University of Southern California Dissertations and Theses

Electrocatalysts for direct liquid-feed fuel cells and advanced electrolytes for lithium-ion batteries

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Electrocatalysts for direct liquid-feed fuel cells and advanced electrolytes for lithium-ion batteries

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Content
ELECTROCATALYSTS FOR DIRECT LIQUID-FEED FUEL CELLS AND ADVANCED
ELECTROLYTES FOR LITHIUM-ION BATTERIES

by

Frederick Charles Krause

A Dissertation Presented to the
FACULTY OF THE USC GRADUATE SCHOOL
UNIVERSITY OF SOUTHERN CALIFORNIA
In Partial Fulfillment of the
Requirements for the Degree
DOCTOR OF PHILOSOPHY
(CHEMISTRY)

August 2012

Copyright 2012 Frederick Charles Krause
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Epigraph

« If you understand, things are just as they are. If you do not understand, things are just as
they are. »
Zen proverb

« Il est vain, si l’on plante un chêne, d’espérer s’abriter bientôt sous son feuillage. »
Antoine de Saint-Exupéry
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Dedication

To my parents, Ann Orcutt Krause and Howard Arthur Krause.
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Acknowledgements

There are many people I would like to thank, without whom this pursuit would
never have been completed. My advisor, Dr. G. K. Surya Prakash, gave me the opportunity,
time, and support to learn new things and explore new ideas; and the guidance to make
sense of everything and keep it going in the right direction. For their contributions to my
work at USC, I am also indebted to all the members of the Olah/Prakash group, and in
particular Dr. Federico Viva, Dr. Bo Yang, Fang Wang, and John-Paul Jones. I must also
thank Dr. Robert Aniszfeld for his constant assistance in keeping our lab running, and
Dr. S. R. Narayanan for his valuable advice and discussions.
At the Jet Propulsion Laboratory, my work would not have been possible without
the guidance of Dr. Marshall Smart who provided the vision for these projects and gave me
the foundations to contribute to them. Dr. Kiah Smith was an invaluable mentor.
Finally, I wish to thank my parents for their lifelong support, and my wife Kate
Loughrey for her never-ending patience and encouragement.

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Table of Contents

Epigraph ................................................................................................................................ ii
Dedication ............................................................................................................................. iii
Acknowledgements .............................................................................................................. iv
List of Tables ........................................................................................................................ ix
List of Figures ....................................................................................................................... x
Abbreviations ....................................................................................................................... xxi
Abstract ................................................................................................................................. xxiii
Chapter 1: Optimized membrane-electrode assemblies based on alkaline anion-
exchange membranes for direct methanol fuel cell applications ................. 1
1.1 Introduction to fuel cells ................................................................................. 1
1.1.1 Historical development ................................................................ 1
1.1.1.1 The Grove gas battery .............................................. 2
1.1.1.2 Alkaline fuel cells (AFC) .......................................... 6
1.1.1.3 Polymer electrolyte membrane fuel cells
(PEMFC) .................................................................... 10
1.1.1.4 Direct methanol fuel cells (DMFC) ........................ 12
1.1.2 Principles of operation ................................................................. 15
1.1.3 Catalyst development ................................................................... 18
1.1.3.1 The anode reaction ................................................... 18
1.1.3.2 The cathode reaction ................................................ 21
1.1.4 Scope of this work ........................................................................ 22
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1.2 Experimental methods .................................................................................... 24
1.3 Results and discussion ..................................................................................... 26
1.3.1 Addition of KOH to the fuel ........................................................ 26
1.3.2 Effects of temperature .................................................................. 26
1.3.3 Effects of oxidant flow rate .......................................................... 27
1.3.4 Effects of methanol concentration .............................................. 28
1.3.5 Effects of potassium hydroxide concentration .......................... 28
1.3.6 Effects of electrode material ........................................................ 30
1.3.7 Comparison of performance of AAEM and Nafion
®
cells ........ 33
1.3.8 Cell durability ................................................................................ 36
1.4 Conclusions ...................................................................................................... 38
1.5 References ........................................................................................................ 39
Chapter 2: Catalysts for direct formic acid fuel cells ........................................................ 46
2.1 Introduction to direct formic acid fuel cells .................................................. 46
2.1.1 Recent developments ................................................................... 48
2.1.2 Mechanistic studies ...................................................................... 50
2.1.3 Scope of this work ........................................................................ 53
2.2 Experimental methods .................................................................................... 54
2.3 Results and discussion ..................................................................................... 58
2.3.1 Palladium-gold .............................................................................. 58
2.3.1.1 Carbon-supported catalysts ..................................... 58
2.3.1.2 Tantalum carbide as a support ................................ 61
2.3.1.3 Formate oxidation .................................................... 69
2.3.2 Palladium-platinum ...................................................................... 71
2.3.3 Palladium and other metals ......................................................... 76
2.4 Conclusions ...................................................................................................... 79
2.5 References ........................................................................................................ 81
Chapter 3: Advanced electrolytes for nonflammable lithium-ion batteries ................... 85
3.1 Introduction to lithium-ion batteries ............................................................ 85
3.1.1 Historical overview ...................................................................... 85
3.1.1.1 Secondary batteries ................................................... 85
3.1.1.2 Lithium metal and lithium-ion ............................... 88
3.1.1.3 Lithium-ion and beyond .......................................... 90
3.1.2 Operating principles and requirements...................................... 92
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3.1.2.1 The “rocking chair” cell ............................................ 92
3.1.2.2 Energy and power density ....................................... 94
3.1.2.3 Cycle and storage life ................................................ 95
3.1.2.4 Operating temperature range .................................. 96
3.1.2.5 Safety .......................................................................... 96
3.1.3 Electrode materials ....................................................................... 98
3.1.3.1 Positive electrode materials ..................................... 98
3.1.3.2 Negative electrode materials .................................... 101
3.1.4 Electrolyte solvents and additives ............................................... 107
3.1.4.1 Carbonates ................................................................. 108
3.1.4.2 Ethers and esters ....................................................... 109
3.1.4.3 Fluorinated solvents ................................................. 111
3.1.4.4 Additives .................................................................... 113
3.1.5 Lithium salts .................................................................................. 120
3.1.5.1 Selection criteria ....................................................... 120
3.1.5.2 LiPF
6
.......................................................................... 120
3.1.5.3 Other salts .................................................................. 122
3.1.6 Scope of this work ........................................................................ 123
3.2 Experimental methods .................................................................................... 126
3.2.1 Electrolyte preparation ................................................................ 126
3.2.2 Cell construction .......................................................................... 126
3.2.3 Electrochemical and electrical characterization ......................... 127
3.3 Results and discussion ..................................................................................... 128
3.3.1 Studies in MCMB/LiNiCoO
2
cells ............................................. 128
3.3.1.1 Incorporation of flame-retardant additives ........... 128
3.3.1.2 Incorporation of fluorinated carbonate
co-solvents ................................................................. 132
3.3.1.3 Reduced-flammability electrolytes in
three-electrode cells .................................................. 136
3.3.2 FEC and LiBOB in MPG-111/LiNiCoAlO
2
cells ...................... 162
3.3.2.1 Electrochemical characterization ............................ 158
3.3.2.2 Effect of FEC on discharge capacity ........................ 162
3.3.3 MPG-111/LiNiCoAlO
2
cells with high additive content ......... 180
3.3.3.1 Formation characteristics ........................................ 180
3.3.3.2 Electrochemical characterization ............................ 183
3.3.3.3 Discharge characteristics .......................................... 197
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3.3.3.4 Rate characterization ................................................ 198
3.3.4 Conclusions ................................................................................... 204
3.3.5 References ..................................................................................... 207

Bibliography ......................................................................................................................... 222

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List of Tables

Table 3-1. Polarization resistance values measured in FEC- and TFEMC-
containing MCMB/NCO cells. ......................................................................... 138
Table 3-2. Polarization resistance values measured in bTFEC-containing
MCMB/NCO cells. ............................................................................................ 138
Table 3-3. MPG-111 anode polarization resistance values recorded for FEC- and
TPPa- containing Graphite/NCA cells at full state-of-charge. ..................... 164
Table 3-4. LiNiCoAlO2 cathode polarization resistance values recorded for FEC-
and TPPa- containing Graphite/NCA cells at full state-of-charge. .............. 164
Table 3-5. Formation characteristics of three-electrode MPG-111/LiNi
x
Co
1-x
AlO
2
cells. ..................................................................................................................... 181
Table 3-6. MPG-111 anode polarization resistance values recorded for FEC- and
LiBOB-containing Graphite/NCA cells at full state-of-charge. .................... 184
Table 3-7. LiNiCoAlO
2
cathode polarization resistance values recorded for FEC-
and LiBOB-containing graphite/NCA cells at full state-of-charge. ............. 184
Table 3-8. Discharge characteristics for FEC- and LiBOB-containing graphite/
NCA cells. ........................................................................................................... 199

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List of Figures

Figure 1-1. Schematic representation of Grove’s 1839 experiment on “the
combination of gases by platinum.” ................................................................ 3
Figure 1-2. Representation of the first working fuel cell constructed by Grove in
1842. Adapted from [5] (arrows were intended to indicate “positive
current” flow). ................................................................................................... 5
Figure 1-3. Cross-sectional diagrams of (a) a membrane-electrode assembly
(not to scale) and (b) an assembled fuel cell. .................................................. 25
Figure 1-4. Effect of temperature: polarization and power density curves, 1 M
methanol + 1 M KOH, 1270 mL min
-1
O
2
, wet-proofed electrodes. .......... 27
Figure 1-5. Effect of methanol concentration: polarization and power density
curves, 1 M KOH, 90 °C, 1270 mL min
-1
O
2
, wet-proofed electrodes. ....... 29
Figure 1-6. Effect of increasing KOH concentration: polarization and power
density curves, 1 M methanol, 90 °C, 1270 mL min
-1
O
2
, wet-proofed
electrodes. .......................................................................................................... 30
Figure 1-7. Effect of electrode wet-proofing: polarization and power density
curves, 2 M KOH + 1 M methanol, 90 °C, 1270 mL min
-1
O
2
. .................... 32
Figure 1-8. Mass transport limitation at low temperature with wet-proofed anode:
polarization and power density curves, 2 M KOH + 1 M methanol,
30 °C, 1270 mL min
-1
O
2
. ................................................................................. 33
Figure 1-9. Comparison of Tokuyama MEA with Nafion MEA: polarization and
power density curves, 2 M KOH + 1 M methanol, 90 °C, 1270 mL min
-1

O
2
(dotted lines represent iR-corrected polarization curves). ..................... 34
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Figure 1-10. Polarization curves showing the effect of reduced oxygen flow rate at
high current densities; 2 M KOH + 1 M methanol, 90 °C; high-flow
represents 1270 mL min
-1
O
2
, low flow represents 200 mL min
-1
O
2
. ..... 35
Figure 1-11. Degradation behavior of a Tokuyama MEA with wet-proofed
electrodes: initial performance, followed by observed degradation after
high concentration fuel experiments, and finally partial recovery after
12h/60 °C water flush; experiments were separated by 48 hour intervals;
polarization and power density curves, 2 M KOH + 1 M methanol,
90 °C, 1270 mL min
-1
O
2
. ............................................................................... 37
Figure 2-1. TEM images of PdAu-A/[C0.5,TaC0.5] at a) 59kx magnification
(scale bar is 100 nm) and b) 160kx magnification (scale bar is 20 nm). ...... 59
Figure 2-2. Current-voltage curves for DFAFCs with palladium and palladium-
gold anode catalysts. Recorded with 2 M HCOOH and pure O
2

at 90 °C. .............................................................................................................. 59
Figure 2-3. Power density curves for DFAFCs with palladium and palladium-
gold anode catalysts. Recorded with 2 M HCOOH and pure O
2

at 90 °C. .............................................................................................................. 60
Figure 2-4. Power density curves normalized to precious metal (Pd and/or Au)
content for DFAFCs with palladium and palladium-gold anode
catalysts. Recorded with 2 M HCOOH and pure O
2
at 90 °C. ...................... 60
Figure 2-5. Chronoamperometric measurements of palladium and palladium-
gold catalysts supported on Vulcan XC-72R and unsupported,
recorded at +0.300 V vs. SHE in 0.5 M H
2
SO
4
/ 0.5 M HCOOH. ................. 63
Figure 2-6. Chronoamperometric measurements of palladium and palladium-
gold catalysts supported on Vulcan XC-72R and tantalum carbide,
recorded at +0.300 V vs. SHE in 0.5 M H
2
SO
4
/ 0.5 M HCOOH. ................ 64
Figure 2-7. Chronoamperometric measurements of palladium-gold catalysts
normalized to initial current, recorded at +0.300 V vs. SHE in
0.5 M H
2
SO
4
/0.5 M HCOOH. ......................................................................... 65
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Figure 2.8. Cyclic voltammograms of palladium catalysts supported on Vulcan
carbon XC-72R and tantalum carbide with currents normalized to Pd
mass. Recorded in 0.5 M H
2
SO
4
/0.5 M HCOOH at 10 mV s
-1
. ................... 66
Figure 2-9. Cyclic voltammograms of palladium-gold catalysts supported on
Vulcan carbon XC-72R and tantalum carbide with currents
normalized to Pd+Au mass. Recorded in 0.5 M H
2
SO
4
/0.5 M HCOOH
at 10 mV s
-1
. ...................................................................................................... 67
Figure 2-10. Cyclic voltammograms of palladium and palladium-gold catalysts
supported on Vulcan carbon XC-72R and tantalum carbide normalized
to Pd and/or Au mass. Recorded in 0.5 M H
2
SO
4
/0.5 M HCOOH
at 10 mV s
-1
. .................................................................................................... 68
Figure 2-11. Current-voltage (dash-dot) and power density (dash) curves for
MEAs with best-performing supported and unsupported palladium-
gold catalysts recorded with 2 M HCOOH and pure O
2
at 90 °C. .............. 69
Figure 2-12. Current-voltage and power density curves for MEAs with best-
performing supported and unsupported palladium-gold catalysts
normalized to Pd/Au mass; recorded with 2 M HCOOH and pure
O
2
at 90 °C. ...................................................................................................... 70
Figure 2-13. Comparison of cyclic voltammograms of PdAu-B/[C0.5+TaC0.5]
in formic acid and potassium formate solutions normalized to Pd+Au
mass. Recorded at 10 mV s
-1
. ......................................................................... 71
Figure 2-14. Comparison of chronoamperometry of PdAu-B/[C0.5+TaC0.5] in
formic acid and potassium formate solutions normalized to Pd+Au
mass. Recorded at 300 rpm. ........................................................................... 72
Figure 2-15. Comparison of chronoamperometry of PdAu-B/[C0.5+TaC0.5] in
formic acid and potassium formate solutions expressed as a percentage
of initial current. Recorded at 300 rpm. ....................................................... 72
Figure 2-16. Comparison of cyclic voltammograms of carbon-tantalum-carbide-
supported PdAu and PdPt catalysts normalized to Pd+Au mass.
Recorded in 0.5 M H
2
SO
4
/0.5 M HCOOH at 10 mV s
-1
. ........................... 74
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Figure 2-17. Chronoamperometry of PdPt/[C0.5,TaC0.5]. Recorded in
0.5 M H
2
SO
4
/0.5 M HCOOH at 300 rpm. ................................................... 75
Figure 2-18. Comparison of chronoamperometry of PdAu-B/[C0.5+TaC0.5] and
PdPt/[C0.5,TaC0.5] expressed as a percentage of initial current.
Recorded in 0.5 M H
2
SO
4
/0.5 M HCOOH at 300 rpm. ............................. 76
Figure 2-19. Current-voltage curves for DFAFCs with palladium-M
x
catalysts.
Recorded with 2 M HCOOH and pure O
2
at 90 °C. .................................... 78
Figure 2-20. Power density curves for DFAFCs with palladium-M
x
catalysts.
Recorded with 2 M HCOOH and pure O
2
at 90 °C. .................................... 78
Figure 2-21. Power density curves for DFAFCs with palladium-M
x
catalysts,
with currents normalized to anode precious metal (Pd and/or Au)
mass. Recorded with 2 M HCOOH and pure O
2
at 90 °C. .......................... 79
Figure 3-1. Schematic diagram of a lithium-ion cell. ....................................................... 93
Figure 3-2. Structures of electrolyte components used in this study. ............................. 125
Figure 3-3. Cyclic voltammetry of FRA-containing electrolytes on a platinum
working electrode. Recorded between 0.01 V – 4.00 V vs. Li/Li
+
,
5 mV s
-1
. ............................................................................................................ 130
Figure 3-4. Cyclic voltammetry of FRA-containing electrolytes on a platinum
working electrode. Recorded between 1.00 V – 5.00 V vs. Li/Li
+
,
5 mV s
-1
. ............................................................................................................ 131
Figure 3-5. Cyclic voltammetry of FRA-containing electrolytes on a platinum
working electrode. Recorded between 2.00 V – 6.00 V vs. Li/Li
+
,
5 mV s
-1
. ............................................................................................................ 132
Figure 3-6. Cyclic voltammetry of fluorocarbonate-containing electrolytes on a
platinum working electrode. Recorded between 0.01 V – 4.00 V vs.
Li/Li
+
, 20 mV s
-1
(ninth cycle shown). .......................................................... 133
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Figure 3-7. Cyclic voltammetry of fluorocarbonate-containing electrolytes on a
platinum working electrode. Recorded between 1.00 V – 5.00 V vs.
Li/Li
+
, 20 mV s
-1
. ............................................................................................. 134
Figure 3-8. Cyclic voltammetry of fluorocarbonate-containing electrolytes on a
platinum working electrode. Recorded between 2.00 V – 6.00 V vs.
Li/Li
+
, 20 mV s
-1
. ............................................................................................. 135
Figure 3-9. Tafel polarization of MCMB electrodes of FEC- and TFEMC-
containing cells at room temperature. ........................................................... 140
Figure 3-10. Tafel polarization of LiNi
x
Co
1-x
O
2
electrodes of FEC- and TFEMC-
containing cells at room temperature. .......................................................... 141
Figure 3-11. Tafel polarization of MCMB electrodes of FEC- and TFEMC-
containing cells at 0 °C. .................................................................................. 141
Figure 3-12. Tafel polarization of LiNi
x
Co
1-x
O
2
electrodes of FEC- and TFEMC-
containing cells at 0 °C. .................................................................................. 142
Figure 3-13. Tafel polarization of MCMB electrodes of FEC- and TFEMC-
containing cells at -20 °C. ............................................................................... 142
Figure 3-14. Tafel polarization of LiNi
x
Co
1-x
O
2
electrodes of FEC- and TFEMC-
containing cells at -20 °C. ............................................................................... 143
Figure 3-15. Tafel polarization of MCMB electrodes of FEC- and TFEMC-
containing cells at -40 °C. ............................................................................... 143
Figure 3-16. Tafel polarization of LiNi
x
Co
1-x
O
2
electrodes of FEC- and TFEMC-
containing cells at -40 °C. ............................................................................... 144
Figure 3-17. Tafel polarization of MCMB electrodes of bTFEC-containing cells
at room temperature. ..................................................................................... 146
Figure 3-18. Tafel polarization of LiNi
x
Co
1-x
O
2
electrodes of bTFEC-containing
cells at room temperature. ............................................................................. 146
Figure 3-19. Tafel polarization of MCMB electrodes of bTFEC-containing cells
at 0 °C. .............................................................................................................. 147
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Figure 3-20. Tafel polarization of LiNi
x
Co
1-x
O
2
electrodes of bTFEC-containing
cells at 0 °C. ...................................................................................................... 147
Figure 3-21. Tafel polarization of MCMB electrodes of bTFEC-containing cells
at -20 °C. .......................................................................................................... 148
Figure 3-22. Tafel polarization of LiNi
x
Co
1-x
O
2
electrodes of bTFEC-containing
cells at -20 °C. .................................................................................................. 148
Figure 3-23. Tafel polarization of MCMB electrodes of bTFEC-containing cells
at -40 °C. Stable measurements were not obtained for cells not shown. .. 149
Figure 3-24. Tafel polarization of LiNi
x
Co
1-x
O
2
electrodes of bTFEC-containing
cells at -40 °C. Stable measurements were not obtained for cells not
shown. .............................................................................................................. 149
Figure 3-25. EIS of MCMB electrodes of FEC- and TFEMC-containing cells at
room temperature. .......................................................................................... 153
Figure 3-26. EIS of LiNi
x
Co
1-x
O
2
electrodes of FEC- and TFEMC-containing cells
at room temperature. ..................................................................................... 153
Figure 3-27. EIS of MCMB electrodes of FEC- and TFEMC-containing cells
at 0 °C. .............................................................................................................. 154
Figure 3-28. EIS of LiNi
x
Co
1-x
O
2
electrodes of FEC- and TFEMC-containing cells
at 0 °C ............................................................................................................... 154
Figure 3-29. EIS of MCMB electrodes of FEC- and TFEMC-containing cells
at -20 °C. .......................................................................................................... 155
Figure 3-30. EIS of LiNi
x
Co
1-x
O
2
electrodes of FEC- and TFEMC-containing cells
at -20 °C. .......................................................................................................... 155
Figure 3-31. EIS of MCMB electrodes of FEC- and TFEMC-containing cells
at -40 °C. .......................................................................................................... 156
Figure 3-32. EIS of LiNi
x
Co
1-x
O
2
electrodes of FEC- and TFEMC-containing cells
at -40 °C. The impedance of FP04 was apparently too high to obtain a
successful measurement. ................................................................................ 156
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Figure 3-33. EIS of MCMB electrodes of bTFEC-containing cells at room
temperature. .................................................................................................... 158
Figure 3-34. EIS of LiNi
x
Co
1-x
O
2
electrodes of bTFEC-containing cells at room
temperature. .................................................................................................... 158
Figure 3-35. EIS of MCMB electrodes of bTFEC-containing cells at 0 °C. ................... 159
Figure 3-36. EIS of LiNi
x
Co
1-x
O
2
electrodes of bTFEC-containing cells at 0 °C.
Stable measurements for GB04 were not obtained. .................................... 159
Figure 3-37. EIS of MCMB electrodes of bTFEC-containing cells at -20 °C. ................ 160
Figure 3-38. EIS of LiNi
x
Co
1-x
O
2
electrodes of bTFEC-containing cells at -20 °C. ...... 160
Figure 3-39. EIS of MCMB electrodes of bTFEC-containing cells at -40 °C. A
stable measurement for GB01 was not obtained. ........................................ 161
Figure 3-40. EIS of LiNi
x
Co
1-x
O
2
electrodes of bTFEC-containing cells at -40 °C.
A stable measurement for GB01 was not obtained ..................................... 161
Figure 3-41. Tafel polarization of MPG-111 electrodes of FEC- and TPPA-
containing cells (MPG-111/NCA) at room temperature. .......................... 166
Figure 3-42. Tafel polarization of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and TPPA-
containing cells (MPG-111/NCA) at room temperature. .......................... 166
Figure 3-43. Tafel polarization of MPG-111 electrodes of FEC- and TPPA-
containing cells (MPG-111/NCA) at 0 °C. ................................................... 167
Figure 3-44. Tafel polarization of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and TPPA-
containing cells (MPG-111/NCA) at 0 °C. ................................................... 167
Figure 3-45. Tafel polarization of MPG-111 electrodes of FEC- and TPPA-
containing cells (MPG-111/NCA) at -20 °C. ............................................... 168
Figure 3-46. Tafel polarization of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and TPPA-
containing cells (MPG-111/NCA) at -20 °C. ............................................... 168
! $#""!
Figure 3-47. Tafel polarization of MPG-111 electrodes of FEC- and TPPA-
containing cells (MPG-111/NCA) at -30 °C. ............................................... 169
Figure 3-48. Tafel polarization of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and TPPA-
containing cells (MPG-111/NCA) at -30 °C. ............................................... 169
Figure 3-49. Tafel polarization of MPG-111 electrodes of FEC- and TPPA-
containing cells (MPG-111/NCA) at -40 °C. ............................................... 170
Figure 3-50. Tafel polarization of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and TPPA-
containing cells (MPG-111/NCA) at -40 °C. ............................................... 170
Figure 3-51. EIS of MPG-111 electrodes of FEC- and TPPa-containing cells
(MPG-111/NCA) at room temperature. ...................................................... 172
Figure 3-52. EIS of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and TPPa-containing cells
(MPG-111/NCA) at room temperature. ...................................................... 172
Figure 3-53. EIS of MPG-111 electrodes of FEC- and TPPa-containing cells
(MPG-111/NCA), 0 °C. .................................................................................. 173
Figure 3-54. EIS of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and TPPa-containing cells
(MPG-111/NCA), 0 °C. .................................................................................. 173
Figure 3-55. EIS of MPG electrodes of FEC- and TPPa-containing cells
(MPG-111/NCA) at -20 °C. ........................................................................... 174
Figure 3-56. EIS of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and TPPa-containing cells
(MPG-111/NCA) at -20 °C. ........................................................................... 174
Figure 3-57. EIS of MPG electrodes of FEC- and TPPa-containing cells (MPG-
111/NCA), -30 °C. ......................................................................................... 175
Figure 3-58. EIS of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and TPPa-containing cells
(MPG-111/NCA) at -30 °C. ........................................................................... 175
Figure 3-59. EIS of MPG electrodes of FEC- and TPPa-containing cells
(MPG-111/NCA) at -40 °C. ........................................................................... 176
! $#"""!
Figure 3-60. EIS of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and TPPa-containing cells
(MPG-111/NCA) at -40 °C. ........................................................................... 176
Figure 3-61. Relative discharge capacities of FEC-containing cells versus EC-
containing cells. .............................................................................................. 179
Figure 3-62. Tafel polarization of MPG-111 electrodes of FEC- and LiBOB-
containing cells (MPG-111/NCA) at room temperature. .......................... 185
Figure 3-63. Tafel polarization of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and LiBOB-
containing cells (MPG-111/NCA) at room temperature. .......................... 185
Figure 3-64. Tafel polarization of MPG-111 electrodes of FEC- and LiBOB-
containing cells (MPG-111/NCA) at 0 °C. ................................................... 186
Figure 3-65. Tafel polarization of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and LiBOB-
containing cells (MPG-111/NCA) at 0 °C. ................................................... 186
Figure 3-66. Tafel polarization of MPG-111 electrodes of FEC- and LiBOB-
containing cells (MPG-111/NCA) at -20 °C. ............................................... 187
Figure 3-67. Tafel polarization of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and LiBOB-
containing cells (MPG-111/NCA) at -20 °C. ............................................... 187
Figure 3-68. Tafel polarization of MPG-111 electrodes of FEC- and LiBOB-
containing cells (MPG-111/NCA) at -30 °C. ............................................... 188
Figure 3-69. Tafel polarization of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and LiBOB-
containing cells (MPG-111/NCA) at -30 °C. ............................................... 188
Figure 3-70. Tafel polarization of MPG-111 electrodes of FEC- and LiBOB-
containing cells (MPG-111/NCA) at -40 °C. ............................................... 189
Figure 3-71. Tafel polarization of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and LiBOB-
containing cells (MPG-111/NCA) at -40 °C. ............................................... 189
Figure 3-72. EIS of MPG-111 electrodes of FEC- and LiBOB-containing cells
(MPG-111/NCA) at room temperature. ...................................................... 191
! $"$!
Figure 3-73. EIS of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and LiBOB-containing
cells (MPG-111/NCA) at room temperature. .............................................. 191
Figure 3-74. EIS of MPG electrodes of FEC- and LiBOB-containing cells
(MPG-111/NCA) at 0 °C. ............................................................................... 192
Figure 3-75. EIS of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and LiBOB-containing
cells (MPG-111/NCA) at 0 °C. ...................................................................... 192
Figure 3-76. EIS of MPG electrodes of FEC- and LiBOB-containing cells
(MPG-111/NCA) at -20 °C. ........................................................................... 193
Figure 3-77. EIS of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and LiBOB-containing
cells (MPG-111/NCA) at -20 °C. ................................................................... 193
Figure 3-78. EIS of MPG electrodes of FEC- and LiBOB-containing cells
(MPG-111/NCA) at -30 °C. ........................................................................... 194
Figure 3-79. EIS of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and LiBOB-containing
cells (MPG-111/NCA) at -30 °C. ................................................................... 194
Figure 3-80. EIS of MPG electrodes of FEC- and LiBOB-containing cells
(MPG-111/NCA) at -40 °C. ........................................................................... 195
Figure 3-81. EIS of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and LiBOB-containing
cells (MPG-111/NCA) at -40 °C. ................................................................... 195
Figure 3-82. Discharge rate study at -20 °C of MPG-111/LiNi
x
Co
1-x
AlO
2
cell
containing 1.0 M LiPF
6
, 0.15 M LiBOB in EC+EMC+TPPa
(20:65:15 v/v). ................................................................................................. 201
Figure 3-83. Discharge rate study at -20 °C of MPG-111/LiNi
x
Co
1-x
AlO
2
cell
containing 1.0 M LiPF
6
, 0.15 M LiBOB in EC+FEC+EMC+TPPa
(10:10:65:15 v/v). ............................................................................................ 201
Figure 3-84. Discharge rate study at -20 °C of MPG-111/LiNi
x
Co
1-x
AlO
2
cell
containing 1.0 M LiPF
6
, 0.15 M LiBOB in FEC+EMC+TPPa
(20:65:15 v/v). ................................................................................................. 202
! $$!
Figure 3-85. Discharge rate study at -20 °C of MPG-111/LiNi
x
Co
1-x
AlO
2
cell
containing 1.0 M LiPF
6
, 0.20 M LiBOB in FEC+EMC+TPPa
(20:65:15 v/v). ................................................................................................. 202
Figure 3-86. Discharge rate study at -20 °C of MPG-111/LiNi
x
Co
1-x
AlO
2
cell
containing 1.0 M LiPF
6
, 0.25 M LiBOB in FEC+EMC+TPPa
(20:65:15 v/v). ................................................................................................. 203
Figure 3-87. Discharge rate study at -20 °C of MPG-111/LiNi
x
Co
1-x
AlO
2
cell
containing 1.0 M LiPF
6
in EC+DEC+DMC (1:1:1 v/v). ............................ 203
Figure 3-88. Anode behavior during -20 °C rate study of FEC- and LiBOB-
containing MPG-111/LiNi
x
Co
1-x
AlO
2
cells. ................................................ 205

! $$"!

Abbreviations

bTFEC bis(trifluoroethyl) carbonate
CA Chronoamperometry
CV Cyclic voltammetry
DAMFC Direct alkaline methanol fuel cell
DEC Diethyl carbonate
DMC Dimethyl carbonate
DMFC Direct methanol fuel cell
E
0
Standard reduction potential referenced to the standard hydrogen electrode
EC Ethylene carbonate
EIS Electrochemical impedance spectroscopy
EMC Ethyl methyl carbonate
FEC Monofluoroethylene carbonate
FAO Formic acid oxidation
HOR Hydrogen oxidation reaction
MCMB Mesocarbon microbeads
! $$""!
MEA Membrane electrode assembly
MOR Methanol oxidation reaction
NCA Lithium nickel cobalt aluminum oxide
NCO Lithium nickel cobalt oxide
ORR Oxygen reduction reaction
PC Propylene carbonate
PEMFC Polymer electrolyte membrane fuel cell
SEI Solid-electrolyte interphase
TFEMC Trifluoroethyl methyl carbonate
VC Vinylene carbonate

! $$"""!

Abstract

Direct alkaline methanol fuel cells were constructed and tested using a hydroxide-
exchange membrane. The operating conditions and material components of the membrane
electrode assemblies (MEAs), such as temperature, oxygen flow rate, fuel composition, and
electrode material, were optimized. These cells were compared to analogous Nafion®-based
MEAs and the power output was found to be comparable under similar conditions.
Direct formic acid fuel cells (DFAFC) were constructed and tested using various
palladium-based anode electrocatalysts. The power output, catalytic activity, and durability
of these MEAs were evaluated. It was found that Pd-Au catalysts supported on a carbon-
TaC blend were more active and durable than unsupported or carbon-supported materials.
New electrolyte formulations for lithium-ion batteries were prepared and tested in
terms of electrochemical stability, lithium intercalation/deintercalation kinetics, and cycling
and rate capabilities with several state-of-the-art electrode systems. Flame-retardant
additives and fluorinated co-solvents were examined in particular, and several formulations
were identified which provide greater safety with comparable or better performance to
known baseline electrolytes.
!
1

Optimized membrane-electrode assemblies
based on alkaline anion exchange membranes
for direct methanol fuel cell applications

1.1 Introduction to fuel cells
1.1.1 Historical development
The fuel cell has existed in various forms since the early nineteenth century. Sir
William Grove, a Scotsman, first demonstrated the concept of the hydrogen-oxygen fuel
cell in 1839 with a pair of platinum wires and a sulfuric acid electrolyte. The first real
application of the fuel cell did not come about until NASA used alkaline hydrogen-oxygen
fuel cells in various missions for producing both power and clean drinking water for
astronauts in manned space missions.
The first generation fuel cells used aqueous electrolytes; in the 1960s, Du Pont (E. I.
Du Pont de Nemours and Company) introduced Nafion
®
-H, an ionomer with a
polytetrafluoroethylene (PTFE) backbone and perfluoroether sulfonic acid side chains, as a
polymer electrolyte membrane (PEM) for chloralkali cell electrolyzers used in chlorine
production. The ability to fabricate ion-conducting membranes led to the invention of the
polymer electrolyte membrane fuel cell, or PEMFC, which eliminated liquid electrolyte
solutions altogether. The first PEMFCs used hydrogen and oxygen; some fuel cells relied on
1
!
2
the decomposition of organic fuels to produce they hydrogen; these were called indirect fuel
cells. The direct methanol fuel cell (DMFC) was co-invented by researchers at the Jet
Propulsion Laboratory and the University of Southern California in the early 1990s. The
DMFC relies on the direct oxidation of methanol provided as an aqueous solution at the
anode. Soon other direct liquid feed fuel cells were explored, including others based on
ethanol, ethylene glycol, and formic acid. A historical overview will be presented in the
following sections. More few more detailed accounts are available [1, 2].
While they possess many strengths, fuel cells possess certain disadvantages that have
kept them from widespread use. Hydrogen, while extremely energetic, is not energy dense
and is difficult to produce, transport, and store. Liquid-feed fuel cells are held back by costs
associated with the Nafion
®
membrane and noble metal electrocatalysts. Much research over
the last several decades has focused on addressing these few, but pressing, issues.

1.1.1.1 The Grove gas battery
Christian F. Schönbein was the first to report in 1838 that “the chemical
combination of hydrogen and oxygen in acidulated (or common) water is brought about by
the presence of platina” [3]. In the following issue of the same journal, William R. Grove
described “the combination of gases by platinum” [4], in what was in essence the first fuel
cell. By inverting tubes of hydrogen and oxygen over platinum wire electrodes submerged in
sulfuric acid, Grove was able to observe an immediate and significant production of current
!
3
(see Figure 1-1). He further observed by the rising liquid level in the tubes that the gases
were consumed during this process, and determined by trial and error of various
substitutions that all three components—hydrogen, oxygen, and platinum—were necessary
for current to flow. The reactions occurring at each of the electrodes were as follows:

H
2
2H
+
+ 2e
-
[E
0
= 0.0 V] (1.1)
½O
2
+ 2H
+
+ 2e
-
H
2
O [E
0
= 1.23 V] (1.2)

Figure 1-1. Schematic representation of Grove’s 1839 experiment on “the combination of gases by
platinum.”

The aqueous sulfuric acid solution provided ionic conductivity between the two
electrodes. Grove wrote that the cell was set-up “so that about half of the platina was
!
4
exposed to the gas, and half to the water.” He was, in fact, describing a concept that is
critical to the operation of a fuel cell, namely the triple-phase boundary. Due to the
interfacial (and interphasial) nature of the fuel cell reactions, it is essential to maintain
mutual contact between the fuel or oxidant, the electrode/catalyst, and the electrolyte. The
fuel—in this case hydrogen—is decomposed on the catalyst to give electrons, which must be
collected by the anode and passed to the load; and protons, which must be transported to the
cathode through the electrolyte. As a result, if Grove’s platinum electrodes exposed to the
gases were allowed to dry, the reaction would take place only at the interface where they
encounter the bulk electrolyte. Indeed Grove observed that the current decreased over time,
but could be recovered by temporarily lifting the tubes, thereby re-covering the platinum
surfaces with electrolyte, and extending the triple-phase boundary when the tubes were
replaced.
In 1842, Grove published a second account of his research [5]. He constructed
batteries consisting of up to fifty hydrogen/oxygen cells in series, as illustrated in Figure 1-2.
He had expressed the desire in his first publication to “effect decomposition of water by its
composition;” Grove was fascinated by the idea that the combination of hydrogen and
oxygen gas into water could provide the energy to electrolytically decompose water into
those same gases, and he accomplished it with a battery of twenty-six cells. It was observed
that the hydrogen was consumed (qualitatively) twice as quickly as the oxygen. In this
!
5
second paper, Grove again expressed ideas central to the understanding of fuel cells that are
still relevant today. The first was that platinum’s role in the cell was catalytic, and he raised
the question of whether other noble metals could fill this role as well. The second was the
relationship between the exposed platinum surface area and the magnitude of the current; in
order to increase the former (and thereby the latter), he electrodeposited platinum onto the
foil electrode surfaces. In addition to electrodeposition, many techniques are still being
developed to this day with the goal of increasing the surface area of platinum (and other)
catalysts so as to reduce the necessary catalyst loading.

Figure 1-2. Representation of the first working fuel cell constructed by Grove in 1842. Adapted from
[5] (arrows were intended to indicate “positive current” flow).

Grove would continue to investigate this “gas battery” in the 1840s. The other
!
6
significant development in the nineteenth century was made by Langer and Mond in 1889,
when they published their results [6] employing a cell resembling the modern membrane-
electrode assemblies of PEM fuel cells (see §1.1.1.3). Porous foils of platinum were affixed on
either side of an inert, non-conducting porous ceramic material impregnated with a liquid
electrolyte. To the backs of these platinum electrodes were affixed strips of lead, forming a
sort of primitive gas diffusion layer and current collector. Alder Wright and Thompson,
also in 1889, and inspired by the success of previous demonstrations, reported their earlier
results using a stack of hydrogen/oxygen “double aeration plate cells,” which bore
resemblance to modern bipolar plates used today [7]. They also expressed a desire to
develop in a practical way “the direct oxidation in this way of alcohol, petroleum, coal and
such like forms of comparatively cheap sources of energy.” While attempts were made at
the direct oxidation of other fuels shortly thereafter, success by practical standards would
not come until the second half of the twentieth century.

1.1.1.2 Alkaline fuel cells (AFC)
Despite Grove’s and others’ early developments, the fuel cell remained obscure
through the end of the nineteenth and the beginning of the twentieth century. This period
was characterized by a proliferation of interest in direct coal-powered fuel cells as a more
efficient alternative to the steam engine. In the 1930s, engineer Francis Thomas Bacon
became interested in Grove’s work and the idea of using the products of water electrolysis as
!
7
fuel in an electrochemical device [8]. He quickly abandoned the platinum/sulfuric acid
system in favor of a more economical nickel catalyst and a potassium hydroxide electrolyte,
which allowed the use of non-noble metals. He planned to compensate for the lower activity
by increasing the operating temperature and pressure. In 1930, Bacon built a cell that
provided 13 mA cm
-2
at 0.89 V operating at 3000 psi [9]. In 1946, Davtyan also reported a
KOH-based H
2
-O
2
fuel cell using a silver anode and a nickel cathode that produced 25 mA
cm
-2
at 0.80 V at ambient temperature and pressure.
Bacon worked on developing alkaline fuel cells throughout the 1940s and 1950s with
a fair amount of success but little commercial interest. Bacon’s self-imposed target was a
current density of 100 mA cm
-2
at an operating voltage of 0.8 V [9]. Porous nickel gas
diffusion electrodes were made from carbonyl nickel powder. To prevent the circulating
aqueous electrolyte from flooding the electrodes, the pore size was made small on the liquid
side and large on the gas feed side; the large pores allowed adequate gas diffusion with less
clogging by the liquid, while the small pores prevented excessive gas leakage into the
electrolyte and vice-versa. Originally both electrodes were made from sintered nickel
powder, but corrosion at the cathode resulted in rapid performance decay. The eventual
solution was to pre-oxidize the nickel for the oxygen electrode to nickel oxide, and dope it
with lithium to improve the conductivity. A six-cell, 150 W stack was shown in 1956, which
used a circulated 45% KOH electrolyte and gases pressurized to 400 psi. Operating at 200 °C,
!
8
it was able to output 230 mA cm
-2
at 0.8 V. In 1959, a similar forty-cell, 6 kW stack with
larger electrodes was demonstrated, the intervening years having been spent on scale-up to
practical levels, as well as designing the internal controls necessary for a commercial product
[9, 10]. Despite several working product demonstrations including a forklift, the technology
was licensed by only a handful of companies and research was halted.
One of these companies was Pratt & Whitney Aircraft Company (now a subsidiary
of United Technologies Corporation), which was contracted to design and build fuel cells
for the Apollo program in 1962 [9]. The hydrogen-oxygen fuel cell represented the most
efficient system able to provide high energy with low weight. Each mission would require
up to 650 kWh for electrical systems during spaceflight, and the 100 kg powerplants were
designed to provide between 563 and 1420 W each during normal operation [11] in a triply
redundant arrangement. The system also provided clean drinking water. Most of Bacon’s
design elements were preserved: aqueous KOH was employed as the electrolyte, without
circulation but at higher concentration; lower pressure and higher temperature were settled
on. Rigorous endurance testing revealed the design to be stable over 360 hours of operation
and two years of storage.
Several other companies worked on alkaline fuel cells in the 1960s. The work of
Kordesch at Union Carbide [2, 12, 13] led to cell with a 9 M KOH electrolyte and carbon
electrodes, which operated at ambient pressure, relatively low temperature (65 °C), and
!
9
produced 100 mA cm
-2
at 0.8 V (the current density was doubled when the cell was
pressurized slightly). Researchers at General Electric [14] reported cells using platinum
black electrodes with loadings between 17 and 45 mg cm
-2
using both alkaline and acidic
electrolytes and producing between 100 and 200 mA cm
-2
at ambient temperature and
pressure. In contrast to Bacon’s dual-pored design, both of these groups used
polytetrafluoroethylene (PTFE) films to control the wetting of the electrodes. Wynveen and
Kirkland [15] with Allis-Chalmers Manufacturing Company demonstrated cells using
platinum-palladium catalysts and an immobilized electrolyte impregnated into an asbestos
matrix, with an output of 108 mA cm
-2
at 65 °C at slightly elevated pressure.
NASA again employed alkaline fuel cells to provide electrical power in the Space
Shuttle Orbiter. The design of these cells was simplified compared to those in the Apollo
program: both the temperature and pressure were reduced, and the KOH was immobilized
in an asbestos matrix rather than circulated by pumps [16]. As a result, platinum black
catalysts supported on gold screens replaced the less active but more economical nickel-
based ones, with loadings of 10 mg cm
-2
at the anode and 20 mg cm
-2
at the cathode (in the
latter electrode it was alloyed with gold). NASA’s specifications for its contractors demanded
a 5000-hour lifetime, 2.5 kW average/5 kW peak power output, and a specific power of 40-
60 lb kW
-1
[17]. The cells that eventually powered the Orbiters had 7/12 kW average/peak
power capacities and a specific power around 20 lb kW
-1
, compared to the Apollo cells’
!
10
200 lb kW
-1
.
In the early 1970s, NASA awarded simultaneous contracts for Orbiter fuel cell
development to Pratt & Whitney and General Electric [17]; the latter had built polymer
electrolyte membrane fuel cells for the Gemini program in the 1960s. Although both met
the stated target requirements, it was Pratt & Whitney’s alkaline system that was eventually
used due to difficulties with water management in early PEM fuel cells (i.e., membrane
humidification vs. electrode flooding [16]). As PEM technology advanced, however, alkaline
KOH-based fuel cells began to fall out of favor.

1.1.1.3 Polymer electrolyte membrane fuel cells (PEMFC)
The polymer electrolyte membrane fuel cell (PEM fuel cell, also called the solid
polymer electrolyte or ion exchange membrane fuel cell) was first developed at General
Electric in 1960 [18]. A commercially available membrane called Amberplex C-1 (Rohm and
Hass Co.) was used, and consisted of crosslinked, sulfonated polystyrene blended with an
inert binder. The conductivity of early PEMs was rather poor compared to concentrated
aqueous electrolytes employed in alkaline cells [19]. Furthermore, while alkaline cells
struggled with the problem of electrode flooding, PEM fuel cells had to use humidified gases
in order to keep the membrane hydrated. The water content of the membrane had been
found to be directly related to its conductivity [19], and is necessary to provide proton
transport between the polymer’s hydrophilic sulfonic acid side-chains.
!
11
The best cells reported in 1960 used platinized nickel mesh electrodes and produced
around 2.4 mA cm
-2
at 0.75 V at ambient temperature, with open circuit voltages between
0.9 and 1.0 V. Despite their low power output compared to Bacon cells, PEM fuel cells were
seen as promising for several reasons. First, they dispensed with the system of pumps
necessary to circulate a liquid electrolyte. Second, the membrane-electrode assemblies
(MEAs) were less than 1 mm thick. Finally, the acidic nature of the proton exchange
membrane meant that the harmful precipitation of carbonate species from CO
2
present in
the gas feeds (air) was not a concern as it is in alkaline fuel cells.
In 1965, GE’s cells were incorporated into Gemini 5 [1]. The cells operated at low
temperature and pressure and were fed with humidified hydrogen and oxygen, and the
platinized titanium electrodes were bound with PTFE to prevent flooding. Each thirty-two-
cell stack was rated for 620/1000 W of average/peak power, or nominally 70 lb kW
-1
.
In the 1960s, Walther Grot of Du Pont [20] designed a new proton exchange
membrane based on a PTFE backbone with sulfonated fluoroether side-chains which would
later be named Nafion
®
. It possessed improved chemical and thermal stability characteristics
compared to the polystyrene-based membranes and also provided better conductivity. It was
Nafion
®
(or possibly a similar sulfonated Teflon derivative) that was used in GE’s next
generation of PEM fuel cells intended for the Space Shuttle Orbiter [21]. GE claimed an
operating lifetime of at least 2000 hours, including at high temperatures, and a specific
!
12
power of 16-25 lb kW
-1
depending on the mode of operation. Despite some disadvantages
(see §1.1.4) Nafion
®
remains the most widely used PEM today.

1.1.1.4 Direct methanol fuel cells (DMFC)
The first direct fuel cells based on fuels other than hydrogen were envisaged in the
late nineteenth and early twentieth centuries. A great deal of effort was put into
electrochemically utilizing the predominant fuel of the time: coal. These investigations
were largely unsuccessful, but the desire to accomplish electrooxidation of other fuels
remained [7]. Indirect fuel cells had been constructed in which various fuels were reformed
into hydrogen and fed into the anode. In the 1960s, work began in earnest to develop fuel
cells which would use hydrocarbon fuels such as methanol directly. Around the same time
that alkaline H
2
/O
2
fuel cells were being incorporated into the Gemini program, several
companies were building direct methanol fuel cells (DMFCs) using both acidic and alkaline
aqueous electrolytes [22-24]. These cells incorporated methanol directly into the circulating
electrolyte. The phenomenon of methanol crossover was first observed in these cells, in
which methanol would leak into the cathode and resulted in a mixed potential. The
formation of carbonates in alkaline electrolytes as a result of formed CO
2
was observed as
well. These first DMFCs were capable of producing current densities on the same order of
magnitude as their contemporary alkaline H
2
/O
2
fuel cells, but generally with much larger
polarization (i.e., lower cell voltage and, consequently, power).
!
13
The catalysts used were predominantly noble metals and alloys thereof in both acidic
and alkaline systems. Researchers at Allis-Chalmers in 1963 reported [22] a DMFC that
used a 6 M KOH, 6 M methanol electrolyte solution with platinum-palladium anode and Ag
or Co
3
O
4
cathode catalysts. A forty-cell stack using oxygen at the cathode and operating at
50 °C was able to produce 151 mA cm
-2
with each cell at 0.25 V (for a maximum power
output of 730 W at 10 V). They observed a steady reduction of the KOH concentration over
time due to reaction with CO
2
. In the same year, Shell [24] started building acidic DMFCs
using 6 N sulfuric acid and 1 M methanol (with a platinum-ruthenium anode catalyst) and
air (with a platinum cathode catalyst); a similar forty-cell stack operating at 60 °C and using
air gave 300 W at 12 V (i.e., an individual cell operating voltage of 0.3 V). While the
problem of carbonate formation is not present in acidic systems, they confirmed the
presence of incomplete methanol oxidation products such as formaldehyde and formic acid,
which poisoned the electrode over time. In a 1965 comparison study of catalysts based on
noble metals, Binder et al. found [25] that platinum-ruthenium was the most durable
catalyst for methanol oxidation in acidic solutions over extended periods.
Research in the next few decades, while not sparse, was comparatively slow moving.
Focus was primarily on catalyst design and preparation [26-31] in parallel with mechanistic
investigations [32, 33]. These latter studies by Watanabe et al. firmly established the
bifunctional model of the oxidation of methanol on platinum-ruthenium catalysts (see
!
14
§1.1.2). Despite these theoretical advances, interest in the DMFC did not take off until the
1990s when researchers at the Jet Propulsion Laboratory and the University of Southern
California introduced a DMFC incorporating a Nafion
®
membrane similar to the PEMFCs
from GE in the 1960s [34]. In place of the circulating sulfuric acid or potassium hydroxide
electrolytes with added methanol, membrane electrode assemblies were constructed and
aqueous methanol was fed to the anode in a manner analogous to the H
2
/O
2
PEMFC. In
cells incorporating the same platinum-ruthenium anode catalyst and fed with 1 M methanol
at 60 °C, the overpotential of the anode was found to be nearly 200 mV less in the cell using
a Nafion
®
-117 membrane versus that using a sulfuric acid electrolyte. The cells were fed
with oxygen at the cathode and aqueous methanol over a range of concentrations at the
anode; dilute solutions were found to suffer from mass transport limitations. Concentrated
(4 M) solutions were detrimentally impacted by crossover, and parasitic methanol oxidation
on the cathode was found to reduce the cell potential. Using moderate conditions (i.e. 2 M
methanol, < 90 °C) the JPL-USC team were able to achieve current densities of 300 mA cm
-2

at 0.5 V.
Over the intervening decades, research into PEM-based DMFCs has exploded [35-
46]. Nafion
®
and other related proton exchange membrane systems with low methanol
crossover have been widely investigated as electrolytes for DMFCs [47-51]. Operating
systems have been demonstrated and a few commercial systems are now available [52-54].
!
15
Recently, there has been renewed interest in alkaline DMFCs using anion exchange
membranes [55-61] because of the kinetic advantages and the use of low-cost non-noble
metal catalysts in an alkaline environment.

1.1.2 Principles of operation
A fuel cell consists of two electrodes separated by an electrolyte, which is ionically
but not electronically conductive. A fuel and an oxidant are supplied to the anode and
cathode, respectively, either through a pump system (active) or in a static reservoir (passive).
Most modern fuel cells are polymer electrolyte membrane fuel cells based on solid polymer
electrolytes such as Nafion
®
. These cells are fabricated in the form of membrane electrode
assemblies, or MEAs, less than 1 mm thick and consisting of a membrane sandwiched by the
two electrodes. The electrodes incorporate a porous gas diffusion layer (GDL) and a catalyst
layer applied to the inner side in contact with the membrane. As the fuel or oxidant is
supplied to the membrane, it diffuses through the GDL toward the active electrode surface
in contact with the electrolyte. The GDL is commonly made of a conductive carbon paper
and also serves as the current collector.
In the case of Nafion
®
and similar PEMs, the charge carrier is H
+
. Protons are
supplied from the oxidation of hydrogen, methanol, or another fuel at the anode and
consumed in the oxygen reduction reaction at the cathode. Electrons travel from the anode
to the cathode through the load. The electrode reactions for a DMFC are:
!
16

CH
3
OH + H
2
O CO
2
+ 6e
-
+ 6H
+
[E
0
= 0.02 V] (1.3)
3
/
2
O
2
+ 6H
+
+ 6e
-
3H
2
O [E
0
= 1.23 V] (1.4)

In a direct alkaline methanol fuel cell, an anion exchange membrane is used and the
charge carrier is OH
-
.

CH
3
OH + 6OH
-
CO
2
+ 5H
2
O + 6e
-
[E
0
= -0.81 V] (1.5)
3
/
2
O
2
+ 3H
2
O +6e
-
6OH
-
[E
0
= 0.401 V] (1.6)

In both cases, the net reaction is the same:

CH
3
OH +
3
/
2
O
2
CO
2
+ 2H
2
O [E
0
= 1.21 V] (1.7)

and is equivalent to the chemical combustion of methanol.
The thermodynamic voltage of a cell is determined by the reversible potentials of its
electrode reactions. Thermodynamically the cell would have the same voltage at any
operating current; in reality, a number of factors introduce overpotentials that reduce this
voltage. These include:
!
17
• fuel crossover: if fuel leaks from the anode to the cathode, its oxidation may compete
with the oxidant reduction reaction and cause a mixed potential at that electrode.
This effect can be seen at open circuit and during operation.
• electrode activation: both the anode and the cathode require catalysts, and less
efficacious catalysts result in additional electrode polarization away from the
thermodynamic potentials (i.e. an overpotential). This effect can be seen at open
circuit or during operation.
• ohmic resistance: according to ohm's law, the voltage loss for a cell with constant
resistance increases with current magnitude; consequently, the voltage of a cell will
steadily decrease as current increases during operation.
• mass transport: because the current density is directly related to the rate of reaction,
the maximum current is determined by the maximum rate at which fuel from the
supply stream can diffuse to the triple-phase boundary. This limitation is seen at
very high current densities.
Because the current is a function of reaction rate, it is also a function of electrode
surface area available for reaction. It is therefore convenient to talk about current density
(i.e. A cm
-2
) in terms of the geometric electrode area in order to compare the performance
of fuel cells of different sizes. Fuel cells can be scaled up, in theory, to any geometry
acceptable for the intended application.
!
18
1.1.3 Catalyst development
The first major area of investigation for improving the commercial prospects of the
DMFC is the discovery of new catalysts. Currently available catalysts suffer from several
issues:
• activity: the overpotential for the methanol oxidation reaction (MOR) is around
250 mV. More efficient catalysts at both electrodes would alleviate the activation
losses described in section 1.1.2.
• durability: metals less stable than platinum may be oxidized and lost during
operation. This is a particular problem with ruthenium in state-of-the-art Pt-Ru
MOR catalysts.
• cost: platinum and other noble metals add significantly to the cost of materials, and
worldwide supply may not be sufficient to meet demand on a global scale.
These problems are being addressed in various ways, including using high surface
area, conductive support materials to lower the overall precious metal loading; alloying
precious metals with other, cheaper metals; and devising noble-metal-free catalysts.

1.1.3.1 The anode reaction
The best catalysts available for methanol oxidation in acidic media are based on
platinum, and the most active of these is a Pt-Ru alloy in a 1:1 atomic ratio [62]. The
reaction mechanism on Pt-Ru materials (and other bimetallic, platinum based catalysts) is
!
19
considered to be bifunctional in nature and is illustrated as follows [62, 63]:

Pt + CH
3
OH Pt–CH
3
OH
ads
(1.8)
Pt–CH
3
OH
ads
Pt–CO
ads
+ 4H
+
+ 4e
-
(1.9)
Ru + H
2
O Ru–H
2
O
ads
(1.10)
Ru–H
2
O
ads
Ru–OH
ads
+ H
+
+ e
-
(1.11)
Pt–CO
ads
+ Ru–H
2
O
ads
Pt + Ru + CO
2
+ 2H
+
+ 2e
-
(1.12)
Pt–CO
ads
+ Ru–OH
ads
Pt + Ru + CO
2
+ H
+
+ e
-
(1.13)

The first faradaic step consists of C-H bond activation, the second of water
activation. Platinum can easily oxidize methanol to carbon monoxide in a four-electron
process; further oxidation is slow and adsorbed CO rapidly poisons an all-platinum catalyst.
Ruthenium by itself is not active for methanol oxidation, but can provide adsorbed oxygen
species more readily than platinum to effect the oxidation of CO, thereby freeing an active
Pt site. For this reason it is widely considered to be necessary to create well-alloyed
materials without disparate platinum or ruthenium regions, and with approximately one Ru
atom for each Pt atom.
A second mode of action of Ru is dubbed the ligand or electronic effect [64-66]: the
donation/back-donation interaction between CO and Pt is weakened when Ru modifies Pt’s
!
20
available electronic states by shifting its Fermi level. Evidence of this effect was seen in
cyclic voltammetry experiments in which the onset potential of methanol oxidation was
reduced. Although the thermodynamic potential of the MOR is only 20 mV positive to that
that of hydrogen oxidation, kinetically the reaction is slower by several orders of magnitude
[67]. Moreover, the catalyst loadings used in DMFCs are an order of magnitude larger than
those used in hydrogen PEMFCs, increasing the cost of the MEA. Finally, the durability of
state-of-the-art MEAs is insufficient, primarily due to anode ruthenium dissolution and
crossover [68], cathode Pt oxidation [69], and carbon degradation [70].
A DMFC run in alkaline medium has the potential to address many of these
problems. The kinetics of the MOR in aqueous hydroxide are known to be more favorable
than in the acidic media [71]. Non-noble metals are also more stable in alkaline
environments, which broadens the range of acceptable materials and alleviates catalyst
durability concerns. Given the success of the alkaline hydrogen-oxygen fuel cell, the
oxidation of methanol in alkaline environment was studied as early as the 1960s [72-74].
The mechanism of methanol oxidation in alkaline media on platinum and platinum-
ruthenium is more complex than that in acid media because it apparently involves more
intermediates. Several mechanisms have been proposed (i.e. [75, 76]) but some features are
generally agreed upon. As in acid media, a bifunctional process is assumed, but due to the
weaker binding of intermediates in alkali [77], this is possible on pure Pt as well as Pt-Ru.
!
21
1.1.3.2 The cathode reaction
Regardless of the fuel, nearly every fuel cell technology under investigation reduces
oxygen, either in pure form or as air, at the cathode. In an acidic system, oxygen reacts with
protons from the electrolyte to form water, while in an alkaline system, it reacts with water
to form hydroxide anions which are transported to the anode. The mechanism of the
oxygen reduction reaction (ORR) on platinum has been studied in both acidic and basic
media [78]. The acidic process is thought to proceed through a hydrogen peroxide
intermediate:

Pt + O
2
+ H
+
+ e
-
Pt–HO
2(ads)
(1.14)
Pt–HO
2(ads)
+ H
+
+ e
-
Pt–H
2
O
2(ads)
(1.15)
Pt–H
2
O
2(ads)
+Pt 2Pt–OH
ads
(1.16)
2Pt–OH + 2H
+
+ 2e
-
2H
2
O + 2Pt (1.17)
Overall: O
2
+ 4H
+
+ 4e
-
2H
2
O (1.18)

Both the initial electron transfer and the O-O bond cleavage have been suggested as
the rate-determining step. In a basic medium, the mechanism proceeds through oxide
species:
2Pt + O
2
2Pt–O
ads
(1.19)
2Pt + 2Pt–O + 2H
2
O 4Pt–OH
ads
(1.20)
!
22
4Pt–OH + 4e
-
4Pt + 4OH
-
(1.21)
Overall: O
2
+ 2H
2
O + 4e
-
4OH
-
(1.22)

Platinum is the most efficient catalyst for oxygen reduction as well as methanol
oxidation; as a result, methanol crossover in Nafion
®
-based MEAs results not only in the
loss of fuel, but also in competitive methanol oxidation at the oxygen electrode. The result is
a mixed cathode potential as well as the formation of poisoning CO. As such, in addition to
cost concerns, novel ORR catalysts should be inactive to methanol oxidation. One strategy
has been to alloy platinum with other metals [79], including Cr [80, 81], Co [81-83], Ni [83,
84], and Fe [85]. Other approaches have involved the complete replacement of precious
metals with transition-metal containing compounds such as chalcogenides [86, 87] and
macrocycles [88-90] that are not active for methanol oxidation.

1.1.4 Scope of this work
An alkaline electrolyte can be advantageous to a direct alcohol fuel cell for several
reasons. It is accepted that the kinetics of electro-oxidation of methanol and other alcohols
are more rapid in alkaline media [77] due to the weaker bonding of chemisorbed
intermediates, such as CO. The oxygen reduction reaction is also more facile in the alkaline
media [91]. The less corrosive environment invites the possibility of using non-noble metal
catalysts at both the anode and the cathode [92]. Furthermore, one of the most problematic
!
23
issues facing Nafion
®
-based DMFCs is methanol crossover through the membrane from the
anode to the cathode compartment, resulting in a mixed potential at the cathode, flooding of
the cathode, and parasitic consumption of fuel. This problem has been partially solved by
using PVDF-PSSA membranes [47, 48]. In alkaline fuel cells, this crossover is likely to be
hindered by the electro-osmotic drag of the hydroxide anions, which travel from the
cathode to the anode. Finally, the cost of Nafion
®
and similar polymers adds significantly to
the cost of materials. The use of non-noble-metal catalysts, which are not compatible with
proton exchange membrane fuel cells, reduces the cost even further.
However, alkaline fuel cells are not without disadvantages. Most alcohol AFCs
include an electrolyte such as KOH in the aqueous fuel to improve the hydroxide ion
conductivity of the membrane; the conductivity of alkaline anion exchange membranes is
generally found to be lower than that of proton exchange membranes [77]. Carbon dioxide
formed at the anode in the presence of a basic electrolyte tends to form insoluble carbonate
and bicarbonate species, simultaneously reducing the electrolyte conductivity and
potentially clogging and damaging the porous electrode structures [93].
In the present study, the alkaline direct methanol fuel cell based on an anion-based
polymer electrolyte membrane (in which quarternary ammonium side chains conduct
hydroxide anions) has been shown to provide similar performance to that of Nafion
®
-based
DMFCs. Specifically, we have investigated the effects of fuel composition, oxidant flow, and
!
24
electrode materials on the performance at a full cell level.

1.2 Experimental methods
Fuel solutions were prepared from HPLC-grade methanol (Sigma-Aldrich) and
reagent grade potassium hydroxide (Mallinckrodt) with deionized water (18 MΩ cm,
Millipore Direct-Q 3).
Membrane electrode assemblies (MEAs) for alkaline DMFCs were prepared using
the Tokuyama A-006 membrane (~50 microns thick, Tokuyama Corp.) and Toray carbon
paper (TGPH-60) with and without 10 w/w% Teflon wet-proofing. Catalyst inks were
prepared from platinum black or platinum-ruthenium (1:1 atom ratio) black (Alfa-Aesar;
Hi-Spec 1000 and Hi-Spec 6000, respectively), an ionomer solution (AS-4, quaternary
ammonium type, Tokuyama Corp.), and deionized water in the mass ratio of 1:1:3
(catalyst:ionomer:water). The inks were painted onto 25 cm
2
carbon paper electrodes to
achieve a loading of 8 mg cm
-2
and the MEAs were hot-pressed at 110 °C. For purposes of
comparison, similar MEAs were prepared with Nafion
®
117 (Ion Power, Inc.) using a 5%
Nafion
®
ionomer solution (Sigma-Aldrich) and hot-pressed at 140 °C. A diagram of the
MEA is shown in Figure 1-3 (a).
!
25

Figure 1-3. Cross-sectional diagrams of (a) a membrane-electrode assembly (not to scale) and (b) an
assembled fuel cell.

Characterization of the MEAs was carried out in a standard cell housing with pin-
type flow fields (Electrochem, Inc.) using a Fuel Cell Test System 890B (Scribner Associates,
Inc.). Fuel solutions were circulated using a Micropump® pump drive, and a Digi-Sense®
temperature controller (Cole-Parmer, Inc.) was used to control the temperature. Pure O
2

was supplied to the cathode at various flow rates using an Accucal® controller (Gilmont,
Inc.). A diagram of the cell housing is shown in Figure 1-3 (b).
Following hydration for a minimum of 6 hours at 60 °C, and MEA conditioning at
90 °C, I-V curve measurements were obtained with 0.5 A steps at 90 s intervals, at 30, 60 or
90 °C with O
2
supplied at 200, 700, or 1270 mL min
-1
.

TO FUEL CELL TESTING STATION
MeOH inlet
MeOH
outlet
(rear)
O2 inlet
O2 outlet
(rear)
a) b)
Membrane
Catalyst Ink
Carbon Paper
GASKETS
MEA
!
26
1.3 Results and discussion
1.3.1 Addition of KOH to the fuel
Initial experiments with aqueous methanol fuel (1-2 M) and no added KOH
electrolyte resulted in no measurable performance at all temperatures. This ultra-low
performance is attributed to the passive polymer electrolyte interface between the electrode
and the alkaline ion exchange membrane, which is not sufficiently conductive to allow facile
hydroxide transfer. Therefore, all further measurements were conducted with 1 M MeOH +
1 M KOH until the fuel composition was optimized (section 1.3.4–1.3.5).

1.3.2 Effects of temperature
MEAs were prepared at first using TGPH-60 that was wet-proofed with 10 w/w%
PTFE. These cells were supplied with 1 M MeOH + 1 M KOH at the anode, and the cathode
at ambient pressure with an oxygen flow rate of 1270 mL·min
-1
. Figure 1-4 shows the
performance of such a cell as a function of temperature. As expected, the current and power
density maxima increase in value with temperature; at 90 °C, the peak power density is
approximately double (63 mW cm
-2
) the value observed at room temperature (27 mW cm
-2
).
The cell behavior at 60 °C is very similar to that at 90 °C until the current density reaches
140 mA cm
-2
(where both cells were at 0.44 V) and mass transport losses due to decreased
diffusivity at lower temperatures take over. This is attributed largely to the Teflon-coated
anode (section 1.3.6); note that at 90 °C, the mass transport limitations are not observed and
!
27
Figure 1-4. Effect of temperature: polarization and power density curves, 1 M methanol + 1 M KOH,
1270 mL min
-1
O
2
, wet-proofed electrodes.

ohmic control persists into high current densities.

1.3.3 Effects of oxidant flow rate
When the oxygen flow rate is decreased from 1270 mL min
-1
to 200 mL min
-1
, the
reduction in performance is typically 5-6% at all three temperatures on either a power
density or a current density basis. For example, at 90 °C, the peak power density decreases
from 63 mW cm
-2
to 59 mW cm
-2
. It is clear from these results that the mass transport of
oxygen is not a limiting factor in the range of current densities investigated.

0!
10!
20!
30!
40!
50!
60!
70!
0.00!
0.10!
0.20!
0.30!
0.40!
0.50!
0.60!
0.70!
0.80!
0.90!
0! 100! 200! 300! 400!
power density (mW cm
-2
)!
voltage (V)!
current density (mA cm
-2
)!
90 ºC!
60 ºC!
30 ºC!
!
28
1.3.4 Effects of methanol concentration
The baseline fuel used in the previously described experiments was 1 M MeOH +
1 M KOH. The fuel composition was optimized at 90 °C and 1270 mL min
-1
O
2
flow rate. At
first, the KOH concentration was held constant and the methanol concentration was
increased. With 1 M KOH, the performance of the MEA decreased with increasing
methanol concentration (Figure 1-5). Although methanol crossover is expected to be less
problematic with alkaline fuel cells, it is likely that the relatively thin Tokuyama membrane
(50 microns) makes these MEAs more susceptible to crossover compared to a thicker 117
Nafion
®
membrane (175 microns). Thus, increasing the concentration of MeOH beyond
1 M is not beneficial. For example, at 50 mA cm
-2
the observed cell voltages for 1 M, 2 M,
and 3 M methanol are 0.59 V, 0.55 V, and 0.39 V, respectively, demonstrating the impact of
crossover. The lack of mass transport limitation at 90 °C using 1 M MeOH (§1.3.2) is further
evidence that 1 M methanol is the optimal fuel concentration.

1.3.5 Effects of potassium hydroxide concentration
The power density was found to increase upon raising the concentration from 1M
KOH to 2 M KOH (Figure 1-6). Also, the performance of the cell with increasing methanol
concentration at 2 M KOH was similar to that observed at 1M KOH. This supports the
hypothesis (1.3.1) that KOH increases the conductivity of the electrode/electrolyte interface.
However, no improvement was observed by increasing the KOH concentration to 3 M; in
!
29
Figure 1-5. Effect of methanol concentration: polarization and power density curves, 1 M KOH,
90 °C, 1270 mL min
-1
O
2
, wet-proofed electrodes.

fact, performance was inferior to that at 2 M KOH and very similar to that at 1 M KOH (not
shown). Based on these data, 1 M MeOH + 2 M KOH was chosen as the ideal fuel
composition to be studied for this system.
There are many factors to consider while examining the effect of electrolyte
concentration on electrode performance, especially the effect of pH on the membrane
potential and carbonate formation and the dependencies could be too complex to be
separable in the present experiment. For example, as CO
2
is formed in the anode gas
diffusion layer (GDL), the rate of the bicarbonate/carbonate formation reaction depends on
the concentration of OH
-
at the interface. The concentration of OH
-
is in turn dependent on
0!
10!
20!
30!
40!
50!
60!
70!
0.00!
0.10!
0.20!
0.30!
0.40!
0.50!
0.60!
0.70!
0.80!
0.90!
0! 100! 200! 300! 400!
power density (mW cm
-2
)!
voltage (V)!
current density (mA cm
-2
)!
1 M MeOH!
2 M MeOH!
3 M MeOH!
!
30
the diffusion rate of OH
-
through the GDL and on the bulk concentration of KOH.
Therefore we suggest that the increase in performance from 1 M to 2 M is due to the
beneficial effect on conductivity, but that when the KOH concentration is further increased
to 3 M, the detrimental effects of increased bicarbonate/carbonate formation begin to
outweigh these benefits.

Figure 1-6. Effect of increasing KOH concentration: polarization and power density curves, 1 M
methanol, 90 °C, 1270 mL min
-1
O
2
, wet-proofed electrodes.

1.3.6 Effects of electrode material
All previous experiments were conducted with standard Toray TGPH-60 with 10%
wet-proofing. Using all the previously optimized conditions, the impact of the Teflon wet-
0!
20!
40!
60!
80!
100!
120!
0.00!
0.10!
0.20!
0.30!
0.40!
0.50!
0.60!
0.70!
0.80!
0.90!
0! 100! 200! 300! 400! 500!
power density (mW cm
-2
)!
voltage (V)!
current density (mA cm
-2
)!
1 M KOH!
2 M KOH!
3 M KOH!
!
31
proofing on the carbon paper electrodes was examined. Figure 1-7 shows the I-V and power
density performance of various MEAs with both wet-proofed and non-wet-proofed Toray
electrodes. Relative to the baseline, replacing the cathode only with non-wet-proofed
carbon paper decreases the maximum current density by nearly 25%, and the maximum
power density by 20%. On the other hand, replacing only the anode with non-wet-proofed
carbon paper more than doubles the maximum current density (from 400 mA cm
-2
to 840
mA cm
-2
) and increases the maximum power density from 101 mW cm
-2
to 168 mW cm
-2
.
This is in line with expectations that the Teflon coating would impede the diffusion of the
aqueous electrolyte into the anode but would aid in water rejection to prevent flooding at
the cathode.
To gauge the performance of the optimized cell at low temperatures, it was also
tested at 30 °C. As expected, the performance is lower than at 60 or 90 °C, but a peak power
density of 42 mW cm
-2
was achieved. Figure 1-8 shows the performance of this cell versus a
teflonized anode, also at 30 °C. The polarization curves confirm that the Teflon coating
introduces a mass transport limitation at low temperature, while the anode without Teflon
appears to be only under ohmic control even at high current densities.
Wet-proofed electrodes are considered to be superior for Nafion
®
cells in order to
hinder crossover due to the electroosmotic drag of water [94]. Unlike Nafion
®
, which has a
very conductive membrane-ionomer interface and can easily transport protons [95], the
!
32
Figure 1-7. Effect of electrode wet-proofing: polarization and power density curves, 2 M KOH + 1
M methanol, 90 °C, 1270 mL min
-1
O
2
.

Tokuyama membrane requires rapid and efficient distribution of electrolyte into the pores
at the electrode/membrane interface in order to ensure good ionic conductivity of the
interface. Finally, the electroosmotic drag of the hydroxide anions toward the anode resists
water (and methanol) flow into the cathode, and consequently the hydrophobic layer is not
necessary.

0!
20!
40!
60!
80!
100!
120!
140!
160!
180!
200!
220!
0.00!
0.10!
0.20!
0.30!
0.40!
0.50!
0.60!
0.70!
0.80!
0.90!
1.00!
0! 100! 200! 300! 400! 500! 600! 700! 800! 900!
power density (mW cm
-2
)!
voltage (V)!
current density (mA cm
-2
)!
both electrodes teflonized!
anode telfonized, cathode non-teflonized!
anode non-teflonized, cathode teflonized!
!
33
Figure 1-8. Mass transport limitation at low temperature with wet-proofed anode: polarization and
power density curves, 2 M KOH + 1 M methanol, 30 °C, 1270 mL min
-1
O
2
.

1.3.7 Comparison of performance of AAEM and Nafion
®
cells
For comparison purposes, a Nafion
®
117 MEA was tested with the electrolyte
previously optimized with the AAEM MEAs (90 °C, with 1 M MeOH and no KOH, with
wet-proofed electrodes). The maximum power density for the Nafion
®
cell was
173 mW·cm
-2
, an improvement of less than 4% over the AAEM cell. Furthermore, the
maximum power density of the Nafion
®
MEA was observed at a lower cell voltage,
suggesting a higher voltage efficiency for the AAEM cell at the peak power density. At the
peak power density voltage of the Tokuyama cell, 0.35 V, the power density of the Nafion
®

0!
10!
20!
30!
40!
50!
0.00!
0.10!
0.20!
0.30!
0.40!
0.50!
0.60!
0.70!
0.80!
0! 100! 200! 300!
power density (mW cm
-2
)!
voltage (V)!
current density (mA cm
-2
)!
non-teflonized anode!
teflonized anode!
!
34
cell is actually 4% lower. Figure 1-9 is a plot of the cell voltage corrected for internal
resistance losses. Tafel slopes calculated for the AAEM and Nafion
®
MEAs were very similar,
consistent with the absence of any difference in the catalyst compositions in both the cells.
Figure 1-9. Comparison of Tokuyama MEA with Nafion MEA: polarization and power density
curves, 2 M KOH + 1 M methanol, 90 °C, 1270 mL min
-1
O
2
(dotted lines represent iR-corrected
polarization curves).

The ohmic resistance measurements made at the start of the tests were used in
calculating the resistance-corrected voltage values. The resistance of the AAEM is quite
similar to the Nafion
®
cell resistance, both less than 10 mΩ (for an electrode active area of
25 cm
2
). Nafion
®
exhibits better performance at high current densities; this may be due to
the higher solubility of O
2
. Further, the cathode in the Tokuyama cell is relatively
0!
20!
40!
60!
80!
100!
120!
140!
160!
180!
200!
0.00!
0.10!
0.20!
0.30!
0.40!
0.50!
0.60!
0.70!
0.80!
0.90!
1.00!
0! 100! 200! 300! 400! 500! 600! 700! 800! 900! 1000!
power density (mW cm
-2
)!
voltage (V)!
current density (mA cm
-2
)!
Nafion MEA!
AAEM MEA!
!
35
unoptimized with respect to the formation of the electrode/electrolyte interface, and thus
O
2
mass transport issues may arise at current densities greater than 500-600 mA·cm
-2
.
Furthermore, the cathode does not have the benefit of increased conductivity because KOH
is added only to the anode. Figure 1-10 shows that the performance decrease observed at
high current densities when the oxygen flow rate is decreased is more pronounced than at
low current densities (§1.3.3). This confirms the limitations to mass transport of oxygen in

Figure 1-10. Polarization curves showing the effect of reduced oxygen flow rate at high current
densities; 2 M KOH + 1 M methanol, 90 °C; high-flow represents 1270 mL min
-1
O
2
, low flow
represents 200 mL min
-1
O
2
.

0.00!
0.10!
0.20!
0.30!
0.40!
0.50!
0.60!
0.70!
0.80!
0.90!
1.00!
0! 100! 200! 300! 400! 500! 600! 700! 800! 900!
voltage (V)!
current density (mA cm
-2
)!
Low-flow oxygen!
High-flow oxygen!
!
36
the AAEM cell. Therefore, it is clear that further optimization in MEA fabrication and
ionomer formulation will result in further improvements to the AAEM cell.

1.3.8 Cell durability
Figure 1-11 shows the performance of a cell over 48-hour intervals. Initially, the cell
was tested repeatedly at 60 °C with low concentration fuel (1 M MeOH + 1 M KOH) over
two days. After these tests, any decrease in performance was completely recovered by
replacing the fuel solution. Subsequently, the temperature was increased to 90 °C and more
concentrated fuels were tested. Figure 1-11 shows the initial performance on the first day of
more extended testing at the optimized fuel concentration (1 M MeOH + 2 M KOH). The
second curve represents the same cell under similar conditions after testing fuels reaching
3 M in MeOH or KOH: the power density is approximately 40% of its previous output.
Finally, the last curve shows the performance after continued testing with concentrated
fuels, but both electrodes were flushed at 60 °C with deionized water overnight before
measuring again at the optimized concentration. The performance recovered significantly,
but not completely (~70% of the initial power density).
Degradation reactions of AAEMs in strongly basic environments (particularly
nucleophilic addition-elimination and Hofmann elimination [96]) are well known. The
increase in temperature to 90 °C may have contributed to the unrecoverable drop in
performance by favoring these or similar processes. However, the reversible reduction in
!
37
performance is more likely due to the formation of bicarbonate species through the reaction
of formed CO
2
with hydroxide anions in the electrolyte. In addition to reducing the
conductivity by decreasing the hydroxide concentration in the fuel, the bicarbonate and
carbonate anions will decrease the conductivity of the membrane by exchanging with the
hydroxide anions in the membrane. Thus, flushing with water over an extended period
restores the membrane sites at least partially.

Figure 1-11. Degradation behavior of a Tokuyama MEA with wet-proofed electrodes: initial
performance, followed by observed degradation after high concentration fuel experiments, and
finally partial recovery after 12h/60 °C water flush; experiments were separated by 48 hour intervals;
polarization and power density curves, 2 M KOH + 1 M methanol, 90 °C, 1270 mL min
-1
O
2
.

0!
20!
40!
60!
80!
100!
120!
140!
160!
180!
200!
0.00!
0.10!
0.20!
0.30!
0.40!
0.50!
0.60!
0.70!
0.80!
0.90!
0! 100! 200! 300! 400!
power density (mW cm
-2
)!
voltage (V)!
current density (mA cm
-2
)!
initial performance!
after concentrated fuel testing!
after water flush!
!
38
1.4 Conclusions
The peak performance obtained from MEAs with the alkaline anion exchange
membrane (Tokuyama A-006) and the Nafion
®
DMFC are comparable when tested under
similar conditions. However, the factors that limit the performance of the two types of cells
are different. In particular, it appears that the oxygen reduction reaction at the cathode
becomes a limiting factor in the Tokuyama cells at high current densities. Also, at high
current densities, the ohmic resistance results in a slightly lower performance with the
AAEM. Using an optimized fuel mixture and choice of electrode materials, we obtained
relatively high power densities and current densities compared to known reports of alkaline
DMFCs with only 1 M methanol as fuel at both high and low temperatures [97]. The most
obvious disadvantage of the cell is that performance degradation is accelerated compared to
that of the Nafion
®
DMFC. Both reversible (due to loss of conductivity through anion
exchange) and irreversible (due to the chemical degradation of the membrane in an alkaline
environment) performance losses were observed. Further improvements in membrane
stability and fabrication techniques will allow higher performance levels to be attained with
the AAEM cells.

!
39
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46

Catalysts for direct formic acid fuel cells

2.1 Introduction to direct formic acid fuel cells
The history of the development of the fuel cell was summarized in section 1.1. The
direct formic acid fuel cell (DFAFC) has experienced increased research in recent years in
parallel with the direct methanol fuel cell. The electrode reactions in a DFAFC are as
follows:

HCOOH CO
2
+ 2H
+
+ 2e
-
[E
0
= -0.25 V] (2.1)
½O
2
+ 2H
+
+ 2e
-
H
2
O [E
0
= 1.23 V] (2.2)

As evidenced by the reaction potentials, the thermodynamic efficiency of the
DFAFC is greater than that of the DMFC. Formic acid is the liquid fuel with the structure
closest to that of carbon dioxide, separated by a two-electron reaction, which in theory is
conducive to complete oxidation. One of the primary concerns with the DMFC is the
crossover of methanol through Nafion
®
membranes. Formic acid undergoes crossover at a
much lower rate than methanol [1-3], and does not require stoichiometric water in the
2
47
anode reaction; as a result, fuel solutions with much higher concentrations can be used in
DFAFCs versus DMFCs. Furthermore, this results in better fuel efficiency by eliminating
parasitic fuel oxidation at the cathode as well as the accompanying loss of voltage. These
factors partially compensate for formic acid's lower energy density, which is less than half
that of methanol (2086 Wh l
-1
vs. 4690 Wh l
-1
). Finally, formic acid is approved as a food
additive by the United States Food and Drug Administration, and is considered
environmentally acceptable.
Formic acid fuel cells share some of the disadvantages of DMFCs when compared to
hydrogen cells. First, the anode kinetics are sluggish compared with the HOR, resulting in
higher catalyst loadings. Second, given that both of these fuels contain carbon, there is
formation of the -CO intermediate which poisons certain catalysts, particularly those
containing platinum. The mechanism of the electrooxidation of formic acid (FAO) is usually
written as a dual-pathway process:

HCOOH
ads
CO
2
+ 2H
+
+ 2e
-
(2.3)
HCOOH
ads
CO
ads
+ H
2
O --> CO
2
+ 2H
+
+ 2e
-
(2.4)

The dehydrogenation or "direct" pathway involves weakly bound intermediates and
produces CO
2
. The dehydration or "indirect" pathway includes a -CO intermediate, the
48
oxidation of which is the rate-determining step. Platinum-based catalysts favor the
dehydration pathway and as a result are quickly poisoned during FAO [4, 5]. Palladium
catalysts, on the other hand, are found to favor the dehydrogenation pathway, and display
faster kinetics as a result. However, they suffer from steady deactivation over the course of
operation. Much research has been devoted to the discovery of a catalyst system similar to
the platinum-ruthenium catalyst for DMFCs which resist poisoning or are able to effectively
oxidize intermediates. In parallel, there is a desire to reduce catalyst cost as in DMFCs,
through the use of lower loadings and cheaper materials.

2.1.1 Recent developments
Both platinum and palladium catalysts have been extensively studied for formic acid
oxidation since the exhaustive work of Capon and Parsons in the 1970s [6-9]. It had been
suggested that the reaction on platinum was poisoned by a non-reactive intermediate [10]
although its exact nature was not known and was the subject of much speculation [11, 12].
In light of other fuel cell systems, much of the focus in early studies of formic acid
oxidation was on platinum [13-15]. During the 1980s, a great deal of effort was put into
formic acid oxidation [16-20] but the construction of a formic acid fuel cell was not
reported until the mid-1990s [21]. It had become apparent that the poisoning of platinum
during FAO was a serious issue and this study reported that PtRu, known to reduce CO
poisoning in DMFCs, was also a superior catalyst for formic acid oxidation. However, the
49
interaction between ruthenium and formic acid around the operating potential of the anode
is too strong [22] and interferes with the efficient oxidation of the CO
ads
formed on
platinum. Subsequent studies have reported that the addition of Ru to Pt for formic acid
oxidation is not beneficial [23, 24].
A great deal of work was put into alloying or otherwise modifying platinum
electrodes with other metals in order to combat the poisoning problem [13, 25-34]. Several
groups identified PtSn as a potential catalyst [23, 24] which functions in a manner
analogous to PtRu in DMFCs, without the problem of Ru deactivation:

HCOOH + Pt Pt-HCOO + H
+
+ e
-
(2.5)
H
2
O + Sn Sn-OH + H
+
+ e
-
(2.6)
Pt-HCOO + Sn-OH CO
2
+ H
2
O + Pt + Sn (2.7)
Pt + HCOOH Pt-CO + H
2
O (2.8)
Pt-CO + Sn-OH CO
2
+ H
+
+ e
-
+ Pt + Sn (2.9)

Contrary to the behavior on pure Pt, the dehydrogenation pathway is favored; for
CO formed in the dehydration pathway, oxygen species formed on Sn help to oxidize and
remove it [35]. In addition to the bifunctional mechanism, a third-body effect for Pt
adatoms has also been postulated in the case of catalytically inactive adsorbed species that
50
block poisoning [30, 36]. A third possible mode of action is an electronic effect [37, 38]. It
has been shown that the strength of interaction with adsorbates is a function of the d-band
center of the metal, and that this is different for strained surfaces, i.e. those that have been
modified by alloying or other means.
Compared to platinum, the oxidation of formic acid on palladium is relatively poorly
understood, in spite of its generally being considered the most active FAO catalyst. Because
palladium oxidizes formic acid through the direct pathway exclusively, the reactive
intermediates involved have not been identified. Nevertheless, the catalyst activity and
therefore cell performance decrease over time [39]. It was observed that a positive potential
pulse was able to oxidize the surface species and reactivate the electrode [40-42].
A number of palladium alloy materials have been investigated as catalysts in recent
years [33, 43-51] as well as carbon-supported materials. A research group from the
University of Illinois in particular has recently demonstrated performance comparable to
DMFCs at room temperature operation using alloyed and supported palladium-based
catalysts, including PdAu, PdSb, and PdBi [52-54].

2.1.2 Mechanistic studies
The "dual path" mechanism has been used to describe the mechanism of formic acid
oxidation, but the precise intermediates involved have not been determined and are the
subject of many sometimes-contradictory studies. On platinum, where both CO poisoning
51
and direct oxidation are known to occur, several theories have been advanced as to the
intermediate or transition state involved in the "direct" process. It has long been speculated
that this intermediate is COOH [6-9].
Samjeské et al. [55, 56] demonstrated more recently via attenuated total reflectance
Fourier transform infrared spectroscopy (ATR-FTIR) the absence of COOH but a
significant presence of formate or HCOO. They determined that formate covers the
platinum electrode surface very quickly during FAO, and that it is consumed and
replenished at a rate consistent with the formic acid oxidation current, and that therefore
formate oxidation is the rate-determining step for FAO on platinum surfaces. They also
reported the concurrent steady build-up of CO, which in addition to reducing available
platinum surface area, apparently decreases the rate of adsorbed formate oxidation. They
further reported that this rate-determining step was also dependent on the degree of
formate coverage, i.e. that formate oxidation is retarded by the presence of additional
adjacent occupied sites.
Following this, similar work was carried out by Behm et al. [57-59] who looked at
current transients of formic acid oxidation under various conditions, as in the previous
studies, in order to correlate FAO activity with surface species coverage. As in the work of
Samjeské, they found that formate coverage builds up very rapidly to a steady-state value,
while poisoning CO builds up more slowly and hinders the reaction. By tracking the CO
52
coverage with time, they determined that at 0.6 V, adsorbed CO is formed through
dehydration and oxidized at approximately the same rate, resulting in a constant low CO
coverage, but that the current associated with this oxidation represents less that 0.01 % of
the total Faradaic current recorded. At the same potential, it was observed that increasing
the formic acid concentration by a factor of one hundred increased the current by a factor of
twenty and the formate coverage only by a factor of five. The researchers concluded
therefore that the oxidation of formate could be, at most, responsible for 25% of the
Faradaic current. Based on isotopic labeling studies, they further determined that the
formate adsorption/desorption equilibrium was much faster than the oxidation of adsorbed
species. They later concluded that the formate pathway is responsible for no more than 15%
of the current observed during formic acid oxidation, and that the bulk of the current
should be due to the direct pathway, which involves a weakly adsorbed formic acid species
not detectable by IR.
Further complicating matters, recent computational studies on the FAO mechanism
[60] confirmed that the COOH intermediate is disfavored compared to other potential
pathways. They calculate a weakly adsorbed activated HCOOH* which interacts with the Pt
surface in a perpendicular position with the carbonyl oxygen down. By incorporating the
effects of solvation (as opposed to earlier calculational studies of FAO in the gas phase) it
was determined that the direct pathway through a short-lived CO
2
* intermediate or
53
transition state, and a multi-step pathway through a bridge-bonded formate intermediate,
are approximately equally favorable.
The pathway of formic acid oxidation on palladium is even less understood. Arenz
et al. [45] also used FTIR to examine FAO on platinum-palladium electrodes to compare
their respective behaviors. First, oxidation was observed on a platinum electrode with a sub-
monolayer of palladium, representing large isolated Pt and Pd regions. They observed no
adsorbed CO on the Pd regions, but did on the Pt regions. They attribute the apparent
poisoning of the Pd regions to the formation of adsorbed OH. Despite parallels to the PtRu
alloy system, Pd-OH at the interface with Pt regions was observed not to contribute to the
oxidation of CO as Ru-OH does during methanol oxidation. It was speculated that the Pd-
OH bond is too strong to fill this role and that it behaves as a poisoning species.

2.1.3 Scope of this work
The direct formic acid fuel cell possesses many characteristics that make it an
interesting target for low-temperature fuel cell research. Among the possible targets for
carbon dioxide reduction, formic acid is the most straightforward, the process requiring
only two electrons and the formation of just an O-H bond and a C-H bond. A system could
be envisioned wherein CO
2
is reduced to formic acid, which is then oxidized in a direct fuel
cell for energy. The CO
2
released by the fuel cell could be captured and the cycle repeated.
54
For this sort of system to be viable, improvements must be made in several areas of
DFAFC research:
• catalyst efficiency: like the DMFC, there is a large overpotential at the anode and
more efficient catalysts must be discovered.
• catalyst cost: the use of platinum-group metals should be reduced, either through
replacement with other metals or through reduced loading.
• catalyst durability: DFAFCs experience anode catalyst deactivation after only a short
period of use. Reactivation methods have been devised, but are not always practical,
and deactivation due to catalyst dissolution likely cannot be recovered.
In order to address these issues, we have prepared catalysts based on palladium
alloyed with other metals, with low metal:support loadings. These various catalysts have
been evaluated for formic acid oxidation activity. Additionally, we have examined the
incorporation of tantalum carbide as a novel component of the catalyst support. Tantalum
carbide is a hard, conductive, and corrosion-resistant material, and is therefore well-suited
for exposure to formic acid in fuel cell electrocatalyst layers. Tantalum metal has a very low
toxicity and has been deemed biocompatible.

2.2 Experimental methods
Fuel and electrolyte solutions were prepared from chemicals supplied by Fluka
Analytical, including formic acid (purum) and sulfuric acid (TraceSELECT), mixed with
55
deionized water (18 MΩ·cm, Millipore Direct-Q 3). Hydrochloric acid (ACS grade) was
obtained from EMD. Ultra High Pure Grade oxygen, prepurified nitrogen, and argon were
obtained from Gilmont Liquid Air Company. Catalyst precursors, including dihydrogen
hexachloroplatinate (Alfa-Aesar, 99.99%), palladium (II) chloride (Alfa Aesar, 99.9%),
ammonium tetrachloropalladate (II) (Alfa-Aesar), hydrogen tetrachloroaurate (III) (Alfa
Aesar, 99.99%, ACS), antimony (III) chloride (Alfa Aesar, 99.0%, ACS), tin (II) sulfate
(Sigma-Aldrich, 95+%), and iron (III) chloride (Sigma-Aldrich). Vulcan XC-72R carbon was
obtained from Electrochem, Inc. and tantalum carbide (100 mesh) was purchased from
Sigma-Aldrich. Catalysts were prepared according to several methods, listed below.

Catalyst preparation
Catalyst precursors in the desired ratios were dissolved in Millipore water. If
required, a support material was added and the mixture was sonicated for 8 minutes. The
precursor-support mixtures were stirred for approximately 2 hours prior to reduction. The
precursors were reduced with excess sodium borohydride with rapid stirring. According to
the first protocol, reduction was carried out at 80 °C with fast addition of NaBH
4
powder.
According to the second protocol, reduction was carried out at room temperature by
dropwise addition of a freshly prepared aqueous solution of NaBH
4
. In both cases, the
reaction mixtures were stirred overnight and collected by successive centrifugation, water
rinsing, and sonication. After a minimum of three rinses, the clean product was collected
56
and dried at 70 °C. The dry powder was crushed with a mortar and pestle prior to use.
Figure 2-1 shows a representative transmission electron microscope image of one of the
catalysts.

Electrochemical characterization
Electrochemical characterization was carried out in a typical three-electrode glass
cell, in which the counter electrode (a platinum wire) was separated from the main
compartment by a glass frit. The working electrode was a rotating disk electrode (RDE) tip
containing a glassy carbon (GC) disk with an exposed area of 0.197 cm
2
, polished with 5 µm
alumina. Catalyst inks were prepared by suspending 1 mg mL
-1
material in a mixture of
isopropanol, water, and Nafion
®
ionomer solution. The suspension was dispersed by
ultrasonication and a 20 µL aliquot was dropped onto the disk. The electrode was dried in
an oven, and then mounted onto a Pine rotator coupled with a Pine speed control.
Chronoamperometric experiments were carried out with the working electrode rotating at
300 rpm, while cyclic voltammetry experiements were conducted without rotation. The
reference electrode was a mercury/mercurous sulfate electrode (MSE) with a measured
potential of +0.625 V vs. SHE. Experiments were carried out on both a Solartron SI 1270
potentiostat with a frequency response analyzer (SI 1230) and a Princeton Applied Research
VersaSTAT 3.

57
Membrane electrode assemblies
Membrane electrode assemblies (MEAs) for formic acid were prepared with
Nafion
®
-117 membranes (7 mil thickness, Ion-Power, Inc.) and Toray carbon paper
(TGPH-60) with 10 w/w% Teflon wet-proofing. Catalyst inks were prepared from catalyst
powders, a 5% Nafion
®
ionomer solution (Sigma-Aldrich), and deionized water in the mass
ratio of 1:1:3 (catalyst:ionomer:water). The inks were painted onto 25 cm
2
carbon paper
electrodes to achieve a catalyst loading of 8 mg cm
-2
and the MEAs were hot-pressed at
140 °C.
Characterization of the MEAs was carried out in a standard cell housing with pin-
type flow fields (Electrochem, Inc.) using a Fuel Cell Test System 890B (Scribner Associates,
Inc.). MEAs were hydrated for a minimum of 6 hours at 60 °C. Fuel solutions were
circulated using a Micropump pump drive, and a Digi-Sense temperature controller (Cole-
Parmer, Inc.) was used to control the temperature. Pure O
2
was supplied to the cathode at
various flow rates using an Accucal controller (Gilmont, Inc.). The setup is illustrated in
diagrams in section 1.2.

58
2.3 Results and discussion
2.3.1 Palladium-gold
2.3.1.1 Carbon-supported catalysts
In order to evaluate their performance as anode catalysts in direct formic acid fuel
cells, a number of palladium and palladium-gold based materials were prepared and tested in
full cells. The series includes supported (60% metal by mass on Vulcan XC-72R carbon) and
unsupported catalysts prepared from precursors by sodium borohydride reduction:
1. Pd
1
Au
1

2. Pd/C
3. Pd
1
Au
1
/C
4. Pd
20
Au
1
/C
Membrane-electrode assemblies were prepared for each of these catalysts with an
anode loading of 200 mg (thus the anode metal loading for supported catalysts was 120 mg).
The MEAs were enclosed in fuel cell housings and characterized with 2 M HCOOH and
pure oxygen gas.
Current-voltage and power density curves of each of the cells are shown in Figures
2-2 and 2-3. The unsupported catalyst unsurprisingly gave the highest overall current, given
its higher metal loading. Among the supported catalysts, the addition of gold results in
improved performance versus palladium alone.
59

Figure 2-1. TEM images of PdAu-A/[C0.5,TaC0.5] at a) 59kx magnification (scale bar is 100 nm)
and b) 160kx magnification (scale bar is 20 nm).

Figure 2-2. Current-voltage curves for DFAFCs with palladium and palladium-gold anode catalysts.
Recorded with 2 M HCOOH and pure O
2
at 90 °C.
60

Figure 2-3. Power density curves for DFAFCs with palladium and palladium-gold anode catalysts.
Recorded with 2 M HCOOH and pure O
2
at 90 °C.

Figure 2-4. Power density curves normalized to precious metal (Pd and/or Au) content for DFAFCs
with palladium and palladium-gold anode catalysts. Recorded with 2 M HCOOH and pure O
2
at
90 °C.
61
Figure 2-4 shows the power density data with the currents normalized to the
precious metal mass in each catalyst. Despite the high surface area afforded by the Vulcan
carbon support, the metal use is apparently more efficient in the case of the unsupported
catalyst.

2.3.1.2 Tantalum carbide as a support
Given the rising costs of precious metals, and recent record prices for both gold and
palladium, a more efficient use of these metals should be achieved by using lower loadings.
After the relatively poor performance of the carbon-supported catalysts mentioned above,
tantalum carbide was incorporated into the catalyst support (either 10% or 50% by mass) of
a series of palladium-gold catalysts, which were prepared in the following compositions:
1. Pd (20%)/C (20% palladium by mass supported on Vulcan carbon)
2. Pd (20%)/[0.1TaC+0.9C]
3. Pd (20%)/[0.5TaC+0.5C]
4. PdAu-A (20%)/C
5. PdAu-A (20%)/[0.1TaC+0.9C]
6. PdAu-A (20%)/[0.5TaC+0.5C]
7. PdAu-B (20%)/[0.5TaC+0.5C]
PdAu-A signifies a palladium:gold atomic ratio of 98.5:1.5 while PdAu-B signifies
95:5. Figure 2-1 shows representative TEM images of catalyst 6; the support structure
62
shows features in the hundreds of nanometers, while confirming that this method produces
distributed metal particles less than 5 nm in diameter. Tantalum carbide was observed by
SEM and EDX; particle sizes observed over a large range, from approximately 100 nm to
30 µm. The Ta:C atomic ratio was measured to be 1.03:1.

Chronoamperometry
The stability of the catalysts was evaluated by measuring their current output during
formic acid oxidation over extended periods of time. The catalysts were deposited onto a
glassy carbon disk and the electrode was rotated in the sulfuric acid/formic acid solution at
300 rpm. The electrodes were held at a constant potential of +0.300 V vs. SHE, in the region
of formic acid oxidation, and the resulting current was measured over thirty minutes. The
degree to which the current decays over this period indicates its propensity to deactivation
during operation. The specific currents measured are shown in Figures 2-5 and 2-6, and in
2-7 as a percentage of initial current. Specific currents are relative to the mass of gold and
palladium present in a given material.
Among the unsupported and carbon-supported catalysts, the gold-decorated
palladium catalyst had the highest specific current density after thirty minutes while the
other unsupported catalyst, palladium-decorated gold, had the worst (Figure 2-5). Although
the gold-decorated palladium maintained a fairly high current during the first third of the
63
experiment, the difference between the performance of the various materials over thirty
minutes was comparatively small.

Figure 2-5. Chronoamperometric measurements of palladium and palladium-gold catalysts
supported on Vulcan XC-72R and unsupported, recorded at +0.300 V vs. SHE in 0.5 M H
2
SO
4
/
0.5 M HCOOH.

The catalysts with tantalum carbide incorporated into the support were tested in the
same way (Figure 2-6). Both of the cells with 10% TaC support performed very poorly,
giving currents an order of magnitude less than the next lowest result, which was palladium
on 50% TaC. The PdAu catalyst supported on 50% TaC gave the largest current by far,
outpacing the other materials by another order of magnitude.
64

Figure 2-6. Chronoamperometric measurements of palladium and palladium-gold catalysts
supported on Vulcan XC-72R and tantalum carbide, recorded at +0.300 V vs. SHE in 0.5 M H
2
SO
4
/
0.5 M HCOOH.

Figure 2-7 shows these same results, but each curve has been adjusted to show the
current at each point as a percentage of its initial current, to reflect durability primarily
rather than catalyst activity. The two catalysts with 50% TaC support provide greatest
current retention after thirty minutes, with the palladium-only catalyst slightly
outperforming the palladium-gold. The two catalysts which retained the least of their initial
current were those incorporating 10% TaC, followed by the unsupported and the carbon-
supported catalysts performing moderately. Based on these results, PdAu appears to be a
more effective catalyst for formic acid oxidation, providing greater currents, and the
65
replacement of 50% carbon support with tantalum carbide is beneficial to the catalyst’s
durability.

Figure 2-7. Chronoamperometric measurements of palladium-gold catalysts normalized to initial
current, recorded at +0.300 V vs. SHE in 0.5 M H
2
SO
4
/0.5 M HCOOH.

Cyclic voltammetry
The activity of the catalysts for the oxidation of formic acid was further
characterized by cyclic voltammetry. Figure 2-8 shows the voltammograms recorded for the
palladium-only catalysts on various supports. All the catalysts show the expected
voltammogram for FAO on Pd in which a forward and reverse scans show broad oxidation
peaks of approximately equal intensity. The carbon-supported catalyst showed the weakest
currents for formic acid oxidation and also the least sharply defined peaks. The catalyst with
66
50% TaC again gave the highest current while the addition of 10% TaC resulted in only
moderate improvement.

Figure 2.8. Cyclic voltammograms of palladium catalysts supported on Vulcan carbon XC-72R and
tantalum carbide with currents normalized to Pd mass. Recorded in 0.5 M H
2
SO
4
/0.5 M HCOOH at
10 mV s
-1
.

Voltammograms of the supported PdAu catalysts are shown in Figure 2-9. Those
with all-carbon or 10% TaC support gave the lowest currents for formic acid oxidation. The
two catalysts with 50% TaC support showed currents several times greater, and increasing
the Au content from 1.5% to 5% resulted in the highest current output. Figure 2-10 shows
67

Figure 2-9. Cyclic voltammograms of palladium-gold catalysts supported on Vulcan carbon XC-72R
and tantalum carbide with currents normalized to Pd+Au mass. Recorded in 0.5 M H
2
SO
4
/0.5 M
HCOOH at 10 mV s
-1
.

the best performing Pd catalyst and the best performing PdAu catalyst in comparison with
the unsupported catalysts. On a precious metal mass basis, the specific currents of the
supported catalysts were much higher than those of the unsupported catalysts, indicating
that the substitution of a large quantity of TaC for the normal carbon support does not
counteract the advantages of the high-surface area Vulcan carbon material.

Fuel cell testing
Earlier tests showed that the unsupported 1:1 PdAu catalysts gave the best fuel cell
performance even when normalized to the precious metal mass compared to the carbon-
68

Figure 2-10. Cyclic voltammograms of palladium and palladium-gold catalysts supported on Vulcan
carbon XC-72R and tantalum carbide normalized to Pd and/or Au mass. Recorded in 0.5 M
H
2
SO
4
/0.5 M HCOOH at 10 mV s
-1
.

supported catalysts. An MEA was constructed using PdAu-A/[C0.5,TaC0.5] as the anode
catalyst and was tested under similar conditions. The performance of this cell compared
with that of the cell with the unsupported catalyst is shown in Figure 2-11. In terms of raw
current density, the performance of the unsupported catalyst is better, as is expected from
using a fivefold greater amount of metal catalyst. However, when the current is normalized
to precious metal content, both the maximum current and the maximum power are
approximately three times greater in the carbon and tantalum carbide supported catalyst
(Figure 2-12). This shows that the replacement of 50% of the carbon support with tantalum
69

Figure 2-11. Current-voltage (dash-dot) and power density (dash) curves for MEAs with best-
performing supported and unsupported palladium-gold catalysts recorded with 2 M HCOOH and
pure O
2
at 90 °C.

carbide results in a much more efficient use of the catalyst and suggests that the activity of
the palladium-gold material is enhanced by its incorporation.

2.3.1.3 Formate oxidation
The oxidation of formate is considered to be more difficult than the oxidation of
formic acid, but a direct formate fuel cell would help to address the problem of the corrosive
nature of formic acid fuels. Consequently, the PdAu-B/[C+TaC] catalyst was evaluated for
the oxidation of formate in alkaline solution.
The cyclic voltammogram recorded for this catalyst in a solution of potassium
70
Figure 2-12. Current-voltage and power density curves for MEAs with best-performing supported
and unsupported palladium-gold catalysts normalized to Pd/Au mass; recorded with 2 M HCOOH
and pure O
2
at 90 °C.

formate is shown in Figure 2-13. Its behavior is quite similar to that of formic acid: a strong
oxidation peak on the forward scan followed by a slightly stronger peak on the reverse scan
in the same region. Interestingly, the reverse peak for formate oxidation is approximately
120 mV cathodic of the peak for formic acid oxidation, which suggests that the kinetics for
formate oxidation are more favorable than those of formic acid on this catalyst.
Chronoamperometric experiments were conducted on samples of the PdAu-
B/[C+TaC] catalyst in the presence of both formic acid and formate at their respective peak
oxidation potentials gleaned from the cyclic voltammtry experiments in order to gauge their
stability when operating in regions of maximum current density (Figures 2-14 and 2-15).
71

Figure 2-13. Comparison of cyclic voltammograms of PdAu-B/[C0.5+TaC0.5] in formic acid and
potassium formate solutions normalized to Pd+Au mass. Recorded at 10 mV s
-1
.

The behavior of the catalyst is very similar in both cases, and both retain approximately the
same relative current after thirty minutes of operation.

2.3.2 Palladium-platinum
Among the most studied palladium alloy catalysts for formic acid oxidation is
palladium-platinum. As described in section 2.1.2, platinum/palladium catalysts show varied
behavior depending on their compositions and morphologies. Surfaces composed of isolated
palladium and platinum regions did not show synergistic behavior, while alloyed materials
behaved differently than the component metals.
72

Figure 2-14. Comparison of chronoamperometry of PdAu-B/[C0.5+TaC0.5] in formic acid and
potassium formate solutions normalized to Pd+Au mass. Recorded at 300 rpm.

Figure 2-15. Comparison of chronoamperometry of PdAu-B/[C0.5+TaC0.5] in formic acid and
potassium formate solutions expressed as a percentage of initial current. Recorded at 300 rpm.
73
In particular, PtPd in a 1:1 molar ratio has been shown to be a more durable catalyst
than palladium black. However, in addition to taking on some of the stability of platinum, it
also oxidized formic acid at a potential intermediate between those of pure palladium and
platinum: this higher overpotential compared to the palladium electrode rendered it
impractical for fuel cell use.
Given the potential for TaC to improve catalytic activity by weakening interactions
with strongly bound intermediates, it may provide a benefit with platinum-containing
catalysts as well. We have therefore evaluated a carbon and tantalum carbide supported
platinum-palladium catalyst similar to the palladium gold catalysts discussed previously. The
catalyst was prepared from precursor salts as described in section 2.2.

Chronoamperometry
The working electrode coated with the catalyst was evaluated at a controlled voltage
of +0.300 V vs. SHE while the current was measured over thirty minutes. This is shown in
Figure 2-16. Compared to the palladium palladium-gold catalysts, the current measured
over this time with the palladium-platinum catalyst was found to be several orders of
magnitude inferior. Pure platinum catalysts oxidize formic acid weakly in the region
characteristic of FAO on palladium, due to CO formation, but show stronger oxidation at
higher overpotentials where CO is readily oxidized. It appeared based on this weaker
current density that the palladium-platinum catalyst was exhibiting more platinum-like
74
behavior. However, it did retain a larger percentage of its initial current after thirty minutes.
Figure 2-17 shows a comparison of the current retention of the 20% PdPt and the 20% Pd
catalysts on TaC+C. The catalyst including platinum reached a steady current much more
rapidly than the pure palladium catalyst. Both of these observations are consistent with
those previously reported for PdPt systems.

Figure 2-16. Comparison of cyclic voltammograms of carbon-tantalum-carbide-supported PdAu and
PdPt catalysts normalized to Pd+Au mass. Recorded in 0.5 M H
2
SO
4
/0.5 M HCOOH at 10 mV s
-1
.

Cyclic voltammetry
In order to better understand the FAO mechanism on the PdPt catalyst, cyclic
voltammograms were recorded in formic acid solution over the relevant potential range
75
(Figure 2-18). As postulated, the catalyst behaves in a manner typical of other platinum-
based materials: a weak, broad oxidation peak on the forward scan, indicating CO
formation, followed by a sharper peak at higher potential indicating CO and oxide removal.
On the reverse scan, direct formic acid oxidation occurs more rapidly on the freshly-cleaned
surface. Comparing the FAO peak potentials on the reverse scan, the palladium-gold catalyst
oxidized formic acid at a potential approximately 200 mV more negative than the platinum-
palladium catalyst, a significantly lower overpotential. The peak for FAO on the forward
scan in this region is very weak with the platinum-palladium catalyst.

Figure 2-17. Chronoamperometry of PdPt/[C0.5,TaC0.5]. Recorded in 0.5 M H
2
SO
4
/0.5 M
HCOOH at 300 rpm.

76

Figure 2-18. Comparison of chronoamperometry of PdAu-B/[C0.5+TaC0.5] and
PdPt/[C0.5,TaC0.5] expressed as a percentage of initial current. Recorded in 0.5 M H
2
SO
4
/0.5 M
HCOOH at 300 rpm.

2.3.3 Palladium and other metals
Further research will be more focused on incorporating cheaper, more abundant
materials in order to reduce catalyst cost. To this end, we have studied the performance of
supported, multi-metallic alloy catalysts including palladium and other metals in direct
formic acid fuel cells. The cataylsts studied, prepared as 60% metal on Vulcan XC-72R
carbon, are as follows:
1. PdSn (1:1)/C
2. PdSn (20:1)/C
77
3. PdFe (1:1)/C
4. PdSb (1:1)/C
5. PdAuSn (1:1:1)/C
6. PdAuSnFe (1:1:1:1)/C
MEAs were constructed with each of these as the anode catalyst with a loading of
200 mg (8 mg cm
-2
) and the fuel cells were characterized at 90 °C with 2 M HCOOH fuel.
Current-voltage and power density curves for these cells are shown in Figures 2-19
and 2-20. The palladium-antimony catalyst, which was recently reported to be very active
for formic acid oxidation [52, 53], gave the weakest performance by far in this series of tests.
Iron and tin, on the other hand, show promise as catalytic components and give reasonable
performance compared to the Pd-only baseline. The ternary material, palladium-gold-tin in
equimolar proportions, gave the best raw maximum power density among the supported
catalysts studied. However, at lower current densities (< 100 mA cm
-2
, representing slower
fuel consumption) both the palladium-tin and palladium-gold catalysts performed better
than the ternary catalyst and nearly as well as the unsupported PdAu.
Figure 2-21 shows the performance with currents normalized to precious metal
mass (i.e., Pd and/or Au); by comparison, the cost of tin, iron, or antimony is negligible
compared to overall cost. The best-performing catalyst on this basis is the ternary
palladium-gold-tin material, which shows the most cost-effective use of precious metals.
78

Figure 2-19. Current-voltage curves for DFAFCs with palladium-M
x
catalysts. Recorded with 2 M
HCOOH and pure O
2
at 90 °C.

Figure 2-20. Power density curves for DFAFCs with palladium-M
x
catalysts. Recorded with 2 M
HCOOH and pure O
2
at 90 °C.
79

Figure 2-21. Power density curves for DFAFCs with palladium-M
x
catalysts, with currents
normalized to anode precious metal (Pd and/or Au) mass. Recorded with 2 M HCOOH and pure O
2

at 90 °C.

2.4 Conclusions
Palladium-gold based catalysts are promising for the oxidation of formic acid in
direct fuel cells because they address some of the major problems of the most active pure
palladium catalysts: catalyst deactivation and dissolution. In order to address the issue of
cost, metal loading should be reduced, less expensive metals should be added, and novel
support materials should be investigated. It has been shown that carbon-supported
palladium-gold catalysts are enhanced by the replacement of 50% of the support with a
commercially available tantalum carbide. PdAu catalysts supported on this blend, in contrast
80
to those supported on carbon alone, showed a greater mass-specific currents than the best-
performing unsupported palladium-gold catalyst. Electrochemical characterization also
showed that the durability and activity of the catalysts on this support during FAO is much
improved.
Palladium-platinum, which is very frequently studied as an anode catalyst for
DFAFCs, showed evidence of CO poisoning and oxidized formic acid at an overpotential
200 mV greater than the PdAu catalysts. The incorporation of other, less expensive metals,
such as iron and tin, is promising based on fuel cell performance and future work will focus
on preparing these less expensive catalysts while incorporating the novel tantalum-
carbide/carbon based support.

81
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85

Advanced electrolytes for nonflammable lithium-
ion batteries

3.1 Introduction to lithium-ion batteries
3.1.1 Historical overview
Over the last two hundred years, many battery technologies have been developed,
reached commercial viability, and been abandoned for a better alternative, particularly in
the last four decades. Many variations on battery chemistries have been devised; what
follows is a selective overview of those that came in and out of favor leading up to the
development of the lithium-ion battery, including both its predecessors and its
contemporary competitors [1].

3.1.1.1 Secondary batteries
The lead-acid battery is the oldest secondary (rechargeable) battery technology in
use today, and is still one of the most common. It was first demonstrated by Raymond
Gaston Planté in 1859, and consisted of two lead electrodes separated by a porous cloth and
submerged in 10% sulfuric acid. The design remains remarkably unchanged today; while
originally a formation process was employed, consisting of the corrosion of the plates to
form lead dioxide, modern lead acid batteries are constructed pre-charged from solid lead
3

86
plates (the negative electrodes) and lead grids impregnated with a lead dioxide-sulfuric acid
slurry (the positive electrodes). Upon discharge, the negative electrode is reduced and the
positive electrode oxidized, both to lead sulfate. The electrolyte, sulfuric acid, participates in
the electrode reactions, and therefore its concentration impacts the cell voltage [2, 3].
The car battery is the best-known example of the lead-acid battery’s continued use;
it owes its longevity in part to its ability to provide short bursts of very high current
necessary to start a combustion engine, which is made possible by the high surface area of
the corroded lead electrodes. While massive lead is detrimental to the cell’s gravimetric
energy density, its power density is competitive thanks to the couple’s relatively high
thermodynamic potential (generally > 2 V). Variants designed for greater depth of discharge
often incorporate thicker, more durable electrodes able to provide more moderate current
over extended periods without loss of structural integrity.
The first challenger to the lead-acid battery in commercial markets was the nickel-
cadmium cell in the late nineteenth and early twentieth centuries, an aqueous alkaline cell
based on a concentrated potassium hydroxide electrolyte. During discharge, the nickel
oxyhydroxide positive electrode is reduced to nickel (II) hydroxide, while the cadmium
negative electrode is oxidized to cadmium (II) hydroxide. (The nickel-iron battery, based on
the same chemistry but with an iron metal negative electrode, has comparatively higher
energy density owing to its higher voltage and iron’s lower mass; however, stability issues

87
related to the corrosion of iron and the consequent evolution of hydrogen led to its fall out
of favor relatively early.) Despite its lower operating voltage, the nickel-cadmium battery
has the advantages of better cycle life, better high-rate capability, and better low-
temperature performance than lead-acid batteries, with comparable-to-higher gravimetric
power densities. These factors led to its continued popularity through the 1990s, where the
nickel-cadmium cell was the most widely used rechargeable alternative to the alkaline
primary cell for consumer applications. However, the cost and toxicity of metallic cadmium
eventually led to their decline in popularity as other technologies were developed [4].
The nickel-metal hydride cell was first commercialized in the late 1980s as an
alternative to Ni-Cd cells with high specific power and energy [5]. Metal hydride batteries
evolved from earlier hydrogen/metal oxide cells, in which pressurized hydrogen was
oxidized at a platinum electrode and a nickel oxyhydroxide electrode was reduced during
discharge. The use of pressurized gaseous H
2
and the concomitant storage vessel drastically
limited the energy density of these cells and introduced a serious risk of explosion. Work on
hydrogen storage alloys had been ongoing since the early 1970s, and several of these
materials were tested in hydrogen batteries. As the negative electrode, the metal alloy must
not only be able to absorb and desorb hydrogen reversibly, but also to oxidize and reduce it
efficiently [6]. The first such application reported used LaNi
5
[7] and the use of other AB
5

and AB
2
type materials soon followed. It was not until the late 1980s and early 1990s,

88
however, that materials with sufficient cycle life and ability to store hydrogen under
pressure suitable for consumer use were developed [8-10]. These materials, allowing the
manufacture of cells with a volumetric energy density twice that of Ni-Cd cells, led to rapid
commercialization and widespread adoption of NiMH over the 1990s. The Toyota Prius
(1997) and the Honda Insight (1999), arguably the most successful mass-market hybrid
electric vehicles (HEV), use NiMH battery packs [11, 12]. Beginning in the 2000s, the
market share of the nickel-metal hydride battery has decreased due in large part to the
advent of the lithium-ion battery, but it still maintains a significant presence among
household secondary batteries and HEVs [13].

3.1.1.2 Lithium metal and lithium-ion
Primary lithium batteries saw growth in the early 1970s [14, 15] when the
lithium/carbon monofluoride primary battery was commercialized by Panasonic
(Matsushita Co.). Lithium ions dissolve from the negative electrode and form LiF and
carbon on the positive electrode during discharge. Lithium is the lightest solid element, and
possesses the most negative standard reduction potential, making the lithium electrode ideal
for both high energy densities and high voltages. The early to mid 1970s also saw the advent
of insertion electrode materials (see §3.1.3), such as manganese dioxide (MnO
2
) and
titanium disulfide (TiS
2
). The first successful implementation of a lithium/manganese
dioxide primary cell was reported between 1973 and 1975 by Ikeda et al. of Sanyo Electric

89
Co., Ltd. [16]; the company subsequently commercialized it in 1976, and a variant of this
chemistry is still used today for general consumer applications where its high voltage (3 V)
and high energy density (280 Wh kg
-1
or 580 Wh l
-1
) are desirable. Whittingham (Exxon)
reported the first lithium secondary cell with a practical energy density in 1976, using TiS
2

as the positive electrode material [17, 18]. In 1980, Goodenough and coworkers described
LiCoO
2
, the first of many positive electrode compounds that were to come out of their work
[19]. The cell was demonstrated to charge to open circuit voltages in excess of 4.2 V vs.
Li/Li
+
with 50% LiCoO
2
lithium utilization and was able to sustain high current densities.
Although reversible lithium intercalation compounds would indeed play a vital role in
rechargeable battery technology, the transition from primary to secondary lithium batteries
was not a simple one.
While the necessary charge reaction at the negative electrode of a secondary lithium
battery—the plating of metallic lithium onto the depleted electrode—is facile, the
morphology of the resulting Li electrode surface is noticeably different from that of an
uncycled one. Instead of a dense, bright, metallic surface, reconstituted Li electrodes are
spongy and characterized by the formation of dendrites. During the subsequent discharge
processes, some of these dendrites may become separated from the bulk during the stripping
process, resulting in a loss of active material in the form of loose, highly chemically reactive
Li pieces. More dangerous still is the potential for the longer dendrites to pierce the

90
electrodes’ separator material and short circuit the cell, causing it to self-heat and
subsequently vent.
The most well known example of the danger of lithium secondary cells was the
Molicel, introduced by Moli Energy, Inc. of Vancouver in 1985. This cell (not to be
confused with the modern Molicel by E-One Moli Energy Corp.) used a molybdenum
disulfide positive electrode which was rated for 250 cycles over a 1.3 − 2.4 V operating
range. These cells were used in mobile phones in Japan for three years, until a failure
incident in 1989 caused them to be discontinued. The practical advantages of the Molicel,
weighed against its catastrophic safety record, was an important impetus toward the
development of lithium-ion as an alternative, and the need to focus on safety as well as
performance considerations [1].

3.1.1.3 Lithium-ion and beyond
The concept of the lithium-ion cell began life in the 1980s [20-23] and was dubbed
the “rocking chair battery” [24] because lithium ions present in the electrolyte are reversibly
transferred back and forth between the two electrodes during charge and discharge. A
material that can reversibly insert or intercalate lithium at a potential near that of the
lithium electrode serves as the negative electrode, and a high potential intercalation
compound forms the positive electrode (see §3.1.2). While the voltage of such a cell is
necessarily lower than one that uses a lithium metal negative electrode, removing the need

91
for lithium plating greatly reduces the risk of internal short circuit. Graphite negative
electrodes [25-28] provide an electrode reaction that is approximately 90 mV more positive
than the Li/Li
+
electrode, maintaining high power while improving safety.
While both positive and negative lithium insertion compounds were known earlier,
the lithium-ion battery was not commercialized until 1991 when Sony [29] put the first
generation cell on the market. It used a lithium cobalt oxide positive electrode and a
petroleum coke negative electrode; the latter would be replaced within a few years by
graphitic carbons, which have a higher capacity and lower average potential, but which
suffered from poor reversibility since their introduction in the 1970s. The discovery by
Dahn’s group of electrolyte systems that form a stable passivating film called the solid-
electrolyte interphase (SEI) finally enabled the use of these materials [30]. Within the last
two decades, lithium ion batteries have become the best-selling secondary cell and are used
in most personal and hand-held electronic devices sold today.
Personal electronics are becoming more complex and are consequently demanding
more power, and users are demanding longer operating times between charges. As
manufacturers look to these and larger-scale applications, such as electric vehicles, a need
for larger-capacity batteries with increased high-power capabilities is becoming clear. In
parallel to the continued improvement of lithium-ion technology, several new technologies
have also been proposed to fill this need. Lithium-air and lithium-sulfur batteries both

92
promise extremely high theoretical energy densities, and have been hailed by some as the
eventual replacement for lithium ion technology; at the present time, however, neither has
been developed commercially.

3.1.2 Operating principles and requirements
3.1.2.1 The "rocking chair" cell
A lithium-ion cell [2, 31, 32] contains no lithium metal electrode, but most
commonly a lithium-transition metal oxide positive electrode and a graphite negative
electrode. The two electrodes are separated by an organic electrolyte containing a lithium
salt. Lithium plating and stripping and the accompanying difficulties are avoided by instead
shuttling lithium cations between two host structures. A graphite-LiCoO
2
cell is
schematized in Figure 3-1. Cells are assembled from these materials in the discharged state.
Lithium cobalt oxide has a layered structure, and during charge, Li
+
is deintercalated from
the structure and enters the electrolyte, while Li+ from the electrolyte is intercalated into
the graphite electrode. During charge, the positive and negative electrodes reach ~4.2 V and
0.10 V vs Li/Li
+
, respectively, for a 4.1 V cell. The capacity of the cell is limited largely by
the positive electrode, which can reversibly relinquish approximately 50% of its lithium
while still maintaining its crystal structure. The graphite electrode can in theory
accommodate lithium up to a final composition of LiC
6
(see §3.1.3.2). During discharge,
lithium is de-intercalated from the graphite electrode and re-intercalated into the oxide

93

Figure 3-1. Schematic diagram of a lithium-ion cell.

94
material. If a cell is not overcharged, the coulombic efficiency of this process can approach
100% (the watt-hour efficiency, however, will be lower).

3.1.2.2 Energy and power density
Currently, lithium-ion batteries have a specific capacity of over 150 Wh kg
-1
. As
devices become more demanding, and especially as electric cars become more and more of a
reality, there is a need to improve the energy density of these batteries to better meet these
applications.
The most serious limitation of electric cars is their range. As of 2012, commercially
available electric vehicles include the Tesla Roadster (EPA: 244 miles), the Nissan Leaf
(EPA: 73 miles), and the Chevrolet Volt (a plug-in hybrid; EPA pure electric range, 35
miles). (Depending on the reservoir size, fuel economy, and driving conditions, gasoline-
powered cars can have a range in excess of 500 miles.) The second issue is the charging
time; as of now, with charge times in excess of 6 hours from a high-voltage power source, it
is only practical to charge an electric vehicle overnight.
Batteries for aerospace applications also have demanding energy density and specific
energy requirements. Mass and volume should be minimized as much as possible for more
efficient launching, and the batteries must conform to the exacting specifications of the
design of the device they are powering. Batteries incorporated on a Mars rover, for example,
must be taken into account during rover design, launch, and landing.

95
3.1.2.3 Cycle and storage life
Unlike primary batteries, whose useful life ends after a single discharge, secondary
battery chemistries have the added requirement of reversibility. An electronic device may
undergo hundreds, if not thousands, of cycles over the course of its lifetime. An electric
vehicle, charged just three times per week over twenty years, will undergo over 3000 cycles,
and in many cases would be cycled more often. The lifetime of space missions whose
operation depends on rechargeable battery power depends on the cycle life of the batteries,
which can never be replaced. The primary mission for the 2003 Mars Exploration Rovers,
for example, was ninety days; nine years later, one of the rovers, Opportunity, continues to
operate. (The second rover, Spirit, fell victim to dust-covered solar panels and could not
recharge its cells.)
In addition to watt-hour efficiency losses, as batteries are cycled, a certain amount of
energy is lost to side reactions during the charging step, known as irreversible capacity.
These reactions consist primarily of decomposition of cell electrolyte components on the
electrodes. Even if the coulombic efficiency of a cell is high, over the course of many
hundreds or thousands of cycles, this loss of active material reduces cell capacity while
increasing electrode impedance, and potentially causes pressure to build up from the gaseous
products of the irreversible reactions.

96
3.1.2.4 Operating temperature range
Necessary operating temperature ranges depend on the application. Consumer
electronics have more forgiving requirements, as they are less often exposed to extreme
environments. Electric vehicles must be able to operate in all climates, and the battery packs
must resist excessive heating during operation. The inability of rechargeable batteries to
operate over a wide temperature range is most often due to the electrolyte freezing at low
temperatures, which can reduce the ionic conductivity to zero; and thermal decomposition
of various components at high temperatures, which leads to impedance growth, loss of
capacity, and possible cell venting.
Aerospace applications require complex thermal management to balance the heat
produced by onboard components and the broad variations in temperature according to
environment, season, or time of day [33]. To enable future planetary rovers, the Jet
Propulsion Laboratory has set a target operable range of −60 to +60 °C for lithium ion
batteries. Human-rated and automotive applications demand narrower ranges of
approximately 0 to +30 °C and -30 to +60 °C, respectively.

3.1.2.5 Safety
To be considered safe for consumer applications, as well as for human-rated
spaceflight applications, lithium-ion batteries need to meet certain safety standards. While
aqueous systems are inherently safer, so far aqueous lithium secondary batteries have

97
inferior energy density and cycle life compared to lithium-ion cells with organic solvents.
Thermal runaway in a cell can be caused by several situations, including internal shorts and
exothermic reactions, improper overcharging, and physical abuse. Changing the cell
chemistry can address the first issue, and the use of additives such as redox shuttles can
address the second. Reducing the flammability of the electrolyte addresses all of these
potential issues by removing or reducing the risk of fire in any situation.
Batteries for cell phones and other personal electronics need to withstand frequent,
high-rate charging alternating with longer periods of low-rate discharge. Batteries for
electric vehicles need to be able to withstand tremendous impacts without catching fire or
exposing passengers to hazardous materials. Rapid overcharge can lead to electrolyte
decomposition and excessive gas evolution, and may cause a dangerous cell rupture even if
the cell components are non-flammable. This applies not only to the engineering of the cell
casing, but to the cell chemistry as well. Batteries for satellites and rovers have little concern
for fire, as there are no passengers and no oxygen, and efforts are focused mainly on
performance. Manned spaceflight, however, demands a rigorous safety standard for a
situation in which the unit may not be replaced and the occupants may not be able to escape
in the case of a failure.
The biggest safety liability in a lithium-ion cell is the hydrocarbon-based electrolyte
solvent system. Consisting mainly of organic carbonates (and in the past, esters and ethers),

98
they are flammable and mostly volatile, and often have flash points at or below room
temperature. Secondly, the most commonly used positive electrode materials,
lithium/transition metal oxides, can decompose to release O
2
at high temperatures,
pressurizing the cell and potentially causing it to vent. Finally, thermal decomposition of
electrode films (see §3.1.3.2) and of the electrolyte salt lead to the formation of highly
reactive products.
There are two major strategies for improving the fire safety of lithium-ion cells:
replacing the hazardous components with less hazardous ones, or blending in additives that
counteract these hazards (such as flame retardants or overcharge protectors). These
approaches will be outlined in sections 3.1.4.3 and 3.1.4.4.

3.1.3 Electrode materials
3.1.3.1 Positive electrode materials
As described in section 3.1.1.2, lithium insertion materials have been paired with
primary and secondary lithium negative electrode cells since the 1970s. Whittingham
examined TiS
2
[34], due to its layered structure, conductivity, low mass, and low cost. He
was able to show [18] that during discharge, lithium inserts into the electrode and forms a
single ternary phase:

xLi + xe
-
+ TiS
2
Li
x
TiS
2
(3.1)

99
The theoretical capacity of the TiS
2
electrode (i.e., for x = 1) is 239 mAh g
-1
.
Observed values for x were above 0.9 when discharged at low rates. Following these reports,
many oxides and chalcogenides were extensively characterized in the late 1970s for battery
applications, with a focus on the structures of the materials and the nature of the metal ion
insertion [35, 36].
In 1980, Goodenough and coworkers reported [19] another layered structure,
LiCoO
2
, in which lithium and cobalt occupy alternating layers, and oxygen atoms form a
cubic close-packed array. During charge, lithium is removed forming a layered CoO
2

structure; during discharge, lithium is reintercalated into the electrode.

LiCoO
2
Li
1-x
CoO
2
+ xLi
+
+ xe
-
(3.2)

It was tentatively concluded that layer structure integrity was maintained, and
therefore that reversible cycling was possible, for x ≤ 0.5. As a result, while LiCoO
2
has a
theoretical capacity of 274 mAh g
-1
, its usable specific capacity is only 140 mAh g
-1
. Three
years later, the same group reported [37] the insertion of lithium into the Mn
2
O
4
and
Mn
3
O
4
spinel structures, in which Li
+
occupies the octahedral holes. Shortly thereafter, they
described the mechanism of lithium insertion and removal [38] against Li and determined
that the Li
x
Mn
2
O
4
spinel structure is stable for compositions between 0.4 ≤ x ≤ 2.0. In 1985,

100
the same group reported [39] a cell based on a lithium-aluminum alloy electrode and
LiNiO
2
, and the latter’s thermal transformation to the spinel LiNi
2
O
4
.
In the 1980s, several systems that could be described as “lithium-ion” were reported
[21, 23, 40, 41] but the first successful commercialization was accomplished by Sony in
1991. The cells were based on LiCoO
2
and petroleum coke, and were charged to 4.1 V. The
breakthrough for secondary lithium batteries came about with the development of positive
insertion electrodes, and the transition to lithium ion was brought about by the successful
replacement of lithium with negative intercalation electrodes (see §3.1.3.2). Subsequent
research and development was focused on improving the specific energy of the lithium ion
cell, which is mostly limited by the positive electrode’s low specific capacity (and that of coke
negative electrodes, before the advent of graphite materials). Consequently, many studies
have focused on increasing the positive electrode’s reversible capacity or its potential. The
spinel LiMn
2
O
4
has received much attention [42] and has a theoretical capacity of
148 mAh g
-1
, but its reversible capacity is less than LiCoO
2
. LiNiO
2
has a higher reversible
capacity but lower a potential. Its ability to form a solid solution with LiCoO
2
and LiAlO
2

has led to extensive work on LiNiCoO
2
(NCO) [43], LiNiAlO
2
[44], and LiNiCoAlO
2

(NCA), materials with higher capacities but lower operating voltages. The cost of cobalt is
of concern; nevertheless, NCO and NCA are widely used in the battery industry and NCA is
considered the state-of-the-art. While polymeric materials have gained some traction as

101
positive electrodes in lithium-ion batteries [45, 46], the vast majority of work is focused on
transition metal oxides and chalcogenides.
In 1997, Goodenough and associates introduced a new polyanion-type lithium
insertion material with an olivine structure, LiFePO
4
[47]. Lithium iron phosphate was
identified as the only example from among several olivine-type materials (i.e. LiMPO
4
)
where lithium extraction was possible: approximately 60% of the lithium could be utilized
reversibly. The material exhibited a very flat discharge profile at around 3.5 V vs. Li/Li
+
,
and a capacity of 135 mAh g
-1
(theoretical capacity is 170 mAh g
-1
). Later developments by
others [48, 49] showed improved performance. Despite its lower operating voltage, the
LiFePO
4
electrode has the advantage of being low cost, relatively benign, and safer than
other materials. Recent claims [50] that LiFePO
4
could be the key to ultra-fast charging
batteries have been controversial [51, 52].
Recently, attention has been given to high voltage cathode materials, such as
LiNiMnCoO
2
[53-55] and LiMnNiO
4
[56-59], with a view to high-power applications but
these are not yet commercially viable.

3.1.3.2 Negative electrode materials
The lithium metal electrode naturally provides the lowest potential among negative
electrodes. It also possesses the highest theoretical specific charge capacity (3862 mAh g
-1
).
Its practical performance, however, is hindered by the formation of dendrites during lithium

102
deposition, which are partially dissolved and lost as inactive material during subsequent
cycles. To compensate for this loss of material, a large excess of lithium is used, to the
detriment of the energy density of the cell. This necessity, as well as the safety
considerations outlined above, allow negative electrode materials with sufficiently low
potentials to be competitive.
Lithium alloys were among the first lithium replacement materials tested. Dey
showed in 1971 that alloys could be formed by reductively polarizing a metal electrode in a
Li
+
electrolyte [60]. In a lithium alloy battery, Li
+
is alloyed into the negative electrode
during charge, and oxidatively removed during discharge.

Li
x
M xLi
+
+ xe
-
+ M (3.3)

During lithiation, a complex series of phases and intermetallic compounds are
formed, causing considerable structural changes in the host material [61]. The highly ionic
character of the resulting material is associated with expansion during lithiation due to the
larger ionic radii compared to the neutral material. Given that ionic materials tend to be
brittle, the repeated expansion and contraction during cycling causes rapid degradation of
mechanical integrity in lithium alloy electrodes. Several attempts have been made to combat
this, including reducing the alloy particle size [61] and incorporating the alloy particles in an

103
inactive matrix [62]. Due to their high theoretical specific capacities, interest and research
into Li alloy materials continue to this day [63].
Tin and silicon are two Li-alloy materials that enjoy much attention today [64].
Silicon is especially interesting owing to its natural abundance and lithium alloying capacity.
On the other hand, like other alloys silicon undergoes a 300% volume expansion upon
lithium insertion, followed by contraction upon lithium removal; the resultant pulverization
and loss of conductivity of the Si structure limits its cycle life, and this challenge remains the
focus of most research on Si negative electrodes [65].
The first truly promising lithium negative electrode insertion compounds were
based on carbon [66]. Both graphitic and non-graphitic carbon materials are known,
although the term “graphite” is commonly used for materials with layered structures lacking
the perfect stacking arrangement required by the rigorous crystallographic definition
[67, 68]. “Intercalation” is a term used to describe the process of a species (i.e., lithium)
inserting into a layered material, and is often accompanied by expansion of the host layer
structure. Lithium has been known to intercalate graphite since the 1950s [25] and can
reach a maximum lithium content of LiC
6
. This equates to a theoretical capacity of
372 mAh g
-1
for carbon. Intercalation typically occurs in stages [28, 69], in which guest
lithium species initially occupy a few spatially distant layers, minimizing the energy required
to expand them; intermittent layers are then filled, progressively reducing the distance

104
between guest-occupied layers until all are filled. At least five stages have been identified
during the intercalation process toward LiC
6
, each distinguished by a characteristic number
of n adjacent unoccupied layers between lithiated layers. Full intercalation results in an
expansion of the graphite interlayer distance from 3.35 Å to 3.7 Å, or about 10% [70], which
is significantly less than other materials.
The intercalation process is reversible, but coulombic efficiency suffers due to side
reactions stemming from reactivity between lithiated carbons and electrolyte components,
which consume lithium and capacity during the first few charge phases. These irreversible
reactions contribute to the “irreversible capacity” of the cell, or capacity which is not
returned during discharge, and which may reach up to 20% for carbon electrodes. Thus, in
the absence of highly reductively stable electrolyte solutions, the viability of carbon
electrode cells depends on the formation of a solid-electrolyte interphase, or SEI, which
passivates the carbon surface. This phenomenon was first proposed by Peled and Dey in the
1970s [71-74] for lithium and other alkali electrodes. In the 1990s, Dahn and coworkers
demonstrated the importance of the same phenomenon on carbon electrodes in non-
aqueous systems [30, 75]. Further studies into the process were carried out by others as well
[76-79]. The formation of an SEI results in a passive layer that protects the electrode while
still remaining permeable to lithium cations. This limits the excessive irreversible capacity
to the first cycle (or first several cycles), allowing coulombic efficiency to approach 100% on

105
subsequent cycles. Ethylene carbonate was found to promote SEI growth (see §3.1.4.1) and
certain additives, known as SEI-promoting or film-forming agents, can aid the process in
electrolyte systems where it is insufficient (see §3.1.4.4). SEI formation will be further
developed in the following section.
Further complicating matters, the ubiquitous solvent used in primary batteries,
propylene carbonate (PC), was found to be electrochemically active during graphite
lithiation as evidenced by a plateau beginning around +0.8 V vs. Li/Li
+
[80]. Not long after,
it was elucidated [81] that in addition to this decomposition reaction, this potential region is
characterized also by the co-intercalation of PC along with Li
+
, which results in much more
severe expansion of the host and subsequent exfoliation of the graphite electrode. This and
other considerations for the selection of electrolyte solvents will be addressed further in
section 3.1.4.
The mesocarbon microbead (MCMB) is a type of graphitic carbon first described by
Honda and Yamada in 1973 [82-85]. MCMBs are low surface area materials consisting of
carbon microspheres (diameter ≤ 40 µm) typically prepared from organic precursors and
graphitized at high temperature (1300 ! 3000 °C). Its use as a lithium intercalation negative
electrode was investigated during the early days of lithium-ion battery research [86-89] and
it remains a commonly used material today (for example, [90-92]).

106
Hard carbons, such as cokes, were incorporated into lithium-ion systems earlier
than graphitic carbons, but suffer from several disadvantages including low density and
large charge-discharge hysteresis [93]. Compared to the layered materials, their more
disordered structure provides a mechanical barrier to solvent coinsertion. As such they
perform well with PC-based electrolytes [75], unlike graphitic carbons. While the specific
capacity of some hard carbons is greater than that of graphite materials, their low density
nullifies this advantage from the standpoint of volumetric capacity [30, 94-96]. Coke
materials have, however, been seen to have high rate capabilities [97] in the absence of the
intercalation process seen in graphitic materials.
Hard carbons, treated between 500 and 1000 °C, retain heteroatoms from their
organic precursors, the cross-linking effect of which contributes to the reduction of
crystallinity and to the solvent insertion barrier [93]. While they tend to show high
reversible capacities, the irreversible capacities on the first cycle tend to be very high; in
addition to SEI formation, which may be magnified by the presence of pores and therefore
greater surface area, reactivity between lithium and some heteroatom groups is likely [98].
Insertion materials, such as oxides [23], chalcogenides [99], and polymers [100],
were also investigated early on as replacements for metallic lithium, but were limited by
their lower capacities. The most promising oxide materials are those based on titanium
oxides [101] and the most important of these so far is Li
4
Ti
5
O
12
(LTO), a spinel structure.

107
Due to high operating potentials (i.e. +1.5 V vs. Li/Li
+
for LTO), common electrolytes are
more stable in titanate-based systems and SEI formation is not necessary. While this reduces
the cell voltage, operating at a potential positive to that of lithium plating reduces the risk of
dendrite formation as well as electrolyte decomposition. Furthermore, lithium intercalation
occurs without added strain, increasing the life of the electrode. LTO-based systems have
been shown to be good performers in terms of power capability, operating temperature
range, and abuse tolerance [102].

3.1.4 Electrolyte solvents and additives
While the number of solvents available in the realm of organic chemistry is
substantial, those suitable for lithium secondary battery systems must meet a stringent set of
requirements that narrow the field considerably. Charged electrodes possess strong
oxidizing (the positive electrode) or reducing (the negative electrode) power, so a solvent
should be stable within an approximate range of 0 − 5 V vs. Li/Li
+
(this criterion alone
eliminates all protic solvents). Good electrolyte conductivity requires good lithium salt
solubility (facilitated by high dielectric solvents) and high lithium ion mobility (facilitated by
low viscosity solvents. If a cell is to be operable in a range of environments, the liquid range
of the solvents should be broad. For safety reasons, its flammability and toxicity should be as
low as possible. As will be shown below, it is impossible to satisfy all these criteria with a

108
single solvent and the use of binary, ternary, and even quaternary mixtures is quite
common.

3.1.4.1 Carbonates
Propylene carbonate (PC) was the solvent of choice for primary lithium batteries,
but it was discovered early on that lithium cycles with very poor reversibility in PC-based
electrolytes [103] despite evidence that PC decomposes to form a protective SEI [104]. After
the transition to lithium-ion and graphite electrodes, a new problem with PC was
discovered: co-intercalation. Solvated Li
+
was found to intercalate into graphite electrodes
and the resulting expansion causes severe exfoliation. Despite this drawback, the first
lithium-ion cells commercialized by Sony used an electrolyte based on PC [105].
PC possesses a high dielectric constant suitable for dissolving salts, and furthermore
is liquid over a wide temperature range (~300 °C). Ethylene carbonate is similar in structure
to PC but has a high melting point (36 °C). The rigid, symmetrical structure that gives it
greater polarity than PC it also renders it solid at room temperature [106, 107]. Because of
the latter property, EC was largely ignored until it was shown that its melting point could be
depressed in the presence of cosolvents [108]. In 1990, Dahn and co-workers demonstrated
EC’s ability to form a stable, protective SEI on graphitic negative electrodes, which is not
seen in PC-based electrolytes [30], and EC remains the most important electrolyte
component for stable SEI formation in lithium-ion cells.

109
Despite its valuable properties, ethylene carbonate requires cosolvents to operate in
any useful temperature range. Linear carbonates possess very low viscosity and freezing
points but also very low dielectric constants, and are miscible with cyclic carbonates such as
EC and PC in any ratios. Thus, in the early 1990s, a cyclic/linear carbonate blend was
proposed that consisted of EC and DMC [109, 110]. It was found to be electrochemically
stable as well as sufficiently solvating of lithium salts. Soon after, other linear carbonates,
such as diethyl carbonate (DEC) [104, 111, 112] and ethyl methyl carbonate (EMC) [113],
were incorporated into similar electrolyte blends. Working toward a wide operating
temperature range system for spaceflight applications, researchers at the Jet Propulsion
Laboratory developed a ternary mixture of EC, DEC, and DMC [114]. Years later, this
formulation remains popular in the lithium-ion literature (i.e., [115, 116]).

3.1.4.2 Ethers and esters
Ethers were investigated for use in lithium secondary cells in the 1980s [117-120]
due to their low viscosity, and were known to promote more stable lithium cycling than PC
[121]. However, dendrite formation was still observed after prolonged cycling [122]. Ethers
were also investigated in lithium ion cells as potential cosolvents for EC during the early
1990s before linear carbonates were adopted [24, 75, 123] but were found to be susceptible
to oxidation at the positive electrode.

110
Esters, on the other hand, are more oxidatively stable. Furthermore, their low
melting points and low viscosities make them attractive choices for pairing with polar,
viscous solvents such as EC and PC for low-temperature applications. To this end, methyl
formate was blended into electrolytes in lithium primary cells [124, 125]. In the mid-to-late
1990s, simple esters such as methyl formate [126-128], methyl acetate [126, 129, 130], ethyl
acetate, and propyl acetate [130] were proposed in blends with EC and/or PC for low-
temperature applications.
Military and space applications in particular have more demanding temperature
requirements [131], and researchers at the Jet Propulsion Laboratory have proposed many
formulations based on blends of carbonates and esters. Quaternary solvent blends were
devised based on previously established ternary all-carbonate electrolytes. The low
molecular weight esters which had seen the most use, methyl and ethyl acetate, were
observed to cause continued SEI growth over long-term cycling owing to their reactivity
toward the graphite electrode [132]. In addition to contributing to wasteful irreversible
capacity, these surface films were observed to increase electrode impedance over time.
Heavier species, such as ethyl propionate and ethyl butyrate, on the other hand, form less
resistive films that more effectively protect the electrode [133]. Further investigations
examined promising formulations with high (80%) ester content for improved performance
below −40 °C [134]. Most recently, however, 20% EC + 60% EMC + 20% ester formulations

111
were demonstrated to provide good rate capability and discharge performance down to
−60 °C [135].
Electrolytes incorporating esters have been reported by workers at many battery
manufacturers as well, including Matsushita Industrial Co., Ltd. [136] (now Panasonic
Corporation), Saft Groupe S.A. [137, 138], Max Power Battery Ltd. [139], and Samsung
Display Devices Co. Ltd. [140].

3.1.4.3 Fluorinated solvents
Research into halogenated electrolyte solvents began with the goal of creating
formulation with better safety characteristics, because aliphatic carbonates, esters, and
ethers are extremely flammable. Halogenated organic compounds are usually considered to
be non-flammable or less flammable than their hydrogenated counterparts, depending on
the compound and the degree of substitution. They also tend to have lower melting points,
greater oxidative stability, and the ability to form stable films on carbon electrodes. One of
the first instances of using halogen-containing compounds was to combat graphite
exfoliation in the presence of PC. Chloroethylene carbonate (4-chloro-1,3-dioxolan-2-one)
was found to decompose on graphite around 1.7 V vs. Li/Li
+
, sufficiently passivating the
electrode and inhibiting PC co-intercalation, which normally occurs at 0.8 V vs Li/Li+
[141]. The authors proposed an optimum content of 30% chloro-EC in PC based on their
results. A subsequent publication described an EC-PC electrolyte with a lower chloro-EC

112
content, which was cycled in graphite/LiCoO
2
cells [142]. The cell had reasonably good
cycle life but low coulombic efficiency.
The fluorinated analogue, fluoroethylene carbonate, was also employed with the aim
of reducing irreversible capacity in the graphite-PC system [143]. Unlike the chloro-EC
containing formulation, cells containing EC+PC+FEC showed near 100% coulombic
efficiency over 200 cycles. Researchers from Kyoto University examined trifluoropropylene
carbonate, which was found to cause high irreversible capacity during graphite lithiation
and exfoliation due to cointercalation [144].
Testing was carried out at Sandia National Laboratory on several fluorinated
carbonates, including FEC, methyl 2,2,2-trifluoroethyl carbonate (TFEMC), and ethyl 2,2,2-
trifluoroethyl carbonate (ETFEC) [145]. Cyclic voltammetry showed ETFEC to be more
stable at negative potentials than DEC, and the conductivities of solutions containing 1.0 M
LiPF
6
in blends of these electrolytes were found to be acceptable. These electrolytes were
tested down to −40 °C.
Work has been ongoing at the Jet Propulsion Laboratory, in collaboration with the
University of Southern California, to create electrolytes that are non-flammable but still
capable of operating over wide temperature ranges. In 2003, JPL and USC researchers first
reported electrolytes incorporating a number of partially fluorinated organic carbonate
solvents in ternary and quaternary blends [146] including methyl-, ethyl-, propyl-, and di-

113
2,2,2-trifluoroethylcarbonate; methyl- and ethyl hexafluoroisopropyl carbonate; and
trifluoroethyl- and hexafluoroisopropyl N,N-dimethyl carbamate. Half-cell (MCMB carbon
vs. lithium metal) and full-cell (LiNiCoO
2
vs. MCMB carbon) tests showed improved
reversible capacity, improved low-temperature performance, more favorable lithium
intercalation kinetics, greater rate capability, better storage life, and better SEI formation as
compared to non-fluorinated baselines.

3.1.4.4 Additives
While novel solvents may provide interesting and useful battery electrolytes,
another approach is to improve the performance or safety of a proven electrolyte system
through the use of an additive. Additives for lithium-ion batteries usually consist of either
lithium salts or organic compounds and are designed to address a variety of limitations
commonly encountered in existing systems. Reactive additives take effect upon
decomposition. While such an additive may provide a desired effect at one electrode,
reaction at the other electrode often cannot be prevented and may be a source of irreversible
capacity; thus, the quantity of these additives is often limited to ~5% or less of the
electrolyte. Flame-retardant additives, on the other hand, provide greater safety benefits
through their presence in larger quantities, and are designed to be stable over the operating
voltage range of the cell.

114
One of the major classes of additives is those that improve the morphology and/or
the composition of the graphite SEI. The electrochemical reduction of carbonate solvents
goes through a radical anion intermediate, and two possible mechanisms have been
identified [31, 78, 147-151]. A fresh graphite surface catalyzes the first pathway, which
results in the formation of lithium carbonate and other inorganic products, and which
occurs at potentials positive to lithium intercalation. These reactions are accompanied by
the formation of gaseous byproducts, such as ethylene, which increase the diffusivity of the
SEI structure. If this reaction is suppressed through surface modification, the SEI formed by
the second pathway nearer the lithium intercalation potential is composed primarily of
carbon- and oxygen-containing organic products, which form a more stable structure
through coordination with lithium [31, 78, 147-151].
As such, film-forming additives that decompose at potentials higher than the
solvents can reduce the catalytic activity of the graphite surface and favor the more desirable
SEI formation mechanism at lower potentials. Polymerizable materials containing ethenyl
groups are among the more common in this category, and vinylene carbonate (VC) in
particular has been extensively studied [137, 152-158]. These compounds can undergo
electrochemical reductive polymerization through radical anion intermediates which are
terminated by reaction with solvent molecules. These reactions form fewer gaseous
byproducts and result in films composed of alkyl carbonates that are more compact and

115
more stable than those consisting primarily of Li
2
CO
3
, formed by solvent decomposition
alone during the first SEI formation step. Further solvent decomposition is effectively
suppressed until the electrode reaches lower potentials near those of lithium intercalation.
Some sulfur-containing compounds such as sulfides and sulfites [159] have a similar effect
by binding to active graphite surfaces that would otherwise catalyze solvent decomposition,
but these compounds tend to produce more undesirable side effects.
Some film-forming additives are not electrochemically reduced on the electrodes,
but rather react with the existing SEI products to form more stable films. These tend to be
species with similar functionality to the SEI, i.e. carbonyl and carboxyl groups. These
include carbon dioxide [77], ketones [160], and esters [161, 162]. Boron-oxygen compounds
have also been found to have an SEI-promoting effect, and lithium bis(oxalato)borate
(LiBOB) [163] has been particularly successful in this regard. The detection of boron on
electrode surfaces showed that LiBOB decomposition products were incorporated into the
SEI through reaction with existing species [164-166]. LiBOB was shown to improve
graphite cycle life in cells both as the primary lithium salt [167-173] and as an additive [171,
174-176].
The positive electrode is susceptible to damage during cycling, but most damage is
due to the presence of acidic impurities, which contribute to its dissolution. Overcharging a
cell can result in oxide decomposition and the release of O
2
, which in turn oxidizes organic

116
solvents and produces water. The presence of even trace water produces HF in LiPF
6
-
containing electrolytes (see §3.1.5.2). Therefore, additives have been conceived that
scavenge acids as well as water. The aforementioned LiBOB has also been used as a cathode
film-forming agent for manganese-containing oxide materials [177-179].
As previously mentioned, carbonate solvents are flammable, and lithium-ion
batteries must meet certain safety requirements for commercial use. The purpose of flame-
retardant additives (FRAs) is to combat this issue, while continuing to use proven solvent
systems known to afford good performance. The majority of FRAs that have been studied
over the last several decades are based on phosphorus or halogens, and much of the work
began with the goal of creating non-combustible polymers [180]. Halogenated solvents
were discussed in the previous section. Two modes of action have been proposed for
phosphorus FRAs. The first is the thermal barrier, a physical process by which the FRA
forms a layer of char that protects the uncombusted condensed phase from the heat
generated in the combusting gaseous phase [180]. The second is a chemical mechanism in
which the FRAs’ decomposition products act as radical scavengers that inhibit the
combustion chain in the gas phase [180, 181]:

RH R· + H· (3.4)
H· + O
2
HO· + O· (3.5)

117
HO· + H
2
H· + H
2
O (3.6)
O· + H
2
HO· + H· (3.7)

FRA
(l)
FRA
(g)
(3.8)
FRA
(g)
[P]· (3.9)
[P]· + H· [P]H (3.10)

Early work in lithium-ion battery FRAs focused on alkyl phosphates, such as
trimethyl and triethyl phosphate, but these were found to be unstable toward the charged
graphite electrodes [181-183]. This instability was found to be inversely related to the
length of the alkyl chain, while flame-retardant ability was found to follow the opposite
trend. Other avenues examined have included halogenated phosphates [184-186], aryl
phosphates [187-190], and phosphites [191, 192].
Trimethyl phosphate (TMP) was first proposed as a nonflammable bulk solvent or
cosolvent by Mitsubishi Chemical Corporation [181, 182] due to its low melting point and
known application as a flame retardant in plastics. Unfortunately, the solvent does not form
a stable anode SEI and extensive decomposition and co-intercalation on graphite electrodes
were observed in pure TMP+LiPF
6
electrolytes. TMP blends with carbonate, ester, or ether
co-solvents were evaluated for non-flammability on the basis of sample self-extinguishing

118
time (SET), and the necessary TMP content varied between 10% and 75% depending on the
co-solvent. Better SEI stability was observed with TMP/carbonate blends, but capacity
retention in full cells was poor when using both graphite and amorphous carbon anodes.
Mitsubishi later proposed the addition of ethylene ethyl phosphate (EEP) to suppress this
decomposition [183].
Investigators at the U.S. Army Research laboratory confirmed [193] that TMP as
well as triethyl phosphate (TEP) were unsuitable for incorporation into lithium-ion cell
electrolytes, and identified hexamethoxycyclotriphosphazene (HMPN) as a promising but
costly alternative previously proposed by others [194]. When incorporated into an LiPF
6
,
EC, EMC electrolyte, 30% HMPN reduced the capacity utilization of a carbon/LiCoO
2
cell
by 10% but had almost no capacity fade over 100 cycles. The SET of HMPN was longer than
either TMP or TEP, but still represented a drastic reduction versus the baseline.
Amine and coworkers at Argonne National Laboratory [187] examined tributyl
phosphate (TBP) and triphenyl phosphate (TPP), an aryl phosphate, which they
incorporated at 5% into LiPF
6
+EC+DEC electrolytes. As measured by accelerated rate
calorimetry, both additives reduced exothermicity and TPP raised the onset temperature to
above 200 °C. In addition, TPP-containing electrolytes demonstrated lower flame
propagation rates as well as comparable cycling performance to the baseline. Subsequently
others have studied electrolytes with TPP [195, 196] and other aryl phosphates [197-200].

119
Dimethyl methyl phosphonate (DMMP), added at 5-20%, was also found to reduce
the SET of LiPF
6
, EC, DEC solutions and improve the cycling efficiency of Li/LiCoO
2
half-
cells [201, 202]. Compatibility with graphite anodes, however, remains an issue [203].
Hexamethylphosphoramide (HMPA) [204, 205] added to LiPF
6
+EC+EMC electrolytes at
20% reduces the SET of the mixture but also decreases conductivity, resulting in lower
capacity in graphite/LiNiCoO
2
cells. HMPA is believed to reversibly form a stable complex
with the highly Lewic acidic PF
5
, preventing it from contributing to thermal decomposition
[205] .
As with solvents, halogenated additives have been investigated to improve the
thermal stability of electrolytes as well. Some chlorinated phosphates have been studied and
been found to reduce the flammability of carbonate-based electrolytes [186, 206].
Fluorinated phosphate additives have been the subject of much research, and tris(2,2,2-
trifluoroethyl) phosphate in particular [184, 185, 207, 208].
Thus far the compounds discussed have been P(V) compounds. P(III) compounds
such as phosphites [192, 199, 209] and fluorinated phosphites [191, 210] which in some
cases have shown superior electrochemical stability compared to the analogous P(V)
phosphates.

120
3.1.5 Lithium salts
3.1.5.1 Selection criteria
Lithium salts are in theory innumerable, but the many criteria necessary for
successful implementation in a lithium-ion cell leave only a handful that are acceptable [31].
The first criterion is solubility (i.e. a high dissociation constant); in spite of the use of cyclic
carbonates, overall the organic solvents employed are low dielectric. This eliminates many
of the simplest salts such as lithium halides and oxides, which possess “hard” counterions.
Solubility and good ionic mobility are necessary to obtain sufficient conductivity. Second,
the salt should not be reactive toward any component of the cell, including the solvents and
the electrode substrates. The anion should be resistant to oxidation at the cathode. It should
ideally be non-toxic and thermally stable.

3.1.5.2 LiPF
6

Lithium hexafluorophosphate [211] is by far the most commonly used lithium salt in
both commercial cells and in the scientific literature. It is, in essence, a compromise: it does
not have the highest conductivity, dissociation constant, ionic mobility, or thermal, anodic,
or chemical stability. It does, however, possess all of these properties to an acceptable degree
without a fatal weakness [31].
The combination of both decent ionic mobility and a decent dissociation constant
means that its conductivity is high, though not quite the highest among lithium salts

121
[212, 213]. It can resist oxidation up to potentials considerably higher than those to which
the positive electrode is normally charged. Its main weakness, however, is related to the fact
that the hexafluorophosphate anion exists in an equilibrium [214]:

LiPF
6
LiF + PF
5
(3.11)

Phosphorus pentafluoride is a highly reactive Lewis acid, and the equilibrium is
shifted to the right at high temperatures [215, 216]. As such, LiPF
6
has relatively poor
temperature stability when compared to some other salts. Furthermore, PF
5
(as well as PF
6
-
)
is susceptible to hydrolysis by even trace amounts of water:

PF
5
+ H
2
O POF
3
+ 2 HF (3.12)

Hydrofluoric acid contributes to the dissolution of positive electrode materials
(which is accompanied by the formation of more water), and lithium fluoride is insoluble in
organic media, effectively reducing the ionic concentration of lithium. It also reacts with SEI
components to form additional LiF, which is more resistive and increases electrode film
impedance [79, 217, 218].

122
3.1.5.3 Other Salts
Lithium tetrafluoroborate (LiBF
4
) has become more popular in recent years, because
it has better high-temperature stability than LiPF
6
and it is also less moisture sensitive. The
reason for a prior lack of interest is its inferior conductivity due to a smaller dissociation
constant compared to LiPF
6
and LiAsF
6
, despite the high mobility of the BF
4
-
anion
[212, 213].
Lithium hexafluoroarsenate (LiAsF
6
) [219] possesses many favorable characteristics,
including higher conductivity than LiPF
6
, good SEI formation, and both cathodic [220] and
anodic [112] stability. It is also much less prone to hydrolysis [79, 218, 221] than LiPF
6
, or
LiBF
4
, and thus ambient moisture is less of a concern. However, the use of even arsenic (V),
which has the potential to form toxic reduction products As(III) and As(0), is undesirable
for commercial use.
Lithium perchlorate (LiClO
4
) was employed in many reports in the literature early
on [222-225]. It has a higher decomposition temperature than LiPF
6
, high anodic stability,
and good conductivity. The formation of HF is not a concern as it is with LiPF
6
. Its strong
oxidizing power, however, makes it quite unstable and could be potentially dangerous for
human-rated applications [226, 227].
Lithium trifluoromethansulfonate (LiSO
3
CF
3
, or triflate) [228] is very stable toward
oxidation, moisture, and temperature, but has poor ionic dissociation and mobility [229].
Related compound lithium bis(trifluoromethanesulfonyl)imide (LiN(SO
2
CF
3
)
2
) was

123
introduced relatively late [75, 230] and possesses all the advantages of lithium triflate with
the added benefit of high conductivity. However, both of these salts proved unsuitable for
battery applications because their anions promoted the corrosion of the positive electrode’s
aluminum substrate. The positive electrode is charged to very high potentials and aluminum
dissolution is promoted by the strong complex of the oxidized Al
3+
with these anions [231].
Similar salts with longer fluorinated alkyl chains have been demonstrated in polyethylene
oxide electrolytes [232].

3.1.6 Scope of this work
To support future NASA applications, many of which are human-rated, there is a
desire to develop advanced electrolytes that have improved safety characteristics and the
ability to operate in high voltage systems. Examples of potential applications that can benefit
from such systems include:
• improved safety: for both aerospace and terrestrial applications, carbonate-based
electrolytes are flammable, and safety should be improved.
• high voltage stability: the energy demand for electronics is constantly increasing,
and both launch vehicles and electric vehicles place a premium on mass and volume.
There has been a push recently to increase power density and capacity by using
higher voltage positive electrode materials, and in such systems more robust
electrolytes are needed.

124
The current baseline formulation for the Space Power Systems program is 1.0 M
LiPF
6
in EC+DEC+DMC (1:1:1 v/v). The desire for improved safety characteristics led to
the creation of the JPL Generation I electrolyte, 1.0 M LiPF
6
in EC+EMC+TPPa+VC
(20:75:5:1.5), which incorporated small quantities of additives and maintained good
performance. Previous work at JPL in collaboration with USC has also included the
formulation of electrolytes for low-temperature applications, which incorporated esters
[233] and fluorinated esters [234] and (2,2,2-trifluoroethyl) butyrate in particular. In 2009,
they examined the impact of the incorporation of various flame-retardant additives
including TEP, TBP, TPP, and bis(2,2,2-trifluoroethyl) methylphosphonate [235]. They
found that the smaller alkyl groups led to better low temperature performance, which was
likely due to viscosity and its effect on conductivity. In contrast, larger alkyl groups led to
better cycle life, in agreement with electrochemical stability trends identified by others. TPP
in particular demonstrated the best capacity retention. A small impact from all the FRAs
was observed on the first cycle irreversible capacity, anode kinetics, and anode impedance of
the cells, which suggested a contribution to anode SEI formation.
In this work, we present the evaluation of novel electrolytes with a view to
addressing all of these criteria. To this end, we have studied the incorporation of higher
quantities of flame retardant additives, such as triphenyl phosphate (TPPa) and triphenyl
phosphite (TPPi); the incorporation of non-flammable fluorinated carbonate cosolvents,

125
such as monofluoroethylene carbonate (FEC), (2,2,2-trifluoroethyl) methyl carbonate
(TFEMC), and bis(2,2,2-trifluoroethyl) carbonate (bTFEC); the incorporation of film-
forming additives, such as vinylene carbonate (VC) and lithium bis(oxalato)borate (LiBOB);
and the electrochemical stability of these components. We have studied these electrolytes in
graphite and LiNiCoO
2
(NCO) and LiNiCoAlO
2
(NCA) cells with a primary goal of
identifying the safest electrolyte formulation that still satisfies cell performance
requirements. The structures of these and other components used are shown in Figure 3-2.

Figure 3-2. Structures of electrolyte components used in this study.

126
3.2 Experimental methods
3.2.1 Electrolyte preparation
Battery grade solvents and electrolyte blends with < 50 ppm water content were
obtained from Novolyte, Inc., including ethylene carbonate, diethyl carbonate, dimethyl
carbonate, ethyl methyl carbonate, methyl butyrate, and monofluoroethylene carbonate.
Battery grade lithium hexafluorophosphate and lithium bis(oxalato)borate were also
purchased from Novolyte. Triphenyl phosphate (>99%) was purchased from Sigma-Aldrich.
Triphenyl phosphite (99%) and vinylene carbonate (98%) were purchased from Acros
Chemicals. Other fluorinated carbonates including trifluoroethyl methyl carbonate and
bis(2,2,2-trifluoroethyl) carbonate were synthesized at the University of Southern California
and used as received. Novel electrolyte blends were prepared and stored in an Argon
glovebox.

3.2.2 Cell construction
Electrode materials were stored under vacuum at 70 °C and spiral-wound, three-
electrode cells were assembled in a dryroom. MCMB carbon electrodes and LiNiCoO
2

positive electrodes were purchased from Yardney Technical Products, Inc. Graphite (N7903
and N7905) and LiNi
x
Co
1-x
AlO
2
(P7983 and P7915) electrodes were purchased from Saft
America, Inc. Reference electrodes were prepared from lithium foil obtained from Foote
Mineral Company. Cathodes and anodes were cut to the desired sizes (1.75” x 6.25” and

127
1.75” x 7.125”, respectively) such that the cathode would be the limiting electrode. A small
amount of material was scraped off to expose a one-eighth-inch strip of bare substrate on
one of the short sides of each electrode, to which aluminum or copper tabs were welded, to
act as current collectors. Electrodes were separated by two layers of polypropylene with 20-
micron pores (Tonen-Setella, Inc.). Wound cells were encased in glass vials and enclosed in
O-ring sealed glass containers. The inner vials were flooded with the appropriate electrolyte
and the O-ring enclosures were sealed in an argon glovebox.

3.2.3 Electrochemical and electrical characterization
Cyclic voltammetry (CV) experiments were carried out in an O-ring sealed glass cell
with a platinum-flag working electrode with a geometric surface area of 1 cm
2
. The cell
included a large lithium foil counter electrode and a smaller lithium foil reference electrode,
each wrapped with porous polypropylene.
CV, dc micropolarization, Tafel polarization, and electrochemical impedance
spectroscopy (EIS) experiments were performed on a Princeton Applied Research 273
potentiostat or VersaSTAT MC potentiostat. Charge, discharge, and cycling experiments
were carried out on an Arbin BT-2040 cycler. For temperature-controlled experiments, the
cells were placed in Tenney environmental chambers (± 1 °C). Electrode potentials given
are referenced to the Li/Li
+
couple.

128
3.3 Results and discussion
3.3.1 Studies in MCMB/LiNiCoO
2
cells
3.3.1.1 Incorporation of flame-retardant additives
The incorporation of flame retardant additives is a common method for improving
the safety of lithium-ion batteries. Film-forming additives are commonly added to improve
the passivation of either of the electrodes. Unlike film-forming additives, which are
designed to decompose during charge in a controlled manner, FRAs are primarily intended
to act in the event of cell failure and fire. As a result, they should be electrochemically stable
over the range of operation of the cell. The safety advantages of an FRA should be weighed
against the impact it has on the performance of the cell, and it should not decompose
appreciably during normal operation.
Phosphorus-based compounds, as discussed earlier, are among the most promising
FRAs, and two in particular, triphenyl phosphate (TPPa) and triphenyl phosphite (TPPi),
were determined to be of interest (see Figure 3-1). In order to gauge their compatibility
with a common electrolyte system, i.e. LiPF
6
in ethylene carbonate (EC) and ethyl methyl
carbonate (EMC), the following formulations were studied by cyclic voltammetry over
various potential ranges:
1. 1.0 M LiPF
6
in EC+EMC (20:80 v/v) (baseline)
2. 1.0 M LiPF
6
in EC+EMC+TPPa (20:75:5 v/v)
3. 1.0 M LiPF
6
in EC+EMC+TPPa (20:75:5 v/v) + 1.5% VC

129
4. 1.0 M LiPF
6
in EC+EMC+TPPi (20:75:5 v/v)
Stability experiments over a typical first-generation lithium-ion battery potential
range are shown in Figure 3-3. The baseline electrolyte showed good oxidative stability up
to +4 V vs. Li/Li
+
; the onset of a small oxidation peak was observed near +3.4 V. The onset
of reductive decomposition was seen at approximately +0.40 V. The introduction of 5%
TPPa increased the observed currents slightly over the range in question, and the onset of
reductive decomposition was comparable. With both 5% TPPa and the further addition of
1.5% vinylene carbonate (a film-forming agent), the mixture showed stability better than the
baseline over the entire range, and the currents observed at the potential extremes were
reduced by an order of magnitude. TPPi, on the other hand, was incredibly electroactive
over the potential range, displaying numerous peaks and currents greater than the baseline
and TPPa-containing mixtures. Reductive decomposition in particular was observed at an
onset potential of +1.0 V.
In light of the general desire to design cells with high voltage cathodes for high-
power applications, the stability of these electrolytes was studied over several ranges. Figure
3-4 shows CV curves up to +5 V vs. Li/Li
+
. The baseline mixture and the one with 5% TPPa
were again quite similar in behavior at the high potential limit. The onset of oxidation again
occurs near +3.4 V and a further increase is seen at approximately +4.8 V. The TPPa/VC-
containing electrolyte again shows only mild oxidation beginning at +3.4 V, but a sharp

130
increase beginning at +4.6 V. Once again, the TPPi-containing electrolyte is oxidized much
more quickly and at lower potentials compared to the others.

Figure 3-3. Cyclic voltammetry of FRA-containing electrolytes on a platinum working electrode.
Recorded between 0.01 V – 4.00 V vs. Li/Li
+
, 5 mV s
-1
.

Stability was further examined up to +6 V (Figure 3-5). As seen in the previous
experiment, the baseline and the TPPa-containing electrolytes are comparable up to +5 V;
just positive of 5 V, the TPPa-containing electrolyte shows a sharp oxidation peak while the
baseline electrolyte shows steady but slower oxidation up to around +5.4 V, at which point
both electrolytes undergo rapid decomposition. The TPPa-containing electrolyte’s onset is
approximately 100 mV positive to that of the baseline, but its current at 6 V is more than
double. The oxidation of the TPPa and VC-containing electrolyte beginning near 4.8 V is
-8.00E-04!
-7.00E-04!
-6.00E-04!
-5.00E-04!
-4.00E-04!
-3.00E-04!
-2.00E-04!
-1.00E-04!
0.00E+00!
1.00E-04!
2.00E-04!
-1.00! 0.00! 1.00! 2.00! 3.00! 4.00! 5.00!
Current (A)!
Potential (V vs. Li/Li+)!
baseline (EC 20%, EMC 80%)!
EL002 (EC 20%, EMC 75%, TPPa 5%)!
EL003 (EC 20%, EMC 75%, TPPa 5%, VC 1.5%)!
EL004 (EC 20%, EMC 75%, TPPi 5%)!

131
observed again and increases to a peak around +5.5 V, decreasing slightly at higher
potentials. Curiously, the TPPi-containing electrolyte undergoes comparably little activity
compared to the other electrolytes in the 4-6 V range, seemingly contradicting the previous
results.

Figure 3-4. Cyclic voltammetry of FRA-containing electrolytes on a platinum working electrode.
Recorded between 1.00 V – 5.00 V vs. Li/Li
+
, 5 mV s
-1
.

Based on these results, it can be concluded that the impact of TPPa on the
electrochemical stability of EC/EMC based electrolytes is minimal in the 0.01 – 4 V range.
TPPa also offers comparable stability comparable to the baseline up to 5 V. Between 5 and 6
V vs. Li/Li+, the apparent decomposition of TPPa is greater than that of the baseline
electrolyte, and both are especially unstable above 5.4 V. The addition of VC is seen to
-1.00E-04!
-5.00E-05!
0.00E+00!
5.00E-05!
1.00E-04!
1.50E-04!
2.00E-04!
2.50E-04!
0.00! 1.00! 2.00! 3.00! 4.00! 5.00! 6.00!
Current (A)!
Potential (V vs. Li/Li+)!
baseline (EC 20%, EMC 80%)!
EL002 (EC 20%, EMC 75%, TPPa 5%)!
EL003 (EC 20%, EMC 75%, TPPa 5%, VC 1.5%)!
EL004 (EC 20%, EMC 75%, TPPi 5%)!

132
increase the stability of the electrolytes at lower potentials, but apparently decomposes at
potentials greater than 4.8 V. TPPi is observed to be unstable compared to other electrolytes
even in the 4 V range. Its apparent stability in 6 V experiments is unexplained.

Figure 3-5. Cyclic voltammetry of FRA-containing electrolytes on a platinum working electrode.
Recorded between 2.00 V – 6.00 V vs. Li/Li
+
, 5 mV s
-1
.

3.3.1.2 Incorporation of fluorinated carbonate co-solvents
The electrochemical stability of various fluorinated carbonate co-solvents was
examined using cyclic voltammetry. The fluorinated electrolytes were compared to a non-
fluorinated EC+EMC+TPPa baseline from the previous section. In the 4 V range (Figure 3-
6), the substitution of 20% TFEMC for a portion of the EMC results in similar behavior to
-1.00E-03!
-5.00E-04!
0.00E+00!
5.00E-04!
1.00E-03!
1.50E-03!
2.00E-03!
2.50E-03!
3.00E-03!
3.50E-03!
1.00! 2.00! 3.00! 4.00! 5.00! 6.00! 7.00!
Current (A)!
Potential (V vs. Li/Li+)!
baseline (EC 20%, EMC 80%)!
EL002 (EC 20%, EMC 75%, TPPa 5%)!
EL003 (EC 20%, EMC 75%, TPPa 5%, VC 1.5%)!
EL004 (EC 20%, EMC 75%, TPPi 5%)!

133
the baseline at the low potential region, but a large oxidation current beginning at 3.3 V vs.
Li/Li
+
. The oxidation of TFEMC is followed on the reverse scan by a reduction current of
similar magnitude, suggesting that some surface species are formed but are unstable. The
mixture with 20% bTFEC, on the other hand, behaves similarly to the baseline at the high
potential limit but shows several oxidation peaks on the forward scan at intermediate
potentials and the reductive decomposition current at the low potential limit is
approximately double that of the baseline. The complete substitution of FEC for EC resulted
in an electrolyte which was remarkably stable over the 0.01 – 4.00 V range, with currents an
order of magnitude smaller than the other electrolytes studied.
Figure 3-6. Cyclic voltammetry of fluorocarbonate-containing electrolytes on a platinum working
electrode. Recorded between 0.01 V – 4.00 V vs. Li/Li
+
, 20 mV s
-1
(ninth cycle shown).
-4.00E-04!
-3.00E-04!
-2.00E-04!
-1.00E-04!
0.00E+00!
1.00E-04!
2.00E-04!
3.00E-04!
4.00E-04!
5.00E-04!
-1.00! 0.00! 1.00! 2.00! 3.00! 4.00! 5.00!
Current (A)!
Potential (V vs. Li/Li+)!
EC 20%, EMC 75%, TPPa 5%!
EC 20%, EMC 50%, TFEMC 20%, TPPa 10%!
EC 20%, EMC 50%, bTFEC 20%, TPPa 10%!
FEC 20%, EMC 70%, TPPa 10%!

134
When the maximum potential was increased to 5.00 V (Figure 3-7), the oxidation of
TFEMC resolves into a broad peak centered at 3.6 V (with the onset at approximately 2.8
V), and the current increases again between 4.6 and 5.0 V. The reduction peak on the
reverse scan was broadened and shifted slightly negative compared to the 4 V experiment.
The 20% bTFEC-containing electrolyte shows onset of oxidation at the same point as
TFEMC, but the magnitude is less by an order of magnitude. Further oxidation of bTFEC
begins at 4.2 V and begins increasing rapidly at 4.8 V, but the current magnitude is
significantly less than TFEMC. FEC undergoes onset of oxidation near 4.3 V and increases
until 5 V, where the current is approximately half that of the oxidation of bTFEC. The
baseline, as seen above, remained relatively flat between 4.0 and 5.0 V.

Figure 3-7. Cyclic voltammetry of fluorocarbonate-containing electrolytes on a platinum working
electrode. Recorded between 1.00 V – 5.00 V vs. Li/Li
+
, 20 mV s
-1
.
-4.00E-04!
-3.00E-04!
-2.00E-04!
-1.00E-04!
0.00E+00!
1.00E-04!
2.00E-04!
3.00E-04!
4.00E-04!
5.00E-04!
0.00! 1.00! 2.00! 3.00! 4.00! 5.00! 6.00! Current (A)!
Potential (V vs. Li/Li+)!
EC 20%, EMC 75%, TPPa 5%!
EC 20%, EMC 50%, TFEMC 20%, TPPa 10%!
EC 20%, EMC 50%, bTFEC 20%, TPPa 10%!
FEC 20%, EMC 70%, 10% TPPa!

135
In the 2.00 to 6.00 V range (Figure 3-8), the oxidation of TFEMC remains relatively
flat compared to the maximum observed in the previous experiment. bTFEC, FEC, and the
baseline maintain lower currents than TFEMC until near 5 V. They all show oxidation
peaks with onset potentials of 4.8 V, 4.9 V, and 4.8 V, respectively. The baseline display a
sharp oxidation peak centered at 5.1 V and then oxidation increases dramatically between
5.5 and 6.0 V. FEC gives a weaker peak initally followed by a broader, larger peak at high
potential. The bTFEC-containing electrolyte had the most positive onset of oxidation and
the smallest current magnitude at the high potential limit. Unlike bTFEC, both FEC and
TFEMC showed hysteresis at the scan reversal around 6.0 V, undergoing significant
oxidation on the reverse scan between 6.0 and 5.25 V.

Figure 3-8. Cyclic voltammetry of fluorocarbonate-containing electrolytes on a platinum working
electrode. Recorded between 2.00 V – 6.00 V vs. Li/Li
+
, 20 mV s
-1
.
-1.00E-03!
-5.00E-04!
0.00E+00!
5.00E-04!
1.00E-03!
1.50E-03!
2.00E-03!
2.50E-03!
3.00E-03!
3.50E-03!
1.00! 2.00! 3.00! 4.00! 5.00! 6.00! 7.00!
Current (A)!
Potential (V vs. Li/Li+)!
EC 20%, EMC 75%, TPPa 5%!
EC 20%, EMC 50%, TFEMC 20%, TPPa 10%!
EC 20%, EMC 50%, bTFEC 20%, TPPa 10%!
FEC 20%, EMC 70%, TPPa 10%!

136
Based on these results, TFEMC is the least oxidatively stable among the fluorinated
carbonates studied, showing significant decomposition compared to the baseline even in the
4 V range. bTFEC is generally the most stable at high potentials (especially up to 6 V), but
undergoes unidentified oxidation reactions at comparatively negative potentials and is the
least stable to reduction near the lithium electrode potential. FEC shows remarkable
stability in the 0 – 4 V range, and outperformed the baseline up to 4.6 V in the 1 – 5 V
experiment, and even up to 6 V in the last experiment.

3.3.1.3 Reduced-flammability electrolytes in three-electrode cells
In order to understand the impact of the addition of these safety-enhancing
components on lithium-ion cells, several electrochemical techniques were employed to
evaluate their performance, including dc micropolarization, Tafel polarization, and
electrochemical impedance spectroscopy (EIS). Three-electrode cells were constructed and
filled with the following electrolytes:
1. 1.0 M LiPF
6
in FEC+EMC+TFEMC+TPPa (20:50:20:10 v/v)
2. 1.0 M LiPF
6
in FEC+EMC+TPPa (20:70:10 v/v)
3. 1.0 M LiPF
6
in EC+EMC+TFEMC+TPPa (20:50:20:10 v/v)
4. 1.0 M LiPF
6
in FEC+EMC+TFEMC+TPPa (20:50:20:10 v/v) +1.5% VC
5. 1.0 M LiPF
6
in EC+DEC+DMC (1:1:1 v/v)

137
A second series of cells was built to evaluate bTFEC:
6. 1.0 M LiPF
6
in EC+EMC (20:80 v/v)
7. 1.0 M LiPF
6
in EC+EMC+TPPa (20:70:10 v/v)
8. 1.0 M LiPF
6
in EC+EMC+bTFEC+TPPa (20:50:20:10 v/v)
9. 1.0 M LiPF
6
in EC+EMC+bTFEC+TPPa (20:30:40:10 v/v)

The electrolytes without fluorinated components or FRAs (5 and 6) served as
baselines.

dc micropolarization experiments
The individual electrodes were subjected to a potential sweep between +5 mV and
−5 mV vs. open circuit potential. The nearly linear behavior at very low overpotentials was
used to extract the polarization resistance of each electrode as the slope of the
current/voltage curve. This was used to evaluate the formation of surface films on the
electrodes and their effect on charge transfer. Additionally, the experiments were repeated
at progressively lower temperatures. The results are shown in Tables 3-1 and 3-2.
All the cells performed comparably at room temperature. Compared to the
EC+EMC baseline, the addition of TPPa decreased the resistance at the cathode at lower
temperatures, but increased the anode resistance at −30 °C and −40 °C. The complete
substitution of FEC for EC led to increased resistance at both electrodes over the range of

138
Table 3-1. Polarization resistance values measured in FEC- and TFEMC-containing MCMB/NCO cells.
Table 3-2. Polarization resistance values measured in bTFEC-containing MCMB/NCO cells.

139
temperatures. The partial substitution of 20% TFEMC for EMC both in the presence and
absence of VC gave inconsistent results, but the overall trend was of increased electrode
resistance, particularly at low temperatures. The electrolyte with both FEC and TFEMC
overall showed similar behavior to that with FEC as the only fluorinated component, but
with much higher resistance at the cathode at −40 °C. The substitution of 20% bTFEC for
EMC had relatively little impact on the cell performance, with somewhat reduced
polarization resistance at the anode at −40 °C. The inclusion of 40% bTFEC, on the other
hand, caused a marked increase in cathode polarization resistance compared to its non-
fluorinated counterpart, which was exacerbated at lower temperatures.
In summary, the use of 20% fluorinated carbonate resulted in an overall modest
increase in polarization resistance, with the exception of bTFEC. The use of 40% fluorinated
carbonate (whether one component or two) resulted in a large increase in resistance,
particularly at lower temperatures. This suggests that bTFEC is the most favorable
fluorinated carbonate solvent at lower temperatures, and that in general the use of
fluorinated solvents may produce less conductive electrode surface films. This could be due
to increased LiF formation.

Tafel polarization experiments
In order to evaluate the lithium intercalation/deintercalation kinetics at each
electrode, the cells were subjected to Tafel polarization experiments in which the electrode

140
potential was swept at a very slow rate at temperatures between 23 C and −40 C. The quasi-
steady-state measurements were conducted from +5 mV to +155 mV vs. open circuit
potential for the anode, and −5 mV to −155 mV vs. open circuit potential for the cathode.
Among the series of cells incorporating FEC and TFEMC (Figures 3-9 – 3-16), at
room temperature, the kinetics of lithium deintercalation were fastest in the baseline
EC+DEC+DMC cell, although the cell with 20% TFEMC was only slightly inferior. The
worst anode kinetics were observed in the cell with 40% fluorinated electrolyte and VC
(20% FEC + 20% TFEMC +1.5% VC). At the cathode, on the other hand, all the cells
containing fluorinated electrolytes outperformed the baseline slightly, and all showed nearly
equivalent lithium intercalation kinetics.
Figure 3-9. Tafel polarization of MCMB electrodes of FEC- and TFEMC-containing cells at room
temperature.
0.100!
0.120!
0.140!
0.160!
0.180!
0.200!
0.220!
0.240!
0.260!
0.280!
0.300!
0.010! 0.100! 1.000!
!"#$%&'#(%")*+&,-&./0&12312
+
)!
Current (A)!
FP01, 1.0 M LiPF6, FEC+EMC
+TFEMC+TPPa (20:50:20:10 v/v%)!
FP02, 1.0 M LiPF6, FEC+EMC+TPPa
(20:70:10 v/v%)!
FP03, 1.0 M LiPF6, EC+EMC
+TFEMC+TPPa (20:50:20:10 v/v%)!
FP04, 1.0 M LiPF6, FEC+EMC
+TFEMC+TPPa (20:50:20:10 v/v%)
+1.5% VC!
FP13, 1.0 M LiPF6, EC+DEC+DMC
(1:1:1 v/v%)!

141
Figure 3-10. Tafel polarization of LiNi
x
Co
1-x
O
2
electrodes of FEC- and TFEMC-containing cells at
room temperature.

Figure 3-11. Tafel polarization of MCMB electrodes of FEC- and TFEMC-containing cells at 0 °C.
4.000!
4.020!
4.040!
4.060!
4.080!
4.100!
4.120!
4.140!
4.160!
4.180!
4.200!
0.01! 0.10! 1.00!
Cathode potential (V vs. Li/Li
+
)!
Current (A)!
FP01, 1.0 M LiPF6, FEC+EMC+TFEMC+TPPa
(20:50:20:10 v/v%)!
FP02, 1.0 M LiPF6, FEC+EMC+TPPa (20:70:10 v/v
%)!
FP03, 1.0 M LiPF6, EC+EMC+TFEMC+TPPa
(20:50:20:10 v/v%)!
FP04, 1.0 M LiPF6, FEC+EMC+TFEMC+TPPa
(20:50:20:10 v/v%) +1.5% VC!
FP13, 1.0 M LiPF6, EC+DEC+DMC (1:1:1 v/v%)!
0.100!
0.120!
0.140!
0.160!
0.180!
0.200!
0.220!
0.240!
0.260!
0.280!
0.010! 0.100! 1.000!
Anode potential (V vs. Li/Li
+
)!
Current (A)!
FP01, 1.0 M LiPF6, FEC+EMC+TFEMC
+TPPa (20:50:20:10 v/v%)!
FP02, 1.0 M LiPF6, FEC+EMC+TPPa
(20:70:10 v/v%)!
FP03, 1.0 M LiPF6, EC+EMC+TFEMC
+TPPa (20:50:20:10 v/v%)!
FP04, 1.0 M LiPF6, FEC+EMC
+TFEMC+TPPa (20:50:20:10 v/v%)
+1.5% VC!
FP13, 1.0 M LiPF6, EC+DEC+DMC
(1:1:1 v/v%)!

142

Figure 3-12. Tafel polarization of LiNi
x
Co
1-x
O
2
electrodes of FEC- and TFEMC-containing cells at
0 °C.

Figure 3-13. Tafel polarization of MCMB electrodes of FEC- and TFEMC-containing cells at -20 °C.
4.000!
4.020!
4.040!
4.060!
4.080!
4.100!
4.120!
4.140!
4.160!
4.180!
0.001! 0.010! 0.100!
Cathode potential (V vs. Li/Li
+
)!
Current (A)!
FP01, 1.0 M LiPF6, FEC+EMC+TFEMC+TPPa
(20:50:20:10 v/v%)!
FP02, 1.0 M LiPF6, FEC+EMC+TPPa (20:70:10 v/v%)!
FP03, 1.0 M LiPF6, EC+EMC+TFEMC+TPPa
(20:50:20:10 v/v%)!
FP04, 1.0 M LiPF6, FEC+EMC+TFEMC+TPPa
(20:50:20:10 v/v%) +1.5% VC!
FP13, 1.0 M LiPF6, EC+DEC+DMC (1:1:1 v/v%)!
0.100!
0.120!
0.140!
0.160!
0.180!
0.200!
0.220!
0.240!
0.260!
0.280!
0.000! 0.001! 0.010! 0.100!
Anode potential (V vs. Li/Li
+
)!
Current (A)!
FP01, 1.0 M LiPF6, FEC+EMC+TFEMC+TPPa
(20:50:20:10 v/v%)!
FP02, 1.0 M LiPF6, FEC+EMC+TPPa (20:70:10 v/v%)!
FP03, 1.0 M LiPF6, EC+EMC+TFEMC+TPPa
(20:50:20:10 v/v%)!
FP04, 1.0 M LiPF6, FEC+EMC+TFEMC+TPPa
(20:50:20:10 v/v%) +1.5% VC!
FP13, 1.0 M LiPF6, EC+DEC+DMC (1:1:1 v/v%)!

143

Figure 3-14. Tafel polarization of LiNi
x
Co
1-x
O
2
electrodes of FEC- and TFEMC-containing cells
at -20 °C.

Figure 3-15. Tafel polarization of MCMB electrodes of FEC- and TFEMC-containing cells at -40 °C.
4.000!
4.020!
4.040!
4.060!
4.080!
4.100!
4.120!
4.140!
4.160!
4.180!
4.200!
0.00! 0.00! 0.01! 0.10!
Cathode potential (V vs. Li/Li
+
)!
Current (A)!
FP01, 1.0 M LiPF6, FEC+EMC+TFEMC+TPPa
(20:50:20:10 v/v%)!
FP02, 1.0 M LiPF6, FEC+EMC+TPPa (20:70:10 v/v
%)!
FP03, 1.0 M LiPF6, EC+EMC+TFEMC+TPPa
(20:50:20:10 v/v%)!
FP04, 1.0 M LiPF6, FEC+EMC+TFEMC+TPPa
(20:50:20:10 v/v%) +1.5% VC!
FP13, 1.0 M LiPF6, EC+DEC+DMC (1:1:1 v/v%)!
0.100!
0.120!
0.140!
0.160!
0.180!
0.200!
0.220!
0.240!
0.260!
0.280!
0.00001! 0.0001! 0.001! 0.01!
Anode potential (V vs. Li/Li
+
)!
Current (A)!
FP01, 1.0 M LiPF6, FEC+EMC+TFEMC+TPPa
(20:50:20:10 v/v%)!
FP02, 1.0 M LiPF6, FEC+EMC+TPPa (20:70:10 v/v%)!
FP03, 1.0 M LiPF6, EC+EMC+TFEMC+TPPa
(20:50:20:10 v/v%)!
FP04, 1.0 M LiPF6, FEC+EMC+TFEMC+TPPa
(20:50:20:10 v/v%) +1.5% VC!
FP13, 1.0 M LiPF6, EC+DEC+DMC (1:1:1 v/v%)!

144

Figure 3-16. Tafel polarization of LiNi
x
Co
1-x
O
2
electrodes of FEC- and TFEMC-containing cells at -
40 °C.

At lower temperatures, the same trend among anode kinetics was observed until
−40 °C, when the cell with 20% TFEMC outperformed the baseline. The cell with 40%
fluorinated content and VC again displayed the weakest kinetics. The cell with 40%
fluorinated content without VC and the cell with 20% FEC only showed very similar anode
behavior, both with inferior kinetics to 20% TFEMC only. The opposite trend was observed
at the cathode: as the temperature was lowered, the baseline showed improved kinetics over
the other cells. The cell with 20% FEC only consistently showed the worst cathode kinetics,
and the incorporation of VC appeared to be beneficial when compared to other fluorinated
electrolytes.
4.000!
4.020!
4.040!
4.060!
4.080!
4.100!
4.120!
4.140!
4.160!
4.180!
0.00001! 0.0001! 0.001! 0.01!
Cathode potential (V vs. Li/Li
+
)!
Current (A)!
FP01, 1.0 M LiPF6, FEC+EMC
+TFEMC+TPPa (20:50:20:10 v/v%)!
FP02, 1.0 M LiPF6, FEC+EMC
+TPPa (20:70:10 v/v%)!
FP03, 1.0 M LiPF6, EC+EMC
+TFEMC+TPPa (20:50:20:10 v/v%)!
FP04, 1.0 M LiPF6, FEC+EMC
+TFEMC+TPPa (20:50:20:10 v/v%)
+1.5% VC!
FP13, 1.0 M LiPF6, EC+DEC+DMC
(1:1:1 v/v%)!

145
Based on these results, electrolytes incorporating 20% TFEMC show similar kinetics
to the baseline under most conditions, with the exception of the cathode intercalation at -
40 °C. The incorporation of FEC was shown to be detrimental to the kinetics at both
electrodes over the range of temperatures examined; poor kinetics of 20% FEC + 20%
TFEMC electrolytes can be attributed primarily to the effects of FEC overwhelmingly. The
presence of VC led to inferior kinetics at the anode but superior cathode performance; it is
possible that VC undergoes reductive polymerization on the cathode during formation,
improving the passive film quality. The maximum currents observed in all cells at −40 °C
were approximately two orders of magnitude less than at room temperature.
In the second series of cells, the effects of TPPa and bTFEC were examined against
an EC+EMC baseline (Figures 3-17 – 3-24). The kinetics at the anode were reduced slightly
upon incorporation of TPPa at all temperatures. The incorporation of 20% and 40% bTFEC
was detrimental, with the 40% mixture performing the worst at all temperatures. The effect
of 20% bTFEC was progressively less at lower temperatures, however, and this mixture
displayed anode kinetics equivalent to the baseline at −20 °C. At the cathode, the addition of
TPPa was beneficial to lithium intercalation at all temperatures, suggesting that
phosphorus-containing species contribute to conductive film formation. The impact of 20%
bTFEC on cathode kinetics at room temperature is minimal, and this electrolyte actually

146

Figure 3-17. Tafel polarization of MCMB electrodes of bTFEC-containing cells at room
temperature.

Figure 3-18. Tafel polarization of LiNi
x
Co
1-x
O
2
electrodes of bTFEC-containing cells at room
temperature.
0.100!
0.120!
0.140!
0.160!
0.180!
0.200!
0.220!
0.240!
0.260!
0.280!
0.300!
0.010! 0.100! 1.000!
Anode potential (V vs. Li/Li
+
)!
Current (A)!
GB01, 1.0 M LiPF6, EC+EMC (20:80 v/v%)!
GB05, 1.0 M LiPF6,EC+EMC+TPPa (20:70:10 v/v%)!
GB03, 1.0 M LiPF6, EC+EMC+bTFEC+TPPa
(20:50:20:10 v/v%)!
GB04, 1.0 M LiPF6, EC+EMC+bTFEC+TPPa
(20:30:40:10 v/v%)!
4.000!
4.020!
4.040!
4.060!
4.080!
4.100!
4.120!
4.140!
4.160!
4.180!
4.200!
0.01! 0.10! 1.00!
Cathode potential (V vs. Li/Li
+
)!
Current (A)!
GB01, 1.0 M LiPF6, EC+EMC (20:80 v/v%)!
GB05, 1.0 M LiPF6,EC+EMC+TPPa (20:70:10
v/v%)!
GB03, 1.0 M LiPF6, EC+EMC+bTFEC+TPPa
(20:50:20:10 v/v%)!
GB04, 1.0 M LiPF6, EC+EMC+bTFEC+TPPa
(20:30:40:10 v/v%)!

147

Figure 3-19. Tafel polarization of MCMB electrodes of bTFEC-containing cells at 0 °C.

Figure 3-20. Tafel polarization of LiNi
x
Co
1-x
O
2
electrodes of bTFEC-containing cells at 0 °C.
0.100!
0.120!
0.140!
0.160!
0.180!
0.200!
0.220!
0.240!
0.260!
0.280!
0.010! 0.100! 1.000!
Anode potential (V vs. Li/Li
+
)!
Current (A)!
GB01, 1.0 M LiPF6, EC+EMC
(20:80 v/v%)!
GB05, 1.0 M LiPF6,EC+EMC
+TPPa (20:70:10 v/v%)!
GB03, 1.0 M LiPF6, EC+EMC
+bTFEC+TPPa (20:50:20:10 v/v%)!
GB04, 1.0 M LiPF6, EC+EMC
+bTFEC+TPPa (20:30:40:10 v/v%)!
4.000!
4.020!
4.040!
4.060!
4.080!
4.100!
4.120!
4.140!
4.160!
4.180!
4.200!
0.00! 0.01! 0.10! 1.00!
Cathode potential (V vs. Li/Li
+
)!
Current (A)!
GB01, 1.0 M LiPF6, EC+EMC (20:80 v/v%)!
GB05, 1.0 M LiPF6,EC+EMC+TPPa
(20:70:10 v/v%)!
GB03, 1.0 M LiPF6, EC+EMC+bTFEC
+TPPa (20:50:20:10 v/v%)!
GB04, 1.0 M LiPF6, EC+EMC+bTFEC
+TPPa (20:30:40:10 v/v%)!

148

Figure 3-21. Tafel polarization of MCMB electrodes of bTFEC-containing cells at -20 °C.

Figure 3-22. Tafel polarization of LiNi
x
Co
1-x
O
2
electrodes of bTFEC-containing cells at -20 °C.
0.000!
0.050!
0.100!
0.150!
0.200!
0.250!
0.300!
0.001! 0.010! 0.100!
Anode potential (V vs. Li/Li
+
)!
Current (A)!
GB01, 1.0 M LiPF6, EC+EMC (20:80 v/v%)!
GB05, 1.0 M LiPF6,EC+EMC+TPPa (20:70:10 v/v%)!
GB03, 1.0 M LiPF6, EC+EMC+bTFEC+TPPa
(20:50:20:10 v/v%)!
GB04, 1.0 M LiPF6, EC+EMC+bTFEC+TPPa
(20:30:40:10 v/v%)!
4.000!
4.020!
4.040!
4.060!
4.080!
4.100!
4.120!
4.140!
4.160!
4.180!
4.200!
0.0001! 0.001! 0.01! 0.1!
Cathode potential (V vs. Li/Li
+
)!
Current (A)!
GB01, 1.0 M LiPF6, EC+EMC (20:80 v/v%)!
GB05, 1.0 M LiPF6,EC+EMC+TPPa (20:70:10 v/v%)!
GB03, 1.0 M LiPF6, EC+EMC+bTFEC+TPPa
(20:50:20:10 v/v%)!
GB04, 1.0 M LiPF6, EC+EMC+bTFEC+TPPa
(20:30:40:10 v/v%)!

149

Figure 3-23. Tafel polarization of MCMB electrodes of bTFEC-containing cells at -40 °C. Stable
measurements were not obtained for cells not shown.

Figure 3-24. Tafel polarization of LiNi
x
Co
1-x
O
2
electrodes of bTFEC-containing cells at -40 °C.
Stable measurements were not obtained for cells not shown.
0.000!
0.050!
0.100!
0.150!
0.200!
0.250!
0.300!
0.000! 0.001! 0.010! 0.100!
Anode potential (V vs. Li/Li
+
)!
Current (A)!
GB01, 1.0 M LiPF6, EC+EMC (20:80 v/v%)!
GB05, 1.0 M LiPF6,EC+EMC+TPPa (20:70:10 v/v%)!
GB03, 1.0 M LiPF6, EC+EMC+bTFEC+TPPa
(20:50:20:10 v/v%)!
GB04, 1.0 M LiPF6, EC+EMC+bTFEC+TPPa
(20:30:40:10 v/v%)!
3.980!
4.000!
4.020!
4.040!
4.060!
4.080!
4.100!
4.120!
4.140!
4.160!
0.00! 0.00! 0.01! 0.10! 1.00!
Cathode potential (V vs. Li/Li
+
)!
Current (A)!
GB01, 1.0 M LiPF6, EC+EMC (20:80 v/v%)!
GB05, 1.0 M LiPF6,EC+EMC+TPPa
(20:70:10 v/v%)!
GB03, 1.0 M LiPF6, EC+EMC+bTFEC+TPPa
(20:50:20:10 v/v%)!
GB04, 1.0 M LiPF6, EC+EMC+bTFEC+TPPa
(20:30:40:10 v/v%)!

150
outperforms the non-fluorinated baseline at lower temperatures. The electrolyte containing
40% shows very poor cathode kinetics overall.
In summary, the negative impact of TPPa on anode kinetics is small but it shows a
significant benefit for the cathode. The addition of 20% bTFEC leads to improved kinetics at
low temperatures. While stable measurements were only obtained for the 20% bTFEC
electrolyte at −40 °C, its anode kinetics were better than those of the 20% TFEMC mixture
from the previous series, and its cathode performance was less than 10% inferior to the
EC+DEC+DMC baseline at −40 °C. It can be concluded that the inclusion of these
fluorinated linear carbonates may be beneficial at low temperatures.

Electrochemical impedance spectroscopy
Electrochemical impedance spectroscopy was used to evaluate the interfacial
properties of the electrodes in the presence of these electrolytes. The respective impedances
of the anodes and cathodes were measured independently against the lithium metal
reference electrodes. Impedance spectra of lithium-ion battery electrodes generally consist
of two inductive loops, one at high frequency (on the left on a Nyquist plot) and one at low
frequency (on the right). The high-frequency Z' intercept corresponds to the ohmic
resistance of the cell, which is the sum of the real resistances of the leads and wires, the
electrode-solution interface, and the solution resistance between the electrodes. The latter is
usually assumed to be the primary contributor. The high-frequency loop is associated with

151
the film impedance of the electrode and describes the ionic conductivity and capacitance of
the SEI. The low-frequency loop is associated with the charge transfer impedance of the
electrode and describes the faradaic process of lithium intercalation/deintercalation. The
diameters of these loops along the Z' axis are taken as the real resistance contributions of
each of these phenomena to the overall electrode impedance. In certain cases, a linear region
in the low-frequency domain is seen and is attributed to the mass-transfer resistance, or
diffusion of lithium within the electrode. By analyzing each electrode, its relative
contribution can be deduced with respect to the overall cell impedance, which previous
studies have shown to very closely match the sum of the individual electrode impedances
[236, 237].
The cells were tested at 23, 0, -20, and -40 °C (Figures 3-25 – 3-32, respectively). At
room temperature, the baseline EC+DEC+DMC formulation had the lowest anode film
resistance versus the cells with FEC and/or TFEMC. The film resistances of the cells with
20% FEC, 20% TFEMC, and 20% FEC + 20% TFEMC all had comparable film resistances;
this suggests the formation of some fluorine-containing species which contribute to a less
conductive SEI, possibly including LiF. As the temperature is lowered, the anode film
resistances of all the cells increase, but become more similar to one another and are
practically indistinguishable at and below - 20 °C. The anode charge transfer resistances,
which were difficult to distinguish at room temperature, become the dominant contributor

152
to electrode impedance as well as the distinguishing factor characterizing the various
electrolytes. The same trend is observed at all temperatures below RT: the cell
incorporating 20% TFEMC has an anode charge transfer resistance which is lower than (or
comparable to) the baseline, while that of the cell with 20% FEC is higher than the baseline.
The cell with 20% FEC + 20% TFEMC lies between the baseline and the cell with FEC only,
and the cell with VC gives the highest anode CT resistance of all. The magnitudes of these
anode charge transfer resistances increase by roughly an order of magnitude with each 20
°C reduction in temperature. This is generally consistent with the anode Tafel polarization
results in the last section.
At the cathode, the opposite trend is observed at room temperature as at the anode.
The ternary baseline had the largest cathode film resistance by far, and the smallest charge
transfer resistance. The cells with 20% FEC only and 20% FEC only had capacitive loops
with approximately equal contribution, and the other cells had very small film resistances.
The cell with FEC, TFEMC, and VC had the lowest overall cathode impedance. This is
consistent with the Tafel polarization results, in which it was among the better performing
cathodes. As the temperature is lowered, the trend starts to mirror that observed at the
anode: the film resistance contributions become dominated by the larger charge transfer
resistances. Still, at low temperatures the baseline and the cell with 20% TFEMC have larger
film resistances than the others, while the cell with FEC only has the smallest. In

153

Figure 3-25. EIS of MCMB electrodes of FEC- and TFEMC-containing cells at room temperature.

Figure 3-26. EIS of LiNi
x
Co
1-x
O
2
electrodes of FEC- and TFEMC-containing cells at room
temperature.
0.00!
0.10!
0.20!
0.30!
0.40!
0.50!
0.60!
0.00! 0.10! 0.20! 0.30! 0.40! 0.50! 0.60!
Z" (Ω)!
Z' (Ω)!
FP01, 1.0 M LiPF6, FEC+EMC+TFEMC+TPPa (20:50:20:10 v/v%)!
FP02, 1.0 M LiPF6, FEC+EMC+TPPa (20:70:10 v/v%)!
FP03, 1.0 M LiPF6, EC+EMC+TFEMC+TPPa (20:50:20:10 v/v%)!
FP04, 1.0 M LiPF6, FEC+EMC+TFEMC+TPPa (20:50:20:10 v/v%)
+1.5% VC!
FP13, 1.0 M LiPF6, EC+DEC+DMC (1:1:1 v/v%)!
-0.10!
0.00!
0.10!
0.20!
0.30!
0.40!
0.50!
0.00! 0.10! 0.20! 0.30! 0.40! 0.50! 0.60!
Z" (Ω)!
Z' (Ω)!
FP01, 1.0 M LiPF6, FEC+EMC+TFEMC+TPPa (20:50:20:10 v/v%)!
FP02, 1.0 M LiPF6, FEC+EMC+TPPa (20:70:10 v/v%)!
FP03, 1.0 M LiPF6, EC+EMC+TFEMC+TPPa (20:50:20:10 v/v%)!
FP04, 1.0 M LiPF6, FEC+EMC+TFEMC+TPPa (20:50:20:10 v/v%)
+1.5% VC!
FP13, 1.0 M LiPF6, EC+DEC+DMC (1:1:1 v/v%)!

154

Figure 3-27. EIS of MCMB electrodes of FEC- and TFEMC-containing cells at 0 °C.

Figure 3-28. EIS of LiNi
x
Co
1-x
O
2
electrodes of FEC- and TFEMC-containing cells at 0 °C
0.00!
0.20!
0.40!
0.60!
0.80!
1.00!
1.20!
1.40!
1.60!
1.80!
2.00!
0.00! 0.50! 1.00! 1.50! 2.00!
Z" (Ω)!
Z' (Ω)!
FP01, 1.0 M LiPF6, FEC+EMC+TFEMC+TPPa (20:50:20:10 v/v%)!
FP02, 1.0 M LiPF6, FEC+EMC+TPPa (20:70:10 v/v%)!
FP03, 1.0 M LiPF6, EC+EMC+TFEMC+TPPa (20:50:20:10 v/v%)!
FP04, 1.0 M LiPF6, FEC+EMC+TFEMC+TPPa (20:50:20:10 v/v%)
+1.5% VC!
FP13, 1.0 M LiPF6, EC+DEC+DMC (1:1:1 v/v%)!
0.00!
0.20!
0.40!
0.60!
0.80!
1.00!
1.20!
1.40!
0.00! 0.20! 0.40! 0.60! 0.80! 1.00! 1.20! 1.40!
Z" (Ω)!
Z' (Ω)!
FP01, 1.0 M LiPF6, FEC+EMC+TFEMC+TPPa (20:50:20:10 v/v
%)!
FP02, 1.0 M LiPF6, FEC+EMC+TPPa (20:70:10 v/v%)!
FP03, 1.0 M LiPF6, EC+EMC+TFEMC+TPPa (20:50:20:10 v/v
%)!
FP04, 1.0 M LiPF6, FEC+EMC+TFEMC+TPPa (20:50:20:10 v/
v%) +1.5% VC!
FP13, 1.0 M LiPF6, EC+DEC+DMC (1:1:1 v/v%)!

155

Figure 3-29. EIS of MCMB electrodes of FEC- and TFEMC-containing cells at -20 °C.

Figure 3-30. EIS of LiNi
x
Co
1-x
O
2
electrodes of FEC- and TFEMC-containing cells at -20 °C.
0.00!
2.00!
4.00!
6.00!
8.00!
10.00!
12.00!
14.00!
0.00! 2.00! 4.00! 6.00! 8.00! 10.00! 12.00! 14.00!
Z" (Ω)!
Z' (Ω)!
FP01, 1.0 M LiPF6, FEC+EMC+TFEMC+TPPa (20:50:20:10 v/v%)!
FP02, 1.0 M LiPF6, FEC+EMC+TPPa (20:70:10 v/v%)!
FP03, 1.0 M LiPF6, EC+EMC+TFEMC+TPPa (20:50:20:10 v/v%)!
FP04, 1.0 M LiPF6, FEC+EMC+TFEMC+TPPa (20:50:20:10 v/v%)
+1.5% VC!
FP13, 1.0 M LiPF6, EC+DEC+DMC (1:1:1 v/v%)!
0.00!
1.00!
2.00!
3.00!
4.00!
5.00!
6.00!
7.00!
8.00!
9.00!
10.00!
0.00! 2.00! 4.00! 6.00! 8.00! 10.00!
Z" (Ω)!
Z' (Ω)!
FP01, 1.0 M LiPF6, FEC+EMC+TFEMC
+TPPa (20:50:20:10 v/v%)!
FP02, 1.0 M LiPF6, FEC+EMC+TPPa
(20:70:10 v/v%)!
FP03, 1.0 M LiPF6, EC+EMC+TFEMC
+TPPa (20:50:20:10 v/v%)!
FP04, 1.0 M LiPF6, FEC+EMC+TFEMC
+TPPa (20:50:20:10 v/v%) +1.5% VC!
FP13, 1.0 M LiPF6, EC+DEC+DMC (1:1:1
v/v%)!
-0.10!
0.00!
0.10!
0.20!
0.30!
0.40!
0.50!
0.30! 0.40! 0.50! 0.60! 0.70!
Z" (Ω)!
Z' (Ω)!

156

Figure 3-31. EIS of MCMB electrodes of FEC- and TFEMC-containing cells at -40 °C.

Figure 3-32. EIS of LiNi
x
Co
1-x
O
2
electrodes of FEC- and TFEMC-containing cells at -40 °C. The
impedance of FP04 was apparently too high to obtain a successful measurement.
0.00!
10.00!
20.00!
30.00!
40.00!
50.00!
60.00!
70.00!
80.00!
90.00!
100.00!
0.00! 20.00! 40.00! 60.00! 80.00! 100.00!
Z" (Ω)!
Z' (Ω)!
FP01, 1.0 M LiPF6, FEC+EMC+TFEMC+TPPa
(20:50:20:10 v/v%)!
FP02, 1.0 M LiPF6, FEC+EMC+TPPa (20:70:10
v/v%)!
FP03, 1.0 M LiPF6, EC+EMC+TFEMC+TPPa
(20:50:20:10 v/v%)!
FP04, 1.0 M LiPF6, FEC+EMC+TFEMC+TPPa
(20:50:20:10 v/v%) +1.5% VC!
FP13, 1.0 M LiPF6, EC+DEC+DMC (1:1:1 v/v%)!
0.00!
5.00!
10.00!
15.00!
20.00!
25.00!
30.00!
35.00!
0.00! 5.00! 10.00! 15.00! 20.00! 25.00! 30.00! 35.00!
Z" (Ω)!
Z' (Ω)!
FP01, 1.0 M LiPF6, FEC+EMC+TFEMC+TPPa
(20:50:20:10 v/v%)!
FP02, 1.0 M LiPF6, FEC+EMC+TPPa (20:70:10 v/
v%)!
FP03, 1.0 M LiPF6, EC+EMC+TFEMC+TPPa
(20:50:20:10 v/v%)!
FP04, 1.0 M LiPF6, FEC+EMC+TFEMC+TPPa
(20:50:20:10 v/v%) +1.5% VC!
FP13, 1.0 M LiPF6, EC+DEC+DMC (1:1:1 v/v%)!
0.00!
0.10!
0.20!
0.30!
0.40!
0.50!
0.60!
0.70!
0.80!
0.70! 0.90! 1.10! 1.30! 1.50!
Z" (Ω)!
Z' (Ω)!

157
terms of charge transfer resistance, the baseline maintains the lowest value as well as the
lowest overall impedance through all temperatures, and the cell with FEC only has the
highest. Nominally, the anode and cathode contributions to the overall cell impedance
appear to be relatively equal.
The series of cells with TPPa and bTFEC was evaluated in the same way (Figures
3-33 – 3-40). At room temperature, the cell with EC and EMC without TPPa had lower
overall anode impedance than the others that included TPPa, which showed approximately
equal performance. The film resistance was the dominant contributor to anode impedance.
As the temperature is lowered, many of the same trends are observed as in the previous
series, i.e. an overall increase in charge transfer resistance. The anode impedance of the cell
with 40% bTFEC increased relative to the others at lower temperatures, while that with 20%
bTFEC maintained the lowest impedance. In agreement with previous results and contrary
to the anode behavior, the addition of TPPa proved beneficial to the cathode, and the EC +
EMC baseline had the highest cathode impedance at room temperature. At lower
temperatures, the same trends were observed as at the anode.
Based on these results, the charge transfer resistance was found to be the primary
contributor to increased electrode impedance at lower temperatures. This effect was also
seen previously in the Tafel polarization data. The ionic conductivity of the electrode
passivating films was not observed to be dramatically affected by low temperature. Between

158

Figure 3-33. EIS of MCMB electrodes of bTFEC-containing cells at room temperature.

Figure 3-34. EIS of LiNi
x
Co
1-x
O
2
electrodes of bTFEC-containing cells at room temperature.
0.00!
0.05!
0.10!
0.15!
0.20!
0.25!
0.30!
0.35!
0.40!
0.45!
0.50!
0.00! 0.10! 0.20! 0.30! 0.40! 0.50!
Z" (Ω)!
Z' (Ω)!
GB01, 1.0 M LiPF6, EC+EMC (20:80 v/v%)!
GB05, 1.0 M LiPF6,EC+EMC+TPPa (20:70:10 v/v%)!
GB03, 1.0 M LiPF6, EC+EMC+bTFEC+TPPa (20:50:20:10 v/v%)!
GB04, 1.0 M LiPF6, EC+EMC+bTFEC+TPPa (20:30:40:10 v/v%)!
0.00!
0.10!
0.20!
0.30!
0.40!
0.50!
0.60!
0.00! 0.10! 0.20! 0.30! 0.40! 0.50! 0.60!
Z" (Ω)!
Z' (Ω)!
GB01, 1.0 M LiPF6, EC+EMC (20:80 v/v%)!
GB05, 1.0 M LiPF6,EC+EMC+TPPa (20:70:10 v/v%)!
GB03, 1.0 M LiPF6, EC+EMC+bTFEC+TPPa (20:50:20:10 v/v%)!
GB04, 1.0 M LiPF6, EC+EMC+bTFEC+TPPa (20:30:40:10 v/v%)!

159

Figure 3-35. EIS of MCMB electrodes of bTFEC-containing cells at 0 °C.

Figure 3-36. EIS of LiNi
x
Co
1-x
O
2
electrodes of bTFEC-containing cells at 0 °C. Stable measurements
for GB04 were not obtained.
0.00!
0.20!
0.40!
0.60!
0.80!
1.00!
1.20!
1.40!
1.60!
0.00! 0.20! 0.40! 0.60! 0.80! 1.00! 1.20! 1.40! 1.60!
Z" (Ω)!
Z' (Ω)!
GB01, 1.0 M LiPF6, EC+EMC (20:80 v/v%)!
GB05, 1.0 M LiPF6,EC+EMC+TPPa (20:70:10 v/v%)!
GB03, 1.0 M LiPF6, EC+EMC+bTFEC+TPPa (20:50:20:10 v/v%)!
GB04, 1.0 M LiPF6, EC+EMC+bTFEC+TPPa (20:30:40:10 v/v%)!
0.00!
0.20!
0.40!
0.60!
0.80!
1.00!
1.20!
0.00! 0.20! 0.40! 0.60! 0.80! 1.00! 1.20!
Z" (Ω)!
Z' (Ω)!
GB01, 1.0 M LiPF6, EC+EMC (20:80 v/v%)!
GB05, 1.0 M LiPF6,EC+EMC+TPPa (20:70:10 v/v%)!
GB03, 1.0 M LiPF6, EC+EMC+bTFEC+TPPa (20:50:20:10 v/v%)!
GB04, 1.0 M LiPF6, EC+EMC+bTFEC+TPPa (20:30:40:10 v/v%)!

160

Figure 3-37. EIS of MCMB electrodes of bTFEC-containing cells at -20 °C.

Figure 3-38. EIS of LiNi
x
Co
1-x
O
2
electrodes of bTFEC-containing cells at -20 °C.
0.00!
1.00!
2.00!
3.00!
4.00!
5.00!
6.00!
0.00! 1.00! 2.00! 3.00! 4.00! 5.00! 6.00!
Z" (Ω)!
Z' (Ω)!
GB01, 1.0 M LiPF6, EC+EMC (20:80 v/v%)!
GB05, 1.0 M LiPF6,EC+EMC+TPPa (20:70:10 v/v%)!
GB03, 1.0 M LiPF6, EC+EMC+bTFEC+TPPa (20:50:20:10 v/v%)!
GB04, 1.0 M LiPF6, EC+EMC+bTFEC+TPPa (20:30:40:10 v/v%)!
0.00!
2.00!
4.00!
6.00!
8.00!
10.00!
12.00!
0.00! 2.00! 4.00! 6.00! 8.00! 10.00! 12.00!
Z" (Ω)!
Z' (Ω)!
GB01, 1.0 M LiPF6, EC+EMC (20:80 v/v%)!
GB05, 1.0 M LiPF6,EC+EMC+TPPa
(20:70:10 v/v%)!
GB03, 1.0 M LiPF6, EC+EMC+bTFEC
+TPPa (20:50:20:10 v/v%)!
GB04, 1.0 M LiPF6, EC+EMC+bTFEC
+TPPa (20:30:40:10 v/v%)!
-0.05!
0.00!
0.05!
0.10!
0.15!
0.20!
0.25!
0.30!
0.35!
0.20! 0.40! 0.60! 0.80! 1.00! 1.20!
Z" (Ω)!
Z' (Ω)!

161

Figure 3-39. EIS of MCMB electrodes of bTFEC-containing cells at -40 °C. A stable measurement
for GB01 was not obtained.

Figure 3-40. EIS of LiNi
x
Co
1-x
O
2
electrodes of bTFEC-containing cells at -40 °C. A stable
measurement for GB01 was not obtained
0.00!
5.00!
10.00!
15.00!
20.00!
25.00!
30.00!
35.00!
0.00! 5.00! 10.00! 15.00! 20.00! 25.00! 30.00! 35.00!
Z" (Ω)!
Z' (Ω)!
GB01, 1.0 M LiPF6, EC+EMC (20:80 v/v%)!
GB05, 1.0 M LiPF6,EC+EMC+TPPa (20:70:10 v/v%)!
GB03, 1.0 M LiPF6, EC+EMC+bTFEC+TPPa (20:50:20:10 v/v%)!
GB04, 1.0 M LiPF6, EC+EMC+bTFEC+TPPa (20:30:40:10 v/v%)!
0.00!
5.00!
10.00!
15.00!
20.00!
25.00!
30.00!
35.00!
0.00! 5.00! 10.00! 15.00! 20.00! 25.00! 30.00! 35.00!
Z" (Ω)!
Z' (Ω)!
GB01, 1.0 M LiPF6, EC+EMC (20:80 v/v%)!
GB05, 1.0 M LiPF6,EC+EMC+TPPa (20:70:10 v/v
%)!
GB03, 1.0 M LiPF6, EC+EMC+bTFEC+TPPa
(20:50:20:10 v/v%)!
GB04, 1.0 M LiPF6, EC+EMC+bTFEC+TPPa
(20:30:40:10 v/v%)!

162
room temperature and -40 °C, the film resistances of the electrodes increased by less than an
order of magnitude, while the charge transfer resistances increased by three orders of
magnitude. bTFEC in particular was observed to have lower impedance at low temperatures
than the EC + EMC baseline.

3.3.2 FEC and LiBOB in MPG-111/LiNiCoAlO
2
cells
3.3.2.1 Electrochemical characterization
Three-electrode cells were constructed from Toda graphite anodes and LiNiCoAlO2
cathodes obtained from Saft America, Inc and filled with the following electrolytes for
study:
1. 1.0 M LiPF
6
in EC+EMC+TPPa (20:70:10 v/v)
2. 1.0 M LiPF
6
in EC+EMC+TPPa (20:65:15 v/v)
3. 1.0 M LiPF
6
in FEC+EMC+TPPa (20:70:10 v/v)
4. 1.0 M LiPF
6
, 0.15 M LiBOB in EC+EMC+TPPa (20:70:10 v/v)
5. 1.0 M LiPF
6
in FEC+EMC+TPPa (20:65:15 v/v)
The first cell, containing EC, EMC, and 10% TPPa, serves as a baseline for
examining the effects of the addition of FEC, LiBOB, and higher TPPa content.

163
dc micropolarization experiments
The polarization resistances of each of the electrodes of fully charged cells were
measured by a linear sweep method at room temperature, 0, −20, and −40 °C. The results
are summarized in Tables 3-3 and 3-4. At the anode, increasing the TPPa component from
10% to 15% resulted in a modest increase in polarization resistance of the cells with EC and
EMC over the range of temperatures. In the cells with FEC and EMC, the cell with 15%
TPPa actually had less anode resistance than the cell with 10%. The substitution of FEC for
EC, on the other hand, caused the anode polarization resistance to increase much more
rapidly with decreasing temperature. The cell with EC, EMC and 10% TPPa had the lowest
anode resistance at all temperatures; the addition of LiBOB to this formulation had little
effect at room temperature, but caused the resistance to increase rapidly at lower
temperatures, at a similar rate to the cells with FEC.
At the cathode, the impact of increasing the TPPa content, substituting FEC for EC,
or adding LiBOB was relatively small and not entirely consistent between cells. A 50%
increase in resistance at −40 °C was seen with the addition of 0.15 M LiBOB, compared with
a 200% increase at the anode. Overall, the cathode polarization resistances increased much
more slowly as the temperature was reduced than did the anodes. The anodes were seen to
be the major contributors to polarization resistance at all temperatures.

164

1.0 M LiPF
6
,
EC+EMC+
TPPa
(20:70:10 v/v%)
1.0 M LiPF
6
,
EC+EMC
+TPPa
(20:65:15 v/v%)
1.0 M LiPF
6
,
FEC+EMC
+TPPa
(20:70:10 v/v%)
1.0 M LiPF
6
,
0.15 M LiBOB,
EC+EMC
+TPPa
(20:70:10 v/v%)
1.0 M LiPF
6
,
FEC+EMC
+TPPa
(20:65:15 v/v%)
23 °C 0.543 0.819 0.829 0.591 0.559
0 °C 1.099 1.370 1.852 2.160 1.873
-20 °C 5.038 6.345 12.180 12.920 10.384
-30 °C 14.129 18.404 41.390 36.630 29.155
-40 °C 47.038 61.761 127.735 153.339 106.230

Table 3-3. MPG-111 anode polarization resistance values recorded for FEC- and TPPa- containing
Graphite/NCA cells at full state-of-charge.

1.0 M LiPF
6
,
EC+EMC
+TPPa
(20:70:10 v/v%)
1.0 M LiPF
6
,
EC+EMC
+TPPa
(20:65:15 v/v%)
1.0 M LiPF
6
,
FEC+EMC+
TPPa
(20:70:10 v/v%)
1.0 M LiPF
6
,
0.15 M LiBOB,
EC+EMC
+TPPa
(20:70:10 v/v%)
1.0 M LiPF
6
,
FEC+EMC+
TPPa
(20:65:15 v/v%)
23 °C 0.174 0.215 0.253 0.403 0.273
0 °C 0.317 0.275 0.249 0.175 0.355
-20 °C 0.928 1.192 1.761 1.156 1.228
-30 °C 1.855 2.007 2.715 4.287 5.065
-40 °C 9.870 8.348 10.997 15.248 13.427

Table 3-4. LiNiCoAlO2 cathode polarization resistance values recorded for FEC- and TPPa-
containing Graphite/NCA cells at full state-of-charge.

Tafel polarization experiments
In order to evaluate the lithium intercalation/deintercalation kinetics at each
electrode, the cells were subjected to Tafel polarization experiments in which the electrode
potential was swept at a very slow rate at temperatures between 23 and −40 °C. The quasi-
steady-state measurements were conducted from +5 mV to +150 mV vs. open circuit

165
potential for the anode, and −5 mV to −150 mV vs. open circuit potential for the cathode
(Figures 3-41 − 3-50).
At room temperature, the kinetics of lithium deintercalation at the anode do not
follow any discernible trend: looking at the cells with EC and EMC, the one with 10% TPPa
content performs better than the one with 15% TPPa. Among the cells with FEC and EMC,
the trend is the opposite. Meanwhile, if one compares the cells with 10% TPPa content, the
cell with EC performs better than the cell with FEC; but looking at the cells with 15% TPPa,
the cell with FEC is the best performing of the series and the cell with EC is the worst.
Meanwhile, the cell with added LiBOB shows nearly identical behavior at the anode to its
counterpart without LiBOB.
At temperatures below ambient, a clearer trend began to emerge. The greatest
impact on anode kinetics was seen in the cell with LiBOB, whose current output was the
weakest over the temperature range. The baseline formulation performed the best at low
temperatures, and the cell with EC, EMC, and 15% TPPa was slightly inferior. The cells
with FEC had poorer kinetics at the anode than the EC-containing cells, and the same small
loss was seen in the case of 15% TPPa.
The cathodes were studied in the same manner. Both the addition of FEC and
increased TPPa content appeared to hinder the lithium intercalation kinetics: the two cells
with FEC were the worst performing at all temperatures. However, at −30 °C and −40 °C,

166

Figure 3-41. Tafel polarization of MPG-111 electrodes of FEC- and TPPA-containing cells (MPG-
111/NCA) at room temperature.

Figure 3-42. Tafel polarization of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and TPPA-containing cells
(MPG-111/NCA) at room temperature.
0.100!
0.120!
0.140!
0.160!
0.180!
0.200!
0.220!
0.240!
0.260!
0.280!
0.010! 0.100! 1.000!
Anode potential (V vs. Li/Li
+
)!
Current (A)!
ST13, 1.0 M LiPF6, EC+EMC
+TPPa (20:70:10 v/v%)!
ST14, 1.0 M LiPF6, EC+EMC
+TPPa (20:65:15 v/v%)!
ST15, 1.0 M LiPF6, FEC+EMC
+TPPa (20:70:10 v/v%)!
ST16, 1.0 M LiPF6, 0.15 M LiBOB,
EC+EMC+TPPa (20:70:10 v/v%)!
ST17, 1.0 M LiPF6, FEC+EMC
+TPPa (20:65:15 v/v%)!
4.000!
4.020!
4.040!
4.060!
4.080!
4.100!
4.120!
4.140!
4.160!
4.180!
4.200!
0.01! 0.10! 1.00!
Cathode potential (V vs. Li/Li
+
)!
Current (A)!
ST13, 1.0 M LiPF6, EC+EMC+TPPa
(20:70:10 v/v%)!
ST14, 1.0 M LiPF6, EC+EMC+TPPa
(20:65:15 v/v%)!
ST15, 1.0 M LiPF6, FEC+EMC+TPPa
(20:70:10 v/v%)!
ST16, 1.0 M LiPF6, 0.15 M LiBOB, EC+EMC
+TPPa (20:70:10 v/v%)!
ST17, 1.0 M LiPF6, FEC+EMC+TPPa
(20:65:15 v/v%)!

167

Figure 3-43. Tafel polarization of MPG-111 electrodes of FEC- and TPPA-containing cells
(MPG-111/NCA) at 0 °C.

Figure 3-44. Tafel polarization of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and TPPA-containing cells
(MPG-111/NCA) at 0 °C.
0.000!
0.050!
0.100!
0.150!
0.200!
0.250!
0.300!
0.001! 0.010! 0.100! 1.000!
Anode potential (V vs. Li/Li
+
)!
Current (A)!
ST13, 1.0 M LiPF6, EC+EMC
+TPPa (20:70:10 v/v%)!
ST14, 1.0 M LiPF6, EC+EMC
+TPPa (20:65:15 v/v%)!
ST15, 1.0 M LiPF6, FEC+EMC
+TPPa (20:70:10 v/v%)!
ST16, 1.0 M LiPF6, 0.15 M LiBOB,
EC+EMC+TPPa (20:70:10 v/v%)!
ST17, 1.0 M LiPF6, FEC+EMC
+TPPa (20:65:15 v/v%)!
3.980!
4.000!
4.020!
4.040!
4.060!
4.080!
4.100!
4.120!
4.140!
4.160!
4.180!
0.01! 0.10! 1.00!
Cathode potential (V vs. Li/Li
+
)!
Current (A)!
ST13, 1.0 M LiPF6, EC+EMC+TPPa
(20:70:10 v/v%)!
ST14, 1.0 M LiPF6, EC+EMC+TPPa
(20:65:15 v/v%)!
ST15, 1.0 M LiPF6, FEC+EMC+TPPa
(20:70:10 v/v%)!
ST16, 1.0 M LiPF6, 0.15 M LiBOB, EC
+EMC+TPPa (20:70:10 v/v%)!
ST17, 1.0 M LiPF6, FEC+EMC+TPPa
(20:65:15 v/v%)!

168

Figure 3-45. Tafel polarization of MPG-111 electrodes of FEC- and TPPA-containing cells (MPG-
111/NCA) at -20 °C.

Figure 3-46. Tafel polarization of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and TPPA-containing cells
(MPG-111/NCA) at -20 °C.
0.000!
0.050!
0.100!
0.150!
0.200!
0.250!
0.300!
0.000! 0.001! 0.010! 0.100!
Anode potential (V vs. Li/Li
+
)!
Current (A)!
ST13, 1.0 M LiPF6, EC+EMC
+TPPa (20:70:10 v/v%)!
ST14, 1.0 M LiPF6, EC+EMC
+TPPa (20:65:15 v/v%)!
ST15, 1.0 M LiPF6, FEC+EMC
+TPPa (20:70:10 v/v%)!
ST16, 1.0 M LiPF6, 0.15 M LiBOB,
EC+EMC+TPPa (20:70:10 v/v%)!
ST17, 1.0 M LiPF6, FEC+EMC
+TPPa (20:65:15 v/v%)!
4.000!
4.020!
4.040!
4.060!
4.080!
4.100!
4.120!
4.140!
4.160!
4.180!
0.00! 0.00! 0.01! 0.10! 1.00!
Cathode potential (V vs. Li/Li
+
)!
Current (A)!
ST13, 1.0 M LiPF6, EC+EMC+TPPa
(20:70:10 v/v%)!
ST14, 1.0 M LiPF6, EC+EMC+TPPa
(20:65:15 v/v%)!
ST15, 1.0 M LiPF6, FEC+EMC
+TPPa (20:70:10 v/v%)!
ST16, 1.0 M LiPF6, 0.15 M LiBOB,
EC+EMC+TPPa (20:70:10 v/v%)!
ST17, 1.0 M LiPF6, FEC+EMC
+TPPa (20:65:15 v/v%)!

169

Figure 3-47. Tafel polarization of MPG-111 electrodes of FEC- and TPPA-containing cells
(MPG-111/NCA) at -30 °C.

Figure 3-48. Tafel polarization of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and TPPA-containing cells
(MPG-111/NCA) at -30 °C.
0.000!
0.050!
0.100!
0.150!
0.200!
0.250!
0.300!
0.000! 0.001! 0.010! 0.100! 1.000!
Anode potential (V vs. Li/Li
+
)!
Current (A)!
ST13, 1.0 M LiPF6, EC+EMC
+TPPa (20:70:10 v/v%)!
ST14, 1.0 M LiPF6, EC+EMC
+TPPa (20:65:15 v/v%)!
ST15, 1.0 M LiPF6, FEC+EMC
+TPPa (20:70:10 v/v%)!
ST16, 1.0 M LiPF6, 0.15 M LiBOB,
EC+EMC+TPPa (20:70:10 v/v%)!
ST17, 1.0 M LiPF6, FEC+EMC
+TPPa (20:65:15 v/v%)!
3.980!
4.000!
4.020!
4.040!
4.060!
4.080!
4.100!
4.120!
4.140!
4.160!
4.180!
0.00! 0.01! 0.10! 1.00!
Cathode potential (V vs. Li/Li
+
)!
Current (A)!
ST13, 1.0 M LiPF6, EC+EMC
+TPPa (20:70:10 v/v%)!
ST14, 1.0 M LiPF6, EC+EMC
+TPPa (20:65:15 v/v%)!
ST15, 1.0 M LiPF6, FEC
+EMC+TPPa (20:70:10 v/v%)!
ST16, 1.0 M LiPF6, 0.15 M
LiBOB, EC+EMC+TPPa
(20:70:10 v/v%)!
ST17, 1.0 M LiPF6, FEC
+EMC+TPPa (20:65:15 v/v%)!

170

Figure 3-49. Tafel polarization of MPG-111 electrodes of FEC- and TPPA-containing cells
(MPG-111/NCA) at -40 °C.

Figure 3-50. Tafel polarization of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and TPPA-containing cells
(MPG-111/NCA) at -40 °C.
0.000!
0.050!
0.100!
0.150!
0.200!
0.250!
0.300!
0.000! 0.000! 0.001! 0.010!
Anode potential (V vs. Li/Li
+
)!
Current (A)!
ST13, 1.0 M LiPF6, EC+EMC
+TPPa (20:70:10 v/v%)!
ST14, 1.0 M LiPF6, EC+EMC
+TPPa (20:65:15 v/v%)!
ST15, 1.0 M LiPF6, FEC+EMC
+TPPa (20:70:10 v/v%)!
ST16, 1.0 M LiPF6, 0.15 M LiBOB,
EC+EMC+TPPa (20:70:10 v/v%)!
ST17, 1.0 M LiPF6, FEC+EMC
+TPPa (20:65:15 v/v%)!
4.000!
4.020!
4.040!
4.060!
4.080!
4.100!
4.120!
4.140!
4.160!
4.180!
4.200!
0.00! 0.00! 0.00! 0.01! 0.10! 1.00!
Cathode potential (V vs. Li/Li
+
)!
Current (A)!
ST13, 1.0 M LiPF6, EC+EMC+TPPa
(20:70:10 v/v%)!
ST14, 1.0 M LiPF6, EC+EMC+TPPa
(20:65:15 v/v%)!
ST15, 1.0 M LiPF6, FEC+EMC+TPPa
(20:70:10 v/v%)!
ST16, 1.0 M LiPF6, 0.15 M LiBOB, EC
+EMC+TPPa (20:70:10 v/v%)!
ST17, 1.0 M LiPF6, FEC+EMC+TPPa
(20:65:15 v/v%)!

171
the cell with EC, EMC, and 15% TPPa outperformed the baseline. The cell with added
LiBOB had the best cathode kinetics at all temperatures except, puzzlingly, for −30 °C. This
point aside, this suggests that LiBOB contributes to favorable film formation at the cathode.
In summary, the relative effects of the substitution of FEC for EC on the anode and
the cathode are comparable and detrimental to the lithium intercalation and deintercalation
kinetics. The effect of increased TPPa concentration is likewise consistent but slight. The
effect of added LiBOB, on the other hand, is beneficial at the cathode but detrimental at the
anode. Because the cathodes were observed to sustain overall greater currents over the
course of these experiments, the net impact of LiBOB may be positive.

Electrochemical impedance spectroscopy
Electrochemical impedance spectroscopy was used to evaluate the interfacial
properties of the electrodes in the presence of these electrolytes (Figures 3-51 – 3-60). The
details of the spectra and their interpretation were explained in section 3.3.1.3. The spectra
observed at room temperature only show one capacitive relaxation loop, as opposed to the
two normally associated with lithium-ion battery electrodes. It is hypothesized that the high
and low frequency separation of the SEI and charge transfer impedances is not as distinct as
in previous studies and that the two contributions are convoluted. The broad, single loops
observed may suggest that the contributions to the overall impedance from each of these
components are approximately equal. At the anode, these overall impedances are very

172

Figure 3-51. EIS of MPG-111 electrodes of FEC- and TPPa-containing cells (MPG-111/NCA) at
room temperature.

Figure 3-52. EIS of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and TPPa-containing cells (MPG-111/NCA)
at room temperature.
0.00!
0.10!
0.20!
0.30!
0.40!
0.50!
0.60!
0.00! 0.10! 0.20! 0.30! 0.40! 0.50!
Z" (Ω)!
Z' (Ω)!
ST13, 1.0 M LiPF6, EC+EMC+TPPa (20:70:10 v/v%)!
ST14, 1.0 M LiPF6, EC+EMC+TPPa (20:65:15 v/v%)!
ST15, 1.0 M LiPF6, FEC+EMC+TPPa (20:70:10 v/v%)!
ST16, 1.0 M LiPF6, 0.15 M LiBOB, EC+EMC+TPPa (20:70:10 v/v%)!
ST17, 1.0 M LiPF6, FEC+EMC+TPPa (20:65:15 v/v%)!
0.00!
0.10!
0.20!
0.30!
0.40!
0.50!
0.60!
0.00! 0.05! 0.10! 0.15! 0.20! 0.25! 0.30! 0.35! 0.40!
Z" (Ω)!
Z' (Ω)!
ST13, 1.0 M LiPF6, EC+EMC+TPPa (20:70:10 v/v%)!
ST14, 1.0 M LiPF6, EC+EMC+TPPa (20:65:15 v/v%)!
ST15, 1.0 M LiPF6, FEC+EMC+TPPa (20:70:10 v/v%)!
ST16, 1.0 M LiPF6, 0.15 M LiBOB, EC+EMC+TPPa (20:70:10 v/v%)!
ST17, 1.0 M LiPF6, FEC+EMC+TPPa (20:65:15 v/v%)!

173

Figure 3-53. EIS of MPG-111 electrodes of FEC- and TPPa-containing cells (MPG-111/NCA), 0 °C.

Figure 3-54. EIS of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and TPPa-containing cells (MPG-111/NCA),
0 °C.
0.00!
0.20!
0.40!
0.60!
0.80!
1.00!
1.20!
1.40!
1.60!
1.80!
0.00! 0.50! 1.00! 1.50! 2.00!
Z" (Ω)!
Z' (Ω)!
ST13, 1.0 M LiPF6, EC+EMC+TPPa (20:70:10 v/v%)!
ST14, 1.0 M LiPF6, EC+EMC+TPPa (20:65:15 v/v%)!
ST15, 1.0 M LiPF6, FEC+EMC+TPPa (20:70:10 v/v%)!
ST16, 1.0 M LiPF6, 0.15 M LiBOB, EC+EMC+TPPa
(20:70:10 v/v%)!
ST17, 1.0 M LiPF6, FEC+EMC+TPPa (20:65:15 v/v%)!
0.00!
0.05!
0.10!
0.15!
0.20!
0.25!
0.30!
0.35!
0.40!
0.00! 0.05! 0.10! 0.15! 0.20! 0.25! 0.30! 0.35! 0.40!
Z" (Ω)!
Z' (Ω)!
ST13, 1.0 M LiPF6, EC+EMC+TPPa (20:70:10 v/v%)!
ST14, 1.0 M LiPF6, EC+EMC+TPPa (20:65:15 v/v%)!
ST15, 1.0 M LiPF6, FEC+EMC+TPPa (20:70:10 v/v%)!
ST16, 1.0 M LiPF6, 0.15 M LiBOB, EC+EMC+TPPa
(20:70:10 v/v%)!
ST17, 1.0 M LiPF6, FEC+EMC+TPPa (20:65:15 v/v%)!

174

Figure 3-55. EIS of MPG electrodes of FEC- and TPPa-containing cells (MPG-111/NCA)at -20 °C.

Figure 3-56. EIS of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and TPPa-containing cells (MPG-111/NCA)
at -20 °C.
0.00!
2.00!
4.00!
6.00!
8.00!
10.00!
12.00!
0.00! 2.00! 4.00! 6.00! 8.00! 10.00! 12.00!
Z" (Ω)!
Z' (Ω)!
ST13, 1.0 M LiPF6, EC+EMC+TPPa (20:70:10 v/
v%)!
ST14, 1.0 M LiPF6, EC+EMC+TPPa (20:65:15 v/
v%)!
ST15, 1.0 M LiPF6, FEC+EMC+TPPa (20:70:10
v/v%)!
ST16, 1.0 M LiPF6, 0.15 M LiBOB, EC+EMC
+TPPa (20:70:10 v/v%)!
ST17, 1.0 M LiPF6, FEC+EMC+TPPa (20:65:15
v/v%)!
0.00!
0.20!
0.40!
0.60!
0.80!
1.00!
1.20!
1.40!
1.60!
1.80!
2.00!
0.00! 0.50! 1.00! 1.50! 2.00!
Z" (Ω)!
Z' (Ω)!
ST13, 1.0 M LiPF6, EC+EMC+TPPa (20:70:10 v/v%)!
ST14, 1.0 M LiPF6, EC+EMC+TPPa (20:65:15 v/v%)!
ST15, 1.0 M LiPF6, FEC+EMC+TPPa (20:70:10 v/v%)!
ST16, 1.0 M LiPF6, 0.15 M LiBOB, EC+EMC+TPPa
(20:70:10 v/v%)!
ST17, 1.0 M LiPF6, FEC+EMC+TPPa (20:65:15 v/v%)!

175

Figure 3-57. EIS of MPG electrodes of FEC- and TPPa-containing cells (MPG-111/NCA), -30 °C.

Figure 3-58. EIS of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and TPPa-containing cells (MPG-111/NCA)
at -30 °C.
0.00!
5.00!
10.00!
15.00!
20.00!
25.00!
30.00!
35.00!
40.00!
0.00! 5.00! 10.00! 15.00! 20.00! 25.00! 30.00! 35.00! 40.00!
Z" (Ω)!
Z' (Ω)!
ST13, 1.0 M LiPF6, EC+EMC+TPPa (20:70:10 v/v%)!
ST14, 1.0 M LiPF6, EC+EMC+TPPa (20:65:15 v/v%)!
ST15, 1.0 M LiPF6, FEC+EMC+TPPa (20:70:10 v/v%)!
ST16, 1.0 M LiPF6, 0.15 M LiBOB, EC+EMC+TPPa
(20:70:10 v/v%)!
ST17, 1.0 M LiPF6, FEC+EMC+TPPa (20:65:15 v/v%)!
0.00!
0.50!
1.00!
1.50!
2.00!
2.50!
3.00!
3.50!
4.00!
4.50!
5.00!
0.00! 1.00! 2.00! 3.00! 4.00! 5.00!
Z" (Ω)!
Z' (Ω)!
ST13, 1.0 M LiPF6, EC+EMC+TPPa (20:70:10 v/v%)!
ST14, 1.0 M LiPF6, EC+EMC+TPPa (20:65:15 v/v%)!
ST15, 1.0 M LiPF6, FEC+EMC+TPPa (20:70:10 v/v%)!
ST16, 1.0 M LiPF6, 0.15 M LiBOB, EC+EMC+TPPa (20:70:10 v/v%)!
ST17, 1.0 M LiPF6, FEC+EMC+TPPa (20:65:15 v/v%)!

176

Figure 3-59. EIS of MPG electrodes of FEC- and TPPa-containing cells (MPG-111/NCA) at -40 °C.

Figure 3-60. EIS of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and TPPa-containing cells (MPG-111/NCA)
at -40 °C.
0.00!
20.00!
40.00!
60.00!
80.00!
100.00!
120.00!
0.00! 20.00! 40.00! 60.00! 80.00! 100.00! 120.00!
Z" (Ω)!
Z' (Ω)!
ST13, 1.0 M LiPF6, EC+EMC+TPPa (20:70:10 v/v%)!
ST14, 1.0 M LiPF6, EC+EMC+TPPa (20:65:15 v/v%)!
ST15, 1.0 M LiPF6, FEC+EMC+TPPa (20:70:10 v/v%)!
ST16, 1.0 M LiPF6, 0.15 M LiBOB, EC+EMC+TPPa (20:70:10 v/
v%)!
ST17, 1.0 M LiPF6, FEC+EMC+TPPa (20:65:15 v/v%)!
0.00!
2.00!
4.00!
6.00!
8.00!
10.00!
12.00!
14.00!
0.00! 2.00! 4.00! 6.00! 8.00! 10.00! 12.00! 14.00!
Z" (Ω)!
Z' (Ω)!
ST13, 1.0 M LiPF6, EC+EMC+TPPa (20:70:10 v/v%)!
ST14, 1.0 M LiPF6, EC+EMC+TPPa (20:65:15 v/v%)!
ST15, 1.0 M LiPF6, FEC+EMC+TPPa (20:70:10 v/v%)!
ST16, 1.0 M LiPF6, 0.15 M LiBOB, EC+EMC+TPPa (20:70:10 v/v%)!
ST17, 1.0 M LiPF6, FEC+EMC+TPPa (20:65:15 v/v%)!

177
comparable from cell to cell; the same is observed at the cathode, except that the cell with
LiBOB has a lower impedance than the others, in agreement with results obtained by
previous techniques. The anode impedance of the cells is greater than that of the cathode by
an order of magnitude.
Beginning at 0 °C, two loops may be seen in some of the anode impedance spectra,
lending credence to the theory that the two were overlapping in the room temperature
spectrum. The magnitude of the film resistance and its evolution compared to the room
temperature value are still difficult to gauge; however, when discernible, the charge transfer
resistance appears to be the larger of the two. The greatest anode impedance is seen in the
cell with 0.15 M LiBOB, which has the most clearly defined low-frequency loop. This
suggests that the charge transfer kinetics are worse at the anode in the presence of LiBOB,
which is consistent with the Tafel results above. The other overall trends from the Tafel
data are borne out as well, i.e. the moderate negative impact of FEC versus EC and the slight
negative impact of increased TPPa content. As the temperature is lowered further, the
charge transfer resistance becomes more dominant, and the distinction between the better
performing cells, i.e. those with EC, EMC, and TPPa only; and the worse performing cells,
i.e. those with FEC or LiBOB, becomes more pronounced.
At the cathode, the spectra still take the form of a single loop at 0 °C. The total
impedances of the cells are approximately equal, with that of the LiBOB-containing cell

178
being slightly less. Evidence of two loops does not appear convincingly until −30 °C,
suggesting that between room temperature and −20 °C, the film and charge transfer
resistances are growing at nearly equal rates. At low temperatures, the EC-containing cells
were again shown to have lower impedance than the FEC-containing cells. However,
contrary to previous findings, the presence of LiBOB increases the cathode charge transfer
resistance dramatically, whereas in the Tafel experiments it seemed to promote higher
currents. The total cathode impedance values are still an order of magnitude less than those
observed at the anode at each temperature.

3.3.2.2 Effect of FEC on discharge capacity
Based on the previous results, the use of FEC as a cosolvent was generally
detrimental to electrode kinetics compared to EC. However, given the ubiquitous nature of
EC as the most important electrolyte component for stable SEI formation, which enables
prolonged cycling, the fact that working cells can be constructed with non-flammable
substitute makes FEC worthy of further study. As such, the cells containing FEC were
subjected to discharge at various rates and temperatures and compared to those containing
EC. Figure 3-61 shows the discharge capacities of these cells at room temperatures, 0 °C,
and −20 °C recorded at C/20, C/10, C/5, and C/3.3. The capacities are expressed as
percentages of the discharge capacity of the final formation cycle, at C/65 and room
temperature.

179

Figure 3-61. Relative discharge capacities of FEC-containing cells versus EC-containing cells.

For both the 10% TPPa and the 15% TPPa cells, the relative capacity of the FEC-
containing cells is generally slightly less than that of the EC-containing cells, but they are for
the most part separated by less than five percentage points. In certain cases, the cells with
FEC retained a greater percentage of their room temperature capacity than did the cells with
EC.

180
3.3.3 MPG-111/LiNiCoAlO
2
cells with high additive content
Cells were constructed from MPG-111 anodes and LiNiCoAlO
2
cathodes from Saft
America, Inc. To build upon the work of the previous series of cells, and elucidate the effects
of higher concentrations of LiBOB and the role of EC and FEC, the cells were filled with the
following electrolytes:
1. 1.0 M LiPF
6
, 0.15 M LiBOB in EC+EMC+TPPa (20:65:15 v/v)
2. 1.0 M LiPF
6
, 0.15 M LiBOB in EC+FEC+EMC+TPPa (10:10:65:15 v/v)
3. 1.0 M LiPF
6
, 0.15 M LiBOB in FEC+EMC+TPPa (20:65:15 v/v)
4. 1.0 M LiPF
6
, 0.20 M LiBOB in EC+EMC+TPPa (20:65:15 v/v)
5. 1.0 M LiPF
6
, 0.25 M LiBOB in EC+EMC+TPPa (20:65:15 v/v)
6. 1.0 M LiPF
6
in EC+DEC+DMC (1:1:1 v/v)

3.3.3.1 Formation characteristics
The cells underwent a five-cycle formation, which consisted of 7.5 mA charge and
discharge to 4.10 V and 2.75 V, respectively (Table 3-5). During these first cycles, the
development of the solid electrolyte interphase (SEI) ensures the stability of the electrolyte
to the electrode surfaces during subsequent operation. The formation of the SEI contributes
to irreversible capacity during the first cycle: for example, ethylene carbonate reacts with
intercalated lithium in the anode during charging, essentially reducing the capacity available
during discharge. As such, a stable SEI should be indicated if the coulombic efficiency (the

181

Table 3-5. Formation characteristics of three-electrode MPG-111/LiNi
x
Co
1-x
AlO
2
cells.

182
percentage of the charge capacity that is reversible) approaches 100% by the end of the
formation. A large irreversible capacity over multiple cycles is evidence that the electrolyte
decomposition products do not form a passivating film. In addition to reducing the
discharge capacity during regular operation, continual loss of lithium and decomposition of
solvents will result in loss of conductivity and hazardous gas build-up.
During the first cycle, the cells had similar coulombic efficiencies ranging from 86-
88%, with the exception of the cell with 0.20 M LiBOB whose coulombic efficiency was
94%. The three cells with EC/EMC/TPPa and LiBOB had lower cumulative irreversible
capacities after all five cycles than the baseline. The two cells with FEC, including the one
with 10% EC, had higher cumulative irreversible capacities and lower fifth-cycle coulombic
efficiencies, indicating that the SEI formed by FEC may not be as stable and/or as
passivating as that formed by EC. The cell with 0.20 M LiBOB had the highest specific
reversible capacity (although it was less than 2% greater than that of the baseline), calculated
from the cathode active material mass and the fifth cycle discharge capacity. The specific
reversible capacities of the FEC-containing electrolytes were approximately 10% less, and
the other cells lay in between. The formation results are summarized in Table 3-5.

183
3.3.3.2 Electrochemical characterization
dc micropolarization experiments
The previous study suggested that the inclusion of LiBOB increases the polarization
resistance of the anode, especially at lower temperatures. This is found to be the case in this
series of cells as well (Table 3-6). The resistances of the cells with more LiBOB increased at
greater rates as the temperature was lowered (an erroneous value for the cell with 0.20 M
LiBOB at −20 °C is ignored). The cells with 0.15 M LiBOB and either 10% or 20% FEC
showed similar trends to the baseline cell with EC only and 0.15 M LiBOB, but with slightly
lower polarization resistance. The ternary baseline without LiBOB had consistently lower
resistance by an order of magnitude.
At the cathode, the polarization resistances are again several orders of magnitude
less than those at the anode (Table 3-7). LiBOB was previously seen to be beneficial at the
cathode, and in this series the impact of 0.20 M LiBOB is negligible compared to 0.15 M and
the negative impact of 0.25 M LiBOB is small but present. The performance of the cells with
EC and up to 0.20 M LiBOB was comparable to the baseline. The substitution of 10% FEC
for half the EC by volume has no effect on cell performance, but the complete substitution
of FEC results in high polarization resistance at −40 °C.

184

1.0 M LiPF
6
,
0.15 M
LiBOB,
EC+EMC
+TPPa
(20:65:15 v/v%)
1.0 M LiPF
6
,
0.15 M LiBOB,
EC+FEC+
EMC+TPPa
(10:10:65:15 v/v%)
1.0 M LiPF
6
,
0.15 M
LiBOB,
FEC+EMC+
TPPa
(20:65:15 v/v%)
1.0 M LiPF
6
,
0.20 M
LiBOB,
EC+EMC
+TPPa
(20:65:15 v/v%)
1.0 M LiPF
6
,
0.25 M LiBOB,
EC+EMC+
TPPa
(20:65:15 v/v%)

1.0 M LiPF
6

EC+DEC
+DMC
(1:1:1 v/v%)
23 °C 1.274 0.733 0.724 0.890 1.994 0.390
0 °C 3.145 2.066 2.252 4.550 5.342 1.159
-20 °C 14.881 8.696 12.937 34364 52.500 3.318
-30 °C 83.193 51.693 55.278 129.221 206.004 8.790
-40 °C 282.764 242.683 275.694 475.888 829.167 39.819

Table 3-6. MPG-111 anode polarization resistance values recorded for FEC- and LiBOB-containing
Graphite/NCA cells at full state-of-charge.

.
1.0 M LiPF
6
,
0.15 M LiBOB,
EC+EMC
+TPPa
(20:65:15 v/v%)
1.0 M LiPF
6
,
0.15 M LiBOB,
EC+FEC+
EMC+TPPa
(10:10:65:15
v/v%)
1.0 M LiPF
6
,
0.15 M LiBOB,
FEC+EMC+
TPPa
(20:65:15 v/v%)
1.0 M LiPF
6
,
0.20 M LiBOB,
EC+EMC
+TPPa
(20:65:15 v/v%)
1.0 M LiPF
6
,
0.25 M LiBOB,
EC+EMC
+TPPa
(20:65:15 v/v%)
1.0 M LiPF
6
,
EC+DEC+
DMC
(1:1:1 v/v%)
23 °C 0.193 0.592 0.201 0.144 0.688 0.477
0 °C 0.202 0.150 0.299 0.546 0.606 0.243
-20 °C 0.900 0.488 1.108 0.993 2.114 0.966
-30 °C 3.168 1.885 4.539 3.568 4.145 2.713
-40 °C 9.878 8.196 29.859 9.579 13.042 10.549

Table 3-7. LiNiCoAlO
2
cathode polarization resistance values recorded for FEC- and LiBOB-
containing graphite/NCA cells at full state-of-charge.

Tafel polarization experiments
Tafel measurements (Figure 3-62 – 3-71) show that the presence of LiBOB is
detrimental to anode kinetics, as expected. The baseline, therefore, shows the best
performance among the low temperatures. For the most part, the current sustained by the

185

Figure 3-62. Tafel polarization of MPG-111 electrodes of FEC- and LiBOB-containing cells
(MPG-111/NCA) at room temperature.

Figure 3-63. Tafel polarization of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and LiBOB-containing cells
(MPG-111/NCA) at room temperature.
0.000!
0.050!
0.100!
0.150!
0.200!
0.250!
0.300!
0.001! 0.010! 0.100! 1.000!
Anode potential (V vs. Li/Li
+
)!
Current (A)!
ST18, 1.0 M LiPF6, 0.15 M LiBOB, EC+EMC
+TPPa (20:65:15 v/v%)!
ST19, 1.0 M LiPF6, 0.15 M LiBOB, EC+FEC
+EMC+TPPa (10:10:65:15 v/v%)!
ST20, 1.0 M LiPF6, 0.15 M LiBOB, FEC+EMC
+TPPa (20:65:15 v/v%)!
ST21, 1.0 M LiPF6, 0.20 M LiBOB, EC+EMC
+TPPa (20:65:15 v/v%)!
ST22, 1.0 M LiPF6, 0.25 M LiBOB, EC+EMC
+TPPa (20:65:15 v/v%)!
ST23, 1.0 M LiPF6, EC+DEC+DMC (1:1:1 v/v%)!
3.980!
4.000!
4.020!
4.040!
4.060!
4.080!
4.100!
4.120!
4.140!
4.160!
4.180!
0.001! 0.010! 0.100! 1.000!
Cathode potential (V vs. Li/Li
+
)!
Current (A)!
ST18, 1.0 M LiPF6, 0.15 M LiBOB, EC
+EMC+TPPa (20:65:15 v/v%)!
ST19, 1.0 M LiPF6, 0.15 M LiBOB, EC
+FEC+EMC+TPPa (10:10:65:15 v/v%)!
ST20, 1.0 M LiPF6, 0.15 M LiBOB, FEC
+EMC+TPPa (20:65:15 v/v%)!
ST21, 1.0 M LiPF6, 0.20 M LiBOB, EC
+EMC+TPPa (20:65:15 v/v%)!
ST22, 1.0 M LiPF6, 0.25 M LiBOB, EC
+EMC+TPPa (20:65:15 v/v%)!
ST23, 1.0 M LiPF6, EC+DEC+DMC
(1:1:1 v/v%)!

186

Figure 3-64. Tafel polarization of MPG-111 electrodes of FEC- and LiBOB-containing cells
(MPG-111/NCA) at 0 °C.

Figure 3-65. Tafel polarization of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and LiBOB-containing cells
(MPG-111/NCA) at 0 °C.
0.000!
0.050!
0.100!
0.150!
0.200!
0.250!
0.300!
0.350!
0.001! 0.010! 0.100! 1.000!
Anode potential (V vs. Li/Li
+
)!
Current (A)!
ST18, 1.0 M LiPF6, 0.15 M LiBOB, EC+EMC+TPPa
(20:65:15 v/v%)!
ST19, 1.0 M LiPF6, 0.15 M LiBOB, EC+FEC+EMC
+TPPa (10:10:65:15 v/v%)!
ST20, 1.0 M LiPF6, 0.15 M LiBOB, FEC+EMC
+TPPa (20:65:15 v/v%)!
ST21, 1.0 M LiPF6, 0.20 M LiBOB, EC+EMC+TPPa
(20:65:15 v/v%)!
ST22, 1.0 M LiPF6, 0.25 M LiBOB, EC+EMC+TPPa
(20:65:15 v/v%)!
ST23, 1.0 M LiPF6, EC+DEC+DMC (1:1:1 v/v%)!
3.980!
4.000!
4.020!
4.040!
4.060!
4.080!
4.100!
4.120!
4.140!
4.160!
4.180!
0.001! 0.010! 0.100! 1.000!
Cathode potential (V vs. Li/Li
+
)!
Current (A)!
ST18, 1.0 M LiPF6, 0.15 M LiBOB, EC+EMC
+TPPa (20:65:15 v/v%)!
ST19, 1.0 M LiPF6, 0.15 M LiBOB, EC+FEC
+EMC+TPPa (10:10:65:15 v/v%)!
ST20, 1.0 M LiPF6, 0.15 M LiBOB, FEC+EMC
+TPPa (20:65:15 v/v%)!
ST21, 1.0 M LiPF6, 0.20 M LiBOB, EC+EMC
+TPPa (20:65:15 v/v%)!
ST22, 1.0 M LiPF6, 0.25 M LiBOB, EC+EMC
+TPPa (20:65:15 v/v%)!
ST23, 1.0 M LiPF6, EC+DEC+DMC (1:1:1 v/v%)!

187

Figure 3-66. Tafel polarization of MPG-111 electrodes of FEC- and LiBOB-containing cells
(MPG-111/NCA) at -20 °C.

Figure 3-67. Tafel polarization of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and LiBOB-containing cells
(MPG-111/NCA) at -20 °C.
0.000!
0.050!
0.100!
0.150!
0.200!
0.250!
0.300!
1.0E-09! 1.0E-08! 1.0E-07! 1.0E-06! 1.0E-05! 1.0E-04! 1.0E-03! 1.0E-02! 1.0E-01! 1.0E+00!
Anode potential (V vs. Li/Li
+
)!
Current (A)!
ST18, 1.0 M LiPF6, 0.15 M LiBOB,
EC+EMC+TPPa (20:65:15 v/v%)!
ST19, 1.0 M LiPF6, 0.15 M LiBOB,
EC+FEC+EMC+TPPa
(10:10:65:15 v/v%)!
ST20, 1.0 M LiPF6, 0.15 M LiBOB,
FEC+EMC+TPPa (20:65:15 v/v%)!
ST21, 1.0 M LiPF6, 0.20 M LiBOB,
EC+EMC+TPPa (20:65:15 v/v%)!
ST22, 1.0 M LiPF6, 0.25 M LiBOB,
EC+EMC+TPPa (20:65:15 v/v%)!
ST23, 1.0 M LiPF6, EC+DEC
+DMC (1:1:1 v/v%)!
3.950!
4.000!
4.050!
4.100!
4.150!
4.200!
0.001! 0.010! 0.100! 1.000!
Cathode potential (V vs. Li/Li
+
)!
Current (A)!
ST18, 1.0 M LiPF6, 0.15 M LiBOB, EC+EMC+TPPa
(20:65:15 v/v%)!
ST19, 1.0 M LiPF6, 0.15 M LiBOB, EC+FEC+EMC
+TPPa (10:10:65:15 v/v%)!
ST20, 1.0 M LiPF6, 0.15 M LiBOB, FEC+EMC
+TPPa (20:65:15 v/v%)!
ST21, 1.0 M LiPF6, 0.20 M LiBOB, EC+EMC+TPPa
(20:65:15 v/v%)!
ST22, 1.0 M LiPF6, 0.25 M LiBOB, EC+EMC+TPPa
(20:65:15 v/v%)!
ST23, 1.0 M LiPF6, EC+DEC+DMC (1:1:1 v/v%)!

188

Figure 3-68. Tafel polarization of MPG-111 electrodes of FEC- and LiBOB-containing cells
(MPG-111/NCA) at -30 °C.

Figure 3-69. Tafel polarization of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and LiBOB-containing cells
(MPG-111/NCA) at -30 °C.
0.000!
0.050!
0.100!
0.150!
0.200!
0.250!
0.300!
1.000E-05! 1.000E-04! 1.000E-03! 1.000E-02! 1.000E-01! 1.000E+00!
Anode potential (V vs. Li/Li
+
)!
Current (A)!
ST18, 1.0 M LiPF6, 0.15 M LiBOB,
EC+EMC+TPPa (20:65:15 v/v%)!
ST19, 1.0 M LiPF6, 0.15 M LiBOB,
EC+FEC+EMC+TPPa
(10:10:65:15 v/v%)!
ST20, 1.0 M LiPF6, 0.15 M LiBOB,
FEC+EMC+TPPa (20:65:15 v/v%)!
ST21, 1.0 M LiPF6, 0.20 M LiBOB,
EC+EMC+TPPa (20:65:15 v/v%)!
ST22, 1.0 M LiPF6, 0.25 M LiBOB,
EC+EMC+TPPa (20:65:15 v/v%)!
ST23, 1.0 M LiPF6, EC+DEC
+DMC (1:1:1 v/v%)!
3.950!
4.000!
4.050!
4.100!
4.150!
4.200!
0.000! 0.001! 0.010! 0.100! 1.000!
Cathode potential (V vs. Li/Li
+
)!
Current (A)!
ST18, 1.0 M LiPF6, 0.15 M LiBOB, EC+EMC+TPPa
(20:65:15 v/v%)!
ST19, 1.0 M LiPF6, 0.15 M LiBOB, EC+FEC+EMC
+TPPa (10:10:65:15 v/v%)!
ST20, 1.0 M LiPF6, 0.15 M LiBOB, FEC+EMC+TPPa
(20:65:15 v/v%)!
ST21, 1.0 M LiPF6, 0.20 M LiBOB, EC+EMC+TPPa
(20:65:15 v/v%)!
ST22, 1.0 M LiPF6, 0.25 M LiBOB, EC+EMC+TPPa
(20:65:15 v/v%)!
ST23, 1.0 M LiPF6, EC+DEC+DMC (1:1:1 v/v%)!

189

Figure 3-70. Tafel polarization of MPG-111 electrodes of FEC- and LiBOB-containing cells
(MPG-111/NCA) at -40 °C.

Figure 3-71. Tafel polarization of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and LiBOB-containing cells
(MPG-111/NCA) at -40 °C.
0.000!
0.050!
0.100!
0.150!
0.200!
0.250!
0.300!
1.000E-07! 1.000E-06! 1.000E-05! 1.000E-04! 1.000E-03! 1.000E-02! 1.000E-01! 1.000E+00!
Anode potential (V vs. Li/Li
+
)!
Current (A)!
ST18, 1.0 M LiPF6, 0.15 M LiBOB, EC
+EMC+TPPa (20:65:15 v/v%)!
ST19, 1.0 M LiPF6, 0.15 M LiBOB, EC
+FEC+EMC+TPPa (10:10:65:15 v/v%)!
ST20, 1.0 M LiPF6, 0.15 M LiBOB, FEC
+EMC+TPPa (20:65:15 v/v%)!
ST21, 1.0 M LiPF6, 0.20 M LiBOB, EC
+EMC+TPPa (20:65:15 v/v%)!
ST22, 1.0 M LiPF6, 0.25 M LiBOB, EC
+EMC+TPPa (20:65:15 v/v%)!
ST23, 1.0 M LiPF6, EC+DEC+DMC
(1:1:1 v/v%)!
3.980!
4.000!
4.020!
4.040!
4.060!
4.080!
4.100!
4.120!
4.140!
4.160!
4.180!
0.000! 0.001! 0.010! 0.100! 1.000!
Cathode potential (V vs. Li/Li
+
)!
Current (A)!
ST18, 1.0 M LiPF6, 0.15 M LiBOB, EC
+EMC+TPPa (20:65:15 v/v%)!
ST19, 1.0 M LiPF6, 0.15 M LiBOB, EC
+FEC+EMC+TPPa (10:10:65:15 v/v%)!
ST20, 1.0 M LiPF6, 0.15 M LiBOB, FEC
+EMC+TPPa (20:65:15 v/v%)!
ST21, 1.0 M LiPF6, 0.20 M LiBOB, EC
+EMC+TPPa (20:65:15 v/v%)!
ST22, 1.0 M LiPF6, 0.25 M LiBOB, EC
+EMC+TPPa (20:65:15 v/v%)!
ST23, 1.0 M LiPF6, EC+DEC+DMC
(1:1:1 v/v%)!

190
electrode decreased with increasing LiBOB content, and in all cases below room
temperature these showed the worst anode kinetics. As expected from the micropolarization
results, the presence of FEC alleviated this effect somewhat but these cells still performed
worse than the LiBOB-free baseline.
Most previous experiments have shown that LiBOB is beneficial to the cathode at
low temperatures, and in most cases the performance of these cells bears that out. The
baseline and the cell with 20% FEC show the poorest cathode kinetics, except for the data
points at room temperature and −20 °C in which the 0.20 M and 0.25 M LiBOB electrolytes
show very weak kinetics compared to the others in the series.

Electrochemical impedance spectroscopy
EIS measurements (Figures 3-72 – 3-81) at room temperature confirmed that the
presence of LiBOB is detrimental to the impedance at the anode. The anode impedance of
the baseline is an order of magnitude less than the others and manifests as a single loop.
Interestingly, among the three EC/EMC/TPPa cells, to the extent that the two loops can be
distinguished from one another, the impact of increasing LiBOB concentration is
manifested in a nearly linear increase in charge transfer resistance, while the film-forming
resistance remains constant. This suggests that LiBOB contributes to less favorable filming
at the anode, as evidenced by the larger high-frequency loop versus the baseline, but that the
concentration does not impact this behavior. This is consistent with the formation data

191

Figure 3-72. EIS of MPG-111 electrodes of FEC- and LiBOB-containing cells (MPG-111/NCA) at
room temperature.

Figure 3-73. EIS of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and LiBOB-containing cells (MPG-
111/NCA) at room temperature.
0.00!
0.10!
0.20!
0.30!
0.40!
0.50!
0.60!
0.00! 0.10! 0.20! 0.30! 0.40! 0.50! 0.60! 0.70! 0.80!
Z" (Ω)!
Z' (Ω)!
ST18, 1.0 M LiPF6, 0.15 M LiBOB, EC+EMC
+TPPa (20:65:15 v/v%)!
ST19, 1.0 M LiPF6, 0.15 M LiBOB, EC+FEC
+EMC+TPPa (10:10:65:15 v/v%)!
ST20, 1.0 M LiPF6, 0.15 M LiBOB, FEC+EMC
+TPPa (20:65:15 v/v%)!
ST21, 1.0 M LiPF6, 0.20 M LiBOB, EC+EMC
+TPPa (20:65:15 v/v%)!
ST22, 1.0 M LiPF6, 0.25 M LiBOB, EC+EMC
+TPPa (20:65:15 v/v%)!
ST23, 1.0 M LiPF6, EC+DEC+DMC (1:1:1 v/v%)!
0.00!
0.01!
0.01!
0.02!
0.02!
0.16! 0.17! 0.18!
0.00!
0.10!
0.20!
0.30!
0.40!
0.50!
0.60!
0.00! 0.05! 0.10! 0.15! 0.20! 0.25! 0.30! 0.35!
Z" (Ω)!
Z' (Ω)!
ST18, 1.0 M LiPF6, 0.15 M LiBOB, EC+EMC+TPPa (20:65:15 v/v%)!
ST19, 1.0 M LiPF6, 0.15 M LiBOB, EC+FEC+EMC+TPPa (10:10:65:15 v/v%)!
ST20, 1.0 M LiPF6, 0.15 M LiBOB, FEC+EMC+TPPa (20:65:15 v/v%)!
ST21, 1.0 M LiPF6, 0.20 M LiBOB, EC+EMC+TPPa (20:65:15 v/v%)!
ST22, 1.0 M LiPF6, 0.25 M LiBOB, EC+EMC+TPPa (20:65:15 v/v%)!
ST23, 1.0 M LiPF6, EC+DEC+DMC (1:1:1 v/v%)!

192

Figure 3-74. EIS of MPG electrodes of FEC- and LiBOB-containing cells (MPG-111/NCA) at 0 °C.

Figure 3-75. EIS of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and LiBOB-containing cells (MPG-
111/NCA) at 0 °C.
0.00!
0.50!
1.00!
1.50!
2.00!
2.50!
3.00!
3.50!
4.00!
4.50!
5.00!
0.00! 1.00! 2.00! 3.00! 4.00! 5.00!
Z" (Ω)!
Z' (Ω)!
ST18, 1.0 M LiPF6, 0.15 M LiBOB, EC+EMC+TPPa (20:65:15 v/v%)!
ST19, 1.0 M LiPF6, 0.15 M LiBOB, EC+FEC+EMC+TPPa (10:10:65:15 v/v%)!
ST20, 1.0 M LiPF6, 0.15 M LiBOB, FEC+EMC+TPPa (20:65:15 v/v%)!
ST21, 1.0 M LiPF6, 0.20 M LiBOB, EC+EMC+TPPa (20:65:15 v/v%)!
ST22, 1.0 M LiPF6, 0.25 M LiBOB, EC+EMC+TPPa (20:65:15 v/v%)!
ST23, 1.0 M LiPF6, EC+DEC+DMC (1:1:1 v/v%)!
0.00!
0.05!
0.10!
0.15!
0.20!
0.25!
0.30!
0.35!
0.40!
0.45!
0.00! 0.10! 0.20! 0.30! 0.40! 0.50!
Z" (Ω)!
Z' (Ω)!
ST18, 1.0 M LiPF6, 0.15 M LiBOB, EC+EMC+TPPa
(20:65:15 v/v%)!
ST19, 1.0 M LiPF6, 0.15 M LiBOB, EC+FEC+EMC
+TPPa (10:10:65:15 v/v%)!
ST20, 1.0 M LiPF6, 0.15 M LiBOB, FEC+EMC
+TPPa (20:65:15 v/v%)!
ST21, 1.0 M LiPF6, 0.20 M LiBOB, EC+EMC+TPPa
(20:65:15 v/v%)!
ST22, 1.0 M LiPF6, 0.25 M LiBOB, EC+EMC+TPPa
(20:65:15 v/v%)!
ST23, 1.0 M LiPF6, EC+DEC+DMC (1:1:1 v/v%)!

193

Figure 3-76. EIS of MPG electrodes of FEC- and LiBOB-containing cells (MPG-111/NCA) at -20 °C.

Figure 3-77. EIS of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and LiBOB-containing cells
(MPG-111/NCA) at -20 °C.
0.00!
5.00!
10.00!
15.00!
20.00!
25.00!
30.00!
35.00!
40.00!
45.00!
50.00!
0.00! 10.00! 20.00! 30.00! 40.00! 50.00!
Z" (Ω)!
Z' (Ω)!
ST18, 1.0 M LiPF6, 0.15 M LiBOB, EC+EMC+TPPa
(20:65:15 v/v%)!
ST19, 1.0 M LiPF6, 0.15 M LiBOB, EC+FEC+EMC
+TPPa (10:10:65:15 v/v%)!
ST20, 1.0 M LiPF6, 0.15 M LiBOB, FEC+EMC+TPPa
(20:65:15 v/v%)!
ST21, 1.0 M LiPF6, 0.20 M LiBOB, EC+EMC+TPPa
(20:65:15 v/v%)!
ST22, 1.0 M LiPF6, 0.25 M LiBOB, EC+EMC+TPPa
(20:65:15 v/v%)!
ST23, 1.0 M LiPF6, EC+DEC+DMC (1:1:1 v/v%)!
0.00!
0.20!
0.40!
0.60!
0.80!
1.00!
1.20!
0.00! 0.20! 0.40! 0.60! 0.80! 1.00! 1.20!
Z" (Ω)!
Z' (Ω)!
ST18, 1.0 M LiPF6, 0.15 M LiBOB, EC+EMC
+TPPa (20:65:15 v/v%)!
ST19, 1.0 M LiPF6, 0.15 M LiBOB, EC+FEC
+EMC+TPPa (10:10:65:15 v/v%)!
ST20, 1.0 M LiPF6, 0.15 M LiBOB, FEC+EMC
+TPPa (20:65:15 v/v%)!
ST21, 1.0 M LiPF6, 0.20 M LiBOB, EC+EMC
+TPPa (20:65:15 v/v%)!
ST22, 1.0 M LiPF6, 0.25 M LiBOB, EC+EMC
+TPPa (20:65:15 v/v%)!
ST23, 1.0 M LiPF6, EC+DEC+DMC (1:1:1 v/v%)!

194

Figure 3-78. EIS of MPG electrodes of FEC- and LiBOB-containing cells (MPG-111/NCA) at -30 °C.

Figure 3-79. EIS of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and LiBOB-containing cells
(MPG-111/NCA) at -30 °C.
0.00!
20.00!
40.00!
60.00!
80.00!
100.00!
120.00!
140.00!
160.00!
180.00!
0.00! 50.00! 100.00! 150.00! 200.00!
Z" (Ω)!
Z' (Ω)!
ST18, 1.0 M LiPF6, 0.15 M LiBOB, EC+EMC+TPPa
(20:65:15 v/v%)!
ST19, 1.0 M LiPF6, 0.15 M LiBOB, EC+FEC+EMC
+TPPa (10:10:65:15 v/v%)!
ST20, 1.0 M LiPF6, 0.15 M LiBOB, FEC+EMC
+TPPa (20:65:15 v/v%)!
ST21, 1.0 M LiPF6, 0.20 M LiBOB, EC+EMC+TPPa
(20:65:15 v/v%)!
ST22, 1.0 M LiPF6, 0.25 M LiBOB, EC+EMC+TPPa
(20:65:15 v/v%)!
ST23, 1.0 M LiPF6, EC+DEC+DMC (1:1:1 v/v%)!
0.00!
0.50!
1.00!
1.50!
2.00!
2.50!
3.00!
3.50!
4.00!
4.50!
5.00!
0.00! 1.00! 2.00! 3.00! 4.00! 5.00!
Z" (Ω)!
Z' (Ω)!
ST18, 1.0 M LiPF6, 0.15 M LiBOB, EC+EMC
+TPPa (20:65:15 v/v%)!
ST19, 1.0 M LiPF6, 0.15 M LiBOB, EC+FEC
+EMC+TPPa (10:10:65:15 v/v%)!
ST20, 1.0 M LiPF6, 0.15 M LiBOB, FEC+EMC
+TPPa (20:65:15 v/v%)!
ST21, 1.0 M LiPF6, 0.20 M LiBOB, EC+EMC
+TPPa (20:65:15 v/v%)!
ST22, 1.0 M LiPF6, 0.25 M LiBOB, EC+EMC
+TPPa (20:65:15 v/v%)!
ST23, 1.0 M LiPF6, EC+DEC+DMC (1:1:1 v/v%)!

195

Figure 3-80. EIS of MPG electrodes of FEC- and LiBOB-containing cells (MPG-111/NCA) at -40 °C.

Figure 3-81. EIS of LiNi
x
Co
1-x
AlO
2
electrodes of FEC- and LiBOB-containing cells
(MPG-111/NCA) at -40 °C.
0.00!
50.00!
100.00!
150.00!
200.00!
250.00!
0.00! 50.00! 100.00! 150.00! 200.00! 250.00!
Z" (Ω)!
Z' (Ω)!
ST18, 1.0 M LiPF6, 0.15 M LiBOB, EC
+EMC+TPPa (20:65:15 v/v%)!
ST19, 1.0 M LiPF6, 0.15 M LiBOB, EC
+FEC+EMC+TPPa (10:10:65:15 v/v%)!
ST20, 1.0 M LiPF6, 0.15 M LiBOB, FEC
+EMC+TPPa (20:65:15 v/v%)!
ST21, 1.0 M LiPF6, 0.20 M LiBOB, EC
+EMC+TPPa (20:65:15 v/v%)!
ST22, 1.0 M LiPF6, 0.25 M LiBOB, EC
+EMC+TPPa (20:65:15 v/v%)!
ST23, 1.0 M LiPF6, EC+DEC+DMC (1:1:1
v/v%)!
0.00!
2.00!
4.00!
6.00!
8.00!
10.00!
12.00!
14.00!
0.00! 2.00! 4.00! 6.00! 8.00! 10.00! 12.00! 14.00!
Z" (Ω)!
Z' (Ω)!
ST18, 1.0 M LiPF6, 0.15 M LiBOB, EC+EMC+TPPa (20:65:15 v/v%)!
ST19, 1.0 M LiPF6, 0.15 M LiBOB, EC+FEC+EMC+TPPa (10:10:65:15 v/v
%)!
ST20, 1.0 M LiPF6, 0.15 M LiBOB, FEC+EMC+TPPa (20:65:15 v/v%)!
ST21, 1.0 M LiPF6, 0.20 M LiBOB, EC+EMC+TPPa (20:65:15 v/v%)!
ST22, 1.0 M LiPF6, 0.25 M LiBOB, EC+EMC+TPPa (20:65:15 v/v%)!
ST23, 1.0 M LiPF6, EC+DEC+DMC (1:1:1 v/v%)!

196
shown above, in which little variation is seen in the first cycle irreversible capacities of the
cells, indicating that a small amount of LiBOB contributes to surface filming. The
relationship between LiBOB concentration and charge-transfer resistance, on the other
hand, may indicate an ionic interaction in which the more massive anion hinders the
intercalation/deintercalation process.
The cathode impedance spectra at room temperature take the form of single semi-
circles. Cells with EC/EMC/TPPa and LiBOB show higher impedance than the baseline,
while cells with FEC and LiBOB have lower or comparable impedance to the baseline. This
beneficial effect of FEC, and the steady increase of impedance with LiBOB concentration, is
similar to the effects seen earlier at the anode. The cathode impedances of the cells with
LiBOB are an order of magnitude smaller than the anode impedances, while their relative
contributions in the baseline cell are approximately equal.
The trend at the anode is the same at lower temperatures: the low-frequency loop
increasingly dominates the impedance of the electrode and the difference between the cells
becomes greater at lower temperature. Cathode behavior at lower temperatures is less
consistent, as the cell with 10% FEC routinely shows the lowest impedance while the cell
with 20% FEC shows the highest. The baseline, meanwhile, at first shows lower impedance
than the EC/EMC/TPPa cells with LiBOB but increases at a higher rate and displays

197
comparable behavior at −30 °C and −40 °C. The anode impedance continues to dominate
over the entire temperature range.
In summary, these electrochemical characterization experiments have confirmed the
detrimental impact of LiBOB on the performance of the anode, seen by the reduced
electrode kinetics in the polarization experiments and by the increased impedance in the EIS
experiments. The impedance data also showed a negative impact on the cathode with
increasing LiBOB concentration, which was not previously seen, but which is likely
overwhelmed by the effect on the anode in full cells. The addition of FEC at either 10% or
20% was seen to alleviate this impact somewhat at both electrodes. As the temperatures
were lowered, the impedance of the baseline increased more slowly than the others, and the
rate at which impedance increased in the other cells was greater for higher LiBOB content
and lesser with FEC.

3.3.3.3 Discharge characteristics
In order to evaluate performance at low temperatures, the cells were charged at
room temperature at 25 mA, and then discharged at various temperatures and rates.
Capacities of the cells are expressed as percentages of their respective capacities measured
during the final discharge of the formation in Table 3-8 (the formation values are collected
in Table 3-5).

198
The rate capabilities of the cells at room temperature are fairly comparable, and the
relative capacities fall between 73% and 79% at C/3.3 (the three cells with 0.15 M LiBOB
outperform the baseline at this rate). At C/20, the cells all perform nearly as well as the
baseline at both 0 °C and −20 °C. At 0 °C, as the rate is increased the baseline retains more
of its room temperature capacity compared to the LiBOB-containing cells. Among these, the
one with 20% FEC is the best performing at very high rate (C/3.3), although the cell with
10% FEC is the worst. At −20 C, the baseline’s performance is more on par with the rest of
the cells, and at the highest rate it is the worst performing in terms of relative capacity. The
cell with 20% FEC again delivers the most capacity (both in absolute and relative terms).
The results of the discharge experiments are summarized in Table 3-8.
Based on these results, the kinetic advantages afforded by FEC at low temperatures
that were identified in the electrochemical characterization experiments appear to be borne
out, especially at high rates. The LiBOB concentration, however, does not appear to impact
the discharge behavior in a significant way, although the baseline did outperform all the
LiBOB-containing cells by a considerable margin at 0 °C and C/3.3.

3.3.3.4 Rate characterization
In a true low-temperature application, cells will have to be resilient to charging as
well as discharging at these low temperatures. Charging at low temperature is avoided, if
possible, due to the risk of plating metallic lithium on the anode, especially at high rates.

199

Table 3-8. Discharge characteristics for FEC- and LiBOB-containing graphite/NCA cells.

200
More robust cell chemistries will ideally be able to cycle at a variety of temperatures and
rates without failure. The cells were therefore subjected to a discharge rate study at 23, 0,
−20, −30, and −40 °C, respectively. The sequence was as follows:
a. 25 mA charge to 4.10 V at room temperature
b. 25 mA discharge to 2.00 V at the given temperature
c. 25 mA charge to 4.10 V at the given temperature
d. 25 mA discharge to 2.00 V at the given temperature
e. 25 mA charge to 4.10 V at the given temperature
f. 50 mA discharge to 2.00 V at the given temperature
g. 25 mA charge to 4.10 V at the given temperature
h. 100 mA discharge to 2.00 V at the given temperature

A comparison of steps b and d show the effects of charging at low temperature
versus room temperature on the discharge capacity of the cell, and subsequent discharges
show the rate capability of the cells when charged and discharged at low temperature. Data
is shown for the −20 °C study as a representation of the overall trends.
Figures 3-82 – 3-87 show the capacities of the four discharges (i.e., b, d, f, and h) as a
percentage of the capacity of the first discharge (b, which was recorded after charging at
room temperature). All the cells show a drop-off in capacity when charged at low

201

Figure 3-82. Discharge rate study at -20 °C of MPG-111/LiNi
x
Co
1-x
AlO
2
cell containing
1.0 M LiPF
6
, 0.15 M LiBOB in EC+EMC+TPPa (20:65:15 v/v).

Figure 3-83. Discharge rate study at -20 °C of MPG-111/LiNi
x
Co
1-x
AlO
2
cell containing
1.0 M LiPF
6
, 0.15 M LiBOB in EC+FEC+EMC+TPPa (10:10:65:15 v/v).
1.00!
1.50!
2.00!
2.50!
3.00!
3.50!
4.00!
4.50!
5.00!
0%! 20%! 40%! 60%! 80%! 100%!
Voltage (V)!
Discharge Capacity (% of room temp. 25 mA capacity)!
25 mA Discharge (charge at 23C)!
25 mA Discharge (charge at -20C)!
50 mA Discharge (charge at -20C) !
100 mA Discharge (charge at -20C)!
1.00!
1.50!
2.00!
2.50!
3.00!
3.50!
4.00!
4.50!
5.00!
0%! 20%! 40%! 60%! 80%! 100%!
Voltage (V)!
Discharge Capacity (% of room temp. 25 mA capacity)!
25 mA Discharge (charge at 23C)!
25 mA Discharge (charge at -20C)!
50 mA Discharge (charge at -20C) !
100 mA Discharge (charge at -20C)!

202

Figure 3-84. Discharge rate study at -20 °C of MPG-111/LiNi
x
Co
1-x
AlO
2
cell containing
1.0 M LiPF
6
, 0.15 M LiBOB in FEC+EMC+TPPa (20:65:15 v/v).

Figure 3-85. Discharge rate study at -20 °C of MPG-111/LiNi
x
Co
1-x
AlO
2
cell containing
1.0 M LiPF
6
, 0.20 M LiBOB in FEC+EMC+TPPa (20:65:15 v/v).
1.00!
1.50!
2.00!
2.50!
3.00!
3.50!
4.00!
4.50!
5.00!
0%! 20%! 40%! 60%! 80%! 100%!
Voltage (V)!
Discharge Capacity (% of room temp. 25 mA capacity)!
25 mA Discharge (charge at 23C)!
25 mA Discharge (charge at -20C)!
50 mA Discharge (charge at -20C) !
100 mA Discharge (charge at -20C)!
1.00!
1.50!
2.00!
2.50!
3.00!
3.50!
4.00!
4.50!
5.00!
0%! 20%! 40%! 60%! 80%! 100%!
Voltage (V)!
Discharge Capacity (% of room temp. 25 mA capacity)!
25 mA Discharge (charge at 23C)!
25 mA Discharge (charge at -20C)!
50 mA Discharge (charge at -20C) !
100 mA Discharge (charge at -20C)!

203

Figure 3-86. Discharge rate study at -20 °C of MPG-111/LiNi
x
Co
1-x
AlO
2
cell containing
1.0 M LiPF
6
, 0.25 M LiBOB in FEC+EMC+TPPa (20:65:15 v/v).
Figure 3-87. Discharge rate study at -20 °C of MPG-111/LiNi
x
Co
1-x
AlO
2
cell containing
1.0 M LiPF
6
in EC+DEC+DMC (1:1:1 v/v).
1.00!
1.50!
2.00!
2.50!
3.00!
3.50!
4.00!
4.50!
5.00!
0%! 20%! 40%! 60%! 80%! 100%!
Voltage (V)!
Discharge Capacity (% of room temp. 25 mA capacity)!
25 mA Discharge (charge at 23C)!
25 mA Discharge (charge at 23C)!
50 mA Discharge (charge at 23C) !
100 mA Discharge (charge at 23C)!
1.00!
1.50!
2.00!
2.50!
3.00!
3.50!
4.00!
4.50!
5.00!
0%! 20%! 40%! 60%! 80%! 100%!
Voltage (V)!
Discharge Capacity (% of room temp. 25 mA capacity)!
25 mA Discharge (charge at
23C)!
25 mA Discharge (charge at
23C)!
50 mA Discharge (charge at
23C) !
100 mA Discharge (charge at
23C)!

204
temperature even when the rate is the same as the first discharge. Comparing the cells with
EC, EMC, TPPa and either 0.15 M, 0.20 M, or 0.25 M LiBOB, the relative capacities were
smaller at all rates with increasing LiBOB content. Comparing the cells with 0%, 10%, and
20% FEC, there is little impact on the relative capacities achieved when charging at low
temperature. The ternary baseline shows excellent relative capacity at 25 mA, but retains the
least capacity when discharging at 100 mA. The cell which gave the best relative discharge
capacity at 100 mA when charged at −20 °C was the cell with 20% FEC.
The anode behavior of these cells over the sequence is shown in Figure 3-88. The
cells with EC, EMC, TPPa and LiBOB experienced a negative anode potential during
charging, while those with FEC and LiBOB and the baseline did not. A potential negative to
that of the lithium couple increases the risk of lithium plating, especially at high rates and
low temperatures, but no direct evidence of lithium stripping is seen in the discharge curves.

3.4 Conclusions
In the preceding work, we have attempted to develop advanced lithium-ion battery
electrolytes with improved safety, temperature, and high-voltage characteristics. To this
end, we have examined flame retardant additives, fluorinated carbonates, and film-forming
additives and their impact on cell chemistry and performance. This was carried out using
electrochemical techniques, including dc micropolarization, Tafel polarization, and

205

Figure 3-88. Anode behavior during -20 °C rate study of FEC- and LiBOB-containing MPG-111/LiNi
x
Co
1-x
AlO
2
cells

206
electrochemical impedance spectroscopy; electrical techniques, including cycling and
individual discharge and charge; and ex-situ techniques such as cyclic voltammetry.
Based on our electrochemical stability studies, electrolytes containing TPPa possess
comparable stability to the baseline all-carbonate electrolyte up to 5 V, while TPPi is
relatively unstable. VC improves electrolyte stability in the 4 V range, but undergoes
oxidation near 4.8 V. FEC and bTFEC showed the potential to have comparable or better
stability than the baseline over a wide potential range. Full cell data further indicated the
viability of both bTFEC and FEC in MCMB-NCO cells, particularly at low temperatures.
Studies in graphite-NCA cells showed that the addition of FEC was kinetically
advantageous, while the addition of LiBOB had a dramatic detrimental effect on the anode.
Nevertheless, discharge data showed that LiBOB did not have a significant impact on cell
capacity, except at very low temperature and high rate of discharge. Furthermore, the use of
higher quantities of the flame-retardant additive TPPa was seen to be a viable approach
toward reducing flammability without severely impacting cell performance.

207
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Abstract (if available)

Abstract Direct alkaline methanol fuel cells were constructed and tested using a hydroxide- exchange membrane. The operating conditions and material components of the membrane electrode assemblies (MEAs), such as temperature, oxygen flow rate, fuel composition, and electrode material, were optimized. These cells were compared to analogous Nafion®-based MEAs and the power output was found to be comparable under similar conditions. ❧ Direct formic acid fuel cells (DFAFC) were constructed and tested using various palladium-based anode electrocatalysts. The power output, catalytic activity, and durability of these MEAs were evaluated. It was found that Pd-Au catalysts supported on a carbon-TaC blend were more active and durable than unsupported or carbon-supported materials. ❧ New electrolyte formulations for lithium-ion batteries were prepared and tested in terms of electrochemical stability, lithium intercalation/deintercalation kinetics, and cycling and rate capabilities with several state-of-the-art electrode systems. Flame-retardant additives and fluorinated co-solvents were examined in particular, and several formulations were identified which provide greater safety with comparable or better performance to known baseline electrolytes.

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Asset Metadata

Creator Krause, Frederick Charles (author)

Core Title Electrocatalysts for direct liquid-feed fuel cells and advanced electrolytes for lithium-ion batteries

School College of Letters, Arts and Sciences

Degree Doctor of Philosophy

Degree Program Chemistry

Publication Date 07/15/2013

Defense Date 06/06/2012

Publisher University of Southern California (original), University of Southern California. Libraries (digital)

Tag electrochemistry,fuel cells,lithium-ion batteries,OAI-PMH Harvest

Language English

Contributor Electronically uploaded by the author (provenance)

Advisor Prakash, G. K. Surya (committee chair), Olah, George A. (committee member), Shing, Katherine (committee member)

Creator Email fkrause@usc.edu,frederick.krause@gmail.com

Permanent Link (DOI) https://doi.org/10.25549/usctheses-c3-60305

Unique identifier UC11288790

Identifier usctheses-c3-60305 (legacy record id)

Legacy Identifier etd-KrauseFred-958.pdf

Dmrecord 60305

Document Type Dissertation

Rights Krause, Frederick Charles

Type texts

Source University of Southern California (contributing entity), University of Southern California Dissertations and Theses (collection)

Access Conditions The author retains rights to his/her dissertation, thesis or other graduate work according to U.S. copyright law. Electronic access is being provided by the USC Libraries in agreement with the a...

Repository Name University of Southern California Digital Library

Repository Location USC Digital Library, University of Southern California, University Park Campus MC 2810, 3434 South Grand Avenue, 2nd Floor, Los Angeles, California 90089-2810, USA