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- Herausgeber: John Wiley & Sons
- Kategorie: Wissenschaft und neue Technologien
- Sprache: Englisch
A guide to lithium sulfur batteries that explores their materials, electrochemical mechanisms and modelling and includes recent scientific developments Lithium Sulfur Batteries (Li-S) offers a comprehensive examination of Li-S batteries from the viewpoint of the materials used in their construction, the underlying electrochemical mechanisms and how this translates into the characteristics of Li-S batteries. The authors - noted experts in the field - outline the approaches and techniques required to model Li-S batteries. Lithium Sulfur Batteries reviews the application of Li-S batteries for commercial use and explores many broader issues including the development of battery management systems to control the unique characteristics of Li-S batteries. The authors include information onsulfur cathodes, electrolytes and other components used in making Li-S batteries and examine the role of lithium sulfide, the shuttle mechanism and its effects, and degradation mechanisms. The book contains a review of battery design and: * Discusses electrochemistry of Li-S batteries and the analytical techniques used to study Li-S batteries * Offers information on the application of Li-S batteries for commercial use * Distills years of research on Li-S batteries into one comprehensive volume * Includes contributions from many leading scientists in the field of Li-S batteries * Explores the potential of Li-S batteries to power larger battery applications such as automobiles, aviation and space vehicles Written for academic researchers, industrial scientists and engineers with an interest in the research, development, manufacture and application of next generation battery technologies, Lithium Sulfur Batteries is an essential resource for accessing information on the construction and application of Li-S batteries.
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Table of Contents
Cover
Preface
Part I: Materials
1 Electrochemical Theory and Physics
1.1 Overview of a LiS cell
1.2 The Development of the Cell Voltage
1.3 Allowing a Current to Flow
1.4 Additional Processes Which Define the Behavior of a LiS Cell
1.5 Summary
References
2 Sulfur Cathodes
2.1 Cathode Design Criteria
2.2 Cathode Materials
2.3 Cathode Processing
2.4 Conclusions
References
3 Electrolyte for Lithium–Sulfur Batteries
3.1 The Case for Better Batteries
3.2 Li–S Battery: Origins and Principles
3.3 Solubility of Species and Electrochemistry
3.4 Liquid Electrolyte Solutions
3.5 Modified Liquid Electrolyte Solutions
3.6 Solid and Solidified Electrolyte Configurations
3.7 Challenges of the Cathode and Solvent for Device Engineering
3.8 Concluding Remarks and Outlook
References
4 Anode–Electrolyte Interface
4.1 Introduction
4.2 SEI Formation
4.3 Anode Morphology
4.4 Polysulfide Shuttle
4.5 Electrolyte Additives for Stable SEI Formation
4.6 Barrier Layers on the Anode
4.7 A Systemic Approach
References
Part II: Mechanisms
5 Molecular Level Understanding of the Interactions Between Reaction Intermediates of Li–S Energy Storage Systems and Ether Solvents
5.1 Introduction
5.2 Computational Details
5.3 Results and Discussions
5.4 Summary and Conclusions
Acknowledgments
References
6 Lithium Sulfide
6.1 Introduction
6.2 Li
2
S as the End Discharge Product
6.3 Li
2
S‐Based Cathodes: Toward a Li Ion System
6.4 Summary
References
7 Degradation in Lithium–Sulfur Batteries
7.1 Introduction
7.2 Degradation Processes Within a Lithium–Sulfur Cell
7.3 Capacity Fade Models
7.4 Methods of Detecting and Measuring Degradation
7.5 Methods for Countering Degradation
7.6 Future Direction
References
Part III: Modeling
8 Lithium–Sulfur Model Development
8.1 Introduction
8.2 Zero‐Dimensional Model
8.3 Modeling Voltage Loss in Li–S Cells
8.4 Higher Dimensional Models
8.5 Summary
References
9 Battery Management Systems – State Estimation for Lithium–Sulfur Batteries
9.1 Motivation
9.2 Experimental Environment for Li–S Algorithm Development
9.3 State Estimation Techniques from Control Theory
9.4 State Estimation Techniques from Computer Science
9.5 Conclusions and Further Directions
Acknowledgments
References
Part IV: Application
10 Commercial Markets for Li–S
10.1 Technology Strengths Meet Market Needs
10.2 Electric Aircraft
10.3 Satellites
10.4 Cars
10.5 Buses
10.6 Trucks
10.7 Electric Scooter and Electric Bikes
10.8 Marine
10.9 Energy Storage
10.10 Low‐Temperature Applications
10.11 Defense
10.12 Looking Ahead
10.13 Conclusion
11 Battery Engineering
11.1 Mechanical Considerations
11.2 Thermal and Electrical Considerations
References
12 Case Study
12.1 Introduction
12.2 A Potted History of Eternal Solar Flight
12.3 Why Has It Been So Difficult?
12.4 Objectives of HALE UAV
12.5 Worked Example – HALE UAV
12.6 Cells, Batteries, and Real Life
12.7 A Quick Aside on Regenerative Fuel Cells
12.8 So What Do We Need from Our Battery Suppliers?
12.9 The Challenges for Battery Developers
12.10 The Answer to the Title
12.11 Summary
Acknowledgments
References
Index
End User License Agreement
List of Tables
Chapter 3
Table 3.1 Electrolyte effects of cell cycling.
Table 3.2 Cell characteristics.
Chapter 5
Table 5.1 Computed solution phase enthalpies of reaction (Δ
H
rxn
) and activatio...
Table 5.2 Computed reaction barriers (kcal mol
−1
) for cleaving the “CO” (...
Chapter 6
Table 6.1 Comparison of electrochemical performance of the recently developed Li
Chapter 8
Table 8.1 Parameters for the 0D model.
Chapter 9
Table 9.1 Specifications of OXIS lithium–sulfur cell [6].
Table 9.2 Rule base of a FIS for Li–S cell SoC estimation.
List of Illustrations
Chapter 1
Figure 1.1 Structure of a LiS cell, indicating the two electrodes, sepa...
Figure 1.2 Scheme for the relevant energy levels in a conductive materi...
Figure 1.3 (a) Connecting two conductors allows electrons to move from ...
Figure 1.4 Diagram of the simplified electrode–electrolyte system. Mobi...
Figure 1.5 (a) Example of an oxidation process, in which an anionic spe...
Figure 1.6 (a) Diagram of the structure of the double layer and (b) rep...
Figure 1.7 (a) Assuming the probes of the voltmeter are electrochemical...
Figure 1.8 According to transition state theory, the oxidant or reducta...
Figure 1.9 Indication of two paths for the internal flow of current thr...
Figure 1.10 Representative resistance variation of a LiS cell as a func...
Chapter 2
Figure 2.1 (a) Scheme of a Li–S cell stack and its components. (b) Sche...
Figure 2.2 (a) Weight distribution of cell components on stack level (e...
Figure 2.3 Morphologies of different carbons: (a) hard carbons, (b) car...
Figure 2.4 Schematic macro‐ and microscopic structure of a carbon/sulfu...
Figure 2.5 Scheme of different polysulfide retention principles: (a) Ni...
Chapter 3
Figure 3.1 Typical discharge–charge curve for a lithium–sulfur cell. ...
Chapter 5
Figure 5.1 Schematic of probable (electro)chemical reactions of polysul...
Figure 5.2 Schematic of the preferred nucleophilic reaction sites of et...
Figure 5.3 Computed activation enthalpies (Δ
H
†
, kcal mol
−1
...
Figure 5.4 Computed geometries of selected complexes formed from the re...
Figure 5.5 Three‐dimensional structures of selected transition state st...
Figure 5.6 Computed enthalpies of activation (kcal mol
−1
) require...
Figure 5.7 Schematic of the fluorine‐substituted dimethoxy ethane (DME)...
Figure 5.8 The computed “CF” and “CO” bond breaking transition state ...
Chapter 6
Figure 6.1 Discharge voltage profile of a Li–S cell with three characte...
Figure 6.2 Voltage profiles (black) and surface temperature (blue) of a...
Figure 6.3 Proposed crystalline structures of the Li
2
S
2
compound and th...
Figure 6.4 Discharge voltage profile of a 3.4 Ah cell (OXIS Energy), cy...
Figure 6.5 Proposed reaction mechanism of a Li–S cell cycled at C/5 rat...
Figure 6.6 Schematic illustration of Li
2
S precipitation depending on th...
Figure 6.7 (a) Schematic representation of Li
2
S deposition on a carbon ...
Figure 6.8 Schematic illustration of the nucleation and growth of Li
2
S
2
Figure 6.9 Evolution of a (111) XRD peak of a Li
2
S compound during its ...
Figure 6.10 Discharge voltage profiles (a) of a 20 Ah Li/S cell (OXIS E...
Figure 6.11 First and second charge (a) and discharge (b) profiles of a...
Figure 6.12 Schematic illustration of the proposed mechanism of the ini...
Figure 6.13 (a–c) GITT results applied to study overpotential phenomeno...
Figure 6.14 (a) Voltage profiles comparison showing the consequence of ...
Figure 6.15 (A) Schematic of the Li
2
S particles coating to obtain Li
2
S@...
Figure 6.16 Voltage behavior measured in a three‐electrode cell with Li...
Chapter 7
Figure 7.1 Discharge capacity of a cell cycled at a constant C‐rate of ...
Figure 7.2 Voltage profiles of certain cycles of the previous cell. (a)...
Figure 7.3 Close‐up of features observed in a charge curve. (a) Shows t...
Figure 7.4 Schematic of degradation processes within a lithium–sulfur c...
Figure 7.5 The effect of sulfur loading on discharge capacity cycle lif...
Figure 7.6 The effect of constant current cycling at different rates on...
Figure 7.7 Evaluation of the discharge capacity (black), charge capacit...
Figure 7.8 Voltage profile of cell with and without shuttle.
Figure 7.9 Sulfur utilization as a function of electrolyte loading [18...
Figure 7.10 Effect of C‐rate on capacity of lithium–sulfur cells over m...
Figure 7.11 Rate capability of different lithium–sulfur cells as report...
Figure 7.12 The rate capability of lithium–sulfur cells with and withou...
Figure 7.13 Charge and discharge capacities for lithium–sulfur cells cy...
Figure 7.14 (a) OXIS prototype cell with greater than 30 layers (b) OXI...
Figure 7.15 Comparison of specific peak power of two cells with similar...
Figure 7.16 A schematic of different types of capacity fade curves obse...
Figure 7.17 Charge and discharge curves of lithium–sulfur cells with va...
Figure 7.18 Evolution of resistance in cells over cycling.
Figure 7.19 Voltage and thickness evolution of a 21 Ah cell cycled in a...
Figure 7.20 Summary of all the methods for improving capacity and perfo...
Chapter 8
Figure 8.1 (a) Simulated constant‐current discharge curves at 1.7 A (0....
Figure 8.2 (a) Cell voltage during a 1 and a 0.1 C charge. (b, c) The c...
Figure 8.3 Voltage and Ohmic resistance during a 0.34 A discharge of a ...
Figure 8.4 (a) Simulated discharge voltage and Ohmic resistance of a Li...
Figure 8.5 Simulated anode potential during 0.1 and 0.3 C discharges.
Figure 8.6 (a) Simulated cell voltage during a 0.2 C discharge and 0.1 ...
Figure 8.7 (a) Simulated discharge curves at 0.2 and 0.6 C with the tra...
Figure 8.8 Top: Measured voltage of an OXIS Energy 3.4 Ah Li–S pouch du...
Figure 8.9 Schematic of a Li–S cell at different scales.
Chapter 9
Figure 9.1 Theoretical capacity and usable capacity. At extremes of sta...
Figure 9.2 Example of variation in capacity with voltage‐based discharg...
Figure 9.3 Determining the state of charge from open‐circuit voltage (O...
Figure 9.4 Li–S cell test equipment; Kepco BOP device and Li–S cell ins...
Figure 9.5 Load current and terminal voltage during a pulse discharge t...
Figure 9.6 Li–S cell discharge test based on UDDS driving cycle.
Figure 9.7 Principle of model‐based estimation. Source: Propp 2017 [38...
Figure 9.8 Input–output relationship of the mechanistic model. Source: ...
Figure 9.9 Terminal voltage predicted by mechanistic model. Source: Pro...
Figure 9.10 Input–output relationship of the ECN model. Source: Propp 2...
Figure 9.11 Terminal voltage predicted by the ECN model.
Figure 9.12 Principle of Kalman filter.
Figure 9.13 Estimation results with EKF, UKF, and PF and two different ...
Figure 9.14 Modeling of a battery using the black‐box technique.
Figure 9.15 Input (current) – output (voltage) recording during a Li–S ...
Figure 9.16 High voltage membership function.
Figure 9.17 Li–S cell's (a) OCV, (b) ohmic resistance, and (c) its deri...
Figure 9.18 Generalized bell‐shaped MF tunable parameters.
Figure 9.19 Battery measurement, identification, and SoC estimation.
Figure 9.20 ANFIS structure including 4 inputs, 15 MFs, and 150 rules. ...
Chapter 11
Figure 11.1 The classic “volcano”‐shaped high‐frequency resistance
R
1
Figure 11.2 The current inhomogeneities caused by three single layer po...
Chapter 12
Figure 12.1 Zephyr 7 in flight.
Figure 12.2 Cloud top height statistics [11].
Figure 12.3 Average wind speeds at 250 mbar (40 000 ft).
Figure 12.4 Average wind speed as a function of time of year and altitu...
Figure 12.5 Average wind speed as a function of time of year and altitu...
Figure 12.6 Minimum sustained altitude required as a function of latitu...
Figure 12.7 Energy available and energy required at winter solstice.
Figure 12.8 Population density as a function of latitude.
Figure 12.9 Indicative cell Wh kg
−1
requirement for year‐round op...
Guide
Cover
Table of Contents
Begin Reading
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Lithium-Sulfur Batteries
Edited by
Mark Wild
OXIS Energy, E1 Culham Science Centre, Abingdon, UK
Gregory J. Offer
Department of Mechanical Engineering, Imperial College London, London, UK
Copyright
This edition first published 2019
© 2019 John Wiley & Sons Ltd
All rights reserved. No part of this publication may be reproduced, stored in a retrieval system, or transmitted, in any form or by any means, electronic, mechanical, photocopying, recording or otherwise, except as permitted by law. Advice on how to obtain permission to reuse material from this title is available at http://www.wiley.com/go/permissions.
The right of Mark Wild and Gregory J. Offer to be identified as the authors of the editorial material in this work has been asserted in accordance with law.
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In view of ongoing research, equipment modifications, changes in governmental regulations, and the constant flow of information relating to the use of experimental reagents, equipment, and devices, the reader is urged to review and evaluate the information provided in the package insert or instructions for each chemical, piece of equipment, reagent, or device for, among other things, any changes in the instructions or indication of usage and for added warnings and precautions. While the publisher and authors have used their best efforts in preparing this work, they make no representations or warranties with respect to the accuracy or completeness of the contents of this work and specifically disclaim all warranties, including without limitation any implied warranties of merchantability or fitness for a particular purpose. No warranty may be created or extended by sales representatives, written sales materials or promotional statements for this work. The fact that an organization, website, or product is referred to in this work as a citation and/or potential source of further information does not mean that the publisher and authors endorse the information or services the organization, website, or product may provide or recommendations it may make. This work is sold with the understanding that the publisher is not engaged in rendering professional services. The advice and strategies contained herein may not be suitable for your situation. You should consult with a specialist where appropriate. Further, readers should be aware that websites listed in this work may have changed or disappeared between when this work was written and when it is read. Neither the publisher nor authors shall be liable for any loss of profit or any other commercial damages, including but not limited to special, incidental, consequential, or other damages.
Library of Congress Cataloging‐in‐Publication Data
Names: Wild, Mark, 1974‐ editor. | Offer, Gregory J., 1978‐ editor.Title: Lithium–Sulfur batteries / edited by Mark Wild, OXIS Energy, E1 CulhamScience Centre, Abingdon, UK, Gregory J. Offer, Mechanical Engineering, ImperialCollege London, London, UK.Description: First edition. | Hoboken, NJ, USA : John Wiley & Sons, Inc.,2019. | Includes bibliographical references and index. |Identifiers: LCCN 2018041630 (print) | LCCN 2018043007 (ebook) | ISBN9781119297857 (Adobe PDF) | ISBN 9781119297901 (ePub) | ISBN 9781119297864(hardcover)Subjects: LCSH: Lithium–Sulfur batteries.Classification: LCC TK2945.L58 (ebook) | LCC TK2945.L58 L57 2019 (print) |DDC 621.31/2424–dc23LC record available at https://lccn.loc.gov/2018041630
Cover design: Wiley
Cover image: Courtesy of OXIS Energy
Preface
In 2014, a team from industry and academia came together to develop a Revolutionary Electric Vehicle Battery and Energy Management System (REVB) funded by the EPSRC and Innovate UK. The project brought together material scientists, electrochemists, physicists, mathematicians, and engineers from OXIS Energy Ltd, Imperial College London, Cranfield University, Lotus Engineering, and Ricardo PLC. In 2016, the team held the first lithium–sulfur conference in the United Kingdom in the Faraday lecture theatre of the Royal Institution in London – a place where scientists, artists, authors, and politicians have shared ideas for over 200 years with the aim of diffusing science for the common purpose of life. The conference was named LiS–M3 and continues annually. M3 stood for materials, mechanisms, and modeling; we also held a fourth session for applications of Li–S technology. Following that first conference, the two conference chairs, Dr. Gregory J. Offer from Imperial College and Dr. Mark Wild from OXIS Energy, were approached to edit this book. The chapters have been provided by those that gave presentations, were invited that day, or were part of the REVB team in the Faraday lecture theatre.
The organization of this book follows the structure of the conference and has the same aim of educating a diverse scientific community about the most promising next‐generation Li–S battery technology, enabling applications requiring batteries with superior gravimetric energy density. Lithium–sulfur batteries are game changers in the world of lightweight energy storage with a theoretical gravimetric energy density of ∼2600 Wh kg−1. Yet, there are challenges, and today the practical energy density target is 500 Wh kg−1.
Materials
.
In Part I we start with basic electrochemical theory to understand the challenges and complexity of lithium–sulfur batteries, and then focus on the approaches by material scientists to overcome those challenges. It soon becomes clear that there are no silver bullets, but that a systems approach is required to increase the areal loading of sulfur in a stable cathode, to increase sulfur utilization through the electrolyte/cathode interface and to reduce degradation at the electrolyte/anode interface. It is also evident that lithium–sulfur technology has reached a point in its development that it can now be tailored to meet the needs of commercial markets such as aviation, marine, or automotive.
Mechanisms
.
Part II considers the current understanding of the complex mechanisms in a lithium–sulfur cell. There remains an incomplete understanding of the mechanism from materials research, and analytical studies only see part of the picture. Elucidating the mechanism of a lithium–sulfur cell is complex and intriguing. There are many studies that have opened windows onto the association and disassociation reactions of the lithium polysulfides and the precipitation and dissolution of solid products at the end of charge and discharge. These underlying mechanisms lead to the unique discharge and charge characteristics and degradation pathways. Included are chapters on polysulfide reactivity and an enlightening look at the lithium–sulfur cell from the perspective of its insoluble end product, lithium sulfide.
Modeling
.
Part III starts by looking broadly at physics‐based models that mimic and predict the performance and degradation of a lithium–sulfur cell under operational conditions. The section concludes with control models used to predict state of charge and state of health in real‐time battery management systems. Modeling requires knowledge of the mechanism (Part II) and performance characteristics of the technology (Part I) and is used to develop the control algorithms and working models required by engineers developing applications (Part IV).
Applications
.
Part IV addresses the commercial application of lithium–sulfur battery technology. It starts with a market analysis, takes in key differences that battery engineers must be aware of in the design of a lithium–sulfur battery, and concludes with the first real‐world application of lithium–sulfur batteries, the
high‐altitude long‐endurance unmanned aerial vehicle
(
HALE–UAV
).
As a guide to access the book, each part begins with its own introduction and also each chapter. If you are looking for a good overview, then start with the part introductions. If you are a material scientist, then start with Part I and continue with mechanisms in Part II; Chapter 10 may also be of interest to identify a target market. If you are an applications engineer you might like to start with modeling in Part III and move to applications in Part IV. If you are interested in modeling lithium–sulfur cells then start with Chapter 2 of Part I and then move to Parts II and III.
Mark Wild
OXIS Energy, Abingdon, UK
Part IMaterials
Lithium–sulfur cells utilize a very similar architecture as today's Li ion pouch cells. Double side coated cathodes (sulfur) are assembled with layers of separators and anodes (lithium foil) either through winding or electrode stacking and subsequently vacuum packaged into a pouch (aluminum/polymer laminate foil). In Li–S cells, cathodes are typically assembled in the charged state with lithium metal as anode.
At first glance, the combination of the lightest, most electropositive metal (lithium) with a safe, abundant (and reasonably light) nonmetal (sulfur) makes good sense as a prospective battery. However, while the lithium–sulfur battery offers a very high theoretical specific energy (∼2600 Wh kg−1) the actual performance delivered is proving to be limited and today a gravimetric energy density target of 500 Wh kg−1 is thought to be an achievable step change in battery performance with this technology.
Materials research lies at the heart of lithium–sulfur cell development and relies on a good understanding of the underlying mechanisms (Part II). The game changer is to achieve a lightweight battery with sufficient cycle life and power performance for relevant applications (Part IV) where the weight of large Li ion battery systems hinders product performance, e.g. aircraft and large vehicles. The goal is to increase the ratio of active sulfur to inactive, yet functional, materials in the cell and to make the best use of this sulfur by achieving the highest sulfur utilization cycle on cycle.
In Chapter 1, we begin with a grounding in basic electrochemical theory. We explore how basic theory translates to a more complex electrochemical system such as a lithium–sulfur cell. The chapter concludes with a theoretical explanation of the main challenges faced by materials research scientists developing commercial lithium–sulfur products.
In Chapter 2, we move on to a discussion of the sulfur cathode, where due to its nonconductive nature, sulfur is most often combined with carbons, additives, and binders to be coated onto a primed aluminum current collector. Even when optimized for high areal sulfur loading, cathode materials contribute to reduced gravimetric energy density and release reactive polysulfides into the electrolyte, leading to degradation.
In Chapter 3, we continue with a discussion of electrolytes and it will become apparent that there is an intimate relationship between sulfur loading and electrolyte loading. Stability of the electrolyte components toward both lithium and polysulfides is also critical to optimizing sulfur utilization and cycle life. A balance is to be struck between trapping polysulfides within the cathode and dissolution of polysulfides into the electrolyte to achieve acceptable energy density, cycle life, and power.
In Chapter 4, we briefly summarize the electrolyte anode interface and the key challenges. This is an area that is poorly covered by the academic literature but is a vital area of research to improve the cycle life of a lithium sulfur cell in tandem with other approaches. Throughout all chapters and in Chapter 7 we make reference to the role of anode in relation to degradation and reduced cycle life. Primarily efforts have included the use of a range of barrier layers either at the cathode surface, as a modification to the separator or as a polymer, or at the ceramic coating on the lithium itself in addition to optimization of cathode and electrolyte formulations.
Not to be lost is the shift in lithium–sulfur cell development to choose the thinnest and lightest components to reduce gravimetric energy density. Such materials require special consideration during scale‐up activities when compared to handling procedures in standard lithium ion manufacture.
1Electrochemical Theory and Physics
Geraint Minton
OXIS Energy, E1 Culham Science Centre, Abingdon, Oxfordshire, OX14 3DB, UK
1.1 Overview of a LiS cell
On discharge, the overall process occurring in a lithium–sulfur (LiS) cell is the reaction of lithium and sulfur to form lithium sulfide, Li2S, according to the reaction shown in Eq. (1.1).
Although both reactants are present in the cell, its design, shown in Figure 1.1, prevents the reaction from taking place directly. The cell comprises a lithium metal electrode and a mixed carbon/sulfur (C/S) electrode. The latter is composed of a mix of highly porous carbon, which provides both electronic conductivity and an electrochemically active surface; sulfur, which is the active material in this electrode; and binder, which holds the structure together. The two electrodes are divided by a separator material, which stops the active materials from making direct contact and also prevents electrons passing internally between the electrodes. Contact between the active materials in each electrode is made indirectly, via an electrolyte, which is in contact with the lithium electrode and permeates the separator and C/S electrode. The electrolyte is composed of a solvent in which a lithium salt, plus any additives, has been dissolved. Adjacent to the C/S electrode is a current collector material to facilitate the flow of electrons to and from and external circuit, a task which is performed by the lithium metal on the other side of the cell.
Figure 1.1 Structure of a LiS cell, indicating the two electrodes, separator, and C/S electrode current collector. Also indicated are the reactions at each electrode and the electron pathway on discharge.
By preventing the direct contact of lithium and sulfur, the design of the cell means that the reaction in Eq. ( 1.1) takes place indirectly: one or both species have to enter the electrolyte in order to react, a process which consumes electrons from (or releases them to) the electrode surfaces. These electrochemical reactions are the fundamental process which must occur in any electrochemical cell, since without them charge would not be replenished on the electrodes when an external current flows, causing the voltage to rapidly decrease to zero. Considering the electrochemical steps occurring in a LiS cell, the overall cell reaction in Eq. ( 1.1) can be split into the following overall half‐reactions occurring at each electrode:
where the equations are written in the standard form Ox+ne−⇌Re, in which Ox is the oxidized (more positively charged) form of the species, Re is the reduced (more negatively charged) form, and e− is an electron. In order for the overall cell reaction to be satisfied, on discharge the lithium reaction must run to the left (electrochemical oxidation), with ions and electrons being formed from the lithium metal, while the C/S reaction must run to the right (electrochemical reduction), with electrons being consumed and ions being formed by the reaction of solid‐phase sulfur. In separating out the reactions in this manner, it is possible to see how the reactions generate the electronic charge in the electrodes which flows as the external current when the electrodes are connected.
The cell half‐reactions might indicate that only electrochemical reactions occur in the cell, but this is not the case. The initial state of the cell includes solid‐phase sulfur, so a dissolution step is required; and to form the precipitated Li2S, a chemical reaction is required to combine the lithium and sulfur ions:
Furthermore, as with the phase change processes at either end of the reaction mechanism, the reactions occurring at the C/S electrode do not appear to take place in a single step, as Eq. (1.2b) may suggest. Instead, a host of intermediate species are produced through both chemical and electrochemical elementary steps [1]. These intermediate species consist of sulfur anions of varying chain lengths, collectively known as polysulfides. They are commonly split into two groups: “high‐order” species are those with chain lengths of five to eight atoms, and “low‐order” species are those with chain lengths of one to four atoms.
Since the exact reaction mechanism is currently unknown, and may even vary under different operational conditions [2–5] or electrolyte compositions, this chapter does not attempt to describe how each individual process affects a LiS cell. Instead, the focus is on how different types of reaction processes contribute to (or detract from) the electrochemical performance of a cell, using a LiS as a reference point.
Before continuing, it will be useful to define how the electrodes are referred to throughout this chapter. The lithium reaction has to generate electrons on discharge, while the polysulfide species have to accept them. Since current flow is in the direction opposite to the flow of electrons, electronic current will flow from the C/S electrode to the lithium electrode, making the lithium electrode the anode on discharge. However, the direction of the current is reversed on charge, which would make the C/S electrode the anode; and if the cell is at rest, then there is no anode. In order to simplify the terminology, and because it is commonly the discharge which is of more interest, the lithium electrode will herein be referred to as the anode, regardless of the direction (or presence) of a current, and the C/S electrode will be referred to as the cathode.
1.2 The Development of the Cell Voltage
The purpose of an electrochemical cell is to drive an electric current through a circuit which is connected across the terminals of the cell. This flow of current occurs spontaneously: naively electrons move from the electrode with the lower electric potential to the one with the higher electric potential, causing the difference between the two electrode potentials to decrease. Thus, maintaining a current requires electrons to be spontaneously generated at the lower potential electrode and consumed at the higher potential electrode; otherwise, the electrode potentials would equilibrate and the flow of electrons would stop.
Thermodynamically, a process will occur spontaneously if it lowers the free energy of a system, where the free energy is the internal energy of the system minus that part of the internal energy which can do no useful work. In the case of an electrochemical cell, the system is the electrochemical cell and the external circuit, and a process is anything occurring within the system, for example the conversion of one species to another via a reaction, or the trend for a species to move in one direction or another. A useful aspect to remember when discussing processes is that we are always considering the net outcome of a very large number of individual random events, some of which increase the free energy and some of which lower it. The “direction” represents how this stochastic process is biased, and this bias is always in the direction which minimizes the free energy, because it is more likely that a particle in a higher energy state will move to a lower energy state than the reverse occurring.
Coupled with the notion of the events occurring at random is the fact that the processes do not stop when equilibrium is reached – for example, individual particles in a gas do not stop moving just because there is no concentration gradient. Instead, equilibrium implies that the bias to the process has been removed – the number of particles in the gas moving to the left is now equal to the number moving to the right, so there is no net change in the concentration.
There are two forms of the free energy commonly used to describe processes in an electrochemical system: the Gibbs free energy G, typically used when discussing the reaction processes, and the Helmholtz free energy F, commonly used when discussing the electrolyte composition [6–9] and species transport. How they are related and defined is beyond the scope of this chapter and, since we will not consider volume or pressure changes in the cell, the two are essentially equivalent. Related to the free energy is the electrochemical potential μi of the species in the system, which is also often more convenient to work with [10,11]. This term represents the change that occurs in the free energy of a system when a particle of type i is added to a point in the system from a point outside of it, and can be used to derive expressions for the system behavior. Under the assumption that there are a large number of particles in the system [12], the electrochemical potential is determined by the following relationships:
where Ni is the number of molecules of type i and the subscripts indicate the properties held constant during the differentiation: p is the pressure, T is the temperature, V is the volume, and Nj≠i is the quantity of all species except species i. The notion of how the electrochemical potential relates to the overall direction a process occurs is summarized in the following two examples:
Particle movement can be thought of as the process of taking a particle of type 1 from point A inside the system, moving it outside of the system, and then placing it back in the system at a different point, B. The removal process changes the free energy of the system by
and the addition step changes it by
. If
, the free energy is reduced by the particle's movement, so the bias for species movement is for particles to move from A to B.
A reaction process effectively removes a particle of type 1 from point A in the system and replaces it with a particle of type 2 (or vice versa). If
, the free energy of the system is reduced by the formation of particle 2 and so this direction is preferred. If not, the reverse process is preferred and particle 1 tends to be formed from particle 2.
There are two types of particle that we are typically interested in when considering an electrochemical system: electrons in the electrodes and particles in the electrolyte, itself comprising ions, neutral species, and solvent molecules. It is common practice when modeling electrochemical systems to treat the solvent molecules as a continuum dielectric background through which the remaining particles in the electrolyte move and interact [13–16], an approach also taken throughout this discussion. The electrochemical potentials of the electrons and the solvated mobile species can be written as [17,18]
where Ef is the energy of the thermodynamic Fermi level of the electrode, discussed subsequently, e0 is the elementary charge, kB is the Boltzmann constant, T is the temperature and, for species i, ci is its concentration, is its excess chemical potential, and zi is its valence. is the chemical potential of species i in the standard state, defined as an ideal solution of that species at the standard concentration c⊖=1 M, at a temperature of 298 K and a pressure of 1 bar. Finally, φ is the local electrostatic potential, defined relative to some fixed reference. For the species in solution, the logarithmic term represents the ideal, or entropic, contribution to the electrochemical potential, the excess chemical potential describes all non‐electrostatic contributions, and the final term describes the contribution from the electrostatic potential.
1.2.1 Using the Electrochemical Potential
To give an example of how the electrochemical potential is the quantity of interest in an electrochemical system, we first look at the behavior of conductive materials. The energy of an electron in a material is represented by the energy of the Fermi level Ef, which is the energy level within the band structure of the material which has a 50% probability of occupancy by electrons. The Fermi level is defined by the kinetic Fermi energy Ek, which is a function of the effective electron masses and their number density, and which sits at a constant position relative to the conduction band edge [17,19]. Because of the differing band structures of different materials, their Fermi level energies will differ.
The other quantity of interest for a material is its work function, which is the amount of energy required to excite an electron from the Fermi level to the vacuum, defined as a point just outside the material where the electric field is zero, such that the electron ends the process with zero kinetic energy. The electric potential at this point is referred to as the vacuum potential. The work function depends on factors such as the crystal face through which the electron exits the material and defects in the surface structure, meaning that its value is different for different surfaces of the material.
In the bulk material, the Fermi level is constant, meaning that an electron added to the material from a vacuum state will have the same final energy, regardless of which surface it enters through. Since the final state must be the same in both cases but the change in energy to get there (the negative of the work function) differs, the implication is that the vacuum potentials at different surfaces are not equivalent. This is represented in Figure 1.2a, showing the positions of the conduction band and Fermi level and the positions of the vacuum potential at two surfaces. Also indicated is the internal potential of the material in the bulk near each surface. We assume a conductive material, which implies that φ1=φ2.
Figure 1.2 Scheme for the relevant energy levels in a conductive material. (a) The Fermi level energy lies a fixed distance above the conduction band edge, CB, but differences in the work functions across different surfaces mean that the vacuum potential is not constant with respect to Ef. Choosing one vacuum potential as the reference point, the energy required to move a charge to the Fermi Level is always the same because of the work which has to be done against the external potential difference between the vacuum potentials, Δφvac. (b) The CB, Fermi level energy, and electrostatic potentials in two disconnected conductors may all be different.
An electron at surface 1, with potential , will lose energy upon moving to the Fermi level, while an electron at surface 2, with potential , will lose a different amount of energy to reach the same final state. Since the final internal states are identical, this means that an electron which starts at surface 1 but enters the material via surface 2 must undergo a change of energy as it moves externally around the material. This change in energy is associated with working against the electric field between the two surfaces, an electric field which must exist because of the differences between the two vacuum potentials.
We now consider the change in the electrochemical potential associated with moving an electron from surface 1 to the Fermi level via each route. Defining the reference electrostatic potential as the vacuum potential at surface 1, the change in the electrochemical potential when the electron moves directly to the Fermi level is μ1=Ef−e0(φ1−0). Alternatively, if the electron takes the indirect path, the electrochemical potential is . Since the Fermi level is constant, and there are no internal variations in the electric potential, the final electrochemical potential in both cases can be seen to be the same and the two routes are equivalent, overall.
We now consider the case of connecting two conductive materials together. Prior to contact, as indicated by Figure 1.2b, the different materials will have different conduction band edges and kinetic Fermi energies Ek, so the Fermi levels will differ. Also, because there may be an arbitrary electric field between the two phases, the internal electrostatic potentials may also be different, and so the electrochemical potential of the electrons in the two materials are not the same. When placed into contact (Figure 1.3), if material 1 has a larger Fermi energy, then there will be electron transfer from material 1 to material 2 (Figure 1.3a). As a result, material 2 will acquire a negative charge and material 1 will be left positively charged.
Figure 1.3 (a) Connecting two conductors allows electrons to move from the material with the larger Fermi level energy to the one with the lower Fermi level energy, resulting in the development of a charge‐separated layer and electric field at the interface. (b) The field defines the Galvani potential difference ΔφG between the materials and causes the conduction band edges to bend near the interface. At equilibrium, no net work is done moving a charge along the path indicated.
This separation of charge generates an electric field , which grows as more charge is transferred. The direction of the field is such that it resists the continued transfer of electrons, so the transfer process is self‐limiting: eventually, the field grows large enough to prevent further electron transfer and an equilibrium state is reached. In this state, the field is such that the potential varies smoothly from φ1, the value in the bulk of material 1, to φ2, the value in the bulk of material 2 (Figure 1.3b). The total potential difference, which can now be explicitly defined in terms of the electric field, is known as the Galvani potential, denoted here as ΔφG.
The other implication of the system being at equilibrium is that an electron moving from material 1 to material 2 should do no overall work. Note that the Fermi levels of the two bulk materials have not changed, so this element still favors the transfer of electrons between materials; it is just that this is now counterbalanced by the electric field. Any energy an electron would be able to lose by moving to the material with the lower Fermi level would be offset by the necessary energy gain required to move up the electric field at the interface.
To state this in terms of the electrochemical potentials of the two phases, a reference point is still required for the electrostatic potential, and here we choose the vacuum potential of one of the external surfaces of material 1. The work function at that surface, being far from the interface region, is largely unchanged by the interfacial field, so the electrochemical potential of an electron in the bulk of material 1 remains as . For material 2, however, the electrochemical potential is now , where φ2 can be written in terms of φ1 as φ2=φ1−ΔφG. Since the electrochemical potential must be spatially invariant at equilibrium, the previous expressions can be equated:
From this, we find that the electrostatic potential difference across the interface equals the difference in the Fermi energies of the two materials.
There are some complications in the interfacial region, where the electric potential varies spatially. In order for the electrochemical potential to be invariant in this region, the definition of the electrochemical potential implies that the Fermi level energies Ef of the two materials must also change. However, Ef is not affected by the field directly. Instead, the electric field changes the shape of the conduction band, causing it to curve toward the vacuum potential on the positively charged side and away from it on the negatively charged side. The Fermi level, which is separated from the conduction band edge by an amount equal to the kinetic Fermi energy (which is unchanged by the electric field), therefore effectively curves in the same direction as the conduction band. The resulting smooth variation in the Fermi level counterbalances the electric potential at all positions. This spatially dependent Fermi energy has been referred to as the thermodynamic Fermi energy [19].
A second observation is that while the electrostatic potential must be continuous across the interface, the changes to the Fermi levels of the two materials do not necessarily cause these to become continuous. The discontinuity is due to the work which has to be done in crossing from the crystal lattice in one material to the crystal lattice in the other, and can be demonstrated by considering the movement of an electron along the path indicated in Figure 1.3b, which must be the equivalent of an electron taking the direct route across the interface.
The potential difference between the phases is unmeasurable in all but a few situations. A voltmeter, for example, is essentially a long conducting wire, and connecting this between conductive materials would simply lead to electron rearrangement at the interface of each end of the wire with the material it is in contact with. The electrochemical potential would equilibrate throughout the entire system, which comprises material one, the wire, and material two, with a Galvani potential developing at each interface, but there would be no continuous net flow of electrons, since the free energy would quickly be minimized. However, a voltmeter requires some flow of electrons, since this does the work to turn the needle of the meter and give a reading. Consequently, no voltage would be measured in the example case.
This lack of a voltage might make us question the use of a voltmeter to measure a potential difference, but what we are generally interested in when measuring a voltage is the propensity of two phases to drive a current between them. In this case, the voltmeter is exactly the device we want; but rather than the difference in the electrostatic potential, what we are actually measuring is a difference in the electrochemical potential across the voltmeter's probes. Such a difference in electrochemical potential implies that the free energy of the system is not minimized, meaning the system is out of equilibrium. The flow of electrons through the voltmeter occurs spontaneously as the system attempts to minimize its free energy, providing the energy to turn the voltmeter's needle while also causing the electrochemical potentials to equilibrate.
Since the driving force for the current is proportional to the difference between the electrochemical potentials in the two materials, connecting the voltmeter will cause the voltage to decrease over time as the materials equilibrate. From the converse of this, we can see that if a constant voltage is measured, processes must be occurring to maintain the electrochemical potential difference across the voltmeter's probes. How the electrochemical potential difference is maintained is at the heart of the function of an electrochemical cell, as discussed in the following sections.
1.2.2 Electrochemical Reactions
The process by which an electrochemical cell is able to develop and maintain a measureable voltage is at the core of what the cell does. As with the behavior of two or more connected metals, the electrochemical potential provides a reasonably intuitive route to understanding the processes driving the cell. However, there is more scope for complexity, even while trying to keep descriptions of individual processes as simple as possible. To begin, we start by considering what happens at the single interface when a conductive electrode material is placed in contact with an electrolyte (Figure 1.4).
Figure 1.4 Diagram of the simplified electrode–electrolyte system. Mobile species in the electrolyte move through a sea of solvent dipoles and may approach the flat, smooth electrode surface.
At the microscopic level, the interfacial region between the phases is highly complex even before any electrochemical reactions have taken place – a vast number of short‐ and long‐range interactions take place between the mobile molecules which form the electrolyte and between these molecules and the surface, as well as other effects like image–charge interactions, surface adsorption, or the simple fact that the surface is not necessarily microscopically flat. In order to help build a simpler description of the system, we make a number of assumptions.
First, we assume that the only explicit surface interaction which occurs is the electrochemical reaction, and that the excess chemical potential in Eq. (1.5b) implicitly describes all other interactions that a particle experiences. Second, we assume that the surface is smooth and flat. Third, we define the reference point for the electric potential as being in the bulk electrolyte. Finally, for simplicity, we assume that the work function of the electrode surface facing the electrolyte is the same as the work function for the electrolyte surface facing the electrode. This means that the potential in the electrode, relative to the bulk electrolyte, is zero. In reality, the ions in the electrolyte would respond to the electric dipole on the electrode surface, compensating for any electric potential difference; but since this potential difference does not contribute to the operation of the cell, it can be ignored.
The electrolyte is formed by the dissolution of the general salts AB and CD into a solvent, forming a homogeneous mix of completely dissociated ions A+, B−, C+, and D−. Also present is a quantity of species E, which has no net charge and is again uniformly distributed. The electrochemical potentials of all species in the electrolyte are given by Eq. (1.5b), with , and zE=0. The species in solution and electrons in the electrode are homogeneously distributed, so there is no initial electric field and the electric potential is zero throughout the system. From Eq. (1.5a), this means that the electrochemical potential of the electrons in the electrode is initially equal to the Fermi level energy of the electrode material, Ef.
As previously stated, the standard form of an electrochemical reaction is
where ne is the number of electrons involved. In the example system, we assume that ions C+ and D− are inert, so there are two possible reactions of the standard type:
For both reactions, reduction makes the electrode positively charged while increasing the amount of negative charge in the electrolyte, either by replacing positive ions with neutral species or replacing neutral species with negative ions. Conversely, oxidation, the process shown in Figure 1.5a, makes the electrode negatively charged while increasing the amount of positive charge in the electrolyte. Note that charge is conserved at the reaction site, which means that the overall electrochemical cell remains charge neutral throughout its operation. However, local charge densities do arise either side of the electrode–electrolyte interface: the electrons have to exist in the electrode phase and ions have to exist in the electrolyte phase, so the reaction creates a local separation of charge across the interface.
Figure 1.5 (a) Example of an oxidation process, in which an anionic species in the solution releases an electron to the electrode surface and becomes a neutral molecule. (b) The response of the ions in solution to the electric field generated by the surface charge on the electrode.
As outlined at the start of this section, a reaction process alters the free energy of the system by removing one or more particles and replacing them with one or more and others; and while the reaction is free to occur in either direction, the direction which minimizes the free energy is dominant. The change in the free energy associated with the electrochemical reaction, ΔGRXN, is the difference between the electrochemical potentials of the active species:
where the terms in the second equality are grouped by species. This thermodynamic expression tells us the overall direction in which the reaction will spontaneously occur: if ΔGRXN>0, the oxidation process will occur, making the electrode more negatively charged; while if ΔGRXN<0, the reduction process occurs, making the electrode more positively charged. As with all thermodynamic expressions, however, it does not tell us the rate of a reaction, although a range of methods are available for determining this 20–23. For our purposes, it is sufficient to know that the reaction rate increases with the magnitude of ΔGRXN. Essentially, the further the system is from equilibrium, the faster it tries to move toward it.
As well as determining the direction that the reaction spontaneously occurs, it can be seen that the magnitude of ΔGRXN will always tend to zero as the reaction proceeds, because the reaction reduces μi for the species being consumed while increasing μi for the species being produced, regardless of the direction. When ΔGRXN reaches zero, the reaction reaches equilibrium. At this point the net reaction rate is zero, although this is not because no reaction events occur, but because the frequency of reaction events in either direction is the same.
While the reaction is taking place, the excess of the reaction product and deficit of the reactant near the surface, relative to the bulk, ordinarily drives a species flux due to diffusion: the product moves away from the surface to the bulk, where the concentration is smaller, and the reactant moves toward the surface from the bulk, where the concentration is larger. While the reaction is occurring, the diffusion component of the flux acts to prevent the accumulation of product or depletion of reactant at the surface, which would otherwise quickly cause ΔGRXN to shrink at the surface, significantly slowing the reaction.
The diffusion flux is countered by a migration flux caused by the electric field generated by the separation of charge at the interface. The electric field acts on all charged particles in the system, causing cations to move down the field (to lower potentials) and anions to move up the field, as shown in Figure 1.5b. It also has a number of other effects, including the alignment of solvent dipoles against the field, which, together with the presence of the ions themselves, alters the permittivity/dielectric constant of the electrolyte and therefore how the field propagates [24–27
