University of Southern California Dissertations and Theses

Studies on lithium-ion battery electrolytes and three component Strecker reaction

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Studies on lithium-ion battery electrolytes and three component Strecker reaction

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STUDIES ON LITHIUM-ION BATTERY ELECTROLYTES AND THREE
COMPONENT STRECKER REACTION

by

Kiah Anton Smith

___________________________________________

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

December 2009

Copyright 2009 Kiah Anton Smith

ii
Epigraph

“If we knew what we were doing, it would not be called research, would it?”
-Albert Einstein

Ad astra per aspera.

iii
Dedication

To my Grandfather, Ralph Harding Smith

iv
Acknowledgements

I would like to thank those who lent their knowledge, support, and imagination to
me as I pursued my doctoral degree.
This work would not have been possible without the support and guidance
provided my advisor Dr. G. K. Surya Prakash who whose steady advice assisted me in
numerous ways, including guiding me toward the field of lithium-ion batteries, instilling
passion in me for research, and providing personal support when I needed it.
Special thanks are also given to those I worked with at the Jet Propulsion
Laboratory. Dr. Marshall C. Smart, in particular, deserves special acknowledgement for
the support he provided on the work performed on Li-ion batteries described within this
work. From the first cell built to the final electrochemical analysis, Marshall’s
willingness to share his knowledge of Li-ion battery development have made my own
research into the subject possible. Others at JPL that deserve special acknowledgement
include Larry Whitcanack, Dr. Kumar Bugga, Dr. S. R. Narayanan, and Dr. Will West,
all of whom contributed to make my time at JPL more fruitful and much more enjoyable.
I would also like to thank members of the Prakash/Olah research group,
particularly Dr. Thomas Mathew and Dr. Chiradeep Panja, who provided assistance to
the development of the PVP-SO
2
Strecker reaction. I would also like to thank group
members Rehana Ismail, Habiba Vaghoo, Dr. Sujith Chacko, and Frederick Krause for
their assistance during my time in the group.
I cannot end without thanking my family, on whose encouragement, humor, and
love I have relied on for so long. To my parents, I am forever thankful to have your
v
support. Thank you for believing in me, at times more than I believed in my own ability.
To my sister, thank you for pushing me to higher standards.
Finally, I thank my fiancée, Elizabeth Burnes, whose patience and love could
never be overstated.

vi
Table of Contents

Epigraph ii

Dedication iii

Acknowledgements iv

List of Tables ix

List of Figures xiv

Abstract xx

CHAPTER 1: An Introduction to Li-Ion Batteries and Their Electrolytes 1
1.1 Introduction 1
1.2 The Development of Li-Ion Batteries 2
1.3 Electrolyte Solvents for Li-Ion Batteries 7
1.4 Chapter 1 References 12

CHAPTER 2: Introduction of Fluorinated Esters to Li-Ion Electrolyte 13
Solvents for Extended Operating Temperatures
2.1 Introduction 13
2.1.1 Electrolytes with Wide Operating 14
Temperatures
2.1.2 Fluorinated Solvents 18
2.1.2.1 Fluorinated Carbonates 19
2.1.2.2 Fluorinated Esters 21
2.2 Experimental Methods 22
2.3 Results and Discussion 25
2.3.1 Experimental Cell Results 28
2.3.1.1 Formation Characteristics 28
2.3.1.2 Low Temperature Discharge 31
Characteristics
2.3.2 Electrochemical Evaluation of Fluorinated Ester 39
Containing Electrolytes
2.3.2.1 Conductivity of Fluoroester Electrolytes 40
2.3.2.2 Cyclic Voltammetry of Fluoroester Electrolytes 42
2.3.2.3 Electrochemical Impedance Spectroscopy 44
2.3.2.4 Tafel Measurements 54
2.3.2.5 DC Micro-Polarization Measurements 58
2.4 Conclusions 60
2.5 Chapter 2 References 62
vii
CHAPTER 3: Improving the Safety Characteristics of Li-Ion Cells via the 64
Use of Flame-Retardant Additives
3.1 Introduction 64
3.1.1 Safety Issues of Li-Ion Batteries 65
3.1.2 Approaches to Improve the Safety of Li-Ion 68
Batteries
3.1.3 Phosphorus Based Flame Retardant Additives 71
3.1.4 Flame Retardant Additives as Lewis Base 73
Additives to Stabilize the LiPF
6
Salt
3.2 Experimental Methods 74
3.3 Results and Discussion 74
3.3.1 Electrolytes Containing Phosphorus Flame 77
Retardant Additives and Fluoroester Co-Solvent
3.3.1.1 Experimental Cell Results 79
3.3.1.1.1 Formation Characteristics 79
3.3.1.1.2 Low Temperature Discharge 83
Results
3.3.1.1.3 Cycle Life Evaluation 88
3.3.1.2 Electrochemical Evaluation 88
3.3.1.2.1 Electrochemical Impedance 89
Spectroscopy
3.3.1.2.2 Tafel Polarization Measurements 97
3.3.1.2.3 DC Micro-Polarization 101
Measurements
3.3.2 Improving Performance of Cells Containing 103
Flame-Retardant Additives through use of an
SEI Enhancing Additive
3.3.2.1 Experimental Cell Results 104
3.3.2.1.1 Formation Characteristics 104
3.3.2.1.2 Low Temperature Discharge 108
Results
3.3.2.1.3 Cycle Life Evaluation 112
3.3.2.2 Electrochemical Evaluation 113
3.3.2.2.1 Electrochemical Impedance 114
Spectroscopy
3.3.2.2.2 Tafel Polarization Measurements 121
3.3.2.2.3 DC Micro-Polarization 125
Measurements
3.3.3 Effects of Branch Substitution and Oxidation 127
State of Phosphorus in Phosphorus Containing
Flame-Retardant Additives in Li-Ion Cells
3.3.3.1 Experimental Cell Results 128
3.3.3.1.1 Formation Characteristics 128
3.3.3.1.2 Low Temperature Discharge 131
Results
viii
3.3.3.1.3 Cycle Life Evaluation 136
3.3.3.2 Electrochemical Evaluation 138
3.3.3.2.1 Conductivity Measurements 139
3.3.3.2.2 Electrochemical Impedance 140
Spectroscopy
3.3.3.2.3 Tafel Polarization Measurements 147
3.3.3.2.4 DC Micro-Polarization 150
Measurements
3.4 Conclusion 152
3.5 Chapter 3 References 154

CHAPTER 4: Electrolyte Salt Modification for Improved Li-Ion Cell 157
Performance
4.1 Introduction 157
4.1.1 Additives for Improving Li-Ion Cell Performance 158
4.1.2 Solid-Electrolyte Interface Enhancing Agents 159
4.1.3 Electrolytes Utilizing Boron Salts 161
4.2 Experimental Methods 163
4.3 Results and Discussion 164
4.3.1 Experimental Cell Results 165
4.3.1.1 Formation Characteristics 165
4.3.1.2 Low Temperature Discharge Characteristics 167
4.3.1.3 High Temperature Resiliency of Electrolytes 170
4.3.2 Electrochemical Characteristics 174
4.3.2.1 Electrochemical Impedance Spectroscopy 174
4.3.2.2 Tafel Polarization Measurements 187
4.3.2.3 DC Micro-Polarization Measurement 191
4.4 Conclusions 194
4.5 Chapter 4 References 195

CHAPTER 5: PVP-SO
2
Complex as a Solid Mild Acid Catalyst for Efficient 197
One Pot, Three Component Synthesis of α-Aminonitriles
5.1 Introduction 197
5.2 Experimental Methods 198
5.2.1 Preparation of the PVP-SO
2
198
5.2.2 One Pot Synthesis of α-Aminonitriles 200
5.3 Results and Discussion 200
5.3.1 NMR Data of the α–Aminonitrile Compounds 206
5.4 Conclusions 212
5.5 Chapter 5 References 213

BIBLIOGRAPHY 215

ix
List of Tables

Table 2.1 Formation data for fluoroester electrolytes 29

Table 2.2 Summary of discharge performance (capacity, % room 35
temperature) of experimental lithium-ion cells at various
low temperatures containing electrolytes consisting of
1.0M LiPF
6
EC+EMC+X (20:60:20 v/v %) (where X=
TFEB, ETFA, TFEA, TFEP, and MPFP)

Table 2.3 Low-Temperature discharge capacity values of experimental 36
lithium-ion cells containing EC-based electrolytes with
various concentrations of 2,2,2-trifluoroethyl butyrate ester
co-solvents at various low temperatures and rates

Table 2.4 Low-Temperature discharge values of experimental 38
lithium-ion cells charged at low temperature

Table 2.5 Conductivity values at depressed temperatures 42

Table 2.6 Cathode LiNi
x
Co
1-x
O
2
Impedance Data 48

Table 2.7 Anode MCMB Impedance Data 51

Table 2.8 Full Cell Impedance Data 53

Table 2.9 DC micro-polarization measurements at various temperatures 59
of anode from lithium-ion cells containing electrolytes
consisting of 1.0M LiPF
6
EC+EMC+X (20:60:20 v/v %)
(where X = TFEB, ETFA, TFEA, TFEP)

Table 2.10 DC micro-polarization measurements at various temperatures 59
of cathode from lithium-ion cells containing electrolytes
consisting of 1.0M LiPF
6
EC+EMC+X (20:60:20 v/v %)
(where X = TFEB, ETFA, TFEA, TFEP)

Table 3.1 Formation Data for the four cells containing flame-retardant 82
additives and baseline electrolyte

Table 3.2 Discharge capacities of cells containing flame-retardant 86
additive electrolytes

x
Table 3.3 Capacity fade of flame-retardant additive cells after low 87
temperature discharge tests

Table 3.4 Anode MCMB impedance data 94

Table 3.5 Cathode LiNi
x
Co
1-x
O
2
impedance data 95

Table 3.6 Full cell impedance data 96

Table 3.7 DC Micro-polarization measurements at various temperatures 102
of anodes from lithium-ion cells containing electrolytes
consisting of 1.0M LiPF
6
EC+EMC+TFEB+X
(20:55:20:5 v/v %) (where X = TPhPh, TBuPh, TEtPh,
and TFMPo)

Table 3.8 DC Micro-polarization measurements at various temperatures 102
of cathodes from lithium-ion cells containing electrolytes
consisting of 1.0M LiPF
6
EC+EMC+TFEB+X
(20:55:20:5 v/v %) (where X = TPhPh, TBuPh, TEtPh,
and TFMPo)

Table 3.9 Formation data for TPhPh/VC cells 107

Table 3.10 Summary of discharge performance (capacity, % room 111
temperature) of experimental lithium-ion cells at low
temperatures of FRA/VC cells at various rates

Table 3.11 Capacity fade of the cells containing flame-retardant additives 112
after low temperature discharge tests

Table 3.12 EIS low temperature data for FRA/VC cathodes 116

Table 3.13 EIS low temperature data for FRA/VC anodes 118

Table 3.14 EIS low temperature data for FRA/VC full cells 120

Table 3.15 DC micro-polarization measurements at various temperatures 126
at anodes from lithium-ion cells containing FRA/VC
additives and baseline electrolytes

Table 3.16 DC micro-polarization measurements at cathodes at various 126
temperatures for Li-ion cells containing FRA/VC additives
and baseline electrolytes

Table 3.17 Formation data for flame retardant additive electrolyte cells 130
xi

Table 3.18 Low temperature capacity of FRA containing electrolytes 134

Table 3.19 Rate capabilities of FRA electrolytes when charged at low 135
temperatures

Table 3.20 Capacity fade of cells containing flame-retardant additives 137
after low temperature discharge tests

Table 3.21 Conductivity measurements at different temperatures for 140
electrolytes containing various flame retardant additives

Table 3.22 Cathode impedance values for electrolytes containing 142
flame retardant additives

Table 3.23 Anode impedance values for electrolytes containing flame 145
retardant additives

Table 3.24 Full cell impedance values for electrolytes containing 146
flame retardant additives

Table 3.25 Current density of the anode of electrolytes with flame 148
retardant additives at various temperatures recorded
via Tafel measurements

Table 3.26 Current density of the anode of electrolytes with flame 150
retardant additives at various temperatures recorded
via Tafel measurements

Table 3.27 Polarization values for the anodes of cells using 151
flame retarding electrolytes

Table 3.28 Polarization values for the anodes of cells using flame 151
retarding electrolytes

Table 4.1 Formation data for cells with new electrolyte salt compositions 166

Table 4.2 Low temperature discharge capacities and capacity retention 169
for cells with new electrolyte salt mixtures

Table 4.3 Discharge capacities and capacity retention for the cells 171
with new electrolyte salt compositions after high
temperature storage periods

xii

Table 4.4 Discharge capacities and capacity retention for the cells 173
with new electrolyte salt compositions after high
temperature storage periods

Table 4.5 Cathode impedance data of new salt composition cells at 178
low temperatures

Table 4.6 Impedance data for cathodes of cells with new electrolyte 179
salt composition after high temperature storage periods

Table 4.7 Impedance data for anodes of cells with new electrolyte 182
salt composition at low temperatures

Table 4.8 Impedance data for cathodes of cells with new electrolyte 183
salt composition after high temperature storage periods

Table 4.9 Impedance data for the full cell with new electrolyte salt 185
composition at low temperatures

Table 4.10 Impedance data for full cells with new electrolyte salt 186
composition after high temperature storage periods

Table 4.11 Current density (A) of the cathodes from cells with new 189
electrolyte salt formulations as a function of temperature

Table 4.12 Current density (A) of the anodes from cells with new 189
electrolyte salt formulation as function of temperature

Table 4.13 Current density (A) of the cathodes from cells with new 190
electrolyte salt formulations after high temperature storage
periods

Table 4.14 Current density (A) of the anodes from cells with new 191
electrolyte salt formulations after high temperature storage
periods

Table 4.15 Polarization resistance values (kΩ) for the cathode of cells 192
containing new electrolyte formulations at reduced
temperatures

Table 4.16 Polarization resistance values (kΩ) for the anodes of cells 193
containing new electrolyte formulation at reduced
temperatures

xiii
Table 4.17 Polarization resistance values (kΩ) for the cathodes of cells 193
containing new electrolyte formulation after high
temperature storage periods

Table 4.18 Polarization resistance values (kΩ) for the anodes of cells 194
containing new electrolyte formulations after high temperature
storage periods

Table 5.1 PVP-SO
2
catalyzed α-aminonitrile synthesis from various 204
aldehydes, anilene, and TMSCN

Table 5.2 PVP-SO
2
catalyzed α-aminonitrile synthesis from various 206
amines

xiv
List of Figures

Figure 1.1 Schematic of Li-ion battery 5

Figure 1.2 Chemical structures of (a) propylene carbonate and 10
(b) ethylene carbonate

Figure 2.1 Chemical structures of fluorinated esters investigated at 26
Li-ion electrolyte co-solvents

Figure 2.2 Fifth discharge of formation cycle as a function of TFEB 30
loading percentage

Figure 2.3 Fifth discharge of formation cycle of electrolytes containing 31
different types of fluorinated ester

Figure 2.4 Discharge capacity (% Room Temp Capacity) of 33
experimental lithium-ion cells at –20
o
C (~ C/16 rate)
containing electrolytes consisting of 1.0M LiPF
6

EC+EMC+X (20:60:20 v/v %) (X = TFEB, ETFA, TFEA,
TFEP, MPFP)

Figure 2.5 Discharge capacity (% Room Temp Capacity) of experimental 33
lithium-ion cells at –40
o
C (~ C/16 rate) containing electrolytes
consisting of 1.0M LiPF
6
EC+EMC+X (20:60:20 v/v %)
(X = TFEB, ETFA, TFEA, TFEP, MPFP)

Figure 2.6 Discharge capacity (% Room Temp Capacity) of experimental 34
lithium-ion cells at –40
o
C (~ C/8 rate) containing electrolytes
consisting of 1.0M LiPF
6
EC+EMC+X (20:60:20 v/v %)
(X = TFEB, ETFA, TFEA, TFEP, MPFP)

Figure 2.7 Discharge capacity (% Room Temp Capacity) of experimental 34
lithium-ion cells at –60
o
C (~ C/80 rate) containing electrolytes
consisting of 1.0M LiPF
6
EC+EMC+X (20:60:20 v/v %)
(X = TFEB, ETFA, TFEA, TFEP, MPFP)

Figure 2.8 Anode potential (v. Li/Li
+
) of cells with 5 and 20% TFEB 39
electrolytes during -20ºC rate study

Figure 2.9 Conductivity comparison of baseline and ester containing 41
Electrolytes

xv
Figure 2.10 Cyclic voltammogram of fluoroester containing electrolytes 43
from 2.0 to 6.0 V

Figure 2.11 Cyclic voltammogram of fluoroester containing electrolytes 44
from 0.1 to 4.0 V

Figure 2.12 Nyquist plot of electrochemical impedance spectroscopy 47
(EIS) measurements at 23ºC of LiNi
x
Co
1-x
O
2
electrodes from
lithium-ion cells containing electrolytes consisting of 1.0M
LiPF
6
EC+EMC+X (20:60:20 v/v %) (where X = TFEB, ETFA,
TFEA, TFEP. and MPFP)

Figure 2.13 Nyquist plot of electrochemical impedance spectroscopy (EIS) 50
measurements at 23ºC of MCMB anodes from lithium-ion
cells containing electrolytes consisting of 1.0M LiPF
6

EC+EMC+X (20:60:20 v/v %) (where X = TFEB, ETFA,
TFEA, TFEP. and MPFP)

Figure 2.14 Nyquist plot of electrochemical impedance spectroscopy (EIS) 52
measurements at 23
o
C of lithium-ion cells containing
electrolytes consisting of 1.0M LiPF
6
EC+EMC+X
(20:60:20 v/v %) (where X = TFEB, ETFA, TFEA, TFEP,
and MPFP)

Figure 2.15 Tafel polarization measurements at 23
o
C of MCMB electrodes 55
from lithium-ion cells containing electrolytes consisting of
1.0M LiPF6 EC+EMC+X (20:60:20 v/v %) (where X = TFEB,
ETFA, TFEA, and TFEP).

Figure 2.16 Tafel polarization measurements at 23
o
C of LiNi
x
Co
1-x
O
2
55
electrodes from lithium-ion cells containing electrolytes
consisting of 1.0M LiPF6 EC+EMC+X (20:60:20 v/v %)
(where X = TFEB, ETFA, TFEA, and TFEB)

Figure 2.17 Tafel polarization measurements at - 40
o
C of MCMB electrodes 56
from lithium-ion cells containing electrolytes consisting of
1.0M LiPF
6
in EC+EMC+X (20:60:20 v/v %) (X = TFEB,
TFEA, and TFEP)

Figure 2.18 Tafel polarization measurements at -40
o
C of LiNi
x
Co
1-x
O
2
57
electrodes from lithium-ion cells containing electrolytes
consisting of 1.0M LiPF
6
EC+EMC+X (20:60:20 v/v %)
(where X = TFEB, ETFA, TFEA, and TFEB).

xvi
Figure 2.19 Tafel polarization measurements at - 60
o
C of MCMB 57
electrodes from lithium-ion cells containing electrolytes
consisting of 1.0M LiPF
6
EC+EMC+X (20:60:20 v/v %)
(where X = TFEB and TFEP)

Figure 2.20 Tafel polarization measurements at - 60
o
C of 58
LiNi
x
Co
1-x
O
2
electrodes from lithium-ion cells containing
electrolytes consisting of 1.0M LiPF6 EC+EMC+X
(20:60:20 v/v %) (where X = TFEB and TFEP)

Figure 3.1 Chemical structures of phosphorus containing 79
flame-retardant additives

Figure 3.2 Fifth discharge of formation cycling for cells containing 81
flame retardant additive and fluoroester cosolvent

Figure 3.3 Discharge of FRA containing electrolytes with fluorinated 84
ester at -20º at 25mA/h

Figure 3.4 Discharge of FRA containing electrolytes at -40º at 25mA/h 85

Figure 3.5 Capacity loss (%) as a function of number of cycles 88

Figure 3.6 Electrochemical Impedance Spectroscopy (EIS) 90
measurements at 23
o
C of MCMB electrodes from
lithium-ion cells containing electrolytes consisting of
1.0M LiPF
6
EC+EMC+TFEB+X (20:55:20:5 v/v %)
(where X = TPhPh, TBuPh, TEtPh, TFMPo)

Figure 3.7 Electrochemical Impedance Spectroscopy (EIS) 93
measurements at 23
o
C of LiNi
x
Co
1-x
O
2
electrodes from
lithium-ion cells containing electrolytes consisting of
1.0M LiPF6 EC+EMC+TFEB+X (20:55:20:5 v/v %)
(where X = TPhPh, TBuPh, TEtPh, TFMPo)

Figure 3.8 Tafel polarization measurements at 23
o
C of LiNi
x
Co
1-x
O
2
98
electrodes from lithium-ion cells containing electrolytes
consisting of 1.0M LiPF6 EC+EMC+TFEB+X
(20:55:20:5 v/v %) (where X = TPhPh, TBuPh, TEtPh, TFMPo)

Figure 3.9 Tafel polarization measurements at -20
o
C of LiNi
x
Co
1-x
O
2
99
electrodes from lithium-ion cells containing electrolytes
consisting of 1.0M LiPF6 EC+EMC+TFEB+X
(20:55:20:5 v/v %) (where X = TPhPh, TBuPh, TEtPh, TFMPo)

xvii
Figure 3.10 Tafel polarization measurements at -40
o
C of Ni
x
Co
1-x
O
2
99
electrodes from lithium-ion cells containing electrolytes
consisting of 1.0M LiPF6 EC+EMC+TFEB+X
(20:55:20:5 v/v %) (where X = TPhPh, TBuPh, TEtPh, TFMPo)

Figure 3.11 Tafel polarization measurements at 23
o
C of MCMB 100
electrodes from lithium-ion cells containing electrolytes
consisting of 1.0M LiPF6 EC+EMC+TFEB+X
(20:55:20:5 v/v %) (where X = TPhPh, TBuPh, TEtPh, TFMPo)

Figure 3.12 Tafel polarization measurements at -20
o
C of MCMB 100
electrodes from lithium-ion cells containing electrolytes
consisting of 1.0M LiPF6 EC+EMC+TFEB+X
(20:55:20:5 v/v %) (where X = TPhPh, TBuPh, TEtPh, TFMPo)

Figure 3.13 Tafel polarization measurements at -40
o
C of MCMB 101
electrodes from lithium-ion cells containing electrolytes
consisting of 1.0M LiPF6 EC+EMC+TFEB+X
(20:55:20:5 v/v %) (where X = TPhPh, TBuPh, TEtPh, TFMPo)

Figure 3.14 Fifth discharge of formation cycling for cells containing 105
TPhPh and/or VC

Figure 3.15 Cell discharge profile at -20 ºC with various electrolytes 109
at 25 mA/h. Cells were charged at room temperature to 4.1 V

Figure 3.16 Cell discharge profile at -40 ºC with various electrolytes 110
at 25 mA/h

Figure 3.17 Cell discharge profile at -60 ºC with various electrolytes 110
at 5 mA/h

Figure 3.18 Capacity fade of cells with FRA/VC electrolytes 113
at room temperature during cycle life testing

Figure 3.19 Nyquist plot of FRA/VC study cathode impedance at 115
room temperature

Figure 3.20 Nyquist plot of FRA/VC study anode impedance at 117
room temperature

Figure 3.21 Tafel polarization measurements at 20ºC of MCMB 122
anodes in FRA/VC study

xviii
Figure 3.22 Tafel polarization measurements at 20ºC of LiNi
0.8
Co
0.2
O
2
122
cathodes in FRA/VC study
Figure 3.23 Tafel polarization measurements at -40ºC of MCMB 123
anodes in FRA/VC study

Figure 3.24 Tafel polarization measurements at -40ºC of LiNi
0.8
Co
0.2
O
2
124
cathodes in FRA/VC study

Figure 3.25 Tafel polarization measurements at -60ºC of MCMB anodes 124
in FRA/VC study

Figure 3.26 Tafel polarization measurements at -60ºC of LiNi
0.8
Co
0.2
O
2
125
cathodes in FRA/VC study

Figure 3.27 Chemical structures of phosphorus-containing flame-retardants 127
examined

Figure 3.28 Fifth discharge of formation cycling for cells containing flame 131
retardant additives

Figure 3.29 Discharge of FRA containing electrolytes at -20º at 25mA/h 133

Figure 3.30 Discharge of FRA containing electrolytes at -40º at 25mA/h 133

Figure 3.31 Life cycle testing of cells containing FRAs 138

Figure 3.32 Anode Tafel measurements of electrolytes containing 148
flame-retardant additives

Figure 3.33 Cathode Tafel measurements of electrolytes containing 149
flame-retardant additives

Figure 4.1 Proposed SEI formation reactions 159

Figure 4.2 Anodic polymerization of additives containing double 161
bonds occurs via a similar mechanism and can help in the
formation of robust SEI

Figure 4.3 Capacity retention of cells with boron salts at -20ºC 168
discharged at 25 mA/h

Figure 4.4 Capacity retention of cells with boron salts at -40ºC 168
discharged at 25 mA/h

xix
Figure 4.5 Room temperature Nyquist impedance data of cathodes of 177
cells containing new electrolyte salt compositions

Figure 4.6 Nyquist impedance plots after third high temperature 180
storage period

Figure 4.7 Nyquist plot of anode impedance data collected for 181
cells with new electrolyte salt composition at room temperature

Figure 4.8 Cathode Tafel measurement data of cells with new 188
electrolyte salt formulations

Figure 4.9 Anode Tafel measurement data of cells with new 189
electrolyte salt formulations

Figure 5.1 SEM images of PVP and PVP-SO
2
complex 199

Figure 5.2 Preparation of PVP-SO
2
complex 201

Figure 5.3 TGA diagram for PVP-SO
2
complex 201

Figure 5.4 PVP-SO
2
catalyzed three component α-aminonitrile synthesis 203
from various aldehydes, aniline and TMSCN

Figure 5.5 PVP-SO
2
catalyzed three component α-aminonitrile synthesis 204
from various amines, 4-chlorobenzaldehyde and TMSCN

xx
Abstract

The first four chapters of this work describes the collaborative effort between the
University of Southern California (Los Angeles, CA) and the Jet Propulsion Laboratory
(Pasadena, CA) focused on developing electrolyte systems to meet the targeted
improvements desired by the United States space program. Within this work effort was
made to explain the effects of electrolyte modification to the overall performance of the
individual electrodes as well as the cell performance on a whole through employment of
electrochemical analysis (impedance spectroscopy, Tafel polarization, DC micro-
polarization, cyclic voltammetry, and conductivity) and electrical measurements (charge-
discharge characteristics, low-operating temperature characterization, high temperature
storage performance, and rate capabilities). Chapter 1 provides a brief discussion of the
background and development of lithium ion batteries. Further description of the
electrolyte systems of such devices is also provided. Developments made to increase the
operational temperature of Li-ion batteries using fluorinated esters are described in
chapter 2 of this document. Of the examined fluorinated esters, 2, 2, 2-trifluoroethyl
butyrate showed the most promised for low-temperature performance. Flame retardant
additives were added to electrolyte formulations to improve the safety characteristics of
the Li-ion cells, and these results are discussed in chapter 3. Within this chapter, a
correlation between flame retardant additive structure and electrochemical stability (and
thus, cell performance) is elaborated upon. Chapter 4 is the final chapter discussing Li-
ion electrolyte development and describes the work performed to extend the cycle life
xxi
and high temperature resiliency of cells using advanced lithium salt electrolytes, lithium
bis(oxolato)borate and lithium tetrafluoroborate.
The final chapter of this dissertation details a modification to the well known
Strecker reaction developed at the University of Southern California employing a solid
poly(4-vinylpyridine)-SO
2
complex as a mild solid acid catalyst for efficient synthesis of
α-aminonitriles in high yield and purity.

1
CHAPTER 1

An Introduction to Lithium-Ion Batteries and Their
Electrolytes

1.1 Introduction

Lithium-ion batteries have grown increasingly prominent since their invention in
1990 because of their high working voltages, energy density, long cycle life, low self-
discharge rates and lack of memory effects. Already common among small portable
electronic applications such as laptop computers and mobile phones, these batteries
continue to garner interest as power sources for applications with higher energy
requirements. An introductory look into the background and development of these
devices is provided. Furthermore, a brief discussion of the electrolytes of these systems
is presented. Though the importance of electrolyte choice within batteries influences the
overall characteristics of a cell to a lesser degree than the choice of electrode materials,
the influence of electrolyte nevertheless plays a critical role in many aspects of cell
performance ranging from lithium cycling efficiency, rate capability, operating
temperature range and abuse tolerance. As such, the University of Southern California
(USC) and Jet Propulsion Laboratory (JPL) have undertaken a collaborative effort to
develop improved electrolytes for use within lithium-ion cells for use within the United
States space program. Efforts on these electrolytes have resulted in numerous targeted
2
improvements to these cells such as wider temperature operation, increased cycle life,
and improved safety characteristics. Promising work to improve these characteristics has
been identified within the current work. Effort was made to explain the effects
modification to the electrolyte system have on the overall performance of the individual
electrodes as well as the cell performance on a whole through employment of
electrochemical analysis (impedance spectroscopy, Tafel polarization, DC micro-
polarization, cyclic voltammetry, and conductivity) and electrical measurements (charge-
discharge characteristics, low-operating temperature characterization, high temperature
storage performance, and rate capabilities). These results will be expanded upon in later
chapters of this work.

1.2 The Development of Li-Ion Batteries

The scientific history of the battery is traced back to the work of Alessandro Volta
in 1800. Using alternating discs of zinc and silver with stacks of brine soaked cardboard
in between, Volta was the first person to generate an electric current.
1
This “wet cell
battery” produced a steady current with an arc that was able to carry electricity over a
measurable distance. His work was instrumental in establishing the electromotive series
of the elements that is still in use today.
Attempts at improving Volta’s design began almost immediately, and throughout
the next century dozens of new systems were developed. Two of the most notable
systems came from W.R. Grove
2
and Georges Leclanche.
3
Grove’s discovery of the
“gas battery” is considered to be the predecessor of the modern day fuel cell, while
3
Leclanche’s discovery had more immediate applications. Within two years, over 20,000
of his zinc/carbon “dry cell” battery forerunners were being used within the telegraph
system worldwide.
The First World War saw a rapid increase in a use of batteries as a power supply
for field torches and radios. As the need for portable power increased, so too did the
advances in battery technology. Batteries became increasingly indispensable in day-to-
day modern life throughout the 20
th
and 21
st
century. As the size of portable electronics
shrink and the desire for specialized, readily available power is growing
The groundwork of the modern secondary (rechargeable) battery was developed
in 1859 by Gaston Planté.
4
For many years the most widely used and cost-effective
rechargeable batteries were based on Planté’s original lead-acid battery design.
Descendants of Planté’s original design are still commonly used in applications where
high power densities are required, such as automobiles.
Aside from their ability to be recharged, secondary batteries have numerous other
characteristics that differentiate them from primary batteries. These include high power
density, flat discharge profiles, and better low temperature performance than primary
batteries. However, primary batteries tend to have higher energy densities and charge
retention.
5

The applications of secondary batteries fall into two major categories:
5

1. The battery is used as an energy storage device to deliver its energy on
demand after being charged by primary energy sources. Examples of these
types of applications include automotive systems and emergency no-fail and
standby power sources.
4
2. The battery is discharged and recharged after use for reasons such as
convenience, cost savings, or are subject to power drains beyond the
capability of primary batteries. These are the batteries that come to mind
immediately when rechargeable batteries are mentioned and supply electrical
power to many of our modern devices including laptop computers and cellular
phones.

Pivotal to the development of secondary battery technology has been the use of
lithium. Because of its low equivalent weight (0.534 g/cm
3
) and high standard potential
(~ -3.0 V vs. SHE) lithium has received considerable attention as an anode material.
Researchers began to explore the uses of lithium for battery applications as early as the
1950’s when it was found that the metal was stable in a number of aprotic organic
solvents despite its high chemical reactivity. However, it wasn’t until Sony
commercialized the lithium ion (Li-ion) battery in 1990
6
that the use of lithium-based
systems began to mushroom. In his comprehensive review of lithium-ion battery
electrolyte technology, Xu detailed many of the challenges that were overcome during
this early growth period for lithium batteries.
7
Throughout the initial development phases,
lithium metal was targeted as an anode material for secondary batteries, but this type of
configuration leads to problems during cycling that make the implementation of such
chemistry difficult, both in terms of cycle life and safety. These obstacles occur due to
the formation of dendrites of lithium at the anode upon charge and discharge that become
electrochemically inactive, but chemically highly reactive due to high surface area of the
dendrites. These conditions lead to capacity fading and eventual cell failure. Potential
5
safety hazards associated with dendrite formation are an even greater concern than cell
failure. Hazards, such as fire and explosions, may occur due to the formation of these
dendrites.

Fig. 1.1 Schematic of Li-ion battery

Unlike previous batteries utilizing lithium, Li-ion batteries do not rely on the use
of lithium metal; instead lithium ions from lithium metal oxide-based cathodes are
“shuttled” through a lithium-ion conducting electrolyte and intercalated into the anode
material at a potential very close to that of lithium metal. The lithium ions “rock” back
and forth between the electrodes during charge and discharge, and no dendritic lithium is
6
ever plated out of the electrolyte solution.
8
The breakthrough for this technology came
when researchers exploited carbonaceous materials as anode intercalation hosts.
9
The
low cost of carbon and high lithium ion activity in the intercalation material made these
types of materials very attractive for further development and graphitic materials
eventually became the standard for Li-ion anode technology. Lithium was long known to
intercalate into graphite forming LiC
6
species with a chemical potential close to that of
metallic lithium. This is important so that the reduction in cell voltage compared with
lithium metal anodes remained small.
8
The half-cell reactions that take place within a Li-
ion cell are listed below.

However, hurdles remained for these systems. Organic electrolyte solvent
decomposition remained a problem, however, until 1990 when it was discovered that
using ethylene carbonate (EC) as a cosolvent helped to form a protective lithium ion
conducting surface film on the carbon anode, much like the layer previously discovered
on metallic lithium anodes.
10
In a highly important publication within the field, Dahn and
coworkers laid the groundwork for demonstrating the importance choice of electrolyte
can have on the performance of Li-ion cells. Within the communication, Dahn
established four important “rules” regarding the role of electrolytes in the formation of
the SEI layer and its influence over the overall performance of the cell. These
conclusions drawn by Dahn regarding ideal SEI formation are as follows:
7

1. The irreversible reaction between the electrolyte and the Li
x
C electrode
occurs only on the first charge of the cell.
2. When all available surface area of the carbon particles is coated with the
film of decomposition products, no further reaction with the electrolyte
solvent occurs. This film prevents further decomposition of the electrolyte
components while acting as an ionic conductor but an electronic insulator.
3. Li/carbon cells can be reversibly cycled many times once the SEI layer on
the carbon surfaces has been formed.
4. The SEI layer is most likely not simply Li
2
CO
3
, although it may be
Li
2
CO
3
based.

These “rules” of electrolyte decomposition and SEI formation played a critical
role in the explosive growth of Li-ion battery research and still exhibit critical influence
on the choice of electrolyte composition to this day.

1.3 Electrolyte Solvents for Li-Ion Batteries

The role of any electrolyte in a secondary battery is to shuttle charges between the
electrode pair. The vast majority of electrolytes consist of salts dissolved in a liquid
electrolyte solvent. Sandwiched between the positive and negative electrodes, the
electrolyte plays an important role in the overall operation of the battery, and any newly
8
developed material for electrolytes requires extensive compatibility testing between the
electrolyte and the electrodes.
Lithium-ion battery electrolytes are most commonly composed of one or more
lithium salts and two or more aprotic organic solvents. Important qualities for the salt
include solubility and disassociation in the chosen electrolyte solvent, inertness against
electrolyte solvents and other cell components, and chemical and electrochemical
stability. Some of the most common salts include LiAsF
6
, LiBF
4
, and LiClO
4
, and most
prominently, LiPF
6
. While the available choices of lithium salts are relatively small
compared to the bounty of choices available for electrolyte solvent, cathode and anode
materials, the choice can exhibit considerable influence upon the cell properties.
While the electrolyte salt plays an important role in determining the properties of
any electrolyte, the electrolyte solvent makes up the bulk of the electrolyte solution. The
properties of a battery can be finely tuned by making adjustments to the electrolyte and
much research has been performed investigating new electrolytes. An ideal electrolyte
for Li-ion batteries should meet the following minimal criteria:
7

1. It should have a high dielectric constant giving it the ability to dissolve
lithium salts to sufficient concentration.
2. To enable facile ion transport, it should be fluid within the operating
temperature range. Preferred solvents will have high boiling points and
low melting points.
3. It should be electrochemically stable in the operating potentials of the cell.
9
4. It should remain inert to all cell components, especially the charged
surfaces of the electrodes, during cell operation.
5. It should be safe, nontoxic, and economical.

The highly active nature of both electrodes (strongly reducing lithiated carbon at
the anode and strongly oxidizing transition metal oxide cathodes) put severe limitations
on the type of electrolytes that can be chosen for the purpose. For example, although
they exhibit excellent solvating properties, electrolytes with active protons are not
suitable solvents because the reduction/oxidation of such protons occur within the
operating potentials of the cells. Conversely, many nonaqueous compounds exhibit the
necessary electrochemical stability, but only those with polar groups can sufficiently
dissolve the lithium salts.
A majority of nonaqueous solvents that have been investigated for use in lithium
batteries have been carbonates (or other carbonyl derivatives such as esters). Ether
solvents garnered substantial consideration in batteries that employed lithium metal due
to desirable characteristics such as low viscosity, high ionic conductivity, and improved
short-term lithium morphology over the short term.
7,11
However, extensive investigation
found these electrolytes employing ether based solvents in Li-ion cells exhibited poor
capacity retention dendritic formation, and poorer than anticipated cathode stability.
7,12

During the 1990s most ethers were phased out of electrolyte systems.
Cyclic carbonates have been critical to the development of suitable Li-ion
electrolytes. These compounds, propylene carbonate (PC) and ethylene carbonate (EC),
have been recognized to play a critical role in the formation of the SEI on the
10
carbonaceous anodes. With its wide liquid range, high dielectric constant, and static
stability, PC, in particular, has a long history of use as an electrochemically stable
solvent.
13
However, despite its static electrochemical stability, PC exhibits dynamic
reactivity in Li-ion cells that leads to lithium loss and eventual disintegration of the
graphene structure of the anode because of PC co-intercalation leading to what Dahn and
coworkers termed “exfoliation” during cycling.
13a

Fig. 1.2 Chemical structures of (a) propylene carbonate and (b) ethylene carbonate

Although it has many favorable properties, EC was not favored as a common
electrolyte solvent for lithium ion batteries due to its high melting point (~36ºC). It was
not until Dahn elicited the SEI forming properties that EC exhibited in the Li-ion cell it
began to gain attention.
10
Although the difference is a single methyl group, EC is able to
form an effective film (SEI) on a graphite anode that PC is unable to form. This film
prevents any sustained electrolyte decomposition on the anode. The key to the film’s
functionality is that is ionically conducting while remaining electronically insulating, i.e.
it allows Li
+
ions through to the electrode without self-discharging the cell. The slight
difference in structure has lead to EC’s ubiquitous use as a cosolvent in Li-ion cells.
11
Solvent systems have been developed that employ the use of linear carbonates as
cosolvents to EC to impart the favorable physical characteristics (melting-point
suppression and lowered viscosity) for wide usage. Linear carbonates have low boiling
points, low viscosity, and much lower dielectric constants than their cyclic counterparts.
They form homogenous solutions with EC at any ratio. When alone, linear carbonates do
not exhibit the necessary oxidative stability at Li-ion cell cathode voltages (~4.0 V vs.
Li); however, when combined with EC a synergistic effect imparts stability upon the
mixture that was first described by Tarascon and Guyamard.
14
It was found that an
electrolyte mixture of EC and dimethyl carbonate (DMC) was stable on the surface of the
spinel cathode out to 5.0 V. Each individual solvent imparts beneficial characteristics on
the mixtures. They are high anodic stability on cathode surfaces, high solvation of
lithium salts by EC, and low viscosity necessary to promote ion transport by DMC. This
combination of EC and linear carbonates has grown to become the bulk of all Li-ion cell
electrolytes to this day with over a billion such cells manufactured each year. Even with
this considerable success, significant research with the field continues to finely tune these
systems to meet the particular needs of specific applications.

12
1.4 Chapter 1 References

1. Volta, A. Phil. Trans. R. Soc. 1800, 2, 403.
2. Grove, W.R. Phil. Mag. 1842, 21, 417.
3. Leclanche, G. US Pat. 55441, 1866; “Electrical battery with primary and
secondary piles combined,” US Pat. 64113, 1867.
4. Planté, G. Compte Rendu 1860, 50, 640.
5. Linden, D. Handbook of Batteries and Fuel Cells McGraw-Hill, Inc.: New York,
1984.
6. Nagaura, T. in: Proc. 5
th
Intern. Seminar on Lithium Battery Technology and
Applications, Deerfield Beach, Fl. March 5-7, 1990.
7. Xu, K. Chem. Rev. 2004, 104, 4303.
8. Brandt, K. Solid State Ionics. 1994, 69, 173.
9. (a) Naguara, T.; Nagamine, M.; Tanabe, I.; Miyamoto, N. Prog. Batteries Sol.
Cells 1989, 8, 84. (b) Naguara, T.; Ozawa, K. Prog. Batteries Sol. Cells, 1990, 9,
209. (c) Nishi, Y.; Azuma, H.; Omarua, A. U.S. Patent 4,959,281, 1990.
10. Fong, A.; von Sacken, U.; Dahn, J. R. J. Electrochem. Soc. 1990, 137, 2009.
11. Koch, V. R.; Young, J. H. J. Electrochem. Soc. 1979, 126, 181.
12. (a) Yohimatsu, I.; Hirai, T.; Yamaki, J. J. Electrochem. Soc. 1988, 135, 2422. (b)
Xu, K., Ding, S. P.; Jow, T.R.; J. Electrochem. Soc. 1999, 146, 4172.
13. (a) Selim, R.; Bro, P. J. Electrochem. Soc. 1974, 121, 1457. (b) Rauh, R. D.;
Brummer, S. B. Electrochim. Acta 1977, 22, 75.
14. (a) Guyomard, D.; Tarascon, J. M. J. Electrochem. Soc. 1993, 140, 3071. (b)
Tarascon, J. M.; Guyomard, D. Solid State Ionics 1994, 69, 293.

13
CHAPTER 2

Introduction of Fluorinated Esters to Li-Ion Electrolyte
Solvents for Extended Operating Temperatures

2.1 Introduction

Collaboration between Jet Propulsion Laboratory and the University of Southern
California resulted in efforts to develop advanced electrolytes to improve the
performance of lithium-ion cells, especially at low temperatures. A number of
fluorinated ester electrolyte co-solvents have been identified that can be incorporated
along with carbonate mixtures into multi-component electrolyte formulations for
enhanced performance over a wide temperature range. In this chapter, work carried out
to investigate the fluorinated ester co-solvents 2,2,2-trifluoroethyl butyrate (TFEB), ethyl
trifluoroacetate (ETFA), 2,2,2-trifluoroethyl acetate (TFEA), 2,2,2-trifluoroethyl
propionate (TFEP), and methyl pentafluoropropionate (MPFP) is reported. The
electrolyte co-solvents were incorporated into multi-component electrolytes of the
following composition: 1.0 M LiPF
6
in ethylene carbonate (EC) + ethyl methyl carbonate
(EMC) + X (20:80-X:X v/v) (where X = fluorinated ester co-solvent volume percentage
ranging between 5 and 60%). Emphasis was placed upon determining the effect of
electrolyte type upon the low temperature performance in experimental MCMB-
14
LiNi
x
Co
1-x
O
2
cells (equipped with reference electrodes) and characterizing the respective
electrolytes influence upon the lithium intercalation/de-intercalation kinetics within the
individual electrodes. These cells were subjected to electrical characterization (charge
and discharge at different temperatures and rates), as well as, electrochemical
characterization (EIS, DC micropolarization and Tafel polarization measurements). Of
these solvents, trifluoroethyl butyrate (TFEB) has been demonstrated to yield the best
performance, with improved low temperature capability and preliminary high
temperature resilience.

2.1.1 Electrolytes with Wide Operating Temperatures

Currently, the state-of-art lithium-ion system has been demonstrated to operate
over a temperature range of -40
o
to +40
o
C; however, the performance is severely limited
beyond these temperature extremes. Limitations at very low temperatures are due to poor
electrolyte conductivity, poor lithium intercalation kinetics at the electrode surface layers,
and limited ionic diffusion within the electrode bulk. Some future applications typically
will require high specific energy batteries that can operate at very low temperatures,
while still providing adequate performance and stability at higher temperatures. Recent
efforts at JPL have resulted in improved low temperature performance with the utilization
of ester-based electrolytes; however, due to their high flammability and reactivity at
higher temperatures improvements are desired with regard to safety and high temperature
resilience.
1

15
Several factors can influence the low temperature performance of lithium-ion
cells, including lithium ion mobility in the electrolyte solution (electrolyte conductivity),
cell design, electrode thickness, separator porosity and separator wetting properties. Of
these parameters, the electrolyte properties are presumably the most dominant, in that
sufficient conductivity is a necessary condition for good performance at low
temperatures. In designing electrolytes with high conductivity at low temperatures, the
solvents should possess a combination of several critical properties, such as: high
dielectric constant, low viscosity, adequate coordination behavior, as well as appropriate
liquid ranges and salt solubilities in the medium. In earlier work, Smart et. al. identified
an optimized electrolyte formulation consisting of lithium hexafluorophosphate (LiPF
6
)
dissolved in a ternary, equi-proportion mixture of ethylene carbonate (EC), dimethyl
carbonate (DMC), and diethyl carbonate (DEC).
2
This electrolyte formulation was
demonstrated to enable improved performance over a wide temperature range and was
incorporated in the battery chemistry ultimately selected and implemented on the 2003
Mars mission MER rovers, consisting of two 8-cell, 8 Ah lithium-ion batteries connected
in series manufactured by Lithion, Inc.
For further improvement of the low temperature performance, a systematic study
of co-solvents possessing low viscosities and freezing points was undertaken. These co-
solvents can be added to the ternary aliphatic carbonate mixture. In such formulations,
the desirable film is expected to form from the aliphatic carbonate components, whereas,
the low temperature conductivity (by reduced viscosity and freezing point) could be aided
by the quaternary co-solvent. Such quaternary mixtures permit an optimization of the
electrolyte properties (i.e., dielectric constant, viscosity, liquid range, coordination
16
properties, and overall stability) and are likely to produce more highly conductive
solutions, especially at low temperatures, due to a disordering effect in the lithium ion
coordination behavior of the solvent medium. This work led to further improvements of
the low temperature performance of lithium ion cells with a quaternary electrolyte
formulation consisting of 1.0 M LiPF
6
EC+DEC+DMC+EMC (1:1:1:2 v/v), which has
been incorporated into cells utilized for Air Force applications.
1b
Further development
has led to the identification of a number of low EC-content quaternary solvent blend
electrolytes, which have enabled excellent performance down to –50
o
C.
In addition to focusing on the development of all-carbonate-based electrolyte
formulations, the use of other low melting, low viscosity co-solvents to further improve
the low temperature conductivity and performance of lithium-ion cells have also been
explored at JPL. Toward this end, the use of alkyl esters, including methyl formate
(MF), methyl acetate (MA), ethyl acetate (EA), ethyl propionate (EP), and ethyl butyrate
(EB) have been used in multi-component electrolyte formulations.
1a
Previous work
showed the higher molecular weight esters (i.e., ethyl propionate and ethyl butyrate)
resulted in both improved low temperature performance and good stability at ambient
temperatures.
2b
Excellent performance was obtained down to –40
o
C with electrolytes
consisting of the following formulations: a) 1.0 M LiPF
6
EC+DEC+DMC+ethyl butyrate
(EB) (1:1:1:1 v/v %) and b) a) 1.0 M LiPF
6
EC+DEC+DMC+ethyl propionate (EP)
(1:1:1:1 v/v %). Although electrolytes with methyl acetate and ethyl acetate (low
molecular weight esters) were shown to result in high conductivity at low temperatures
and good cell performance at low temperature initially, their high reactivity toward the
anode led to continued cell degradation and poor cell performance over extended cycling.
17
Thus, subsequent work was focused upon developing low temperature electrolytes that
contain high molecular weight ester co-solvents (which inherently possess greater
stability) incorporated into electrolyte formulation in large proportion to further extend
the low temperature performance down to -70
o
C. Toward this end, workers at JPL
demonstrated improved performance with multi-component electrolytes of the following
composition: 1.0 M LiPF
6
in ethylene carbonate (EC) + ethyl methyl carbonate (EMC) +
X (1:1:8 v/v %) (where X = methyl butyrate (MB), ethyl butyrate (EB), methyl
propionate (MP), and ethyl valerate (EV).
Other researchers have also investigated some ester-based co-solvents and have
recognized the potential for improved low temperature performance. For example,
researchers at Matsushita
3
have investigated electrolyte of the following compositions: a)
1.5 M LiPF
6
in EC+DEC+MA (1:2:2), b) 1.5 M LiPF
6
in EC+DEC+MP (1:2:2), and c)
1.5 M LiPF
6
in EC+DEC+EP (1:2:2). Although promising performance was reported,
the incorporation of a large proportion of diethyl carbonate (DEC) is not preferred due to
the undesirable effects that this solvent has upon the surface films of carbon anodes. In
another study, researchers from Saft have investigated electrolytes containing ethyl
acetate (EA) and methyl butyrate (MB). More specifically, the group investigated the
following electrolyte formulations: a) 1.0 M LiPF
6
in EC+DMC+MA , b) 1.0 M LiPF
6
in
EC+DMC+MB, c) 1.0 M LiPF
6
in EC+PC+MB and d) 1.0 M LiPF
6
in EC+DMC+EA
(solvent ratios not provided).
4
The group reported good low temperature performance
with the methyl butyrate-based electrolyte; however, they did not investigate the
performance at temperatures below –40
o
C. Other researchers have also investigated the
use of methyl acetate and ethyl acetate in ternary mixtures with and without blending
18
with toluene in an attempt to obtain improved performance to temperatures as low as –50
o
C.
5
Work at Samsung investigated the performance of a number of electrolyte
formulations performance at low temperatures, including: a) 1.0 M LiPF
6
in
EC+EMC+EA (30:30:40), b) 1.0 M LiPF
6
in EC+DMC+MA (30:35:35), c) 1.0 M LiPF
6

in EC+DEC+EP (30:35:35), and d) 1.0 M LiPF
6
in EC+EMC+EP (30:30:40). Although
good performance was demonstrated at –20
o
C, the performance attributes under more
extreme temperature demands were not investigated.
6

2.1.2 Fluorinated Solvents

With increased interest in developing Li-ion battery electrolytes with improved
safety characteristics that maintain desired stability and performance for the cells,
researchers have turned some of their attention to halogenated solvents. The hope was
realize some safety improvements to the cells because of the generally accepted idea that
solvent flammability is decreased as the degree of halogenation is increased.
7
While
chlorinated solvents have been investigated
8
, the vast majority of research on halogenated
Li-ion co-solvents has focused on fluorine to impart these desired characteristics.
Numerous varieties of fluorine containing compounds have been tested as co-
solvents in Li-ion batteries with a wide range of results. Nakajima et al. tested the
fluorinated ethers 2,2,3,3,3-pentafluoropropyl methyl ether, 1,1,3,3,3-pentafluoro-2-
trifluoromethyl ether, 2,2,3,3-tetrafluoropropyl difluoromethyl ether, 1,1,2,3,3,3-
hexafluoropropyl ethyl ether, and 1,1,2-trifluoro-2-chloroethyl 2,2,2-trifluoroethyl ether
at 10% v/v loading level to 1.0 M solutions of EC.DEC (50.50 v/v%).
9
This group
19
discovered that modifying the structure of the ethers can influence the stability of
electrolyte in Li-ion cells leading to increased reductive stability of the electrolytes and
higher capacities for these cells compared to electrolytes containing the non-fluorinated
analogous ethers. However, because results were obtained using LiClO
4
, an electrolyte
salt that readily reacts with many organic species in a violent nature under high rate or
extreme temperature conditions, the test system is difficult to extrapolate beyond
laboratory conditions. Methyl nanofluorobutyl ether, with its nonflammable
characteristics, was investigated as a co-solvent to improve the safety of Li-ion cells with
promising results.
10
In this system, cycle life and SEI impedance data are comparable to
non-halogenated carbonate mixtures, but the fluorinated ether deleteriously affects the
rate capability and contributed to a significant voltage drop leading to substantial power
loss from cells containing such an electrolyte. Additionally, electrolytes containing the
fluorinated amide N,N-dimethyl trifluoroacetamide exhibited attractive properties
including beneficial filming behavior, oxidative stability cathode potentials, and desirable
physical properties such as moderate viscosity, low melting point and comparatively high
boiling point and flash point.
11
However, the data in this study was collected against a
lithium metal anode with PC as a co-solvent, thereby no data exists with regards to the
actual feasibility of such a system in a practical Li-ion cell.

2.1.2.1 Fluorinated Carbonates

The most heavily studied of the fluorinated co-solvent materials are the
fluorinated carbonates. Fluoroethylene carbonate
12
and trifluoropropylene carbonate
13
20
have been studied extensively. In the case of the former compound, the introduction of
FEC reduces the irreversible capacity loss in the first cycle and increases the reversible
capacity of the cell. Overall results indicate that the fluorinated EC significantly
outperforms the chlorinated version. With trifluoropropylene carbonate researchers
demonstrated that lithium can be intercalated into graphite electrodes when the carbonate
is the only electrolyte solvent, though results significantly improved by employing is as a
co-solvent.
Collaboration between USC and JPL has netted interesting results on numerous
fluorocarbonates and carbamates that demonstrated good low temperature performance
with excellent reversibility in many cases.
14
Unlike previously investigated
fluoroethylene carbonate and trifluoropropylene carbonate electrolyte, which are used to
completely replace the EC vital to SEI formation, the fluorinated linear carbonates were
used to serve as diluents to the high-melting and viscous ethylene carbonate while
imparting desirable properties to the electrolyte solution including lower melting points,
higher anodic stability, increased safety, and favorable SEI-forming characteristics.
Blending them as co-solvents (20-50 v/v%) in ternary and quaternary electrolyte
formulations with traditional cyclic and linear aliphatic carbonates, Smart et al.
14

incorporated eight carbonate and carbamates into Li-ion cells: methyl-2,2,2-trifluoroethyl
carbonate, ethyl-2,2,2-trifluoroethyl carbonate, propyl-2,2,2-trifluoroethyl carbonate,
methyl-2,2,2,2’,2’,2’-hexafluoro-i-propyl carbonate, ethyl-2,2,2,2’,2’,2’-hexafluoro-i-
propyl carbonate, di-2,2,2-trifluoroethyl carbonate, 2,2,-trifluoroethyl N,N-dimethyl
carbamate and hexafluoro-i-propyl N,N-dimethyl carbamate. In additions to examining
the physical properties and their viability in full Li-ion cell, JPL researchers employed
21
electrochemical techniques such as Tafel polarization and DC micro-polarization to
confirm the facile kinetics of lithium ion intercalation/de-intercalation made possible by
the new SEI chemistry imparted by the fluorinated co-solvents. Interfacial resistance and
utilized capacity were investigated using EIS. It was found that the fluorinated
carbonates outperformed the fluorinated carbamates. In particular, the cells containing
ethyl-2,2,2-trifluoroethyl carbonate showed improved performance, in terms of cell
capacity at room temperature and at -20 ºC as well as exhibiting desirable film forming
properties.

2.1.2.2 Fluorinated Esters

Researchers have studied some fluoroester-based co-solvents in the context of
lithium-ion battery electrolytes as well. For example, researchers at Kyoto University
15

have investigated electrolytes containing lithium perchlorate (LiClO
4
) in solutions of
ethylene carbonate (EC) and diethyl carbonate (DEC) with a number of fluoroesters,
including: a) methyl difluoroacetate (CHF
2
COOCH
3
), b) ethyl pentafluoropropionate
(CF
3
CF
2
COOCH
2
CH
3
), c) methyl 2,2,2,2’,2’,2’-hexafluoro-iso-butyrate
((CF
3
)
2
CHCOOCH
3
), d) methyl septafluorobutyrate (F(CF
2
)
3
COOCH
3
), e) ethyl
1,1,2,2,3,3,4,4-octafluorovalerate (H(CF
2
)
4
COOCH
2
CH
3
), f) methyl 7-fluorooctanoate
(F(CH
2
)
7
COOCH
3
), and g) ethyl 7-fluorooctanoate (F(CH
2
)
7
COOCH
2
CH
3
). In terms of
performance, the best electrolyte formulation investigated was 1.0M LiClO
4
EC+DEC+
CHF
2
COOCH
3
(47.6 : 47.6 : 4.8 v/v %), which exhibited increased charge capacity and
coulombic efficiencies with graphite electrodes at –4
o
C compared to the baseline
22
solution. Researchers also found that the fluoroesters are reduced at potentials higher
than EC. Although researchers performed no studies to confirm the finding, their higher
reduction potential may allow for adaptation of the SEI layer properties through
incorporation into the film. However, all other fluoroester electrolyte solutions examined
in the study resulted in poorer coulombic efficiencies and reduced capacities compared
with the EC+DEC solution. In a more recent study, Nakajima et al. investigated
electrolyte formulations using LiClO
4
dissolved in

EC+PC+DEC with addition of methyl
difluoroacetate or methyl 2,2,2,2’,2’2’-hexafluoro-iso-butyrate.
9
In some cases, a modest
increase in the delivered capacities of graphite electrodes at slightly depressed
temperatures (0 to –10
o
C) was observed. Yamaki and coworkers
16
have investigated the
efficacy of using methyl difluoroacetate as an electrolyte solvent and have demonstrated
enhanced thermal stability of solutions based upon LiPF
6
. The same researchers
suggested that these systems might enable the use of lithium metal as an anode material
in rechargeable lithium batteries.

2.2 Experimental Methods

Cell performance, aging tests, and electrochemical measurements including cyclic
voltammetry, conductivity measurements, Tafel, DC micro-polarization, and impedance
testing was performed on three electrode cells. These ~400 mAh experimental Li-ion
cells are constructed within O-ring sealed, glass cells containing jelly roll configuration
of MCMB carbon anode electrodes, LiNi
0.8
Co
0.2
O
2
cathode electrodes, and Li metal
reference electrodes. Construction of cells is performed in dry room with dew points
23
typically -50 ºC +/-10 ºC. Anodes are MCMB carbon with polyvinyldifluoride (PVDF)
binder and conductive diluent (high surface area carbon or graphite) on 1 mil copper foil
and are supplied by Lithion Inc. (formerly Yardney Technical Products) of Pawtucket,
CT. Cathodes are LiNi
0.8
Co
0.2
O
2
, mixed with binder PVDF and conductive diluent (high
surface area carbon or graphite) on aluminum foil and are supplied by Lithion Inc.
Lithium metal references are fabricated from lithium metal foil obtained from Foote
Mineral Co of Frazer, PA.
To prepare the electrodes for use, the anodes are trimmed to 1 ¾ inches by 7 1/8
inches long. 1/8 inch wide of active material is scraped from anodes to expose bare
copper to which a 1/8 inch copper foil tab is attached via spot welding to allow
connection to the glass cell. A nearly identical process is performed on cathodes with the
exception that the electrodes are trimmed to 6 3/8 inches and tabbed with aluminum foil.
Lithium metal reference electrodes were prepared using nickel webbing approximately 1
¾ inches x ¼ inch tabbed with nickel foil with lithium metal wrapped around the nickel
webbing. Each electrode (anode, cathode, reference) is rolled flat with a hand-held roller.
Electrodes are individually sealed in porous polypropylene supplied by Tonen-Setella.
Electrodes are rolled around a PTFE mandible in jellyroll configuration and connected to
glass cell via tabs using aluminum setscrew shaft couplers.
Electrolyte solutions were prepared in an argon glove box for examination with
O
2
typically below 2 ppm. Stock solution carbonate mixtures of battery-grade purity
from Mitsubishi Chemicals were used as base components to all mixed solutions. These
stock solutions are certified to contain <50 ppm H
2
O content and come prepared at 1.0 M
LiPF
6
. The fluorinated esters tested as electrolyte co-solvents were procured two ways.
24
Initially, the esters were synthesized at the University of Southern California according to
known methods of esterification, typically reaction between corresponding acid chloride
and alcohol. Reaction products were worked up and purified via distillation under inert
gas. Later tests used fluorinated esters purchased directly from Synquest Labs, Inc. and
distilled under inert gas at the University of Southern California. Esters were transported
to Jet Propulsion Laboratory in Pasadena, CA in airtight vessels. Fluorinated esters were
added to the electrolyte solution in contents ranging from 5-60% while all other additives
were generally kept at <5% or lower. Additional LiPF
6
salt was added to bring the final
concentration of the solution to 1.0 M.
Electrochemical measurements were made using an EG&G
potentiostat/galvanostat interfaced with an Windows PC, using Softcorr 352 initially and
converting to CorrWare 2 for later testing. Care was taken to ensure both programs gave
results consistent with one another. A Solartron 1255 Frequency Response Analyzer was
coupled with this potentiostat in order to perform impedance measurements. M388
software was used to control the instrument during the early stages of testing, and
PowerSuite software was used to control during later testing. Charge-discharge
measurements and cycling tests were performed with an Arbin battery cycler. The cells
were charged to 4.10V at 25mA/h, followed by a tapered charge period at constant
potential until the current dropped below 2 mA/h. This was followed by discharge to 25
mA/h top 2.75V. This formation cycling is repeated 5 times. Discharge tests were
performed using the same apparatus.
The specific conductivities of the electrolyte solution were measured between –
60
o
to 25
o
C using a conductivity cell, which consists of two platinized electrodes
25
immobilized in a glass apparatus and separated by a fixed distance. The cell constant of
the conductivity cell was determined using a standard 0.1 M KCl solution. A Tenney
environmental chamber was used to maintain the desired temperature within + 1
o
C for
the cells. Samples were subjected to a minimum of two hours of soak at the desired
temperature prior to measurements. Cyclic voltammetric measurements were carried out
in an O-ring sealed glass cell, with a 1 cm
2
platinum working electrode against a lithium
counter and reference electrodes. To ensure reproducible conditions of the platinum
electrode, it was etched in aqua-regia before each experiment.

2.3 Results and Discussion

In this study we focused our attention on a number of fluorinated ester co-solvents
with the intent of extending the temperature range of state-of-practice electrolyte
formulations and improving the inherent safety characteristics of the resultant lithium-ion
cells. The fluorinated ester co-solvents investigated include: trifluoroethyl butyrate
(TFEB), 1, ethyl trifluoroacetate (ETFA), 2, trifluoroethyl acetate (TFEA), 3,
trifluoroethyl propionate (TFEP), 4, and methyl pentafluoropropionate (MPFP), 5, (Fig.
2.1).

26
Fig. 2.1 Chemical structures of fluorinated esters investigated at Li-ion electrolyte co-
solvents

A number of experimental lithium-ion cells, consisting of MCMB carbon anodes
and LiNi
0.8
Co
0.2
O
2
cathodes, were fabricated to study the electrolytes described below.
These cells served to verify and demonstrate the reversibility, low temperature
performance, and electrochemical aspects of each electrolyte when used in a cell.
Described below are the formulations of the tested electrolytes. Two baseline
electrolytes containing only carbonate electrolyte solvents were also tested as references
(Cells K and L).

A. 1.0 M LiPF
6
in EC+EMC+TFEB (20:75:5 v/v %)
B. 1.0 M LiPF
6
in EC+EMC+TFEB (20:60:20 v/v %)
C. 1.0 M LiPF
6
in EC+EMC+TFEB (20:40:40 v/v %)
D. 1.0 M LiPF
6
in EC+EMC+TFEB (20:20:60 v/v %)
E. 1.0 M LiPF
6
in EC+EMC+ETFA (20:60:20 v/v %)
O
O C
H
2
N
3
C
H
2
C
CF
3
O
O H
3
C
H
2
C
CF
3
O
O C
H
2
H
2
C
H
2
C
CF
3
H
3
C
O
O F
3
C
H
2
C
CH
3
O
O C
F
2
F
3
C CH
3
2
3
4
5
1
27
F. 1.0 M LiPF
6
in EC+EMC+TFEA (20:60:20 v/v %)
G. 1.0 M LiPF
6
in EC+EMC+ TFEA (20:40:40 v/v %)
H. 1.0 M LiPF
6
in EC+EMC+TFEP (20:60:20 v/v %)
I. 1.0 M LiPF
6
in EC+EMC+TFEP (20:40:40 v/v %)
J. 1.0 M LiPF
6
in EC+EMC+MPFP (20:60:20 v/v %)
K. 1.0 M LiPF
6
in EC+EMC (20:80 v/v %)
L. 1.0 M LiPF
6
in EC+EMC+DMC (33:33:33 v/v %)

Various electrochemical methods were employed to evaluate the viability of the
fluorinated ester-containing electrolytes described, including electrolyte conductivity
measurements, cyclic voltammetric measurements to evaluate the electrochemical
stability of the investigated electrolytes, determination of the kinetics of lithium
intercalation and de-intercalation through Tafel and DC micro-polarization
measurements. Additionally, the relative interfacial stabilities of the candidate electrolyte
with MCMB-carbon anodes and LiNi
0.80
Co
0.2
O
2
cathodes were studied by using EIS
measurements. These methods will be described more fully in the following sections.
Discharge characterization at various temperatures and rates was also performed.
Of the electrolytes listed above, the formulations consisting of 1.0 M LiPF
6
in
EC+EMC+TFEB (20:60:20 v/v %) and 1.0 M LiPF
6
in EC+EMC+TFEA (20:60:20 v/v
%) yielded the greatest improvements to cell performance, especially at low
temperatures. These electrolyte formulations were compared with baseline solutions in
order to discern the beneficial effect of the fluoroesters, namely: 1.0 M LiPF
6
in
EC+DEC+DMC (1:1:1 v/v %) and 1.0 M LiPF
6
in EC+EMC (20:80 v/v %).
28
2.3.1 Experimental Cell Results

2.3.1.1 Formation Characteristics

In order to assess the viability of the fluoroester-based electrolytes, the
charge/discharge characteristics of MCMC-LiNiCoO
2
experimental cells were tested
using the various electrolytes described. The irreversible and reversible capacities were
monitored as a function of electrolyte type during the initial five cycles. For clarity sake,
these first 5 cycles will be referred to as the formation cycles because the formation of the
SEI layer occurs at this period. In most cases, the electrolytes containing fluoroesters did
not perform as well as those containing carbonate-only mixtures during the formation
cycles (Table 2.1), exhibiting greater irreversible capacity loss after the first cycle. For
example, the cell containing 20% MPFP had an irreversible capacity loss double to both
of the baseline mixtures. This capacity loss is exacerbated as the percentage content of
fluoroester is increased.
However, large irreversible capacity loss was not observed for all fluorinated ester
containing cells, as the cells containing the fluorinated ester TFEB at 5, 20 and 40 v/v%
performed comparably to the carbonate-only cells. In fact, the cells containing the
electrolyte 1.0 M LiPF
6
EC+EMC+TFEB (20:60:20 v/v %) and 1.0 M LiPF
6

EC+EMC+TFEB (20:75:5 v/v %) were observed to have increased performance in terms
of initial irreversible capacity loss than the cell containing EC:EMC (20:80 v/v%) as
electrolyte solvents (Fig. 2.2). These results suggest that the TFEB solvent has adequate
29
electrochemical stability and/or decomposes in such a way as to form robust electrode
surface films inhibiting further electrolyte decomposition from occurring.

Table 2.1 Formation data for fluoroester electrolytes. Generally, cells with fluoroester
electrolytes did not perform as well as those with baseline carbonate only electrolytes.
Cells (A) and (B) exhibit improved formation characteristics when compared to the
baseline cells.

Electrolyte Type
Charge Capacity
(mAh) 1st Cycle
Discharge
Capacity (mAh)
1st Cycle
Irreversible
Capacity (1st
Cycle)
Coulombic
Efficiency
(1st Cycle)
Charge Capacity
(mAh) 5th Cycle
Reversible
Capacity (mAh)
5th Cycle
Cumulative
Irreversible Capacity
(1st-5th Cycle)
Coulombic
Efficiency
(5th Cycle)
1.0 M LiPF6
EC+EMC+TFEB
(20:75:5 v/v%)
523.11 447.61 75.50 85.57 445.88 436.13 123.82 97.81
1.0 M LiPF6
EC+EMC+TFEB
(20:60:20 v/v%)
517.93 442.02 75.91 85.34 438.26 431.23 114.02 98.40
1.0 M LiPF6
EC+EMC+TFEB
(20:40:40 v/v%)
530.78 442.52 88.26 83.37 444.40 432.90 146.49 97.41
1.0 M LiPF6
EC+EMC+TFEB
(20:20:60 v/v%)
427.51 305.52 121.99 71.46 296.36 286.16 187.47 96.56
1.0 M LiPF6
EC+EMC+ETFA
(20:60:20 v/v%)
484.59 356.29 128.30 73.52 348.77 337.26 142.97 96.70
1.0 M LiPF6
EC+EMC+TFEA
(20:60:20 v/v%)
395.76 292.02 103.74 73.79 310.29 303.45 279.04 97.80
1.0 M LiPF6
EC+EMC+TFEA
(20:40:40 v/v%)
423.94 307.87 116.07 72.62 325.46 311.74 151.79 95.78
1.0 M LiPF6
EC+EMC+TFEP
(20:60:20 v/v%)
494.19 412.87 81.32 83.54 418.66 403.41 160.00 96.36
1.0 M LiPF6
EC+EMC+TFEP
(20:40:40 v/v%)
458.09 372.93 85.16 81.41 375.25 358.96 160.45 95.66
1.0 M LiPF6
EC+EMC+MPFP
(20:60:20 v/v%)
480.23 294.15 186.08 61.25 212.81 188.44 325.48 88.55
1.0 M LiPF6
EC+EMC (20:80
v/v%)
468.16 404.43 63.73 86.39 401.26 391.36 113.58 97.53
1.0 M LiPF6
EC+DEC+DMC
(33:33:33 v/v%)
478.80 410.60 68.20 85.76 411.70 397.90 131.34 96.65
30

Fig. 2.2 Fifth discharge of formation cycle as a function of TFEB loading percentage.

The cells containing fluorinated ester co-solvents compare more favorably to the
baseline with regard to reversible capacity. After five cycles, nearly all fluoroester
containing electrolytes have coulombic efficiencies greater than 95%. The exception is
the cell containing MPFP, which falls behind all other cells by a wide margin (coulombic
efficiency <89%). Ranking the other electrolyte solvent systems by their coulombic
efficiencies after 5 cycles follows the following trend: EC+EMC+TFEB (20:60:20)
(98.40%) > EC+EMC+TFEA (20:60:20) (97.80%) = EC+EMC+TFEB (20:75:5) (97.8
%) > EC+EMC (20:80) (97.76%) > EC+EMC+TFEB (20:40:40) (97.41%) >
EC+EMC+ETFA (20:60:20) (96.70%) > EC+DEC+DMC (33:33:33) (96.65%) >
EC+EMC+TFEB (20:20:60) (96.56%) > EC+EMC+TFEA (20:40:40) (95.78%) >
EC+EMC+MPFP (20:60:20) (88.55%) (Fig 2.3).
31

Fig. 2.3 Fifth discharge of formation cycle of electrolytes containing different types of
fluorinated ester.

Initial formation cycling of the cells consisted of charging and discharging at 25 mA
(~ C/16 rate) over a voltage range of 2.75V-4.10V at 23
o
C. After initial formation data is
collected, the cells were subjected to electrochemical evaluation (EIS, Tafel and DC
micro-polarization measurements) over a wide temperature range (20 to -60
o
C), which
will be discussed in later sections.

2.3.1.2 Low Temperature Discharge Characteristics

When the cells containing 20% v/v of the fluorinated esters were evaluated at
reduced temperature (-20
o
C) at moderate rates (~ C/16 discharge rate, 25mA/h), as shown
32
in Fig. 2.4, good low temperature performance was observed in many of the cells
containing these co-solvents, with >80% of the room temperature capacity being
delivered in the fluoroester cells containing TFEB, ETFA, and TFEA. The best
performing fluoroester cells were the cells containing TFEB and ETFA, each performing
as well or better than the baseline carbonate electrolytes. As the temperature of the tests
was lowered to -40ºC at similar rates, cells with TFEB and ETFA continue to outperform
the other fluorinated ester cells (Fig. 2.5). The ETFA cell outperformed both baseline
carbonate mixtures by a small margin at this rate and temperature while the TFEB and
TFEP cells outperformed the ternary carbonate mixture. The cell containing 20% TFEB
electrolyte was the best performing of the tested electrolytes as the temperature was
depressed and/or the rate was increased at these reduced temperatures (Fig. 2.6). When
the rate is doubled to ~C/8 at -40ºC the ETFA cell performance fell considerably while
the TFEB cell maintained a relatively high level of performance. Discharge of the cells
at 5mA at -60ºC supported the conclusion that TFEB was the best performing of the
series of tested fluoroester cells (Fig. 2.7). Under these conditions the TFEB based
fluoroester mixture retains ~10% more of its room temperature capacity than does the
next best performing fluoroester electrolyte, the cell with TFEP (63.15 to 53.59%). Table
2.2 displays all of the tested low temperature discharge results as a function of
fluoroester. For these low temperature discharge evaluations, all cells were charged at
room temperature then soaked at the depressed testing temperature for at least 6 hours.
33

Fig. 2.4 Discharge capacity (% Room Temp Capacity) of experimental lithium-ion cells
at –20
o
C (~ C/16 rate) containing electrolytes consisting of 1.0M LiPF
6
EC+EMC+X
(20:60:20 v/v %) (X = TFEB, ETFA, TFEA, TFEP, MPFP).

Fig. 2.5 Discharge capacity (% Room Temp Capacity) of experimental lithium-ion cells
at –40
o
C (~ C/16 rate) containing electrolytes consisting of 1.0M LiPF
6
EC+EMC+X
(20:60:20 v/v %) (X = TFEB, ETFA, TFEA, TFEP, MPFP).
34

Fig. 2.6 Discharge capacity (% Room Temp Capacity) of experimental lithium-ion cells
at –40
o
C (~ C/8 rate) containing electrolytes consisting of 1.0M LiPF
6
EC+EMC+X
(20:60:20 v/v %) (X = TFEB, ETFA, TFEA, TFEP, MPFP).

Fig. 2.7 Discharge capacity (% Room Temp Capacity) of experimental lithium-ion cells
at –60
o
C (~ C/80 rate) containing electrolytes consisting of 1.0M LiPF
6
EC+EMC+X
(20:60:20 v/v %) (X = TFEB, ETFA, TFEA, TFEP, MPFP).
35

Table 2.2 Summary of discharge performance (capacity, % room temperature) of experimental lithium-ion cells at various
low temperatures containing electrolytes consisting of 1.0M LiPF
6
EC+EMC+X (20:60:20 v/v %) (where X = TFEB,
ETFA, TFEA, TFEP, and MPFP). Cells were charged at room temperature prior to discharge.

36

Table 2.3 Low-Temperature discharge capacity values of experimental lithium-ion cells containing EC-based
electrolytes with various concentrations of 2,2,2-trifluoroethyl butyrate ester co-solvents at various low temperatures
and rates. Cells charged at room temperature.

37
Concentration effects were monitored as the amount of TFEB is varied from 5 to
60% in cells. As shown in Table 2.3, the 20% is the best performing electrolyte solution
at nearly every temperature and rate examined. Cells with 5 and 20 v/v% TFEB
performed considerably better than cells with higher loading contents at all depressed
temperatures indicating a limit to how much fluorinated ester can be added to an
electrolyte solution before severe deleterious effects are imparted onto the cell.
After all low temperature testing performed with room temperature charging, the
effects of charging the cells at low temperatures was also examined (Table 2.4). Cells are
charged at room temperature initially followed by temperature depression to the testing
temperature. Subsequent charges and discharges was performed at the various testing
temperatures. Again the cell with 20 v/v% TFEB outperforms all other TFEB containing
cells by a wide margin at all rates and temperatures. It is also appropriate to note that
these tests were performed after the room temperature charge/ low temperature discharge
tests were performed indicating the 20% TFEB mixture exhibits a certain amount of
robustness as well.
Figure 2.8 tracks the potentials of the anodes of the cells containing 5 and 20%
TFEB. During low temperature charging (and especially at higher rates of charge) the
anode plunges negative (vs. Li/Li
+
), which can lead to lithium metal plating onto the
electrode. If this were the case, the discharge profile for such cells would have
demonstrated a characteristic bump that is caused by the lithium re-dissolving into the
electrolyte solution. However, such evidence was not seen (not shown). Electrolytes
with higher TFEB content suffered from the same phenomenon, but the cells became
polarized so quickly that the graphical data was not useful to collect.
38

Table 2.4 Low-Temperature discharge values of experimental lithium-ion cells charged at low temperature. Cells
contain EC-based electrolytes with 2,2,2-trifluoroethyl butyrate ester co-solvents at various concentrations.
39

Fig. 2.8 Anode potential (v. Li/Li
+
) of cells with 5 and 20% TFEB electrolytes during
-20ºC rate study. Anode plunges into negative potentials when charged at low
temperatures.

2.3.2 Electrochemical Evaluation of Fluorinated Ester Containing Electrolytes

Although the reason for low temperature performance of the fluorinated ester-
based solutions has not been entirely elucidated, we that suspect the enhanced
performance is primarily due to improved mass transfer characteristics in the electrolyte
(higher ionic conductivity) and/or facile kinetics of lithium intercalation/de-intercalation
at the interface due to favorable film formation behavior at the electrode surfaces. To
enhance this understanding, we have assessed the characteristics of the systems using a
number of electrochemical techniques to evaluate the cells, including Tafel polarization
measurements, electrochemical impedance spectroscopy (EIS), and DC micro-
40
polarization measurements. It should be noted that in some cases electrochemical
performance of the electrolyte under initial testing conditions was deemed poor enough to
discontinue further measurements.

2.3.2.1 Conductivity of Fluoroester Electrolytes

The specific conductivities of three electrolyte formulations were measured at
different temperatures, from 25
o
C to -60
o
C (Fig. 2.9). The electrolytes were formulated
to analyze the effect of the addition of a fluorinated ester to the electrolyte system. For
comparison sake, the non-fluorinated ester, ethyl butyrate (EB) was incorporated into an
electrolyte formulation as well. The electrolyte compositions tested were as follows:

1. 1.0 M LiPF
6
in EC:EMC (20.80 v/v%)
2. 1.0 M LiPF
6
in EC:EMC:EB (20.60.20 v/v%)
3. 1.0 M LiPF
6
in EC:EMC:TFEB (20.60.20 v/v%)

Of the electrolytes investigated, solutions containing lower molecular weight
esters displayed highest ionic conductivities at low temperatures, as may be expected
from the reduced viscosity and freezing point of these compounds. At nearly all
temperatures, the electrolyte containing ethyl butyrate exhibited the highest conductivity
values (Fig. 2.9). Only at room temperature did the baseline electrolyte display superior
results. At all temperatures, the fluorinated ester containing electrolytes exhibited the
poorest conductive properties. The baseline mixture responded most dramatically to
41
changes in temperature. At room temperature and moderately depressed temperatures,
the carbonate only mixture exhibited good conductivity, giving results nearly to the
electrolyte containing the low viscosity, low freezing point ester. As the temperature was
depressed, conductivity values of the baseline mixture sharply declined approaching that
of the electrolyte containing the fluorinated ester (Table 2.5). At –60
o
C, the conductivity
was observed to vary with the following trend EC + EMC + EB (20:60:20 v/v%) (0.342
mS/cm) > EC + EMC + TFEB (20:60:20 v/v%) (0.196 mS/cm) > EC + EMC (20:80
v/v%) (0.195 mS/cm). Examining these results lead to the conclusion that any beneficial
results the fluorinated ester may contribute to the overall cell performance most likely
comes from favorable film forming properties and not improved mass-transfer effects.

Fig. 2.9 Conductivity comparison of baseline and ester containing electrolytes.

42
Table 2.5 Conductivity values at depressed temperatures. Values are given in mS/cm.

2.3.2.2 Cyclic Voltammetry of Fluoroester Electrolytes

Cyclic voltammetric measurements were made to determine the oxidative and
reductive stability of solutions containing the ester co-solvents relative to the baseline
formulations. 1.0 M LiPF
6
electrolyte mixtures were measured in a cell using a platinum
working electrode and lithium metal counter and reference electrodes. At high potentials
it was generally found that the addition of fluoroesters to carbonate mixtures reduced the
oxidative stability of the electrolyte (Fig. 2.10). When greater amounts of fluoroester co-
solvents were incorporated into the solvent mixtures, lower oxidative stability of the
electrolytes were observed. Of the fluoroester containing electrolytes, the mixture
containing 20% TFEB exhibited the greatest oxidative stability at higher potentials,
behaving similarly to the baseline at potentials greater than 4.0V.

Electrolyte Type
1.0M LiPF
6
in EC.EMC
(20.80 v/v%)
1.0M LiPF
6
in
EC.EMC.TFEB
(20.60.20 v/v%)
1.0M LiPF
6
in
EC.EMC.EB
(20.60.20 v/v%)
25 8.7 6.83 8.59
10 6.2 5.18 6.26
0 4.89 4.11 5.08
-10 3.7 3.12 3.99
-20 2.62 2.24 2.97
-30 1.685 1.481 2.04
-40 0.975 0.883 1.29
-50 0.487 0.457 0.72
-60 0.195 0.196 0.342
Temperature
43

Fig. 2.10 Cyclic voltammogram of fluoroester containing electrolytes from 2.0 to 6.0 V.
Electrolytes containing fluorinated esters exhibit reduced oxidative stability compared to
the baseline.

At reductive potentials within the operating potentials of the cell, none of the
fluorinated esters exhibited the electrochemical stability of the baseline (fig. 2.11). This
result is consistent with the findings published by Nakajima and coworkers, who
demonstrated that the fluorinated esters were reduced at higher potentials than
carbonates.
11
It is believed that their higher reductive potentials may lead to intriguing
film forming properties of electrolytes that contain fluorinated ester.

44

Fig. 2.11 Cyclic voltammogram of fluoroester containing electrolytes from 0.1 to 4.0 V.
Electrolytes containing fluorinated esters exhibited reduced reductive stability compared
to baseline.

2.3.2.3 Electrochemical Impedance Spectroscopy

Generally, the effects of the cathode dominate the interfacial impedance of the
total cell. These effects have been noted in previous reports by JPL and others and are
supported by our Tafel polarization studies of the same cells where it is demonstrated that
the MCMB anodes typically sustain higher currents than their LiNiCoO
2
cathodic
counterparts.
1b,17
It is also partly attributable to the cathode/anode balance with the
corresponding electrode loadings.
To make comparison of the co-solvent performance more straightforward, 20%
solutions of the fluorinated esters were used for analysis. EIS data were obtained at open
45
circuit voltages of fully charged cells over a frequency range of approximately 10
5
Hz to
~5 mHz, generated with an ac amplitude of 10 mV. The impedance pattern typically
contained two relaxation loops at room temperature, one at high frequency in the range of
10
5
to 10
2
Hz and the other in the low frequency range, in the range of 10
2
to 10
-1
Hz.
Two equivalent circuits were evaluated to interpret the data collected at low
temperatures collected data, R(CR)(CR) and R(CR)(C(RW)). The two relaxation loops
of each of these circuits were composed of a series resistance, R
s
, which represents an
algebraic sum of the electronic resistance from the electrodes, leads and the electrolyte
ionic resistance, a parallel resistor-capacitor network (C
hf
and R
hf
) in series for the high
frequency relaxation loop representing the film resistance-capacitance of the individual
electrodes. Another resistor-capacitor parallel network in series for the low frequency
relaxation loop, (C
lf
and R
lf
) explained the resistance-capacitance of the charge-transfer
mechanism. Additionally, the second evaluated equivalent circuit contained a Warburg
impedance term (W), to represent the solid-state diffusion of lithium inside the cathode or
anode. The impedance data were analyzed using the above equivalent circuits with Z
Simwin modeling software from Princeton Applied Research.
EIS of Cathodes
As it can be seen from Fig. 2.12, the nature of the electrolyte greatly affected the
magnitude of the first and second relaxation loops of the cathode Nyquist plots
suggesting that electrolyte solvent played a significant role determining the interfacial
properties at the cathode, affecting both the film and charge transfer resistances. For
example, the cathode film resistance for the EC:EMC 20:80 v/v% baseline was 0.08289
Ω while the electrolytes containing fluorinated esters varied from 0.0665 Ω (20% MPFP)
46
to 0.9049 Ω (40% TFEA). The 20% MPFP electrolyte was the only electrolyte that
demonstrated lower film impedance than the baseline. This seems to suggest that the
fluorinated materials degrade more readily on the cathode surface during the formation
process to form a resistive film. The charge transfer impedance for these electrolytes at
room temperature also varied significantly between the cells. Values range from 0.01892
Ω (5% TFEB) to 0.439 Ω (40% TFEB).
As the temperature of the cell is lowered, the values for both the film and charge-
transfer impedance increased, though not nearly at the same rate. The charge-transfer
impedance increases at a much greater rate than the film impedance. Table 2.6 compares
the cathodes of all the fluorinated ester electrolytes at 20 v/v% compared to the baseline.
Of these, the 20% TFEB maintains the lowest impedance at reduced temperatures which
may explain some of its improved discharge performance at reduced temperatures
compared to the other fluorinated ester electrolytes may arise from its decreased
impedance at this electrode leading to better cell performance at /.these temperatures.

47
Fig. 2.12 Nyquist plot of electrochemical impedance spectroscopy (EIS) measurements at
23ºC of LiNi
x
Co
1-x
O
2
electrodes from lithium-ion cells containing electrolytes consisting
of 1.0M LiPF
6
EC+EMC+X (20:60:20 v/v %) (where X = TFEB, ETFA, TFEA, TFEP.
and MPFP).

0.00
0.30
0.60
0.90
1.20
1.50
0.00 0.50 1.00 1.50
Z' (Ohms)
1.00 M LiPF6 EC+EMC+TFEB (20:60:20 v/v %)
1.00 M LiPF6 EC+EMC+ETFA (20:60:20 v/v %)
1.00 M LiPF6 EC+EMC+TFEA (20:60:20 v/v %)
1.00M LiPF6 EC:EMC:TFEP (20:60:20 v/v%)
1.00 M LiPF6 EC+EMC+MPFP (20:60:20 v/v %)
1.00 M LiPF6 EC+EMC (20:80 v/v %)
Temperature = 23
o
C
Li-ion Experimental Three Electrode Cell
MCMB Carbon - LiNi
x
Co
1-x
O
2
Lithium metal Reference Electrode
48

Table 2.6 Cathode LiNi
x
Co
1-x
O
2
Impedance Data. Baseline EC:EMC electrolyte solvent compared to solvents containing 20%
of various fluorinated ester. Values given in Ohms.

49
EIS of Anodes
As with the reported cathode data, electrolyte composition played a vital role in
influencing both of the relaxation loops that comprise the Nyquist impedance plots (Fig.
2.13). All of the electrolytes containing fluorinated esters at 20 v/v% exhibited some
deleterious effects at the anode at room temperatures. Film resistances for all fluoroester
cells increased relative to the EC:EMC baseline, although the cells that contain TFEB,
TFEA, and TFEP come close suggesting the structure of these electrolytes heavily
influences the degree to which these materials decompose at the anode. These same
fluorinated esters also perform better with regards to the charge-transfer impedance (the
second of the two relaxation loops).
As the temperature of the cells is reduced the values for both the film and charge-
transfer impedance values increase (Table 2.7), but at rates that attest to the influence of
electrolyte solvent structure has on the decomposition of the individual electrolytes. The
three (2,2,2-trifluoroethyl) esters had anode impedances that grow at relatively similar
rates as the temperature was decreased to -50ºC. Between the temperatures of 0ºC and -
50ºC, these electrolytes have film and charge-transfer resistance values that are close to,
and in some cases lower than the values collected for the baseline electrolyte under the
same conditions. At -60º the skewed values probably have more to do with the
mathematical calculations performed by the simulation software than the actual
impedance of the electrodes, and cannot be trusted to be a completely accurate
description of the impedance of the electrodes.
50
Fig. 2.13 Nyquist plot of electrochemical impedance spectroscopy (EIS) measurements at
23ºC of MCMB anodes from lithium-ion cells containing electrolytes consisting of 1.0M
LiPF
6
EC+EMC+X (20:60:20 v/v %) (where X = TFEB, ETFA, TFEA, TFEP. and
MPFP).

The electrolytes composed of 20% ETFA and MPFP grow at a much faster rate as
the temperature is decreased compared to the 2,2,2-trifluoroethyl varieties, or begin with
a much higher inherent values. This inherent high impedance at room temperature is the
reason that limited low temperature characterization did not occur.

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' (Ohms)
1.00 M LiPF6 EC+EMC+TFEB (20:60:20 v/v %)
1.00 M LiPF6 EC+EMC+ETFA (20:60:20 v/v %)
1.00 M LiPF6 EC:EMC:TFEP (20:60:20 v/v %)
1.00 M LiPF6 EC+EMC+TFEA (20:60:20 v/v %)
1.00 M LiPF6 EC+EMC+MPFP (20:60:20 v/v %)
1.00 M LiPF6 EC+EMC (20:80 v/v %)
Temperature = 23
o
C
Li-ion Experimental Three Electrode Cell
MCMB Carbon - LiNi
x
Co
1-x
O
2
Lithium metal Reference Electrode
51

Table 2.7 Anode MCMB Impedance Data. Baseline EC:EMC electrolyte solvent compared to solvents containing 20% of various
fluorinated ester. Values given in Ohms.

52
EIS of Full Cells
Typically, the full cell impedance trends were predicted from those found in
individual electrode trends (Fig. 2.14). As previously noted, the 2,2,2-trifluoroethyl-
based electrolytes had lower impedance than other fluoroester electrolytes evaluated at
low temperatures (Table 2.8). Although it is difficult to draw exact conclusions solely
from full cell impedance data, when collected in tandem with individual electrode
impedance data, it is apparent a large portion of the impedance for the entire cell comes
from the impedance growth on the anode for fluoroester containing electrolyte solutions,
unlike some data that was previously reported. This information is consistent with CV
data that shows the fluorinated esters degrade at working anode potentials.

Fig. 2.14 Nyquist plot of electrochemical impedance spectroscopy (EIS) measurements
at 23
o
C of lithium-ion cells containing electrolytes consisting of 1.0M LiPF
6

EC+EMC+X (20:60:20 v/v %) (where X = TFEB, ETFA, TFEA, TFEP, and MPFP).
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' (Ohms)
1.00 M LiPF6 EC+EMC+TFEB (20:60:20 v/v %)
1.00 M LiPF6 EC+EMC+ETFA (20:60:20 v/v %)
1.00 M LiPF6 EC+EMC+TFEA (20:60:20 v/v %)
1.00 M LiPF6 EC:EMC:TFEP (20:60:20 v/v %)
1.00 M LiPF6 EC+EMC+MPFP (20:60:20 v/v %)
1.00 M LiPF6 EC+EMC (20:80 v/v %)
Temperature = 23
o
C
Li-ion Experimental Three Electrode Cell
MCMB Carbon - LiNi
x
Co
1-x
O
2
Lithium metal Reference Electrode
53

Table 2.8 Full Cell Impedance Data. Baseline EC:EMC electrolyte solvent compared to solvents containing 20% of various
fluorinated ester. Values given in Ohms.

54
2.3.2.4 Tafel Measurements

Evaluation of the lithium intercalation/de-intercalation electrode kinetics was
collected using Tafel polarization measurements of the MCMB-Li
x
Ni
y
Co
1-y
O
2
cells
employing the described electrolyte formulations. The measurements were performed on
the cells containing 20 v/v% fluoroester while they were fully charged (OCV= ~4.095V)
at temperatures ranging from room temperature down to -40ºC. The data was generated
under potentiodynamic conditions, but at slow scan rates, approximating steady-state
conditions.
Measurement of the anode at room temperature showed two of the fluorinated
ester formulations, those containing TFEB and TFEA, perform with slightly improved
kinetics over the baseline while another containing TFEP performed similarly (Fig. 2.15).
The electrolyte mixture containing ETFA demonstrated substantially worse kinetics than
all other electrolytes tested on the anodes. At the cathode, lithium intercalation is most
rapid for the baseline mixtures, though electrolytes containing TFEB and TFEP
performed nearly as well. Both ETFA and TFEA containing mixtures demonstrated
considerably poorer kinetic capabilities (Fig. 2.16). In all cases, the limiting current
densities of the anodes were much greater (≥2x) than they were for the cathode under
identical conditions, consistent with previous findings indicating cathode-limited design
of these cells.

55

Fig 2.15 Tafel polarization measurements at 23
o
C of MCMB electrodes from lithium-ion
cells containing electrolytes consisting of 1.0M LiPF6 EC+EMC+X (20:60:20 v/v %)
(where X = TFEB, ETFA, TFEA, and TFEP).

Fig. 2.16 Tafel polarization measurements at 23
o
C of LiNi
x
Co
1-x
O
2
electrodes from
lithium-ion cells containing electrolytes consisting of 1.0M LiPF6 EC+EMC+X
(20:60:20 v/v %) (where X = TFEB, ETFA, TFEA, and TFEB).
56

Fig. 2.17 Tafel polarization measurements at - 40
o
C of MCMB electrodes from lithium-
ion cells containing electrolytes consisting of 1.0M LiPF
6
in EC+EMC+X (20:60:20 v/v
%) (X = TFEB, TFEA, and TFEP).

As expected, both the anode and cathode lithium intercalation/de-intercalation
kinetics are reduced as the temperature of the cells is lowered. On the anode side, the
greatest depreciation in performance takes place qualitatively with TFEA electrolyte. At
room temperature it is one of the highest performing; however, upon reduction in
temperature it becomes the worst kinetic performance anode (Fig. 2.17). It should be
noted that the cell with ETFA as a co-solvent was not tested at such reduced temperatures
due to presumed poor performance and the TFEA anode likely would have outperformed
it from a kinetic standpoint at these temperatures. At the cathode the TFEB and TFEA
electrolytes perform within 10% of the baseline mixture at -40ºC (Fig. 2.18).

57

Fig. 2.18 Tafel polarization measurements at -40
o
C of LiNi
x
Co
1-x
O
2
electrodes from
lithium-ion cells containing electrolytes consisting of 1.0M LiPF
6
EC+EMC+X (20:60:20
v/v %) (where X = TFEB, ETFA, TFEA, and TFEB).

Fig. 2.19 Tafel polarization measurements at - 60
o
C of MCMB electrodes from lithium-
ion cells containing electrolytes consisting of 1.0M LiPF
6
EC+EMC+X (20:60:20 v/v %)
(where X = TFEB and TFEP).
58

Fig. 2.20 Tafel polarization measurements at - 60
o
C of LiNi
x
Co
1-x
O
2
electrodes from
lithium-ion cells containing electrolytes consisting of 1.0M LiPF
6
EC+EMC+X (20:60:20
v/v %) (where X = TFEB and TFEP).

Further depressing the temperature to -60º provides interesting results. Although
only the fluoroester electrolytes tested are those with TFEB and TFEP, both outperform
the baseline on the anode by a considerable amount (Fig. 2.19). At the cathode at -60ºC
TFEB electrolyte exhibits the least impressive kinetics overall (Fig. 2.20).

2.3.2.5 DC Micro-Polarization Measurements

DC micro-polarization tests were performed on the cells to study the charge
transfer behavior of the passivating films affected by the addition of fluorinated esters
and low temperature on the cells. The polarization resistance values of the electrodes
59
were calculated from the slopes of the linear plots generated under potentiodynamic
conditions of scan rates of 0.2 mV/s and measured over a +/-5mV from the equilibrium
potential of each electrode. Polarization values of cells using fluorinated ester
electrolytes increased in response to reduced temperatures in a manner similar to the
baseline (Tables 2.9 and 2.10). The TFEB cell anode polarization values increase most
slowly for the cells examined, giving results even more favorable than the baseline. The
TFEB cell has an anomalous value at -40º, which was disregarded in this study. The
baseline performs better than all fluoroester cells at the cathode. Results support the claim
that TFEB is the best performing of the fluoroesters at low temperatures.

Table 2.9 DC micro-polarization measurements at various temperatures of anode from
lithium-ion cells containing electrolytes consisting of 1.0M LiPF
6
EC+EMC+X (20:60:20
v/v %) (where X = TFEB, ETFA, TFEA, TFEP).

Table 2.10 DC micro-polarization measurements at various temperatures of cathode from
lithium-ion cells containing electrolytes consisting of 1.0M LiPF
6
EC+EMC+X (20:60:20
v/v %) (where X = TFEB, ETFA, TFEA, TFEP).
Room Temp At 0ºC At -20ºC At -30ºC At -40ºC At -50ºC At -60ºC
0.621 0.869 2.907 10.000 139.860 195.652 368.421
1.382 1.351 2.825 12.225 27.027 190.476 327.103
1.664 5.236 33.333 108.696 303.030 900.000 2000.000
0.329 0.939 3.327 14.660 36.000 318.792 1055.556
0.302 0.517 3.546 8.382 35.842 57.143 138.675
1.168 1.675 1.779 21.845 36.101 0.000 630.631
1.0M LiPF
6
in EC.EMC.TFEA
(20.60.20)
1.00 M LiPF
6
EC+EMC+TFEB
(20:60:20)
1.0M LiPF
6
in EC.EMC.TFEP
(20.60.20)
1.0 M LiPF 6 EC+DEC+DMC
(1:1:1)
1.0 M LiPF
6
EC+EMC+ETFA
(20:60:20)
1.0M LiPF
6
in EC.EMC
(20.80)
Cathode Linear Polarization Data
Room Temp At 0ºC At -20ºC At -30ºC At -40ºC At -50ºC At -60ºC
0.353 0.724 2.174 7.387 113.233 55.556 280.000
0.963 1.334 5.476 * 38.508 130.137 475.000
0.507 1.034 3.507 13.194 25.000 105.556 416.667
0.357 1.343 2.819 11.059 24.611 99.581 416.667
0.349 1.014 2.987 13.552 27.378 114.458 492.228
0.479 0.918 2.738 12.516 29.412 * 1000.000
1.00 M LiPF
6
EC+EMC+TFEB
(20:60:20)
1.0 M LiPF
6
EC+EMC+ETFA
(20:60:20)
1.0M LiPF
6
in EC.EMC.TFEA
(20.60.20)
1.0M LiPF
6
in EC.EMC.TFEP
(20.60.20)
1.0M LiPF
6
in EC.EMC
(20.80)
1.0 M LiPF
6
EC+DEC+DMC
(1:1:1)
Anode Linear Polarization Data
60
2.4 Conclusions

The cell performance characteristics of numerous Li-ion cells with various
electrolyte mixtures containing fluorinated ester co-solvents were extensively studied and
reported within. A three-electrode configuration with MCMB anodes, LiNi
0.8
Co
0.2
O
2

cathodes and Li-metal reference electrodes was utilized within these cells. The discharge
characteristics were extensively studied over a range of temperatures with special
attention given to the low temperature characteristics of these cells. In addition to the
discharge characteristics of the cells, we examined the electrochemical properties of the
individual electrodes using electrochemical impedance spectroscopy, Tafel polarization
techniques and DC micro-polarization methods. These methods give us insight into
electrochemical characteristics such as film formation properties, lithium
intercalation/de-intercalation kinetics, and polarization resistances imparted onto the
electrodes by the various investigated electrolytes.
1.0 M LiPF
6
solutions containing 5-60 v/v% of numerous fluorinated esters were
evaluated. Of the investigated fluorinated ester solutions, the fluorinated ester co-solvent
found to impart the most desirable properties upon the cell was 2,2,2-trifluoroethyl
butyrate (TFEB). The electrolyte mixture of 1.0 M LiPF
6
in ethylene carbonate/ethyl
methyl carbonate/TFEB (20.60.20 v/v%) was found to exhibit the greatest capacity
retention at reduced temperatures and various discharge rates. This electrolyte was the
best of the tested electrolytes at reduced temperatures regardless of the charging
temperature. Reasons for this improved performance were elucidated using the
electrochemical methods described above. The films produced at the anode due to the
61
addition of TFEB were less resistive at low temperatures than those of the baseline
carbonate-only solutions. The electrolyte with TFEB imparted improved lithium
intercalation/de-intercalation kinetics at both electrodes at many temperatures tested.
Polarization resistance growth of the electrodes was also reduced on the cell containing
TFEB and confirmed by DC micro-polarization analysis.

62
2.5 Chapter 2 References

1. (a) Smart, M. C.; Ratnakumar, B. V.; Surampudi, S. J. Electrochem. Soc. 2002,
149, A361. (b) Smart, M. C.; Ratnakumar, B. V.; Whitcanack, L. D.; Chin, K. B.;
Surampudi, S.; Croft, H.; Tice, D.; Staniewicz, R. J. Power Sources 2003, 119,
349. (c) Smart, M. C.; Whitacre, J. F.; Ratnakumar, B. V.; Amine, K. J. Power
Sources 2007, 168, 501.
2. (a) Surampudi, S.; Smart, M. C.; Ratnkumar, B. V.; Huang, C.-K. U.S. Patent
6,492,064, December 10, 2002. (b) Smart, M. C.; Ratnakumar, B. V.; Surampudi,
S. J. Electrochem. Soc., 1999, 146, 486.
3. Ohta, A.; Koshina, H.; Okuno, H.; Murai, H. J. Power Sources, 54 (1), 6-10,
(1995).
4. (a) Herreyre, S.; Huchet, O.; Barusseau, S.; Perton, F.; Bodet, J. M.; Biensan, P.
J. Power Sources, 2001, 97-98, 576. (b) Herreyre, S. et al. U. S. Patent
6,399,255, June 4, 2002.
5. Shiao, H. –C.; Chua, D.; Lin, H. –P.; Slane, S.; Solomon, M.; J. Power Sources.,
2000, 87, 167.
6. Sazhin, S. V.; Khimchenko, M. Y.; Tritenichenko, Y. N.; Lim, H. S. J. Power
Sources, 2000, 87, 112.
7. (a) Robin, M. L. in Hudlicky, M.; Pavlath, A.E. (Eds.) Chemistry of Organic
Fluorine Compounds. II. A Critical Review, ACS Monograph 187, American
Chemical Society, Washington, D.C., 1995. (b) Banks, R. E.; Tatlow, J. C. in
Banks, R. E. (Ed.) Organofluorine Chemistry: Principles and Commercial
Applications, Plenum Press, New York, 1994.
8. (a) Shu, Z. X.; McMillan, R. S.; Murray, J. J.; Davidson, I. J. J. Electrochem.
Soc., 1996, 143, 2230. (b) Shu, Z. X.; McMillan, R. S.; Murray, J. J.; Davidson, I.
J. J. Electrochem. Soc., 1995, 142, L161.
9. Nakajima, T.; Dan, K.; Koh, M.; Ino, T.; Shimizu, T. J. Fluorine Chemistry,
2001, 111, 167.
10. Arai, J. J. Electrochem. Soc., 2003, 150, A219.
11. Moller, K. –C.; Hodal, T.; Appel, W. K.; Winter, M.; Besenhard, J. O. J. Power
Sources, 2001, 97-98, 595.
12. McMillan, R.; Slegr, H.; Shu, Z. X.; Wang, W. J. Power Sources, 1999, 81-82,
20.
63
13. (a) Inaba, M.; Kawatate, Y.; Funabiki, A.; Jeong, S. –K.; Abe, T.; Ogumi, Z.
Electrochim. Acta, 1999, 45, 99. (b) Katayama, H.; Arai, J.; Akahoshi, H. J.
Power Sources, 1999, 81-82, 705. (c) Arai, J.; Katayama, H.; Akahoshi, H. J.
Power Sources, 2002, 149, A217.
14. Smart, M. C.; Ratnakumar, B. V.; Ryan-Mowrey, V. S.; Surampudi, S.; Prakash,
G. K. S.; Hu, J.; Cheung, I. J. Power Sources, 2003, 119-121. 359.
15. Nakajima, T.; Dan, K.; Koh, M. J. Fluorine Chemistry, 1998, 87, 221.
16. Yamaki J. –I.; Yamazaki, I.; Egashira, M,; Okada, S. J. Power Sources, 2001,
102, 288.
17. (a) Smart, M. C.; Lucht, B. L.; Ratnakumar, B. V. J. Power Sources, 2008, 155,
A557. (c) Nagasubramanian, G.; Jungst, R. G.; Doughty, D. H. J. Power Sources,
1999, 83, 193. (d) Li, J.; Murphy, E.; Winnick, J.; Kohl, P. A. J. Power Sources,
2001, 102, 294.

64
CHAPTER 3

Improving the Safety Characteristics of Li-Ion Cells via the
Use of Flame-Retardant Additives

3.1 Introduction

With their high working voltages and energy density, long cycle life, and low
self-discharge rates compared to other secondary battery technologies, lithium-ion
batteries are an attractive option for applications requiring high power. While these
batteries have been implemented into many small portable electronic applications such as
laptop computers and cell phones, they have yet to see significant incorporation into
many applications with higher power requirements, specifically electric vehicles (EV)
and hybrid electric vehicles (HEV). The main reasons for their slow inception into these
applications are the safety concerns for such batteries upon scale up required for EV use.
When exposed to high temperature, overcharge, over-discharge, and short circuit
conditions, not only is cell failure a distinct possibility, the conditions may lead fire or
even explosion. Therefore, it is critical to develop new methods for improving the safety
and usability of lithium-ion batteries before these batteries can be incorporated into a
wider range of applications. One method for improving the safety characteristics of Li-
ion cells is via the incorporation of flame-retardant additives into the electrolyte medium,
which serve to reduce the flammability of the organic electrolyte. Herein, numerous
65
studies were performed to monitor the effects these additives have on not only the
discharge characteristics of such cells at various rates and temperatures, but also the
electrochemical effects the additives produce on the individual electrodes of the cells.
Numerous studies were performed to more fully understand the effects of the additives.
In the first study, four flame retardant additives were combined with a fluorinated ester.
This is done to take advantage to the inherent reduction in flammability present in
halogenated materials. The second study combined a flame retardant additive with an
SEI enhancing agent in an attempt to reduce the deleterious effects that flame retardant
additives impart on Li-ion cell performance. The third and final study investigated the
effect of substitution pattern at the phosphorus atom, oxidation state of the phosphorus
atom and fluorinated substitution of the branches of the flame retardant molecules.

3.1.1 Safety Issues of Li-Ion Batteries

Insuring the safety of Li-ion batteries is one of the most important areas in the
development of such devices. Early safety issues stemmed from the use of a Li-metal
anode, but even after lithium-intercalating graphitic materials were substituted for
metallic lithium, Li-ion cells continued to experience growing pains to overcome the
safety issues inherent in dealing with both energetic devices and lithium. Numerous
examples from the early days of these cells exist that speak to the dangers these materials
can present. In 1995, numerous lithium-ion batteries were overcharged during in-house
testing at Apple causing the company to remove all lithium-ion power packs from their
product line at the time. Around same time, Ericsson moved away from the batteries for
66
mobile phones and other portable devices. In 2000, Dell voluntarily recalled 27,000 Li-
ion batteries and Compaq, 55,000 batteries from their notebook computers. EV Global
Motors Company recalled 2,000 batteries from their electric bicycles the same year. In
2006, Dell issued another recall, this one even larger, of 4.1 million notebook computer
batteries because of the possibility of the devices erupting in flames. Apple recalled
128,000 Li-ion batteries from their line of notebook computers the same year for similar
reasons. 2006 also saw several incidents of fire on airplanes that were linked to the
batteries leading to a ban on packing spare Li-ion batteries in checked luggage of
airplanes.
Strict luggage requirements, withdrawal of products, loss of market share, and
outright bans on Li-ion batteries are examples of some of the backlash prompted by such
incidents. However, due to their high energy density, excellent cycle life, and lack of
memory effect Li-ion batteries are still viewed as the future of the battery industry, with
the most recent figures have shown revenues of $5.89 billion on 1.76 billion units and
estimates that this figure will grow to 3.19 billion units by 2013.
1
Therefore,
improvement of safety characteristics while maintaining their desirable properties
remains of paramount importance.
Li-ion batteries pose their greatest safety risks when they are subject to abuse
conditions, such as exposure to high temperatures, short circuiting (internal or external),
or crushing. These conditions can lead to spontaneous heat-evolving reactions involving
the highly oxidizing and reducing materials present in the electrodes and highly
flammable electrolyte solvents compounded by a cell design that is ill fitted for heat
dispersal.
67
A closer inspection of Li-ion cell operation reveals context to their safety risks.
The cells operate at potentials (~4 V) beyond the thermodynamic stability window of
electrolytes. This leads to decomposition of the electrolytes at the charged electrodes.
Additionally, the interface at the cathode is complicated by partial dissolution of the
positive active materials. At high potentials and elevated temperatures this leads to
accelerated electrolyte oxidation.
Thermal runaway reactions pose a significant threat to the safety characteristics of
any Li-ion system. These processes occur when exothermic chemical reactions take
place in the cell without the cell being properly able to dissipate the heat causing the
cell’s internal temperature to continue to rise and accelerate the chemical reactions.
These reactions can cause mechanical failures in the cell, which may lead to short
circuits, premature failure of the cell, distortion, swelling, explosion of the cell and
ignition of the electrolyte. Because the ability to dissipate heat is tied to overall mass of
the cell, which is inversely related to energy density, low thermal capacity is an
unavoidable penalty of higher energy density batteries.
Thermal (overheating), electrical (overcharge, high pulse power) or mechanical
(crushing, internal or external short circuit) abuse conditions can lead to thermal runaway
in Li-ion cells. Additionally, exothermic reactions within the cell can lead to thermal
runaway as well.
2
These reactions include thermal decomposition of the electrolyte,
reduction of the electrolyte at the anode, oxidation by the cathode, thermal decomposition
of the anode or cathode, and melting of the separator material that can cause an internal
short.
2a

68
3.1.2 Approaches to Improve the Safety of Li-ion Batteries

In order make Li-ion batteries safer and more feasible for consumer use,
numerous mechanisms have been developed to improve the safety characteristics of such
devices. Balakrishnana et al. sorted these mechanisms into 6 categories.
3
: (1)
conventional safety devices, (2) self-resetting devices, (3) shutdown separators, (4) active
materials, (5) coatings, and (6) electrolytes. A brief introduction into each of these is
given below.

1. Conventional safety devices- These devices terminate the flow of
electricity in cells where safety limits are exceeded. Examples include
excess current or temperature. Devices in this category include traditional
thermal fuses, circuit breakers, overcharge/discharge protection systems,
and safety vents for release of pressure within the cell. Devices in this
category are external to the Li-ion cell and thereby reduce the energy
density and increase the manufacturing costs associated with Li-ion cells.
Additionally, after the devices are triggered, the devices or the entire
battery must be replaced introducing additional costs.
2. Self-resetting devices- Self-resetting, fuse-like devices make attractive
alternatives because of the inconvenience of resetting conventional safety
devices. These positive temperature coefficient devices have based on
materials whose resistances increase dramatically with a rise in
temperature. These devices also protect against large current flows
69
because Joule heat evolution occurs during at high current flux. Devices
in this category are based upon conductive polymers that become
amorphous above their glass transition temperatures. These devices
become economically attractive only when used in costly equipment or
long-term warranties are demanded.
3. Shutdown separators- Polyethylene and polypropylene microporous
separator films are used to keep the electrodes from contacting one
another while allowing ionic flow between the cathode and anode.
Shutdown separators respond to abnormal rise in temperature typically
caused by electrical by closing micropores shutting down the
electrochemical reaction. The drawback for these safety devices is that the
shutdown is irreversible and once actuated, the melted separators render
the battery useless.
4. Active materials- the stability of commercial Li-ion batteries is severely
compromised above 60
0
C.
4
Above these temperatures, reactions between
the electrolyte and the active electrode materials occur. Anode and
electrolyte reactions occur initially,
5
with cathode/electrolyte reactions
dominating heat-evolution processes at higher temperatures.
6
Because the
graphitic structure of carbon must be maintained in the anodes, little work
can be done to change the nature of the materials used; however,
significant work has been performed to investigating the safety
characteristics of cathode materials. The most promising of the new
70
cathode materials is LiFePO
4
, which exhibits low reactivity with
electrolytes and no heat evolution below 200
o
C.
6

5. Coatings- Several electrode coating methods have been developed to
increase the stability between the electrolyte and electrodes. Chemical
additives, such as γ-butyrolactone, encapsulate the cathode upon
decomposition reducing direct reaction with electrolytes.
7
Other coating
methods include coating carbon on graphite anodes to suppress propylene
carbonate decompositions
8
and micro-encapsulation of graphite with
nanosized nickel-composite particles to diminish solvent co-intercalation,
subsequent exfoliation and gas evolution.
9

6. Electrolytes- Lithium is intrinsically unstable with the most common
electrolytes. For example, alkyl carbonates are known to react with
lithiated graphite and delithiated cathode materials. Much research has
been performed to develop new means to improve safety characteristics of
Li-ion electrolytes to suppress these reactions.
5,6,10
These reactions are
exacerbated at elevated temperatures when the SEI on the graphite anode
begins to decompose, allowing rapid and direct reaction with the
electrode. Delithiated cathodes are highly oxidizing and react
exothermically under these conditions as well. To reduce the likelihood of
the reactions between the electrode active materials and the electrolyte,
numerous methods of improving electrolyte safety have been developed.
Some methods include redox shuttles, which provide overcharge
protection by being reversibly oxidized at potentials slightly higher than
71
normal end-of-charge potentials.
11
Shutdown additives which act
similarly, but shut the cell down permanently when activated.
12
Ionic
liquids have also been proposed to enhance the safety characteristics of Li-
ion batteries due to their thermal stability, non-flammability, non-
volatility, and low heat of reaction with active materials.
12,13
Several new
lithium salts have been invented with safety characteristics greater than
LiPF
6
including LiPF
3
(C
2
F
5
)
3
(LiFAP),
14
LiN(SO
2
CF
3
)
2
(LiTFSI),
15
and
LiBC
4
O
8
(LiBOB).
16
Flame-retardant additives have also been added to
electrolytes to reduce some the dangers associated with cell abuse. Many
of these additives are phosphorus-containing compounds. Searching for
the additives that produce minimal performance sacrifices while
maintaining the desired safety characteristics is an area that has demanded
much focus.

3.1.3 Phosphorus Based Flame Retardant Additives

Suppression of the electrolyte flammability is common method used to improve
the safety characteristics of these batteries. These safety characteristics must be
improved before they can be incorporated into high power applications like hybrid
electric vehicles. To this end, numerous types of additives have been investigated for
their ability to impart flame-retarding characteristics on to lithium-ion cells. Nearly all
flame retardant additives (FRAs) contain phosphorus,
17-25
although substitution at the
72
phosphorus atom and oxidation state of the phosphorus varies substantially among
investigated compounds.
Aromatic phosphates,
17
alkyl phosphate relatives,
18
and compounds that contain
both aromatic and alkyl branches have been extensively investigated.
19
Among
phosphates that contain only hydrocarbon branches, those with smaller methyl and ethyl
branches exhibit better flame retarding properties; however, a tradeoff exists between
flame retarding capabilities and electrochemical stability.
17f, 18a,b
While their higher
phosphorus content makes short branched phosphates better suited to reducing electrolyte
flammability, it also causes reduced reductive stability upon the graphitic anode.
Halogenated phosphorus compounds have also been investigated. While
chlorinated phosphates have been studied,
20
fluorinated phosphates make up the bulk of
halogenated phosphates examined with tris(2,2,2-trifluoroethyl)phosphate as one of the
most promising FRAs found to date.
21
It has been reported that when this compound is
added to an electrolyte mixture at 20 wt.% complete non-flammability of the electrolyte
is achieved without sacrificing performance from the anode or cathode in the cell.
Other phosphorus compounds investigated as FRAs include phosphites.
4,22
The
P(III) oxidation state of the phosphorus is thought to possibly lead to improved
electrochemical and thermal stabilities of electrolytes via improved SEI facilitation and
PF
5
deactivation. A potential drawback of these materials is their reduced solubility in
many electrolyte solvent compounds. Phosphonates,
23
phosphoramides,
24
and
phosphazenes
25
have also been studied on a limited basis for use as FRAs within lithium-
ion batteries.

73
3.1.4 Flame Retardant Additives as Lewis Base Additives to Stabilize LiPF
6
Salt

As previously mentioned, LiPF
6
is presently the most popular choice among Li-
ion battery salts. Its popularity can be traced to a combination of well-balanced
properties with a relatively few compromises mixed in. The limitations of LiPF
6
include:
(1) a high equilibrium constant for the decomposition shown below which results in the
gaseous PF
5
being generated, and (2) high reactivity of the Lewis acid PF
5
with organic
solvents.
26

LiPF
6 PF
5
+ LiF

The generation of gaseous PF
5
drives the equilibrium towards the products,
especially at elevated temperatures and in the presence of moisture. The resulting PF
5

reacts with the electrolyte and SEI components to deteriorate the overall stability of the
system. Early commercialization attempts of LiPF
6
also misjudged the importance of
purity and shipped salts containing HF that magnified the deficiencies of the salt.
One of the means investigated to alleviate these problems has been to incorporate
weak Lewis bases into the electrolyte solution in small amounts. These compounds
stabilize the LiPF
6
salt by forming a weak 1:1 complex with PF
5
. Sufficiently weak base
must be chosen as not to drive the equilibrium of the LiPF
6
too far towards the products.
Example of such weak bases include tris(2,2,2-trifluoroethyl) phosphite
4
and amide-
based compounds.
27

74
3.2 Experimental Methods

Cell preparation and all performance and electrochemical tests were performed
according to details outlined in section 2.2 of this manuscript. Flame-retardant additive
candidate compounds were purchased from various sources and used without any further
purification. In cases where flame-retardant additive material is added to base electrolyte
mixtures, additional LiPF
6
is charged to bring the final concentration of the solution to
1.0M LiPF
6
.

3.3 Results and Discussion

While extensive work has been performed observing the capacity and
electrochemical properties of full Li-ion cells containing flame retardant additives, little
to no work has been performed studying the effects electrolytes containing these
phosphorus compounds have upon the individual electrode performance and
electrochemical characteristics within these cells. In this study, numerous variations of
electrolytes containing flame retardant additives (FRAs) were investigated. In the first
set of electrolytes investigated, four potential flame retardant additives were added to
electrolyte solutions containing the most promising fluoroester electrolyte formulation,
1.0M LiPF
6
in EC.EMC.TFEB (20.60.20 v/v%). It was thought that the employment of a
halogenated solvent might suppress the flammability of the electrolytes within the cells,
thereby decreasing the total amount of phosphorus FRA needed. This in turn should lead
to improved cell cyclability and performance. The four flame retardant additives chosen
75
for this study are triphenyl phosphate (TPhPh), tributyl phosphate (TBuPh), triethyl
phosphate (TEtPh), and bis(2,2,2-trifluoroethyl) methyl phosphonate (TFMPo). These
additives were selected to compare the effect of chain length, substitution around the
P(V) atom, and partial halogenation on electrochemical stability. These additives are
blended with EC, EMC and TFEB in the following ratios:

A. 1.0 M LiPF
6
in EC:EMC:TFEB:TPhPh (20:55:20:5 v/v%)
B. 1.0 M LiPF
6
in EC:EMC:TFEB:TBuPh (20:55:20:5 v/v%)
C. 1.0 M LiPF
6
in EC:EMC:TFEB:TEtPh (20:55:20:5 v/v%)
D. 1.0 M LiPF
6
in EC:EMC:TFEB:TFMPo (20:55:20:5 v/v%)

Of the cells tested in this study, the electrolyte employing TPhPh as an FRA
demonstrated the greatest long-term performance and cyclability at room temperature, an
unsurprising result given that previous studies had demonstrated aromatic phosphates to
exhibit greater electrochemical stability than their alkyl counterparst. Testing was also
performed at reduced temperatures where electrolytes containing the less viscous alkyl
phosphates outperformed bulkier the triphenyl phosphate-containing electrolyte. A
baseline mixture containing no flame-retarding additive was included with the study for
ease of comparison.
In the second series of FRA tests, the effects of the most promising flame
retardant additive, TPhPh, were monitored against a standard carbonate only EC:EMC
(20:80 v/v%) baseline. Additionally, the film forming additive, vinylene carbonate (VC),
was added to a cell containing TPhPh to observe if deleterious effects imparted on the
76
performance of the cell by the flame retardant additive can be reduced by use of such an
additive. These additives were blended with carbonate-based electrolytes in the following
ratios:
A. 1.0 M LiPF
6
in EC:EMC (20:80 v/v%)
B. 1.0 M LiPF
6
in EC:EMC (20:80 v/v%) w/ 1.5% VC
C. 1.0 M LiPF
6
in EC:EMC:TPhPh (20:75:5 v/v%)
D. 1.0 M LiPF
6
in EC:EMC:TPhPh (20:75:5 v/v%) w/ 1.5% VC

As anticipated, improved cycle life was imparted onto cells containing the flame
retardant additive when VC was included as part of the electrolyte composition. The
reason for this improved performance comes from the protective layer VC forms upon
the anode when it is reduced at working anode potentials, thereby protecting the flame
retardant additive from direct contact with the anode.
A third and final series of flame retardant additives was tested in Li-ion cells.
These additives compose of phosphorus base flame-retardants that are chosen to
demonstrate the effect of substitution pattern at the phosphorus atom, oxidation state of
the phosphorus atom, and fluorine substitution on the branches of the FRA molecule
around the phosphorus atom upon the performance of a Li-ion cell. The flame retardant
additives examined were tris(2,2,2-trifluoroethyl) phosphate (TFPa), tris(2,2,2-
trifluoroethyl) phosphite (TFPi), triphenylphosphite (TPPi), diethyl ethyl phosphonate
(DEP), and diethyl phenyl phosphate (DPP). The additives were added to carbonate
electrolytes in the following ratios:

77
A. 1.0 M LiPF
6
in EC:EMC (20:80 v/v%)
B. 1.0 M LiPF
6
in EC:EMC (20:80 v/v%) w/ 1.5% VC
C. 1.0 M LiPF
6
in EC:EMC:TPhPh (20:75:5 v/v%)
D. 1.0 M LiPF
6
in EC:EMC:TPhPh (20:75:5 v/v%) w/ 1.5% VC

To monitor the influence each additive imparted onto the system electrochemical
measurements on the individual electrodes of the cell, low temperature testing and 100
cycle life testing were performed. It was found that structure of the additive influenced
cell performance to a considerable degree. Throughout the low temperature testing of the
cells, electrolytes containing flame retardant additives with small branches exhibited the
best performance. Cycle testing performed at room temperature demonstrated
electrochemical stability of the additives imparts the greatest influence upon the cycle life
of the cells. It is also observed that the electrolytes that employ the use of phosphites
(where the P is in the +3 valence state) exhibit improved performance in agreement with
previous findings demonstrating these molecules can act as Lewis-base scavengers of
acidic impurities during cell cycling that otherwise deleteriously affect the performance
of the cell.

3.3.1 Electrolytes Containing Phosphorus Flame-Retardant Additives and Fluoroester Co-
Solvent

Four experimental lithium-ion cells with MCMB carbon anodes and
LiNi
0.8
Co
0.2
O
2
cathodes were constructed to verify and demonstrate the reversibility,
78
electrochemical aspects, low temperature performance, and long-term charge-discharge
cycling performance of the chosen electrolyte compositions containing flame retardant
additives in combination with the fluoroester. Various electrochemical methods were
used to evaluate the series of flame-retardant additive containing electrolytes, including
conductivity studies and determination of the lithium intercalation and de-intercalation
kinetics through Tafel and DC micro-polarization studies. Additionally interfacial
characterization of each candidate electrolyte and MCMB-carbon anodes and
LiNi
0.8
Co
0.2
O
2
cathodes was analyzed by EIS. Following the electrochemical evaluation
of the cells the charge-discharge characteristics of the cells were made at various
temperatures and rates after all other testing was completed. Long-term charge-discharge
cycling was also evaluated for the cells. These evaluations will be discusse in the
following section. To make comparison easier, FRA electrolyte cell data was compared
to a baseline mixture containing no flame retardant additive, 1.0 M LiPF
6
in
EC.EMC.TFEB (20.60.20 v/v%). Data from this cell is actually the results collected
from the cell described in section 2.3.1 of this text.

79

Fig. 3.1 Chemical structures of phosphorus containing flame-retardant additives:
1) triphenylphosphate, 2) tributylphosphate, 3) triethylphosphate, 4) bis(2,2,2-
trifluoroethyl) methylphosphonate

3.3.1.1 Experimental Cell Results

3.3.1.1.1 Formation Characteristics

In order to determine the viability of electrolytes containing flame retardant
additives for use in lithium ion cells, the charge/discharge characteristics of these
MCMB/LiNo
x
Co
1-x
O
2
experimental cells were investigated. During the initial five
formation cycles the effects of electrolyte type on the irreversible and reversible
capacities were monitored. During this initial cycling the cell was charged and
discharged between 4.1V and 2.75V.
80
The fifth discharge capacities of the formation cycles are displayed in Fig. 3.2. It
is clearly visible that the baseline mixture of 1.0M LiPF
6
in EC:EMC:TFEB (20:60:20
v/v%) displays greater discharge capacity than the cells with FRA electrolytes. However,
this figure only tells part of the story because it does not account for discrepancies in
electrode weight and other influences on the total capacity of the cell such as packing
thickness and electrolyte loading. In order to accurately compare these cells, irreversible
capacity loss and coulombic efficiency during the formation cycling were evaluated.
Table 3.1 displays the cell data for electrolytes containing FRA. Data from these cells
compared favorably with the baseline during the initial formation cycles of these cells,
both in terms of irreversible capacity loss and coulombic efficiency. The order of the
cells in terms of cumulative irreversible capacity loss is (from least to greatest) 1.0 M
LiPF
6
in EC.EMC.TFEB.TPhPh (20.55.20.5 v/v%), 1.0 M LiPF
6
in EC.EMC.TFEB
(20.60.20 v/v%), 1.0 M LiPF
6
in EC.EMC.TFEB.TEtPh (20.55.20.5 v/v%), 1.0 M LiPF
6

in EC.EMC.TFEB.TBuPh (20.55.20.5 v/v%), 1.0 M LiPF
6
in EC.EMC.TFEB.TFMpo
(20.55.20.5 v/v%). Coulombic efficiency of these cells after five cycles is very good.
The baseline and TPhPh electrolyte maintain >98% efficiency and the other three cells
just slightly below 97%. These values are all acceptable and indicate the cells were
healthy and suitable SEI layers had been formed.
81

Fig. 3.2 Fifth discharge of formation cycling for cells containing flame retardant additive
and fluoroester cosolvent. Discharge at 25 mA/h at room temperature.

82

Table 3.1 Formation Data for the four cells containing flame retardant additives and baseline electrolyte.

83
3.3.1.1.2 Low Temperature Discharge Results

The cells were tested at a range of low temperatures and rates to assess their
performance under these conditions. These cells are charged at room temperature and
soaked at the reduced temperatures for >4 hours to ensure uniform temperature through
the entire cell. As demonstrated in Figs. 3.3 and 3.4, these electrolytes suffer a
significant decline in capacity compared to their performance at room temperature. The
capacity loss at reduced temperatures for these cells was significantly more than the
baseline electrolyte cell, which maintained >86% of its room temperature capacity at -
20ºC and >73% at -40ºC (25 mA/h discharge). In contrast, even the best performing
FRA electrolyte, the electrolyte with TBuPh, retained only slightly more than 68% of its
original room temperature capacity under similar conditions. At -40ºC the TEtPh cell ,
retained slightly more than 52% of its room temperature capacity. This discharge
capacity was a significantly higher percentage than any other FRA electrolyte cells at -
electrolyte cells at -40ºC but still lagged far behind the baseline. It should be noted the
tests performed at -40ºC were the first of the low temperature evaluations performed on
the cells, and while the cell with TEtPh electrolyte performed exceedingly well during
these initial low temperature evaluations, upon further testing the cell exhibited
significant aging and capacity fading over the life of the cell. We conclude that the
stability of this material within Li-ion cells is less than desired and will discuss this
finding further in the following section. The room temperature discharges of these cells
were compared to their respective 5
th
cycle formation discharges. In these tests, the cells
were only discharged to 2.75V, unlike the low temperature discharge data, which was
84
collected from cells discharged to 2.00 V. This fact further demonstrated the reduced low
temperature capabilities of such cells. A complete list of low temperature evaluations is
provided in Table 3.2. There were no conditions tested under which the cells containing
FRAs outperformed the baseline at reduced temperatures.

Fig. 3.3 Discharge of FRA containing electrolytes with fluorinated ester at -20º at
25mA/h. Values compared to percentage of original room temperature discharge.

85

Fig. 3.4 Discharge of FRA containing electrolytes at -40º at 25mA/h. Values compared
to percentage of original room temperature discharge.

86

Table 3.2 Discharge capacities of cells containing flame retardant additive electrolytes. Capacities are compared to the
final formation cycle discharge capacity of respective cells.

87
3.3.1.1.3 Cycle Life Evaluation

After low temperature assessment of the cells is performed, the capacity retention
of the cells is measured (room temperature, 25mA/h discharge). Table 3.3 shows the
cells lost a considerable amount of their discharge capacity during reduced temperature
testing demonstrating poor cycle life by the electrolytes. This trend is most recognizable
in the electrolytes containing phosphates with shorter straight chain hydrocarbon
branches consistent with previously reported studies.

Electrolyte Composition
Fifth Cycle of Formation
(mAh)
After Low Temperature
Discharge Tests (mAh)
% of Original
EC:EMC:TFEB:TPhPh (20:55:20:5) 399.01 272.79 68.37
EC:EMC:TFEB:TBuPh (20:55:20:5) 395.25 219.14 55.44
EC:EMC:TFEB:TEtPh (20:55:20:5) 384.01 182.61 47.55
EC:EMC:TFEB:TFMPo (20:55:20:5) 385.46 220.35 57.17

Table 3.3 Capacity fade of flame-retardant additive cells after low temperature discharge
tests

The cells were cycled at 25mA/h and monitored for capacity retention. Figure 3.5
further demonstrates the capacity loss of cells with FRAs. After 100 cycles the cell with
TPhPh additive maintains the highest capacity retention of the tested FRA cells. This
result is further exacerbated when the capacity lost during the low temperature testing is
included. When the cumulative losses of the low temperature discharge tests and 100
charge-discharge cycles are combined the cell with TPhPh maintained the highest
percentage of original room temperature at 24.2% followed by TBuPh at 16.2% and
TEtPh at 3.9%. Cell TFMPo maintained 28.8% of its original room temperature capacity
88
after only 52 cycles. If the trends in Fig. 3.5 were continued, it is assumed this cell would
have performed somewhere between cells with TPhPh and TBuPh.

Fig. 3.5 Capacity loss (%) as a function of number of cycles. Cells are taken to start at
100% of starting capacity even though significant capacity loss has already occurred
during low temperature discharge testing.

3.3.1.2 Electrochemical Evaluation

It is obvious from the low temperature discharge and long term charge-discharge
cycling that the flame-retardant additives imparted some deleterious effects onto the cells
to which they were added. Several electrochemical test techniques were employed to
more fully understand the means by which the FRAs impart these effects. To this end,
electrochemical impedance spectroscopy was used to study the effects the additives had
on the filming characteristics at both electrodes. Tafel polarization and DC micro-
89
polarization were used to determine the effects the additives have on the lithium
intercalation/de-intercalation kinetics at the individual electrodes.

3.3.1.2.1 Electrochemical Impedance Spectroscopy

Generally, the effects of the cathode dominate the interfacial impedance of the
total cell. These effects have been noted in previous reports by us and others

and are
supported by our Tafel polarization studies of similar cells where it is demonstrated that
the MCMB anodes typically sustain higher currents than their NiCoO
2
cathodic
counterparts.
28
It is also partly attributable to the cathode/anode balance with the
corresponding electrode loadings.
Our EIS data is obtained at open circuit voltages of fully charged cells over a
frequency range of 10
5
Hz to ~10 mHz, generated with an ac amplitude of 10 mV. The
impedance pattern typically contains two relaxation loops at room temperature, a high
frequency loop occurs in the range of 10
5
to 10
2
Hz and the other in the low frequency
range, between 10
2
Hz to 10
-1
Hz. The impedance data was analyzed using Z Simwin
modeling software from Princeton Applied Research. For in-depth information regarding
how the collected data was interpreted, refer to the impedance discussion in the
fluorinated ester electrolyte section of this text, 2.3.2.3.
EIS of Anodes
Electrolyte composition plays an important role in influencing both relaxation
loops that make up the Nyquist impedance plot at the anode (Fig. 3.6). Data from this
plot generated at room temperature revealed series resistance to be the value that was
90
most heavily influenced by electrolyte composition. As previously noted, the series
resistance is the sum total of electronic resistances from multiple sources, including
electrolyte ionic resistance which influences the total value of the series resistance
significantly though contribution from the leads, connections, and cell components
cannot be ignored. The film resistance (R
f
)

of the baseline was highest for the baseline
anode at room temperature, though not by a significant amount. The charge-transfer
resistance (R
CT
) of the baseline was similar to the anodes of the cells mixed with FRAs.
Fig. 3.6 Electrochemical Impedance Spectroscopy (EIS) measurements at 23
o
C of
MCMB electrodes from lithium-ion cells containing electrolytes consisting of 1.0M
LiPF
6
EC+EMC+TFEB+X (20:55:20:5 v/v %) (where X = TPhPh, TBuPh, TEtPh,
TFMPo). A baseline mixture without FRA is included as well.

0.00
0.05
0.10
0.15
0.20
0.25
0.30
0.35
0.40
0.00 0.10 0.20 0.30 0.40
Z' (Ohms)
1.00 M LiPF6 EC:EMC:TFEB:TPhPh (20:55:20:5 v/v %)
1.00 M LiPF6 EC:EMC:TFEB:TBuPh (20:55:20:5 v/v %)
1.00 M LiPF6 EC:EMC:TFEB:TEtPh (20:55:20:5 v/v %)
1.00 M LiPF6 EC:EMC:TFEB:TFMPo (20:55:20:5 v/v %)
1.00 M LiPF6 EC+EMC+TFEB (20:60:20 v/v %) (Standard)
Li-ion Experimental Three Electrode Cell
MCMB Carbon - LiNi
x
Co
1-x
O
2
Lithium metal Reference Electrode
91
As the temperature was decreased, all impedance values increased (Table 3.4);
however, the baseline value begins to consistently outperform the anode resistance values
for those cells containing FRAs. At all reduced temperatures the film impedance at the
anode of this cell was lower than any of the FRA cells. It was obvious the additives
formed a film with deleterious effects at reduced temperatures. This is supported from
CV data generated by others which found phosphate FRA compounds suffer from
reductive decomposition on natural graphite electrodes at potentials below of ~1.2 V vs.
Li/Li
+
.
18b
Specifically, these researchers found that increased amounts of the structurally
similar compound, trimethyl phosphate, lead to greater impedance values upon lithium
electrodes.
The nature of the individual phosphates induces some influence on the nature of
the impedance growth on the anodes. Although not completely consistent, the impedance
values for cell containing TPhPh generally produced the highest impedance values at the
anode. In contrast, anode impedance values of the TEtPh electrolyte were consistently
the lowest of the FRA cells indicating the nature of the side chain significantly influences
the impedance values of the cells, especially at the anode.

92

Table 3.4 Anode MCMB impedance data. Baseline EC:EMC:TFEB electrolyte solvent compared to solvents containing
flame retardant additives. Values given in ohms.

93
EIS of Cathodes

Similar to the effect on the anode, the nature of the electrolyte exhibited
considerable influence upon the impedance growth of both relaxation loops of the
cathode Nyquist plots (Fig. 3.7). Similar to the anodes, the series resistance at room
temperature was lowest for the baseline cathode; however, the FRA electrolytes exhibited
film and charge-transfer resistance values similar to the baseline at the same temp.

Fig. 3.7 Electrochemical Impedance Spectroscopy (EIS) measurements at 23
o
C of
LiNi
x
Co
1-x
O
2
electrodes from lithium-ion cells containing electrolytes consisting of 1.0M
LiPF6 EC+EMC+TFEB+X (20:55:20:5 v/v %) (where X = TPhPh, TBuPh, TEtPh,
TFMPo). A baseline mixture without FRA is included as well.

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' (Ohms)
1.00 M LiPF6 EC:EMC:TFEB:TPhPh (20:55:20:5 v/v %)
1.00 M LiPF6 EC:EMC:TFEB:TBuPh (20:55:20:5 v/v %)
1.00 M LiPF6 EC:EMC:TFEB:TEtPh (20:55:20:5 v/v %)
1.00 M LiPF6 EC:EMC:TFEB:TFMPo (20:55:20:5 v/v %)
1.00 M LiPF6 EC+EMC+TFEB (20:60:20 v/v %) (Standard)
Li-ion Experimental Three Electrode Cell
MCMB Carbon - LiNi
x
Co
1-x
O
2
Lithium metal Reference Electrode
Temperature = 23
o
C
94
At -30ºC and below, the cathode of the baseline exhibited the most favorable
impedance characteristics, but it is intriguing to note that at temperatures above this the
FRA containing cells compared quite favorably with the baseline at the cathode,
indicating that the additives influence cathode impedance to a lesser degree than was seen
at the anode. Cyclic voltammetry data reported by others supports this notion. Hyung
and coworkers
17f
found no evidence of oxidative decomposition for similar flame
retardant additives until potentials above 5.0V were reached, well within the working
voltage of typical lithium-ion cells.
Although it has previously been found that the cathode typically plays a dominant
role in the determination of the total cell impedance
67
the addition of FRAs changed the
balance of these contributions. Whereas the cathode typically exhibits greater film and
charge-transfer impedance than the anode, the addition of FRAs increase the film
resistance disproportionately at the anode compared to the cathode. Coupled with
previous data demonstrating reductive decomposition below 1.2 V for phosphate FRAs, it
is little surprise that these materials increased the film resistance at the anodes in greater
proportion than they did at the cathodes (Table 3.5).

95

Table 3.5 Cathode LiNi
x
Co
1-x
O
2
impedance data. Baseline EC:EMC:TFEB electrolyte solvent compared to solvents
containing flame retardant additives. Values given in ohms.
96
EIS of Full Cells
Table 3.6 Full cell impedance data. Baseline EC:EMC:TFEB electrolyte solvent compared to solvents containing flame
retardant additives. Values given in ohms.

97

As expected, the full cell impedance trends can be predicted from those found in
the individual electrode trends (Table 3.6). As with the impedance values of the
individual electrodes, the nature of substitution around the phosphorus atom of the FRAs
influenced the impedance values of the full cell to a substantial degree. Shorter, alkyl
carbon branched FRA exhibited lower impedance values for the cell. This trend is
exacerbated at reduced temperatures.

3.3.1.2.2 Tafel Polarization Measurements

Investigation of the lithium intercalation/de-intercalation electrode kinetics was
performed using Tafel polarization measurements of the MCMB-Li
X
Ni
y
Co
1-y
O
2
cells
with the described electrolyte formulations. The measurements were performed on the
fully charged cells (OCV= ~4.095 V) with data collected between room temperature and
-40ºC. The data was generated under potentiodynamic conditions but slow scan rates
approximated steady-state conditions within the cells.
The electrolytes were examined at room temperature, -20ºC, and -40ºC to analyze
the effects reduced temperatures have on the lithium kinetics of the electrodes. As
expected, reduced temperatures hinder the kinetic performance at both electrodes of the
cell, greatly depressing current density under these conditions. At the cathode, the FRA
electrolytes exhibited lithium kinetics similar to the baseline at all tested temperatures
(Figs. 3.8-3.10). However, the FRAs extended significant deleterious effects upon the
anodes of the electrolytes to which they are added when the results were compared to the
98
baseline electrolyte (Figs. 3.11-3.13). A likely source of this current density depression
of these electrodes is the film growth observed using impedance spectroscopy in the
previous section. Whereas previous studies found that the anode typically exhibits better
lithium kinetics than does the cathode, two of the FRA electrolytes tested here, the cells
with TBuPh and TEtPh electrolytes, exhibited better lithium kinetic performance at their
cathode than they did at their anodes. Even the FRA electrolytesthat maintained better
performance at their anodes (those with TPhPh and TFMPo) did so to a significantly
lesser degree than did the baseline, which had an anode current density roughly double of
its cathode.

Fig. 3.8 Tafel polarization measurements at 23
o
C of LiNi
x
Co
1-x
O
2
electrodes from
lithium-ion cells containing electrolytes consisting of 1.0M LiPF6 EC+EMC+TFEB+X
(20:55:20:5 v/v %) (where X = TPhPh, TBuPh, TEtPh, TFMPo).
99

Fig. 3.9 Tafel polarization measurements at -20
o
C of LiNi
x
Co
1-x
O
2
electrodes from
lithium-ion cells containing electrolytes consisting of 1.0M LiPF6 EC+EMC+TFEB+X
(20:55:20:5 v/v %) (where X = TPhPh, TBuPh, TEtPh, TFMPo).

Fig. 3.10 Tafel polarization measurements at -40
o
C of Ni
x
Co
1-x
O
2
electrodes from
lithium-ion cells containing electrolytes consisting of 1.0M LiPF6 EC+EMC+TFEB+X
(20:55:20:5 v/v %) (where X = TPhPh, TBuPh, TEtPh, TFMPo).
100

Fig. 3.11 Tafel polarization measurements at 23
o
C of MCMB electrodes from lithium-ion
cells containing electrolytes consisting of 1.0M LiPF6 EC+EMC+TFEB+X (20:55:20:5
v/v %) (where X = TPhPh, TBuPh, TEtPh, TFMPo).

Fig. 3.12 Tafel polarization measurements at -20
o
C of MCMB electrodes from lithium-
ion cells containing electrolytes consisting of 1.0M LiPF6 EC+EMC+TFEB+X
(20:55:20:5 v/v %) (where X = TPhPh, TBuPh, TEtPh, TFMPo).
101

Fig. 3.13 Tafel polarization measurements at -40
o
C of MCMB electrodes from lithium-
ion cells containing electrolytes consisting of 1.0M LiPF6 EC+EMC+TFEB+X
(20:55:20:5 v/v %) (where X = TPhPh, TBuPh, TEtPh, TFMPo).

3.3.1.2.3 DC Micro-Polarization Measurements

DC micro-polarization techniques were employed on these cells to further study
how the charge transfer behavior of the electrodes were affected by the addition of FRAs
as well as the temperature change upon the cells. The polarization resistance values of
the electrodes were calculated from the slopes of the linear plots generated under
potentiodynamic conditions of scan rates of 0.2 mV/s and measured over a +/- 5mV
range from the equilibrium potential of each electrode. Polarization values measured at
the anode for FRA electrolyte cells did not perform as well as the baseline at all
examined temperatures supporting the claim that the reduced performance of the cells
102
containing these electrolytes is a result of their deleterious effects on the anode (Table
3.7). At the cathode cells containing FRA perform comparably to the baseline cathode
indicating the FRAs have little effect on the performance of this electrode (Table 3.8).

Table 3.7 DC Micro-polarization measurements at various temperatures of anodes from
lithium-ion cells containing electrolytes consisting of 1.0M LiPF
6
EC+EMC+TFEB+X
(20:55:20:5 v/v %) (where X = TPhPh, TBuPh, TEtPh, and TFMPo).

Table 3.8 DC Micro-polarization measurements at various temperatures of cathodes from
lithium-ion cells containing electrolytes consisting of 1.0M LiPF
6
EC+EMC+TFEB+X
(20:55:20:5 v/v %) (where X = TPhPh, TBuPh, TEtPh, and TFMPo).
103
3.3.2 Improving Performance of Cells Containing Flame Retardant Additives through use
of an SEI Enhancing Additive

When flame retardant additives are charged to Li-ion electrolytes, the
performance of the cells is significantly compromised, especially with regard to capacity
retention over the life of the cell. In an attempt to reduce the deleterious effects imparted
on the electrolyte by the addition of the flame retardant additive an SEI enhancing agent,
vinylene carbonate (VC), was added to the system. The four experimental Li-ion cells
with MCMB carbon anodes and LiNi
0.8
Co
0.2
O
2
cathodes were constructed to verify and
demonstrate the reversibility, electrochemical aspects, low temperature performance, and
long-term charge-discharge cycling performance of the chosen electrolyte compositions.
The four cells were filled with electrolytes designed to monitor the effects of triphenyl
phosphate (TPhPh) and VC, both individually and in combination. The electrolyte
formulations are as follows:

A. 1.0 M LiPF
6
in EC.EMC (20.80 v/v%)
B. 1.0 M LiPF
6
in EC.EMC (20.80 v/v%) w/ 1.5% VC
C. 1.0 M LiPF
6
in EC.EMC.TPhPh (20.75.5 v/v%)
D. 1.0 M LiPF
6
in EC.EMC.TPhPh (20.75.5 v/v%) w/ 1.5% VC

Various electrochemical methods were used to evaluate the series of flame-
retardant additive containing electrolytes, including conductivity and determination of the
lithium intercalation and de-intercalation kinetics through Tafel and DC micro-
104
polarization studies. Additionally, the relative interfacial stability of each electrolyte
with MCMB-carbon anodes and LiNi
0.8
Co
0.2
O
2
cathodes was analyzed by performing
EIS. Following the electrochemical evaluation of the cells, the charge-discharge
characteristics of the cells were made at various temperatures and rates. Long-term
charge-discharge cycling was also evaluated for the cells. These electrical evaluations
will be discussed in the following section.

3.3.2.1 Experimental Cell Results

3.3.2.1.1 Formation Characteristics

In order to determine the viability of electrolytes containing triphenyl phosphate
and/or vinylene carbonate for use in lithium ion cells, the charge/discharge characteristics
of these MCMB/LiNo
x
Co
1-x
O
2
experimental cells were investigated. Discharge studies
of these cells help determine the effects TPhPh and VC have on the performance of the
cell. During the initial five formation cycles, the irreversible and reversible capacities of
each cell were monitored. During this initial cycling the cell was charged and discharged
between 4.1V and 2.75V.

105

Fig. 3.14 Fifth discharge of formation cycling for cells containing TPhPh and/or VC.
Discharge at 25 mA/h at room temperature.

The fifth discharge capacities of the formation cycles are displayed in Fig. 3.14.
The cell utilizing the electrolyte of 1.0M LiPF
6
in EC:EMC:TPhPh (20:75:5 v/v%)
exhibits the greatest discharge capacity after 5 cycles. However, this only reveals part of
the story as it does not account for discrepancies in electrode weight and other influences
on the capacity such as packing thickness and electrolyte loading. In order to accurately
compare the formation characteristics of these cells, irreversible capacity loss and
coulombic efficiency during the formation cycling were evaluated. Table 3.9 shows the
cells containing FRA/VC electrolytes compare favorably with the baseline during the
initial formation cycles of these cells, both in terms of irreversible capacity loss and
coulombic efficiency. It should be noted that the coulombic efficiency of the cell
106
containing the electrolyte, 1.0 M LiPF
6
in EC:EMC:TPhPh (20.75.5) with 1.5% VC was
considerably lower than all other cells after the first cycle. At first glance it was
considered that the reduced coulombic efficiency stem from the fact that both additives
may be breaking down upon the first charge leading to a higher charge capacity than the
electrodes actually accepted. However, further reviews showed neither cell employing
electrolytes that contain TPhPh or VC individually exhibit any reduced coulombic
efficiency that would be anticipated if this were the case. Formation data for a cell
employing the electrolyte 1.0M LiPF
6
in EC+EMC+TFEB (20:60:20 v/v%) is also
presented for the sake of comparison.
The cells were initially subjected to five formation cycles at room temperature.
During these cycles the cells were charged at 25 mA/h to a 4.1V cut-off and discharged at
25 mA/h (~C/16 rate) to 2.75 V. All cells exhibited stable capacity over the first five
formation cycles with results comparable to the baseline carbonate solution. After
completing the five formation cycles, the characterization of the cells at various
temperatures was performed using numerous electrochemical techniques (EIS, Tafel, DC
micro-polarization measurements) and then subjected to discharge evaluation at low
temperatures at various rates. Following these tests cycle life performance evaluations
were performed to monitor the long-term cycling effects the FRA/VC additives had upon
the cells.

107

Table 3.9 Formation data for TPhPh/VC cells. Cells with flame retardant and/or SEI enhancing additives
compare favorably to the baseline formulations.

108
3.3.2.1.2. Low Temperature Discharge Results

The cells were tested at a range of low temperatures and rates to analyze their
performance under these conditions. The cells were charged at room temperature and
soaked at the reduced temperatures for >4 hours to ensure the cells reached the test
conditions uniformly through the entire cell. Figures 3.15-3.17 show the discharge
profiles for the cells at declining temperatures. At -20 ºC, as exhibited in Fig. 3.15, all
cells maintained a similar degree of discharge capacity retention (within 3%) relative to
their room temperature capacity at similar rates. The formulation of 1.0M LiPF
6

EC:EMC:TPhPh (20:75:5) retained the highest percentage of its capacity, although no
discernable pattern exists with regard to additive content and capacity retention exhibited
under these conditions. As the temperature was lowered and/or the rate of discharge
increased, a discernable pattern emerges. Fig. 3.16 shows the 25 mA/h discharge profile
for the cells at -40 ºC. The cells exhibit similar discharge capabilities to one another, but
begin to show the trend that can be more fully viewed in Fig. 3.17. Under the conditions
of this test (-60 ºC and 5 mA/h discharge rate) it is visible that capacity loss of these cells
had a pattern that was easily explained. The baseline retains the highest percentage of
room temperature capacity (54.04%), followed by the cell with 1.5% VC (51.9%), 1.5%
VC and 5% TPhPh (42.88%), and finally, the FRA only (38.67%). Not surprisingly, the
baseline exhibits the least amount of capacity loss. Trailing only slightly behind the
baseline, it was obvious VC offers some degree of protection against the deleterious
effect the addition of TPhPh has upon the cell at low temperature discharge. Table 3.10
shows the capacity retention of the investigated cells under all test conditions. At nearly
109
every rate and temperature combination below -30ºC the flame retardant electrolyte
mixed with VC exhibited better low temperature characteristics than does the FRA cell
without the SEI enhancing additive.

Fig. 3.15 Cell discharge profile at -20 ºC with various electrolytes at 25 mA/h. Cells
were charged at room temperature to 4.1 V.

110

Fig. 3.16 Cell discharge profile at -40 ºC with various electrolytes at 25 mA/h. Cells
were charged at room temperature to 4.1 V.

Fig. 3.17 Cell discharge profile at -60 ºC with various electrolytes at 5 mA/h. Cells were
charged at room temperature to 4.1 V.
111

Table 3.10 Summary of discharge performance (capacity, % room temperature) of experimental lithium-ion cells at low
temperatures of FRA/VC cells at various rates.

112
3.3.2.1.3. Cycle Life Testing

After electrochemical characterization and low temperature discharge testing, the
experimental cells were charged and discharged at 25 mA/h for 100 cycles at room
temperature. The capacity of the cells is initially compared to the room temperature
capacity of the fifth formation charge. Table 3.11 displays this data. It is clear that the
cells in this test suffered similar capacity loss during low temperature testing. The cell
with electrolyte 1.0 M LiPF
6
in EC:EMC with 1.5% VC exhibited the largest amount of
capacity fade during low temperature testing, which is surprising because the additive is
thought to improve cyclability of cells by forming a protective layer at the anode.
However, this result is easily within an accepted margin and the small difference can be
accounted for by other cell variability factors.

Electrolyte Composition
Fifth Cycle of
Formation (mAh)
After Low Temperature
Discharge Tests (mAh)
% of Original
Discharge Capacity
1.0 M LiPF6 EC+EMC (20:80 v/v %) 378.09 281.54 74.46
1.0 M LiPF6 EC+EMC (20:80 v/v %)
w/ 1.5% VC
390.29 279.74 71.67
1.0 M LiPF6 EC+EMC+TPhPh (20:75:5
v/v %)
397.36 301.93 75.98
1.0 M LiPF6 EC+EMC+TPhPh (20:75:5
v/v %) w/ 1.5% VC
377.43 289.93 76.82

Table 3.11 Capacity fade of the cells containing flame-retardant additives after low
temperature discharge tests

Figure 3.18 shows the capacity fading of the cells over life cycle testing. The
addition of VC to electrolyte containing an flame-retardant additive enhanced the cycle
life of the cell compared to the cell that contains the FRA without VC added. However,
the fact that the cell with FRA and VC outperformed both the baseline and VC only is
113
puzzling. The predicted trend would have the VC only cell and (possibly) the baseline
outperforming any cell that contains a FRA. It does appear that these results are within
the margin for cell-to-cell variation and at minimum, this test proves that the addition of
VC to a cell containing FRA exhibits no negative effect on the cell and probably
enhances the cycle life of the cell.

Fig. 3.18 Capacity fade of cells with FRA/VC electrolytes at room temperature during
cycle life testing.

3.3.2.2 Electrochemical Evaluation

The addition of the flame retardant additive and/or film-enhancing additive to the
electrolyte mixture extends significant influence upon the low temperatures discharge and
life cycle characteristics of the cells to which they are employed. In order to elucidate the
114
mode of action these additives had on the performance, electrochemical techniques were
used to evaluate the cells. Use of a reference electrode in the cells allows us to monitor
the individual characteristics of the active electrodes using Tafel polarization
measurements, electrochemical impedance spectroscopy, and DC micro-polarization. The
following sections describe the results found through the use of these tests and
hypothesizes on how these findings affect the performance of the cells.

3.3.2.2.1 Electrochemical Impedance Spectroscopy Results

Impedance results were collected in the same manner as has previously been
described within this text. For further description of the methodology refer to section
2.3.2.3 of this text. Impedance data was particularly of interest in this study because of
its ability to probe the nature of the anode passivating film, where the additives were
expected to exhibit the greatest effect.
EIS of Cathodes
Figure 3.19 and Table 3.12 demonstrate the effect of the electrolyte composition
on the nature of the interfacial properties of the electrodes, even at room temperature.
The relaxation loops corresponding to both the film and charge transfer resistances of the
cathode Nyquist plots are both significantly altered. The film resistance for the baseline
(0.2423 Ω) is in line with the both electrolytes that contain only one additive (0.202 Ω for
1.5% VC electrolyte and .2702 Ω for 5% TPhPh electrolyte, respectively). However, the
electrolyte that contains both the FRA and VC exhibited a significantly lower film
impedance value at the cathode (0.006863 Ω). However, it should be noted that none of
115
the impedance values for the FRA/VC cell seem to correlate well with the other cells or
other tests performed on this cell and may need to be disregarded.

Fig. 3.19 Nyquist plot of FRA/VC study cathode impedance at room temperature.

As the temperature is decreased, all cells suffered similar film and charge-transfer
impedance growth indicating the additives exhibit little influence upon the filming
properties of the cathode. This was expected as all previous references indicate that both
VC and FRA affect the anode much more heavily than the cathode.

0.00
0.10
0.20
0.30
0.40
0.50
0.60
0.70
0.80
0.90
1.00
0.00 0.20 0.40 0.60 0.80 1.00
Z' (Ohms)
1.00 M LiPF6 EC:EMC (20:80 v/v %)
1.00 M LiPF6 EC:EMC (20:80 v/v %) w/ 1.5% VC
1.00 M LiPF6 EC:EMC:TPhPh (20:75:5 v/v %)
1.00 M LiPF6 EC:EMC:TPhPh (20:75:5 v/v %) w/ 1.5% VC
Li-ion Experimental Three Electrode Cell
MCMB Carbon - LiNi
x
Co
1-x
O
2
Lithium metal Reference Electrode
Temperature = 23
o
C
116

Table 3.12 EIS low temperature data for FRA/VC cathodes. Impedance growth is affected similarly for all cells at reduced
temperatures.

117
EIS of Anodes
Electrolyte composition significantly influences both relaxation loops comprising
the Nyquist impedance plots of the anode at all temperatures (Fig. 3.20). At room
temperature, the cells containing VC extended deleterious effects upon the film resistance
that were exacerbated as the temperature is lowered (Table 3.13). These results are
expected given the mechanism for the additive involves degradative reductive
polymerization at the anode. However, the film resistance data for the cell containing 5%
TPhPh was surprising because the degree of film resistance growth was considerably
lower.
Fig. 3.20 Nyquist plot of FRA/VC study anode impedance at room temperature.

0.00
0.10
0.20
0.30
0.40
0.50
0.60
0.70
0.80
0.90
1.00
0.00 0.20 0.40 0.60 0.80 1.00
Z' (Ohms)
1.00 M LiPF6 EC:EMC (20:80 v/v %)
1.00 M LiPF6 EC:EMC (20:80 v/v %) w/ 1.5% VC
1.00 M LiPF6 EC:EMC:TPhPh (20:75:5 v/v %)
1.00 M LiPF6 EC:EMC:TPhPh (20:75:5 v/v %) w/ 1.5% VC
Temperature = 23
o
C
Li-ion Experimental Three Electrode Cell
MCMB Carbon - LiNi
x
Co
1-x
O
2
Lithium metal Reference Electrode
118

Table 3.13 EIS low temperature data for FRA/VC anodes. Charge-transfer impedance growth is affected similarly for all
cells containing additives as a function of reduced temperatures. However, the TPhPh containing electrolyte exhibits better
film impedance trends
119
When the temperature is lowered, cells containing either type of additive exhibit
increased impedance growth. This is especially true for charge-transfer impedance
growth in all cells. Film resistance in cells containing VC is consistently higher than in
cells without, and this trend is exacerbated as the temperature is lowered. However, in
the cell utilizing TPhPh, film resistance grows only noticeably higher than the baseline
when the temperature is lowered to -40ºC or lower, which indicated a degree of
temperature sensitivity. These results are unsurprising as some performance loss is
anticipated at the anode of cells containing either type of additive.

EIS of Full Cells
In most cases, the full cell impedance trends can be predicted from the individual
electrode trends. (Table 3.14). Unsurprisingly, the full cell impedance data indicates that
the additives contribute to significant impedance growth in the entire cells with
exacerbated low temperature results.

120

Table 3.14 EIS low temperature data for FRA/VC full cells. Cells containing electrolyte additives exhibit deleterious effects
with these effects exaggerated at reduced temperatures.

121
3.3.2.2.2 Tafel Polarization Measurements

Evaluation of the lithium intercalation/de-intercalation electrode kinetics was
gathered using Tafel polarization measurements of the MCMB-LiNi
0.8
Co
0.2
O
2
cell with
the described electrolytes previously discussed. Measurements were performed on the
cells while they were fully charged (OCV= ~4.095V) at room temperature. Low
temperature evaluations were also performed on the cells. Data were generated under
potentiodynamic conditions with slow scan rates approximating steady-state conditions.
Interpretation of the measurement of the anodes at room temperature is difficult.
The cells with 1.5% VC and 5% TPhPh were both able to pass more current than the
baseline under these condition, but the cell containing both additives performed the
poorest of the entire set (Fig. 3.21). Addition of the additives is unlikely to improve the
performance of the cell, especially at the anode and no justification can be developed to
explain these results. The degree to which the limiting current densities varied was also
puzzling.
Interpretation of the cathode polarization data is difficult as well. Each of the
cells containing additives exhibited higher limiting current densities than the baseline
(Fig. 3.22). The cathode current densities were nearly equal to the anode current
densities of the examined cells, in contrast to typical cells wherein the current density is
limited heavily by the cathode. This provides strong evidence of the additive
deleteriously affecting the performance of the anode.

122

Fig. 3.21 Tafel polarization measurements at 20ºC of MCMB anodes in FRA/VC study.

Fig. 3.22 Tafel polarization measurements at 20ºC of LiNi
0.8
Co
0.2
O
2
cathodes in FRA/VC
study

123
As expected, both the anode and cathode suffer significantly reduced lithium
kinetics when the temperature was dropped. At the anode, cells containing additives
suffer from the greatest loss in current density. At -40ºC the baseline cell performed the
best (Fig. 3.23). As the temperature is lowered, the baseline continues to out perform the
additive containing cells to greater and greater degrees (Fig. 3.25). These results further
show the additives negatively affected the performance of the cells primarily due to their
affect on the anode and support the findings collected from impedance spectroscopy.
At the cathode, results were much murkier. Although the data indicates that the
additives examined in this study improve the current density at all test temperatures (Figs.
3.24 and 3.26), no mechanism can be proposed to support a reason for this phenomenon.
Presently, all that is claimed is that little evidence exists to support any notion that the
additives extend deleterious effects upon the current density at the cathode.

Fig. 3.23 Tafel polarization measurements at -40ºC of MCMB anodes in FRA/VC study.
124

Fig. 3.24 Tafel polarization measurements at -40ºC of LiNi
0.8
Co
0.2
O
2
cathodes in
FRA/VC study

Fig. 3.25 Tafel polarization measurements at -60ºC of MCMB anodes in FRA/VC study.
125

Fig. 3.26 Tafel polarization measurements at -60ºC of LiNi
0.8
Co
0.2
O
2
cathodes in
FRA/VC study

3.3.2.2.3 DC Micro-Polarization Measurements

DC micro-polarization techniques were employed to the cells to further study the
charge transfer behavior of the passivating films affected by the addition of the additives
to the electrolytes as well as the effects of temperature upon the cells. The polarization
resistance values of the electrodes were calculated from the slopes of the linear plots
generated under potentiodynamic conditions of scan rates of 0.2 mV/s and measured over
a +/-5 mV from the OCV of each electrode. When the polarization values are measured
at the anode, polarization values for cells containing either flame-retardant or VC
additive consistently measured higher than the baseline further supporting claims that the
126
additives negatively affect the performance of cells due to their action at the anode (Table
3.15). In contrast, the linear polarization values at the cathode for the cells further the
claim that the additives impart no deleterious effects on upon the cathode (Table 3.16). It
must be noted that the polarization resistance for all cells performed at -20ºC appear out
of line and much lower than we would anticipate for unknown reasons. All other test
results appear to be consistent.

Table 3.15 DC micro-polarization measurements at various temperatures at anodes from
lithium-ion cells containing FRA/VC additives and baseline electrolytes. (Results in kΩ)

Table 3.16 DC micro-polarization measurements at cathodes at various temperatures for
Li-ion cells containing FRA/VC additives and baseline electrolytes. (Results in kΩ)

127
3.3.3 Effects of Branch Substitution and Oxidation State of Phosphorus in Phosphorus-
Containing Flame-Retardant Additives in Li-Ion Cells

Fig. 3.27 Chemical structures of phosphorus-containing flame retardants examined: 1)
tris(2,2,2-trifluoroethyl) phosphate, 2) tris(2,2,2-trifluoroethyl) phosphite, 3)
triphenylphosphite, 4) diethyl ethylphosphonate, 5) diethyl phenylphosphonate

Electrolytes containing various flame retardant additives at 5 v/v% were charged
to five experimental lithium-ion cells with MCMB carbon anodes and LiNi
0.8
Co
0.2
O
2

cathodes. The flame retardant additives in this study were chosen to demonstrate the
effects branch substitution at the phosphorus atom and oxidation state of the phosphorus
atom on the performance of the Li-ion cells. The flame retardant additives examined in
this study were tris(2,2,2-trifluoroethyl) phosphate (TFPa), tris(2,2,2-trifluoroethyl)
phosphite (TFPi), triphenylphosphite (TPPi), diethyl ethylphosphonate (DEP), and
diethyl phenylphosphonate (DPP). The electrolytes tested in this study were as follows:

128
A. 1.0 M LiPF
6
in EC.EMC.TFPa (20.75.5 v/v%)
B. 1.0 M LiPF
6
in EC.EMC.TFPi (20.75.5 v/v%)
C. 1.0 M LiPF
6
in EC.EMC.TPPi (20.75.5 v/v%)
D. 1.0 M LiPF
6
in EC.EMC.DEP (20.75.5 v/v%)
E. 1.0 M LiPF
6
in EC.EMC.DPP (20.75.5 v/v%)

To verify and demonstrate the viability of the electrolytes in Li-ion cells, the
reversibility, low-temperature performance, and long-term charge-discharge cycling
performance for each individual electrolyte/cell combination were examined. The
electrochemical measurements were used in the determination of the lithium intercalation
and de-intercalation kinetics through Tafel and DC micro-polarization studies.
Additionally, the relative interfacial stability between candidate electrolyte and MCMB-
carbon anodes and LiNi
0.8
Co
0.2
O
2
cathodes was analyzed using EIS. Following the
electrochemical evaluation of the cells, the charge-discharge characteristics of the cells
were made as a function of temperature and rate. Evaluation of the charge and discharge
performance characteristics of the cells will be discussed in the following section.

3.3.3.1 Experimental Cell Results

3.3.3.1.1 Formation Characteristics

MCMB-LiNi
0.8
Co
0.2
O
2
experimental cells are constructed and filled with
candidate flame-retardant additive electrolyte in a glove box. The irreversible and
129
reversible capacities of these cells are closely monitored during the initial five formation
cycles. The cells in this study exhibited excellent reversibility and coulombic efficiency
during these cycles and compare quite favorably to some of the baseline electrolytes
described in previous sections of this text. Ranking the electrolytes in order of the
irreversible capacity loss over the course of the first five cycles from highest to lowest
follows the trend (from worst to best) EC+EMC+TFPA (0.3410 Ah), EC+EMC+TFPi
(0.1353 Ah), EC+EMC+TPPi, (0.1283 Ah), EC+EMC+DPP (0.1027 Ah), and
EC+EMC+DEP (0.0703). The coulombic efficiencies of these cells follow the same
order as well (Table 3.17). The data collected with regard to coulombic efficiency and
irreversible capacity loss for the cell containing the additive TFPa is measurably worse
than all other cells within the study and may signal that this additive does not
demonstrate the electrochemical stability necessary for long cycle life in a Li-ion cell.
This finding is curious, as numerous other researchers have singled out the additive as a
potentially exciting flame retardant additive with excellent kinetic, filming and cycle life
properties. One explanation may be that the additive is sacrificially participating in the
formation of the SEI, however, this conclusion cannot be drawn from the formation data
alone. Electrochemical evaluation performed on the cell gave further insight into
whether this conclusion is possible to draw.
Initial formation cycling of the cells consisted of charging and discharging of the
cells at 25 mA (~C/16 rate) over a voltage range of 2.75V-4.10V at 23ºC. After initial
formation data were collected, the cells were subjected to electrochemical evaluation over
a wide temperature range followed by low temperature discharge (room temperature
130
charge) studies, and cycle life testing to determine the long-term stability of the
electrolytes in question.

Table 3.17 Formation data for flame retardant additive electrolyte cells. Generally, cells with the phosphorus-
containing flame retardants exhibited excellent reversibility and efficiency.

131

Fig. 3.28 Fifth discharge of formation cycling for cells containing flame retardant
additives. Discharge at 25 mA/h at room temperature.

3.3.3.1.2 Low-Temperature Discharge Results

When the cells containing the FRA electrolytes were monitored for their low-
temperature performance, some distinct trends made themselves apparent among the
electrolytes. The cells were charged at room temperature and at -20ºC using moderate
~C/16 discharge rates (fig. 3.29), the capacity retention of the cells rank (from highest to
lowest): EC+EMC+DEP (87.5% of room temperature discharge) > EC+EMC+DPP
(86.5%) > EC+EMC+TFPi (86.4%) > EC+EMC+TFPa (84.1%) > EC+EMC+TPPi
(75.3%). When the cells were discharged at the same rate at -40ºC, the order of capacity
retention remains the same (Fig. 3.30). In fact, under all tested reduced temperature
132
discharge experiments, the cells retained the same order of discharge capacity retention
(Table 3.18). There is a link between the steric bulk of the FRA molecule branches and
the reduced temperature capabilities of the electrolytes. This is because the increased
bulk at the branched site lead to reduced conductivity and kinetic capabilities and will be
discussed further in later sections. Limitation of conductivity and kinetics exacerbate
themselves at lowered temperatures. The trend linking branch bulk size and low
temperature performance was also present in the electrolytes discussed in section 2.3.1 of
this chapter.
After all of the room temperature charge/low temperature discharge testing was
performed, the effect of charging the cells at low temperatures was also examined.
Initially, the cells were charged at room temperature, then soaked for at least 4 hours at
the testing temperature. The effects of charge and discharge rates on cell performance
were closely monitored in these tests. The results of these tests mirrored the tests that
utilized room temperature charge, although the degree to which the other electrolytes
outperformed the cell using TPPi was not as large. It is thought that the electrochemical
stability of this particular additive/electrolyte lead to improved performance relative to
the other electrolytes over the life of the individual cells. This was confirmed during
cycle life tests of the cells and will be discussed in the following section of this text.

133

Fig. 3.29 Discharge of FRA containing electrolytes at -20º at 25mA/h. Values compared
to percentage of original room temperature discharge.

Fig. 3.30 Discharge of FRA containing electrolytes at -40º at 25mA/h. Values compared
to percentage of original room temperature discharge.

134

Table 3.18 Low temperature capacity of FRA containing electrolytes.

135

Table 3.19 Rate capabilities of FRA electrolytes when charged at low temperatures

136
3.3.3.1.3 Cycle Life Evaluation

After low temperature assessment of the cells is performed, the capacity of the
cells is compared to the starting capacity under the same conditions (room temperature,
25mA/h discharge). Table 3.20 shows the cells lost a considerable amount of their
discharge capacity after being tested at reduced temperatures demonstrating the cells
exhibit poor cycle life. This trend is most recognizable in the alkyl phosphates with
shorter straight chain hydrocarbon branches consistent with previously reported studies.
100 charge/discharge cycles at room temperature was used to assess the long-term
cycle life stability of the electrolytes within the cells. Cells were charged/discharged at
moderate 25 mA/h rates. Table 3.20 shows the amount of irreversible capacity loss the
individual cells lost during low temperature evaluation prior to cycle life testing. In
addition to playing an integral role in the low-temperature discharge profiles, the branch
constitution of the flame retardant additives also significantly influenced the long-term
stability of the electrolytes. Of note is the fact the lifetime stability capabilities of the
electrolytes were found to be nearly exactly opposite to the low temperature performance.
While short branches of the flame retardant additives improve conductivity and kinetic
rate capability, bulky branches at the phosphorus atom favor long-term electrochemical
stability. This has been reported previously by others and discussed in section 3.3.1 of
this text. Further evidence of this phenomenon is illustrated in Fig. 3.31. The impressive
cycling performance exhibited by the electrolyte containing 5% triphenyl phosphite (P
ox.
state
= +3) is more impressive than the cycle performance of the cell containing triphenyl
phosphate (P
ox. state
= +5) shown in previous sections of this text. Previous studies
137
showed impressive electrochemical stability for flame retardant electrolytes containing
additives that feature phosphorus in the +3 oxidation state because these compounds act
as Lewis-base scavengers of acidic contaminants that manifest themselves during
breakdown of the electrolyte during cycling of the cells. Such a phenomenon is not
observed for the additives featuring 2,2,2-trifluoethyl substituents because of their
diminished Lewis basicity

Electrolyte Composition
Fifth Cycle of
Formation (mAh)
After Low Temperature
Discharge Tests (mAh)
% of Original
Discharge Capacity
1.0 M LiPF6 EC+EMC+TFPa
(20:75:5 v/v %)
387.04 349.67 90.34
1.0 M LiPF6 EC+EMC+TFPi
(20:75:5 v/v %)
409.70 320.80 78.30
1.0 M LiPF6 EC+EMC+TPPi
(20:75:5 v/v %)
402.16 360.49 89.64
1.0 M LiPF6 EC+EMC+DEP
(20:75:5 v/v %)
398.67 287.63 72.15
1.0 M LiPF6 EC+EMC+DPP
(20:75:5 v/v %)
386.16 308.29 79.84

Table 3.20 Capacity fade of cells containing flame-retardant additives after low
temperature discharge tests

Figure 3.31 further demonstrates the capacity loss of cells with FRAs. After 100
cycles, the cell with TPPi additive demonstrates excellent capacity retention most likely
due to the flame retardant acting as an electrolyte stabilizer by scavenging acidic
decomposition products from the electrolytes. The poorest performing flame retardant
additives were the ethyl diethylphosphonate and phenyl diethylphosphonate. These
structurally similar molecules further demonstrate the poor capacity retention of flame
retardant electrolytes containing small branched additives.

138

Fig. 3.31 Life cycle testing of cells containing FRAs. Cells are discharged and charged
at room temperature at moderate rates (25mA/h) for 100 cycles.

3.3.3.2 Electrochemical Evaluation

To elucidate the way flame retardant additives affect the performance of the Li-
ion cells, a series of electrochemical tests was performed on the cells, including
conductivity measurements, Tafel polarization measurements, EIS, and linear micro-
polarization measurements. These tests will help us to determine if the effect these
additives were due to their impact on the mass transfer characteristics in the electrolyte
(change in conductivity) and/or charge transfer characteristics due to the modification of
the intercalation/de-intercalation characteristics at the SEI layer caused by introduction of
the flame retarding additives.
139
3.3.3.2.1 Conductivity Measurements

The specific conductivities of electrolytes containing flame retardant additives
were measured over a range of temperatures from -60ºC to 25ºC. Of interest was the
effect of structural modifications to the FRA on conductivity. At all temperatures, the
electrolyte with triethyl phosphate exhibits the highest conductivity, surpassing even the
baseline, which contained no flame retardant additive (Table 3.21). At room
temperatures, the conductivity was observed to follow the trend (all solutions consist of
1.0 M LiPF
6
dissolved in the following solvent mixtures): EC.EMC.triethyl phosphate
(TEtPh) (20.75.5 v/v) (8.94 mS/cm) > EC.EMC (20.80 v/v) (8.70 mS/cm) >
EC.EMC.tris(2,2,2-trifluoroethyl) phosphate (TFEPh) (20.75.5) (8.18 mS/cm) >
EC.EMC.triphenyl phosphite (TPhPi) (20.75.5 v/v) (7.89 mS/cm) > EC.EMC.triphenyl
phosphate (TPhPh) (20.75.5 v/v) (7.88 mS/cm). As indicated, reasonable conductivity
was realized when short chained phosphorus containing flame retardants are used in the
electrolyte, though moderate loss in conductivity is realized when the branches are
fluorinated. A significant loss of conductivity was realized when aromatic substituents
are placed around the phosphorus atom, presumably because of increased viscosity
imparted on the electrolyte by the additive. As mentioned, the oxidation state of the
phosphorus atom within the flame retardant molecules does not influence the
conductivity of the solutions as the solutions containing triphenyl phosphate (where P =
+5) and triphenyl phosphite (P = +3) exhibited nearly identical results at all temperatures.

140

Table 3.21 Conductivity measurements at different temperatures for electrolytes
containing various flame retardant additives

3.3.3.2.2 Electrochemical Impedance Spectroscopy

Similar to the previous sections that discussed the impedance growth for
electrolytes containing flame retardant additives, electrolytes examined in this study
exhibited a higher degree of impedance growth at the anode than they did at the cathode.
Data was obtained in the exact fashion as previous impedance measurements on fully
charged cells. Data was also processed using Z Simwin modeling software from
Princeton Applied Research as before. Further details regarding the testing can be found
in section 2.3.2.3 of this text.
EIS of Cathodes
As evidenced is Table 3.22, the nature of the electrolyte significantly influenced
the values of both relaxation loops for the cathode impedance. At room temperature, the
electrolyte employing TPPi exhibited significantly higher film impedance (0.3047 Ω)
than do any of the other electrolyte solutions (0.02715 Ω, 0.05291 Ω, 0.0661 Ω, and
0.08001 Ω for the electrolytes with DPP, TFPa, TFPi, and DEP respectively) (Table
Electrolyte Type
1.0M LiPF
6
in EC.EMC
(20.80 v/v%)
1.0M LiPF
6
in
EC.EMC.TEtPh
(20.75.5 v/v%)
1.0M LiPF
6
in
EC.EMC.TFEPh
(20.75.5 v/v%)
1.0M LiPF
6
in
EC.EMC.TPhPh
(20.75.5 v/v%)
1.0M LiPF
6
in
EC.EMC.TPhPi
(20.75.5 v/v%)
25 8.7 8.94 8.18 7.88 7.89
10 6.2 6.51 5.86 5.56 5.61
0 4.89 5.23 4.63 4.45 4.45
-10 3.7 4.05 3.45 3.36 3.37
-20 2.62 2.96 2.5 2.42 2.39
-30 1.685 2.06 1.69 1.588 1.56
-40 0.975 1.18 1.06 0.931 0.905
-50 0.487 0.698 0.61 0.47 0.453
-60 0.195 0.325 0.19 0.1893 0.137
Temperature
141
3.22). As the temperature was depressed, the TPPi electrolyte’s film impedance grew at
a slower rate than the other electrolytes though it is not until the temperature is lowered to
-40ºC that film impedances for the other electrolytes grew larger than that of the TPPi
electrolyte. Under all tested conditions, the charge transfer resistance of the TPPi was
higher than the other electrolytes. The TFPA electrolyte consistently demonstrated the
next highest charge transfer impedance values. The other electrolytes have similar
charge transfer impedance values at all examined temperatures. The link between
cathode impedance values and discharge performance is inconsistent. It is quite possible
that impedance growth at the cathode has little influence on performance of the cell. It
certainly has little to do with the cycle life of the cells as the electrolyte with the most
resistive film and charge transfer values (TPPi) exhibits the best capacity retention over
cycle life testing. It is possible that a weak link exists between the charge transfer
resistance of the electrolytes within a cell and the cell’s low temperature performance as
the TPPi electrolyte exhibits high charge transfer values and noticeably poor low
temperature performance. However, this link is not definitive as the DEP electrolyte
exhibited the best low temperature performance without the lowest charge transfer
cathode impedance.

142

Table 3.22 Cathode impedance values for electrolytes containing flame retardant additives

143
EIS of Anode
As expected, electrolyte composition significantly influenced both relaxation
loops at the anode. In contrast to its performance at the cathode, the electrolyte
containing TPPi exhibited the lowest film impedance at room temperature with a value of
0.02330 Ω. The order of anode film impedance was (from lowest to highest) TPPi, TFPa
(0.02466 Ω), DEP (0.03955 Ω), TFPi (0.05288 Ω), DPP (0.08105 Ω) (table 3.23). A link
exists between anode film impedance and cycle life performance as the electrolytes
employing TPPi and TFPa exhibit the lowest anode film impedance and retain the
greatest percentage of their starting capacities over cycle life testing. This is somewhat
expected as these two additives were predicted to exhibit the greatest electrochemical
stability and interfering least with the formation of the SEI film to the lowest degree of
the tested FRAs. In contrast to the findings at the cathode, there seemed to be little
relation between the impedance growth at the anode and the low temperature
performance of the electrolyte.

144

Table 3.23 Anode impedance values for electrolytes containing flame retardant additives

145
EIS of Full Cell
The impedance data for the full cell can be predicted the individual electrode
trends (Table 3.24). However, because the individual electrode data for the cathode and
anode are somewhat reversed in terms of which electrolytes perform the best for each
electrode, the full cell results were muddled and difficult to interpret. This fact illustrates
the benefits of incorporating a reference electrode into the system, which allows the
electrodes to be monitored individually.

146

Table 3.24 Full cell impedance values for electrolytes containing flame retardant additives

147
3.3.3.2.3 Tafel Polarization Measurements

Information regarding the lithium intercalation/de-intercalation kinetics of the
electrodes was gathered using Tafel polarization measurements. The fully charged
MCMB and LiNiCoO
2
electrodes were tested at temperatures ranging as low as -40 ºC.
The data was generated under potentiodynamic conditions, but at slow scan rates in order
to approximate steady-state conditions within the cells.
Room temperature measurements of the anodes at room temperature show that the
electrolyte with TFPa exhibits the highest current density (0.411 A) with the electrolyte
with TPPi following close behind (0.397 A). The current densities for the other
electrolyte follow in the order DEP (0.378 A), TFPi (0.329 A) and DPP (0.287 A). When
the temperature is depressed, the anode kinetics suffered dramatically with nearly a 90%
loss of current density for electrolytes containing DEP and DPP at -20 ºC (Table 3.25).
Relatively speaking, the electrolyte containing TPPi retains a significantly higher degree
of its charge density (~20%). Such a significant difference in current density retention
further substantiates the claim that TPPi affected the formation of the SEI at the anode in
the least deleterious manner.

148

Fig. 3.32 Anode Tafel measurements of electrolytes containing flame retardant additives

Table 3.25 Current density of the anode of electrolytes with flame retardant additives at
various temperatures recorded via Tafel measurements.

The current densities of the cathodes were lower for anodes of the same cells.
This was not unexpected as the cells were designed to be cathode limited to inhibit
lithium plating out of solution (Table 3.26). However, it is interesting to note that the
performance order of current density was nearly exactly opposite those at the anode
149
according to electrolyte with the DPP electrolyte exhibiting the highest current density
(0.254 A) and the TPPI electrolyte demonstrating the lowest (0.143 A). It is also
interesting to note that the cathodes do not lose nearly the current density with
temperature reduction as the anodes of the same cells. At -40ºC the cathode current
densities are quite similar, and in some cases, higher than their corresponding anode
current densities.

Fig. 3.33 Cathode Tafel measurements of electrolytes containing flame retardant
additives

150

Table 3.26 Current density of the anode of electrolytes with flame retardant additives at
various temperatures recorded via Tafel measurements.

3.3.3.2.4 DC Micro-Polarization Measurements

To further understand the charge transfer behavior of the electrodes in cells
containing flame retardant additive electrolytes, polarization resistance values of the
electrodes were calculated from the slopes of the linear plots generated under
potentiodynamic conditions of scar rates of 0.2 mV/s and measured over a +/-5mV range
from the equilibrium potential of each electrode. At the anode, the electrolyte exhibiting
least polarization growth as the temperature was depressed was EC.EMC.TPPi, followed
by the electrolyte with TFPa, TFPa, DPP. DEP exhibited the greatest polarization growth
at the anode indicating the passivating film is quite resistive compared to the other
electrolytes (Table 3.27). Values for the polarization at the cathodes of the same cells
further demonstrate that the electrolyte with TPPi did not form an easily passivated film
at the positive electrode (Table 3.28). This supports the data collected using EIS and
Tafel polarization measurements discussed in previous sections of this text. The other
151
electrolytes exhibit relatively similar polarization values at the cathode at room
temperature to one another. As the temperature was lowered, the electrolyte with TFPi
exhibited a considerably higher polarization at -40ºC, which may contribute to the
decreased discharge performance at reduced temperatures.

Table 3.27 Polarization values for the anodes of cells using flame retarding electrolytes

Table 3.28 Polarization values for the anodes of cells using flame retarding electrolytes

152
3.4 Conclusions

Numerous efforts were made to comprehend the effects of flame retardant
additives on electrolyte solutions. In order to accomplish this goal, discharge
performance data under a wide variety of conditions are collected. Additionally,
electrochemical methods were employed to gain insight on the effects these additives
impart on the individual electrodes of the Li-ion cells to which they are incorporated.
Following performance and electrochemical characterization, these cells were tested for
their cyclability over 100 charge/discharge cycles.
The success of TFEB as a co-solvent prompted us to evaluate the co-solvent into
electrolytes investigated to improve the safety characteristics of standard Li-ion cells via
introduction of flame retarding additives. In addition to TFEB, which should exhibit
desirable non-flammable characteristics due to halogenation present, four phosphorus-
containing compounds were examined at 5 v/v% in the electrolyte solutions. The branch
constitution around the phosphorus atom of the flame retardant molecules greatly
affected the performance of the cells at low temperature and over the life cycle of the
entire cell. Generally, electrolytes containing shorter alkyl-substituted phosphates
performed better at lower temperatures than electrolytes that contain the larger aromatic
phosphate additives. However, these trends were reversed when the cell was cycled at
room temperature with the smaller phosphates exhibiting poorer performance over time.
Through employment of electrochemical testing techniques it was found that the likely
source of the reduced performance of these electrolytes generally resulted from the
deleterious effects the additives imparted on the anodes, likely due to their incorporation
153
in the SEI at the electrode interface or reduced reductive stability at the anode. To
improve the cycle life of cells with flame retardant additives, vinylene carbonate was
added the electrolyte solution. Addition of this additive was found to improve the low
temperature performance of cells containing the flame retardant additive, triphenyl
phosphate. This additive imparted improved cycle life characteristics as well, reducing
the capacity fading over the life of the cell. Increased impedance growth at the anode
demonstrated VC decomposed at the electrode to form a layer that protected the flame
retardant additive from the active material. A final study of flame retardant additives was
performed to analyze the effect branch substitution around and oxidation state of the
phosphorus atom within flame retardant additives had on the electrochemical and
performance characteristics of the cell. Aromatic branches and P(III) oxidation states for
the flame retardant additives increase the performance of the cell over the long term, but
the same bulky substituents for these molecules led to reduced low temperature
performance. This reduced performance at low temperatures was most likely due to
deleterious mass transfer effects imparted upon the electrolyte by the addition of bulky
flame retardant additives.

154
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157
CHAPTER 4

Electrolyte Salt Modification for Improved Li-Ion Cell
Performance

4.1 Introduction

Additives incorporated into the electrolyte mixture at levels typically <5% by
volume are often able to impart significant positive influence on the performance of Li-
ion cells. Numerous types of Li-ion electrolyte additives exist and exhibit influence on a
wide variety of properties important to these cells, including: SEI improvements and
Lewis-base salt stabilizers. Other types of additives exist including flame-retarding
additives that were discussed in the previous chapters of this text. Here, we discuss the
effect of adding of small amounts of boron salts to the total electrolyte salt composition.
The two salts, lithium bis(oxalato) borate (LiBOB) and LiBF
4
affect the performance in
different ways with LiBOB being found to improve the stability of the SEI, while LiBF
4
leads to improved extreme temperature performance. In order to evaluate the validity of
both salts as additives, performance characteristics of cells employing the electrolyte salts
at 0.2M of each of these salts in combination with 1.0M LiPF
6
are monitored for their
low temperature performance and high temperature resiliency. Additionally, the
electrochemical properties these salts imparted onto their corresponding electrodes is also
described and discussed.
158
4.1.1 Additives for Improving Li-ion Cell Performance

Methods exist to adjust the properties of the electrolyte, beyond modifying the
bulk electrolyte. A concept that has received significant study is the use of electrolyte
additives. The term “additive” generally refers to any compound that is charged to the
total electrolyte at less than 5% by volume. At such low concentrations, these materials
often do not significantly alter the physical properties, yet even at these low
concentrations these materials are capable of imparting significant influence upon the
properties and performance of the cell as a whole. Most commonly, additives are chosen
to improve efficiency and cycle life of cells. The method of action for these additives
designed to improve cell performance generally fall into one of the following categories:
1

1. Additives that facilitate improved SEI at the surface of the anode
2. Additives that reduce irreversible capacity loss and gas generation during
the SEI formation and over long-term cycling
3. Additives that enhance thermal stability of the LiPF
6
toward the organic
electrolyte solvents
4. Additives that protect cathode material from dissolution and overcharge
5. Additives that improve properties such as ionic conductivity, viscosity,
wetting ability to the separator, and ion solvation capability
Additionally, several additives have been designed and tested to improve the
safety characteristics of Li-ion cells. These will be addressed in a later section.

159
4.1.2 Solid-Electrolyte Interface Enhancing Additives

Since the serendipitous discovery by Dahn et al.
2
regarding EC’s ability to form a
stable protective layer on a graphitic anode, researchers have thrown themselves toward
understanding the nature of this layer and how to improve it. Investigation of the SEI
using various spectroscopic techniques has identified the components of this layer to
include decomposition products of the electrolyte solvents and salts, such as Li
2
CO
3
,
lithium alkyl carbonates, and lithium alkoxides. Other salts such as LiF were also found
in electrolytes that employed LiPF
6
as the salt.
3
With the knowledge that EC plays an
integral role in the formation of the SEI, two mechanisms have been proposed for
reduction of carbonate-based solvents (fig. 4.1). Spectroscopic analysis indicates that
both mechanisms participate in the formation of the SEI in a competing fashion.
Mechanism (I) generates a Li
2
CO
3
dense SEI and more gaseous products. Production of
such by-products results in a SEI that is less stable than the more compact layer produced
in mechanism (II), which is composed primarily of insoluble materials.
1

Fig. 4.1 Proposed SEI formation reactions. Mechanism (II) leads to a more stable SEI.
160
Numerous SEI modifying additives have been investigated in order to produce a
more stable layer. One type of SEI additive that has been investigated can be considered
a reduction-type additive. These additives work by chemically coating an organic film
onto the surface of the graphitic anodes through electrochemical reduction of the
additives to produce a more robust SEI layer. These additives have higher reduction
potentials than the electrolyte solvents to which they are added in order to assure they are
reduced to form an insoluble product to cover the surface of the anode. This film acts as
a protective layers against the catalytic activity of the graphite to the electrolyte solvent
leading to reduced gas generation and increased overall SEI stability due to the inclusion
of the additive moieties into the SEI. Reduction-type additives can further be broken
down into two subgroups, reductive agents and polymerizable monomers.
Reductive agents assist SEI formation via adsorption of their reduced products
onto the active graphitic sites. Many of these additives contain sulfur and their efficacy
seems to increase with higher sulfur content, indicating a mode of action similar to the
sulfur poisoning effect that typically deactivate many catalysts. Examples of such
catalysts include SO
2
, CS
2
, polysulfide, cyclic alkyl sulfites and aryl sulfites.
4
Restraint
must be demonstrated when determining the loading content of these additives because
they are anodically unstable at high potentials and their presence may result in a high
self-discharge rate as a result of the internal redox shuttle.
Polymerizable monomers used as SEI additives contain one or more double
bonds. Examples of these additives include vinylene carbonate, vinyl ethylene carbonate,
allyl ethyl carbonate, vinyl acetate, divinyl adipate, acrylic acid nitrile, 2-vinylpyridine,
succinic anhydride, maleic anhydride, vinyl phosphonates, vinyl–containing silane
161
compounds, and furan derivatives.
5
The effectiveness of such additives at improving the
SEI layer stems from the reduction of gas generation that leads to irreversible capacity
loss, and stabilization of the SEI against extended cycling. Not surprisingly, the
mechanism of these additives is based on the electrochemically-induced polymerization
of the double-bond moieties where the radical anion can be terminated by a solvent
molecule to form an insoluble and stable product as the preliminary SEI layer. This
reductive polymerization is quite well known in synthetic chemistry and very effective.
Care must be taken when developing electrolyte solutions that employ the use of such
additives as loading such molecules at a high level can lead to oxidative polymerization
at the cathode, which leads to increased impedance and irreversibility of the electrode
performance.

Fig. 4.2 Anodic polymerization of additives containing double bonds occurs via a similar
mechanism and can help in the formation of robust SEI.

4.1.3 Electrolytes Utilizing Boron Salts

Another method of modifying the electrolyte system to improve the SEI layer is
to include boron-based salts in the electrolyte salt formulations. Boron-based compounds
improve the characteristics of cells in numerous ways, including increasing the cycle life
of Li-ion cells by stabilizing the SEI. Additionally, boron compounds have been found to
n
162
improve the rate capability and low temperature performance in some Li-ion cells.
6

Numerous boron salts have been used in this manner, although they have multiple means
of action. One such salt in particular, lithium bis(oxalato) borate (LiBOB), has received
significant attention.
7
It was initially investigated as a salt to improve the high
temperature resiliency of Li-ion cells; however, cells containing LiBOB were also found
to have a stabilized SEI layer for extended cycling and excellent overcharge tolerance. It
has been proposed that the LiBOB breaks down on the anode to form oligomers and
polymers by reacting with other SEI components such as lithium bicarbonate and lithium
alkoxide to form the more stable SEI. Spectroscopic methods verified that B-O moieties
were incorporated into the SEI. Not all data pertaining to LiBOB is positive, and some
drawbacks arise when LiBOB is incorporated into a Li-ion cell. The protective
oligomeric layer LiBOB forms on the anode and stabilize the SEI is quite resistive and
consequently, reduces power and rate capability of the cell, especially at reduced
temperatures.
The structurally similar lithium oxalyldifluoroborate (LiODFB) works with a
mechanism similar to LiBOB, but exhibits preferable qualities that make it a favorable
alternative to LiBOB.
8
These characteristics include improved carbonate solubility, better
rate capability and low temperature performance. LiBOB and LiDFOB are example of a
series of additives that work to improve the performance of the cell without electronic
transference. These additives, “reactive-type additives” as Zhang calls them
1
, react with
intermediates or radicals of the solvent reduction or combine with the final products that
compose the SEI layer. The end product in all cases leads to a more stable SEI layer that
leads to improved cycling.
163
Another boron-based lithium salt, LiBF
4
, has significantly different properties.
The salt was long disregarded for use in lithium batteries due to poor ionic conductivity
resultant from a smaller dissociation constant than the widely used LiPF
6
9
; however, due
to the thermal instability and moisture sensitivity of LiPF
6
, LiBF
4
has received attention
and demonstrated improved performance over a wider temperature range, up to 50
0
C,
improved performance at low temperatures due to low charge-transfer resistance, and
reduced water sensitivity.
10
Drawbacks remain in utilizing LiBF
4
. Because of the
stability of the BF
4
-
anion it is unable to participate in the formation of the SEI. Also,
electrolytes with LiBF
4
salt have an inability to form a homogenous solution at very low
temperatures (-60
0
C) because of the symmetry and small size of the anion.
10b,c

In the same way LiODFB has been examined as an evolution of LiBOB, LiBF
3
Cl
has been investigated as an evolutionary LiBF
4
. Zhang has claimed that the asymmetry
present in LiBF
3
Cl leads to improved solubility compared to LiBF
4
. Additionally, the
weaker coordination of the chlorine atom to boron results in a chemical equilibrium
shown below that can participate in SEI formation.
11

BF
3
Cl
-
BF
3
+ Cl
-

4.2 Experimental Methods

Cell preparations and all performance and electrochemical tests were performed
according to details outlined in section 2.2 of this manuscript. Boron salt candidate were
164
purchased from suppliers at battery grade from various sources and used without any
further purification.

4.3 Results and Discussion

As previous studies have reported promising properties for electrolytes that
employed boron salts, we sought to investigate for ourselves the affects the addition of
these salts have upon the reversibility, low temperature performance and high
temperature resiliency of Li-ion cells. Additionally, the MCMB carbon anodes and
LiNi
0.8
Co
0.2
O
2
cathodes were monitored to determine the electrochemical aspects the
addition of boron salts had on the individual electrodes before and after high temperature
storage of the cells. The electrolytes selected for evaluation were a baseline solution that
employs only the common electrolyte salt, LiPF
6
, at 1.0 M concentration and two
electrolytes that had boron salts added to the baseline solution. The exact formulation of
each electrolyte was as follows:

1. 1.0 M LiPF
6
in EC.EMC.MB (20.60.20 v/v %)
2. 1.0 M LiPF
6
& 0.2 M LiBOB in EC.EMC.MB (20.60.20 v/v %)
3. 1.0 M LiPF
6
& 0.2 M LiBF
4
in EC.EMC.MB (20.60.20 v/v %)

Methyl butyrate was incorporated into the electrolyte solvent system because
previous studies at JPL found the addition of the ester improved the low temperature
performance of Li-ion cells. Electrochemical evaluation methods employed in this study
165
were Tafel polarization and DC micro-polarization methods to determine the lithium
intercalation and de-intercalation kinetics and EIS measurements to study the relative
interfacial stability of the candidate electrolytes, which were of particular interest due to
inclusion of boron salts into the electrolyte mixture. All analyses were performed before
and after high temperature storage periods.

4.3.1 Experimental Cell Results

4.3.1.1 Formation Characteristics

Initial assessment of charge/discharge characteristics of the cells was performed
through monitoring the irreversible and reversible capacities during the critical formation
cycles. While the cumulative irreversible capacity loss for the cells containing the boron
salts in the electrolyte systems was considerably higher than for the baseline, most of the
difference was accounted for in the first cycle. The baseline irreversible capacity loss
was only 39.8 mA while the capacity loss was 75.5 mA and 79.7 mA for the cells using
LiBF
4
and LiBOB, respectively (Table 4.1). As this is the cycle where the greatest
percentage of SEI formation takes place it is likely that the electrolytes plays a significant
role in the formation of the interfacial layer. By the fifth cycle, the irreversible capacity
loss for the cells is similar, and the cells exhibit nearly identical coulombic efficiencies.
After the initial formation data of the cells were collected, they were subjected to
electrochemical evaluation over a wide temperature range (25ºC to -60ºC) followed by
low temperature discharge testing.
166

Table 4.1 Formation data for cells with new electrolyte salt compositions

167
4.3.1.2 Low Temperature Discharge Characteristics

The cells containing the various electrolyte salt formulations were evaluated at
reduced temperatures with a variety of rates. Cells are charged at room temperature and
soaked at the testing temperature for >4 hours. As shown in Fig.4.3, both cells
containing boron salt in their electrolyte slightly outperformed the baseline at moderately
reduced temperatures and rate (-20 ºC at 25 mA/h). All cells perform quite well under
these conditions retaining >85% of their room temperature capacity. When the
temperature was lowered to -40ºC, the LiBOB electrolyte maintained better performance
than the baseline (75.23% v. 72.46% of room temperature capacity, respectively), but the
cell with BF
4
(72.04% of room temperature capacity) no longer exhibited any
performance enhancement over the baseline (Fig. 4.4). In fact, while the BF
4
electrolyte
maintains relatively good performance down to -20ºC, the cell using the electrolytes
performed the worst of the tested cells at temperatures below this mark. The cell with the
LiBOB electrolyte consistently maintained the greatest capacity retention (Table 4.2).
168

Fig. 4.3 Capacity retention of cells with boron salts at -20ºC discharged at 25 mA/h

Fig. 4.4 Capacity retention of cells with boron salts at -40ºC discharged at 25 mA/h

169

Table 4.2 Low temperature discharge capacities and capacity retention for cells with new electrolyte salt mixtures.

170
4.3.1.3 High Temperature Resiliency of Electrolytes

Both boron salts investigated in this study have been hypothesized to lead to
greater high temperature resiliency of their respective electrolytes, albeit through
different mechanisms. LiBOB is thought to polymerize upon decomposition at the anode
leading to a film that protects the electrolyte from decomposition at the negative
electrode. LiBF
4
simply has better thermal stability characteristics than does the common
LiPF
6
electrolyte.
In addition to monitoring the electrochemical characteristics of the cells after high
temperature exposure (discussed in later sections), discharge studies of the cells charged
at various rates at room temperature and at -20ºC after successive high temperature
storage periods were performed. The cells were stored for 10-day periods at 55ºC, 60ºC,
and finally, 65ºC. Table 4.2 shows the results of the discharge tests performed after each
of the storage tests. Prior to storage, each of the cells exhibited similar discharge
characteristics, retaining a similar percentage of their starting capacity regardless of
charge or discharge rates. However, with each successive high temperature storage
period, the electrolyte with LiBOB and (especially) LiBF
4
exhibited poorer and poorer
discharge characteristics relative to the baseline cell (Table 4.2).

171

Table 4.3 Discharge capacities and capacity retention for the cells with new electrolyte salt compositions after high
temperature storage periods. Cells were charged and discharged at various rates at room temperature.

172
Results were somewhat mixed when the tests were performed at reduced
temperatures (Table 4.4). Cells were initially charged at room temperature then lowered
to -20ºC and subsequent discharging and charging were performed at this temperature.
The LiBOB electrolyte cell somewhat increased the low temperature performance of the
cells relative to the baseline after high temperature storage. After the initial 55ºC, the
LiBOB electrolyte gave improved capacity retention relative to the baseline, especially
when the cells are charged at the reduced temperature. This trend was present after the
third high temperature storage period at 65ºC; although the capacity retention was so
small it hardly seems important in the overall scheme. After the second, 60ºC storage,
the baseline performed the best of tested electrolytes, though only by a small margin.
The electrolyte with LiBF
4
performed poorest after high temperature storage. After the
third storage period the cell had degraded to such an extent that it became polarized
instantly and incapable of discharging any capacity under the tested conditions.

173

Table 4.4 Discharge capacities and capacity retention for the cells with new electrolyte salt compositions after high
temperature storage periods. Cells were charged and discharged at various rates at -20ºC.

174
4.3.2 Electrochemical Characteristics

The addition of boron salts to the electrolyte system had a dramatic effect
on the cell. To more fully understand the nature of these effects, various electrochemical
analyses were performed. It was believed that the addition of the salts would have a
significant effect upon the electrode characteristics both in terms of lithium
intercalation/de-intercalation kinetics and film formation behavior. The electrochemical
techniques such as electrochemical impedance spectroscopy (EIS), Tafel polarization and
DC micro-polarization measurements were used to more fully understand the effects
addition of boron salts had on the cell.

4.3.2.1 Electrochemical Impedance Spectroscopy

As previously stated within this text, generally the effects of the cathode dominate
the interfacial impedance of the total cell. However, the addition of boron salt was
expected to have a greater effect on the anodes of the cell. Electrochemical impedance
spectroscopy was employed to monitor the impedance growth caused by the addition of
the salts. In particular, the effects of low temperature and high temperature storage
periods on the film and charge-transfer resistances were examined. EIS data was
collected in the same manner previously described in this text (section 2.3.2.3)
EIS of Cathodes
Addition of boron salts significantly affects the first and second relaxation loops
of the cathode Nyquist plots suggest that the electrolyte plays a significant role
175
determining the interfacial properties as can be seen in Fig. 4.5. This translates into
significant influence of the electrolyte salt upon the film and charge transfer resistances
of the electrode by the salts. As can be seen in Fig. 4.5, not only was the magnitude of
the relaxation loops of the cathodes affected by the addition of the boron salts, but so too
was the shape significantly altered. In particular, the electrolyte with LiBOB had a
relaxation loop that loops below zero. It is unknown what phenomenon would lead to
such results, and further investigation is needed for full understanding. At room
temperature, the LiBOB electrolyte gave the most favorable impedance results in terms
of both film resistance and charge-transfer resistance. The film resistance for this
electrode was ~0.035 Ω. (This value is an estimate as it was manually due to the unique
shape of the impedance curve rather than through the use of Z Simwin computer program
used to calculate other impedance values.) The film resistance of the baseline mixture
was 0.07674 Ω, and the LiBF
4
containing electrolyte exhibited the highest film resistance
characteristics on the cathode at 0.1655 Ω. The LiBOB electrolyte exhibited the lowest
charge transfer resistance of the three cells (0.01915 Ω). The baseline and LiBF
4

electrolytes exhibited similar charge transfer resistances to one another (0.03441 Ω and
0.03451 Ω, respectively). As the temperature of the cells was lowered, the impedance
trends remained consistent with the LiBOB cell exhibiting the lowest film and charge
transfer resistances at the cathode. The baseline exhibited lower film resistances than the
LiBF
4
electrolyte, but their charge transfer impedance values remain similar no matter the
testing temperature (Table 4.5).
The impedance values were monitored after each high temperature storage period.
The values of the film and charge transfer impedance grew significantly during these
176
tests. After each high temperature storage periods, the impedance values were
significantly higher for each of the cells, but the order of performance remained the same
as room temperature. The LiBOB cell exhibited the lowest film and charge transfer
resistances (2.663 Ω and 0.3964 Ω) followed by the baseline (3.292 Ω and 1.173 Ω) and
the LiBF
4
cell (7.333 Ω and 3.439 Ω). The film resistances for each cell grew at a rate
significantly higher than the charge transfer impedances (Table 4.6). The shape of the
relaxation loops of the Nyquist changed after the cells under went exposure to high
temperatures, (Fig. 4.6). It seems that not only does the film on the cathode grow
significantly as expected, but also its composition is changed in some unknown matter
(composition, morphology); however, we do not have the capability to test this
hypothesis.
177
Fig. 4.5 Room temperature Nyquist impedance data of cathodes of cells containing new
electrolyte salt compositions
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' (Ohms)
1.00 M LiPF6 EC:EMC:MB (20:60:20 v/v %)
1.00 M LiPF6 & 0.2 M LiBOB in EC:EMC:MB (20:60:20 v/v %)
1.00 M LiPF6 & 0.2 M LiBF4 in EC:EMC:MB (20:60:20 v/v %)
Temperature = 23
o
C
Li-ion Experimental Three Electrode Cell
MCMB Carbon - LiNi
x
Co
1-x
O
2
Lithium metal Reference Electrode
178

Table 4.5 Cathode impedance data of new salt composition cells at low temperatures

179

Table 4.6 Impedance data for cathodes of cells with new electrolyte salt composition after high temperature
storage periods

180
Fig. 4.6 Nyquist impedance plots of cathodes after third high temperature storage period.
The shape of the curves changed significantly after storage.

EIS of Anodes
Anode impedance data is also significantly influenced by electrolyte composition.
From the room temperature data, the baseline electrolyte exhibits the lowest film
impedance (0.02033 Ω) and charge transfer impedance (0.04267 Ω), though the LiBF
4

anode trailed only by a small of margin for both impedance values (0.02178 Ω and
0.04848 Ω). Both impedance values for the LiBOB electrolyte are significantly higher
than the other electrolytes. This is not unexpected as LiBOB is thought to decompose
and polymerize at the anode, though the degree to which the impedance values grow
compared to the baseline is somewhat staggering (Fig. 4.7).
0.00
1.00
2.00
3.00
4.00
5.00
6.00
7.00
8.00
0.00 2.00 4.00 6.00 8.00
Z' (Ohms)
1.00 M LiPF6 EC+EMC+MB (20:60:20 v/v %)
1.00 M LiPF6 & 0.2 M LiBOB in EC+EMC+MB (20:60:20 v/v %)
1.00 M LiPF6 & 0.2 M LiBF4 in EC+EMC (20:80 v/v %)
Temperature = 23
o
C
Li-ion Experimental Three Electrode Cell
MCMB Carbon - LiNi
x
Co
1-x
O
2
Lithium metal Reference Electrode
181
The trends that were observed at room temperature are also seen as the
temperature was depressed, though exacerbated as expected. At all tested temperatures,
the baseline and LiBF
4
electrolyte had much lower values both in terms of film and
charge transfer impedances than the LiBOB anode (Table 4.7). The data supported the
claim that LiBOB decomposes at the anode causing significant film growth.
When the anodic data is analyzed after the cells are subjected the high
temperature storage periods it was obvious that the addition of LiBOB offered some
degree of protection in terms of impedance growth. Both the baseline and LiBF
4
have
impedance values that grew at a significantly faster rate than did the LiBOB cell,
especially regarding to the film resistance (Table 4.8).

Fig. 4.7 Nyquist plot of anode impedance data collected for cells with new electrolyte salt
composition 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' (Ohms)
1.00 M LiPF6 EC:EMC:MB (20:60:20 v/v %)
1.00 M LiPF6 & 0.2 M LiBOB in EC:EMC:MB (20:60:20 v/v %)
1.00 M LiPF6 & 0.2 M LiBF4 in EC:EMC:MB (20:60:20 v/v %)
Temperature = 23
o
C
Li-ion Experimental Three Electrode Cell
MCMB Carbon - LiNi
x
Co
1-x
O
2
Lithium metal Reference Electrode
182

Table 4.7 Impedance data for anodes of cells with new electrolyte salt composition at low temperatures

183

Table 4.8 Impedance data for anodes of cells with new electrolyte salt composition after high temperature storage
periods

184
EIS of Full Cells
The impedance data collected for the full cell can be predicted from the trends
found at the individual electrodes, though the anode is observed to influence the
impedance data to a greater degree. At low temperature, the baseline had much lower
impedance than either of the cells containing a boron salt, a trend easily predicted from
the anode data (Table 4.9). After high temperature storage, the impedance properties of
the LiBOB electrolyte are more favorable than the baseline, and significantly better than
the cells with LiBF
4
electrolyte (Table 4.10).

185

Table 4.9 Impedance data for the full cell with new electrolyte salt composition at low temperatures

186

Table 4.10 Impedance data for full cells with new electrolyte salt composition after high temperature storage periods
187
4.3.2.2 Tafel Polarization Measurements

Tafel polarization measurements of the MCMB-LiNiCoO
2
cells with the
described electrolyte compositions was performed the temperature range of 23ºC to -60ºC
and again after each of the high temperature storage periods. Generated under
potentiodynamic conditions, but at slow scan rates, the data collected supports
conclusions drawn from the EIS data. At room temperature, the cathode of the cell with
the LiBOB electrolyte demonstrated the more favorable lithium intercalation kinetics
with higher current density (0.299 A) than the two other electrolytes (0.209 A for the
baseline and 0.188 A for the LiBF
4
electrolyte, respectively) examined in this study (Fig.
4.8). As with the EIS data, the LiBOB cell continues to have the most preferable
characteristics at the cathode when compared to the other electrolytes, as the cathode
from this cell demonstrates higher current density all the way to -60ºC (Table 4.11). This
data demonstrates a link between the film growth on the cathode and the lithium
intercalation kinetics at the same electrode.
EIS trends also manifested themselves at the Li kinetics of the anode where the
baseline cell demonstrated the greatest lithium de-intercalation kinetics of the examined
cells (0.368 A current density), a trend that continued as the temperatures is depressed.
Unlike the EIS studies where the LiBF
4
cell demonstrated the more preferable
characteristics at room temperature of the two cells with boron salts added (Fig. 4.9),
Tafel studies at the anode show the LiBOB cell to demonstrate superior lithium de-
intercalation kinetics at room temperature (0.3 A and 0.24 A for LiBOB and LiBF
4
,
respectively). However, the LiBOB anode demonstrates greater sensitivity to low
188
temperatures, and when the temperature was depressed demonstrated the poorest lithium
kinetics of any tested cell at reduced temperatures. This fact was supported by high
impedance at reduced temperatures using EIS (Table 4.12).

Fig. 4.8 Cathode Tafel measurement data of cells with new electrolyte salt formulations

Table 4.11 Current density (A) of the cathodes from cells with new electrolyte salt
formulations as a function of temperature

189

Fig. 4.9 Anode Tafel measurement data of cells with new electrolyte salt formulations

Table 4.12 Current density (A) of the anodes from cells with new electrolyte salt
formulation as function of temperature

Tafel measurements were made after each high temperature storage period. As
expected, the lithium kinetics greatly suffered at both electrodes after the cells are
subjected to elevated temperatures. Table 4.13 shows the current density of each cell’s
cathode after each 10-day storage period. It is clear that after each period, the lithium
190
intercalation kinetics were reduced considerably. The cell that loses the greatest
percentage of its cathode current density after the first high temperature storage period
was the cell with the LiBOB electrolyte. Initially, this cell demonstrates the highest
current density at the anode, but it suffers from the greatest reduction to its intercalation
kinetics, losing 76.3% of its current density after the first storage period. The baseline
also suffers greatly in terms of cathode kinetics, losing 70.7% of its current density over
the same period. By comparison, the LiBF
4
cell retained relatively decent lithium
kinetics at its cathode, losing only 58.7% of its current density after the first storage
period (table 4.13). Not surprisingly, subsequent high temperature storage periods lead to
further reduction in kinetic capabilities at the cathode, and after the third high-
temperature storage period all of the cells lose approximately 90% of the current density
they displayed prior to being exposed to high temperatures.

Table 4.13 Current density (A) of the cathodes from cells with new electrolyte salt
formulations after high temperature storage periods

Results at the anode vary significantly from the cathode (Table 4.14). As
expected, addition of LiBOB to the cell’s electrolyte offered a degree of protection
against degradation at the anode. After the first high temperature storage, the current
191
density at the anode was reduced by only 31.8% compared to 48.7% for the baseline and
69.9% for the electrolyte with LiBF
4
. This is further evidence that the LiBOB
decomposes to form a protective layer at the anode. Though evidence of this protective
layer is evident, there does appear to be limits to its robustness as subsequent high-
temperature storage period cause the lithium anode kinetics to degrade to nearly the level
of the other tested cells. After the third storage period, the LiBOB anode lost 89.2% of
its original current density compared to 92.7% for the baseline and 97.0% for the LiBF
4

cell.

Table 4.14 Current density (A) of the anodes from cells with new electrolyte salt
formulations after high temperature storage periods

4.3.2.3 DC Micro-Polarization Measurements

The changes in charge transfer behavior of the passivating films of cells with
boron salts was studied using DC micro-polarization techniques. The polarization
resistance values of the electrodes were calculated from the slopes of the linear plots
generated under potentiodymic conditions of scan rates of 0.2 mV/s and measured over a
+/- 5mV range from the equilibrium potential of each electrode as in previous analyses.
192
As with previous electrochemical evaluation techniques, the effect of low temperature
and high temperature storage periods upon the individual electrodes was monitored.
When the polarization values are measured at the cathode, the results supported
previously observed data from the EIS and Tafel polarization studies. Improved kinetic
performance was seen from the cathode of the cell with LiBOB electrolyte. At room
temperature, the polarization resistance of the cathode was 0.351 kΩ, compared to 0.455
kΩ for the baseline and 0.638 kΩ for the electrolyte containing LiBF
4
. Similar trends at
the cathode existed as the temperature was lowered (Table 4.15). The anode performance
of the baseline was the most favorable of the tested cells with the lowest polarization
resistance at ambient and depressed temperatures. The LiBF
4
electrolyte cell performed
nearly as well as the baseline under these conditions, while the LiBOB demonstrated
considerable polarization resistance growth confirming the development of a resistive
film due to the presence of the boron salt (Table 4.16).

Table 4.15 Polarization resistance values (kΩ) for the cathode of cells containing new
electrolyte formulations at reduced temperatures

193

Table 4.16 Polarization resistance values (kΩ) for the anodes of cells containing new
electrolyte formulation at reduced temperatures

After each high temperature storage period, the polarization resistance of each
electrode grew considerably, as expected. At the cathode, the LiBOB electrolyte
exhibited the greatest growth of 1.404 kΩ after a single high temperature storage period.
However, the rate of polarization growth at the cathode was inconsistent for every cell
and during subsequent high temperature storage periods, both the baseline and LiBF
4

cathodes exhibited higher growth rates (Table 4.17). At the anode, as expected, the
addition of the LiBOB salt imparted some degree of protection upon the electrode, and
although the LiBOB anode began with the highest polarization resistance (0.805 kΩ), this
value grew at a significantly lower rate than the other cell’s values of the same electrode
(Table 4.18).

Table 4.17 Polarization resistance values (kΩ) for the cathodes of cells containing new
electrolyte formulation after high temperature storage periods

194

Table 4.18 Polarization resistance values (kΩ) for the anodes of cells containing new
electrolyte formulations after high temperature storage periods

4.4 Conclusions

To understand the effect the addition of boron salts to an electrolyte solution, both
0.2 M of LiBOB and LiBF
4
were added to a standard electrolyte solution of 1.0 M LiPF
6

in EC.EMC.methyl butyrate (20.60.20 v/v%). Improved performance due to the addition
of these boron salts following high temperature storage periods was of particular interest.
It was found that addition of LiBOB to the electrolyte solution offered improved high
temperature resiliency due to polymerization of the salt’s anion at the anode forming a
protective layer impeding further electrolyte decomposition.

195
4.5 Chapter 4 References

1. Zhang, S. S. J. Power Sources, 2006, 162, 1379.
2. Fong, A.; von Sacken, U.; Dahn, J. R. J. Electrochem. Soc. 1990, 137, 2009.
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Electrochem. Soc., 1997, 144, L180. (b) Aurbach, D.; Ein-Eli, Y. J. Electrochem.
Soc., 1995, 142, 1746,
4. (a) Ein-Eli, Y.; Thomas, S. R.; Koch, V. R. J. Electrochem. Soc., 1996, 143,
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W.; Liebenow, C.; Besenhard, J. O. J. Power Sources, 1997, 68, 328. (e)
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Mogi, R.; Inaba, M.; Jeong, S. K.; Iriyama, Y.; Abe, T.; Ogumia, Z. J.
Electrochem. Soc., 2002, 149, A1578. (g) Wrodnigg, G. H.; Besenhard, J. O.;
Winter, M. J. Power Sources, 2001, 97-98, 592. (h) Wrodnigg, G. H.; Wrodnigg,
T. M., Besenhard, J. O.; Winter, M.; Electrochem. Comm., 1999, 1, 148.
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Acta, 2007, 53, 650. (b) Shim, E. G.; Nam, T. H.; Kim J. G.; Kim, H. –S.; Moon,
S. –I. J. Power Sources, 2007, 172, 901. (c) Nam, T. H.; Shim, E. G.; Kim J. G.;
Kim, H. –S.; Moon, S. –I. J. Electrochem. Soc., 2007, 154, A957. (d) Wang, Z.;
Hu, Y.; Chen, L. J. Power Sources, 2005, 146, 51. (e) Abe, K.; Yoshitake, H.;
Kitakura, T.; Hittori, T.; Wang, H.; Yoshio, M. Electrochim. Acta., 2004, 49,
4613. (f) Hu, Y.; Kong, W.; Wang, Z. X.; Li, H.; Huang, X.; Chen, L.
Electrochem. Comm., 2004, 6, 126. (g) Aurbach, D.; Gnanaraj, J. S.; Geissler, W.;
Schmidt, M. J. Electrochem. Soc., 2004, 151, A23. (h) Aurbach, D.; Markovsky,
B.; Rodkin, A.; Levi, E.; Cohen, Y. S.; Kim, H. –J.; Schmidt, M. Electrochim.
Acta, 2002, 47, 4291. (i) Lee, J. T.; Lin, Y. W.; Jan, Y. S. J. Power Sources, 2004,
132, 244. (j) Santner, H. J.; Moller, K. C.; Ivanco, J. Ramsey, M. G.; Netzer, F.
P.; Yamaguchi, S.; Besenhard, J. O.; Winter, M. J. Power Sources, 2003, 119-
121, 368. (k) Komaba, S.; Itabashi, T.; Ohtsuka, T.; Groult, H.; Kumagai, N.;
Kaplan, B.; Yashiro, H. J. Electrochem. Soc., 2005, 152, A937. (l) Lee, H.; Choi,
S.; Choi, S.; Kim, H. –J.; Choi, Y.; Yoon, S.; Cho, J. –J. Electrochem. Comm.,
2007, 9, 801. (m) Ufheil, J..; Baertsch, M. C.; Wursig, A.; Novak, P. Electrochim.
Acta, 2005, 50, 1733. (n) Schroeder, G.; Gierczyk, B.; Waszak, D.; Kopczyk, M.;
Walkowiak, M. Electrochem. Comm., 2006, 8, 523. (o) Korepp, C.; Santner, H. J.;
Fuji, T.; Ue, M.; Besenhard, J. O.; Moller, K. C.; Winter, M. J. Power Sources,
2006, 158, 578.
196
6. Contestabile, M.; Morselli, M.; Praventi, R.; Neat, R. J. J. Power Sources, 2003,
119-121, 943.
7. (a) Xu, K.; Zhang, S. S.; Jow, T. R.; Xu, W.; Angell, C. A. Electrochem. Solid-
State Lett., 2002, 5, A26. (b) Xu, K.; Lee, U.; Zhang, S. S.; Allen, J. L.; Jow, T. R.
Electrochem Solid-State Lett., 2004, 7, A273. (c) Xu, K.; Lee, U.; Zhang, S. S.;
Jow, T. R. J. Electrochem. Soc., 2004, 151, A2106. (d) Xu, K.; Zhang, S. S.; Lee,
U.; Allen. J. L.; Jow, T. R. J. Power Sources, 2005, 146, 79.
8. (a) Zhang, S. S. Electrochem. Comm., 2006, 8, 1423. (b) Zhang, S. S. J. Power
Sources, 2007, 163, 713.
9. (a) Ue, M. J. Electrochem. Soc., 1994, 141, 3336. (b) Ue, M.; Mori, S. J.
Electrochem. Soc., 1995, 142, 2577.
10. (a) Zhang, S. S.; Xu, K.; Jow, T. R. Electrochem. Comm., 2002, 4, 928. (b)
Zhang, S. S.; Xu, K.; Jow, T. R. J. Solid State Electrochem., 2003, 7, 147. (c)
Zhang, S. S.; Xu, K.; Jow, T. R. J. Electrochem. Soc., 2002, 149, A586.
11. Zhang, S. S. J. Power Sources, 2008, 180, 586.
12. Zhang, S. S.; Xu, K.; Jow, T. R. Electrochem Solid State Lett, 2002, 5, A206.
13. (a) Wang, X.; Naito, H.; Sone, Y.; Segami, G.; Kuwajima, S. J. Electrochem.
Soc., 2005, 152, A1996. (b) Li, W.; Campion, C.; Lucht, B. L.; Ravdel, B.;
DiCarlo, J.; Abraham, K. M. J. Electrochem. Soc., 2005, 152, A1361. (c) Appel,
W.; Pasenok, S. U.S. Patent 6,159,640, 2000.

197
CHAPTER 5

PVP-SO
2
Complex as a Solid Mild Acid Catalyst for Efficient
One Pot, Three Component Synthesis of α α α α-Aminonitriles

5.1. Introduction

Polymer supported reagents are becoming important and useful tools in synthetic
organic chemistry.
1
The solid support, not only makes the reactions simple, easy and
environmentally safer, but also helps sometimes to fine-tune the reactivity of the reagents
towards various synthetic transformations. Recently, we were interested in using poly(4-
vinylpridine) (PVP) as a solid support for various acidic gaseous and liquid reagents and
catalysts. During the course of our studies, we found that solid PVP and SO
2
gas form a
stable, solid 1:1 complex (one SO
2
per one pyridine unit), which can be used as a solid
supported SO
2
equivalent.
Strecker reaction
2
is a useful way to synthesize biologically important α-
aminonitriles, which has been extensively studied using a variety of catalysts such as
metal triflates, NiCl
2
, BiCl
3
, ZnX
2
, RuCl
3
, LiClO
4
, etc.
3
The majority of these reactions
involve the use of expensive reagents and catalysts, harsher reaction conditions, longer
reaction time and tedious work up or purification methods to afford the α-aminonitriles in
analytically pure form. Recently, Strecker reaction using solid supported catalysts has
drawn significant attention.
4
Realizing the importance of the Strecker reaction and the
198
role of a solid catalyst in organic synthesis, we decided to explore the potential use of the
solid PVP-SO
2
complex as a mild acidic catalyst in the reaction. Herein, we disclose our
studies on PVP-SO
2
complex and its use as an efficient catalyst in one pot, three
component, Strecker reaction for the synthesis of α-aminonitriles.

5.2 Experimental Methods

All chemicals were purchased from commercial sources and were used as such. PVP-
SO
2
complex was prepared by following the procedure described below.
1
H,
13
C and
19
F
NMR spectra were recorded on Varian NMR spectrometers at 400 MHz and 300 MHz,
respectively.
1
H NMR chemical shifts were determined relative to internal
tetramethylsilane at δ 0.0 ppm.
13
C NMR chemical shifts were determined relative to
internal tetramethylsilane at δ 0.0 ppm or to the
13
C signal of CDCl
3
at δ 77.0 ppm.
19
F
NMR chemical shifts were determined relative to internal CFCl
3
at δ 0.0 ppm.

5.2.1 Preparation of the PVP-SO
2
Complex

In a Nalgene bottle, 2% cross-linked poly(4-vinylpyridine) was taken and the
container was cooled to –78
o
C under dry nitrogen atmosphere. Keeping the temperature
constant, SO
2
gas was passed slowly through the polymer with vigorous shaking of the
mixture. The morphology of the complex changed during the course of the addition and
formed a bright yellow fine powder. The addition of SO
2
was continued till the molar
199
ratio of poly(4-vinyl pyridine to SO
2
reached 1:1. Any excess amount of SO
2
could be
removed by warming the powder to room temperature. The complex was stored in a
freezer (–20
o
C) under dry conditions.

Fig. 5.1 SEM images of PVP (top) and PVP-SO
2
complex (bottom)

200
5.2.2 One Pot Synthesis of α-Aminonitriles

To a solution of the aldehyde (1 mmol) and the amine (1 mmol) in dichloromethane
(4 mL) taken in pressure tube, PVP-SO
2
complex (0.1 g) was added followed by TMSCN
(2 mmol) and sealed. The mixture was stirred under heating at 50
o
C for the specified
amount of time. NMR and TLC were used to monitor the progress of the reaction at
different time intervals. After completion of the reaction, the mixture was filtered and the
residue was washed several times with dichloromethane. The solvent from the collected
filtrate was then removed under reduced pressure to get the crude product. Further
purification was carried out by titration of the products with excess hexanes and removal
of the hexanes under reduced pressure. This process was repeated several times till the
product achieved desired purity. All the products were characterized by analyzing their
spectral data (
1
H,
13
C,
19
F). The residue (PVP) left after washing and filtration was dried
and recycled for further complexation with SO
2
.

5.3 Results and Discussion

PVP is well known to form complexes, which have a wide range of uses such as
universal surface modifier for immobilization of nano particles, in phase change type
liquid crystal device, charge generators in heterolamellar multilayer thin films, in poly(4-
vinyl pyridine) electrolytes etc.
5
Our group has extensively used PVP as an efficient
reservoir and carrier with effective release and retaking ability for acids and acid
catalysts. Solid PVP-(HF)
n
catalysts, developed by Olah and co-workers, have gained use
201
as solid HF equivalents and have been applied as strong, but environmentally safer
catalysts for alkylation and fluorination reactions.
6
Much work was carried out on the
complexation of the amine compounds, such as trimethylamine, triethylamine, aniline
derivatives, etc. with SO
2
gas and the properties and behavior of these solid complexes
were well studied.
7
Our group has also explored the synthetic utility of such complexes
for many organic transformations.
8

Fig 5.2 Preparation of PVP-SO
2
complex

Fig. 5.3 TGA diagram for PVP-SO
2
complex
N
+
-78
o
C
PVP-SO
2
( )
n
N
SO
2
S
O O
Bright yellow solid
( )
n
202
Because cross linked poly(4-vinylpyridine) as mentioned can act as a good solid
Lewis base support for acidic reagents, we prepared polymer bound SO
2
complex with
poly(4-vinylpyridine) as potential mild acid catalyst for organic synthetic
transformations. Recently, Chanda et al. have used PVP-Cu(II) complexes for the
oxidation of aqueous SO
2
.
9
However, to our best knowledge, preparation, properties, and
use of PVP-SO
2
complex have not been reported in the literature. When we passed SO
2

gas through solid 2% cross linked poly(4-vinyl pyridine) at -78
0
C, we observed the
formation of a very bright yellow solid fine powder (Scheme 1). From the weight
increment, this yellow solid was found to be a 1:1 stoichiometric adduct. SEM studies
show that there is no considerable change in the surface morphology of PVP, which
indicates the formation of only a coordinated 1:1 PVP-SO
2
complex (Figure 1a and b).
Similar information was also obtained when we examined the thermal stability of this
complex by TGA. TGA data shows that the complex starts releasing SO
2
at ∼ 50
o
C (fig.
5.3) and the release was completed at ∼130
o
C. Therefore PVP can act as an efficient
recyclable solid support for SO
2
and the complex can be efficiently used as an SO
2

source.
In order to examine the potential of the PVP-SO
2
complex as a catalyst for the
Strecker reaction, we first attempted the reaction of 4-chlorobenzaldehyde with aniline
and TMSCN at room temperature using excess amount of the solid PVP-SO
2
complex
(500 mg for 1mmol of substrates). We observed the formation of a mixture of
corresponding imines and α-aminonitriles, under the reaction conditions. We further
optimized the reaction conditions (temperature and the amount of PVP-SO
2
) and
successfully achieved a clean Strecker reaction (Fig. 5.4) for a variety of aldehydes. The
203
results are summarized in Table 5.1. The amount of PVP-SO
2
used, was reduced from
500 mg to 100 mg to perform the reaction on a 1 mmol scale, but lesser amount of PVP-
SO
2
complex (60 mg) was also effective albeit under longer reaction time. Simple
filtration, removal of the solvent and titration of the product with excess hexanes afforded
the corresponding α-aminonitriles in high yields and purity. No tedious work up or
purification process is necessary. PVP can be recovered and recycled. Therefore PVP can
be considered as a reversible solid support for SO
2
.

R H
O
TMS-CN
+ +
PVP-SO
2
50
o
C, 6h
NH
2
H
N
CN
R
H
CH
2
Cl
2

Fig. 5.4 PVP-SO
2
catalyzed three component α-aminonitrile syntheses from various
aldehydes, aniline and TMSCN

We found that aldehydes, bearing electron withdrawing group as well as electron
donating group on the aromatic ring, react with almost equal efficiency under similar
reaction conditions (Table 5.1, entry 1 to 7). Other aldehydes, such as 2-naphthaldehyde,
9-anthraldehyde and cinnamaldehyde also gave clean Strecker products in good yield and
purity (Table 5.1, entry 8, 9 and 10). In the case of cinnamaldehyde, the attack of
TMSCN to the internal imine intermediate occurred in a 1,2 fashion, thus leaving the
double bond of the cinnamyl group intact in the final product (Table 5.1, entry 10).
However, attempts to perform this reaction with ketones failed. We decided to further
explore this methodology by using different amines with 4-chlorobenzaldehyde and
204
TMSCN using PVP-SO
2
as the catalyst (fig. 5.5). We found that both aromatic and
aliphatic amines provided the corresponding α-aminonitriles in good yields.

Table 5.1 PVP-SO
2
catalyzed α-aminonitrile synthesis from various aldehydes, anilene,
and TMSCN

Fig. 5.5 PVP-SO
2
catalyzed three component α-aminonitrile synthesis from various
amines, 4-chlorobenzaldehyde and TMSCN

205
We observed significant substituent effects on the rate of the reaction for the amines
used. Aliphatic, benzylic and aromatic amines with electron donating groups, react much
faster than the aromatic amines bearing electron withdrawing groups on the aromatic ring
(Table 5.1). For example, the reaction of isopropyl amine was completed within an hour
and afforded the corresponding aminonitrile in 91% yield (Table 5.1, entry 1). Similarly, 4-
methoxyaniline took two hours to provide the corresponding three component product in
98% yield and high purity. On the other hand, halogenated anilines showed low reactivity
towards this three component reaction and took 12 hours for the completion of the reaction
(Table 5.1, entry 7, 8 and 9).
Majority of these reactions were very clean. Progress of the reaction was monitored by
NMR, which shows that the reaction proceeds through the formation of the imine
intermediate followed by the addition of cyanide to provide the final product. Furthermore,
the recovered PVP could be recycled to form the PVP-SO
2
complex thus making this
method more efficient and useful.

206
Table 5.2 PVP-SO
2
catalyzed α-aminonitrile synthesis from various amines

5.3.1 NMR Data of the α–Aminonitrile Compounds

2-Phenyl-2-(phenylamino)acetonitrile (Table 1, entry 1)
1
H NMR: δ 4.12 (brs, 1H), 5.42 (s, 1H), 6.77 (d, J = 7.69 Hz, 2H), 6.90 (t, J = 7.41
Hz, 1H), 7.24-7.29 (m, 2H), 7.40-7.46 (m, 3H), 7.58-7.60 (m, 2H);
13
C NMR: δ 50.1,
114.1, 118.2, 120.2, 127.2, 129.3, 129.50, 129.53, 133.9, 144.6

207

2-(4-Fluorophenyl)-2-(phenylamino)acetonitrile (Table 1, entry 2)
1
H NMR: δ 4.03 (brs, 1H), 5.41 (s, 1H), 6.76 (d, J = 7.69 Hz, 2H), 6.90 (t, J = 7.51
Hz, 1H), 7.11-7.16 (m, 2H), 7.24-7.30 (m, 2H), 7.56-7.60 (m, 2H);
13
C NMR: δ 49.5,
114.2, 116.3 (d,
2
J
C-F
= 22.13 Hz ),118.0, 120.4, 129.1 (d,
3
J
C-F
= 9.16 Hz), 129.6, 129.7
(d,
4
J
C-F
= 3.05 Hz ), 144.4, 163.2 (d,
1
J
C-F
= 249.48 Hz); 19F NMR: δ -112.0 (m).

2-(4-Bromophenyl)-2-(phenylamino)acetonitrile (Table 1, entry 3)
1
H NMR: δ 4.03 (brs, 1H), 5.38 (s, 1H), 6.73 (d, J = 7.69 Hz, 2H), 6.90 (t, J = 7.41
Hz, 1H), 7.23-7.28 (m, 2H), 7. 46 (d, J = 8.0 Hz, 2H), 7. 57 (d, J = 8.61 Hz, 2H);
13
C
NMR: δ 49.7, 114.3, 117.7, 120.6, 123.7, 128.83, 129.6, 132.5, 132.9, 144.3.

2-(4-Methoxyphenyl)-2-(phenylamino)acetonitrile (Table 1, entry 4)
1
H NMR: δ 3.84 (s, 3H), 4.0 (brs, 1H), 5.36 (s, 1H), 6.77 (d, J = 8.61 Hz, 2H), 6.89 (t,
J = 7.42 Hz, 1H), 6.96 (d, J = 8.79 Hz, 2H), 7.25-7.30 (m, 2H), 7.51 (d, J = 8.69 Hz, 2H);
13
C NMR: δ 49.6, 55.4, 114.1, 114.6, 118.4, 120.2, 125.9, 128.6, 129.5, 144.7, 160.4.

2-(2-Methylphenyl)-2-(phenylamino)acetonitrile (Table 1, entry 5)
1
H NMR: δ 2.36 (s, 3H), 3.80 (brs, 1H), 5.45 (s, 1H), 6.77 (d, J = 8.97 Hz, 2H), 6.89
(t, J = 7.41 Hz, 1H), 7.23-7.34 (m, 5H), 7.69 (dd, J = 7.32 Hz, J = 1.83 Hz, 1H);
13
C
NMR: δ 18.6, 48.0, 113.8, 118.3, 120.1, 126.9, 127.5, 129.6, 129.7, 131.3, 132.0, 136.5,
144.8.

208

2-(4-Ethylphenyl)-2-(phenylamino)acetonitrile (Table 1, entry 6)
1
H NMR: δ 1.23 (t, J = 7.69 Hz, 3H), 2.65 (q, J = 7.69 Hz, 2H), 4.08 (brs, 1H), 5.33
(s, 1H), 6.74 (d, J = 7.69 Hz, 2H), 6.86 (t, J = 7.41 Hz, 1H), 7.21-7.25 (m, 4H), 7.46 (d, J
= 8.06 Hz, 2H);
13
C NMR: δ 15.4, 28.4, 49.8, 114.0, 118.3, 120.0, 127.2, 128.7, 129.4,
131.1, 144.7, 145.7.

2-(Phenylamino)-2-(3-trifluoromethylphenyl)acetonitrile (Table 1, entry 7)
1
H NMR: δ 4.06 (d, J = 6.59 Hz, 1H), 5.49 (d, J = 6.23 Hz, 1H), 6.76 (d, J = 7.69 Hz,
2H), 6.91 (t, J = 7.41 Hz, 1H), 7.23-7.29 (m, 2H), 7.58 (t, J = 7.87 Hz, 1H), 7.69 (d, J =
7.87 Hz, 1H), 7.81 (d, J = 7.69 Hz, 1H), 7.86 (s, 1H);
13
C NMR: δ 49.8, 114.3, 117.6,
120.7, 123.6 (q,
1
J
C-F
= 273.14 Hz), 124.1 (q,
3
J
C-F
= 3.81 Hz), 126.3 (q,
3
J
C-F
= 3.81 Hz),
129.6, 129.9, 130.5, 131.3 (q,
2
J
C-F
= 32.8 Hz), 134.9, 144.2;
19
F NMR: δ -63.2.

2-(9-Anthryl)-2-(phenylamino)acetonitrile (Table 1, entry 8)
1
H NMR: δ 4.29 (d, J = 4.21 Hz, 1H), 6.66 (d, J = 4.40 Hz, 1H), 6.93-6.97 (m, 3H),
7.33-7.37 (m, 2H), 7.52-7.56 (m, 2H), 7.61-7.66 (m, 2H), 8.11 (d, J = 8.42 Hz, 2H), 8.43
(d, J = 8.97 Hz, 2H), 8.60 (s, 1H);
13
C NMR: δ 44.2, 113.5, 119.0, 120.0, 123.1, 123.9,
125.5, 127.8, 129.2, 129.69, 129.72, 130.6, 131.5, 145.3.

2-(2-Naphthyl)-2-(phenylamino)acetonitrile (Table 1, entry 9)
1
H NMR: δ 4.10 (d, J = 6.05 Hz, 1H), 5.86 (d, J = 5.86 Hz, 1H), 6.79 (d, J = 8.51 Hz,
2H), 6.90 (t, J = 7.41 Hz, 1H), 7.25-7.30 (m, 2H), 7.53-7.56 (m, 2H), 7.59 (dd, J = 8.60
209
Hz, J = 1.83 Hz, 1H), 7.85-7.87 (m, 2H), 7.90 (d, J = 8.60 Hz, 1H), 8.10 (s, 1H);
13
C
NMR: δ 50.2, 114.1, 118.2, 120.2, 124.3, 126.5, 126.9, 127.1, 127.7, 128.2, 129.3, 129.5,
131.0, 133.0, 133.4, 144.6.

4-Phenyl-2-phenylamino-but-3-enenitrile (Table 1, entry 10)
1
H NMR: δ 3.90 (brs, 1H), 5.05 (d, J = 4.03 Hz, 1H), 6.27 (dd, J = 15.93 Hz, J = 5.13
Hz, 1H), 6.77 (d, J = 8.51 Hz, 2H), 6.90 (t, J = 8.40 Hz, 1H), 7.04 (dd, J = 15.94 Hz, J =
1.47 Hz, 1H), 7.25-7.29 (m, 2H), 7.30-7.39 (m, 3H), 7.41-7.44 (m, 2H);
13
C NMR: δ
47.7, 114.3, 117.7, 120.3, 120.9, 126.9, 128.8, 128.9, 129.6, 134.9, 135.1, 144.4.

2-(4-Chlorophenyl)-2-(isopropylamino)acetonitrile (Table 2, entry 1)
1
H NMR: δ 1.12 (d, J = 6.41 Hz, 6H), 1.45 (brs, 1H), 3.19 (heptet, J = 6.23 Hz, 1H),
4.76 (s, 1H), 7.37 (d, J = 8.43 Hz, 2H), 7.47 (d, J = 8.42 Hz, 2H);
13
C NMR: δ 21.3, 23.5,
47.2, 51.6, 118.7, 128.6, 129.1, 133.9, 134.9.

2-(4-Chlorophenyl)-2-(cyclopentylamino)acetonitrile (Table 2, entry 2)
1
H NMR: δ 1.34-1.43 (m, 2H), 1.55-1.66 (m, 3H), 1.68-1.76 (m, 2H), 1.82-1.92 (m,
2H), 3.42 (quintet, J = 6.23, 1H), 4.68 (s, 1H), 7.35 (d, J = 8.43 Hz, 2H), 7.45 (d, J = 8.42
Hz, 2H);
13
C NMR: δ 23.75, 23.80, 32.1, 33.5, 52.6, 57.7, 118.9, 128.5, 128.9, 133.8,
134.6.

210
2-(4-Chlorophenyl)-2-(benzylamino)acetonitrile (Table 2, entry 3)
1
H NMR: δ 1.89 (brs, 1H), 3.91 (d, J = 13.0 Hz, 1H), 4.01 (d, J = 13.0 Hz, 1H), 4.70
(s, 1H), 7.28-7.38 (m, 7H), 7.46 (d, J = 8.60 Hz, 2H);
13
C NMR: δ 51.1, 52.7, 118.3,
127.7, 128.3, 128.6, 128.9, 129.1, 133.1, 134.9, 137.8.

2-(4-Chlorophenyl)-2-(phenylamino)acetonitrile (Table 2, entry 4)
1
H NMR: δ 4.07 (s, 1H), 5.40 (s, 1H), 6.75 (d, J = 8.51 Hz, 2H), 6.90 (t, J = 7.41 Hz,
1H), 7.24-7.29 (m, 2H), 7.41 (d, J = 8.60 Hz, 2H), 7.52 (d, J = 8.42 Hz, 2H);
13
C NMR: δ
49.5, 114.2, 117.8, 120.5, 128.5, 129.4, 129.60, 132.3, 135.5, 144.3.

2-(4-Chlorophenyl)-2-(4-methylphenylamino)acetonitrile (Table 2, entry 5)
1
H NMR: δ 2.27 (s, 3H), 3.92 (d, J = 6.77 Hz, 1H), 5.37 (d, J = 7.87 Hz, 1H), 6.67 (d,
J = 8.43 Hz, 2H), 7.07 (d, J = 8.42 Hz, 2H), 7.41 (d, J = 8.61 Hz, 2H), 7.52 (d, J = 8.24
Hz, 2H);
13
C NMR: δ 20.5, 50.0, 114.6, 117.9, 128.5, 129.4, 129.98, 130.04, 132.5,
135.4, 142.0.

2-(4-Chlorophenyl)-2-(4-methoxyphenylamino)acetonitrile (Table 2, entry 6)
1
H NMR: δ 3.75 (s, 3H), 3.83 (brs, 1H), 5.31 (s, 1H), 6.73 (d, J = 8.97 Hz, 2H), 6.83
(d, J = 8.97 Hz, 2H), 7.40 (d, J = 8.61 Hz, 2H), 7.52 (d, J = 8.42 Hz, 2H);
13
C NMR: δ
50.9, 55.6, 114.9, 116.5, 118.1, 128.5, 129.4, 132.5, 135.4, 138.1, 154.2.

211
2-(4-Chlorophenyl)-2-(4-fluorophenylamino)acetonitrile (Table 2, entry 7)
1
H NMR: δ 3.98 (d, J = 8.24 Hz, 1H), 5.31 (d, J = 8.24 Hz, 1H), 6.69-6.73 (m, 2H),
6.95-7.0 (m, 2H), 7.42 (d, J = 8.60 Hz, 2H), 7.53 (d, J = 8.79 Hz, 2H);
13
C NMR: δ 50.4,
115.8 (d,
3
J
C-F
= 7.63 Hz), 116.2 (d,
2
J
C-F
= 22.89 Hz), 117.7, 128.5, 129.5, 132.1, 135.6,
140.5 (d,
4
J
C-F
= 2.3 Hz), 157.4 (d,
1
J
C-F
= 239.56 Hz);
19
F NMR: δ -124.0 (m).

2-(4-chlorophenyl)-2-(4-chlorophenylamino)acetonitrile (Table 2, entry 8)
1
H NMR: δ 4.12 (s, 1H), 5.36 (s, 1H), 6.67 (d, J = 8.79 Hz, 2H), 7.21 (d, J = 8.97 Hz,
2H), 7.42 (d, J = 8.61 Hz, 2H), 7.51 (d, J = 8.42 Hz, 2H);
13
C NMR: δ 49.8, 115.6, 117.7,
125.5, 128.7, 129.60, 129.69, 132.0, 135.8, 143.0.

2-(4-Bromophenylamino)-2-(4-chlorophenyl)acetonitrile (Table 2, entry 9)
1
H NMR: δ 4.12 (d, J = 8.60 Hz, 1H), 5.37 (d, J = 8.42 Hz, 1H), 6.63 (d, J = 8.97 Hz,
2H), 7.35 (d, J = 8.97 Hz, 2H), 7.42 (d, J = 8.61 Hz, 2H), 7.51 (d, J = 8.61 Hz, 2H);
13
C
NMR: δ 49.5, 112.6, 115.9, 117.4, 128.5, 129.6, 131.8, 132.4, 135.7, 143.3.

212
5.4 Conclusions

In summary, we have prepared solid PVP-SO
2
complex and explored it as an
efficient, mild and safe catalyst for the one pot three-component Strecker reaction of
aldehydes, amines and TMSCN to synthesize the corresponding α-aminonitriles in
excellent yields and high purity. Readily available starting materials, mild, simple and
clean reaction conditions, minimal work up, simple purification and recycling of PVP are
some of the salient features of this methodology. Widening the scope of this useful solid
catalyst to various other synthetic transformations is currently underway.

213
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Abstract (if available)

Abstract The first four chapters of this work describes the collaborative effort between the University of Southern California (Los Angeles, CA) and the Jet Propulsion Laboratory (Pasadena, CA) focused on developing electrolyte systems to meet the targeted improvements desired by the United States space program. Within this work effort was made to explain the effects of electrolyte modification to the overall performance of the individual electrodes as well as the cell performance on a whole through employment of electrochemical analysis (impedance spectroscopy, Tafel polarization, DC micro-polarization, cyclic voltammetry, and conductivity) and electrical measurements (charge-discharge characteristics, low-operating temperature characterization, high temperature storage performance, and rate capabilities). Chapter 1 provides a brief discussion of the background and development of lithium ion batteries. Further description of the electrolyte systems of such devices is also provided. Developments made to increase the operational temperature of Li-ion batteries using fluorinated esters are described in chapter 2 of this document. Of the examined fluorinated esters, 2, 2, 2-trifluoroethyl butyrate showed the most promised for low-temperature performance. Flame retardant additives were added to electrolyte formulations to improve the safety characteristics of the Li-ion cells, and these results are discussed in chapter 3. Within this chapter, a correlation between flame retardant additive structure and electrochemical stability (and thus, cell performance) is elaborated upon. Chapter 4 is the final chapter discussing Li-ion electrolyte development and describes the work performed to extend the cycle life and high temperature resiliency of cells using advanced lithium salt electrolytes, lithium bis(oxolato)borate and lithium tetrafluoroborate.

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Asset Metadata

Creator Smith, Kiah Anton (author)

Core Title Studies on lithium-ion battery electrolytes and three component Strecker reaction

School College of Letters, Arts and Sciences

Degree Doctor of Philosophy

Degree Program Chemistry

Publication Date 09/11/2009

Defense Date 04/06/2009

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

Tag batteries,electrolytes,flame retardant,flame retardant additives,fluorinated electrolytes,lithium,lithium ion,OAI-PMH Harvest,Strecker reaction

Language English

Contributor Electronically uploaded by the author (provenance)

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

Creator Email kiah.a.smith@gmail.com,kiah.smith@gotoltc.edu

Permanent Link (DOI) https://doi.org/10.25549/usctheses-m2600

Unique identifier UC1313628

Identifier etd-Smith-2887 (filename),usctheses-m40 (legacy collection record id),usctheses-c127-264708 (legacy record id),usctheses-m2600 (legacy record id)

Legacy Identifier etd-Smith-2887.pdf

Dmrecord 264708

Document Type Dissertation

Rights Smith, Kiah Anton

Type texts

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

Repository Name Libraries, University of Southern California

Repository Location Los Angeles, California

Repository Email cisadmin@lib.usc.edu