the rational design of selective electrocatalysts for renewable …349502/fulltext.pdf · the...

The Rational Design of Selective Electrocatalysts for Renewable Energy Devices

by Daniel F. Abbott

B.S. in Chemistry, Framingham State University

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

March 24, 2015

Dissertation directed by

Sanjeev Mukerjee

Professor of Chemistry and Chemical Biology

ii

Acknowledgements

Firstly, I would like to thank Sanjeev Mukerjee for taking me into his lab as a young and

inexperienced researcher. Over the past five years I have had many incredible opportunities and

experiences due to your resourcefulness. You developed my skills as a researcher by challenging

my train of thought and you simultaneously provided me with opportunities and tools to better

myself. I am particularly grateful for the experiences provided at the Army Research Laboratory,

the many trips to Brookhaven National Labs, and for opportunity to study abroad in the Czech

Republic.

My deepest thanks to Petr Krtil for providing me with the opportunity to complete part of my

Ph.D. studies at the J. Heyrovský Institute in Czech Republic. Not only did you continuously

challenge my fundamental understanding of electrocatalysis, but you also helped me make the

most of my experience in the Czech Republic by understanding and supporting my desire to

travel and experience the culture while the opportunity was there. I would also like to extend my

thanks to Hana Hoffmannová and Antonín Trojánek among others for making my experience at

the Heyrovský Institute so memorable.

My most sincere gratitude to my parents, Leon and Cynthia Abbott. You have always been

supportive of my decisions throughout life and have always displayed such pride in my

achievements. I regret that my father is not here to see me today; I know how proud he would

have been. I would also like to thank my brother and sister, James and Emily Abbott, for being

the best siblings I could ask for.

iii

I would like to thank my dissertation committee members, David Budil and Max Diem, for their

time and advice over the years. Many thanks are also extended to the Department of Chemistry

and Chemical Biology, including Cara Shockley, Andrew Bean, and Alex Henriksen for their

continued help.

I would like to give special thanks and recognition to Kara Strickland, Michael Bates, and

Urszula Tylus. We've shared each other's highs and lows over the past 5 years and if it weren't

for the understanding and support of each other then I'm sure we would have had several (more)

mental breakdowns by now. I would also like to thank the rest of my colleagues, past and

present, for the their knowledge, help, and support, including Qingying Jia, Iromie Gunasekara,

Mehmet Nurullaah Ates, Gizem Aysal, Ryan Pavlicek, Shraboni Ghoshal, Jingkun Li, Nagappan

Ramaswamy, Christopher Allen, Matthew Trahan, Jaehee Hwang, William Fowle, and Robert

Allen.

I would like to thank my incredible roommates Chrystine Reilly, Louis DeVita, Noah Grabeel,

Karen Turner, and Rose Fieschko. You've made my living experience in Boston the best that it

could have been. Never did I imagine that I would be living with such a group of fun and caring

people. I'll always be grateful for our family dinners, nights on the porch, Sunday fundays, and

miscellaneous shenanigans on Forbes Street and around JP.

Without the support of my friends I'd be lost. I would like to express my deepest gratitude to my

closest friends Jordan Guertin, Steve Boulay, Kevin Leis, Andrew Fuhrman, David Manchester,

Sean Milan, Robert Duca, Patrick Flaherty, and Dawn Mahoney, for their continued friendship

iv

and unwavering support over the past decade and more. You've all been there to help me

celebrate the best of times and to pick me up during the worst of times. I cannot express how

much I appreciate you all.

I express my deepest gratitude to my friends abroad. Grant Philips and his lovely wife, Kateřina

Phillips, for their loving friendship, generosity, and hospitality during my stay in Czech

Republic. I met you both towards the beginning of my stay and you helped create the lasting

experience I returned with. I'm ever grateful for our many trips to Znojmo, whether for the

Vinobraní or for Christmas; you and your family always made me feel welcome. Anna Havlová

for her love and support throughout my last months in Prague and for joining me in Boston for as

long as you could. My travel companions, Esther Serrano, Liina Ränkel, and Pavel Havlíček for

the many adventures we shared throughout Czech Republic and beyond - and also for those still

to come! I also can't forget my good friends Anna Fodor, Jack Hudson, Peterz Gerhard, Phil

D'Esposito, Veronika Hanakova, and Cesar Figueroa for our numerous dinner dates, game

nights, and late nights.

Finally, I would like to recognize the influence of my former mentors. My highschool chemistry

teacher, Susan Seery, who inspired me pursue my B.S. in Chemistry with her extreme

enthusiasm for chemistry and the physical sciences. My undergraduate research advisor,

Catherine Dignam, for taking the time to work with me and exposing me to the world of

research. Carol Russell for convincing me to pursue my Ph.D. Her genuine concern for her

students and her love of physical chemistry undoubtedly helped lead me to where I am today.

v

Abstract of Dissertation

The rational design of electrocatalysts is paramount to the development of

electrochemical devices. In particular, modifications to the structure and electronic properties of

a particular catalyst can have a strong influence on the activity and selectivity towards various

electrochemical reactions or pathways. In many cases this can lead to a particular reaction

pathway being opened or closed, the formation of intermediates being stabilized or inhibited, the

adsorption of poisonous species being mitigated, or the removal of poisonous species being

promoted. In the this dissertation the design and characterization of catalysts for electrochemical

devices (fuel cells, electrolyzers, and hydrogen pumps) will be discussed with regards to

tailoring the selectivity in order to promote or inhibit certain electrochemical reactions. The

electrochemical reactions of primary interest will include the methanol oxidation reaction

(MOR), the oxygen reduction reaction (ORR), oxygen evolution reaction (OER), and hydrogen

oxidation reaction (HOR).

Chapter 2 introduces and discusses specific issues hindering the advancement of anion-

exchange membrane (AEM)-based direct alcohol fuel cells (DAFCs). Specifically, alkaline

DAFCs experience severe potential losses at low current densities in comparison to the proton

exchange membrane (PEM)-based analogues. Quaternary ammonium cations are known to

specifically adsorb to the catalyst surface and induce electrostatic effects that result in a loss of

active surface area, which limits OH- adsorption and leads to a significant drop in methanol

oxidation current. Characterization of the anode electrode-polymer electrolyte interface is

performed with specific attention paid to the adsorption of quaternary ammonium cations present

in the polymer electrolyte. This research also aims to develop and investigate catalysts that

exhibit improved methanol oxidation activity and high resistance to specific adsorption of

vi

quaternary ammonium in alkaline media. The developed Pt/NiPb/C catalyst shows a significantly

higher electrochemical activity of than that of the commercial Pt/C and PtRu/C electrocatalysts.

In Chapter 3 nanocrystalline ruthenium dioxide and doped ruthenia of the composition

Ru1-xMxO2 (M = Co, Ni, Zn) are prepared and the corresponding oxygen reduction activity and

selectivity is evaluated in alkaline media. In general, the ruthenium based oxides show a strong

preference towards the 2-electron oxygen reduction pathway to hydrogen peroxide at low

overpotentials. However, the selectivity shifts at higher overpotentials towards the complete 4-

electron reduction pathway to H2O. It is shown that Ni- and Co-doped ruthenia continue to

produce significant amounts of peroxide at high overpotentials whereas the Zn- and nondoped

materials prefer the 4-electron reduction pathway. DFT-based analyses on ruthenium based

oxides show that the suppression of the 4-electron reduction pathway on Ni and Co-doped

catalysts can be accounted for by the presence of the Ni and Co cations in the cus binding sites.

Chapter 4 again examines the selectivity of ruthenium based oxides. In this chapter

nanocrystalline Mg-doped ruthenium dioxide catalysts with the formula Ru1-xMgxO2 are

synthesized. The chlorine and oxygen evolution (CER/OER) selectivity of the synthesized

materials in chloride containing media is then related to the catalyst local structure. Through X-

ray absorption spectroscopy (XAS) it is shown that the magnesium ions are not distributed

homogeneously in the material but exist in Mg-rich clusters. The refinement of the Mg EXAFS

functions shows that the Mg-rich clusters contain Mg in a highly strained environment similar to

that of the rutile-type structure at low Mg concentrations (< 10%). As the Mg content is

increased (> 10%) the Mg environment shifts to an ilmenite-type inclusion. Although the Mg

containing catalysts show lower overall CER/OER activities compared with the non-doped

ruthenia, they exhibit a preference for the chlorine evolution process. The observed shift in

vii

selectivity is primarily attributed to the opening of a reaction pathway for chlorine evolution

associated with presence of Mg modified active sites.

In Chapter 5 the reformate tolerance of carbon-supported Pt-based alloys is investigated

under electrochemical hydrogen pump conditions. The CO and CO2 tolerance of the

electrocatalysts under HOR conditions is evaluated and correlated to the composition of the

catalyst. Specifically, the incorporation of sites more oxophilic than Pt at the catalyst surface

promote the activation of water at lower overpotentials. This leads to the removal of adsorbed

CO at lower overpotentials, therefore allowing for higher rates of hydrogen oxidation in the

presence of CO/CO2 and improving the overall cell performance.

Chapter 6 will summarize the conclusions of each research chapter while providing

further insight. Additional methods of experimentation capable of providing valuable

information and extending the current findings will be discussed.

viii

Table of Contents

Acknowledgements ii

Abstract of Dissertation v

Table of Contents viii

List of Figures xiii

List of Tables xvii

List of Abbreviations and Symbols xviii

Chapter 1 Introduction 1

1.1 Importance of Renewable Energy Technology 1

1.2 Electrochemical Energy Conversion and Purification Devices 2

1.2.1 Fuel Cells 2

1.2.2 Electrolyzers 5

1.2.3 Electrochemical Hydrogen Pump 7

1.3 Electrocatalysis 8

1.3.1 Electrochemistry Fundamentals 8

1.3.2 Electrical Double Layer 10

1.3.3 Cyclic Voltammetry 14

1.3.4 Rotating Ring Disk Electrode (RRDE) Method 16

1.4 Differential Electrochemical Mass Spectrometry (DEMS) 18

1.5 X-Ray Absorption Spectroscopy (XAS) 20

1.6 Density Functional Theory (DFT) 26

1.7 Scope of Dissertation 31

ix

1.8 References 33

Chapter 2 Analysis of Double Layer and Adsorption Effects at the Alkaline 38

Polymer Electrolyte-Electrode Interface and the Development of a

Quaternary Ammonium Poisoning Resistant Electrocatalyst for

Methanol Oxidation

2.1 Introduction 38

2.2 Experimental 41

2.2.1 Electrochemical Characterization of the Anode-Polymer 41

Electrolyte Interface

2.2.2 Preparation of Pt/NiPb/C Electrocatalyst 42

2.2.3 XRD and TEM Characterization 43

2.2.4 Electrochemical Characterization of Pt/NiPb/C 43

2.3 Results and Discussion 44

2.3.1 Electrochemical Characterization of the Anode-Polymer 44

Electrolyte Interface

2.3.2 XRD and TEM Characterization 52

2.3.3 Electrochemical Characterization of Pt/NiPb/C 54

2.4 Conclusions 61

2.5 Acknowledgements 62

2.6 References 63

Chapter 3 Oxygen Reduction on Nanocrystalline Ruthenia - Local 68

Structure Effects

3.1 Introduction 68

x

3.2 Experimental 70

3.2.1 Materials Preparation 70

3.2.2 XRD, XPS, and SEM Characterization 71

3.2.3 Electrochemical Measurements 72

3.2.4 DFT Analysis of Oxygen Reduction 73


3.3.1 XRD and SEM Characterization 73

3.3.2 Electrochemical Measurements 76

3.3.3 DFT Analysis of Oxygen Reduction 83

3.4 Conclusions 89


3.6 References 91

Chapter 4 Selective Chlorine Evolution Catalysts Based on Mg-Doped 95

Nanoparticulate Ruthenium Dioxide

4.1 Introduction 95

4.2 Experimental 97

4.3.1 Materials Preparation 97


4.3.3 Local Structure Characterization 99

4.3.4 Electrochemical and DEMS Measurements 100



4.3.2 Local Structure Characterization 105

xi

4.3.3 Electrochemical and DEMS Measurements 111

4.3.4 Discussion 116

4.4 Conclusions 118


4.6 References 120

Chapter 5 Reformate Tolerant Pt-based Catalysts for the 123

Electrochemical Hydrogen Pump

5.1 Introduction 123

5.2 Experimental 125

5.2.1 Catalyst Preparation 125

5.2.2 Physical Characterization 126

5.3.2 Electrochemical Cell Polarization Measurements 127


5.3.1 Physical Characterization 128

5.3.2 Cell Polarization Measurements 131

5.4 Conclusions 139


5.6 References 141

Chapter 6 Dissertation Summary, Conclusions, and Future Directions 146

6.1 Summary 146

6.2 Chapter Synopses and Future Directions 146

6.2.1 Chapter 2 - Analysis of Double Layer and Adsorption 146

Effects at the Alkaline Polymer Electrolyte-Electrode

xii

Interface and the Development of Quaternary

Ammonium Poisoning Resistant Electrocatalyst for

Methanol Oxidation

6.2.2 Chapter 3 - Oxygen Reduction on Nanocrystalline 148

Ruthenia - Local Structure Effects

6.2.3 Chapter 4 - Selective Chlorine Evolution Catalysts 149

Based on Mg-Doped Nanoparticulate Ruthenium

Dioxide

6.2.4 Chapter 5 - Reformate Tolerant Pt-based Catalysts 150

for the Electrochemical Hydrogen Pump

6.3 Concluding Remarks 151

6.4 References 153

xiii

List of Figures

Figure 1.2.1. (a) Schematic of a PEM-based H2/O2 fuel cell. (b) Schematic of an AAEM-based

H2/O2 fuel cell. .................................................................................................................................4

Figure 1.2.2. Schematic of an AAEM-based MeOH/O2 fuel cell. ..................................................5

Figure 1.2.3. Schematic of an acid-based water electrolysis cell. ..................................................6

Figure 1.2.4. Schematic of a PEM-based electrochemical hydrogen pump with a contaminated

H2 feed to the anode and producing purified H2 at the cathode. ......................................................8

Figure 1.3.1. Representation of the double-layer interface based upon the Gouy-Chapman-Stern

model showing the electrode surface (M), the Inner Helmholtz Plane (IHP), Outer Helmholtz

Plane (OHP), and Diffuse Layer. ..................................................................................................11

Figure 1.3.2. Cyclic voltammogram of a polished polycrystalline Pt disk electrode in Ar

saturated 0.1 M KOH. Scan rate of 20 mV s-1

. The anodic scan is shown in blue; the cathodic

scan is shown in red. ......................................................................................................................15


saturated 0.1 M KOH + 10 mM K3Fe(CN)6. Scan rate of 20 mV s-1

. ...........................................16

Figure 1.3.4. Depiction of RRDE tip. ...........................................................................................17

Figure 1.4.1. (a) Schematic of an electrochemical cell designed for use with a DEMS apparatus,

(b) Photo of DEMS apparatus used, and (c) Flow chart of DEMS................................................20

Figure 1.5.1. Schematic of the typical XAS experimental setup. .................................................21

Figure 1.5.2. (a) Example of a normalized XAS spectrum for RuO2 showing both the XANES

and EXAFS regions; (b) EXAFS spectrum of RuO2 shown in k-space; (c) Fourier transform of k-

space EXAFS spectrum to R-space; (d) XANES spectrum of RuO2 showing the pre-edge and

absorption edge regions; (e) Local structure of RuO2 ((110) plane shown in blue) used to model

EXAFS data. ..................................................................................................................................23

Figure 1.5.3. Example of a single photoelectron backscattering event and a multiple

backscattering event that contribute to the oscillating EXAFS signal. ..........................................24

Figure 1.6.1. Example of free energy diagrams for the reduction of O2 on RuO2. The dotted line

represents the equilibrium potential of the reduction of O2 to H2O2. ............................................29

Figure 1.6.2. Volcano plot for the 2-electron (blue) and 4-electron (red) reduction of O2 to H2O2

and H2O, respectively, using the binding energy of OOH as a descriptor. The dotted lines

represent the equilibrium potentials for the reduction products. ...................................................31

Figure 2.1. Percent loss in current density as a function of contaminant concentration in 0.1 M

KOH + 0.5 M MeOH + x mM contaminant. Glassy carbon disk electrode (0.247 cm2) with 5 μL

xiv

of 30% Pt/C ink deposited (Loading = 7.5 ugPt/cm2). Steady state currents obtained from

chronoamperometry at 0.6 V vs. RHE at 900 seconds. .................................................................45

Figure 2.2. Expected potential profile of anode double layer interface in NaOH, TMAOH, and

polymer electrolyte solutions when the electrode potential is more negative than PZC[16]. .......47

Figure 2.3. CV of Fe(CN)63-/4-

redox couple in Ar saturated 0.1M NaOH + 10mM K3Fe(CN)6

solution. Scans performed on glassy carbon disk electrode with 5 μL of 5 wt.% AS4 ionomer

deposited on surface. Geometric surface area = 0.247 cm2. Scan rate = 20 mV/s. ......................49

Figure 2.4. CV of Co(NH3)62+/3+

redox couple in Ar saturated 0.1M NaOH + 10mM

Co(NH3)6Cl3 solution. Scans performed on glassy carbon disk electrode with 5 μL of 5 wt.%

AS4 ionomer deposited. Geometric surface area = 0.247 cm2. Scan rate = 20 mV/s. .................50

Figure 2.5. Expected potential profile of anode double layer interface in the presence and

absence of anion exchange ionomer when the electrode potential is more positive than PZC. ....51

Figure 2.6. Powder XRD patterns for nanocrystalline Pt/C and Pt/NiPb/C samples....................53

Figure 2.7. TEM images of the Pt/NiPb/C catalyst. Insert is higher magnification. ....................54

Figure 2.8. Electrochemical measurement of Pt/C, PtRu/C, and Pt/NiPb/C catalysts in 0.1 M

KOH and 1.0 M methanol at 298 K. Current densities are normalized to the geometric surface

area. (a) CV measurements at a scan rate of 20 mV/s. (b) Chronoamperometric tests with 50 mV

potential steps and displaying the steady-state current value after 180 s. ....................................56

Figure 2.9. Cyclic voltammograms of 40% Pt/C and Pt/NiPb/C catalysts with Nafion or AS4

ionomer used as a catalyst binder. Scans were taken in Ar saturated 0.1 M KOH + 0.5 M

methanol at 298 K at a scan rate of 10 mV/s. Cyclic voltammograms collected at a scan rate of

10 mV/s in Ar saturated 0.1 M KOH are shown in the insert. ......................................................60

Figure 2.10. Electrochemical impedance spectroscopy of 40% Pt/C and Pt/NiPb/C with either

Nafion or AS4 ionomer used as a catalyst binder in 0.1 M KOH + 0.5 M MeOH. EIS data was

collected at E = 0.70 V vs. RHE from 32 kHz to 1 Hz with a amplitude of 10 mV. The insert

shows the high frequency resistance. .............................................................................................60

Figure 3.1. Powder X-ray diffraction patterns for nanocrystalline RuO2 and Ru1-xMxO2 (M = Ni,

Co, Zn) after annealing in air for 1 hour at 400 °C ........................................................................74

Figure 3.2. SEM images of nanocrystalline (a) RuO2, (b) Ru0.90Zn0.10O2, (c) Ru0.80Zn0.20O2, (d)

Ru0.95Ni0.05O2, (e) Ru0.90Ni0.10O2, and (f) Ru0.90Co0.10O2 after annealing at 400 °C in air for 1

hour. ..............................................................................................................................................75

Figure 3.2. ORR polarization curves and ring currents at 1600 rpm for RuO2 and Ru1-xMxO2 (M

= Ni, Co, Zn) electrodes at 20 mV s-1

in O2 saturated 0.1 M NaOH. Ering = 1.1 V vs. RHE. ........77

Figure 3.4. Potential dependence of the average number of electrons transferred during oxygen

reduction on RuO2 and Ru1-xMxO2 (M = Ni, Co, Zn) electrodes. Presented data were calculated

using Koutecky-Levich equation. ..................................................................................................79

xv

Figure 3.5. Phenomenological mechanism of oxygen reduction according to reference[41]. ......80

Figure 3.6. |ID/IR| vs. ω-1/2

plots for (a) RuO2, (b) Ru0.80Zn0.20O2, and (c) Ru0.90Ni0.10O2.

Presented data were extracted from RRDE experiments carried out in O2 saturated 0.1 M NaOH.

........................................................................................................................................................80

Figure 3.3. Potential dependence of the rate constants for the reduction of O2 to H2O (k1), of O2

to H2O2 (k2), and H2O2 to O2 (k3) on nanocrystalline ruthenia based catalysts. The presented data

correspond to experiments carried out in O2 saturated 0.1 M NaOH at 1600 rpm. .......................81

Figure 3.4. The potential of equal rate in 2- and 4-electron reduction for different ruthenia based

catalysts. .........................................................................................................................................82

Figure 3.5. Surface Pourbaix diagram for RuO2. Detailed description of the diagram

construction is given in the supplementary information. ...............................................................84

Figure 3.6. Surface Pourbaix diagram for Ni-doped RuO2. Detailed description of the diagram/s

construction is given in the supplementary information. ..............................................................85

Figure 3.11. Free energy diagrams for the reduction of O2 on three catalytic active sites, the Ru

cus site on conventional ruthenia (green), the cus Ru site on Ni doped RuO2 (magenta) and the

cus Ni site on Ni doped RuO2 (blue). The dotted line represents the equilibrium potential of the

reduction O2 to H2O2. The key difference is the binding of O on the Ni cus site compared to the

Ru cus sites. ...................................................................................................................................86

Figure 3.7. Volcano plot for the 2-electron (blue) and 4-electron (red) reduction of O2 to H2O2


represent the equilibrium potentials for the reduction products. In the case of the Ni-doped

ruthenia the limiting over-potential for both possible reaction sites (Rucus and Nicus) are shown

along with that of conventional ruthenia. ......................................................................................87

Figure 4.1. Powder X-ray diffraction patterns for nanocrystalline RuO2 and Ru1-xMxO2 after

annealing in air for 1 hour at 400 °C. ..........................................................................................102

Figure 4.2. Survey scans (a) of Ru0.8Mg0.2O2 before (1) and after (2) electrochemical

experiments. Panes (b) and (c) contain fitted high resolution spectra of Ru 3d + C 1s (b) and Ru

4s + Mg 2s photoelectrons, respectively. The high resolution spectra of the Ru0.8Mg0.2O2 after

electrochemical experiments do not show significant differences from those plotted in panes (b)

and (c). .........................................................................................................................................103

Figure 4.3. SEM images of nanocrystalline (a) RuO2, (b) Ru0.95Mg0.05O2, (c) Ru0.90Mg0.10O2, and

(d) Ru0.80Mg0.20O2 after annealing in air for 1 hour at 400 °C. ....................................................105

Figure 4.4. EXAFS functions extracted from the X-ray absorption spectra of the Ru1-xMgxO2

(0<x<0.2) measured on the Ru K edge (a) and Mg K edge (b). Actual Mg content is shown in the

Figure legend. ..............................................................................................................................106

Figure 4.5. (a) A typical example of the non-linear least square fit of the Ru EXAFS function of

Ru0.95Mg0.05O2; (b) A typical example of the non-linear least square fit of the Ru EXAFS

xvi

function of Ru0.90Mg0.10O2. The square symbols represent the experimental data, the red line

denotes the best fit. .....................................................................................................................107

Figure 4.6. Linear scan voltammograms of the oxygen evolution on MgxRu1-xO2electrodes

(0<x<0.2) recorded in 0.1 M HClO4 at a polarization rate of 5 mV s-1

. The curve assignment is

given in the Figure legend............................................................................................................111

Figure 4.7. Chloride concentration dependence of the oxygen evolution (a) and chlorine

evolution (b) contributions to the overall current response of MgxRu1-xO2 electrodes to anodic

polarization in chloride containing acid media. The presented values correspond to potentiostatic

experiments at 1.25 V vs. Ag/AgCl. ...........................................................................................112

Figure 4.8. Composition dependence of the Mg doped ruthenia selectivity towards chlorine

evolution. The data correspond to potentiostatic experiments at 1.25 V (Ag/AgCl). The actual

chloride concentrations are given in the Figure legend. ..............................................................113

Figure 4.9. Potential dependence of the Mg doped ruthenia selectivity towards chlorine

evolution. The data were extracted from potentiostatic experiments in 0.1M HClO4 containing 10

mM NaCl (a) and 50 mM NaCl (b). ............................................................................................114

Figure 4.10. Time course of DEMS-based signals of potentiostatically generated oxygen (blue)

and chlorine (red) for RuO2 (a), Ru0.95Mg0.05O2 (b), Ru0.90Mg0.10O2 (c), and Ru0.80Mg0.20O2 (d).

Signals were recorded in 0.1 M HClO4 containing10 mM NaCl; the potential perturbation

corresponded to a step from 0.70 V to 1.25 V vs. Ag/AgCl. .......................................................115

Figure 5.1. Powder XRD patterns for electrocatalyst samples. ..................................................129

Figure 5.2. HRTEM images of the synthesized PtNi/C (a) and PtNiRu/C materials. ................131

Figure 5.3. Anode half-cell polarization curves of an electrochemical hydrogen pump with the

anode and cathode being fed pure H2. Polarization measurements were collected at a cell

temperature of 85°C, 100% relative humidity, and 25/30psi backpressure on the anode/cathode,

respectively. .................................................................................................................................132

Figure 5.4. Anode half-cell polarization curves in an electrochemical hydrogen pump with the

cathode being fed hydrogen and the anode being fed by hydrogen containing 300ppm CO (filled

circles) or a reformate gas mixture containing 100 ppm CO, 15% CO2, 1% CH4, 45% H2, and N2

balance (empty circles). Polarization measurements were collected at a cell temperature of 85°C,

100% relative humidity, and 25/30psi backpressure on the anode/cathode, respectively. ..........135


cathode being fed pure hydrogen and the anode being fed by hydrogen containing 50% molar

CO2. Polarization measurements were collected at a cell temperature of 85°C, 100% relative

humidity, and 25/30psi backpressure on the anode/cathode, respectively. .................................138

xvii

List of Tables

Table 1.2.1. Typical operating conditions and properties for various types of fuel cells. .............2

Table 2.1. XRD and electrochemical results for various catalysts. ...............................................53

Table 3.1. Results of the analysis of the powder diffraction data for RuO2 and doped RuO2

samples. ..........................................................................................................................................75

Table 4.1. Results of the analysis of the powder diffraction data of the MgxRu1-xO2 catalysts .102

Table 4.2. XPS based surface metal content in Ru0.8Mg0.2O2. ....................................................103

Table 4.3. (a) Results of the NLLS fit of the EXAFS functions obtained from the Ru Kedge X-

ray absorption spectra of the MgxRu1-xO2 (0 < x < 0.2) catalysts. CN denotes the coordination

number and d stands for the bonding length; (b) Theoretically conceived local structures of the

Mg in rutile and ilmenite type MgRuO oxides and results of the NLLS fit of the EXAFS

functions obtained from the Mg K edge X-ray absorption spectra of the MgxRu1-xO2 (0<x<0.2)

catalysts. The theoretically conceived data are given in italics. Symbol assignment are the same

as in the case of Table 3a. ............................................................................................................108

Table 5.1. Results of XRD and HRTEM physical characterization. ...........................................130

Table 5.2. Performance and operating conditions of ECHPs reported in the literature ..............133

xviii

List of Abbreviations and Symbols

α Charge Transfer Coefficient, Position of Transition

State Along Reaction Coordinate

βi Width of Diffraction Peak at Half Max Intensity

η Electrochemical Overpotential

ω Rotation Rate

υ Kinematic Viscosity of the Electrolyte

ϕM Inner-potential of the Metal Electrode

ϕS Potential at the Diffuse Layer

ϕ1 Potential at the Inner-Helmholtz Plane

ϕ2 Potential at the Outer-Helmholtz Plane

θ Angle of Incidence/Coverage

λ Wavelength of incident radiation

μ Absorption Coefficient

μ(E) Absorption Coefficient at energy E

μ0(E) Absorption Coefficient at energy E0

χ(E) EXAFS function at energy E

χ(k) EXAFS function at wavenumber k

μm Micrometer (10-6

meters)

σ2 Debye-Waller Factor

Ω Ohms

δ Phase Shift

∆E Reaction Energy

ΔEp Peak Separation Potential

∆G°, ∆G0 Standard Gibbs Free Energy

∆G Gibbs Free Energy

∆GpH Gibbs Free Energy at pH, pH

∆GE Gibbs Free Energy at an electrode potential E

∆H° Standard Enthalpy Change

∆H Enthalpy of Reaction

∆S Entropy Change

∆ZPE Zero Point Energy Difference

Å Angstroms (10-10

meters)

A Geometric Electrode Area, Pre-exponential Factor

atm Atmospheric Pressure

aO Chemical Activity of Oxidant Species

aR Chemical Activity of Reductant Species

C Coulombs

°C Degrees Celsius

CO Concentration of Molecular Oxygen

Concentration of Hydroxide at OHP

Concentration of Hydroxide in bulk solution

cm centimeter

xix

D Diffusion Coefficient

Di Crystallite Domain Size

e Charge of an Electron

e- Electron

eV Electron Volts

E° Standard Electrode Potential

E Electrode Potential, or Energy

E0 Binding Energy, Activation Energy of Reference

Reaction

Ea Anodic Peak Potential, Activation Energy

Ec Cathodic Peak Potential

Eonset Onset Potential

Epeak Peak Potential

Erxn Standard Electrode Potential of Reaction

EPZC Electrode Potential at Potential of Zero Charge

ERing Ring Electrode Potential

EDisk Disk Electrode Potential

f Scattering Amplitude

F Faraday’s Constant

ħ Planck’s Constant

Hz Hertz

i Current Density

ik Kinetic Current Density

ilim Limiting Current Density

ipeak Peak Current Density

i0 Exchange Current Density

it0 True Exchange Current Density

iR Ring Current

iD Disk Current

I0 Intensity of Incident X-ray Beam

It Intensity of Transmitted X-ray Beam

Ir Intensity of Reference X-ray Beam

J Current Density

k photoelectron wavenumber, Boltzmann constant,

or rate constant

k0 Apparent Rate Constant

kt0 True Rate Constant

M Molarity

mA Milliamp

mV Millivolt

me Mass of an Electron

n Number of Electrons Transferred

nm Nanometer (10-9

meters)

N Coordination Number

ppm Part per Million

xx

psi Pounds per Square Inch

qM

Charge on the Metal Electrode

R Universal Gas Constant, Bond Length

T Temperature

V Volts

Welec Electrical Work

x Thickness of Sample

z Charge on an Ion

ads adsorbate

AAEM Alkaline Anion Exchange Membrane

AAEMFC Alkaline Anion Exchange Membrane Fuel Cell

AEI Anion Exchange Ionomer

AFC Alkaline Fuel Cell

ASE Atomic Simulation Environment

BET Brunauer–Emmett–Teller

CA Cyclic Amperometry

CE Counter Electrode

CER Chlorine Evolution Reaction

cus Coordinated Unsaturated Site

CV Cyclic Voltammetry

DEMS Differential Electrochemical Mass Spectrometry

DFT Density Functional Theory

DMFC Direct Methanol Fuel Cell

ECHP Electrochemical Hydrogen Pump

EDAX/EDS Energy Dispersive Analysis of X-rays

EIS Electrochemical Impedance Spectroscopy

EOR Ethanol Oxidation Reaction

EXAFS Extended X-ray Absorption Fine Structure

FC Fuel Cell

fcc Face Centered Cubic

FRA Frequency Response Analyzer

GC Glassy Carbon

GDE Gas Diffusion Electrode

GDL Gas Diffusion Layer

GCS Guoy-Chapman-Stern

hcp Hexagonal Close Packed

HOR Hydrogen Oxidation Reaction

HRTEM High Resolution Transmission Electron Microscope

HT Heat Treated

IHP Inner-Helmholtz Plane

MCFC Molten Carbonate Fuel Cell

MEA Membrane Electrode Assembly

MeOH Methanol

MOR Methanol Oxidation Reaction

MSCV Mass Spectrometric Cyclic Voltammogram

xxi

MWCNT Multi-wall Carbon Nanotubes

NASA National Aeronautics and Space Administration

NHE Normal Hydrogen Electrode

NSLS National Synchrotron Light Source

OCP Open Circuit Potential

OER Oxygen Evolution Reaction

OHP Outer-Helmholtz Plane

ORR Oxygen Reduction Reaction

PAFC Phosphoric Acid Fuel Cells

PBI Polybenzimidazole

PSA Pressure Swing Adsorption

PTFE Polytetrafluoroethylene

PZC Potential of Zero Charge

PZFC Potential of Zero Free Charge

PZTC Potential of Zero Total Charge

PEM Proton Exchange Membrane

PEMFC Proton Exchange Membrane Fuel Cell

QA+ Quaternary Ammonium Ion

RH Relative Humidity

RHE Reversible Hydrogen Electrode

RPM Revolutions Per Minute

RE Reference Electrode

RRDE Rotating Ring-Disk Electrode

SEM Scanning Electron Microscopy

SHE Standard Hydrogen Electrode

SOFC Solid Oxide Fuel Cell

TEM Transmission Electron Microscopy

upd Underpotential Deposition

WE Working Electrode

XANES X-ray Absorption Near Edge Spectroscopy

XAS X-ray Absorption Spectroscopy

XRD X-ray Diffraction

XPS X-ray Photoelectron Spectroscopy

1

Chapter 1

Introduction

1. Introduction

1.1 Importance of Renewable Energy Technology

The production of energy from clean and renewable sources is crucial to the future

development of human civilization. Given the fact that the burning of traditional fossil fuels is

contributing to the ever-growing issue of climate change and is generally unsustainable, the

adoption of renewable and sustainable energy sources is of paramount importance[1]. Hydrogen

is an alternative source of energy that can be obtained from several resources, such as the

electrolytic splitting of water or the reforming of hydrocarbons (e.g. natural gas, propane,

methanol). When hydrogen is produced and reclaimed using the energy obtained from renewable

sources such as wind, solar, hydro, etc. it offers a means of escaping the paradigm of burning

carbon-based fossil fuels. The clean energy expended to produce hydrogen can then later be

reclaimed using fuel cell technology for anything from portable electronics, to automotive

vehicles, to grid-scale power. This can ultimately lead to a reduction of CO2 emissions and oil

dependency. To develop and influence a hydrogen economy free of fossil fuels, efficient and cost

effective methods for hydrogen production, purification, storage, transport, and conversion are

imperative.

The following introductory sections provide a brief overview of the electrochemical

technologies targeting the production, purification, and consumption of hydrogen (electrolyzers,

electrochemical hydrogen pumps, and fuel cells). Moreover, they introduce the fundamental

concepts and techniques required for the development and characterization of electrocatalysts for

2

these devices. The primary focus of subsequent chapters will be on the rational design of

electrocatalysts tailored to address specific issues hindering the development of hydrogen-based

technologies.

1.2 Electrochemical Energy Conversion and Purification Devices

1.2.1 Fuel Cells

Fuel cells are electrochemical devices that convert chemical energy into electrical energy.

There are several major types of fuel cells, including Proton Exchange Membrane Fuel Cells

(PEMFCs), Phosphoric Acid Fuel Cells (PAFCs), Solid Oxide Fuel Cells (SOFCs), and Alkaline

Anion Exchange Membrane Fuel Cells (AAEMFCs). The electrolyte, charge carrier, and

operating temperature of a fuel cell are dependent upon the type of fuel cell being considered.

The major types of fuel cells are summarized in Table 1.2.1[2].

Table 1.2.1. Typical operating conditions and properties for various types of fuel cells.

PEMFC AAEMFC AFC PAFC MCFC SOFC

Electrolyte Polymer

membrane

Polymer

membrane

Liquid

KOH

Liquid

H3PO4

Molten

carbonate Ceramic

Operating Temp (°C) 80 60 60-220 200 650 600-1000

Charge Carrier H+ OH

- OH

- H

+ CO3

2- O2

-

One focus of the work presented in the following chapters is the electrocatalysis

encountered in direct methanol-based AAEMFCs. AAEMFCs offer several advantages over the

PEMFC analogues, including the use of non-noble metals, reduced electro-osmotic drag,

improved kinetics in the case of ORR and MOR, and the lack of carbonate precipitation issues[3-

6]. The progressive development of alkaline anion exchange polymers has led to a growing

interest in the development of AFCs that use an AEM in the past 10 years. This technology

3

clearly surpasses the liquid KOH electrolyte used in previous systems[7, 8], including those used

by NASA in the Gemini space programs[9, 10].

Both PEMFCs and AAEMFCs can use H2 (or small alcohols) as a fuel and O2 or air as an

oxidant. The difference lies within the charge carrier and the electrochemistry at both the anode

and the cathode. In a PEMFC, hydrogen is oxidized to protons at the anode, which are then

transported across the PEM to the cathode where they combine with reduced oxygen to form

H2O (See Figure 1.2.1a). In an AAEMFC, hydroxide ions are produced at the cathode instead of

producing protons at the anode. In this case molecular oxygen is reduced to hydroxide ions,

which are then transported across the AAEM to the anode where they combine with hydrogen to

form water (See Figure 1.2.1b). The half cell reactions for anode and cathode for the PEMFC

and AAEMFC can be seen below[11]:

PEMFC:

Anode: H2 → 2H+ + 2e

- E° = 0.000 V vs. SHE (1.2.1)

Cathode: 1/2O2 + 2H+ + 2e

- → H2O E° = 1.229 V vs. SHE (1.2.2)

AAEMFC:

Anode: H2 + 2OH- → 2H2O + 2e

- E° = -0.828 V vs. SHE (1.2.3)

Cathode: 1/2O2 + H2O + 2e- → 2OH

- E° = 0.401 V vs. SHE (1.2.4)

4

Figure 1.2.1. (a) Schematic of a PEM-based H2/O2 fuel cell. (b) Schematic of an AAEM-based

H2/O2 fuel cell.

A schematic of an AEM-DMFC can be seen in Figure 1.2.2. In this case, the methanol

oxidation reaction (MOR) is the primary reaction that occurs at the anode (Eq. 1.2.8)[12], while

oxygen reduction (Eq. 1.2.4) remains the primary reaction at the cathode. It should be noted that

partial oxidation of methanol can occur, which produces unwanted intermediates such as formate

and carbon monoxide as shown in Eq. 1.2.5 and Eq. 1.2.6, respectively. Alkaline Exchange

Membrane Direct Methanol Fuel Cells (AEM-DMFCs) will be discussed in more detail in

Chapter 2.

CH3OH + OH- → CH3Oads + H2O + e

- (1.2.5)

CH3Oads + 3OH- → COads + 3H2O + 3e

- (1.2.6)

COads + 2OH- → CO2 + H2O + 2e

- (1.2.7)

CH3OH + 6OH- → CO2 + 5H2O + 6e

- E° = -0.810 V vs. SHE (1.2.8)

5

Figure 1.2.2. Schematic of an AAEM-based MeOH/O2 fuel cell.

1.2.2 Electrolyzers

Water electrolysis essentially reverses the direction of the electrochemical process

employed in H2/O2 fuel cells. Electrolyzers use electrical energy to drive chemical reactions, i.e.

convert electrical energy to chemical energy. Water electrolysis, the splitting of water into its

constituent elements, H2 and O2, has been known for over 200 years and has been used as a

method to produce hydrogen industrially since 1888. Different electrolytic processes are also

widely used, such as HCl electrolysis to produce hydrogen and chlorine from concentrated HCl

solutions, or the chlor-alkali process used to produce NaOH and Cl2 from concentrated brine

(NaCl) solutions. In water electrolysis, molecular oxygen is produced at the anode via the

oxidation of water. This in turn generates protons, which then cross the membrane to the cathode

where they recombine to form hydrogen gas (See Figure 1.2.3). The half cell reactions are

presented below:

6

Anode: H2O → 1/2O2 + 2H+ + 2e

- E° = 1.229 V vs. SHE (1.2.9)

Cathode: 2H+ + 2e

- → H2 E° = 0.00 V vs. SHE (1.2.10)

It should be noted that water electrolysis is sensitive to the presence of contaminants in

solution, particularly that of chloride. Given the relatively close standard electrode potentials of

both oxygen evolution (Eq. 1.2.9) and chlorine evolution (Eq. 1.2.11) and the fact that OER

requires a significant overpotential, one must anticipate the simultaneous production of both

oxygen and chlorine at reasonable operating potentials when chloride is present. Since chloride

contamination is one of the more serious issues in conventional water electrolysis, the

development of selective catalysts for OER/CER will be discussed in Chapter 4.

Anode: 2Cl- → Cl2 + 2e

- E° = 1.358 V vs. SHE (1.2.11)

Figure 1.2.3. Schematic of an acid-based water electrolysis cell.

7

1.2.3 Electrochemical Hydrogen Pump

Hydrogen purification is crucial to the advancement and integration of hydrogen-based

fuel cells. When hydrogen is produced via gassification of coal or natural gas, it typically

requires several stages of purification before it can be consumed commercially. Common

contaminants from steam reforming processes include CO2, CO, and methane. CO is particularly

problematic when considering the Pt-based catalysts typically employed in fuel cell applications

since it binds strongly to the Pt surface, effectively blocking the surface sites available for

hydrogen oxidation. CO2 is also a problematic contaminant in fuel cell applications as it can be

converted to CO via the reverse water gas shift reaction (See Eqs. 1.2.12 and 1.2.13). Current

state of the art technologies for hydrogen purification include pressure swing adsorption (PSA),

palladium membrane filters, cryogenic separation, and chemical adsorption methods based on

alkaline or ammine scrubbing[13-15]. The focus here will be on the electrochemical purification

of hydrogen, which was first pioneered by Sedlak in 1981[16].

CO2 + H2 ⇌ CO +H2O (1.2.12)

CO2 + 2M-Hads ⇌ 2M-COads + H2O + M (1.2.13)

An electrochemical hydrogen pump (ECHP) is very similar to traditional proton

exchange membrane fuel cells (PEMFCs) in terms of electrochemistry and in terms of hardware.

The electrochemical reaction at the anode is hydrogen oxidation (Eq. 1.2.14), while the reaction

at the cathode is hydrogen evolution (Eq. 1.2.15) instead of ORR. An electrochemical hydrogen

pump operates on the principle that the hydrogen oxidation and reduction reactions are extremely

facile on Pt, requiring a very small overpotential to drive the overall reaction. This may offer a

cost effective means of purifying hydrogen from various sources given that PSA and cryogenic

8

separation methods are typically energy intensive and require several steps. A schematic of a

hydrogen pump can be seen in Figure 1.2.4 in which a contaminated stream of hydrogen is fed to

the anode while purified hydrogen is produced at the cathode. The focus in subsequent chapters

will be on the development of CO and CO2 tolerant Pt-based catalysts to be used at the anode in

an ECHP.

Anode H2 → 2H+ + 2e

- E° = 0.000 V vs. SHE (1.2.14)

Cathode 2H+ + 2e

- → H2 E° = 0.000 V vs. SHE (1.2.15)

Figure 1.2.4. Schematic of a PEM-based electrochemical hydrogen pump with a contaminated

H2 feed at the anode and producing purified H2 at the cathode.

1.3 Electrocatalysis

1.3.1 Electrochemistry Fundamentals

Electrochemistry is a branch of chemistry that is largely concerned with the

thermodynamics and kinetics of chemical redox reactions (i.e. the Gibb's free energy associated

9

with the oxidation or reduction of a chemical species) in relation to the electrical work associated

with the addition or removal of electrons. The Gibb's Free Energy of a system can be described

by the following equation[2]:

(1.3.1)

where ∆G° is the standard Gibb's Free Energy, R is the ideal gas constant, T is the absolute

temperature, and aO and aR are the chemical activity of oxidant and reductant species in a

reaction such as that shown below:

O + n → R (1.3.2)

The change in Gibb's Free Energy of a system is also equal to the electrical work performed by

the system, or the negative of the work done on the system:

(1.3.3)

where n is the number of electrons transferred, F is Faradays constant (96,485 C/mol), and Erxn is

the standard electrochemical potential associated with a given reaction. By substituting Eq. 1.3.3

into Eq. 1.3.1, it is possible to derive the Nernst equation (See Eq. 1.3.4), which relates the half-

cell reduction potential to the redox composition of the system at a given temperature and

pressure.

(1.3.4)

where E is the half-cell reduction potential and E° is the standard electrode potential.

The kinetics of a standard one electron transfer reaction, such as that shown in Eq. 1.3.2,

can be described using the Butler-Volmer equation:

(1.3.5)

10

where i0 is the exchange current density, α is the transfer coefficient, F is Faraday's constant, η is

the electrochemical overpotential, R is the ideal gas constant, and T is the absolute temperature.

The Butler-Volmer is a fundamental equation in electrochemistry that describes the relationship

between the current and overpotential in the absence of mass-transport limitations. This equation

can be simplified in high overpotential regions to yield the Tafel equation:

(1.3.6)

A plot of η vs. ln(i) then allows for valuable information regarding the charge transfer coefficient

and exchange current density to be extracted.

1.3.2 Electrical Double Layer

The structure of the electrical double layer and how it is altered at different electrolyte/Pt

interfaces will be considered in Chapter 2. The development of a suitable model to accurately

describe the interface between the electrode surface and the electrolyte has taken place over the

past 150 years. The structure of the double layer is generally believed to consist of three distinct

regions: the Inner Helmholtz Plane (IHP), the Outer Helmholtz Plane (OHP), and the diffuse

layer (See Figure 1.3.1). Species located in the IHP are said to be specifically adsorbed and can

only approach the electrode surface to a finite distance, x1. This distance is defined by the locus

of the centers of the specifically adsorbed ions. Specifically adsorbed ions in the IHP experience

strong short-range interactions with the charged metal surface. The Outer Helmholtz plane

(OHP) is defined at a distance, x2, which is determined by how close ions or molecules can

approach the electrode before desolvating and adsorbing to the surface[11]. This is considered

the plane of closest approach and is limited by the thickness of the compact layer. Ions found in

the OHP are said to be nonspecifically adsorbed since their interaction with the charged metal

11

surface occurs only through long-range electrostatic forces. Beyond the OHP, nonspecifically

adsorbed ions are distributed throughout a three dimensional region known as the diffuse layer,

which extends from the OHP to the bulk solution. Figure 1.3.1 shows the three regions of the

double layer for a negatively charged electrode surface.

Figure 1.3.1. Representation of the double-layer interface based upon the Gouy-Chapman-Stern

model showing the electrode surface (M), the Inner Helmholtz Plane (IHP), Outer Helmholtz

Plane (OHP), and Diffuse Layer.

The point at which an electrode surface is free of adsorbates and free of excess charge

(i.e. completely neutral) is defined as the potential of zero free charge (PZFC). The potential at

which any excess charge on the electrode surface is balanced by ions in solutions (specifically

12

adsorbed or not) is defined as the potential of zero total charge (PZTC). It is important to note

that most electrode surfaces are not well defined (i.e. polycrystalline electrodes having a certain

roughness). In this case it is generally accepted that the potential of zero charge (PZC) is the

potential at which the differential capacitance experiences a minimum in dilute electrolyte

solutions[11].

In the absence of specific adsorption, the voltage drop between the metal electrode and

the solution occurs across the compact and diffuse layers, as described by the Gouy-Chapman-

Stern model, with the OHP defined by the locus of centers of hydrated ions[11]. It is important to

note that the thickness of the double layer is dependent upon the concentration of the electrolyte

solution. At high electrolyte concentrations, the double layer becomes more compact because

there is a higher concentration of ions within close proximity of the surface to form the compact

layer. The effect is that the potential drop across the double layer occurs primarily in the compact

layer with little contribution from the diffuse layer. At more dilute electrolyte concentrations, or

ones where the ions have limited mobility, a greater thickness is needed in order to accumulate

enough charge to counterbalance the charge of the metal surface. The effect is that the potential

drop across the diffuse layer becomes important. In this case the electrode reaction is no longer

driven by the potential difference ϕM - ϕS, where ϕM is the potential of the metal electrode and

ϕS is the potential of the solution. Instead, the electroactive species at the OHP experiences a

potential difference (ϕM - ϕ2), where ϕ2 is the potential at the OHP. In addition, charged

electroactive species migrate under the influence of the electric field within the diffuse layer. The

concentration of electroactive species located at the OHP will then differ from that immediately

outside the in the diffuse layer by a factor of exp(-zϕ2F/RT), where z is the charge on the ion, F is

Faraday’s constant, R is the ideal gas constant, and T is the absolute temperature.

13

In general, cations will typically be drawn to the electrode surface when the surface has a

negative charge (E < EPZC) whereas anions will be attracted to a positively charged electrode

surface (E > EPZC). Similarly, cations will be repelled from a positively charged surface (E >

PZC) and anions will be repelled from a negatively charged surface (E < PZC). Consequently, in

the case of E < EPZC the positive charge contributed by cations in the IHP will create a potential

shift in ϕ2 such that species in the OHP plane will experience a potential more positive than ϕ1

(See Figure 1.3.1). Similarly, at positive potentials anions will populate the IHP creating a

negative shift in ϕ2. These effects can have a large influence on the kinetics of a given reaction

since ϕ2 varies with the electrode potential and electrolyte concentration.

The described effect of the double layer on kinetics is often referred to as either the

Frumkin effect or the ϕ2 effect. In general, the overall result is that i0 and k

0 are dependent upon

the variation of ϕ2 with E - EPZC and concentration[17-19]. The Frumkin correction can be used

to relate the apparent rate constant with the true rate constant:

(1.3.6)

where k0 is the apparent rate constant,

is the true rate constant, α is the transfer coefficient, ϕ2

is the potential at the OHP, F is Faraday's constant, R is the ideal gas constant, T in the absolute

temperature, and z is the charge on the ion. The potential driving the electrode reaction at the

OHP is corrected for by the first exponential term. This is also valid for outer sphere electron

transfer reactions. The concentration profile for any charged species is corrected for by the

second exponential term. As follows from Eq. 1.3.6, the true exchange current density can also

be determined:

(1.3.7)

where i0 is the apparent exchange current density and

is the true exchange current density.

14

While ϕ2 can be predicted using the Gouy-Chapman-Stern model in the absence of

specific adsorption, it becomes more complex in the presence of specific adsorption of charged

species e.g. quaternary ammoniums. The adsorption of electro-inactive species create significant

changes in the structure of the double layer that differ from the non-adsorbing case[11, 17, 19].

In this case the actual potential at the OHP cannot be clearly defined. Instead a qualitative

assessment must be made. Furthermore, the adsorption of electro-inactive species results in

blockage of the electrode surface. This inhibits the reaction rate independent of ϕ2. In general,

the blockage effect only becomes dominant at high concentrations of the electro-inactive species.

At low concentrations of the adsorbing species the reaction rate is affected by the ϕ2 effect.

1.3.3 Cyclic Voltammetry (CV)

Cyclic voltammetry is perhaps one of the most widely used techniques used for

electrochemical characterization. Cyclic voltammetry is commonly used in the characterization

of electrocatalytic materials and the study of oxidation/reduction processes that the occur at the

electrode-electrolyte interface. Consider a polished polycrystalline Pt electrode in Ar saturated

0.1 M KOH solution, such as that shown in Figure 1.3.2. Saturating the electrolyte with argon, or

any inert gas, allows for the Pt surface to be studied in detail without the interference of

dissolved oxygen. The cathodic scan in the cyclic voltammogram reveals information regarding

the reduction processes occurring at the Pt surface while the anodic scan reflects the oxidation

processes. Generally speaking, a typical CV of a polycrystalline surface can be divided into three

regions. The features observed in Region I are characteristic of the adsorption and desorption of

hydrogen. Region II is referred to as the "double-layer" region. This is due to the fact that there

are no faradaic reactions occurring in this potential regime; only the double-layer capacitance is

15

observed. Finally, Region III is defined by the formation and reduction of surface oxides. As one

scans anodically, features corresponding to the formation of PtOH and PtO2 are observed and

eventually the evolution of O2 can be seen at potentials positive of 1.6 V vs. RHE. On the

cathodic scan, features corresponding to the reduction of surface oxides can be seen in Region

III. Again Region II corresponds to the absence of faradaic processes on Pt. Continuing to

cathodically scan into Region I, the formation of Pt-H is observed before the eventual formation

and desorption of H2 at potentials negative of 0.05 V vs. RHE.


saturated 0.1 M KOH. Scan rate of 20 mV s-1

. The anodic scan is shown in blue; the cathodic

scan is shown in red.

Cyclic voltammetry is also used in the study of electroactive species dissolved in solution

at the electrode surface. Consider again a polycrystalline Pt electrode submerged in Ar saturated

0.1 M KOH, however, this time with the addition of 10 mM K3Fe(CN)6. The outer-sphere

oxidation/reduction behavior of this species can be examined by scanning the appropriate

16

potential regime as shown in Figure 1.3.3. The peak corresponding to the oxidation of

[Fe(CN)6]4-

to [Fe(CN)6]3-

can be seen at 1.2210 V. Similarly, the peak corresponding to the

reduction of [Fe(CN)6]3-

to [Fe(CN)6]4-

can be seen at 1.1398 V. The separation of peaks is

approximately 81 mV, which corresponds a one-electron transfer reaction. This redox reaction is

said to exhibit a Nernstian behavior. It should be noted that a redox couple displaying ideal

Nernstian behavior would have a peak separation of 56.5 mV/n, where n is the number of

electrons transferred.


saturated 0.1 M KOH + 10 mM K3Fe(CN)6. Scan rate of 20 mV s-1

.

1.3.4 Rotating Ring Disk Electrode (RRDE) Method

The Rotating Ring Disk Electrode (RRDE) is a common method used in electrochemistry

to study the products and intermediate species generated during an electrochemical reaction

under convective diffusion as the potential of the working electrode is varied. This technique has

been used extensively to study the oxygen reduction reaction (ORR) due to its ability to detect

17

any hydrogen peroxide intermediate species produced[20-23]. A schematic of the rotating ring

disk apparatus can be seen in Figure 1.3.2. Essentially, the electrode tip consists of a disk

electrode (carbon, platinum, gold, or other material) embedded in the center of a Teflon housing.

Surrounding the disk electrode is a secondary ring electrode, the potential of which is controlled

separately. The ring electrode material is typically either gold or platinum. The tip is fitted into

the shaft of the rotator, which controls the speed at which the electrode can be rotated and thus

controls the transport of electroactive species to the electrode surface. The electrodes are

connected to a potentiostat/galvanostat that is used to control the potential of each electrode

while monitoring the current or vice versa. The tip of the electrode is submerged in an electrolyte

solution containing the supporting electrolyte and electroactive species. Ink suspensions

containing the catalyst of interest can be made and deposited on the disk of the electrode to

provide a thin film of the catalyst under investigation.

Figure 1.3.4. Depiction of RRDE tip.

18

The advantage of the RRDE setup is that it allows an electroactive species to be

oxidized/reduced at the disk electrode and then the intermediate species generated are

reduced/oxidized at the ring electrode. The transport of species to the disk electrode is controlled

by varying the rotation rate, which also controls the rate at which species are swept away radially

from the disk to the secondary ring electrode. By controlling the rotation rate of the electrode and

monitoring the disk current, one can apply the Koutecký-Levich equation to obtain information

about the number of electrons transferred in a reaction or the diffusion coefficient of the

electroactive species in solution[11]:

(1.3.8)

where ik represents the kinetic current, n is the number of electrons transferred, F is Faraday's

constant, A is the electrochemically active surface area, DO is the diffusion coefficient of the

reactant species, ω is the angular rotation rate of the electrode, ν is kinematic viscosity of the

electrolyte, and CO is the concentration of the reactive species.

1.4 Differential Electrochemical Mass Spectrometry (DEMS)

Differential Electrochemical Mass Spectrometry (DEMS) is a unique technique that

allows for the on-line detection and identification of volatile products and intermediate species of

faradaic reactions. When used in parallel with cyclic voltammetry, DEMS is able to provide on-

line detection of species generated at the electrode surface as a function of potential, resulting in

a mass spectrometric cyclic voltammogram (MSCV)[24]. Other electrochemical techniques,

such as chronoamperometry or chronopotentiometry, can also be combined with DEMS.

Reaction products can be detected somewhat quantitatively, allowing for the distinction between

the main and side reactions products.

19

The DEMS setup is fairly simple, consisting of a 3-electrode electrochemical cell (See

Figure 1.4.1a) separated via a porous PTFE membrane from the apparatus (See Figure 1.4.1b),

which consists of turbomolecular dragging pump station, a rotary vane vacuum pump, and a

quadrupole mass spectrometer. A flow chart of the typical DEMS setup can be seen in Figure

1.4.1c.

Species generated at the electrode surface diffuse through a porous PTFE membrane via

needle valve to a vacuum environment where they are then transported to the mass spectrometer.

The chemical species are ionized in the ion source before being separated by mass and

transported to the detector. Simultaneously, the potential of the working electrode can be held

constant or varied to provide the necessary electrochemical perturbation. DEMS has been used

among others to study the electrochemical decomposition of alcohols[25, 26], quantifying

simultaneous oxygen and chlorine evolution[27-30], and the exchange of oxygen between the

electrode catalyst material and solvent[31].

20

Figure 1.4.1. (a) Schematic of an electrochemical cell designed for use with a DEMS apparatus,

(b) Photo of DEMS apparatus used, and (c) Flow chart of DEMS.

1.5 X-ray absorption Spectroscopy (XAS)

X-ray absorption spectroscopy is a technique that relies on the use of high intensity X-

rays, typically generated by a synchrotron light source, to excite the core level electrons (K, L,

and/or M shells) of a particular element in a sample from the ground state into an excited

electronic state or into the continuum[32]. Promoting an electron to an excited state leaves an

empty core shell hole behind, which is quickly filled (~10-15

s) by an electron from a higher

energy orbital[33]. The relaxation of an electron from a higher energy orbital to a lower one

results in the release of energy in the form of fluorescence. The amount of energy released is

equal to the energy difference between the two orbitals and is element specific, therefore the

21

fluorescence radiation emitted as a result of the transition is characteristic of the absorbing

element.

Not all of the incident X-ray radiation is absorbed by the sample and the remainder is

transmitted through the sample. Absorption is a process of finite probability and the amount of

X-ray radiation transmitted through the sample is exponentially dependent upon the thickness of

the sample and the X-ray absorption coefficient:

(1.5.1)

where It is the intensity of the transmitted radiation, I0 is the intensity of the incident radiation, x

is the thickness of the sample, and μ is the X-ray absorption coefficient. The X-ray absorption

coefficient describes the probability that an X-ray will be absorbed by a given element or sample.

This is an element-specific quantity that increases with the binding energy of core-level electrons

and thus increases with atomic number.

The typical setup for an XAS experiment is shown in the schematic below:

Figure 1.5.1. Schematic of the typical XAS experimental setup.

The incident X-ray radiation passes through an ionization chamber that measures the radiation

flux, or initial intensity, I0. The radiation then hits the sample, producing photoelectrons (X-ray

fluorescence) that can be detected by a fluorescence detector set up at a 45° angle from the

sample. The remaining radiation that is transmitted through the sample then enters a second

22

ionization chamber to measure the transmitted radiation, It. Typically a reference foil is placed

after the secondary ionization chamber and a third ionization chamber is used to measure the

intensity of the X-rays transmitted through the reference material, Ir. The reference signal is

important in ensuring that the position of the monochromator does not shift during the

measurement. By rearranging Eq. 1.5.1, the X-ray absorption coefficient can be determined

based upon the measured intensity of the incident radiation and the intensity of the of radiation

transmitted:

(1.5.2)

Similarly, μ can also be determined from the intensity of the X-ray fluorescence since X-ray

absorption is proportional to the generation of a core hole and the subsequent relaxation process:

(1.5.3)

As can be inferred from the above equation, a sharp increase in μ is observed when the incident

radiation is equal to the energy required to excite promote a core level electron. This can be

observed in a typical XAS spectrum (See Figure 1.5.2b) and the feature is referred to as the

absorption edge. The absorption edge energy is specific to the binding energy of the electrons for

each element.

23

Figure 1.5.2. (a) Example of a normalized XAS spectrum for RuO2 showing both the XANES

and EXAFS regions; (b) EXAFS spectrum of RuO2 shown in k-space; (c) Fourier transform of k-

space EXAFS spectrum to R-space; (d) XANES spectrum of RuO2 showing the pre-edge and

absorption edge regions; (e) Local structure of RuO2 ((110) plane shown in blue) used to model

EXAFS data.

XAS spectra are generally divided into two regions: the X-ray Absorption Near-Edge

Structure (XANES) and the Extended X-Ray Absorption Fine Structure (EXAFS) (See Figure

1.5.2). The XANES portion of the spectrum can be used to obtain information regarding the

oxidation state and coordinate chemistry of the absorbing atoms. EXAFS is used to obtain

information about the local molecular structure of the element of interest within a sample. The

XANES region ranges from approximately -50 to +100 eV of the absorption edge.

The EXAFS region begins at approximately 50 eV+ of the absorption edge. The EXAFS

region is sensitive to the surrounding environment because at excitation energies higher than the

24

absorption edge the photoelectron is ejected from the absorbing atom into the continuum. The

ejected photoelectrons are of dual nature and can be described as spherical waves propagating

from the absorbing atom that scatter off of the surrounding atoms. The phase of the scattered

photoelectron wave is determined by its wavelength and the distance between the propagating

atom and the scattering atoms. The minima and maxima observed in the oscillating EXAFS

signal are dependent upon the phase of the outgoing photoelectron wave and the scattered wave

at the absorbing atom. A minimum occurs when the waves are out of phase; a maximum occurs

when the waves are in phase. There are generally many scattering paths a photoelectron can take

depending upon the nature of neighboring atoms and the distances between them, although some

pathways are degenerate (See Figure 1.5.3). Neighboring atoms at the same interatomic distance

from the absorbing atom all contribute to the same component of the EXAFs signal. Any

neighboring atoms at the same distance from the absorbing atom are all part of the same shell

and the total number of atoms in a shell determines the coordination number.

Figure 1.5.3. Example of a single photoelectron backscattering event and a multiple

backscattering event that contribute to the oscillating EXAFS signal.

25

The EXAFS function, χ(E), describes the oscillatory part of the of the absorption

spectrum (see Figure 1.5.2) and is defined as:

(1.5.4)

where µ(E) is the observed absorption coefficient and µ0(E) is the absorption coefficient at an

energy equivalent to the absorption edge, E0. EXAFS is more commonly interpreted in terms of

the emitted photoelectron energy rather than the energy of the incident X-ray radiation. The

transformation to k-space (See Figure 1.5.2b), where k is the wave number of the photoelectrons,

takes place through the following relationship:

(1.5.5)

where me is the mass of an electron and is Planck's constant. The EXAFS can then be

expressed as the sum of the contribution from all scattering paths of the photoelectron:

(1.5.6)

where Nj is the coordination number of the neighboring atom, Rj is the distance to the

neighboring atom, σj2 is the mean-square disorder of the neighboring distance, k is the photo-

electron wavenumber, λ(k) is the wavelength of the radiation, fj(k) is the scattering amplitude,

and δj(k) is the phase shift.

A Fourier transform of the data expressed in k-space (momentum space) leads to an

expression of the data in R-space (direct space), where R represents the radial distance of the

neighboring atom from the absorbing atom. An example of the EXAFS spectrum in R-space can

be seen in Figure 1.5.2c. Theoretical EXAFS spectra can be generated from a chosen structural

model (see Figure 1.5.4e) using software packages such as IFFEFIT[34]. The theoretical EXAFS

26

spectra can then be fitted to the measured spectra as the parameters of the model are refined. The

model shown in Figure 1.5.4e is a representation of the RuO2 local structure that was used to fit

the EXAFS spectrum shown in Figure 1.5.2a.

1.6 Density Functional Theory (DFT)

Density Functional Theory (DFT) is a quantum mechanical computational approach used

in many branches of chemistry. In electrocatalysis, DFT serves a great purpose in terms of

modeling electrode surfaces and the electrocatalytic reactions that take place at the electrode-

electrolyte interface. In this respect, DFT is a valuable tool for understanding, predicting, and

designing catalyst materials. It has already been shown that DFT can be used to successfully

model electrocatalytic reactions such as ORR, OER, and CER on a variety of metal and metal

oxide surfaces[35-40].

The activity of different catalysts towards various chemical reactions can be predicted

computationally by considering the relationship between the adsorption energies of adsorbates in

addition to the Brønsted-Evans-Polanyi relation between reaction energies and the reaction

barriers. Classically, the Brønsted-Evans-Polanyi (BEP) relation describes the linear trend

between the activation energy and the reaction energy for reactions belonging to the same

class[41, 42]. The relationship can be expressed as:

(1.6.1)

where Ea is the activation energy, E0 is the activation energy of a reference reaction of the same

class, ΔH is the enthalpy of reaction, and α characterizes the position of the transition state along

the reaction coordinate (0 ≤ α ≤ 1). This linear relationship serves as a convenient way to

calculate the activation energy of reactions within a distinct family. The activation energy can

27

then be used to extract information regarding the kinetics of a particular reaction by applying the

Arrhenius equation:

(1.6.2)

where k is rate constant, A is pre-exponential factor, R is the Ideal Gas Constant, and T is the

absolute temperature.

In the context of heterogeneous catalysis, it was recognized that BEP relations exist

between the activation energy for the dissociative chemisorption of a number of molecules and

the reaction energy[43-45]. Essentially, this links the height of the activation barrier for a given

reaction to the driving force of the reaction. DFT calculations have been able to take advantage

of BEP relations such that a single parameter can be used to describe the catalytic activity of

many different materials towards a particular reaction.

In the case of the ORR, the binding energies of O*, OH*, and OOH* intermediate species

to the metal/metal oxide surface were found to serve as reliable descriptors for the

electrochemical activity. However, it is important to note that the binding energies of these

intermediates (*OH and *OOH) cannot be varied independently due to the scaling relationships

between the binding of related intermediates[40]. The thermodynamic model employed for the

ORR allows free energy of reaction intermediates to be shifted with respect to several

parameters, including pH, potential, concentration, and electric field strength. Essentially, the

Gibb's Free Energy of the reaction intermediates can be calculated by[35]:

(1.6.3)

where ∆E is the reaction energy calculated via DFT, ∆ZPE is the difference in zero point energy,

T is the absolute temperature, and ∆S is the change in entropy. Both ∆ZPE and ∆S are calculated

by DFT using the vibrational frequencies and standard tables for gas molecules. The typical

28

method employed for such calculations sets the reference potential to be that of the standard

hydrogen electrode (SHE). This allows for the chemical potential of H+ + e

- to be related to that

of 1/2H2 in the gas phase. Considering the pH of the electrolyte to be pH 0, the reaction free

energy of 1/2H2 → H+ + e

- is also zero for 1 bar of H2 at 298 K and an electrode potential of E =

0. Therefore, the free energy of the reaction *AH → A + 1/2H2 can be calculated since ∆G0 =

∆G at standard conditions. At a pH other than pH 0, the free energy of H+ ions can be corrected

by considering the concentration dependence of entropy:

(1.6.4)

From Eq. 1.6.4, the reaction free energy can then be calculated by:

(1.6.5)

where ∆GE = -eE is used to account for the effect of a bias on all states involving an electron at

the electrode by shifting the energy of this state.

Modeling the ORR at the electrode surface follows the associative mechanism of oxygen

reduction (Eqs. 1.6.6a-1.6.6d), which divides the 4-electron transfer reaction into a series of four

consecutive, concerted single electron and proton transfer steps[46, 47]. This follows from the

Marcus theory of electron transfer given the fact that it is generally more energetically favorable

to proceed through a series of single electrons transfers in which intermediates are formed than it

is to simultaneously transfer two electrons[40].

O2(g) + 4H+ + 4e

- + * → OOH* + 3H

+ + 3e

- (1.6.6a)

OOH* + 3H+ + 3e

- → O* + H2O(l) + 2H

+ + 2e

- (1.6.6b)

O* + H2O(l) + 2H+ + 2e

- → OH* + H2O(l) + H

+ + e

- (1.6.6c)

OH* + H2O(l) + H+ + e

- → * + 2H2O(l) (1.6.6d)

29

The calculated free energy associated with each individual electron/proton transfer step can be

used to construct a free energy diagram, which shows the free energy of each step vs. the

reaction coordinate as shown in Figure 1.6.1[48].

Figure 1.6.1. Example of free energy diagrams for the reduction of O2 on RuO2. The dotted line

represents the equilibrium potential of the reduction of O2 to H2O2.[48]

In general, the diagrams can be interpreted by considering the sequence of free energies

associated with each electron/proton transfer at the given driving force (E). According to the free

energy diagrams shown in Figure 1.6.1[48], at a potential of E = 0.8 V the reduction of O2 is

unfavorable. This is evidenced by the fact that not all steps are downhill in free energy

(following the reaction coordinate from right to left), specifically the final reduction of *OH to a

free surface site * (Step 1.6.6d). This indicates that the oxygen species are bound too strongly to

the RuO2 surface. It can also be observed that the first reduction step (Step 1.6.6a) forms a

strongly bound *OOH intermediate that continues to be more stable than the formation of H2O2

at most reasonable potentials. It allows to predict that the formation of H2O2 is not favorable

until the free energy associated with the formation of *OOH is higher than that of equilibrium

potential for the reduction of O2 to H2O2, EO2/H2O2 (i.e. the step associated with *OOH must be

above the dotted line in Figure 1.6.1). It is not until lower potentials (higher overpotentials) that

all steps are downhill in free energy and the complete 4-electron reduction of O2 can proceed, for

30

instance, at E = 0.5 V. It is not until E = 0.14 V that the formation of H2O2 becomes

thermodynamically possible, as can be seen in the free energy diagram for E = 0 V.

An ideal ORR catalyst should facilitate the complete 4e- reduction to H2O just below the

equilibrium potential of 1.23 V. Similarly, an ideal catalyst for O2 reduction to H2O2 would

facilitate the reaction just below the equilibrium potential of 0.68 V. Since the binding of *OH

and *OOH are related by a constant value (scaling relationship) and therefore cannot be varied

independently[40], it follows that there cannot be a single ideal catalytic site that binds all

oxygen intermediates optimally. A consequence of this scaling relationship is that ORR is

limited by either one of two steps: the activation of O2 to *OOH (Step 1.6.6a) or the removal of

adsorbed *OH (Step 1.6.6d). In the event that a catalyst binds oxygen intermediates too strongly

Step 1.6.6d will be the limiting step, whereas binding oxygen intermediates too weakly results in

Step 1.6.6a being the limiting step. This is essentially a formulation of the Sabatier principle[49],

which concludes that an ideal interaction of the adsorbed intermediate with the catalyst (neither

too weak or too strong) is desired.

An activity plot based on the DFT thermodynamic analysis is shown in Figure 1.6.2. It is

well known for heterogeneous catalysis that a volcano-type relationship is established when the

activity of a catalyst is plotted as a function of some property of that catalyst[50-52]. This

quantitative representation of the Sabatier principle has been reported in the literature for the

HER[53] and ORR[38, 54, 55]. In this case, the catalyst activity for the 2e- and 4e

- ORR

pathways is calculated in terms of (over)potential and the ability of the catalyst to bind *OOH

was used as the activity descriptor (note that the scaling relationship between oxygen

intermediates allows for the activity to be described using a single parameter, which again in this

case is the *OOH binding energy[56]) . The right (weak binding) leg of the volcano represents

31

the activation of O2 to *OOH (Step 1.6.6a). This step is the same in both the 4e- reduction of O2

to H2O and the 2e- reduction to form H2O2. The left (strong binding) leg of the volcano

represents the removal of *OH (Step 1.6.6d) in the 4e- reduction to H2O and the removal of

*OOH for the 2e- reduction to H2O2[38]. As follows from the Sabatier principle, an ideal catalyst

displaying an ideal binding (neither too strong nor too weak) for the oxygen intermediates would

be found at the apex of the volcano. Considering the ORR on RuO2 in Figure 1.6.2, it is shown

that RuO2 suffers from binding *OOH too strongly. This indicated that Step 1.6.6a is the limiting

step for ORR on RuO2. The ORR on doped and non-doped RuO2 will be considered in greater

detail in Chapter 3.

Figure 1.6.2. Volcano plot for the 2-electron (blue) and 4-electron (red) reduction of O2 to H2O2


represent the equilibrium potentials for the reduction products.

1.7 Scope of Dissertation

This dissertation focuses on various aspects of electrocatalytic devices for hydrogen

energy storage and purification. In Chapter 2, alkaline direct methanol fuel cells will be

discussed in regards to the potential losses encountered in an AEM environment. This chapter

32

identifies the structure of the double layer, electrostatic effects, and the adsorption of quaternary

ammonium species as causes of the observed potential losses which need to be removed. It also

aims at developing a supported Pt-based electrocatalyst for these applications. In Chapters 3 and

4 the effect of the electrocatalyst local structure on the selectivity in electrocatalytic reactions

relevant to electrolysis and fuel cells (i.e. oxygen evolution, chlorine evolution, and oxygen

reduction) will be discussed. Particularly, correlations between the variations of the RuO2 local

structure induced by transition metal doping and the electrochemical selectivity are examined on

the basis of XAS, DFT, and electrochemical data. Finally, Chapter 5 addresses the issue of

hydrogen purification in an electrochemical hydrogen pump. Pt-based catalysts tailored to

withstand CO/CO2 poisoning associated with the use of reformate gas mixtures are treated in

detail. Chapter 6 expands upon the research presented and offers insight to future directions of

hydrogen based technologies.

33

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9. Kinoshita, K., Electrochemical Oxygen Technology. The ECS Series of Texts and

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13. Liu, K., C. Song, and V. Subramani, Hydrogen and Syngas Production and Purification

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16. Sedlak, J.M., J.F. Austin, and A.B. LaConti, Hydrogen Recovery and Purification Using

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18. Frumkin, A.N., Influence of cation adsorption on the kinetics of electrode processes.

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22. Markovic, N.M., H.A. Gasteiger, and P.N. Ross, Oxygen Reduction on Platinum Low-

Index Single-Crystal Surfaces in Alkaline Solution - Rotating Ring Disk Studies. J. Phys.

Chem., 1996. 100: p. 6715-6721.

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23. Paulus, U.A., et al., Oxygen reduction on a high-surface are Pt/Vulcan carbon catalyst: a

thin-film rotating ring-disk electrode study. J. Electroanal. Chem., 2001. 495: p. 134-145.

24. Baltruschat, H., Differential electrochemical mass spectrometry. J. Am. Soc. Mass.

Spectrom., 2004. 15: p. 1693-1706.

25. Wang, H., Z. Jusys, and R.J. Behm, Ethanol oxidation on carbon-supported Pt, PtRu,

and Pt3Sn catalysts: A quantitative DEMS study. J. Power Sources, 2006. 154: p. 351-

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26. He, Q., et al., Highly Stable Pt−Au@Ru/C Catalyst Nanoparticles for Methanol Electro-

oxidation. J. Phys. Chem. C, 2013. 117: p. 1457-1467.

27. Arikawa, T., Y. Murakami, and Y. Takasu, Simultaneous determination of chlorine and

oxygen evolving at RuO2/Ti and RuO2-TiO2/Ti anodes by differential electrochemical

mass spectroscopy. J. Appl. Electrochem., 1998. 28: p. 511-516.

28. Macounova, K., et al., Parallel oxygen and chlorine evolution on Ru1−xNixO2−y

nanostructured electrodes. Electrochim. Acta, 2008. 53: p. 6126-6134.

29. Petrykin, V., et al., Local structure of Co doped RuO2 nanocrystalline electrocatalytic

materials for chlorine and oxygen evolution. Catal. Today, 2012. 202(63-69).

30. Petrykin, V., et al., Zn-doped RuO2 electrocatalysts for selective oxygen evolution:

relationship between local structure and electrocatalytic behavior in chloride containing

media. Chem. Mater. , 2011. 23: p. 200-207.

31. Macounova, K., M. Makarova, and P. Krtil, Oxygen evolution on nanocrystalline RuO2

and Ru0.9Ni0.1O2 electrodes – DEMS approach to reaction mechanism determination.

Electrochem. Commun., 2009. 11: p. 1865-1868.

32. Yano, J. and V.K. Yachandra, X-ray Absorption Spectroscopy. Photosynth. Res., 2009.

102: p. 241-254.

33. Kelly, S.D., D. Hesterberg, and B. Ravel, Analysis of Soils and Minerals Using X-Ray

Absorption Spectroscopy. Methods of Soil Analysis: Part 5. Mineralogical Methods.

2008, Madison, WI USA: Soil Science Society of America, Inc.

36

34. Newville, M., IFEFFIT : interactive XAFS analysis and FEFF fitting. J. Synchrotron

Radiat., 2001. 8: p. 322-324.

35. Rossmeisl, J., et al., Electrolysis of water on oxide surfaces. J. Electroanal. Chem., 2007.

607: p. 83-89.

36. Hansen, H.A., et al., Electrochemical chlorine evolution at rutile oxide (110) surfaces.

Phys. Chem. Chem. Phys., 2010. 12: p. 283-290.

37. Man, I.C., et al., Universality in Oxygen Evolution Electrocatalysis on Oxide Surfaces.

Chem. Cat. Chem., 2011. 3: p. 1159-1165.

38. Viswanathan, V., et al., Unifying the 2e– and 4e– Reduction of Oxygen on Metal

Surfaces. The Journal of Physical Chemistry Letters, 2012. 3(20): p. 2948-2951.

39. Halck, N.B., et al., Beyond the volcano limitations in electrocatalysis - oxygen evolution

reaction. Phys. Chem. Chem. Phys., 2014. 16: p. 13682.

40. Koper, M.T.M., Thermodynamic theory of multi-electron transfer reactions: Implications

for electrocatalysis. J. Electroanal. Chem., 2011(660): p. 254-260.

41. Evans, M.G. and M. Polanyi, Further considerations of the thermodynamics of chemical

equilibria and reaction rates. Trans. Faraday Soc., 1936. 32: p. 1333-1360.

42. Brønsted, J.N. and K.J. Pedersen, in Zeitschrift für Phys. Chemie, Stöchiometrie und

Verwandtschaftslehre. 1924. p. 185–235.

43. Nørskov, J.K., et al., Universality in Heterogeneous Catalysis. J. Catal., 2002. 209: p.

275-278.

44. Liu, Z.-P. and P. Hu, General trends in CO dissociation on transition metal surfaces. J.

Chem. Phys., 2001. 114: p. 8244-8247.

45. Logadottir, A., et al., The Bronsted-Evans-Polanyi Relation and the Volcano Plot for

Ammonia Synthesis over Transition Metal Catalysts. J. Catal., 2001. 197: p. 229-231.

37

46. Nørskov, J.K., et al., Origin of the Overpotential for Oxygen Reduction at a Fuel-Cell

Cathode. J. Phys. Chem. B, 2004. 108: p. 17886−17892.

47. Karlberg, G.S., J. Rossmeisl, and J.K. Nørskov, Estimations of electric field effects on the

oxygen reduction reaction based on the density functional theory. Phys. Chem. Chem.

Phys., 2007. 9: p. 5158−5161.

48. Abbott, D.F., et al., Oxygen reduction on nanocrystalline ruthenia – local structure

effects. RSC Adv., 2015. 5(1235-1243).

49. Sabatier, F., La catalyse en chimie organique. 1920, Berauge, Paris.

50. Boudart, M., Handbook of Heterogeneous Catalysis. 1997, Weinheim: Wiley-VCH.

51. Boudart, M. and G. Djéga-Mariadassou, Kinetics of Heterogeneous Catalytic Reactions.

1984, Princeton, NJ: Princeton Univ. Press.

52. Dumesic, J.A., et al., The Microkinetics of Heterogeneous Catalysis. 1993, Washington,

D.C.: Am. Chem. Soc.

53. Greeley, J., et al., Computational High-throughput Screening of Electrocatalytic

Materials for Hydrogen Evolution. Nat. Mater., 2006. 5: p. 909-913.

54. Stamenkovic, V., et al., Changing the Activity of Electrocatalysts for Oxygen Reduction

by Tuning the Surface Electronic Structure. Angew. Chem., Int. Ed., 2006. 45: p. 2897-

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55. Greeley, J., et al., Alloys of Platinum and Early Transition Metals as Oxygen Reduction

Electrocatalysts. Nat. Chem., 2009. 1: p. 552-556.

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Surfaces. ACS Catal., 2012. 2: p. 1654-1660.

38

Chapter 2

Analysis of Double Layer and Adsorption Effects at the Alkaline Polymer Electrolyte-

Electrode Interface and the Development of a Quaternary Ammonium Poisoning Resistant

Electrocatalyst for Methanol Oxidation

2.1 Introduction

Over the years methanol electro-oxidation in alkaline media has been the subject of

considerable research interest for applications such as anion exchange membrane direct methanol

fuel cells (AEM-DMFCs)[1-6]. The use of alkaline media in direct methanol fuel cell

applications provides several advantages over the proton exchange membrane (PEM)-based

analogues. It is well known that the process of methanol dehydrogenation and the removal of

COads intermediate species in alkaline media is far more facile in comparison with acidic

electrolytes due to the abundance of hydroxide anions[3, 7-9]. Furthermore, the use of alkaline

media opens up the possibility of using transition metal based electrocatalysts that would

otherwise be unstable and subject to dissolution in acidic media[10, 11].

A major obstacle for the advancement of AEM-DMFCs involves the development of

stable anion-exchange membranes and ionomers with sufficient ionic conductivity[12, 13].

Recent efforts typically rely on the use of quaternary ammonium species, e.g.

tetramethylammonium, as charge transfer groups for conducting hydroxide ions between the

cathode and anode[13-15]. Although progress has been made over the past 5-10 years, the ionic

conductivity and stability of anion exchange polymers still needs considerable improvement. The

addition of alkali metal hydroxide salts to the liquid fuel feed has been reported to significantly

improve ionic conductivity and thus the cell performance, although this is highly undesirable as

39

it ultimately leads to carbonate precipitation issues[12]. In addition to low ionic conductivity and

low thermal stability, the quaternary ammonium groups used to conduct hydroxide ions in AEMs

can also induce potential losses through electrostatic interactions and adsorption, i.e. poisoning,

of the catalyst surface, especially in the case of Pt[16].

The quaternary ammonium poisoning issue that arises with the use of Pt stems from both

the solvation energy of the quaternary ammonium species and the potential of zero charge (PZC)

of the Pt surface. Quaternary ammonium species are characterized as weak adsorbates that are

poorly solvated in water[17]. The low solvation energy of quaternary ammonium species leads to

favorable electrostatic interactions with a surface of the opposite charge, in contrast to alkali

metal cations in which the solvation energy is higher than that of electrostatic adsorption. The

potential of zero charge of Pt also plays an important role in the adsorption of such species. The

PZC of Pt has been reported as -0.4 V vs. NHE (0.37 V vs. RHE) at pH = 12[18] and -0.11 V vs.

NHE (0.55 V vs. RHE) at pH = 11.2[19, 20]. Below the PZC of Pt, the negatively charged

electrode surface favors the adsorption of weakly solvated and positively charged quaternary

ammonium species. The electrostatic adsorption of the charge transfer groups ultimately blocks

the catalyst surface, both decreasing the active surface area and preventing the adsorption of OH-

. It is not until higher potentials (E > PZC) that the positively charged moieties are repelled from

the electrode surface and sufficient adsorption of OH- begins to occur. This has a considerable

effect on the electrocatalysis of small alcohol molecules, i.e. MeOH[16], therefore anode

catalysts for use in an AEM environment must exhibit a high tolerance to quaternary ammonium

poisoning in order to provide high performance.

Studies have shown that alloying Pt with other metals can significantly improve the

catalytic activity and CO tolerance in comparison with pure Pt catalysts[21-23]. The increase in

40

activity is generally associated with increasing the oxophilicity of the catalyst surface, allowing

for the adsorption of hydroxide species at lower overpotentials to aid in the removal of adsorbed

CO from adjacent Pt sites, i.e. the bifunctional mechanism. Several metal oxide supports have

been proposed in the past, including those based on Ni[24, 25], Sn[26-28], Ce[24], Mn[24, 29],

Ti[30-32], Ru[33, 34], and Nb[33-35]. Among the materials investigated, Ni has been studied

extensively. In the case of Pt-Ni alloys the enhanced activity was explained in terms of the bi-

functional mechanism and was also supported by DFT calculations[36]. It has been shown that

Ni serves to optimize the electronic structure of Pt sites via the ligand effect, thus lowering the

binding energy of COads and promoting its removal. The former enhancement is likely due to the

fact that Ni forms a Ni(OH)2 passivation layer in the alkaline environment[25]. Pb is also a

promising candidate as it offers a significant enhancement in the electro-oxidation of alcohols

(MeOH, EtOH, etc.) and other organic fuels[23, 29, 37-41]. It was suggested that ordered

intermetallic phases, such as Pt-Pb, exhibited improved catalytic activity for MOR as evidenced

by a lower onset potential[42].

This chapter will begin with a review of the research conducted to characterize the

adsorption issues arising at the anode-polymer electrolyte interface[16]. In the following sections

the extent of ammonium poisoning in alkaline media will be addressed both quantitatively and

qualitatively. Furthermore, the electrostatic effects of adsorbed quaternary ammonium species

will be examined through the use of transition metal redox probes. In response to the poisoning

issues encountered in an alkaline AEM environment, we then report herein the characterization

and electrochemical methanol oxidation performance of Pt deposited on the NiPb/C support

(denoted as Pt/NiPb/C) in alkaline media. This catalyst exhibits high electrocatalytic methanol

oxidation activity and a high resistance to quaternary ammonium poisoning.

41

2.2 Experimental

2.2.1 Electrochemical Characterization of the Anode-Polymer Electrolyte Interface

Ammonium contamination studies were conducted on Pt/C modified glassy carbon

electrodes. A catalyst ink was prepared from 25 mg of 30% Pt/C (BASF) in 10 mL Millipore

H2O and 10 mL isopropyl alcohol. The ink was sonicated for 30 minutes. Exactly 5 μL of ink

was deposited on a glassy carbon rotating disk electrode (Ageo = 0.247 cm2) to yield a total

catalyst loading of 7.5 μgPt/cm2. Prior to ink deposition, the electrode was polished with 0.05 μm

alumina slurry and sonicated twice in a sonic bath for 30 seconds each in a 50:50 solution of

Millipore H2O and isopropyl alcohol. The reference electrode used was a reversible hydrogen

electrode (RHE) prepared from a solution of 0.1 M KOH. At the beginning of each experiment,

the electrode was cycled 20 times from 0.05 V to 1.2 V at 50 mV/s in 0.1 M KOH followed by

five cycles at 20 mV/s. Absolute methanol was then added to the KOH solution to obtain a total

MeOH concentration of 0.5 M. Methanol oxidation was then performed by holding the potential

at 0.6 V vs. RHE for 900 seconds. Subsequently, an aliquot of a given contaminant was added to

bring the total contaminant concentration to 1 mM and the potential was again held at 0.6 V for

900 seconds. This procedure was repeated for the remaining contaminant concentrations of 5, 10,

20, 40, 60, and 120 mM. A new ink coating was used in the study of each contaminant.

Contaminants investigated included tetramethylammonium hydroxide (TMAOH, 25% w/w aq) ,

tetraethylammonium hydroxide (TEAOH, 35% w/w aq), tetra-n-propylammonium hydroxide

(TPAOH, 40% w/w aq.), and benzyltrimethylammonium hydroxide (BTMAOH, 25% w/w aq.)

all from Alfa Aesar. Electrochemical measurements were made using an Autolab

Potentiostat/Galvanostat Model PGSTAT30 (Metrohm USA).

42

The electrochemical behavior of the transition metal complexes was studied at an alkaline

ionomer modified glassy carbon electrode. A glassy carbon rotating disk electrode (Ageo = 0.247

cm2), after polishing and pretreatments described above, was allowed to dry after deposition of 5

μL of 5 wt.% anion exchange ionomer (AS4, Tokuyama Corp., Japan) on the surface. The

electrode was placed in an argon saturated 0.1 M NaOH solution and subjected to 20 cycles

from 0.05 V to 1.2 V at 50 mV/s and five cycles at 20 mV/s. 10 mM of K3Fe(CN)6 was then

added to the solution and the solution was purged with argon for an addition 10 minutes before

collecting data. This experiment was repeated using 10 mM Co(NH3)6Cl3 in place of K3Fe(CN)6.

2.2.2 Preparation of Pt/NiPb/C Electrocatalyst

The layered Pt/NiPb/C electrocatalyst reported herein was synthesized via the sequential

aqueous impregnation method using NaBH4 as a reducing agent. The catalyst consists of 20

wt.% Pt deposited on a composite support consisting of 60 wt.% Ni-Pb (2:1 atomic ratio Ni:Pb)

on carbon (Black Pearls Carbon-2000, Cabot Corp.). The carbon black was dispersed in ~200

mL solution containing stoichiometric amounts of Ni(NO3)2 and Pb(NO3)2 (Alfa Aesar, 99.99%

metal basis) metal salt precursors in ultra-pure water (18.2 MΩ, Milli-Pore filtration system).

The suspension was homogenized for 30 min before drop-wise addition of an aqueous NaBH4

solution (3 mL, 2.5 molar excess) and subsequent stirring for 1 h. The product was filtered and

rinsed 3 times with Milli-Pore water. The resulting NiPb/C powder was then dried in a vacuum-

oven at 80 oC overnight. Platinum was deposited on the NiPb/C support by repeating the above

procedure: The NiPb/C powder was dispersed in ~100 mL of aqueous H2PtCl6 (Alfa Aesar,

99.99% metal basis) solution and stirred for 30 min, followed by drop-wise addition of NaBH4

solution (3 mL, 2.5 molar excess). The final product was stirred for 1 h, filtered and rinsed with

43

ultra-pure water, and then dried in the vacuum-oven at 80 oC overnight. Commercial 46.2 wt.%

Pt/C (Tanaka Corp, Japan), 73.5 wt.% PtRu/C (Tanaka Corp, Japan), and 40 wt.% Pt/C (ETEK,

USA) electrocatalysts were used as received.

2.2.3 XRD and TEM Characterization:

The dispersion of Pt on the NiPb/C support was observed using transmission electron

microscopy (TEM, JEOL 2100). Powder X-ray diffraction (XRD) analysis of catalyst samples

was carried out with a Rigaku Ultima IV diffractometer using CuKα radiation (λ = 1.542 Å)

operating at 40 kV and 44 mA with a 0.05o step and 5 sec hold per step in order to obtain high

resolution XRD signals. XRD analysis was performed using Rigaku PDXL software.

2.2.4 Electrochemical Characterization of Pt/NiPb/C

Electrochemical measurements were performed in a conventional single-compartment,

three-electrode cell at ambient temperature (298 K). The glassy carbon working electrode (0.247

cm2 surface area) was polished with 0.05 µm alumina paste and rinsed with Milli-Pore water

prior to catalyst ink deposition. Catalyst ink suspensions, which consisted of an appropriate

amount of catalyst, isopropyl alcohol, ultra-pure water, and 5 wt.% AS4 anion exchange ionomer

solution (Tokuyama Corp., Japan), were ultrasonically mixed for at least 30 min. Exactly 5 μL of

ink was drop-cast on the mirror-polished glassy carbon working electrode to achieve a mass

loading of 8 μgPt/cm2. The reference electrode was a reversible hydrogen electrode (RHE) in 0.1

M KOH solution and Pt mesh was used as the counter electrode. All potentials are reported

versus the RHE scale. Before measurements, the 0.1 M KOH electrolyte was purged with Ar for

30 min.

44

Cyclic voltammetry (CV), chronoamperometry (CA), and electrochemical impedance

spectroscopy (EIS) measurements were performed using an AutoLab PGSTAT30 (Metrohm,

USA) and the Nova 1.10 software package. Cyclic voltammograms were collected within a

potential range of 0.05 V to 1.00 V (PtRu/C was measured from 0.05 V to 0.80 V) at 20 mV s-1

.

The CA results were generated within a potential range of 0.40 V to 0.80 V using 50 mV step

size and 180 s hold time. Additional CV measurements were made to investigate the effect of the

anion exchange ionomer on the MOR activity. In this case, 50 μL of either 5 wt.% Nafion

solution or 5 wt.% AS4 solution was used in the catalyst ink suspensions described above. Cyclic

voltammograms were collected in Ar purged 0.1 M KOH and then in 0.1 M KOH + 0.5 M

MeOH at a scan rate of 10 mV/s. The EIS measurements were recorded within a 32 kHz to 1 Hz

frequency range at E = 0.70 V with an amplitude of 10 mV. EIS data was analyzed using non-

linear least-squares fit of Nyquist plot to simple Randles cell circuit.

2.3 Result and discussion

2.3.1 Electrochemical Characterization of the Anode-Polymer Electrolyte Interface

In order to address and quantify the issue of quaternary ammonium poisoning on

platinum, the oxidation of methanol on platinum was studied as a function of the ammonium

group adsorbed on the surface of the electrode. The first three cations studied were quaternary

ammoniums with different length alkyl chains (methyl, ethyl and propyl). The fourth one was

benzyltrimethylammonium hydroxide. The hydroxide salt of these quaternary ammonium cations

was added to a 0.5 M MeOH solution. The percent loss in current density after 900 sec potential

control at 600 mV (vs. RHE) in 0.1 M KOH was then recorded as a function of quaternary

ammonium as shown in Figure 2.1. In general, all ammonium cations showed a significant

45

adsorption effect on the platinum surface resulting in a severe loss of methanol oxidation

activity. The decrease in current became more pronounced at higher concentrations of the added

cations. TMA+ showed the smallest drop in current, whereas the other three cations with bulkier

substitute groups resulted in substantially greater loss in current. This observation is consistent

with previous literature that showed a higher surface inhibition with the longer alkyl length on

the ammonium group[43].

Conc. of Contaminant [mM]

0 20 40 60 80 100 120

% L

oss in

Cu

rre

nt

De

nsity

-100

-80

-60

-40

-20

0

TMAOH

TEAOH

TPAOH

BTMAOH

Figure 2.1. Percent loss in current density as a function of contaminant concentration in 0.1 M

KOH + 0.5 M MeOH + x mM contaminant. Glassy carbon disk electrode (0.247 cm2) with 5 μL

of 30% Pt/C ink deposited (Loading = 7.5 ugPt/cm2). Steady state currents obtained from

chronoamperometry at 0.6 V vs. RHE at 900 seconds.

These results with different quaternary ammonium cations in aqueous electrolytes show

that quaternary ammonium ion adsorption on Pt surfaces lowers the rate of methanol oxidation

by blockage of the active catalyst surface area. However, the solid polymer electrolyte, ion-

conducting medium in a polymer-based fuel cell is different from that of free ions in an aqueous

46

solution. A fuel cell electrode consists of quaternary ammonium cations with little chain segment

mobility tethered to a polymer backbone. A more realistic representation of the reaction medium

in a fuel cell on a planar electrode surface was investigated by Unlu et al.[16], in which an anion

exchange polymer with quaternary ammonium ion sites, poly-tetramethylammonium hydroxide

(PTMAOH), was used as the polymer electrolyte without additional alkali metal hydroxide. This

polymer electrolyte provides high anion conductivity with high viscosity in the electrolyte phase,

as would occur in a fuel cell electrode.

Methanol oxidation is an inner-sphere electron transfer reaction and methanol is a neutral

molecule. Therefore, the Frumkin effect described in Chapter 1 (See Section 1.3.2) does not

directly affect the methanol species. Instead, the oxidation of alcohols on platinum requires the

presence of adsorbed hydroxide species to proceed through a Langmuir-Hinshelwood

mechanism. Transport of hydroxide anions across the diffuse layer and subsequent specific

adsorption on the electrode surface are necessary for alcohol oxidation. The concentration of

hydroxide at the OHP, OHP

OHC , available for adsorption is controlled by the bulk concentration,

charge of the ion, and ϕ2, as shown in below in Eq. 2.1.

2expOHP bulk

OH OH

zFC C

RT

(2.1)

The double layer structure of an electrode-electrolyte interface significantly differs in

polymer electrolytes and aqueous electrolytes. The double layer structures in NaOH, TMAOH,

and polymer electrolyte surfaces are depicted for a negatively charged electrode surface in

Figure 2.2. When the electrode potential is negative of PZC in a NaOH solution, there will be a

potential gradient from the electrode to the bulk solution potential. In TMAOH solution, TMA+

species are adsorbed on the Pt surface, particularly because ϕM < ϕS, forming a compact inner

47

layer of positive charges. In contrast, TMA+ moieties tethered to a polymer chain are not as

mobile as free TMA+ causing the excess charge in polymer electrolytes to be distributed across

the diffuse layer. The mobile TMA+ cations can form a more compact double layer due to their

mobility. These variations in the structure of the double layer result in large differences in ϕ2.

The order of ϕ2 would be 2 2 2

T N P in TMAOH, NaOH, and Polymer electrolytes,

respectively. However, even though the relatively more positive value of 2

T would favor the

transport of hydroxide ions to the OHP, this factor is outweighted by the blockage of the

electrode surface, which decreases the available catalyst surface area and therefore the oxidation

rate.

Figure 2.2. Expected potential profile of anode double layer interface in NaOH, TMAOH, and

polymer electrolyte solutions when the electrode potential is more negative than PZC[16].

48

When the TMAOH and polymer electrolyte solutions are considered, the TMA+ cation is

more mobile than the tethered polymer electrolyte cation which could lead to a greater degree of

adsorption and loss of catalyst surface area, however, in the case of Unlu et al.[16] the polymer

electrolyte cations show smaller currents. The lower reaction rate in polymer electrolyte solutions

is probably due to the more negative value of 2

P than that of 2

T . The negative value of 2

P

inhibits the transport of hydroxide to the electrode surface, limiting the reaction rate. This effect

on hydroxide transport and adsorption explains the enhancement in the performance when alkali

metal hydroxide is added to the fuel, as shown in the literature[16]. The addition of NaOH to

polymer electrolytes introduces two effects; (i) the bulk concentration of OH- ions increases in

Eq. 2.1 and (ii) the mobile Na ions lead to a decrease in the thickness of diffuse layer, resulting

in lower ϕ2. These changes lead to higher apparent reaction rates.

The proposed ϕ2 effect on hydroxide transport and subsequent adsorption are difficult to

quantify. Instead, charged redox couples were added to the electrolyte to help identify changes

and probe the double layer structure and consequent ϕ2 effect. In this study, the Fe(CN)63-/4-

and

Co(NH3)62+/3+

redox couples were employed in order to probe the electrode surface for evidence

of the ϕ2 effect in anion exchange ionomers and membranes. Cyclic voltammetry was used to

observe the redox behavior of each couple on a glassy carbon electrode in the presence and

absence of an anion exchange ionomer film (Tokuyama, AS4 which contains quaternary

ammonium cation ion-exchange groups). Figure 2.3 displays the cyclic voltammograms for the

Fe(CN)63-/4-

redox couple on a GC electrode with and without a deposited film of AS4 ionomer.

Typically, the current magnitude cannot be used to make a direct comparison between solution

and AEM environment because both concentration and diffusion coefficients change. However,

the difference between the oxidation and reduction peaks can be evaluated to define a change in

49

reaction rates. From Figure 2.3, it can be seen that the peak separation (ΔEp) decreases from

279.5 mV in the absence of AS4 to 131.0 mV in presence of the AS4 ionomer film. In contrast,

Figure 2.4 shows the CV profiles for the Co(NH3)62+/3+

redox couple on a GC electrode with and

without a deposited film of AS4 ionomer. The current densities are lower with AS4 ionomer than

without ionomer because AS4 ionomer mostly excludes positively charged species, decreasing

the concentration on the electrode surface, i.e. lower current. However, it is critical to note that

there is a slight increase in ΔEp from 65.3 mV to 83.9 mV in the presence of ionomer.

Figure 2.3. CV of Fe(CN)63-/4-

redox couple in Ar saturated 0.1M NaOH + 10mM K3Fe(CN)6

solution. Scans performed on glassy carbon disk electrode with 5 μL of 5 wt.% AS4 ionomer

deposited on surface. Geometric surface area = 0.247 cm2. Scan rate = 20 mV/s.

50

Figure 2.4. CV of Co(NH3)62+/3+

redox couple in Ar saturated 0.1M NaOH + 10mM

Co(NH3)6Cl3 solution. Scans performed on glassy carbon disk electrode with 5 μL of 5 wt.%

AS4 ionomer deposited. Geometric surface area = 0.247 cm2. Scan rate = 20 mV/s.

For these redox couples, the electrode potentials are more positive than PZC (ca. 0.55 V vs.

RHE[19, 20]) at the electron transfer potentials and the reaction is believed to proceed through

an outer-sphere electron transfer. The double layer structure differs from the methanol oxidation

region where the electrode potential is more negative than PZC. Figure 2.5 shows the double

layer structure when the electrode potential is greater than PZC in the presence and absence of

AEM ionomer. As a distinct difference, the diffuse layer in an AEM becomes thicker relative to

that in the absence of an AEM. Consequently, ϕ2 is more positive in the presence of AEM. For a

negatively charged redox couple, both exponential terms favor the electron transfer, i.e. higher

kinetics. This is consistent with the lower peak splitting for Fe(CN)63-/4-

in Figure 2.3. However,

for positively charged species, the driving force still favors the reaction but the transport of these

species is inhibited. This inhibition factor is greater than the enhancement in driving force,

51

resulting in lower reaction kinetics. This is consistent with the greater peak splitting in Figure

2.4.

Figure 2.5. Expected potential profile of anode double layer interface in the presence and

absence of anion exchange ionomer when the electrode potential is more positive than PZC[16].

Now that the issues surrounding the specific adsorption of quaternary ammonium species

have been addressed both qualitatively and quantitatively we will consider the rational design of

a Pt-based catalyst capable of delivering high methanol electro-oxidation activity while showing

a high resistance to quaternary ammonium poisoning. The following sections will detail the

synthesis and characterization of a highly oxophilic NiPb/C support for Pt nanoparticles.

52

2.3.2 XRD and TEM Characterization

The powder X-ray diffraction patterns of the layered Pt/NiPb/C and commercial Pt/C

samples are shown in Figure 2.6. Both samples display the anticipated peaks at approximately

39.5 and 46.0 degrees, which are characteristic of the Pt fcc (111) and (220) crystallite

reflections. The absence of any significant change in the lattice parameter (See Table 2.1) or shift

in the primary fcc diffracting regions of the layered Pt/NiPb/C material relative to that of Pt/C

suggests that the deposited Pt exists as a separate phase and is not alloyed with the NiPb/C

support to any significant extent. In addition, the layered Pt/NiPb/C sample exhibits features

characteristic of a hydro-cerussite Pb3(CO3)2(OH)2 phase. Although crystalline Ni/Ni-Ox phases

were not detected in the XRD analysis, the slight shift of the characteristic hydro-cerussite

reflections to higher angles suggests that some Ni is likely incorporated into the Pb crystallite

phase. Additionally, some Ni may also exist as an amorphous surface oxide layer. It should also

be noted that a PtNiPb/C material was synthesized and characterized by Chen et al.[44] that did

not contain hydro-cerrusite reflections in the diffractograms but did show a shift in the Pt fcc

reflections (and thus incorporation of Ni and Pb into the platinum crystal lattice).

The corresponding coherent domain sizes as calculated by the Scherrer equation are

shown in Table 2.1. From Table 2.1 it can be seen that the calculated crystallite domain sizes are

nearly the same for both the commercial Pt/C and the PtRu/C material (1.7 and 1.5 nm,

respectively), while the size of the Pt/NiPb/C crystallites are more than twice the size of the

commercial samples. This indicates that the surface area of the Pt/NiPb/C sample is likely much

less than that of the commercial samples. As a result, any enhancement in MOR activity for

Pt/NiPb/C with respect to that of Pt/C or PtRu/C is unlikely to be a result of increased surfaces

area.

53

Figure 2.6. Powder X-ray diffraction patterns for nanocrystalline Pt/C and Pt/NiPb/C samples.

Table 2.1. XRD and electrochemical results for various catalysts.

XRD results Electrochemical results

Catalyst Lattice

parameter (nm)

Coherent domain

size (nm)

Eonset

(V)

Epeak

(V)

iPeak

(mA/cm2)

iat 0.7V

(mA/cm2)

Pt/C 0.3918 1.7 0.356 0.959 12.98 2.20

PtRu/C 0.3868 1.5 0.299 0.761 2.72 2.40

Pt/NiPb/C 0.3912 4.4 0.252 0.918 15.73 8.61

Representative TEM images of the layered Pt/NiPb/C electrocatalyst are shown in Figure

2.7. From Figure 2.7 it can be seen that the nanoparticles are well dispersed on the carbon

support. The average size of the supported nanoparticles obtained from TEM analysis is less than

5 nm, which is in agreement with the coherent domain size calculated from the Scherrer

equation. The high dispersion of nanoparticles is likely a major factor contributing to the high

MOR activity, though it is not the only source responsible for the enhancement in activity.

54

Figure 2.7. TEM images of the Pt/NiPb/C catalyst. Insert is higher magnification.

2.3.3 Electrochemical Characterization of Pt/NiPb/C

The electrochemical methanol oxidation response of the corresponding catalysts in 0.1 M

KOH and 1.0 M methanol solution are shown in Figure 2.8a. The MOR activity of the layered

Pt/NiPb/C is compared with the commercial Pt/C and PtRu/C catalysts. The layered Pt/NiPb/C

material displays the highest overall MOR activity among these catalysts. It shows the highest

peak current (ipeak) and the lowest onset potential (Eonset) (See Table 2.1). Additionally, the

chronoamperometric response was measured, which provides further data on catalyst activity and

stability. The steady state current measured at 180 s for each potential step can be seen in Figure

2.8b. The same general trend in methanol oxidation activity that is shown by the cyclic

voltammograms is observed again here, with the Pt/NiPb/C material showing higher oxidation

currents across the entire potential range examined. From Figure 2.8a and Table 2.1 it can be

seen that the layered Pt/NiPb/C and PtRu/C have a similar onset potential, with methanol

oxidation beginning at roughly 250 and 300 mV vs. RHE, respectively, which is approximately

55

50 to 100 mV lower than Pt/C. Although the onset of methanol oxidation occurs at roughly the

same potential for Pt/NiPb/C and PtRu/C, the methanol oxidation current is considerably higher

for Pt/NiPb/C across the entire MOR potential range. In general, the low activity of PtRu/C can

be attributed to the weak adsorption of methanol on Ru sites and the dilution of Pt surface sites

required for methanol dehydrogenation[8]. Furthermore, it is interesting to note that the peak

current potential (Epeak) for Pt/NiPb/C is positioned nearly 80 mV beyond the Epeak for PtRu/C

even though Eonset for the two materials is about the same. On the basis of the bifunctional

mechanism, this suggests that OH adsorption (and subsequent COads removal) begins at roughly

the same potential for both Pt/NiPb/C and PtRu/C. However, at more positive potentials it is

likely that the higher surface coverage and formation of more strongly bound oxygenated species

begins to inhibit further removal of the COads intermediate on PtRu/C, thus decreasing the

availability of active sites for MOR[8]. In the case of Pt/NiPb/C, however, it appears that

sufficient surface coverage and binding of OHads exists over a wider potential range and that the

formation of strongly bound oxo species does not begin to inhibit MOR catalysis until higher

potentials.

56

Figure 2.8. Electrochemical measurement of Pt/C, PtRu/C, and Pt/NiPb/C catalysts in 0.1 M

KOH and 1.0 M methanol at 298 K. Current densities are normalized to the geometric surface

area. (a) CV measurements at a scan rate of 20 mV/s. (b) Steady state currents obtained after 180

s from chronoamperometric tests ranging from 400 to 800 mV with 50 mV steps.

Enhancement of MOR activity has been previously reported for PtNi-based catalysts[36,

45, 46]. Sun et al.[36] reported higher specific activities for PtmNin/C and decreased CO

adsorption when compared to Pt/C, which was attributed to modification of the Pt electronic

structure and weakened interactions between PtmNin clusters and adsorbed CO. Wu et al.[45]

reported lower onset potentials (by approx. 50 mV) and higher peak current densities (approx.

2x) for PtNi (4:1) supported on MWCNTs over Pt/MWCNTs. Furthermore, Jiang et al. have

shown that Ni-Ox electrodes can oxidize MeOH at much higher anodic potentials, with the

Ni2+/3+

redox transition occurring at nearly the same potential as the MOR[36]. Lyons et al. have

also explained that in the alkaline environment, Ni and other transition metals assume a hydrous

oxy-hydroxide layer during potential cycling[46]. This hydrous oxy-hydroxide layer should

exhibit sufficient chemical reversibility (similar to the Ti, Nb & W bronzes described by Jaksic

et al.[47]) to facilitate spillover of OHad – thus promoting complete MOR.

57

Even greater enhancement of alcohol oxidation activity has been reported for Pb-

containing catalysts[37, 38, 40, 41, 44, 48, 49]. Sun et al.[38] reported that the Pt mass

normalized peak current density for PtPb/C was nearly 2.1 times that of Pt/C. Furthermore,

Matsumoto et al.[37] reported BET surface area normalized currently densities for PtPb/C that

were nearly 8 times higher than those of PtRu/C at the same potential. In both the cases the MOR

onset potential for PtPb was reported to be significantly lower (by 100-150 mV) than those in the

case of Pt/C or PtRu/C. Chen et al.[44] reported a Pt-Ni-Pb/C catalyst that reached peak current

densities in acidic media that were twice as high as Pt/C, though the dramatic shift to a more

negative MOR onset potential relative to Pt/C was not observed as in alkaline media[37, 38].

Table 2.1 shows that the layered Pt/NiPb/C material displays a peak current density

approximately 1.2 times that of Pt/C, although the current density at intermediate potentials, i.e.

0.7 V vs. RHE (see Table 2.1) is up to 4 times higher than that of Pt/C. The enhancements

reported for Pt-Pb catalysts have been described in terms of the bifunctional effect and the

electronic effect, although Sun et al.[38] claimed that the bifunctional effect is unlikely to play a

major role and attributed the observed enhancements to a partial charge transfer from Pb to Pt,

resulting in a modification of the Pt d-orbitals and a subsequent decrease in CO adsorption. It is

important to note that these reports have dealt mostly with single-phase bitmetallic PtPb or PtNi

alloys, which is not the case for the layered Pt/NiPb/C material described herein that includes a

hydro-cerussite (Pb3(CO3)2(OH)2 phase in addition to the Pt fcc crystal phase. The effect of Pb in

the Pt/NiPb/C catalyst has not been fully understood yet and further evaluations of the

synergistic effects of Ni and Pb will likely require the employment of DFT calculations.

The interfacial effects of the anion exchange ionomer (AEI) and the extent of poisoning

resulting from the presence of quaternary ammonium species were investigated by monitoring

58

the electrochemical response of each catalyst in the presence and absence of AEI (AS4, which

contains quaternary ammonium cation ion-exchange groups) via cyclic voltammetry and

electrochemical impedance spectroscopy (EIS). The methanol electro-oxidation curves for 40%

Pt/C and Pt/NiPb/C with and without AEI in 0.1 M KOH + 0.5 M methanol solution are shown

in Figure 2.9. A 5 wt.% Nafion solution was used as a catalyst binder in the absence of AEI.

From Figure 2.9 it can be seen that the Pt/NiPb/C catalyst with AEI exhibits a significant

increase in methanol oxidation current when compared to the MOR activity in the absence of

AEI, indicating that the presence of quaternary ammonium facilitates the diffusion of OH- to the

Pt surface. This also indicates that the AEI is not blocking the active catalyst surface area. This is

in contrast to what is observed for Pt/C. In Figure 2.9 it is shown that the presence of AEI

decreases the overall magnitude of the methanol oxidation current for Pt/C. The insert in Figure

2.9 shows the cyclic voltammograms of both Pt/C and Pt/NiPb/C in 0.1 M KOH without the

addition of MeOH. From Figure 2.9 it can be seen that the addition of AEI to the catalyst ink

decreases the active catalyst surface area on Pt/C, as evidenced by the loss of charge in the Hupd

region. Although there are no strong Hupd features for Pt/NiPb/C, there is an increase in the

hydrogen oxidation current in the presence of AEI following evolution of hydrogen from the

cathodic scan.

It is interesting to note that the formation of surface oxides is also enhanced in the

presence of the ionomer for Pt/C. This is likely due to the fact that as one goes positive of the

PZC on Pt (E > ~0.55 V vs. RHE) the adsorbed quaternary ammonium are repelled by the

increasing positive charge of the catalyst surface. As quaternary ammonium are repelled from the

Inner Helmholtz Plane (IHP), it allows OH- (which is now attracted by the positively charged

surface) to access the IHP. The potential regime for MOR on Pt/C only ranges from

59

approximately 0.4 to 0.9 V vs. RHE. The onset of MOR therefore coincides with the adsorption

region of OH- while the formation of the surface oxide layer at more positive potentials (E > 0.9

V) inhibits the oxidation MeOH, resulting in a sharp decrease in the MOR activity[8, 50].

Therefore, when the potential is sufficiently positive to repel adsorbed quaternary ammonium (E

> 0.55 V) the electrode potential is already nearing the regime where it begins to bind oxygen too

strongly to allow for optimal MeOH coverage. In the case of Pt/NiPb/C, the support has a very

high oxo/hydrophilic character that likely inhibits quaternary ammonium adsorption in this

potential regime while allowing for enhanced OH- adsorption and transport to adjacent Pt MOR

active sites. The effects can also be seen in the EIS (See Figure 2.10), where the presence of AEI

decreases both the high frequency resistance and the charge transfer resistance on Pt/NiPb/C.

The accompanying decrease in the charge transfer resistance suggests that not only does the

Pt/NiPb/C catalyst have a high resistance to quaternary ammonium poisoning, but the AEI in fact

exhibits a significant promoting effect on MOR activity. The opposite case is observed for Pt/C,

where the high frequency resistance and charge transfer resistance are increased. In the case of

Pt/C this suggests that the AEI is blocking the active surface area and decreasing the number of

sites available for OH- adsorption. This effect is in agreement with Unlu et al.[16], where it was

determined that the specific adsorption of quaternary ammonium and possible ϕ2 effects all work

to inhibit MOR until the anode potential becomes positive of PZC.

60

Figure 2.9. Cyclic voltammograms of 40% Pt/C and Pt/NiPb/C catalysts with Nafion or AS4

ionomer used as a catalyst binder. Scans were taken in Ar saturated 0.1 M KOH + 0.5 M

methanol at 298 K at a scan rate of 10 mV/s. Cyclic voltammograms collected at a scan rate of

10 mV/s in Ar saturated 0.1 M KOH are shown in the insert.

Figure 2.10. Electrochemical impedance spectroscopy of 40% Pt/C and Pt/NiPb/C with either

Nafion or AS4 ionomer used as a catalyst binder in 0.1 M KOH + 0.5 M MeOH. EIS data was

collected at E = 0.70 V vs. RHE from 32 kHz to 1 Hz with a amplitude of 10 mV. The insert

shows the high frequency resistance.

61

The dramatic improvement in MOR activity of the Pt/NiPb/C over the commercial Pt/C

and PtRu/C is likely due to the morphology of the catalyst itself. The NiPb/C support appears to

improve the quaternary ammonium poisoning resistance due to the high oxo/hydrophilicity and

preference of both Ni and Pb to form surface hydroxide and oxyhydroxide species in alkaline

media. Pt is known to be poisoned by quaternary ammonium species in addition to the COads

intermediate in MOR. Therefore, the presence of an oxophilic metal in close proximity to Pt sites

is necessary to aid not only in the removal of COads but also to inhibit the adsorption of

quaternary ammonium species. Although PtRu/C is well known to promote the removal of CO,

the alloyed material may not provide enough Ru sites in close proximity to Pt sites to effectively

inhibit quaternary ammonium adsorption and provide sufficient OH- spillover. We suggest that

by depositing highly dispersed Pt nanoparticles onto the highly oxophilic NiPb/C support, as

opposed to alloying or depositing oxophilic metals onto Pt particles, that we effectively inhibit

the specific adsorption of quaternary ammonium on Pt active sites while providing a high rate of

OH- transport from the surrounding support, thus significantly improving the MOR activity in

alkaline media.

2.4 Conclusion

Quaternary ammonium cations in solution were shown to inhibit the oxidation of

methanol through specific adsorption. Specific adsorption, migration in the diffuse layer due to

hydroxide repulsion away from the electrode, and possible ϕ2 effects, all work against an

efficient electrode structure until the anode potential becomes positive of PZC. In this study, the

effect of PZC on anode operation was explored. Transition metal redox complexes in a polymer

electrolyte showed behavior consistent with the ϕ2 effect.

62

The deposition NiPb onto a carbon support prior to Pt deposition can significantly

improve the Pt electrochemical performance for methanol oxidation versus Ru-alloying. TEM

images of Pt/NiPb/C catalyst indicate that the Pt nanoparticles are well dispersed on the NiPb/C

support. In the case of Pt/NiPb/C, the addition of anion exchange ionomer to the catalyst ink

suspension is an important factor for enhancing the electrochemical performance. The CV, CA,

and impedance results with/without AEI indicate that NiPb/C support imparts a strong resistance

to quaternary ammonium poisoning of the Pt catalyst while dramatically improving MOR

activity. It is proposed that the AEI facilitates both diffusion of OH- from the bulk electrolyte

through the ionomer and surface diffusion of OHads from the NiPb interlayer. Further analysis to

elucidate the origin of the synergistic effects of NiPb/C on Pt is currently in progress.

2.5 Acknowledgments

The authors wish to thank the Army Research Office for the financial support under a

single investigator grant. The authors gratefully acknowledge the financial support from the U.S.

Army and DuPont Corporation especially Dr. Deryn Chu (US Army) and Shoibal Banerjee

(DuPont). Materials support from Tokuyama Corp. is also gratefully acknowledged. Murat Ünlü

and Paul A. Kohl of Georgia Institute of Technology and Nagappan Ramaswamy of

Northeastern University are gratefully acknowledged for their significant contributions to the

development of interfacial model of the anode-polymer electrolyte interface. Additionally,

Myoungseok Lee and Michael Bates of Northeastern University are gratefully acknowledged for

their assistance in the development and electrochemical characterization of the Pt/NiPb/C

catalyst.

63

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methanol electro-oxidation. Int. J. Hydrogen Energy, 2012. 37: p. 1263-1271.

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PteTiO2/ITO electrode under UV illumination. Int. J. Hydrogen Energy, 2010. 35: p.

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32. Yang, L., W. Yang, and Q. Cai, Well-Dispersed PtAu Nanoparticles Loaded into Anodic

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68

Chapter 3

Oxygen Reduction of Nanocrystalline Ruthenia - Local Structure Effects

3.1 Introduction

The fuel cell related electrocatalytic processes based on controlled hydrogen oxidation

and oxygen reduction have recently gained importance mainly in connection with the increasing

utilization of renewable energy sources. Despite efforts devoted to the optimization of existing

systems, the performance of real fuel cells still lags behind the expectations and the cathodic

oxygen reduction is seen as the performance limiting process. Simultaneously, the

electrochemical fuel cell reactions can also be generally employed in the energy storage mode

using the excess electricity or solar energy to generate energetically useful hydrogen (produced

along with the oxygen). This leads to the introduction of the regenerative fuel cell concept[1, 2].

It also needs to be stressed that the regenerative fuel cell applications have sparked extensive

catalyst development primarily for the oxygen evolution/reduction processes.

Oxygen electrochemistry, including oxygen evolution as well as reduction, represents the

simplest example of multiple electron charge transfer processes which have been extensively

studied both experimentally as well as theoretically[3, 4]. In contrast to the development of

suitable catalysts for independent oxygen evolution (OER) or oxygen reduction (ORR), the

catalysts' application for regenerative fuel cells faces significant restrictions in terms of

minimizing the energetic barriers of both kinetically irreversible processes. The fact that the

oxygen evolution reaction proceeds solely on oxide covered surfaces disfavors the use of metal

catalysts which are reported to be of superior activity in oxygen reduction. The oxide activity in

69

the oxygen evolution was traditionally investigated in both acidic as well as alkaline media on

various systems based on oxide of ruthenium[5-8], iridium[5, 7-9], cobalt[5, 10] or

manganese[5]. Oxygen reduction studies on oxides are less frequent and are generally restricted

to alkaline media. Oxygen reduction has been studied on rutile[11, 12], spinel[13, 14],

perovskite[15] and pyrochlore[16, 17] structural types based on ruthenium, manganese, nickel,

cobalt and iridium oxides. The investigated oxide catalysts were the subject of electrochemical

characterization which was phenomenologically analyzed in order to explain the possible

reaction pathways leading to both 4-electron and 2-electron oxygen reduction processes. In

contrast to the studies of oxygen evolution, no detailed investigations aiming at the role of the

catalyst structure, including the local structure of the oxygen reduction active site, have been

reported so far.

The theoretical approach allowing for the generalization of oxygen electrochemistry on

oxides based on DFT calculations was recently reported [4, 18, 19]. The DFT calculations

identify the active sites for oxygen activation and the charge transfer to so-called coordination

unsaturated sites (cus), the surface population of which is a function of the surface orientation.

The cus surface sites feature (n - 1) oxygen bonding partners, where n is equal to number of

metal-oxygen bonds present in the oxide bulk. It is believed that only cus sites can form the atop

reaction intermediate(s), which are essential to the oxygen electrochemistry. In this respect one

can easily predict that the catalytic activity and selectivity of oxide catalysts may be altered if

one controls the population and stacking of the cus sites at the oxide surface. This trend has been

shown for oxygen evolution on heterostatically doped ruthenia when the incorporation of lower

valency cations, such as Ni[20-23], Co[24-27], Fe[28] or Zn[29, 30] into ruthenia framework

70

resulted in changes to both the activity and selectivity of anodic processes including oxygen and

chlorine evolution. Similar systematic studies focused on other oxide systems are, so far, lacking.

This paper focuses on the role of the local structure of the oxide catalysts in the oxygen

reduction reaction. We report on the ORR activity of model nanocrystalline ruthenia catalysts

with local structure controlled by doping with Ni, Co, and Zn. The observed electrocatalytic

activity and selectivity are related to the actual local structures and rationalized using DFT-based

thermodynamic analysis of the oxygen reduction process.

3.2 Experimental

3.2.1 Materials Preparation

Ruthenium dioxide and doped samples of the composition Ru1-xMxO2 (M = Co, Ni, Zn)

were synthesized using the spray-freezing freeze-drying method as described in references[30,

31]. Generally, an 8 mM solution was prepared by dissolving the appropriate amount of

Ru(NO)(NO3)3 (31.3% Ru, Alfa Aesar) in 100 mL of Millipore H2O. In the case of doped

materials, a stoichiometric amount of the appropriate transition metal salt was added to the

solution. Zinc-doped samples were prepared from the acetate precursor, Zn(C2H3O2)2 • 2H2O

(99.5% ACS reagent grade, Fluka). Cobalt- and nickel-doped samples were prepared from the

nitrate salts, Co(NO3)2 • 6H2O and Ni(NO3)2 • 6H2O (99.999% trace metal basis, Sigma Aldrich),

respectively. The starting solution was then sprayed into liquid N2. The resulting ice slurry was

collected in an aluminum tray pre-cooled with liquid N2 and quickly transferred to a freeze-dryer

(FreeZone Triad Freeze Dry System 7400030, Labconco) pre-cooled to -30°C. The frozen

solvent was sublimated at reduced pressure (≈1.0 Pa) while the temperature was ramped

71

according to the following program: -30°C (1h), -25°C (5h), -20°C (4h), -15°C (6h), 30°C (4h).

After drying, the resulting powder was annealed in air at 400°C for 1 hour.

3.2.2 XRD, XPS, and SEM Characterization

The crystallinity of sample powders was characterized using a Rigaku Miniflex 600

powder X-ray diffractometer with CuKα radiation operating at 40 kV and 15 mA. The average

sample compositions were evaluated with X-ray energy dispersive spectroscopy using a Hitachi

S4800 scanning electron microscope (SEM) equipped with a Nanotrace EDX detector (Thermo

Electron). Sample compositions did not deviate significantly from the projected ones. Particle

size was evaluated by analyzing SEM images and averaging the size of 300 randomly chosen

particles. The X-ray photoelectron spectra (XPS) of the prepared materials were measured using

a modified ESCA 3 MkII multitechnique spectrometer equipped with a hemispherical electron

analyzer operating in the fixed transmission mode. Al Kα radiation was used for electron

excitation. The binding energy scale was calibrated using the Au 4f7/2 (84.0 eV) and Cu 2p3/2

(932.6 eV) photoemission lines. The spectra were collected at a detection angle of 45° with

respect to the macroscopic surface normal. The studied materials were characterized using

survey scan spectra and high resolution spectra of overlapping Ru 3d + C 1s photoelectrons, Ru

4s, Zn 2s and O 1s photoelectrons. The spectra were curve fitted after subtraction of Shirley

background using the Gaussian–Lorentzian line shape and nonlinear least-squares algorithms.

Quantification of the elemental concentrations was accomplished by correcting the photoelectron

peak intensities for their cross sections and for the analyzer transmission function. The typical

error of quantitative analysis by XPS is ~10%.

72

3.2.3 Electrochemical Measurements

The electrochemical oxygen reduction activity of the prepared materials was assessed in a

three-electrode single-compartment cell with a rotating ring-disk electrode (RRDE) setup (Pine

Instruments, USA). The potential was controlled using an Autolab PGSTAT30 (EcoChemie, The

Netherlands). Catalyst ink suspensions were prepared by sonicating 9.8 mg RuO2 or Ru1-xMxO2

(M = Co, Ni, Zn) with 5.00 mL Millipore water, 4.95 mL isopropyl alcohol, and 50 μL of 5 wt.%

Nafion® ionomer solution as a binder until the suspension was well dispersed. A 10.0 μL aliquot

of the ink was drop cast on a 0.196 cm2 glassy carbon disk electrode equipped with a platinum

ring to yield a total catalyst loading of approximately 50 μg cm-2

. All experiments were

conducted at room temperature in 0.1 M NaOH prepared from sodium hydroxide pellets

(semiconductor grade, 99.99%, Sigma-Aldrich). A platinum wire served as the counter electrode

and a saturated calomel electrode (SCE) served as the reference electrode. All potentials reported

are quoted against RHE. Electrolyte solutions were saturated with O2 for 30 minutes prior to

oxygen reduction measurements. The measured oxygen reduction currents were corrected for the

contribution of the capacitive current by subtracting the cyclic voltammograms obtained under

identical conditions in Ar saturated solution. Cyclic voltammograms were recorded at a scan rate

of 20 mV s-1

and the potential of the platinum ring electrode was held at 1.1 V vs. RHE during all

measurements. The ring collection efficiency was determined to be 0.275 according to the

procedure described in reference[32].

73

3.2.4 DFT Analysis of Oxygen Reduction

The thermodynamic analysis of the ORR on ruthenia based [110] surfaces was addressed

using GPAW (grid-based projector-augmented wave) DFT based code[33] together with the

ASE (atomic simulation environment)[34]. For all surfaces the exchange correlation functional,

revised Perdew Burke Ernzerhof[35], was used. The grid spacing selected was 0.18 and the

Brillouin zone was sampled using a 4 x 4 x 1 Monkhorst–Pack grid. The two model systems, the

non-doped and Ni-doped ruthenia [110], were approximated using a 2 x 1 and a 3 x 1 supercell,

respectively, with four atomic trilayers and with the bottom two trilayers fixed. The remaining

layers and adsorbates were relaxed until the residual forces in all directions were less than 0.05

eV Å-1

. The positions of the Ni atoms were modeled using the approach described in

reference[33]. The calculations containing Ni were spin-polarized.

3.3 Results and Discussion

3.3.1 XRD and SEM Characterization

X-ray diffraction patterns of all studied materials are shown in Figure 3.1. In all cases the

recorded patterns conform to a single phase tetragonal structure of the rutile type identical with

that of RuO2 (PDF file #431027). The average size of coherent crystallite domains was evaluated

using the Scherrer formula:

(3.1)

74

where Di is the size of the crystallite domain, λ is the wavelength of the incident radiation (CuKα

= 1.540598 Å), βi is the width of the diffraction peak at half maximum intensity measured in

radians, and θi is the angle of the hkl reflection.

Figure 3.8. Powder X-ray diffraction patterns for nanocrystalline RuO2 and Ru1-xMxO2 (M = Ni,

Co, Zn) after annealing in air for 1 hour at 400 °C

The average coherent domain size ranged between 4.3 and 5.7 nm (see Table 3.1).

Representative SEM images of the doped ruthenia are summarized in Figure 3.2. The particle

sizes evaluated from SEM micrographs agree with the coherent domain size values (See Table

3.1). Average sample compositions did not deviate from the projected ones and are listed in

Table 3.1.

The surface composition of all doped samples reflects the metastable character of the

materials and previous thermal treatment which result in a dopant enrichment of the surface

layer.[21, 36] This effect is most pronounced in the case of the Zn doped materials when the

75

actual surface compositions of both studied materials correspond to Ru0.73Zn0.27O2.43 and to

Ru0.63Zn0.37O2.23 for the materials with nominal Zn contents of 0.1 and 0.2, respectively. It needs

to be noted that the apparent excess of the oxygen in the surface composition can attributed to

surface OH groups as well as to adsorbed water.

Table 3.1. Results of the analysis of the powder diffraction data for RuO2 and doped RuO2

samples.

Actual Composition Coherent Domain

Size [nm]

Strain

[%] a [Å] c [Å]

Particle

Size [nm]

RuO2 5.7 0.46 4.470 3.120 7.6 ± 2.2

Ru0.9Zn0.1O2-z 4.9 0.61 4.526 3.108 8.9 ± 2.2

Ru0.82Zn0.18O2-z 5.5 0.00 4.519 3.099 5.8 ± 1.5

Ru0.95Ni0.05O2-z 5.3 0.00 4.515 3.096 7.2 ± 1.4

Ru0.91Ni0.09O2-z 5.0 0.00 4.501 3.079 7.9 ± 2.3

Ru0.90Co0.10O2-z 4.3 0.21 4.505 3.081 7.4 ± 1.8

Figure 3.2. SEM images of nanocrystalline (a) RuO2, (b) Ru0.90Zn0.10O2, (c) Ru0.80Zn0.20O2, (d)

Ru0.95Ni0.05O2, (e) Ru0.90Ni0.10O2, and (f) Ru0.90Co0.10O2 after annealing at 400 °C in air for 1

hour.

76

3.3.2 Electrochemical Measurements

All prepared ruthenia materials are active ORR catalysts in alkaline media. The ORR

polarization curves for RuO2 and Ru1-xMxO2 (M = Ni, Co, Zn) samples are shown in Figure 3.3.

The disk current (iD) which reflects the oxygen reduction shows a pronounced peak at

approximately 0.40 V to 0.55 V before approaching a mass transport controlled region. The disk

current feature can be tentatively associated with a change in the Ru oxidation state from Ru(IV)

to Ru(III)[11, 37]. This process is usually connected with cation insertion into the oxide structure

to balance the charge in cationic and anionic sub-lattices[36]. The behavior giving rise to the

peak in the disk current is also manifested in the ring current, indicating a pronounced formation

of hydrogen peroxide in this potential region. The hydrogen peroxide formation in the 0.40 V to

0.55 V interval seems to be unaffected by the chemical composition of the catalyst. In addition,

the formation of hydrogen provide seems to be suppressed with increasing rotation rate. The

precise mechanism of this reduction process is, however, not evident.

77

Figure 3.9. ORR polarization curves and ring currents at 1600 rpm for RuO2 and Ru1-xMxO2 (M

= Ni, Co, Zn) electrodes at 20 mV s-1

in O2 saturated 0.1 M NaOH. Ering = 1.1 V vs. RHE.

The overall ORR activity of the doped ruthenia catalysts is lower than that of the non-

doped ruthenia. The ORR activity as reflected in the disk currents (iD) generally decreases for the

Co- and Ni-doped samples. The corresponding ring currents (iR) are, however, higher than that of

the non-doped ruthenia, particularly at high overpotentials (i.e. at potentials negative to 0.4 V vs.

78

RHE). This shows a pronounced tendency of Co and Ni-doped materials to produce H2O2

namely at high overpotentials (n ranging between 3.0 and 3.4). In contrast, the Zn-doped

materials show a preference for the 4-electron reduction pathway with n values ranging between

approximately 3.6 to 3.8 while the activity remains comparable to that of the non-doped ruthenia.

Also the selectivity of the doped ruthenia in oxygen reduction is controlled by the doping process

itself rather than by the actual dopant content.

The observed behavior reflects the surface sensitivity of oxygen reduction on oxide

surfaces, which can be related to the surface local structure. Quantitative visualization of this

behavior is shown in Figure 3.4, which plots the potential dependence of the average number of

electrons transferred to an oxygen molecule on different doped ruthenium dioxide materials as

calculated from the Koutecky-Levich equation[38]:

(3.2)

where F is Faraday's constant, A is the geometric area of the electrode, D is the diffusion

coefficient (1.90 × 10-5

cm2 s

-1), ν is the kinematic viscosity (8.70 × 10

-3 cm

2 s

-1), and C is the

bulk concentration of O2 (1.22 × 10-6

mol cm-3

)[39].

79

Figure 3.4. Potential dependence of the average number of electrons transferred during oxygen

reduction on RuO2 and Ru1-xMxO2 (M = Ni, Co, Zn) electrodes. Presented data were calculated

using Koutecky-Levich equation.

It has to be stressed that in contrast to the behavior known for metal electrocatalysts in

acid media, the oxygen reduction on ruthenia based catalysts apparently forms primarily

hydrogen peroxide, namely at low overpotentials. The observed selectivity of ruthenia-based

catalysts in ORR shows a more complex potential dependence which can be treated either by a

phenomenological or a local structure sensitive approach.

Assuming a general phenomenological model of the oxygen reduction mechanism as

proposed previously (See Fig. 3.5)[40, 41], oxygen can be reduced to water (4-electron

pathway) either directly or sequentially with H2O2 as the main adsorbed intermediate.

In principle, H2O2 either desorbs and can be detected on the ring or can be further

reduced to water in the second 2-electron reduction process. The measured disk current

80

summarizes the current contributions from the complete 4-electron reduction to H2O and the 2-

electron reduction to H2O2 while the recorded ring current is proportional only to the amount of

oxygen reduced to H2O2. In this respect a ratio of iD/iR can be used as an indicator of the actual

mechanism which should yield a straight line proportional to k1/k2 when plotted against ω-1/2

(See

Fig. 3.6)[41].

Figure 3.5. Phenomenological mechanism of oxygen reduction according to reference[41].

Figure 3.6. |ID/IR| vs. ω-1/2

plots for (a) RuO2, (b) Ru0.80Zn0.20O2, and (c) Ru0.90Ni0.10O2.

Presented data were extracted from RRDE experiments carried out in O2 saturated 0.1 M NaOH.

The actual iD/iR data deviate from linearity (see Figure 3.6) as can be expected since the

formalism incorporated in the scheme depicted in Fig. 3.5 disregards the nature of the individual

reaction steps composing both 2- and 4-electron reduction pathways and their different

dependence on the electrode potential.

81

The individual rate constants k1, k2, and k3 were evaluated from ORR data assuming that

all three processes proceed simultaneously and that the values of k-1, k-2, and k-3, corresponding

to reversed reactions, are negligible. The adsorption of oxygen on the electrode surface is also

assumed to proceed sufficiently fast. The potential dependence of the rate constants for all

considered catalysts in the overall oxygen reduction mechanism is shown in Figure 3.7.

Figure 3.10. Potential dependence of the rate constants for the reduction of O2 to H2O (k1), of O2

to H2O2 (k2), and H2O2 to O2 (k3) on nanocrystalline ruthenia based catalysts. The presented data

correspond to experiments carried out in O2 saturated 0.1 M NaOH at 1600 rpm.

It seems that the conversion of H2O2 to H2O through the series pathway (k3) is negligible

on all electrode surfaces at high overpotentials. It has to be noted, that although the k3 values are

negligible with respect to k1 and k2 there is a significant difference between k3 of the Zn-doped

and non-doped samples and those obtained for the Ni- and Co-doped samples. The values of k3

observed for Ni- and Co-doped samples are approximately one order of magnitude lower and

82

seem to correspond to decreased tendency of these materials to reduce oxygen through the 4-

electron pathway.

As follows from Figure 3.7, the conversion of O2 to H2O2 appears to be the dominant

process on ruthenium based oxides at low overpotentials. In this respect the reduction behavior

of the ruthenia differs significantly from that of metals which prefer the 4-electron reduction at

low overpotentials. The role of the chemical composition in selectivity of doped catalysts

towards 2- and 4-electron reduction pathways can be visualized by the potential at which the

catalytic system shows the same preference for the 4-electron and 2-electron reaction pathways,

i.e. potential at which k1/k2 = 1 (see Figure 3.8).

A fundamental description of the oxygen reduction on oxide surfaces can be based on the

thermodynamic analysis of the observed trend, which highlights the enhanced tendency of the

Ni- and Co-doped materials to form hydrogen peroxide compared with non-doped and Zn-doped

ruthenia and apparently reflects the local structure of the doped ruthenium oxides.

Figure 3.11. The potential of equal rate in 2- and 4-electron reduction for different ruthenia

based catalysts.

83

3.3.3 DFT Analysis of Oxygen Reduction

A fundamental description of the oxygen reduction on oxide surfaces can be based on the

thermodynamic analysis of the process utilizing the DFT modeling. Reverting to the formalism

used for the oxygen evolution reaction we can describe the overall reduction process as a

sequence of four consecutive concerted electron/proton transfers - if one aims for the complete 4-

electron reduction - or of two consecutive electron/proton transfers if hydrogen peroxide is

considered as the reaction product. The results of the DFT investigations of ORR reduction on

ruthenium dioxide based catalysts are summarized in Figures 3.9-3.12.

A systematic description of the stable surface structures at different potentials represents

a prerequisite step in the theoretical investigation of oxygen reduction on an oxide surface, which

in this case is the [110] rutile surface of ruthenia. This procedure results in computational

Pourbaix diagrams where the stable surface at any given potential features the highest

stabilization (i.e. the most negative surface energy) of the system.

Bearing in mind that the [110] oriented surface of a rutile type oxide features the

transition metal cations in two local environments, cus and bridge, one can visualize the surface

of non-doped ruthenium dioxide as changing from the surface structure characterized by

protonated oxygen on cus sites and deprotonated oxygens in bridge sites (region C in Figure 3.9)

to the surface featuring vacant cus and protonated oxygens in bridge sites (region A in Figure

3.9). Since the ORR was not observed at potentials positive to 0.7 V (vs. RHE) one can restrict

the DFT investigations of the oxygen reduction on conventional RuO2 to the surface stable in the

region (A). In the case of doped ruthenia (as shown in the case of the Ni doped material

presented in Figure 3.10) one needs to consider the complexity arising from the chemical

composition when both types of transition metal cations enter the cus and bridge positions. This

84

variability in the chemical composition also increases the number of distinctive oxygen atoms

available at the surface, the binding energy of which depends on their nearest neighbors.

Although the electrode potential dependent variability of the surface structure is more

pronounced in this case, the stable structures predicted from the potential range in which the

ORR proceeds is qualitatively the same and corresponds to vacant cus sites complemented by

protonated oxygen atoms connecting the bridge sites (see structure D in Figure 3.10).

Figure 3.12. Surface Pourbaix diagram for RuO2. Detailed description of the diagram

construction is given in the supplementary information.

85

Figure 3.13. Surface Pourbaix diagram for Ni-doped RuO2. Detailed description of the

diagram/s construction is given in the supplementary information.

The DFT models predict that the entire process begins with oxygen adsorption at

coordination unsaturated (cus) cationic sites. The behavior of both the non-doped and the doped

ruthenia catalysts is controlled by the local structure and depends on the nature of the cation

residing in the cus site as well as on the electrode potential. In the case that the cus site is

occupied by a ruthenium cation (which are present on all investigated catalysts) the first electron

reduction forms a rather strongly bound *OOH intermediate, which is more stable than the

hydrogen peroxide at most reasonable electrode potentials (see Figure 3.11). Consequently, the

further reduction of the *OOH intermediate located on Ru cus site cannot form hydrogen

peroxide unless one uses a rather strong external electric field to weaken the *OOH binding to

the surface. The actual potential(s) at which hydrogen peroxide formation becomes

thermodynamically allowed are given in the legend of the Figure 3.11.

86

Figure 3.11. Free energy diagrams for the reduction of O2 on three catalytic active sites, the Ru

cus site on conventional ruthenia (green), the cus Ru site on Ni doped RuO2 (magenta) and the

cus Ni site on Ni doped RuO2 (blue). The dotted line represents the equilibrium potential of the

reduction O2 to H2O2. The key difference is the binding of O on the Ni cus site compared to the

Ru cus sites.

In the case that the cus site is occupied by an heteroatom, e.g. Ni or Co, (see Figure 3.11)

one observes a significantly weaker binding of the *OOH and *O compared with the Ru

occupied cus sites. This fact decreases the potential at which the reduction on Nicus starts to

contribute to the overall reduction process. The weak interaction of the *OOH with the

heteroatom-containing cus site restricts the presence of such an adsorbate in the potential region

with low total surface coverage, i.e. to relatively high over-potentials. It needs to be noted

though, that the formation of hydrogen peroxide from *OOH confined on a heteroatom occurs at

much more positive potentials than in the case of *OOH confined to Ru-containing cus sites and

further reduction of the *OOH intermediate can proceed via the 4-electron or 2-electron

reduction pathway with approximately the same probability.

Figure 3.12 shows the dependence of the electrode potential needed to drive the oxygen

reduction on oxide based surfaces either via the 4-electron (red) or 2-electron (blue) reaction

pathway as a function of the reaction descriptor - i.e. adsorption energy of the *OOH

intermediate. It needs to be noted that in a similar manner one may describe the reaction with the

87

adsorption of *OH due to the interdependence of the adsorption energies of the intermediate

formed in the first and third charge transfer step[3, 33].

Figure 3.14. Volcano plot for the 2-electron (blue) and 4-electron (red) reduction of O2 to H2O2


represent the equilibrium potentials for the reduction products. In the case of the Ni-doped

ruthenia the limiting over-potential for both possible reaction sites (Rucus and Nicus) are shown

along with that of conventional ruthenia.

Such a dual volcano plot has been used with great success in literature[42, 43]. The

volcano curves presented in Figure 3.12 clearly show a quantitative prediction of the

thermodynamic preference of the 4-electron reduction pathway over the 2-electron reduction on

strongly adsorbing cus sites. As follows from Figure 3.11, oxygen reduction on Ni-doped

ruthenia should proceed at slightly more positive potentials compared with the non-doped

ruthenia as long as the cus sites are occupied with Ru cations. The reduction process on Ru

occupied cus sites should show a pronounced preference for 4-electron reduction and the

formation of hydrogen peroxide should be excluded for potentials positive of 0.14 V or 0.43 V

(vs. RHE) for non-doped ruthenia and the Ni-doped material, respectively (see Figure 3.11). The

88

easier formation of hydrogen peroxide on the Ni-doped material should be compensated by an

earlier onset of the oxygen reduction process as predicted for Ni-doped material. In the case of

weakly adsorbing sites, e.g. in the case of Ni cus sites - there is no apparent thermodynamic

preference for either the 4- or 2-electron reduction pathway. The DFT model predicts the onset

of the oxygen reduction process to occur at potentials comparable with the ORR on non-doped

ruthenia. The formation of hydrogen peroxide is possible at significantly more positive potentials

(see Figure 3.11)

Analyzing the experimental behavior of the ruthenium dioxide based catalysts for the

oxygen reduction process in the light of the DFT results one can qualify the existence of two

classes of catalysts – one favoring the 4-electron reduction (non-doped RuO2 and Zn-doped

RuO2) and another showing significant activity in hydrogen peroxide production (Ni- and Co-

doped ruthenia). Realizing that the Zn present in the Zn-doped ruthenia is itself redox inactive

one can assume that the catalysts in the first group have all active cus sites occupied by Ru

regardless of the actual chemical composition. The confinement of the catalytic activity on Ru

itself justifies the selectivity towards 4-electron reduction pathway as it is shown in Figure 3.4

and Figure 3.8. In the case of the Co- and Ni-doped ruthenia the significant amount of hydrogen

peroxide formed in the process can be attributed primarily to the Ni/Co cus sites although the Ru

cus sites also contribute to the hydrogen peroxide formation at lower potentials. In contrast to the

complementary oxygen evolution process, the Ni (or Co) ions located in the bridge sites, which

play a crucial role in the complementary anodic process[33], apparently have no effect on the

oxygen reduction activity of these materials. A different role of the catalysts local structure in

oxygen reduction is not entirely surprising given the irreversibility of oxygen

evolution/reduction.

89

The DFT calculations, however, fail to explain pronounced formation of the hydrogen

peroxide on all ruthenium based catalysts at low overpotentials (0.55-0.40 V) when the hydrogen

peroxide on Ru cus sites should be thermodynamically excluded. Given the relatively short

timescale of the RRDE experiments one may therefore suggest that the system fails to reach the

thermodynamically stable surface structure on the experimental timescale and the hydrogen

peroxide is released from meta-stable intermediates not reflected in the DFT calculations.

3.4 Conclusions

Nanocrystalline ruthenia based electrocatalysts offer a convenient model for investigating

the role of the local structure in the oxygen reduction on oxide electrodes. The oxygen reduction

related activity of RuO2 is comparable with that of the doped ruthenia . The selectivity of doped

ruthenia catalysts differs from that of the RuO2 in which the non-doped as well as Zn-doped

catalysts prefer 4-electron oxygen reduction while the Ni- and Co-doped ruthenia produce

significant amount of hydrogen peroxide. The observed selectivity trends can be rationalized

using a thermodynamic analysis of the oxygen reduction process based on DFT calculations.

The DFT based analysis confines the oxygen reduction activity to cus sites, the

occupancy of which controls the selectivity of the oxygen reduction process. Oxygen reduction

on non-doped ruthenium dioxide is controlled by the fourth electron transfer. Doping the

ruthenium dioxide shifts the potential control to the first electron transfer. This trend can be

attributed to decreasing occupancy of the cus sites with ruthenium. The strong adsorption of the

*OOH intermediate on the Ru cus site steers the reaction mechanism towards 4-electron

reduction pathway. Incorporation of reactive transition metal cations into bridge sites has a

90

negligible effect on the ORR activity. A confinement of the reactive transition metal into cus

sites weakens the adsorption of the reaction intermediates and opens the 2-electron reaction

pathway at relatively low overpotentials.

3.5 Acknowledgments

This work was supported by the European Commission within the Initial Training

Network ELCAT (Project No. 214936). The support of the Danish Ministry of Science,

Technology and Innovation though the CASE is also gratefully acknowledged. Niels Bendtsen

Halck and Jan Rossmeisl of Technical University of Denmark are gratefully acknowledged for

performing all DFT calculations and for their assistance in the subsequent analysis. The authors

would like to thank Valery Petrykin of the J. Heyrovský Institute of Physical Chemistry for his

significant contributions to the EXAFS-based analysis of the doped ruthenia and Zdeněk Bastl,

also from the J. Heyrovský Institute of Physical Chemistry, for his assistance with the XPS

measurements.

91

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95

Chapter 4

Selective Chlorine Evolution Catalysts Based on Mg-Doped Nanoparticulate Ruthenium

Dioxide

4.1 Introduction

Understanding the electrocatalytic behavior of transition metal oxides for the oxygen

evolution (OER) and chlorine evolution reactions (CER) is fundamental to the chlor-alkali

industry as well as water electrolysis. Iridium and ruthenium based oxides have long been the

center of attention in this respect since they represent the industrial benchmark catalysts due to

their exceptional activity and stability for these reactions[1-3]. Although the traditional approach

in the development of oxide electrocatalysts aims primarily at the optimization of the oxide’s

activity, the issue of systematic selectivity control receives more attention in connection with the

rational design of oxygen evolution catalysts and with the rational control of the electrocatalytic

processes. Both anodic gas evolution processes are among the least thermodynamically favored

redox reactions as shown by the rather high standard electrode potential of each reaction:[4]

O2 + 4H+ + 4e

− ⇌ 2H2O E° = 1.229V (4.1.1)

Cl2(g) + 2e− ⇌ 2Cl

− E° = 1.358V (4.1.2)

Accepting the fact that the industrially relevant production of each gas requires a non-zero

overpotential, one has to anticipate a simultaneous production of oxygen and chlorine at most

potentials, making chloride contamination one of the more serious issues in water electrolysis.

The activity of ruthenium dioxide in both gas evolving reactions was systematically

investigated by means of the density functional theory (DFT) based thermodynamic analysis.

96

DFT analysis has been able to offer some insight regarding the competitive nature of both

processes, confirming that the OER and CER are strongly correlated on rutile-type oxides[5].

According to the DFT analyses, the oxygen evolution reaction proceeds primarily through the

formation of surface confined hydroxo, oxo, or peroxo intermediate species on pentacoordinated

Ru sites, also denoted as coordination unsaturated sites (cus)[6, 7]. Similarly, chlorine evolution

on ruthenium dioxide is most likely to proceed via the formation of peroxo species bridging two

adjacent cus sites (O2,cus) which are subsequently subject to oxidative attack by Cl−[5]. This

theoretical description was recently challenged by Over et. al which assumes that cus confined

oxo species are the most stable surface structure at the given conditions[8].

The theoretical analysis also shows that the activity in chlorine evolution can be

described with the same reaction descriptor as the oxygen evolution process, which in this case is

the oxygen adsorption energy[5]. Despite the less favorable standard potentials, the theory

confirms that the CER process is more facile than oxygen evolution and that the theoretical

limiting overpotential required to drive CER is lower than that of OER in the whole range of

oxygen binding energies[6]. This fact should lead to a complete suppression of the oxygen

evolution activity on oxide surfaces even at moderate chlorine concentrations.

Keeping in mind the interdependence of the oxygen and chlorine evolution processes one

may anticipate that the catalyst selectivity may be affected by controlling the oxygen binding

energy to the oxide surface in a manner similar to that employed for controlling the oxygen

evolution activity. Cationic substitution turns out to be a versatile approach in controlling the

oxygen binding energy and has been demonstrated in the case of ruthenium dioxide modified

with Ni[9-12], Co[11-17], Zn[18], Sn[19-21], and Fe[22]. Experiments combining voltammetric

methods with differential electrochemical mass spectrometry (DEMS) have shown that the

97

incorporation of the heterovalent cations generally affects the selectivity of the ruthenia-based

catalyst in oxygen and chlorine evolution[9, 23, 24]. A general description of the relationships

between the nature of the modified ruthenia-based catalyst and its selectivity is so far missing.

The issue of the oxide catalyst selectivity is addressed in this paper which also extends

the previous studies on the selectivity of rutile-type oxides. In the current work we describe the

synthesis of Mg-modified ruthenium dioxide containing various amounts of Mg and match their

activity and selectivity in parallel oxygen and chlorine evolution with detailed structural

correlations using extended X-ray absorption fine structure (EXAFS).

4.2 Experimental

4.2.1 Materials Preparation

Ruthenium dioxide and Ru1-xMgxO2 materials were synthesized using the spray-freezing

freeze-drying method as outlined in reference [25]. Aqueous solutions (8 mM) were prepared

from ruthenium (III) nitrosylnitrate (31.3% Ru, Alfa Aesar) and magnesium acetate tetrahydrate

(Puratronic®, 99.997% metals basis, Alfa Aesar) in 100 mL of Millipore H2O. The solutions

were then sprayed into liquid N2 to create fine ice particles. The resulting ice slurry was

transferred to an aluminum tray precooled with liquid N2 and placed in the freeze dryer

(FreeZone Triad Freeze Dry System 7400030, Labconco) precooled to −30 °C. The pressure was

decreased to approximately 1.0 Pa and the temperature was ramped according to the following

program: −30 °C (2h), −25 °C (5h) −20 °C (6h), −15 °C (5h), 30 °C (4h). Afterwards, the

resulting powder was annealed in the furnace at 400 °C for one hour.

98


Crystallinity and phase purity of the prepared materials was checked using a Rigaku

Miniflex 600 powder X-ray diffractometer with CuKα radiation. Morphology of the synthesized

catalysts was characterized using a Hitachi S4800 scanning electron microscope (SEM) and a

Nanotrace EDX detector (Thermo Electron) was used to evaluate the average sample

compositions by X-ray energy dispersive spectroscopy. Particle size was evaluated by analyzing

SEM images and averaging the size of 300 randomly chosen particles.

The X-ray photoelectron spectra (XPS) of the prepared materials and of the electrodes

which were used in electrochemical experiments were measured using a modified ESCA 3 MkII

multitechnique spectrometer equipped with a hemispherical electron analyzer operating in the

fixed transmission mode. Al Kα radiation was used for electron excitation. The binding energy

scale was calibrated using the Au 4f7/2 (84.0 eV) and Cu 2p3/2 (932.6 eV) photoemission lines.

The spectra were collected at a detection angle of 45° with respect to the macroscopic surface

normal. The studied materials were characterized using survey scan spectra and high resolution

spectra of overlapping Ru 3d + C1s photoelectrons, Ru 4s, Mg 2s, O 1s photoelectrons and Mg

KL23L23 Auger electrons. The spectra were curve fitted after subtraction of Shirley background

using the Gaussian−Lorentzian line shape and nonlinear least-squares algorithms. Quantification

of the elemental concentrations was accomplished by correcting the photoelectron peak

intensities for their cross sections and for the analyzer transmission function. The typical error of

quantitative analysis by XPS is ∼10%.[26]

99

4.2.3 Local Structure Characterization

EXAFS spectra were measured at the Photon Factory synchrotron of the High Energy

Accelerator Organization (KEK) in Tsukuba, Japan. The spectra were measured in transmission

mode at the Ru K edge at beam-line AR-NW10A (Si(311) monochromator); the Mg K

absorption edge was measured in total electron yield mode at the BL-11A beam-line (grazing

incidence monochromator). Ru K scans extended to 20 Å−1

and Mg K data were limited to 15

Å−1

. Each spectrum was recorded at four different scanning step sizes: the preedge region from

500 to 50 eV was scanned in 6.5 eV steps to enable background subtraction; in the 50 eV pre-

edge and 100 eV post-edge range a step size of 0.4–0.5 eV was used to acquire the XANES part

of the spectra, while a 2.5–3.0 eV and 7.0 eV scanning step was maintained in the post-edge

regions of 100–500 eV and above 500 eV, respectively.

All data handling pre-requisite to the local structure refinement of the extended X-ray

absorption fine structure (EXAFS) functions (i.e., normalization, smoothing and background

subtraction, the Fourier transforms of the spectra and windowing) was done in the IFEFFIT

software package.[27] The photoelectron wave vector k for the Fourier transform of spectra was

kept within the range of k = 3–14Å−1

for Ru-EXAFS and k = 3–12 Å−1

for Mg-EXAFS. The k-

weighting factor of 2 was applied. For the analysis of the local structure of Ru1-xMgxO2 materials

a full-profile refinement of the EXAFS spectra by non-linear least squares (NLLS) minimization

in the R-space with a k-weighting factor equal to 2 was carried out using the Artemis NLLS

module of the IFEFFIT package. The theoretical model was generated using FEFF6.2 library.

100

4.2.4 Electrochemical and DEMS Measurements

The electrocatalytic activity and selectivity in the oxygen evolution and chlorine

evolution reactions were evaluated using a combination of potentiostatic experiments with

differential electrochemical mass spectrometry (DEMS).Working electrodes were prepared from

water and isopropanol based catalyst suspensions by sedimentation on Ti mesh (Goodfellow,

20% open area). Catalyst suspensions were pre pared by mixing 10.0 mg RuO2 or Ru1-xMgxO2

with 2.0 mL millipore H2O and 2.0 mL isopropyl alcohol with subsequent homogenization in an

ultrasonic bath and drop cast onto the Ti mesh current collector. The electrodes were dried in air

at 60 °C repeatedly until the desired catalyst loading (1–2 mg cm−2

) was obtained. Electrodes

were annealed at 400 °C in air for 2 hours to stabilize the catalyst layer. The measured current

values were normalized with respect to the actual surface area based on the known mass of the

catalyst and particle size based specific surface area.

Electrochemical measurements were performed in a three electrode arrangement in a

home-made Kel-F single compartment cell using a Pt auxiliary electrode and sat. Ag/AgCl

reference electrode in 0.1 M HClO4. The reference electrode itself was placed outside the cell to

avoid chloride leakage; the conductive connection of the reference electrode was achieved by a

Luggin capillary. The potential of the working electrode in all experiments was controlled using

a PAR 263A potentiostat. The DEMS apparatus consisted of a PrismaTM QMS200 quadrupole

mass spectrometer (Balzers) connected to a TSU071E turbomolecular drag pumping station.

101



X-ray diffraction patterns for the synthesized materials are shown in Figure 4.1. The

recorded patterns conform to a single phase tetragonal structure of the rutile-type as in RuO2

(PDF file #431027). The incorporation of Mg into the rutile structure results in a change of the

unit cell parameters which can be evaluated from the diffraction patterns. Assuming the

oxidation state of Mg in the prepared material to be II, its incorporation into the RuO2 framework

should results in a slight expansion of the unit cell volume in the Mg-doped ruthenia

(anticipating Mg to maintain octahedral coordination) due to the slightly larger ionic radius of

the divalent magnesium (0.72 Å) with respect to that of the tetravalent ruthenium (0.62 Å)[26].

This trend, however, is not experimentally encountered (see Table 4.1). Incorporation of

magnesium leads to an expansion of the lattice parameter a (which increases monotonously with

increasing Mg content) and is accompanied by a significant shortening of the lattice parameter c,

which further decreases with increasing Mg content. Although the unit cell volume of Mg-

modified materials increases with increasing Mg content, its concentration dependent increase

does not compensate for the initial decrease in unit cell volume associated with the incorporation

of Mg. Similar unit cell shortening in the [001] direction has been previously reported for Ni-

doped ruthenia and can be taken as an indicator that Mg is not evenly distributed in the resulting

material[10]. The observed trend in the unit cell volume may also be interpreted in terms of the

Mg local environment, specifically that of the oxygen coordination number with increasing Mg

content. Such an interpretation should be, however, supported by local structure data based, e.g.

on X-ray absorption spectroscopy. It needs to be noted that this change should be most

pronounced in materials with a low magnesium content (x < 0.1). The single phase nature of the

102

recorded XRD patterns suggests that the actual chemical composition agrees with the projected

composition. The magnesium seems to be homogeneously distributed between bulk and surface

of the nanoparticles as evidenced by the XPS-based surface composition and projected average

composition (see Figure 4.2 and Table 4.2). This agreement is, however, weakened by the fact

that the Mg XPS data are not available for samples with a Mg content lower than 0.2.

Figure 4.1. Powder X-ray diffraction patterns for nanocrystalline RuO2 and Ru1-xMxO2 after

annealing in air for 1 hour at 400 °C.

Table 4.1. Results of the analysis of the powder diffraction data of the MgxRu1-xO2 catalysts

Sample a [Å] c [Å] V [Å3]

Coherent domain

size [nm]

Particle size

[nm] Strain (%)

RuO2 4.454(15) 3.21(8) 63.68 4.3 9.6 ± 3.0 1.0(3)

Ru0.95Mg0.05O2 4.478(13) 3.12(2) 62.56

5.9 5.3 ± 1.4 0.7(4)

Ru0.90Mg0.10O2 4.505(3) 3.105(4) 63.02

4.6 5.0 ± 1.3 1.0(5)

Ru0.80Mg0.20O2 4.509(3) 3.108(3) 63.19

12.0 5.5 ± 1.8 0.3(2)

103

Figure 4.2. Survey scans (a) of Ru0.8Mg0.2O2 before (1) and after (2) electrochemical

experiments. Panes (b) and (c) contain fitted high resolution spectra of Ru 3d + C 1s (b) and Ru

4s + Mg 2s photoelectrons, respectively. The high resolution spectra of the Ru0.8Mg0.2O2 after

electrochemical experiments do not show significant differences from those plotted in panes (b)

and (c).

Table 4.2. XPS based surface metal content in Ru0.8Mg0.2O2.

Sample XRu XMg

As synthesized 0.80 0.20

After Electrochemistry 0.79 0.21

104

The relatively broad diffraction peaks are compatible with the anticipated nanocrystalline

character of the materials. The average coherent domain size of the prepared materials can be

approximated using the Scherrer formula:

(4.3.1)

where Di is the size of the crystallite domain, λ is the wavelength of the incident radiation (CuKα

= 1.540598 Å), βi is the width of the diffraction peak at half maximum intensity measured in

radians, and θi is the diffraction angle. The diffraction peak analysis yields an average coherent

domain size ranging between 4 and 12 nm (see Table 4.1). The nanocrystalline nature of the

prepared materials is also confirmed in the SEM micrographs (see Fig. 4.3). Regardless of the

actual composition, the Mg-doped ruthenia forms isometric crystals with an average size of

approximately 5 nm (see Table 4.1) complemented with larger particles of approximately 40 nm.

The coarse 40 nm particles are, however, significantly less represented. Despite the actual

underrepresentation of the coarse nanoparticles, they form the principal diffracting regions

controlling the shape of the diffraction peaks, particularly in the material containing 20% Mg.

This fact results in a significant discrepancy between the coherent domain size and the actual

particle size such that the coherent domain size exceeds the particle size based on analysis of the

micrographs.

105

Figure 4.3. SEM images of nanocrystalline (a) RuO2, (b) Ru0.95Mg0.05O2, (c) Ru0.90Mg0.10O2, and

(d) Ru0.80Mg0.20O2 after annealing in air for 1 hour at 400 °C.

4.3.2 Local Structure Characterization

Information complementary to the diffraction is extractable from the X-ray absorption

spectra measured on the Ru and Mg K edges (See Figure 4.4). The experimentally observed Ru

K edge position of the Mg doped ruthenia is not sensitive to the actual Mg content and is equal to

22133 eV. This value agrees well with that reported for ruthenium dioxide in the literature[9,

13]. The corresponding EXAFS functions as they are presented in Figure 4.4a are also

insensitive to the magnesium content and agree qualitatively with the Ru EXAFS functions

106

reported for doped ruthenia previously[10]. The observed behavior is generally compatible with

the non-homogeneous distribution of magnesium in the RuO2 matrix, resulting in the formation

of Mg enriched clusters, the nature of which may depend on the overall Mg content. The

magnesium K edge position in the X-ray absorption spectra of the Mg doped ruthenia shifts with

increasing Mg content from ca. 1310 eV (material containing 5% Mg) to 1308 eV (material

containing 20% Mg). The Mg content-dependent shift of the Mg K edge position may be

interpreted either in terms of a change in oxidation state or in terms of a change in local

structure. The observed Mg absorption edge values, primarily those for the materials with low

total Mg content, exceed that observed for the MgO standard (1309 eV) and could be formally

assigned to an increase in the oxidation state. Keeping in mind the alkaline earth nature of

magnesium, however, the assignment of the observed trend in the position of the Mg absorption

edge energy to the overall effects of bonding arrangements in the vicinity of Mg seems to be

more likely. The Mg EXAFS functions obtained from the Mg K edge X-ray absorption spectra,

in contrast to the Ru EXAFS functions, are affected by the actual Mg content (see Figure 4.4b).

Figure 4.4. EXAFS functions extracted from the X-ray absorption spectra of the Ru1-xMgxO2

(0<x<0.2) measured on the Ru K edge (a) and Mg K edge (b). Actual Mg content is shown in the

Figure legend.

107

The Ru EXAFS functions can be refined assuming a structural model based on ruthenium

dioxide with a rutile structure. A typical example of the refinement for the Ru EXAFS functions

is plotted in Figure 4.5a; the results of the EXAFS function refinement are summarized in Table

4.3a. It may be concluded that the bonding arrangements in the ruthenium environment are

essentially unperturbed by the incorporation of Mg. The Ru-O as well as the Ru-Me (Me = Ru or

Mg) bonding distances and the Ru coordination arrangements do not differ significantly from

those of pure ruthenium dioxide. The presence of Mg in the direct neighbor metal positions with

respect to the absorbing ruthenium is lower than that corresponding to the average chemical

composition.

Figure 4.5. (a) A typical example of the non-linear least square fit of the Ru EXAFS function of

Ru0.95Mg0.05O2; (b) A typical example of the non-linear least square fit of the Ru EXAFS

function of Ru0.90Mg0.10O2. The square symbols represent the experimental data, the red line

denotes the best fit.

108

Table 4.3. (a) Results of the NLLS fit of the EXAFS functions obtained from the Ru Kedge X-

ray absorption spectra of the MgxRu1-xO2 (0 < x < 0.2) catalysts. CN denotes the coordination

number and d stands for the bonding length; (b) Theoretically conceived local structures of the

Mg in rutile and ilmenite type MgRuO oxides and results of the NLLS fit of the EXAFS

functions obtained from the Mg K edge X-ray absorption spectra of the MgxRu1-xO2 (0<x<0.2)

catalysts. The theoretically conceived data are given in italics. Symbol assignment are the same

as in the case of Table 3a.

(a)

(b)

Sample Bond CN Error d [Å] Error Sample Bond CN Error d [Å] Error

Ru0.95Mg0.05O2 Ru-O 2.0 0.2 1.937 0.006 Rutile Mg-O 2 1.94

4.0 0.4 1.982 0.006 4 1.98

Ru-Ru 1.8 0.4 3.100 0.010 Mg-Ru 2 3.14

Ru-O 4.0 0.4 3.410 0.007 Mg-O 4 3.41

Ru-Ru 7.1 1.5 3.552 0.006 Mg-Ru 8 3.55

Ru-Mg 0.8 0.9 3.552 0.006 Ilmenite Mg-O 3 1.95

Ru0.90Mg0.10O2 Ru-O 2.0 0.2 1.938 0.007 3 2.17

4.0 0.3 1.982 0.007 Mg-Ru 1 3.01

Ru-Ru 2.1 0.6 3.101 0.009 3 3.05

Ru-O 4.0 0.3 3.408 0.007 Mg-O

3,40

3 3.40

Ru-Ru 6.3 2.0 3.556 0.008

Mg-Ru 6 3.36

Ru-Mg 1.7 1.5 3.556 0.008 Ru0.90Mg0.10O2 Mg-O 6.1 0.5 2.07 0.02

Ru0.80Mg0.20O2 Ru-O 2.0 0.3 1.937 0.008 Mg-Ru 2.9 0.9 3.08 0.03

4.0 0.6 1.981 0.008 Mg-Mg 4.0 1.4 3.38 0.03

Ru-Ru 2.2 0.8 3.101 0.010 Mg-O 5.1 1.2 3.315 0.08

Ru-O 4.0 0.6 3.407 0.008 Mg-Ru 4.0 0.5 3.55 0.03

Ru-Ru 6.5 2.1 3.557 0.009 Ru0.80Mg0.20O2 Mg-O 6.5 1.2 2.00 0.05

Ru-Mg 1.5 1.5 3.667 0.009 Mg-Ru 2.2 0.7 2.88 0.06

Mg-Mg 1.8 0.5 3.07 0.04

Mg-O 4.0 1.0 3.52 0.07

Mg-Ru 1.5 0.8 3.37 0.04

-

M

g

M

g

-

M

g

M

g

-

M

g

M

m

m

m

m

2.0 0.8 3.56 0,04

The refinement of the Mg EXAFS functions is only possible for the materials containing

at least 10% Mg in the cationic positions. The quality of the XAS spectra for materials with

lower Mg content was not sufficient for the full refinement. The refinement of the Mg local

structure using a rutile-type structural model does not give convergent results as can be seen by

109

the absence of the characteristic cation-cation interactions corresponding to bonding distances of

ca. 3.1 and 3.5 Å. It ought to be stressed that resolved scattering due to the cation-cation

interactions with a bonding distance in the range of 3 to 4 Å is characteristic for most of the

binary and ternary oxide structural types and its absence suggests a rather complex and

disordered nature of the Mg local environment. The actual Mg-O bonding distances in the first

coordination shell increase with increasing Mg content from ca. 2.00 (x = 0.1) to 2.07 Å (x =

0.2). This trend reflects a gradual relaxation of the Mg environment from the bonding

arrangements native to rutile framework of RuO2. The trend in the Mg-O bonding distance

reflects the general incompatibility of Mg with the RuO2 environmental confinement which

brings substantial strain to the material and has to be stabilized by the small size of the Mg

cluster as well as the overall particle size. The Mg-O bond length observed in materials with low

overall Mg content is comparable with that of the Ru-O bond in the ruthenium dioxide. It is,

however, also approximately 0.1 Å shorter that the Mg-O bonding distance in the

thermodynamically stable cubic magnesium oxide (2.12 Å). The refined Mg-O bonding distance

trend seems to be in accordance with the observed shift in the Mg absorption edge toward higher

energies. It is, however, unrealistic to assign the observed behavior to the removal of a third

electron from Mg and we attribute the experimental behavior rather to the strain imposed by the

adjacent Ru-rich rutile-type matrix.

A full refinement of the Mg EXAFS functions requires the formulation of a convenient

structural model reflecting the overall chemical composition of the prepared materials. Since Ru

and Mg do not form stable double oxides one has to base an applicable local structure model on

the double oxides existing in the Mg-Ti-O ternary system[28]. Although magnesium and

titanium form two stable ternary oxides conforming to the spinel and ilmenite structural models,

110

the absence of the pronounced scattering between 3 and 4 Å seems to disagree with both possible

structural models for the Mg EXAFS function refinement (see Figure 4.5b).

A satisfactory fit of the experimental data can be achieved assuming that the local

environment of Mg bears features of both the ruthenia host (rutile structural model) and of the

stable ternary Mg-Ru-O oxide conforming to an ilmenite structural model. Typical result of the

full refinement for the Mg EXAFS functions is shown in Figure 4.5b; the parameters of the

refinement are included in Table 4.3. As shown by the parameters of the refinement summarized

in Table 4.3, the observed increase in the Mg-O bonding distance with increasing overall Mg

content from ca. 2.00 to ca. 2.07 Å can therefore be viewed as a relaxation from when the Mg

originally confined in the rutile-like oxygen coordination gradually relaxes to a local structure

similar to that in stable Mg oxide(s)[29]. The refined EXAFS functions do not reflect any change

in the Mg-O coordination suggested by the composition dependence of the unit cell volume. The

possibility of a change in the Mg-O coordination, specifically at low Mg contents, however,

cannot be ruled out since the EXAFS functions of Ru1-xMgxO2 with x < 0.10 are not available.

It needs to be stressed that the Mg-O bonding distance in all prepared materials remains

shorter than that in Mg-O or Mg-Ti-O and the Mg local environment remains rather strained. The

cationic arrangement also shows a gradual development between bonding distances and

coordination numbers characteristic for a rutile-like structure (Ru0.9Mg0.1O2) and those

approaching the values expected for ilmenite (Ru0.8Mg0.2O2). It needs to be stressed that although

the nearest cationic coordination in the latter material starts to resemble the ilmenite, there still

remain observable cation–cation interactions characteristic of a rutile-like structure. This

behavior is not surprising since both structural types are related as it was shown, e.g. in reference

[18].

111

4.3.3 Electrochemical and DEMS Measurements

The Mg doped materials are active catalysts for both oxygen and chlorine evolution. The results

of the electrochemical characterization are summarized in Figures 4.6-4.10. The incorporation of

Mg into the rutile structure apparently decreases the activity of the ruthenia-based catalysts for

the oxygen evolution process (see Figure 4.6). Such behavior contradicts that of the Ni or Co

containing analogs[10, 13]. The observed behavior seems to reflect the fact that the cation

perturbing the local structure of the catalyst is itself catalytically inactive (i.e. cannot enter redox

reactions and become a binding site of the reaction intermediates). The decrease in the oxygen

evolution activity shows, however, a non-monotonous dependence on the actual Mg content with

a maximum activity observed for the material with an overall Mg content of x = 0.10. This

behavior cannot be directly linked with the refined local structure and its explanation will most

likely need to employ advanced theoretical approaches.

Figure 4.6. Linear scan voltammograms of the oxygen evolution on MgxRu1-xO2electrodes

(0<x<0.2) recorded in 0.1 M HClO4 at a polarization rate of 5 mV s-1

. The curve assignment is

given in the Figure legend.

112

The overall activity of the Mg modified ruthenia and pure ruthenia toward parallel

oxygen and chlorine evolution can be seen in Figure 4.7. The measured faradaic current shows a

linear dependence on the chloride concentration in the electrolyte solution. The increase in the

overall activity is also accompanied with a significant change in the material’s selectivity. While

approximately 30% of the charge being passed at a concentration of 300 mM Cl− can be

attributed to oxygen evolution on the non-doped ruthenia (see Figure 4.8), in the case of the Mg

doped materials the yield of the oxygen drops below 20% for a chloride concentration of 50 mM

and the chlorine evolution becomes practically quantitative at higher chloride concentrations.

Figure 4.7. Chloride concentration dependence of the oxygen evolution (a) and chlorine

evolution (b) contributions to the overall current response of MgxRu1-xO2 electrodes to anodic

polarization in chloride containing acid media. The presented values correspond to potentiostatic

experiments at 1.25 V vs. Ag/AgCl.

113

Figure 4.8. Composition dependence of the Mg doped ruthenia selectivity towards chlorine

evolution. The data correspond to potentiostatic experiments at 1.25 V (Ag/AgCl). The actual

chloride concentrations are given in the Figure legend.

The selectivity of the Mg modified ruthenia toward chlorine evolution seems to be little

affected by the catalyst’s composition; at higher chloride concentrations the selectivity is also

independent of the electrode potentials. A potential dependence of the catalyst’s selectivity

remains, however, observable at low to medium chloride concentrations (10 to 50 mM) when all

catalysts seem to favor the chlorine evolution at lower overpotentials. The selectivity of non-

doped catalysts, however, changes toward oxygen evolution at more positive potentials (see

Figure 4.9). Detailed information on the nature of the selectivity of both Mg doped as well as

non-doped ruthenia can be obtained by analyzing the time courses of the oxygen and chlorine

evolution at constant potential in solution containing 10 mM of sodium chloride (see Figure

4.10). The DEMS signals corresponding to both anticipated reaction products start to rise

simultaneously after the application of the potential step and the chlorine evolution signal attains

114

a steady state value after 10 and 40 s. These steady state values reflect equalization of the rate of

chlorine production and its removal into the DEMS apparatus. The fact that the chlorine

evolution signal attains a steady state value before the oxygen evolution signal indicates that the

oxygen evolution process does not affect the rate of chlorine evolution. This type of behavior

contradicts the assumed strict competition of both anodic processes for the same active sites at

the catalyst’s surface. The chlorine and oxygen evolution processes apparently use different

surface structures. The obtained data also suggest that the chlorine evolution pathway is

kinetically less hindered than that used for oxygen evolution. It needs to be noted that the active

sites used in chlorine evolution can also catalyze the oxygen evolution process if chlorine

evolution becomes transport limited.

Figure 4.9. Potential dependence of the Mg doped ruthenia selectivity towards chlorine

evolution. The data were extracted from potentiostatic experiments in 0.1 M HClO4 containing

10 mM NaCl (a) and 50 mM NaCl (b).

115

Figure 4.10. Time course of DEMS-based signals of potentiostatically generated oxygen (blue)

and chlorine (red) for RuO2 (a), Ru0.95Mg0.05O2 (b), Ru0.90Mg0.10O2 (c), and Ru0.80Mg0.20O2 (d).

Signals were recorded in 0.1 M HClO4 containing 10 mM NaCl; the potential perturbation

corresponded to a step from 0.70 V to 1.25 V vs. Ag/AgCl.

The conditions of the electrochemical experiments are generally incompatible with the

sample composition in general. To establish a general validity of the electrochemical data all

experiments were run with fresh electrodes and the duration of the experiments was kept short

and never exceeded 1000 s. The surface composition after the electrochemistry as reflected in the

XPS spectra agrees well with the initial one (see Tables 4.1 and 4.2). Although this fact validates

the relevance of the presented electrochemical data it does not allow to make any predictions

related to the practical suitability of the prepared materials in industry-like applications where

the catalyst should exhibit stability on significantly longer time-scales (~108 s).

116

4.3.4 Discussion

Traditionally, the theoretical description of the oxygen evolution process on rutile-type

ruthenium dioxide designates the active sites as transition metal atoms in so-called cus positions.

The significant enhancement of the oxygen evolution activity achieved by the modification of

the rutile structure with Ni or Co is then attributed to the activation of the bridge sites for a

proton transfer by incorporation of the additional transition metal[30]. The incorporation of

alkaline earth metal cations, such as Mg, however, is unlikely to improve the catalytic activity of

the rutile structure even if present in the cus positions. The catalytic cycle for oxygen evolution

formally requires the active metal cations to be oxidizable if they should contribute to the overall

activity. This can be hardly achievable given the oxidation state of Mg in the prepared materials.

This fact is reflected in the drop in oxygen evolution activity which therefore results from a

decrease in the available active Ru cus sites via dilution effect.

The link between the selectivity of parallel oxygen and chlorine evolution with the

catalyst’s local structure has not been established. Although the selectivity of similar systems,

e.g. Co[13] or Ni[9], doped ruthenia has been experimentally assessed in parallel oxygen and

chlorine evolution, the presented results were interpreted mainly in terms of morphology rather

than local structure[24]. Although the theory suggests that superiority of the chlorine evolution

should be structure insensitive, the experimental studies on Ni, Co, and Zn[18] doped ruthenia

are rather equivocal. In this regard one may consider two types of structures attributable to

dopant incorporation. The first type represents Ni and Co doped ruthenia which generally form

surface structures characterized by isolated or stacked Ni and Co cations which reside in a

strained rutile-like environment[9, 13]. The stacking of the dopant atoms along the c axis shifts

the selectivity toward chlorine evolution[9]. Breaking the dopant arrangements along the c axis,

117

e.g. due to shear plane formation[10], shifts the selectivity toward oxygen evolution[9]. The Zn

doped materials, on the other hand, form local structures characterized by partially or fully

broken stacking of cus and bridge sites[18] and their unusual selectivity toward oxygen

evolution was previously attributed to the structural hindrance of forming peroxo bridges

between adjacent cus sites.

The Mg doped materials, from a structural point of view, represent a transition in the

local structure from where the dopant stacking in the bridge and cus sites is essentially

unperturbed (characteristic for pure RuO2 and for Ni and Co doped ruthenia) to a structure

containing dopant enriched ilmenite-like inclusions characteristic for Zn doped ruthenia. A

purely structural comparison, therefore, suggests the electrocatalytic behavior of the Mg doped

ruthenia to fall between that of the non-doped and Zn doped ruthenia. This trend seems to be

fully reflected in oxygen evolution data when both Zn2+

and Mg2+

suppress the activity in the

oxygen evolution. The adherence of the behavior of Mg doped ruthenia observed in chloride

containing solutions to the theory predictions is less pronounced. The Mg doped materials show

a preference for chlorine evolution which peaks for the material featuring about 5% Mg in

cationic positions (see Figure 4.8). Although the detailed local structure information for this

material is missing it may be envisaged that this material shows the closest similarity to the non-

doped ruthenia. The moderate selectivity shift toward oxygen evolution observed for the

materials with higher Mg content reflects the development of the catalysts’ local structure toward

ilmenite-type inclusions. Breaking the stacking of cationic positions decreases the possibility to

form peroxo bridges between adjacent cus sites which were proposed to be primary chlorine

evolution active sites[18]. The shift of the selectivity of Mg doped ruthenia toward oxygen

evolution, however, lags behind that of the Zn doped ruthenia for chlorine evolution. The precise

118

nature and functionality of the chlorine evolution active sites in Mg doped ruthenia cannot be,

however, described experimentally with sufficient precision and needs to be backed with

complementary DFT-based studies modeling both surface stability and activity of doped ruthenia

based catalysts.

4.4 Conclusions

Nanocrystalline Mg doped ruthenium dioxide catalysts were prepared by the spray-

freezing freeze-drying technique. Regardless of the average chemical composition, the prepared

materials are of nanocrystalline character with an average particle size of approximately 5 nm.

Despite the larger radius of the Mg cation compared with the tetravalent ruthenium cation, the

observed decrease in the unit cell volume suggests a nonhomogeneous distribution of Mg in the

structure with a local environment corresponding to that of a disordered ilmenite structure. Such

an arrangement is also supported by the analysis of the Mg EXAFS functions. The incorporated

Mg apparently exists in a significantly strained environment as shown by the shift in the Mg K

absorption edge energies as well as the unusually short Mg-O bonding distances which do not

exceed 2.07 Å. The prepared materials are active in both oxygen as well as in chlorine evolution

reaction. While the oxygen evolution activity is suppressed on Mg doped ruthenia, the chlorine

evolution is enhanced compared to the conventional ruthenia.

4.5 Acknowledgements

This work was supported by the European Commission within the Initial Training

Network ELCAT (Project No. 214936). The synchrotron measurement time was provided by the

KEK of Japan within the projects 2013R-35 and 2009R-29. Valery Petrykin of the J. Heyrovský

119

Institute of Physical Chemistry and Maki Okube of the Tokyo Institute of Technology are

gratefully acknowledged for their assistance with the synchrotron measurements and additional

thanks to Valery Petrykin for performing the EXAFS data analysis. The authors would also like

to thank Zdeněk Bastl of the J. Heyrovský Institute of Physical Chemistry for his assistance with

the XPS measurements and analysis.

120

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electrocatalytic properties of Ti/RuO2(x)+Co3O4(1-x) electrodes. J. Electroanal. Chem.,

2002. 532: p. 141-150.

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(Ti,Ru)O2 electrode coatings. Mater. Sci. Eng., B, 2013. 178: p. 1515-1522.




19. Gaudet, J., et al., Physicochemical Characterization of Mixed RuO2-SnO2 Solid

Solutions. Chem. Mater., 2005. 17: p. 1570-1579.

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reaction in proton exchange membrane water electrolysers. Int. J. Hydrogen Energy,

2011. 36: p. 14796-14804.

21. Xiong, K., et al., Sn and Sb co-doped RuTi oxides supported on TiO2 nanotubes anode

for selectivity toward electrocatalytic chlorine evolution. J. Appl. Electrochem., 2013. 43:

p. 847-854.

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Materials. J. Electrochem. Soc., 2008. 11: p. F27-F29.

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particle shape and surface composition. Electrochim . Acta, 2008. 53: p. 2656-2664.

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Publishing Corporation.



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Chapter 5

Reformate-Tolerant Pt-based Catalysts for the Electrochemical Hydrogen Pump

5.1 Introduction

Efficient and cost-effective methods for hydrogen purification are important for several

applications and the development of such methods is critical to the overall advancement of

hydrogen fuel-based technologies. Many industrial processes, such as steam reforming of

gasified coal and natural gas in conjunction with the water gas shift reaction, produce gas

mixtures of CO2, H2, water vapor, CO, and other trace impurities. Current state of the art

technologies for hydrogen purification and separation include chemical adsorption methods

based on alkaline or ammine scrubbing, pressure swing adsorption (PSA), membrane separation

purifiers (polymer, ceramic, or palladium based), and cryogenic separation. Pressure swing

adsorption is the most widely used method in the large scale separation and purification of

hydrogen[1-3]. Although this method is able to produce hydrogen at high purity (>99.9%) with

moderate to high recovery rates (70-95%), it can also require several wash columns to remove

CO and CO2 [3-6]. Similarly, cryogenic separation also has the advantage of providing hydrogen

at a high purity (90-99%) with an even higher recovery rate (95%), but this method requires a

high energy input to condense out impurities in the contaminated gas stream[5, 7]. Membrane

separation technologies are generally used for the final purification step and require a relatively

clean hydrogen stream (>98%) to produce highly purified hydrogen (99.9999%), The downside

to membrane separation techniques is that they generally require the use of elevated temperatures

(≥ 300 °C) to operate. Furthermore, the membrane separators often require high pressure inlet

gas feeds, requiring additional energy input, and can experience hydrogen embrittlement[3, 5].

124

An attractive alternative to the conventional hydrogen separation and purification

technologies is the electrochemical hydrogen pump (ECHP). The concept of hydrogen

purification using a PEM-based ECHP was first demonstrated by Sedlak et al.[8] in the early

1980's. Since then it has been shown by Gardner et al. that it was possible to recover 80% of

hydrogen with an energy efficiency of 80% from H2/CO2 mixtures in a single cell while pulsing

the anode potential[9]. More recently, Benziger et al. reported that a multistage electrochemical

hydrogen pump operating with programmed voltage pulsing was able to recover >98% of

hydrogen from H2/CO2 gas mixtures with energy efficiencies >92%[10]. In addition to being an

efficient means of obtaining highly purified H2, it has also been demonstrated that

electrochemical hydrogen pumps offer an efficient means of hydrogen compression [11-14] and

hydrogen recirculation in a fuel cell stack[15]. The revitalized interest in this technology within

the past few years has led to large advancements of the ECHP and it is likely that higher

recovery rates and energy efficiencies could be achieved with further development.

There have been several studies recently focusing on high temperature hydrogen pumps

using polybenzimidazole (PBI) membranes[16, 17]. Some investigations have also looked at

phosphoric acid [5], molten carbonate[18, 19], and solid-acid membrane[11] analogues.

However, there is a lack of literature regarding low temperature (< 100 °C) hydrogen pump

systems utilizing CO-tolerant electrocatalysts at the anode[9, 20]. Over the last several decades

there has been a surge of literature exploring the use of reformate-tolerant catalysts in proton-

exchange membrane fuel cells (PEMFCs), including both gas-fed and liquid-fed fuel cells such

as direct-methanol fuel cells[21-28]. The current work aims to extend the previous findings and

demonstrate the use of Pt-based alloys as efficient reformate-tolerant catalysts suitable for the

electrochemical hydrogen pump.

125

5.2 Experimental

5.2.1 Catalyst Preparation

The PtRu/C (41.4 wt.% Pt, 32.1 wt.% Ru) and Pt/C (46.2 wt.% Pt) catalysts used in this

work were obtained from Tanaka Corporation, Japan. The PtSn/C (20 wt.%) with an atomic ratio

of 3:1 (Pt:Sn) was obtained from E-TEK Corporation.

PtNi/C (40 wt. %) with an atomic ratio of 1:1 (Pt:Ni) was synthesized by aqueous

precipitation of Ni2+

onto Pt/C. The Pt/C precursor was synthesized by dissolving an appropriate

amount of H2Pt(OH)6 (DF Goldsmith Chemical and Metal, USA) in a 6% H2SO3 solution. The

solution was stirred at room temperature overnight and then diluted with Millipore H2O to bring

the total volume to 50 mL. Approximately 700 mg of Vulcan XC-72R (Cabot Corp., USA) was

added to the flask and ultrasonically dispersed for 30 minutes before adjusting the pH to 3.0 with

1.0 M NaOH solution. A stoichiometric amount of 30% H2O2 relative to H2Pt(OH)6 was then

added dropwise to the slurry and the pH was periodically adjusted to 3.0 with 1.0 M NaOH three

more times. Afterwards, the slurry was boiled for one hour and cooled to room temperature

before filtering and washing with 500 mL of Millipore H2O. The resulting powder was dried

overnight at 110 °C and then ultrasonically dispersed in 50 mL of Millipore H2O. An appropriate

amount of Ni(NO3)2•6H2O (Alfa Aesar, USA) was added to the slurry and stirred for 30 minutes.

The Ni was then precipitated onto the Pt/C support in the form of an oxy/oxy-hydroxide layer by

dropwise addition of 1.0 M NaOH until the pH was adjusted to 8.5. The slurry was boiled for 30

minutes then cooled to room temperature, filtered, and washed with 500 mL Millipore H2O. The

resulting powder was dried at 110 °C overnight and then annealed at 800 °C under 5% H2/Ar for

two hours. An acid wash treatment (stirring in 1.0 M HNO3 at 70 °C overnight) was used to

leach out any excess Ni.

126

Pt3Mo (40 wt. %) was synthesized by aqueous impregnation of Pt/C with MoO3. The

Pt/C precursor was synthesized according to the above procedure. A solution of H2MoO4 was

prepared and an appropriate amount was added to a slurry of Pt/C in Millipore H2O. The slurry

was stirred for 24 hours and then allowed to settle before filtering, washing 200 mL Millipore

H2O, and drying at 110 °C overnight. The resulting powder was annealed in 5% H2/Ar at 500 °C

for 6 hours.

Ternary alloy catalyst PtRuNi/C was synthesized in-house starting with the acid washed

PtNi/C precursor. Ruthenium was deposited by dispersing the PtNi/C in Millipore H2O and

dissolving an appropriate amount of RuCl3•xH2O. The ruthenium was then reduced by passing a

stream of hydrogen through the stirred solution for approximately 2 hours. Afterwards, the

catalyst powder was filtered, washed, and dried at 80 °C overnight. The dried powder was

annealed in 5% H2/Ar for 3 hours at 300 °C. The final metal loading was calculated to be 49

wt.% with an atomic ratio of 1:1:1 (Pt:Ru:Ni).

5.2.2 Physical Characterization

The crystallinity and particle size of the carbon-supported nanoparticles were

characterized using a Rigaku Ultima IV X-ray diffractometer with Cu Kα radiation (λ = 1.5418

Å) operating at 44 kV and 40 mA. Samples were scanned from 20° < 2θ < 90° and the recorded

patterns were matched against the PDXL database. The average sample composition of

synthesized binary and ternary alloy nanoparticles was confirmed with X-ray dispersive

spectroscopy using a Hitachi S-4800 scanning electron microscope (SEM) equipped with an

EDAX Sapphire Si(Li) detector. In addition, HRTEM images of the synthesized PtNi/C and

PtNiRu/C materials were obtained with JEOL JEM-2010F Field Emission Electron Microscope.

127

5.2.3 Electrochemical Cell Polarization Measurements

Membrane electrode assemblies (MEAs) were prepared using the corresponding anode

catalyst and a Pt/C cathode. A total metal loading of approximately 0.5 mg cm-2

was used for all

anode and cathode electrodes. To prepare the anode and cathode electrodes, catalyst ink

suspensions were prepared in a water-alcohol mixture with the required amount of ionomer

solution (Nafion D-521, 5 wt.%) and sprayed onto a Sigracet gas diffusion layer (GDL).

Electrodes were coated with an interfacial layer of 0.5 mgNafion cm-2

before hot pressing. Hot

pressing the electrodes with a Nafion 212 membrane was carried out at 500 psi and 135 °C for

four minutes. MEA testing was performed in a single cell with graphite serpentine flow fields.

All MEAs were humidified with N2 (100% RH) at a cell temperature of 85 °C for at least 90

minutes before activating with H2/H2 (25/30 psi backpressure on anode/cathode) for a minimum

of 90 minutes. All subsequent measurements were taken at a cell operating temperature of 85 °C

and 100% RH with a backpressure of 25 psi on the anode and 30 psi on the cathode. Hydrogen

pump polarization curves were recorded using an Autolab Potentiostat/Galvanostat Model

PGSTAT30 with Booster 20A module (Metrohm, USA) and the Nova 1.10 software package.

Cell activation was carried out by holding the cell potential at 25 mV for one hour and

then holding at 50 mV for an additional 30 minutes. Galvanostatic polarization was performed

by ramping the cell current to 7.0 A at 70 mA steps. After recording the cell polarization curve

under H2, the anode gas feed was switched to a reformate gas mixture (100 ppm CO, 15% CO2,

1% CH4, 45% H2, balance N2) for 30 minutes before repeating the cell polarization curve

measurement. The anode and cathode gas feeds were switched to nitrogen and hydrogen,

respectively, between polarization measurements to allow for half-cell polarization and to

remove any adsorbed CO. Finally, a hydrogen gas mixture containing 300 ppm CO was passed

128

over the anode for 30 minutes before collecting the final polarization measurements. A similar

set of cell measurements was made using a new set of MEAs to investigate the CO2 tolerance of

each catalyst. The cells were first polarized while flowing H2 over both the anode and the

cathode. Afterwards, the anode gas feed was switched to 50:50 (molar ratio) H2:CO2 and the cell

polarization was recorded again.

Impedance measurements were recorded between polarization measurements in order

determine the internal cell resistance. Impedance measurements were made potentiostatically at

the open circuit potential (OCP) from 20 kHz to 0.1 Hz with a 10 mV amplitude. All reported

potentials are quoted in the reversible hydrogen electrode (RHE) scale and have been corrected

for iR losses.


5.3.1 Physical Characterization

The powder XRD patterns for each electrocatalyst sample tested are shown in Figure 5.1.

The Pt/C (Tanaka, Japan) sample shows the anticipated diffraction pattern regions characteristic

of the fcc (111) and (220) reflections appearing at 39.51 and 46.05 degrees, respectively. The

expected shift in the primary diffracting regions due to the incorporation of dopant atoms is

observed in the commercially available Pt3Sn/C and Pt2Ru3/C samples. The peaks in the Pt3Sn/C

diffraction pattern shift to lower 2θ values as expected due to the larger atomic radius of Sn

relative to Pt. Similarly, the peaks in the Pt2Ru3/C sample were shown to shift to higher 2θ

values due to the smaller atomic radius of Ru. In addition to shifting to higher angles, some Ru

may exist in the hexagonal phase as evidenced by the broad overlap between the peaks at 40.48

and 44.87 degrees. This is in agreement with Tripkovic et al.[29]. Due to the similar size of the

129

Pt (1.39 Å) and Mo (1.39 Å) atomic radii, no significant shift was observed in the diffraction

peaks for this sample. A significant difference was observed, however, in the diffraction patterns

of the synthesized PtNi/C and PtNiRu/C materials. The broad shoulder towards higher angles

associated with each peak indicates the presence non-alloyed platinum in addition to the alloyed

PtNiRu or PtNi supported particles. In the case of PtNiRu/C, a shoulder at 38.48 degrees and a

small peak at 43.91 degrees indicates the presence of some Ru in the hexagonal phase, similar to

the commercial Pt2Ru3/C material.

Figure 5.1. Powder XRD patterns for electrocatalyst samples.

130

Table 5.1. Results of XRD and HRTEM physical characterization.

Electrocatalyst

Sample

Lattice

Type

Atomic

Ratio

XRD Analysis

TEM

Particle Size

/ Å

Lattice

Parameter

/ Å

d(111)

/ Å Particle

Size / Å

d-spacing

/ Å

Pt/C fcc - 17.3 3.9184 2.279

Pt2Ru3/C fcc 2:3 15.1 3.8679 2.226

Pt3Mo/C fcc 3:1 40.3 3.9209 2.265

PtNi/C fcc 1:1 44.7 3.8250 2.208

56.2 ± 16.9 2.25 ± 0.03

Pt3Sn/C fcc 3:1 15.6 3.9916 2.308

PtRuNi/C fcc 1:1:1 40.7 3.8269 2.204

59.2 ± 15.8 2.20 ± 0.11

The coherent domain size of each sample as calculated by the Scherrer equation can be

seen in Table 5.1. Representative HRTEM images of the PtNi/C and PtNiRu/C electrocatalysts

synthesized in-house can be seen in Figure 5.2. From Figure 5.2 it can be seen that the

nanoparticles are well dispersed on the carbon support. The average particle sizes of these

samples was also calculated by measuring over 300 randomly selected particles and are shown in

Table 5.1. The average size of the supported nanoparticles calculated from TEM analysis is

between 5.6 and 6.0 nm, which is only slightly larger than those calculated from the Scherrer

equation.

131

Figure 5.2. HRTEM images of the synthesized PtNi/C (a) and PtNiRu/C materials.

5.3.2 Cell Polarization Measurements:

Shown in Figure 5.3 are the half-cell anode polarization curves for hydrogen oxidation.

All polarization curves are normalized to the Pt metal-loading, since in most cases Pt provides

the only catalytically active site for hydrogen oxidation. In addition, all polarization curves were

corrected for iR-losses using the high frequency resistance measurement determined via

electrochemical impedance spectroscopy (EIS). This configuration allows for a direct

comparison of the different anode catalyst activities toward hydrogen oxidation for an

electrochemical cell operating in the hydrogen pump mode. As expected, pure Pt/C shows one of

the highest activities for the hydrogen oxidation reaction (HOR). It should be noted that the

PtRu/C and PtNiRu/C show a higher activity due to the fact that both samples contained some

Ru in the hcp phase (as evidenced by XRD in Figure 5.1), which is known to show sufficient

HOR activity between 0 to 300 mV vs. RHE (below the point of OHads nucleation on Ru)[30].

The other materials investigated showed a lower HOR activity in the presence of pure hydrogen.

132

This is due to the fact that the other metals present in the Pt-based alloy are not HOR active and

thus dilute the sites available for hydrogen oxidation. Although the polarization curves are

normalized to the Pt-loading, one cannot assume a completely homogeneous distribution of each

metal within the nanoparticles or at the particle surface, therefore the number of surface sites

available for HOR may be less than anticipated. In all cases it can be seen that high specific

currents are obtained at very low overpotentials due to the extremely facile kinetics of the HOR

on Pt and Pt-based catalysts[31].

Figure 5.3. Anode half-cell polarization curves of an electrochemical hydrogen pump with the

anode and cathode being fed pure H2. Polarization measurements were collected at a cell

temperature of 85°C, 100% relative humidity, and 25/30psi backpressure on the anode/cathode,

respectively.

Other PEM-based ECHPs have been reported previously. Shown in Table 5.2 are the

various operating conditions and current densities for ECHPs operating with pure H2 supplied to

the anode. Comparatively, the current density reported for the Pt anode in Figure 5.3 is

133

considerably higher than those reported in the literature. However, it is important to note that in

all cases the operating temperature used in this study was higher than the others reported.

Additionally, the use of a thinner membrane (50.8 μm for Nafion 212 vs. 127 μm for Nafion

115) also contributes significantly considering that the reported current densities were not iR-

corrected.

Table 5.2. Performance and operating conditions of various ECHPs reported in the literature

Reference Operating

Temp / °C

Anode Loading /

mgPt cmgeo-2

Membrane

J @ 150 mV /

A cmgeo-2

Figure 5.3 (Pt/C) 85 0.5 Nafion 212 1.5

Gardner et al.[9] 20 1.0

(1.5 mgPtRu cm2)

Nafion 115 0.31

Benziger et al.[10] 50 0.4 Nafion 115 0.65

Fateev et al.[13] 75 0.8 Nafion 117 0.51

Barbir et al.[15] 30 0.8 3M PFSA 1000 EW (28 μm) 0.60

Shown in Figure 5.4 are the anode half-cell polarization curves for hydrogen oxidation

with an anode gas feed of either H2 containing 300 ppm CO or a reformate gas mixture

containing 100 ppm CO, 15% CO2, 1% CH4, 45% H2, and N2 to balance. PtRu/C and PtNiRu/C

again show the highest activity, and thus the highest tolerance to CO/CO2 poisoning. The high

activity of these two materials can be attributed to the presence of Ru. It has been reported by

Tripković et al. that Pt2Ru3/C is generally a poor catalyst for methanol oxidation, but shows an

extremely high CO-tolerance[32]. They attributed the excellent CO tolerance not to the twinned

structure of Pt2Ru3 but rather to the presence of a Ru-rich (hcp) phase in the catalyst, such as that

observed in the XRD profiles in Figure 5.1. It is widely accepted that the mechanism for CO-

removal is dependent upon water activation at Ru sites to form OHads, otherwise known as the

bifunctional mechanism[33, 34]. PtMo/C also shows a relatively high tolerance to CO and the

134

reformate mixture. It has been shown in previous studies that PtMo/C has a comparable CO-

tolerance to PtRu/C and in some cases even shows a higher tolerance[21]. The mechanism for

PtMo/C, although similar to that of PtRu/C, is more dependent on the turnover of the Mo(IV/VI)

redox couple. The lower activity of PtNiRu/C compared to that of PtRu/C is likely due to

presence of Ni, which appears to dilute the surface Pt sites required for hydrogen adsorption in

addition to diluting the Ru surface sites needed for OH adsorption to remove COads. This dilution

effect is also evidenced by the lower activity of PtNi/C vs. Pt/C, in which case the surface Ni

atoms are inactive for hydrogen oxidation and offer no means of promoting COads removal via

the bifunctional mechanism. It has been reported by Mohsen et al.[35] that the presence of Ni

lowers the d-band center of Pt and aids in CO2 tolerance, though no correlation was found

between the CO tolerance and Pt d-band center and PtNi/C. As was also reported by Mohsen et

al.[35], Pt-based anode electrocatalysts that rely on the bifunctional mechanism for CO removal,

such as PtRu/C, PtMo/C and PtSn/C, typically show a high CO tolerance but show a low

tolerance to CO2. In contrast, Pt-based catalysts that drastically improve CO2 tolerance by

altering the electronic properties of Pt, such as PtNi/C or PtCo/C, generally show a very poor

tolerance to CO poisoning. The behavior observed in Figure 5.4 seems to indicate that the anode

electrocatalyst activity is heavily dependent upon the presence of CO and that the presence of

CO2 or CH4 in the reformate mixture is not the dominant poisoning species.

135


cathode being fed hydrogen and the anode being fed by hydrogen containing 300ppm CO (filled

circles) or a reformate gas mixture containing 100 ppm CO, 15% CO2, 1% CH4, 45% H2, and N2

balance (empty circles). Polarization measurements were collected at a cell temperature of 85°C,

100% relative humidity, and 25/30psi backpressure on the anode/cathode, respectively.

It is well known that CO2 can be converted to CO via the reverse water gas shift reaction

by either of the following two pathways[36-40]:

CO2 + H2 ⇌ CO +H2O (5.1)

CO2 + 2M-Hads ⇌ 2M-COads + H2O + M (5.2)

Such a conversion occurs at the electrode interface when CO2 adsorption occurs adjacent to H2

adsorption. It has been shown that CO2 can poison Pt-based catalysts that otherwise show a high

tolerance to CO poisoning beyond the levels accounted for by simple dilution[37, 38].

136

Furthermore, the trends observed for CO tolerance in Pt-based alloys do not necessarily apply

when considering CO2 tolerance[41-43].

The hydrogen oxidation activity of four different Pt-based catalysts in the presence of an

H2-CO2 (50:50 molar ratio) gas mixture at the anode can be seen in Figure 5.5. It can be seen

from Figure 5.5 that the extent of catalyst deactivation induced by CO2 poisoning is less than that

observed for CO (See Figure 5.4). Although the extent of poisoning is less in the case of CO2,

the same trend observed for CO tolerance is observed again here, with Pt2Ru3/C showing the

highest HOR activity and thus the highest CO2 tolerance followed by PtRuNi/C, Pt/C, and

PtNi/C. This is in contrast to the results reported by Mohsen et al.[35], where PtNi/C was shown

to have the highest tolerance to CO2 and PtRu/C was shown to have one of the lowest tolerances.

It is important to note that Strasser et al.[44] has shown that unsupported PtxNiy nanoparticles

undergo selective dissolution when exposed to corrosive anodic conditions. Although it was

concluded that the resulting Pt-enriched core-shell structures show higher catalytic activity for

the oxygen reduction reaction (ORR), this may not be the case for HOR. It is well known that the

surface structure of PtxNiy is heavily dependent on the Ni content, the material preparation

method, and the electrode potential[44-46]. It has been shown by Hoffmannová et al.[46] that

hydrogen adsorption on PtNi triggers Ni segregation to the surface while on the other hand the

Ni is confined primarily to the subsurface when the electrode is polarized to the double-layer

region. The discrepancies observed in the CO2 tolerance between the PtNi/C sample reported

herein and that reported by Mohsen et al.[35] are likely due to differences in the surface

composition of the catalysts, which may be due to synthetic route chosen or the electrode history

prior to polarization measurements. As mentioned previously, Ni segregation occurs when the

electrode is polarized to the hydrogen adsorption region, which would decrease HOR Pt-based

137

specific activity (hence the lower activity than Pt) as Ni is significantly less catalytically active

for the HOR than Pt. It is likely that the nonhomogeneous distribution of Ni under these

circumstances would also mitigate any promotional effects on the Pt d-band center that are

believed to increase the CO2 tolerance[35], as such effects (in the case of ORR) are known to

occur when the Ni is located subsurface and a Pt skin is formed on the surface[46]. It should be

noted that two other PtxNiy/C samples were also prepared in-house using the same method

reported above and were subsequently tested in the ECHP (not shown). The additional samples

displayed nearly identical HOR activities and CO/CO2 tolerances, indicating that the poor CO2

tolerance is not batch-specific.

The improved CO2 tolerance of Pt2Ru3/C, similar to the improved CO and reformate

tolerance observed in Figure 5.4, is attributed to the activation of water via the bifunctional

mechanism. The presence of Ru lowers the potential at which water is activated, decreasing the

onset of surface OH formation, which subsequently removes adsorbed CO formed by the reverse

water gas shift reaction on adjacent Pt sites. This is also the case for the PtNiRu/C material,

although again the Ni dilutes the key catalytic sites on the particle surface reducing the overall

activity and CO2 tolerance. Ideally, the presence of Ni would lower the Pt d-band center aiding in

the removal of CO2 while the presence of surface Ru sites adjacent to Pt would aid in the

removal of CO via the bifunctional mechanism. It possible that with more development an

improved ternary alloy (PtNiRu/C) could be formed that offers superior tolerance to both CO and

CO2.

138


cathode being fed pure hydrogen and the anode being fed by hydrogen containing 50% molar

CO2. Polarization measurements were collected at a cell temperature of 85°C, 100% relative

humidity, and 25/30psi backpressure on the anode/cathode, respectively.

Although a direct comparison cannot be made between the PEM-based ECHPs reported

in the literature for CO/CO2 tolerance due to the large difference in gas mixture composition,

there have been reports of PBI membrane-based ECHPs tested in the presence of reformate gas

mixtures similar to those employed here. For instance, Thomassen et al.[17] reported an ECHP

operating with a PBI membrane at 160 °C while the anode was exposed to an unhumidified gas

mixture consisting of 44% H2, 35% N2, 21% CO2, 100 ppm CO. A current density of

approximately 1.25 A cm-2

was obtained at an anode overpotential of 180 mV, compared to an

overpotential of 240 mV (360 mV uncorrected) at the same current density for the cell operating

with a Pt2Ru3/C anode reported in Figure 5.4. Additionally, Perry et al.[16] reported a PBI

membrane-based ECHP with a 1.0 mgPt cm-2

loading at the anode operating at 160 °C that

139

attained a current density of approximately 2.0 A cm-2

at 150 mV under a humidified pure

hydrogen feed or 0.95 A cm-2

at 150 mV under unhumidified conditions. This is comparable to

the performance of the Pt2Ru3/C catalyst presented in Figure 5.3, which reached a current density

of 1.5 A cm-2

at 150 mV (uncorrected). In either case, the higher current densities attained by the

PBI-based ECHPs indicate that the elevated operating temperature (160 °C in PBI vs. 80 °C in

PEM) may offer reformate tolerances that are beyond what is attainable with reformate-tolerant

catalysts in low temperature ECHPs. This is largely due to the fact that CO adsorption on Pt is

thermodynamically unfavorable at elevated temperatures as evidenced by the rather negative

standard entropy value[47, 48].

5.4 Conclusions

Of the Pt-based electrocatalysts reported herein, Pt2Ru3/C has proven to be a promising

CO/CO2 tolerant anode catalyst for a PEM-based cell operating in the ECHP mode. This agrees

well with previous findings in the literature that investigated the use of PtRu/C as a reformate

tolerant catalyst in cells operating in the PEMFC mode. The high tolerance of Pt2Ru3/C is

attributed to the activation of water on Ru sites at much lower overpotentials, which then permits

the removal of CO adsorbed on adjacent Pt sites, i.e. the bifunctional effect. The presence of Ni

in the PtNi/C and PtNiRu/C samples was not shown to enhance the CO2 tolerance as was

reported previously. However, it is possible that with more development and fine tuning of the

Ni distribution that ternary alloys such as PtNiRu/C could offer superior reformate tolerances

that exceed Pt2Ru3/C.

140

5.5 Acknowledgements

The authors gratefully acknowledge the financial support from the Department of Energy,

under the auspices of a Small Business Innovative Research grant lead by Proton Onsite. The

authors would also like to thank Shraboni Ghoshal of Northeastern University for her assistance

with the hydrogen pump cell measurements.

141

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146

Chapter 6

Dissertation Summary, Conclusions, and Future Directions

6.1 Summary

Several topics have been introduced and discussed throughout the course of this

dissertation. Although each topic is somewhat unique, the underlying theme of electrocatalytic

selectivity is present. Chapter 2 considered the rational design of catalysts showing high MOR

activity while inhibiting quaternary ammonium poisoning, whereas Chapters 3 and 4 examined

the local structure effects of RuO2 on the selectivity of oxygen and chlorine electrocatalysis, and

finally Chapter 5 considered optimizing the reformate tolerance of Pt catalysts through alloying

with transition metals. All of the topics introduced address various issues encountered in

electrocatalysis, though the primary objective was to understand and develop ways of altering the

selectivity of the electrocatalyst to promote (or inhibit) specific electrochemical processes.

6.2 Chapter Synopses and Future Directions

6.2.1 Chapter 2 - Analysis of Double Layer and Adsorption Effects at the Alkaline Polymer

Electrolyte-Electrode Interface and the Development of a Quaternary Ammonium

Poisoning Resistant Electrocatalyst for Methanol Oxidation

The aim of Chapter 2 was to address the issues that arise when transitioning from an acidic,

PEMFC environment to an alkaline AEMFC. The development of stable and highly conductive

membranes is still a challenge, although the development of electrocatalysts that operate well

with the current state-of-the-art anion-exchange polymers is imperative[1, 2]. Here the issue of

quaternary ammonium poisoning that arises with the use of AEMs was addressed while

147

considering the nature of the catalyst itself. The specific adsorption of quaternary ammonium

species on Pt-based anode catalysts is a serious issue that results in significant performance

losses[3]. This is largely due to a combination of the favorable electrostatic interactions of poorly

solvated quaternary ammonium species with the anode surface and the potential of zero charge

of Pt. Here we introduced the idea of tailoring a highly oxophilic transition metal support to help

mitigate this issue. It was shown that the deposition of oxophilic metals, i.e. Ni and Pb, to the

underlying carbon support can drastically improve the MOR activity while inhibiting the specific

adsorption of quaternary ammonium species on the supported Pt nanoparticles.

The high MOR activity and resistance to quaternary ammonium poisoning of the

Pt/NiPb/C was demonstrated in Chapter 2. A more in-depth analysis of the structure, however, is

so far lacking. In order to truly elucidate the nature of the electrocatalytic enhancements

observed it may be necessary to employ a combined EXAFS and DFT approach. The detailed

structural information obtained from EXAFS would offer valuable insight regarding the

distribution of Ni and Pb in the highly oxophilic support. Currently it is known that the Pb exists

in a hydro-cerrusite phase that is separate from the Pt fcc phase, as evidenced by XRD analysis,

but the location of Ni is uncertain and can only be speculated. Furthermore, the interaction and

distribution of Pt on the NiPb support is still not completely understood. DFT analysis would

offer valuable information regarding the synergystic effects between Ni and Pb in addition to

information regarding the binding energies of quaternary ammonium and MeOH species with the

supported Pt nanoparticles and NiPb support.

148

6.2.2 Chapter 3 - Oxygen Reduction on Nanocrystalline Ruthenia - Local Structure Effects

Chapters 3 and 4 both examine the impact of doping transition metals into the rutile

ruthenia on the selectivity towards electrocatalytic reactions. In the case of Chapter 3, the effect

of doping varying amounts of Ni, Co, or Zn into RuO2 was shown to induce changes on the

selectivity of the ORR. Through the interpretation of electrochemical data and DFT calculations,

it was determined that the incorporation of Ni or Co into the cus active site positions results in a

shift of the ORR selectivity towards the 2-electron reduction pathway to H2O2. By lowering the

binding energy of the *OOH intermediate species, the generation of H2O2 becomes more

thermodynamically favorable at lower overpotentials. In the case of Zn doped ruthenia, the

disruption of the surface to an ilmenite-type facet disrupts the ordered stacking of bridge and cus

sites such that the 4-electron reduction pathway to H2O is favored. Considering the fact that Zn is

itself redox inactive, the Zn-doped material most closely resembles the non-doped RuO2 since

the catalytic activity is confined to Ru active sites in the cus position. The effect of the dopant

metals in the bridge positions, however, was shown to have a negligible effect. This is in contrast

to what is observed for the anodic process of oxygen evolution, where the presence of Ni or Co

in the bridge position was found to have a profound effect[4].

Although the doped ruthenia examined in Chapter 4 are by no means ideal ORR

catalysts, the results of this study can still be extended to other catalytic systems. The ORR is a

widely studied reaction of paramount importance for many electrochemical systems. The

opening and closing of catalytic pathways through the optimization of the local structure is

something to be considered in the rational design of any catalyst. Here it was shown that

introducing disruptions in the ordered stacking of active sites is enough to guide the direction of

the overall reduction process.

149

6.2.3 Chapter 4 - Selective Chlorine Evolution Catalysts Based on Mg-Doped Nanoparticulate

Ruthenium Dioxide

Chapter 4 again dealt with the doped ruthenia. In this case the effects of Mg-doping on

the selectivity of ruthenia towards the oxygen and chlorine evolution reactions was examined. It

was shown through X-ray diffraction that the synthesized materials are single phase and conform

to a tetragonal oxide of the rutile structural type. The use of DEMS and EXAFS was able to

provide valuable information regarding the selectivity with respect to the structural changes

induced with varying Mg-content. The refinement of the EXAFS data was able to confirm that

the Mg is not homogeneously distributed throughout the materials and that Mg-rich clusters are

formed. Furthermore, the Mg resides in a rather strained rutile-type structure at low Mg

concentrations while shifting to an ilmenite-type inclusion at higher Mg concentrations (Mg >

10%). All Mg modified materials are active in the oxygen evolution and chlorine evolution

reactions. Although the Mg containing catalysts show lower overall activities compared with the

non-doped ruthenia, they feature enhanced selectivity toward the chlorine evolution process,

which is attributed primarily to the opening of a reaction pathway for chlorine evolution

associated with presence of Mg modified active sites.

The results of this study offer insight regarding the fundamental nature of the Mg-doped

ruthenia in relation to the local structure. Similar to the doping effects discussed in Chapter 3, the

results here indicate that creating changes in the local structure can modify the selectivity of the

catalyst. Chapter 3 focused more on steering the selectivity of a single reaction towards the

production of the intermediate species. Here it is shown that the competitive nature of two

separate reactions can be controlled by altering the structure of the active sites involved.

150

6.2.4 Chapter 5 - Reformate Tolerant Pt-Based Catalysts for the Electrochemical Hydrogen

Pump

The reformate tolerance of Pt and various Pt-based alloys was examined in Chapter 5. In

this chapter special attention was paid to the dopant metal introduced. Prior studies have

indicated that the introduction of transition metals such as Ni, Co, and Fe improve the CO2

tolerance of Pt. Similarly, prior studies have concluded that the CO tolerance of Pt can be greatly

enhanced through the introduction of oxophilic Ru sites (via bifunctional mechanism) or Mo

sites (via turnover of the Mo4+/6+

redox couple). On this basis, the original goal was to develop a

Pt-based ternary metal catalyst that would take advantage of the bifunctional effect as well as the

electronic effects in order to offer an unprecedented tolerance to reformate gas mixtures

containing both CO and CO2. This concept has been demonstrated previously, though in a H2/O2

PEMFC environment as opposed to a H2/H2 ECHP setting[5]. Throughout this project various

catalysts were synthesized and tested, including PtMoNi/C, PtMoCo/C, PtRuNi/C, PtMoRu/C,

and PtRuCo/C. Of all these attempts, the PtRuNi/C showed the highest hydrogen oxidation

activity in the presence of CO/CO2 and the smallest potential losses. Therefore this was the only

synthesized ternary material discussed in detail in Chapter 5. However, even the reformate

tolerance of PtRuNi/C fell short of that measured for the commercial Tanaka PtRu/C sample.

The structure of the Pt-based alloy is believed to have a significant impact on the

reformate tolerance. Although previous studies have indicated that PtNi/C shows a high

tolerance to CO2, the results of the hydrogen pump cell testing conducted in Chapter 5 show that

PtNi/C has of the lowest CO2 tolerances. A review of the literature indicates that the PtNi

structure is highly dependent upon the preparation of the material, Ni content, the electrode

151

history, and the operating potential[6-8]. Of the catalysts examined, the commercial PtRu/C was

shown to have the highest CO tolerance and the highest tolerance to CO2. Although detailed

structural information is not available for the prepared materials, it can be speculated that the

small particle size and high dispersion of Ru on Pt is responsible for the high reformate tolerance

observed.

It is possible that with more development a Pt-based ternary catalyst could show superior

CO/CO2 tolerance with respect to the PtRu/C sample tested. Although many materials were

synthesized and developed, an in-depth analysis of the catalyst structure was never conducted.

With the aid of XAS it is possible to obtain detailed structural information that could be

correlated with the reformate tolerance. Optimization of the Ni and Ru content and distribution

are necessary to gain the highest catalytic enhancement.

6.3 Concluding Remarks

The development of electrochemical technologies targeting the production, purification,

and consumption of hydrogen (electrolyzers, electrochemical hydrogen pumps, and fuel cells)

are crucial to future of sustainable energy. Electrochemical hydrogen purification is still in its

infancy and has not yet received considerable research interest. With the development of low-

loading, high surface area, reformate tolerant catalysts it is possible that this technology may be

able to compete with the current state-of-the-art hydrogen purification methods. Similarly, AEM-

based direct methanol fuel cells also require a significant amount of research interest and

improvement before commercialization can be considered. We have demonstrated that

quaternary ammonium resistant catalysts can be developed, however, the in situ cell performance

has yet to be demonstrated. In addition, improvements to the ionic conductivity and stability of

152

AEMs still remains a major challenge. Although great strides have been made over the years,

there is still much work to be done before the widespread adoption of electrochemical hydrogen

pumps and AEM-based fuel cells can occur. From an electrocatalysis standpoint, it will be

interesting to see how these technologies develop and progress over the years to come.

153

6.4 References





1735.

3. Unlu, M., et al., Analysis of Double Layer and Adsorption Effects at the Alkaline Polymer

Electrolyte-Electrode Interface. J. Eletrochem. Soc., 2011. 158: p. B1423-B1431.



5. Ehteshami, S.M.M., et al., The role of electronic properties of Pt and Pt alloys for

enhanced reformate electro-oxidation in polymer electrolyte membrane fuel cells.


6. Tuaev, X., et al., In Situ Study of Atomic Structure Transformations of Pt-Ni Nanoparticle

Catalysts during Electrochemical Potential Cycling. ACS Nano, 2013. 7: p. 5666-5674.

7. Stamenkovic, V., et al., Surface segregation effects in electrocatalysis: kinetics of oxygen

reduction reaction on polycrystalline Pt3Ni alloy surfaces. J. Electroanal. Chem., 2003.

554: p. 191-199.

8. Hoffmannov , H., et al., Surface Stability of Pt3Ni Nanoparticulate Alloy

Electrocatalysts in Hydrogen Adsorption. Langmuir, 2013. 29: p. 9046-9050.

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