Copyright by Venkatraman Sivaramakrishnan 2004 · Venkatraman Sivaramakrishnan, Ph.D. The...

Copyright

by

Venkatraman Sivaramakrishnan

2004

The Dissertation Committee for Venkatraman Sivaramakrishnan certifies that this

is the approved version of the following dissertation:

STRUCTURAL AND CHEMICAL CHARACTERIZATIONS OF

DELITHIATED LAYERED OXIDE CATHODES OF LITHIUM-ION

CELLS

Committee:

Arumugam Manthiram, Supervisor

Desiderio Kovar

Harovel G. Wheat

David L. Bourell

Peter F. Green



CELLS

by

Venkatraman Sivaramakrishnan, B. Tech, M.S.E.

Dissertation

Presented to the Faculty of the Graduate School of

The University of Texas at Austin

in Partial Fulfillment

of the Requirements

for the Degree of

Doctor of Philosophy

The University of Texas at Austin

May 2004

Dedication

To Appa, Amma, Raghu, Manni, Meena and Yagnya!

v

Acknowledgements

I express my deep gratitude to my supervisor Dr. Arumugam Manthiram for his

constant support and encouragement during the course of this work. But for his evergreen

enthusiasm in this project, I could not have completed the work. My special thanks to my

committee members Dr. Desiderio Kovar, Dr. Harovel G. Wheat, Dr. David L. Bourell,

and Dr. Peter F. Green for their useful suggestions. I am grateful to all the MS&E area

faculty and staff.

I would also like to thank all the past and present members of Prof. Manthiram

research group, particularly Dr. Ramanan V. Chebiam, Dr. Fernando Prado, Dr. A. M.

Kannan, Dr. Y. Shin, Dr. R. Vadari and Mr. T. A. Arun kumar for their support. I am

indebted to Ramanan and Fernando for being so patient and supportive during my first

year in graduate school.

Financial support by the Welch Foundation Grant F-1254, Center for Space

Power at the Texas A&M University (a NASA Commercial Space Center), and NASA

Glenn Research Center is gratefully acknowledged.

vi



CELLS

Publication No._____________

Venkatraman Sivaramakrishnan, Ph.D.

The University of Texas at Austin, 2004

Supervisor: Arumugam Manthiram

Commercial lithium-ion batteries use the layered LiCoO2 cathode, but only 50 %

of its theoretical capacity could be practically utilized (140 mAh/g). In contrast, 65 and

55 % of the theoretical capacities of the analogous LiNi0.85Co0.15O2 (180 mAh/g) and

LiNi0.5Mn0.5O2 (160 mAh/g) could be practically utilized. However, the reason for the

differences in capacities among the various layered LiMO2 cathodes has not been fully

understood in the literature. With an aim to understand the factors that control the

reversible capacity limits of the layered oxide cathodes, this dissertation focuses on the

structural and chemical characterizations of the Li1-xMO2 (M = Co, Co1-yNiy, Co1-yAly,

Co1-yMgy, and Ni1-yMny, and 0 ≤ (1-x) ≤ 1) phases obtained by chemically extracting

lithium from LiMO2.

vii

Chemical extraction of lithium was accomplished by stirring the layered LiMO2

oxide powders with an acetonitrile solution of the powerful oxidizer NO2BF4 under argon

atmosphere, followed by a careful handling and storage of the products to avoid reaction

with the ambient. Structural characterization of the various cathodes indicate that while

Li1-xCoO2 and the cobalt-rich Li1-xCo1-yMyO2 (M = Ni, Al, and Mg) compositions

generally show the formation of P3 and O1 type phases, the nickel-rich Li1-xNi1-yMyO2

(M = Co and Mn) compositions tend to maintain the initial O3 type structure, but with a

smaller c lattice parameter. Factors such as the nature of Mn+, cation disorder between the

Mn+ and Li+ planes, and the presence of residual Li+ ions in the Li+ plane influence the

structure of the phases formed and the phase relationships. Additionally, chemical lithium

extraction technique has been identified as a convenient and faster method to obtain

qualitative information on the degree of cation disorder in layered oxides.

Oxygen content analysis of the various delithiated Li1-xMO2-δ samples indicate

that the systems tend to lose oxygen from the lattice at deep lithium extraction due to

chemical instability. The chemical instability decreases in the order Li1-xCoO2-δ > Li1-

xNi0.5Mn0.5O2-δ > Li1-xNi0.75Mn0.25O2-δ ˜ Li1-xNi0.85Co0.15O2-δ, which is consistent with the

observed charge voltage profiles. The observed loss of oxygen from the Li1-xMO2-δ is due

to an overlap of the M3+/4+:3d band with the top of the O2-:2p band and a consequent

oxidation of the oxide ions at deep lithium extraction. The oxygen loss from the lattice is

accompanied by a decrease in the c parameter or the formation of new phases. More

importantly, the lithium content at which oxygen loss begins to occur correlates well with

the reversible limit of lithium extraction and the practical capacities, suggesting that the

chemical instabilities may play a crucial role in determining the reversible capacity limits

of lithium-ion battery cathodes.

viii

Table of Contents

List of Tables .........................................................................................................xii

List of Figures .......................................................................................................xiii

CHAPTER 1

INTRODUCTION 1

1.1 Electrochemical power sources..........................................................................1

1.2 Battery ...............................................................................................................2

1.3 Lithium-ion battery............................................................................................3

1.3.1 Anodes ...................................................................................................7

1.3.2 Electrolyte ..............................................................................................9

1.3.3 Separator ..............................................................................................11

1.4 Cathode ............................................................................................................11

1.4.1 Layered structure..................................................................................14

1.4.2 Spinel structure ....................................................................................17

1.4.3 Olivine type LiFePO4 cathodes............................................................20

1.4.4 Other oxide cathodes............................................................................20

1.5 Structural nomenclature of layered oxides.......................................................22

1.6 Chemical delithiation.......................................................................................24

1.7 Objective ..........................................................................................................26

CHAPTER 2

GENERAL EXPERIMENTAL PROCEDURES 28

2.1 Materials synthesis ...........................................................................................28

2.2 Chemical delithiation.......................................................................................29

2.3 Materials characterization................................................................................30

2.3.1 X-ray powder diffraction (XRD) .........................................................30

2.3.2 Rietveld refinement of the X-ray diffraction data................................30

ix

2.3.3 Atomic absorption spectroscopy (AAS) ..............................................31

2.3.4 Redox titration......................................................................................32

2.3.5 Fourier transform infrared spectroscopy (FTIR) .................................33

2.3.6 Thermo gravimetric analysis (TGA)....................................................33

2.3.7 Scanning electron microscopy (SEM) .................................................33

2.4 Electrochemical characterization.....................................................................33

CHAPTER 3

STRUCTURAL AND CHEMICAL CHARACTERIZATIONS OF CHEMICALLY DELITHIATED LAYERED Li1-XCOO2-δ (0 ≤ (1-X) ≤ 1.0) 34

3.1 Introduction......................................................................................................34

3.2 Experimental....................................................................................................36

3.3 Results and discussion .....................................................................................37

3.3.1 Structural analysis ................................................................................37

3.3.2 Oxygen content analysis ......................................................................46

3.3.3 Difference between chemical and electrochemical lithium

extraction.......................................................................................................52

3.3.4 Lithium re- insertion into the P3-type CoO2-δ................................................... 53

3.3.5 Infrared spectra ....................................................................................56

3.3.6 Comparison of non-aqueous and aqueous chemical delithiation.........58

3.3.7 Particle morphology.............................................................................58

3.3.8 Cobalt dissolution ................................................................................59

3.4 Conclusions ......................................................................................................63

x

CHAPTER 4

PHASE RELATIONSHIPS AND FACTORS INFLUENCING THE CHEMICAL LITHIUM EXTRACTION RATE FROM LAYERED LiNi1-yCoyO2 (0 ≤ y ≤ 1.0) CATHODES 64

4.1 Introduction......................................................................................................64

4.2 Experimental....................................................................................................67


4.3.1 LiNi1-yCoyO2 (0 ≤ y ≤ 1.0) system.......................................................68

4.3.2 Structural analysis ................................................................................74

4.3.3 Structural instabilities in Li1-xCoO2 and Li1-xNi0.85Co0.15O2

cathodes.........................................................................................................79

4.3.4 Chemical instabilities in Li1-xCoO2 and Li1-xNi0.85Co0.15O2

cathodes ........................................................................................................................................... 81

3.4 Conclusions ......................................................................................................88

CHAPTER 5

STRUCTURAL AND CHEMICAL CHARACTERIZATIONS OF CHEMICALLY DELITHIATED LAYERED Li1-xNi1-yMnyO2-δ (y = 0.25 and 0.5 and 0 ≤ (1-x) ≤ 1) OXIDES 89

5.1 Introduction......................................................................................................89

5.2 Experimental....................................................................................................92


5.3.1 Crystal chemistry .................................................................................92

5.3.2 Oxygen content analysis ....................................................................103

5.3.3 Chemical lithium extraction rate from LiNi1-yMnyO2........................108

5.3.4 Infrared spectra ..............................................................................................................110

5.3.5 Comparison of non-aqueous and aqueous chemical delithiation.......111

5.4 Conclusions ....................................................................................................119

xi

CHAPTER 6

COMPARISON OF THE PHASE RELATIONSHIPS OF CHEMICALLY DELITHIATED LAYERED Li1-xCo1-yMyO2 (M = Al and Mg) OXIDES 121

6.1 Introduction...................................................................................................121

6.2 Experimental..................................................................................................123

6.3 Results and discussion ...................................................................................123

6.3.1 Crystal chemistry ...............................................................................123

6.3.2 LiCo1-yAlyO2 (y = 0.1 and 0.25) system ............................................127

6.3.3 LiCo1-yMgyO2 (y = 0.06 and 0.1) system...........................................135

6.3.4 Comparison of the phase relationships of various Li1-xMO2

oxides ..............................................................................................................................................141

6.4 Conclusions ....................................................................................................144

CHAPTER 7

SUMMARY 146

References...........................................................................................................151

Vita .....................................................................................................................161

xii

List of Tables

Table 1.1: Capacities and volume changes of various anode materials ...............8

Table 1.2: Room temperature conductivities of various lithium-ion battery

electrolytes ........................................................................................10

Table 1.3: Cathode materials of interest and their intercalation voltage

range..................................................................................................12

Table 3.1: Refinement model of O3-type LiCoO2............................................................ 40

Table 3.2: Refinement model of P3-type CoO2-δ...............................................40

Table 3.3: Refinement model of O1-type CoO2-δ ............................................................. 40

Table 4.1: Time required to extract all the lithium from LiCo1-yNiyO2 and the

structure of the end members............................................................70

Table 4.2: Degree of cation disorder, transition metal slab thickness d(MO2),

lithium inter-slab thickness d(LiO 2), and lithium thermal parameter BLi

obtained from the Rietveld refinement of the X-ray diffraction data for

LiCo1-yNiyO2 .....................................................................................75

Table 5.1: Structural parameters of LiNi0.75Mn0.25O2 ........................................94

Table 5.2: Structural parameters of LiNi0.5Mn0.5O2...........................................94

Table 6.1: Refinement parameters of the parent LiCo1-yAlyO2 (y = 0.1 and 0.25)

and LiCo1-yMgyO2 (y = 0.06 and 0.1) oxides ..................................124

xiii

List of Figures

Figure 1.1: Comparison of the volumetric and gravimetric energy densities of

various rechargeable battery systems..................................................3

Figure 1.2: Production trends of small rechargeable battery systems ...................4

Figure 1.3: Schematic illustration of the charge-discharge process in a lithium-ion

cell.......................................................................................................5

Figure 1.4: Qualitative energy diagram of a transition metal oxide ....................13

Figure 1.5: Crystal structure of an ideal layered LiMO2 oxide ...........................14

Figure 1.6: Crystallographic hexagonal unit cell of an ideal layered LiMO2

oxide..................................................................................................16

Figure 1.7: Crystal structure of the spinel LiMn2O4 cathode ..............................18

Figure 1.8: Relative formation of layered versus spinel structure as a function of

M3+ radius .........................................................................................19

Figure 1.9: Crystal structure of LiFePO4.............................................................21

Figure 1.10: Crystal structures of O3-type LiCoO2, O1-type CoO2, and P3-type

CoO2, viewed along the (100) plane .................................................23

Figure 1.11: (a) Estimated redox potentials of various oxidizing reagents in

acetonitrile medium, and (b) charging voltage profile of LiCoO2

cathode ..............................................................................................25

Figure 3.1: Observed (•) and calculated (−) X-ray diffraction data and the difference

between them for LiCoO2. The positions of the reflections are also

indicated ............................................................................................38

xiv

Figure 3.2: Evolution of the X-ray diffraction patterns of Li1-xCoO2 with lithium

content (1-x). The arrow indicates the appearance a shoulder to the (003)

reflection and the formation of a new phase .....................................39

Figure 3.3: Variations of the a and c lattice parameters with lithium content (1-x) in

chemically delithiated Li1- xCoO2......................................................42

Figure 3.4: X-ray diffraction patterns of the end member CoO2-δ obtained by

reacting LiCoO2 with NO2BF4 for (a) 1 h, (b) 3 h, (c) 2 days, and (d) 7

days. The patterns in (a), (b), and (c) correspond to single P3-type phase

and that in (d) corresponds to a mixture of P3-type and O1-type

phases................................................................................................44

Figure 3.5: Rietveld refinement data of the end member CoO2-δ that was obtained by

reacting LiCoO2 with NO2BF4 for (a) 3 h and (b) 7 days. Circles and

lines correspond, respectively, to the observed and calculated intensities.

The differences between the observed and calculated patterns and the

peak positions corresponding to the P3-type and O1-type phases are also

shown ................................................................................................45

Figure 3.6: Variations of (a) oxidation state of cobalt and (b) oxygen content with

lithium content for the Li1-xCoO2-δ samples obtained after a reaction time

of 1 h (closed symbols) and 2 days (open symbols) .........................47

Figure 3.7: Variations of (a) oxygen content and (b) inter-slab distance, c/3, with

reaction time t for CoO2-δ..................................................................48

Figure 3.8: Qualitative energy diagram of Li1-xCoO2..........................................50

Figure 3.9: Schematic illustration of the transformation of O3-type phase into P3

and O1-type phases ...........................................................................52

xv

Figure 3.10: X-ray diffraction patterns of (a) O3-type LiCoO2, (b) P3-type CoO1.88

(obtained after a reaction time of 1 h), (c) lithium re- inserted

Li0.80CoO1.88 (O3-type), and (d) lithium re- inserted

Li0.41CoO1.72 (O3-type) .....................................................................54

Figure 3.11: Comparison of the sharing of the [LiO 6] octahedra/prism and [MO6]

octahedra in the O3-type and P3-type LiMO2 oxides. While the O3-type

structure involves only edge sharing, the P3-type structure involves both

edge and face sharing........................................................................55

Figure 3.12: FTIR spectra of (a) O3-type LiCoO2, (b) P3-type CoO1.88, and (c)

lithium re-inserted Li0.80CoO1.88 (O3-type).......................................57

Figure 3.13: FTIR spectra of Co(OH)2, HCoO2, and acid delithiated

Li0.3H0.1CoO2.....................................................................................60

Figure 3.14: SEM micrographs of (a) LiCoO2, (b) delithiated Li0.5CoO2, (c)

delithiated CoO2-δ, and (d) lithium re- inserted Li0.8CoO2................61

Figure 3.15: Variations of the amount of cobalt dissolution during chemical

delithiation as a function of lithium content in Li1-xCoO2 ................62

Figure 4.1: Crystal structure of the ideal O3-type layered LiMO2 oxide. The

distances d(MO2) and d(LiO 2) refer, respectively, to the thicknesses of

the transition metal sheet and the inter-slab space for lithium

ions ...................................................................................................66

Figure 4.2: Variations of the unit cell parameters of LiNi1-yCoyO2 (0 ≤ y ≤ 1.0) with

cobalt content y.................................................................................69

Figure 4.3: Variations of the lithium content (1-x) achieved with reaction time for

the nickel-rich Li1-xNi0.7Co0.3O2 system............................................71

xvi


the nickel-rich Li1-xNi0.85Co0.15O2 system.........................................72


the nickel-rich Li1-xNiO2 system.......................................................73

Figure 4.6: Comparison of the SEM micrographs of (a) LiCoO2 at low

magnification, (b) LiCoO2 at high magnification, (c) LiNiO 2 at low

magnification, and (d) LiNiO 2 at high magnification.......................77

Figure 4.7: Evolution of the X-ray diffraction patterns of Li1-xNi0.85Co0.15O2 with

lithium content (1-x). The arrow indicates the appearance a shoulder to

the (003) reflection and the formation of a new phase .....................80

Figure 4.8: Variations of the a and c lattice parameters with lithium content (1-x) in

chemically delithiated (a) Li1- xCoO2 and (b) Li1-xNi0.85Co0.15O2 .....82

Figure 4.9: Variations of the (a) oxidation state of (Co1-yNiy) and (b) oxygen content

with lithium content (1-x) in chemically delithiated Li1-xCoO2-δ (open

symbols) and Li1-xNi0.85Co0.15O2-δ (closed symbols) ........................83

Figure 4.10: Comparison of the first charge profiles of the LiCoO2 and

LiNi0.85Co0.15O2 systems ...................................................................86

Figure 4.11: Qualitative energy diagrams for Li0.5CoO2 and

Li0.5Ni0.85Co0.15O2 .............................................................................87

Figure 5.1: Crystallographic hexagonal unit cell of an ideal layered LiMO2

oxide..................................................................................................91


between them for LiNi0.75Mn0.25O2 and LiNi0.5Mn0.5O2. The positions of

the reflections are also indicated .......................................................95

xvii

Figure 5.3: Evolution of the X-ray diffraction patterns of Li1-xNi0.75Mn0.25O2-δ

for 0 ≤ (1-x) ≤ 1.................................................................................96

Figure 5.4: Evolution of the X-ray diffraction patterns of Li1-xNi0.5Mn0.5O2-δ

for 0 ≤ (1-x) ≤ 1.................................................................................98

Figure 5.5: Variations of the unit cell parameters of Li1-xNi0.75Mn0.25O2-δ with

lithium content (1-x). The closed (•) and open (O) symbols refer,

respectively, to the O3 and O3’ phases...........................................100

Figure 5.6: Variations of the unit cell parameters of Li1-xNi0.5Mn0.5O2-δ with lithium

content (1-x) ....................................................................................102

Figure 5.7: Variations of the (a) average oxidation state of (Ni1-yMny) and (b)

oxygen content with lithium content (1-x) for Li1-xNi1-yMnyO2-δ

(y = 0.25 and 0.5) ............................................................................104

Figure 5.8: Comparison of the first charge profiles of the LiCoO2, LiNi0.85Co0.15O2,

LiNi0.75Mn0.25O2, and LiNi0.5Mn0.5O2 cathodes. The data were collected

at C/100 rate ....................................................................................106

Figure 5.9: Qualitative energy diagrams for Li0.5CoO2, Li0.5Ni0.85Co0.15O2,

Li0.5Ni0.75Mn0.25O2, and Li0.5Ni0.5Mn0.5O2 ......................................107

Figure 5.10: Variations of the lithium content (1-x) achieved with reaction

time for the LiNi1-yMnyO2 system...................................................109

Figure 5.11: FTIR spectra of (a) Li1-xNi0.75Mn0.25O2-δ and

(b) Li1-xNi0.5Mn0.5O2-δ .....................................................................111

Figure 5.12: Comparison of FTIR spectra of Ni(OH)2, Ni0.5Mn0.5O2-δ,

H0.5Li0.3Ni0.5Mn0.5O2 and the parent LiNi0.5Mn0.5O2. Arrow indicates the

position of characteristic absorption band (~3500 cm-1) corresponding to

O-H groups......................................................................................113

xviii

Figure 5.13: FTIR spectra of Li1-xMn2O4 oxides.................................................114

Figure 5.14: Comparison of the TGA plots of Li1-xMn2O4

and LixHyNi0.5Mn0.5O2....................................................................114

Figure 5.15: FTIR spectra of Li2-xMnO3 oxides ..................................................117

Figure 5.16: X-ray diffraction patterns of Li2-xMnO3 oxides. The arrow indicates the

appearance a shoulder to the (003) reflection and the formation of a new

phase for the sample synthesized in the acid medium ....................118


between them for (a) LiCo0.9Al0.1O2, and (b) LiCo0.75Al0.25O2. The

positions of the reflections are also indicated .................................125


between them for (a) LiCo0.94Mg0.06O2, and (b) LiCo0.9Mg0.1O2. The

positions of the reflections are also indicated .................................126

Figure 6.3: Evolution of the X-ray diffraction patterns of Li1-xCo0.9Al0.1O2 with

lithium content (1-x). The arrow indicates the appearance a shoulder to

the (003) reflection and the formation of a new phase ...................129

Figure 6.4: Evolution of the X-ray diffraction patterns of Li1-xCo0.75Al0.25O2 with

lithium content (1-x) .......................................................................130

Figure 6.5: Variations of the unit cell parameters of Li1-xCo1-yAlyO2 with lithium

content (1-x): (a) y = 0.1 and (b) y = 0.25 ......................................131

Figure 6.6: Comparison of the amount of cobalt dissolved during chemical

delithiation between Li1-xCoO2 and Li1-xCo0.75Al0.25O2..................133

Figure 6.7: FTIR spectra of Li1-xCo1-yAlyO2 and Li1-xCo1-yMgyO2 ...................134

xix

Figure 6.8: Evolution of the X-ray diffraction patterns of Li1-xCo0.94Mg0.06O2 with

lithium content (1-x). The arrow indicates the appearance of a shoulder

to the (003) reflection and the formation of a new phase ...............136

Figure 6.9: Variations of the unit cell parameters of Li1-xCo0.94Mg0.06O2 with lithium

content (1-x) ....................................................................................138

Figure 6.10: Variations of the oxidation state of cobalt and oxygen content with

lithium content (1-x) in Li1-xCo0.94Mg0.06O2 ...................................139

Figure 6.11: Evolution of the X-ray diffraction patterns of Li1-xCo0.9Mg0.1O2 with

lithium content (1-x) .......................................................................140

Figure 6.12: Phase relationships of the Li1-xMO2 samples obtained by chemical

delithiation. The hatched region refers to the inaccessibility of the

phases..............................................................................................142

Figure 6.13: Comparison of the sharing of the [LiO 6] octahedra/prism and [MO6]

octahedra in the O3-type and P3-type LiMO2 oxides. While the O3-type

structure involves only edge sharing, the P3-type structure involves both

edge and face sharing......................................................................143

1

CHAPTER 1

Introduction

1.1 ELECTROCHEMICAL POWER SOURCES

In the quest for alternative energy resources to replace the fossil fuels,

electrochemical power sources take an important place as they are environmentally

benign, clean energy technologies compared to fossil fuels. The electrochemical power

sources differ from the conventional power sources such as the power plants since they

convert chemical energy directly into electrical energy without any intermediate step and

hence are not subjected to Carnot cycle limitations.1

The electrochemical power sources in general can be classified into the following

three types: batteries, fuel cells, and super-capacitors.1,2 Both the batteries and super-

capacitors are energy storage devices, in which the active materials are an integral part of

the device. On the other hand, fuel cell is an energy conversion device, in which the fuel

and oxidant (active materials) are supplied to the electrochemical cell from an external

source as and when required. The fundamental driving force for all the electrochemical

reactions is the reaction free energy change ∆G, which is related to the equilibrium cell

voltage E by the following equation:

∆G = − nFE (1-1)

where n and F refer, respectively, to the number of moles of electron exchanged during

the electrochemical reaction and Faraday constant (96,485 C/mol).

It would be desirable if all the available energy ∆G could be converted to useful

work. However, losses due to polarization always occur when a current is passed through

the electrochemical cell. These losses arise due to activation polarization, concentration

2

polarization, and ohmic polarization (IR drop) in the electrodes and electrolyte which

lower the available energy.1

1.2 BATTERY

The major components of a battery are anode, cathode, and electrolyte. During the

normal battery operation (discharge), the anode gets electrochemically oxidized and

releases the electron, which travels via the external circuit (doing useful work) to the

cathode. When the electron reaches the cathode, an electrochemical reduction reaction

takes place. To maintain the overall electrical neutrality during this process, ions

generated at the anode travel to the cathode through the electrolyte and take part in the

electrochemical reaction at the cathode surface.

Batteries can be broadly classified into two types: primary and secondary

batteries. Primary batteries are non-rechargeable. The electrochemical reaction occurring

in a primary battery is either irreversible or not fully reversible. Secondary batteries are

rechargeable. The electrochemical reaction in this case is fully reversible.1 Examples of

primary batteries are Lechlanche, alkaline MnO2, silver oxide, and zinc/air batteries.

Examples of secondary batteries are lead-acid, nickel-cadmium, nickel-metal hydride,

and lithium-ion batteries.1,2 The typical applications of batteries are SLI batteries (starter-

light- ignition) in automobiles and as power sources in portable electric devices. However,

the exponential growth in portable electronic devices such as cell phones, notebook

computers, and camcorders in the last decade has created an ever increasing demand for

light-weight, compact batteries. The world market for battery has reached 50 billion

dollars in 2002 and continues to grow.1

3

1.3 LITHIUM-ION BATTERY

Among the various rechargeable battery systems available, lithium-ion batteries

offer the highest energy density (~ 200 Wh/Kg) and volumetric density (~ 400 Wh/L).3-6

Lithium-ion batteries are the lightest and most compact battery system available and are

preferred over other secondary batteries. Figure 1.1 compares the gravimetric and

volumetric energy density of various secondary batteries. The improvements in the

energy density of the lithium-ion batteries from 1996 to 2003 have resulted mainly due to

the developments in the anode and better packaging of commercial cells. 1

Figure 1.1 Comparison of the volumetric and gravimetric energy densities of various rechargeable battery systems.1

0 100 200 300 400

50

100

150

200

250

2003

1996

Lead-Acid

Nickel-CadmiumNickel Metal Hydride

Lithium-ion

Smaller

Ligh

ter

Gra

vim

etric

Ene

rgy

Den

sity

(W

h/K

g)

Volumetric Energy Density (Wh/L)

4

Figure 1.2 compares the cell production rates of the small secondary batteries,

namely nickel-cadmium (Ni-Cd), nickel-metal hydride (Ni-MH), and lithium-ion

batteries. The lithium-ion battery market has grown from 100 million units in 1996 to 700

million units in 2001 and is projected to grow to 1 billion units in 2005, mostly fueled by

the explosive growth in cell phones and notebook computers.7 Figure 1.2 clearly indicates

that the lithium-ion batteries are taking over the small rechargeable battery market from

Ni-Cd and Ni-MH batteries.

Figure 1.2 Production trends of small rechargeable battery systems.7

Lithium-ion batteries are comprised of cells that employ lithium intercalation

compounds as the positive and negative electrode materials. As the battery is cycled, the

lithium ions shuttle between the electrodes via the electrolyte and for this reason they are

sometimes referred to as rocking-chair batteries. Figure 1.3 shows the schematics of the

electrochemical processes in a lithium-ion cell. The cathode is a layered transition metal

1992 1994 1996 1998 2000 2002

0

400

800

1200

1600

Li-ion

Ni-MH

Ni-Cd

Uni

t cel

ls, m

illio

ns

Year

5

Figure 1.3 Schematic illustration of the charge-discharge process in a lithium-ion cell.

Anode Cathode

Load e- e-

LixC6 Li1-xCoO2 Electrolyte

Charge Li+

Li+ Discharge

charge Cathode: LiCoO2 Li1-xCoO2 + x Li

+ + x e-

discharge charge Anode: C + x Li+ + x e- LixC discharge charge Overall: LiCoO2 + C LixC + Li1-xCoO2 discharge

6

oxide (LiCoO2), and the anode is carbon, typically graphite. When a lithium-ion cell is

charged, the cathode is oxidized and the anode gets reduced. During this process, lithium

ions are deintercalated from the cathode, travel through the electrolyte, and get

intercalated into the graphite anode. Exactly the reverse process occurs on discharge.

Figure 1.3 also indicates the over all cell reaction occurring during the operation of the

lithium-ion cells.

In actual cell fabrication, the active materials are adhered to a metal foil current

collector with a binder, usually polyvinyldiene fluoride (PVDF), and conductive diluents,

typically a high surface area carbon black or graphite. The anode and cathode are

electrically separated by a micro-porous polyethylene or polypropylene film.

The commercial cells have safety features such as shutdown separators, pressure

vents, current interrupt devices (CID), and positive temperature coefficient (PTC)

resistors to protect against abusive conditions from the external short circuit.5 The PTC

changes resistance at a set temperature or current flow and stops a thermal runaway

condition from developing. The CID is incorporated into the cell cap and it interrupts the

electrical connection between the cathode tab and positive terminal when the internal

pressure in the cell reaches a critical value. The CID device is activated from the vapor

pressure of the electrolyte solvent. In addition, electronic power management and safety

circuitry measures temperature, voltage, and current during the cell operation. Lithium-

ion batteries are prone to self-destruction, if the operation is not controlled.

7

1.3.1 Anodes

In the decade before 1990, lithium batteries were pursued with lithium metal as

the anode. The low atomic weight of lithium metal coupled with its high oxidation

potential led to good energy density. However, lithium metal anodes were discarded after

the failure of the Li/MoS2 cells in the operating devices.8 Safety problems arose due to

the reaction of lithium anode with the electrolyte and the subsequent formation of passive

lithium alkyl carbonate on the electrode surface. Lithium deposition occurs non-

epitaxially over this film, leading to lithium dendrite growth, which internally short

circuits the cell. The dendrite growth problem was solved by the use of lithium

intercalating graphite as the anode instead of metallic lithium. The graphite met the major

requirements for an anode, namely fast insertion kinetics and a low redox potential for

lithium intercalation to provide a sufficiently large cell voltage. The carbon anode based

lithium-ion battery was first commercialized by Sony Energytech in 1991.

The low inherent cost of carbon is one of the major reasons for the commercial

success of carbon-based anodes. Also, when graphite electrode is polarized to negative

potential during initial lithium intercalation in an ethylene carbonate (EC) based

electrolyte, the organic compound decomposes to form a stable surface film over the

anode, called solid electrolyte interface (SEI).9 This SEI film effectively passivates the

anode surface and prevents further co- intercalation of solvent, allowing only lithium ion

migration.

However, the major disadvantage of carbon anode is that it suffers from a

limitation in volumetric capacity (theoretical limit of 833 Ah/L) and a larger irreversible

capacity (IRC) loss in the first cycle. Also, a slight increase in anode operating voltages

above that of the currently used carbon based materials is highly desirable in commercial

cells to inhibit lithium metal deposition, especially at high rates.

8

The alternative anodes being explored are the alloys and composites of Al, Si, Ge,

Sn, Pb, Sb, and Bi.10 Table 1.1 lists the various alloy anodes and their electrochemical

properties. As shown in Table 1.1, all these alloys show large volume expansion upon

lithium intercalation that results in a pulverization of the anode and a loss of electrical

contact between grains during prolonged charge-discharge. Further developments in the

chemistry of these alloys are necessary before they can be commercialized.

Table 1.1 Capacities and volume changes of various anode materials.

Anode

Lithiated

phase

Theoretical

specific capacity (mAh/g)

Theoretical volumetric

capacity (Ah/l)

Volume change

(%)

C

LiC6

372

833

12

Al

Li9Al4

2235

6035

238

Si

Li21Si5

4010

9340

297

Sn

Li17Sn4

959

7000

257

Bi

Li3Bi

385

3773

115

9

1.3.2 Electrolyte

The liquid electrolytes used in lithium batteries are solutions of lithium salt in

organic solvents, typically based on carbonates. The electrolytes used in lithium-ion

batteries should satisfy the following criteria:

• Good ionic conductivity (> 10-3 S/cm at -40 to 90 °C) to minimize internal

resistance

• Li+ ion transference number approaching unity to limit concentration

polarization and avoid electronic conduction leading to short circuit

• Wide electrochemical stability window (0 to 5 V)

• Thermal stability up to 70 °C

• Compatibility with other cell components

Table 1.2 lists the various lithium salt/solvent combinations used in lithium-ion

cells and their conductivities.1 Majority of the commercial cells utilize LiPF6 as the salt

as its solutions have high conductivity and good safety properties. However, one of the

major drawbacks with LiPF6 salt is the formation of HF in the electrolyte upon exposure

to traces of moisture.11-13 Other salts such as LiBF4, LiAsF6, and LiClO4 are also used in

lithium batteries. The most recent lithium salt being investigated to replace LiPF6 and

reduce the HF formation problem in the electrolyte is the lithium bisoxalatoborate

(LiBOB).14 A wide variety of solvents including carbonates, ethers, and acetates have

been evaluated for non-aqueous electrolytes. The industry is now focused on the

carbonates as they offer excellent stability, good safety properties, and compatibility with

electrode materials.

10

Table 1.2 Room temperature conductivities of various lithium-ion battery electrolytes.

Electrolyte formulations in current lithium-ion cells typically utilize two to four

solvents. Formulations with multiple solvents provide better cell performance, higher

conductivity, and a broader temperature range of operation. The ethylene carbonate (EC)

solvent provides low irreversible capacity loss and low capacity fade as it forms an SEI

(solid electrolyte interface) film with graphitic anodes and hence it is used in most of the

solvent formulations. However, EC cannot be used alone as a solvent as it is a solid at

room temperature. EC is mixed with other solvents to bring down the viscosity and

freezing point of the overall mixture.1

Lithium

salt

Solvent*

Solvent

volume ratio

Conductivity at

20 °C (mS/cm)

LiPF6

EC/PC EC/DMC EC/DME EC/DEC

1 : 1 1 : 2 1 : 2 1 : 2

6.5 10.0 18.1 7.0

LiAsF6

EC/DME PC/DME

1 : 1 1 : 1

14.5 13.1

LiClO4

EC/DMC EC/DEC EC/DME

1 : 2 1 : 2 1 : 2

8.4 5.2 16.5

*EC: ethylene carbonate, PC: propylene carbonate, DMC: dimethyl carbonate, DEC: diethyl carbonate, DME: dimethoxyethane.

11

1.3.3 Separator

Lithium-on batteries use thin (10 to 30 µm), micro-porous films to electrically

isolate the anode and the cathode. Most of the commercial cells use polyolefin based

separator materials. The separators used in lithium-ion batteries must have high machine

direction strength, should not yield or shrink in width, must be resistant to puncture by

electrode materials, must have effective pore size of < 1 µm, must be easily wetted by the

electrolyte, and must be compatible with other cell components.

1.4 CATHODE

Commercial lithium-ion batteries use the layered LiCoO2 as the cathode, which

intercalates lithium ions reversibly at a high voltage of 4 V versus Li/Li+ and delivers a

practical capacity of 140 mAh/g. The general thermodynamic and kinetic criteria for an

efficient cathode in a lithium-ion battery are given below:

• Ability to intercalate/deintercalate large amount of lithium reversibly

• Lithium intercalation at a high voltage versus metallic lithium

• Limited variation in the voltage profile as a function of Li+ content during

the operation of the cell

• Good structural and chemical stabilities during

intercalation/deintercalation

• Little variation in unit cell volume during intercalation/deintercalation

• Fast enough Li+ and electron diffusion to allow high rate capability in

order to deliver high power density

• Low cost, low or no toxicity, ease of bulk synthesis, and ease of

fabrication

• Formation of a stable interface with the electrolyte

12

Table 1.3 lists the various cathodes being studied in the literature and their

intercalating voltages.6 Layered LiMO2 type cathodes typically deliver capacities in the 4

V region. The doped LiMn2-yMyO4 (y > 0.5) spinels show capacities in the 5 V region as

well.

Table 1.3 Cathode materials of interest and their intercalation voltage range.

The intercalation voltage of the transition metal oxide based cathodes depends on

the position of the Mn+/(n+1)+ couple in the electronic structure. Figure 1.4 shows the

qualitative energy diagram for a typical transition metal oxide.15 During the operation of

a lithium-ion cell for topotactic lithium intercalation, the following three main

contributions to the energy levels are involved:

• an increase/decrease in the number of electrons in the transition metal

electronic bands

Voltage range

vs. Li+/Li (V)

Cathode material

5

LiMn2-yMyO4 (y > 0.5 and M = Cr, Fe, Co, Ni, and Cu)

4

LiCoO2, LiNiO2, LiCo1-y-zNiyMnzO2, LiMn2O4

4-3

LiMnO2, LixMn1-yMyO2 ( M = Li, Ni, and Co)

3.5

LiFePO4

3

LiMn2O4, LixMnO2, LixVyOz

13

• modification of the shape of the bands due to changes in chemical

interactions as a result of lithium intercalation

• a change in the electrostatic energy of the material resulting from the

change in composition

The position of the transition metal:3d band and the changes in its position

relative to the position of the oxygen:2p band upon lithium deintercalation/intercalation

has serious consequences in determining the chemical stabilities of the cathode. For

example, if the M3+/4+ redox couple falls into the top of the O2-:2p band during charging,

then the O2-:2p band could also be depopulated upon deintercalation. This could lead to

the oxidation of O2- ions and the consequent loss of oxygen from the lattice.

Figure 1.4 Qualitative energy diagram of a transition metal oxide.

O 2 - :2p

Density of states N(E)

E

E F M 3+/4+ :3d

14

1.4.1 Layered structure

The most popular cathode hosts investigated for the lithium-ion batteries are the

layered LiMO2 (M= Co, Ni, Mn, V) type oxides. In an ideal layered structure, the Li+

ions and the transition metal M3+ ions occupy the alternate (111) planes of the rock salt

structure.2 This arrangement leads to an ordered strictly two dimensional structure as

shown in Figure 1.5. The MO2 slab-thickness d(MO2), and the lithium interlayer spacing

d(LiO2) are also marked in Figure 1.5.

Figure 1.5 Crystal structure of an ideal layered LiMO2 oxide.

d(MO2)

d(LiO2)

(MO2)n

Li+

15

Figure 1.6 shows the crystallographic unit cell for the layered LiMO2 type oxides.

It is a hexagonal unit cell with R m3 space group. The lithium ion is located in the

crystallographic 3a site, the transition metal M3+ ion is located in the 3b site, and the

oxide ion resides in the 6c site to give a ionic distribution of [Li]3a(M)3b{O2}6c. In this

structure, both the lithium and transition metal ions occupy the octahedral sites. In a ‘non-

ideal’ layered LiMO2 cathode, there will be some transition metal ions present in the

lithium layer due to cation disorder. This can have serious implications on the rate

capabilities of the cathode and the observed phase relationships on delithiation.

LiCoO2

The commercially used LiCoO2 is the most popular cathode in the layered oxide

family because it forms a strictly two-dimensional layered structure shown in Figure 1.6

and is easy to synthesize. However, the major drawbacks of LiCoO2 are toxicity, high

cost, and the limited capacity of 140 mAh/g. The limited capacity arises from the

limitations in the reversible intercalation range for lithium to 0.5 ≤ (1-x) ≤ 1.0 in Li1-

xCoO2. The reasons for the failure of the cathode beyond this intercalation limit has been

attributed to phase transitions, large changes in unit cell volume, and loss of oxygen from

the lattice.16-19

16

Figure 1.6 Crystallographic hexagonal unit cell of an ideal layered LiMO2 oxide.

a

c M Li O

17

LiNiO2

The nickel based LiNiO 2 offers a less expensive alternative to LiCoO2, and it can

deliver a capacity of around 190 mAh/g. However, LiNiO 2 cathode suffers from a few

drawbacks: difficulty in synthesizing in the ideal layered structure and the consequent

presence of nickel ions in the lithium plane, Jahn-Teller distortion associated with the

presence of low-spin Ni3+:d7 ion, irreversible phase changes upon delithiation, and the

loss of oxygen at elevated temperatures.20,21 However, partially cobalt substituted

LiNi0.85Co0.15O2 cathode alleviates some of the problems and delivers a better, stable

reversible capacity of 180 mAh/g. The partial substitution of cobalt has the effect of

stabilizing the layered struc ture with better ordering as well.22

LiCo1-y-zNiyMnzO2

The manganese based LiMnO2 cannot be directly synthesized in the ideal layered

structure by conventional high temperature synthesis. However, the solid solutions of Ni,

Co, and Mn (LiCo1-y-zNiyMnzO2) can be synthesized in the layered structure. Having

multiple 3d metal ions for M in the layered LiMO2 gives one the opportunity to overcome

some of the drawbacks encountered when having a single M3+ ion. Among the LiCo1-y-

zNiyMnzO2 type cathodes, LiMn0.5Ni0.5O2 delivers a capacity of 160 mAh/g and

LiCo1/3Ni1/3Mn1/3O2 delivers a capacity of 170 mAh/g.23,24

1.4.2 Spinel structure

Electrode materials with the spinel type A[B2]X4 structure are attractive because

the spinel structure is a thermodynamically highly stable structure and many compounds

in nature exist with this structure. Figure 1.7 shows the crystal structure of the LiMn2O4

spinel cathode, showing the existing three dimensional diffusion paths for the lithium

18

ions. In this compound, the Li+ and Mn3+/4+ ions occupy the 8a tetrahedral and 16c

octahedral sites respectively, and the oxygen occupies the 32e site to give (Li)8a

[Mn2]16c{O4}32e with the Fd3m space group.25 In the LiMn2O4 cathode, lithium

deintercalation takes place at around 4 V versus Li/Li+, and the cathode maintains the

cubic symmetry during this 4 V process. However, in this structure, an additional lithium

can also be inserted into the LiMn2O4 spinel network by reducing Mn4+ to Mn3+ to give

Li2Mn2O4. This reaction occurs at around 3 V and is accompanied by a cubic to

tetragonal transition. The cubic to tetragonal transition is due the creation of a higher

amount (> 50%) of the Jahn-Teller ion Mn3+:3d4 (t2g3eg1), which leads to a cooperative

distortion of the MnO6 octahedra.

Figure 1.7 Crystal structure of the spinel LiMn2O4 cathode.

Li+

MnO6

19

The LiMn2O4 spinel delivers a capacity of around 120 mAh/g in the 4 V region.

Manganese is relatively inexpensive and environmentally benign when compared to

cobalt and nickel. However, the LiMn2O4 cathode shows slow capacity fade on cycling in

the 4 V region. The major failure mechanisms proposed for this capacity fade are

manganese dissolution, Jahn-Teller effect leading to the formation of tetragonal phase on

the surface of the particles, instability of the manganese oxides with the organic

electrolyte at high voltages, and the microstrain caused by a large lattice parameter

difference between the two cubic phases formed during the charge-discharge process.26-28

The relative stability of the layered versus spinel structure in lithiated metal oxides

depends to some extent on the size of the M3+ ions . Figure 1.8 compares the radius of

the various M3+ ions of the 3d series and the resulting structures.29 The smaller ions such

as Co3+ and Ni3+ tend to adopt the two-dimensional layered LiMO2 structure, while the

larger Ti3+ and Mn3+ ions tend to adopt the three-dimensional spinel AB2O4 structure.

Figure 1.8 Relative formation of layered versus spinel structure as a function of M3+ radius.29

0.52

0.56

0.60

0.64

0.68

0.72

0.76

r M 3+

Ti

r Li+ = 0.74 Å

NiCoFeMnCrV

Spinel Layered

(Å)

20

1.4.3 Olivine type LiFePO4 cathodes

The olivine structure shown in Figure 1.9 has a hexagonal oxygen packing in

which Li+ and Fe2+ ions occupy half of the octahedral sites and P5+ ions occupy 1/8th of

the tetrahedral sites.30 The Li+ ion conductivity mostly proceeds only in the [001]

direction, and not in a three dimensional fashion as in spinels. Additionally, LiFePO4

suffers from poor electronic conductivity. The theoretical capacity of 170 mAh/g, a flat

potential plateau at 3.3 V, inexpensive nature of iron based composition and the high

chemical stability makes this cathode attractive for large scale battery applications such

as in electrical vehicles (EV). However, the rate capability is really poor in this material,

which prevents its application for electric vehicles. The current research focus is on

making LiFePO4 - carbon composites to alleviate this problem. Very recent studies claim

to have drastic improvement in the electronic conductivity of LiFePO4 cathodes by

doping Fe partially with Nb or Zr, but further work is needed to fully verify the claim.31

1.4.4 Other oxide cathodes

Amorphous manganese oxides

Some manganese based amorphous oxides exhibit remarkably high capacities of

300 mAh/g, but delivers this capacity over a wide voltage range of 4.3 to 1.5 V with a

sloping discharge profile.32 The amorphous nature of these materials allow a smooth

accommodation of the Jahn-Teller lattice distortion associated with the Mn3+:3d4 (t2g3eg1)

ions compared to the crystalline counterparts. Also, these amorphous oxides do not

transform to spinel-type phases upon repeated cycling unlike the manganese-based

crystalline layered oxide cathodes.

21

Figure 1.9 Crystal structure of LiFePO4.

Vanadium oxides

A variety of vanadium oxides such as VO2 (B) and V6O13 show high capacity with

extended cyclability.33,34 The V6O13 oxide structure consists of distorted VO6 octahedra

sharing corners and edges. The structure contains tricapped cavities joined through shared

square faces with the open faces of the cavity permitting lithium ion diffusion along

(010). The major disadvantage of these oxides is that they do not contain lithium in the

discharged state, and hence cannot be coupled with the carbon anodes currently used in

FeO6

PO4

Li+

22

commercial lithium-ion cells. These cathodes require the development of lithium

containing anodes in order for them to be successful in lithium-ion batteries.

1.5 STRUCTURAL NOMENCLATURE OF LAYERED OXIDES

Metal oxides with general formula AxMO2 (A = alkali metal ion and M =

transition metal ion) crystallize in layer structures in which the alkali metal ions reside in

between the (MO2)n sheets formed by edge-shared MO6 octahedra. Delmas et al.35

classified such layered compounds according to the coordination environment for the

alkali metal ion (prismatic, tetrahedral or octahedral) and the number of MO2 sheets per

unit cell. Figure 1.10 shows the projection of O3, P3, and O1-type structures.

O3-type AxMO2

In the O3 structure, the An+ ions occupy the octahedral sites with three MO2

sheets per unit cell with an oxygen stacking sequence of …ABCABC… along the c-axis.

LiCoO2, which is used as a cathode material in commercial lithium-ion cells, has this O3-

type structure. In this structure, the AO6 octahedra and the MO6 octahedra share only

edges between them.

P3-type AxMO2

In the P3 structure, the An+ ions occupy the prismatic sites with three MO2 sheets

per unit cell with an oxygen stacking sequence of …ABBCCA… along the c-axis. In this

structure, the AO6 prisms share one face with one MO6 octahedron and three edges with

three MO6 octahedra.

23

Figure 1.10 Crystal structures of O3-type LiCoO2, O1-type CoO2, and P3-type CoO2, viewed along the (100) plane.

24

O1-type AxMO2

In the O1 structure, the An+ ions occupy the octahedral sites, but with only one

MO2 sheet per unit cell. This structure has an oxygen stacking sequence of

…ABABAB… along the c-axis. In this structure also, the AO6 octahedra share both

edges as well as faces with the MO6 octahedra.

Various other structures such as T1, P2, and O2 are also known to exist with

differing oxygen stacking sequences.35 In the case of T1 structure, the An+ ions occupy

the tetrahedral sites.

1.6 CHEMICAL DELITHIATION

The charging of the lithium-ion battery cathodes involves the forced removal of

electrons from the transition metal:3d band and extraction of lithium (deintercalation)

from the LiMO2 cathode. One can also simulate this electrochemical charging reaction

outside the battery environment by oxidizing M3+ to M4+ in LiMO2 by using a chemical

reagent and obtain Li1-xMO2 phases. However, the bulk synthesis of delithiated cathodes

have rarely been pursued, as it is difficult to oxidize M3+ to M4+ with the commonly used

oxidizing agents like I2 or Br2.

However, Wizansky et al.36 showed that powerful oxidizing agents such as

NO2BF4 and NO2PF6 can oxidize M3+ to M4+ in non-aqueous medium enabling one to

synthesize Li1-xMO2 for the whole range of 0 ≤ (1-x) ≤ 1.0. They estimated the redox

potential of different oxidizing couples in acetonitrile medium, and Figure 1.11a

compares the redox potentials of various reagents versus normal hydrogen electrode

(NHE) reference potential. For a comparison with the electrochemical charging process

and potential, Figure 1.11b shows the full charging curve for the Li1-xCoO2 cathode. For

example, the Br2/Br- couple with an oxidizing power of 1.1 V versus NHE can extract

25

only a maximum of ~ 0.5 lithium from Li1-xCoO2. On the other hand, the NO2+/NO2

redox couple with an oxidizing power of 2.1 V versus NHE, can extract all the lithium

from Li1-xCoO2. The major advantage of chemical delithiation is the ability to access bulk

Li1-xMO2 phases without getting contaminated with carbon and binder present in the

actual battery cathodes.

Figure 1.11 (a) Estimated redox potentials of various oxidizing reagents in acetonitrile medium, and (b) charging voltage profile of LiCoO2 cathode.

Li+/Li

+4.4

n BuLi

NO2+/NO2

NO+/ NO

Br2 / Br -

I2 / I -

MoF6/MoF6-

PtF6/PtF6-

+1.7

+ 1.1

+ 0.5

0.0 V vs. NHE

- 2.0

- 3.0

+2.4 +2.1

0.2 0.4 0.6 0.8 1.00.0

0.4

0.8

1.2

1.6

2.0

Vol

tage

(V) v

s. N

HE

Lithium content, (1-x)

Li1-xCoO2

(a) (b)

26

1.7 OBJECTIVE

Commercial lithium-ion cells presently use the layered LiCoO2 as the cathode, but

only 50% of its theoretical capacity (140 mAh/g) could be practically utilized. In

contrast, the layered lithium nickel oxide with a partial substitution of Co for Ni

(LiNi0.85Co0.15O2) shows a much higher reversible capacity of 180 mAh/g, which

corresponds to 65% of its theoretical capacity.22 Also, the layered LiNi0.5Mn0.5O2 and

LiCo1/3Ni1/3Mn1/3O2 have been found to exhibit higher capacities of 160 and 170

mAh/g.23,24 Although the cationic substitutions lead to some improvement in the

reversible capacity, the factors that control the reversible capacity limits of the layered

LiMO2 cathodes are not fully understood in the literature.

Most of the studies in this regard have focused invariably on the structural

characterization of the electrochemically charged cathodes.16,17 Despite the recognition

that the highly oxidized redox couples such as Co3+/4+ and Ni3+/4+ are characterized by a

near-equivalence of the metal:3d and O2-:2p energies particularly in the case of

perovskite oxides, little attention has been paid in the literature to the possible oxidation

of O2- ions during the charge-discharge process and the consequent chemical instability

leading to oxygen loss from the lattice. One of the reasons for the lack of such an

information is the contamination of the electrochemically charged samples by carbon,

binder, electrolyte, and the consequent difficulty in analyzing the oxidation states and

oxygen contents by wet-chemical analysis. However, one can synthesize bulk samples of

Li1-xMO2 free from carbon, binder, and electrolyte by chemically extracting lithium from

LiMO2 with an oxidizer in non-aqueous media and address this issue.18,19

The objective of this dissertation is to understand the factors that govern the

reversible capacity limits of the layered LiMO2 cathodes by systematically investigating

the structural and chemical instabilities of the layered Li1-xMO2 (M = Co, Co1-yNiy, Co1-

27

yAly, Co1-yMgy, Ni1-yMny) cathode materials. Bulk Li1-xMO2 (0 ≤ 1-x ≤ 1.0) samples are

synthesized by non-aqueous chemical delithiation with a powerful oxidizer NO2BF4 and

analyzed for structural and chemical instabilities.

With an introduction to lithium ion batteries in chapter 1, chapter 2 presents the

general details of the experimental procedures. Chapter 3 details the chemical delithiation

behavior of LiCoO2, characterization of Li1-xCoO2 (0 ≤ (1-x) ≤ 1.0) phases formed, and

their chemical and structural instabilities. Chapter 4 investigates the kinetics of chemical

delithiation in the nickel-cobalt solid solution series LiCo1-yNiyO2 (0 ≤ y ≤ 1.0) and their

relationship to cation disorder in the parent cathode. Chapter 5 explores the structural and

chemical instabilities of LiNi1-yMnyO2 (y = 0.25 and 0.5) cathodes. The observed oxygen

loss from the lattice upon chemical delithiation is explained on the basis of an overlap of

the metal: 3d band with the O2-:2p band and is correlated to the charging voltage profile

of the cathodes.

Chapter 6 investigates the influence of aluminum and magnesium substitutions on

the phase relationships of chemically delithiated in Li1-xCo1-yAlyO2 and Li1-xCo1-yMgyO2.

Chapter 6 also summarizes and compares the results of all the Li1-xMO2 (M = Co, Co1-

yNiy, Co1-yAly, Co1-yMgy, Ni1-yMny) cathodes investigated in this dissertation, and the

phase relationships as a function of lithium content for the various Li1-xMO2 are

rationalized on the basis of electrostatic considerations. Summary and recommendations

for future work are given in Chapter 7.

28

CHAPTER 2

General Experimental Procedures

2.1 MATERIALS SYNTHESIS

All the layered LiMO2 (M = Co, Co1-yNiy, Co1-yAly, Co1-yMgy, Ni1-yMny) oxides

used in this study were prepared either by conventional solid-state methods or by co-

precipitation procedures.

LiCoO2 was synthesized by solid-state reaction between Li2CO3 (99.0% purity,

Alfa Aesar) and Co3O4 (reagent grade, GFS chemicals) at 900 °C for 24 h in air. The

reaction mixture consisted of 2 atom % excess lithium to compensate for any

volatilization of lithium that may occur during the high temperature firing. The cobalt-

rich LiCo1-yNiyO2 (y < 0.5) oxides were also synthesized by a similar procedure

incorporating required amounts of green NiO (reagent grade, GFS chemicals), but firing

at 850 °C for 24 h under flowing oxygen. The precursor materials were typically ground

for 1 h with a mortar and pestle and fired in alumina boats in tubular furnaces. The

typical heating and cooling rates of the furnace were 2°C/min and 1°C/min respectively.

The nickel-rich LiCo1-yNiyO2 (y ≥ 0.5), LiNi1-yMnyO2, LiCo1-yAlyO2, and LiCo1-

yMgyO2 samples were synthesized by a co-precipitation procedure. Required amounts of

the metal acetates and/or nitrates (nickel (II) acetate (99+%, Alfa Aesar), manganese (II)

acetate (99+%, Acros organics), cobalt (II) acetate (98+%, Alfa Aesar), aluminum nitrate

(reagent grade, Spectrum Chemicals), and magnesium nitrate (reagent grade, Acros

Organics)) were first dissolved in deionized water. This solution was then added drop by

drop into a 0.1 M KOH (laboratory grade, Fischer scientific) solution to co-precipitate the

29

metal ions as fine hydroxides. The co-precipitate was then filtered, washed with de-

ionized water, and dried overnight at 100 °C in an air oven. The co-precipitate of metal

ions was then ground with a required amount of LiOH.H2O (laboratory grade, Fischer

scientific) for 1 h and fired at a specific temperature and atmosphere mentioned in the

later chapters for respective compounds.

2.2 CHEMICAL DELITHIATION

Chemical extraction of lithium was carried out by stirring the LiMO2 powders in

an acetonitrile solution of NO2BF4 for 2 days under argon atm using a Schlenk line:

LiMO2 + x NO2BF4 → Li1-xMO2 + x NO2 + x LiBF4 (2-1)

Li1-xMO2 compositions with various values of lithium contents (1-x) could be obtained by

controlling the molar ratio of LiMO2:NO2BF4 in the initial reaction mixture. Due to the

high reactivity of NO2BF4 and the possibilities of its decomposition prior to use and side

reactions, the experiments invariably required excess amounts of the oxidizer than that

would be expected based on reaction 2-1 to achieve a specific value of lithium content (1-

x) in Li1-xMO2. The products formed after the reaction were washed three times with

acetonitrile under argon to remove LiBF4 and dried under vacuum at ambient

temperature. After drying, the reaction flasks were opened in an argon-filled glove box.

For the chemical lithium extraction reaction, typically 300 mg of the LiMO2

samples dried overnight was weighed into a clean and dry custom made chemical

extraction flask. The flask was then taken into the glove box filled with dry argon and a

required amount of the oxidizer NO2BF4 (95+%, Aldrich) was weighed and added into

30

the reaction flask inside the glove box. The reaction flask closed tightly under argon was

taken into the Schlenk line, where typically 20 mL of acetonitrile was transferred into the

reaction flask under a positive argon pressure. The reaction mixture was then stirred

continuously for 2 days on a magnetic stirrer before filtering.

2.3 MATERIALS CHARACTERIZATION

All the layered Li1-xMO2 (0 ≤ 1-x ≤ 1.0) samples synthesized were analyzed by

the following characterization techniques.

2.3.1 X-ray powder diffraction (XRD)

Structural characterizations were carried out with X-ray powder diffraction using

Cu Kα radiation. The X-ray diffraction data were collected from 2θ = 10 to 80° with a

counting time of 10 s per 0.02° using a Philips 3550 diffractometer. Samples were

typically prepared by adding a few drops of amyl acetate to the ground powder on a

microscopic glass slide and spread uniformly. The recorded X-ray diffraction patterns

were then compared with the electronic JCPDS files using the JADE software.

2.3.2 Rietveld refinement of the X-ray diffraction data

In order to determine the accurate lattice parameters and to determine any subtle

changes in crystal structures, the X-ray diffraction data of the Li1-xMO2 samples were

analyzed with the Rietveld method using the DBWS-9411 PC program.1

In the Rietveld method, the least-squares refinements were carried out until the

best fit is obtained between the entire observed powder diffraction pattern taken as a

whole and the entire calculated pattern based on the simultaneously refined models for

31

crystal structure, diffraction optics effects, instrumental factors, and other specimen

charactersitics.2 The quantity minimized in the least-squares refinements is the residual,

Sy:

where wi = 1/yi and yi is the observed intensity at ith step, and yci is the calculated

intensity at the ith step.

The typical refinable parameters for each phase are scale factors, atomic

coordinates, specimen-profile parameters, lattice parameters, thermal parameters, peak

widths, and preferred orientation. The global refinable parameters are 2?-zero error,

instrumental profile, profile asymmetry, background, specimen displacement, and

specimen transparency and absorption.

2.3.3 Atomic absorption spectroscopy (AAS)

The lithium contents of the Li1-xMO2 (0 ≤ 1-x ≤ 1.0) samples were determined

with a Perkin-Elmer 1100 atomic absorption spectrometer. The AAS solutions were

typically prepared by dissolving about 40 mg of the Li1-xMO2 sample in concentrated

HCl acid in a closed Teflon vial and heating it in the oven at 100 °C. The dissolved

sample solutions were diluted to the required concentrations with deionized water. A

lithium standard solution of 2 mg/L was prepared with lithium carbonate Li2CO3. AAS

was also used to determine transition metal ion concentration in the filtrate obtained after

chemical extraction. The standard solutions for the various transition metal ions were

prepared using their corresponding oxides.

Sy = ∑ wi (yi - yci )2 (2-2) i

32

2.3.4 Redox titration

Iodometric titrations were performed to determine the average oxidation state of

the transition metal M(2+y)+ ion in Li1-xMO2.3 Typically about 40 mg of the sample was

dissolved in 15 mL of freshly prepared 10 % KI solution in the presence of 10 mL of 3.5

N HCl under constant stirring. The liberated iodine was then titrated with 0.03 N sodium

thiosulfate solution until the color of the mixture changed to golden yellow. Few drops

of 1 % starch solution was added at this point as an indicator, and this turned the color of

the solution into dark blue due to the formation of a molecular complex between I2 and

starch. The end-point for the titration was the change of this dark blue color to a clear

solution. The following equations summarize the overall reactions:

M(2+y)+ + y I- → y/2 I2 + M2+ (2-3)

y/2 I2 + y (S2O3)2- → y I- + y/2 (S4O6)2- (2-4)

The equivalence of thio consumed, y, is calculated as follows:

y = (N1 × V1 × F.wt)/ (W) (2-5)

where, N1 and V1 are, respectively, the normality and volume in mL of the thiosulfate

consumed in the titration. F.wt and W refer, respectively, to the formula weight of the

sample and weight of the sample (in mg) taken for the titration. From the value of y, the

oxidation state of M(2+y)+ could be determined. By combining the oxidation state of the

transition metal ion M(2+y)+ and the lithium content (1-x), the oxygen content value in Li1-

xMO2-δ was calculated using the charge neutrality principle:

(2-δ) = {(2+y) + (1-x)}/2 (2-6)

33

2.3.5 Fourier transform infrared spectroscopy (FTIR)

Fourier transform infrared (FTIR) spectra were recorded with pellets made with

moisture free KBr and the Li1-xMO2 sample using a Nicolet AVATAR 360 FTIR

spectrometer.

2.3.6 Thermo gravimetric analysis (TGA)

A Perkin-Elmer series 7 thermo gravimetric analyzer was used to study the

thermal behavior (change in mass) of the samples.

2.3.7 Scanning electron microscopy (SEM)

The morphology of the Li1-xMO2 samples were determined with a JEOL JSM-

5610 scanning electron microscope.

2.4 ELECTROCHEMICAL CHARACTERIZATION

Electrochemical extraction of lithium was achieved by charging the CR2032 coin

cells assembled with the LiMO2 cathode, lithium anode, and 1 M LiPF6 in ethylene

carbonate (EC) and diethyl carbonate (DEC) electrolyte at C/100 rate. The cathodes were

fabricated by mixing 75 wt% active material powder with 20 wt% acetylene black and 5

wt% of polytetrafluoroethylene (PTFE) binder, rolling into thin sheets of about 0.2 mm

thick, and cutting into circular electrodes of 0.65 cm2 area. The cells were fabricated in an

argon-filled glove box. The C/100 rate means an extremely slow charging of the cells,

requiring 100 hours for one charge. The slow charging rate was used to ensure a situation

close to equilibrium.

34

CHAPTER 3

Structural and Chemical Characterizations of Chemically Delithiated Layered Li1-xCoO2-δ (0 ≤ (1-x) ≤ 1.0)

3.1 INTRODUCTION

The layered LiMO2 oxides exhibit facile lithium intercalation/deintercalation

properties and have become attractive candidates as cathodes for lithium-ion cells. The

lithium extraction properties of such oxides at ambient temperatures provide a convenient

route to access Li1-xMO2 (0 ≤ (1-x) ≤ 1) phases that are otherwise inaccessible by

conventional high temperature synthetic procedures. The Li1-xMO2 oxides consist of

unusually high oxidation states such as Fe3+/4+, Co3+/4+, and Ni3+/4+ and are unstable at

high temperatures, decomposing to lower valent oxides. Such highly oxidized redox

couples are characterized by a near-equivalence of the metal:3d and oxygen:2p energies

and can exhibit interesting electronic properties.1-3

Among the AMO2 type compounds, both LiCoO2 and LiNiO2 are the most widely

studied because of their use as cathodes in lithium-ion batteries. These compounds have

the O3-type structure with an oxygen stacking sequence of …ABCABC… along the c-

axis. The lithium intercalation properties of these compounds have mainly been studied

by electrochemical charge/discharge procedures. The electrochemical lithium extraction

reactions of LiCoO2 have been reported to yield the end member CoO2 either as a single

O1-type phase or as a mixture of two O1-type phases.4-6 The O1-type structure has an

oxygen stacking sequence of …ABABAB… along the c-axis with a single CoO2 sheet

per unit cell.

35

However, the synthesis of bulk CoO2 samples free from carbon and binder used to

fabricate the electrodes for lithium cells has rarely been pursued as it is difficult to

oxidize Co3+ to Co4+ by the commonly used oxidizing agents such as I2 or Br2.7 Several

years ago, Wizansky et al.8 showed that powerful oxidizing agents such as NO2PF6 can

be used to oxidize Co3+ to Co4+. Also, Chebiam et al.9,10 showed recently that bulk

samples of CoO2-δ free from carbon and binder can be synthesized successfully by

chemically extracting lithium from LiCoO2 at ambient temperatures with NO2PF6 in

acetonitrile medium. The CoO2-δ sample obtained by such a chemical lithium extraction

procedure involving the stirring of the LiCoO2 powder with the oxidizer for 2-3 days was

found to consist of a mixture of predominantly a P3-type phase and a small amount of

O1-type phase, which is in contrast to the O1-type phases reported for the

electrochemically prepared CoO2 sample. The P3-type structure has an oxygen stacking

sequence of …ABBCCA… along the c-axis. Both the P3-type and O1-type phases can

form from the parent O3-type LiCoO2 phase by a gliding of the CoO2 sheets during

lithium extraction (Fig. 1.10). Such a gliding of sheets involves very low reaction

energies without the breaking of any Co-O bonds, and therefore, it can occur at room

temperature.11

However, the P3-type structure with a stacking sequence of …ABBCCA…

without any alkali metal ions in between the MO2 sheets would be expected to be

unstable due to the electrostatic repulsion between the negatively charged oxide ions that

lie directly one above the other across the van der Waals gap. In fact, the P3- type

structure has usua lly been observed for predominantly covalent compounds containing

small amounts of large alkali metal ions such as K+.1 The formation of P3-type phase has

not been reported before for MO2 phases without any alkali metal ions in between the

sheets. This chapter presents the synthesis of single phase P3-type CoO2-δ and its

36

characterization by X-ray diffraction, wet-chemical analysis, and infrared spectroscopy.

The investigation focuses on the influence of chemical lithium extraction time on the type

of phases formed and on their chemical and structural stabilities. It also addresses why

the P3-type phase is formed during chemical lithium extraction, but not during

electrochemical lithium extraction.

3.2 EXPERIMENTAL

Synthesis of LiCoO2 and chemical extraction of lithium from LiCoO2 with

NO2BF4 to obtain Li1-xCoO2 for 0 ≤ (1-x) ≤ 1 were carried out as described in section 2.2

in chapter 2. Lithium re- insertion into the deintercalated CoO2-δ was carried out by

stirring the CoO2-δ powders with an acetonitrile solution consisting of excess anhydrous

LiI (99% purity, Alfa Aesar) for 3 days under argon:

Co (3+x)+O2-δ + x LiI → LixCo3+O2-δ + x/2 I2 (3-1)

To avoid reaction with ambient air, the CoO2-δ powder and LiI were taken in the

reaction flask inside the argon-filled glove box and then acetonitrile was added into the

flask in the Schlenk line. After stirring, the product formed was filtered, washed

repeatedly with acetonitrile under argon atmosphere to remove I2, and dried under

vacuum at ambient temperature in the Schlenk line.

The lithium and oxygen contents in the samples were determined, respectively, by

atomic absorption spectroscopy and iodometric titration12 as described in chapter 2.

Structural characterizations were carried out by a Rietveld analysis of the X-ray

diffraction data using the DBWS-9411 PC program.13 The samples were also

37

characterized by FTIR as described in chapter 2. The FTIR data provide qualitative

information regarding the electrical conduction behavior.

3.3 RESULTS AND DISCUSSION

3.3.1 Structural Analysis

Figure 3.1 shows the Rietveld refinement of the X-ray diffraction data of the as-

prepared LiCoO2. The refinements were carried out on the basis of the α–NaFeO2 type

(O3-type) structure with the rhombohedral R m3 (space group: 166) symmetry. Table 3.1

shows the refinement model for R m3 space group. The results presented in Figure 3.1

were obtained by fitting the X-ray diffraction data with a strictly two-dimensional crystal

structure model, [Li]3a(Co)3b{O2}6c, without allowing any cation disorder between the

lithium and transition metal planes. A good matching between the observed and

calculated patterns coupled with low Rwp values and satisfactory goodness of fits confirm

the strictly two-dimensional structure without any detectable cation disorder.

Figure 3.2 shows the X-ray diffraction patterns of the Li1-xCoO2 samples that

were obtained after a reaction time of 2 days with NO2BF4. The patterns in Figure 3.2

were recorded immediately after removing the samples from the glove box. In order to

check the sensitivity of the samples to air, the X-ray patterns were also recorded by

covering the samples with a 6 µm thick Mylar film (before removing from the glove

box). No difference in the X-ray patterns was noticed between the covered and uncovered

samples.

38

Figure 3.1 Observed (•) and calculated (−) X-ray diffraction data and the difference between them for LiCoO2. The positions of the reflections are also indicated.

10 20 30 40 50 60 70 80

LiCoO2

Inte

nsity

(arb

itrar

y un

it)

Cu Kα 2θ (degree)

39

Figure 3.2 Evolution of the X-ray diffraction patterns of Li1-xCoO2 with lithium content (1-x). The arrow indicates the appearance a shoulder to the (003) reflection and the formation of a new phase.

10 20 30 40 50 60 70 80

CoO2

Inte

nsity

(arb

itrar

y un

it)

Cu Kα 2θ (degree)

Li0.35CoO2 Li0.39CoO2

Li0.52

CoO2

Li0.66CoO2

LiCoO2

(00

3)

(012

)

(113

)(1

10)

(018

)

(107

)

(015

) (1

04)

(006

) (1

01)

18 20 22

40

Table 3.1 Refinement model of O3-type LiCoO2 (Space group: R m3 (S.G 166))

Atom

Site

x

y

z

Li

3a

0

0

0

Co 3b 0 0 ½ O 6c 0 0 zox

Table 3.2 Refinement model of P3-type CoO2-δ (R3m space group (S.G 160))

Atom

Site

x

y

z

Co

3a

0

0

0

O (1) 3a 0 0 zox O (2) 3a 0 0 zox'

Table 3.3 Refinement model of O1-type CoO2-δ ( 13mP space group (S.G 164))

Atom

Site

x

y

z

Co

1a

0

0

0

O 2d 1/3 2/3 zox

41

After the 2 days of reaction time, the initial O3-type structure is maintained for

the range 0.5 ≤ (1-x) ≤ 1 in Li1-xCoO2. At around (1-x) = 0.45, another phase is formed as

indicated by the shoulder on the right side of the strong (003) peak. This new phase

grows with further lithium extraction and in the region 0 < (1-x) ≤ 0.45, the two phases

coexist. The X-ray diffraction data in the region 0 < (1-x) ≤ 0.45 could be indexed as a

mixture of O3- and P3-type phases and the data of the end member CoO2-δ could be fitted

on the basis of a single P3-type phase using Rietveld analysis. Moreover, the (003)

reflection of the end member CoO2-δ occurs at a slightly higher 2θ value than that of the

new phase formed at (1-x) = 0.45, which could be due to a small lithium solid solubility

range for the P3-type phase and/or changes in the oxygen content of the P3-type phase

with the overall lithium content (see later). The refinement model for the P3-type phase is

shown in Table 3.2.

Figure 3.3 shows the variations of a and c lattice parameters with lithium content

(1-x) in Li1-xCoO2. The c parameter increases with decreasing lithium content in the

region where the initial O3-type structure is maintained, namely for 0.5 ≤ (1-x) ≤ 1 in Li1-

xCoO2. The P3 phase has lower c parameters than the O3 phase, signaling a stronger O-O

interaction across the van der Waals gap between the CoO2 sheets. Although the overall

trend in the c parameter variation is in general agreement with that reported for the

electrochemically charged samples, the type of phase observed are different for the two

delithiation processes.4-6 The a parameters decrease slightly with decreasing lithium

content in the O3-phase region, and the P3 phase has larger a parameters or smaller c/a

ratio than the O3 phase.

42

Figure 3.3 Variations of the a and c lattice parameters with lithium content (1-x) in chemically delithiated Li1- xCoO2.

Additionally, the formation of the monoclinic phase

Copyright by Venkatraman Sivaramakrishnan 2004 · Venkatraman Sivaramakrishnan, Ph.D. The...

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