrolyte at Ru K-edge.

The extent in adsorption of oxygen onto Ru surfaces at elevated potentials is probed by

the surface-sensitive Δμ analysis as displayed in Figures2.11(a). The ascendingΔμ amplitude

with increasing potentials indicates that the Ru sites are progressively occupied by oxygen

adsorbates with increasing potentials. This trend is expected, given the strong oxophilic character

of Ru which causes formation of surface oxides at even low potentials. Therefore, a large

fraction of the Ru active sites starts to get poisoned by oxygen species at relatively low

68

potentials, and this site-blocking effect limits the catalytic performance of the Ru/C catalysts. In

contrast, the corresponding Δμ signals for RuSeMo/C catalyst are minimal (Figure 2.11(b)). This

indicates that the oxidation of Ru does not occur till 0.9V. The significant delay of the Ru

oxidation in RuSeMo/C catalyst compared to Ru/C greatly alleviates the site-blocking effect, and

hereby accounts partially for the enhanced ORR activity. Figure 2.11(c) explicitly brings out the

oxygen coverage difference represented by Δμ amplitude as a function of applied potential in the

surface Ru in Ru/C and RuSeMo/C. In addition, the delay of the Ru oxidation till potential

beyond 0.9 V reflects the weakening of the Ru-O binding energy, which may also improve the

ORR activity given that the binding energy between the Ru/C and oxygen is overly strong.50

Figure 2.10 In situ (a) XANES and (b) FT-EXAFS of RuSeMo/C (HT) catalyst collected in deaerated, Ar saturated 0.1M HClO4 electrolyte at Ru K-edge.

To unravel the structural origin of the suppression of the Ru oxidation or the weakening

of the Ru-O binding energy in the RuSeMo/C catalyst, EXAFS fitting analysis were conducted at

the Ru and Se K-edge simultaneously. The fitting results are listed in Table 2.4. As seen, the Ru-

Ru coordination number (NRu-Ru) is much higher than the Ru-Se coordination number (NRu-Se),

and the Se-Ru coordination number (NSe-Ru) is also moderately higher than the Se-Se

coordination number (NSe-Se). These results indicate that Ru and Se are surrounded mostly by Ru

69

atoms. In addition, while the total coordination number of Ru (NRu-Ru+NRu-Se) in RuSeMo/C is

slightly higher than that in Ru/C, the coordination number of Se (NSe-Ru+NSe-Se) in RuSeMo/C is

extremely low. This shows that most of the Se atoms are significantly under-coordinated.

Compiling these results together, it can bededucedthat the Se atoms are dispersed at the edges

and/or corners (under-coordinated positions) of the RuSeMo/C nanoparticles. According to the

particle size effect, the under-coordinated metal atoms in the edges/corners are more vulnerable

to oxygen adsorption compared to the atoms in intact facets (such as (111), (100)) due to the

stronger binding energy with oxygen.50Despite this, Se atoms are not oxidized till 0.8 V as

evidenced by the unchanged Se K-edge XANES up to 0.8V (Figure 2.12 (a)). The intensity of

XANES gets slightly reduced at 1.0V, which can be ascribed to the insertion of the oxygen

atoms into the NPs that reduces the Se-Se/Ru scattering. Such tolerance against oxidation can be

explained by looking at the higher electron affinity of Se compared to Ru, making the latter more

vulnerable for oxidation. Therefore, the suppression of the Ru oxidation for the RuSeMo/C

catalyst is partially ascribed to the replacements of the vulnerable Ru atoms in the edges/corners

with oxidation-tolerant Se atoms.

70

Figure 2.11 Δμ signatutes using XANES Ru K- edge for (a) Ru/C (HT) catalyst and (b) RuSeMo/C (HT) catalysts (scale is same as in Figure 7(a) to compare the Δμ amplitudes) (c) plot showing variation of Δμ amplitude as a function of applied potential for Ru/C (HT) and RuSeMo/C (HT) catalysts. Data were collected in Ar- saturated 0.1 M HClO4 electrolytes.

71

Figure 2.12 In situ (a) XANES and (b) FT- EXAFS of RuSeMo/C (HT) catalyst in deaerated Ar saturated 0.1M HClO4 electrolyte at Se K- edge.

The high NSe-Ru given by EXAFS fitting results given in Table 2.4 directly shows the

interaction between the Ru and Se. Given that the electron affinities between the two elements

are drastically different, the ligand effect induced by the charge transfer between them is

expected. Indeed, the white line intensity of Se in the RuSeMo/C catalyst is much less than that

of Se bulk metal (Figure 2.12 (a)), which indicates a substantial charge transfer from Ru to Se as

expected from their electron affinity values. Accordingly, the Ru XANES of the RuSeMo/C

catalyst is also different from that in Ru/C, indicating modifications in the electronic

configuration by the ligand effect, but the difference is less intense compared to Se because Ru is

still mostly surrounded by Ru atoms and the ligand effect is therefore relatively weak. The

deficient charge of Ru caused by the charge transfer from Ru to Se hinders the charge transfer

from Ru to O, thereby suppressing oxygen adsorption. This was also observed by Babu et al. on

Ru/Se systems using EC-NMR and XPS studies.51

In addition to the favorable ligand effect induced by Se, compressive-strain effects

induced by the insertion of Se is also expected given that the radius of Se is much smaller than

that of Ru. Indeed, the Ru-Se bond distance (2.44 Å) in the RuSeMo/C catalyst is much less than

72

the Ru-Ru bond distance (2.69 Å). On the other hand, the Se-Se bond distance (2.52 Å) and Se-

Ru bond distance (2.44 Å) are significantly longer than that of the Se reference foil fitted with

the same parameters (2.36 Å) (Figure 2.13).

Figure 2.13 FT EXAFS for Se reference foil

Table 2.4 In situ EXAFS Fit results for RuSeMo/C – HT catalyst Obtained from Experiments

Performed at the Ru K- edge and Se K- edge as a function of potential in Ar saturated 0.1M

HClO4Electrolytea

aS02 fixed at 0.88 and 0.945 for Ru and Se, respectively as obtained by fitting the reference foils. Fits were

done in R−space, k1,2,3 weighting. For Ru, 1.23< R < 3.00 Å and Δk = 3.153 – 13.954 Å-1were used; for Se, 1.369< R < 2.802 Å and Δk = 2.794 – 13.134 Å-1 were used. Values in parentheses represent the largest statistical errors of all of the least−squares fits determined by ARTEMIS.

According to the d-band theory developed by Norskovet al.,52 the compressive-strain weakens

the Ru-O binding energy via downshift of the Ru d-band relative to the Fermi level, thereby

suppressing the Ru oxidation. Overall, the subdued Ru oxidation can be attributed to the

favorable strain and ligand induced by the decoration of the Se atoms onto the edges/corners of

the RuSeMo/C NPs.It should be noted that the FT-EXAFS of both Se and Ru K-edge data can be

well fitted without including the scattering from the Mo neighbors. In addition, the Mo K-edge

Catalyst Ru/C RuSeMo/C

Parameters N R(Å) N R(Å) Ru-Ru 9.2(1.2) 2.66(3) 7.8(6) 2.69(1) Ru-Se --- --- 1.9(2) 2.44(1) Se-Ru --- --- 2.1(1.0) 2.44(1) Se-Se --- --- 0.5(1.0) 2.52(5)

73

spectra do not change with applied potentials. These results suggest that Mo does not actively

participate in electrocatalysis;this is in congruence with the previous results reported by

Ramaswamy et al.,29in an analogous system studied in alkaline media. Mo mainly contributes in

spilling over peroxide intermediates as shown in the Levich analysis of RuSeMo/C catalyst.

This, however, cannot be directly observed by in situ XAS.

2.3.4 Solid Acid Fuel Cell Testing

Figure 2.14(a) demonstrates the steady state polarization result and corresponding power

density curves obtained using 30% RuSeMo/C and 40% Pt/C, in a solid acid fuel cell. In general,

small leaks across the cell, micro-cracks in the electrolyte and electrolyte resistance mostly

account for the performance losses in a solid acid fuel cell. SAFC electrodes requires high CDP

to catalyst ratio to sustain optimum protonic conductivity, thus introducing additional

performance losses. Moreover, the interaction between CDP and the catalyst plays a vital role in

determining the ohmic losses. CDP lacks the ability to percolate the catalyst as much as

conventional electrolytes such as Nafion®.As a result, both Pt/C and RuSeMo/C catalyst are

subjected to substantial losses in the activation region. A standard state of the art 40% Pt/C

generates 100 mA/cm2 current density at 669 mV iR free, while the chalcogenide catalyst executes

the same current density at 552 mV iR free. The chalcogenide catalyst reaches 67% the power

density of the standard Pt/C cathode run under the same conditions. When subjected to long time

operation, the chalcogenide catalyst exhibits remarkable stability along with a gradual

improvement in the performance from beginning of life (BOL) till 120 hours (Figure 2.15).

74

Metal loading and carbon content play pivotal role in determining the catalyst

performance in a single cell. It is intuitive that a lower carbon content and higher active metal

loading in a catalyst will result in superior performance. This in fact, holds true for Pt/C catalyst.

As shown in Figure 2.14(b), a 10% Pt/C catalyst can be seen to be losing performance rapidly as

compared to 40% Pt/C. Factors such as low platinum content and aggravated carbon corrosion in

10% Pt/C lead to such low-grade performance. However, remarkably different trends are

obtained when chalcogenide catalysts with higher Ru loading are tested. Table 2.5 compares the

performances of 50% RuSeMo/C and 30% RuSeMo/C in RDE experiment in 0.1M HClO4 at

room temperature and under SAFC operating conditions. Under RDE experimental conditions,

50% RuSeMo/C has an ORR onset value of 900 mV, compared to 30% RuSeMo/C with an onset

at 860 mV (Figure 2.6). This enhanced performance arises form a higher loading of the ORR

active Ru metal in 50% RuSeMo/C resulting in more abundant ORR active sites. Compared to

30% RuSeMo/C, 50% RuSeMo/C shows a positive shift of 120 mV in half wave potential, with

the latter having a well-defined limiting current region at 0.6 V vs RHE as shown in Figure 2.6.

However, SAFC polarization results give an inverse trend. While both 30% RuMoSe/C and 50%

RuSeMo/C exhibit a take-off voltage of 820 mV,the latter catalyst dramatically loses

performance and fails to generate current beyond 80 mA/cm2.A higher metal loading may ensure

significant enhancement in performance under RDE experimental conditions, but SAFC

environment is exceedingly different. Other negating factors such as reactivity profile of CDP

may become prominent and show detrimental effect on the catalytic performance. Numerous

metals oxides, including Ru and Mo, tend to react with CDP and form non-conducting phases- a

phenomenon that is prevalent with SAFC operation. Therefore, a higher loading of Ru and Mo in

75

50% RuSeMo/C is most plausible cause for the poor performance of the chalcogenide catalyst,

suggesting that an optimum balance of Ru and Mo is crucial for SAFC application.

Table 2.5 Comparison between RDE experimental results (in 0.1M HClO4 electrolyte)towards ORR

and solid acid fuel cell performance results as obtained from 30% RuSeMo/C and 50% RuSeMo/C

Figure 2.14 iR corrected SAFC test results collected in H2 and Air, under ambient pressure at a cell temperature of 250ºC. (a) Comparison of 40% Pt/C and 30% RuSeMo/C catalysts – polarization curves and corresponding power density curves. Loading of Pt = 1.8 mg/cm2; loading of Ru = 1.2 mg/cm2 (b) Polarization curves obtained from Pt/C and Ru based catalysts with various metal loadings.

As a control experiment, unmodified 40% Ru/C catalyst was tested in SAFC. As

displayed in Figure 2.14(b), the polarization trend is analogous to that of 50% RuSeMo/C,

further indicating that a higher Ruthenium content is detrimental to the catalytic performance in

presence of CDP. The disparate features present in the polarization curve for Ru/C as compared

to a conventional polarization curve are probably due to the spontaneous reaction occurring

between Ru and CDP. Thus, Se modification not only impedes the oxidation of active metal, but

also acts as a protective shield against CDP and arrests any possible reaction with Ru. TGA data

further establishes the importance of Se modification on Ru/C at high temperature (Figure 2.16).

Catalyst

Atomic Ratio

(Ru:Se:Mo)

RDE experiment results SAFC test results

ORR Onset (V)

E ½ (V)

Take off voltage (V)

Voltage @ 200 mA/cm2 (V)

30% RuSeMo/C 1:0.5:0.2 0.86 0.623 0.82 0.479

50% RuSeMo/C 1:0.5:0.4 0.90 0.743 0.76 ---

76

Figure 2.15 Durability test for 30% RuSeMo/C catalyst in SAFC operated at 250oC, with H2 and air as anode and cathode gas feeds, respectively. Loading of Ru in the MEA: 1.2 mg/cm2

Figure 2.16 TGA profiles of Ru/C, RuSe/C and RuSeMo/C catalysts when heated till 250oC under air. (a) Catalytic weight loss/gain when heated till 250ºC and (b) Catalytic weight loss/gain when held at 250ºC for 30 minutes

While Ru/C rapidly loses catalyst weight when held at 250ºC in air, RuSe/C and

RuSeMo/C show no such loss. This happens because in the absence of Se, Ru is subjected to

oxidation, which makes the Ru moieties electron deficit. This results in rapid carbon corrosion,

hence the observed catalytic weight loss. But in presence of Se which has an ideal

electronegativity, Ru oxidation is significantly mitigated, and optimum electron density in Ru is

restored. As a result, no catalytic weight loss occurs in the case of RuSe/C and RuSeMo/C. The

77

slight increase in catalyst weight for the modified catalysts can be due to traces of oxide

formation.

2.4 Conclusion

In this report, we have critically examined the prospects of RuSeMo/C as cathode

catalyst in SAFC, supported by a complete structural and electrochemical characterization.

Combined XRD and XAS results prove that chalcogenide catalyst exists as a Ru core with Se

atoms residing at the periphery within the Ru lattice. In situ XAS experiment validates the

stability of Ru in 0.1M HClO4 electrolyte beyond 0.9V, which is essential for fuel cell

application. Chalcogen modification plays a pivotal role in the success of RuSeMo/C catalyst in

three ways: (i) Se located at the under-coordinated sites suppresses Ru oxidation by inducing

ligand and strain effects. Hence, the ORR active Ru sites are well preserved (ii) Since Se

modification restricts formation of RuO, CDP does not react with RuO to form non- conducting

phases. This helps to sustain the activity of the catalyst over prolonged hours of operation (iii)

Under harsh oxidizing atmosphere at 250ºC, Se protects Ru/C from corrosion, which otherwise

fails to survive and loses catalytic weight due to carbon oxidation, which is explicitly proven by

TGA results. Ternary metal molybdenum mostly acts towards shifting the ORR pathway towards

4 electron route, which is proven by K-L plot analysis. Compared to Pt/C, RuMoSe/C catalyst

shows higher tolerance towards phosphate anion poisoning. Semi quantitative analyses of the

Tafel slopes indicate that phosphate ions alter the platinum surfacedue to strong adsorption, a

phenomenon which is much alleviated in RuMoSe/C. For the first time, SAFC operation using

Ru based catalyst is reported and steady state polarization curves are collected. Although the

activity of the chalcogenide catalyst is inferior to Pt/C, the chalcogenide catalyst displays high

78

stability under continuous operation. Furthermore, Ru is considerably cheaper compared to Pt,

establishing Ru based chalcogenides as impressive candidates for SAFC.

Since SAFC is a recent discovery, there arevoids in the technology thatneed to be

addressed. Further research on the interaction of CDP with catalysts can give crucial information

about predicting catalyst behavior in SAFC. Intensive analytical tools such as in situ Raman

spectroscopy and in situ X- ray crystallographymay be employed to understand the interaction of

super-protonic CDP and catalyst at 250°C degrees, which will eventually help us to build

improvised catalyst models for SAFC.

2.5 Acknowledgement

The authors deeply appreciate the financial support from U.S Department of Energy

funded ARPA–ENERGY (Award number DE-FOA-0001026). Authors would like to thank

Stanford linear Accelerator Center for allowing them to work at the Stanford Synchrotron

Radiation Laboratory and perform the X–ray absorption experiments. Use of synchrotron

facilities weresupported by the US department of Energy, Office of Science, Office of Basic

Energy Sciences, former under the Contract No. DE-SC0012704 and DE-AC02-76SF00515;

respectively. The valuable assistance of Dr. Syed Khalid and Dr. NebojsaMarinkovic(beam 2-2,

SSRL) is gratefully appreciated.

2.6 References

(1) Debe, M. K. Nature2012, 486, 43. (2) Gasteiger, H. A.; Kocha, S. S.; Sompalli, B.; Wagner, F. T. Appl. Catal. B-Environ.2005,

56, 9. (3) Kunz, H. R. J. Electrochem. Soc.1977, 124, C137. (4) Macdonald, J.; Stonehart, P.; Baris, J. J. Electrochem. Soc.1980, 127, C342. (5) McBreen, J.; Olender, H.; Srinivasan, S.; Kordesch, K. V. J. Appl. Electrochem.1981, 11,

787. (6) Xiao, G.; Li, Q. F.; Hjuler, H. A.; Bjerrum, N. J. J. Electrochem. Soc.1995, 142, 2890.

79

(7) He, Q.; Mukerjee, S.; Zeis, R.; Parres-Esclapez, S.; Illan-Gomez, M. J.; Bueno-Lopez, A. Applied Catalysis a-General2010, 381, 54.

(8) Haile, S. M.; Boysen, D. A.; Chisholm, C. R. I.; Merle, R. B. Nature2001, 410, 910. (9) Boysen, D. A.; Uda, T.; Chisholm, C. R. I.; Haile, S. M. Science2004, 303, 68. (10) Boysen, D. A.; Haile, S. M.; Liu, H. J.; Secco, R. A. Chem. Mater.2003, 15, 727. (11) Botez, C. E.; Hermosillo, J. D.; Zhang, J. Z.; Qian, J.; Zhao, Y. S. J. Chem. Phys.2007,

127. (12) Haile, S. M.; Chisholm, C. R.; Sasaki, K.; Boysen, D. A.; Uda, T. Faraday Discuss.2007,

134, 17. (13) Uda, T.; Haile, S. M. Electrochem. Solid State Lett.2005, 8, A245. (14) Kiho, M.; Matsunaga, K.; Morikawa, S.; Kato, O. Electrochemistry2001, 69, 580. (15) Lefevre, M.; Proietti, E.; Jaouen, F.; Dodelet, J. P. Science2009, 324, 71. (16) Proietti, E.; Jaouen, F.; Lefevre, M.; Larouche, N.; Tian, J.; Herranz, J.; Dodelet, J. P.

Nature Communications2011, 2. (17) Jaouen, F.; Proietti, E.; Lefevre, M.; Chenitz, R.; Dodelet, J. P.; Wu, G.; Chung, H. T.;

Johnston, C. M.; Zelenay, P. Energy & Environmental Science2011, 4, 114. (18) Kim, J. H.; Ishihara, A.; Mitsushima, S.; Kamiya, N.; Ota, K. I. Electrochim. Acta2007,

52, 2492. (19) Nam, K. D.; Ishihara, A.; Matsuzawa, K.; Mitsushima, S.; Ota, K.; Matsumoto, M.; Imai,

H. Electrochim. Acta2010, 55, 7290. (20) Maekawa, Y.; Ishihara, A.; Kim, J. H.; Mitsushima, S.; Ota, K. I. Electrochem. Solid

State Lett.2008, 11, B109. (21) Liu, Y.; Ishihara, A.; Mitsushima, S.; Kamiya, N.; Ota, K. Electrochem. Solid State

Lett.2005, 8, A400. (22) Fischer, C.; Alonsovante, N.; Fiechter, S.; Tributsch, H. J. Appl. Electrochem.1995, 25,

1004. (23) Vante, N. A.; Jaegermann, W.; Tributsch, H.; Honle, W.; Yvon, K. J. Am. Chem.

Soc.1987, 109, 3251. (24) Inukai, J. J.; Cao, D. X.; Wieckowski, A.; Chang, K. C.; Menzel, A.; Komanicky, V.;

You, H. Journal of Physical Chemistry C2007, 111, 16889. (25) Li, J.; Ghoshal, S.; Liang, W.; Sougrati, M. T.; Jaouen, F.; Halevi, B.; McKinney, S.;

McCool, G.; Ma, C.; Yuan, X. Energy & Environmental Science2016. (26) Zitolo, A.; Goellner, V.; Armel, V.; Sougrati, M.-T.; Mineva, T.; Stievano, L.; Fonda, E.;

Jaouen, F. Nature materials2015, 14, 937. (27) Ishihara, A.; Lee, K.; Doi, S.; Mitsushima, S.; Kamiya, N.; Hara, M.; Domen, K.;

Fukuda, K.; Ota, K. Electrochem. Solid State Lett.2005, 8, A201. (28) Vante, N. A.; Tributsch, H. Nature1986, 323, 431. (29) Ramaswamy, N.; Allen, R. J.; Mukerjee, S. Journal of Physical Chemistry C2011, 115,

12650. (30) Stolbov, S. Journal of Physical Chemistry C2012, 116, 7173. (31) Ziegelbauer, J. M.; Gulla, A. F.; O'Laoire, C.; Urgeghe, C.; Allen, R. J.; Mukerjee, S.

Electrochim. Acta2007, 52, 6282. (32) Ramaswamy, N.; Mukerjee, S. Journal of Physical Chemistry C2011, 115, 18015. (33) Arruda, T. M.; Shyam, B.; Lawton, J. S.; Ramaswamy, N.; Budil, D. E.; Ramaker, D. E.;

Mukerjee, S. Journal of Physical Chemistry C2010, 114, 1028. (34) Jia, Q.; Liang, W.; Bates, M. K.; Mani, P.; Lee, W.; Mukerjee, S. Acs Nano2015, 9, 387.

80

(35) Newville, M. Journal of Synchrotron Radiation2001, 8, 322. (36) Ravel, B.; Gallagher, K. Phys. Scr.2005, T115, 606. (37) Ankudinov, A. L.; Ravel, B.; Rehr, J. J.; Conradson, S. D. Physical Review B1998, 58,

7565. (38) Jia, Q.; Caldwell, K.; Ziegelbauer, J. M.; Kongkanand, A.; Wagner, F. T.; Mukerjee, S.;

Ramaker, D. E. J. Electrochem. Soc.2014, 161, F1323. (39) Teliska, A.; O'Grady, W. E.; Ramaker, D. E. J. Phys. Chem. B2005, 109, 8076. (40) Ziegelbauer, J. M.; Gatewood, D.; Gulla, A. F.; Guinel, M. J. F.; Ernst, F.; Ramaker, D.

E.; Mukerjee, S. Journal of Physical Chemistry C2009, 113, 6955. (41) Arruda, T. M.; Shyam, B.; Ziegelbauer, J. M.; Mukerjee, S.; Ramaker, D. E. Journal of

Physical Chemistry C2008, 112, 18087. (42) Rao, C. V.; Viswanathan, B. The Journal of Physical Chemistry C2007, 111, 16538. (43) Kelechukwu, B. O.; Israel, O. O.; Emeka, E. O.; Ihebrodike, M. M.; Lebe, N.; Joseph, E.

International Letters of Chemistry, Physics and Astronomy2014, 5, 17. (44) Zaikovskii, V. I.; Nagabhushana, K. S.; Kriventsov, V. V.; Loponov, K. N.;

Cherepanova, S. V.; Kvon, R. I.; Bönnemann, H.; Kochubey, D. I.; Savinova, E. R. The Journal of Physical Chemistry B2006, 110, 6881.

(45) Alonso-Vante, N.; Tributsch, H.; Solorza-Feria, O. Electrochim. Acta1995, 40, 567. (46) Malakhov, I. V.; Nikitenko, S. G.; Savinova, E. R.; Kochubey, D. I.; Alonso-Vante, N. J.

Phys. Chem. B2002, 106, 1670. (47) He, Q. G.; Yang, X. F.; Chen, W.; Mukerjee, S.; Koel, B.; Chen, S. W. PCCP2010, 12,

12544. (48) Park, H.-Y.; Yoo, S. J.; Kim, S. J.; Lee, S.-Y.; Ham, H. C.; Sung, Y.-E.; Kim, S.-K.;

Hwang, S. J.; Kim, H.-J.; Cho, E. Electrochem. Commun.2013, 27, 46. (49) He, Q. G.; Shyam, B.; Nishijima, M.; Ramaker, D.; Mukerjee, S. Journal of Physical

Chemistry C2013, 117, 4877. (50) Stampfl, C.; Schwegmann, S.; Over, H.; Scheffler, M.; Ertl, G. Phys. Rev. Lett.1996, 77,

3371. (51) Babu, P. K.; Lewera, A.; Chung, J. H.; Hunger, R.; Jaegermann, W.; Alonso-Vante, N.;

Wieckowski, A.; Oldfield, E. J. Am. Chem. Soc.2007, 129, 15140. (52) Norskov, J. K.; Bligaard, T.; Rossmeisl, J.; Christensen, C. H. Nat. Chem.2009, 1, 37.

81

Chapter 3

Multifunctional Pt-Nb Electrocatalysts Facilitated by Multiphase Niobium

This Chapter is based on a submitted paper with the same title

ACS Catalysis

Shraboni Ghoshal1, Qingying Jia1, Michael K. Bates1,2, Jingkun Li1, Chunchuan Xu3,Kerrie Gath3, Jun Yang3,James Waldecker3, Haiying Che4, Wentao Liang5,Guangnan Meng6, Zi-Feng Ma4, and Sanjeev Mukerjee1

1Northeastern University for Renewable Energy Technology, Department of Chemistry & Chemical Biology, Northeastern University, Boston, MA 02115, United States

2Present address: Nano Terra, Inc., Cambridge, MA 02138, United States

3Ford Motor Company, Dearborn, MI 48124, United States

4Shanghai Electrochemical Energy Devices Research Center, Department of Chemical Engineering, Shanghai Jiao Tong University, Shanghai 200240, People's Republic of China

5Department of Biology, Northeastern University, Boston, MA 02115, United States

6ULVAC Technologies, Inc., 401 Griffin Brook Drive, Methuen, Massachusetts 01844, United States

82

Chapter 3

Multifunctional Pt-Nb Electrocatalysts Facilitated by Multiphase Niobium

3.1 Introduction

Evolution of electrocatalytic energy conversion processes is crucial to mitigate global

warming and curtail our need of fossil energy. This demands smart design of electrocatalysts that

are economical, durable and highly efficient. More than two decades of research1,2 on

modification of Pt surfaces (alloying, ad-metal, nanoparticle, core shell etc.,) have provided

significant improvements in activity (ORR and HOR) and durability while lowering the Pt

loading. The choice of binary metal is dictated by several factors principle among which is the

extent of solid solution formation (i.e., substitutional effect) and its concomitant effect on the

electronic states of Pt. Extensive studies on Pt alloys such as alloying with first row transition

elements1,3 have clearly alluded to correlation of both electronic (density of states), Pt skin

formation and lattice contraction as reasons for modification of inherent activity of Pt for ORR.

The enhanced activity of ORR has also been explained as a function of lattice strain4, ligand

effect5-7, though depending on the transition metal and its specific segregation potential8,9 various

degree of Pt skin effects were determined resulting in attempts to engineer Pt in terms of various

core shell structures.10,11 It should be noted that in most of these prior works alloying elements

with smaller atomic sizes were used (such as the first-row transition elements). As shown by

several prior reports the resulting alloy therefore exhibits lattice contraction and a concomitant

charge transfer from Pt to the alloying element.12,13 At the anode, composites of Pt where

bifunctional role was rendered primarily using surface specific oxophillic moieties such as Ru,

Mo etc., was invoked in the case of anodic reformate tolerance as well as direct oxidation of

liquid fuels such as methanol etc. In the acidic pH stable presence of oxophillic moieties on the

83

surface of Pt being the premium feature of its electrocatalysis as opposed to those manifest for

ORR (i.e., lattice strain, ligand effect etc.,).

In this report we probe both ORR (acidic pH) and HOR (alkaline pH), both kinetically

challenged in their own aqueous domains with the choice of Nb as an alloying element. Nb

being larger in atomic size compared to Pt has the potential of causing lattice strain in the

opposite direction compared to the wealth of prior Pt alloys where lattice contraction occurs. As

this would constitute charge transfer in the opposite direction (i.e., towards Pt) the concomitant

effect on ORR in acidic and HOR in alkaline pH would be of interest. With properties such as

high oxophilicity14, exceptional stability in acid15,16, wide electrochemical window and

availability of multiple oxidation states17, Nb constitutes an interesting candidate. A vast

literature on Pt-Nb bimetallic systems exists,mostly based on niobium oxide as stable supports

for Pt (Adzicet al.,18,19 and Mitlinet al.,20).Relative few reports have specifically examined the

correlation of Nb used as an alloying element and its role in changing the inherent

electrocatalytic behavior of Pt. Due to strong oxophilicity and high negative reduction potential,

there is a thermodynamic barrier which restricts the formation of zero valent Nb under regular

(< 1000◦C) heat treatment temperatures, and hence formation of alloy between Pt and Nb is

scarce in the literature.21

Considering ORR; in alkaline systems there is nominal overpotential losses due to the

occurrence of an outer- sphere electron transfer which enables the use several non-platinum

catalysts as cathodes.22-24 Meanwhile in acid electrolyte, high performance losses and

requirement of platinum catalysts for ORR brings out challenges pertaining to the cost

effectiveness of the system. Niobium oxide when present as a support in platinum catalysts has

been reported to enhance the ORR catalytic performance in acidic pH despite having no intrinsic

84

catalytic activity of its own.25 For example, the hydrophilic NbxOy materials are known to offer

better platinum mass utilization towards ORR.20 In addition to this, the existence of a synergistic

co-catalytic effect between platinum and niobium oxide support prevents aggregation of

platinum particles during fuel cell operation.26In the case of PtNb alloys, Norskovet al.27 have

theoretically predicted the adsorption strength of oxygen species on Pt3Nb alloy to be in the

favorable range predicting superior intrinsic activity towards ORR as compared to Pt/C.

Nonetheless, this has not been verified experimentally due to the lack of proper platinum

niobium alloys hitherto.

As shown earlier, in alkaline electrolyte HOR kinetics is the challenge,28,29 resulting in an

overvoltage of several hundred millivolts at moderate current densities. Here Pt is not a non-

polarizable electrode in contrast to acidic pH.30,31 The observed sluggish HOR kinetics in

alkaline electrolyte compared to acid electrolyte is explained by two different theories namely (a)

binding energy theory32-34 and (b) reactive OHad species theory.35Theory (a) states that an

increased in M-Had bond strength at higher pH decelerates the HOR rate.36 This theory takes into

account that the Volmer reaction (M-Had + OH-� M+ H2O + e-) is most likely the rate

determining step of HOR,37 thereby putting the onus on M-Had binding strength being the sole

descriptor of the HOR rate. However, the availability of OH- can potentially influence the rate of

the Volmer step as evident from the reaction, this gave rise to the second school of thought.

Theory (b) states that the non-availability of reactive OHad species at the electrical double layer

causes slow alkaline HOR kinetics.35 As a matter of fact, faster HOR kinetics in alkaline

electrolyte is seen when Pt is alloyed with “oxophilic” metals such as Ru and Ni that are believed

to act as the source for furnishing reactive OHad species at the electrode interface.35No prior

HOR study on Pt-Nb system exists, but the aforesaid interesting properties of a Pt-Nb alloy-

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NbOx composite such as oxophilicity and alloy interaction between Pt and Nb are expected to

have conducive effect on the alkaline HOR catalysis.

In this report, we present a carbon supported platinum niobium alloy catalyst

(PtNb/NbOx-C) synthesized by an innovative route. We have strategically designed the

composite to manipulate the base metal Nb in two ways: (a) as Nb oxide to circumvent catalyst

degradation, (b) as an alloying metal for tweaking the electronic and geometric properties of

platinum. With the help of X- ray diffraction (XRD) study, X-ray photoelectron spectroscopy

(XPS) and in situ X- ray absorption spectroscopic (XAS) experiments, we have validated the

structure of the catalyst in detail. In addition to ORR experiments where PtNb/NbOx-C catalyst

exhibited superior performance as compared to Pt/C, we have performed CO adsorption

experiments to provide insight on the role of Nb atoms in terms of electronic and bifunctional

effects. Our PtNb/NbOx-C catalyst have been compared to controls i.e., Pt/NbOx (where Nb is

present in form of oxide support) as well as pure Pt/C to delineate the contributions of binding

energy and reactive OHad effects toward alkaline HOR. To the authors’ knowledge, this is the

first time a Pt-Nb alloy interaction has been established by structure sensitive experiments in a

Pt-Nb electrocatalyst system having multifaceted catalytic properties.

3.2 Experimental Section

3.2.1 Synthesis of PtNb/NbOx-C Catalyst

The PtNb/NbOx-C catalyst was prepared per the scheme shown in this report. In a typical

synthesis, 532 mg of H2Pt(OH)6 was dissolved in 30 mL 6% SO2 solution and stirred under room

temperature for 24 hours to form H3Pt(SO3)2(OH). After 24 hours, a clear yellowish solution was

obtained, that was transferred into a round bottomed flask containing 50 mL water and 750 mg

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Vulcan carbon XC- 72R, resulting in a black slurry. In a separate flask, 200 mg of NbCl5 was

dissolved in 1 mL ethanol to obtain a clear solution. This solution was further treated with 8 mL

0.3M ammonia solution and 10 mL 30% H2O2 solution to obtain peroxoniobic acid.52 This

resulting solution was then transferred to the black slurry via slow, dropwise addition under

controlled pH of 3.5. Once transferred, the slurry was heated at 85ºC for 30 minutes, and

eventually cooled down to room temperature. It was then filtered, dried and heated at 900ºC

under 5% H2/ Ar mixture for 8 hours to get the final catalyst.

3.2.2 Physicochemical Characterization

X- ray diffraction characterization was done using a RigakuUltima IV XRD with Cu Kα

source (lambda= 1.541 Å) operated at 40 kV and 44 mA. 2θ/θ scans were conducted using a 0.05

step size and 5 sec hold per step. Scanning electron microscopy (SEM) characterization was

done using a Hitachi S- 4800 FE- SEM. For validating sample elemental composition, EDAX

Genesis on the same SEM instrument was used. Transmission electron microscopy (TEM) and

selected area electron diffraction (SAED) data were collected using a JEOL 2010 field emission

gun at 200 kV accelerating voltage. The samples were deposited on a holey carbon film on a 300

mesh copper grid.X-ray photoelectron spectroscopy (XPS) data were acquired on a Kratos Axis

Ultra DLD X-ray photoelectron spectrometer usingan Al Ka source monochromatic operating at

150 W. Acquisitiontimes were 4 min for survey spectra, 14 min for Pt4f spectra, and 16 min for

Nb3d spectra. Data analysis and quantificationwere performed using XPSpeak41 software.

3.2.3 Electrochemical Characterization

All electrochemical measurements were done at room temperature using a rotating ring

disc electrode (RDE) from Pine Instruments connected to an Autolab (Ecochemie Inc., model

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PGSTAT 30) potentiostat/ galvanostat. Tests were conducted in a 50 ml jacketed three electrode

cell in 0.1 M HClO4or 0.1M KOH. For each test, reversible hydrogen electrode (RHE), freshly

made, was used as the reference electrode and a Pt mesh wire was used as a counter electrode. A

glassy carbon disc (geometrical area of 0.245 cm2)with a gold ring was used as working

electrode. All the results shown here were collected after conditioning the electrodes at a scan

rate of 50 mV/sec for 30 scans between 0.05 to 1.0 V vs RHE, or until stable features were

achieved.

In CO stripping experiment, CO was purged into the electrolyte for 10 minutes while holding the

potential at 0.05 V vs. RHE. Then argon was purged for 20 minutes to remove any dissolved CO

in the electrolyte and to ensure that CO is present only at the electrode surface adsorbed on the

catalyst. A linear scan was then started from 0.05V at a scan rate of 10 mV/sec.

3.2.4 Durability Measurement

To test the durability of the catalysts, a square wave procedure was followed where the

potential of the working electrode was stepped from 0.1V to 1.0V (vs. RHE) and held for 3

seconds. The potential of the working electrode was then stepped down to 0.1V and held for

another 3 seconds. Thus, each cycle would last for 6 seconds, and data was collected at

beginning of life (BOL), after 12500 cycles and 25000 cycles.

3.2.5 X-Ray Absorption Spectroscopy Measurements

In situ synchrotron XAS studies were conducted at the Pt (L3-edge, 11564 eV)and Nb (K-

edge, 18986 eV) binding energies at National Synchrotron Light Source (NSLS) at Brookhaven

National Laboratories. Detailed description of the spectro-electrochemical cell is given in a

previous publication from our group.53 All the data at the Pt L3-edge and Nb K-edge were

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collected in transmission mode. For transmission data, typically three gas ionization detectors (Io,

It / If and I ref) setup was used with 10% photon absorption in I0 and 50%-70% in It and Iref.

Typical loading of the electrode used in these studies were based on the transmitted x-ray

absorption cross section designed to provide a step height of unity. Nitrogen saturated 0.1 M

HClO4 was used as the electrolyte along with an RHE preparedusing the same electrolyte.

Complete details of EXAFS analysis have been described in detail previously.54 Briefly,

IFEFFIT suite 1.2.9 was used for background subtraction using AUTOBK algorithm and

normalization. Typical K-range window used was 3-14.0 Å-1. The data were processed and

fitted using the Athena55and Artemis56 programs. The χ(R) transforms were modeled using single

scattering paths calculated by the FEFF6 code.57 In addition to the bulk averaged EXAFS

analysis for obtaining short range atomic order around Ru (bond distance, coordination number,

Debye Waller factor etc.) surface specific information was obtained using a previously

developed subtractive method referred to as delta mu (Δμ)58 involving the near edge part of the

spectrum (X-ray absorption near edge structure, XANES), where the effect of the invariant bulk

signal was removed from the surface by subtracting XANES measured at a reference potential

(0.1 V vs. RHE) from other potentials of interest. Data analysis for delta mu (Δμ) studies at Pt

L3-edge X- ray absorption near-edge structure (XANES) spectra involved specific normalization

procedures detailed elsewhere).59 This involved careful calibration of edge energy (Pt L3-edge,

11564 eV) and alignment to the standard reference scan to account for any drift in the beam

energy. A post edge normalization procedure was then applied to the aligned scans via a cubic

spline function, which normalizes the oscillations over a specific energy range (typically 25-200

eV with respect to E0) thus enabling normalized data on a per-atom basis. Difference spectra

were obtained using the equation Δμ = μ(V) - μ (0.54 V), where μ(V) is the XANES at various

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potentials and μ(0.54 V) is the reference signal at 0.54 V, where Pt surface is considered clean

(i.e., free of surface adsorbates). These experimental difference signature profiles were then

compared with theoretically generated profiles constructed using cluster models and the FEFF

8.0 code. Detailed modeling of the clusters in the context of surface adsorbed species was not

necessary in this study as the primary focus was to examine the extent of susceptibility towards

anion adsorption. Prior studies of such nature have been reported earlier.60

3.3 Results and Discussions

3.3.1 Catalyst Synthesis

Scheme 3.1 illustrates the synthesis of the PtNb/NbOx-C catalyst. One of the most critical

factors that determines the state of Nb atoms (as an alloy component or as an oxide support) is

the sequence of deposition of Pt and Nb (in form of oxides) from their respective precursor

solutions. Co-deposition of the metal oxides (Pt and Nb) is crucial to maintain an even

distribution among Pt and Nb atoms, which results in a better alloy formation. H3Pt(SO3)2OH is

a chemically stable solution, and is commonly used as a precursor to obtain PtO2

electrochemically or chemically under oxidizing conditions.38 In this modified method, the use of

H3Pt(SO3)2OH and H2O2 ensures the formation of monodispersed, small particles.39 The peroxo-

niobic acid, along with calculated excess of hydrogen peroxide, is used to decompose

H3Pt(SO3)2OH into PtO2 and while Nb is precipitated as Nb2O5. The final heat treatment at high

temperature under reductive atmosphere is an essential step as it promotes the reduction of PtO2

to Pt, and Nb2O5 to Nb (enabling alloy formation with Pt) and Nb oxides where Nb attains

multiple oxidation states.

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Scheme 3.1 Synthesis of PtNb/NbOx-C catalyst

3.3.2 Physicochemical Characterizations

Figure 3.1(a) shows the Transmission Electron Microscopy (TEM) image of the

synthesized PtNb/NbOx-C catalyst as dispersed on a Cu grid under low and high resolution

(inset) modes. The catalyst particles are found to be homogeneously dispersed with an average

particle size of 4-5 nm. The well-defined lattice fringes (Figure 3.1(a), inset) confirm the high

crystallinity of the supported metal clusters. The Scanning Electron Microscopy (SEM) and

Electron Dispersive X- ray Spectroscopy (EDX) images in Figure 3.1(b) give a general

morphological idea

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of the catalyst. The SEM-EDX elemental mapping images of C, Pt and Nb are juxtaposed on

Figure 3.1(b). They show that Nb and Pt are evenly distributed on the carbon support.

Figure 3.1 (a) High resolution and low resolution TEM of PtNb/NbOx-C catalyst. (b) SEM and elemental mapping results obtained using SEM-EDX (c) SAED image of PtNb/NbOx-C and (d) X ray diffraction pattern of Pt/C (standard) and PtNb/NbOx-C catalyst

The selected area electron diffraction (SAED) pattern in Figure 3.1 (c) shows concentric

rings consisting of discrete diffraction spots, indicating the high crystallinity of the material in

agreement with the TEM results. The lattice spacing of 2.29Å and 1.98Å can be indexed to

{111} and {200} planes of Pt cubic lattice, respectively. Elemental analysis of the catalyst

revealed that the loading of Pt and Nb were 20% and 24%, by weight, respectively. The

representative XRD of PtNb/NbOx-C and Pt/C (as standard) collected from 2θ values of 30º- 90º

are shown in Figure 3.1(d). Based on Pt {111} crystallite line broadening, the Pt crystallite size

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in PtNb/NbOx-C was found to be 3.3 nm.PtNb/NbOx-C exhibits pattern corresponding to the Pt

cubic phase with the peaks shifted towards lower 2θ values. The lattice constant for the

PtNb/NbOx-C was found to be 3.942, whereas for Pt/C, the lattice constant value was 3.912.

Considering the larger atomic radius of Nb (145 pm) compared to that of Pt (135 pm), such shift

in 2θ value is expected per Vegard’s law. This observation in addition to the changed lattice

constant values strongly suggests the formation of a Pt-Nb “alloy” phase which has not been

reported experimentally so far. Since catalysis is a surface phenomenon it is crucial to evaluate

the surface species present in PtNb/NbOx-C catalyst, this is conducted using high resolution XPS

spectra collected at the relevant Pt and Nb binding energy regions (Figure 3.2). Pt 4f spectra

strongly suggests that Pt mostly exists as metal on the surface, along with some traces of PtO. Nb

3d spectra brings out the multiple oxidation states of Nb present on the surface. While Nb(V)

oxide and metallic Nb(0) are the most dominant species detected, traces of Nb(IV) and Nb(II)

oxides can also be identified at the surface. Hence XPS results confirm the presence of multiple

oxidation states of Nb including metallic Nb on the surface. XPS spectra of a niobia supported Pt

catalyst was found to be devoid of any Nb(0) peak (Figure 3.3). XAS experiments (vide infra)

further ascertain the formation of Pt-Nb alloy and gives vital information on the electrochemical

properties of the PtNb/NbOx-C catalyst.

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Figure 3.2 XPS high resolution spectra of PtNb/NbOx-C catalyst for (a) Pt 4f and (b) Nb 3d binding energyregions

Figure 3.3XPS high resolution spectra of NbOx supported Pt catalyst for Nb 3d binding energy region

3.3.3 X- Ray Absorption Spectroscopy

In situ XAS measurements were conducted at both Pt L3 and Nb K edges as a function of

applied potentials to quantitatively determine the local structure of the PtNb/NbOx-C catalyst.

The Nb data remain unchanged within a wide potential range from 0.1 to 1.5 V in both N2 or O2

saturated electrolyte (Figure 3.4), which indicates that Nb does not undergo a redox transition

within the entire fuel cell operating potential range. In situX-ray absorption near edge structure

Figure 3.4In situ XANES of the PtNb/NbOx-C catalyst at the Nb K edge. Data were collected in 0.1 M HClO4 electrolyte in a flow cell

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(XANES) at the Nb K-edge (Figure 3.5(a)) reveals that the bulk-average oxidation state of the

Nb in PtNb/NbOx-C is close to +2 as illustrated by overlapping XANES spectra of aNbO

standard. However, the Fourier Transform of the extended X-ray absorption fine structure

(EXAFS) of the catalyst is drastically different from that of the NbO standard (Figure 3.5(b)).

PtNb/NbOx-C exhibits a prominent peak in the Nb-Nb/Pt scattering region (~2.5 Å, without

phase correction) but the Nb-O peakaround 1.5 Å (without phase correction) has a relatively low

intensity. This suggests that the Nb in the catalysts can be divided into two categories: one

alloyed with Pt, and one present in the form of Nb oxides. The coexistence of the multiple phases

of Nb makes the FT-EXAFS fitting of the Nb data infeasible. On the other hand, the FT-EXAFS

fitting of the Pt edge data can be well fitted, which provides direct evidence of the Pt-Nb

interaction as shown below (Figure 3.6).

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Figure 3.5(a) In situ XANES and (b) FT-EXAFS data of the PtNb/NbOx-C catalyst. The data were collected at the Nb K edge at 0.54 V vs. RHE in N2-saturated 0.1 M HClO4 electrolyte in a flow cell. The XAS data of several standard Nb oxides collected at the same beamline are also included as references.

Figure 3.6In situ FT-EXAFS data and the corresponding fitting results of the (a) PtNb/NbOx-C and (b) Pt/C catalysts. The data were collected at the Pt L3 edge at 0.54 V vs. RHE in N2-saturated 0.1 M HClO4 electrolyte. The fitting parameters are given in Table 1.

The abundance of Pt-Nb bonds in the catalyst is quantitatively evaluated by EXAFS analysis

(Table 3.1) at the Pt L3 edge. By including a Pt-Nb scattering path into the fitting model a

reasonably good fit was achieved, providing coordination number of Pt-Nb (NPt-Nb)as 2.3 which

is significant compared to the fitting uncertainty (0.3) (Table 3.1),In addition, the measured Pt-Pt

and Pt-Nb bond distances in PtNb/NbOx-C were longer than that of Pt/C (Table 3.1), in

agreement with the larger lattice constant given by XRD (Figure 3.1(d)). These results further

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provide conclusive evidence for the alloying of Pt-Nb in the PtNb/NbOx-C catalyst. To the

authors’ best knowledge, this is the first time the Pt-Nb alloying phase has been unambiguously

demonstrated. Since the Nb atoms present in the alloying phase are present in zero oxidation

state, the bulk average oxidation state of +2 (estimated from XANES at the Nb K-edge) points to

the coexistence of Nb0 and NbOx, where x≥1. This has already been demonstrated in the XPS

data where multiple oxidation states of Nb were identified.

Table 3.1Summary of EXAFS fitting results of the data of the PtNb/NbOx-C and Pt/C catalysts collected at the Pt L3 edge at 0.54 V vs. RHE in N2-saturated 0.1 M HClO4 electrolyte. *

Pt-Pt scattering Pt-Nb scattering E0 (eV)

R (Å) N σ2×10-3 (Å2) R (Å) N σ2×10-3 (Å2) PtNb/NbOx-C 2.771(2) 8.1(7) 4.6(5) 2.792(7) 2.3(3) 4.6(5) 7.2(5) Pt/C 2.753(3) 8.4(9) 5.0(5) - - - 7.5(9)

*S02 fixed at 0.84 for Pt as obtained by fitting the reference foil. Fits were done in R-space,

k1,2,3weighting. For Pt, 1.23<R < 3.09 Å and Δk= 3.1 - 13.4 Å-1 were used; for PtNb/NbOx-C, 1.3 <R < 3.06 Å and Δk= 2.539 – 11.646 Å-1 were used. Values in the parentheses indicate the largest statistical error of all the least squares fit determined by ARTEMIS.

Figure 3.7 In situ XANES data of the (a) Pt/C and (b) PtNb/NbOx-C catalysts collected at the Pt L3 edge at various potentials in N2-saturated 0.1 M HClO4 electrolyte. (c) Changes in absorption peaks μ(V) – μ(0.54 V) derived from (a) and (b), as a function of applied potentials.

In situ XAS measurements show that the Pt white line intensities for both PtNb/NbOx-C and Pt/C

increase with increasing potentials (Figure 3.7). This trend is caused by the charge transfer from

the Pt surface to the oxygenated adsorbates, and thus reflects the progressive accumulation of the

oxygenated adsorbates onto the Pt surface with increasing potentials. Assuming the Pt surface is

free of adsorbates at 0.54 V (the double layer region), the relative oxygen coverage at the

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potential V can be represented by the subtractive Δμ(V)=μ(V)-μ(0.54V).40This subtractive

technique referred to as the ∆µ method involved subtraction of XANES spectra measured under

controlled conditions such as 0.54 (Ar purged) where minimum surface adsorbates are expected

minus spectra where surface adsorbed species were expected. The comparison of the Δμ(V)

between PtNb/NbOx-C and Pt/C as displayed in Figure 3.7(c) clearly shows the late onset of

adsorption of oxygenated speciesvia water activation on PtNb/NbOx relative to Pt/C, this is

closely related to an enhanced ORR activity as illustratedin the next section. The suppression of

the formation of surface adsorbed oxygenated species cannot be ascribed to the compressive-

strain effect as the Pt-Pt bond distance in PtNb/NbOx-C is largeras compared to Pt/C, and the

resultant tensile-strain presumably facilitates oxygen adsorption by strengthening the Pt-O

binding energy.41,42It should be noted, Norskovet al.27 have predicted that the Pt-O binding

energy of the Pt3Nb alloy is weaker than that of pure Pt, which therefore, has to be attributed to

the ligand effect induced by Nb. It is important to emphasize that there is a key difference

between the Pt-Nb system and the widely studied Pt alloys with first row transition series such as

the Pt-(Co/Ni)catalytic system. Surface Co and Ni undergo rapid dissolution in acid and are

selectively leached out from the catalyst particles in acid media. As a result, the alloys eventually

form a core-shell structure with pure Pt overlayers wherein the ligand effect disappears or are

severely attenuated.43,44 On the contrary, Nb is insoluble in acid, which ensures that the ligand

effect induced by Nbstays germane during the entire period of operation. Our results provide the

first experimental proof for superior ORR activity of Pt as a result of the ligand effect of Nb, a

fact which was predicted theoretically.27It should be noted thatsuch suppressed adsorption of

oxygenated species was also observed on the Pt/NbOx/C system wherein the Pt was supported

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onto Nb which was ascribed active site support synergy and the ability of Nb to getter oxygen in

preference to adjacent Pt.5,18

3.3.4 Oxygen Reduction Reaction in acid on PtNb/NbOx-C catalyst

We benchmarked the electrocatalytic properties of the PtNb/NbOx-C catalyst using a

commercial Pt/C catalyst (40% by weight, 3 nm Pt particles supported on Vulcan XC-72R). Pt/C

exhibitsa typical cyclic voltammogram (CV)with distinct HUPD adsorption/ desorption peaks (Pt+

H3O+↔Pt-Had + H2O) between 0.05V and 0.35V (Figure 3.8(a); inset). The same features are

also present in the CV of PtNb/NbOx-C, though slightly diminished. The electrochemical surface

area (ECSA) derived from the HUPDwere found to be 45m2/gPt for PtNb/NbOx-C, as compared to

65 m2/gPt for Pt/C. The lower ECSA of the PtNb/NbOx-C catalyst is most likely due to presence

of catalyst agglomerates caused by the heat treatment during the synthesis in addition to the

presence of niobium species on surface. The formation of Pt-OHad starts at ~0.6V in Pt/C catalyst

because of water activation (2H2O +Pt(M) → Pt(M)OHad + H3O+ + e-). As shown earlier45 this

reaction impedes ORR by blocking ORR active sites, and is one of the major reasons of ORR

activity loss in Pt based catalysts. This process, however, is delayed in the case of PtNb/NbOx-C

as suggested by the anodic shift of the redox peak in the CV (Figure 3.8(a), inset) indicating

suppressed oxygen coverage (see Δµ(v) as shown in Figure 4(c). As expected from the

diminished site-blocking effect, PtNb/NbOx-C exhibits superior ORR activity compared to Pt/C

as shown in Figure 3.8(a). Mass transport corrected Tafel plots for PtNb/NbOx-C and Pt/C

(Figure 3.8(b)) and indicate no change in ORR mechanism in the two catalysts. The total number

of electrons transferred during ORR was determined from the K-L plots (inset of Figure 3.8(b)),

and was found to be ~4 for both Pt/C and PtNb/NbOx-C. The higher ORR activity of

PtNb/NbOx-C despite having a lower ECSA can be attributed to higher specific activity enabled

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by the ligand effect originating from the Pt-Nb alloy formation (as predicted by Norskov27 based

on a computational screening in 2009).

Figure 3.8 (a) ORR polarization curves for Pt/C and PtNb/NbOx-C recorded under room temperature in O2- saturated 0.1M HClO4 at 1600 rpm at 20 mV/sec and (inset) CV of Pt/C and PtNb/NbOx-C in Ar- saturated 0.1M HClO4 solution collected at sweet of 20 mV/sec (b) Mass transport corrected Tafel plots of Pt/C and PtNb/NbOx-C catalyst using the results in (a) and (inset) K-L plots of Pt/C and PtNb/NbOx-C (c) ORR polarization curves of PtNb/NbOx-C collected under room temperature in O2 saturated 0.1M HClO4 with various amount of H3PO4 doped into the electrolyte (d) Comparison of mass activity between Pt/C and PtNb/NbOx-C with cycling in Ar- saturated 0.1M HClO4 solution

As mentioned earlier, Nb is known to have exceptional stability in acids and its oxides tend to

influence the electronic features of Pt metal via strong metal- metal oxides interaction. Interplay

of various effects, such as electrostatic phenomena (work function shift), band structure effects

(d band center shift) and conductivity (density of states at Fermi level) are expected to affect the

activity of Pt in the presence of Nb oxides.46 These facts largely work towards increasing the

electrochemical stability of the noble metal catalyst and making it more durable and tolerant

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Figure 3.9ORR polarization curves of Pt/C collected under room temperature in O2 saturated 0.1M HClO4 with various amount of H3PO4 doped into the electrolyte

towards anion poisoning and catalyst degradation. To test this hypothesis, PtNb/NbOx-C was

subjected to anion poisoning by introducing calculated amounts of phosphoric acid as dopants to

O2 saturated 0.1M HClO4 solution (Figure 3.8(c)). With an increasing dopant concentration, the

half wave potential shifts cathodically, resulting in a 55mV shift from 0 mM H3PO4 to 100 mM

H3PO4, whereas a net cathodic shift of 110 mV was observed for Pt/C under the same

experimental conditions (Figure 3.9). Hence, modifying the Pt structure with Nb mitigates

phosphate anion poisoning to a substantial extent. This trend can be envisioned in terms of the

surface niobium species that impart electrostatic repulsion to phosphate anions, and alleviates the

poisoning of Pt active sites by the anions. Furthermore, the presence of surface Nb species

significantly dilutes the fcc sites of Pt (one surface Nb atom can at most destroy four Pt fcc

sites), and in turn inhibits attachment of PO43- that selectively binds to the Pt fcc sites. A

somewhat similar effect has been reported by us earlier using a PtNi alloy.47

Figure 3.8(d) compares the results of durability measurement obtained from RDE experiment in

O2 saturated 0.1M HClO4 for Pt/C and PtNb/NbOx-C. For the testing, the catalysts were

subjected to a square wave process between at 0.1V and 1.0V, while holding at each potential for

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3 seconds. ORR data were collected at beginning of life (200 cycles), after 12500 cycles and

25000 cycles. As expected, the presence of Nb increases the endurance of the Pt catalyst.

PtNb/NbOx-C, which has a twofold higher mass activity as compared to Pt/C, retains the activity

with a loss of 12.5% after 12500 cycles and 17.5% after 25000 cycles. However, Pt/C lost 15%

activity after 12500 cycles and 34% of activity after 25000 cycles. Such high tolerance towards

catalyst degradation can be related to the presence of NbOx moieties. Also, under high

operational voltage and oxidative environment, carbon is vulnerable to oxidation especially in

presence of Pt that can destroy the integrity of the catalyst frame. Due to a stronger interaction

between Pt and NbOx, the mobility of the Pt atoms is lower on NbOx as compared to C, which in

turn reduces the probability of catalyst degradation and particle agglomeration. It is also

plausible that an electronic effect due to the existence of the Nb oxide dispenses extra stability to

the platinum catalyst.

3.3.5 CO Stripping Experiments on PtNb/NbOx-C catalyst

CO stripping reaction can lead to a deeper insight into the surface properties of an

electrocatalyst, which is especially relevant in terms of investigating Pt-M bimetallic surfaces.

This commonly studied reaction on Pt-M surfaces (M= binary metal) can be represented as:

CO+ Ptsurf →PtsurfCO

Msurf + H2O → MsurfOHad + H++ e-

MsurfOHad + PtsurfCO → Msurf + Ptsurf + CO2 + H+ + e-

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Figure 3.10CO stripping data for (a) Pt/C, and (b) PtNb/NbOx-C in 0.1M HClO4 (dotted) and 0.1M KOH (dashed) collected at 10 mV/sec sweep rate. Loading of Pt = 20 µg/cm2.

In Figure 3.10, we have compared the CO stripping profiles of Pt/C and PtNb/NbOx-C catalysts

under acidic and alkaline environments. In acid electrolyte, the potential of zero charge (PZC) of

Pt/C is ~0.63V, beyond which the Pt surface begins to develop positive charge promoting

adsorption of OHads species on Pt surface as indicated in Figure 3.8(a) inset. This phenomenon,

in fact, matches with the onset of CO oxidation on Pt/C, and is accompanied by a small pre-peak

at 0.72V (Figure 3.10(a)). This pre-peak at 0.72V has been explained in terms of a sudden rise in

coverage of OHad species formed via activation of water, which in turn, expedites CO

oxidation.48 At 0.8V, the OHad current reaches a peak as a consequence of increase in coverage

and stabilization of OHad on the Pt surface.49 As expected, the CO oxidation peak on Pt/C is

therefore, located around 0.8V as well. However, in PtNb/NbOx-C catalyst, the adsorption of

OHad species is significantly delayed as a consequence of shift of water activation to higher

potentials (Figure 3.8(a)) and thus the onset of CO oxidation is shifted anodically to 0.72V

(Figure 3.10(b)). Therefore, CO stripping peak on PtNb/NbOx-C is centered at 0.8V and is

devoid of any pre-peak unlike Pt/C.

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In alkaline electrolyte, the electrode- electrolyte interface is dramatically changed. Formation of

OHads on pure Pt surface in alkaline media takes place as per the equation:

Pt-OH2∙∙∙∙∙∙∙ OH-(HOH)2 ↔Pt-OH∙∙∙∙∙∙H-OH(HOH)2 + e- (i)

The reversible potential associated with equation (i) is ~0.17V, indicating that it is

thermodynamically possible for OHad to adsorb on Pt at such low potentials.50 This phenomenon

is facilitated by the presence of Had on Pt at the lower potentials, that impart ∂+ charge to Pt sites

allowing H2O molecules to bond on Pt surface by lone pair donation. The adsorbed water

molecules then undergo reorientation and transfer of electrons and subsequently form Pt-OHad

that are primarily responsible for the electro-oxidation of CO in alkaline electrolyte at lower

potentials (~0.3V), as displayed in Figure 3.10(a). Beyond 0.6V, the Pt surface is strongly

adsorbed by OH- anions, that aid in oxidation of adsorbed CO, thereby explaining the sharp peak

at 0.68V.

From the above discussions, it can be inferred that in alkaline media, strongly adsorbed H atoms

on Pt surface play a major role in facilitating electro-oxidation of CO at lower potentials.

Interestingly, PtNb/NbOx-C, does not exhibit CO stripping onset till 0.4V and features a single,

sharper peak at 0.7V which corresponds to the adsorption of OH- anions on the surface (Figure

3.10(b). This indicates that at the lower potentials, Hads may be weakly attached on the Pt surface

in case of PtNb/NbOx-C catalyst, resulting in the anodic shift in the buildup of ∂+ charge on Pt.

As mentioned above in acidic pH, suppression of water activation provides for lower reactive

OHadspecies thereby lower site blocking in the context of ORR providing a beneficial role. CO

stripping difference at this pH between Pt/C and PtNb/NbOx-C is therefore consistent with this

notion. In alkaline pH, however the transition to potentials positive to zero charge brings

serious competition with OH- species. The CO stripping behavior of PtNb/NbOx-C catalyst in

104

alkaline media brings up several questions regarding the Hads- Pt interface in the low potential

region that can be resolved using alkaline HOR experiments. Moreover, considering the unique

structure of our catalyst which has both metallic Nb alloyed with Pt as well as Nb oxides as

support, it would be interesting to study HOR in light of debates surrounding the mechanism of

alkaline HOR (vide infra).

Figure 3.11(a)Cyclic voltammogram of Pt/C and PtNb/NbOx-C catalysts in deaerated 0.1M KOH at10 mV/sec sweep rate and (b) HOR polarization curves collected for Pt/C and PtNb/NbOx-C in H2- saturated 0.1M KOH; (inset) mass transport corrected kinetic current densities of Pt/C and PtNb/NbOx-C

3.3.6 Alkaline Hydrogen Oxidation Reaction on PtNb/NbOx-C catalyst

Figure 3.11(a) compares the CV of Pt/C and PtNb/NbOx-C in 0.1M KOH electrolyte. Per

the binding energy theory, a difference in Pt-H binding energy is reflected in the position of Pt-

HUPD desorption (M-Had� M+ H+ + e-) peak. However, Pt/C and PtNb/NbOx-C catalysts exhibit

the HUPD desorption peak at the same potentials. While hydrogen binding energy theory

emphasizes on the effect of metal- adsorbate bond strength, reactive OHad theory relates the

kinetics of alkaline HOR to the availability of reactive OHad species.34 To test the validity and

the contribution of these two theories in our system, we collected HOR polarization curves in

alkali. Figure 3.11(b) shows the alkaline HOR activity of PtNb/NbOx-C and Pt/C in 0.1M KOH.

The inset of Figure 3.11(b) shows the mass transport corrected kinetic current densities of Pt/C

105

and PtNb/NbOx-C. The exchange current densities extracted from this plot clearly validates that

HOR kinetics on PtNb/NbOx-C is significantly faster than on Pt/C. The quantitative comparison

of the HOR performance of the two catalysts are illustrated in Table 3.2. At an overpotential of

50 mV, the specific activity of PtNb/NbOx-C is 4 mA/cm2 compared to Pt/C, which executes a

specific activity of only 0.5mA/cm2(Figure3.12). Nb is a non–noble transition metal that is

essentially

Table 3.2Quantitative comparison of Pt/C and PtNb/NbOx-C catalysts in terms of ECSA, specific activities (js) and exchange current densities (i0) in 0.1M KOH toward alkaline HOR at 1600 rpm *

Catalyst ECSA (m2/gPt) jo(mAcm-2Pt) js @ 50mV(mAcm-2

Pt) Pt/C 65 0.56 .5

PtNb/NbOx-C 45 0.8 4

*The ECSA values were calculated from HUPDdesorption area in 0.1M HClO4 electrolyte

Figure 3.12Mass transport andiR corrected specific activities of Pt/C and PtNb/NbOx-C in H2 saturated 0.1M KOH collected at a scan rate of 10 mVsec-1 at 1600 rpm

passivated with an oxide or hydroxide layer at alkaline electrochemical environment. As per

reactive OHad theory, such enhanced HOR activity of PtNb/NbOx-C is due to the presence of

surface oxides/ hydroxides on Nb, which facilitate the extraction of hydrides from platinum

surface to form water. It is known that the Nb surface is passivated and forms oxy-hydroxides at

HOR potentials.51 Following this hypothesis, it can be envisaged that the alloy formation

106

between Pt and Nb is not mandatory for better HOR activity; rather a close proximity of Pt and

Nb may just be adequate to improve the HOR kinetics. To corroborate this, we further tested the

HOR activity of some NbOx- supported Pt catalysts where the interaction between Pt and Nb

does not constitute a strong metal-metal alloy

Figure 3.13 HOR polarization curves collected for Pt/C and NbOx supported Pt/C catalyst in H2- saturated 0.1M KOH at 1600 rpm at a scan rate of 10 mVsec-1

interaction as compared to Pt-Nb alloys (PtNb/NbOx-C) (Figure 3.13). As shown in figure 3.13,

the Niobia supported Pt catalysts are much less active as compared to PtNb/NbOx-C. Hence, the

reactive OH species theory does not seem to be decisive in appraising the HOR activity of Pt- Nb

systems in alkaline conditions. The fact that there is an alloy formation in PtNb/NbOx-C catalyst,

leads us to consider the electronic changes in Pt to be a strong contributor via engendering the

weakening of the binding of hydrogen to the Pt surface. This however, may not be discernible in

the CV (Figure 3.8(a)). These results in conjunction to the CO stripping results obtained in the

previous section, strongly imply that binding energy of hydrogen on Pt is altered in PtNb/NbOx-

C catalyst that in turn, influences the mechanism of alkaline HOR. However, we also believe that

neither reactive OHad theory nor HBE theory are sole descriptors of the activity of PtNb/NbOx-C

catalyst towards HORalkali; rather it is a combination of both that determines the catalytic profile

of PtNb/NbOx-C.

107

3.4 Conclusion

In this report, we have assessed the catalytic profile of a platinum niobium alloy

supported on NbOx and C which was synthesized in house. XRD and XAS experiments confirm

the formation of alloy between Pt and Nb. XPS data further affirm the presence of metallic Nb

and Nb oxides on the surface of the catalyst as opposed to all oxides in the case of Pt supported

on NbOx. In situ XAS results illustrate the effect of Nb moieties on the electrochemical behavior

of Pt. These results prove that Nb suppresses the adsorption of OH on Pt due to strong ligand

effect, and therefore enhances its catalytic activity toward ORR. While the improved ORR

activity and durability of PtNb/NbOx-C catalyst is attributed to coalition of different interactions

between Pt, Nb, NbOx and C, the explanation of the superior HORalkali activity seems to be more

intricate. By studying the CO stripping profiles of Pt/C and PtNb/NbOx-C catalysts, we have

tried to deduce the nature of Nb in PtNb/NbOx-C catalyst under operating electrochemical

environment and have extended it to perceive the rationale behind the superior HORalkali

performance. We believe the enhancement in HORalkali activity occurs because of interplay of

both modified electronic structure (because of Pt-Nb interactions) and presence of reactive OH

species (supplied by oxophilic Nb moieties). However, it is still unclear whether any of these two

factors have dominance towards the observed HORalkali catalytic performance.

This work is an effort to encompass the wide spectrum of catalytic activity offered by platinum

surface modified by multi-phase niobium species, and to reach a consensus on the ongoing

debate regarding HORalkali activity of Pt alloys. To the authors knowledge, this is the first

reported work on Pt-Nb alloys, an interaction that has been mostly restricted to theoretical works.

PtNb/NbOx-C catalyst is a unique material that contains Nb in form of oxides and metal. Such a

108

singular frame aids in bringing the catalytic activity of Pt-Nb systems to the forefront for

application in both acidic and alkaline energy conversion devices.

3.5 Acknowledgement

The authors thank Ford Motor Co. for a University Research Program Award.Use of the

synchrotron facilities at the National Synchrotron Light Source, beamlines X19A and X3B at

Brookhaven National Laboratory, Upton, NY is supported by the U.S. Department of Energy,

Office of Science, Office of Basic Energy Sciences, former under Contract No. DE-AC02-

98CH10886.Authors would like to thank Brookhaven National laboratories for allowing to work

at National Synchrotron Light Source- I at beamlines X19A and X3B.This work was funded by

US department of Energy- Office of Basic Energy Sciences. The valuable assistance of Dr. Eric

Farquhar and Dr. Syed Khalid is gratefully appreciated.

3.6 References

(1) Mukerjee, S.; Srinivasan, S.; Soriaga, M.; McBreen, J. Journal of Physical Chemistry1995, 99.

(2) Paulus, U.; Wokaun, A.; Scherer, G.; Schmidt, T.; Stamenkovic, V.; Radmilovic, V.; Markovic, N.; Ross, P. The Journal of Physical Chemistry B2002, 106, 4181.

(3) Mukerjee, S.; Srinivasan, S.; Soriaga, M. P.; McBreen, J. Journal of the Electrochemical Society1995, 142, 1409.

(4) Strasser, P.; Koh, S.; Anniyev, T.; Greeley, J.; More, K.; Yu, C.; Liu, Z.; Kaya, S.; Nordlund, D.; Ogasawara, H. Nature chemistry2010, 2, 454.

(5) Scott, F. J.; Mukerjee, S.; Ramaker, D. E. The Journal of Physical Chemistry C2009, 114, 442.

(6) Kitchin, J. R.; Nørskov, J. K.; Barteau, M. A.; Chen, J. Physical review letters2004, 93, 156801.

(7) Bligaard, T.; Nørskov, J. K. Electrochimica Acta2007, 52, 5512. (8) Kitchin, J.; Nørskov, J. K.; Barteau, M.; Chen, J. The Journal of chemical physics2004,

120, 10240. (9) Ruban, A.; Skriver, H. L.; Nørskov, J. K. Physical Review B1999, 59, 15990. (10) Wang, J. X.; Inada, H.; Wu, L.; Zhu, Y.; Choi, Y.; Liu, P.; Zhou, W.-P.; Adzic, R. R.

Journal of the American Chemical Society2009, 131, 17298. (11) Zhang, J.; Lima, F.; Shao, M.; Sasaki, K.; Wang, J.; Hanson, J.; Adzic, R. The Journal of

physical chemistry B2005, 109, 22701.

109

(12) Toda, T.; Igarashi, H.; Uchida, H.; Watanabe, M. Journal of The Electrochemical Society1999, 146, 3750.

(13) Mukerjee, S. Journal of Applied Electrochemistry1990, 20, 537. (14) Kepp, K. P. Inorganic Chemistry2016, 55, 9461. (15) Tanabe, K. Catalysis Today2003, 78, 65. (16) Reichman, B.; Bard, A. J. J. Electrochem. Soc1980, 127, 241. (17) Özer, N.; Rubin, M. D.; Lampert, C. M. Solar Energy Materials and Solar Cells1996, 40,

285. (18) Sasaki, K.; Zhang, L.; Adzic, R. Physical Chemistry Chemical Physics2008, 10, 159. (19) Sasaki, K.; Adzic, R. Journal of the Electrochemical Society2008, 155, B180. (20) Zhang, L.; Wang, L.; Holt, C. M.; Navessin, T.; Malek, K.; Eikerling, M. H.; Mitlin, D.

The Journal of Physical Chemistry C2010, 114, 16463. (21) Ziolek, M. Catalysis Today2003, 78, 47. (22) Ramaswamy, N.; Mukerjee, S. Advances in Physical Chemistry2012, 2012. (23) Ramaswamy, N.; Tylus, U.; Jia, Q.; Mukerjee, S. Journal of the American Chemical

Society2013, 135, 15443. (24) Yang, W.; Fellinger, T.-P.; Antonietti, M. Journal of the American Chemical

Society2010, 133, 206. (25) Blizanac, B. B.; Pylypenko, S.; Olson, T. S.; Konopka, D.; Atanassov, P. Journal of The

Electrochemical Society2011, 158, B485. (26) Chun, H.-J.; Kim, D. B.; Lim, D.-H.; Lee, W.-D.; Lee, H.-I. International Journal of

Hydrogen Energy2010, 35, 6399. (27) Greeley, J.; Stephens, I.; Bondarenko, A.; Johansson, T. P.; Hansen, H. A.; Jaramillo, T.;

Rossmeisl, J.; Chorkendorff, I.; Nørskov, J. K. Nature chemistry2009, 1, 552. (28) Sheng, W.; Gasteiger, H. A.; Shao-Horn, Y. Journal of The Electrochemical Society2010,

157, B1529. (29) Sheng, W.; Myint, M.; Chen, J. G.; Yan, Y. Energy & Environmental Science2013, 6,

1509. (30) Conway, B.; Tilak, B. Electrochimica Acta2002, 47, 3571. (31) Schmidt, T.; Ross, P.; Markovic, N. Journal of Electroanalytical Chemistry2002, 524,

252. (32) Durst, J.; Siebel, A.; Simon, C.; Hasche, F.; Herranz, J.; Gasteiger, H. Energy &

Environmental Science2014, 7, 2255. (33) Sheng, W.; Zhuang, Z.; Gao, M.; Zheng, J.; Chen, J. G.; Yan, Y. Nature

communications2015, 6. (34) Wang, Y.; Wang, G.; Li, G.; Huang, B.; Pan, J.; Liu, Q.; Han, J.; Xiao, L.; Lu, J.;

Zhuang, L. Energy & Environmental Science2015, 8, 177. (35) Strmcnik, D.; Uchimura, M.; Wang, C.; Subbaraman, R.; Danilovic, N.; van der Vliet,

D.; Paulikas, A. P.; Stamenkovic, V. R.; Markovic, N. M. Nature chemistry2013, 5, 300. (36) Schwämmlein, J. N.; El-Sayed, H. A.; Stühmeier, B. M.; Wagenbauer, K. F.; Dietz, H.;

Gasteiger, H. A. ECS Transactions2016, 75, 971. (37) St. John, S.; Atkinson III, R. W.; Unocic, R. R.; Zawodzinski Jr, T. A.; Papandrew, A. B.

The Journal of Physical Chemistry C2015, 119, 13481. (38) Thompson, S.; Jordan, L.; Shukla, A.; Forsyth, M. Journal of electroanalytical

chemistry2001, 515, 61. (39) Stonehart, P. Berichte der Bunsengesellschaft für physikalische Chemie1990, 94, 913.

110

(40) Arruda, T. M.; Shyam, B.; Ziegelbauer, J. M.; Mukerjee, S.; Ramaker, D. E. The Journal of Physical Chemistry C2008, 112, 18087.

(41) Greeley, J.; Mavrikakis, M. Nature materials2004, 3, 810. (42) Kitchin, J.; Nørskov, J. K.; Barteau, M.; Chen, J. Journal of Chemical Physics2004, 120,

10240. (43) Bing, Y.; Liu, H.; Zhang, L.; Ghosh, D.; Zhang, J. Chemical Society Reviews2010, 39,

2184. (44) Jia, Q.; Caldwell, K.; Strickland, K.; Ziegelbauer, J. M.; Liu, Z.; Yu, Z.; Ramaker, D. E.;

Mukerjee, S. ACS catalysis2014, 5, 176. (45) Teliska, M.; Murthi, V. S.; Mukerjee, S.; Ramaker, D. E. Journal of the Electrochemical

Society2005, 152, A2159. (46) Zhang, L.; Wang, L.; Holt, C. M.; Zahiri, B.; Li, Z.; Malek, K.; Navessin, T.; Eikerling,

M. H.; Mitlin, D. Energy & Environmental Science2012, 5, 6156. (47) He, Q.; Shyam, B.; Nishijima, M.; Ramaker, D.; Mukerjee, S. The Journal of Physical

Chemistry C2013, 117, 4877. (48) Marković, N.; Schmidt, T.; Grgur, B.; Gasteiger, H.; Behm, R.; Ross, P. The Journal of

Physical Chemistry B1999, 103, 8568. (49) Michaelides, A.; Hu, P. The Journal of Chemical Physics2001, 114, 513. (50) Anderson, A. B. Electrochimica acta2002, 47, 3759. (51) Robin, A. Electrochimica acta2004, 49, 1915. (52) Uekawa, N.; Kudo, T.; Mori, F.; Wu, Y. J.; Kakegawa, K. Journal of Colloid and

Interface science2003, 264, 378. (53) Arruda, T. M.; Shyam, B.; Lawton, J. S.; Ramaswamy, N.; Budil, D. E.; Ramaker, D. E.;

Mukerjee, S. The Journal of Physical Chemistry C2009, 114, 1028. (54) Jia, Q.; Liang, W.; Bates, M. K.; Mani, P.; Lee, W.; Mukerjee, S. ACS nano2015, 9, 387. (55) Newville, M. Journal of synchrotron radiation2001, 8, 322. (56) Ravel, B.; Gallagher, K. Physica Scripta2005, 2005, 606. (57) Ankudinov, A.; Ravel, B.; Rehr, J.; Conradson, S. Physical Review B1998, 58, 7565. (58) Jia, Q.; Caldwell, K.; Ziegelbauer, J. M.; Kongkanand, A.; Wagner, F. T.; Mukerjee, S.;

Ramaker, D. E. Journal of The Electrochemical Society2014, 161, F1323. (59) Arruda, T.; Shyam, B.; Ziegelbauer, J.; Ramaker, D.; Mukerjee, S. ECS

Transactions2007, 11, 903. (60) Ziegelbauer, J. M.; Gatewood, D.; Gullá, A. F.; Guinel, M. J.-F.; Ernst, F.; Ramaker, D.

E.; Mukerjee, S. The Journal of Physical Chemistry C2009, 113, 6955.

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Chapter 4

Investigation into the Alkaline Hydrogen Oxidation Reaction Mechanism for

Bimetallic Platinum Composites

112

Chapter 4

Investigation into the Alkaline Hydrogen Oxidation Reaction Mechanism for

Bimetallic Platinum Composites

4.1 Introduction

To address the twin challenges of energy crisis and climate change, development of zero

emission energy conversion devices is a potential solution. Proton exchange membrane fuel cells

(PEMFCs) and anion exchange membrane fuel cells (AEMFCs) are such devices that operate in

acidic and alkaline environments.1 PEMFCs face major challenge in power market owing to their

dependence on platinum for electrocatalysis of oxygen reduction reaction (ORR).2 With the

recent advancements in alkaline exchange membranes technology, AEMFCs have emerged as

strong contenders to PEMFCs, especially given the fact that AEMFCs can operate with non-

precious metal catalysts for the electrochemical reduction of oxygen at the cathode side of the

device.3 Additionally, the selective hydroxide ion conductivities of these alkaline exchange

membranes are reported to be in the range of 20- 40 mS cm-1, which provides the avenue for

commercialization of the AEMFCs to a fuller extent.4 However, a key question that remains

unanswered is regarding the electrocatalysis of hydrogen oxidation reaction (HOR) in alkaline

electrolyte.5 HOR suffers from slow kinetics and can be characterized by a reduction in the

exchange current density up to two orders of magnitude as compared to HOR in acidic media.5

Such sluggish kinetics leads to severe performance losses in AEMFCs, and a demands higher

loading of platinum at the anode side.6

To understand the origin of such sluggish kinetics, the mechanism of HOR in both acid

and alkaline electrolytes need to be investigated. In acid, HOR involves dissociative adsorption

of hydrogen on Pt surface (Tafel step: ½ H2 + Pt→ Pt-Had) which is then followed by the Volmer

113

step (Pt-Had → Pt + H+ + e-) involving direct oxidation of adsorbed hydrogen. Though Tafel step

remains unchanged in alkaline media, the Volmer step is modified significantly since the

adsorbed hydrogen reacts with hydroxide ion to form water. Thus, the Volmer step in alkaline

media can be represented as: Pt-Had + OH- → Pt + H2O + e-. Volmer step is the rate determining

step and hence, the overall kinetics of the alkaline HOR is dependent on the Volmer step.7 There

is an ongoing debate on the origin of slow Volmer step in alkaline HOR and two theories exist

that address the question in different ways. The first theory, named the binding energy theory

states that a stronger binding of hydrogen to the platinum surface in alkaline media leads to a

slow Volmer step.8 Whereas, proponents of the other theory, called reactive-OHads theory believe

that the unavailability of OH- species at the electrode- electrolyte interface gives rise to a slower

Volmer reaction.9 This debate is extended to explain HOR in bimetallic system of Pt-Ru alloys

where the alkaline HOR kinetics are significantly superior as compared to Pt/C. Binding energy

theory explains this phenomenon in terms of “reduced binding energy” of hydrogen to platinum

induced by Ru atoms.8 On the other hand, reactive OHads theory interprets it as a consequence of

higher availability of hydroxide species furnished by the oxophilic Ru atoms at the electrode-

electrolyte interface.9 Despite extensive research work on fundamentals of alkaline HOR, this

question regarding the origin of slow HOR kinetics is still unanswered.

114

Scheme 4.1 Illustration of HOR mechanism in alkaline electrolyte

To find a plausible explanation of the slow alkaline HOR kinetics which can bring a

consensus among the researchers, alternative Pt-Ru systems (other than only PtRu alloys) need to

be explored. Electrocatalysis being a surface phenomenon, the surface-active species that

predominantly affect the alkaline HOR kinetics need to be studied carefully. This can be done by

adapting different techniques to form a HOR active bimetallic Pt system- such as

electrodeposition, formation of core shell structured nanoparticles and exploring alternative

bimetallic systems such as Pt-Mo. In this work, we have explored alkaline HOR activity in all

these three different types of bimetallic Pt systems. Electrodeposition technique has been

reported by many researchers such as Adzicet al.,10 where Pt was deposited on a Ru/C surface.

Such formation of surface metal-metal bond is known to form near surface alloys (NSA) that are

known to induce significant modifications in the electronic properties of the metal overlayer.11,12

In this work, a modified version of the same technique has been tried to control the coverage of

Ru on Pt surface by the under- potential deposition of Cu on Pt, followed by the displacement of

Cu by Ru on Pt surface. Such an arrangement can give vital information on how Ru surface

coverage affects the rate of alkaline HOR and other reactions such as oxidative CO stripping. In

a core shell structured Pt-Ru system, we have developed a Ru/C core with uniform coating of Pt

on it, with variable thickness. Such a system ensures that there is no Ru on the surface exposed to

the electrolyte under ex situ conditions. This structural arrangement is expected to provide

insight into the dependence of HOR on surface availability of Ru and may shed some light onthe

dominance of binding energy effect or the reactive OH effect. Recently, Gasteiger et al.13 have

reported similar Ru core@ Pt shell system and its effect on CO stripping reaction and alkaline

HOR. However, there is a lack of evidence whether Ru is indeed located only at the

115

core.Moreover, the absence of electrochemical cyclic voltammogram (CV) data in alkaline

electrolyte compels us to investigate further into such core shell structure electrocatalysts.

Scheme 4.2 Illustration of the electrodeposition method to obtain controlled coverage of Ru on Pt surface

In the final part of the work, we have composed a PtMo(Ox)/C system to study alkaline

HOR and oxidative CO stripping reaction. Mo offers a rich electrochemistry where Mo can

attain multiple oxidation states ranging from -II to +VI, and thus Mo has a far-reaching

significance in electrochemical energy devices.14 However, the key factor in the electrochemical

behavior of Mo is the stability and the nature of the oxides, which are known to be MoO2 or

hydrated MoO2 in aqueous environment. In alkaline environment, Mo exists as MoO4-, whereas

in acid it exists as Mo3+ (low potentials) and MoO2 or MoO3 (at higher potentials).15 Such

oxophilic nature of Mo can be a source of reactive OHads to facilitate alkaline HOR, as shown in

scheme 4.3. PtMo systems have been widely applied in low temperature acid electrolyte systems

such as PEMFCs to combat CO poisoning and allow relaxation on the purity of the anode gas

feed. The CO stripping mechanism, like alkaline HOR, is widely debated and two theories try to

explain the increased tolerance of PtMo systems. The first theory is based on ligand effect that

suggests that metals like Mo, Ru tweak the electronic properties of Pt that weaken the Pt-CO

bond.16 The other theory is based on bifunctional effect that states that the binary metals M (Ru,

116

Mo) activate water (forming M-OHads) thatfacilitate oxidation of CO adsorbed on the

neighboring Pt sites.17

Scheme 4.3 The schematic representation of electrode- electrolyte interface for PtMo(Ox)/C electrode

In this report, an attempt has been made to answer the key questions: (1) what is the role

of Ru or Mo toward alkaline HOR mechanism (2) how does Ru and Mo influence the CO

stripping processes in acid and alkali electrolytes. While electrodeposition technique ensures the

surface localization of Ru, core- shell morphology assures that Ru atoms are segregated at the

core of the structure, while Pt atoms are confined at the outer shell. We have tested two core

shell structures with variable Pt layer thicknesses, and studied their morphology and

electrochemical behavior separately. In the last part, we have investigated a PtMo(Ox)/C catalyst

toward alkaline HOR and CO stripping reaction to elucidate the interrelationship between

electronic effects and bifunctional effects.

4.2 Experimental Section

117

4.2.1 Synthesis of Electrocatalysts

(a)Pt-Ru system prepared by electrodeposition: Pt/C (commercially acquired from Tanaka,

Japan) was deposited on glassy carbon electrode (GCE). After depositing Cu on Pt/C at certain

potential, the electrode was washed with deionized water and then immersed in a 1 mM solution

of RuCl3 for 5 minutes. In this process Ru displaces Cu from the GCE and forms a uniform layer.

The coverage of Ru is dictated by the coverage of Cu on the Pt, which in turn, is controlled by

potential at which Cu is deposited.

(b)Ru/C- Pt core @ shell nanoparticles: The synthesis recipe was adapted from a similar process

reported by Papandrewet al.18 40% Ru/C was mixed with calculated amount of Pt(acac)2 and put

inside a vacuum furnace, along with a separate vial containing distilled water. After purging Ar

and evacuating the furnace three times, the furnace was sealed under vacuum at a pressure of

0.03kPa and kept at a temperature of 230 oC for 15 hours. This metal organic chemical vapor

deposition (MOCVD) method ensures that Pt(acac)2 undergoes chemical deposition and forms a

coating of zero valent Pt on Ru/C. By varying the amount of Pt(acac)2, the thickness of the

coating can be changed.

Scheme 4.4 Illustration of the MOCVD procedure to obtain Ru/C@ Pt catalyst

(c) Synthesis of PtMo(Ox)/C catalyst: H2Pt(OH)6 was dissolved in 6% SO2 solution and stirred

for 24 hrs. to obtain a clear yellow solution of H3Pt(SO3)2OH. The solution was then diluted with

118

water and Vulcan XC-72R carbon was added into it to obtain a slurry. 30% H2O2 solution was

added dropwise into the slurry while maintaining a pH of 3.5, to obtain PtO2/C. This product was

filtered, washed and dried in vacuum oven overnight. Calculated amount of Molybdic acid was

mixed with water and PtO2/C. the reaction was stirred at room temperature for 24 hrs. The

resulting solid was then filtered, washed and dried in a vacuum oven. The solid was heat treated

at 550oC under 5% H2-Ar mixture to obtain the final catalyst, represented as PtMo(Ox)/C.

4.2.2Structural Characterizations

X- ray diffraction characterization was done using a RigakuUltima IV XRD with Cu Kα

source (lambda= 1.541 Å) operated at 40 kV and 44 mA. 2θ/θ scans were conducted using a 0.05

step size and 5 sec hold per step. Scanning electron microscopy (SEM) characterization was

done using a Hitachi S- 4800 FE- SEM. For validating sample elemental composition, EDAX

Genesis on the same SEM instrument was used. Transmission electron microscopy (TEM) and

selected area electron diffraction (SAED) data were collected using a JEOL 2010 field emission

gun at 200 kV accelerating voltage.

4.2.3Electrochemical Characterizations

All electrochemical measurements were done at room temperature using a rotating ring

disc electrode (RDE) from Pine Instruments connected to an Autolab (Ecochemie Inc., model

PGSTAT 30) potentiostat/ galvanostat. Tests were conducted in a 50 ml jacketed three electrode

cell in 0.1 M HClO4or 0.1M KOH. For each test, reversible hydrogen electrode (RHE), freshly

made, was used as the reference electrode and a Pt mesh wire was used as a counter electrode. A

glassy carbon disc (geometrical area of 0.2457 cm2)with a gold ring was used as working

electrode (WE). All the results shown here were collected after conditioning the electrodes at a

119

scan rate of 50 mV/s between 0.05 to 0.8 V vs RHE for Ru containing samples, and between

0.05 to 1.2 V for Pt/C, for 20 scansor until stable features were achieved.

For the under-potential deposition of Cu on Pt/C, the Pt/C catalyst deposited on the WE

was immersed in a solution of 2 mM of CuSO4 in 0.1M HClO4. The WE was polarized at 1.0V

for 2 mins to ensure the oxidation of any pre-deposited Cu on Pt/C. Following this step, the WE

was held at desirable potentials to control the deposition of Cu on Pt/C. For stripping experiment,

the WE electrode was then subjected to a cyclic volammetric process starting from the potential

it was held at.

For CO stripping experiment, CO was purged into the electrolyte for 10 minutes while

holding the potential at 0.05 V vs. RHE in 0.1M HClO4 or, at 0.2 V vs. RHE in 0.1M KOH.

After that, argon was purged for 20 minutes to remove any dissolved CO in the electrolyte and to

ensure that CO is present only at the electrode surface adsorbed on the catalyst. The scan was

then continued at a scan rate of 10 mV/sec.

4.3 Results and Discussions

4.3.1 Pt/C-Ru catalysts obtained by electrodeposition

(a) Cyclic Voltammogram and HOR activity: Under potential Cu deposition (CuUPD) is a

nondestructive and simple electroanalytical technique that helps to determine the electrochemical

surface area of a catalyst.19,20 Such underpotential deposition causes metal atoms to deposit onto

an electrode forming monolayer, or sub monolayer at a potential more positive than the required

for the bulk deposition. Cu is used for this method involving Pt and Ru systems, owing to the

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Figure 4.1CuUPD stripping experiments on Pt/C catalyst in 2mM CuSO4/0.1M HClO4 electrolyte at a scan rate of 10 mVsec-1 under room temperature

similarity in the atomic radii of Cu, Ru and Pt. Displayed in Figure 4.1 is the CuUPD stripping

curves obtained on Pt/C catalyst deposited on a GCE. As evident from the figure, the intensity of

the peak reduces with the increase in the potential at which Cu was deposited. For example, the

peak for CuUPD strip at 0.3V has a higher area compared to that at 0.5V, and the peak for 0.5V

has a higher area compared to that for 0.7V. The area under the peak gives a value of the total

electronic charge associated with the stripping process, which in turn, can determine the

electrochemical area associated with the CuUPD process. Deposition of Ru was done by

displacing the Cu from the Pt surface, by immersing the WE in a RuCl3 solution. Therefore,

coverage of Ru on Pt surface can be controlled by controlling the coverage of Cu monolayer by

applying different UPD potentials.

The as prepared Pt/C-Ru catalysts were then tested in alkaline electrolyte for collecting

CV and assaying their HORcatalytic activity (Figure 4.2). From the CV, the presence of Ru on

Pt/C is verified by the characteristic HUPD adsorption and desorption peaks at 0.15VRHE. It is

interesting to note that inPtRu/C alloy catalyst predominantly exhibits HUPD stripping peak

characteristic on Ru surface, whereas the Pt/C-Ru displays the HUPD peaks characteristic of both

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Ru and Pt surfaces. This can be ascribed to the presence of thin layer of Ru on the Pt surface,

where Hads atoms are oxidized on both Pt and Ru. In PtRu/C alloy, the catalyst surface contains

both Pt and Ru atoms,

Figure 4.2 (a) CV of Pt/C-Ru catalysts along with Pt/C and PtRu/C in Ar purged 0.1M KOH under room temperature at 10 mVsec-1 scan rate, (b) HOR polarization curves collected at 1600 rpm in H2purged 0.1M KOH under room temperature at a scan rate of 10 mVsec-1.

Figure 4.3 Plot of kinetic current density for electrodeposited Pt/C-Ru catalysts as a function of potentials toward HOR in 0.1M KOH

and the strong HUPD stripping peak from Ru surface diminishes the HUPD stripping feature

emanating from Pt surface. Therefore, the ultra-low coverage of Ru on Pt/C in Pt/C-Ru samples

results in prominent HUPD peaks from both Pt and Ru surfaces. A Pt/C-Ru BLANK sample was

prepared by directly immersing the Pt/C electrode into RuCl3 solution skipping the CuUPD step.

This blank sample showed minor HUPD peak characteristic to Ru surface, thus indicating that Ru

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gets self- adsorbed onto Pt surface even without displacing Cu atoms. This similar study has

been reported by Wieckowskiet al.11 where they had reported ion (self) adsorption of Ru ions

onto Pt surface. It is important to note that these self- adsorbed anions are not zero valent Ru

metal deposits, and hence should not be considered as a system with deposition of Ru. In fact, in

order to achieve a layer of metallic Ru from the adsorbed anions, they system needs to be

discharged at a low potential, a step which was not carried out in this work. HOR polarization

curves were then collected and are shown in Figure 4.2(b). While Pt/C catalyst exhibits slower

kinetics and reaches limiting current at a potential of 350 mVRHE, all the Ru- modified Pt/C

catalysts exhibited superior performance by reaching limiting current before 100 mVRHE. The

image in the inset of Figure 4.2 (b) qualitatively displays the catalysts performance at the micro-

polarization region. It is unlikely that the weakly adsorbed layer of Ru ions would impart

electronic changes in Pt structure, or form NSA moieties. Hence the fact that the blank sample

shows superior performance than Pt/C indicates strongly that the enhancement in performance in

bimetallic PtRu catalysts are influenced by the availability of oxophilic Ru moieties on the

catalyst surface. In Figure 4.3, a plot of kinetic current density vs potential is displayed, using

which the exchange current densities were derived that compares the kinetic performance of the

catalysts. The exchange current densities for Pt/C-Ru samples @ 0.3V, 0.5V and Pt/C-Ru

BLANK samples were found to be same as that of PtRu/C catalyst- 0.14mA/cm2Pt. On the other

hand, Pt/C-Ru @ 0.7V was found to be 0.9mA/cm2Pt, a value much closer to that for pure

Pt(0.56mA/cm2Pt). As seen clearly in Figure 4.3, the surface coverage of Ru atoms on Pt has a

direct influence on the HOR kinetics. Interestingly, the Pt/C-Ru BLANK sample also exhibited

fast kinetics, even greater than Pt/C-Ru@ 0.7V, where Ru has partial surface coverage. It is

important to state here that in blank sample; the Ru ions form a thin film on Pt surface held by

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weak forces. But when a layer of Cu is present on Pt surface, Ru atoms specifically displace the

Cu atoms during the 2 min deposition process forming a layer a Ru atoms on Pt.

Figure 4.4 oxidative CO stripping profiles of Pt/C-Ru electrodeposited catalysts in (a) 0.1M HClO4 and (b) 0.1M KOH collected under room temperature at a scan rate of 10 mVsec-1. (c) and (d) represent CO stripping curves of Pt/C and PtRu/C in 0.1M HClO4(solid) and 0.1M KOH(dashed) collected under room temperature at a scan rate of 10mVsec-1.

(b)Oxidative CO stripping experiments: The electrodeposited Pt/C-Ru samples were then tested

for oxidative CO stripping reaction in acid and alkaline electrolytes (Figure 4.4). In acid, the

onset of the CO stripping process is governed by the formation of OHads(from water activation)

on the catalyst surface. For example, in pure Pt/C, the formation of Pt-OHads(from water

activation) begins at 0.62VRHE, and this is reflected clearly in the onset of CO stripping peak

(Figure 4.4 (a)). If Ru is present in the catalyst, then Ru is expected to initiate the formation of

OHads at a much lower potential, and indeed, the onset of CO stripping in PtRu/C occurs at 0.45V

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(Figure 4.4 (d)). However, in the electrodeposited samples, the coverage of Ru is seen to have a

direct effect on the CO stripping profiles in 0.1M HClO4. The CO stripping peak potential

undergoes cathodic shifts once a higher coverage of Ru is achieved as in Pt/C-Ru @0.4VRHE and

0.5VRHE. Therefore, the population of surface Ru atoms can be seen to have a direct influence on

the CO stripping potential. The unavailability of reactive OHads moieties in acidic electrolyte

(sole source of reactive OHads is via water activation) causes the CO oxidation reaction to rely

solely on the hydroxyl species furnished by the surface Ru atoms. It is important to note that the

CO stripping peak appears at an even lower potential in PtRu/C catalyst (0.55VRHE), where the

population of Ru is significantly higher than the Pt/C-Ru samples. These observations point to an

important fact on the mechanism of CO stripping in acid. If ligand effect were solely responsible

for CO stripping, then a trend as in Figure 4.4 (a) would not have been obtained. Therefore, it

can be stated that in acid electrolyte, the bifunctional effect has a strong impact on CO stripping.

However, in alkaline medium, the CO stripping potential for all the Pt/C-Ru samples are

positioned at 0.55V and are independent of the surface coverage of Ru atoms on Pt, thereby

proving that in alkaline medium OH- ions (from the electrolyte) ensure adequate reactive OHads

to facilitate CO oxidation.

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Figure 4.5 (a) X-ray diffraction pattern of Ru/C@ Pt catalysts; (b) and (c) display the HRTEM image of Ru/C@ Pt thin shell and thick shell catalysts, respectively

4.3.2 Ru/C @Pt core shell electrocatalysts

(a) Structural Characterization: The core @ shell catalysts were obtained by depositing metallic

platinum on Ru/C using MOCVD method. This type of structure is supposed to have only Pt

atoms on the surface, but its catalytic properties may be substantially modified due to the

interaction of the Pt shell atoms with Ru core atoms. Two different catalysts were synthesized

with different coating thicknesses of Pt, while preserving the underlying Ru/C frame. The

representative XRD crystallographic structures are illustrated in Figure 4.5 (a), along with the

corresponding diffraction patterns for Ru hcp and Pt fcc phases. In case of Ru/C @Pt thin shell,

the diffraction pattern shows distinctive diffraction from both Ru and Pt phases. On the other

hand, the thick shell catalyst predominantly exhibits Pt fcc, diffraction pattern showing that a

higher coating of Pt results in the formation of a “Pt-Ru” phase and Ru no longer exists as a

separate phase. The TEM images in Figure 4.5(b), (c) show the images of the core shell catalysts

under transmission electron microscope (TEM). It can be clearly seen that in Ru/C @Pt samples,

there is a significant variation in the particle size in the two samples.While the thin shellcatalyst

hasan average particle size

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Figure 4.6 (a) HOR polarization curves of Pt/C, PtRu/C, Ru/C@ Pt thin shell and thick shell catalysts collected at 1600 rpm in H2 purged 0.1M KOH at a scan rate of 10 mVsec-1 under room temperature and (b) Plot of kinetic current density as a function of potential in the micro polarization region drawn from the right panel

of 2.5nm, the average particle size was 6 nm for the thick shell catalysts. The morphologies of

these two catalysts are dramatically different that greatly affects their electrochemical active

surface areas.

(b) HOR activity in 0.1M KOH: The core@ shell catalysts with various thicknesses were tested

for HOR activity in 0.1M KOH electrolyte. As seen from the TEM images, the two catalysts

have discernible differences particle size, and hence the catalytic surface area is expected to be

significantly different. Indeed, the electrochemical surface area (ECSA) of Ru/C@ Pt- thin shell

was found to be 30 m2/gPt, whereas ECSA of the thick shell catalyst was found to be 10 m2/gPt. It

should be noted that the comparatively lower ECSA of both core@ shell catalysts as compared to

commercial Pt/C catalysts (usually in the range of 45-65 m2/gPt) arises due to the coating process,

where metallic Pt coats the carbon support as well, thereby lowering active catalyst surface area.

A direct consequence of this feature can be seen in the HOR polarization profile of the catalysts

(Figure 4.6 (a)). While kinetic plot shows equal performance in the micro-polarization region,

the limiting current displayed by the two catalysts varies significantly in the two catalysts. This is

because the thick shell Ru/C@ Pt catalyst has significantly lower ECSA that limits the number of

accessible active sites for HOR, thereby affecting the limiting current. To achieve a deeper

insight into the role of Ru core, the core shell catalysts were tested for CO stripping reactions.

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Figure 4.7 CO stripping profiles of Ru/C@ Pt thin shell catalyst in (a) 0.1M HClO4, (b) 0.1M KOH and Ru/C@ Pt thick shell catalysts in (c) 0.1M HClO4 and (d) 0.1M KOH collected under room temperature at a scan rate of 10 mVsec-1.

(c) CO stripping experiment on Ru/C@ Pt core shell catalysts: the CO stripping profiles of the

core shell catalysts are displayed in Figure 4.7. Both the thin and thick shell catalysts exhibit

similar CO stripping profiles suggesting that the thickness of the Pt coating does not impact the

oxidation of adsorbed CO on Pt. In fact, the onset of the oxidative process occurs at similar

potentials in all the cases, both in acidic and alkaline media. Considering the Ru atoms are

localized in the core and there is no OHads facilitated by surface Ru, we can say that the Ru atoms

weakens the energy of Pt-CO bond by inducing ligand effect. This is different than the results

obtained in the previous section, where we showed the surface Ru atoms have direct effect on the

CO stripping process. In a prior study Eichhornet al.,21 they claimed that bifunctional mechanism

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in such Ru@ Pt core-shell cannot be implicated due to the absence of Ru in the surface layer, and

hence they explained the CO stripping behavior of Ru@ Pt catalysts due to electronic

modifications on the basis of DFT calculations. However, to be absolutely sure about the

mechanism, some in situ XAS experiments need to be done to detect any morphological changes

or segregation effect on the catalyst during in situ operating conditions. Therefore, it would be

safe to say that both bifunctional effect and electronic effect play a role while oxidative stripping

of CO in PtRu catalysts, however, depending on the structure and composition of the catalyst,

one effect may have dominance over other as per the studies carried out in this work.

4.3.3 PtMo(Ox)/C catalyst

(a) Structural Characterization: The PtMo(Ox)/C catalyst synthesized was tested under TEM and

the catalyst particles were found to be 3-5nm in size (Figure 4.8 (a). The inset shows a high

resolution TEM image of the PtMo(Ox)/C catalyst which establishes the high crystallinity of the

nanoparticles. The XRD data is illustrated in figure 4.8 (b) and is juxtaposed on the XRD pattern

of a Pt/C catalyst. The diffraction peaks are at same position and it implies that no geometric

change has occurred in the Pt lattice due to the inclusion of Mo in the catalyst. The crystallite

size of the PtMo(Ox)/C catalyst was found to be 3.4 nm, whereas the Pt/C (made in house for the

preparation of PtMo(Ox)/C) catalyst had a crystallite size of 3.3nm.

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Figure 4.8 TEM image of PtMo(Ox)/C catalyst with a HRTEM image of a single particle in inset and (b) X-ray diffraction profile of Pt/C and PtMo(Ox)/C catalysts

(b) Electrochemical characterization of PtMo(Ox)/C catalyst: Figure 4.9(a) shows the CV of the

PtMo(Ox)/C catalyst in acid and alkaline electrolyte. In acid electrolyte, there is a distinct redox

peak discernible at 0.45VRHE due to transition of Mo from HMoO4- to MoO2. This redox

transition does not occur in alkaline electrolyte since Mo exists as HMoO4- throughout the

experimental potential window in alkaline medium and no additional redox process occurs

within the potential range. Therefore, in alkaline electrolyte, the CV of the PtMo(Ox)/C catalyst

strongly resembles that of Pt/C. The PtMo(Ox)/C catalyst was then tested for alkaline HOR and

polarization curves were collected at room temperature (Figure 4.9(b)). Compared to Pt/C

catalyst, PtMo(Ox)/C catalyst exhibited superior catalytic activity. However, compared to PtRu/C

catalyst, PtMo(Ox)/C catalyst exhibited lower kinetics. The reason may be explained in terms of

two factors, namely:

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Figure 4.9 (a) CV of PtMo(Ox)/C catalyst in Ar purged 0.1M HClO4 and 0.1M KOH electrolytes collected under room temperature at a scan rate of 10 mVsec-1 and (b) HOR polarization curves of Pt/C, PtRu/C and PtMo(Ox)/C collected at 1600 rpm in H2 purged 0.1M KOH at a scan rate of 10 mVsec-1 under room temperature

1) In PtRu/C catalyst, the extent of alloy formation between Pt and Ru is larger than in

PtMo(Ox)/C, that causes weakening of Pt-Had bond in PtRu catalyst, hence, HOR is more

facile.

2) Ru has higher oxophilicity compared to Mo,22 and hence, can furnish more reactive OH-

for HOR compared to Mo, and hence, PtRu/C has faster rate of HOR.

While factor (1) is based on the ligand effect theory, factor (2) is based on the

bifunctional mechanism theory, and it is hard to tell which of these two factors are dominant in

deciding alkaline HOR activity. The exchange current density derived from the HOR

polarization curve was found to 1.1 mA/cm2Pt, compared to 1.4 mA/cm2

Pt in the case of PtRu/C

catalyst. The fact that PtMo(Ox)/C has a catalytic activity significantly better than Pt/C without

having geometric interaction with Pt lattice indicates the importance of surface oxides in alkaline

HOR, where the oxides provide the reactive hydroxyl moieties to facilitate HOR. However, Mo

can potentially form thin NSA layer on Pt overlayer, thereby tweaking the electronic properties

of Pt toward H2 adsorption.

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(c) CO stripping experiment on PtMo(Ox)/C catalyst: To appraise the role of Mo atoms, CO

stripping experiments were performed (Figure 4.10). In Pt/C, the CO stripping process starts at

0.7VRHE, in acid. In case of PtMo(Ox)/C, onset of CO stripping occurs at a much lower potential

of 0.55VRHE, just after the inherent redox transition of HMoO4- to MoO2. This study strongly

suggests that the surface oxides play a crucial role in determining the CO stripping process. Ross

et al.14 have attributed the CO oxidation property of Pt70Mo30 bulk alloy based on availability of

Oads and OHads species on Mo, as observed using spectroscopic techniques. However, there is

lack of proof whether these adsorbed species are reactive enough to facilitate CO oxidation.

Therefore, they stated both bifunctionality and modified electronic properties of Pt70Mo30 alloy

as factors affecting CO stripping reaction on the alloy. In alkaline medium, however, both Pt/C

and PtMo(Ox)/C exhibit very similar CO stripping profiles. The cathodic shift in the CO

stripping onset

Figure 4.10 CO stripping profile of (a) Pt/C and (b) PtMo(Ox)/C catalyst in 0.1M HClO4 (dashed) and 0.1M KOH (dotted) collected at a scan rate of 10 mVsec-1.

of Pt/C catalysts have been explained in terms of shift in the potential of zero charge to 0.17VRHE

in alkali. In such low potential where the Pt has Hads on its surface, water can attach weakly on Pt

due to the presence of partial positive ∂+ charge on Pt imparted by HUPD. Comparing the CO

stripping profiles of PtMo(Ox)/C and Pt/C in alkaline electrolyte, it seems that Hads on Pt is same

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in both the catalysts, suggesting that there is no change in H2 binding energy on Pt in

PtMo(Ox)/C catalyst. This observation strongly points towards the dominance of reactive OHads

effect on PtMo(Ox)/C toward alkaline HOR.

4.4 Conclusion

This work encompasses a detailed study of three different kind of systems: Pt/C-Ru,

which was obtained by electrodeposition of Ru on Pt/C; Ru/C@Pt core shell catalysts; which

were obtained by water vapor assisted chemical vapor deposition; and PtMo(Ox)/C; which was

synthesized in house by deposition of Mo oxide on Pt/C catalyst. In case of the electrodeposited

catalysts, it was found that a thin monolayer, or sub monolayer of Ru atoms on Pt/C is sufficient

to induce superior alkaline HOR activity. Even a blank sample where Ru ions were loosely

attached to Pt surface via weak forces showed superior HOR activity, thereby strongly indicating

that the availability of surface oxides play a major role in determining the performance of PtRu

systems toward HOR. CO stripping experiments also provided insight into the different

mechanisms which are followed in acid and alkaline electrolytes. While in acid, where the

population of reactive OHads is scarce, there is a clear dependence of CO stripping potential on

the coverage of Ru on Pt surface. However, in alkaline electrolyte, the CO stripping potential

was found to be independent of Ru concentration, and hence suggests that the CO stripping is

facile due to the opulent availability of OH species in alkaline electrolyte.

In the Ru/C@ Pt core shell catalysts, the Ru atoms are expected to be segregated at the

catalyst core, and covered by a shell of metallic Pt. In case of a thin shell, the Ru and Pt phases

can be distinguished by XRD, suggesting that the Ru and Pt phases retain their separate entities.

However, in the thick Pt shell catalyst, separate Ru and Pt phases could not be identified and a

Pt-Ru alloy phase was detected by XRD, indicating that Ru atoms have migrated into the Pt shell

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to form an intermetallic Pt-Ru alloy. Both the thick and thin Pt shell catalysts exhibit superior

activity compared to Pt/C, with their performance matching with that of commercial PtRu/C

catalyst. These results in conjunction to the XRD reports strongly suggest that a PtRu alloy phase

has been formed that facilitate HOR. From the CO stripping results, there is a clear indication

that a higher Pt content anodically shifts the CO stripping potential in acidic medium, this shows

that for CO stripping reaction, a higher Ru content is required to oxidatively remove CO at a

lower potential.

Lastly, PtMo(Ox)/C catalyst was studied for its alkaline HOR activity. Owing to its high

oxophilicity and ability to achieve multiple oxidation states, Mo was selected for combining with

Pt and study the resulting catalytic activity. XRD data results suggest that Mo does not modify

the Pt lattice. However, it is outside the scope of this work to predict whether a thin layer of NSA

was formed on Pt surface. Toward alkaline HOR, PtMo(Ox)/C catalyst exhibits significantly

better activity than Pt/C. The reason behind this performance is ascribed to both oxophilic nature

of Mo, where the surface Mo oxides provide OH species to facilitate HOR and ligand effect of

Mo which weakens the Pt-H bond to a significant extent. CO stripping experiments further

suggest that the potential at which CO is oxidatively stripped is a direct function of the redox

transition from HMoO4- to MoO2.

Hence, in this work, we have tried to elucidate the predominant effect that decide alkaline

HOR mechanism and CO stripping mechanism. Further experiments such as in situ synchrotron

spectroscopy and XPS can reveal the nature of oxides and how they affect the electrochemical

reactions occurring at the electrode- electrolyte interface. Most of the previously reported studies

on Pt-Mo systems have been reported in acid media before owing to the importance of proton

exchange membranes as electrolytes in fuel cells. However, as the AEMFC technology has also

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developed in recent times, it is important to know in detail the mechanism of HOR in alkaline

medium, to develop more efficient and cost effective electrocatalysts.

4.5 Acknowledgement

The author would like to thank U.S. Department of Energy, Office of Science (Contract

No. DE-AC02-98CH10886) for providing financial assistance to carry out this project.

4.6 References

(1) Zhang, H.; Shen, P. K. Chemical Society Reviews2012, 41, 2382. (2) Arges, C. G.; Ramani, V. K.; Pintauro, P. N. The Electrochemical Society Interface2010,

19, 31. (3) Kim, O.-H.; Cho, Y.-H.; Chung, D. Y.; Kim, M. J.; Yoo, J. M.; Park, J. E.; Choe, H.;

Sung, Y.-E. Scientific reports2015, 5. (4) Lin, B.; Qiu, L.; Lu, J.; Yan, F. Chemistry of materials2010, 22, 6718. (5) Sheng, W.; Gasteiger, H. A.; Shao-Horn, Y. Journal of The Electrochemical Society2010,

157, B1529. (6) Koper, M. T. Nature chemistry2013, 5, 255. (7) Durst, J.; Siebel, A.; Simon, C.; Hasche, F.; Herranz, J.; Gasteiger, H. Energy &

Environmental Science2014, 7, 2255. (8) Sheng, W.; Myint, M.; Chen, J. G.; Yan, Y. Energy & Environmental Science2013, 6,

1509. (9) Strmcnik, D.; Uchimura, M.; Wang, C.; Subbaraman, R.; Danilovic, N.; Van Der Vliet,

D.; Paulikas, A. P.; Stamenkovic, V. R.; Markovic, N. M. Nature chemistry2013, 5, 300. (10) Yang, L.; Vukmirovic, M. B.; Su, D.; Sasaki, K.; Herron, J. A.; Mavrikakis, M.; Liao, S.;

Adzic, R. R. The Journal of Physical Chemistry C2013, 117, 1748. (11) Chrzanowski, W.; Wieckowski, A. Langmuir1997, 13, 5974. (12) Ticanelli, E.; Beery, J.; Paffett, M.; Gottesfeld, S. Journal of electroanalytical chemistry

and interfacial electrochemistry1989, 258, 61. (13) Schwämmlein, J. N.; El-Sayed, H. A.; Stühmeier, B. M.; Wagenbauer, K. F.; Dietz, H.;

Gasteiger, H. A. ECS Transactions2016, 75, 971. (14) Grgur, B.; Markovic, N.; Ross, P. The Journal of Physical Chemistry B1998, 102, 2494. (15) Saji, V. S.; Lee, C. W. ChemSusChem2012, 5, 1146. (16) Ochal, P.; de la Fuente, J. L. G.; Tsypkin, M.; Seland, F.; Sunde, S.; Muthuswamy, N.;

Rønning, M.; Chen, D.; Garcia, S.; Alayoglu, S. Journal of Electroanalytical Chemistry2011, 655, 140.

(17) Roth, C.; Papworth, A.; Hussain, I.; Nichols, R.; Schiffrin, D. Journal of Electroanalytical Chemistry2005, 581, 79.

(18) St. John, S.; Atkinson III, R. W.; Unocic, K. A.; Unocic, R. R.; Zawodzinski Jr, T. A.; Papandrew, A. B. ACS Catalysis2015, 5, 7015.

(19) Green, C. L.; Kucernak, A. The Journal of Physical Chemistry B2002, 106, 1036.

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(20) Green, C. L.; Kucernak, A. The Journal of Physical Chemistry B2002, 106, 11446. (21) Alayoglu, S.; Nilekar, A. U.; Mavrikakis, M.; Eichhorn, B. Nature materials2008, 7, 333. (22) Kepp, K. P. Inorganic Chemistry2016, 55, 9461.

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Chapter 5

Engendering Non-Precious Catalysts for Oxygen Depolarized Cathodes in

Chlor Alkali Electrolyzer Cells

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Chapter 5

Engendering Non-Precious Catalysts for Oxygen Depolarized Cathodes in

Chlor Alkali Electrolyzer Cells

5.1 Introduction

Chlorine generation is an integral part ofchemical industry as it is a building block of

several polymers including polyvinyl chloride (PVC).1 There are two major sources of chlorine

that are used for chlorine generation- sea water (brine) and hydrochloric acid (HCl) which is a

byproduct obtained from the polymer industries.2,3 Though the vast availability of brine ensures

low cost, the electrolysis process requires extensive electric power that restrains the

establishment of brine electrolysis as the dominant chlorine generating route. In fact, 2% of total

USA electric power, which amounts to 87600 GWh/year, is dedicated towards meeting the

power requirements of brine electrolysis.4 The common electrolytic processes associated with

chlorine generation from brine include diaphragm cell process and the membrane -cell process,

both of which involve production of hypochlorite as a byproduct that reduces the efficiency of

the plant.5,6 Albeit such side reaction can be avoided in a mercury cell electrolyzer7,8, the

presence of mercury species impose threat on the practical applicability of this process. The

introduction of efficient ion exchange membranes has mitigated the high cost to some extent, but

membrane technology has reached a saturation point that can no further reduce the power

requirement. An alternative route of lowering the power requirement has emerged by employing

oxygen depolarized cathode (ODC).9-11 This way, the unwanted formation of hydrogen is

prohibited and more importantly, the theoretical voltage requirement is reduced drastically owing

to the redox potential of oxygen reduction reaction. In other words, the thermodynamic potential

difference associated with the more classic route involving hydrogen generation is about 2.2V,

which can be reduced by 1.23V with the incorporation of ODC. However, ODC brings up a

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major challenge, i.e. sluggish kinetics of oxygen reduction reaction which causes huge

overpotential losses.12-14 The state of the art catalyst Pt/C is known to suffer significant

overpotential losses due to strong adsorption of hydroxyl anions on the Pt ORR active sites,

resulting in poisoning of the catalyst.15-17 Additionally, Pt is known to aggravate carbon

corrosion, which is of relevance in electrolyzer cells due to the oxidative environment of

ODC.18,19 Therefore, this report will focus on the applicability of non- platinum electrocatalysts

as ODC in chlor alkali electrolyzers.

Among the non-precious electrocatalysts, the metal-nitrogen-carbon (M-N-C) are one of

the most active ORR catalysts reported.20 Typical metal centers consist of transition metals such

as Fe, Ni and Co embedded in a carbon matrix with nitrogen coordination centers.21-23 General

synthesis procedure involves combining the metal, nitrogen and/ or carbon precursors followed

by one or two pyrolysis steps, which gives the catalyst its final structure.20,21,24,25 In some cases,

nitrogen precursors may be absent, and pyrolysis under reactive ammonia is generally used to

incorporate nitrogen species in the catalyst framework. It is believed that the transition metal

constitutes the active site of the catalyst and participates in ORR.20,25 In alkaline electrolyte,

these M-N-C catalysts are known to have even better performance than Pt/C. This is due to the

outer sphere electron charge transfer occurring in alkaline electrolyte that allows a non-

specificity of the underlying electrode surface, facilitating ORR by a two electron route.24,26 Such

phenomenon opens the gate to non-precious catalysts to be used in alkaline electrolyte,

minimizing overpotential loss and avoiding the cost associated with precious metals. In this

work, we have studied Fe-N-C catalyst, synthesized using two very different methods. The first

method consisted of using metal organic framework as a template and support to insert active

sites, giving rise to FeMOF catalyst. This method does not include a separate precursor for

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carbon, as carbon is synthesized in situ during the pyrolysis step.21 The second method used a

polymer based approach, where polyaniline was used as the nitrogen precursor. In this case,

carbon black was added separately during synthesis, and Fe-PEI catalyst was obtained.25 It is

understood that despite the similar nature of the active sites generated by these two routes, the

carbon structure is very different. While the MOF based catalyst has a higher surface area and

porous framework, the second method has the carbon morphology resembling that of carbon

black. In addition to this, the extent of graphitization in the carbon matrix of the two catalyst is

very different. This has a direct impact on the durability and corrosion resistance of the catalyst,

since higher graphitization provides better tolerance against corrosion. In this study, we have

first compared the ORR catalytic profiles of FeMOF and Fe-PEI catalysts in alkaline electrolyte

using rotating ring disk electrode experiments. We have then applied these catalysts in an

electrolyzer cell to appraise the performance in an actual cell. Preliminary results have revealed

that both these catalysts outperform the state of the art catalyst Pt/C, in RDE and chlor alkali

electrolyzer cell. The two catalysts perform similarly with some minor differences which we

believe, are due to the very different structure of the carbon matrix in the two catalysts.

5.2. Experimental Section

5.2.1 Synthesis of catalysts

(a) Synthesis of Fe- PEI catalyst: The FePEI catalyst was synthesized from branched

polyethyleneimine (50–100 kMW) and ferric chloride as nitrogen and metal source, respectively.

Initially, the branched PEI was dissolved in water to obtain 10 wt.% PEI solution, and then ferric

chloride solution was added drop wise to the PEI solution while stirring. The metal-polymer

network was then supported on high surface area carbon (Ketjen Black 600 JD: Black Pearl 2000

= 1:1 by weight) through sonication. The mixture was left stirring over 12 h to allow full

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complexation. After purging Ar for half an hour, diluted H2O2 solution was added to cross-link

the metal-polymer/C network through fenton reaction. Then the mixture was filtered, followed

by evaporation of solvent in vacuum oven at 80 oC over period of 12 h. The supported polymer-

carbon hybrid materials were pyrolyzed at 900oC in argon atmosphere for one hour. The

morphology of the resulting catalyst is displayed in Scheme 1.

(b) Synthesis of FeMOF catalyst: The detailed synthesis recipe is described in a previous

publication from our group.21Commercially procured zinc oxide was treated at 400oC in air prior

to use. The treated zinc oxide, 2- methylimidazole and ammonium sulfate (along with trace

amount of methanol) were ball milled for an hour using methacrylate balls as grinders.

Following this, Fe(II) acetate and 1,10- phenanthroline monohydrate were added and ball milled

for another two hours. The resulting pinkish powder was then heat treated at 1050oC for 1 hr

under Ar. A second heat treatment at 1050oC was then conducted under ammonia atmosphere.

The morphology of the resulting catalyst is displayed in Scheme 5.1. As can be clearly seen, the

morphology of Fe-MOF is remarkably different than Fe-PEI catalyst. The high porosity and

homogeneous distribution of the pores can be seen in the SEM image.

Scheme 5.1Synthesis of Fe-PEI and FeMOF catalysts along with their corresponding SEM images

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5.2.2 X-ray absorption spectroscopy

The XAS studies were conducted at Fe K-edge (7112 eV). The in situ XAS spectra was

collected using a specially designed spectro- electrochemical flow cell. Ar purged 0.1M HClO4

electrolyte was constantly circulated through the flow cell.27 The ex situ spectra were collected in

transmission mode and in situ spectra were collected in fluorescence mode, using a 32 element

Ge detector. Detailed explanation on XAS data analysis and EXAFS fitting are described in a

previous publication. Measurements were collected under in situ conditions at various potentials.

Before each measurement, the cell was held for 5 minutes to reach steady state. The data was

processed and fitted using Athena and Artemis software programs. The calibration, background

removal and normalization were done using IFEFFIT suite.28-30

5.2.3 Fabrication of gas- diffusion electrode(GDE)

A semi solid paste ofCarbon (Ketjan Black) in water, isopropanol and 60% PTFE

(carbon: PTFE= 7:3 by weight) was made and coated on a commercial carbon cloth using a

doctor blade. The final loading of carbon was 10 mg/cm2 and that of PTFE was 0.3 mg/cm2.

After drying the carbon layer, the carbon cloth was sintered in a muffle furnace at 400oC for 15

mins to ensure an even distribution of PTFE throughout the gas diffusion layer (GDL). The

active catalyst layer was then sprayed on the GDL using a catalyst ink comprising water,

isopropanol and 5% Nafion®solution. The final loading of Pt/C catalyst and non PGM catalysts

were 1 mg/cm2 and 3 mg/cm2, respectively. A final layer of 5% Nafion in isopropanol was

sprayed for creating a hydrophilic layer to facilitate smoother transition of the ions to/ from the

electrode surface.

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5.2.4 Rotating disc electrode experiment

All electrochemical measurements were done at room temperature using a rotating ring

disc electrode (RRDE) from Pine Instruments connected to an Autolab (Ecochemie Inc., model

PGSTAT 30) potentiostat/ galvanostat. Tests were conducted in a 50 ml jacketed three electrode

cell in the desirable electrolytes. For each test, freshly made, mercury mercuric oxide (Hg/HgO)

reference electrode was used with a Pt mesh wire serving as a counter electrode. A glassy carbon

disc (geometrical area of 0.247 cm2) was used as working. All the results shown here were

collected after conditioning the electrodes at a scan rate of 50 mV/sec for 50 scans in a voltage

range between (0.05V – 1.0V), or until stable features were achieved.

5.2.5Electrolyzer cell testing

The half-cell (obtained from DeNora Tech) comprised a Teflon body with compartment

designed for the flow of 5M NaOH electrolyte. While a nickel foil was used as the electrode

surface for oxygen evolution reaction (OER) at the anode terminal, the catalyst sprayed GDE

was attached to the nickel foil on the cathode side as the electrode for ORR. Oxygen gas was

flown in the cathode side while 5M electrolyte was continuously circulated inside the cell

compartment. With the help of metal plates at the exterior, the cell compartment was kept at a

temperature of 60oC. All the data collected using the half- cell were recorded at 60oC. The full

cell data was collected at DeNora Tech using a brine electrolyzer. 5MNaCl solution was used as

the electrolyte. A dimensionally stable anode (consisting of Ti, Ru and Ni) was used for chlorine

evolution, while the desired catalyst was used as the cathode catalyst serving as ODC. Data were

recorded at 60oC.

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Scheme 5.2Representation of the half-cell chlor alkali electrolyzer (with OER or Cl2 evolution as the anodic process) with description of the ODC in enlarged form

5.3. Results and Discussion

Figure 5.1 FT-EXAFS of (a) FeMOF and (b) Fe-PEI catalysts under ex situ and in situ conditions collected at the Fe K-edge in Ar saturated 0.1M HClO4 electrolyte.

5.3.1 Structure of iron centers

Iron centers that constitute the active center of Fe based non PGM catalysts were studied

using in situ XAS. Fourier transform (FT) of the extended XAS fine structure (EXAFS) helps to

identify the coordination number of the iron center and the gives details about the identity of the

immediate neighboring species. While a distinct scattering peak corresponding to the Fe-

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Nxspecies is present, the Fe-Fe scattering peak at 2.1Å (without phase corrections) is nondescript

(Fig 5.1(a)). This shows that a minimal amount of inorganic Fe moieties is present in the FeMOF

catalyst. This can further be corroborated by Mössbauer spectroscopy as reported in a previous

work.21In accordance to the earlier studied metal coordinated functionalities in the high

temperature treated materials, two types of metal centers can be identified in Fe-PEI catalyst (Fig

5.1(b))- namely metal nitrogen coordinated centers (Fe-N4) and Fe nanoparticles (Fe/FexOy). The

Fe-Fe FT XAS peaks perfectly match with the spectra of a standard Fe foil, indicating the

presence of Fe nanoparticles. The Fe-Nx can be fitted using a Fe-N4 model, proof of which have

been reported in our earlier reported works.20, 21, 25 On the other hand, the FT XAS spectra of

FeMOF catalyst is significantly different. Thus Fe-PEI catalyst contains both Fe-Nx centers and

Fe nanoparticles, while FeMOF mostly has Fe-N4 species and a very low amount of Fe

nanoparticles. These two catalysts were then further tested for ORR activity in alkaline

electrolyte to test their ability to perform as a ODC in an actual chlor alkali electrolyzer cell.

5.3.2Oxygen reduction reaction activity at high pH electrolyte

Prior to accessing the catalytic activity of the Fe- N-C catalysts, it is important to

optimize the loading that must be maintained during the experiments, since the catalyst loading

plays a major role in ORR catalysis, especially for non PGM catalysts.31 A high loading of

porous catalyst layer hinders mass transport and traps the reactive intermediates (peroxides) in

the catalyst frame.20 On the other hand, a very low loading does not deliver a high limiting

current, and hence can give faulty Koutecky Levich results. Therefore, it is of utmost importance

to optimize the loading of catalyst material in order to get the highest performance. In fact, when

ORR polarization curves were collected for a non PGM catalyst (Fe-MOF) at different loadings

(Fig 5.2(a)). While lower loadings of the catalyst (200 µg/cm2 and 400 µg/cm2) did not exhibit

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much variations in the ORR features, 600 µg/cm2 displayed significant improvement in terms of

onset potential, half wave

Figure 5.2 ORR polarization of (a) FeMOF catalyst in 0.1M NaOH with various mass loadings of the catalyst, (b) Pt/C, FeMOF and Fe-PEI catalysts in 0.1M NaOH, where loading of Pt was 20µg/cm2

geo and the loadings of FeMOF and Pe-PEI catalysts were 600µg/cm2

geo, (c) FeMOF catalyst in different concentration of NaOH and (d) Pt/C, FeMOF and Fe-PEI catalysts in 5M KOH. All data were collected at 1600 rpm at a scan rate of 20 mVsec-1 in O2- saturated electrolytes under room temperature

potential and limiting current densities. Such trend can be expected since a higher loading

increases the number of active sites. However, a further increase in the catalyst loading may limit

mass transport limitations imposed by the thick catalyst layer. Therefore, for all the RDE

experiments, the non PGM catalyst loading was maintained at 600 µg/cm2. Once the optimum

loading of non PGM catalyst was identified, it was further compared with the state of the art Pt/C

catalyst as shown in Fig 5.2(b). Fe-MOF catalyst and Fe-PEI catalyst facilitated ORR at different

146

potentials. While the difference in the limiting current densities can be attributed to the different

nature, density and availability of the active sites, the variation in onset potential is difficult to

interpret. As discussed above, both the catalysts have very similar structures of active sites.

While in Fe-PEI catalyst, there are some Fe nanoparticles along with Fe-N4 coordination, the

FeMOF catalyst contains only Fe-N4 as the active site. The major reason behind such marked

difference in the ORR onset potential can be due to the very different morphologies of the two

catalysts. While Fe-PEI catalyst is supported on carbon black, the carbon in FeMOF is generated

in situ from the precursor materials (phenanthroline, imidazole). As seen in the SEM image in

Scheme 5.1, the morphology of the FeMOF catalyst is highly porous which provides the

reactants better accessibility to the active sites that are situated inside the macropores. As a

further study, we have carried out ORR measurements in NaOH electrolyte with different

electrolyte strength (Figure 5.2(c)). As the concentration of alkali is increased, several

mechanistic changes happen at the electrode- electrolyte interface. The primary effect is

diminished availability of protons to form peroxide from superoxide ions. This causes a change

in shift in the number of electrons transferred from 2 to 1. Moreover, the high concentration of

NaOH reduces the solubility of oxygen drastically, a direct consequence of which can be seen in

Figure 5.2(c), where a gradual decrease in limiting current density occurs on increasing the

electrolyte concentration. The non PGM catalysts are well known for being tolerant towards OH-

adsorption, and is clearly illustrated in Figure 5.2 (d). While the onset potential of Pt/C

cathodically shifts by 50 mV when electrolyte is changed from 0.1M NaOH to 5M NaOH, no

significant change happens for the non PGM Fe- based catalysts. This is the primary reason of

using the non PGM catalysts as oxygen depolarized cathodes in chlorine generation plants.

5.3.3Electrolyzer cell polarization data- Half Cell and Full Cell

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Half-cell polarization curves were collected at 60oC in 5M NaOH, with pure oxygen gas

supplied at the cathode chamber. The platinum metal loading was kept at 1 mg/cm2 while the

non PGM catalyst loading was maintained at 3 mg/cm2. In the half cell set up, oxygen was

generated via oxygen evolution reaction on thenickel foil surface. The gas diffusion cathode for

oxygen

Figure 5.3 Steady state polarization curves collected in (a) half-cell electrolyzer for Pt/C, unsupported Ag, Fe-PEI and FeMOF catalysts in 5M NaOH at 60oC, (inset) Steady state measurement of voltage at a current density of 300 mAcm-2 for Pt/C, Fe-PEI and FeMOF catalysts; (b) full cell brine electrolyzer in 5M NaCl at 60oC for Fe-PEI and unsupported Ag catalysts

reduction reaction was attached to the nickel foil on the cathode side. Fig 5.3(a) displays the

polarization curve collected in the half cell electrolyzer using non PGM catalysts and Pt/C

catalyst. Over the entire operating current density, the non PGM catalysts exhibited better

performance with a reduced overpotential of 150 mV in comparison to the Pt/C catalyst. As a

result of high concentration of OH-, the active sites on Pt surface are blocked due to the strong

adsorption of OH- ions. This phenomenon is heavily attenuated in non PGM catalysts as

observed in the RDE results. Hence, the activity of the non PGM catalysts are restored and the

electrolyzer cell can be operated at considerably lower voltages. In a prior report by Kiroset al.,32

it was observed that non-precious metal catalysts suffered from activity loss at the beginning of

test owing to difficulties related to mass transfer due to the porous structure of the catalyst.

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However, no such trend was observed in our system, and the performances exhibited by all the

catalysts were very consistent throughout the operating conditions. For a better comparison, an

unsupported Ag catalyst designed by DeNora Tech was used as a secondary standard. Both the

non PGM catalysts had better performance compared to the precious metal catalysts. Specially,

FeMOF displayed significant reduction in voltage requirement. This observation can be directly

related to the RDE results where FeMOF had more positive onset of ORR. At a current density

of 300 mA/cm2, Fe-PEI and FeMOF operate at 120 mV and 410 mV lower voltages compared to

Pt/C. The performance of unsupported Ag, on the other hand, was found to be close to that of Fe-

PEI. To estimate the stability and activities of the catalysts, measurements were then taken under

a constant load of 300 mA/cm2 as shown in Figure 5.3 (a), inset at 60oC. Both the non-precious

catalysts exhibited exceptional stability as shown in the data which was collected for over 8

hours in contrast to Pt/C, which suffered from higher voltage losses beyond 2 hrs. of cell

operation. This studies attest the fact that the non-precious catalysts are suitable for use in

electrolyzer cells and can survive the strongly alkaline environment inside an actual plant. Hence

the catalysts were tested in an actual brine electrolyzer cell at DeNora Tech facility. Both

FeMOF and Fe-PEI catalysts were tested against unsupported Ag as the standard catalyst (Figure

5.3(b)). The result obtained from FeMOF catalyst is excluded since this catalyst faced severe

issues regarding flooding in an actual brine electrolyzer. The data for Fe-PEI illustrates the

success of the non PGM catalyst, with the latter operating at a voltage 100 mV lower than that

for unsupported Ag catalyst. This unusual behavior of FeMOF catalyst can be ascribed to the

different nature of the carbon. In a brine electrolyzer cell, the electrolyte (5M NaCl) can

percolate through the highly porous carbon frame of the FeMOF catalyst, resulting blocking

access of O2 to the active sits, which are located inside the pores. In addition to this, the anode

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and cathode chambers are filled with different electrolytes in a brine electrolyzer. An

overpressure on the anode side and strong caustic environment in the cathode side act in a

consolidated manner and aggravate the flooding in the gas diffusion electrode in the cathode

side.33

5.4 Conclusion

In order to establish oxygen depolarized cathodes for chlorine generation plants, it is

desirable to have highly active and robust oxygen reducing catalysts. Here, we have tested two

iron based non-precious catalysts as oxygen depolarized cathodes in brine electrolyzer cells.

Preliminary characterizations indicate that FeMOF catalyst has significantly different

morphology with a very porous carbon matrix. On the other hand, Fe-PEI catalyst has a carbon

base similar to commercial carbon black. These two catalysts, which are synthesized using

completely different approaches, perform differently in RDE experiments. While FeMOF

catalyst has a higher ORR onset potential than Fe-PEI, both exhibit similar kinetics. We believe

that the porous structure of FeMOF catalyst facilitates easier access of reactants and reaction

intermediates to the active sites. Both these catalysts show extreme tolerance to OH- anions in

5M NaOH electrolyte compared to Pt/C catalyst, which undergoes significant cathodic shift in

the ORR onset potential. The two catalysts were then tested in a half cell at 60oC comprising 5M

NaOH as electrolyte. Both the catalysts displayed superior performances compared to Pt/C, or

Ag/C, and were very stable under steady state experiments conducted over a period of 8 hours. It

is intriguing that half-cell electrolyzer using FeMOF operated at a lower voltage requirement

than that of Fe-PEI, which can be directly deduced from the RDE experimental observations.

However, the trends changed once these non-precious catalysts were tested in a full cell brine

electrolyzer. Due to heavy flooding issues, the FeMOF catalyst could not survive in a brine

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electrolyzer. As discussed earlier, we believe this is due to the percolation of 5M NaCl into the

pores of FeMOF catalyst, that practically leads to flooding of the active sites. In RDE and half-

cell electrolyzer, FeMOF performed better than Fe-PEI due to the high dispersity of the active

sites and high porosity. On the other hand, the same factor caused FeMOF catalyst to degrade in

a brine electrolyzer owing to flooding of electrode.

5.5 Acknowledgement

The author would like to thank the collaborator DeNora Tech, OH, USA for their

financial support and assistance with catalyst testing in the brine electrolyzer unit.

5.6 References

(1) Stringer, R.; Johnston, P. Chlorine and the environment: An overview of the chlorine industry; Springer Science & Business Media, 2001.

(2) Ye, L.; Ramström, O.; Mosbach, K. Analytical Chemistry1998, 70, 2789. (3) Ye, J.; Yang, H.; Kim, H. K.; Li, Y. Journal of food science2001, 66, 729. (4) S J, H. Engineering2011, 2011. (5) Gerhard, G.; Google Patents: 1975. (6) Snyder, L. D.; Miller, N. H.; Stavins, R. N. The American Economic Review2003, 93,

431. (7) Beer, H. B. Journal of the Electrochemical Society1980, 127, 303C. (8) Angel, G.; LundÉn, T. Journal of The Electrochemical Society1952, 99, 435. (9) Gullá, A. F.; Gancs, L.; Allen, R. J.; Mukerjee, S. Applied Catalysis A: General2007,

326, 227. (10) Moussallem, I.; Jörissen, J.; Kunz, U.; Pinnow, S.; Turek, T. Journal of Applied

Electrochemistry2008, 38, 1177. (11) Johnson, H. B.; Chamberlin, R. D.; Google Patents: 1981. (12) Subbaraman, R.; Danilovic, N.; Lopes, P.; Tripkovic, D.; Strmcnik, D.; Stamenkovic, V.;

Markovic, N. The Journal of Physical Chemistry C2012, 116, 22231. (13) Schmidt, T.; Stamenkovic, V.; Arenz, M.; Markovic, N.; Ross, P. Electrochimica

Acta2002, 47, 3765. (14) Gong, K.; Du, F.; Xia, Z.; Durstock, M.; Dai, L. science2009, 323, 760. (15) Yu, E. H.; Scott, K.; Reeve, R. W. Journal of Electroanalytical Chemistry2003, 547, 17. (16) Tidswell, I.; Markovic, N.; Ross, P. Journal of Electroanalytical Chemistry1994, 376,

119. (17) Schmidt, T.; Ross, P.; Markovic, N. The Journal of Physical Chemistry B2001, 105,

12082. (18) Roen, L.; Paik, C.; Jarvi, T. Electrochemical and Solid-State Letters2004, 7, A19.

151

(19) Bagotzky, V.; Khrushcheva, E.; Tarasevich, M.; Shumilova, N. Journal of Power Sources1982, 8, 301.

(20) Strickland, K.; Miner, E.; Jia, Q.; Tylus, U.; Ramaswamy, N.; Liang, W.; Sougrati, M.-T.; Jaouen, F.; Mukerjee, S. Nature communications2015, 6.

(21) Li, J.; Ghoshal, S.; Liang, W.; Sougrati, M.-T.; Jaouen, F.; Halevi, B.; McKinney, S.; McCool, G.; Ma, C.; Yuan, X. Energy & Environmental Science2016, 9, 2418.

(22) Aijaz, A.; Masa, J.; Rösler, C.; Xia, W.; Weide, P.; Botz, A. J.; Fischer, R. A.; Schuhmann, W.; Muhler, M. Angewandte Chemie International Edition2016.

(23) Kattel, S.; Wang, G. Journal of Materials Chemistry A2013, 1, 10790. (24) Ramaswamy, N.; Tylus, U.; Jia, Q.; Mukerjee, S. Journal of the American Chemical

Society2013, 135, 15443. (25) Tylus, U.; Jia, Q.; Strickland, K.; Ramaswamy, N.; Serov, A.; Atanassov, P.; Mukerjee,

S. The Journal of Physical Chemistry C2014, 118, 8999. (26) Ramaswamy, N.; Mukerjee, S. Advances in Physical Chemistry2012, 2012. (27) Arruda, T. M.; Shyam, B.; Lawton, J. S.; Ramaswamy, N.; Budil, D. E.; Ramaker, D. E.;

Mukerjee, S. The Journal of Physical Chemistry C2009, 114, 1028. (28) Newville, M. Journal of synchrotron radiation2001, 8, 322. (29) Ravel, B.; Gallagher, K. Physica Scripta2005, 2005, 606. (30) Ankudinov, A.; Ravel, B.; Rehr, J.; Conradson, S. Physical Review B1998, 58, 7565. (31) Ma, S.; Goenaga, G. A.; Call, A. V.; Liu, D. J. Chemistry–A European Journal2011, 17,

2063. (32) Kiros, Y.; Pirjamali, M.; Bursell, M. Electrochimica Acta2006, 51, 3346. (33) Morimoto, T.; Suzuki, K.; Matsubara, T.; Yoshida, N. Electrochimica acta2000, 45,

4257.

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Chapter 6

Summary and Future Directions

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Chapter 6

Summary and Future Directions

This chapter is based on the study of various types of electrocatalytically active metallic

structures and metal surfaces along with their effect on some selected electrochemical reactions,

such as oxygen reduction reaction (ORR) and hydrogen oxidation reaction (HOR). Both these

reactions hold major importance in terms of sustaining electrochemical energy devices such as

fuel cells and electrolyzers. This work encompasses several metal systems such as metal-metal

alloys, metal- metal oxide supports, electrodeposited bimetallic surfaces, core@ shell bimetallic

catalysts and nitrogen coordinated metal catalysts. Such customized metal systems induce

perturbations in the inherent catalytic properties of the metal and this trait has been exploited to

design sustainable electrocatalysts for various energy devices, which are mentioned in the

following sections.

Solid acid fuel cells (SAFC)s are a new class of fuel cells that comprise ceramic proton

conductors as membranes. These ceramic materials attain high protonic conductivity at elevated

temperatures, typically around 250oC. Obtaining an oxygen reducing catalyst becomes

challenging especially at such high temperature. In addition to this, the chemical reactivity

profile of the commonly used proton conductor, i.e., Cesium Dihydrogen Phosphate (CDP) puts

restrictions on the choice of catalyst materials. SAFC relies mainly on unsupported platinum as

catalyst that causes extravagant cost and low mass utilization owing to the terribly low surface

area of unsupported platinum. In Chapter 2, Ru based chalcogenide catalyst supported on carbon

black has been demonstrated as an ORR catalyst that can potentially replace platinum at the

cathode. This strategically synthesized catalyst has a Ru core and Se- modified Ru surface that

preserves the ORR activity of the Ru centers while preventing the Ru core from getting

154

passivated. A ternary addition of Mo metal was found to further enhance the catalytic activity of

the chalcogenide catalyst toward ORR. Detailed structural, electrochemical and in situ

spectroscopic experiments have been studied to elucidate the structure and catalytic mechanism

of this catalyst toward ORR under room temperature as well as at high temperature in an

operating SAFC. Steady state polarization curves collected under operating SAFC exemplify the

potential replacement of unsupported Pt catalysts with the cost effective, durable and efficient Ru

based chalcogenide catalyst.

Formation of alloys with Pt has been vastly studied and different types of catalytic trends

are obtained when Pt is alloyed with different transition materials. These alloys tweak the

structure and properties of Pt in diverse ways that give rise to a wide spectrum of catalytic

activities. Nb is one such transition metal that has been studied widely, though no report of Pt-Nb

alloy catalyst exists. Due to the high negative reduction potential of Nb, formation of Nb metal

or Nb alloys is futile using conventional wet chemistry synthesis routes. In Chapter 3, a highly

active carbon supported Pt-Nb alloy has been reported where Nb is partly present as a zero-

valent metal forming alloy with Pt, and partly present in form of oxides that serve as additional

support material alongside carbon. A detailed structural and mechanistic study of this material

(labelled as PtNb/NbOx-C) has been provided. The unique structure of this catalyst bestows

myriad catalytic activities toward ORR and HOR which have been studied in detail.

In Chapter 4, several Pt based bimetallic systems have been explored to examine their

alkaline HOR activities in alkaline electrolyte. The major debate surrounding the mechanism of

HOR in alkaline electrolytes consist of two theories- namely hydrogen binding energy (HBE)

theory and reactive OHads theory. In order to understand the origin of the slower kinetics of

alkaline HOR, variety of Pt catalysts have been probed mainly to understand the influence of the

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transition metal on the catalytic activity of Pt. Two different types of Pt-Ru systems- (1)

electrodeposited Pt/C-Ru; where Ru resides at the surface and (2) core@ shell Ru/C@ Pt; where

Ru is localized at the core of the catalyst have been studied. Such different types of Ru

distributions are expected to alter the Pt structure and surface in diverse ways. While the

electrodeposited Pt-Ru systems show clear dependence of HOR activity on the surface coverage

of Ru, the core@ shell Pt-Ru catalysts show same activity despite having disparate Pt shell

thicknesses. These results indicate that both HBE and reactive OHads have impact on HOR

activity, and the dominance of one effect on the other is decided by the structure and morphology

of the catalyst. CO stripping reaction further helps to establish the relation between the potential

of zero charge (PZC), role of water activation on Pt surface and the corresponding effects on

hydrogen oxidation. Another transition metal, Mo has been studied by combining with Pt. This

catalyst exhibits superior HOR activity in alkaline electrolyte as compared to Pt/C. By merging

the catalytic activities of Pt-Ru and Pt-Mo systems toward alkaline HOR and CO stripping

processes, the effect of transition metal on Pt has been investigated. It is strongly believed that

the contribution of HBE and reactive OHads is structure dependent and it varies from one system

to other.

In Chapter 5, a brief study on Fe-N-C catalysts has been carried out in alkaline media to

design a non-precious catalyst for chlor alkali electrolyzer cells. Chlorine generation is an energy

intensive industry and in order to curtail the energy requirement, Oxygen Depolarized Cathode

(ODC)s are introduced. However, a real challenging task is to design an ORR active catalyst that

is cheap, efficient and durable in extreme harsh condition of a chlorine generator. Two types of

Fe-N-C catalysts with different morphologies and disparate active sites have been tested under

rotating disk electrode (RDE) conditions to achieve fundamental kinetic detail of the

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electrocatalysts and trends in their catalytic activities have been discussed. The catalysts have

been tested in half-cell electrolyzers and full cell chlorine generator where the overall

performances of the two catalysts have been explained based on their active sites and

morphologies. Both types of Fe-N-C catalysts exhibit superior performance and durability than

state of the art catalyst Pt/C. However, in an actual electrolyzer cell, their performance is mainly

dictated by their morphologies. FeMOF catalyst despite having better activity than Fe-PEI

catalyst, fails to survive inside the brine electrolyzer due to the porous structure that aggravates

flooding.

Surface characterization constitutes an integral part of understanding the fundamentals of

electrocatalysis. For example, the critical interactions between catalyst and membranes need to

be elucidated under operating conditions. Therefore, as an extension of this work, high

temperature Raman spectroscopy should be used to investigate the interactions between catalyst,

support materials and the electrolytes. This would give vital clues that can be used to modify

catalyst structures and membranes. As a continuation of the study of Pt-Ru systems, in situ

spectroscopy experiments can be employed to understand the surface configuration of the

catalyst and identify any changes in the catalyst framework that may affect the catalytic rate.

Such advanced characterization tools will further help to design model catalysts for customized

use in various electrochemical energy devices. In general, basic understanding of catalyst

structure and its interaction with the electrochemical environment is critical in terms of designing

efficient, cost effective and stable catalysts that can be applied in the commercial market. This

will help toward developing a green economy for a longer survival of life on earth.