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The modification of catalysts and their supports for use in various fuel cells
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The modification of catalysts and their supports for use in various fuel cells
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i
THE MODIFICATION OF CATALYSTS AND THEIR SUPPORTS FOR USE IN VARIOUS
FUEL CELLS
by
Dean Glass
A Dissertation Presented to the
FACULTY OF THE GRADUATE SCHOOL
UNIVERSITY OF SOUTHERN CALIFORNIA
In Partial Fulfillment of the
Requirements of the Degree
DOCTOR OF PHILOSOPHY
(CHEMISTRY)
May 2019
Copyright 2019 Dean Glass
ii
“... the journey is more important than the end or the start.”
iii
DEDICATION
To all that seek knowledge and truth.
iv
ACKNOWLEDGEMENTS
Looking back, there are so many people that have worked to help make my graduate
experience and research possible. First of all I would like to thank my advisor Prof. Prakash, who
took me in the first semester with no questions asked and took me out of my comfort zone in
electrochemistry. His hands-off style in guidance allowed me to work at my own pace and forced
me to figure things out for myself, which has been crucial to my growth and development as a
chemist and researcher. I would also like to thank Prof. Narayan for his knowledge and expertise
in electrochemistry and for inviting me to his group meetings to further broaden my knowledge of
electrochemistry and applications. I would also like to acknowledge the late Prof. Olah for his
vision and profound wisdom and knowledge, both in his stories and lectures in group meeting, and
for opening up the field of methanol fuel cells, which was what my first few projects focused on.
He will be greatly missed by many.
I have been fortunate to be part of a great and talented research group that has made my
graduate studies enjoyable. I would first like to acknowledge the people in my immediate group
in the fuel cell lab. First, I would like to thank Dr. Marc Iulliucci for showing me the ropes and
training me on some of the instruments and showing me the techniques to get me started when I
knew very little about electrochemistry. His companionship for 4 years made my time in the lab
enjoyable. I would also like to thank Mr. Vicente Galvan for all of his help in running the tests
and cooperation these past two years as my future protégé in the fuel cell lab. I also wish to thank
Dr. John-Paul Jones for his technical expertise in the lab including troubleshooting and helping
refine various techniques and overall lab practices. Other people from the fuel cell lab I wish to
acknowledge are Dr. Bo Yang for his insights in chemistry and research, Mr. Eugene Kong, one
of our collaborators from chemical engineering, who always made me think about the
v
fundamentals of electrochemistry with his questions, Mr. Adam Ung and Dr. Sergey Muzkin of
the Hogen-Esch group for their collaboration, and Dr. Akihisa Saitoh for his friendship and
kindness. I would also like to thank the rest of the members of the Prakash group for their help
and kindness during my time here. I would first like to thank Dr. Alain Goeppert and Dr. Micklos
Czaun for their help with various instrumentation and techniques in the lab, and finally others
including Kavita who came in the same year as me, Patrice, Ralph, Fang, Shao, Attila, Arjun,
Hema, Socrates, Laxman, Sankar, Jothee, Huong, Alex, Archith, Sahar, Vinayak, Colby, and
everyone else for providing a warm and friendly work environment.
Along with the Prakash group, I would like to thank the members of the Narayan group
who have opened up their lab to me to feel like a “third” home. I would like to thank Dr. Lena
Hoober-Burkhardt for all of the talks and advice back and forth about research, Dr. Souridip
Malkhandi, Dr. Phong Trinh, Dr. Derek Moy, Ms. Dan Fang, and Mrs. Buddhine Jayathilake for
their help with some of the electrochemical instrumentation and discussions about research. I
would also like to thank Debanjin, Advaith, and others from the group.
Furthermore, I would also like to thank the members of the Loker Hydrocarbon Research
Institute for their assistance to help fund the research and deal with the day-to-day aspects of
running an institute. A special thanks to Dr. Robert Anszifeld, the glue that holds Loker together.
His hard work with the business aspects and everything else I can’t begin to think about is truly
appreciated. Also, to Jessie, whose talks about movies and TV shows has helped and made the
graduate school experience fun and enjoyable as well as Gloria, who works hard to make Loker
shine both literally and figuratively. I would also like to extend a special thanks to Michele, the
chemistry department guru for helping me with all of my questions and reminding me to schedule
vi
for classes, turn in forms, and graduate. I would also like to thank Prof. Jessica Parr for her
mentorship with teaching and lecturing the General Chemistry II class.
I would also like to thank the people at CEMMA for helping me with the various
characterization instrumentation. I would like to thank Mr. Andrew Clough for training me
(continuously) on the XPS, Dr. Leonard Velasco for his training on the SEM, and Mr. Lang Shen
for his assistance in running the TEM as well as John Curulli for helping coordinate everything at
CEMMA.
I would also like to thank the people at my undergraduate institution, Baldwin-Wallace
(University) for help preparing me, first for industry, then for graduate school. I would especially
like to thank my undergraduate advisor, Prof. James McCargar for inspiring me and pushing me
to do better, as well as for getting me to apply to graduate schools. He truly knew me and my
skills better than I knew myself. I would also like to thank my undergraduate research advisor,
the late Prof. Joseph Gorse, for taking me on to do research my final year of undergrad and for
helping me become independent in doing research. I would also like to thank the other professors
and staff there, who initially helped broaden my horizons both in science as well as the world.
The last group of people I wish to thank is my family. I would like to thank my husband,
Ryan, for all of his support through my graduate school experience. I have been truly fortunate to
come home to a warm house every night with his support and without any drama. I would finally
like to thank my parents Doug and Darlene, who have been there from the very beginning. Their
support and encouragement in putting my education above much else has really gone a long way
to me getting to the various stages of my education and to where I am today.
To the numerous other people that have encouraged and inspired me to continue along this
path that I cannot remember, thank you as well.
vii
TABLE OF CONTENTS
Dedication.......................................................................................................................................iii
Acknowledgements.........................................................................................................................iv
List of Figures.................................................................................................................................ix
List of Schemes...........................................................................................................................xviii
List of Tables.................................................................................................................................xix
Abstract.........................................................................................................................................xx
Chapter 1 Introduction............................................................................................................1
1.1 The Current State of Energy.............................................................................1
1.2 History of Fuel Cells.........................................................................................3
1.3 Fuel Cell Basics.................................................................................................4
1.4 Types of Fuel Cells...........................................................................................7
1.4.1 Proton Exchange Membrane Fuel Cells (PEMFCs)………..……......7
1.4.2 Alkaline Fuel Cells (AFCs)………………………………………...9
1.4.3 High Temperature Fuel Cells……………………………………...10
1.5 Fuel Cell Utilization………………………………………………………....11
1.6 The Oxygen Reduction Reaction…………………………………………....14
1.6.1 ORR in Acidic Media…………........……………………………....15
1.6.2 ORR in Basic Media…………………………………………........16
1.7 Overview of The Thesis………………....………………….....………….....17
1.8 References……………………………………………………………….......18
Chapter 2 CFx as a Catalyst Support for Oxygen Reduction in Acidic and Alkaline
Media.....................................................................................................................24
2.1 High Surface Area Carbon-based Supports and Catalyst Preparation……....24
2.2 The Effect of Fluorination of Carbon Catalyst Support on ORR
Kinetics in Acidic Media…………………………………………….…………...27
2.2.1 Experimental Methods………………………….……………….....27
2.2.2 Results……………………………………………………………..30
2.3 Effect of Microwave-Assisted Polyol Reduction of Fluorinated Carbon
Supported Platinum Catalysts for ORR in Acidic Media…...............52
2.3.1 Experimental Methods………………………………………….....52
2.3.2 Results…………………………………………………………......53
2.4 The Effect of Fluorinated Carbon Supports for Various Manganese
Dioxide Phases on ORR Kinetics in Basic Media…………………………….....73
2.4.1 Experimental Methods………………………………………….....73
2.4.2 Results…………………………………………………………......75
2.5 Conclusions…………………………………………………………..……...88
2.6 References……………………………………………………………….......90
viii
Chapter 3 The Applications of Graphene as Both a Catalyst and Support for Fuel Cells
3.1 Graphene Characteristics and Applications………………………………....99
3.2 The Effect of pH on the Reduction of Graphene Oxide and its
Catalytic Properties Towards ORR in the Basic Medium…………………….....102
3.2.1 Experimental Methods…………………………………………...102
3.2.2 Results…………………………………………………………....105
3.3 A Facile One-Pot Synthesis of Platinum on Electrochemically
Exfoliated Graphite for ORR in Acidic Media………………………………....125
3.3.1 Experimental Methods…………………………………………...125
3.3.2 Results…………………………………………………………....126
3.4 Reduced Graphene Oxide Emulsions as a Catalyst Support for Various Metals
for Fuel Cell Catalysts………………………………………………………...136
3.4.1 Experimental Methods…………………………………………...136
3.4.2 Results…………………………………………………………....137
3.5 Conclusions………………………………………………………………...150
3.6 References……………………………………………………………….....152
Chapter 4 The Modification of Nickel and Supports for Urea Oxidation
4.1 Urea Oxidation for Use in Direct Urea Fuel Cells........................................159
4.2 Effect of Fluorinated Carbon Supports and Annealing in Ni/C
Catalysts for Urea Oxidation…………………………………………………...161
4.2.1 Experimental Methods……………………………………….......161
4.2.2 Results………………………………………………………........164
4.3 Effect of Annealing Temperature on Ni/rGO Catalysts on the
Structure and Urea Oxidation...............………………………………………....187
4.3.1 Experimental Methods……………………………………….......187
4.3.2 Results…………………………………………………………....187
4.4 Effect of Annealing Temperature on NiO@rGO Catalysts on
the Structure and Urea Oxidation...............…………………………………......207
4.4.1 Experimental Methods...................................................................207
4.4.2 Results………………………………………………………........208
4.5 Effect of First Row Transition Metal Doping of Nickel Oxides
on the Urea Oxidation..................…………………………………………….....228
4.5.1 Experimental Methods……………………………………….......228
4.5.2 Results…………………………………………………………....229
4.6 Conclusions…………………………………………………………….......237
4.7 References……………………………………………………………….....239
ix
LIST OF FIGURES
Figure 1.1 History of the global energy consumption from the early 19
th
Century to the 21
st
Century.
Figure 1.2 Breakdown of the world energy consumption in 2010 by energy source.
Figure 1.3 Schematic of the original “gas battery” by William Grove.
Figure 1.4 Schematic of a typical fuel cell currently in use.
Figure 1.5 Typical polarization curve of a fuel cell
Figure 1.6 Summary of various types of fuel cells.
Figure 1.7 Schematic of a proton exchange membrane fuel cell (PEMFC) and an alkaline
fuel cell (AFC).
Figure 1.8 Power output ranges of various types of fuel cells.
Figure 1.9 a) Chart with the difference in vehicles. b) Ragone plot of various power
sources.
Figure 1.10 Utilization of the H2 fuel cell in “The Hydrogen Society”.
Figure 1.11 Utilization of the DMFC in “The Methanol Economy”.
Figure 2.1 TGA curves of the Pt/Ccatalysts with the: a) CFx supports and; b) XC72 supports.
Figure 2.2 SEM images of the Pt/C catalysts: a) Pt-2C and; b) Pt-2F.
Figure 2.3 TEM images of the Pt/C catalysts with the particle size distribution charts of: a)
Pt-1F; b) Pt-2F; c) Pt-3F and; d) Pt-4F. The measuring bar indicates 20 nm.
Figure 2.4 TEM images of the Pt/C catalysts (top) with the particle size distribution charts
(bottom): a) Pt-1C; b) Pt-2C; c) Pt-3C and; d) Pt-4C. The measuring bar indicates
20 nm.
Figure 2.5 XRD patterns of the Pt/C catalysts on the: a) CFx support and; b) XC72 support.
Figure 2.6 XPS spectra of the Pt-2 catalysts: a) the Pt 4f high resolution region and; b) the C
1s high resolution region.
Figure 2.7 CV scans of the Pt/C catalysts in 0.5 M H2SO4 solution at 20 mV s
-1
. The solid
lines indicate the CFx support; the dashed lines indicate the XC72 support.
x
Figure 2.8 CO stripping CVs for the Pt/C catalysts. The solid lines indicate the CFx support;
the dashed lines indicate the XC72 support.
Figure 2.9 The ECSAs of the Pt/C catalysts for the hydrogen adsorption (solid line) and CO
stripping (dashed line).
Figure 2.10 The LSV scans of the Pt/C catalysts at 1600 RPM under O2 at a) normal scan
window and; b) zoom-in of the onset potential region. The solid lines indicate the
CFx supported catalysts; the dashed lines indicate the XC72 supported catalysts.
Figure 2.11 Summary of the onset potentials of the Pt/C catalysts from Figure 2.10.
Figure 2.12 Typical series of LSV scans for the Pt1-F catalyst.
Figure 2.13 Koutecky-Levich plot of the current of the Pt/C catalysts at 0.6 V. The solid
markers indicate the CFx support; the hollow markers indicate the XC72 support.
Figure 2.14 Tafel plot for the Pt/C catalysts. The solid markers indicate the CFx support; the
hollow markers indicate the XC72 support.
Figure 2.15 Polarization curves for the Pt/C catalysts at ambient temperature. The solid lines
indicate the 0.5 mg cm
-2
loadings; the dashed lines indicate the 0.1 mg cm
-2
loadings.
Figure 2.16 Polarization curves for the Pt/C catalysts at 50
O
C. The solid lines indicate the 0.5
mg cm
-2
loadings; the dashed lines indicate the 0.1 mg cm
-2
loadings.
Figure 2.17 EIS curves for the Pt/C catalysts at ambient temperature. The solid markers
indicate the 0.5 mg cm
-2
loadings; the hollow lines indicate the 0.1 mg cm
-2
loadings; the lines indicate the fitted circuit.
Figure 2.18 EIS curves for the Pt/C catalysts at 50
O
C. The solid markers indicate the 0.5 mg
cm
-2
loadings; the hollow lines indicate the 0.1 mg cm
-2
loadings; the lines
indicate the fitted circuit.
Figure 2.19 The equivalent circuit for the EIS spectra.
Figure 2.20 SEM images of the CFx supported catalysts synthesized by a) impregnation and;
b) microwave reduction.
Figure 2.21 TEM images with the corresponding particle size distribution charts of the XC72
supported catalysts synthesized by: a) impregnation and; b) microwave reduction.
Figure 2.22 TEM images with the corresponding particle size distribution charts of the CFx
supported catalysts synthesized by: a) impregnation and; b) microwave reduction.
Figure 2.23 XRD patterns of the μwave Pt/C catalysts.
xi
Figure 2.24 XPS Pt 4f spectra of the a) uwave reduced catalysts and; b) sodium borohydride
reduced catalysts. The (1) spectrum indicates the XC72 supported catalysts; the
(2) spectrum indicates the CFx supported catalysts.
Figure 2.25 XPS C 1s spectra of the a) uwave reduced catalysts and; b) sodium borohydride
reduced catalysts. The (1) spectrum indicates the XC72 supported catalysts; the
(2) spectrum indicates the CFx supported catalysts.
Figure 2.26 CV scans of the μwave Pt/C catalysts in 0.5 M H2SO4 under argon. The solid
lines indicate the microwave reduction; the dotted lines indicate the impregnation.
Figure 2.27 CO stripping curves for the μwave Pt/C catalysts. The solid lines indicate the
microwave reduction; the dashed lines indicate the impregnation.
Figure 2.28 Summary of the ECSAs for the μwave Pt/C catalysts from the CO stripping and
hydrogen adsorption.
Figure 2.29 The LSV scans of the μwave Pt/C catalysts at 1600 RPM under O2 at a) normal
scan window and; b) zoom-in of the onset potential region. The solid lines
indicate the microwave reduction; the dashed lines indicate the impregnation.
Figure 2.28 Zoom-In of the LSV curves highlighting the onset potential of the μwave Pt/C
catalysts. The solid lines indicate the microwave reduction; the dashed lines
indicate the impregnation.
Figure 2.30 The Koutecky-Levich plot of the μwave Pt/C catalysts at 0.6 V vs. RHE. The
solid markers indicate the microwave reduction; the hollow markers indicate the
impregnation.
Figure 2.31 Tafel curves for the μwave Pt/C catalysts. The solid markers indicate the
microwave reduction; the hollow markers indicate the impregnation.
Figure 2.32 Polarization curves for the H2/O2 fuel cells of the μwave Pt/C catalysts at: a)
ambient temperature and; b) 50
O
C. The solid lines indicate the microwave
reduced catalysts; the dashed lines indicate the impregnated catalysts.
Figure 2.33 EIS curves of the μwave Pt/C catalysts at 0.4 V vs. RHE at: a) ambient
temperature and; b) 50
O
C. The solid markers indicate the microwave reduced
catalysts; the hollow markers indicate the impregnated catalysts; the lines indicate
the fitted circuit.
Figure 2.34 CV scan for the fuel cells of the μwave Pt/C catalysts. The solid lines indicate the
microwave reduction; the dashed lines indicate the impregnation.
Figure 2.35 Polarization curves for the H2/Air fuel cells for the μwave Pt/C catalysts at: a)
ambient temperature and; b) 50
O
C with air flowing in the cathode. The solid
lines indicate the microwave reduced catalysts; the dashed lines indicate the
impregnated catalysts.
xii
Figure 2.36 SEM images of: a) α-MnO2; b) on XC72 support and; c) on CFx support.
Figure 2.37 SEM images of a) β-MnO2; b) on XC72 support and; c) on CFx support.
Figure 2.38 SEM images of: a) γ-MnO2; b) on XC72 support and; c) on CFx support.
Figure 2.39 EDAX spectrum of α-MnO2/CFx.
Figure 2.40 a) SEM image of α-MnO2/CFx with corresponding EDAX elemental mapping
images of: b) carbon; c) fluorine, d) manganese and; e) oxygen.
Figure 2.41 XRD patterns of the: a) α-MnO2; b) β-MnO2 and; c) γ-MnO2 phase catalysts.
Figure 2.42 CV scans of the MnO2/C catalysts. The solid lines indicate the samples with CFx
support; the dotted lines indicate the samples with XC72 support.
Figure 2.43 LSV scans of the MnO2/C catalysts in 0.1M KOH solution. a) -MnO2/CFx scans
from 100 RPM to 1600 RPM; b) scans of all the catalysts at 1600 RPM. The
solid lines indicate the catalysts with CFx support; the dotted lines indicate the
catalysts with the XC72 support.
Figure 2.44 Koutecky-Levich plot of the MnO2/C catalysts at -0.4 V vs MMO. The solid
markers indicate the catalysts with the CFx support; the hollow markers indicate
the catalysts with the XC72 support.
Figure 2.45 Tafel plots of the MnO2/C catalysts at 1600 RPM. The solid markers indicate the
catalysts with the CFx support; the hollow markers indicate the catalysts with the
XC72 support.
Figure 2.47 Polarization curves of the alkaline DMFCs for the MnO2/C catalysts at: a)
ambient temperature and; b) 50
O
C.
Figure 3.1 Number of publications based on graphene by year.
Figure 3.2 Various implementations of 2D graphene.
Figure 3.3 TGA curves of the rGO catalysts.
Figure 3.4 SEM images of the rGO catalysts: a) GO; b) rGO-1; c) rGO-4; d) rGO-7; e) rGO-
10 and; f) rGO-13.
Figure 3.5 TEM images of the rGO catalysts: a) GO; b) rGO-1; c) rGO-4; d) rGO-7; e) rGO-
10 and; f) rGO-13. The measuring bar indicates 100 nm.
Figure 3.6 XRD patterns of the rGO catalysts.
Figure 3.7 Summary of the XRD characteristics of the rGO catalysts from Figure 3.6.
xiii
Figure 3.8 FTIR spectra of the rGO catalysts.
Figure 3.9 Raman spectra of the rGO catalysts.
Figure 3.10 Sumary of the Raman spectra of the rGO catalysts from Figure 3.9.
Figure 3.11 C1s XPS Spectra of the rGO catalysts: a) GO; b) rGO-1; c) rGO-4; d) rGO-7; e)
rGO-10 and; f) rGO-13.
Figure 3.12 a) Titration curve of the rGO catalyst solutions in 0.1 M NaOH and; b)
concentration of ionized species in the rGO catalysts at certain pH values.
Figure 3.13 The pKa distributions of the acid groups on the rGO catalysts from the
concentrations in Figure 3.12 with Gaussian distribution curves for: a) rGO-1; b)
rGO-4; c) rGO-7; d) rGO-10 and; e) rGO-13.
Figure 3.14 CV scans of the rGO catalysts in 0.1 M KOH at 20 mV s
-1
.
Figure 3.15 LSV scans of the rGO catalysts in 0.1 M KOH under O2 at 1600 RPM.
Figure 3.16 Summary of the LSV scans of the rGO catalysts from Figure 3.15.
Figure 3.17 Koutecky-Levich plot of the LSV scans of the rGO catalysts.
Figure 3.18 Tafel curves of the rGO catalysts in 0.1 M KOH under O2 at 1600 RPM.
Figure 3.19 Chronoamperommetry plots of the rGO catalysts in at -0.5 V vs. MMO at 1600
RPM.
Figure 3.20 SEM images of the Pt/exrGO catalysts at: a) 2.5 V; b) 5 V; c) 8 V and; d) 10 V.
Figure 3.21 TEM images of the Pt/exrGO catalysts at: a) 2.5 V; b) 5 V; c) 8 V; d) 10 V. The
measuring bar indicates 100 nm.
Figure 3.22 Characterization of the exfoliated graphite rods: a) XRD patterns and; b) Raman
spectra.
Figure 3.23 Characterization of the Pt/exrGO catalysts: a) XRD patterns and; b) Raman
spectra.
Figure 3.24 CV scan of the Pt/exrGO catalysts under argon.
Figure 3.25 LSV scan of the Pt/exrGO catalysts under O2 at 1600 RPM.
Figure 3.25 CO Stripping of the Pt/exrGO catalysts.
Figure 3.26 LSV of the Pt/exrGO catalysts under O2 at 1600 RPM.
Figure 3.27 SEM images of the rGOem catalysts: a) GO; b) GOem; c) rGO and; d) rGOem.
xiv
Figure 3.28 SEM images of the Pt/rGOem catalysts: a) Pt/rGO impreg; b) Pt/rGOem impreg;
c) Pt/rGO uwave and; d) Pt/rGOem uwave.
Figure 3.29 XRD patterns of the rGOem catalysts: a) rGOem and; b) Pt/rGOem.
Figure 3.30 Raman spectra of the rGOem catalysts: a) rGOem and; b) Pt/rGOem.
Figure 3.31 XPS C1s spectra of the rGOem catalysts: a) GOem; b) rGOem; c) Pt/rGOem
impreg and; d) Pt/rGOem uwave. The (1) spectrum indicates GOem supported
catalysts; the (2) spectrum indicates the regular GO supported catalysts.
Figure 3.32 XPS Pt 4f spectra of the Pt/rGOem catalysts: a) Pt/rGOem impreg and; b)
Pt/rGOem uwave. The (1) spectrum indicates GOem supported catalysts; the (2)
spectrum indicates the regular GO supported catalysts.
Figure 3.33 CV scans of the Pt/rGOem catalysts in 0.5 M H2SO4 under argon.
Figure 3.34 CO Stripping scans of the Pt/rGOem catalysts.
Figure 3.35 LSV curves of the Pt/rGOem catalysts in 0.5 M H2SO4 under O2 at 1600 RPM.
Figure 3.36 Koutecky-Levich plot of the Pt/rGOem catalysts.
Figure 3.37 Tafel curves of the Pt/rGOem catalysts in 0.5 M H2SO4 under O2 at 1600 RPM.
Figure 4.1 TGA curves of the Ni/C catalysts.
Figure 4.2 SEM images of: a) Ni; b) Ni-400; c) Ni/CFx; d) Ni/CFx-400; e) Ni/XC72 and; f)
Ni/XC72-400.
Figure 4.3 TEM images of the catalysts with particle size distribution analysis of: a) Ni; b)
Ni-400; c) Ni/XC72; d) Ni/CFx; e) Ni/XC72-400 and; f) Ni/CFx-400.
Figure 4.4 XRD patterns of the: a) Ni/C catalysts and; b) carbon supports.
Figure 4.5 Raman spectra of the: a) Ni/C catalysts and; b) carbon supports.
Figure 4.6 XPS spectra of Ni 2p of the Ni/C catalysts: a) Ni; b) Ni/XC72 and; c) Ni/CFx.
The (1) spectrum indicates the catalyst annealed at 400
O
C; the (2) spectrum
indicates the unannealed catalyst.
Figure 4.7 XPS C1s spectra of the Ni/C catalysts: a) Ni/CFx; b) Ni/XC72; c) CFx and; d)
XC72. The (1) spectrum indicates the catalyst annealed at 400
O
C; the (2)
spectrum indicates the unannealed catalyst.
Figure 4.8 CV scans of the Ni/C catalysts in: a) 1.0 M KOH + 0.33 M urea solution and; b)
1.0 M KOH solution at 20 mV s
-1
.
xv
Figure 4.9 CV scans of the carbon supports in: a) 1.0M KOH + 0.33M urea solution and; b)
1.0M KOH solution at 20 mV s
-1
.
Figure 4.10 Step CV of the Ni/C catalysts in 1.0 M KOH + 0.33 M urea solution.
Figure 4.11 a) CV of the Ni/CFx-400 catalyst at different scan rates; b) the relationship
between the square root of the scan rate and the current density of the Ni/C
catalysts.
Figure 4.12 a) CV of Ni/XC72 catalyst cycled 50 times for stability; b) stability plots at 0.65
V for the Ni/C catalysts.
Figure 4.13 Chronoamperommetry graph of the Ni/C catalysts at 1000 RPM.
Figure 4.14 Diagram of the micro direct urea/hydrogen peroxide fuel cell.
Figure 4.15 Polarization curves for the Ni/C catalysts in the micro direct urea fuel cells.
Figure 4.16 EIS curves for the Ni/C catalysts in the micro direct urea fuel cells.
Figure 4.17 Equivalent circuit for the EIS spectra for the Ni/C catalysts.
Figure 4.18 SEM images of the Ni/rGO catalysts: a) Ni; b) Ni/rGO; c) Ni/rGO-300; d)
Ni/rGO-400; e) Ni/rGO-500; f) Ni/rGO-600 and; g) Ni/rGO-700.
Figure 4.19 TEM images of the Ni/rGO catalysts: a) Ni; b) Ni/rGO; c) Ni/rGO-300; d)
Ni/rGO-400; e) Ni/rGO-500; f) Ni/rGO-600 and; g) Ni/rGO-700. The scale
measures 500 nm.
Figure 4.20 XRD patterns for the Ni/rGO catalysts.
Figure 4.21 Summary of the XRD patterns for the Ni/rGO catalysts from Figure 4.20.
Figure 4.22 Raman spectra for the Ni/rGO catalysts.
Figure 4.23 Summary of the raman spectra for the Ni/rGO catalysts from Figure 4.22.
Figure 4.24 XPS Ni 2p spectra of the Ni/rGO catalysts: a) Ni/rGO; b) Ni/rGO-300; c)
Ni/rGO-400; d) Ni/rGO-500; e) Ni/rGO-600 and; f) Ni/rGO-700.
Figure 4.25 XPS C 1s spectra of the Ni/rGO catalysts: a) Ni/rGO; b) Ni/rGO-300; c) Ni/rGO-
400; d) Ni/rGO-500; e) Ni/rGO-600 and; f) Ni/rGO-700.
Figure 4.26 CV scans of the Ni/rGO catalysts in: a) 1.0 M KOH + 0.33 M urea and; b) 1.0 M
KOH only.
Figure 4.27 Summary of the CV scans for the Ni/rGO catalysts from Figure 4.26a.
Figure 4.28 Step CV for the Ni/rGO catalysts in 1.0 M KOH + 0.33 M urea.
xvi
Figure 4.29 a) CV of the Ni/rGO-500 catalyst at different scan rates; b) the relationship
between the square root of the scan rate and the current density of the Ni/rGO
catalysts.
Figure 4.30 a) CV of Ni/rGO-500 catalyst cycled 50 times for stability; b) stability plots at
0.65 V for the Ni/rGO catalysts.
Figure 4.31 Chronoamperommetry graph of the Ni/rGO catalysts at 1000 RPM.
Figure 4.32 Polarization curves for the Ni/rGO catalysts in the micro direct urea fuel cells.
Figure 4.33 EIS curves for the Ni/rGO catalysts in the micro direct urea fuel cells.
Figure 4.34 SEM images of: a) GO; b) NiO; c) Ni@rGO-300; d) Ni@rGO-400; e) Ni@rGO-
500; f) Ni@rGO-600 and; g) Ni@rGO-700.
Figure 4.35 TEM images of a) GO; b) NiO; c) Ni@rGO-300; d) Ni@rGO-400; e) Ni@rGO-
500; f) Ni@rGO-600; and g) Ni@rGO-700. The measuring bar indicates 500 nm.
Figure 4.36 XRD patterns of the NiO@rGO catalysts.
Figure 4.37 Summary of the XRD patterns of the NiO@rGO catalysts from Figure 4.36.
Figure 4.38 Raman spectra of the NiO@rGO catalysts.
Figure 4.39 Summary of the Raman spectra of the NiO@rGO catalysts from Figure 4.38.
Figure 4.40 XPS Ni 2p spectra of the NiO@rGO catalysts: a) NiO; b) NiO@rGO; c)
Ni@rGO-300; d) Ni@rGO-400; e) Ni@rGO-500; f) Ni@rGO-600 and; g)
Ni@rGO-700.
Figure 4.41 XPS C 1s spectra of the NiO@rGO catalysts: a) NiO@rGO; b) Ni@rGO-300; c)
Ni@rGO-400; d) Ni@rGO-500; e) Ni@rGO-600 and; f) Ni@rGO-700.
Figure 4.42 CV scans of the NiO@rGO catalysts in: a) 1.0M KOH + 0.33M urea and; b) 1.0M
KOH only.
Figure 4.43 Summary of the CV scans of the NiO@rGO catalysts from Figure 4.42a.
Figure 4.44 Step CV for the NiO@rGO catalysts in 1.0 M KOH + 0.33 M urea.
Figure 4.45 a) CV of the NO@rGO-700 catalyst at different scan rates; b) the relationship
between the square root of the scan rate and the current density of the NiO@rGO
catalysts.
Figure 4.46 a) CV of NiO@rGO-600 catalyst cycled 50 times for stability; b) stability plots at
0.65 V for the NiO@rGO catalysts.
Figure 4.47 Chronoamperommetry graph of the NiO@rGO catalysts at 1000 RPM.
xvii
Figure 4.48 Polarization curves for the NiO@rGO catalysts in the micro direct urea fuel cells.
Figure 4.49 EIS curves for the NiO@rGO catalysts in the micro direct urea fuel cells.
Figure 4.50 Representative SEM image and EDS mapping of the NixZn1-xO catalyst: a) SEM
image at low magnification; b) SEM/EDS image at high magnification; c) EDS
mapping of nickel at high magnification and; d) EDS mapping of zinc at high
magnification.
Figure 4.51 XRD patterns of the: a) NixM1-x(OH)2 and; b) NixM1-xO catalysts.
Figure 4.52 CV scans of the NixM1-xO catalysts in: a) 1.0 M KOH + 0.33 M urea and; b) 1.0
M KOH solution.
Figure 4.52 Summary plots of the CV scans for the NixM1-xO catalysts for: a) onset potential
and; b) current density from Figure 4.52a.
xviii
LIST OF SCHEMES
Scheme 1.1. Oxygen reduction reaction mechanism in: a) acidic media and; b) basic media.
Scheme 3.1. Overview of the one-pot exfoliation/reduction synthesis.
xix
LIST OF TABLES
Table 2.1 Summary of the TGA results for the platinum loadings of the Pt/C catalysts.
Table 2.2 Summary of the RDE electrochemical tests of the Pt/C catalysts.
Table 2.3 Summary of the fuel cell tests of the Pt-2F and Pt-2C catalysts.
Table 2.4 Summary of the RDE electrochemical tests of the μwave Pt/C catalysts.
Table 2.5 Summary of the H2/O2 fuel cell tests of the μwave Pt/C catalysts.
Table 2.6 Summary of the RDE electrochemical tests of the MnO2/C catalysts.
Table 2.7 Summary of the fuel cell tests of the MnO2/C catalysts.
Table 3.1 Summary of the deconvoluted C1s XPS spectra of the rGO catalysts.
Table 3.2 Summary of the characterization and electrochemical tests of the rGO catalysts.
Table 3.3 Summary of the synthesis and characterization tests of the Pt/exrGO catalysts.
Table 3.4 Summary of the electrochemical tests of the Pt/exrGO catalysts.
Table 3.5 Summary of the characterization tests of the rGOem catalysts.
Table 3.6 Summary of the characterizations of the Pt/rGOem catalysts.
Table 3.7 Summary of the electrochemical tests of the Pt/rGOem catalysts.
Table 4.1 Summary of the XPS C1s spectra of the Ni/C catalysts.
Table 4.2 Summary of the characterization and electrochemical tests of the Ni/C catalysts.
Table 4.3 Summary of the fuel cell tests of the Ni/C catalysts.
Table 4.4 Summary of the electrochemical and fuel cell tests of the Ni/rGO catalysts.
Table 4.5 Summary of the electrochemical and fuel cell tests of the NiO@rGO catalysts.
Table 4.6 Summary of the characterization tests of the Ni-doped catalysts.
xx
ABSTRACT
In order for the further commercialization of the myriad of fuel cell technologies that are
present today, considerable work still needs to be done to optimize the various parameters of the
fuel cell. One of the parameters is the catalyst design including the modification of both catalyst
and catalyst support morphologies and properties.
In Chapter 2, partially fluorinated high surface area carbon (CFx) supports for platinum
were assessed and compared to the widely implemented Vulcan Carbon (XC72) supports. The
Pt/CFx catalysts displayed enhanced oxygen reduction reaction (ORR) kinetics at the lower
platinum loadings where the CFx properties were more visible. The platinum supported catalysts
were also synthesized via a microwave-assisted polyol reduction method. The microwave
reduction method yielded enhanced ORR kinetics compared with the impregnated catalysts while
the CFx supported platinum yielded higher kinetics than the XC72 supported platinum. In both
instances, the Pt/CFx displayed higher hydrogen fuel cell performances than the Pt/XC72 catalyst.
Various phases of manganese dioxide were prepared and grafted onto the aforementioned carbon
supports. Again, the CFx supported MnO2 displayed higher kinetics than the XC72 supported
MnO2 catalysts in both the half-cell and alkaline direct methanol fuel cell testing. The enhanced
conductivity from the fluorination in the CFx with the Nafion® ionomer and Tokuyama anionomer
in the catalyst layers in the fuel cells led to enhanced performance in the fuel cells in both acidic
and alkaline media, respectively.
In Chapter 3, reduced graphene oxide (rGO) was synthesized by various routes including
Hummer’s method and electrochemical exfoliation. The graphene oxide (GO) synthesized from
the Hummer’s method was dispersed in various aqueous solutions with pH values of 1-13 before
being reduced by sodium borohydride. The lower pH values led to less repair of the sp
2
graphitic
xxi
lattice as confirmed by Raman and X-Ray Photoelectron Spectroscopy (XPS) due to the increased
degradation of sodium borohydride in acidic solutions. This led to lower ORR kinetics in alkaline
half-cell testing. Reduced graphene oxide was also synthesized from electrochemical exfoliation
from graphite rods and deposited on platinum in a one pot procedure. The potential of exfoliation
affected the lattice of the rGO impacting the ORR activity of the platinum supported catalysts.
Graphene oxide was also kept in an emulsion form from Hummer’s method without drying. The
GOem displayed similar structural characteristics in both the oxidized and reduced forms to that
of the dried form as well as with platinum reduced on it by impregnation and microwave reduction.
In Chapter 4, nickel was deposited onto various supports to examine the effects on urea
electrooxidation. As in Chapter 2, nickel was impregnated onto both CFx and XC72 to examine
to effects of fluorination. The nickel displayed smaller particle sizes on the CFx compared with
the XC72. This led to enhanced urea electrooxidation kinetics in both the half-cell and micro fuel
cell testing. Mild oxidative annealing was also applied to the catalysts, which resulted in lower
activity. Nickel was also reduced onto GO and annealed at temperatures from 300 – 700
O
C under
argon to study the effects. The rGO displayed greater disorder from the Raman and XPS spectra
from the removal of the oxygen functional groups in the lattice. This led to the partial oxidation
of the nickel nanoparticles, leading to decreased kinetics in both the half-cell and micro fuel cell
tests. Nickel oxide was also grafted onto rGO and annealed at the same temperatures under argon.
The same effects were viewed in the graphitic lattice. However, the catalysts showed an increase
in urea electrooxidation activity at the higher annealing temperatures in the half-cell and micro
fuel cell tests. Finally, nickel was co-doped with all of the third row transition metal elements at
a 7%-wt and annealed under air to form nickel oxides. This doping led to decreased urea
electrooxidation activity from the strain in the nickel lattice.
1
Chapter 1: Introduction
1.1: The Current State of Energy
Ever since the industrial revolution in the mid-19
th
Century, humankind has been
consuming energy, mainly in the form of fossil fuels and coal, at an ever-increasing rate. This
trend is illustrated in Figure 1.1. Presently, the global energy consumption is around 10 billion
gigawatts, up from only 1 billion gigawatts in 1990 [1]. This is expected to continue to increase
due to the increasing human population as well as the continuing industrialization and
modernization of many 3
rd
world countries with large populations such as China and India [2].
The use of fossil fuels, including coal and natural gas, while a relatively simple combustion process
that has been used for centuries, has many drawbacks. One main drawback is the rising level of
CO2 emissions, a greenhouse gas, which has recently risen to global levels above 400 ppm [3].
This marks a feat not seen for hundreds of thousands of year on this planet [4] and is widely viewed
as one of the leading causes of global warming and climate change [5]. Moreover, the burning of
fossil fuels can also emit other polluting gases such as methane, NOx, SOx, and particulate matter,
which can be detrimental to the environment and human health [6]. Another main drawback is the
finite availability of fossil fuels and the relatively unstable regions in which many reserves are
located [7]. Some of the newer methods and fuel sources such as fracking for the extraction of
natural gases from the ground can also be potentially harmful to the environment [8]. Furthermore,
the likelihood of spills and accidents in both extraction such as the gulf oil spill or transportation
such as the Exxon Valdez can also be considerably hazardous to the environment and human health.
2
Figure 1.1. History of the global energy consumption from the early 19
th
Century to the 21
st
Century. Adapted from [9].
In the past half century, however, the development and use of alternative and renewable
energy has also greatly increased. As shown in Figure 1.1, the amount of energy consumed
worldwide from renewable sources exceeded 100 exajoules in 2000. However, fossil fuels still
comprise about 80% of the energy sources worldwide as shown in Figure 1.2. Various devices
such as solar cells, wind power, hydroelectric power, and nuclear energy have the potential to help
alleviate the use of non-sustainable energy sources, but also have their own drawbacks. Solar cells
fail to produce energy when the sun is not shining and wind turbines fail to produce energy when
there is no wind. Another such drawback is the need for energy storage, when the energy demand
and the supply occur at different times. Hydroelectric power can be detrimental to the surrounding
environment, often displacing people and wildlife. Numerous nuclear power plants are aging and
pose the risk of a nuclear meltdown which poses a great radiation risk to the environment. Even
if the safety measures are upgraded, there is still a stigma from the catastrophes from the Chernobyl
and Fukushima meltdowns. Another alternative energy source that has been gaining increased
attention recently that can operate in a manner similar to fossil fuel and coal combustion without
many of the hazardous side effects is fuel cells.
3
Figure 1.2. Breakdown of the world energy consumption in 2010 by energy source. Adapted
from [10].
1.2: History of Fuel Cells
In the past few decades, fuel cells have emerged as an increasingly viable source of
alternative energy, both for commercial and personal use. This technology however has been
around much longer, like that of other energy sources such as solar cells and batteries. The concept
of a fuel cell was first developed by William Grove in 1839 [11]. Originally, it was called the “gas
battery” because it involved gases being flowed into the cell, which were connected by electrodes.
A schematic of this device is shown in Figure 1.3 and is shown in both a single cell (left) and a
series of cells (right).
The device wasn’t called a “fuel cell” until 1889 when Ludwig Mond and Charles Langer
built the first practical device [12]. Another half a century passed with minimal progress until
Francis Bacon replaced some key elements of the device and coined the term “Bacon Cell” in 1932
[13]. Still, the fuel cell lacked practical applications until alkaline fuel cells (AFCs) were used in
4
the 1960s and 1970s for NASA’s early spacecraft missions [14]. It wasn’t until the energy crisis
and oil embargo of the 1970s that hydrogen fuel cell research was geared toward commercial
application in automobile use. More recently, direct methanol fuel cells (DMFCs) started by
researchers at the University of Southern California and NASA’s Jet Propulsion Lab in the early
1990s have gained much attention as both a portable power source and for commercial applications
[15]. Other fuel cells that have been utilized run at higher temperatures such as the solid oxide
fuel cells and molten carbonate fuel cells, will be briefly discussed in Section 1.4.3.
Figure 1.3. Schematic of the original “gas battery” by William Grove. Adapted from [16].
1.3: Fuel Cell Basics
Fuel cells rely on the energy from the catalyzed chemical reactions to produce electricity
and the potential difference to drive the current. Fuel cells, on average, are more efficient than the
5
typical combustion engine (around 12%) [17] used in most vehicles and machinery being operated
today at around 40% [18,19] depending on the type of fuel cell and the operating parameters . A
basic schematic of most of the fuel cells being utilized presently is shown in Figure 1.4. The fuel
cell consists of two main compartments, where the main reactions are carried out. The oxidation
reaction of the fuels takes place in the anode compartment while the reduction reaction is carried
out at the cathode compartment. The two compartments are separated usually by a semi-permeable
membrane that allows certain charged species to pass from one compartment to another depending
on the fuel cell. In many low temperature fuel cells (which is what the bulk of this thesis will be
focused on), the catalyst is usually applied as a thin layer either directly onto the membrane or on
to the gas diffusion layers (GDLs). The electrons produced by the oxidation reaction at the anode
flow through the external wires to the “load” or source to be powered and to the cathode for the
reduction reaction.
Figure 1.4. Schematic of a typical fuel cell currently in use.
Current Collectors
Load
Membrane
Flow Fields
Gas Diffusion
Layer (GDL)
Catalyst Layer
6
The main technique to assess the performance of a fuel cell is a polarization curve shown
in Figure 1.5. The polarization curve is usually plotted with the cell voltage or potential on the y-
axis and the current density on the x-axis. The open circuit voltage (OCV) of a fuel cell is lower
than the theoretical OCV due to factors including parasitic losses from fuel mixture across the
membrane [20]. In the region of low current density, the initial decrease in potential is from the
activation of the catalyst [21]. The next region, characterized by a linear decrease, is the ohmic
region characterized by ohmic (IR) losses from the internal resistance of the fuel cell. The high
current density region is characterized by mass transport losses and limitations. A power density
curve can also be formulated from integrating the polarization curve with respect to current density,
resembling an inverse parabola shape. Most of the fuel cells (specifically stationary) operate at a
continuous voltage, which usually corresponds to the peak power density derived from the
polarization curve. Since the OCV is usually less than 1.5 V, fuel cell stacks comprised of
individual fuel cells in series and parallel are required to power various vehicles and machinery.
Figure 1.5. Typical polarization curve of a fuel cell. Adapted from [22].
7
1.4: Types of Fuel Cells
There are a variety of fuel cells that are being developed based on the type of chemical
reactions, membrane/electrolyte, setup, and temperature of operation. A summary of some of the
different types of fuel cells is shown in Figure 1.6. Fuel cell operating temperatures can range
from around room temperature to over 1000
O
C depending on the type of fuel cell and the particular
operating parameters.
Figure 1.6. Summary of various types of fuel cells. Adapted from [23].
1.4.1: Proton Exchange Membrane Fuel Cells (PEMFCs)
One of the most widely utilized low temperature fuel cell is the PEM fuel cell. The main
focus in the PEMFC is the membrane, which conducts protons from the anode to the cathode
compartment as shown in Figure 1.7a. In this case, the reactants such as hydrogen are flowed
8
through the GDL of the anode compartment, where they are oxidized on the catalyst. The protons
formed cross through a proton conducting membrane to the cathode, where the reactants are flowed
through and are reduced. Most of the PEMs have a tetrafluoroethylene backbone with a sulfonate
group attached to it as the proton conductor. The basic cell reactions are shown below [24]:
Anode: 2H2 → 4H
+
+ 4e
-
(1)
Cathode: O2 + 4H
+
+ 4e
-
→ 2H2O (2)
Overall: 2H2 + O2 → 2H2O (3)
The only emission product associated with the hydrogen fuel cell is pure water. There are a few
drawbacks associated with PEMFCs. Some include the requirement of noble metal catalysts on
the anode as well as for the oxygen reduction reaction on the cathode in an acidic environment,
the low tolerance of CO and other impurities leading to catalyst poisoning, and the storage and
transportation of hydrogen gas [25,26].
Figure 1.7. Schematic of: a) a proton exchange membrane fuel cell (PEMFC) and; b) an
alkaline fuel cell (AFC).
Reactants Flowing
Into Anode
Excess
Reactants/
Products Flowing
Out of Anode
Reactants Flowing
Into Cathode
Excess
Reactants/
Products Flowing
Out of Cathode
Anode
Cathode
e
-
e
-
H
+
H
+
H
+
H
+
a)
Reactants Flowing
Into Anode
Excess Reactants/
Products Flowing
Out of Anode
Reactants Flowing
Into Cathode
Excess Reactants/
Products Flowing
Out of Cathode
Anode
Cathode
e
-
e
-
OH
-
OH
-
OH
-
b)
9
Ever since their invention in the 1990s, direct methanol fuel cells (DMFCs), another type
of PEMFC, have gained importance and are being widely studied [15]. One of the appealing
aspects of DMFCs is the portability as well as the ease of fuel storage and transportation. DMFCs
are compact, lightweight, and can be quickly reloaded as well as being able to operate at relatively
low temperatures (around 30 C to 60 C) [27–30]. The main reactions are shown below:
Anode: CH3OH + H2O → CO2 + 6H
+
+ 6e
-
(4)
Cathode: 3/2O2 + 6H
+
+ 6e
-
→ 3H2O (5)
Overall: CH3OH + 3/2O2 → 2H2O (6)
There are some factors however, that can hinder DMFCs performance. Some of these include
methanol crossover in the membrane [31] and CO poisoning on the platinum catalyst surface [32].
1.4.2: Alkaline Fuel Cells (AFCs)
Alkaline fuel cells use KOH as the main electrolyte for the reactions and show great
potential for the fuel cell industry. A schematic is shown in Figure 1.7b showing the similarities
and differences between the PEMFCs and the AFCs. In the case of the AFCs, hydroxide ions
travel across the membrane from the cathode compartment to the anode compartment, where they
are involved in the oxidation of the reactants. Many of the AEMs involve the use of quaternary
ammonium or imidazolium hydroxides to conduct the hydroxide ions [33,34]. One of the
advantages of the AFC is the ability to employ non-noble metal catalysts for the oxidation and
reduction reactions such as nickel [35]. The main reactions in alkaline fuel cells are shown in the
equations below [36]:
Anode: 2H2 + 4OH
-
→ 4H2O + 4e
-
(7)
10
Cathode: O2 + 2H2O + 4e
-
→ 4OH
-
(8)
Overall: 2H2 + O2 → 2H2O (9)
Some of the drawbacks associated with AFCs is the low conductivity and membrane durability in
basic environment [37,38]. Another drawback is the carbonate that can build up in the anode
compartment which is detrimental to fuel cell activity [39,40]. Further optimization of both the
system and AEM is required for widespread commercialization.
1.4.3: High Temperature Fuel Cells
Other fuel cells currently in use are run at high temperatures relative to that of the PEMFCs
and AFCs (>500
O
C). Two of the main types of high temperature fuel cells are the molten
carbonate (MCFC) and solid oxide (SOFC) fuel cells. MCFCs were first introduced in the 1960s
and developed by Broers and Ketelaar [41]. Carbonate (CO3
2-
) is formed at the cathode from the
oxygen and carbon dioxide and travels to the anode via a molten carbonate at high temperatures
in the following reactions [42]:
Anode: 2H2 + 2CO3
2-
→ 2H2O + 2CO2 + 4e
-
(10)
Cathode: O2 + 2CO2 + 4e
-
→ 2CO3
2-
(11)
Overall: 2H2 + O2 → 2H2O (12)
The main catalyst used in MCFCs is nickel oxide usually at both the anode and cathode [43]. There
are certain advantages with MCFCs including very high operating efficiencies, the use of non-
noble metal catalysts, the recyclability of CO2 in the system, and the use of various types of fuels
[44]. However, there are issues with the corrosivity of the electrolyte at the operating temperature
as well as the necessity of continuous CO2 supply at the cathode [45].
11
SOFCs are the highest temperature fuel cells which are usually run at temperatures of
around 700
O
C to 1000
O
C. In the cell, O
2-
is formed at the cathode and travels through a dense
oxygen ion conducting ceramic membrane to the anode. One of the preferred catalysts for these
high temperature fuel cells is nickel either alone or on yttrium stabilized zirconia [46,47]. They
are based on the following redox reactions [48]:
Anode: 2H2 + O
2-
→ 2H2O + 4e
-
(13)
Cathode: 2O2 + 4e
-
→ 2O
2-
(14)
Overall: 2H2 + O2 → 2H2O (15)
Some of the advantages of SOFCs are long life duration, low emissions of CO, flexibility of fuels,
and the use of non-noble metals to operate [11]. However, SOFCs require a long start up time to
achieve the high temperatures, require high temperatures to operate, and have issues with the
thermal cycling needed to operate [49].
Each type of fuel cell has various advantages and disadvantages corresponding to the power
output, catalyst requirements, operating requirements, and various other factors related to the
overall performance and operation. These conditions are being studied and next generation fuel
cells are being developed to increase performance and efficiency while lowering material and
manufacturing costs for the potential of increased commercialization.
1.5: Fuel Cell Utilization
There are numerous applications involved with the variety of fuel cells in use, from
powering portable devices to powering large-scale factories as shown in Figure 1.8. The low
temperature fuel cells such as DMFCs and PEMFCs have lower power outputs, which are suitable
12
for powering portable devices such as laptops, small gadgets, and even vehicles. High temperature
fuel cells such as MCFCs and SOFCs have higher power outputs overall, which are suitable for
stationary devices and large power plants that can accommodate the high temperatures required to
operate the fuel cells.
Figure 1.8. Power output ranges of various types of fuel cells. Adapted from [50].
Since the focus of this thesis will be mainly on PEMFCs and AFCs, the remainder of this
section will be dedicated to the utilization of these particular fuel cells. In the past few years, the
hydrogen fuel cell vehicle has gained in popularity with the release of the Toyota Mirai, Hyundai
Tucson, and most recently the Honda Clarity. Considerable research from these and other major
car manufacturing companies has gone into the development of a fuel cell car as well as larger
vehicles. Figure 1.9a shows a comparison of traveling distance and vehicle size of various fuel
powered vehicles. The hydrogen fuel cell is mostly suitable for the larger trucks/vehicles as well
as mid-size cars due to its necessity for more space to accommodate the various components
including pressurized tanks to store the hydrogen at high pressures. The travel distance is derived
mainly from the energy density of the system, which is shown in the Ragone plot in Figure 1.9b.
13
While the fuel cell is among the lowest in power density output, it is the highest in energy density,
meaning longer travel times before refilling. Furthermore, the fueling time for the hydrogen fuel
cell cars on the market are around 5-10 minutes for a range of over 350 miles, making the refueling
times more appealing than the current electric vehicles.
Figure 1.9. a) Chart with the difference in vehicles. Adapted from [51]. b) Ragone plot of
various power sources. Adapted from [52].
Along with the utilization in medium-sized and large vehicles, fuel cells can also play a
part in more expansive, interconnected systems and processes. Various complex, interwoven
“societies” have been developed that would utilize various energy sources and deal with the
recycling and regeneration of CO2 and other emissions from fossil fuels and energy sources.
a)
b)
14
Figure 1.10 shows the “Hydrogen Society” which involves the implementation of fossil fuels and
renewable energies to both extract and recycle energy through means of a smart grid. In this
society, hydrogen fuel cells play a role of transforming the hydrogen gas formed by means of
electrolysis or other conversion into water, which would then by recycled and regenerated.
Another such society or economy, proposed by the late Nobel Laureate George Olah and Professor
Prakash, is the “Methanol Economy” shown in Figure 1.11. In this case, fossil fuels and CO2 from
the atmosphere and power plants is electrochemically converted into methanol, which can be used
in combustion engines, DMFC vehicles, and stationary devices.
Figure 1.10. Utilization of the H2 fuel cell in “The Hydrogen Society”. Adapted from [53].
Figure 1.11. Utilization of the DMFC in “The Methanol Economy”. Adapted from [54].
15
1.6: The Oxygen Reduction Reaction
The oxygen reduction reaction (ORR) is one of the most widely studied electrochemical
reactions in chemistry and most often used in the fuel cells at the cathode. This is often regarded
as the limiting factor in the chemical reaction pathway, yielding slow kinetics, when compared to
the various oxidation reactions carried out at the anode compartment. The kinetics of this process
is sluggish, mainly due to the difficulty in breaking of the strong oxygen double bond [55] leading
to lower onset potentials, lower current densities, and decreased fuel cell performance [56,57].
Numerous studies have been carried out on the various pathways of the ORR, however, the detailed
mechanism is still being debated.
1.6.1: ORR in Acidic Media
The main pathway for ORR in the acidic media is the 4 e
-
pathway shown in Scheme 1.1a
[58]. This leads to the formation of water and is the preferred pathway. Another pathway is the 2
e
-
pathway which is less desirable because it forms hydrogen peroxide, which can corrode parts of
the fuel cell [59]. Platinum is viewed as the prime metal catalyst for ORR in many types of fuel
cells. It often exhibits the highest ORR activity, when compared to other metal catalysts, most
notably in the acidic media [60,61]. Platinum, however, suffers from CO poisoning at low levels
and can be unstable during potential cycling and operation [62,63]. It is also rare and expensive,
and its usage should be kept to a minimum for mass production [64]. Some of the synthetic
avenues employed to reduce the amount of incorporated platinum include core-shell catalyst
design [65,66] and platinum alloyed catalysts [67,68]; particularly PtCo alloyed catalysts
16
employed in present fuel cell vehicles. Other avenues utilize catalyst supports in order to maximize
the dispersion, surface area, stability, and synergistic effects of the catalyst being studied [69].
Scheme 1.1. Oxygen reduction reaction mechanism in: a) acidic media and; b) basic media.
Adapted from [70] and [71], respectively.
1.6.2: ORR in Basic Media
The ORR in basic media has been demonstrated to be achieved without the use of platinum
or noble metals, often times efficiently on solely carbon-based catalysts [35]. One of the proposed
pathways is shown in Scheme 1.1b. Similar to the ORR in acidic media, there are two main
a)
b)
17
products of the ORR in the alkaline media. The 2 e
-
pathway produces the peroxide anion which
can be detrimental to the fuel cell [72] while the 4 e
-
pathway produces the hydroxide which travels
across the AEM and reacts with the anode compartment reactants.
1.7: Overview of the Thesis
The main objective of this thesis is to explore various catalysts and their supports to be
used in a variety of fuel cells, mainly PEMFCs and AFCs. Chapter 2 covers the use of partially
fluorinated high surface area conductive carbon supports (CFx) as a support for both platinum and
manganese dioxide catalysts. The various effects of the fluorination was assessed and compared
to the widely-employed Vulcan Carbon (XC72) in characterization tests. The ORR kinetics was
assessed in half-cell and full fuel cell environments in both acidic and alkaline media. Chapter 3
covers the modification of reduced graphene oxide as both a catalyst and catalyst support
synthesized both by Hummer’s Method and electrochemical exfoliation. Reduced graphene oxide
both as a stand-alone catalyst and catalyst support for platinum was assessed for the ORR
properties in both acidic and alkaline media. Finally, Chapter 4 covers the modification of nickel
supports and doping for urea electrooxidation. CFx was studied as a support for nickel
nanoparticles and compared to Vulcan XC72 carbon support. The effect of annealing temperature
with nickel and nickel oxides on reduced graphene oxide was also be investigated through both
the half-cell and micro fuel cell testing. Finally, the doping effects on nickel from the third row
transition metal elements on urea electrooxidation was assessed.
18
1.8: References
[1] S. Bilgen, Structure and environmental impact of global energy consumption, Renew.
Sustain. Energy Rev. 38 (2014) 890–902. doi:10.1016/j.rser.2014.07.004.
[2] A. Boudghene Stambouli, E. Traversa, Fuel cells, an alternative to standard sources of
energy, Renew. Sustain. Energy Rev. 6 (2002) 295–304. doi:10.1016/S1364-
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24
Chapter 2:
CFx as a Catalyst Support for Oxygen Reduction in Acidic and Alklaine Media
2.1: High Surface Area Carbon-based Supports and Catalyst Preparation
For the implementation of catalysts for use in fuel cells, the catalyst support can be just as
critical for fuel cell performance optimization as the catalyst itself. Various types of carbon
supports for metal catalysts include undoped graphene [1–4], heteroatom (nitrogen, sulfur,
phosphorous) doped graphene [5–8], carbon nanotubes [9–12], and carbon nanofibers [13,14].
One of the most widely employed catalyst supports in use is carbon blacks. These types of carbon
supports are high in surface area and electrically conductive. Some of the various types of carbon
blacks that are being manufactured include: Denka Black, Black Pearls 2000, and Ketjen Black
[15,16]. One of the most widely used carbon blacks is XC72, or more widely known as “Vulcan
Carbon”. This type of carbon black is synthesized by pyrolysis in an oil furnace, where a minute
amount of air is supplied to the carbon starting material and annealed, usually at 1400
O
C [17].
This method is relatively cheap and can be reliably conducted in large quantities. This carbon
support consists of nanoscale electrically conductive carbon particles in the size range of around
30 nm to 60 nm giving it a high surface area and sufficient pore size for the catalysts to adhere
onto [18]. Vulcan Carbon and other carbon blacks can also be used as stand-alone catalysts and
additives used in various applications such as batteries [19–21] and solar cells [22–24]. Recent
studies have involved doping the carbon blacks with p-block elements, most often nitrogen, to
25
improve the catalytic function towards oxygen evolution, oxygen reduction, water splitting, and
various other catalytic processes [25–29].
Partially fluorinated high surface area electrically conductive carbon, or CFx, is a similar
species to the carbon blacks. In most fluorination procedures, especially in the one used in this
thesis, the degree of fluorination is around 10%-wt. The fluorination is carried out by plasma
annealing fluorine gas into the nano-sized carbon structure. This loading percentage can be
controlled and tailored to control the specific qualities being researched. There have been
numerous studies on CFx as a cathode material for batteries [30–32]. The fluorinated carbon has
been demonstrated to lead to higher discharge capacities in the cells [33,34]. The first accounts of
fluorinating carbon blacks for stand-alone catalysts for ORR in acidic and alkaline media was
carried out by Sun et al. with both Vulcan Carbon and Black Pearls [28,29]. Other groups have
since doped other carbon black supports with heteroatoms such as nitrogen and fluorine [5,35,36]
and other halogen species [27,37]. While these metal-free catalysts showed promise, they still
pale in comparison to the platinum and other noble metal-based catalysts. Prior research has
implemented fluorinated carbon supports with polytetrafluoroethylene (PTFE) in the gas diffusion
and catalyst layers of the cathode, which resulted in higher performance from the decrease in
flooding and better water management [38].
Another technique in catalyst preparation used to enhance performance is through the
reduction technique of the metal salts onto the support. Some of these methods include
impregnation with a reducing agent such as sodium borohydride [39–41] or hydrazine [42–44],
electrochemical deposition [45,46], and UV irridation [47–49]. Another reduction method that is
facile and efficient is the microwave-assisted polyol reduction method [50]. Metal salts are
dissolved in various polyols, mainly ethylene glycol, and reduced at high temperatures (>150
O
C)
26
and pressures (>10 bar) under microwave irridation [51,52]. In this method, the polyol solution
acts as both a solvent and reducing agent ensuring a homogenous reduction leading to a better
overall dispersion of the metal nanoparticles [53,54], preventing nanoparticle agglomeration on
the catalyst support [55]. This can yield high surface areas of the catalyst and smaller particle
sizes for increased activity [56,57].
The catalyst preparation conditions can also effect the phase of various transition row metal
catalyst, which in turn effects the structure and catalytic properties. Previous studies have shown
manganese dioxide to be a sufficient catalyst for oxygen reduction in alkaline media [58–60].
Manganese dioxide is one such catalyst in which the precursors and preparation can greatly affect
the overall phase and morphology [61,62]. The varying phases have been demonstrated to have
different catalytic properties towards oxygen reduction [63–65]. Furthermore, the morphology of
the various manganese dioxide catalysts can affect the ORR activity in the various phases whether
it is a nanoparticle [66], nanowire [67], or other morphology. Previous studies gave varying results
and found manganese dioxide phases to give superior activity [68,69]. Various carbon-based
supports have also been shown to enhance the catalytic activity by increasing the surface area and
active sites of the manganese dioxide [70].
In this chapter, commercially manufactured CFx was employed as a carbon support for
platinum and manganese dioxide catalysts and compared with the widely-used commercial XC72
to study the effects of the fluorination of the support on the ORR properties of the catalyst. Various
characterization methods were employed to assess the changes in catalyst morphology and ORR
performance. The platinum catalysts were tested in acidic media in both half-cell and a hydrogen-
based PEM fuel cell, while the manganese dioxide catalysts were tested in alkaline media in both
half-cell and in an alkaline DMFC.
27
2.2: The Effect of Fluorination of Carbon Catalyst Support on ORR Kinetics in Acidic Media
2.2.1: Experimental Methods
Both CFx and the widely used Vulcan Carbon (XC72, Cabot) were utilized as the carbon
support for the platinum catalyst. The platinum salts were deposited onto the carbon support by
means of impregnation in different loading amounts in order to assess the effect of support
percentage as well as fluorination of the carbon support on the activity of the platinum towards
oxygen reduction catalytics. All of the water used in the syntheses was of ultra-high purity
(Millipore, Direct-Q UV, 18.2 MΩ). An appropriate amount of XC72 and CFx were dissolved in
aqueous solutions by sonication and vigorous stirring. It is important to note that due to the
hydrophobic nature of the fluorination of the carbon, the CFx was harder to homogenize in the
aqueous solution and thus had to undergo additional cycles of sonication and stirring compared to
the XC72. Once a homogenized aqueous solution of carbon support was obtained, a certain
amount of aqueous solution of the platinum salts (H2PtCl6, 2g/100mL, Alfa Aeasar) was added to
the aqueous carbon solution. The pH was then adjusted from around 2 from the addition of the
acidic platinum salt solution to 9 from the addition of 1.0 M sodium hydroxide solution. The
solution was then heated in an oil bath to 80
O
C. Due to its relative instability in water, a freshly
made solution of sodium borohydride was made and quickly added to the solution with a molar
excess of around ten times the amount needed to reduce the platinum. The catalyst solutions was
stirred at 80
O
C for 1 hour after the addition and then at room temperature overnight. The catalyst
solutions were then washed by Millipore water and centrifuged at 6000 rpm. This process was
28
repeated until the pH of the supernatant had reached a value of 7. The catalysts were then dried in
an oven at 65
O
C overnight.
The catalysts were characterized by a variety of methods. The transmission electron
microscope (TEM) images were taken on a JEOL JEM-2100F with an acceleration voltage of 200
keV. The catalyst particle size was analyzed from the TEM images from the ImageJ software.
The scanning electron microscope (SEM) images were taken on a JEOL JSM-7001F with an
acceleration voltage of 20 kV. Thermogravimetric analysis (TGA) was performed on a TGA-50
Thermogravimetric Analyzer (Shimadzu). The samples were heated up at a rate of 10
o
C per
minute. X-ray diffraction (XRD) patterns were taken on a Rigaku Ultima Diffractometer with a
Cu-Kα (0.154056 nm) radiation source from a 2θ value of 15
O
to 75
O
. X-Ray Photoelectron
Spectroscopy (XPS) measurements were obtained on a Kratos Axis Ultra with a Monochromatic
X-Ray radiation source.
The half-cell electrochemical testing was carried out in a three cell rotating disk electrode
(RDE) setup. The setup comprised of a Teflon wrapped glassy carbon RDE (0.195 cm
2
working
area) as the working electrode, a platinum wire as the counter electrode, and a mercury sulfate
electrode (MSE) (0.5 M H2SO4 filling soln.) as the reference electrode. 1 mg of the catalyst was
added to 1 mL of a Nafion® binder solution comprising of 10 mg of liquid Nafion® ionomer
solution (5% wt in isopropanol, Alfa Aesar), 90% Millipore water, and 10% isopropanol to make
a catalyst ink solution. The catalyst ink was mixed via both sonication and vortex mixing until the
catalyst was homogenously dispersed. Then, 20 μL of ink was micropipetted onto the RDE and
dried in an oven at 65
O
C. The electrode was then connected to the three cell setup and all the tests
were performed in 0.5 M H2SO4 solution. Before the tests, the cell was purged with argon (ultra-
high pure grade, 99.999%) for 15 minutes and the cell was cycled from 0.0 V to 1.28 V vs. RHE
29
to remove any impurities from the catalyst. For the CO stripping tests, CO (CP Grade 99.5%) was
bubbled in for 15 minutes followed by argon for 30 minutes. For the ORR testing, the cell was
purged with oxygen (ultra-high pure grade, 99.994%) for 15 minutes and the cell was cycled from
0.0 V to 1.28 V vs. RHE. The half-cell tests were measured in triplicate.
For the fuel cell tests, a separate catalyst ink solution was made for the 25%-wt platinum
on the carbon supports in a 1:3:5 ratio of catalyst: Millipore water: 5% Nafion® ionomer solution.
This solution was sonicated for 8 minutes and then manually painted on a 5 cm
2
piece of Toray
Carbon paper (TGH-060, 10% wet proofing) and dried in an oven at 110
O
C. The platinum
loadings for both the XC72 and CFx supports were controlled and set to two different loadings for
each: 0.1 mg cm
-2
and 0.5 mg cm
-2
. This served as the cathode electrode while the anode electrode
was made from platinum black inks with the same composition as previously mentioned and
painted on the same type of carbon paper with loadings of 2.0 mg cm
-2
to ensure the limiting
reaction was the ORR on the cathode and dried in the same oven. The membrane electrode
assembly (MEA) was then made by sandwiching a protonated Nafion®-211 membrane between
the electrodes and pressing at 0.5 tons of force at 140
O
C for 5 minutes. The MEA was then put
in a fuel cell system composed of conducting graphite plates and brass current collectors and
hydrated with Millipore water at 60
O
C overnight before testing. The MEA was tested by flowing
humidified hydrogen (ultra-high pure grade, 99.999%) through the anode at 150 mL min
-1
and
humidified oxygen (ultra-high pure grade, 99.994%) through the cathode at 200 mL min
-1
.
Polarization curves were taken by scanning the potential at 50 mV every 30 seconds from open
circuit voltage (OCV) to 0.1 V. Electrochemical impedance spectroscopy (EIS) measurements
were also taken at 0.4 V vs. reference with an AC amplitude of 10 mV.
30
2.2.2: Results
In order to quantify the platinum loading on each of the catalysts, TGA curves were
performed and shown in Figures 2.1a&b. The percent weights were calculated and subtracted
from the corresponding CFx and XC72 “blanks” which were also run. Table 2.1 shows the loading
targets with the corresponding actual loadings from the TGA curves. The weight loss from both
the XC72 supported “Pt-C” and CFx supported “Pt-F” catalysts starts to occur at around the same
temperature signifying the similar stabilities between the CFx and XC72 supported platinum
catalysts at high temperatures. Due to the hydrophobic nature of the CFx, it was harder to disperse
in the aqueous solutions. This most likely contributes to the higher platinum loadings and greater
deviations from the targets when compared to the XC72 catalysts.
0%
20%
40%
60%
80%
100%
120%
0 100 200 300 400 500 600 700 800
Weight (% of Initial)
Temperature (
o
C)
Pt-1F
Pt-2F
Pt-3F
Pt-4F
CFx
a)
31
Figure 2.1. TGA curves of the Pt/Ccatalysts with the: a) CFx supports and; b) XC72 supports.
Table 2.1. Summary of the TGA results for the Platinum loadings of the Pt/C catalysts.
Catalyst Pt-1F Pt-1C Pt-2F Pt-2C Pt-3F Pt-3C Pt-4F Pt-4C
Target Loading 10% 10% 25% 25% 50% 50% 75% 75%
Actual Loading 15.9% 10.3% 30.0% 30.6% 68.4% 49.9% 83.2% 73.8%
Figures 2.2a&b shows the SEM images of the Pt-2C and Pt-2F catalysts that were studied
in the fuel cell tests. The catalysts have similar morphologies with no macroscopic differences
between them. The other synthesized catalysts show very similar images. This shows that the
fluorination of the carbon support does not affect the overall macroscopic differences nor does the
amount of platinum loading on the overall structure. The particle size of the XC72 and CFx
supports is also visible showing about a 30 nm diameter. In order to effectively analyze the
platinum particle size, TEM images were taken and shown in Figures 2.3a-d and Figures 2.4a-d
0%
20%
40%
60%
80%
100%
120%
0 100 200 300 400 500 600 700 800
Weight (% of Initial)
Temperature (
o
C)
Pt-1C
Pt-2C
Pt-3C
Pt-4C
XC72
b)
32
with the corresponding platinum particle size distributions for the Pt-F and the Pt-C catalysts,
respectively. The images show a clearer representation of the platinum particles on the carbon
supports with the sizes and morphologies visible. There is minimal apparent difference in the
overall platinum particle morphologies between the CFx and XC72 supports at each of the
respective loadings indicating the fluorination of the support again does not affect the overall
platinum particle morphology. The average particle size and distributions show a smaller platinum
particle size for the CFx support at each of the catalyst loadings. This could be due to the better
dispersion of the platinum salts in the aqueous solution before the sodium borohydride reduction
due to the hydrophobic nature of the fluorination in the support. The fluorination of the CFx could
possibly act as a dispersant to keep the platinum nanoparticles from agglomerating during the
reduction step. There is no apparent trend in platinum particle size for either support with respect
to the platinum loading.
Figure 2.2. SEM images of the Pt/C catalysts: a) Pt-2C and; b) Pt-2F.
b)
a)
33
0
5
10
15
20
25
30
1 2 3 4 5 6 7 8 9 10
Frequency (%)
Pt Particle Size (nm)
a)
3.712±1.832 nm
0
5
10
15
20
25
1 2 3 4 5 6 7 8 9 10
Frequency (%)
Pt Particle Size (nm)
b)
3.58±1.322 nm
a)
b)
34
Figure 2.3. TEM images of the Pt/C catalysts with the particle size distribution charts of: a) Pt-
1F; b) Pt-2F; c) Pt-3F and; d) Pt-4F. The measuring bar indicates 20 nm.
0
5
10
15
20
25
30
35
1 2 3 4 5 6 7 8 9 10
Frequency (%)
Pt Particle Size (nm)
c)
4.127±1.256 nm
0
5
10
15
20
25
30
1 2 3 4 5 6 7 8 9 10
Frequency (%)
Pt Particle Size (nm)
d)
3.798±1.493 nm
c)
d)
35
0
5
10
15
20
25
1 2 3 4 5 6 7 8 9 10
Frequency (%)
Pt Particle Size (nm)
a)
4.640±1.626 nm
0
5
10
15
20
25
30
35
1 2 3 4 5 6 7 8 9 10
Frequency (%)
Pt Particle Size (nm)
b)
3.654±1.222 nm
a)
b)
36
Figure 2.4. TEM images of the Pt/C catalysts (top) with the particle size distribution charts
(bottom): a) Pt-1C; b) Pt-2C; c) Pt-3C and; d) Pt-4C. The measuring bar indicates 20 nm.
In order to gauge the crystallographic nature of both the carbon support and platinum
nanoparticles, XRD patterns were taken and shown in Figures 2.5a&b. Each of the patterns display
peaks at 40
O
, 46
O
, and 68
O
signifying the (111), (200), and (220) Miller Indices of fcc platinum
[71]. There is also a peak at 22
O
signifying the (002) phase of the carbon support. As the platinum
loadings increase, the carbon peak height decreases while the platinum peak heights increase due
0
5
10
15
20
25
30
1 2 3 4 5 6 7 8 9 10
Frequency (%)
Pt Particle Size (nm)
c)
4.241±1.650 nm
0
5
10
15
20
25
1 2 3 4 5 6 7 8 9 10
Frequency (%)
Pt Particle Size (nm)
d)
5.076±1.856 nm
d)
c)
37
to higher ratios of platinum particles to the carbon supports at higher loadings. The fluorination
of the carbon support has minimal effect on the overall platinum crystallographic orientation.
Figure 2.5. XRD patterns of the Pt/C catalysts on the: a) CFx support and; b) XC72 support.
15 25 35 45 55 65 75
Intensity (a.u.)
2 θ (degrees)
CFx
Pt-1F
Pt-2F
Pt-3F
Pt-4F
(111)
(200) (220)
(002)
a)
15 25 35 45 55 65 75
Intensity (a.u.)
2 θ (degrees)
XC72
Pt-1C
Pt-2C
Pt-3C
Pt-4C
(111)
(200) (220)
(002)
b)
38
The surface properties of the catalysts were analyzed with XPS. The Pt 4f of the Pt-2
catalysts are shown in Figure 2.6a. Two main peaks are visible in the Pt 4f spectra at around 72
eV and 69 eV, which agree with the ones previously reported in the literature [72]. These were
further deconvoluted into four peaks signifying the Pt(O) 7/2 and 5/2 peaks and the Pt(II) 7/2 and
5/2 peaks. The Pt(II) peak arises from the chemadsorption of the oxygen on the surface [45].
There is minimal difference evident in the Pt 4f peaks signifying the fluorination of the carbon
support has minimal effects on the platinum surface binding. The C 1s spectra are shown in Figure
2.6b. There is one main peak visible that was deconvoluted into peaks representing the C=C, C-
C, C-O, C=O, and -COOH binding with an extra C-F peak for the fluorine binding [73] in the CFx
support with minimal difference in the other types of peak binding.
Figure 2.6. XPS spectra of the Pt-2 catalysts: a) the Pt 4f high resolution region and; b) the C 1s
high resolution region. The (1) spectrum indicates the XC72 supported catalysts; the (2)
spectrum indicates the CFx supported catalysts.
The intrinsic catalytic properties of the catalysts were then determined via half-cell testing
in 0.5 M H2SO4. Figure 2.7 shows the characteristic CV scans of the catalysts under argon. The
65 70 75 80
Intensity (a.u.)
Binding Energy (eV)
Raw
Pt(O) 4f 7/2
Pt(O) 4f 5/2
Pt(II) 4f 7/2
Pt(II) 4f 5/2
Envelope
a)
(2)
275 280 285 290 295
Intensity (a.u.)
Binding Energy (eV)
Raw
C=C
C-C
C-O
C=O
C-F
-COOH
Envelope
b)
(1)
(2)
(1)
39
hydroxide adsorption region in 1.0 V to 1.28 V, the oxygen reduction region in 0.5 V to 0.9 V, and
the hydrogen adsorption/desorption region in 0.0 V to 0.3 V are found in all of the catalysts and
are indicative of a CV of fcc platinum. The catalysts with the XC72 support display an increase
in catalytic activity with the increase in platinum loadings. This trend has been well documented
in previous studies [74]. There are a few methods employed to calculate the electrochemically
active surface areas (ECSAs) of the catalysts. One such way is from the charge in the hydrogen
adsorption region from the area under the curve. The equation for the ECSA from hydrogen
adsorption is:
ECSA = QPt / ΓH-adsL (1)
Where QPt is the integration of the charge density over the hydrogen adsorption area measured,
ΓH-ads is the charge required to strip a monolayer of protons from the platinum surface (210 μC cm
-
2
Pt) and L is the platinum loading in the sample measured (gPt cm
-2
electrode) [75]. Another
calculation of the ECSA is from the CO stripping of the catalyst. Figure 2.8 shows the CO
stripping CVs of the catalysts. The ECSA is calculated by taking the area under the CO stripping
curve of the poisoned catalysts from the first scan and subtracting the area under the second scan.
The ECSA is then calculated by the equation:
ECSA = QPt / ΓCOL (2)
Where QPt is the integration of the charge density under the CO stripping peak, ΓCO is the charge
required to oxidize a monolayer of CO from the platinum surface (420 μC cm
-2
Pt) and L is the
platinum loading in the sample measured (gPt cm
-2
electrode) [76]. For both types of catalyst
supports, the ECSA increases as the platinum loadings increase. The normalized ECSAs of both
the hydrogen adsorption and CO stripping for the catalysts is shown in Figure 2.9. In most of the
determinations, the hydrogen adsorption ECSA was calculated to be higher than the CO stripping.
40
This is most likely due to the fact that the CO does not for a uniform monolayer on the platinum
surface whereas hydrogen does [77]. For both of the supports, the normalized ECSA decreases
for both types calculations as the platinum loadings for the catalysts increases. The normalized
ECSAs of the CFx supported catalysts are higher than those of the XC72 supported catalysts at the
lower loadings and but lower at the higher loadings. This is most likely due to the increased
agglomeration of platinum particles on the CFx at higher loadings due to the hydrophobic nature
of the doped fluorine compared to that of the XC72. There is enhanced dispersion for the platinum
nanoparticles on the CFx at the lower loadings which is evident in the TEM images. At the higher
platinum loadings, the ECSA values tend to converge in both the platinum support catalysts and
ECSA determination method. The size and morphology of the platinum nanoparticles can greatly
affect the ECSA of the catalysts. Previous studies have shown the optimal platinum nanoparticle
diameter for ORR is between 2-4 nm [78,79]. Most of the CFx supported catalysts synthesized
have size distributions in this range while only the Pt-2C and Pt-3C catalysts fall in this range.
Another important aspect worth discussing in the graph is the onset potential of CO oxidation. The
overall trend from the onset potential of the CO stripping is the CFx supported platinum has a
lower onset potential of oxidation suggesting the fluorine might have an effect to help mitigate the
backbonding effects of CO on platinum.
One aspect worth consideration pertaining to the catalyst inks in the electrochemical tests
is that whereas the XC72 powder had better dispersion in aqueous solution during the impregnation
process, the CFx supported catalysts had better dispersion in the catalyst ink solutions for the
electrochemical tests. This is due to the interactions with the fluorine present in CFx with the
tetrafluoroethylene backbone in Nafion®. The CFx was considerably hydrophobic and would not
disperse well in the aqueous solutions for reduction.
41
Figure 2.7. CV scans of the Pt/C catalysts in 0.5 M H2SO4 solution under argon at 20 mV s
-1
.
The solid lines indicate the CFx support; the dashed lines indicate the XC72 support.
Figure 2.8. CO stripping CVs for the Pt/C catalysts. The solid lines indicate the CFx support;
the dashed lines indicate the XC72 support.
-2
-1.5
-1
-0.5
0
0.5
1
1.5
0 0.2 0.4 0.6 0.8 1 1.2 1.4
Current Density (mA cm
-2
)
Potential (V vs. RHE)
Pt-1
Pt-2
Pt-3
Pt-4
-2
-1
0
1
2
3
4
5
6
0 0.2 0.4 0.6 0.8 1 1.2 1.4
Current Density (mA cm
-2
)
Potential (V vs. RHE)
Pt-1
Pt-2
Pt-3
Pt-4
42
Figure 2.9. The ECSAs of the Pt/C catalysts for the hydrogen adsorption (solid line) and CO
stripping (dashed line).
The ORR catalytic properties of the catalysts were investigated mainly with linear sweep
voltammetries (LSVs) found in Figure 2.10a for the catalysts at a rotation rate of 1600 RPM. A
more accurate view of the onset potentials is found in Figure 2.10b. A summary of the onset
potential and limiting current densities the LSV scans are shown in Figure 2.11. There is some
variation in the onset potential for the LSVs in all of the catalysts loadings between the two types
of carbon supports. This signifies that the CFx support has some synergistic effects on the intrinsic
ORR activities of the platinum catalyst compared to that of the XC72. There is, however, a
difference in the limiting current density trends between the catalyst supports. The limiting current
density of the XC72 supported catalysts increases as the platinum loading increases. This stems
from the overall increase in particle size and agglomerations reported in previous studies [80,81].
Other studies have also shown in higher loadings of platinum on carbon (>60%) that the smaller
particle sizes have greater mass and specific activity for ORR [74]. The limiting current densities
for the CFx supported catalysts, however, decrease overall as the platinum loading increases. This
could possibly stem from the enhanced oxygen diffusion properties of the CFx in comparison. At
0
30
60
90
120
150
180
0% 20% 40% 60% 80% 100%
ECSA (m
2
g
-1
Pt)
Pt Loading (%-wt)
CFx
XC-72R
43
the lower platinum catalyst loadings, there is an overall higher amount of CFx for the support in
comparison for enhanced oxygen diffusion to the platinum nanoparticles, whereas at the 75%
loading, the synergistic effects of the CFx is considerably reduced.
LSV scans were also performed by varying rotation rates of 400, 800, 1200, 1600, 2000,
and 2400 RPM. One such plot is shown in Figure 2.12 for the Pt-1F catalyst. Figure 2.13 shows
a Koutecky-Levich plot created from plotting the negative inverse limiting current against the
rotation rate raised to the negative one-half power from the current density values in Figure 2.12
at 0.6 V. From the slopes of the lines, the Koutecky-Levich equation was used to calculate the
electron number [71]:
2 / 1 6 / 1 3 / 2
2 2
62 . 0
1 1 1 1 1
b
O O k d k
c v nFAD I I I I
−
+ = + = (3)
where I is the measured current, Ik is the kinetic limitation current, Id is the diffusion limitation
current, n is the electron transfer number, F is the Faraday constant, A is the electrode area, DO2 is
the diffusion coefficient of oxygen in the bulk phase, v is the kinematic viscosity of the solution,
cO2 is the concentration of oxygen in the bulk phase, and is the rotation rate in rad s
-1
. The
average electron transfer numbers from the plot were calculated and shown in Table 2.2. Overall,
the average electron transfer numbers of the CFx supported catalysts were higher than the XC72
supported catalysts (except for Pt-4) but still signifies a general 4 electron transfer process for each
catalyst.
44
Figure 2.10. The LSV scans of the Pt/C catalysts at 1600 RPM under O2 at a) normal scan
window and; b) zoom-in of the onset potential region. The solid lines indicate the CFx
supported catalysts; the dashed lines indicate the XC72 supported catalysts.
-5
-4
-3
-2
-1
0
0 0.2 0.4 0.6 0.8 1 1.2
Current Density (mA cm
-2
)
Potential (V vs. RHE)
Pt-1
Pt-2
Pt-3
Pt-4
-0.5
-0.4
-0.3
-0.2
-0.1
0
0.8 0.84 0.88 0.92 0.96 1
Current Density (mA cm
-2
)
Potential (V vs. RHE)
Pt-1
Pt-2
Pt-3
Pt-4
b)
a)
45
Figure 2.11. Summary of the onset potentials of the Pt/C catalysts from Figure 2.10.
Figure 2.12. Typical series of LSV scans for the Pt1-F catalyst.
0.8
0.83
0.86
0.89
0.92
0.95
0 20 40 60 80 100
Onset Potential (V vs. RHE)
Pt Loading (%-wt)
CFx
XC-72
-6
-5
-4
-3
-2
-1
0
0 0.2 0.4 0.6 0.8 1 1.2
Current Density (mA cm
-2
)
Potential (V vs. RHE)
400 RPM
800 RPM
1200 RPM
1600 RPM
2000 RPM
2400 RPM
46
Figure 2.13. Koutecky-Levich plot of the current of the Pt/C catalysts at 0.6 V. The solid
markers indicate the CFx support; the hollow markers indicate the XC72 support.
Tafel plots are ways of showing the intrinsic activity of a catalyst at low current densities
close to the OCV. Figure 2.14 shows the Tafel plots of the Pt/C catalysts with the Tafel Slopes
reported in Table 2.2. For the CFx supported catalysts, there is an increase in the Tafel slope as
the platinum loading increases. This is due to the larger potential drop through the pores which
increases the Tafel slope as was reported in previous studies [82]. This trend is also displayed in
the XC72 supported catalyst runs except for the Pt-1C catalyst. The Tafel slopes at each respective
loading for the XC72 catalysts are lower than the CFx supported catalysts suggesting that the CFx
has a slight detrimental effect on the intrinsic oxygen reduction capabilities of the platinum at the
low current densities [83]. The slopes for each of the catalysts are close to the measured values of
-60 mV dec
-1
for platinum in acidic media as has been reported numerous times before [84,85].
Another aspect of the Tafel plots worth examining is the overpotential of the system. The
overpotentials of the CFx supported catalysts increase with increased platinum loadings again
1000
1500
2000
2500
3000
3500
4000
0 0.05 0.1 0.15 0.2
Negative Inverse Current (-A
-1
)
ω
-½
Pt-1
Pt-2
Pt-3
Pt-4
47
suggesting that the fluorination of the carbon support could have positive effects on the ORR
activity in low current regions.
Figure 2.14. Tafel plot for the Pt/C catalysts. The solid markers indicate the CFx support; the
hollow markers indicate the XC72 support.
Table 2.2. Summary of the RDE electrochemical tests of the Pt/C catalysts.
The ORR activity was further measured in fuel cells with the polarization curves shown in
Figures 2.15 for the cells tested at ambient temperature (~21
O
C) and Figure 2.16 for the cells
0.7
0.75
0.8
0.85
0.9
0.01 0.1 1 10
Potential (V vs. RHE)
Current Density (mA cm
-2
)
Pt-1
Pt-2
Pt-3
Pt-4
Catalyst Onset Potential
(V vs RHE)
Limiting Current
Density (mA cm
-2
)
Electron
Trans
Tafel Slope
(mV dec
-1
)
Pt-1F 0.880.02 4.490.09 3.970.02 -774
Pt-1C 0.880.03 4.170.54 3.840.08 -834
Pt-2F 0.900.03 4.360.13 3.950.04 -814
Pt-2C 0.880.01 4.310.23 3.790.03 -7510
Pt-3F 0.890.03 4.320.39 3.970.01 -887
Pt-3C 0.890.02 4.100.05 3.630.05 -784
Pt-4F 0.860.01 3.890.12 3.580.21 -916
Pt-4C 0.900.01 4.390.21 3.850.01 -839
48
tested at 50
O
C with the power densities and OCV values reported in Table 2.3. The MEAs with
the fluorinated carbon supported catalysts showed an overall enhancement in peformance at both
the 0.1 mg cm
-2
and 0.5 mg cm
-2
loadings and at ambient temperatures and 50
O
C. More
specifically, at the 0.5 mg cm
-2
loadings, the CFx supported catalyst MEAs shows a 23% and 18%
enhancement in peak power density over the XC72 supported catalyst at ambient temperature and
50
O
C, respectively. This could most likely be due to the enhanced diffusion of oxygen to the
catalyst as well as better water management leading to less flooding from both the fluorinated
support and the teflonied gas diffusion layer (GDL). One of the outlying factors to the
enhancement of the power density from the CFx support is the lower power density compared to
the XC-27R support in the 0.1 mg cm
-1
loadings at ambient temperature. A possible explaination
for this could be the considerably low amount of catalyst and support in the catalyst layers to
effectively promote the oxygen diffusion and water alleviation at the cathode under the ambient
temperatures. This could have possibly caused some flooding in the cathode in comparison to the
higher loadings of CFx supported catalyst leading to lower performances compared to the XC72
supported catalysts. The OCV is another part of the polarization curve that has a considerable
effect on overall performance with the results reported in Table 2.3. At the 0.1 mg cm
-2
loadings,
the CFx supported catalyst MEAs display a 98 mV and 82 mV increase in OCV compared to the
XC-72R supported catalysts at ambient temperature and 50
O
C, respectively, whereas, at the 0.5
mg cm
-2
loadings, the OCVs of both MEAs are about equal.
Another effective way of probing the reactions in the fuel cells is through EIS. EIS spectra
were taken of the MEAs at 0.4 V and shown in Figure 2.17 and Figure 2.18 corresponding to the
cells tested at ambient temperature and 50
O
C, respectively. An equivalent circuit was constructed
based on the overall shape and particular processes involved and shown in Figure 2.19. All of the
49
curves resemble depressed semicircles in which L1 and L2 represent the inductance of the system
arising from the wires and equipment, Rmem and Rct represent the membrane and charge transfer
resistances respectively, CPE represents the pseudo-capacitive nature of the catalyst layers, and
Ws represents the Warburg element for the finite diffusion of reactants and products through the
porous catalyst layers. From the equivalent circuit, the charge transfer resistances were calculated
and are reported in Table 2.3. The charge transfer resistances are lower for the catalysts at the 0.5
mg cm
-2
loadings due to the higher current flowing in the fuel cells from the same cell potentials
at which the MEAs were tested at. The charge transfer resistances of the MEAs at ambient
temperature are also higher than those at 50
O
C from the faster kinetics at elevated temperatures
and relatively lower cell resistances. CFx has been shown to display relatively high conductivity
even though there is saturated C-F bonds in the carbon support [86]. The CFx supported catalysts
display a 11% and 33% decrease in charge transfer resistance at 0.5 mg cm
-2
and 0.1 mg cm
-2
loadings at 50
O
C, respectively. The charge transfer resistance is also lower for the MEAs with
the CFx support than the XC72 support. This is most likely from the increased oxygen diffusion
and hydrophobicity of the fluorocarbon surface [87]. Furthermore, there could be better
synergistic connection between the fluorine backbone of the CFx and the Nafion® ionomer
solution ink creating an enhance network to facilitate charge transfer. Another aspect worth
mentioning is the high frequency x-intercept, which is indicative of the membrane resistance, Rmem,
of the MEAs. The Rmem of the CFx MEAs is lower than that of the XC72 MEAs meaning higher
conductivity, possibly stemming from better water management at the membrane and catalyst
layers from the hydrophobicity of the CFx as well as the triple-phase boundary connection between
the catalyst, ionomer, and the membrane.
50
Figure 2.15. Polarization curves for the Pt/C catalysts at ambient temperature. The solid lines
indicate the 0.5 mg cm
-2
loadings; the dashed lines indicate the 0.1 mg cm
-2
loadings.
Figure 2.16. Polarization curves for the Pt/C catalysts at 50
O
C. The solid lines indicate the 0.5
mg cm
-2
loadings; the dashed lines indicate the 0.1 mg cm
-2
loadings.
0
100
200
300
400
500
600
0
0.2
0.4
0.6
0.8
1
1.2
0 500 1000 1500 2000 2500
Power Density (mW cm
-2
)
Potential (V)
Current Density (mA cm
-2
)
Pt-2C
Pt-2F
0
200
400
600
800
1000
0
0.2
0.4
0.6
0.8
1
0 500 1000 1500 2000 2500 3000
Power Density (mW cm
-2
)
Potential (V)
Current Density (mA cm
-2
)
Pt-2C
Pt-2F
51
Figure 2.17. EIS curves for the Pt/C catalysts at ambient temperature. The solid markers
indicate the 0.5 mg cm
-2
loadings; the hollow lines indicate the 0.1 mg cm
-2
loadings; the lines
indicate the fitted circuit.
Figure 2.18. EIS curves for the Pt/C catalysts at 50
O
C. The solid markers indicate the 0.5 mg
cm
-2
loadings; the hollow lines indicate the 0.1 mg cm
-2
loadings; the lines indicate the fitted
circuit.
-0.2
-0.1
0
0.1
0.2
0 0.1 0.2 0.3 0.4 0.5 0.6
-Z" (Ohms)
Z' (Ohms)
Pt-2C
Pt-2F
-0.2
-0.1
0
0.1
0.2
0 0.1 0.2 0.3 0.4 0.5
-Z" (Ohms)
Z' (Ohms)
Pt-2C
Pt-2F
52
Figure 2.19. The equivalent circuit for the EIS spectra.
Table 2.3. Summary of the fuel cell tests of the Pt-2F and Pt-2C catalysts.
Catalyst MEA OCV (V) Peak Power
Density (mW cm
-2
)
Charge Transfer
Resistance (mΩ)
Ambient 50
O
C Ambient 50
O
C Ambient 50
O
C
Pt-2C (0.1 mg cm
-2
) 0.894 0.903 306 382 411 284
Pt-2F (0.1 mg cm
-2
) 0.992 0.985 264 457 282 256
Pt-2C (0.5 mg cm
-2
) 1.01 1.01 415 737 152 110
Pt-2F (0.5 mg cm
-2
) 1.00 1.00 510 868 107 83
2.3: Effect of Microwave-Assisted Polyol Reduction of Fluorinated Carbon Supported Platinum
Catalysts for ORR in Acidic Media
2.3.1: Experimental Methods
In this section, both CFx and XC72 were again used as the carbon supports for the platinum
catalysts. The platinum salts were deposited onto the support by means of a microwave-assisted
polyol reduction. The carbon supports were first homogenously dispersed in a 10 mL ethylene
glycol solution in a 20 mL microwave vial (Biotage) by means of sonication and vigorous stirring
for 1 hour. The CFx more readily dissolved into ethylene glycol than Millipore water in this case.
The same platinum salt solution as described in Section 2.2.1 was added to the solution to make a
20%-wt platinum solution. The pH was adjusted with 1.0 M aqueous sodium hydroxide solution
L1 Rmem Rct
CPE W1
L2
53
to a value of 7. The solution was again stirred and sonicated to ensure homogeneity and the vial
was tightly sealed and placed in a microwave reactor (Biotage). The solution was heated to 180
O
C for 3 minutes via microwave irradiation with a working power of 55 W. The solution was then
repeatedly washed with a 1:1 Millipore water:acetone solution and centrifuged until the pH of the
supernatant was 7 and dried in an oven at 65
O
C overnight. 20%-wt platinum on both CFx and
XC72 supports were also synthesized by impregnation in the same manner as described in the
Section 2.2.1 for comparison.
The characterization and half-cell electrochemical tests were carried out in the same
fashion as in Section 2.2.1 for both the microwave reduced and impregnated catalysts. Fuel cell
tests were also carried out in the same manner as described in Section 2.2.1 for both the microwave
reduced and impregnated catalysts. ECSA measurements of the cathode compartment of the
MEAs were carried out by flowing hydrogen through the anode compartment at 100 mL min
-1
and
Millipore water through the cathode compartment at 200 mL min
-1
with a cell temperature of 60
o
C. A CV scan was taken at 50 mV s
-1
from 0.0 to 1.0 V vs. RHE.
2.3.2: Results
Figure 2.20 shows the SEM images of the CFx supported catalysts. As discussed in Section
2.2.2, the SEM images show very similar catalyst morphologies for both the impregnation and
microwave reduction methods for the CFx supported catalysts. TEM images are shown in Figures
2.21a&b for the XC72 supported catalysts and in Figures 2.22a&b for the CFx supported catalysts
with the corresponding platinum particle size averages and distributions. This again shows a
clearer depiction of the catalyst particles on the carbon support. The particle size histograms show
54
the average platinum particle size and standard deviation. These particle sizes are consistent with
those previously reported [88]. The catalysts from the microwave reduction synthesis show
considerably smaller particle sizes than the ones from the impregnation synthesis. This is due to
the better overall dispersion of the platinum salts in the ethylene glycol and the homogenous
reduction from the ethylene glycol in the microwave reactor. In the microwave reduction catalysts,
the platinum particles reduced on the CFx support have a smaller average diameter than the ones
reduced on the XC72 support. This can be due to the better dispersion capabilities of the fluorine
in the carbon support in ethylene glycol.
Figure 2.20. SEM images of the CFx supported catalysts synthesized by a) impregnation and; b)
microwave reduction.
a)
b)
55
Figure 2.21. TEM images with the corresponding particle size distribution charts of the XC72
supported catalysts synthesized by: a) impregnation and; b) microwave reduction.
0
5
10
15
20
25
1 2 3 4 5 6 7 8 9 10 11
Frequency (%)
Pt Particle Size (nm)
a)
5.61±1.87 nm
0
5
10
15
20
25
30
35
40
1 2 3 4 5 6 7 8 9 10
Frequency (%)
Pt Particle Size (nm)
b)
2.82±1.27 nm
a)
b)
56
Figure 2.22. TEM images with the corresponding particle size distribution charts of the CFx
supported catalysts synthesized by: a) impregnation and; b) microwave reduction.
XRD patterns of the catalysts were taken are shown in Figure 2.23. Each pattern has peaks
at 2θ values of 39.9
O
, 46.6
O
, 68.2
O
, 81.9
O
, and 87.1
O
signifying the crystallographic indices of
(111), (200), (220), (311), and (222) for the fcc platinum particles. There is also a peak at 22
O
signifying the (002) index of the carbon support. The XRD patterns of the microwave reduced
catalysts exhibit peak broadening when compared to the impregnated catalysts signifying a smaller
0
5
10
15
20
25
30
35
40
1 2 3 4 5 6 7 8 9 10
Frequency (%)
Pt Particle Size (nm)
a)
4.21±1.43 nm
0
5
10
15
20
25
30
35
40
45
1 2 3 4 5 6 7 8 9 10
Frequency (%)
Pt Particle Size (nm)
b)
2.76±1.26 nm
a)
b)
57
particle size as was shown from the TEM images. The average particle sizes were calculated from
the Scherrer equation [89]:
τ = Kλ/βcosθ (4)
where τ is the average particle size, K is the dimensionless shape factor (usually 0.9), λ is the
wavelength of the incident radiation, and θ is the angle of the peak that is being studied. The
average crystallite sizes of the catalysts were determined from the (111) peak in the catalysts to
be: 5.7 nm and 4.8 nm for the microwave reduced XC72 and CFx catalysts, respectively and 8.3
nm and 8.0 nm for the impregnated XC72 and CFx, respectively. This decrease in crystalline size
is to be expected due to the smaller overall particle sizes of the microwave reduced compared with
the impregnated catalysts. The crystalline sizes for the CFx supports were also slightly smaller
than for the XC72 supports which can be due to the increased dispersion and decreased
agglomeration of the platinum nanoparticles as was confirmed in the TEM images. The graphite
peaks of the carbon supports are also less pronounced in the microwave reduction catalyst patterns
further signifying a greater coverage of platinum on the support from the homogenous reduction.
XPS spectra were also taken of the catalysts and shown in Figures 2.24a&b for the Pt 4f
spectra. In each spectra, the same two peaks are visible as was discussed in Section 2.2.2. Overall,
the two sets of the peaks are very similar signifying a minimal change to the intrinsic binding
properties of platinum in the catalysts from the two types of reduction. The area under the curve
for the Pt(II) 7/2 peak is larger for the microwave reduced catalysts than for the impregnated
catalysts. This could arise from the smaller particle sizes in the microwave reduced catalysts that
have a larger surface area and higher chemadsorption of oxygen on the platinum surface. The
same trend is seen for the CFx supported catalysts compared to the XC72 supported catalysts for
both the impregnated and microwave reduced catalysts. Again, the particle size of CFx is smaller
58
for both sets of catalysts than the XC72, leading to a larger surface area and more sites on the
platinum for the chemadsorption of oxygen. The C 1s spectra of the catalysts are also shown in
Figures 2.25a&b. One main peak is visible for each of the spectra which was deconvoluted in the
same manner as was described in Section 2.2.2. There are similar features to the spectra between
both reduction techniques showing minimal effects on the carbon support from the microwave
reduction compared to the impregnation.
Figure 2.23. XRD patterns of the μwave Pt/C catalysts.
20 30 40 50 60 70 80 90
Intensity (a.u.)
2 θ (degrees)
Pt/XC-72R uwave
Pt/CFx uwave
Pt/XC-72R impreg
Pt/CFx impreg
(002)
(111)
(200)
(220)
(331) (222)
59
Figure 2.24. XPS Pt 4f spectra of the a) uwave reduced catalysts and; b) sodium borohydride
reduced catalysts. The (1) spectrum indicates the XC72 supported catalysts; the (2) spectrum
indicates the CFx supported catalysts.
Figure 2.25. XPS C 1s spectra of the a) uwave reduced catalysts and; b) sodium borohydride
reduced catalysts. The (1) spectrum indicates the XC72 supported catalysts; the (2) spectrum
indicates the CFx supported catalysts.
Figure 2.26 shows the characteristic CV scans of the catalysts under argon at 20 mV s
-1
for
the half-cell electrochemical tests. The ECSAs of the catalysts were calculated from the hydrogen
adsorption curves in the same manner as in Section 2.2.2 and are shown in Figure 2.28. The
65 70 75 80
Intensity (a.u.)
Binding Energy (eV)
Raw
Pt O 4f 7/2
Pt II 4f 7/2
Pt O 4f 5/2
Pt II 4f 5/2
Envelope
(2)
65 70 75 80
Intensity (a.u.)
Binding Energy (eV)
Raw
Pt O 4f 7/2
Pt II 4f 7/2
Pt O 4f 5/2
Pt II 4f 5/2
Envelope
(1)
(2)
b)
275 280 285 290 295
Intensity (a.u.)
Binding Energy (eV)
Raw
C=C
C-C
C-O
C=O
C-F
-COOH
Envelope
(1)
(2)
a)
275 280 285 290 295
Intensity (a.u.)
Binding Energy (eV)
Raw
C=C
C-C
C-O
C=O
C-F
-COOH
Envelope
(1)
(2)
b)
(1)
a)
60
microwave reduced catalysts display a much larger ECSA than the impregnated catalysts. This is
due to the smaller, more highly dispersed platinum particles as shown in the TEM images. The
CFx supported catalysts also show higher ECSA values than the XC72 supported catalysts in both
types of reduction methods. Figure 2.27 shows the CO stripping scans for the catalysts. The
ECSA of the catalysts was calculated in the same manner as in Section 2.2.2 and shown in Figure
2.28. Again, the hydrogen adsorption ECSA values are higher than the CO stripping values for
the same reasons as previously discussed in Section 2.2.2. The ECSA values for the CO stripping
concur with those from the hydrogen adsorption measurements. The onset potential of CO
oxidation was also roughly the same for all of the catalysts showing no apparent increased CO
tolerance of the platinum for either support or reduction method. The microwave reduced catalysts
fall in this range whereas the impregnated catalysts are slightly larger.
Figure 2.26. CV scans of the μwave Pt/C catalysts in 0.5 M H2SO4 under argon. The solid lines
indicate the microwave reduction; the dotted lines indicate the impregnation.
-1.2
-0.8
-0.4
0
0.4
0.8
0 0.2 0.4 0.6 0.8 1 1.2 1.4
Current Density (mA cm
-2
)
Potential (V vs. RHE)
Pt/XC-72R
Pt/CFx
61
Figure 2.27. CO stripping curves for the μwave Pt/C catalysts. The solid lines indicate the
microwave reduction; the dashed lines indicate the impregnation.
Figure 2.28. Summary of the ECSAs for the μwave Pt/C catalysts from the CO stripping and
hydrogen adsorption.
The LSV scans for the catalysts at 1600 RPM as displayed in Figure 2.29a with the limiting
current densities and onset potentials shown in Table 2.4. In both reduction methods, the CFx
-4
-2
0
2
4
6
8
10
0 0.2 0.4 0.6 0.8 1 1.2 1.4
Current Density (mA cm
-2
)
Potential (V vs. RHE)
Pt/XC-72R
Pt/CFx
0
30
60
90
120
150
180
XC72 uwave CFx uwave XC72 impreg CFx impreg
ECSA (m
2
g
-1
Pt)
CO strip
H-Adsorp
62
support shows higher limiting current densities due to the increased oxygen diffusion capabilities
from the fluorine in the carbon backbone of CFx as was discussed in Section 2.2.2. The microwave
polyol reduced catalysts also display higher limiting current densities compared with those of the
impregnated catalysts. This can arise from the increased ECSA of the microwave reduced catalysts
compared to the impregnated catalysts especially since the limiting current density for the XC72
microwave reduced catalyst is greater than that of the CFx impregnated catalyst. A close-up view
of the onset potential region of the LSVs is shown in Figure 2.29b. The onset potentials of the
ORR are also slightly higher for the CFx catalysts in both reduction methods compared to that of
the XC72. This shows a slight synergistic effect CFx has on the platinum for ORR as was
demonstrated in Section 2.2.2. The microwave reduced catalysts also have a slightly higher onset
potential than the impregnated catalysts meaning the microwave reduction process has some effect
on the ORR capabilities of the platinum due to the size of the particles [90]. It has been well
demonstrated that the decrease in the nanoparticle size via the microwave reduction method
correlates to an increase in catalytic activity with various types of reactions, mainly methanol
oxidation and oxygen reduction [91–93].
A Koutecky-Levich plot was made in the same manner as in Section 2.2.2. The
corresponding electron transfer numbers are shown in Table 2.4. The electron transfer numbers
for each catalyst is close to four resembling a four electron process for each. The electron transfer
numbers for the CFx catalysts are slightly higher than those of the XC72 catalysts meaning less
hydrogen peroxide on average was produced in the ORR. Also the electron transfer numbers are
higher for the microwave reduced catalysts than for the sodium borohydride reduced catalysts.
63
Figure 2.29. The LSV scans of the μwave Pt/C catalysts at 1600 RPM under O2 at a) normal
scan window and; b) zoom-in of the onset potential region. The solid lines indicate the
microwave reduction; the dashed lines indicate the impregnation.
-6
-5
-4
-3
-2
-1
0
0 0.2 0.4 0.6 0.8 1 1.2
Current Density (mA cm
-2
)
Potential (V vs. RHE)
Pt/XC-72R
Pt/CFx
-0.5
-0.4
-0.3
-0.2
-0.1
0
0.8 0.82 0.84 0.86 0.88 0.9
Current Density (mA cm
-2
)
Potential (V vs. RHE)
Pt/XC-72R
Pt/CFx
b)
a)
64
Figure 2.30. The Koutecky-Levich plot of the μwave Pt/C catalysts at 0.6 V vs. RHE. The solid
markers indicate the microwave reduction; the hollow markers indicate the impregnation.
Figure 2.31 shows the Tafel curves for the catalysts with the Tafel slopes shown in Table
2.4. All of the Tafel slopes are similar and around the expected value of 60 mV dec
-1
for platinum
as found in previous reports [84,94]. Since all of the platinum loadings were at around 20%-wt,
this further confirms that the Tafel Slopes can be greater affected by the platinum loadings on the
carbon supports compared to the support or the reduction procedure. The catalysts synthesized
with the microwave reduction show lower overpotentials than those synthesized via the
impregnation method. The CFx supported catalysts also show lower overpotentials than do those
of the XC72 supported catalysts for the respective reduction method.
500
1000
1500
2000
2500
0 0.05 0.1 0.15 0.2
Negative Inverse Current (-A
-1
)
ω
-1/2
Pt/XC-72R
Pt/CFx
65
Figure 2.31. Tafel curves for the μwave Pt/C catalysts. The solid markers indicate the
microwave reduction; the hollow markers indicate the impregnation.
Table 2.4. Summary of the RDE electrochemical tests of the μwave Pt/C catalysts.
Catalyst Onset Potential
(V vs RHE)
Limiting Current
Density (mA cm
-2
)
Electron
Transfer
Tafel Slope
(mV dec
-1
)
XC72 µwave 0.894±0.008 4.98±0.39 3.91±0.19 -80±2
CFx µwave 0.907±0.001 5.20±0.15 4.04±0.14 -79±1
XC72 impreg 0.872±0.006 4.72±0.24 3.79±0.10 -80±1
CFx impreg 0.891±0.005 4.88±0.12 3.84±0.17 -78±1
Figures 2.32a&b show the polarization curves for the fuel cell tests with the open circuit
voltage (OCV) and maximum power density results shown in Table 2.5. The ambient temperature
was measured at the same temperature and conditions as reported in Section 2.2.2. From the curves,
the OCVs of the microwave reduction catalysts were slightly higher than the impregnated catalysts.
The activation overpotentials for the impregnated catalysts were considerably greater in the
ambient temperature polarizations causing a decrease in power density compared to the microwave
0.78
0.8
0.82
0.84
0.86
0.88
0.9
0.92
0.01 0.1 1 10
Potential (V vs. RHE)
Current Density (mA cm
-2
)
Pt/XC-72R
Pt/CFx
66
reduced catalysts. The CFx supported catalysts also displayed a lower IR drop in the ohmic region
of the curve compared to the XC72 catalysts leading to an 8% and 24% increase in power density
compared with the XC72 supported catalysts at 50
O
C in the microwave reduced and impregnated
catalysts, respectively. This could be due to the enhanced oxygen diffusion capabilities and the
hydrophobic nature of the C-F bond in the CFx support and enhanced dispersion in the Nafion®
ionomer-based catalyst ink, most likely leading to an enhanced conductive network in the catalyst
layer, especially at the triple phase boundary as was also discussed in Section 2.2.2. In both of the
curves, the catalysts synthesized via the microwave reduction method show higher power densities
than the ones synthesized via impregnation method. There is a 49% and 30% increase in power
density at 50
O
C for the XC72 and CFx supported catalysts synthesized via the microwave
reduction, respectively. This can be due to the smaller, more dispersed platinum nanoparticles on
the carbon support leading to lower overpotentials and higher performance.
Electrochemical Impedance Spectroscopy (EIS) was performed on the fuel cells at the
same temperatures and shown in Figures 2.33a&b. The curves were fit to the circuit shown in
Figure 2.19 and the charge transfer resistances were calculated in the same manner as in Section
2.2.2. There is a decrease in the charge transfer resistance in the microwave reduced catalysts
compared to the impregnated cataysts due to the increase in current at the potential that the MEAs
were measured at. There is also a decrease in charge transfer resistance in the catalysts with the
CFx support compared to the XC72 support with the respective reduction method. The catalysts
synthesized by the microwave reduction displayed a 60% and 39% decrease in the charge transfer
resistance at 50
O
C for the XC72 and CFx supported catalysts, respectively. There was also an 8%
and 33% decrease in the charge transfer resistance between the CFx and XC72 supported catalysts
at 50
O
C for the microwave reduced and impregnated catalysts, respectively. A few factors can be
67
attributed to this decrease in charge transfer resistance. One factor could be the enhanced
dispersion and catalyst connective network with the Nafion® ionomer-based ink and the CFx
along with the inherent conductivity of CFx in the catalyst layers. This leads to a decrease in the
charge transfer resistance of the CFx supported catalysts in both types of reductions compared to
the XC72 which has less synergetic interaction with the Nafion® ionomer as was discussed in
Section 2.2.2. Another factor is the reduction in platinum particle size and enhanced dispersion
on the carbon-based support with the microwave reduction compared to the impregnation. This
also enhances the triple-phase boundary network in the catalyst layers of the cathode leading to a
decrease in charge transfer resistance.
CV scans were also taken of the MEAs in the fuel cell and shown in Figure 2.34. The
ECSA was calculated from the hydrogen adsorption peak in the same manner as from the CVs
from the electrochemical tests previously discussed in Section 2.2.2. The microwave reduced
catalyst MEAs displayed a 20% increase in ESCA compared to the impregnated catalyst MEAs.
This confirms the increased particle dispersion and ECSA of the microwave reduction catalysts in
the electrochemical tests translates over to MEA and fuel cell fabrication. The ECSAs for the
particular reduction methods for the catalysts however, have different results. For the microwave
reduced catalysts, the CFx supported catalyst displayed a 7% increase in ECSA compared to the
XC72 calalyst. However, for the impregnated catalysts, the XC72 supported catalyst displayed a
3% increase in ECSA compared to the CFx catalyst. This demonstrates that while the ECSA is
important for catalytic activity for fuel cells, other factors are required for enhanced fuel cell
performance such as minimal charge transfer resistance, enhanced catalyst network properties,
enhanced triple-phase boundary, intrinsic catalytic ORR activity, and other such properties.
68
In commercial fuel cells in operation presently, purified air instead of ultra-high pure
oxygen is pumped through the cathode compartment. In order to further test the efficacy of both
the fluorinated carbon support and the microwave reduction in commercial type settings, the fuel
cells were also tested with air flowing through the cathode compartment at the same flow rate as
the oxygen in Figures 2.32a&b. The polarization curves for the H2/air fuel cells are displayed in
Figures 2.35a&b. The results mimic the polarization curves for the fuel cells with the oxygen
flowing through the cathode compartment. For the microwave reduced catalysts, the CFx
supported catalyst displayed a 15% and 6% increase in power density compared to the XC72
supported catalyst at ambient temperature and 50
O
C, respectively. Furthermore, for the
impregnated catalysts, the CFx supported catalyst displayed a 16% and 28% increase in power
density compared with the XC72 supported catalyst at ambient temperature and 50
O
C,
respectively. Also, for the CFx supported catalysts, the microwave reduced catalyst displayed a
65% and 74% increase in power density compared with the impregnated catalyst at ambient
temperature and 50
O
C, respectively. Again, the fluorination of the carbon support aids in the
oxygen dissolution in the cathode catalyst layers as well as the enhanced connection in the catalyst
layers and triple-phase boundary as previously discussed. These factors hold true at the lower
stoichiometric amounts of oxygen present in air, most often used as the cathodic fuel in PEM-
based fuel cell vehicles and stationary devices.
69
Figure 2.32. Polarization curves for the H2/O2 fuel cells of the μwave Pt/C catalysts at: a)
ambient temperature and; b) 50
O
C. The solid lines indicate the microwave reduced catalysts;
the dashed lines indicate the impregnated catalysts.
0
100
200
300
400
0
0.2
0.4
0.6
0.8
1
0 200 400 600 800 1000 1200 1400 1600
Power Density (mW cm
-2
)
Potential (V)
Current Density (mA cm
-2
)
Pt/XC-72R
Pt/CFx
a)
0
100
200
300
400
500
600
0
0.2
0.4
0.6
0.8
1
0 500 1000 1500 2000 2500
Power Density (mW cm
-2
)
Potential (V)
Current Density (mA cm
-2
)
Pt/XC-72R
Pt/CFx
b)
70
Figure 2.33. EIS curves of the μwave Pt/C catalysts at 0.4 V vs. RHE at: a) ambient temperature
and; b) 50
O
C. The solid markers indicate the microwave reduced catalysts; the hollow markers
indicate the impregnated catalysts; the lines indicate the fitted circuit.
-0.2
-0.1
0
0.1
0.2
0 0.1 0.2 0.3 0.4
-Z" (Ohms)
Z' (Ohms)
Pt/XC-72R
Pt/CFx
a)
-0.2
-0.1
0
0.1
0.2
0 0.1 0.2 0.3 0.4
-Z" (Ohms)
Z' (Ohms)
Pt/XC-72R
Pt/CFx
b)
71
Figure 2.34. CV scan for the fuel cells of the μwave Pt/C catalysts. The solid lines indicate the
microwave reduction; the dashed lines indicate the impregnation.
-0.12
-0.08
-0.04
0
0.04
0.08
0 0.2 0.4 0.6 0.8 1
Normalized Current (A mg
-1
Pt)
Potential (V vs. RHE)
Pt/XC-72R
Pt/CFx
0
40
80
120
160
200
0
0.2
0.4
0.6
0.8
1
0 100 200 300 400 500 600 700
Power Density (mW cm
-2
)
Potential (V)
Current Density (mA cm
-2
)
Pt/XC-72R
Pt/CFx
a)
72
Figure 2.35. Polarization curves for the H2/Air fuel cells for the μwave Pt/C catalysts at: a)
ambient temperature and; b) 50
O
C with air flowing in the cathode. The solid lines indicate the
microwave reduced catalysts; the dashed lines indicate the impregnated catalysts.
Table 2.5. Summary of the H2/O2 fuel cell tests of the μwave Pt/C catalysts.
Catalyst MEA OCV (V) Peak Power
Density (mW cm
-2
)
Charge Transfer
Resistance (mΩ)
ECSA (m
2
g
-1
Pt)
Ambient 50
o
C Ambient 50
o
C Ambient 50
o
C 60
o
C
Pt/XC72 uwave 0.902 0.914 292 475 102.4 62.4 121.37
Pt/CFx uwave 0.891 0.902 311 514 80.7 57.1 128.91
Pt/XC72 impreg 0.892 0.924 191 318 143.6 105.1 100.50
Pt/CFx impreg 0.859 0.937 234 395 102.6 79.7 98.04
0
50
100
150
200
250
0
0.2
0.4
0.6
0.8
1
0 100 200 300 400 500 600 700 800
Power Density (mW cm
-2
)
Potnetial (V)
Current Density (mA cm
-2
)
Pt/XC-72R
Pt/CFx
b)
73
2.4: The Effect of Fluorinated Carbon Supports for Various Manganese Dioxide Phases on ORR
Kinetics in Basic Media
2.4.1: Experimental Methods
Three different phases of manganese dioxide: α, β, and γ were synthesized in order to
effectively test the effect of fluorination on the carbon support for oxygen reduction in alkaline
media. For α-MnO2, 1.014 g MnSO4·H2O was added to 20 mL Millipore (Direct-Q UV, 18.2 MΩ)
and the solution stirred and sonicated until homogenous. 0.95 g KMnO4 was then dissolved in 20
mL Millipore water, added slowly to the solution, then again stirred and sonicated for 30 min. The
solution was then transferred to an oil bath and refluxed at 100
O
C for 20 hrs. Afterwards, the
solution was washed and centrifuged with Millipore water until the pH of the supernatant was
neutral and dried in an oven.
For β-MnO2, 1.2676 g MnSO4·H2O in 150 mL Millipore water was mixed with 1.7115 g
(NH4)2S2O8 in an oil bath at 30
O
C for 30 min. It was then heated to 80
O
C and mixed for 2 hrs.
The solution was vacuum filtered and dried in an oven overnight. The resulting powder was then
annealed in a tube furnace under air for 1 hr at 400
O
C with the temperature heated up and cooled
down at 10
O
C min
-1
.
For γ-MnO2, 0.81 g (NH4)2S2O8 in 10 mL Millipore water was added to 1.1 g MnSO4·H2O
in 20 mL Millipore water. The solution was stirred and sonicated for 30 min and then sealed in a
glass-lined autoclave and heated to 140
O
C for 1 hr. It was taken out and washed and centrifuged
with Millipore water until the pH of the supernatant reached neutral.
Each of the synthesized phases of manganese dioxide was stirred and sonicated with the
XC72 and CFx carbon support in an aqueous solution until the solution was well dispersed. The
74
manganese dioxide loading in each sample was set to 20%. The solution was then centrifuged and
dried in an oven at 110
O
C overnight.
The catalysts were characterized from a variety of methods. The scanning electron
microscope (SEM) images were taken on a JEOL JSM-7001 electron microscope with an
acceleration voltage of 10 keV. Elemental dispersion spectroscopy (EDS) measurements were
also carried out on the EDAX Genesis at the same acceleration voltage as the SEM images. X-ray
diffraction (XRD) patterns were taken on a Rigaku Ultima Diffractometer from a 2θ value of 10
O
to 90
O
.
The electrochemical testing was carried out in a three cell rotating disk electrode (RDE)
setup. The setup comprised of a Teflon wrapped glassy carbon RDE (0.195 cm
2
working area) as
the working electrode, a platinum wire as the counter electrode, and a mercury mercury oxide
(MMO) (4.24 M KOH filling soln.) as the reference electrode. 1 mg of the catalyst was added to
1 mL of a Nafion® binder solution comprising of 10 mg of liquid Nafion® solution (5% wt, Alfa
Aesar), 90% Millipore water, and 10% isopropanol to make a catalyst ink solution that would be
applied to the electrode. The catalyst ink was mixed via both sonication and vortex mixing until
the catalyst was homogenously dispersed. Then, 20 μL of ink was micropipetted onto the RDE
and dried in an oven at 65
O
C. The electrode was then connected to the three cell setup and all the
tests were performed in 0.1 M KOH solution. Before the tests, the cell was purged with argon
(ultra high pure grade, 99.999%) for 15 minutes and the cell was cycled from -1.0 V to 0.5 V vs.
MMO to remove any impurities from the catalyst. For the ORR testing, the cell was purged with
oxygen (ultra high pure grade, 99.994%) for 15 minutes and the cell was cycled between the same
potentials as previously mentioned.
75
For the alkaline direct methanol fuel cell tests, separate catalyst ink solutions were prepared
of the α-MnO2/C catalysts with a 1:5:2 ratio of catalyst: Millipore water: 5% Tokuyama anionomer.
This solution was sonicated for 8 minutes and then manually painted on a 4 cm
2
piece of Toray
Carbon paper (TGH-060, 10% wet proofing) with a catalyst loading of 2.0 mg cm
-2
and dried in
an oven at 110
O
C. This served as the cathode electrode while the anode electrode was made from
platinum/ruthenium black inks with the same composition as previously mentioned and painted on
the same type of carbon paper with loadings of 4.0 mg cm
-2
to ensure the limiting reaction was the
ORR on the cathode and dried in the same oven. The membrane electrode assembly (MEA) was
then made by sandwiching a Tokuyama A201 membrane between the electrodes and pressing at
0.5 tons of force at 110
O
C for 5 minutes. The MEA was then put in a fuel cell system composed
of conducting graphite plates and brass current collectors. The MEA was tested by flowing a 1.0
M methanol and 2.0 M KOH solution through the anode at 5 mL min
-1
and humidified oxygen
(ultra-high pure grade, 99.994%) through the cathode at 100 mL min
-1
. Polarization curves were
taken by scanning the potential at 50 mV s
-1
.
2.4.2: Results
The various structures of the MnO2 phases were first characterized by SEM imaging shown
in Figures 2.37a-c for the α-MnO2 catalysts, Figures 2.38a-c for the β-MnO2 catalysts, and Figures
2.39a-c for the γ-MnO2 catalysts. For each of the phases, the unsupported MnO2 catalyst has a
very similar structure as the supported catalyst signifying little change in the structure when
grafting the MnO2 catalysts onto the carbon supports. Each phase has a distinctive morphology
which maintains its structure on both the fluorinated and non-fluorinated carbon support.
76
To further confirm the deposition of the MnO2 on the supports and the fluorine in the CFx,
EDAX measurements were taken and the one for the α-MnO2/CFx is shown in Figure 2.39. The
manganese present in the EDAX spectrum was quantified at around 28%-wt which was observed
for the rest of the catalysts in the different phases and supports. There is also the fluorine peak
present indicating the fluorine backbone of the CFx was still intact after the synthesis. Further
elemental mapping was obtained to show the distribution of the particular elements from the SEM
image shown in Figures 2.40a-e.
Figure 2.36. SEM images of: a) α-MnO2; b) on XC72 support and; c) on CFx support.
a)
b)
c)
77
Figure 2.37. SEM images of a) β-MnO2; b) on XC72 support and; c) on CFx support.
a)
b)
c)
a)
b)
78
Figure 2.38. SEM images of: a) γ-MnO2; b) on XC72 support and; c) on CFx support.
Figure 2.39. EDAX spectrum of α-MnO2/CFx.
c)
79
Figure 2.40. a) SEM image of α-MnO2/CFx with corresponding EDAX elemental mapping
images of: b) carbon; c) fluorine, d) manganese and; e) oxygen.
a)
b)
c)
d)
e)
80
Further characterization on the crystallinity of the various MnO2 phases was carried out by
XRD patterns and are displayed in Figures 2.42a-c. Figure 2.42a shows the XRD pattern of the α-
MnO2 phase (PDF#01-072-1982) [95] with peaks signifying the particular phase. The intensity of
the peaks is slightly reduced with the carbon supported catalysts, however, there is no apparent
shift in the angles indicating no loss of crystallinity in the catalyst after the grafting. Figure 2.42b
shows the XRD pattern of the β-MnO2 phase (PDF#03-024-0735) [69]. Again, these peaks are
visible with a lower intensity shown for the carbon supported catalysts. Figure 2.42c shows the
XRD pattern of the γ-MnO2 phase (PDF#01-14-0644) [96]. The same characteristic is again seen
in the spectra for the carbon supports as in the other two manganese dioxide phases. In each of
the patterns, the peaks corresponding to the various MnO2 phases are visible in the carbon
supported catalysts signifying a preservation of the crystallinity during the grafting process. After
the grafting, the intensities of the manganese dioxide phases are also reduced with a larger FWHM
indicating the decrease in crystallite size. The peak at around 23
O
signifies the (002) Miller Index
of the CFx and XC72 carbon supports.
10 20 30 40 50 60 70 80 90
Intensity (a.u.)
2 θ (degrees)
aMnO2
aMnO2/XC72
aMnO2/CFx
(220
(310)
(121)
(301
(411)
(600)
(521
(002)
(451)
(312)
(002)
a)
(200)
81
Figure 2.41. XRD patterns of the: a) α-MnO2; b) β-MnO2 and; c) γ-MnO2 phase catalysts.
Figure 2.43 shows a regular CV scan of the catalysts in 0.1 M KOH solution. There are
three main types of curves present in the CVs each indicating the three various phases of
10 20 30 40 50 60 70 80 90
Intensity (a.u.)
2 θ (degrees)
b-MnO2
b-MnO2/XC72
b-MnO2/CFx
(101)
(111)
(211)
(220)
(002)
(301)
(002)
b)
10 20 30 40 50 60 70 80 90
Intensity (a.u.)
2 θ (degrees)
g-MnO2
g-MnO2/XC72
g-MnO2/CFx
(111)
(211)
(212)
(511)
(002)
c)
(110)
(101)
82
manganese dioxide. For each of the respective phases, the CFx supported catalyst shows slightly
higher currents most likely signifying a greater active surface area. To assess the ORR kinetics of
the catalysts, LSV scans were performed and shown in Figure 2.44. A summary of the onset
potential and limiting current densities are reported in Table 2.6. In each of the LSV scans, the
manganese dioxide catalyst grafted onto the CFx support displays increased limiting current
densities compared to the XC72 supported catalysts for each of the respective phases. Further, the
CFx supported catalysts displayed increased onset potentials for each of the phases compared to
the XC72 supported catalysts. The certain structures of the various phases of the manganese
dioxide can significantly affect the ORR properties of the catalyst.
The mechanism of ORR in alkaline media has been studied previously on the manganese
dioxide catalysts. As was previously reported, the MnOOH species is the species present at the
surface in the alkaline media and is responsible for the ORR taking place [97]. The oxygen binds
to the MnOOH (Mn
3+
) in the various steps shown in the following equations [68,69,98]:
Mn
4+
+ e
-
↔ Mn
3+
(5)
O2 + Mn
3+
→ Mn
3+
-O2,ads (6)
Mn
3+
-O2,ads → Mn
4+
-O2
-
ads (7)
Mn
4+
-O2
-
ads + H2O + e
-
→ Mn
4+
-HO2
-
ads + OH
-
(8)
A Koutecky-Levich plot was also made in the same manner as in Section 2.2.2 and is shown
in Figure 2.44. The electron transfer numbers were extracted from the Koutecky-Levich equation
in the same manner as in Section 2.2.2 and the electron transfer numbers are reported in Table 2.6.
The electron transfer numbers follow the same trend for overall activity as the LSV curves. For
each of the manganese dioxide phases, the CFx supported catalysts display higher electron
numbers than the XC72. In addition, the α-MnO2/C catalysts display the higher electron number
83
while the β-MnO2/C catalysts display, on average, the lowest electron transfer number. These
numbers are similar to ones reported previously for carbon supported manganese dioxide catalysts
[99].
Tafel curves were also taken for the catalysts and are shown in Figure 2.45. The Tafel
slopes are in the range of one reported in previous studies with manganese dioxide in alkaline
media [100]. For each Tafel curve, the slope of the CFx supported manganese dioxide catalyst is
lower than the XC72 supported catalyst. This signifies higher intrinsic activity for the CFx
supported catalyst than the XC72. Furthermore, each of the phases of the manganese dioxide with
the Tafel slopes: α-MnO2/C > γ-MnO2/C > β-MnO2/C.
Figure 2.42. CV scans of the MnO2/C catalysts. The solid lines indicate the samples with CFx
support; the dotted lines indicate the samples with XC72 support.
-1
-0.75
-0.5
-0.25
0
0.25
0.5
0.75
1
-1 -0.8 -0.6 -0.4 -0.2 0 0.2 0.4 0.6
Current Density (mA cm
-2
)
Potential (V vs. MMO)
a-MnO2
b-MnO2
g-MnO2
84
Figure 2.43. LSV scans of the MnO2/C catalysts in 0.1M KOH solution at 1600 RPM. The
solid lines indicate the catalysts with CFx support; the dotted lines indicate the catalysts with the
XC72 support.
Figure 2.44. Koutecky-Levich plot of the MnO2/C catalysts at -0.4 V vs MMO. The solid
markers indicate the catalysts with the CFx support; the hollow markers indicate the catalysts
with the XC72 support.
-3
-2.5
-2
-1.5
-1
-0.5
0
-0.8 -0.6 -0.4 -0.2 0 0.2
Current Density (mA cm
-2
)
Potential (V vs. MMO)
a-MnO2
b-MnO2
g-MnO2
1000
2000
3000
4000
5000
6000
7000
8000
0 0.05 0.1 0.15 0.2 0.25 0.3 0.35
Neg. Inverse Current (-A
-1
)
ω
-½
a-MnO2
b-MnO2
g-MnO2
85
Figure 2.45. Tafel plots of the MnO2/C catalysts at 1600 RPM. The solid markers indicate the
catalysts with the CFx support; the hollow markers indicate the catalysts with the XC72 support.
Table 2.6. Summary of the RDE electrochemical tests of the MnO2/C catalysts.
Catalyst Onset Potential
(V vs MMO)
Limiting Current
(mA cm
-2
)
Tafel Slope
(mV dec
-1
)
Electron
Transfer
α-MnO2/XC72 -0.023±0.016 2.49±0.36 117±9 3.7
α-MnO2/CFx -0.018±0.007 2.63±0.38 111±7 3.75
β-MnO2/XC72 -0.094±0.028 2.04±0.12 77±2 2.64
β-MnO2/CFx -0.082±0.014 2.28±0.24 69±1 3.33
γ-MnO2/XC72 -0.075±0.035 2.11±0.23 97±13 2.47
γ-MnO2/CFx -0.033±0.021 2.80±0.15 89±8 3.43
In order to further assess the ORR activity of the CFx and XC72 supports, MEAs were
constructed and run in alkaline DMFCs and the polarization curves shown in Figures 2.47a&b
with the α-MnO2 catalysts with the results reported in Table 2.7. At both ambient temperature and
65
O
C, the catalysts with the carbon supports give higher performances than the unsupported
manganese dioxide catalyst. At ambient temperature, the CFx supported catalyst displayed a 81%
-0.2
-0.15
-0.1
-0.05
0
0.01 0.1 1
Potential (V vs. MMO)
Current Density (mA cm
-2
)
a-MnO2
b-MnO2
g-MnO2
86
and 167% increase in power density compared with the XC72 supported catalyst and bare catalyst,
respectively. At 65
O
C, there is a 87% and 127% increase in power density with the CFx supported
catalyst compared to the XC72 supported catalyst and bare catalyst, respectively. Another aspect
of the polarization curves is the OCV of the cells. The CFx supported catalyst MEA displayed a
slightly increase in OCV compared with the XC72 supported catalyst and the bare catalyst. This
could arise from the increased onset potential from the half-cell LSV tests in Figure 2.43.
The CFx could also have an enhanced conductive network in the Tokuyama ionomer
solution similar to that with the Nafion® ionomer discussed in Sections 2.2.2 and 2.3.2. In addition
to the increased oxygen dispersion, the CFx support and been demonstrated to have is the increased
water dispersion from the hydrophobic fluorinated carbon support. This could help increase the
kinetics of the oxygen reduction in alkaline conditions, specifically since the ORR in alkaline
conditions requires both the oxygen and water as reactants.
0
1
2
3
4
5
6
0
0.2
0.4
0.6
0.8
1
0 10 20 30 40 50 60
(Power Density (mW cm
-2
)
Potential (V)
Current Density (mA cm
-2
)
a-MnO2
a-MnO2/XC72
a-MnO2/CFx
a)
87
Figure 2.47. Polarization curves of the alkaline DMFCs for the MnO2/C catalysts at: a) ambient
temperature and; b) 65
O
C.
Table 2.7. Summary of the fuel cell tests of the MnO2/C catalysts.
Catalyst MEA OCV (V) Peak Power Density (mW cm
-2
)
Ambient 65
O
C Ambient 65
O
C
α-MnO2 0.81 0.84 2.1 5.5
α -MnO2/XC72 0.77 0.84 3.1 6.7
α -MnO2/CFx 0.81 0.86 5.6 12.5
2.5: Conclusions
In order to reduce the amount of noble metal catalyst usage in various fuel cell systems to
make them more commercially viable, the catalyst support can play a crucial role. When screening
catalysts for use in fuel cells and other applications, other factors must be considered outside of
the catalyst size, morphology, support etc. Previous studies have shown that the larger the particle
0
3
6
9
12
15
0
0.2
0.4
0.6
0.8
1
0 20 40 60 80 100
Power Density (mW cm
-2
)
Potential (V)
Current Density (mA cm
-2
)
a-MnO2
a-MnO2/XC72
a-MnO2/CFx
b)
88
size of the catalyst, particularly platinum, the greater the specific ORR activity. This isn’t always
the case as was shown in Section 2.2.2. At each of the loadings, the platinum particle size was
smaller on CFx, yet many of the platinum loadings had higher specific ORR activities compared
to the XC72 support. One of the aspects of half-cell or fuel cell testing that is overlooked is the
interaction of the catalyst with the ionomer in the ink solution. This aspect can affect the overall
properties of the catalyst layer and performance of the MEA. It seems there is a greater symbiotic
effect of the fluorinated carbon backbone from the CFx that allows better catalyst-ionomer
dispersion and intercalation than with the non-fluorinated carbon backbone of XC72. This was
shown to translate into higher fuel cell performances and lower charge transfer and membrane
resistances even though the fluorination of carbon has been shown to reduce the overall
conductivity of the catalyst network in the cathode compartment of the fuel cell.
In addition to the particle size is the myriad of synthetic routes that can be undergone to
reduce metal salts onto carbon supports. More facile reduction methods such as the microwave-
assisted polyol method as investigated in Section 2.3.2 showed decreased particle size and
increased dispersion when compared to the impregnation route in Section 2.2.2. This enhanced
dispersion led to an enhanced ECSA in both the half cell and fuel cell testing. The CFx supported
catalysts displayed enhanced ORR activity in the half cell tests which translated into enhanced
performance in the fuel cell tests for both reduction types. This is due to factors including the
improved oxygen dispersion and water management properties of the CFx support as well as the
enhanced catalyst dispersion and conductivity in the Nafion® ionomer-based ink solution used in
both the half-cell and fuel cell tests. The nature of the catalyst and ionomer-based ink interaction
is a factor in catalyst synthesis that is often overlooked and can have a considerable role in the
performance of various fuel cells.
89
This enhancement in ORR activity translates not just in the acidic media, but in the basic
media. The Nafion® ionomer has been shown to be an effective binder for alkaline-based half-
cell RDE testing which doesn’t appear to affect the catalytic properties of the catalyst ink solution.
This can potentially be carried over to the MEA fabrication, yielding similar results. In Section
2.4.2, the CFx support was shown to help increase the intrinsic ORR properties of the manganese
dioxide catalysts at different phases. In addition to help disperse the catalyst, the fluorination also
assisted in the overall oxygen dispersion in the humidified basic environment and the enhanced
catalyst ink network in the MEAs with the Tokuyama anionomer.
Numerous studies involving catalysts only deal with the half-cell studies and neglect the
MEA fuel cell studies. The full fuel cell studies integrate numerous other variables that go beyond
that of the half-cell studies. Sufficient connectivity between the catalyst and the binder/conductive
solution need to be ascertained and along with the triple boundary aspect. The fluorination of the
high surface area conductive carbon supports helps interact with the fluorine present in the
Nafion® ionomer and Tokuyama anionomer backbone to help create and enhanced conductive
network in the catalyst ink which translates into higher performances in fuel cells both in acidic
and alkaline media.
90
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Chapter 3:
The Applications of Graphene as both a Catalyst and Support in Fuel Cells
3.1: Graphene Characteristics and Applications
No other materials have been as extensively studied in the 21
st
Century as graphene and its
constituents. As shown in Figure 3.1, the number of publications involving graphene have
increased exponentially since it’s first “official” discovery by Novoselov and Geim in 2004 [1].
Graphene has a two dimensional graphitic structure comprised of sp
2
hybridized carbons that form
a long sheet. Graphene in its basic form can be used to synthesize various carbon structures as
shown in Figure 3.2. Over the past decade, graphene has attracted much attention in the scientific
and industrial communities. Due to its high surface area [2], charge carrier mobility [3],
capacitance [4], functionalization ability [5], and other aspects, graphene has been sought after for
use in batteries [6,7], capacitors [8,9], sensors [10,11], and many other devices. Graphene is also
a relatively versatile material that can be synthesized via multiple routes from graphene oxide
(GO), graphite, and other small organic molecules. A few of these routes include hydrothermal
[12], microwave [13], ultraviolet irradiation [14], and electrochemical [15] methods. Each
synthetic route affects the size and morphology, which impacts certain chemical and electrical
properties of graphene. In the hydrothermal route, GO is usually dispersed in an aqueous solution
or other solvents such as ethylene glycol. A reducing agent is then used to remove the oxygen
functional groups such as epoxides, carboxylic acids and hydroxides while repairing the sp
2
carbon
100
lattice. Unlike hydrazine, which is toxic, sodium borohydride is one of the widely used reducing
agents due to its strong reducing ability and use in a wide range of conditions.
Figure 3.1. Number of publications based on graphene by year. Adapted from [16].
Figure 3.2. Various implementations of 2D graphene. Adapted from [17].
101
Much work has been carried out to both enhance the electrochemical ORR kinetics as well
as to eliminate the use of noble metal catalysts such as platinum in fuel cell and battery applications.
Recently, graphene and its derivatives have shown great promise as substitutes for noble metal
catalysts for ORR especially in the alkaline medium [18]. Much research has been done on
graphene to enhance its ORR capabilities such as doping p-block elements into the structure
[19,20], graphene quantum dots [21], graphene flakes [22], and other moieties as both a stand-
alone catalyst and support for other metal catalysts.
One of the most widely implemented synthetic routes to synthesize GO is Hummer’s
method [23]. After the oxidation step, the GO is often washed with hydrochloric acid and water
in order to remove the residual ions present. This creates a colloidal solution that is very stable
[24]. One of the potential drawbacks of the method is the drying step, which often occurs over
numerous days with a dialysis bag or requires several solvents to help remove the water [25,26].
In many of the studies involving the hydrothermal reduction of GO, the pH of the solution
is usually adjusted to around 9-10 before adding the reducing agent [27–29]. This is also the case
with many metal salts impregnated onto supports with sodium borohydride [30]. There have been
other studies, wherein graphene oxide has been reduced in an autoclave at different pH levels.
However, different compounds were used for the particular reduction [31].
A more recent approach in synthesizing GO and few layer graphene materials has been the
electrochemical exfoliation of graphite rods [32–34], foils [35], sheets [36], pencil cores [37–39],
and highly ordered pyrolytic graphite (HOPG) [40]. In this technique, the graphite is connected
to a power supply or potentiostat and immersed in a solution. A potential is applied, exfoliating
the graphene sheets by the intercalation of the various ions of the solution in the sheets of graphene.
The exfoliated sheets can then be easily separated and cleaned. Some of the solutions that have
102
yielded high quality, few layered graphene have been ionic liquids [39,41,42] and by regulating
the pH of the solution [40,43,44]. Electrochemically exfoliated graphene has been used as a
platinum nanoparticle support in previous research [43,45] However, no previous research has
been performed on the exfoliation and platinum deposition in a truly “one-pot” synthesis from start
to finish in an aqueous platinum salt solution.
In this chapter, reduced graphene oxide (rGO) has been synthesized via a variety of routes,
including the widely used Hummer’s method and electrochemical exfoliation. Different synthetic
parameters were used in each of these routes to study the effects on the structure and the various
catalytic properties, mainly ORR, of the reduced graphene oxide as both a stand-alone catalyst and
a catalyst support for various metal catalysts.
3.2: The Effect of pH on the Reduction of Graphene Oxide and its Catalytic Properties Towards
ORR in the Basic Medium
3.2.1: Experimental Methods
Graphene oxide was synthesized from the modified Hummer’s method [23]. In a typical
synthesis, 2.0 g of graphite powder (Alfa Aesar, 99.8%, 325 mesh) was added to 50 mL anhydrous
sulfuric acid (Marcon) and the mixture stirred for 30 minutes in an ice bath. 5.0 g potassium
permanganate (Macron) was then added and the solution was stirred in the ice bath for 1 hour
before being transferred to an oil bath at 40
O
C and stirred for 1 hour. In order to terminate the
reaction, 50 mL of Millipore water was added followed by 15 mL of 30% hydrogen peroxide
103
(Macron) was added. The solution was then vacuum filtered and washed and centrifuged 3 times
with 10% hydrochloric acid followed repeated washing with Millipore water and then acetone.
The resulting solution was filtered and the obtained solid then dried in an oven at 80
O
C overnight
to obtain the GO powder.
Reduced graphene oxide was synthesized at various pH values from the sodium
borohydride reduction method. 0.1 g of the previously synthesized GO powder was stirred and
sonicated in 100 mL of Millipore water until it was well dispersed. The pH of the solution was
constantly monitored by a pH meter (Mettler Toledo). The initial pH was around 3.5 and was
adjusted to the values of 1, 4, 7, 10, and 13 by either hydrochloric acid or sodium hydroxide. The
solution was then heated to 95
O
C in an oil bath. Freshly made aqueous solutions of sodium
borohydride were added in 3 aliquots each containing 0.4 g sodium borohydride in order to reduce
the GO. The solution was stirred at 95
O
C for 1 hour then at room temperature overnight. The
rGO was then vacuum filtered and washed and centrifuged with Millipore water until the pH value
of the supernatant was 7 and then dried in an oven at 80
O
C overnight.
Thermogravemetric analysis (TGA) was performed on a TGA-50 Thermogravimetric
Analyzer (Shimadzu) with the heat rate of 10
O
C per minute. The scanning electron microscope
(SEM) images were taken on a JEOL JSM-7001 electron microscope with an acceleration voltage
of 10 keV. The transmission electron microscope (TEM) images were obtained on a JEOL JEM
2100F electron microscope with an acceleration voltage of 200 keV. X-ray diffraction (XRD)
patterns were obtained on a Rigaku X-Ray Diffractometer from a 2θ value of 10
O
to 90
O
. Fourier
Transform Infrared Spectroscopy (FTIR) spectra were taken on a Jasco FT/IR-4600 from 4000
cm
-1
to 400cm
-1
. Raman spectra were obtained on a Horiba XploRA ONE Raman Microscope
with a laser excitation wavelength of 532 nm. The pKa of the catalyst surface species were also
104
determined by acid-base potentiometric titration. 10 mg of catalyst was sonicated in 10 mL of 0.1
M NaOH solution. The pH of the solution was constantly monitored with a digital pH meter
(Mettler Toledo) while 100 μL aliquots of 0.1 M HCl solution were pipetted with the pH
equilibrating between each addition. A blank solution of 10 mL 0.1 M NaOH was also titrated for
comparison.
The electrochemical testing was carried out in a three cell rotating disk electrode (RDE)
setup. The setup comprised of a Teflon wrapped glassy carbon RDE (0.195 cm
2
working area) as
the working electrode, a platinum wire as the counter electrode, and mercury/mercury oxide
(MMO) (Koslow, 4.24 M KOH filling soln.) as the reference electrode. 1 mg of the catalyst was
added to 1 mL of a Nafion binder solution comprising of 10 mg of liquid Nafion solution (5% wt,
Alfa Aesar), 90% Millipore water, and 10% isopropanol to make a catalyst ink solution that would
be applied to the electrode. The catalyst ink was mixed via both sonication and vortex mixing
until the catalyst was homogenously dispersed. Then, 20 μL of ink was micropipetted onto the
RDE and dried in an oven at 65
O
C. The electrode was then connected to the three cell setup and
all the tests were performed in 0.1 M KOH solution. Before the tests, the cell was purged with
argon (ultra-high pure grade, 99.999%) for 15 minutes and the cell was cycled from -1.0 V to 0.5
V vs. MMO to remove any impurities from the catalyst. For the ORR testing, the cell was purged
with oxygen (ultra-high pure grade, 99.994%) for 15 minutes and then cycled from -1.0 V to 0.5
V vs. MMO.
105
3.2.2: Results
The reduction of GO was first confirmed through TGA curves shown in Figure 3.3. The
GO catalyst displays three main weight loss drops: at 50
o
C from the loss of water adhering to the
surface, at 200
o
C from the loss of the hydroxide, carboxyl, and epoxide groups, and at 600
O
C
from the loss of the graphitic lattice [46]. There is only one major weight loss for the rGO catalysts
at around 400-600
O
C depending on the catalyst signifying the removal of the graphitic lattice.
Figures 3.4a-f shows the SEM images of the rGO catalysts. Each of the rGO catalysts show
increased wrinkles, when compared with the GO. There is no apparent macro morphological
differences between the rGO sheets. Figures 3.5a-f shows the TEM images of the rGO catalysts.
The images also show an increase of wrinkles and folding in the rGO sheets, when compared to
the GO.
Figure 3.3. TGA curves of the rGO catalysts.
0%
20%
40%
60%
80%
100%
120%
0 100 200 300 400 500 600 700 800
weight (% of initial)
Temperature (
o
C)
rGO-1
rGO-4
rGO-7
rGO-10
rGO-13
GO
106
Figure 3.4. SEM images of the rGO catalysts: a) GO; b) rGO-1; c) rGO-4; d) rGO-7; e) rGO-
10 and; f) rGO-13.
a) b)
c) d)
e)
f)
107
a) b)
c) d)
108
Figure 3.5. TEM images of the rGO catalysts: a) GO; b) rGO-1; c) rGO-4; d) rGO-7; e) rGO-10
and; f) rGO-13. The measuring bar indicates 100 nm.
In order to confirm the reduction of the catalysts, XRD patterns were obtained and are
shown in Figure 3.6. The GO catalyst displays a sharp peak at a 2θ value of 11.0
o
signifying the
(100) crystallographic orientation of GO with a d-spacing of 8.01 Å. In each of the reductions, the
main peak shifts to around 23
o
for the (002) Miller Index peak of the rGO species. The d-spacing
is the average distance between the rGO layers and is calculated from the equation:
nλ = 2dsin θ (9)
Figure 3.7 shows the d-spacing of the rGO catalysts and the crystalline sizes as calculated from
the full width at height maximum (FWHM). As the pH of the solution decreases, the d-spacing
of the catalysts decreased except for the rGO-13 catalyst. This could stem from the decrease in
concentration of oxygen containing species on the graphitic lattice from the reduction in the more
acidic solutions. As the pH of the solution becomes more alkaline, the crystalline size of the rGO
sheets increases and the amount of defects in the graphitic structure decrease [47,48].
e) f)
109
The FTIR spectra displays the functional groups of the rGO catalysts are shown in Figure
3.8. The spectrum of the GO has the characteristic peaks at 3342 cm
-1
and 1414 cm
-1
representing
the OH and C-OH stretching, respectively, 1715 cm
-1
indicating the carboxyl stretch, 1612 cm
-1
denoting the aromatic sp
2
carbon stretch, 1218 cm
-1
representing the C-O epoxy stretch, and 1026
cm
-1
corresponding to the alkoxy C-O stretch [49,50]. These major peak intensities are greatly
reduced in the rGO catalysts signifying the loss of many of the functional groups after the reduction
[51]. In the rGO-13 and rGO-10 catalysts, there is a slight presence of the hydroxide stretching
peaks at 3342 cm
-1
and the C-OH stretching at 1414 cm
-1
and 1419 cm
-1
, most likely due to the
basic environment that the GO was reduced in.
Raman spectroscopy plays a vital role in the characterization and measuring disorder of the
graphitic plane in both GO and rGO catalysts. Figure 3.9 shows the Raman spectra of the catalysts.
There are two characteristic peaks in the spectra, one at around 1370 cm
-1
signifying the D-band
and the other at around 1580 cm
-1
signifying the G-band. The D-band measures the lattice
distortion of the graphene from various functional groups and defects, while the G-band measures
the first order scattering of the E2g phonon from the sp
2
hybridized carbon lattice [52]. The ID/IG
ratio reveals the extent of defects and size of the sp
2
graphitic domain of the catalyst [53]. A
summary of the ID/IG ratios is reported in Figure 3.10. After the reduction, the ID/IG ratios of the
rGO catalysts increase due to the smaller overall graphitic domain size compared to GO. The
ratios for the rGO catalysts also increase as the pH of the solution decreases due to the defects in
the graphite lattice from the removal of some of the oxygen groups [54]. This could be due to the
harsher acidic environment of the reduction for sodium borohydride, where many of the oxygen
groups are removed but the sp
2
domain is not restored. Another aspect of the ID/IG ratio is the
calculation of the in-plane sp
2
grain size of the rGO catalyst particles from the equation [55]:
110
La = C(λ) * (ID/IG)
-1
(8)
where La is the in-plane grain size (in nm) and C(λ) is the incident wavelength (532 nm)
corresponding to a value of 4.4 nm. The in-plane grain sizes of the rGO catalysts are shown in
Figure 3.10. As the pH of the solution increases, the in-plane grain size of the catalyst increases,
indictive of the better restoration of the sp
2
graphitic lattice from the sodium borohydride. Another
aspect from the Raman spectra is the shifting of the D and G band peaks. The red shift in the G-
band from the GO to rGO catalysts can be attributed to a more ordered aggregation of the sp
2
hybridized carbons from the reduction.[56] The rGO-13, rGO-10, and rGO-1 catalysts all display
the same amount of redshifting for both the D and G bands of the Raman spectra. The red shift in
the D band signifies the amount of strain induced on the graphitic lattice [57,58]. The rGO-7 and
rGO-4 catalysts display a higher red shift in the D and G bands indicating more strain in the
graphitic lattice. It is unclear why the rGO-1 catalyst exhibits the same red shift for the D and G
bands as the rGO-13 and rGO-10 catalysts instead of the rGO-7 and rGO-4 catalysts. One possible
explanation could be the addition of the HCl in the aqueous solution has an effect on the graphitic
structure. The shift in the G band could also be due to the presence of a second peak, the D’ band,
at 1620 cm
-1
[59]. However, the presence of the D’ band is evident in all of the rGO catalysts as
well as the GO catalyst and does not explain why some of the peaks are more red shifted than
others.
111
Figure 3.6. XRD patterns of the rGO catalysts.
Figure 3.7. Summary of the XRD characteristics of the rGO catalysts from Figure 3.6.
10 20 30 40 50 60 70 80 90
Intensity (a.u.)
2 θ (degrees)
GO
rGO-13
rGO-10
rGO-7
rGO-4
rGO-1
(002)
(100)
40
45
50
55
60
3.5
3.6
3.7
3.8
3.9
4
0 2 4 6 8 10 12 14
Crystallite Size (nm)
d-spacing (Å)
pH
d-spacing
crystallite size
112
Figure 3.8. FTIR spectra of the rGO catalysts.
XPS is also a powerful technique for characterizing the different functional groups and
structures of the rGO catalysts. XPS C1s spectra were obtained and the deconvoluted spectra are
shown in Figure 3.11a-f with the deconvoluted peak results shown in Table 3.1. The spectrum for
GO in Figure 3.11a displays two peaks: one comprised of the sp
2
and sp
3
hybridized carbon regions
of the graphitic lattice at 284.4 eV and 284.8 eV, respectively and the other at 286.8 eV for the
epoxide C-O peak. The size of the C-O peak is greatly reduced in each of the rGO catalysts
showing an effective reduction of GO across the pH spectrum. The C=O peak showed little
decrease in area due to sodium borohydride’s relative inability to reduce carboxyl groups [60].
Another aspect worth noting is the size of the sp
2
and sp
3
deconvoluted peaks in the spectra shown
in Table 3.1. As the pH of the solution increases, the ratio in size between sp
2
and sp
3
hybridized
carbon increases. This can be seen as the effective repairing of the graphitic lattice in the reduction
process [61] as the solution becomes more alkaline. This concurs with the ID/IG ratios from the
0 500 1000 1500 2000 2500 3000 3500 4000
Intensity (a.u.)
Wavenumber (cm
-1
)
GO
rGO-1
rGO-4
rGO-7
rGO-10
rGO-13
113
Raman spectra, where the ratio decreases as pH increases showing larger graphitic sp
2
domains
and less disorder in more basic pH reduction environments. Overall, the sodium borohydride is
successful at removing much of the C-O groups during the reduction at all pH values in the solution,
however, it seems that the acidity of the solution hinders the ability to repair the sp
2
lattice resulting
in smaller overall graphitic domains.
Surface area determination of the catalysts was also carried out using aqueous Methylene
Blue (MB) solutions to adsorb onto the surface of the catalysts [62,63] and are reported in Table
3.2. The rGO-10 catalyst displayed the highest surface area, whereas the surface area decreases
as the pH of the reduction solution decreases. This trend closely mimics the D-spacing trend of
the catalysts in the XRD measurements.
Figure 3.9. Raman spectra of the rGO catalysts.
1000 1200 1400 1600 1800 2000
Intensity (a.u.)
Raman Shift (cm
-1
)
GO
rGO-13
rGO-10
rGO-7
rGO-4
rGO-1
D
G
114
Figure 3.10. Sumary of the Raman spectra of the rGO catalysts from Figure 3.9.
3
3.5
4
4.5
5
0.5
0.75
1
1.25
1.5
0 5 10 15
In-Plane Grain Size (nm)
I
D
/I
G
pH
Id/Ig
Grain Size
280 285 290 295 300
CPS
Binding Energy (eV)
Raw Data
C=C
C-C
C-O
C=O
O-C-OH
280 285 290 295 300
CPS
Binding Energy (eV)
Raw Data
C=C
C-C
C-O
C=O
pi-pi*
b)
a)
115
Figure 3.11. C1s XPS Spectra of the rGO catalysts: a) GO; b) rGO-1; c) rGO-4; d) rGO-7; e)
rGO-10 and; f) rGO-13.
Table 3.1. Summary of the deconvoluted C1s XPS spectra of the rGO catalysts.
GO rGO-1 rGO-4 rGO-7 rGO-10 rGO-13
C=C 10.48% 29.33% 32.90% 34.52% 34.10% 28.85%
C-C 43.82% 42.80% 39.76% 38.74% 31.66% 17.61%
C-O 37.83% 21.93% 20.88% 13.36% 23.96% 32.79%
C=O 6.13% 4.01% 3.60% 6.19% 7.43% 14.18%
π-π* 1.74% 1.93% 2.86% 7.20% 2.86% 6.57%
sp
2
/sp
3
0.24 0.69 0.83 0.89 1.07 1.63
280 285 290 295 300
CPS
Binding Energy (eV)
Raw Data
C=C
C-C
C-O
C=O
pi-pi*
280 285 290 295 300
CPS
Binding Energy (eV)
Raw Data
C=C
C-C
C-O
C=O
pi-pi*
280 285 290 295 300
CPS
Binding Energy (eV)
Raw Data
C=C
C-C
C-O
C=O
pi-pi*
280 285 290 295 300
CPS
Binding Energy (eV)
Raw Data
C=C
C-C
C-O
C=O
pi-pi*
c) d)
e)
f)
116
The pKa of the oxygen containing functional groups were also analyzed via acid-base
potentiometric titrations. The catalysts and the blank were dissolved in a NaOH solution and
titrated with HCl. The titration curves of the catalysts and the blank are shown in Figure 3.12a.
The concentration of the ionized species were calculated from subtracting the amount of HCl added
at each of pH values and are plotted in Figure 3.12b. Figures 3.13a-e were derived from
differentiating the curves in Figure 3.12b with the changing pH values and were deconvoluted with
Gaussian distributions to show the pKa values of the acid groups on the catalysts [64]. Three pKa
peaks arise at around 4.0, 6.2, and 10, signifying two carboxylic acid groups and one hydroxide
group respectively in all of the catalysts except rGO-13 [65]. In rGO-13, a fourth peak arises at
around 7.4 most likely signifying a carboxylate ion present [66]. The hydroxide group arises from
the rGO in alkaline solution [67]. Overall, the pKa of the ionized species from the functional
groups on the graphene lattice of the catalysts is minimally affected by the different pH values in
the reduction.
0
2
4
6
8
10
12
14
0 2000 4000 6000 8000 10000 12000
pH
Amount of HCl Added ( μL)
Blank
rGO-1
rGO-4
rGO-7
rGO-10
rGO-13
Δ
a)
117
Figure 3.12. a) Titration curve of the rGO catalyst solutions in 0.1 M NaOH and; b)
concentration of ionized species in the rGO catalysts at certain pH values.
0
0.03
0.06
0.09
0.12
0.15
0 2 4 6 8 10 12 14
Concentration of Ionized Species (mol L
-1
)
pH
rGO-1
rGO-4
rGO-7
rGO-10
rGO-13
b)
0
0.5
1
1.5
2
2 4 6 8 10 12
f(pKa)
pKa
Exp Data
pKa 1
pKa 2
pKa 3
0
0.25
0.5
0.75
1
2 4 6 8 10 12
f(pKa)
pKa
Exp Data
pKa 1
pKa 2
pKa 3
a)
b)
118
Figure 3.13. The pKa distributions of the acid groups on the rGO catalysts from the
concentrations in Figure 3.12 with Gaussian distribution curves for: a) rGO-1; b) rGO-4; c)
rGO-7; d) rGO-10 and; e) rGO-13.
A regular CV of the catalysts under argon at 20 mV s
-1
in the alkaline media is shown in
Figure 3.14. The relative electrochemically active surface area (ECSA) of the catalysts in solution
can be found in the double layer capacitance region around -0.1 V to 0.3 V in the CV. This is due
to the sorption of the hydroxide groups onto the graphitic carbon lattice edges in the alkaline media
[67]. The active surface area appears to decrease as the pH value of the solution of reduction
0
0.25
0.5
0.75
2 4 6 8 10 12
f(pKa)
pKa
Exp Data
pKa 1
pKa 2
pKa 3
0
0.5
1
1.5
2
2.5
3
2 4 6 8 10 12
f(pKa)
pKa
Exp Data
pKa 1
pKa 2
pKa 3
0
0.5
1
1.5
2 4 6 8 10 12
f(pKa)
pKa
Exp Data
pKa 1
pKa 2
pKa 3
pKa 4
c)
d)
e)
119
decreases. This trend mimics the one found with the surface area determinations with the MB
adsorption.
Figure 3.15 shows LSV scans of the catalysts at 1600 RPM in the alkaline media. The
current densities at -0.5 V vs. MMO as well as onset potentials are shown in Figure 3.16. The
ORR kinetics are enhanced from the reduction process, when the rGO catalysts are compared to
GO. Furthermore, the ORR activity tends to decrease as the pH value of the reduction solution
decreases. This could be due to the smaller active surface areas and less ORR active sites on the
catalysts as shown in Figure 3.14. The greater strain could play a role leading to lower current
densities in the rGO-7 and rGO-4 catalysts. Another trend in the LSV curves among the catalysts
is the onset potential of the catalysts increases as the pH of the reducing solution increases. This
could also be the result of the increased repair of the rGO catalyst and the increase in surface edges
from the higher active surface areas. This could result from the decrease in repair of the sp
2
graphitic lattice as was shown in the XPS and Raman spectra.
There have been conflicting results pertaining to the mechanism and nature of the reactive
sites for undoped graphene structures for ORR. Some studies have concluded that the sp
2
hybridized carbon edges of the graphene are the most active for oxygen reduction in the alkaline
media with zig-zag edges being more ORR active than the armchair region [68,69]. Other studies
have concluded that the functional groups on the graphene sheets such as carbonyl and quinones
are actually responsible [70,71]. From the results of the XPS spectra in Figures 3.11 and Table
3.1 along with the pH titration results in Figures 3.13a-e, the trends from our half-cell
electrochemical tests would indicate the former explanation to be applicable in this case. The XPS
spectra displayed no direct correlation between the pH of the reducing solution and the amount of
surface functional groups found on the rGO catalysts. The pKa titration curves also displayed
120
minimal difference in the surface functional groups on the rGO catalysts except for the rGO-13.
There seems to be a correlation between the extend of repair of the graphene sheets, which is
contrary to the amount of functional groups present and ORR performance. In such a case, the sp
2
carbon edges would seem to be the active sites for ORR. The pH of the solution in the pre-
reduction could possibly have an effect of breaking apart the graphene sheets.
Previous studies have shown that the half-life of sodium borohydride decreases in an order
of magnitude with each decreasing pH unit from the equation [72]:
Log10t1/2(mins) = pH-(0.034T-1.92) (10)
where t1/2 is the half-life of the sodium borohydride in minutes and T is the temperature of the
solution in Kelvin. This could be a reason that there is more repair to the graphitic lattice at higher
pH values: the sodium borohydride has a greater concentration and time to repair the graphitic
lattice before it decomposes. The longer lifetime of the sodium borohydride, the smaller the rGO
sheets are with increased surface area as demonstrated from the MB adsorption tests and CVs.
A Koutecky-Levich plot was made from the current densities of the LSVs at a potential of
-0.5 V vs MMO and is shown in Figure 3.17. The slopes from the plot were taken and the electron
transfer values are reported in Table 3.2, calculated from the Koutecky-Levich equation discussed
in Section 2.2.2. The electron transfer numbers for all the catalysts are all close to 2 signifying a
2 electron transfer process resulting in mainly the peroxide anion being produced. This is typical
of other undoped, bare rGO catalysts previously studied in an alkaline medium [71,73].
Tafel curves of the catalysts were taken and shown in Figure 3.18 and the Tafel slopes
reported in Table 3.2. There are two main different slope values for the catalysts. For rGO-13,
rGO-10, and rGO-7 the slope is in the low to mid 80 mVs dec
-1
. These slopes are similar to ones
reported for doped graphene [74] and platinum [75]. For GO, rGO-4, and rGO-1 the slope is in
121
the lower 100 mVs dec
-1
which is similar to other rGO catalysts reported elsewhere [76]. The
overpotentials in the Tafel plots for rGO-13, rGO-10, and rGO-7 are also around the almost same
value, whereas the overpotentials for GO and rGO-4 are exact and the overpotential for rGO-1 is
the highest. Chronoamperommetry measurements were obtained and are shown in Figure 3.19.
All of the curves have similar slopes signifying similar long-term oxygen reduction stabilities. The
current densities resemble the ones shown in the LSVs with GO being the lowest and rGO-13
being the highest.
Figure 3.14. CV scans of the rGO catalysts in 0.1 M KOH at 20 mV s
-1
.
-1.2
-0.8
-0.4
0
0.4
0.8
-0.8 -0.6 -0.4 -0.2 0 0.2 0.4 0.6
Current Density (mA cm
-2
)
Potential (V vs. MMO)
GO
rGO-13
rGO-10
rGO-7
rGO-4
rGO-1
122
Figure 3.15. LSV scans of the rGO catalysts in 0.1 M KOH under O2 at 1600 RPM.
Figure 3.16. Summary of the LSV scans of the rGO catalysts from Figure 3.15.
-2.5
-2
-1.5
-1
-0.5
0
-0.75 -0.5 -0.25 0 0.25
Current Density (mA cm
-2
)
Potential (V vs. MMO)
GO
rGO-13
rGO-10
rGO-7
rGO-4
rGO-1
-0.16
-0.12
-0.08
-0.04
0
-1.8
-1.5
-1.2
-0.9
-0.6
-0.3
0
0 5 10 15
Onset Potential (V vs. MMO)
Limiting Current Density (mA cm
-2
)
pH
Limiting Current Density
Onset Potential
123
Figure 3.17. Koutecky-Levich plot of the LSV scans of the rGO catalysts.
Figure 3.18. Tafel curves of the rGO catalysts in 0.1 M KOH under O2 at 1600 RPM.
0
2000
4000
6000
8000
10000
12000
0 0.05 0.1 0.15 0.2 0.25 0.3 0.35
Neg. Inverse Current (-A
-1
)
ω
-1/2
rGO-13
rGO-10
rGO-7
rGO-4
rGO-1
-0.2
-0.15
-0.1
-0.05
0
0.001 0.01 0.1
Potential (V vs. MMO)
Current Density (mA cm
-2
)
GO
rGO-13
rGO-10
rGO-7
rGO-4
rGO-1
124
Figure 3.19. Chronoamperommetry plots of the rGO catalysts in at -0.5 V vs. MMO at 1600
RPM.
Table 3.2. Summary of the characterization and electrochemical tests of the rGO catalysts.
Catalyst D-band
position (cm
-1
)
G-Band
position (cm
-1
)
Surface Area
(m
2
g
-1
)
Electron
Transfer Number
Tafel Slope
(mV dec
-1
)
GO 1349.85 1591.27 381 n/a 102±9
rGO-13 1347.34 1578.70 397 2.34 85±4
rGO-10 1347.34 1578.70 418 2.41 86±1
rGO-7 1344.82 1576.18 286 2.19 81±3
rGO-4 1344.82 1581.21 247 2.21 111±5
rGO-1 1347.34 1578.70 117 2.43 98±15
-2
-1.5
-1
-0.5
0
0 500 1000 1500 2000 2500 3000
(Current Density (mA cm
-2
)
Time (sec)
GO
rGO-13
rGO-10
rGO-7
rGO-4
rGO-1
125
3.3: A Facile One-Pot Synthesis of Platinum on Electrochemically Exfoliated Graphite for ORR
in Acidic Media
3.3.2: Experimental Methods
A simple one-cell was used for the graphite exfoliation and subsequent reduction shown in
Scheme 3.1. 20 mL of a 10 mM aqueous platinum salt solution, (H2PtCl6 •H2O, Alfa Aesar) was
added to a 25 mL beaker. Two graphite rods (Unocal Poco, 0.35 mm diameter) were connected
to an external power supply at the anode and cathode end and submerged in the platinum solution
at a distance of 16 mm from each other. Different potentials of 2.5, 5, 8, and 10 V for 1 hour, 15
minutes, 10 minutes, and 5 minutes, respectively, were applied between the wires causing
exfoliation of the graphite at the anode side. These samples will be referred to at Pt/exrGO-2.5, 5,
8, or 10 for the potential of exfoliation. The rods were then removed and the pH of the solution
adjusted to 12 with 3.0 M NaOH. 5 mL of hydrazine hydrate (Sigma-Aldrich, 35 wt-% in H2O)
was added to facilitate the reduction. After an hour of mixing, the solution was washed and
centrifuged with Millipore water (Direct-Q UV, 18.2 MΩ) until chloride ions were no longer
detectible with silver nitrate and the solid residue was dried in an oven at 80
O
C.
The catalysts were characterized by the same methods as were described in Section 3.2.1.
Half-cell measurements were carried out in a three cell rotating disk electrode (RDE, Pine
Instruments) setup. The catalyst inks were prepared by adding 3 mg of catalyst to a 1 mL solution
of 10% isopropyl alcohol, 90% Millipore water, and 10 mg of 5% Nafion® solution (Ion Power).
The solution was sonicated and 40 μL pipetted onto the glassy carbon electrode of the RDE with
an active area of 0.195 cm
2
. The electrode was dried in an oven and placed in the three cell testing
system with the RDE as the working electrode, a platinum wire as the counter electrode, and a
mercury sulfate (MSE, 0.5 M H2SO4 filling solution) reference electrode. The cell was purged
126
with ultra-pure argon (99.999%) before the CV testing and ultra-pure oxygen (99.994%) before
the oxygen reduction reaction (ORR) tests for 15 minutes each, respectively. The CO stripping
tests were carried out by purging the cell with carbon monoxide (99.5%) for 15 minutes followed
by ultra-pure argon for 30 minutes.
3.3.3: Results
Scheme 3.1 shows the one-pot synthesis procedure for the Pt/exrGO catalysts. The
resulting currents measured during the exfoliation process steadily increased over time and are
reported in Table 3.3. The overall conductivity of the ions present in the platinum salt solution
allowed for the exfoliation to take place similar to various sulfuric and hydrochloric acid solutions
in previous studies [36,40]. It was discovered that 2.5 V was the lowest potential, where noticeable
exfoliation took place in the system since the current readings on the external power supply
registered down to 10 mA. The graphite rods were exfoliated at different times due to the platinum
deposition that occurred at the cathode rod and are reported in Table 3.3. Originally, the rods were
exfoliated at 8 V and 10 V for 15 minutes each, however leading to the exhaustion of platinum in
the solution. Thus, the exfoliation times at the higher potentials were reduced to ensure ample
platinum present in the solution in order to be reduced onto the exfoliated graphene. Due to the
latent heat of the solution from the exfoliation, there was no external heat applied to the solution
during the addition of sodium hydroxide and hydrazine.
Scanning Electron Microscope (SEM) images were obtained and shown in Figures 3.20a-
d. EDAX was also taken and the platinum loadings are reported in Table 3.3. The platinum
nanoparticle sizes for each of the catalysts are around 100 nm with minor variation in size. This
127
is most likely due to the same type of reduction conditions in the synthesis. In the Pt/exrGO2.5
catalyst, there is considerable platinum coverage compared to the exfoliated graphite. This is due
to the high concentration of platinum left in the solution after the exfoliation process in comparison
to the amount of exfoliated graphite. TEM images are shown in Figures 3.21a-d and display an
enhanced resolution of the platinum nanoparticles with a “nanoflower” type morphology [43].
Here, the individual nanoparticles have a size of around 15 nm and are agglomerated in clusters of
around 100 nm.
XRD patterns were taken for the unreduced exfoliated graphite and are shown in Figure
3.22a. In all the samples, except the 2.5 V exfoliated sample, there is a slight (100) peak at around
10.5
O
indicating the presence of graphene oxide. There is also a (002) peak at around 24
O
indicating the presence of reduced graphene oxide. In the 10 V and 8 V samples, there is also a
sharp peak at 26
O
indicating the (002)* index of graphite. This can be due to the strong oxidation
and exfoliation of the graphite sheets at the higher exfoliation potentials [32]. Large intact pieces
of the graphite were removed from the rod at the higher potentials leading to the intact (002)
graphite phase. XRD patterns of the catalysts are shown in Figure 3.23a for the Pt/exrGO catalysts.
In each pattern, there are peaks at 40
O
, 46
O
, 68
O
, 81
O
, and 85
O
signifying the (111), (200), (220),
(331), and (222) Miller Indices of fcc platinum, respectively. There is also a peak at around 24
O
representing the (002) index of the reduced graphene oxide. This peak is small in the 2.5 V
exfoliated catalyst due to the relatively low amount of graphite exfoliated in comparison to
platinum reduced from the solution. The crystallite sizes of the platinum were calculated from the
Scherrer equation [77] and are reported Table 3.3. The Pt/exrGO2.5 displays the largest size most
likely due to the high platinum loading present. Aside from this, the crystallite size increases as
the exfoliation potential increases.
128
Figure 3.22b shows the Raman spectra of the exfoliated graphite. The spectra show two
main peaks at around 1360 cm
-1
and 1590 cm
-1
signifying the D-band and the G-band, respectively
as was discussed in Section 3.2.2. The ID/IG ratios for the exfoliated graphite is 1.05, 1.19, 1.03,
and 1.12 for the graphite exfoliated at 2.5, 5, 8, and 10 V, respectively. There is more disorder in
the graphitic lattice as the potential is increased from 2.5 V to 5 V and from 8 V to 10 V. One
explanation for this is the intact pieces of graphite from the XRD patterns in Figure 3.22a at the
higher potentials could result in less disorder in the overall lattice in the 8 V and 10 V exfoliated
graphite in comparison to the 5 V exfoliated graphite. Figure 3.23b shows the Raman spectra of
the Pt/exrGO catalysts with the ID/IG ratios reported in Table 3.3. As the exfoliation potential is
increased, the ID/IG ratio increases due to the stronger oxidative conditions present on the graphitic
lattice. For the Pt/exrGO10 catalyst, however, it decreases most likely due to the mixture of intact
graphite layers again confirmed by the XRD pattern in Figure 3.22a for the bare exfoliated graphite.
Scheme 3.1. Overview of the one-pot exfoliation/reduction synthesis.
129
Figure 3.20. SEM images of the Pt/exrGO catalysts at: a) 2.5 V; b) 5 V; c) 8 V and; d) 10 V.
a)
b)
c) d)
a)
b)
130
Figure 3.21. TEM images of the Pt/exrGO catalysts at: a) 2.5 V; b) 5 V; c) 8 V; d) 10 V. The
measuring bar indicates 100 nm.
0 10 20 30 40 50
Intensity (a.u.)
2 θ (degrees)
10V
8V
5V
2.5V
c) d)
a)
(100)
(002)
(002)*
131
Figure 3.22. Characterization of the exfoliated graphite rods: a) XRD patterns and; b) Raman
spectra.
0 500 1000 1500 2000 2500 3000 3500
Intensity (a.u.)
Raman Shift (cm
-1
)
2.5V
5V
8V
10V
b)
D G
0 10 20 30 40 50 60 70 80 90
Intensity (a.u.)
2 θ (degrees)
2.5V
5V
8V
10V
a)
(002)
(111)
(200)
(220)
(311)
(222)
132
Figure 3.23. Characterization of the Pt/exrGO catalysts: a) XRD patterns and; b) Raman
spectra.
Table 3.3. Summary of the synthesis and characterization tests of the Pt/exrGO catalysts.
In order to gauge the catalytic properties of the catalysts, CVs were obtained under argon
and displayed in Figure 3.24. The Pt/exrGO5 catalyst displays the highest current activity even
though the Pt/exrGO2.5 contained the highest amount of platinum per weight. CO stripping CV
curves were also performed and shown in Figure 3.25 and are consistent with the regular CV
curves. Again, the catalyst exfoliated at 5 V displays the highest current as well as the highest
electrochemically active surface area (ECSA) from the area under the CO stripping curve as shown
0 500 1000 1500 2000 2500 3000 3500
Intensity (a.u.)
Raman Shift (cm
-1
)
2.5V
5V
8V
10V
b)
Exfoliation
Potential (V)
Exfoliation
Current (A)
Exfoliation
Time (min)
Platinum Loading
(%-wt) from EDAX
Crystallite Size
(nm) from XRD
ID/IG
2.5 0.15-0.17 60 73.55 7.0 0.991922
5 0.8-1.1 15 11.62 6.1 1.294515
8 1.6-2.1 10 22.55 6.1 1.292967
10 2.8-4.0 5 31.57 6.5 1.16503
D
G
133
in Table 3.4. ECSAs for the hydrogen adsorption from the CV scans were also calculated and
reported in Table 3.4 and displayed the same trend as with the CO stripping. One reason for the
high ECSA in the Pt/exrGO5 catalyst could be the presence of both the GO and rGO from the
XRD patterns leading to a higher surface area for the platinum to be reduced onto compared with
the intact graphite layers of the higher potentially exfoliated catalysts. LSVs of the catalysts under
oxygen are shown in Figure 3.26 at 1600 RPM with the onset potential and current density values
reported in Table 3.4. Again, the catalyst exfoliated at 5 V displays superior onset potential and
current density values while the catalyst exfoliated at 8 V displays the lowest current densities and
onset potentials. The catalysts exfoliated at the lower potentials overall (2.5 and 5 V) displayed
higher activity than the catalysts exfoliated at the higher potentials (8 and 10 V). This could
possibly arise from the structure of the exfoliated graphite supports. From the XRD patterns in
Figure 3.22a, the graphite exfoliated at the higher potentials still have the graphitic phase along
with the graphene oxide and reduced graphene oxide phase whereas the graphite exfoliated at the
lower potentials only has the graphene oxide and reduced graphene oxide phase present.
134
Figure 3.24. CV scan of the Pt/exrGO catalysts under argon.
Figure 3.25. CO Stripping of the Pt/exrGO catalysts.
-1
-0.5
0
0.5
0 0.2 0.4 0.6 0.8 1 1.2 1.4
Current Density (mA cm
-2
)
Potential (V vs. RHE)
2.5V
5V
8V
10V
-1.5
-1
-0.5
0
0.5
1
1.5
0 0.2 0.4 0.6 0.8 1 1.2 1.4
Current Density (mA cm
-2
)
Potential (V vs. RHE)
2.5V
5V
8V
10V
135
Figure 3.26. LSV of the Pt/exrGO catalysts under O2 at 1600 RPM.
Table 3.4. Summary of the electrochemical tests of the Pt/exrGO catalysts.
Exfoliation
Potential (V)
ECSA CO striping
(m
2
g
-1
Pt)
ECSA H-
adsorption (m
2
g
-1
Pt)
Onset Potential
(V vs. RHE)
Current Density
(mA cm
-2
)
2.5 2.94254 1.784597 0.774±0.026 2.13±0.47
5 23.24208 21.96951 0.798±0.013 2.89±0.19
8 2.550934 3.31834 0.710±0.037 1.05±0.46
10 1.291322 1.68537 0.766±0.018 1.98±0.20
-2.5
-2
-1.5
-1
-0.5
0
0 0.2 0.4 0.6 0.8 1
Current Density (mA cm
-2
)
Potential (V vs. RHE)
2.5V
5V
8V
10V
136
3.4: Reduced Graphene Oxide Emulsions as a Catalyst Support for Various Metals for Fuel Cell
Catalysts
3.4.1: Experimental Methods
Graphene oxide emulsions (GOem) were synthesized via the Hummer’s method as was
described in Section 3.2.1. In this instance, the graphene oxide was not dried in an oven after the
washing step but kept in a sealed container as an emulsion. Dried graphene oxide as synthesized
in Section 3.2.1 was used in comparison.
Various reduction methods were performed on both the graphene oxide emulsions and
dried graphene oxide powder. For the sodium borohydride-based reduction, 5 mL of the graphene
oxide emulsion (40 mg dry graphene oxide) was stirred and heated to 80
O
C in an oil bath. 1 g of
sodium borohydride powder was then added to the solution and stirred for 1 hour at 80
O
C. It was
then washed and centrifuged with Millipore water four times until the pH of the supernatant
reached 7. The dried graphene oxide was also reduced the same manner. Briefly, 40 mg of the
graphene oxide powder added to 5 mL Millipore water and was vigorously stirred and sonicated.
The solution was then heated up to 80
O
C in an oil bath where 1 g sodium borohydride powder
was added and the solution stirred for 1 hour at 80
O
C followed by washing with Millipore water
and drying in an oven. The platinum supported catalysts were synthesized in a similar manner.
Briefly, 2.65 mL of H2PtCl6 solution (1 g/100 mL Millipore water) was added to 5 mL of the
graphene oxide emulsion (20%-wt Pt loading). The solution was stirred and sonicated and heat to
80
O
C where 1 g sodium borohydride was added. It was stirred at 80
O
C for 1 hour and then
washed and centrifuged with Millipore water before being dried in an oven. The catalysts were
137
also synthesized with 40 mg of dried graphene oxide powder in 5 mL of Millipore water with the
same amounts of platinum solution added as was previously described.
The catalysts were also synthesized via the microwave-assisted polyol reduction method.
5 mL of the graphene oxide emulsion solution was added to a 20 mL microwave reaction vial
(Biotage). The same amounts of the platinum solutions as previously discussed were then added
followed by 5 mL of ethylene glycol and the solution was stirred and sonicated. 0.5 mL of 3 M
aqueous sodium hydroxide solution was then added to adjust the pH to 12. The vial was then
sealed and placing in a microwave reactor (Biotage) where it was heated up to 180
O
C for 3
minutes. The solution was then washed and centrifuged with ethanol four times and dried in an
oven. The catalysts were also synthesized with 40 mg dried graphene oxide powder in 5 mL
Millipore water with the same amounts of platinum solution added as was previously described.
The characterization methods were carried out as described in Section 3.2.1. The
electrochemical half-cell tests were carried out in the same manner as described in Section 3.3.1
for the ORR.
3.4.2: Results
The impregnation-reduction of the graphene oxide compounds were carried out with
minimum amount of water in order to study the effects of the drying step for Hummer’s method.
This was to ensure minimal solvent effects for both the GO and GOem. In this respect, the only
DI water added to the dried catalyst was equivalent to the volume of the GOem for the
corresponding weight. The emulsion was determined to be 19%-wt GO from drying tests in an
oven at 110
O
C.
138
SEM images were taken for the bare GO catalysts and are shown in Figures 3.27a&b.
Similar structures for the GO and GOem catalysts are evident with long wrinkled sheets. There
appears to be more wrinkles in the dried GO catalyst possibly due to the extra drying step carried
out during the synthesis. The images for the reduced catalysts are shown in Figures 3.27c&d. In
this instance, there appears to be more wrinkles in the rGOem catalyst in comparison. This could
be due to the harsh reduction conditions of the sodium borohydride. SEM images were also taken
for the platinum supported catalysts with the impregnation method shown in Figures 3.28a&b and
the microwave assisted polyol method in Figures 3.28c&d. Both sets of images display similar
characteristics in both the rGO structure for each respective reduction as well as platinum particle
size and morphology. The rGO supports synthesized via the impregnation method display a more
wrinkled structure with the individual sheets more visible, especially with the rGOem, much
similar to that of the bare reduction. The microwave assisted catalysts however display a more
“block”-like structure indicating that the microwave assisted polyol reduction might not have been
efficient in dispersing the rGO sheets and reducing the platinum, like that of the impregnated
synthesis.
a)
b)
139
Figure 3.27. SEM images of the rGOem catalysts: a) GO; b) GOem; c) rGO and; d) rGOem.
Figure 3.28. SEM images of the Pt/rGOem catalysts: a) Pt/rGO impreg; b) Pt/rGOem impreg;
c) Pt/rGO uwave and; d) Pt/rGOem uwave.
c)
d)
a)
b)
c)
d)
140
XRD patterns for both the bare rGOem and Pt/rGOem catalysts are shown in Figure 3.29a
and Figure 3.29b, respectively. In Figure 3.29a, there is one peak visible for the GOem catalysts
at 12.54
O
and 11.52
O
signifying the (100) Miller Index peak for the GO and GOem catalysts,
respectively. The d-spacing corresponding to the peaks is reported in Table 3.5. The d-spacing
for the GOem is greater than for the GO most likely due to the dispersive nature of the emulsion
as well as water molecules being trapped due to the polarity of the surface groups in the GO sheets
[78]. After drying, the GO sheets stack closer together. After the reduction, the peak shifts to
24.08
O
and 23.76
O
signifying the (002) Miller Index peak for the rGO and rGOem catalysts,
respectively. The d-spacing corresponding to the peaks is reported in Table 3.5. The d-spacing is
considerably more similar between the reduced catalysts compared to that of the unreduced
catalysts. Still, the d-spacing of the emulsion catalyst is still greater than that of the dried catalyst.
Figure 3.30b shows the XRD patterns for the platinum supported catalysts. Each pattern has the
representative peaks for fcc platinum as was discussed in Section 2.3.2 with the platinum crystallite
sizes calculated from the Scherrer equation also discussed in Section 2.3.2 and are reported in
Table 3.6. The crystallite sizes are the same for the corresponding reduction method signifying
minimal differences between the rGO and rGOem supports for platinum reduction. There is also
a peak at 22
O
signifying the (002) index of the rGOem as previously discussed. The exact 2θ
values and corresponding d-spacing is reported in Table 3.6. The d-spacing for the rGOem catalyst
is higher than the d-spacing for the rGO catalyst which mimics the trend in the d-spacing observed
in the bare GO and rGO catalysts.
In order to gauge the relative disorder of the rGOem catalysts, Raman spectra were taken
and shown in Figure 3.30a for the bare rGOem catalysts and Figure 3.30b for the platinum
supported rGOem catalysts. As was discussed in Section 3.2.2, there are two main peaks, the d-
141
band and g-band, present. The d-and g-band peak positions as well as the ID/IG ratios for the
rGOem catalyst are reported in Table 3.5. For the GOem catalyst, the d-band peak location is
slightly blue shifted from the regular GO catalyst signifying a slight strain of the graphitic lattice
in comparison. The g-band location for the rGOem catalyst is slightly blue shifted, signifying a
slight strain in the ordered graphitic lattice in comparison. This trend is displayed in the Pt/rGOem
catalysts for both reduction methods. Another aspect of the Raman spectra is the ID/IG ratios of
the catalysts. The ratio is higher in the GOem compared with the GO. This could arise from more
wrinkles in the sheets as was shown in the SEM images. After the reduction, the ratio for GOem
is lower signifying a more complete repair of the graphitic lattice. This could be due to the
increased d-spacing present and the properties of the emulsion itself, allowing for better reduction
kinetics of the sodium borohydride. For each of the platinum catalysts, however, the ratio is higher
for the Pt/rGOem catalysts than the respective Pt/rGO.
10 20 30 40 50 60
Intensity (a.u.)
2 θ (degrees)
GO
GOem
rGO impreg
rGOem impreg
a)
(100)
(002)
142
Figure 3.29. XRD patterns of the rGOem catalysts: a) rGOem and; b) Pt/rGOem.
10 20 30 40 50 60 70 80 90
Intensity (a.u.)
2 θ (degrees)
Pt/rGO impreg
Pt/rGOem impreg
Pt/rGO uwave
Pt/rGOem uwave
b)
(002)
(111)
(200)
(220)
(311)
(222)
0 500 1000 1500 2000 2500 3000 3500
Intensity (a.u.)
Raman Shift (cm
-1
)
GO
GOem
rGO impreg
rGOem impreg
a)
D
G
143
Figure 3.30. Raman spectra of the rGOem catalysts: a) rGOem and; b) Pt/rGOem.
Figures 3.31a&b show the XPS C1s spectra for the rGOem catalysts while Figures 3.31c&d
show the XPS C1s spectra for the Pt/rGOem catalysts with the (1) spectra indicating the GOem
catalysts and the (2) spectra indicating the GO catalysts. Similar to the results displayed in Figures
3.11a-f, the GO catalysts display two main peaks, deconvoluted into the C=C/C-C peak and the
C-O peak. After the reduction, the C-O is significantly reduced. In all of the peaks, the GOem
support displays a greater sp
2
/sp
3
carbon ratio compared to the GO support, signifying a greater
C=C intact lattice as was discussed in Section 3.2.2. Another trend seen is the C-O peak for the
GOem is larger signifying a greater presence of C-O species in the graphitic lattice. Some of the
oxygen species present could be removed from the graphitic lattice during the drying step. The
XPS Pt 4f spectra for the Pt/rGOem catalysts are displayed in Figures 3.32 a&b with the (1) spectra
indicating the GOem catalysts and the (2) spectra indicating the GO catalysts. In each spectra,
there are two peaks evident as was discussed in Section 2.2.2. In both types of reduction, the Pt
0 500 1000 1500 2000 2500 3000 3500
Intensity (a.u.)
Raman Shift (cm
-1
)
Pt/rGO impreg
Pt/rGOem impreg
Pt/rGO uwave
Pt/rGOem uwave
b) D
G
144
(II) 7/2 peaks are larger in the regular Pt/rGO catalysts compared to the Pt/rGOem catalysts. This
is due to the increased removal of the oxygen containing functional groups in the rGO and the
anchoring of metallic Pt
O
species on the catalyst surface [79,80]. This could explain the increase
in the ID/IG ratio from the Raman spectra with the Pt/rGOem compared to the Pt/rGO catalysts.
Figure 3.31. XPS C1s spectra of the rGOem catalysts: a) GOem; b) rGOem; c) Pt/rGOem
impreg and; d) Pt/rGOem uwave. The (1) spectrum indicates GOem supported catalysts; the (2)
spectrum indicates the regular GO supported catalysts.
275 280 285 290 295
Intensity (a.u.)
Binding Energy (eV)
Raw
C=C
C-C
C-O
C=O
-COOH
Envelope
(1)
(2)
a)
275 280 285 290 295
Intensity (a.u.)
Binding Energy (eV)
Raw
C=C
C-C
C-O
C=O
-COOH
Envelope
(1)
(2)
b)
275 280 285 290 295
Intensity (a.u.)
Binding Energy (eV)
Raw
C=C
C-C
C-O
C=O
-COOH
Envelope
(1)
(2)
c)
275 280 285 290 295
Intensity (a.u.)
Binding Energy (eV)
Raw
C=C
C-C
C-O
C=O
-COOH
Envelope
(1)
(2)
d)
145
Figure 3.32. XPS Pt 4f spectra of the Pt/rGOem catalysts: a) Pt/rGOem impreg and; b)
Pt/rGOem uwave. The (1) spectrum indicates GOem supported catalysts; the (2) spectrum
indicates the regular GO supported catalysts.
Table 3.5. Summary of the characterization tests of the rGOem catalysts.
Catalyst 2θ d-spacing
(nm)
d-band
(cm
-1
)
g-band
(cm
-1
)
ID/IG
GO 12.54 0.704 1347.34 1591.44 .984
GOem 11.52 0.768 1352.37 1591.27 .996
rGO 24.08 0.369 1349.85 1583.73 1.223
rGOem 23.76 0.374 1349.85 1584.24 1.179
Table 3.6. Summary of the characterizations of the Pt/rGOem catalysts.
Catalyst 2θ d-spacing
(nm)
Pt Crystallite
Size (nm)
d-band
(cm
-1
)
g-band
(cm
-1
)
ID/IG
Pt/rGO impreg 24.16 0.368 5 1347.34 1576.18 1.129
Pt/rGOem impreg 24.08 0.369 5 1349.85 1588.75 1.131
Pt/rGO uwave 24.08 0.369 6 1347.34 1591.23 1.106
Pt/rGOem uwave 23.76 0.374 6 1349.85 1596.3 1.147
65 70 75 80
Intensity (a.u.)
Binding Energy (eV)
Raw
Pt (O) 4f 7/2
Pt (O) 4f 5/2
Pt (II) 4f 7/2
Pt (II) 4f 5/2
Envelope
(1)
(2)
a)
65 70 75 80
Intensity (a.u.)
Binding Energy (eV)
Raw
Pt (O) 4f 7/2
Pt (II) 4f 7/2
Pt (O) 4f 5/2
Pt (II) 4f 5/2
Envelope
b)
(1)
(2)
146
Figure 3.33 displays the electrochemical CV scans of the Pt/rGOem catalysts in 0.5 M
H2SO4. The minor peak at around 0.4 V is due to the surface oxygen groups present in the rGO
support which is more evident in the emulsion catalysts [81]. Overall, the size of the curves for
the impregnated catalyst displays a larger activity and a more defined “typical” platinum CV shape
compared to that of the microwave assisted catalysts. Furthermore, the rGOem supported catalysts
display a higher current compared to the dry catalysts. This could be due to the more dispersed
platinum nanoparticles on the supports. In order to further analyze the ECSA of the platinum, CO
stripping CVs were taken and displayed in Figure 3.34. The ECSA values for the CO stripping
were calculated in the same manner as in Section 2.2.2 and are reported in Table 3.7. The
impregnated catalysts display higher ECSA values than the microwave assisted catalysts. In
addition, the rGOem supported catalysts display a higher ECSA than the microwave reduced
catalysts. This could result from the increased dispersion of the platinum nanoparticles present in
the rGOem catalysts as well as the enhanced rGO sheet dispersion from the reduction as displayed
in the SEM images. This could also cause the higher ID/IG ratios in the Pt/rGOem catalysts in
comparison.
The ORR activity was analyzed by LSVs shown in Figure 3.35 with the onset potentials
and limiting current densities reported in Table 3.7. For the impregnated catalysts, the onset
potential is very similar indicating minimal change in the synergistic effects of the rGO support to
the platinum nanoparticles on the ORR activity. The microwave reduced catalysts displayed
different onset potentials indicating a change in the synergistic effects of the rGO support with the
platinum nanoparticles. The limiting current densities are also lower than the impregnated
catalysts possibly due to the overall structure of the catalyst as seen in the SEM images in Figures
3.2c&d. The “block”-like structure of the rGO support could be impeding the reactant diffusion
147
to the platinum nanoparticles as well as decrease the ECSA as was displayed in the CO stripping
CVs in Figure 3.34. Koutecky-Levich plots were also created from the LSV curves for the
Pt/rGOem catalysts shown in Figure 3.36. The average electron transfer for each of the catalysts
was calculated from the Koutecky-Levich equation discussed in Section 2.2.2. For all of the
catalysts, the average electron transfer was all around 4, which has been shown to be typical of the
platinum catalyst in acidic media.
Tafel plots were also created and shown in Figure 3.37 with the Tafel slopes reported in
Table 3.7. As with the Koutecky Levich plot, there are two types of Tafel plots shown for the
catalysts: one for the impregnated catalysts and one for the microwave assisted catalysts. The
Tafel slopes for the impregnated catalysts are lower than the ones for the microwave assisted
catalysts signifying lower intrinsic kinetic activity at lower currents.
Figure 3.33. CV scans of the Pt/rGOem catalysts in 0.5 M H2SO4 under argon.
-6
-4
-2
0
2
4
6
0 0.2 0.4 0.6 0.8 1 1.2 1.4
Current Density (mA cm
-2
)
Potential (V vs. RHE)
Pt/rGO impreg
Pt/rGOem impreg
Pt/rGO uwave
Pt/rGOem uwave
148
Figure 3.34. CO Stripping scans of the Pt/rGOem catalysts.
Figure 3.35. LSV curves of the Pt/rGOem catalysts in 0.5 M H2SO4 under O2 at 1600 RPM.
-3
0
3
6
0 0.2 0.4 0.6 0.8 1 1.2 1.4
Current Density (mA cm
-2
)
Potential (V vs. RHE)
Pt/rGO impreg
Pt/rGOem impreg
Pt/rGO uwave
Pt/rGOem uwave
-5
-4
-3
-2
-1
0
0 0.2 0.4 0.6 0.8 1 1.2
Current Density (mA cm
-2
)
Potential (V vs. RHE)
Pt/rGO impreg
Pt/rGOem impreg
Pt/rGO uwave
Pt/rGOem uwave
149
Figure 3.36. Koutecky-Levich plot of the Pt/rGOem catalysts.
Figure 3.37. Tafel curves of the Pt/rGOem catalysts in 0.5 M H2SO4 under O2 at 1600 RPM.
0
500
1000
1500
2000
2500
3000
3500
0 0.05 0.1 0.15 0.2
Neg. Inverse Current (-A
-1
)
ω
-1/2
Pt/rGO impreg
Pt/rGOem impreg
Pt/rGO uwave
Pt/rGOem uwave
0.7
0.8
0.9
1
0.001 0.01 0.1 1
Potential (V vs. RHE)
Current Density (mA cm
-2
)
Pt/rGO impreg
Pt/rGOem impreg
Pt/rGO uwave
Pt/rGOem uwave
150
Table 3.7. Summary of the electrochemical tests of the Pt/rGOem catalysts.
Catalyst CO Strip
ECSA (m
2
g
-1
)
Onset Potential
(V vs. RHE)
Limiting Current
(mA cm
-2
)
Tafel Slope
(mV dec
-1
)
Pt/rGO impreg 53.16 0.917±0.023 4.66±0.29 64±5
Pt/rGOem impreg 78.14 0.917±0.005 4.52±0.27 66±4
Pt/rGO uwave 15.08 0.831±0.010 2.43±0.23 108±4
Pt/rGOem uwave 20.87 0.906±0.015 3.51±0.51 84±6
3.5: Conclusions
Graphene is another viable catalyst support that can be used for various electrochemical
reactions for fuel cells. The versatility of graphene makes it an attractive compound that can be
easily modified and tuned to serve a wide range of functions. The various techniques to synthesize
and prepare graphene can affect its overall structure and kinetic properties as a catalyst for various
reactions. In Section 3.2.2, the pH of the aqueous-based reduction was shown to affect the
reducing agents ability to repair the carbon lattice of graphene. This led to a change in the overall
catalytic activity of the rGO for oxygen reduction. The rGO catalysts synthesized in solutions at
higher pH values showed a better repair of the sp
2
graphitic lattice and a smaller lattice size with
decreasing pH values. This enhanced repair results in increased ORR kinetics even though the
catalysts synthesized at the lower the pH values display smaller particle sizes and more edges,
however, there is lesser repair of the lattice with defects most likely around the edges. Much work
has been done to optimize stand-alone graphene for use as an efficient ORR catalyst but there is
still more needed if it is to be a viable alternative to replace many noble metals for use in fuel cells
in the future.
151
Another aspect of catalyst synthesis is the facile or non-facile nature of the synthetic route.
If a catalyst is to be commercialized, the fewer steps for synthesis, the more viable it is for large
scale synthesis. In numerous studies, catalysts that display enhanced activity or often synthesized
in numerous steps, impedes on the feasibility for large scale manufacturing and commercialization.
In Section 3.3.2, platinum on graphene was synthesized in a facile one pot step from the exfoliation
of inexpensive graphite rods. Although the platinum particle size was rather large in comparison
to previous ones reported in the literature, the size can most likely be tuned by adding surfactants
or other solvents such as ionic liquids in the platinum solution. This would also work to affect the
size and properties of the exfoliated graphite, further changing the catalytic activity. Many other
aspects of this synthesis can also be carried out for catalyst optimization.
Along with the commercial viability of the synthetic route is the reduction of the steps used
to synthesize the catalyst. The drying process associated with Hummer’s method is often arduous,
either requiring many days to dry or to be dried through a dialysis bag among other methods. In
Section 3.4.2, the drying step of Hummer’s method was forgone and the characteristics of the rGO
supports with platinum were investigated. Overall, the emulsion catalysts displayed similar ORR
kinetics compared to the dried GO catalysts signifying the drying process that is employed in the
Hummer’s method could be eliminated. This in effect could cut synthesis time and costs for the
commercialization of graphene or metal supported graphene catalysts.
The use of graphene as a catalyst support was further investigated in Chapter 4 as a support
for various nickel catalysts for urea oxidation in alkaline media.
152
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159
Chapter 4:
The Modification of Nickel and Supports for Applications in Urea Oxidation
4.1: Urea Oxidation for Use in Direct Urea Fuel Cells
One particular energy generation method is from fuel cells utilizing liquid-based reactants
such as formic acid [1,2], methanol [3–5], and hydrazine [6]. Compared to fuel cells that employ
gas-based reactants, liquid-based fuel cells can efficiently store and transport reactants with
minimal changes required to the existing infrastructure [7]. One particular type of fuel that has
recently attracted attention for use in liquid-based fuel cells is solid urea, which is soluble in
aqueous medium. Urea is most commonly found in human and animal waste at around 2 wt% [8],
and is currently being manufactured in large quantities from ammonia as a fertilizer. Its high
energy density (16.9 MJ L
-1
) and relatively high hydrogen content (10.1% weight) [8,9], make it
an enticing candidate as both a potential hydrogen carrier source and fuel.
The Direct Urea Fuel Cell (DUFC) has attracted much attention as a potential new type of
fuel cell for the oxidation of urea under alkaline conditions. The redox reactions are shown below:
[8,10,11]
Anode: CO(NH2)2 + 6OH
-
→ N2 + 5H2O + CO2 + 6e
-
E
O
= -0.746V (1)
Cathode: 2H2O + O2 + 4e
-
→ 4OH
-
E
O
= +0.40V (2)
Overall: CO(NH2)2 + H2O → N2 + 3H2 + CO2 E
O
= +1.146V (3)
One of the attractive features of DUFCs is the ability to use non-noble metal catalysts in both the
anode and cathode sides such as nickel, cobalt, and manganese. This differs from some of the
other acid-based fuel cells such as Direct Methanol Fuel Cells (DMFCs), which require platinum
and ruthenium [3,12], and Direct Formic Acid Fuel Cells (DFAFCs), which require palladium at
160
the anode [13]. Moreover, the alkaline environment in which DUFCs are run in do not require the
use of platinum for the oxygen reduction reaction (ORR) at the cathode as was discussed in Section
1.6.2 [14].
There are some drawbacks to alkaline fuel cells as was discussed in Section 1.4.2. such as
the anion exchange membranes (AEMs), which are not as optimized and commercially
manufactured as the present proton exchange membranes (PEMs), as well as carbonate build up
in the anode compartment during CO2 release [5,15]. More recently, membraneless micro fuel
cells have been studied to eliminate the use of a membrane [16]. Paper-based fuel cells have
recently emerged as a more cost effective alternative to traditional membrane electrode assembly
(MEA)-based fuel cells for testing catalysts [17]. The intrinsic properties of paper/porosity and
capillary action remove the need for external pumps and membranes. Moreover, the use of paper
allows for facile manipulation of the cell design. Another advantage is the use of small volumes
of fuel, allowing for urine to be a more practical source of urea. This can further be implemented
for urea sensors in urine testing. A more recent form of the DUFC is the direct urea/hydrogen
peroxide fuel cell [18–20]. The hydrogen peroxide is reduced at the cathode compartment from
the following equation [20]:
H2O2 + 2e
-
→ 2OH
-
(4)
This type of fuel cell has just recently started to be studied for the urea oxidation performance and
displays much promise.
Nickel has been demonstrated in numerous studies to be the prime catalyst for urea
electrooxidation. The structure and morphology of nickel can considerably affect the
electrooxidation performance of urea. Some of these forms include nickel metal nanoparticles
[9,21], oxides [22,23], hydroxides [24,25], perovskites [26], as well as varying structural
161
morphologies including foams [27,28] and nanowires [29,30]. Nickel oxides have been shown to
be useful for a variety of other applications such as supercapacitors [31], lithium-ion batteries [32],
and photocathodes [33]. One of the ways to enhance the oxidation properties of nickel is to alloy
it with various other elements such as platinum group metals [34,35], manganese [36], iron [37,38],
and most often cobalt [39–41]. There are still a vast number of elements on the periodic table that
can be screened with nickel with hopes of further enhancing the urea oxidation properties. Another
way to enhance the activity of metal catalysts such as nickel is by employing a catalyst support.
Some of the catalyst supports that are currently being explored are carbon nanotubes [42,43],
carbon blacks [44], graphene [21,45], and metal supports [46].
In this chapter, various forms of nickel were deposited onto certain carbon-based supports.
CFx was again employed and compared to XC72 to investigate the effects of the C-F bond in the
support on the urea oxidation kinetics. Both nickel and nickel oxides were also grafted onto
reduced graphene oxide to study the effects of annealing temperature on the urea oxidation
properties of nickel. Finally, nickel oxides were doped with the third row transition metals to study
the effect of doping on the urea oxidation properties.
4.2: Effect of Fluorinated Carbon Supports and Annealing in Ni/C Catalysts for Urea Oxidation
4.2.1: Experimental Methods
Nickel catalysts both by itself and on the carbon supports (Ni/C) were prepared via an
impregnation reduction method using hydrazine hydrate as the reducing agent. Briefly, 0.1 g of
either XC72 or CFx carbon support was added to 100 mL Millipore water (Direct-Q UV, 18.2
162
MΩ) and vigorously stirred and sonicated until homogenously dispersed. 0.66 g NiCl2•6H2O (Alfa
Aesar, 99.95% metals basis) was then dissolved in 50 mL Millipore water and added to the solution.
10 mL of hydrazine hydrate (Sigma-Aldrich, 35% in H2O) was slowly pipetted and the pH adjusted
to 10 by adding a 3.0 M sodium hydroxide solution. The solution was then transferred to an oil
bath and heated to 80
o
C for 1 hour. It was then washed and centrifuged until the pH of the
supernatant reached 7 and the obtained solid was dried in an oven, overnight. Nickel was also
reduced in the same manner without the carbon support.
The catalysts were also annealed in a tube furnace under air at 400
o
C for 5 hours as well
as the bare XC72 and CFx catalyst supports to study the effects of heat treatment on them.
SEM measurements were performed on a JEOL JSM-7001F electron microscope with an
acceleration voltage of 15 kV. TEM measurements were performed on a JEOL JEM-2100F
electron microscope with an acceleration voltage of 200 kV. XRD measurements were performed
on a Rigaku X-Ray diffractometer using Cu-Kα (0.15458 nm) radiation with a scan rate of 6
o
min
-
1
from 10
O
to 90
O
. Raman spectra were taken on a Horiba Raman Microscope with a laser
wavelength of 532 nm. XPS measurements were made on a Kratos Ultra XPS with a mono
aluminum radiation source. Thermogravemetric analysis (TGA) was performed on a TGA-50
Thermogravimetric Analyzer (Shimadzu). The samples were heated up at a rate of 10
o
C per
minute. The weight loss was analyzed from the graph from a thermogravimetric analysis program.
The half-cell electrochemical measurements were carried out on a standard rotating disk
electrode (RDE) setup. The measurements were taken on a Solartron SI 1287 Potentiostat and
analyzed using Corrware Software. Briefly, 5 mg of catalyst was added to a 1 mL solution
consisting of 10% isopropyl alcohol, 90% Millipore water, and 10 mg of 5% Nafion® solution
(Ion Power). The solution was sonicated and 20 μL was pipetted onto a glassy carbon electrode
163
(Pine Instruments) with a surface area of 0.195 cm
2
. The electrode was dried and placed into a
three cell testing system with the RDE as the working electrode, a platinum wire as the counter
electrode, and a mercury oxide (MMO, 4.24 M KOH filling solution) reference electrode. The
cell was purged with ultra pure argon (99.999%) for 15 min before the cyclic voltammetry (CV)
scans were performed from 0 V to 0.7 V vs. MMO.
The micro fuel cell was fabricated by cutting strips of filter paper (Whatman) in the “Y”
shape as shown in Figure 4.14. Catalyst ink solutions were formulated with a 1:9:1 composition
of catalyst: Millipore water: anionomer solution (Tokuyama anionomer, 5% weight). Two filter
paper strips were aligned and the catalyst ink was sonicated and painted on both sides of the anode
side of the electrode with an active surface area of 0.125 cm
2
. Colloidal graphite was used as the
cathode catalyst and prepared with the same catalyst ink composition and surface area as
aforementioned. Strips of stainless steel mesh were wrapped around the catalysts ink and the micro
fuel cell briefly pressed to ensure efficient bonding between the catalyst ink and the potentiostat
wires, and thus good electrical conductivity.
The micro fuel cells were tested by immersing the anode compartment into a solution of
6.0 M KOH and 0.33 M urea and the cathode compartment into a 30% hydrogen peroxide solution
(Millipore). The polarization curves were taken on a Solartron SI1287 potentiostat with the
potential scanned at 5 mV s
-1
from the open circuit voltage (OCV) to the short circuit current (JSC).
Electrochemical impedance spectroscopy (EIS) measurements were performed on a Solartron
SI1287 and SI1260 Impedence Phase Analyzer at a potential of 0.1 V vs. reference with an
amplitude of 5 mV.
164
4.2.2: Results
The temperature for the annealing of the catalysts was chosen based on the TGA curves of
the Ni/C catalysts and are shown in Figure 4.1. As the temperature increases past 300
O
C, the
weight loss curve rises to above 100% catalyst weight. This is the point where the oxidation of
the nickel starts to occur. The Ni/C curves peak at around 450
O
C and decrease. This is where the
carbon support starts to oxide and burn off. From this, the Ni/C-400 catalysts were determined to
be annealed at 400
O
C in order to oxide the nickel while keeping the carbon support intact.
Figure 4.1. TGA curves of the Ni/C catalysts.
SEM images were taken and shown in Figures 4.2a-f for the bare nickel and Ni/C catalysts.
There is minimal difference evident in the morphologies of the XC72 and CFx carbon supports
carbon supports from the images. The partial fluorination minimally affects the macro
0%
20%
40%
60%
80%
100%
120%
0 100 200 300 400 500 600 700 800
Weight (% of Initial)
Temperature (
O
C)
Ni/XC72
Ni/CFx
165
morphology of the particles, which still display an average size of around 30 nm for both the CFx
and XC72 as evident from the images. Moreover, there is minimal difference in the structure of
the unannealed and annealed bare carbon supports. EDAX spectra (not shown) were also taken
from the SEM images in Figures 4.2c-f and the nickel loading weights are reported in Table 4.2.
The size of the nickel particles from the SEM images greatly decreased from around 1 μm in both
the Ni and Ni-400 synthesized catalysts to around 200 nm for the carbon supported catalysts. The
nickel particles are slightly visible for the CFx supported catalysts and more visible for the XC72
supported catalysts displaying a considerably smaller particle size for the CFx supported catalysts.
For enhanced clarity of the nickel nanoparticles, TEM images were also taken and are
displayed in Figure 4.3a-f with the nickel particle size distributions shown in Figures 4.3c-f and
reported in Table 4.2. The particle sizes for the Ni/XC72 are similar to those found in earlier
reported works, where there is substantial aggregation of nickel nanoparticles without any
dispersing agent employed such as sodium citrate [47]. Compared with the bare nickel catalyst,
the particle size of the nickel is greatly reduced on the carbon supports. Moreover, the Ni/CFx
catalyst displayed a 3-fold decrease in nickel nanoparticle size compared to the Ni/XC72. This
could result from the hydrophobic nature of the fluorine-doped carbon support preventing the
nickel nanoparticle agglomeration. As was discussed in Section 2.2.2, during the synthesis, the
CFx powder was considerably more difficult to disperse in the aqueous solution than the XC72.
This could potentially have a dispersant effect on the nickel salts leading to smaller nickel
nanoparticle sizes There is also minor differences in nano particle size and morphology after the
annealing step most likely due to sintering of the nanoparticles [48] in both the carbon supported
and bare nickel catalysts showing little effect on the macro morphology of the nickel catalyst and
support from the annealing.
166
Figure 4.2. SEM images of: a) Ni; b) Ni-400; c) Ni/CFx; d) Ni/CFx-400; e) Ni/XC72 and; f)
Ni/XC72-400.
c)
d)
a) b)
e)
f)
167
100 150 200
0
200
400
600
800
1000
1200
1400
1600
1800
2000
0
5
10
15
20
25
100 110 120 130 140 150 160 170 180 190
Frequency (%)
Ni Particle Size (nm)
c)
a)
b)
c)
168
100 150 200 250
0
5
10
15
20
100 110 120 130 140 150 160 170 180 190
Frequency (%)
Ni Particle Size (nm)
10 30 50 70 90
0
200
400
600
800
1000
1200
1400
1600
1800
2000
0
5
10
15
20
25
30
10 20 30 40 50 60 70 80 90 100
Frequency (%)
Ni Particle Size (nm)
d)
e)
d)
e)
169
Figure 4.3. TEM images of the catalysts with particle size distribution analysis of: a) Ni; b) Ni-
400; c) Ni/XC72; d) Ni/XC72-400; e) Ni/CFx and; f) Ni/CFx-400. The scale measures 500 nm.
In order to gauge the effects of the carbon supports and annealing on the crystallite
structures of the nickel nanoparticles, XRD patterns were obtained and displayed in Figure 4.4a.
The nickel patterns show three main peaks denoted by the squares at values of 45
O
, 52
O
, and 77
O
,
signifying the (111), (200), and (220) Miller Indices of fcc nickel (PDF #04-0850) [49]. The Ni/C-
400 patterns show major peaks at the same values as well as slight peaks at 37
O
, 43
O
, and 63
O
denoted by the triangles indicating a slight presence of the (111), (200), and (220) Miller Indices,
for the fcc nickel oxide phase (PDF #73-1519) [50]. In the Ni/C catalysts, the Miller Indices are
visible for the fcc nickel phase as well as a slight (002) phase at around 22
O
signifying the graphitic
lattice of the carbon support. The annealed Ni/C catalysts display the nickel fcc peaks and slight
nickel oxide fcc peaks as was shown in the bare nickel annealed peaks. The crystallite sizes of the
nickel nanoparticles were calculated by the Scherrer equation [51] from the fcc nickel (200) peak
at 52
O
and are reported in Table 4.2. The crystallite sizes are usually calculated from the (111)
peak in fcc metal catalyst systems, however, the fcc (111) peaks for the annealed catalysts display
10 30 50 70 90
0
200
400
600
800
1000
1200
1400
1600
1800
2000
0
5
10
15
20
25
30
10 20 30 40 50 60 70 80 90 100
Frequency (%)
Ni Particle Size (nm)
f)
f)
170
the slight presence of the (111) fcc nickel oxide peak that would skew the overall results. Thus,
the next largest diffraction peak, the (200) peak of the fcc nickel was chosen. For each of the
catalysts, the crystallite size of the (200) peak increases after the oxidative annealing. This could
be due to the larger overall particle sizes of the Ni/CFx as displayed in the TEM images due to the
sintering of the particles, as well as the larger particle size distribution as seen in the Ni/XC72.
The Ni-400 pattern displays very slight NiO peaks as evident in the annealed Ni/C catalysts. This
could be due to the size of the nanoparticles from the TEM and SEM images which are
considerably larger than the Ni/C catalysts. This shows only a slight oxidation of the nickel after
the annealing. XRD patterns were also obtained for both the regular and annealed carbon supports
and are shown in Figure 4.4b. These patterns are very similar, showing little effect on the d-
spacing or crystallite size of the graphite phase of the carbon supports after the annealing process.
Raman spectra were taken of the Ni/C catalysts and bare carbon supports both before and
after annealing and are shown in Figures 4.5a&b. In each of the spectra, there arises two
characteristic peaks, one at around 1370 cm
-1
signifying the D-band and the other at around 1550
cm
-1
signifying the G-band as was discussed in Section 3.2.2. The ID/IG ratio of XC72 and XC72-
400 are 1.00 and 1.03, respectively, while the ID/IG ratio for CFx and CFx-400 are 0.97 and 0.98,
respectively. This is possibly due to a partial decomposition and removal in the epoxy and
hydroxyl groups on the carbon lattice during the oxidative annealing step without repairing the sp
2
hybridized carbons at the sites [52]. There is also a decrease in intensity in the spectra between
the normal and the annealed carbon supports showing an effect on the support from the annealing
process. The XC72 spectrum displays a greater ID/IG ratio difference than the CFx after the
annealing step. The ID/IG ratio of the Ni/CFx and the Ni/CFx-400 catalysts are 1.02 and 1.03,
respectively, while the ID/IG ratio of the Ni/XC72 and the Ni/XC72-400 is 1.03 and 1.07,
171
respectively. Again, this shows very minimal change in the disorder of the structure of the CFx
and XC72 supports after the slight annealing step.
Figure 4.4. XRD patterns of the: a) Ni/C catalysts and; b) carbon supports.
10 20 30 40 50 60 70 80 90
Intensity (a.u.)
2θ (degrees)
Ni
Ni-400
Ni/XC72
Ni/CFx
Ni/XC72-400
Ni/CFx-400
a)
10 20 30 40 50 60 70 80 90
Intensity (a.u.)
2θ (degrees)
XC72
CFx
XC72-400
CFx-400
b)
172
Figure 4.5. Raman spectra of the: a) Ni/C catalysts and; b) carbon supports.
The state of the nickel nanoparticle oxidation was investigated through XPS. The Ni 2p
deconvoluted XPS spectra are shown in Figures 4.6a-c. In each of the spectra there are roughly
four main peaks visible: two corresponding to the Ni 2p 3/2 and Ni 2p 1/2 at around 852 eV and
0 500 1000 1500 2000 2500 3000 3500
Intensity (a.u.)
Raman Shift (cm
-1
)
Ni/CFx
Ni/CFx-400
Ni/XC72
Ni/XC72-400
a)
G
0 500 1000 1500 2000 2500 3000 3500
Intensity (a.u.)
Raman Shift (cm
-1
)
CFx
CFx-400
XC72
XC72-400
b)
D G
D
173
871 eV, respectively, and the other two corresponding to shake-up satellite peaks at around 859
eV and 879 eV [53,54]. The nickel catalysts with no oxidative annealing display only one main
peak for the Ni 2p 3/2 binding peak at around 852 eV. The annealed catalysts display peak splitting
that was further deconvoluted into a Ni
O
and NiO peak with the Ni
O
shifting to a lower binding
energy of around 850 eV [55,56]. Whereas the XRD patterns between the Ni and Ni-400 displayed
only slight nickel oxide peaks, the nickel oxide binding is more evident in the XPS spectra. This
could arise from the size of the nickel particles. It is possible that only the surface of the nickel
particles was oxidized, which was detected since XPS is a surface analysis technique. The
oxidation peaks did not appear in the XRD because the mild oxidation did not produce highly
ordered nickel oxide indices leaving the fcc nickel phase intact overall. The Ni/XC72 and Ni/CFx
have smaller nickel particle sizes which have a larger surface area present for oxidation both at the
surface as well as to create the partially ordered indices. XPS spectra were also taken for the C 1s
peaks and are shown in Figures 4.7a-d. Each of the C 1s spectra displays one major peak, which
was deconvoluted to C=C, C-C, C-O, C=C, and –COOH with the CFx supported catalysts
displaying a C-F peak. The deconvolution results are reported in Table 4.1. A few trends in the
catalysts arise from the C 1s spectra. One trend is the decrease in the C-O peak after the annealing
step. The oxidative annealing helps to remove the C-O species present in the graphitic lattice,
which can be viewed in the Raman spectra as the increase in the ID/IG ratio. Another trend with
the oxygen species is the slight increase in the C=O peak after the annealing step for the catalysts.
This could result in the further oxidation of the C-O functional groups causing the decrease in the
C-O peak. The nickel could also have been oxidized from the partial removal of the oxygen groups
in the C-O in the annealing step [57]. The sp
2
/sp
3
ratios of the carbon lattice were also calculated
and reported in Table 4.1. In each of the catalysts as well as the bare supports, the sp
2
/sp
3
ratio
174
and decreased after the annealing process indicating an increase in the damage in each of the
supports.
Figure 4.6. XPS spectra of Ni 2p of the Ni/C catalysts: a) Ni; b) Ni/XC72 and; c) Ni/CFx. The
(1) spectrum indicates the catalyst annealed at 400
O
C; the (2) spectrum indicates the unannealed
catalyst.
840 850 860 870 880 890
Intensity (a.u.)
Binding Energy (eV)
Raw
Ni 2p 3/2
NiO
Ni 2p 1/2
Satellite
Satellite
Envelope
a)
(2)
(1)
840 850 860 870 880 890
Intensity (a.u.)
Binding Energy (eV)
Raw
Ni 2p 3/2
NiO
Ni 2p 1/2
Satellite
Satellite
Envelope
b)
(1)
(2)
840 850 860 870 880 890
Intensity (a.u.)
Binding Energy (eV)
Raw
Ni 2p 3/2
NiO
Ni 2p 1/2
Satellite
Satellite
Envelope
c)
(2)
(1)
175
Figure 4.7. XPS C1s spectra of the Ni/C catalysts: a) Ni/CFx; b) Ni/XC72; c) CFx and; d)
XC72. The (1) spectrum indicates the catalyst annealed at 400
O
C; the (2) spectrum indicates the
unannealed catalyst.
Table 4.1. Summary of the XPS C1s spectra of the Ni/C catalysts.
Binding Ni/CFx Ni/CFx-400 Ni/XC72 Ni/XC72-400 CFx CFx-400 XC72 XC72-400
C=C 26.41 29.18 25.02 29.49 31.09 32.39 32.84 26.1
C-C 19.48 28.31 21.4 29 26.33 30.55 26.32 24.38
C-O 36.62 27.9 32.73 17.79 18.95 4.7 28.2 26.34
C=O 2.05 2.67 17.35 17.6 6.14 8.49 9.19 16.67
C-F 10.59 3.14 N/A N/A 8.86 12.71 N/A N/A
-COOH 4.85 8.81 3.5 6.12 8.62 11.15 3.27 6.51
sp
2
/sp
3
1.36 1.03 1.17 1.02 1.18 1.06 1.25 1.07
275 280 285 290 295
Intensity (a.u.)
Binding Energy (eV)
Raw
C=C
C-C
C-O
C=O
C-F
H-C-OOH
Envelope
a)
(2)
(1)
275 280 285 290 295
Intensity (a.u.)
Binding Energy (eV)
Raw
C=C
C-C
C-O
C=O
-COOH
Envelope
b)
(1)
(2)
275 280 285 290 295
Intensity (a.u.)
Binding Energy (eV)
Raw
C=C
C-C
C-O
C=O
C-F
-COOH
Envelope
c)
(1)
(2)
275 280 285 290 295
Intensity (a.u.)
Binding Energy (eV)
Raw
C=C
C-C
C-O
C=O
-COOH
Envelope (1)
(2)
d)
176
Half-cell testing was employed in order to gauge the urea electrooxidation activity of the
Ni/C catalysts. Standard CV measurements were taken under argon purging as shown in Figures
4.8a&b. Figure 4.8a shows the CVs with urea while Figure 4.8b shows the CVs with only the
potassium hydroxide solution. In the potassium hydroxide solution, the OH
-
species adsorb onto
the nickel surface forming a layer of Ni(OH)2 species. At the onset potential region, the Ni(OH)2
is further oxidized to NiOOH, which has been demonstrated as being the nickel species involved
in urea oxidation [58]. In the cathodic scan, the Ni
3+
species is reduced back into the Ni
2+
species
at around 0.4 V -0.5 V [59]. The charge from the area above this redox curve was integrated for
the CV in Figure 4.8b and divided by the coulombic charge density of 514 μC cm
-2
required to
reduce Ni
3+
to Ni
2+
based on the catalyst loading [60] in order to obtain the electrochemically
active surface area (ECSAs) for the catalysts. The ECSAs were calculated and reported in Table
4.2. The ECSAs of the bare nickel and nickel-400 catalysts display very similar values signifying
a minimal change in the surface area after the annealing. This is most likely due to the large size
of the nickel particles and relatively minor influence of the oxidative annealing had on the overall
structure and surface properties. The Ni/C catalysts displayed an overall enhancement in ECSA
compared to the bare nickel catalysts due to the smaller particle size and enhanced nanoparticle
dispersion on the carbon supports. The Ni/XC72 catalysts displayed a 21% decrease in ECSA
after the slight oxidative annealing while the Ni/CFx catalysts only displayed a 6% decrease in
ECSA. This could arise from the slightly larger nickel nanoparticle size of the annealed Ni/CFx
as confirmed by the TEM images as well as the slight oxidation of the catalyst surface as shown
in the XPS spectra.
The current densities and onset potentials of urea oxidation from Figure 4.8a are shown in
Table 4.2. The urea electrooxidation activity is higher for each of the catalysts on carbon supports
177
before the annealing process. The onset potentials are also more favorable for the non-heat treated
catalysts than for the annealed catalysts for both types of carbon supports. However, the CV curves
for the unannealed and annealed nickel catalysts are very similar in both onset potential and overall
shape and current density showing minimal effect of the heat treatment on catalytic activity. This
could mean that the carbon supports act both to support the nickel catalyst and to assist in the
eletrooxidation of the urea. From the XPS spectra, the carbon supports displayed minor damage
from the annealing process and this lead to a decrease in activity for the Ni/C-400 catalysts for
both the current density and onset potential. Another interesting aspect to note is the CV curves
for the carbon supports without the nickel shown in Figures 4.9a&b. While the XC72 and XC72-
400 carbon supports showed no activity for urea oxidation, both the CFx and CFx-400 supports
display some activity with about 4.45 mA cm
-2
and 1.53 mA cm
-2
current density at 0.65 V vs
MMO respectively and an onset potential of 0.51 V and 0.52 V vs. MMO, respectively. This could
indicate a possible effect on the fluorine doping of the carbon support with activity for urea
electrooxidation. Figure 4.10 shows the potential stair step of the catalysts in the 1.0M KOH and
0.33M urea solution. The overall trend of the potential stair step follows that of the CV charts for
the catalysts.
178
Figure 4.8. CV scans of the Ni/C catalysts in: a) 1.0 M KOH + 0.33 M urea solution and; b) 1.0
M KOH solution at 20 mV s
-1
.
-4
0
4
8
12
16
20
24
0 0.1 0.2 0.3 0.4 0.5 0.6 0.7
Current Density (mA cm
-2
)
Potential (V vs. MMO)
Ni
Ni-400
Ni/XC72
Ni/XC72-400
Ni/CFx
Ni/CFx-400
a)
-4
-2
0
2
4
6
8
10
0 0.1 0.2 0.3 0.4 0.5 0.6 0.7
Current Density (mA cm
-2
)
Potential (V vs. MMO)
Ni
Ni-400
Ni/XC72
Ni/XC72-400
Ni/CFx
Ni/CFx-400
b)
179
Figure 4.9. CV scans of the carbon supports in: a) 1.0M KOH + 0.33M urea solution and; b)
1.0M KOH solution at 20 mV s
-1
.
-1
0
1
2
3
4
5
6
0 0.1 0.2 0.3 0.4 0.5 0.6 0.7
Current Density (mA cm
-2
)
Potential (V vs. MMO)
XC72
CFx
XC72-400
CFx-400
a)
-2
0
2
4
6
8
0 0.1 0.2 0.3 0.4 0.5 0.6 0.7
Current Density (mA cm
-2
)
Potential (V vs. MMO)
XC72
CFx
XC72-400
CFx-400
b)
180
Figure 4.10. Step CV of the Ni/C catalysts in 1.0 M KOH + 0.33 M urea solution.
CVs were also taken at various scan rates and shown in Figure 4.11a for Ni/CFx-400 and
in Figure 4.11b as a summary for all the catalysts. All of the curves in Figure 4.11b are linear
showing a diffusion controlled reaction with the urea for all of the catalysts. Stability
measurements were also carried out by scanning the catalysts from 0 V to 0.7 V vs. MMO at 100
mV s
-1
. A regular stability scan is shown in Figure 4.12a with the stability curves in Figure 4.12b.
The stability values after 50 scans are reported in Table 4.2. Again, the stability results for the
bare nickel catalysts are almost identical in nature showing little to no effect of the mild heat
treatment. Chronoamperommetry scans were also taken of the catalysts at 0.65 V and are
displayed in Figure 4.13. The curves follow the same trend as the CV curves with the Ni/CFx
catalyst displaying the highest urea electrooxidation currents followed by the Ni/XC72 and the
bare nickel with the annealed catalysts displaying lower current densities. The overall decay
displayed in the stability curves are also around the same signifying no effect of the catalyst support
on the long term stability of the catalyst.
1.2
1.3
1.4
1.5
1.6
1.7
0
0.5
1
1.5
2
2.5
3
0 100 200 300 400 500
Potential (V)
Current Density (mA cm
-2
)
Time (sec)
Ni
Ni-400
Ni/XC72
Ni/XC72-400
Ni/CFx
Ni/CFx-400
Potential
181
Figure 4.11. a) CV of the Ni/CFx-400 catalyst at different scan rates; b) the relationship
between the square root of the scan rate and the current density of the Ni/C catalysts.
-2
0
2
4
6
8
0 0.1 0.2 0.3 0.4 0.5 0.6 0.7
Current Density (mA cm
-2
)
Potential (V vs. MMO)
5
10
20
50
100
a)
0
5
10
15
20
25
30
0 5 10 15 20
Current Density (mA cm
-2
)
ω
1/2
Ni
Ni-400
Ni/XC72
Ni/XC72-400
Ni/CFx
Ni/CFx-400
b)
182
Figure 4.12. a) CV of Ni/XC72 catalyst cycled 50 times for stability; b) stability plots at 0.65 V
for the Ni/C catalysts.
0
3
6
9
12
15
18
0.4 0.5 0.6 0.7
Current Density (mA cm
-2
)
Potential (V vs. MMO)
1st
50th
a)
70%
80%
90%
100%
0 10 20 30 40 50
Current Density (% of Initial at 0.65V vs.
MMO)
Cycle Number
Ni
NiO
Ni/XC72
Ni/XC72-400
Ni/CFx
Ni/CFx-400
183
Figure 4.13. Chronoamperommetry graph of the Ni/C catalysts at 1000 RPM.
Table 4.2. Summary of the characterization and electrochemical tests of the catalysts.
Catalyst TEM Ni
Particle
Size
(nm)
Percent
Nickel
EDAX
Ni
Crystallite
Size (nm)
Onset
Potential (V
vs. MMO)
Current
Density @
0.65 V
(mA cm
-2
)
Stability
(% of
Initial at
0.65 V)
ECSA
(m
2
g
-1
)
Ni ~500 100 4.3 0.507±0.013 4.31±1.23 83.9 21.69
Ni-400 ~500 86.2 10.7 0.512±0.010 4.28±0.97 84.9 22.73
Ni/XC72 137±24 57.6 8.6 0.485±0.011 13.34±3.26 77.2 85.69
Ni/CFx 43±14 43.8 7.3 0.502±0.008 18.22±3.90 97.9 116.82
Ni/XC72-400 137±42 82.8 9.3 0.501±0.006 6.70±1.50 78.2 69.83
Ni/CFx-400 47±13 54.4 7.7 0.511±0.019 7.70±2.63 81.2 109.76
The efficacy of the Ni/C catalysts was further tested with direct micro urea/hydrogen
peroxide fuel cells shown in Figure 4.14. The polarization curves are displayed in Figure 4.15
with the OCV and power density results reported in Table 4.3. The Ni/C catalysts displayed both
0
1
2
3
4
0 1000 2000 3000
Current Density (mA cm
-2
)
Time (s)
Ni
NiO
Ni/XC72
NiO/XC72
Ni/CFx
NiO/CFx
184
enhanced OCV and power density values compared to the bare nickel catalysts with the Ni/CFx
and Ni/XC72 displaying a 63% and 25% increase in power density, respectively. This could be
due to the increased dispersion and smaller sizes of the nickel nanoparticles leading to an enhanced
ECSA. The Ni/CFx also displayed a 30% increase in power density compared with the Ni/XC72.
This could arise from the decrease in the nickel nanoparticle size with the CFx supports compared
with the XC72 supports and possible synergistic effects of the fluorine doping in the carbon
support. The urea oxidation activity follows the same trend as was observed in the half-cell testing.
Urea oxidation activity was further analyzed in the micro fuel cells from the EIS spectra
shown in Figure 4.16 with charge and mass transfer resistances reported in Table 4.3. Each curve
resembles a depressed semicircle that was fitted to a circuit shown in Figure 4.17. From the
equivalent circuit, L1 represents the inductance from the equipment and wires, Rel is the
electrolyte resistance of the two solutions in the anode and cathode, Rkin is the charge transfer
resistance, Rmas is the mass transfer resistance, CPE1 and CPE2 are the pseudo-capacitive nature
of the catalyst layers, and W1 is the Warburg element for the finite diffusion of reactant and
products through the filter layers. The spectrum for the bare nickel catalyst displays a considerably
higher semi-circle signifying higher charge and mass transfer resistances as reported in Table 4.3.
The Ni/CFx and Ni/XC72 catalysts displayed an 82% and 65% decrease in charge transfer
resistance compared with the bare nickel catalyst. This reduction in charge and mass transfer
resistance could arise from the higher current densities of the Ni/C catalysts compared to the bare
nickel. Another reason could be a better conducting network with the carbon supported catalysts
and the anionomer ink particularly with the CFx support as was discussed in Section 2.4.2 in
alkaline conditions. Due to the properties of the paper-based micro fuel cell, external heating could
cause the premature evaporation of the reactant fluids and loss of connectivity and conductivity so
185
the fuel cell tests were carried out at room temperature. The OCV and the power density would
be expected to increase due to the more favorable kinetics at higher temperatures [9,61].
Figure 4.14. Diagram of the micro direct urea/hydrogen peroxide fuel cell.
Figure 4.15. Polarization curves for the Ni/C catalysts in the micro direct urea fuel cells.
0
0.05
0.1
0.15
0
0.1
0.2
0.3
0.4
0 0.5 1 1.5 2 2.5
Power Density (mW cm
-2
)
Potential (V)
Current Density (mA cm
-2
)
Ni
Ni/CFx
Ni/XC72
186
Figure 4.16. EIS curves for the Ni/C catalysts in the micro direct urea fuel cells.
Figure 4.17. Equivalent circuit for the EIS spectra for the Ni/C catalysts.
Table 4.3. Summary of the fuel cell tests of the Ni/C catalysts.
Catalyst OCV Power Density
(mW cm
-2
)
Charge Transfer
Resistance (Ohms)
Mass Transfer
Resistance (Ohms)
Ni 0.21 0.08 437.5 755.5
Ni/CFx 0.33 0.13 78.19 431.3
Ni/XC72 0.28 0.10 149.5 524.7
-50
0
50
100
150
200
250
300
350
0 200 400 600 800 1000 1200 1400
-Z" (Ohms)
Z' (Ohms)
Ni
Ni/CFx
Ni/XC72
L1 Rel Rkin
CPE1
Rmas
CPE2 W1
Element Freedom Value Error Error %
L1 Free(+) 4.2007E-06 7.4337E-08 1.7696
Rel Fixed(X) 66 N/A N/A
Rkin Free(+) 436.4 12.679 2.9054
CPE1-T Free(+) 3.1656E-05 1.6266E-06 5.1384
CPE1-P Free(+) 0.77836 0.014206 1.8251
Rmas Free(+) 49.48 8.8648 17.916
CPE2-T Free(+) 2.4075E-05 1.7184E-05 71.377
CPE2-P Free(+) 0.72589 0.075791 10.441
W1-R Free(+) 1.752 2.1462 122.5
W1-T Free(+) 1.0571E-05 5.0856E-06 48.109
W1-P Free(+) 0.66385 0.22003 33.145
Chi-Squared: 0.00064179
Weighted Sum of Squares: 0.061612
Data File: E:\Nickel Files\20170518-NirGO300-03.z
Circuit Model File: C:\Dean\DMFC\AnodeLayers\AnodewInductorWarburg.mdl
Mode: Run Fitting / Selected Points (10 - 62)
Maximum Iterations: 100
Optimization Iterations: 0
Type of Fitting: Complex
Type of Weighting: Calc-Modulus
187
4.3: Effect of Annealing Temperature on Ni/rGO Catalysts on the Structure and Urea Oxidation
4.3.1: Experimental Details
Graphene oxide (GO) was synthesized via the modified Hummer’s method [62] as was
described in Section 3.2.1. The nickel on reduced graphene oxide (Ni/rGO) catalysts were
synthesized via an aqueous-based reduction method. 0.75 g of the previously synthesized GO was
dissolved into 200 mL Millipore water (Direct-Q UV, 18.2 MΩ) and vigorously stirred and
sonicated for one hour. Then, 3.0374 g of NiCl2·6H2O was dissolved into 100 mL Millipore water
and added to the GO solution in order to obtain a 50%-wt nickel catalyst loading. 30 mL of
hydrazine hydrate solution (24-26%, Sigma-Aldrich) was added dropwise to the solution and the
pH was adjusted to 10 by adding a 3.0 M aqueous sodium hydroxide solution. The solution was
then heated to 80
o
C in an oil bath for one hour and then centrifuged and washed with Millipore
water until the supernatant reached a pH of 7 and the obtained solid dried in an oven. The Ni/rGO
catalysts were then annealed at 300, 400, 500, 600, and 700
o
C, respectively under argon for five
hours. An unsupported nickel catalyst was also synthesized in the same manner from the
NiCl2·6H2O salts.
All of the characterization, half-cell electrochemical, and micro fuel cell testing methods
were carried out as previously described in Section 4.2.1.
188
4.3.2: Results
The structure of the Ni/rGO catalysts was analyzed by SEM images shown in Figures
4.18a-g. The particle size for the bare nickel particles is around 900 nm and displays the typical
“sea urchin” type morphology as found in previous studies [63]. The nickel nanoparticles in the
Ni/rGO catalysts, are considerably smaller in size, mostly around 100 nm, supported directly onto
the graphene, while the unsupported nickel nanoparticle agglomerates are around 200 nm in
diameter. There appears to be a slight difference in the nickel nanoparticle morphology between
the different annealing temperatures. At the unannealed and lower annealing temperatures, the
nickel nanoparticles display a rough texture, similar to the bare nickel particles. At the higher
annealing temperatures (600 and 700
O
C), the nanoparticle morphology has a smoother appearance
due to sintering. This aspect is more visible in the TEM images shown in Figure 4.19a-g. TEM
images also show two distinct nickel particle sizes as well as the change in nickel nanoparticle
morphology at the higher annealing temperatures (600 and 700
O
C).
a)
b)
189
Figure 4.18. SEM images of the Ni/rGO catalysts: a) Ni; b) Ni/rGO; c) Ni/rGO-300; d) Ni/rGO-
400; e) Ni/rGO-500; f) Ni/rGO-600 and; g) Ni/rGO-700.
c)
d)
e)
f)
g)
190
a) b)
c)
d)
191
Figure 4.19. TEM images of the Ni/rGO catalysts: a) Ni; b) Ni/rGO; c) Ni/rGO-300; d) Ni/rGO-
400; e) Ni/rGO-500; f) Ni/rGO-600 and; g) Ni/rGO-700. The scale measures 500 nm.
XRD patterns were obtained and are shown in Figure 4.20. The peak at 11
O
is indicative
of the (100) Miller Index of graphene oxide. After reduction, the peak shifts to 23
O
signifying the
(002) Miller Index of the reduced graphene oxide found in all of the Ni/rGO catalysts. As the
annealing temperature increases in the Ni/rGO catalysts, the (002) peak sharpens signifying the
e)
f)
g)
192
decrease of the crystallite size in the graphitic lattice. The peak location also shifts to a higher
degree indicating a smaller d-spacing in the graphitic sheets as displayed in Figure 4.21. This is a
result of the gradual removal of the oxygen functional groups from the graphitic lattice during the
annealing step [64,65]. There are also peaks present at 45
O
, 52
O
, and 77
O
, signifying the (111),
(200), and (220) Miller Indices of fcc nickel as was discussed in Section 4.2.2. As the annealing
temperature increases, there is a sharpening of the fcc nickel peaks signifying a larger crystallite
size of the nickel nanoparticles. The crystallite sizes for the nickel are shown in Figure 4.21.
The Raman spectra for the Ni/rGO catalysts are shown in Figure 4.22. Two main peaks
arise in the graphene spectra: the D- and G-bands as was discussed in Section 4.2.2. A summary
of the ID/IG ratios of the Ni/rGO catalysts is found in Figure 4.23. The ID/IG ratio increases as the
annealing temperature is increased due to the increased disorder in the graphene lattice from the
removal of some of the oxygen functional groups after the reduction [64].
Figure 4.20. XRD patterns for the Ni/rGO catalysts.
10 20 30 40 50 60 70 80 90
Intensity (a.u.)
2θ (degrees)
GO
Ni
Ni/rGO
300
400
500
600
700
193
Figure 4.21. Summary of the XRD patterns for the Ni/rGO catalysts from Figure 4.20.
Figure 4.22. Raman spectra for the Ni/rGO catalysts.
3.2
3.3
3.4
3.5
3.6
50
75
100
125
150
0 200 400 600 800
rGO d-spacing (Å)
Ni Crystallite Size (nm)
Annealing Temperature (
O
C)
Crystal. Size
d-space
Ni/rGO
0 500 1000 1500 2000 2500 3000 3500
Intensity (a.u.)
Raman Shift (cm
-1
)
GO
Ni/rGO
300
400
500
600
700
194
Figure 4.23. Summary of the raman spectra for the Ni/rGO catalysts from Figure 4.22.
Figures 4.24a-g display the Ni 2p XPS spectra. Two main peaks arise as was discussed in
Section 4.2.2. As the annealing temperature is increased, the Ni 2p 3/2 peak broadens and
eventually splits into two noticeable peaks from 500-700
O
C. These peaks can be deconvoluted
into a Ni
O
and a Ni
2+
(NiO) [66,67] peak at around 849 eV and 852 eV, respectively, signifying a
greater NiO characteristic at higher annealing temperatures. This could arise from the nickel being
oxidized from the oxygen present in the oxygen functional groups in the graphitic lattice at the
higher temperatures [68]. Nickel oxide peaks, however, were absent from the XRD patterns most
likely because the oxidized nickel was amorphous and not crystalline in nature.
C1s XPS spectra for the graphene lattices were also taken and are shown in Figures 4.25a-
f. Each of the spectra displays one major peak that was deconvoluted to the C=C, C-C, C-O, and
C=O binding species. There is minimal difference in the C-O peak among the different annealing
temperatures signifying minimal change to the C-O functional groups in the graphitic lattice.
There are, however, some differences in the C=O peak across the range of annealing temperatures.
At the higher annealing temperatures of 600 and 700
O
C, the C=O peak is greatly reduced,
1.1
1.15
1.2
1.25
1.3
1.35
1.4
0 200 400 600 800
I
D
/I
G
Annealing Temperature (
O
C)
GO Ni/rGO
195
signifying the loss of some of the oxygen groups. This corroborates the results of the XRD and
Raman spectra, where the loss of oxygen functional groups is evident, leading to a smaller d-
spacing and greater disorder of the graphitic lattice.
845 855 865 875 885
Intensity (a.u.)
Binding Energy (eV)
Raw
Ni 2p 3/2
Satellite
Ni 2p 1/2
Satellite
Envelope
a)
845 855 865 875 885
Intensity (a.u.)
Binding Energy (eV)
Raw
Ni 2p 1/2
Satellite
Satellite
Ni 2p 3/2
Envelope
b)
845 855 865 875 885
Intensity (a.u.)
Binding Energy (eV)
Raw
Ni 2p 1/2
Satellite
Ni 2p 3/2
Satellite
Envelope
c)
845 855 865 875 885
Intensity (a.u.)
Binding Energy (eV)
Raw
Ni 2p 1/2
Satellite
Ni 2p 3/2
Satellite
Envelope
d)
196
Figure 4.24. XPS Ni 2p spectra of the Ni/rGO catalysts: a) Ni/rGO; b) Ni/rGO-300; c) Ni/rGO-
400; d) Ni/rGO-500; e) Ni/rGO-600 and; f) Ni/rGO-700.
845 855 865 875 885
Intensity (a.u.)
Binding Energy (eV)
Raw Ni 2p 1/2
Satellite NiO
Satellite Ni 2p 3/2
Envelope
e)
845 855 865 875 885
Intensity (a.u.)
Binding Energy (eV)
Raw Ni 2p 1/2
Satellite Ni 2p 3/2
Satellite NiO
Envelope
f)
845 855 865 875 885
Intensity (a.u.)
Binding Energy (eV)
Raw Ni 2p 1/2
Satellite Ni 2p 3/2
NiO Satellite
Envelope
g)
280 285 290 295
Intensity (a.u.)
Binding Energy (eV)
a)
C=C
C-C
C-O
C=O
Envelope
280 285 290 295
Intensity (a.u.)
Binding Energy (eV)
b)
C=C
C-C
C-O
C=O
Envelope
197
Figure 4.25. XPS C 1s spectra of the Ni/rGO catalysts: a) Ni/rGO; b) Ni/rGO-300; c) Ni/rGO-
400; d) Ni/rGO-500; e) Ni/rGO-600 and; f) Ni/rGO-700.
Electrochemical half-cell tests of the Ni/rGO catalysts were carried out and the CVs are
displayed in Figure 4.26a with urea and in Figure 4.26b without urea. In the alkaline medium, the
hydroxide ions adsorb onto the nickel surface to form a Ni(OH)2 layer, which is oxidized by the
hydroxide ions to NiOOH at the onset potential region for urea oxidation as was discussed in
Section 4.2.2. Also, there is a peak present at around 0.45 V in the reverse scan signifying the
Ni
3+
/Ni
2+
redox couples on the catalyst surface [59]. The electrochemically active surface areas
(ECSAs) of the catalysts were also calculated from the CV in Figure 4.26b in the same manner as
280 285 290 295
Intensity (a.u.)
Binding Energy (eV)
c)
C=C
C-C
C-O
C=O
Envelope
280 285 290 295
Intensity (a.u.)
Binding Energy (eV)
d)
C=C
C-C
C-O
C=O
Envelope
280 285 290 295
Intensity (a.u.)
Binding Energy (eV)
e)
C=C
C-C
C-O
C=O
Envelope
280 285 290 295
Intensity (a.u.)
Binding Energy (eV)
f)
C=C
C-C
C-O
C=O
Envelope
198
in Section 4.2.2 by integrating under the cathodic sweep for the Ni
3+/2+
peak at around 0.4-0.5 V
and the results are reported in Table 4.4. Another aspect of the CV scan is the oxidation current
of the urea at a specific potential of 0.65 V and onset potentials of the catalysts in the urea solution
are displayed in Figure 4.27. Each of the annealed Ni/rGO catalysts displayed an enhanced urea
oxidation current and more favorable onset potential compared to that of the bare nickel catalyst.
The unannealed Ni/rGO catalyst displays an 87% increase in the current density compared to the
bare nickel catalyst due to the smaller particle size and higher surface area. As the annealing
temperature of the catalysts is increased, the catalytic activity of the Ni/rGO catalysts decreases.
One explanation for this could be the increasing nickel oxide characteristics of the catalyst surface
as displayed in the Ni 2p XPS spectra even though the nickel crystallite sizes are larger. As the
annealing temperature of the catalysts increases, the onset potential for urea oxidation shifts to
more negative potentials. One explanation for this could be the higher surface area and larger
crystallite size of the nickel nanoparticles for the Ni/rGO at higher temperatures as was displayed
in the XRD patterns. This increase in active surface area in addition to the larger crystallite sizes
could help aid in the initial oxidation of the Ni(OH)2 layer to NiOOH, followed by the oxidation
of urea.
Figure 4.28 shows the potential stair-step of the catalysts in urea solution. These scans
overall mimic the catalytic activity as displayed in Figure 4.26a. The effect of the scan rate on
urea oxidation kinetics has also been investigated and the results are displayed in Figures 4.29a&b.
Each of the Ni/rGO catalysts displays a linear function as plotted in Figure 4.29b indicating a
diffusion limiting process in the oxidation [69]. Cycling stability tests were also performed on the
Ni/rGO catalysts at the same scan window as the regular CV tests in Figures 4.26a&b and are
shown in Figures 4.30a&b and reported in Table 4.4. Each of the Ni/rGO catalysts display similar
199
cycling stabilities indicating that the continued oxidation and reduction of the Ni
2+
/Ni
3+
active
species isn’t greatly affected from repeated cycling. Chronoamperommetry studies were also
performed by holding the potential at 0.65 V at a rotation rate of 1000 RPM and shown in Figure
4.31. The urea oxidation current densities align with the CVs scans with minimal difference in the
stability of the catalysts over an extended period of time. Based on both the cycling and the
chronoamperommetry tests, it is reasonable to assume that the nickel catalysts underwent very
minor changes in their structure and catalytic properties.
-2
0
2
4
6
8
0 0.1 0.2 0.3 0.4 0.5 0.6 0.7
Current Density (mA cm
-2
)
Potential (V vs. MMO)
Ni
Ni/rGO
300
400
500
600
700
a)
200
Figure 4.26. CV scans of the Ni/rGO catalysts in: a) 1.0 M KOH + 0.33 M urea and; b) 1.0 M
KOH only.
Figure 4.27. Summary of the CV scans for the Ni/rGO catalysts from Figure 4.26a.
-2
0
2
4
6
0 0.1 0.2 0.3 0.4 0.5 0.6 0.7
Current Density (mA cm
-2
)
Potential (V vs. MMO)
Ni
Ni/rGO
300
400
500
600
700
b)
0.45
0.46
0.47
0.48
0.49
0.5 3
4
5
6
7
0 100 200 300 400 500 600 700
Onset Potential (V vs. MMO)
Current Density (mA cm
-2
)
Annealing Temperature (
O
C)
Current Density
Onset Potential
Ni/rGO
201
Figure 4.28. Step CV for the Ni/rGO catalysts in 1.0 M KOH + 0.33 M urea.
1.2
1.3
1.4
1.5
1.6
1.7
0
0.4
0.8
1.2
1.6
2
0 100 200 300 400 500
Potential (V)
Current Density (mA cm
-2
)
Time (s)
Ni
Ni/rGO
300
400
500
600
700
Potential
-2
0
2
4
6
0 0.1 0.2 0.3 0.4 0.5 0.6 0.7
Current Density (mA cm
-2
)
Potential (V vs. MMO)
5
10
20
50
100
a)
202
Figure 4.29. a) CV of the Ni/rGO-500 catalyst at different scan rates; b) the relationship
between the square root of the scan rate and the current density of the Ni/rGO catalysts.
1
2
3
4
5
6
0 2 4 6 8 10 12
Current Density (mA cm
-2
)
ω
1/2
Ni
Ni/rGO
300
400
500
600
700
b)
0
1
2
3
4
0.5 0.55 0.6 0.65 0.7
Current Density (mA cm
-2
)
Potential (V vs. MMO)
a)
203
Figure 4.30. a) CV of Ni/rGO-500 catalyst cycled 50 times for stability; b) stability plots at 0.65
V for the Ni/rGO catalysts.
Figure 4.31. Chronoamperommetry graph of the Ni/rGO catalysts at 1000 RPM.
0.95
0.96
0.97
0.98
0.99
1
0 10 20 30 40 50 60
Current Density (% of initial)
Cycle Number
Ni/rGO
300
400
500
600
700
b)
0
1
2
3
0 500 1000 1500 2000 2500 3000
Current Density (mA cm
-2
)
Time (s)
Ni/rGO
300
400
500
600
700
204
The urea oxidation activity was further assessed using a micro direct urea/hydrogen
peroxide fuel cell in the same manner as described in Section 4.2.2. Polarization curves for the
micro fuel cells were obtained and are shown in Figure 4.32 at ambient temperature with a
summary of the OCVs and peak power densities reported in Table 4.4. There is a 100% and 125%
enhancement in power density in the Ni/rGO and Ni/rGO-300 catalysts compared to the bare
nickel catalyst, respectively. This could arise from the increased surface area of the nickel
nanoparticles on the reduced graphene oxide support. The annealing of the Ni/rGO catalyst also
affects the performance. The Ni/rGO-300 catalyst displayed a 13% increase in peak power density
when compared to the unannealed catalyst while the Ni/rGO-700 catalyst displayed a 31%
decrease in comparison. This trend in performance mimics the one displayed in the half-cell
testing with decreasing current density as the annealing temperature of the catalysts was increased.
Another aspect of the polarization curves is the OCV of the catalysts. Two sets of OCV values
arise for the catalysts in the polarization curves. The bare nickel and unannealed Ni/rGO catalysts
have similar OCV values while the Ni/rGO catalysts annealed at 300 and 700
O
C have similar but
higher values in comparison. This trend mimics the onset potentials of the catalysts in the half cell
tests. A few factors could be responsible for this trend. The annealing of the catalyst was shown
to increase the OCV compared to the unannealed and bare nickel catalysts. This could arise from
the larger crystallite size of the annealed catalyst as was indicated by the half-cell measurements.
The Ni/rGO-700, however, displays a slightly lower OCV compared to the Ni/rGO-300 even
though the crystallite size is larger. Upon annealing, the nickel present in the Ni/rGO-700 catalyst
was partially oxidized by the rGO support. This in effect could lower the OCV of the fuel cell.
EIS curves of the micro fuel cell tests were also obtained and are shown in Figure 4.33
with the data reported in Table 4.4. All of the curves resemble depressed semi-circles and were
205
fitted with an equivalent circuit shown in Figure 4.17. Each of the Ni/rGO catalysts displays a
lower charge and mass transfer resistance than the bare Ni catalyst. This is most likely due to the
smaller nickel nanoparticle size and larger surface area present for urea oxidation. The Ni/rGO-
300 catalyst displays the lowest charge and mass transfer resistances, most likely due to having
the highest current flowing through the system at the measured potential. The Ni/rGO-300 catalyst
shows a very similar EIS spectrum to that of the unannealed Ni/rGO catalyst meaning similar
kinetics are occurring at the electrode. This is due to the similar structure and properties of the
nickel present in the catalyst annealed at lower temperatures. The Ni/rGO-700 catalyst displays a
larger charge and mass transfer resistance, most likely stemming from the low current densities as
measured in the polarization curves in comparison at the measured potential.
Figure 4.32. Polarization curves for the Ni/rGO catalysts in the micro direct urea fuel cells.
0
0.05
0.1
0.15
0.2
0
0.1
0.2
0.3
0.4
0 0.5 1 1.5 2 2.5 3
Power Density (mW cm
-2
)
Potential (V)
Current Density (mA cm
-2
)
Ni
Ni/rGO
300
700
206
Figure 4.33. EIS curves for the Ni/rGO catalysts in the micro direct urea fuel cells.
Table 4.4. Summary of the electrochemical and fuel cell tests of the Ni/rGO catalysts.
Catalyst ECSA
(m
2
g
-1
)
Stability
(% of
Initial)
OCV Power
Density
(mW cm
-2
)
Mass Transfer
Resistance
(Ohms)
Charge Transfer
Resistance
(Ohms)
Ni 7.761 N/A 0.21 0.08 1262 103.7
Ni/rGO 8.725 96.4 0.26 0.16 427.4 75.1
Ni/rGO-300 12.702 97.1 0.36 0.18 436.4 49.5
Ni/rGO-400 13.457 96.4 N/A N/A N/A N/A
Ni/rGO-500 15.012 97.3 N/A N/A N/A N/A
Ni/rGO-600 15.466 95.6 N/A N/A N/A N/A
Ni/rGO-700 16.668 96.1 0.33 0.11 500.5 97.5
-100
0
100
200
300
400
500
0 200 400 600 800 1000
-Z" (Ohms)
Z' (Ohms)
Ni
Ni/rGO
300
700
207
4.4: Effect of Annealing Temperature on NiO@rGO Catalysts on the Structure and Urea
Oxidation
4.4.1: Experimental Methods
The nickel oxide catalyst used was synthesized by a facile aqueous-based reduction and
oxidative annealing. Bare nickel nanoparticles were first synthesized from 2.0 g of NiCl2·6H2O
(Alfa Aesar, 99.95% metals basis) was added to 100 mL Millipore water (Direct-Q UV, 18.2 MΩ).
The solution was vigorously stirred and sonicated until the nickel solution was homogenously
dispersed. The solution was then heated to 80
o
C in an oil bath and ten times excess hydrazine
hydrate (Sigma-Aldrich, 35%) was added to reduce the nickel. The nickel solution was stirred for
1 hour at 80
o
C and then at room temperature overnight. The solution was then washed and
centrifuged until the pH of the supernatant reached 7 and the obtained solid dried in an oven at 80
O
C overnight. The resulting nickel powder was then annealed in a furnace under air at 500
o
C for
5 hours to form nickel oxide.
GO was synthesized via the modified Hummer’s method as previously described in Section
3.2.1. 1.0 g of the previously synthesized GO was vigorously stirred and sonicated for 1 hour in
250 mL Millipore water followed by adding 0.5 g of the previously synthesized nickel oxide. The
solution was then stirred and sonicated for another hour until it was homogenously dispersed. 30
mL of hydrazine hydrate solution was added dropwise to the solution and the pH was adjusted to
13 by adding 3.0 M aqueous sodium hydroxide solution. The solution was then heated to 80
o
C in
an oil bath for 1 hour and then washed and centrifuged with Millipore water until the supernatant
reached a pH of 7 and dried in an oven overnight. Under the reduction conditions, the NiO is not
208
reduced to Ni
O
(vide infra). The NiO@rGO powders were then annealed at 300, 400, 500, 600,
and 700
o
C under argon for 4 hours.
All of the characterization, electrochemical half-cell, and micro fuel cell tests were carried
out as previously described in Section 4.2.1.
4.4.2 Results
Nickel oxides were synthesized separately from the reduction of graphene instead of nickel
salts being reduced along with graphene oxide in the same step followed by oxidative annealing
under air. Previous synthetic attempts of the aforementioned method produced catalysts with
different levels of oxidized nickel morphologies with varying nickel and nickel oxide phases
confirmed by XRD measurements. Furthermore, the rGO supports would be oxidized starting at
temperatures above 400
O
C and fully burned off at temperatures above 500
O
C. It was decided
that in order to preserve the overall integrity of the graphene supports, the nickel oxide would be
synthesized in a separate step than the graphene reduction and the catalysts annealed under argon.
SEM images were taken and are displayed in Figures 4.32a-g. In each of the annealing
images, the nickel catalysts maintained their overall shape and morphology throughout the
different annealing temperatures. The GO reduction step with the deposited nickel oxide was
shown to have minimal effects on the nickel oxide morphology with the hydrazine hydrate. There
is also small difference between the rGO sheets and the majority of the GO sheets being smooth
while the rGO sheets display wrinkles after the reduction and annealing processes. TEM images
were also taken and shown in Figures 4.33a-g. These images show the same overall morphology
as was viewed in the SEM images.
209
a)
b)
c)
d)
e)
f)
210
Figure 4.34. SEM images of: a) NiO; b) NiO@rGO; c) Ni@rGO-300; d) Ni@rGO-400; e)
Ni@rGO-500; f) Ni@rGO-600 and; g) Ni@rGO-700.
g)
a)
b)
211
c)
d)
e) f)
212
Figure 4.35. TEM images of a) NiO; b) NiO@rGO; c) Ni@rGO-300; d) Ni@rGO-400; e)
Ni@rGO-500; f) Ni@rGO-600 and; g) Ni@rGO-700. The measuring bar indicates 500 nm.
To gauge the effects of the reduction and annealing on the NiO@rGO catalysts, XRD
patterns were obtained and are shown in Figure 4.34. The graphene oxide pattern shows a peak at
11
O
signifying the (100) Miller Index. For all of the NiO@rGO catalysts, the peak shifts to around
a 2θ value of 23
O
signifying the (002) Miller Index indicating a successful reduction of the
graphene oxide by the hydrazine hydrate. These (002) peaks of the reduced graphene oxide
however, display a change in both the overall shape and peak position. As the annealing
temperature increases, the (002) peak becomes sharper and shifts to a higher 2θ degrees as was
discussed in Section 4.3.2. This shifting indicates a decrease in the d-spacing of the reduced
graphene oxide sheets and are reported in Figure 4.35. This arises from the removal of the oxygen
functional groups from the lattice at the higher annealing temperatures [64,65]. The NiO pattern
shows peaks at 2θ values of 37
O
, 43
O
, 63
O
, 75
O
, and 80
O
signifying the (111), (200), (220), (311),
and (222) Miller Indices of nickel oxide in the fcc phase. The NiO@rGO catalyst displays the
same peak positions showing no reducing effects on the fcc nickel oxide structure from the
g)
213
hydrazine hydrate reduction of graphene oxide. As the temperature is increased, peak sharpening
is evident in the nickel oxide peaks indicating an increase in crystallite size [70,71], as reported in
Figure 4.35. As the annealing temperature increases, there is also another small peak appearing at
45
O
indicative of fcc nickel.
In order to further investigate the annealing effects on the rGO support, Raman spectra
were obtained and are shown in Figure 2b. A summary of the ID/IG ratios is displayed in Figure
4.37. After annealing, the ID/IG ratios of the NiO@rGO catalysts increased due to the smaller
overall graphitic domain size compared to GO. As the annealing temperature increased, the ID/IG
ratio also increased. This could also be due to the removal of the oxygen functional groups at
higher annealing temperatures as was evidenced by the decrease in d-spacing in the XRD patterns
[57].
Figure 4.36. XRD patterns of the NiO@rGO catalysts.
10 20 30 40 50 60 70 80 90
Intensity (a.u.)
2θ (degrees)
GO
NiO
NiO@rGO
300
400
500
600
700
(200)
(222)
(220)
(311)
a)
(002)
(111)
(100)
214
Figure 4.37. Summary of the XRD patterns of the NiO@rGO catalysts from Figure 4.36.
Figure 4.38. Raman spectra of the NiO@rGO catalysts.
3.25
3.35
3.45
3.55
3.65
3.75
125
135
145
155
165
175
0 100 200 300 400 500 600 700 800
rGO d-spacing (Å)
Ni Crystallite Size (nm)
Annealing Temperature (
O
C)
Crystal. Size
d-space
NiO@rGO
0 500 1000 1500 2000 2500 3000 3500
Intensity (a.u.)
Raman Shift (cm
-1
)
GO
NiO
NiOatrGO
300
400
500
600
700
D
G
215
Figure 4.39. Summary of the Raman spectra of the NiO@rGO catalysts from Figure 4.38.
The XPS Ni 2p spectra are displayed in Figures 4.40a-g. In each of the spectra, there are
two main nickel peaks along with the corresponding satellite shake up peaks as was discussed in
Section 4.3.2. For each of the catalysts, the Ni 2p 3/2 peak was further deconvoluted into two
peaks with the Ni
O
peak around 850 eV and the Ni
2+
(NiO) peak at around 852 eV [66,67]. As the
annealing temperature is increased, the intensity of the Ni
O
relative to the Ni
2+
increases signifying
a greater amount of nickel on the surface of the catalyst compared to nickel oxide. The annealing
step could be responsible for removing some of the oxygen from the catalyst surface for both the
nickel particles and the rGO support.
1.1
1.15
1.2
1.25
1.3
1.35
0 200 400 600 800
I
D
/I
G
Annealing Temperature (
O
C)
GO NiO@rGO
216
845 855 865 875 885
Intensity (a.u.)
Binding Energy (eV)
Raw Ni 2p 3/2
NiO Ni 2p 1/2
Satellite Satellite
Envelope
a)
845 855 865 875 885
Intensity (a.u.)
Binding Energy (eV)
Raw Ni 2p 1/2
Satellite Satellite
NiO Ni 2p 3/2
Envelope
b
845 855 865 875 885
Intensity (a.u.)
Binding Energy (eV)
Raw Satellite
Ni 2p 1/2 NiO
Ni 2p 3/2 Envelope
Satellite
c)
845 855 865 875 885
Intensity (a.u.)
Binding Energy (eV)
Raw Ni 2p 1/2
Satellite Ni 2p 3/2
Satellite NiO
Envelope
d)
845 855 865 875 885
Intensity (a.u.)
Binding Energy (eV)
Raw Ni 2p 1/2
Satellite NiO
Satellite Ni 2p 3/2
Envelope
e)
845 855 865 875 885
Intensity (a.u.)
Binding Energy (eV)
Raw Ni 2p 1/2
Satellite NiO
Satellite Ni 2p 3/2
Envelope
f)
217
Figure 4.40. XPS Ni 2p spectra of the NiO@rGO catalysts: a) NiO; b) NiO@rGO; c)
Ni@rGO-300; d) Ni@rGO-400; e) Ni@rGO-500; f) Ni@rGO-600 and; g) Ni@rGO-700.
XPS C1s spectra were also taken in order to analyze the surface components of the rGO
supports and are displayed in Figures 4.41a-f. One main peak is observed for each spectra and
was deconvoluted to four peaks representing the C=C, C-C, C-O, and C=O bonding of the carbon
lattice in the rGO supports. As the annealing temperature is increased, a few trends appear. One
trend is the gradual decrease in the C=O peak from 300-600
O
C with a sharp increase at 700
O
C.
This is from the removal of the C=O functional group at higher temperatures. There is also a less
pronounced decrease in the C-O peak also resulting from the slight removal of the C-O groups
from the annealing step [57]. This corroborates with the XRD patterns and the Raman spectra
showing the removal of some of the oxygen functional groups without the repair of the graphitic
lattice as was discussed in Section 4.3.2. Another trend observed is the gradual increase in the C-
C peak in comparison to the C=C peak. This shows the slight degradation of the sp
2
hybridized
carbon lattice in the rGO support mostly likely due to the removal of the oxygen functional groups
without the subsequent repair of the graphitic carbon lattice. This corroborates the ID/IG ratios
845 855 865 875 885
Intensity (a.u.)
Binding Energy (eV)
Raw Ni 2p 1/2
Satellite Satellite
Ni 2p 3/2 NiO
Envelope
g)
218
from the Raman spectra showing increased disorder in the graphitic lattice under higher
temperature annealing.
280 285 290 295
Intensity (a.u.)
Binding Energy (eV)
a)
Raw
C=C
C-C
C-O
C=C
Envelope
280 285 290 295
Intensity (a.u.)
Binding Energy (eV)
b)
Raw
C=C
C-C
C-O
C=O
Envelope
280 285 290 295
Intensity (a.u.)
Binding Energy (eV)
c)
Raw
C=C
C-C
C-O
C=O
Envelope
280 285 290 295
Intensity (a.u.)
Binding Energy (eV)
d)
Raw
C=C
C-C
C-O
C=O
Envelope
219
Figure 4.41. XPS C 1s spectra of the NiO@rGO catalysts: a) NiO@rGO; b) Ni@rGO-300; c)
Ni@rGO-400; d) Ni@rGO-500; e) Ni@rGO-600 and; f) Ni@rGO-700.
The urea electrooxidation properties of the NiO@rGO catalysts were first assessed in half-
cell studies. Basic CV scans are shown in Figure 4.42a with urea and Figure 4.42b without urea.
The onset potentials and urea electrooxidation currrents at 0.65 V are reported in Figure 4.43.
When the nickel oxide catalyst is in the alkaline medium, a layer of hydroxide forms on the surface.
It has been previously reported that the Ni
3+
species is involved from the nickel oxide surface in
the oxidation of urea from the following redox equations [22]:
6NiO + 6OH
-
→ 6NiOOH + 6e
-
(5)
6NiOOH + CO(NH2)2 → 6NiO + CO2 + N2 + 5H2O (6)
The electrochemically active surface areas (ECSAs) for the catalysts were calculated by taking the
area above the curves of the Ni
3+
/Ni
2+
peak from Figure 4.42b at around 0.5-0.6 V and are reported
in Table 4.6. The charge was divided by the coulombic charge density of 514 μC cm
-2
required to
reduce the Ni
3+
species to Ni
2+
species and multiplied by the catalyst loading on the electrode [60].
As the annealing temperature increases, the ECSA of the catalyst decreases. This could be due to
the decrease in nickel oxide characteristic on the catalyst surface at higher temperatures as was
280 285 290 295
Intensity (a.u.)
Binding Energy (eV)
e)
Raw
C-C
C=C
C-O
C=O
Envelope
280 285 290 295
Intensity (a.u.)
Binding Energy (eV)
f)
Raw
C=C
C-C
C-O
C=O
Envelope
220
displayed in the XPS Ni 2p spectra. One of the trends shown is the increase in current density of
the catalysts as the annealing temperature increases. This could also be due to the increased Ni
O
characteristic on the surface of the catalyst from the removal of some of the oxides on the nickel
and oxygen functional groups from annealing at higher temperatures. Even though there is a lower
ECSA present in the catalyst surface at higher annealing temperatures, there is a more favorable
catalyst morphology at the surface for higher urea oxidation activity. Another aspect worth noting
is that the current density is higher for the unsupported nickel oxide catalyst due to the fact there
is only around 30%-wt nickel oxide loading on the rGO supported catalysts. When the current
densities are normalized for the nickel oxide catalyst loading, the current densities become 20.5
mA cm
-2
mg
-1
for the NiO catalyst and 41.7, 55.0, 57.0, 62.0, 64.0, and 68.7 mA cm
-2
mg
-1
for the
NiO@rGO unannealed and annealed at 300, 400, 500, 600 and 700
O
C, respectively. There is a
65% increase in urea electrooxidation currents from the 700
O
C annealed catalyst compared to the
unannealed catalyst. The reduced graphene oxide helps disperse the agglomerated nickel oxide
nanoparticles as shown from the SEM images, leading to higher overall current densities. Another
aspect is the onset potential of urea electrooxidation. As the annealing temperature increases, the
onset potential decreases. This could be due to the increase in the crystallite lattice size as was
displayed in the XRD patterns [57].
221
Figure 4.42. CV scans of the NiO@rGO catalysts in: a) 1.0M KOH + 0.33M urea and; b) 1.0M
KOH only.
-2
0
2
4
6
8
0 0.1 0.2 0.3 0.4 0.5 0.6 0.7
Current Density (mA cm
-2
)
Potential (V vs. MMO)
NiO
NiO@rGO
300
400
500
600
700
-0.2
0
0.2
0.4
0 0.1 0.2 0.3 0.4 0.5 0.6 0.7
Current Density (mA cm
-2
)
Potential (V vs. MMO)
NiO
NiO@rGO
300
400
500
600
700
b)
a)
222
Figure 4.43. Summary of the CV scans of the NiO@rGO catalysts from Figure 4.42a.
The potential stair-step graph of the catalysts in the urea solution is displayed in Figure
4.44. These scans display the overall same trend as was shown in the CV scans in Figure 4.42a.
The effect of the scan rate on the urea electrooxidation was also investigated and the data are
shown in Figure 4.45a for the NiO@rGO-700 and a summary of the current densities vs. scan rates
are shown in Figure 4.45b. Each of the catalysts displays a linear dependency versus the scan rate
signifying a diffusion limiting process occurring in the system [72]. Stability CV scans were also
taken by scanning the catalysts at the same potential window as in Figures 4.42a&b over fifty
cycles. The stability CV scan for the NiO@rGO-600 catalyst is shown in Figure 4.46a with a
summary of the stabilities for each catalyst in Figure 4.46b. There is little relative decrease for
each of the catalysts compared to the initial scan meaning each of the catalysts are relatively stable
in the potential window and the annealing temperature has minimal effects on the stability. Figure
4.47 displays the chronoamperommetry curves of the catalysts under a rotation rate of 1000 RPM
at 0.65 V vs. the reference. Each of the curves displays around the same decay slope signifying
similar long term stabilities of the catalysts.
0.45
0.47
0.49
0.51
0.53
0.55 1
1.5
2
2.5
3
0 100 200 300 400 500 600 700
Onset Potential (V vs. MMO)
Current Density (-mA cm
-2
)
Annealing Temperature (
O
C)
Current Density @ 0.65V
Onset Potential
NiO NiO@rGO
223
Figure 4.44. Step CV for the NiO@rGO catalysts in 1.0 M KOH + 0.33 M urea.
1.2
1.3
1.4
1.5
1.6
1.7
0
0.5
1
1.5
2
2.5
0 100 200 300 400 500
Potential (V)
Current Density (mA cm
-2
)
Time (s)
NiO
NiO@rGO
300
400
500
600
700
Potential
-1
0
1
2
3
4
5
0 0.1 0.2 0.3 0.4 0.5 0.6 0.7
Current Density (mA cm
-2
)
Potential (V vs. MMO)
5
10
20
50
100
a)
224
Figure 4.45. a) CV of the NO@rGO-700 catalyst at different scan rates; b) the relationship
between the square root of the scan rate and the current density of the NiO@rGO catalysts.
0
1
2
3
4
0 5 10 15
Current Density (mA cm
-2
)
ω
1/2
NiO
NiO@rGO
300
400
500
600
700
b)
0
2
4
6
0.5 0.55 0.6 0.65 0.7
Current Density (mA cm
-2
)
Potential (V vs. MMO)
a)
1st
50th
225
Figure 4.46. a) CV of NiO@rGO-600 catalyst cycled 50 times for stability; b) stability plots at
0.65 V for the NiO@rGO catalysts.
Figure 4.47. Chronoamperommetry graph of the NiO@rGO catalysts at 1000 RPM.
90
92
94
96
98
100
0 10 20 30 40 50 60
Current Density (% of Initial)
Cycle Number
NiO
NiO@rGO
300
400
500
600
700
b)
0
0.5
1
1.5
2
0 500 1000 1500 2000 2500 3000
Current Density (mA cm
-2
)
Time (s)
NiO
NiO@rGO
300
400
500
600
700
226
In order to further ascertain the urea oxidation properties of the catalysts, micro-direct
urea/hydrogen peroxide fuel cells were fabricated as was described in Section 4.2.1 and the
polarization curves are displayed in Figure 4.48 with a summary of the OCV and maximum power
densities reported in Table 4.5. The catalysts annealed at 300 and 700
O
C displayed a 52% and
119% increase in power density compared to the unannealed catalyst, respectively. This mimics
the results displayed from the half-cell tests. The unannealed catalyst displayed an 11% decrease
in power density compared with the bare nickel oxide catalyst, most likely due to the fact that there
was only around 30%-wt nickel oxide catalyst in comparison. When normalized for nickel oxide
loading, the power densities of the unannealed catalyst is 0.583 mW cm
-2
mg
-1
versus 0.196 mW
cm
-2
mg
-1
for the nickel oxide.
EIS measurements of the catalysts in the fuel cells were also obtained and displayed in
Figure 4.49 with the equivalent circuit in Figure 4.17. A summary of the charge and mass transfer
resistances derived from the EIS spectra are reported in Table 4.5. From the results in Table 4.5,
the NiO@rGO catalyst displayed a 355% decrease in charge transfer resistance compared to the
plain NiO catalyst. This is most likely due to the increased dispersion of the catalyst particles on
the reduced graphene oxide. The catalysts annealed at 300 and 700
O
C display a 21% and 132%
decrease in charge transfer resistance compared to the unannealed catalyst, respectively. This can
be due to the both the higher currents at the measured potential as well as the larger crystallite
sizes of the nickel as demonstrated in the XRD patterns.
227
Figure 4.48. Polarization curves for the NiO@rGO catalysts in the micro direct urea fuel cells.
Figure 4.49. EIS curves for the NiO@rGO catalysts in the micro direct urea fuel cells.
0
0.02
0.04
0.06
0.08
0.1
0
0.05
0.1
0.15
0.2
0.25
0.3
0 0.2 0.4 0.6 0.8 1 1.2 1.4
Power Density (mW cm
-2
)
Potential (V)
Current Density (mA cm
-2
)
NiO
NiOrGO
300
700
-50
0
50
100
150
200
250
300
0 250 500 750 1000 1250
-Z" (Ohms)
Z' (Ohms)
NiO
NiO@rGO
300
700
228
Table 4.5. Summary of the electrochemical and fuel cell tests of the NiO@rGO catalysts.
Catalyst ECSA
(m
2
g
-1
)
Stability
(% of
Initial)
OCV
(V)
Power
Density
(mW cm
-2
)
Charge Transfer
Resistance
(Ohms)
Mass Transfer
Resistance
(Ohms)
NiO 0.581 96.0 0.27 0.047 270 1090
NiO@rGO 3.995 93.7 0.26 0.042 59.3 475.4
NiO@rGO-300 3.203 95.1 0.25 0.064 49 893
NiO@rGO-400 2.398 96.5 N/A N/A N/A N/A
NiO@rGO-500 2.044 94.7 N/A N/A N/A N/A
NiO@rGO-600 1.587 98.0 N/A N/A N/A N/A
NiO@rGO-700 1.243 96.4 0.29 0.092 25.5 614
4.5: Effect of First Row Transition Metal Doping of Nickel Oxides on the Urea Oxidation
4.5.1: Experimental Methods
First row transition metal doped nickel oxide catalysts, NixM1-xO, (M = Sc, Ti, V, Cr, Mn,
Fe, Co, Cu, Zn) were synthesized via the precipitation method. All of the metal precursors were
derived from transition metal chloride salts (Alfa Aesar) with oxidation states of (II) except for the
titanium and vanadium, whose oxidation states were (III). It is important to note that great care
was taken when handling the TiCl3 salts in the fume hood. Briefly, 1.3747 g NiCl2•6H2O was
dissolved in 100 mL of Millipore water (Direct-Q UV, 18.2 MΩ) under vigorous stirring and
sonication. A certain amount of transition metal chloride salts with a 7%-at ratio to nickel was
then added to the solution under vigorous stirring. The pH was adjusted to 13 by adding a 3.0 M
aqueous sodium hydroxide solution dropwise while visually monitoring the precipitation of the
229
metal hydroxides. The solution was then heated to 50
O
C in an oil bath and stirred for one hour
before being washed and centrifuged with Millipore water. This process was repeated until no
more chlorine ions were detected in the supernatant from the silver nitrate test. The catalyst was
then dried in an oven at 60
O
C overnight and ground up by mortar and pestel before being annealed
under air at 500
o
C for 5 hours in a tube furnace.
All of the characterization and electrochemical half-cell tests were carried out as previously
described in Section 4.2.1.
4.5.2: Results
Figure 4.50a shows the SEM image of the NixZn1-xO catalyst under low magnification.
This is a representative image of all of the doped NixM1-xO catalysts. Figure 4.50b shows the high
magnification SEM image with energy dispersive spectroscopy (EDS) mapping images of the
nickel and zinc elements shown in Figures 4.50c&d. The zinc atoms are well dispersed from
Figure 4.50d indicating a thorough intercalation of the doping atoms in the co-precipitation and
annealing processes. The rest of the transition metal catalyst dopings show very similar EDS
mapping images. EDS spectra were also gathered to quantify the extent of the doping in each of
the catalysts (not shown). The percent doping of the nickel for each transition metal is reported in
Table 4.6. Almost all of the metal dopings fell around the target of 7% except for the titanium and
vanadium. This could be due to the difficulty of handling their chloride salts in the co-precipitation
step.
XRD patterns are shown in Figure 4.51a for the NixM1-x(OH)2 catalysts and Figure 4.51b
for the NixM1-xO catalysts. In each of the patterns in Figure 4.51a, there are peaks at around 19
O
,
230
34
O
, 39
O
, 52
O
, 60
O
, 63
O
, 70
O
, and 74
O
representing the (001), (100), (101), (102), (110), (111),
(013), and (112) Miller Indices of nickel hydroxide phase (Theophrastite, JCPDS #01-073-6992)
[73]. There appears to be minimal difference between the doped and undoped nickel spectra
meaning the absence of impurities in the phases from the doping. In each of the patterns in Figure
4.51b, peaks arise at 37
O
, 43
O
, 63
O
, 75
O
, and 80
O
which represent the (111), (200), (220), (311),
and (222) Miller Indices of the fcc phase nickel oxide (Bunsenite, JCPDS #03-065-2901) [74,75].
There was also an absence of impurity phases in the patterns, meaning the doping doesn’t affect
the overall crystallographic structure of the nickel oxide. The positions of the major peaks vary
by less than 0.1
O
showing minimal variation in the d-spacing of the lattices from the doping. The
average crystallite sizes for the NixM1-xO catalysts were calculated from the Scherrer Equation
[76] from the (200) peak and are shown in Table 4.6. There is peak broadening for the first half
(Sc, Ti, V, Cr, Mn) of the transition metals relative to the second half (Fe, Co, Ni, Cu, Zn). There
appears to be two distinct sets of crystallite sizes arising from the doping. The second half of the
transition metal doped catalysts are around twice the average crystallite size as the first half. The
decreasing of the size is due to the strain to the lattice structure caused by small disorders in the
structure of the doping [77]. The difference in the ionic radius from the first half of the row is
greater to that of nickel than the second half and has a more impeding effect on the crystallite
growth and order. This has been shown to be the case in previous research with nickel oxide
doping [78]. The titanium doping shows the largest crystallite size in the first half of elements,
however, it is possible it arises from the lower doping percentage than the rest of the elements.
231
Figure 4.50. Representative SEM image and EDS mapping of the NixZn1-xO catalyst: a) SEM
image at low magnification; b) SEM/EDS image at high magnification; c) EDS mapping of
nickel at high magnification and; d) EDS mapping of zinc at high magnification.
c) d)
b)
a)
232
Figure 4.51. XRD patterns of the: a) NixM1-x(OH)2 and; b) NixM1-xO catalysts.
Figure 4.52a shows the cyclic voltammetry (CV) curves of the NixM1-xO catalysts in a 1.0
M KOH and 0.33 M urea solution. Figures 4.53a&b shows the plot of both the onset potential and
current density versus the doping element number. Each curve in Figure 4.52a shows a redox peak
10 20 30 40 50 60 70 80 90
Intensity (a.u.)
2θ (degrees)
Zn Cu Ni
Co Fe Mn
Cr V Ti
Sc
a)
10 20 30 40 50 60 70 80 90
Intensity (a.u.)
2θ (degrees)
Zn Cu Ni
Co Fe Mn
Cr V Ti
Sc
b)
b)
233
at around 0.45 V which is the reversible Ni
2+
/Ni
3+
peak [59] as was discussed in Section 4.2.2.
This is more apparent in Figure 4.52b for the catalysts in the 1.0 M KOH solution. It has been
shown that in alkaline conditions, nickel and nickel oxide converts to nickel hydroxide and
oxyhyroxide as was discussed in Section 4.2.2.
Nickel has been shown to be the only active catalyst in first row transition metals, for urea
oxidation. The doping of other transition metals to nickel serves to alter the state of the nickel. In
a previous study, cobalt was found to promote the oxidation state of nickel for electron transfer in
urea oxidation [79]. It is doubtful that the first half of the transition metals would have this effect
and are only hindering the reaction. From Figures 4.52a and 4.53b, the second half of the first row
transition metal doping show an average of three times higher current densities compared to the
first half. This trend mirrors the trend shown from the crystallite sizes from the XRD patterns.
The same type of trend is found in Figure 4.53a with the onset potentials of the catalysts. The
overall crystallite size of the catalysts seems have an effect on the urea oxidation capabilities. The
addition of other metals in the past has been shown to break apart the nickel lattices creating larger
surface areas with active sites for urea oxidation [80]. This has been shown before in numerous
reports with cobalt being one of the preferred first row transition metals to be added with nickel
[41,61,81]. However, the strain induced from the doping of non-compatible elements such as the
first half of the transition metals could impede the active nickel sites for urea oxidation leading to
lower performance. The second half transition metals seem to be more suitable, causing less strain
in the lattice which can possibly lead to higher oxidation performances. Other transition metals
such as rhodium, platinum, and iridium have also been shown to enhance the urea oxidation
properties of nickel [34,35]. These metals are found in the nickel rows (platinum) and cobalt rows
234
(rhodium and iridium) in the transition metals, whose oxidized state ionic radii are very similar to
that of nickel [82].
Other works have shown an increase in urea oxidation activity when alloying the nickel
with another first row transition metal such as cobalt. There was, however a difference in catalyst
preparation whereas the catalysts were prepared via the impregnation method in Xu’s study [61]
or electrodeposition as in Guo’s study [39]. Moreover, in Yan’s study, the onset potential shifted
to lower values when the atomic percentage of cobalt in the nickel-cobalt catalyst increased from
0% to 43% [41]. It is possible that the doping percentage in this study is not high enough to yield
synergistic effects as seen in Yan’s or other studies. The small doping of 7 atomic % could be
slight enough to disrupt the urea oxidation performance of the nickel even with the more
compatible elements to nickel in the second half of the transition metals as carried out in the
aforementioned studies.
-5
0
5
10
15
20
25
30
0 0.1 0.2 0.3 0.4 0.5 0.6 0.7
Current Density (mA cm
-2
)
Potential (V vs. MMO)
Zn Cu
Ni Co
Fe Mn
Cr V
Ti Sc
a)
235
Figure 4.52. CV scans of the NixM1-xO catalysts in: a) 1.0 M KOH + 0.33 M urea and; b) 1.0 M
KOH solution.
-4
-2
0
2
4
6
8
10
0 0.1 0.2 0.3 0.4 0.5 0.6 0.7
Current Density (mA cm
-2
)
Potential (V vs. MMO)
Zn Cu
Ni Co
Fe Mn
Cr V
Ti Sc
b)
0.45
0.5
0.55
0.6
20 21 22 23 24 25 26 27 28 29 30
Onset Potential (V vs. MMO)
Doping Element Number
a)
236
Figure 4.53. Summary plots of the CV scans for the NixM1-xO catalysts for: a) onset potential
and; b) current density from Figure 4.52a.
Table 4.6. Summary of the characterization tests of the Ni-doped catalysts.
Doped
Element
Doping (%
atomic)
Crystallite Size
(nm)
Sc 6.64 4.5
Ti 1.49 5.6
V 3.42 3.9
Cr 6.46 3.0
Mn 6.70 3.9
Fe 5.12 6.4
Co 6.98 8.2
Ni N/A 11.3
Cu 7.99 7.1
Zn 9.26 9.7
0
5
10
15
20
20 21 22 23 24 25 26 27 28 29 30
Current Density (mA cm
-2
)
Doping Element Number
b)
237
4.6: Conclusions
Different types of catalyst supports can significantly affect the catalyst performance for
various types of reactions. The supports can offer greater dispersion of metal nanoparticles as well
as electronic effects to the catalyst particles themselves. In Section 4.2.2, CFx was further
demonstrated to have superior effects as a catalyst support compared with XC72 regarding urea
electrooxidation. The partially fluorinated carbon acted to considerably decrease the nickel
nanoparticles size from a facile aqueous-based reduction with hydrazine, when compared to the
non-fluorinated commercial XC72 carbon support and the bare nickel nanoparticles. Other effects
such as annealing were shown to have an effect on the catalyst supports, which can affect the
performance of the catalysts. Annealing was shown to lead to the slight degradation of the carbon
supports as well as slight oxidation of the nickel nanoparticles leading to lower oxidation activity
in the half-cell tests. The bare nickel catalyst showed minimal change in performance after the
annealing step in comparison most likely due to the large diameters of the particles. The stability
of the Ni/C catalysts however wasn’t as significantly affected.
The supports can also affect the catalyst in different ways via the various post treatment
techniques involved. Annealing was also shown to have varying effects on the rGO supported
nickel-based catalysts depending on the morphology of the nickel. In the case of the Ni/rGO
catalysts in Section 4.3.2, the annealing displayed a detrimental effect on the urea oxidation
kinetics of the nickel. This was from the partial oxidation of the nickel nanoparticles from the
oxygen functional groups in the rGO lattice even though the catalysts were annealed under argon.
However, the opposite trend was displayed in the NiO@rGO catalysts in Section 4.4.2 where the
urea oxidation kinetics increased with increasing annealing temperatures. In this case, the nickel
(O) properties of the nickel oxide increased from the removal of the oxygen functional groups in
238
the support. In both sets of catalysts, the crystallite size increased from the sintering which led to
enhanced onset potentials for urea oxidation. The dispersion on graphene, both from the direct
reduction of nickel salts for the nickel nanoparticles and the prior synthesis and post grafting for
the nickel oxide led to increased urea oxidation properties most likely due to the increased surface
area of the dispersion and possible enhanced charge transfer network present in the system.
The doping effects of other metals can also affect the activity of the catalysts. In Section
4.5.2, the 7% doping of the third row transition metal catalysts were shown to have detrimental
effects on the urea oxidation properties of nickel oxides. The greater the difference in the atomic
radii from that of nickel led to a lower crystallite size of the nickel oxide nanoparticles. This led
to decreases in both the onset potential and urea oxidation currents in the doped nickel catalysts.
Overall, there are numerous factors that can affect the urea electrooxidation on nickel
catalysts. There appears to be a correlation between nickel crystallite size and onset potential of
urea electrooxidation. In Sections 4.2.2, 4.3.2, and 4.4.2, the larger the crystallite size, the more
enhanced the onset potential in the half cells. However, this factor does not always correlate with
higher urea electrooxidation currents and enhanced fuel cell performance. Another factor to
consider is the particle size and dispersion as was demonstrated in Section 4.2.2. Even though the
XC72 supported nickel catalysts displayed larger crystallite sizes, the particle size was larger and
more agglomerated leading to a significantly reduced ECSA and decreased performance in the fuel
cell compared with the CFx supported nickel catalysts. Another factor is the oxidation of the
nickel in the Ni/rGO catalysts at the higher temperatures in Section 4.3.2. Even though the nickel
crystallite size and ECSA had increased, the oxidation of the nickel from the rGO support had
hindered the overall urea electrooxidation, leading to decreased fuel cell performance.
239
As society progresses further into the 21
st
Century, it is increasingly vital that problems
with the increasing demand of energy be solved before it is too late. Various alternative energy
sources, including fuel cells, are already being implemented, however slowly, and with some push
back from society. In order to help make these energy sources more viable, the catalysts of the
future will need to display enhanced kinetics, lower synthetic/raw material costs, scalability,
stability, and ease of implementation into the existing infrastructure. The use of varying catalyst
modification techniques is vital to optimizing catalysts for use in various types of fuel cells,
batteries, sensors, and other energy devices to help power the society of the future.
240
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Abstract (if available)
Abstract
In order for the further commercialization of the myriad of fuel cell technologies that are present today, considerable work still needs to be done to optimize the various parameters of the fuel cell. One of the parameters is the catalyst design including the modification of both catalyst and catalyst support morphologies and properties. ❧ In Chapter 2, partially fluorinated high surface area carbon (CFx) supports for platinum were assessed and compared to the widely implemented Vulcan Carbon (XC72) supports. The Pt/CFx catalysts displayed enhanced oxygen reduction reaction (ORR) kinetics at the lower platinum loadings where the CFx properties were more visible. The platinum supported catalysts were also synthesized via a microwave-assisted polyol reduction method. The microwave reduction method yielded enhanced ORR kinetics compared with the impregnated catalysts while the CFx supported platinum yielded higher kinetics than the XC72 supported platinum. In both instances, the Pt/CFx displayed higher hydrogen fuel cell performances than the Pt/XC72 catalyst. Various phases of manganese dioxide were prepared and grafted onto the aforementioned carbon supports. Again, the CFx supported MnO₂ displayed higher kinetics than the XC72 supported MnO₂ catalysts in both the half-cell and alkaline direct methanol fuel cell testing. The enhanced conductivity from the fluorination in the CFx with the Nafion® ionomer and Tokuyama anionomer in the catalyst layers in the fuel cells led to enhanced performance in the fuel cells in both acidic and alkaline media, respectively. ❧ In Chapter 3, reduced graphene oxide (rGO) was synthesized by various routes including Hummer’s method and electrochemical exfoliation. The graphene oxide (GO) synthesized from the Hummer’s method was dispersed in various aqueous solutions with pH values of 1-13 before being reduced by sodium borohydride. The lower pH values led to less repair of the sp² graphitic lattice as confirmed by Raman and X-Ray Photoelectron Spectroscopy (XPS) due to the increased degradation of sodium borohydride in acidic solutions. This led to lower ORR kinetics in alkaline half-cell testing. Reduced graphene oxide was also synthesized from electrochemical exfoliation from graphite rods and deposited on platinum in a one pot procedure. The potential of exfoliation affected the lattice of the rGO impacting the ORR activity of the platinum supported catalysts. Graphene oxide was also kept in an emulsion form from Hummer’s method without drying. The GOem displayed similar structural characteristics in both the oxidized and reduced forms to that of the dried form as well as with platinum reduced on it by impregnation and microwave reduction. ❧ In Chapter 4, nickel was deposited onto various supports to examine the effects on urea electrooxidation. As in Chapter 2, nickel was impregnated onto both CFx and XC72 to examine to effects of fluorination. The nickel displayed smaller particle sizes on the CFx compared with the XC72. This led to enhanced urea electrooxidation kinetics in both the half-cell and micro fuel cell testing. Mild oxidative annealing was also applied to the catalysts, which resulted in lower activity. Nickel was also reduced onto GO and annealed at temperatures from 300–700℃ under argon to study the effects. The rGO displayed greater disorder from the Raman and XPS spectra from the removal of the oxygen functional groups in the lattice. This led to the partial oxidation of the nickel nanoparticles, leading to decreased kinetics in both the half-cell and micro fuel cell tests. Nickel oxide was also grafted onto rGO and annealed at the same temperatures under argon. The same effects were viewed in the graphitic lattice. However, the catalysts showed an increase in urea electrooxidation activity at the higher annealing temperatures in the half-cell and micro fuel cell tests. Finally, nickel was co-doped with all of the third row transition metal elements at a 7%-wt and annealed under air to form nickel oxides. This doping led to decreased urea electrooxidation activity from the strain in the nickel lattice.
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Glass, Dean Edson
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Core Title
The modification of catalysts and their supports for use in various fuel cells
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College of Letters, Arts and Sciences
Degree
Doctor of Philosophy
Degree Program
Chemistry
Publication Date
01/30/2019
Defense Date
09/19/2018
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catalysts,CFx,fuel cells,graphene,graphene oxide,nickel,OAI-PMH Harvest,oxygen reduction,platinum,reduced graphene oxide,urea oxidation,Vulcan carbon,XC-72
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Prakash, Surya (
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Tags
catalysts
CFx
fuel cells
graphene
graphene oxide
nickel
oxygen reduction
platinum
reduced graphene oxide
urea oxidation
Vulcan carbon
XC-72