January 31, 2017

The Incorporation of Lithium Alloying Metals into Carbon Matrices for Lithium Ion Battery Anodes

by Kevin A. Hays

B.S in Chemistry, May 2009, Salisbury University

A Dissertation submitted to

The Faculty of The Columbian College of Arts and Sciences of The George Washington University in partial fulfillment of the requirements for the degree of Doctor of Philosophy

January 31, 2017

Dissertation directed by

Michael J. Wagner Associate Professor of Chemistry

The Columbian College of Arts and Sciences of The George Washington University certifies that Kevin A. Hays has passed the Final Examination for the degree of Doctor of Philosophy as of September 27, 2016. This is the final and approved form of the dissertation.

The Incorporation of Lithium Alloying Metals into Carbon Matrices for Lithium Ion Battery Anodes

Kevin A. Hays

Dissertation Research Committee:

Michael J. Wagner, Associate Professor of Chemistry, Dissertation Director

Vladislav Sadtchenko, Associate Professor of Chemistry, Committee Member

Claire Besson, Assistant Professor of Chemistry, Committee Member

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Dedication

To my parents William and Marianne Hays. Thanks for getting me this far. To my brother Adam you are an inspiration to all. And to my wife Anna, I am so happy to share this adventure with you.

Acknowledgments

The author wishes to acknowledge the Department of Chemistry at George Washington

University. Everyone in this department has made my time in Washington, DC a pure joy and have expanded my education greatly. My research advisor Professor Michael Wagner has taught me and given me room to teach myself both academically and technically.

I would also like to thank my lab mates, we have worked on many projects together in and out of lab. Thank you to Jon Cox, Nathan Banek, Ming Zhang, Ryan Kopreski, Minji

Li, Sumin Li, Yan Chao, Clifford Cook, and Kwasi Osae-Kwapong.

To everyone at the George Washington Nanofabrication and Imaging Center,

Christine Brantner and especially Anastas Popratilof who has help me forward my career greatly.

I would like to thanks my examination committee Vladislav Sadtchenko, Claire Besson,

Adelina Voutchkova-Kostal, and Michael King, the chair.

Finally to all my friends in the Department of Chemistry, Robin Samuel, Paula Binari, Paula

Cantos, Ashley Mills, Camille Lombard-Banek, Nick Driefel, Jennifer Herdman, Deepanjan

Bhattacharya, and Badri Shyam.

Abstract of Dissertation

The Incorporation of Lithium Alloying Metals in Carbon Matrices for Lithium Ion Battery Anodes

An increased interest in renewable energies and alternative fuels has led to recognition of the necessity of wide scale adoption of the electric vehicle. Automotive manufacturers have striven to produce an electric vehicle that can match the range of their petroleum-fueled counterparts. However, the state-of-the-art lithium ion batteries used to power the current offerings still do not come close to the necessary energy density. The energy and power densities of the lithium ion batteries must be increased significantly if they are going to make electric vehicles a viable option.

The chemistry of the lithium ion battery, based on lithium cobalt oxide cathodes and graphite anodes, is limited by the amount of lithium the cathode can provide and the anode will accept. While these materials have proven themselves in portable electronics over the past two decades, plausible higher energy alternatives do exist. The focus is of this study is on anode materials that could achieve a capacity of more than 3 times greater than that of graphite anodes.

The lithium alloying anode materials investigated and reported herein include tin, arsenic, and gallium arsenide. These metals were synthesized with nanoscale dimensions, improving their electrochemical and mechanical properties. Each exhibits their own benefits and challenges, but all display opportunities for incorporation in lithium ion batteries. Tin is incorporated in multilayer graphene nanoshells by introducing small amounts of metal in the core and, separatedly, on the outside of these spheres.

Electrolyte decomposition on the anode limits cycle life of the tin cores, however, tin

vi oxides introduced outside of the multilayer graphene nanoshells have greatly improved long term battery performance.

Arsenic is a lithium alloying metal that has largely been ignored by the research community to date. One of the first long term battery performance tests of arsenic is reported in this thesis. Anodes were made from nanoscale arsenic particles that were synthesized on multiwall carbon nanotubes by akalide reduction. The performance of these anodes proved sensitive to electrolyte composition, which was significantly improved by using fluorinated ethylene carbonate. Additionally, further gains in capacity retention can be made by limiting the loading voltage to 0.75 V vs lithium metal. The arsenic and multiwall carbon nanotube composite was found to have excellent cycle life and capacity at high mass loading (80% arsenic) when the nanoparticles were directly synthesized on the multiwall carbon nanotubes.

Gallium arsenide is well known for its semiconducting properties, but its performance as in Li-ion battery anodes is first reported here. Gallium is a metal with a low melting point that has been touted as a possible self-healing material for lithium ion anodes. Alone, gallium proves to be unstable as a lithium ion battery anode, but when synthesized as gallium arsenide nanoparticles and mixed with multiwall carbon nanotubes it can charge and discharge in a battery 100 times with approximately twice the capacity of graphite anodes. This first study of gallium arsenide shows dramatic cycle life improvements by using nanoscale rather that micron size gallium arsenide.

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Table of Contents

Dedication…………………………………………………………………….………… iv Acknowledgements………………………………………………………………..…...…v Dedication ...... iv Acknowledgments...... v Abstract of Dissertation ...... vi Table of Contents ...... viii List of Figures ...... xi List of Tables ...... xiv Glossary of Terms ...... xv Foreword ...... xvii 1. Introduction: Advanced Li-ion Batteries, A Solution for Next Generation Energy Storage ...... 1 1.1 The Demand for Energy Storage ...... 1 1.2 Background: The Li-ion Battery ...... 4 1.3 Anodes for Li-ion Batteries ...... 13 1.3.1 Carbon based Anode Materials ...... 14 1.3.2 Lithium Alloying Anode Materials ...... 17 1.3.3 Conversion Anode Materials ...... 25 1.4 References ...... 28 2. Multilayer Graphene Nanoshells Synthesis Modifications and Analysis ...... 33 2.1 Carbon Nanomaterials: A Background ...... 33 2.1.1 Fullerenes ...... 33 2.1.2 Graphene ...... 34 2.1.3 Carbon Nanotubes ...... 36 2.1.4 Carbon Nano Onions...... 37 2.2 Multilayer Graphene Nanoshells: A Description ...... 39 2.3 Multilayer graphene nanoshells: Synthesis ...... 41 2.3.1 Physical Characterization...... 46 2.3.2 Mixing of Salt and Cellulose ...... 46 2.3.3 Effects of Charring ...... 47 2.3.4 Pellet Rotation Speed for Laser Pyrolization ...... 50 2.4 Purification of Ni@MGNS ...... 52 2.4.1 Air Oxidation ...... 53

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2.4.2 Potassium Permanganate Oxidation ...... 56 2.4.3 Steam Oxidation...... 59 2.4.4 Nitric Acid Reflux Oxidation...... 62 2.5 Electrochemical Characterization ...... 66 2.6 Conclusions ...... 66 2.7 References ...... 67 3. Incorporation of Tin into Multiwall Graphene Nanoshells...... 71 3.1 Introduction ...... 71 3.2 Experimental ...... 73 3.2.1 Synthesis – Sn@MGNS ...... 73

3.2.2 Synthesis – SnO2/MGNS ...... 74 3.3 Characterization ...... 75 3.3.1 Mixed Metal MGNS ...... 75

3.3.2 SnO2@MGNS ...... 84 3.3.3 Sn@MGNS ...... 87 3.4 Results and Discussion ...... 90

3.5 Introduction: The SnO2/MGNS composite ...... 100

3.6 Characterization: The SnO2/MGNS composite ...... 100

3.7 Results and Discussion: The SnO2/carbon composite ...... 105

3.8 Conclusions: The SnO2/carbon composite...... 116 3.9 References ...... 117 4. High Performance Arsenic – Multiwall Carbon Nanotube Composite Anodes for Li-ion Batteries ...... 119 4.1 Introduction ...... 119 4.2 Experimental ...... 122 4.3 Physical Characterization...... 123 4.4 Electrochemical Characterization ...... 124 4.5 Results and Discussion ...... 125 4.6 Conclusion ...... 139 4.7 References ...... 142 4.8 Supporting Information ...... 144 5. A Room Temperature Study of Gallium Arsenide as a New Multi Alloying Anode for Li-Ion Batteries ...... 151 5.1 Introduction ...... 151 5.2 Experimental ...... 153

5.3 Results and Discussion ...... 156 5.4 Conclusions ...... 169 5.5 References ...... 170 5.6 Supplementary Material ...... 172 6. Conclusions ...... 176 7. Future Work ...... 178 7.2 References ...... 179

List of Figures

Chapter 1 Figure 1.1 Modified Ragone Chart…………………………………………….…3 Figure 1.2 Schematic of Li-ion battery……………………………………….…..6 Figure 1.3 Voltage profile of Li-ion battery……………………………………...7 Figure 1.4 Cycle life of Li-Graphite half cell..……………………………….…..9 Figure 1.5 Rate performance of LiFePO4………………………………...……..11 Figure 1.6 Voltage profiles of Li half cells and full cell……………………...…13 Figure 1.7 Comparison of hard carbon and graphite half cell voltage profile…..15 Figure 1.8 Capacity of various Li alloying anodes……………………………...18 Figure 1.9 SEM of Sn thin film electrode…………………………………...…...19 Figure 1.10 Voltage profile of Li-Sn half cell………………………..……….....24 Figure 1.11 Voltage profile of Li-FeO half cell………………………………….27

Chapter 2 Figure 2.1 HRTEM of MGNS……………………...………………...…………38 Figure 2.2 Powder XRD of MGNS……………………………………………...39 Figure 2.3 Schematic of laser pyrolization setup………………………………..41 Figure 2.4 TEM of MGNS before and after HNO3 reflux………………………43 Figure 2.5 Powder XRD of Ni@MGNS...... 46 Figure 2.6 Raman spectrograph of MGNS ……………………………………..47 Figure 2.7 Scraping yield of Ni@MGNS based on rotation time…….…………49 Figure 2.8 TEM of MGNS at different rotation times…………………….…….50 Figure 2.9 TGA of Ni@MGNS in air……………………………………….…..52 Figure 2.10 Powder XRD of MGNS purified by air oxidation………………….53 Figure 2.11 TEM of MGNS purified by air oxidation…………………………..54 Figure 2.12 TEM of MGNS purified by KMnO4 oxidation…………………….56 Figure 2.13 Powder XRD of MGNS purified by KMnO4 oxidation……………57 Figure 2.14 Photograph of steam purification setup…………………………….58 Figure 2.15 Powder XRD of MGNS purified by steam…………………………59 Figure 2.16 TEM of MGNS purified by steam………………………………….59 Figure 2.17 TGA of HNO3 oxidized MGNS in air…………………………...... 62 Figure 2.18 HRTEM of HNO3 oxidized MGNS……………………………...... 63

Chapter 3 Figure 3.1 Schematic of Lithiation in core shell nanomaterials………………...70 Figure 3.2 TEM mixed metal MGNS………………………………………...... 75 Figure 3.3 Powder XRD of 9:1 Ni:Sn MGNS…………………………………..76 Figure 3.4 Powder XRD of 8:2 Ni:Sn MGNS…………………………………..77 Figure 3.5 Powder XRD of 7:3 Ni:Sn MGNS…………………………………..78 Figure 3.6 Powder XRD of 6:4 Ni:Sn MGNS…………………………………..79 Figure 3.7 Powder XRD of 5:5 Ni:Sn MGNS…………………………………..80 Figure 3.8 Bar graph of the change in Sn mass through purification…………...81 Figure 3.9 Bar graph of the change ID/IG ratio in the Raman signal of MGNS throughout purification…………………………………..……….…82 xi

Figure 3.10 TEM of SnO2@MGNS …………………………………………….85 Figure 3.11 TEM of Sn@MGNS...... 87 Figure 3.12 HRTEM of Sn@MGNS...... 88 Figure 3.13 TGA of Sn@MGNS in air………………………………………….89 Figure 3.14 Voltage profile of Sn@MGNS...... 93 Figure 3.15 Differential capacity of Sn@MGNS...... 93 Figure 3.16 Cycle life of Sn@MGNS...... 94 Figure 3.17 Coulombic efficiency of Sn@MGNS...... 95 Figure 3.18 Voltage profile of 36.8% Sn@MGNS...... 98 Figure 3.19 Nyquist plot of Sn@MGNS...... 98 Figure 3.20 Powder XRD of SnO2/MGNS …………………………………….102 Figure 3.21 TGA of SnO2/MGNS in Air……………………………………….103 Figure 3.22 TEM of SnO2/MGNS……………………………………………...104 Figure 3.23 SAD of SnO2/MGNS………………………………………………105 Figure 3.24 Powder XRD of Sn/MGNS………………………………………..106 Figure 3.25 Cycle life of Sn/MGNS-Li half cell.....…………………………....107 Figure 3.26 Cycle life of SnO2/MGNS-Li half cell…………………………….108 Figure 3.27 Voltage profile and differential capacity of SnO2/MGNS-Li half cells……………………………………………………………..…110 Figure 3.28 Cycle life of SnO2/MGNS-Li half cell with PAA binder……….…112 Figure 3.29 Voltage profile of SnO2/MGNS-Li half cell with PAA binder……113 Figure 3.30 Coulombic efficiency of SnO2/MGNS half cell with PAA binder..114 Figure 3.31 Rate performance of SnO2-Li half cell………………………….…115 Figure 3.32 Cycle life of SnO2/MGNS half cell in PC electrolyte……………..116

Chapter 4 Figure 4.1 Powder XRD of micron and nano Arsenic……………………….…127 Figure 4.2 TEM of As nanoparticles……………………………………………128 Figure 4.3 Voltage profile of micron As – Li halfcell………………………….129 Figure 4.4 Cycle life of various As/MWCNT composites – Li half cells……...131 Figure 4.5 Coulombic efficiency of As/MWCNT composites – Li half cells….131 Figure 4.6 Voltage profile of Asnano/MWCNT – Li half cell to 0.02V……..…..132 Figure 4.7 Voltage profile of Asnano/MWCNT – Li half cell to 0.75V…………133 Figure 4.8 Nyquist plots of Asnano/MWCNT – Li half cells……………………135 Figure 4.9 SEM of Asnano mixed with MWCNTs………………………………139 Figure 4.10 SEM of Asnano reduced on MWCNTs……………………………..139 Figure 4.11 TEM and EDS of Asnano reduced on MWCNTS…………………..140 Figure 4.12 Voltage profile of synthesized Asnano/MWCNT – Li half cell…….141 Figure 4.13 Differential capacity of As/MWCNT – Li………………………...142 Figure 4.S1 Deconvolution of graphite & MWCNTs powder XRD…………...147 Figure 4.S2 SEM of MWCNTs/graphite……………………………………….148 Figure 4.S3 Voltage profile of MWCNT – Li half cell………………………...148 Figure 4.S4 EDS of Asnano……………………………………………………...149 Figure 4.S5 SEM of Micron As………………………………………………...150 Figure 4.S6 Voltage profile of Asnano/MWCNT – Li half cell in EC:DMC……151 Figure 4.S7 EDS line scan of Asnano/MWCNT ………………………………...152

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Figure 4.S8 Rate performance of Asnano/MWCNT – Li half cell………………153

Chapter 5 Figure 5.1 Voltage profile of GaAsmicron – Li half cell…………………...…….159 Figure 5.2 Cycle life of GaAsmicron – Li half cell……………………………….160 Figure 5.3 Differential capacity of GaAsmicron – Li half cell……………………161 Figure 5.4 Powder XRD of GaAs electrode……………………………………162 Figure 5.5 STEM and HRTEM of GaAs electrode……………………………..163 Figure 5.6 Powder XRD of GaAsnano, GaAsmicron, Asnano, and Ganano……..……164 Figure 5.7 TEM of GaAsnano……………………………………………………165 Figure 5.8 HRTEM of GaAsnano……………………………………………...... 165 Figure 5.9 EDS mapping of GaAsnano…………………………………………..166 Figure 5.10 Powder XRD of GaAsnano electrode……………………………….168 Figure 5.11 Cycle life of GaAsnano/MWCNTs – Li half cells…………………..169 Figure 5.12 Voltage profile of GaAsnano/MWCNTs – Li half cells…………….170 Figure 5.13 Differential capacity of GaAsnano/MWCNTs – Li half cells………171 Figure 5.14 Coulombic efficiency of GaAsnano/MWCNTs – Li half cells……...172 Figure 5.S1 UV/Vis spectra of GaAsnano in MeOH…………………………….175 Figure 5.S2 Deconvolution of graphite & MWCNTs powder XRD…………...176 Figure 5.S3 SEM of MWCNTs/graphite…………………………………….…177 Figure 5.S4 Voltage profile of MWCNT – Li half cell………………………...177 Figure 5.S5 Voltage profile comparison of micron and nano GaAs – Li half cells……………………………………………….178

Chapter 6 None.

Chapter 7 None.

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List of Tables

Chapter 1 None.

Chapter 2 Table 2.1 Comparison of rotation speed for MGNS…………………………....49 Table 2.2 Comparison of HNO3 reflux time on MGNS……………………..….61

Chapter 3 Table 3.1 Capacity of Sn@MGNS based on Ni:Sn ratio………………………..91 Table 3.2 Fitted values of EIS data found in figure 19………………………..…99 Table 3.3 Capacity of SnO2/MGNS based on SnO2 content…………………..108

Chapter 4 Table 4.S1 EIS fit data………………………………………………………….152

Chapter 5 None.

Chapter 6 None.

Chapter 7 None.

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Glossary of Terms

Active material: The electrochemically active portion of the electrode in a battery that provide the capacity

Binder: A polymer material used to adhere electrode particles together and to the current collector.

Coin Cell: A circular metal casing used for batteries, which is crimped shut to seal the system. Often referred to as 2016 or 2032 relating to the thickness of the cell.

Conductive additive: An additive, usually carbon, that is used to increase the electrical conductivity of the electrode.

Coulombic Efficiency: The ratio of the charge removed from the battery relative to the charge applied to the battery.

C Rate: The current required to complete charge/discharge in some fraction of an hour (i.e. 1C = 1 charge per hour, 2C = 1 charge per ½ hour)

Cycle: A complete charge and discharge of a secondary battery

Cycler: A potentiostat used to charge and discharge batteries under programmed regimes.

Full Cell: A Li-ion battery consisting of a Li-ion cathode (e.g. LiCoO2) and Li-ion anode (e.g. graphite)

Half Cell: A Li battery consisting of a Li-ion cathode or anode material with a Li metal anode.

Inactive material: The portions of the electrode that are not part of providing the capacity (e.g. binder and conductive additives)

Load/Unload: Refers to the load and unload of Li in secondary battery materials. Analogous with charge/discharge, but is preferred when used in discussion of Li anode half cells.

Open circuit voltage: Voltage on battery when no current or voltage control is applied

Pouch cell: A plastic, vacuum sealed pouch with exposed tabs as current collectors, used for batteries.

Power density: The power output of the battery relative to the volume of the battery (W l-1). Often ignores case volume unless specified.

Practical capacity: The experimental value measured for the capacity of a battery.

Primary battery: A battery designed to only provide a single discharge

Secondary Battery: A rechargeable battery that can provide multiple charge and discharges

Separator: Li-ion conductive membrane between cathode and anode that is electrically insulating

Solid electrolyte interphase (SEI): A decomposition product of the electrolyte that forms on the electrode surface during the charge and discharge of the Li-ion battery. This product is electrically insulating but ionically conductive.

State-of-Charge: The relative amount of capacity at a given voltage for a battery.

Specific Capacity: The capacity of a battery electrode relative to the mass of the active material (mAh g-1).

Specific energy density: The energy output of a battery relative to the mass of the battery (Wh kg-1). Often ignores casing mass unless specified.

Specific power: The power output of a battery relative to the mass of the battery (W kg-1). Ignores casing mass unless specified.

Theoretical capacity: The amount of capacity based on complete utilization of batteries active materials

Voltage Profile: A plot of the voltage vs capacity of an electrochemical cell, at a constant applied current.

Volumetric capacity: The capacity of a battery electrode relative to the volume of the active material (mAh ml-1).

Volumetric energy density: The energy output of a battery relative to the volume of the battery (Wh ml-1). Often ignores casing volume unless specified

xvi

Foreword

This work is a culmination of four research projects all based around nanomaterial alloy incorporation for lithium alloying anodes ion batteries. Several chapters are arranged in the form of manuscripts, which are to be submitted for publications in the journals mentioned at the beginning of the chapters. Each chapter contains an introduction, experimental, results, and discussion pertinent to the individual project. While the subjects may vary from chapter to chapter they all fall in line with the development of nanomaterials for lithium ion batteries. Chapter 2 is an introduction to multilayer graphene nanoshells. Although electrochemical properties of this nanomaterial has been investigated previously, an in depth report of the synthesis is covered here. Chapter 3 is the inclusion of Sn, a metal that alloys with Li, into the multilayer graphene nanoshells.

The synthesis and electrochemical properties are examined. Chapter 4 details the synthesis and electrochemical properties of As nanoparticles. As, another Li-alloying anode material, is presented as an alternative to P, another high capacity pnictide. This chapter is a manuscript to be submitted to Journal of Power Sources. Nathan Banek and

Michael Wagner are co-authors on this work. Finally, chapter 5 investigates the synthesis and electrochemical properties of GaAs nanoparticles as a new Li-alloying anode material. This chapter is a manuscript to be submitted to Journal of the

Electrochemical Society. Nathan Banek and Michael Wagner are co-authors on this work.

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1. Introduction: Advanced Li-ion Batteries, A Solution for Next Generation Energy Storage

1.1 The Demand for Energy Storage

The Industrial Revolution began in the 18th century. Manufacturing, mining, metallurgy, and infrastructure development all took off in Great Britain and later significant portions of Europe and the New World. This rise in production demanded a vast amount of energy. Coal became the primary fuel for steam engines used in factories, locomotives, and boats to produce and move much of the goods, supplies, and people around the world.[1] Even today coal remains a vital energy source; during 2015, over

200 years later, 800 million tons of coal was consumed in United States alone.[2]

The 19th and 20th centuries brought the rise of the oil industry with new techniques for refinement of petroleum. Petroleum proved to provide very useful forms of fuel for a variety of new means of transportation including automobiles and aircrafts.

While new energy sources have been developed, a large portion of the population still relies heavily on petroleum to fuel their everyday transportation. The US Department of

Energy states that world petroleum consumption has increased from 2.3 billion barrels per year in 1980, to 3.3 billion barrels per year in 2014.[3] New forms of oil production, such as hydraulic fracturing has made it easier to obtain previously difficult to access oil and gas,[4] but as they are only formed on a geological timescale, fossil fuels are considered non-renewable and thus limited resources.[5]

Fossil fuel use has contributed greatly to man-made climate change, resulting in global temperature rise [6] and sea level rise.[7] CO2 and N2O are both byproducts of fossil fuel combustion. The atmospheric concentration of these “greenhouse gases” have

1 increased drastically since the industrial revolution, in large part causing the average global temperature rise. [8] In 2015, President Obama announced the Global Climate

Change Action Plan to battle climate change.[9] A major part of this plan was to increase renewable energy generation 30% by 2030. Renewable energy such as solar (e.g. photovoltaic, solar thermal), wind, and geothermal have the potential to leave a dramatically reduced carbon footprint on this planet.[10] These forms of electricity production can be used to fuel electric vehicles and decrease the dependence on fossil fuels.

In order for large scale electrified transportation to become reality, greatly improved electrical storage is needed. Renewable energy has the capacity to produce large amounts of electricity, but much like a gas tank for petroleum, there is a need for somewhere to store this energy. Mechanical storage is possible (e.g. pumped hydroelectric), but generally require tremendous amounts of land and is not appropriate for transportation. Capacitors are more appropriate for mobile applications, but have a relatively low energy density. Electrochemical cells, in the form of batteries, have the ability to store large amounts of energy in confined spaces and thus are currently the most popular alternative energy storage devices for transportation.

A major interest in the rechargeable battery lays in the creation of long-range electric vehicles. In order for this mode of transportation to be successful, a battery needs to hold a large amount of energy, while being as small and light as possible. These needs can be characterized in terms of volume and weight. The volumetric energy and power densities are measured in units of Wh L-1 and W L-1, respectively. The specific energy and power are measured in units of Wh kg-1 and W kg-1, respectively. Production

2 of high energy, high power batteries is no easy task, and while there is a wide array of battery chemistries, not all are suitable for electric vehicles (Figure 1). On top of having high energy and power, these batteries must be discharged and charged thousands of times in order last the life time of the vehicle. Other important factors include charge rate, operating temperature, and safety. No battery fills all these roles, but the Li-ion battery shows the most promise to be a successful candidate for the electric vehicle.

Figure 1 Modified Ragone chart showing volumetric energy density versus specific energy. Lead acid batteries are heavy and large compared to Li-ion batteries, which can put out significantly more energy for the same sized or weight cell. (Adapted from [11]) As of 2016, Tesla, the leading electric vehicle maker in the USA, claimed an EPA range of 270 miles on a single overnight charge for their premier model, the Model S

P90D.[12] With its 90 kWh battery, this 532 hp vehicle can achieve speeds of 155 mph.

The EPA range of the Model S is well over the mean commute distance in the USA, 18.8 miles,[13] but the EPA tests are performed under very controlled conditions.[14] These

EPA tests do not take into account factors such as passenger weight, temperature, elevation change, and additional interior electronics, all of which can reduce the vehicle range. The 2016 Chevrolet Malibu Premier, a sedan of equivalent size that is fueled by petroleum, has an EPA highway range of 521 miles.[15] While the range of petroleum fueled vehicles will also be negatively affected by the same external factors as battery electric vehicles, a petroleum vehicle can be refueled in a matter of minutes, where a battery electric vehicle takes no less than 75 minutes to get to a fully charged state.[16]

The battery powered electric vehicle is at an inherent disadvantage to petroleum based travel. An infrastructure built around travel distances easily achieved by petroleum fueled vehicles has set the bar for market entry high for electric vehicles. Significant improvements in the capabilities of the batteries that power these electric vehicles must be made in order for them to perform in a fashion that is acceptable to the average person.

Several new battery chemistries are in development, but at this time and for the near future, Li-ion is the most likely to meet the needs for electric vehicles.

1.2 Background: The Li-ion Battery

The Li-ion battery as we know it today is the culmination of over 50 years of research and development from both academic, government and industrial laboratories from all around the world. Its forerunner, the Li battery, was first developed as a primary battery (non-rechargeable) in the 1970’s,[17] but it soon became evident that with the correct electrolyte system, the Li battery could be discharged and charged multiple times.

Comprised of a Li anode and a TiS2 cathode,[18] this design showed promise within the laboratory, but ultimately had very little commercial success. Over multiple charge/discharge cycles, dendritic Li buildup on the surface of the Li metal resulting in

4 internal shorts, which caused the cells to be compromised by thermal runaways leading to fires.[19]

It became evident that Li metal could not safely be used for secondary batteries without a dramatic breakthrough. Li insertion type materials, primarily LiCoO2, were developed as cathodes and as the source of Li+ ions, allowing the removal of Li metal from the battery.[20] In 1990, Sony commercialized a Li-ion battery with a LiCoO2 cathode and petroleum coke anode, to power their handheld camcorders.[21] The last major improvement in this time period was the introduction of the graphite anode to replace coke.[22] This change greatly improved the operating voltage and cycling lifetime of the Li-ion battery.

Small modifications and incremental improvements have been made to Li-ion battery chemistry, but they have largely remained unchanged over the past 20 years. A schematic of the Li-ion battery can be seen in Figure 2. The cathode consists of a LiMO2

(M = Co, Mn, Ni, or some combination of the three) framework on an aluminum current collector. The anode is a synthetic graphite on a copper current collector. These metal current collectors are chosen so they will not be oxidized during the charge/discharge process. The electrolyte utilizes LiPF6 salt in non-aqueous carbonate solvent, usually a combination of ethylene carbonate (EC), dimethyl carbonate (DMC), ethylmethyl carbonate (EMC), propylene carbonate (PC), vinylene carbonate (VC), and/or fluoro ethylene carbonate(FEC), because water is not electrochemically stable over the entire anode and cathode potential ranges. A porous polypropylene separator is placed between the cathode and anode, which allows Li-ion transport while preventing electron transport.

Figure 2 Schematic of a conventional Li-Ion battery. During discharge Li-ions deintercalate from graphite, move through the electrolyte and separator and reinsert in the metal oxide. As a result, electrons move from the anode, through the current collector, through the load, and back into the cathode. (adapted from [23] and [24]) The Li-ion battery is assembled in a discharged state. Upon charging, Li+ migrate from the cathode into the electrolyte and Li+ to intercalate from the electrolyte into the graphite anode. The reverse occurs on discharge. This cell potential is as high as ~ 4.2 V in the fully charged state and ~ 3.0 V when discharged. The variation of the cell potential as the state of charge changes is conventionally plotted as a “voltage profile” (Figure 3).

Generally, the plot is read from left to right, with charge and discharge curved both on the same plot. The discharge and charge curves can be distinguished by the trend in voltage.

Charge curves increase in voltage as capacity increase, while discharge curves decrease in voltage as capacity increases. On occasion the charge or discharge curve may be plotted in reverse, allowing for easy comparison of any hysteresis in the voltage between the two curves. Out of convenience, the voltage is referenced against Li/Li+ instead of the 6 standard hydrogen electrode, which displaces the voltage by -3.05 V vs SHE. Capacity is in terms of mAh g-1 (specific or gravemetric capacity), providing an easy means of defining the capacity of a material per unit mass. Volumetric capacity can also be given

(mAh ml-1), but measuring the volume of an electrode material proves more difficult than weight. Plotting the potential against the fraction of total capacity, as is done in Figure 3, makes it possible to easily compare the charge and/or discharge curves of materials with different capacities.

Figure 3 Voltage profile of LiCoO2 – graphite Li- ion battery.

Ideally, rechargeable batteries should be capable of being charged and discharged many times with little loss in capacity. Each charge/discharge regime is defined as a single cycle. The capacity after each discharge is collected and plotted versus cycle number (Figure 4), allowing one to see the capacity fade of the battery; battery manufacturers generally refer to the number of cycles it takes to fade to 80% of the original capacity as the cycle life. In a perfect system, 100% of the electrons put into the battery would be retrieved each cycle, but there are a variety of side reactions that can take place lowering the overall efficiency. The amount of capacity discharged divided by the charge capacity is defined as the Coulombic efficiency, generally given as a percentage. Electrolyte decomposition is one of the primary causes of less than 100%

Coulombic efficiency and capacity fade. Similar to aqueous electrolyte systems, nonaqueous systems are thermodynamically unstable, particularly susceptible to reduction at

Li-ion anodes which are very near Li/Li+ potential when fully charged. However, nonaqueous solvents used in Li-ion batteries all decompose to form passivation films on the electrode surfaces, called the solid electrolyte interphase (SEI).[25] The SEI is ionically conductive and electronically insulating, slowing or preventing further electrolyte decomposition. The exact composition of the SEI is a topic of active debate, but has both organic and inorganic components.

Figure 4 Cycling performance of TIMREX® SFG44 with 10% PVDF vs Li metal, in 1 M LiPF6 in EC/DMC (1:1 w:w), cycling rate C/7. (Adopted from [37]). A Li-ion battery relies on diffusion of Li+ ions through both liquid electrolyte and solid electrode materials. Solid-state diffusion is notoriously slow, relative to liquid diffusion, so the rate which these batteries can be charged and discharged rely heavily on how quickly the Li+ ions can pass into and out of the electrodes. Additionally, the material must have good electrical conductivity in order to pass electrons at the same rate as the Li+ ions. The rate of charge and discharge can be measured in terms of current per mass, using units of A g-1. Alternatively, the rate can be defined in terms of C-rate, using the current required to complete charge/discharge in some fraction of an hour (i.e. 1C = 1 charge per hour, 2C = 1 charge per ½ hour). Both terms provide much of the same

9 information, but C-rate is easier to translate into the time to charge or discharge, although one must know the cell’s total capacity in order to obtain the current density from it.

Figure 5 shows how C-rate can affect the capacity of a LiFePO4 cathode, a material considered to perform at higher rates than most other Li cathode materials.[26, 27] As the rate increased the material capacity decreases, a trend that is true for all known cathode and anode materials. Additionally, LiFePO4 performance is limited by its electrical conductivity. This is evident by differences in the potentials of the charge and discharge curves (indicated by the arrows in Figure 5). When the rate is increased, an internal resistance within the cell shifts the potentials away from equilibrium. In turn, the material is discharged at lower potentials and charged at higher potentials as the rate increases.

Figure 5 Voltage profiles of LiFePO4 cathode at various rates. The electrodes are comprised of active materials (compounds that readily react with Li) and conductive carbons reinforced with polymer binders and supported by current collectors. The conductive carbons are crucial to ensuring the active electrode material maintains electrical conductivity with the current collector. The binder keeps the active material and conductive carbon adhered together and to the current collector.

Depending on the active material, electrochemical addition of Li can cause expansion of

10 to 300% of the original size,[28] thus the binder plays an important role in the long term cycling of these batteries. These inactive components (binders and conductive carbons) must be added at a sufficient concentration to perform their role, but also minimized in order to optimize the energy density of the electrode.

It quickly becomes clear that these Li-ion batteries are very complex systems. A change to one component of the cell or electrode can play a major role on how the entire battery functions. One experimental approach to find improved materials is to change as few variables at once in order to gain better knowledge of a single process. While commercial batteries cannot risk using Li-metal, due to dendrite formation, laboratory studies will switch out the cathode or anode with Li metal in order to more easily study the electrochemical processes of the individual electrodes (Figure 6), commonly referred to as half-cells. Lithium metal is a readily prepared, highly reproducible standard electrode with excellent performance. As such, it does not limit the performance of the electrode under study and allows direct comparison of performance to other similar studies reported in literature. In addition, the Li metal electrode is generally in large excess and as such does not contribute to capacity fade or significantly to Coulombic efficiency, allowing both to be characterized for the electrode under study throughout its cycle life. One must exercise caution, however, when extrapolating half-cell data to potential full cell performance. For instance, the cathode is the only source of Li for reversible electrochemical reaction in a commercial full-cell. Any Coulombic losses result in decreased cell capacity, an effect that would go unnoticed in a half-cell due to its effectively inexhaustible Li metal electrode. Coulombic efficiency of even 99% could lead to a shortened cycle life of the cell, because 1% of the Li is consumed every cycle

(Li → Li+ + e−).

Capacity % of full charge/discharge

Figure 6 Voltage profiles of a LiCoO2-graphite full-cell (black line) and half-cells with a commercial anode (blue line) and cathode (red line)

1.3 Anodes for Li-ion Batteries

Quality cathodes and anodes are necessary for high-energy Li-ion batteries.

Continuous efforts are made to increase the potential and capacity of cathodes for these batteries. Advances tend to be small due to voltage restrictions based on the oxidation of non-aqueous electrolytes and relatively low capacities due to the need for ternary transition metal materials that allow for charge balance and structural integrity during Li insertion/deinsertion, making them relatively heavy. On the other hand, there are wide arrays of materials that can reversible store Li at appropriate potentials that are suitable for Li-ion battery anodes.

1.3.1 Carbon based Anode Materials Since the introduction of the commercial Li-ion battery by Sony in 1990, carbonaceous materials have filled the primary role as the active anode material to replace the less safe Li metal. Carbon can be used in its natural form as found in the earth as graphite, but it only performs well after considerable processing, so generally, synthetic graphites made from petroleum (most often coke) are used. Heating of these petroleum precursors in inert atmosphere results in high purity carbon with varying degrees of order.[22] When heated, these organic compounds lead to two general categories of carbon; hard or soft. Carbons fall into one category or another by their ability to reorder at high temperatures. Hard carbons have structures with sufficient crosslinking that does not allow sufficient flow and reordering for graphite formation, instead having “pinned” features that leads to the formation of a porous structure when heated, which is often describes as a “house of cards” arrangement of bound graphene sheets. Soft carbons form precursors that can flow and reorder when heated, and align to form graphitic structures, becoming progressively less defective as the heating temperature is increased.

The structures of these carbons play an important role in how Li reacts under electrochemical conditions. Figure 7 compares porous hard carbon vs a graphite synthesized from soft carbon. Hard carbons have a theoretical capacity of up to ~740

-1 mAh g , based on the formation of Li2C6.[29] Li is absorbed on either side of the carbon sheets in a wide array of sites, leading to a sloping potential in the voltage profile. The sloping potential leads to a lower cutoff potential when hard carbons are used as anode active materials in full-cells. Hard carbons were initially used in commercial cells but

14 were ultimately abandoned for graphite anodes because of the flatter voltage profile at lower potentials.

Figure 7 Voltage profiles of carbon – Li half-cells with 1M LiAsF6 in 1:1 EC:PC electrolyte. The top plot show Li loading and unloading into graphite synthesized from soft carbon while the bottom plot is from hard carbon. These voltage profiles show the differences of how the ordering of the carbon plays important roles in the potential and capacity for these materials as anodes. The insets are representations of the ordering in the sp2 carbon layers of each respective structure(Adopted from [22, 76]) Graphite, which has a theoretical capacity roughly half that of hard carbons (372 mAh g-1), has a voltage profile with well-defined steps, referred to as staging, upon

15 loading and unloading of Li. Unlike hard carbon, graphite can only lithiate to LiC6, because the relatively small spacing between the graphite layers and limited gallery space for Li. In highly crystalline graphite, Li+ enters only through the edge plane,[30] loading at potentials very close to that of Li metal.[31] First cycle Coulombic efficiency is improved over hard carbons because electrolyte decomposition (SEI formation) primarily takes place over this edge plane, opposed to hard carbons were SEI forms from all directions. However, using graphite limits the use of a number of inexpensive, low melting electrolytes in which hard carbons cycle well. For instance, early work on Li-ion cells were primarily completed using propylene carbonate (PC) based electrolytes, but was later realized to be cause of graphite exfoliation by co-intercalation of the solvent molecule.[32] Once this was determined, electrolytes were changed to use solvent combination of ethylene carbonate (EC), dimethyl carbonate (DMC), and diethylene carbonate (DEC).

While loading of Li through only the edge plane of graphite prevents excess SEI, it also limits the rate that Li can be loaded and unloaded. Li can be inserted and extracted in between graphite layers only as quickly as the previous Li can diffuse (~10-1 cm2 s-1).[30, 33] The limited diffusivity of Li in graphite leads to certain fundamental drawbacks. First, charge times for Li-ion batteries are extended. These batteries must be charged over multiple hours in order to fully intercalate the graphite. If they are charged too fast Li can plate on the graphite rather than intercalate in the graphite.[34] At low temperatures[35] after multiple cycles, this plating can lead to dendritic Li growth and battery failure.[36] Second, the power of the battery is limited by

16 relatively slow Li de-intercalation.[37] High power applications, such as electric vehicles, require more cells connected in parallel in order to acquire the needed current.

1.3.2 Lithium Alloying Anode Materials

Graphite has been the “go-to” material for Li-ion battery anodes since soon after its introduction over 20 years ago. It provides reliable cycling at a modest capacity, and a low potential near Li metal. That being said, if the capacity of Li-ion batteries is to be greatly increased, anode materials with greater capacity must be developed.

There are a variety of elements that electrochemically alloy with Li. Dey first illustrated this in 1971,[38] naming the phenomenon, “spontaneous electrochemical alloying.” Li was shown to alloy with Sn, Pb, Al, Au, Pt, Zn, Cd, Ag, and Mg in 1M

LiClO4-PC electrolyte. However, it wasn’t until the introduction of the Stalion® Li-ion cell, which used a LiCoO2 cathode and amorphous tin oxide anode, from Fuji Photo Film

Cellttec Co. in 1996, that research in this field gained momentum.[39]

Alloying anode materials are of interest because they can form intermetallics that store much larger quantities of Li than graphitic carbon anodes.[40] Of the elements that alloy with Li, Si has the greatest theoretical gravimetric capacity. Upon full lithiation of

-1 Si, Li22Si5 is formed with a theoretical capacity of 4200 mAh g , although it is thought that only the Li15Si4 phase can be achieved at room temperature.[41] Sn, which has also been extensively studied,[39] lithiates to form Li17Sn4,[42] but only has a theoretical capacity of 960 mAh g-1 because of its higher atomic mass. The theoretical volumetric capacities of these alloys are much closer at 8500 mAh cm-3 and 7400 mAh cm-3 for Si and Sn respectively.[40]

Of course these are not the only Li-alloying elements, there is a broad range of elements that form alloys (Figure 8).[43] Similar to Sn, Pb forms the Li4.25M (M = Sn, or

Pb) alloy,[39] while Ge has the same lithiated stoichiometry as Si at Li3.75M. P, As, Sb, and Bi all form the Li3M alloy,[44] Ga and Ag forms Li2M,[45, 46] and In is slightly less at Li1.5M.[46] Finally, Al and Zn alloy in a 1:1 molar ratio.[47, 48]

4500

4000

3500

3000

2500

2000

1500

1000

500

0 C In Bi Zn Te Pb Sb Ga Sn Al As Ge P Si

Figure 8 Specific (blue bars – mAh g-1 ) and volumetric (red bars – mAh ml-1 ) capacity of various examples of Li-alloying anode elements. (Adopted from [40]) Li-alloying anodes demonstrate enormous capacities, but are not used extensively in commercial batteries. This is because of a sizeable volumetric expansion that must occur in order to accommodate the large amount of Li added during the alloying reaction.

Volume changes upwards of 300% upon full lithiation have been documented for Si[49] and Sn[50]. Such an expansion is dramatic, especially when compared to graphitic carbon, which only experiences a 10% change after lithiation.[51] The change in volume becomes a problem in secondary Li-ion batteries over multiple cycles. The alloying and de-alloying of the material leads to mechanical degradation of the electrode, resulting in

18 loss of electrical and ionic conductivity (Figure 9).[52] Essentially, the electrode falls apart under the force of the material expansion and shrinking.

Figure 9 SEM by Beaulieu et. al., showing formation of cracks in a Sn thin film electrode after Li alloying/dealloying.[28]

1.3.2.1 Restriction of the Size Domain

It was recognized early on that if Li-alloying anode materials had any chance of moving from laboratory experiments to commercialization, the particle size would have to decrease from micron to nanometer scale in order to mitigate volume changes.[46]

Smaller particle sizes reduce the mechanical stresses experienced in the larger morphologies, as the overall dimension changes are still very small.

Smaller particle size alone does not greatly improve the cycle life of Li-alloys.

Particles that are too close together will agglomerate during the alloying process because 19 of the resulting expansion. This agglomeration leads to particles that are larger in size, thus more susceptible to the mechanical stresses, which in turn leads to the pulverization of the electrode.[53] To prevent agglomeration, a number of strategies have been employed involving the synthesis of intermetallics or the mixing of the particles in inactive matrixes (components that do not react with lithium).

1.3.2.2 Intermetallic Alloying Anode Materials

Intermetallic anodes are materials composed of one or more non-alloying metals with a Li-alloying element. Conceptually, these materials allow one component to alloy with

Li while the other component(s) creates a strong backbone to prevent the overall expansion of the particle.[54] An example of this is Cu2Sb.[55] Li inserts itself into the

Cu2Sb lattice at 0.7 volts and displaces 50% of the Cu, yielding Li2CuSb (Equation 1).

Upon further lithiation at 0.61 volts additional Cu is displaced and Li3Sb is formed

(Equation 2).

1. 2Li + Cu2Sb →Li2CuSb + Cu

2. Li + Li2CuSb + Cu → 2Cu + Li3Sb

While this Cu2Sb showed stable cycling performance over 25 cycles, the inactive Cu lattice leads to low specific capacity (250 mAh g-1). The SnSb anode is similar to this concept, but it uses two active materials in the intermetallic.[56] The Sb first reacts with

Li at 800 mV (Equation 3), while the Sn provides a matrix to buffer the expansion of the alloying. When further lithium is loaded into the anode, the roles are reversed; the Li3Sb acts as the matrix as Li17Sn5 is formed (Equation 4).

3. Li + SnSb → Li3Sb + Sn

4. Li + Li3Sb + Sn → Li3Sb + Li17Sn5

SnSb anodes were shown to cycle 200 times with little capacity fade. However, in order to achieve this cycle life the cell capacity was restricted to 360 mAh g-1, roughly the capacity of graphite.

1.3.2.3 Conductive Carbon Matrixes

The intermetallic anode’s fundamental drawback is the mass of the inactive metal in the composite, drastically limiting the possible specific capacity of the material.

Conductive carbons can also be used to form a matrix that prevents agglomeration from particle expansion. These carbons are considerably lighter and act as a good electrically conductive backbone for the Li-alloying nanoparticles. For example, when Si nanoparticles were mixed with multiwall carbon nanotubes in a 1:1 (w:w) ratio, the electrodes were able maintain 64% of the original capacity (2013 mAh g-1) after 50 cycles.[57]

In order for these anodes to be practical there must be a high enough loading of the active alloying material in the conductive matrix. Too much conductive carbon will lead to an electrode with an energy density no greater than the commercial graphite electrode.[58] An amorphous P/ carbon black electrode with a 7:3 P:C mass loading was reported to maintain a capacity of 2170 mAh g-1 of the active P for 100 cycles.[59] The increase mass loading surely is an improvement, but still is far from commercial graphite cells, which have no less than 90% active material.

1.3.2.4 Complex Architectures

The materials reviewed up to this point in this chapter could all be applied by traditional slurry and doctor blade methods. Advances in materials science have led to some unique architectures to mitigate the expansion of Li-alloying materials. Uniform, vertically-aligned pillars of various elemental compositions can be deposited on substrates by sputtering techniques. Theoretically, such pillars should be able to expand without agglomeration of the neighboring pillars, while maintaining electrical contact with the current collector. An example of this is Si-Al-C “nanoscoops,” which are pillars with C bases, Si tops, and an Al component connecting the two sections.[60] The Al was intended to act as an intermediate between the large expansion Si and small expansion C.

The resulting anodes had a capacity not much larger than graphite, 412 mAh g-1. The volumetric and areal (per electrode unit area) energy densities of this material was in fact lower that that of graphite due to the vacant space between the pillars and the limited pillar height.

Novel architectures such as the Si-Al-C nanoscoops show the extremes of what can be designed and chemically engineered based on the challenges presented. These materials are academically interesting, but would prove difficult to translate to large-scale production. Such industrial plants use hoppers and mixers to apply slurries of electrode materials to ribbons of current collector (usually thin Cu foil for anodes). Material that require little change to the traditional production methods, keeping costs low for the manufacturers, have a significantly lower barrier to market entry and thus are more likely to succeed commercially.

1.3.2.5 Tin as a Lithium Alloying Anode Material

Sn and its oxides have been thoroughly explored and the most widely commercialized of the Li-alloying anodes. Matsushita-Panasonic was the first to utilize

Sn in button cells produced in 1979.[39] Fuji Photo Celltec Co. released the Stalion® in

1996, followed most recently by the Nexelion® battery by Sony in 2005.[61]

However, these products are no longer sold, as they fell out of favor, losing their market share to more traditional graphite anode with a more reliable cycle life.

Sn electrochemically lithiates to the Li17Sn5 phase. There are multiple other

LixSny phases before complete lithiation, as seen in Equations 5 - 11.

5. 2 Li + 5 Sn → Li2Sn5

6. 3 Li + Li2Sn5 → 5LiSn

7. 4 Li + 3 LiSn → Li7Sn3

8. Li + 2 Li7Sn3 → 3 Li5Sn2

9. 11 Li + 5 Li5Sn2 → 2 Li13Sn5

10. 9 Li + 2 Li13Sn5 → 5 Li7Sn2

11. 3 Li + Li7Sn2 → Li17Sn5

Each of these phase changes occur at plateaus in the voltage profiles of a Li-Sn half-cell.

Figure 10 shows the loading and unloading of Li into a Sn anode (adapted from [39, 62]).

This is in contrast to the Si and Ge anodes, which undergoes amorphization during lithiation before sudden crystallization to form their respective end products. This

23 difference can be explained by Sn’s low melting point relative to Si and Ge. This allows for greater mobility of the Sn atoms which results in phases that crystalize more easily at room temperature.[43] At high temperature, distinct plateaus are visible in the Si voltage profile.[63]

Figure 10 Voltage profiles of Sn and Si Li-half cells. (Adapted from [5] and [29]).

SnO2 and SnO also electrochemically react with Li in non-aqueous electrolytes. Li storage with these oxides was demonstrated first by Fuji Film,[64] and very soon after confirmed by Dahn et. al.[65] They follow a alloying mechanism similar to that of Sn metal, but first must undergo a conversion reaction, forming Li2O at 0.8V vs Li

(Equations 12 – 13), before Li alloys with Sn.

12. 6.25 Li + SnO → Li2O + 4.25 Li + Sn ↔ Li2O + Li4.25Sn

13. 8.25 Li + SnO2 → Li2O + 4.25 Li + Sn ↔ Li2O + Li4.25Sn

The formation of Li2O from SnO and SnO2 requires the presence of additional Li, resulting in a greatly diminished first cycle columbic efficiency. First cycle efficiencies for these oxides are regularly reported in a range from 30 to 65%.[64-69] The reaction of

Li with SnOx to form Li2O and Sn is reversible,[70] but kinetically unfavorable. Low intensity, broad peaks, at 1.3 V vs Li in cyclic voltammograms of SnO2 show the conversion back from Li2O,[69] but an overpotential up to 3V vs Li is necessary to complete the reaction. Unloading the anode to such a high potential leads to an overall reduced cell potential and increases the rate of capacity fade.[71]

1.3.3 Conversion Anode Materials

As seen with SnO2, it is possible to electrochemically react Li with a metal oxide in order to form Li2O. This reaction is not limited to the oxides of alloying materials. A large array of metal compounds undergo conversion reactions.[72] This reaction can be generalized as:

14. M푎X푏 + (푏 ∙ 푛) Li ↔ 푎 M + 푏 Li푛X where M = metal, X = anion, and n = oxidation state. Conversion reactions for Li-ion batteries have been identified for the transition metals Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Mo,

W, and Ru, with the anions O, S, N, P, and F.[72] The oxides have been the most thoroughly studied because of their ease of synthesis. Of these, MnO2 has the highest theoretical capacity (1233 mAh g-1). A capacity of ~600 mAh g-1, at 0.4V vs Li, was achieved after 100 cycles using a thin film of MnO2 nanowires on a Ni substrate.[73]

Theoretically, nitrides can have an even higher specific capacity due to their lower mass. CrN thin films were reported to have a specific capacity of ~1200

25 mAh g-1.[74] This capacity faded rapidly and could only be obtained when the Li unloading potential was extended to 3.5 V, even though Li was loaded at 0.02 V. This exemplifies one of the largest drawbacks of these conversion materials. LinX frameworks with metal nanoparticles are formed upon reacting Li (Figure 11). Micron scale particles

(e.g. FeO, Figure 11-A) first undergo a certain amount of Li insertion (Figure 11-B) before bonding between the anion and metal is broken to form LinX and the metal nanoparticle (Fe and Li2O, Figure 11-C). As the potential is further decreased electrolyte decomposition takes place (Figure 11-D), contributing to a poor first cycle Coulombic efficiency. The available surface area between the LinX particles formed and the metal nanoparticles causes excessive SEI formation. The LinX compounds formed also have very poor electrical conductivity and because of this, large overpotentials must be applied in order to insert and remove Li from these compounds. The large hysteresis between charge and discharge greatly reduces the roundtrip efficiency (the energy output relative to the energy input) of these materials. Upon unloading Li, the deconversion reaction results in the formation of nanosized particles of the FeO (Figure 11–A’), rather than reforming micron scale particles. This causes a significant increase in the loading potential of Li between the 1st and 2nd cycle, and consequently a decrease in hysteresis.

Figure 11 Generalized 1st load/unload (solid lines) and is the 2nd load (dashed line) voltage profiles for the conversion process of FeO with Li. Various notable points of the reaction are labeled as A) micron FeO particles, B) insertion of Li into FeO resulting in Li-Fe-O, C) conversion of Li-Fe-O to Li2O and Fe, D) formation of SEI on Li2O/Fe matrix, A’) nano FeO after deconversion and unloading of Li, B’) insertion of Li into nano FeO, C’) conversion of Li-Fe-O to Li2O and Fe nano, and D’) formation of SEI on Li2O/Fe matrix (Adopted from [75]). Binary compounds that undergo conversion reactions hold promise as anodes for

Li-ion batteries. That said, these materials are the least researched of the possible anodes available, probably because of their large hysteresis and less favorable voltage plateaus relative to carbons and alloying compounds. Considerable effort in this area of research could lead to gains that may make conversion materials practical alternatives to graphite.

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2. Multilayer Graphene Nanoshells Synthesis Modifications and Analysis

2.1 Carbon Nanomaterials: A Background

This section is intended to provide a brief background discussion of various carbon nanomaterials. Because of the broad array of carbon allotropes, this will not be an all-inclusive review, but instead a primer to help understand multilayer graphene nanoshells.

2.1.1 Fullerenes

2 Fullerenes are a carbon allotrope that consist completely of sp hybridized 5 and 6 member rings in the form of spherical carbon nanoparticles. The most common forms of fullerenes are C60 and C70, each consisting of 12 pentagonal rings and 20 or 25 hexagonal rings respectively. A C60 molecule is approximately 1 nm in diameter.[1]

Richard Smalley and Harold Kroto first discovered C60, also known as

Buckminsterfullerene, in 1985.[2] Laser evaporation of graphite created a plume of molecular species, and using mass spectrometry they were able to identify C60 as the primary product. Five years later, macroscopic amounts of C60 and C70 were synthesized by resistive heating of graphite in inert atmosphere.[3] The product labeled “fullerene soot,” was further purified by chromatographic methods to obtain the fullerenes.[4]

The physical properties and chemistry of fullerenes have been well characterized.

C60 gradually oxidizes at 200C and decomposes at 400C.[5] The solubility of C60 can be varied widely by functionalization it with various groups, and is generally found to be

33 increased by fluorination.[6] A wide variety of reactions including, but not limited to, halogenation, organometallic and polymerization have been studied.[7]

In addition to these reactions, fullerenes have been investigated as anode materials for Li-ion batteries. Li12C60 was shown to form using cyclic voltammetry, but due to a lack of large scale production of pure fullerene at reasonable prices and showing little or no advantage over graphite, further testing has been limited.[8] While showing poor reversible capacity as an active material for Li-ion batteries,[9] fullerene soot showed improvements as an additive to graphite electrodes. Addition of 10% mass fullerene soot to graphite electrodes extended cycle life in half cell configuration versus Li.[10]

2.1.2 Graphene

Since the realization that it’s potential for breakthrough applications in 2004,[11] graphene has been a hot topic in the world of chemistry and material science.[12-16]

Because of the electronic,[17] thermal,[18] and mechanical[19] properties graphene displays, it has been explored for use in energy storage,[12, 20] electronic devices,[21] and catalysis.[22, 23] It has even been manufactured into carbon fiber bicycle race wheels with claims of increased stiffness, reduced weight, and reduced temperature build-up on the brake track.[24]

Graphene was first isolated using the Scotch tape method.[11] Very small quantities of graphene were obtained by pealing back Scotch tape multiple times of a piece of pyrolytic graphite, for which Novoselov and Geim later received the Nobel Prize in physics. In practice, this method and most other methods of production actually produce a mix of graphene and few layer graphene, which are essentially graphite crystals that are only a few layers thick.[25] Fundamentally, the structure of graphene is

34 a two-dimensional, single layer of sp2 hybridized carbon atoms., but both single and few layer graphene are commonly called “graphene”, although only for the former is name rigorously true.

The most popular synthesis for graphene is mechanical exfoliation of graphite oxide by sonication.[26] Hummers’ method can be used to oxidize graphite to graphite oxide.[27] Graphite oxide disperses well in H2O because of its polar groups. Sonication of the dispersion overcomes the van der Waals interactions holding the graphite oxide sheets together, resulting in graphene oxide sheets.[28] Graphene oxide has low electrical conductivity compared to graphene so further modification is necessary.

Chemical reduction using hydrazine hydrate reduces the graphene oxide and improves the electrical conductivity of the resulting product to 2420 S m-1.[26] While reduced graphene oxide has improved conductivity, it is still lower than what is expected for graphene, because of defects in the structure. To produce graphene with fewer defects, graphite can be directly exfoliated by sonication in N-methyl-2-pyrrolidone (NMP).[29]

Conductivity of this product is near 6500 S m-1, but yield is only 1 wt%.

Graphene has been proposed as a “high performance” energy storage material because of its superb electrical conductivity and relative low production cost.[12] Its use for electrochemical storage in Li-ion batteries[30] and supercapacitors,[31] as well as chemical storage of molecular hydrogen have been explored.[32] As an anode material for Li-ion batteries, graphene has a theoretical capacity of 744 mAh g-1, under the assumption it forms LiC3. This capacity is twice that of graphite and is based on the assumption that Li can be adsorbed on both sides of the graphene sheet. An experimental

-1 reversible capacity of 848 mAh g was reported, higher than what is expected in the LiC3

35 model.[33] Such high capacity could be due to Li storage in defects of the graphene.

Additionally, these anodes were unloaded up to 3.5V vs Li. Unlike graphite, which intercalates and deintercalates at one potential, graphene’s adsorption of Li occurs over a large range of potentials,[34] leading to a sloping voltage profile. Also, like many carbon nanomaterials,[34] graphene anodes suffer from poor first cycle columbic efficiency because of their high surface area.[35]

2.1.3 Carbon Nanotubes

Carbon nanotubes are classified as either single wall carbon nanotubes

(SWCNTs) and multiwall carbon nanotubes (MWCNTs). SWCNTs can be envisioned as seamless cylindrical graphene sheets and can be terminated by fullerene caps on one or both ends. Typically, they are 1 to 5 nm in diameter and can vary in length up to several

m.[36] Physical and electronic properties of these nanotubes are greatly affected by their chirality.[37] This determines the exact diameter of the SWCNT along with its metallic or semiconducting nature. MWCNTs have a larger diameter as they are made of multiple concentric graphene sheets. Sheet spacing is similar to graphite at 0.34 nm and wall thickness can vary greatly.[38]

Common synthesis techniques for carbon nanotubes include chemical vapor deposition,[38] laser vaporization,[39] and arc-discharge,[40] the last being the preferred method for bulk production. These methods start with a carbon source and catalyzing metal. After formation, purification is necessary to remove unreacted amorphous carbon and catalyzing metals.[41] Nitric acid oxidation is one common method of purification, which oxidizes the amorphous carbon and dissolves the metal catalysts.[42] As a result of this process, defects on the surface of the carbon nanotubes are converted to carboxylic

36 acid groups. The carboxylic acid groups can be used as a starting point for various functionalization routes.[43] Much like fullerenes, functionalization of carbon nanotubes is well studied.[44]

The electronic properties of SWCNTs are of particular interest to researchers.

SWCNTs with the “armchair” conformation are always metallic.[45] SWCNTs that show some form of chirality can be either metallic or semiconducting based how they are rolled up.[45] Most reactions produce multiple different types of carbon nanotubes so it is important to separate them to obtain material with uniform properties. Separation can be carried out by physical methods such as electrophoresis and centrifugation or chemical methods such as chromatography, selective solubility, and selective reactions.[37]

Carbon nanotubes are of interest for Li-ion anode materials because of their high electrical conductivity and Li+ uptake in the form of surface reactions in SWCNTs and intercalation in MWCNTs. Reversible capacity of 625 and 340 mAh g-1 for SWCNTs and

MWCNTs, respectively, have been reported when unloaded at voltage of 3V vs Li.[46]

SWCNTs and MWCNTs can be fabricated as freestanding electrodes reducing the overall weight of the anode. While this might increase the overall energy density, these carbon nanotube electrodes suffer from poor first cycle columbic efficiency, greatly reducing the real world practicality of this material as an anode alternative.

2.1.4 Carbon Nano Onions

Carbon nano-onions is a term for a broad range of carbon nanostructures, which are spherical in shape with multiple layers of concentric carbon shells up to 100 nm in diameter.[47] The name derives from the similarity to an onion. These nano-onions were discovered in 1980 by transmission electron microscopy, five years before the

37 fullerene, as a byproduct of carbon black synthesis.[48] It was found that graphitization of amorphous carbon into these small onion like particles occurred due to the high intensity of the electron beam.

Larger quantities of carbon nano onions can be synthesized by vacuum annealing nano diamond precursor at 1400C, resulting in 5 to 10 nm particles.[49] Onion like carbons with a size closer to 50 nm were synthesized using metals as catalysts under a variety of conditions. Examples of this include catalytic decomposition of methane on

NiO,[50] calcination of a resorcinol-formaldehyde-iron polymeric complex,[51] and chemical vapor deposition of ethane gas and liquid iron carbonyl.[52]

Carbon onions have been studied as catalytic supports and supercapacitors. As a supercapacitor, charge storage on carbon onions increase as their diameter decreases.[53]

Certain carbon onions showed to have 10 times the power of the more conventional activated carbons, but ultimately a lower energy density due to their low available surface area.[54]

Carbon onions have been largely ignored for use in Li-ion batteries, while other carbon materials have been extensively explored for use as anode materials and conductive additives.[34, 55] Publications on this subject are rare, but one paper reported that 30 nm carbon onions have a reversible specific capacity of 391mAh g-1 when cycled between 0.02 and 3V.[56] While stable out to 60 cycles, much of the capacity is discharged between 2 and 3V. This potential range is not realistic for practical Li-ion batteries as it is too close to potential at which Li+ is loaded and unloaded from the cathode.

2.2 Multilayer Graphene Nanoshells: A Description

Multilayer graphene nanoshells (MGNS) is a term employed by this lab to describe nanoscale graphitic like spheres that have a void space in the middle.[57]

Average size is 50 nm and wall thickness is ~10nm, as seen in the TEM in Figure 1A, although these parameters appear to be dependent on synthetic details. Graphitic layering is confirmed by the presence of the 002 reflection at 26.4 2Ɵ in the powder XRD

(Figure 2). Furthermore, graphitic layering, at 3.42 Å, is visible in HRTEM (Figure 1B) and polycrystalline ring SAD matches the 002 reflection for graphite (Figure 1C).

Defects in the graphitic layers allowing curling, leading to the “sphere-like” shape.

Major defects in the form of a hole in one side of the particle are formed during synthesis or purification.

Figure 12A) TEM micrograph of MGNS particle at 150kX magnification B) HRTEM of interlayer spacing of MGNS particle C)SAED of MGNS particles at camera length of 50cm.

Figure 13 Powder XRD of MGNS

2.3 Multilayer graphene nanoshells: Synthesis

The starting materials for MGNS were ground cellulose (Avicel PH-105) and

NiCl26H2O (Aldrich, ≥98%). These materials were mixed together in a ratio of 4:1

(w:w). The cellulose:NiCl26H2O mixture was pressed into pellets, 25.4 mm in diameter, at 20,000 PSI. A hole was drilled into the center of each pellet so that they could later be skewered with a 6 mm stainless steel rod. They were then charred in a tube furnace under a 13 mL s−1 argon flow at 375 C for ½ h. This was followed up by skewering the pellets and placing them in a custom 4-cross vacuum chamber (Figure 3). The reaction

41 chamber consisted of a ¼”. ultratorr inlet for the sample rod, a ZnSe window for the laser to pass through, a beam stop and a vacuum port. The chamber was then evacuated and a very low flow of He allowed to leak in, bringing the pressure up to between 700 and 1000 mTorr, preventing accumulation of soot on the ZnSe window. The pellets were rotated at

0.0125 rotations/s with an Applied Motion Products STM 23S stepping motor while the pellet’s outer edge was pyrolized with a 64 W, 10.6 m CO2 laser (Synrad Firestar t60), operating at 95% power. Pellets were collected and their edges were scraped off with a razor blade. The pyrolized edge material from the pellets was ground in a 250 mL stainless steel bowl using 4 stainless steel balls (1.27 cm in diameter) with a Fritsch

Pulverisette 6 planetary ball mill at 5000 rpm for 10 min.

Figure 14 Schematics of Laser Pyrolization table and 4-cross reaction chamber.

At this point in the synthesis, the sample consisted of Ni encapsulated in MGNS

(Ni@MGNS) and amorphous carbon (Figure 4A). The sample was mixed with concentrated HNO3 at a ratio of 1 g:10 mL (w:V) in a round bottom flask. The mixture was refluxed for 4 hours, cooled, and quenched with equal volume distilled water. The

MGNS/water mixture was then collected and washed with water and separated by centrifuged multiple times until the supernate was clear. Initially, the supernate is dark green or brown as a result of the dissolved Ni and carbon. Finally, the MGNS was allowed to air-dry overnight resulting in the final product (Figure 4B).

Figure 15 TEM image acquired at 100 kX of A) Ni@MGNS before HNO3 reflux and B) MGNS after HNO3 reflux.

2.3.1 Physical Characterization

Powder x-ray diffraction (XRD) patterns were obtained with a Rikagu

Miniflex+ diffractometer (Cu Kα radiation) maintained in the air-free environment of a N2 filled glovebox (< 1 ppm H2O and O2). Samples were scanned on a zero background quartz plate from 10° to 90° 2θ, at 0.5° min-1. Morphology was examined using a 100 kV JEOL 1200EX TEM. Raman spectra were collected on a

Horiba LabRam HR Raman microscope, with a 532 nm source. After baseline corrections, the intensities at 1330 and 1560 wavenumbers were measured to acquired ID/Ig ratio. Thermogravemetric analysis was performed on a Perkin –

Elmer Pyris 1 with air or N2 flow of 20 ml min-1 and a ramp rate of 20°C min-1.

2.3.2 Mixing of Salt and Cellulose

Two methods were explored for making metal halide/ cellulose mixtures. For

Method 1, the salt and cellulose were mixed by planetary ball mill, using an 80 mL stainless steel cup with 4, 10 mm, stainless steel balls at 200 RPM for 10 minutes.

Method 2 differed in that NiCl22H2O was dissolved in 250 mL of deionized water, and cellulose was then soaked in the solution for 1h, followed by drying using a rotary evaporator. Both methods produced good distributions of the NiCl22H2O in the cellulose. No differences were found in the final MGNS product by changing between these methods.

2.3.3 Effects of Charring

The second step in MGNS production is charring of the cellulose/NiCl2 pellets under argon at 375C for a half hour. Herring et. al. states that the purpose of this step is to decompose the cellulose to a network of polycyclic aromatic hydrocarbons (PAHs) around NiCl2 crystallites.[57] They continued by writing that this method has been used to produce carbon nanoparticles and is not well understood but is “strongly influenced” by the presences of the templated species.

To study the influence charring has on the production MGNS, we synthesized

MGNS with and without the charring step. Even without charring, Ni@MGNS was produced after pyrolization of the cellulose/NiCl2 pellet. Based on this observation, it is possible that laser pyrolization sequentially forms the PAH networks, reduces the NiCl2, and forms MGNS. Both powder XRD diffractograms and TEM image show no discernable difference in the post-pyrolysis product. The peaks in the powder XRD

(Figure 5) pattern are good matches to the 002 reflection of graphite and 111 reflection of

Ni, while TEM images showed agglomeration of roughly 50nm Ni@MGNS particles

(Figure 4A).

Figure 16 Powder XRD diffractogram of Ni@MGNS without charring acquired at a scan speed of at 1°/min.

Following pyrolysis, the Ni@MGNS made without charring was refluxed in

HNO3 for 4 hours. Raman spectra were collected on a Horiba LabRam HR Raman microscope, with a 532 source, of the final MGNS product analyzed to confirm any

3 2 changes in the sp to sp carbon ratio. A comparison of the ID/IG band ratio of charred and non-charred showed no change (Figure 6A&B). Both samples had a 0.74 ID/IG band ratio, showing the graphitic nature of MGNS was not affected by the elimination of the charring step.

Figure 17 Raman spectra of A) MGNS made with charring and B) MGNS made without charring.

Finally, a comparison in the yields of charred and non-charred MGNS showed little change. The yield of the product after HNO3 reflux, based on pre-refluxed mass of

Ni@MGNS, was ~25% whether charred or not, indicating no further cellulose was converted to MGNS as a result of the charring step.

The biggest difference in charred vs non-charred pellets was the mass of the pellet prior to being laser pyrolized. The mass of a charred pellet was 42% that of a non- charred cellulose/NiCl2 pellet. The majority of this mass loss is due to water and bio-oils.

According to Herring, the tar is made up of levoglucosan and other similar higher molecular weight species.[58] This is important because these byproducts are removed either during the charring step or the pyrolysis step. When the cellulose pellets are not

49 charred these byproduct cool and contaminate the vacuum line of the pyrolysis reaction chamber, which proves to be more of an engineering issue than a chemistry one.

Additionally, when collected properly, these bio-oils have potential to be alternative, green energy sources.[59] With these factors in mind it may be worth having further investigation of charring and collection of its byproducts.

2.3.4 Pellet Rotation Speed for Laser Pyrolization

The effect of cellulose/NiCl2 pellet rotation speed on MGNS synthesis in the reaction chamber was studied. By changing the rotation speed we controlled how long the pellet was under laser beam irradiation. The beam had a spot size of 2.2 mm and the pellet’s circumference was 79.8 mm. The pellet was pyrolized at 4 different rotation speeds (Table 1). The time any one portion of the pellet was under the beam (irradiation time) was calculated using the pellet circumference, spot size, and rotation speed. The pellets were scraped and product was collected and weighed. A portion (0.04g) of that product was refluxed in HNO3 for 4 hours and the final MGNS was weighed. TEM was used to analyze the Ni@MGNS and MGNS to verify any morphological changes. Raman spectroscopy was used to determine the ratio of sp3 carbon to sp2 carbon.

Table 1 – The effect of rotation speed for pyrolysis on MGNS Rotation Full Irradiation Scraping Reflux Reflux Raman Speed (RPS) Rotation Time (sec) Yield (g) Yield (g) % Yield ID/IG Time (sec) Ratio 0.05 20 0.55 0.0409 0.0011 2.59 0.41 0.0208 50 1.38 0.1088 0.0192 17.37 0.32 0.0125 80 2.21 0.1629 0.0283 17.65 0.29 0.0083 120 3.32 0.2654 0.0852 32.10 0.19

As the rotation speed decreased and the irradiation time increased, the penetration depth of the laser also increased. This was determined by the amount of product that flaked easily off the cellulose pellet with light scraping. This yield followed a linear trend with the irradiation time (Figure 7). Each sample was refluxed to determine the amount of the graphitic MGNS produced in the pyrolysis based on the percent yield. The percent yield nearly doubled between 80 and 120 second rotation times from 17.65 to

32.10%. The product also seemed to get more graphitic as irradiation time increased.

The ID/IG band ratio of the Raman spectra decreased from 0.41 to 0.19 with increasing irradiation time. The decrease in this ratio would indicate less sp3 carbon, which could be due to a larger number of graphene layers per nanoshell (less curvature) or fewer defect in the MGNS structure. TEM images showed no apparent difference in the morphology of the MGNS. Agglomerations of multiple aspherical particles are readily visible regardless of pyrolysis time (Figure 8).

Figure 18 Scraping yield of Ni@MGNS (circles) vs rotation time and linear least squares fit of the data (line).

Figure 19 TEM micrograph obtained at 200 kV of MGNS agglomerates made with pyrolysis full rotation times of A) 20 seconds and B) 120 seconds

2.4 Purification of Ni@MGNS

As previously mentioned in section 2.3, the resulting product of the laser pyrolysis is Ni@MGNS and amorphous carbon. In order to isolate the MGNS, we must remove the amorphous carbon and Ni cores. Purification of MGNS is similar to that of

MWCNTs. Most techniques employ a route in which competitive oxidation between the amorphous carbon and graphitic carbon takes place. The key to such methods is that the oxidation of amorphous carbon is kinetically much quicker than graphitic carbon due to the excess of available dangling bonds. Additionally, these techniques dissolve the Ni cores. The following sections discuss four routes for purifying MGNS. While we are not limited to these methods, they are the ones that showed the most success for MGNS.

2.4.1 Air Oxidation

Ni@MGNS/amorphous carbon was heated in a tube furnace at 500C with a 20 ml min-1 air flow for 300 min. The sample was collected and washed in concentrated HCl to remove the now oxidized NiO cores. The process followed reactions 1 through 3.

(1) C + O2  CO2

(2) 2C + O2  2 CO

(3) NiO + 2 HCl  NiCl2(aq) +H2O

TGA was used to determine the optimal time of oxidation. Temperature was ramped at

20C min-1 to 500C and held for 300 minutes. Two distinct slopes in the profile of the

TGA curve can are apparent (Figure 9). There is a quick descending slope to 75% mass in the first 27 minutes. This portion of the curve is due to amorphous carbon oxidation.

Nickel oxidation causes a hump in the profile due additional oxygen mass. After 27 minutes the slope changes rapidly to a much shallower loss in mass due to MGNS graphitic carbon oxidation. Powder XRD was collected following HCl wash (Figure 10).

As expected, there is a reflection at 26.2 for the graphitic MGNS along with small broad peaks for some leftover NiO. An ID/IG Raman band ratio of 0.71 indicates the sample has an increased level of defects from this high temperature oxidation. TEM showed definitive evidence of additional unoxidized amorphous carbon (Figure 11). It is possible that oxidation does not occur homogenously and sections of the sample that are not in contact with gas flow oxidize at a slower rate leading to excess amorphous carbon. In

53 order for this method to be effective, the time must be increased and over-oxidation must take place, which in turn may lead to a more defective MGNS structure.

Figure 20 TGA thermogram of Ni@MGNS heated under air at a ramp rate of 20C min-1.

Figure 21 Powder XRD pattern of MGNS purified by air oxidation and HCl wash collected at a scan rate 0.5 °C/min.

Figure 22 TEM micrograph obtained at 200 kV of MGNS oxidized by air then washed by HCl. Excess amorphous carbon is visible intermixed within the MGNS particles.

2.4.2 Potassium Permanganate Oxidation

526.3 mg KMnO4 was dissolved in 50 mL of 0.5 M H2SO4 to make a 0.3 M

KMnO4 solution. 100 mg Ni@MGNS/amorphous carbon was heated in this solution at

80C for 6 hours until the solution turned from violet to black, at which point it was allowed to heat for an additional 1 hour. The sample was collected and separated by

56 centrifuge at 5000 rpm for 5 minutes. The solid was washed with 50 ml of concentrated

HCl, and then with 50 ml of 0.5 M HCl, and separated by centrifuged after each wash.

Finally, the sample was allowed to dry overnight in air.

- This purification uses MnO4 to oxidize carbon to CO2 as seen in reaction 4.

+ + (4) 3 C + 4 KMnO4 + 4 H  4 MnO2 + 3 CO2 + 4 K + 2 H2O

The solution turned from violet to black once all the KMnO4 reacted to MnO2. The resulting product was Ni@MGNS and NiO@MGNS coated with small MnO2 crystallites

(Figure 12). Concentrated HCl was used to dissolve the Ni and NiO cores (reaction 3) while the dilute HCl removed the MnO2 (reaction 5)

(5) MnO2 + 4 HCl(aq)  MnCl2(aq) + 2 H2O + Cl2(g)

Figure 23 TEM at 100 kV of MnO2 coated Ni@MGNS 75kX.

The powder XRD pattern (Figure 13) of the purified MGNS consisted of a broad

(002) graphitic reflection and small peaks consistent with Ni and NiO. Additional investigation by Raman spectroscopy show a high ID/IG band ratio of 0.80, characteristic of increased defects in the structure.

Figure 24 Powder XRD pattern of KMnO2 oxidized and HCl washed MGNS obtained at a scan rate 0.5°/min.

2.4.3 Steam Oxidation

The steam purification setup can be seen in Figure 14. A ceramic boat holding

0.2 g Ni@MGNS/amorphous carbon was placed in a 25 mm diameter quartz heating tube inside a tube furnace. The heating tube was purged with Ar for 0.5 h. Ar gas was bubbled, via sparger, through distilled water held at 98C. The tube furnace was set to

950 C and the water saturated Ar flow at 200 ml/min, for 3 hours. After the sample cooled, it was washed with concentrated HCl to remove the Ni cores.

Figure 25 Steam purification setup. Ar gas is bubbled through distill water heated to 98C. Ar/steam then flows through the copper tube into a quartz tube within furnace held at 950C and vented to a fume hood.

The steam oxidizes the carbon, following reaction 6, to CO and is further oxidized to

CO2 (reaction 7).

(6) C + H2O  CO + H2

(7) CO + H2O  CO2 + H2

The powder XRD pattern of the MGNS showed it was free of Ni after HCl wash (Figure

15). The ID/IG band ratio of the Raman spectra, 0.68, was lower than that of both the air and KMnO4 oxidation. The TEM micrograph of the material displayed no abnormalities in the MGNS morphology, with very little observable amorphous carbon (Figure 16).

Figure 26 Powder XRD pattern of steam purified MGNS after HCl wash obtained at a scan rate of 1°/min.

Figure 27 TEM micrograph obtained at 200 kV of MGNS oxidized by steam at 200 °C followed by HCl washing to remove Ni cores.

2.4.4 Nitric Acid Reflux Oxidation

The Ni@MGNS was mixed with concentrated HNO3 at a ratio of 1 g:10 mL

(w:V) in a round bottom flask. The mix was refluxed for up to 6 hours, cooled, and quenched with equal volume distilled water. The MGNS/water mix was then collected and washed with water 3 times followed by separation by centrifuge.

HNO3, a strong oxidizing, easily oxidized amorphous carbon at elevated temperatures, following reaction 8. Additionally, because it is a strong acid, it dissolves the Ni cores at the same time (reaction 9)

(8) C + 4 HNO3  CO2 + 2 H2O + 4 NO

(9) Ni + HNO3  Ni(NO3)2(aq) + H2

Ni@MGNS/amorphous carbon oxidation by refluxing HNO3 was halted after ½, 2, 4, and

6 h. The products were analyzed for crystal structure with powder XRD, surface area using BET, and ID/IG band ratio with Raman spectroscopy. Results are shown in Table 2.

Table 2 – The effect of HNO3 reflux time on MGNS Reflux time (Hours) % Yield ID/IG ratio BET Surface Area (m2/g) 0.5 37 0.80 23 2 34 0.52 56 4 37 0.61 82 6 23 0.66 107

The yield of the samples oxidized for ½ - 4 hours was very similar, between 34 and 37%. Most of the mass is lost very quickly because the Ni is dissolved out rapidly.

TGA thermograms collected with a 20 ml/min airflow (Figure 17) shows after all the carbon is removed by oxidation, there is only 3% of the initial mass remaining. This is consistent for all the samples, independent of the reflux time, and is due to NiO that was not removed. After 6 hours the yield drops down to 23%. At this point the reflux may not just be removing amorphous carbon, but a significant portion of the graphitic carbon of the MGNS as well. The surface area MGNS washed by in refluxing HNO3 for 6 h was

2 -1 107 m g , much higher than any of the other samples. The ID/IG band ratio is lowest for the sample treated for 2 h, 0.52, and largest, 0.80, for the sample treated for the shortest time (½ h). This indicates that the brief treatment left a large amount of disordered carbon, most likely because of unoxidized amorphous carbon. TEM micrographs (Figure

18 A-D) show a decrease in amorphous carbon around the MGNS as reflux time increased. Furthermore, HRTEM images of MGNS after a 2 h treatment (Figure 18E) show the expected graphitic layering with little evidence of disordered carbon around the edges. Based on these results, HNO3 reflux time of 2 hours was determined to be the most appropriate method for purifying the MGNS.

Figure 28 TGA thermogram of HNO3 oxidized MGNS collected at a flow rate 20 ml air/min.

Figure 29 TEM micrographs of MGNS refluxed in HNO3 for A) ½ h B) 1 h C) 2 h and D) 6 h. E) HRTEM of MGNS after 2 h HNO3 reflux, showing graphitic layering at the edges.

2.5 Electrochemical Characterization

Detailed electrochemical characterization was previously reported by Jonathan

Cox.[60] MGNS was showed to have a nominal capacity of 220 mAh g-1 with little capacity fade over 200 cycles when cycled from 0.02 to 1.5V vs Li, in half cell configuration. It has high rate capability, charging/discharging as high as 16C for 200 cycles with ~10% capacity fade. Additionally, MGNS can cycle in low temperature solvents such as PC without any co-intercalation, unlike graphite, which suffers from exfoliation in these solvents.[61]

2.6 Conclusions

Multilayer graphene nanoshells have been developed for use in Li-ion anodes.

The synthesis was investigated further for optimization. It was determined that while charring may provide useful byproducts, it is not crucial to the synthesis of the MGNS nanoparticles. Increased pyrolysis times were found to increase the overall yield.

Several oxidative purification techniques were investigated including oxidation by air, steam, KMnO4, and HNO3. HNO3 proved to be the most successful at removing amorphous carbon and Ni metal. While HNO3 proved excellent at removing amorphous carbon, it could also destroy the graphitic MGNS during prolonged reflux times. An

HNO3 reflux time of 2 hours was found to be optimal to remove amorphous carbon without being too detrimental to the MGNS. Several aspects of this synthesis were investigated in this study to provide a foundation for the carbon backbone needed for

Sn@MGNS studies reported in the subsequent chapter.

2.7 References

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41. Hou, P.X., C. Liu, and H.M. Cheng, Purification of carbon nanotubes. Carbon, 2008. 46(15): p. 2003-2025. 42. Shao, L., et al., Removal of amorphous carbon for the efficient sidewall functionalisation of single-walled carbon nanotubes. Chemical Communications, 2007(47). 43. Wu, H.-C., et al., Chemistry of carbon nanotubes in biomedical applications. Journal of Materials Chemistry, 2010. 20(6): p. 1036-1052. 44. Tasis, D., et al., Chemistry of carbon nanotubes. Chemical Reviews, 2006. 106(3): p. 1105-1136. 45. Dai, H.J., Carbon nanotubes: opportunities and challenges. Surface Science, 2002. 500(1-3): p. 218-241. 46. Landi, B.J., et al., Carbon nanotubes for lithium ion batteries. Energy & Environmental Science, 2009. 2(6): p. 638-654. 47. McDonough, J.K. and Y. Gogotsi, Carbon Onions: Synthesis and Electrochemical Applications. Interface, 2013. 22(3): p. 61-65. 48. Iijima, S., Direct Observation Of The Tetrahedral Bonding In Graphitized Carbon-Black By High-Resolution Electron-Microscopy. Journal of Crystal Growth, 1980. 50(3): p. 675-683. 49. Cebik, J., et al., Raman spectroscopy study of the nanodiamond-to-carbon onion transformation. Nanotechnology, 2013. 24(20). 50. He, C.N., et al., Low-temperature synthesis of carbon onions by chemical vapor deposition using a nickel catalyst supported on aluminum. Scripta Materialia, 2006. 54(4): p. 689-693. 51. Zhang, C., et al., Hollow graphitic carbon nanospheres: synthesis and properties. Journal of Materials Science, 2013. 49(5): p. 1947-1956. 52. Wang, J.N., et al., Synthesis of high surface area, water-dispersible graphitic carbon nanocages by an in situ template approach. Chemistry of Materials, 2007. 19(3): p. 453-459. 53. Huang, J.S., et al., Curvature effects in carbon nanomaterials: Exohedral versus endohedral supercapacitors. Journal of Materials Research, 2010. 25(8): p. 1525-1531. 54. Pech, D., et al., Ultrahigh-power micrometre-sized supercapacitors based on onion-like carbon. Nature Nanotechnology, 2010. 5(9): p. 651-654. 55. Su, D.S. and R. Schlogl, Nanostructured Carbon and Carbon Nanocomposites for Electrochemical Energy Storage Applications. Chemsuschem, 2010. 3(2): p. 136-168. 56. Han, F.D., B. Yao, and Y.J. Bai, Preparation of Carbon Nano-Onions and Their Application as Anode Materials for Rechargeable Lithium-Ion Batteries. Journal of Physical Chemistry C, 2011. 115(18): p. 8923-8927. 57. Herring, A.M., et al., A novel method for the templated synthesis of homogeneous samples of Multilayer graphene nanoshells from cellulose chars. Journal of the American Chemical Society, 2003. 125(33): p. 9916-9917. 58. Lee, J.G., et al., Effect of metal doping on the initial pyrolysis chemistry of cellulose chars. Energy & Fuels, 2008. 22(4): p. 2816-2825. 59. Mohan, D., C.U. Pittman, and P.H. Steele, Pyrolysis of Wood/Biomass for Bio- oil: A Critical Review. Energy & Fuels, 2006. 20(3): p. 848-889. 69

60. Cox, J.D., Electrochemical study of Multilayer graphene nanoshells as high-rate and low temperature negative electrodes for lithium ion batteries, 2015, The George Washington University, Ph.D. 61. Aurbach, D., et al., On the correlation between surface chemistry and performance of graphite negative electrodes for Li ion batteries. Electrochimica Acta, 1999. 45(1-2): p. 67-86.

3. Incorporation of Tin into Multiwall Graphene Nanoshells 3.1 Introduction

Intercalation based carbons have been a primary component of Li-ion battery anodes since the early 90’s, but they have some fundamental drawbacks. One of the largest of these is the limitation of Li intercalation to 1 Li+ per 6 C in graphite, the active material in commercial Li-ion battery anodes. This results in a relatively low theoretical specific capacity of 372 mAh/g, and an even lower practical capacity (~340 mAh/g) due to defects in the graphitic crystal structure.[1] Another drawback is the electrochemical potential at which Li+ intercalates into graphite, 0.02 V vs Li, very close to the potential that Li plates (0 V vs Li). This can become problematic at fast loading rates and low temperatures where Li can locally plate on the graphite because of the slow intercalation kinetics of graphite.[2] Li plating can lead to dendrite formation and consequential internal shorts.[3] As a result of these issues with graphite, enormous research effort has been put into alternative materials, with Li-alloying anode materials being chief among these (Section 1.3.2).

Encapsulation of Li-alloying nanoparticles by materials that do not react with Li+ has been explored by many research groups.[4-7] The most commonly studied systems consist of a Li alloying metal surrounded by a carbon shell, allowing enough void space for the expansion that goes along with the alloying reaction (Figure 1). The nanoparticles inside the shell can undergo a certain degree of pulverization while maintaining good electrical contact with the walls of the carbon shells. Agglomeration is prevented because the nanoparticles are confined inside the shells. An example of this is the “yolk-shell” Si-

71 nano morphology reported by Cui’s group that cycled stably for 600 cycles with over

1500 mAh g-1.[8]

Figure 1 Schematic of the lithiation process of encapsulated Li-alloy based materials. Multiwall graphene nanoshells (MGNS) are composed of multilayered curved graphene with an inner void (chapter 2, Figure 1a).[9] Lithium can intercalate between the graphene layers, akin to the storage mechanism of graphite. However, its storage capacity, ~220 mAh g-1,[10] could be greatly improved by making use of the empty interior space. An appropriate loading of Sn nanoparticles on the inside of MGNS would allow for expansion without mechanical stress and prevent agglomeration. Using MGNS as an encapsulating material would differ from other similar carbon coatings, as it would contribute to Li storage capacity as well as acting as physical confinement.

In addition to the interior space that MGNS could provide for Sn nanoparticles, the spaces between packed MGNS could provide space for additional Sn or other alloying metal(s). While MGNS might be simplistically envisioned as nanospheres, electron microscopy observations demonstrate that they are more often highly distorted shells, consistent with considerable wall flexibility. This flexibility may provide stress relief to allow alloying metal(s) in between, rather than or in addition to inside, the spheres to reversibly expand/contract during lithiation/delithiation while maintaining electrode mechanical integrity and metal nanoparticle separation.

In this chapter, separate studies of MGNS anodes with Sn inside (Sn@MGNS) and outside (Sn/MGNS) the shells are described. While separation of the studies of these two Sn-MGNS composite options simplified the research, the final anode material envisioned would be a combination of the two; MGNS modified by partially filling both the insides of the shell and the interstices between the shells. This would maximize use of the vacant space inside and around the electrically conductive, lithium intercalating

MGNS “scaffolding”, optimizing both gravimetric and volumetric energy density.

Furthermore, the Sn-MGNS, may require little or no additional processing in order to be applied to traditional Li-ion battery manufacturing techniques. Sn-MGNS similarities to graphite should allow it to directly replace graphite in Li-ion battery manufacturing lines.

3.2 Experimental 3.2.1 Synthesis – Sn@MGNS

The synthesis of Sn@MGNS is accomplished by a simple modification of the synthesis of MGNS. Various ratios of the NiCl26H2O (Aldrich, ≥98%), and SnCl22H2O

(Fisher Scientific, ≥95%) salts were mixed with ground cellulose (Avicel PH-105) in a ratio of 1:4 (w:w). The salt:cellulose mixtures were pressed into pellets, 25.4 mm in diameter, at 20,000 PSI. A hole was drilled in the center of each pellet and they were skewered with a 6mm stainless steel rod which was placed in a custom 4-cross vacuum chamber (Chapter 2, Figure 3). The pellets were rotated at 0.0125 rotations per second with an Applied Motion Products STM 23S stepping motor while the pellet’s outer edge was pyrolized with a 64 W, 10.6 m CO2 laser (Synrad Firestar t60), at 95% power. The pyrolized material was scraped from the edges by razor blade and ground for 10 min in a

250 mL stainless steel bowl using 4 stainless steel balls (1.27 cm in diameter) with a

Fritsch Pulverisette 6 planetary ball mill operating at 5000 rpm.

The product consisted of a NixSny intermetallic encapsulated MGNS (mixed metal

MGNS), NixSny intermetallic exterior to the MGNS, and amorphous carbon. The exterior intermetallic was first removed by washing the sample with concentrated HCl with ultrasonic agitation for 2 hours. The sample was then washed with deionized water followed by refluxing concentrated HNO3 for 4 hours (1g:10mL sample:HNO3), oxidizing and removing the amorphous carbon. In addition, the HNO3 leached the Ni from the intermetallic cores, leaving SnO2 encapsulated by MGNS (SnO2@MGNS). The product was dispersed in water with ultrasonic agitation and separated by centrifugation

(5000 RPM, 5 min) multiple times until the supernate appeared clear. The SnO2 cores were reduced back to Sn by H2 reduction, using a 5%H2/95%N2 gas mixture at a flow rate of 20 ml min-1 through a tube furnace at 800⁰C, leaving Sn@MGNS.

3.2.2 Synthesis – SnO2/MGNS

A 0.1852 M NaBH4 solution was prepared by dissolving and slowly diluting 0.70 g NaBH4 in 100 ml of 14 M NaOH at 0 ˚C. The borohydride solution was stored in an ice/water bath to slow hydrolysis.[11] MGNS (0.15 g) was dispersed in a 25 ml solution of SnCl2·2H2O in deionized water by magnetic stirring. A range of SnCl2·2H2O concentrations were examined from 12 mM to 56 mM in order to adjust the final SnO2 ratio relative to the MGNS. The NaBH4 solution was slowly added dropwise by pipet to the MGNS dispersed in SnCl2 solution. The reaction occurred quickly, with a few seconds vigorous bubbling (likely the release of H2 gas), after which the solution was stirred for an additional 5 min. The basic solution was neutralized with 6 M HCl to help

74 precipitate the product (SnO2/MGNS). The product was washed and separated by centrifugation 3 times with water and then once with ethanol and then dried in air overnight.

3.3 Characterization 3.3.1 Mixed Metal MGNS

After pyrolization, the product consisted of NixSny particles in the range of 50 to

200 nm, along with the encapsulated mixed metal MGNS. Washing with concentrated

HCl removed the NixSny exterior but not that encapsulated by the MGNS. TEM of 90:10

Ni:Sn mixed metal MGNS before and after the HCl wash shows the loss of the exterior metal after the wash (Figure 2A&B). A 200 nm intermetallic particle can be seen prior to HCl wash along with other smaller agglomerates of metal nanoparticles, in addition to mixed metal MGNS. After the HCl wash, 50nm metal cores are visible surrounded by

MGNS and amorphous carbon. The metal content was independent of the HCl wash time, showing a mass loss of ~40% after wash times of 1 and 2 hours. Powder XRD diffractograms, Figure 3A&B – 7A&B, display a decrease in the intensity of the NixSny reflections, but the peak positions remain the same. X-ray fluorescence was used to monitor the Ni to Sn mass ratio by relative peak intensities of the Kα1 peaks of Ni (7.48 keV) and Sn (25.27 keV). The mass ratio was not significantly changed by pyrolization or HCl washing (Figure 8). HCl washing did result in an increase in the Raman spectroscopy ID/IG band ratio across all Ni:Sn mass ratios (Figure 9). This suggests that

HCl washing results in an increase in defects in the MGNS as has been shown to occur in carbon nanotubes, while having very little effect on the amorphous carbon content.[12]

Figure 2 TEM micrograph at 75kX, 100kV of 9:1 Ni:Sn mixed metal MGNS A) prior to HCl wash B) after HCl wash.

Figure 3 Powder XRD diffractograms collected at 0.5°/min of 9:1 Ni:Sn mixed metal MGNS A) after pyrolysis B) after HCl wash C) after HNO3 reflux D) after H2 reduction. Reflections are indicated as MGNS (circles), Ni (triangles), SnO2 (squares), Sn (diamonds), unidentified (asterisks).

Figure 4 Powder XRD diffractograms collected at 0.5°/min of 8:2 Ni:Sn mixed metal MGNS A) after pyrolysis B) after HCl wash C) after HNO3 reflux D) after H2 reduction. Reflections are indicated as MGNS (circles), Ni (triangle), Ni3Sn (empty triangles) SnO2 (Square), Sn (Diamond), unidentified (asterisks).

Figure 5 Powder XRD diffractograms collected at 0.5°/min of 7:3 Ni:Sn mixed metal MGNS A) after pyrolysis B) after HCl wash C) after HNO3 reflux D) after H2 reduction. Reflections are indicated as MGNS (circle), Ni (triangle), Ni3Sn (empty triangle), Ni3Sn2 (pentagon), SnO2 (square), Sn (diamond), unidentified (asterisks). The subwindow is a zoomed in area from (D) showing the MGNS peak.

Figure 6 Powder XRD diffractograms collected at 0.5°/min of 6:4 Ni:Sn mixed metal MGNS A) after pyrolysis B) after HCl wash C) after HNO3 reflux D) after H2 reduction. Reflections are indicated as MGNS (circle), Ni (triangle), Ni3Sn (empty triangle), Ni3Sn2 (pentagon), SnO2 (square), Sn (diamond), unidentified (asterisks). The subwindow is a zoomed in area from (D) showing the MGNS peak.

Figure 7 Powder XRD diffractograms collected at 0.5°/min of 5:5 Ni:Sn mixed metal MGNS A) after pyrolysis B) after HCl wash C) after HNO3 reflux D) after H2 reduction. Reflections are indicated as MGNS (Circle), Ni (triangle), Ni3Sn (empty triangle), Ni3Sn2 (pentagon), SnO2 (square), Sn (diamond), unidentified (asterisks). The subwindow is a zoomed in area from (D) showing the MGNS peak.

100

Mass % Mass Sn 40

0 90:10 80:20 70:30 60:40 50:50 Starting Mass Ratio

Figure 8 Sn mass percent, as measured by XRF relative to Ni mass, after each synthesis step and for each initial Ni:Sn ratio. Pyrolization (red bar), HCl wash (blue bar), HNO3 reflux (purple bar), and H2 reduction (orange bar).

0.9

0.8

0.7

0.6

0.5 I (D/G) I 0.4

0.3

0.2

0.1

0 90:10 80:20 70:30 60:40 50:50 Starting Mass Ratio

Figure 9 The Raman ID/IG ratio after each synthesis step for each Ni:Sn ratio. Pyrolization (red bar), HCl wash (blue bar), HNO3 reflux (purple bar), and H2 reduction (orange bar).

3.3.2 SnO2@MGNS

Refluxing the mixed metal MGNS in concentrated HNO3 resulted in the oxidation and removal of the amorphous carbon and Ni metal. However, low intensity XRD reflections between 42 and 45° 2Ɵ, corresponding to Ni (PDF#04-0850), Ni3Sn

(PDF#35-1362), or Ni3Sn2 (PDF#06-0414) indicate that a very small fraction of the Ni was not removed (Figures 3-7). In addition, hot HNO3 readily oxidizes Sn to SnO2, an insoluble species in an acidic environment. This can be seen by the increase in the intensities of the SnO2 (110) and (101) reflections at 26.6 and 33.9° 2Ɵ. Peak analysis using the Scherrer equation indicates that the SnO2 crystallite size is ~5 nm.

XRF spectra are consistent with the XRD results, showing up to 52% (m/m) of the metal remaining (Sn + Ni) in 9:1 Ni:Sn MGNS is Ni, with lower mass ratios showing much more complete Ni removal. For example, ~20% of the metal remaining in the 7:3,

8:2, and 9:1 Ni:Sn@MGNS samples is Ni. It is unclear if the Ni remains as partially oxidized NixSny/SnO2@MGNS particles or as completely unoxidized NixSny cores in some portion of the MGNS particles. It should be noted that a small fraction of the Ni cores of MGNS synthesized with only Ni survive HNO3 reflux, suggesting that the same protection might be provided to some of the NixSny cores of the mixed metal material.

Refluxing resulted in a decrease in the Ramen ID/IG band ratio across all samples, except the 50:50 Ni:Sn MGNS sample, consistent with the oxidation and removal of the amorphous carbon. The ID/IG band ratio increases with increasing Sn content, suggesting that the defects present in the MGNS increase with decreasing Ni content during pyrolysis. This is not particularly surprising considering that MGNS cannot by synthesized from Sn in the absence of Ni.

TEM images show SnO2@MGNS particles ranging in size from 50 to 150 nm that do not completely fill MGNS, consistent with Ni removal from the cores (Figure

10A ). The SnO2 nanoparticles are far bigger than the XRD peak widths indicate, suggesting that they are agglomerates of much smaller nanocrystals. In addition, the images show that amorphous carbon has been nearly completely removed from the material. The SAD pattern of the partially filled MGNS (Figure 10B) shows point reflections corresponding to SnO2 and MGNS.

Figure 300 A) TEM image of SnO2@MGNS made from 6:4 Ni:Sn mixed metal MGNS obtained at 100 kx magification, 100 kV. B) SAD pattern of area in (A) Camera length 50cm. Indices (hkl) are indicated for SnO2 (white font, PDF#41-1445) and MGNS (red font) reflections.

3.3.3 Sn@MGNS

SnO2 was reduced to Sn, according to Equation 1, at 800°C.

∆ 1. SnO2 + 2H2(g) → Sn((s, l)) + 2H2O(g)

The reaction is thermodynamically favorable at temperatures above 550°C,[13] but reduction was not observed in XRD patterns of these samples below 750 °C. Following reduction, the broad SnO2 peaks are no longer evident, having been replaced by sharper peaks that are consistent with Sn. Analysis of the peak widths by the Scherrer equation yields average crystallite sizes from 20 to 30 nm, in good agreement with the size of the

Sn nanoparticles seen in TEM images (Figure 11).

Taken together, the XRD and TEM results are consistent with the transformation of the

SnO2 nanocrystal agglomerates into Sn nanoparticles consisting of single or few crystalline domains in the interior of the MGNS (Sn@MGNS). However, it is clear that the size distribution of nanocrystals in the Sn@MGNS is broad. Some MGNS appear more filled than others, while some are not filled with metal at all. High resolution TEM shows the MGNS maintains its graphitic nature around the metal cores (Figure 12). The

Raman ID/IG generally decreased after H2 reduction (Figure 9), indicating a reduction in the defects in the MGNS caused by HNO3 oxidation The mass percent of the metal in the

Sn@MGNS samples was determined by measuring the mass loss during oxidation by

TGA, assuming complete oxidation to SnO2 and NiO (Figure 13). After H2 reduction, the relative amount of Ni and Sn remained largely unchanged in the Sn@HCNS.

(Figure 8). Leftover Ni is not immediately evident in the powder XRD patterns of the H2 reduced Sn@MGNS particles; however, the (111) reflection for Ni occurs at 44.5° 2Ɵ, overlapped by the (211) reflection of Sn at 44.9° 2Ɵ. A detailed investigation of Figures

3D through 7D show a small shoulder on the lower angle side of the Sn (211) reflection at 44.9, likely the result of residual Ni.

Figure 11 TEM image of Sn@MGNS produced from 7:3 Ni:Sn mixed metal MGNS. 75kX, 100kV.

Figure 12 HRTEM image of Sn@MGNS made from 7:3 Ni:Sn mixed metal MGNS. 400kX, 125kV.

Figure 13 TGA thermogram of Sn@MGNS samples oxidized in air at a ramp rate of 20 °C min-1, to 900 ˚C and held at 900 °C for 0.5 h.

3.4 Results and Discussion

The mass percentages of Sn, Ni and MGNS were calculated from TGA and XRF results for the Sn@MGNS, which in turn enabled us to calculate the theoretical specific capacity of each sample (Table 1). Again, it should be noted that not all of the Ni could be removed. Ni accounts for ~7% of the mass in all the samples, regardless of composition. Also, Sn content did not increase beyond ~36% for the 7:3, 6:4, and 5:5 samples. This is consistent with the Sn mass percent in Ni3Sn, after Ni has been removed, suggesting that MGNS will not grow around nanoparticles with higher Sn fraction, leaving nanoparticles with excess Sn without protection from the initial HCl wash. The theoretical capacities are calculated according to Equation 2:

mAh 2. Specific Capacity ( ) = %Sn(960) + %Ni(O) + %MGNS(200) Sn@MGNS g

The majority of the contribution of the capacity of each of the compositions is a result of the Sn metal, while Ni does not form Li alloys and thus does not add to the capacity. A portion of the capacity is due to the intercalation of Li into MGNS. The sum of each of these contributions is the total theoretical capacity of the Sn@MGNS. Based on the contents of each of these components, the max specific capacity that can be obtained is

~470 mAh g-1.

Table 1 – Capacity of Sn@MGNS based on Ni:Sn ratio

Sn@MGNS Initial % Sn % Ni % C Theoretical Experimental Composition Mass Mass Mass Mass Capacity Capacity Label Ratio (mAh g-1 (mAh g-1 (Ni:Sn) Composite) Composite) Sn@MGNS- 9:1 8.2 7.6 84.3 247 191 A Sn@MGNS- 8:2 16.3 7.3 76.4 309 246 B Sn@MGNS- 7:3 35.9 7.4 56.7 458 300 C Sn@MGNS- 6:4 37.2 6.0 56.7 471 306 D Sn@MGNS- 5:5 36.8 6.5 56.7 467 306 E

Anodes with a composition of 80% Sn@MGNS, 10% Timical Super P carbon black and 10% Kureha 1100 PVDF binder were prepared for electrochemical testing.

Active loading was ~3 mg cm-2 on copper foil (0.127 mm thick, 98%, Alfa Aesar). The electrodes were cycled against Li metal (99.9%, MTI Corp.) with Celguard 2400 separator and 1M LiPF6 (99.999%, Novolyte) in 1:1 EC:DMC (v:v) electrolyte (99.99%,

Novolyte), in 2016 style coin cells(MTI Corp.). The cells were cycled at a rate of C/10

(based on theoretical capacities stated in Table 1), between 0.02 and 1.5V vs Li metal, unless otherwise noted. The cells were allowed to rest for 1 h at open circuit voltage between each load and unload step.

Figure 14 shows the voltage profiles of the first cycle of the five compositions of

Sn@MGNS samples. Sn@MGNS-C, Sn@MGNS-D and Sn@MGNS-E had nearly identical load and unload curves resulting in a reversible capacity of 306 mAh g-1.

Sn@MGNS with lower Sn content (Sn@MGNS-A and Sn@MGNS-B) had reversible

91 capacities of 191 and 246 mAh g-1, respectively. The dealloying plateaus coordinating with Li7Sn3, LiSn, and Li2Sn5 are visible at 0.58, 0.7, and 0.78V vs Li, respectively, for all Sn@MGNS samples. This can be better seen in the differential capacity plot, making the unload of lithium at 0.43V vs Li more obvious. This broad hump is at potentials consistent with the Li17Sn5, Li7Sn2, Li13Sn5, and Li5Sn2 phases. In addition, the differential capacity exemplifies the harder to see Li loading plateaus. Sharp anodic peaks at 0.62 and 0.4 V match loading potentials for LiSn and Li7Sn3. The anodic peaks for Li2Sn3, Li5Sn2, Li13Sn5, Li7Sn2, and Li17Sn5 are less apparent. The anodic peak for

Li2Sn3 sits at ~0.68V, but on the first cycle is overwhelmed by a broad anodic peak ranging 0.84 V to 0.68 V due to SEI formation. The SEI primarily forms on the first cycle and the anodic peak of Li2Sn3 is clearly visible on the second cycle. Formation of the Li5Sn2, Li13Sn5, Li7Sn2, and Li17Sn5 phases produces much smaller anodic peaks due the considerably less Li added in each respective phase. Observing their formation in the load curve is further complicated by Li intercalation into the graphitic layers of MGNS, which begins to take place at similar potentials.[14] Small anodic peaks at 0.1 and 0.07 V vs Li occur where the stage-2 and stage-1 intercalated compounds LiC12 and LiC6 formation takes place, respectively. The unload curves of Sn@MGNS appear to be superpositions of those of MGNS and Sn. Sn@MGNS-A has a nearly identical profile to

MGNS below 0.5V vs Li. As the mass percentage of Sn is increased in the following samples, the unload curves appear more and more like that of Sn; the deintercalation from MGNS becomes difficult to discern in the profiles.

Figure 14 First cycle voltage profile of various Sn@MGNS compositions cycled at C/10.

Figure 15 Differential capacity of the first (blue line) and second (red line) cycles of Sn@MGNS-D.

Capacity retention with long term cycling of Sn@MGNS was tested by performing 100 cycles at C/10 (Figure 16). The capacity faded rapidly over the first 20 cycles for all compositions and became more stable for the remaining cycles. The samples with higher loadings, Sn@MGNS-C, Sn@MGNS-D, and Sn@MGNS-E, lost the most capacity, fading from ~300 to ~135 mAh g-1. Sn@MGNS-A had considerably less fade, losing only ~20 mAhg, while Sn@MGNS-B was between the other two compositions losing 80 mAh g-1. Sn@MGNS-B had a rapid drop off of capacity at the

20th cycle but regained it after 10 more cycles. This was likely an anomaly due to an inconsistent connection to the potentiostatic cycler. This cell regained capacity that could be extrapolated from its initial behavior and continued to cycle without fade for the next

70 cycles.

Figure 16 Long term cycling of various Sn@MGNS compositions at C/10. The Coulombic efficiencies of these cells are shown in Figure 17. The first cycle

Coulombic efficiency increased with mass percent of Sn in the Sn@MGNS formulations, from 48 to 66%. Initial Coulombic efficiency is generally inversely proportional to the surface area of a material; the carbon nanomaterials used in this study have large surface 94 areas that must be passivated and thus have low initial efficiencies.[1] The Coulombic efficiencies increase to ~95% on the second cycle, showing that the initial passivation film was largely effective. However, it is important to maximize the Coulombic efficiency to restrict Li and electrolyte depletion over multiple cycles. Sn@MGNS-A achieved 99% Coulombic efficiency after 13 cycles while Sn@MGNS compositions with higher Sn content took upwards of 27 cycles.

Figure 17 Coulombic efficiency of Sn@MGNS composites.

The voltage profiles over 100 cycles of Sn@MGNS-C was closely examined to investigate this capacity fade further (Figure 18). The primary loss in the unload capacity occurred between 0.4 and 0.9 V, seen as a shortening of the LixSny delithiation plateaus. Corresponding decreases in the LixSny lithiation plateaus in the load curves demonstrates that the degree of Sn lithiation decreases with increasing cycle number. The capacity contribution of Sn continuously decreased until the ~50th cycle, where little fade occurred afterwards. Even after 50 cycles, there was still a small plateau at ~0.6V vs Li indicating further electrochemical

95 activity of Sn. Unload capacity contribution below 0.4V vs Li did not change significantly. The majority of this capacity comes from Li deintercalation from

MGNS. Based on these observations, it appears the MGNS remained structurally and electrochemically intact over 100 cycles, while Sn became largely electronically or ionically inaccessible. The exact reason for this is unclear, but could be due to passivation of the Sn inside the Sn@MGNS particles by an excessively thick SEI. The

Columbic efficiency data is consistent with this explanation; it reaches its maximum during the cycle at which the capacity of the Sn@MGNS electrodes stabilizes, when little of the Sn remains electrochemically active.

Further evidence of excessive SEI formation comes from electrochemical impedence spectroscopy (EIS). EIS was collected on Sn@MGNS-C half-cells after the

1st, 10th, 20th, and 50th unload using a Solartron Analytical 1260 impedance/gain- phase analyzer and 1287A potentiostat/galvanostat. EIS was collected after the cells were unloaded and rested for 5 hours. An excitation voltage of 10mv was applied to the cells and the complex impedance measured as a function of frequency using a logarithmic sweep from 10 kHz to 5mHz. The data and fits of the data to an equivalent circuit are shown in a Nyquist plot (Figure 19). The mid-frequency semicircle (2nd from left) of these Nyquist plots is due to SEI impedance, represented by a RC circuit in our equivalent circuit. There little change in the resistance of the SEI between the 1st and 20th cycle only increasing from 11.13 to

15.62Ω, but it increased to 26.11 Ω between the 21st and 50± cycles. This SEI could block the electron pathways to the Sn making this portion of the material of the non- conductive, however, as the Sn is in intimate contact with the MGNS, one would also

96 expect decreased capacity from the MGNS, which is not seen in our experiments. It is possible that SEI builds up around the Sn inside the MGNS. Repeated cycles of expanision/contraction upon loading/unload the Sn could mechanically disrupt the local SEI, causing further SEI growth, and low Coulombic efficiency until it is electrically isolated from the MGNS. Alternatively, Sn may maintain electrical contact with the MGNS, but its lithiation may be kinetically inhibited by SEI buildup.

SEI has low Li+ ionic conductivity. Progressive SEI buildup around the Sn nanoparticles would increase Li+ loading impedance, and eventually the Li+ diffusion rate would be too slow to load at the same rate that Li+ intercalates the

MGNS. Regardless of whether the fade mechanism is electrical or ionic in nature, or some combination of the two, it is clear that the Sn inside the becomes electrochemically inaccessible after repeated cycling. If the alloying metal is not electrochemically accessible in the core, it acts as dead weight in the battery only decreasing the overall specific energy density of the anode.

Figure 18 Voltage profiles of Sn@MGNS-C, at C/10, at the 1st, 10th, 20th, 50th, and 100th cycle.

Figure 19 Nyquist plot of Sn@MGNS-C/Li halfcell. The EIS spectra was collected after the unload step of the 1st, 10th, 20th, and 50th cycle. The subwindow shows the equivalent circuit used to model the EIS spectra. This circuit includes the resistance due to cell components and leads (R cell), the resistance of SEI (R SEI), the capacitance of SEI interphase (C SEI), the resistance of charge transfer (R CT), the capacitance of the double layer (C DL), and a Warburg component for diffusion (W diff).

Table 2 - Fitted values of EIS data found in figure 19. 1st cycle 10th cycle 20th cycle 50th cycle Element Value Standard Value Standard Value Standard Value Standard deviation deviation deviation deviation Rcell 4.6 0.1 4.9 0.3 3.2 0.1 3.7 0.1 RSEI 11 1 13 1 16 1 26 1 CSEI - T 1.5 0.1 E-4 6.0 1.2 E-4 1.4 0.4 E-4 7.5 0.1 E-5 E-4 E-4 E-4 E-5 CSEI - P 0.86 0.01 0.40 0.02 0.83 0.01 0.91 0.01 RCT 9.9 0.4 8.6 0.4 12 1 16 1 CDL - T 2.1 0.4 E-4 1.3 0.1 E-4 1.3 0.1 E-4 2.7 0.2 E-4 E-4 E-4 E-4 E-4 CDL - P 0.54 0.02 0.91 0.01 0.53 0.01 0.47 0.01 Wdiff - R 4.7 0.6 18 1 19 1 20 1 Wdiff - T 0.45 0.09 3.3 0.2 3.5 0.2 3.8 0.2 Wdiff - P 0.33 0.01 0.37 0.01 0.37 0.01 0.34 0.01

3.5 Introduction: The SnO2/MGNS composite

MGNS has an electrical conductivity of ~10 S cm-2. Because of this, MGNS could be considered as a carbon to be combined with an alloying nanomaterial to form a composite to help buffer expansion, while preventing agglomeration. While the alloying nanoparticles can be simply mixed with the carbon, by going one step further and seeding the particle growth directly on the carbon, we can better disperse the nanoparticles. From this, agglomeration can be better prevented,[15] leading to increased cycling stability.[16] Nanoscale SnO2 particles can be synthesized by a variety of means,[17] but none are more simple than reducing aqueous SnCl2 with NaBH4,[18] following

Equation 3.

3. SnCl2(aq) + NaBH4(aq) + NaOH(aq) → Sn(s, nano) + 2NaCl(aq) +

B(OH)3(aq) + H2(g)

∆ 4. Sn(s, nano) + O2(aq) → SnO2(s, nano)

The Sn2+ is reduced to Sn(0) by borohydride reduction in the presence of suspended

MGNS, but because of dissolved O2 in the water, and heat produced the exothermic reaction, SnO2 is the final product. This borohydride reduction is a room temperature method that allows for a homogenous reduction of SnO2 on its carbon support with little additional clean up necessary.

3.6 Characterization: The SnO2/MGNS composite

The presence of SnO2 in the composite was verified by powder XRD. The 101 reflection (39.8° 2Ɵ) of SnO2 increased relative to that of the MGNS as the mass fraction of SnO2 increased in the composite (Figure 20). The 110 reflection (26.4° 2Ɵ) of SnO2 falls under the graphitic 002 reflection (26.2° 2Ɵ) of MGNS, so its intensity increase is

100 more difficult to discern. The relative mass of the components in the composite was determined by oxidation of the MGNS in air to CO2 using TGA, with the mass remaining after heating to 900 ˚C being SnO2 (Figure 21). SnO2/MGNS was synthesized with mass ratios of 20% to 50% SnO2, showing good control of the reaction based on the starting materials. TEM images show small SnO2 particles, no more than a few nm in size, decorating the MGNS (Figure 22). SAD confirmed that the material was SnO2

(Figure 23), in agreement with powder XRD results.

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Figure 20 a) Powder XRD patterns of SnO2/MGNS with the follow SnO2 mass percentages A) 0% B) 19% C) 31% D) 42% and E) 52%. Collected at 0.5° min-1. F) Powder diffraction files for MGNS (blue), SnO2 (red), and NiO (green).

102

Figure 21 TGA analysis of SnO2/MGNS at various mass compositions. Samples were heated in air at a flowrate of 20 °C min-1 to 900 °C and then held at that temperature for 0.5 h.

103

Figure 22 TEM micrograph of 42% SnO2/MGNS (75 kx magnification, 100 keV accelerating potential).

104

Figure 23 SAD of the area shown in Figure 23, of 43%SnO2/MGNS (Camera length 50 cm, 100 keV accelerating potential).

3.7 Results and Discussion: The SnO2/carbon composite

The reduction of SnCl2·2H2O was performed with an excess of NaBH4.

SnCl2·2H2O was the limiting reactant and the yield was 90% or better for all ratios of

SnO2 to MGNS. Reaction temperature proved to be important to obtaining nanoparticles of SnO2. When this reaction was performed at 0 °C the Sn particles formed did not oxidize to SnO2. Powder XRD of Sn/MGNS produced at low temperature shows Sn particles ~20 nm in size based on Scherrer’s equation (Figure 24). When the reaction temperature is not controlled, heat from reaction oxidizes the Sn to SnO2 (Equation 4) with far less particle growth. Sn nanoparticles may prove to perform better than SnO2, because there is no formation of Li2O on the first cycle, but the size of these particles is

105 so large that their expansion during lithiation might mechanically degrade the electrode, limiting the cycle life of the Sn/MGNS composite. When 30% Sn/MGNS was cycled against Li (Figure 25), the cell faded quickly to 150 mAh g-1, a capacity below that of the standard graphite electrode.

Figure 24 Powder XRD of 30% Sn/MGNS collected at 0.5° min-1.

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Figure 25 Reversible capacity of 30% Sn/MGNS. 80% Sn/MGNS, 10% Super P, and 10% PVDF (2 mg cm-2 electrode loading, cycled at C/10 between 0.02 – 1.5V vs Li).

SnO2/MGNS electrodes with SnO2 (20 to 50 % w/w) were cycled from 1.5V to

0.02V vs Li metal. Electrodes made with 2 different binder systems were investigated with compositions of 80% SnO2/MGNS, 10% Super P, and 10% Polyvinylidene fluoride

(PVDF) or 70% SnO2/MGNS, 10% Super P, 15% Carboxymethylcellulose (CMC), and 5%

Styrene-Butadiene Rubber (SBR). The active material loading was between 1 and

2 mg cm-2. Initial capacities for these electrodes are very close to the theoretical capacity of Sn based on the SnO2 loading (Figure 26). The 52% SnO2/MGNS (42% Sn by mass including binder and conductive additive), has an initial unload capacity of 490 mAh g-1.

A theoretical capacity of 403 mAh g-1 would be expected accounting only for the Sn alloying of Li+, but the MGNS also intercalates Li+. The MGNS, which constitutes 38%

107 of the electrode mass, could contribute ~90 mAh g-1 (based on MGNS obtaining

-1 200 mAh g ) to the composite capacity. Thus, the capacity of the 52% SnO2/MGNS material is very close to its theoretical capacity. Table 3 shows theoretical and experimental capacities for the SnO2/MGNS composites.

Figure 26 Reversible capacities of SnO2/MGNS composite – Li half-cells.

Table 3 – Capacity of SnO2/MGNS based on SnO2 content SnO2/MGNS SnO2 % w/w Theoretical Capacity Experimental Capacity (mAh g-1 composite) (mAh g-1 composite) 19 309 260 31 378 308 43 448 385 52 499 490

Although the SnO2/MGNS has a high initial capacity, it fades within 50 cycles to

-1 ~250 mAh g , regardless of SnO2 loading. The voltage profile of 31% SnO2/MGNS

108

(Figure 27a) shows a change in the discharge curve over the 50 cycles. This can be better seen in the differential capacity of these electrodes, (Figure 27b). The unload curves show the development of additional cathodic peaks after 20 cycles at 0.46V, 0.62V, and

0.74V and the disappearance of broad peaks at 0.6V and 1.21V on the first few cycle. In the early cycles the SnO2 is partially reversible as the Sn and Li2O reform SnO2 at 1.21V.

The three new peaks on the unload align with the formation of Li7Sn3, LiSn, and Li2Sn5.

These new phases arise throughout cycling due to formation of crystalline larger Sn particles caused by the particle growth as a result of agglomeration. The higher mass loadings of SnO2 fade at a quicker rate, due to a greater rate of agglomeration, eventually leading to mechanical degradation of the electrode. Lower loadings of SnO2, such as the

19% SnO2/MGNS, show little fade. The initial unloaded capacity of 19% SnO2/MGNS is 260 mAh g-1 and only fades ~20 mAh g-1 over 50 cycles. The stability of these electrodes are better than the higher loadings, but there is very little overall capacity gain over MGNS without SnO2.

109

Figure 31 a) Voltage profiles of 31% SnO2/MGNS cycled at a C/10 rate. b) Differential capacity curves of 31%SnO2/MGNS calculated from the voltage profiles shown in Figure 28A

110

Switching the binder in these SnO2/MGNS electrodes to CMC/SBR greatly increases the cycle life greatly. Figure 28 shows that cycle life of the 31% SnO2/MGNS electrodes is tripled, reaching 160 cycles before fading to 250 mAh g-1. Similar to the electrodes made with PVDF, the higher mass ratios fade more rapidly. Even so, their cycle life is much better than seen with PVDF binder; the 43% and 52% SnO2/MGNS reach ~80 cycles before fading to 250 mAh g-1. While the elasticity of the CMC/SBR binder system in considerably less than PVDF,[19] the cycling performance is dramatically better. This may seem counterintuitive because the dramatic volume changes of the SnO2 particles would be more likely to pulverize the more brittle

CMC/SBR binder. Similar cycle life enhancement were observed for Si anodes, for which it was proposed that the increase in performance is a result of a chemical compatibility between the metal oxide particles and carboxylic acid groups of the

CMC.[20] Terminal OH groups on the surface of the SnO2 and the carboxylic acid may undergo a condensation reaction, resulting in “esterlike” covalent bonds. These in turn keep the SnO2 particles in place within the MGNS matrix, slowing agglomeration.

Looking closely at the voltage profiles of the 31%SnO2/MGNS electrodes prepared with

CMC (Figure 29), it is clear that despite the change in binder, the SnO2 follows the same pattern of partial reversibility of SnO2 in early cycles and Sn(0) particle formation in later cycles, presumably due to irreversible phase separation from Li2O.

111

Figure 28 Reversible capacities as a function of cycle number for various SnO2/MGNS compositions using CMC/SBR binder cycled at a C/10 rate.

Figure 29 Voltage profiles of 31% SnO2/MGNS anodes using CMC/SBR binder cycled at a C/10 rate for the indicated number of cycles.

112

First cycle Coulombic efficiencies of these SnO2/MGNS electrodes are ~45%

(Figure 30) rising to ~90% for the second cycle. By the tenth cycle the efficiencies of these cells are all above 99%. The low first cycle Coulombic efficiency of these cells can be attributed the partially irreversible formation of Li2O from SnO2 and formation of SEI.

According to Equation 13 from Chapter 1.3.2, upwards of 8.25 mol of Li is consumed as

Li2O. Assuming 12.5 mol of Li is consumed in the first cycle (8.25 mol of Li to Li2O and

4.25 mol of Li to Li17Sn5) and only 4.25 mol of Li is reversible, we expect a Coulombic efficiency of 34% as we get 1 e- transferred for each Li+ oxidized. This greater

Coulombic efficiency than expected is consistent with partial reversibility of the Li2O to

SnO2 reaction. Additionally, SEI formation as a result of electrolyte reduction also contributes to the first cycle Coulombic losses. While we do not know the exact contribution, we would expect a large portion of this low first cycle Coulombic efficiency to be a result of the SEI due to the high surface area of the material that must be passivated. MGNS electrodes without SnO2 have a Coulombic efficiency of ~35%. Thus, if the surface of the MGNS is still equally accessible by the electrolyte, this SEI likely still forms. Thus, in agreement with the voltage profiles, the Li2O to SnO2 reaction is largely reversible, becoming less so with increasing cycles, probably as a result of Sn aggregation into larger particles.

113

Figure 30 Coulombic efficiencies of various SnO2/MGNS compositions.

The rate performance of these SnO2/MGNS electrodes was also tested. Figure 31 shows the change in capacity of these electrodes, as a percentage of their first cycle capacities, as the rate of load and unload of Li+ is increased. The decrease in cell capacities is independent of SnO2 concentration as the cycling rate is increased. At 2C, the SnO2/MGNS electrodes have 45% of their C/10capacity. At 10C, they have less than

5% their original capacity, while at 20C essentially no reversible capacity is observed.

After cycling at 20C, the cells were returned to a rate of C/5 where the cells recovered their initial capacity, demonstrating that the decrease in capacity with increasing rate is a result of kinetic limitations and is not because of capacity fade.

114

Rate

Figure 31 Rate performance of SnO2/MGNS compositions as a percentage of their first cycle reversible capacity.

While the SnO2/MGNS electrodes discussed thus far were cycled with a 1 M

LiPF6 in EC:DMC (v:v) electrolyte, there are other electrolyte systems that could prove beneficial. Propylene carbonate (PC) is a solvent that has been extensively studied for

Li-ion batteries.[21] It is a less expensive solvent that has a lower freezing point (-

50.3°C) than most other solvents employed for Li-ion battery electrolytes (-7°C for 1:1

EC:DMC). PC was ultimately abandoned because it co-intercalates with the Li+ ions into graphite anodes.[21] This leads to exfoliation of the graphite and rapid cell failure, with no reversible capacity for even a single cycle. It was previous shown that PC is compatible with MGNS and no exfoliation takes place[10]. With this in mind,

SnO2/MGNS was cycled with 1M LiPF6 in PC. Figure 32 shows little difference in the capacity fade of 43% SnO2/MGNS cycled in PC or EC:DMC.

115

Figure 32 Reversible capacity of 41% SnO2/MGNS, cycled at a C/10 rate in PC (circles) and EC:DMC (squares).

3.8 Conclusions: The SnO2/carbon composite

Sn@MGNS was produced using a modified MGNS synthesis. Both Ni and Sn salts were mixed with cellulose, and pyrolized, to make mixed metal MGNS. HNO3 was used to dissolve Ni and oxidize Sn to SnO2, which was subsequently reduced with hot H2 gas. A maximum Sn content of 37.2% w/w was achieved for Sn@MGNS. The resulting product was used to make anodes for Li half cell batteries. Initial capacities of 307 mAh g-1 were achieved but faded to 131 mAh g-1 after 50 cycles. The exact cause of capacity fade is unclear, passivation of the Sn inside the MGNS by an excessively thick

SEI appears to be the primary reason for this steady fade.

116

A composite of nanoscale SnO2 on the exterior of MGNS was formed by reduction of SnCl2 with NaBH4. Various compositions of SnO2/MGNS were synthesized with upwards of 52% w/w SnO2. The SnO2/MGNS was tested as anode materials for

Li-ion batteries. The SnO2 improved the capacity and Coulombic efficiency of MGNS significantly. While large loadings of SnO2 produced initial capacities as high as 550 mAh g-1 composite, these electrodes rapidly faded over multiple cycles due to agglomeration and pulverization of the alloying nanoparticles. SnO2/MGNS (31% w/w

SnO2) anodes prepared with CMC:SBR binder showed little fade out to 160 cycles. The

SnO2/MGNS electrodes maintained 45% of their low rate capacity at rates as high as 2C, and proved to be compatible in PC based electrolyte.

Combining the two Sn-MGNS composite options, inside (Sn@MGNS) and outide

(SnO2/MGNS) deposition, appears to be of little value due to the poor capacity retention of the Sn@MGNS. However, with further optimization, SnO2/MGNS by itself could prove to be a useful material for advanced Li-Ion anodes.

3.9 References

1. Winter, M., P. Novak, and A. Monnier, Graphites for lithium-ion cells: The correlation of the first-cycle charge loss with the Brunauer-Emmett-Teller surface area. Journal of the Electrochemical Society, 1998. 145(2): p. 428-436. 2. Zinth, V., et al., Lithium plating in lithium-ion batteries at sub-ambient temperatures investigated by in situ neutron diffraction. Journal of Power Sources. 271: p. 152-159. 3. Steiger, J., D. Kramer, and R. Monig, Mechanisms of dendritic growth investigated by in situ light microscopy during electrodeposition and dissolution of lithium. Journal of Power Sources. 261: p. 112-119. 4. Cui, G.L., et al., A one-step approach towards carbon-encapsulated hollow tin nanoparticles and their application in lithium batteries. Small, 2007. 3(12): p. 2066-2069. 5. Liu, N., et al., A pomegranate-inspired nanoscale design for large-volume-change lithium battery anodes. Nat Nano, 2014. 9(3): p. 187-192.

117

6. Lee, K.T., Y.S. Jung, and S.M. Oh, Synthesis of tin-encapsulated spherical hollow carbon for anode material in lithium secondary batteries. Journal of the American Chemical Society, 2003. 125(19): p. 5652-5653. 7. Zhang, W.M., et al., Tin-nanoparticles encapsulated in elastic hollow carbon spheres for high-performance anode material in lithium-ion batteries. Advanced Materials, 2008. 20(6): p. 1160-+. 8. Liu, N., et al., A Yolk-Shell Design for Stabilized and Scalable Li-Ion Battery Alloy Anodes. Nano Letters, 2012. 12(6): p. 3315-3321. 9. Herring, A.M., et al., A novel method for the templated synthesis of homogeneous samples of hollow carbon nanospheres from cellulose chars. Journal of the American Chemical Society, 2003. 125(33): p. 9916-9917. 10. Cox, J.D., Electrochemical study of hollow carbon nanospheres as high-rate and low temperature negative electrodes for lithium ion batteries. 2015, The George Washington Univeristy, Ph.D. 11. Rohm and Hass, Sodium Borohydride Digest. 2003. 12. Datsyuk, V., et al., Chemical oxidation of multiwalled carbon nanotubes. Carbon, 2008. 46(6): p. 833-840. 13. Kim, B.S., et al., Reduction of SnO2 with Hydrogen. Materials Transactions, 2011. 52(9): p. 1814-1817. 14. Levi, M.D. and D. Aurbach, Simultaneous measurements and modeling of the electrochemical impedance and the cyclic voltammetric characteristics of graphite electrodes doped with lithium. Journal of Physical Chemistry B, 1997. 101(23): p. 4630-4640. 15. Du, Z.F., et al., In situ synthesis of SnO2/graphene nanocomposite and their application as anode material for lithium ion battery. Materials Letters, 2010. 64(19): p. 2076-2079. 16. Yuan, L., et al., Nano-structured SnO2-carbon composites obtained by in situ spray pyrolysis method as anodes in lithium batteries. Journal of Power Sources, 2005. 146(1-2): p. 180-184. 17. Chen, J.S. and X.W. Lou, SnO2-Based Nanomaterials: Synthesis and Application in Lithium-Ion Batteries. Small, 2013. 9(11): p. 1877-1893. 18. Markova-Deneva, I., T. Petrov, and I. Denev, Template Synthesis Of Cu And Cu- Sn Nanoparticles Using Carbon Foam As a Support, in METAL 2012 - 21st International Conference on Metallurgy and Materials. 2012: Czech Republic. p. 1-5. 19. Li, J., R.B. Lewis, and J.R. Dahn, Sodium Carboxymethyl Cellulose: A Potential Binder for Si Negative Electrodes for Li-Ion Batteries. Electrochemical and Solid- State Letters, 2007. 10(2): p. A17-A20. 20. Hochgatterer, N.S., et al., Silicon/Graphite Composite Electrodes for High- Capacity Anodes: Influence of Binder Chemistry on Cycling Stability. Electrochemical and Solid-State Letters, 2008. 11(5): p. A76-A80. 21. Xu, K., Nonaqueous liquid electrolytes for lithium-based rechargeable batteries. Chemical Reviews, 2004. 104(10): p. 4303-4417.

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4. High Performance Arsenic – Multiwall Carbon Nanotube Composite Anodes for Li-ion Batteries

For submition to the Journal of Power Sources* 4.1 Introduction

Lithium ion batteries have set the standard in the portable electronics market, keeping our phones, tablets and laptops powered over the last two decades.

While these batteries have thus far met consumer needs for portable electronics, these needs are ever growing and emerging markets, such as electric vehicles (EVs) would benefit greatly from improved batteries. Current consumer EVs generally have sufficient energy storage for local use, however, their short driving range puts them at a considerable disadvantage to automobiles driven by internal combustion engines.

In 2012 the Department of Energy set a goal of increasing the energy density of batteries to 400 Wh L-1 while keeping the cost down to $125/kWh.[1] The United

States Advanced Battery Consortium (USABC), a collaborative organization between major automotive companies Fiat Chrysler Automobiles, Ford Motor Company,

General Motors, has established a goal to commercialize a 750Wh L-1 battery by

2020.[2] In order to reach these values, new, inexpensive materials must be developed that have sufficient volumetric capacity and conductivity for high areal loading electrodes.

Graphite, the current standard for Li-ion anodes, is limited by a low specific and volumetric capacity (372 mAh g-1 and 833 mAh cm-3).[3] While increasing the

119 thickness of the anode coating can increase areal loading, it does not address the problems of low gravimetric and volumetric capacity and results in an unacceptably high increase in cell impedance.[4, 5] In addition, graphite exhibits a build up of the solid electrolyte interface (SEI) due to surface reactions with electrolyte that leads to capacity fade[6] and short cycle life.[7, 8] Li plating on the graphite anode is of also of concern because of safety issues due to possible dendrite formation and consequent cell shorting.[9] This phenomena is most prevalent in graphitic anodes because the intercalation potential is near the Li deposition potential. High charge rates and low temperatures lead to polarized conditions where Li intercalation kinetics cannot keep up with the reduction of the Li+ ions at the anode.[10] While this plating can be reversible if it is of insufficient quantity to cause a short circuit, it is still a significant issue as it increases SEI formation and thus capacity fade.[11]

The shortcomings of graphite have lead to great interest in higher capacity alternatives, with most of this effort put toward developing a viable alloying metal anode. Considerable progress has been made on Li-alloying anode materials such as

Si and Sn, which have high specific capacities of 4200 and 960 mAh g-1 respectively.[2, 12-15] It is now well understood that the extraordinary capacities these materials achieve comes at the cost of large volumetric expansions upon Li uptake.[16] Mechanical degradation then occurs upon agglomeration of these particles after multiple cycles of the material.[17] A vast array of strategies has been explored to mitigate the detrimental effect of volumetric expansion.[14, 18, 19]

Mixing the active alloying metals in carbon matrices is a well studied means of limiting capacity fade due to expansion and agglomeration,[20-25] while

120 additionally increasing the conductivity of the material.[26, 27] Although this method has shown some success in increasing specific capacity over many of cycles, long cycle life with anodes with the high mass ratio of the alloying metal necessary to achieve substantially increase volumetric energy density is an outstanding challenge.

While carbon based anode materials form a stable SEI within the first few cycles,[6] forming a stable SEI layer on alloying anode materials is more challenging.[28] The SEI on an alloying anode material cracks during the large volumetric changes that take place during lithiation/delithiation, exposing the anode surface to electrolyte. Passivation of this newly exposed surface by the formation of additional SEI increases cell impedance with every cycle. Cycling stability of both carbonaceous[29] and alloying anode materials[30] can be increased by using electrolyte additives such as vinlyne carbonate (VC) and fluoroethlylene carbonate (FEC), better controlling the growth of the SEI.

The pnictides (P, As, Sb, and Bi) have garnered much attention because of high theoretical capacities based on the formation of Li3X (X = P, As, Sb, or Bi). P,[27,

31-34] Sb,[35, 36] and even Bi[37] have been studied as anode materials for reversible Li or Na cells, but As has largely been ignored. Alloying of Li with As was first demonstrated by Besenhard et. al. in 1975.[38] Only recently,[39] 40 years after Besenhard’s initial report, has another study of elemental As based Li-ion anodes appeared. The absence of As anode research is presumably because of environmental concerns. However, each year thousands of metric tons of As are used industrially in wood preservatives, agricultural chemicals, lead acid batteries,

121 and semiconductors in the United States.[40, 41] While As is toxic, its EPA drinking water maximum contaminant level is 10 µg L-1, nearly the same as Pb (15 µg L-

1),[42] which is present in large quantities in the ubiquitous Pb-acid motor vehicles battery. Effective recycling, as is the case for Pb-acid batteries, could make the use of As based batteries environmentally acceptable.

The theoretical reversible specific capacity of As, 1073 mAh g-1 is the fourth largest of any Li alloying element, trailing only Ge, P, and Si. Its volumetric capacity,

2057 mAh L-1, is also among the highest, trailing additionally Au. Unlike Ge and Au,

As is inexpensive with a bulk cost that is approximately the same as Si.[12] Being a semi-metal, its electrical conductivity (3.3x106 S m-1,) is 1000 times greater than

Si,[43] and is far more air stable. Here we report on its outstanding performance as a Li alloying anode material when combined with multiwalled carbon nanotubes, rivaling that of other alloying elements and making As a potential alternative for inexpensive, high capacity Li-ion anodes.

4.2 Experimental

Preparation of Nanoparticulate As: AsCl3 (99.999%, Alfa Aesar) was used without further purification. Tetrahydrafuran (THF, Fisher Scientific, 99.9+% HPLC grade, inhibitor-free) was purified by stirring over NaK alloy until a persistent blue solution was obtained, then vacuum transferred to storage bottles. Solution preparation and alkalide reduction[44] were performed in open vessels at room temperature in an N2 glovebox (< 1 ppm H2O and O2). K+(15-crown-5)2Na- solution was prepared by stirring NaK alloy (0.376 g) and 15-crown-5 (4.41g, 98%, Alfa

Aesar) in 30 ml of THF. After complexation, the akalide solution was swiftly added

122 to a rapidly stirring solution of 0.484 AsCl3 in 20ml of THF. A dark brown precipitate of As nanoparticles (Asnano) formed immediately. Multiwalled carbon nanotubes (MWCNT, 99+%, >20 nm diameter, 1-12 µm length, Cheap Tubes Inc.) were added to the AsCl3 solution with vigorous stirring prior to addition of alkalide solution to affect direct reduction onto the tubes (Asnano/MWCNT).

Following reduction, methanol (80 ml, 99.9%, Fisher Scientific, dried on 3Å sieves) was added to increase the rate of precipitation. The product was separated from the solution by centrifuge (5min, 8300 x g). The product was washed 3 times with N2 purged methanol in a N2 purged glove bag, each time followed by separation of the product from the wash by centrifuge. Finally, the samples were dried under vacuum and stored under N2. Final yield, based on As content, was typically ~ 73.%, with the vast majority of the losses coming during washing of the products of the small scale synthesis.

4.3 Physical Characterization

Powder x-ray diffraction (XRD) patterns were obtained with a Rikagu

Samples were examined using a FEI Talos 200X TEM with high-angle, annular dark field (HAADF) scanning TEM and Super X electron dispersive spectroscopy (EDS). EDS mapping was analyzed by Bruker ESPRIT ™ software.

123

Additionally, electrodes were examined on a FEI Teneo LV SEM with In-lens secondary electron detector.

Elemental analysis to determine As content was obtained by X-ray fluorescence (Kratos Analytical EDX-700). Known amounts of SnO2 were added to the samples to qualitatively determine ratio of As and Sn, yielding data that was in turn used to calculate the As mass fraction in each sample.

4.4 Electrochemical Characterization

Electrodes were prepared by a slurry and doctor blade method. Three types of As electrode active materials were used in this study; Asnano mixed with MWCNTs,

Asnano directly reduced on MWCNTs (Asnano/MWCNT) and bulk As (metal lump, 2 –

6 mm, 99.9999%, Alfa Aesar) mixed with MWCNT’s. The bulk As was hand ground with mortar and pestle into a fine powder (Asmicro) prior to mixing with MWCNTs.

Dispersion of the Asnano and Asmicro amongst the MWCNTs was accomplished by combining As and MWCNTs and adding an additional 15% by mass polyacrylic acid

(PAA, MW ~1 x 106, Polysciences, Inc.) followed by milling for 1 minute in a Wig-L-

Bug® grinding mill (Reflex Analyical Corp.). N-methyl-2-pyrrolidone (NMP, 99+%,

Sigma Aldrich, dried on 4Å sieves) was then added to produce a slurry. This slurry was homogenized using an ultrasonic bath for 15 minutes and applied to pre- weighed carbon coated copper current collectors (1 µm conductive carbon on 9 µm copper foil, MTI Corp.). The electrodes were heated at 150°C under active vacuum

(10-3 torr) for 1 hour and then reweighed to obtain electrode mass loading. Active material (As) loading was 1.5 to 2 mg cm-2.

124

All As anode testing was performed in a half-cell configuration with a Li foil

(99.9%, MTI Corp.) counter electrode in 2016 coin cells (MTI Corp.) with Celgard®

2325 separator (Celgard LLC) and 1M LiPF6 (99.99%, Novolyte) in 1:1 FEC:DEC

(V:V, >99%, Solvay) or 1:1 EC:DMC (V:V, 99.9%, Novolyte) electrolyte. The As half- cell battery performance was tested using Arbin BT2000 cyclers. The cells were cycled under constant current from 0.02 to 2.0V vs Li at a C/10 rate, unless otherwise noted. They were rested for 1 hour between each charge and discharge.

Impedance testing was performed using a Solartron Analytical 1260 impedance/gain-phase analyzer with 1287A potentiostat/galvanostat. Electrodes for EIS testing were made with a 80:10:10 (As:MWCNT:PAA) by mass composition ratio. EIS was collected after the cells were unloaded and rested for 5 hours. An excitation voltage of 10mv was applied to the cells and the complex impedance measured as a function of frequency using a logarithmic sweep from 10 kHz to

5mHz.

4.5 Results and Discussion

The MWCNTs were examined by powder XRD (Figure 1). The presence of a sharp peak at d = 3.36 Å with a broad shoulder at higher d-spacing is indicative of the presence of graphite in addition to the MWCNTs in the material as received.

Deconvolution using Pearson VII peak shapes (Jade 6.1, Materials Data Inc.) resulted in an estimate of the graphite content of 21.6% (Figure S1). The presence of such a significant fraction of graphite in the “99% pure” MWCNTs was confirmed by SEM observation (Figure S2). Further confirmation of the presence of graphite can be found in the charge/discharge curves electrodes consisting of 90% MWCNTs and

125

10% PAA cycled in half cell configuration vs lithium metal which shows clear plateaus in the discharge curves, characteristic of graphite (Figure S3).

Asmicro

Asnano

Asnano/MWCNTs

MWCNTs

10 20 30 40 50 60

2q Figure 1. Powder x-ray diffractograms of Asmicro, Asnano, Asnano reduced on MWCNTs, and MWCNTs. The asterisk indicates the (002) reflection of graphite.

Following washing, Asnano synthesized by akalide reduction slowly oxidizes in air, so exposure time was minimized. The powder XRD pattern (Figure 1) shows three broad humps at 15°, 31°, and 53° 2θ, that are similar to those seen for amorphous red phosphorous[27, 33] and suggest that the Asnano is either

126 amorphous or nanocrystalline. Examination of the material by TEM showed that it consists of irregularly shaped particles with diameters in the range of 10 – 20 nm

(Figure 2). Selected area diffraction (SAD) shows 3 broad polymorphous rings at diffraction angles consistent with the powder XRD results, indicating that the Asnano is amorphous. Electron dispersive spectroscopy (EDS) (Figure S4) confirmed the presence of As along with a small amount of residual K and Cl, presumably trace byproducts from the synthesis that washing failed to remove. In contrast, Asmicro is a crystalline (Figure 1) material with a grain size of ~ 1 to 5 µm (Figure S5).

Figure 2. TEM image of Asnano synthesized by akalide reduction. Inset: SAD pattern.

Asnano and Asmicro were dispersed amongst the MWCNTs in mass ratios of

79:11:10 and 80:10:10 As:MWCNT:PAA, respectively, and made into electrodes.

Direct comparison of the voltage profiles of Asnano with Asmicro (Figure 3) showed similar Li+ loading (analogous to the charge process in a full cell) curves, but the nanoparticulate form displays vastly improved reversibility. In addition, while

127

Asmicro loading displays a distinct plateau at ~ 0.8 V, Asnano loading occurs at higher potentials over a gently sloping plateau. The lack of a distinct plateau in the loading profile is consistent with the Asnano being amorphous.

2.5

As Micro AsNano

2.0

e m

1.5

e g

a 1.0

o V

0.5

0.0 0 400 800 1200 1600 -1 Capacity/mAh g As

Figure 3 Load and unload voltage profiles of anodes of MWCNT mixed with Asmicro (dashed lines) and Asnano (solid lines).

Li-As alloying of the Asnano begins at ~ 1.1 V and continues with the voltage profile gently sloping between ~ 0.97 and 0.8 V. Loading capacity obtained above

1.1V can be attributed to reduction of the carbonate electrolyte associated with SEI formation,[45] but the majority of first cycle losses occurred at potentials lower than ~0.7 V due to SEI formation on the MWCNTs.[23] Evidence of lithiation of the

MWCNTs/graphite can be seen in the profile below 0.3V vs Li.

The first unload (analogous to the discharge process in a full cell) of Asnano occurs at significantly lower potential than Asmicro. A small fraction of its capacity at

128 potentials less than 1.0 V is due to MWCNT/graphite capacity, the majority of its

1076 mAhg-1 capacity coming from the dealloying reaction of Li and As over a gently sloping profile from ~ 1 – 1.15 V, a plateau at ~ 1.2V and a steeper unload thereafter. The capacity is somewhat larger than the theoretical capacity of As

(1073 mAhg-1) due to MWCNT/graphite storage (~80 – 100 mAh g-1).

The differences in the lithiation of Asnano and Asmicro were further investigated by powder XRD. Lithiated electrodes were disassembled and their diffraction patterns collected in inert atmosphere. Diffraction patterns show that while Asmicro lithiated to form the Li3As phase, Asnano remains amorphous with no detectable crystalline Li – As phase formed.

The cycling performance of Asnano is highly dependent upon electrolyte composition. When cycled in EC:DMC, the electrode had a high initial reversible capacity of 810 mAh g-1, but capacity faded to below 100 mAh g-1 within 40 cycles

(Figure 4). The Coulombic efficiency, initially 65%, does not increase beyond 95%

(Figure 5), indicative of significant continuous SEI formation. In addition, the voltage profiles of these cells show an increase in internal impedance with cycle number, indicated by a dramatic increase in the polarization in the charge and discharge curves over the first 20 cycles (Figure S6).

129

1000

800

h 600

t i

c 400

p a C 200

0 0 50 100 150 200 Cycle number

Figure 4. Cycling performance in 1M LiPF6 in 1:1 FEC:DEC w:w of Asnano mixed with MWCNTs (crosses, 57:30:13 As:MWCNT:PAA w:w) loaded to 0.02V, Asnano reduced on MWNCTs (triangles, 70:17:13 As:MWCNT:PAA w:w), Asnano mixed with MWCNTs (squares, 57:30:13 As:MWCNT:PAA w:w) loaded to 0.75V, and Asnano mixed with MWCNTs (circles, 57:30:13 As:MWCNT:PAA w:w) loaded to 0.80V. Cycling performance of Asnano mixed with MWCNTs (saltire crosses, 57:30:13 As:MWCNT:PAA w:w) loaded to 0.02V, in 1M LiPF6 in 1:1 EC:DMC (w:w).

100

y c

n 90 100

c i

f 98

c 80 96

i b

m 94

o l

u 70

o 92 C 90 60 5 10 15 20

0 50 100 150 200 Cycle number

Figure 5. Dependence of the Coulombic efficiency in 1:1 FEC:DEC w:w of Asnano mixed with MWCNTs (crosses, 57:30:13 As:MWCNT:PAA w:w) loaded to 0.02V, Asnano reduced on MWNCTs (triangles, 70:17:13 As:MWCNT:PAA w:w), Asnano mixed with MWCNTs (squares, 57:30:13 As:MWCNT:PAA w:w) loaded to 0.75V, and Asnano mixed with MWCNTs (circles, 57:30:13 As:MWCNT:PAA w:w) loaded to 0.80V. Cycling performance of Asnano mixed with MWCNTs (saltire crosses, 57:30:13 As:MWCNT:PAA w:w) loaded to 0.02V, in 1M LiPF6 in 1:1 EC:DMC (w:w).

Using an FEC:DEC based electrolyte markedly improved cycling performance.

Asnano electrodes cycled in 1M LiPF6 FEC:DEC, at C/10, showed an initial reversible

130 capacity of 920 mAh g-1 that only faded 35% over 95 cycles (Figure 4). Coulombic efficiency on the first cycle was 60%, increasing rapidly to 98% over the first 10 cycles (Figure 5). Dramatic improvements in performance observed in other alloying anode electrodes cycled in FEC has been attributed to the building of a more stable SEI.[30] Examination of the voltage profiles (Figure 6) shows that, while slowed considerably, the polarization still increased from 0.4V to 0.6V between the

1st and 60th cycle. The improvement in Coulombic efficiency and polarization when cycled in FEC is consistent with a decrease in, but still significant, SEI buildup.

60th 30th 1st 2.0

1.5

s 1.0

o V 0.5

0.0 0 400 800 1200 1600

-1 Capacity/mAh g As

Figure 6. Voltage profiles of 1st, 30th, and 60th cycles of 57:30:13 Asnano:MWCNT:PAA (w:w)halfcells in 1M LiPF6 in 1:1 FEC:DEC w:w loaded to 0.02V.

Better electrochemical performance can be achieved by limiting the loading cutoff voltage to 0.75 V vs Li. By doing so, irreversible losses due to SEI buildup on

131 the MWCNT can largely be avoided, with minimal impact on Li-As alloying as it almost exclusively occurs at higher potentials. Limiting the loading potential results in stable cycling of Asnano mixed with MWCNTs, yielding an initial reversible capacity of 594 mAh g-1 with negligible fade to 50 cycles, slowly fading over subsequent cycles to 444 mAh g-1 after 200 cycles at a C/10 rate (Figure 4). Additionally, the

Coulombic efficiency rises from 76% on the first cycle to over 99% after 15 cycles

(Figure 5). In contrast to Asnano electrodes loaded down to 0.02V, there is very little increase in the load/unload hysteresis over 200 cycles (Figure 7).

200th 100th50th1st 2.0

1.8

l a

t 1.6

i 1.4

e 1.2

t l

o 1.0 V 0.8

0.6 0 200 400 600 800 -1 Capacity/mAh g As

Figure 7. Voltage profiles of 57:30:13 Asnano:MWCNT:PAA half cells loaded to 0.75V vs Li. The first cycle load and unload is presented as a dashed line for clarity.

The stability of the SEI layer imparted by using a 0.75 V rather than a 0.02 V cutoff is further demonstrated by EIS measurements. Spectra were collected after the 1st, 10th, and 20th cycles with a 0.02 V (Figure 8a) and 0.75 V (Figure 8b) load cutoff, and fit with an equivalent circuit model (Figure 8c). The model includes elements for the resistance in electrolyte and cell components (Rcell), SEI resistance

132

(RSEI), charge-transfer resistance (RCT), constant phase elements for the SEI capacitance(CSEI), the double layer capacitance (CDL), and a Warburg element for diffusion (Wdiff). The constant phase elements and the Warburg element defined by equations (1) and (2) respectively.

Z = (1/[Q0ωn])e-([π/2]n*i) (1)

Z = R * ctnh([ I*T*ω]P)/(I*T*ω)P (2)

Fits of the impedance data to this model shows that the resistance due to SEI dramatically increases from 3.6Ω on the 1st cycle, to 74Ω on the 10th, and to

1.2 x 10-2 Ω on the 20th cycle for electrodes loaded down to 0.02V. In contrast, RSEI is essentially constant when loading is cutoff at 0.75 V, less than half as large in magnitude as that found with a 0.02 V cutoff after 20 cycles. A complete listing of the fitted values can be found in Table S1.

133

300

a 250

200 "

Z 150

100

0 0 100 200 300 400

140 b

120

100

" Z

0 0 20 40 60 80 100 120 140

Figure 8. Electrochemical impedance spectra of Asnano half cells, on the 1st (circles), 10th (squares), and 20th (triangles) cycles, loaded to a) 0.02V vs Li and b) to 0.75V vs Li. Solid lines are fits of the data to the equivalent circuit model (c).

134

Cycling losses can be further reduced by an additional 5 mV increase of the loading cutoff voltage to 0.80 V, albeit at a loss of initial capacity from 594 to ~ 512 mAh g-1 (Figure 4). While capacity fade over 200 cycles are approximately 25% when loaded to 0.75 V, less than 18% is lost when the cutoff is raised to 0.80 V, the reversible capacities becoming nearly identical after 200 cycles.

In an attempt optimize the active material loading and capacity further while decreasing impedance,[46] the Asnano was reduced directly on the nanotubes. Up to

80% weight loading was added by akalide reduction of AsCl3 in the presence of suspended MWCNTs. The high loading of As on the MWCNTs largely obscured the

MWCNT/graphite peaks in the powder XRD pattern, with only a small peak at 26.2°

2θ discernable (Figure 1). Reducing the Asnano in the presence of MWCNT dramatically changes the morphology and distribution of the As nanoparticles interspersed amongst the nanotubes. Mixing the Asnano results in a broad distribution of As agglomerate sizes, up to several hundred nm in diameter and a highly inhomogeneous distribution of them among the MWCNTs (Figure 9). In contrast, Asnano made by reduction in the presence of MWCNTs (Asnano/MWCNT) appears as much smaller agglomerates that are far more evenly distributed amongst and coating the MWCNTs (Figure 10). It is also clear that the Asnano are in much better contact with the MWCNTs in the Asnano/MWCNT material; the Asnano agglomerates show significantly less charging in SEM images than seen with the mixed material.

135

Figure 9. SEM image of Asnano mixed with MWCNTs. Agglomerates of Asnano are bright areas as a result of sample charging due to poor electrical contact.

Figure 10. SEM image of Asnano reduced on MWCNTs. Note the lack of charging due to better electrical contact.

Further examination of Asnano/MWCNT with TEM (Figure 11a), shows that it consists of ~ 10 – 20 nm As nanoparticle agglomerates that adhere to and coat the

MWCNTs. EDS mapping confirmed the location and composition of the

Asnano/MWCNT composite (Figure 11b-e). Oxygen is present throughout the

Asnano/MWCNT sample: on the Asnano, MWCNTs and carbon-coated TEM grid. Line scans (Figure S7) show that the oxygen content does not change dramatically from the background to the As particles; there with no evidence of an oxide shell.

136

Furthermore, crystalline As2O3 cannot be detected in the powder XRD or selected area diffraction (SAD) patterns of the Asnano/MWCNT. Thus, while it seems highly likely that the Asnano, having been exposed to air, contain some oxide content or surface bound species (such as CO2), it is clear that the content is low.

Figure 11. a) TEM image of Asnano/MWCNT composite synthesized by akalide reduction. b) HAADF STEM image of Asnano/MWCNT composite synthesized by akalide reduction. EDS maps of c) As (green), d) O (blue), and e) C (red).

The first cycle Coulombic efficiency and initial reversible capacity of

Asnano/MWCNT electrodes is significantly higher than those of Asnano mixed with

MWCNTs, increasing from 76% and 594 mAh g-1 to 83% and 746 mAh g-1 when discharged to 0.75 V vs Li (Figure 12). The better Coulombic efficiency despite the higher assessable Asnano surface area may be a result of better mechanical stability allowing less initial loss of As capacity due to electrical isolation caused by the large expansion during lithiation, more complete coverage around the MWCNTs by the

Asnano, passivating functional groups on the MWCNTs, or the reduction of those

137 functional groups by alkalide solution during synthesis. Reversible capacity retention is also greatly improved, retaining more than 620 mAh g-1 over 200 cycles

(17% loss) when cycled to 0.75V (Figure 4). The dip in capacity that appears after

140 cycles is due to power fluctuations in the laboratory; the cell fully recovered by the 160th cycle.

200th100th10th1st 2.0

1.8

l a

t 1.6

L 1.4

e 1.2

t l

o 1.0 V 0.8

0.6 0 200 400 600 800 1000 -1 Capacity/mAh g

Figure 12. Voltage profiles of Asnano/MWCNT composite loaded to 0.75 V vs Li (70:17 :13 Asnano:MWCNT:PAA). Cycle number of each unload curve indicated at the top of the graph; corresponding load curves are in the same order.

The differential capacity curve of Asnano/MWCNTs shows a broad anodic peak centered at ~ 0.78 V, present in the corresponding Asnano mixed with MWCNTs curve as a much broader feature, appearing as a shoulder(Figure 13). The feature occurs at a potential that is consistent with the formation of the LiAs, LiAs2, and

LiAs3 phases and is much sharper in the Asnano/MWCNT composite curve as a result of better electrical conductivity.

138

Asnano/MWCNT Composite

A m

/ *

/ Q

D Asnano mixed with MWCNTs

0.6 0.8 1.0 1.2 1.4

Voltage vs Li metal

Figure 13. Differential capacity plots of anodes composed of Asnano and MWCNTs. The Asnano/MWCNT composite displays an additional anodic peak centered at ~ 0.78 V (indicated by the asterisk).

Rate performance of the Asnano/MWCNT composite electrodes were examined from C/10 to 10C (Figure S8). The cells maintained ~90% and ~60% the original unload capacity at C/5 and C/2, respectively, under symmetric load/unload rates. When the load rate was decreased to C/10, the unload could be performed at

2C with ~85% the original unload capacity. Thus, high output power density is possible if charge rates are kept modest.

4.6 Conclusion

The performance of Asmicro/MWCNT composite anodes is significantly better than that of micron size As anodes previously reported.[39][23] While the load/unload curves were very similar, the initial capacity was found to more than

139 double from 89 mAh g-1 to ~220 mAh g-1. This increase in reversibility can be attributed to the addition of the MWCNTs, probably by decrease anode impedance, providing electrical contact for a much larger fraction of the As in the anode.

Mixing Asnano, rather than Asmicro, with MWCNTs vastly improved the reversible capacity to 1076 mAhg-1 and decreased the load/unload hysteresis.

However, the capacity fades rapidly with cycle number if cycled in 1M LiPF6 in 1:1

EC:DMC. Capacity fade is markedly improved by replacing the electrolyte with 1M

LiPF6 in 1:1 FEC:DEC, probably due to the formation of a more effective SEI layer.

Capacity fade can be nearly eliminated by increasing the load cutoff potential from

0.02 to 0.85 V, albeit at the cost decreasing the reversible capacity to ~ 600 mAh g-1.

Much of that capacity can be regained by direct reduction of the Asnano on the

MWCNTs, further decreasing anode impedance and load/unload hysteresis, and resulting in a high capacity As anode that is demonstrated to cycle stably over 200 cycles.

The performance of the Asnano/MWCNT composite anodes surpassed the best results previously reported for As anodes and even surpassed that reported for some P anodes, when the voltage is confined to the same range.[27, 34, 47] The reversible capacity demonstrated here is more than twice that of the best previously reported (based on full electrode formulation mass),[39] rivaling that of other alloying elements.[12] This improved reversible capacity, hysteresis and capacity retention can be accounted for by the improved dispersion of the As in the conductive matrix. Direct reduction allows for higher loadings with good electrically conductive pathways, improved mechanical stability during the

140 expansion/contraction of loading/unloading, and avoiding agglomeration during cycling. High energy mechanical milling techniques cannot provide as evenly coated surfaces on conductive carbons. These materials are hindered by agglomeration at high loadings thus are limited to their Li2X phase (X = P or As).

One might argue that pnictides can not be practical anodes because of their high unload potential compared to graphite, which delithiates at potentials very close to that of Li metal. However, it should be noted that all high capacity alloying elements have higher unload potentials than graphite, albeit many are lower than

As. In addition, the low unload potential of graphite and other alloying elements, including Si, comes with the requirement of lithiation very near to the Li metal plating potential, giving rise to serious safety issues. The higher charging potential of As eases electrolyte electrochemical stability requirements, potentially simplifying their design. Finally, even with its lower average unload voltage of 1.2V vs Li, the Asnano/MWCNT anode could meet the demands set by the USABC in a full cell configuration. Based on calculations laid out by Karulkar et al.,[48] when paired with a traditional NCA cathode, the full cell voltage would be 2.55V and the energy density 795 Wh L-1 (rate – C/10, anode active mass 2 mg cm-2, cathode active mass

10.6 mg cm-2). Under this charge/discharge procedure, the Asnano/MWCNT anode moved closer to the USABC 2020 target for power density with a value of 821 W L-1 while hindering the energy density slightly at 720 Wh L-1. So while the unload potential of As is somewhat higher than some other alloying elements and its toxicity must be mitigated by an effective recycling strategy, its combination of low

141 cost, high capacity and stable cycling make it an attractive material for further study as a Li-ion anode active material.

4.7 References [1] United States Department of Energy. EV Everywhere Grand Challenge. http://energy.gov/eere/vehicles/downloads/ev-everywhere-grand-challenge-road-success, 2014 (accessed 15. 08.01) [2] D. Larcher, S. Beattie, M. Morcrette, K. Edstroem, J.C. Jumas, J.M. Tarascon, J. Mater. Chem., 17 (2007) 3759-3772. [3] J.R. Dahn, T. Zheng, Y.H. Liu, J.S. Xue, Science, 270 (1995) 590-593. [4] M.D. Levi, D. Aurbach, J. Phys. Chem. B, 101 (1997) 4630-4640. [5] H. Buqa, D. Goers, M. Holzapfel, M.E. Spahr, P. Novak, J. Electrochem. Soc., 152 (2005) A474-A481. [6] D. Aurbach, B. Markovsky, I. Weissman, E. Levi, Y. Ein-Eli, Electrochim. Acta, 45 (1999) 67-86. [7] D. Aurbach, B. Markovsky, A. Rodkin, M. Cojocaru, E. Levi, H.-J. Kim, Electrochim. Acta, 47 (2002) 1899-1911. [8] J. Wang, P. Liu, J. Hicks-Garner, E. Sherman, S. Soukiazian, M. Verbrugge, H. Tataria, J. Musser, P. Finamore, J. Power Sources, 196 (2011) 3942-3948. [9] J. Steiger, D. Kramer, R. Monig, J. Power Sources, 261 (2014) 112-119. [10] V. Zinth, C. von Luders, M. Hofmann, J. Hattendorff, I. Buchberger, S. Erhard, J. Rebelo-Kornmeier, A. Jossen, R. Gilles, J. Power Sources, 271 (2014) 152-159. [11] Z. Li, J. Huang, B.Y. Liaw, V. Metzler, J.B. Zhang, J. Power Sources, 254 (2014) 168-182. [12] N. Nitta, G. Yushin, Part. Part. Syst. Charact., 31 (2013) 317-336. [13] M.T. McDowell, S.W. Lee, W.D. Nix, Y. Cui, Adv. Mater., 25 (2013) 4966-4984. [14] J.O. Besenhard, J. Yang, M. Winter, J. Power Sources, 68 (1997) 87-90. [15] C.M. Park, J.H. Kim, H. Kim, H.J. Sohn, Chem. Soc. Rev., 39 (2010) 3115-3141. [16] M.T. McDowell, S.W. Lee, I. Ryu, H. Wu, W.D. Nix, J.W. Choi, Y. Cui, Nano Lett., 11 (2011) 4018-4025. [17] I.A. Courtney, W.R. McKinnon, J.R. Dahn, J. Electrochem. Soc., 146 (1999) 59-68. [18] R. Krishnan, T.-M. Lu, N. Koratkar, Nano Lett., 11 (2013) 377-384. [19] N. Liu, Z. Lu, J. Zhao, M.T. McDowell, H.-W. Lee, W. Zhao, Y. Cui, Nat Nano, 9 (2014) 187-192. [20] D.S. Su, R. Schlogl, ChemSusChem, 3 (2010) 136-168. [21] W.W. Sun, Y. Wang, Nanoscale, 6 (2014) 11528-11552. [22] S. Goriparti, E. Miele, F. De Angelis, E. Di Fabrizio, R. Proietti Zaccaria, C. Capiglia, J. Power Sources, 257 (2014) 421-443. [23] B.J. Landi, M.J. Ganter, C.D. Cress, R.A. DiLeo, R.P. Raffaelle, Energy Environ. Sci., 2 (2009) 638-654. [24] H. Kim, E.-J. Lee, Y.-K. Sun, Mater. Today, 17 (2014) 285-297. [25] M.L. Terranova, S. Orlanducci, E. Tamburri, V. Guglielmotti, M. Rossi, J. Power Sources, 246 (2014) 167-177.

142

[26] J.E. Trevey, K.W. Rason, C.R. Stoldt, S.-H. Lee, Electrochem. Solid-State Lett., 13 (2010) A154-A157. [27] Y. Kim, Y. Park, A. Choi, N.-S. Choi, J. Kim, J. Lee, J.H. Ryu, S.M. Oh, K.T. Lee, Adv. Mater., 25 (2013) 3045-3049. [28] N. Liu, H. Wu, M.T. McDowell, Y. Yao, C. Wang, Y. Cui, Nano Lett., 12 (2012) 3315-3321. [29] D. Aurbach, J.S. Gnanaraj, W. Geissler, M. Schmidt, J. Electrochem. Soc., 151 (2004) A23-A30. [30] V. Etacheri, O. Haik, Y. Goffer, G.A. Roberts, I.C. Stefan, R. Fasching, D. Aurbach, Langmuir, 28 (2012) 965-976. [31] C. Marino, A. Debenedetti, B. Fraisse, F. Favier, L. Monconduit, Electrochem. Commun., 13 (2011) 346-349. [32] J. Sun, G. Zheng, H.-W. Lee, N. Liu, H. Wang, H. Yao, W. Yang, Y. Cui, Nano Lett., 14 (2014) 4573-4580. [33] J. Qian, X. Wu, Y. Cao, X. Ai, H. Yang, Angew. Chem., Int. Ed., 52 (2013) 4633- 4636. [34] J.F. Qian, D. Qiao, X.P. Ai, Y.L. Cao, H.X. Yang, Chem. Commun., 48 (2012) 8931-8933. [35] C.-M. Park, S. Yoon, S.-I. Lee, J.-H. Kim, J.-H. Jung, H.-J. Sohn, J. Electrochem. Soc., 154 (2007) S19. [36] M. He, K. Kravchyk, M. Walter, M.V. Kovalenko, Nano Lett., 14 (2014) 1255- 1262. [37] C.-M. Park, S. Yoon, S.-I. Lee, H.-J. Sohn, J. Power Sources, 186 (2009) 206-210. [38] J.O. Besenhard, H.P. Fritz, Electrochim. Acta, 20 (1975) 513-517. [39] C.-M. Park, J. Solid State Electrochem., 20 (2015) 517-523. [40] U.S. Geological Survey, Mineral commodity summaries 2016: U.S. Geological Survey, 2016, 26-27. [41] J. R. Loebenstein, Materials Flow Analysis of Arsenic in the United States in Bureau of Mines Information Circular. IC 9382, 1994. [42] United States Environmental Protection Agency, Arsenic and Clarifications to Compliance and New Source Monitoring Rule: A Quick Reference Guide, EPA 816-F- 01-004, 2001. [43] R. C. Weast, Handbook of chemistry and physics. A ready reference book of chemical and physical data, 45 ed., The Chemical Rubber Co. Cleveland, Ohio, 1964. [44] M.J. Wagner, Inorganic Nanomaterials Synthesis Using Alkalide Reduction, in: C.M. Lukehart, R.A. Scott (Eds.) Encyclopedia of Inorganic Chemistry, John Wiley & Sons, Ltd, Chinchester, UK, 2008, pp. 357-368. [45] K. Xu, Chem. Rev., 104 (2004) 4303-4417. [46] Z.X. Yang, G.D. Du, Z.P. Guo, X.B. Yu, S.A. Li, Z.X. Chen, P. Zhang, H.K. Liu, Nanoscale, 2 (2010) 1011-1017. [47] C.-M. Park, H.-J. Sohn, Adv. Mater., 19 (2007) 2465-+. [48] M. Karulkar, R. Blaser, B. Kudla, J. Power Sources, 273 (2015) 1194-1201.

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4.8 Supporting Information

Figure S1. Deconvolution of the overlapping graphite and MWCNT XRD peaks into the individual contributions using Person VII functions. The broad peak (green) can be assigned to the MWCNTs while the sharp peak (red) is due to graphite.

144

Figure S2. SEM of the MWCNT material used in this study. The inset shows a zoomed in area to highlight MWCNT content. Large graphite crystals can be seen throughout but are especially evident in the image center to lower and upper right where they have very little MWCNT coverage.

2.0

l a t 1.5

e 2nd

i Load 1st Load

v 1.0

a t

l 1st & 2nd

o 0.5 unload V

0.0 0 200 400 600 800 1000 1200 1400

-1 Capacity (mAh g ) Figure S3. Voltage profiles of the first cycles of a MWCNT - Li halfcell.

145

Figure S4. EDS spectrum of Asnano. The large C, O, and Cu peaks are from the C coated Cu grid.

146

Figure S5. SEM micrograph of Asmicro.

147

2.0

1.5

L 1.0

a t

l 0.5 o

V 20th 10th 1st

0 400 800 1200 1600

-1 Capacity mAh g As Figure S6. Voltage profiles of Asnano anode - Li metal half-cell cycled between 0.02 and 2.0 V in EC:DMC.

148

8a 1st cycle 10th cycle 20th cycle Element Value Standard Value Standard Value Standard deviation deviation deviation Rcell 8.5 1.5 4.5 0.4 3.2 0.3 RSEI 3.6 2.7 74 27 120 5 CSEI - T 1.6 E-5 3.1 E-5 2.5 E-4 0.4 E-4 1.5 E-4 0.1 E-4 CSEI - P 0.80 0.26 0.55 0.04 0.61 0.01 RCT 26 1 40 8 44 10 CDL - T 1.1 E-3 0.1 E-3 4.6 E-3 0.8 E-3 2.0 E-3 0.2 E-3 CDL - P 0.50 0.03 0.71 0.05 0.90 0.04 Wdiff - R 12 5 5.0 62 170 50 Wdiff - T 1.9E-2 1.2E-2 0.030 0.550 2.2 1.2 Wdiff - P 3.5E-2 0.2 E-2 0.33 0.01 0.29 0.01

8b 1st cycle 10th cycle 20th cycle Element Value Standard Value Standard Value Standard deviation deviation deviation Rcell 7.5 0.2 7.4 0.4 7.3 0.1 RSEI 30 6 48 14 46 1 CSEI - T 1.9 E-5 3.6 E-5 1.8 E-4 0.4 E-4 3.7 E-5 0.1 E-5 CSEI - P 0.62 0.03 0.51 0.04 0.73 0.01 RCT N/A N/A 19 6 7.3 0.9 CDL - T N/A N/A 1.4 E-4 0.2 E-4 1.4 E-4 0.1 E-4 CDL - P N/A N/A 0.92 0.05 1.0 0.1 Wdiff - R 5.0 16 6.0 59 40 1 Wdiff - T 1.4 E-2 6.9 E-2 0.02 0.27 0.53 0.02 Wdiff - P 0.35 0.01 0.31 0.01 0.32 0.01 Table S1. Fitted values of EIS data found in 8a and 8b. The use of separate RC elements for the SEI and CT was not justified by the fit of the first cycle data of 8b. Values listed for the first cycle RCT , CDL – T , and CDL – P are those of a single RC circuit representing the combined SEI and CT contributions.

149

Figure S7. EDS line scan across Asnano reduced on a MWCNT, perpendicular to the principal axis of the MWCNT. There is no significant increase of the oxygen content scanning from the carbon grid across the Asnano on the MWCNT, indicating little or no oxidation layer on the Asnano.

)

0 100

t a

(

i c

a 60

i f

i 40

e v

i 20

e R 0 C/10 C/5 C/2 1C 2C 5C 10C C/10

Unload C-Rate Figure S8. Rate performance of Asnano reduced on MWCNTs (squares, 70:17:13 As:MWCNT:PAA w:w, Loaded at C/10, unloaded at specified rate) and Asnano reduced on MWCNTs (circles, 70:17:13 As:MWCNT:PAA w:w, Symetric Load/Unload) cycled in 1:1 FEC:DEC w:w with 1M LiPF6. Lines between data points are to guide the eye.

150

5. A Room Temperature Study of Gallium Arsenide as a New Multi Alloying Anode for Li-Ion Batteries

For Submission to the Journal of the Electrochemical Society* 5.1 Introduction

The increased demand for Li-ion battery performance has led to significant efforts in anode material research, with great gains being made by using Li-alloying elements that have large specific capacities. Alloying materials such as Si, Ge, and Sn have extremely large electrochemical Li storage capabilities, achieving specific capacities of

3579, 1384, and 960 mAh g-1 respectively.[1] At first glance, the numbers seem spectacular when compared to the industry standard graphite anode (376 mAh g-1) used in Li-ion batteries, but obtaining these values over many charge/discharge cycles at practical current densities and electrode loadings has thus far proven to be difficult. For example, theoretically, 1 mole of Si can alloy with 4.4 moles of Li, yielding Li22Si5.[2]

However, the crystal structure is expanded upwards of 400% in order to accommodate this large amount of Li.[3] Upon unloading of Li, the crystal structure of Si is left with defects and voids. After multiple charge/discharge cycles cracks form, resulting in loss of electrical contact and capacity fade.

Size reduction from the micron to nano domain improves the cycling characteristics of alloying materials, as small particles undergo the large volume change with less mechanical degradation, but agglomeration after lithiation of the alloying nanoparticles leads to larger particles and consequent capacity fade.[4] The agglomeration of Li-alloying nanomaterials can be mitigated to some extent by dispersion in conductive carbons.[5, 6] This technique has proven successful in

151 prolonging capacity retention, but addition of too much non-active material results in energy densities (specific, volumetric and/or areal) with little advantage over graphite.[7,

8] Because of this, other strategies should be developed utilizing less commonly researched alloying anode materials.

Gallium is a Li-alloying metal with a theoretical specific capacity of 769 mAh g-

1 ,[9] based on the formation of Li2Ga.[10] It has a low melting point of 29°C, allowing it to exist as a liquid in the operational temperature range of Li-ion batteries, a property that could make it a “self-healing” anode material.[11] Voids caused by the expansion/contraction that occurs during cycling could be healed by operating in a liquid state, avoiding mechanical degradation. Deshpande et. al. [12] reported that despite the lack of mechanical degradation, capacity fade still occurred, most likely due to the lack of a stable SEI formation. Commercial anodes form a stable SEI within the first few cycles, but liquid Ga reforms its surface resulting in additional SEI formation during each cycle.

To better stabilize the Ga, binary compounds such as CuGa2,[13] Ga2S3,[14] and Ga2Se3

[15] have been investigated. The addition of these electrochemically inactive components result in improved capacity retention, but reduced energy density.

Here we report on the electrochemical performance of GaAs, which has a number of advantages over previously studied binary Ga compounds. The addition of As to stabilize Ga improves, rather than decreases, its capacity because As has the third largest specific capacity of any Li alloying element, 1073 mAh g-1, higher than Sn or Ga,[1] As is earth abundant and inexpensive, and GaAs has very good electrical conductivity.

GaAs itself is inexpensive, it is a waste product of the GaAs semiconductor wafer manufacturing, which consumes ~75% of all Ga produced in the United States.[16] The

152 yield of this process is small; less than 6 percent of the original grown GaAs bullion becomes usable wafers after wafer slicing, backthinning, dicing, and packaging.[17]

Thus, GaAs is a readily available material with the potential to become an economical, high capacity anode material. Furthermore, while Ga and As have individually been investigated as Li-ion anode materials,[11, 18-21] studies of the alloy itself are absent from the literature. Herein we present the first study of GaAs anodes, in both bulk and nanomaterial form, for Li-ion batteries.

5.2 Experimental

GaAs nanoparticles were synthesized by akalide reduction.[22] GaI3 (Strem, 99%, anhydrous) and AsI3 (Strem, 98%, anhydrous) were used without further purification.

Tetrahydrafuran (THF, Fisher Scientific, 99.9+% HPLC grade, inhibitor-free) was purified by stirring over NaK alloy until a persistent blue solution was obtained, then vacuum transferred to storage bottles. Solution preparation and alkalide reduction were performed in open vessels at room temperature in an N2 glovebox (< 1 ppm H2O and O2).

+ - K (15-crown-5)2Na solution was prepared by stirring NaK alloy (0.31 g) and 15-crown-5

(2.21 g, Alfa Aesar) in 30 ml of THF. After complexation, the solution was swiftly added to a rapidly stirring solution of 0.304g AsI3 and 0.302g GaI3 in 30ml of THF. A black precipitate of GaAs nanoparticles formed immediately.

Following reduction, methanol (80 ml, Fisher Scientific, 99.9%, dried on 3Å sieves) that was added to increase the rate of precipitation. The product was separated from the solution by centrifuge (5 min, 8300 x g). The product was washed 3 times with

N2 purged methanol in a N2 purged glove bag, each time followed by separation of the product from the wash by centrifuge. Finally, the samples were dried under vacuum and

153 stored under N2. The collected yield was ~ 58.%, with the majority of the losses occurring in the washing. Ga and As nanoparticles were synthesized following the same procedure by reduction of the individual salts rather than co-reduction.

Powder x-ray diffraction patterns were obtained with a Rikagu Miniflex+ diffractometer (Cu Kα radiation) maintained in the air-free environment of a N2 filled glovebox (< 1 ppm H2O and O2). Samples were scanned on a zero background quartz plate from 10° to 90° 2θ, at 0.5° min-1.

Inductively coupled plasma atomic emission spectroscopy was used to determine the atomic ratio of Ga and As in the synthesized GaAs nanoparticles with a Shimadzu ICPE-9800. GaAsnano was dissolved in nitric acid and the obtained concentration (mg/L) were converted to molar values.

Samples were examined using a FEI Talos 200X TEM with high-angle, annular dark field scanning TEM and Super X electron dispersive spectroscopy (EDS). EDS mapping was analyzed by Bruker ESPRIT™ software.

UV-Vis absorbance spectroscopy of the GaAsnano was collected on a Thermo

Scientific Evolution 600 UV-visible spectrometer. Spectra of a colloidal suspension of

GaAsnano in methanol were collected from 200 to 800 nm using a 10mm path length quartz cuvette. The average of 10 scans was used to perform a linear least-squares fit of the shoulder of the absorption band to extrapolate the band gap value and estimate the

GaAsnano particle size.[23]

Electrodes were prepared by slurry and doctor blade method in a N2 purged glovebag. Bulk GaAs (99.999%, 25mm polycrystalline, Alfa Aeser) was ground from 5 mm pieces down to a fine powder (GaAsmicron) with a Fritsch Mini-Mill Pulverisette 23

154 ball mill located in a N2 purged glovebag, using a 15 ml stainless steel grinding bowl and one 15 mm diameter stainless steel ball at 40 Hz for 10 min. GaAsmicron was then mixed with carbon black (Super C45, Imerys Graphite and Carbon) and polyacrylic acid (PAA,

6 MW ~1 x 10 , Polysciences, Inc.) in a 80:10:10 w:w, GaAsmicron:carbon black:PAA ratio and milled for an additional 1 minute. GaAsnano, multiwall carbon nanotubes (MWCNT,

99+%, >20 nm diameter, 1-12 µm length, Cheap Tubes Inc.), and PAA were mixed in

68:22:10 or 45:45:10 ratios (GaAsnano:MCWCNT:PAA w:w) for 1 minute in a Wig-L-

Bug® grinding mill (Reflex Analyical Corp.) located inside a N2 purged glovebag. N- methyl-2-pyrrolidone (NMP, Sigma Aldrich, 99+%, dried on 4Å sieves) was added to the

GaAsmicron and GaAsnano mixtures to produce slurries. The slurries were homogenized using an ultrasonic bath for 15 minutes and applied to pre-weighed carbon coated copper current collectors (1 µm conductive carbon on 9 µm copper foil, MTI Corp.). The electrodes were dried under active vacuum (10-3 torr) for 6 hour and then reweighed to obtain the electrode mass loading. Active material (GaAs) loading was 1.5 to 2 mg cm-3.

All GaAs anodes testing was performed in a half-cell configuration with a Li

(99.9%, MTI Corp.) counter electrode in 2016 coin cells (MTI Corp.) with Celgard®

2325 separator (Celgard LLC) and 1M LiPF6 in 1:1 FEC:DEC (V:V, >99%, Solvay) electrolyte. The GaAs half-cell battery performance was tested using Arbin BT2000 cyclers. The cells were cycled under constant current from 0.01 to 2.0V vs Li at C/10, unless otherwise noted. The cells were rested under open voltage after each load/unload cycle for 1 hour.

Ex-situ powder XRD was performed by loading the GaAs anodes with a 90:10

GaAs:polyvinylidene fluoride (PVDF, Kureha KF 1100) binder ratio to a given potential

155 at constant current (C/20), followed by constant voltage until current decreased to 1x10-

7A. The GaAs electrodes were extracted from their cells and their powder x-ray diffraction (XRD) patterns obtained at a scan rate of 0.15° min-1 with a Rikagu Miniflex+ diffractometer (Cu Kα radiation) maintained in the air-free environment of a N2 filled glovebox (< 1 ppm H2O and O2).

5.3 Results and Discussion

The voltage profiles of the first and second cycles of GaAsmicron electrodes

(80:10:10 GaAs:Super C45:PAA) are shown in Figure 1. During the first cycle, the majority of the lithium loaded at ~0.4 V, while the corresponding unload curve was steep with a decreasing slope between ~0.3 and 1.1 V, increasing thereafter. Second cycle load and unload curves were similar but with steeper features, nearly linear at potentials below ~ 0.8 V and increasingly steeply sloping at higher potentials, consistent with an increase in cell impedance. The GaAs anode displayed a large initial reversible capacity,

650 mAh g –1, however, the capacity faded rapidly, declining to ~460 mAh g –1 in the second cycle and losing ~ 2/3 of its initial capacity over the first 10 cycles (Figure 2).

Coulombic efficiency (CE) increased only slightly from 1st to 2nd cycle, increasing gradually from the 3rd to the 10th cycle but still remaining below 99%, indicative of ineffective SEI layer passivation.

156

Figure 32 Voltage profiles for the 1st (solid line) and 2nd (dotted line) load/unload cycles of a GaAsmicron electrode.

Figure 33 The reversible capacity (open circles, left axis) and Coulombic efficiency (filled circles, right axis) of GaAs plotted as a function of cycle number. The lines between the data points are to guide the eye.

Plotting the differential capacity of the first cycle clarifies features not easily identifiable in the voltage profile (Figure 3). The small shoulder in the unload curve at

0.6 V followed by a broad cathodic peak at 0.8 V is similar to that observed in Li-Ga cells for delithiation from LiGa to Li2Ga7 and Li2Ga7 to Ga.[8] Two additional broad cathodic peaks at ~ 1.1 and 1.35 V are similar to those observed in Li - As cells, likely a

157 result of Li3As dealloying to Li2As and LiAs dealloying to As.[20] The small broad peak at ~0.4 V is consistent with delithiation of the carbon black [24].

Figure 34 Differential capacity plot of the first cycle of a GaAsmicron electrode.

The fully lithiated state of GaAsmicron was investigated by ex situ powder XRD

(Figure 4). Patterns were collected of a freshly prepared GaAsmicron anode prior to cycling and again after loading it to 0.01V, and finally after unloading to 2.0V. The pattern of the freshly prepared anode shows sharp reflections of crystalline GaAs. These reflections are not present in the pattern for the loaded material, instead displaying broad reflections corresponding to LiGa (PDF# 09-0043). The LiGa crystallite size was found to be ~7 nm using Scherrer’s equation.[25] Although crystalline Li-As could not be identified in the pattern, it is clear that it must be lithiated because the reversible capacity is too large to be accounted for by Ga lithiation alone. The formation of LiGa could account for a capacity of 385 mAh g-1, just 60% of the reversible capacity obtained. After unloading the Li to

2.0V the XRD patterns show no evidence of any crystalline phase, including the GaAs seen in the pristine anode.

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Figure 4 Ex situ powder XRD of freshly prepared, uncycled (top), loaded to 0.01 V (middle) and loaded to 0.01 V and unloaded to 2.0 V (bottom) GaAsmicron electrodes. Reflections from the Cu metal and carbon coating on Cu current collector are indicated by pound sign and asterisk respectively.

Significant pulverization of the GaAsmicron anodes is clear from HAADF-STEM images taken before and after a single cycle (Figure 5A & B). Initially, particle sizes ranged from ~100 nm to over 2 µm. After a single load/unload cycle, the GaAs particles were generally ~ 20 nm or less in their largest dimension. Evidence for the presence of nanocrystalline GaAs in the unloaded anode can be seen in HRTEM images that display lattice fringes consistent with the (111) plane of GaAs (Figure 5C). Taken together, powder XRD and STEM results are consistent with the decomposition and pulverization of the GaAsmicron to separate nanoscale LiGa and Li-As alloys and the formation of

159 nanocrystalline GaAs upon unloading. Furthermore, the pulverization from the large expansion/contraction of the load/unload cycle results in the formation an unstable SEI with consequently poor CE and rapid capacity fade.

Figure 5 HAADF STEM images of GaAsmicron electrode (A) prior to Li loading (B) after a single Li load/unload cycle and (C) a HRTEM image showing (111) lattice spacing of GaAs after a single Li load/unload cycle.

Reducing crystallite size to nanoscale domains may alleviate electrode stress caused by the volume changes that occur during lithiation/delithiation.[26] GaAs nanoparticles have been synthesized by a number of methods,[27-29] but all of these routes involve multi-day reactions at temperatures above 100°C. Here we present results for GaAs nanocrystals (GaAsnano) synthesized in a few minutes at room temperature by akalide reduction.[22] The resulting product changed from black to brown upon prolonged exposure to air, indicating slow oxidation. Thus, exposure of the product to oxygen was minimized. Powder XRD diffraction patterns for GaAsnano display three broad humps at 13, 31, and 53°, 2Ɵ (Figure 6A), consistent with amorphous nanoparticles or very small nanocrystals. As (Figure 6B) and Ga (Figure 6C) nanoparticles synthesized individually by akalide reduction also resulted in material that appears amorphous by powder XRD, however, the pattern for GaAsnano does not appear to be the sum of the individual As and Ga nanoparticle patterns, consistent with the 160 product being an intimate mixture of Ga and As at the atomic level. ICP-AES confirmed an atomic ratio of 1:1.1 (Ga:As).

Figure 6 Powder XRD pattern of (A) GaAsnano, (B) As nanoparticles (C) Ga nanoparticles and (D) GaAsmicron.

TEM micrographs of GaAsnano show that it consists of agglomerates of nanoparticles (Figure 7). Fast Fourier transformation of HRTEM images of these agglomerates reveal a ring consistent with polycrystalline domains with 3.3 Å lattice fringe spacing in good agreement with the (111) planes of GaAs (Figure 8). EDS mapping, along with HAADF-STEM, confirmed uniform distribution of As and Ga throughout the sample, consistent with GaAs formation rather than separate Ga and As

161 particles (Figure 9). Oxygen was also present throughout EDS map, most likely due to oxidation during transfer to TEM.

Figure 7 TEM micrograph of GaAsnano agglomerates.

Figure 8 HRTEM of a GaAs nanoparticle agglomerate. Inset: FFT of HRTEM image showing a ring that is consistent with the (111) reflection of GaAs.

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Figure 9 (A) HAADF STEM of GaAs nanoparticles, EDS maps of GaAsnano showing the distribution of (B) Ga, (C) O, (D) As and (E) line scan of 100 nm section of GaAsnano.

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The UV-Vis absorption spectra of GaAsnano dispersed in methanol are consistent with GaAs nanocrystals, (Figure S1) with the bandgap blue shifted from that expected for bulk GaAs. The size of the nanocrystals of GaAs is estimated from the magnitude of the blue shift [30] to be 5 nm in good agreement with HRTEM observations.

MWCNTs were chosen as the conductive carbon for GaAsnano anodes because of their excellent electrical conductivity, relatively low price and high surface area for alloy coating.[31] The “99% pure” MWCNTs used in this study were examined by powder

XRD (Figure S2). A sharp peak at 3.36 Å and broad shoulder at a slightly high d-spacing provide evidence of graphite in addition to the MWCNTs. Graphite content is estimated to be ~21.6% (w:w) by deconvolution of the reflection using a Pearson VII peak shape (,

Jade 6.1, Materials Data Inc.). This graphite content is evident in SEM images (Figure

S3), although it is also clear that the MWCNTs vastly outnumber the graphite crystallites and are largely responsible for the materials surface area. Examination of the voltage profiles of 90:10 w:w MWCT:PAA Li half cells clearly show plateaus, characteristic of graphite intercalation(Figure S4). Thus, although the conductive carbon used in this study was a mixture of MWCNT and graphite, it will be referred as MWCNTs in the text for the sake of brevity.

st th The 1 and 10 cycle of the voltage profiles of GaAsnano:MWCNT:PAA

(45:45:10 by mass) electrodes are similar to those of GaAsmicron, particularly the unload curves which are nearly identical, with the exception of the greater hysteresis between load and unload (Figure S5). This is indicative of more resistive electrical contact to the

GaAsnano than was present in the GaAsmicron electrodes, possibly due to less effective contact to the conductive matrix, grain boundary resistance between the nanoparticles or

164 surface oxidation that occurred during electrode preparation. The majority of GaAsnano loading occured at 0.4 V and lower potentials, however, with a much steeper slope than displayed by GaAsmicron. This behavior is consistent with inhomogeneous conductivity; nanoparticles in closest contact with the MWCNTs having a relatively low impedance conductive path to the current collector and less conductive paths through the GaAsnano agglomerates. It should be noted that a small fraction of the capacity (~60 mAh g-1) is due to MWCNT and graphite intercalation, occurring at low potential.[31] The powder

XRD pattern of a GaAsnano electrode loaded to 0.01V and disassembled is similar to that found for GaAsmicron, displaying a small peak at 24.8° 2θ matching the (111) reflection of

LiGa, but absent of reflections matching any crystalline Li-As phases (Figure 10). It should be noted that while no evidence of crystalline Li-As phases are observed, it must be present to account for the high, and stable, reversible capacity, perhaps as amorphous

Lix-As and/or Lix-GaAs phases.

Figure 10 Ex Situ powder XRD of Li - loaded and unloaded GaAsnano electrodes. Electrodes comprised of 90:10 GaAs:PVDF on carbon coated Cu current collectors. Reflections from the Cu metal, carbon coating on Cu current collector, and PVDF are indicated by pound sign and asterisk respectively.

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Electrodes with a 45:45:10 (GaAsnano:MWCNT:PAA w/w) display excellent capacity retention (Figure 11). Capacity initially increased from 560 to 714 mAh g-1 over the first 12 cycles, indicating that the expansion/contraction of cycling created conductive paths to initially isolated GaAs nanoparticles, possibly through physical rearrangement or cracking/flaking of surface oxidation. Examination of the load/unload curves (Figure

12 A) shows that the increase in capacity over the first 10 cycles is primarily due to the development of sloping plateaus at ~ 0.85 and ~ 1.05 V in the load and unload curves, respectively, which appear as broad peaks in the differential capacity curves (Figure 13).

The development of the plateaus is accompanied by a small decrease unload and a larger increase in the load potentials, significantly decreasing the hysteresis, consistent with the development of better conductive networks to the GaAsnano over the first 10 cycles.

Capacity began to fade after ~ 18 cycles, with a monotonic and nearly linear loss or an average of 0.16 % of initial capacity per cycle, retaining 610 mAh g-1 after 100 cycles.

Figure 11 Specific capacity retention plotted as a function of cycle number for electrodes with GaAsnano:MWCNT w/w ratios of 50:50 (squares) and 75:25 (circles).

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Figure 12 Voltage profiles of GaAsnano electrodes with 50:50 (A) and 75:25 (B) w/w GaAsnano:MWCNTs compositions.

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Figure 13 Differential capacity plot for the 10th cycle of the 50:50 GaAsnano:MWCNT electrode.

Electrodes with a higher fraction of alloying active material (68:22:10

-1 GaAsnano:MWCNT:PAA w/w) achieved higher initial capacity, 850 mAh g , however, capacity fade was much more dramatic, losing 76% of the original capacity to retain only

200 mAh g-1 after 80 cycles (Figure 11). In contrast to the development of plateaus and decrease in hysteresis found with 45:45:10 electrodes, no significant change in the functional form of the load and unload curves occurs and hysteresis monotonically increases with increasing cycle number (Figure 12B). This behavior is consistent with continual degradation of the electrode with consequent impedance growth.

The first cycle CE of the GaAsnano electrodes increase with decreasing MWCNT content (Figure 14), rising from 36 and 58% with an GaAsnano:MWCNT ratio from 50:50 to 75:25 w/w (45:45:10 and 68:22:10 GaAsnano:MWCNT:PAA w/w). Carbon nanotubes have excellent conductivity, but have high surface area that must be passivated in the first cycle.[32] Subsequent cycles have much higher CE, ~ 90% on second cycle and rising above 99% during extended cycling. While both electrode compositions obtain very high

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CE, the 45:45:10 electrode’s CE is far more stable, consistently above 99%, while that of the 68:22:10 electrode is more erratic, decreasing to as low as ~95% on some cycles.

Taken together with the discharge curves, it seems clear that the more highly loaded

GaAsnano electrode exhibits far more significant mechanical degradation due to the stress of the expansion/contraction cycle of loading/unloading Li, resulting in increased SEI growth during extended cycling and loss of active material due to electrical isolation. In contrast, the more lightly loaded electrode provides a sufficient network of conductive material and void space to allow for more reversible, non-destructive expansion during loading.

Figure 14 Coulombic efficiency plotted as a function of cycle number for electrodes with GaAsnano:MWCNT w/w ratios of 50:50 (squares) and 75:25 (circles).

5.4 Conclusions

Micron and nanoscale GaAs was examined for its electrochemical properties as an alternative anode material for Li ion batteries. The low melting point of Ga results in a material that, alone, is unsuitable for Li ion anodes, because of an inadequate formation of the SEI. GaAs may benefit from the “self healing” properties of Ga along with the higher capacity of As, producing a stable Li alloying anode material that can be charged and discharged over many cycles. Micron GaAs, while displaying good first cycle

169 capacity (~660 mAh g-1) , faded rapidly likely due to mechanical degradation in the electrode. GaAs nanoparticles were synthesized by akalide reduction and incorporated into MWCNTs to improve the conductivity and the capacity retention of this material.

The conductive network of the MWCNTs proved crucial to prolonging the cycle life of

GaAsnano . A 50:50 mixture of GaAsnano and MWCNTs was determined to be the optimal ratio to obtain 100 cycles over 600 mAh g-1, maxing out at 715 mAh g-1. This work may provide a foundation gain a better understanding of the exact method for lithiation of

GaAs. Furthermore, the excellent capacity and cycling retention makes GaAs an enticing material for further investigation in Li-ion batteries.

5.5 References

1. N. Nitta and G. Yushin, Part. Part. Sys. Charac., 31, 317 (2013). 2. M. T. McDowell, S. W. Lee, W. D. Nix and Y. Cui, Adv. Mater., 25, 4966 (2013). 3. M. N. Obrovac and L. Christensen, Electrochem. Solid-State Lett., 7, A93 (2004). 4. I. A. Courtney, W. R. McKinnon and J. R. Dahn, J. Electrochem. Soc., 146, 59 (1999). 5. W. W. Sun and Y. Wang, Nanoscale, 6, 11528 (2014). 6. S. Goriparti, E. Miele, F. De Angelis, E. Di Fabrizio, R. Proietti Zaccaria and C. Capiglia, J. Power Sources, 257, 421 (2014). 7. B. P. N. Nguyen, S. Chazelle, M. Cerbelaud, W. Porcher and B. Lestriez, J. Power Sources, 262, 112 (2014). 8. M. Karulkar, R. Blaser and B. Kudla, J. Power Sources, 273, 1194 (2015). 9. J. Saint, M. Morcrette, D. Larcher and J. M. Tarascon, Solid State Ion., 176, 189 (2005). 10. C. J. Wen and R. A. Huggins, J. Electrochem. Soc., 128, 1636 (1981). 11. K. T. Lee, Y. S. Jung, T. Kim, C. H. Kim, J. H. Kim, J. Y. Kwon and S. M. Oh, Electrochemical and Solid-State Letters, 11, A21 (2008). 12. R. D. Deshpande, J. C. Li, Y. T. Cheng and M. W. Verbrugge, J. Electrochem. Soc., 158, A845 (2011). 13. K. T. Lee, Y. S. Jung, J. Y. Kwon, J. H. Kim and S. M. Oh, Chem. Mat., 20, 447 (2008). 14. H. Senoh, H. Kageyama, T. Takeuchi, K. Nakanishi, T. Ohta, H. Sakaebe, M. Yao, T. Sakai and K. Yasuda, J. Power Sources, 196, 5631 (2011). 15. J.-J. Ding, Y.-N. Zhou, Y.-H. Cui and Z.-W. Fu, ECS Electrochem. Lett., 1, A7 (2012). 16. U.S. Geological Survey, Mineral commodity summaries 2016: U.S. Geological Survey, 26 (2016).

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17. K. Torrance and H. E. Keenan, Semiconductor Today, 4, 66 (2009). 18. W. Liang, L. Hong, H. Yang, F. Fan, Y. Liu, H. Li, J. Li, J. Y. Huang, L.-Q. Chen, T. Zhu and S. Zhang, Nano Lett., 13, 5212 (2013). 19. M. Yarema, M. Worle, M. D. Rossell, R. Erni, R. Caputo, L. Protesescu, K. V. Kravchyk, D. N. Dirin, K. Lienau, F. von Rohr, A. Schilling, M. Nachtegaal and M. V. Kovalenko, J. Am. Chem. Soc., 136, 12422 (2014). 20. C.-M. Park, J. Solid State Electrochem., 20, 517 (2015). 21. K. Hays, N. Banek and M. J. Wagner, Unpublished Results (2016). 22. M. J. Wagner, in Encyclopedia of Inorganic Chemistry, 2nd ed., C. M. Lukehart and R. A. Scott Editors, p. 357, John Wiley & Sons, Ltd, Chinchester, UK (2008). 23. J. Tauc, R. Grigorovici and A. Vancu, physica status solidi (b), 15, 627 (1966). 24. R. M. Gnanamuthu and C. W. Lee, Mater. Chem. Phys., 130, 831 (2011). 25. P. Scherrer, Nachr. Ges. Wiss. Göttingen 26, 98 (1918). 26. J. O. Besenhard, J. Yang and M. Winter, J. Power Sources, 68, 87 (1997). 27. S. S. Kher and R. L. Wells, Chem. Mat., 6, 2056 (1994). 28. M. A. Malik, P. O'Brien, S. Norager and J. Smith, J. Mater. Chem., 13, 2591 (2003). 29. H. Uchida, C. J. Curtis and A. J. Nozik, J. Phys. Chem., 95, 5382 (1991). 30. L. E. Brus, J. Chem. Phys., 80, 4403 (1984). 31. B. J. Landi, M. J. Ganter, C. D. Cress, R. A. DiLeo and R. P. Raffaelle, Energy Environ. Sci., 2, 638 (2009). 32. M. Winter, P. Novak and A. Monnier, J. Electrochem. Soc., 145, 428 (1998).

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5.6 Supplementary Material

Figure S1 UV/Vis spectrograph of GaAs nanoparticles dispersed in methanol. The dotted line represents the linear fit used to extrapolate the particle size from bandgap energy.

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Figure S2. Deconvolution of the overlapping graphite and MWCNT XRD peaks into the individual contributions using Person VII functions. The broad peak (green) can be assigned to the MWCNTs while the sharp peak (red) is due to graphite.

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Figure S3. SEM of the MWCNT material used in this study. The inset shows a zoomed in area to highlight MWCNT content. Large graphite crystals can be seen throughout but are especially evident in the image center to lower and upper right where they have very little MWCNT coverage.

Figure S4. Voltage profiles of the first cycles of a MWCNT - Li halfcell.

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th Figure S5 1st (solid lines) and 10 (dashed lines) cycle voltage profiles of a GaAsnano electrode.

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6. Conclusions

The state of Li-ions batteries is at a turning point that leaves the scientific community an opportunity to make large gains in the battery’s power and energy by utilizing nanomaterials, especially in the anode. Alloying anodes, which previously proved problematic due to expansion and mechanical degradation, are now approaching the cycle life and exceeding the capacity of traditional graphite intercalation anodes.

Multilayer graphene nanoshells have the ability to charge and discharge much more rapidly than graphite, but at a lower capacity (~220 mAh g-1). In an attempt to improve the capacity, Sn metal was encapsulated into the hollow cores by synthesizing

MGNS with a mix of Ni and Sn salts and selectively removing the Ni by HNO3 reflux.

This Sn@MGNS improved capacity to 300 mAh g-1, but faded quickly to ~160 mAh g-1 with repeated charge/discharge cycles. Passivation of the metal cores lead to excessive capacity fade, however, SnO2 added to the outside of MGNS by NaBH4 reduction of the metal halide and subsequent oxidation cycled very well with an initial capacity of

-1 400 mAh g . SnO2/MGNS composites showed much improvement in capacity retention, fading to only 250 mAh g-1 over 160 cycles.

While incorporation of Sn into MGNS showed capacity improvements, loading of greater than ~30% by mass Sn into the MGNS resulted in poor cycle life. Further gains in the capacity of the anodes were desired so other alloying anode materials were tested.

Asnano/MWCNT composite electrodes were cycled against Li for 200 cycles, maintaining a capacity of 444 mAh g-1. In order to obtain such along cycle life, FEC proved to be a necessary additive to the electrolyte. Since As loads and unloads Li at 0.8V and 1.1V respectively, the load potential was restricted to 0.75 V, avoiding much of the SEI

176 formation that occurs at lower potential and essentially eliminating the possibility of Li plating, a significant safety hazard for conventional Li-ion cells. The As was reduced directly on the MWCNTs with alkalide reduction to maximize the capacity. By doing so the capacity was increased to 626 mAh g-1 over 200 cycles. These results are the best reported for Li-As half-cells to date.

To expand upon this further, GaAsnano was synthesized by alkalide reduction and compared with micron GaAs. While GaAs has a slightly lower capacity that elemental

As, Ga has been explored as a self-healing Li alloying anode material, which may lead to more stable cycling performance. Although micon sized GaAs faded rapidly,

-1 GaAsnano:MWCNT:PAA (45:45:10) electrodes maintained 605 mAh g over 100 cycles.

Compositions with less MWCNTs (68:22:10), faded over 80 cycles to 218 mAh g-1.

In order to see long term improvements in Li-ion batteries incorporation of nanomaterials seems crucial. It is well known that size restriction is an effective way to allow alloying anode nanoparticles to function with long term capacity retention.

Inclusion of alloying anode nanoparticles such as Sn, As, and GaAs into carbon materials such as MGNS or MWCNTs can lead to higher capacity, longer cycle life, and better rate capabilities. Based on the results described herein, these materials should be considered for further investigation.

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7. Future Work

While considerable gains were made with the addition of Sn, As, and GaAs to carbon matrices, such as MGNS and MWCNTs, there is more to be learned in these

-1 materials. SnO2@MGNS faded from a composite specific capacity of 390 mAh g to

250 mAh g-1. The buildup of SEI likely contributed to the capacity loss in this material.

The incorporation of fluorinated ethylene carbonate in electrolyte showed improvement in the capacity retention in Asnaoo/MWCNT electrodes and could be applied to this system as well. In addition to using a more advanced electrolyte for this material, a combination of the synthesis of Sn@MGNS and SnO2/MGNS could lead to a material with a higher volumetric energy density utilizing both the space inside and outside the

MGNS. This could be achieved by first synthesizing Sn@MGNS and then performing a

NaBH4 reduction of SnCl2 in their presence, resulting in a SnO2/Sn@HCNS composite.

The study of nanoparticulate As synthesized by akalide reduction was the first of its kind. The performance of the nanoscale material reported here far exceeds that of the only As anode previously reported, made by mechanically milling of bulk As followed by mixing with carbon additives. Even so, there is considerably more to be learned about

As as an alloying anode material. As is known to form Li3As upon complete electrochemical loading of Li, but little is known of the other phases formed during the electrochemical process. A study by Park et al.[1] gave some insight into the formation of Li2As, but there is still little evidence of this occurring with nanoscale material because of the amorphous nature of As when reduced to small size. The ability to perform uninterrupted in situ studies during the charge/discharge process of this material

178 would be of huge significance. Performing elemental analysis or powder XRD during galvanostatic cycling could give elemental and structure data not observed upon opening the cell. In addition to amorphous As, other allotropes do exist. “Black” As, a puckered layered structure that is isostrutural with black P, exists under the right conditions. Since this material has a layered structure similar to graphite, there is possibilities for increased rate performance due to ease of lithiation. Furthermore, the electrical conductivity of the material increases along these puckered planes. In addition, the puckering results in layer spacing of black As (~0.5nm) [2] that is greater than that of graphite (~0.33nm). The larger interlayer spacing may make it possible to use As for Na+ ion batteries. Na+ is roughly 40 pm larger than Li+ and will not intercalate into graphite, but may intercalate into the large spacing of black As. The synthesis of a black AsP alloy has also been demonstrated, and the amount of As can be tuned in order to vary properties such as bandgap.[3] This material could also be applied to the Na+ batteries in order to limit the toxicity of the system by reducing the overall amount of As in the battery.

7.2 References

1. Park, C.M. and H.J. Sohn, Quasi-Intercalation and Facile Amorphization in Layered ZnSb for Li-Ion Batteries. Advanced Materials, 2010. 22(1): p. 47. 2. Liu, B., et al., Black Arsenic–Phosphorus: Layered Anisotropic Infrared Semiconductors with Highly Tunable Compositions and Properties. Advanced Materials, 2015. 27(30): p. 4423-4429.

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