Metal-Air Batteries E-Book

• Herausgeber: Wiley-VCH Verlag GmbH & Co. KGaA
• Kategorie: Wissenschaft und neue Technologien
• Sprache: Englisch

Seitenzahl: 841

Veröffentlichungsjahr: 2018

Beschreibung

A comprehensive overview of the research developments in the burgeoning field of metal-air batteries

An innovation in battery science and technology is necessary to build better power sources for our modern lifestyle needs. One of the main fields being explored for the possible breakthrough is the development of metal-air batteries. Metal-Air Batteries: Fundamentals and Applications offers a systematic summary of the fundamentals of the technology and explores the most recent advances in the applications of metal-air batteries. Comprehensive in scope, the text explains the basics in electrochemical batteries and introduces various species of metal-air batteries.

The author-a noted expert in the field-explores the development of metal-air batteries in the order of Li-air battery, sodium-air battery, zinc-air battery and Mg-O2 battery, with the focus on the Li-air battery. The text also addresses topics such as metallic anode, discharge products, parasitic reactions, electrocatalysts, mediator, and X-ray diffraction study in Li-air battery. Metal-Air Batteries provides a summary of future perspectives in the field of the metal-air batteries. This important resource:

-Covers various species of metal-air batteries and their components as well as system designation
-Contains groundbreaking content that reviews recent advances in the field of metal-air batteries
-Focuses on the battery systems which have the greatest potential for renewable energy storage

Written for electrochemists, physical chemists, materials scientists, professionals in the electrotechnical industry, engineers in power technology, Metal-Air Batteries offers a review of the fundamentals and the most recent developments in the area of metal-air batteries.

Leseprobe

Table of Contents

Cover

Preface

Chapter 1: Introduction to Metal–Air Batteries: Theory and Basic Principles

1.1 Li–O

2

Battery

1.2 Sodium–O

2

Battery

References

Chapter 2: Stabilization of Lithium‐Metal Anode in Rechargeable Lithium–Air Batteries

2.1 Introduction

2.2 Recent Progresses in Li Metal Protection for Li–O

2

Batteries

2.3 Challenges and Perspectives

Acknowledgment

References

Chapter 3: Li–Air Batteries: Discharge Products

3.1 Introduction

3.2 Discharge Products in Aprotic Li–O

2

Batteries

3.3 Discharge Products in Li–Air Batteries

Acknowledgment

References

Chapter 4: Electrolytes for Li–O

2

Batteries

4.1 General Li–O

2

Battery Electrolyte Requirements and Considerations

4.2 Future Outlook

References

Chapter 5: Li–Oxygen Battery: Parasitic Reactions

5.1 The Desired and Parasitic Chemical Reactions for Li–Oxygen Batteries

5.2 Parasitic Reactions of the Electrolyte

5.3 Parasitic Reactions at the Cathode

5.4 Parasitic Reactions on the Anode

5.5 New Opportunities from the

Parasitic

Reactions

5.6 Summary and Outlook

References

Chapter 6: Li–Air Battery: Electrocatalysts

6.1 Introduction

6.2 Types of Electrocatalyst

6.3 Research of Catalyst

6.4 Reaction Mechanism

6.5 Summary

References

Chapter 7: Lithium–Air Battery Mediator

7.1 Redox Mediators in Lithium Batteries

7.2 Selection Criteria and Evaluation of Redox Mediators for Li–O

2

Batteries

7.3 Charge Mediators

7.4 Discharge Mediator

7.5 Conclusion and Perspective

References

Chapter 8: Spatiotemporal Operando X‐ray Diffraction Study on Li–Air Battery

8.1 Microfocused X‐ray Diffraction (μ‐XRD) and Li–O

2

Cell Experimental Setup

8.2 Study on Anode: Limited Reversibility of Lithium in Rechargeable LAB

8.3 Study on Separator: Impact of Precipitates to LAB Performance

8.4 Study on Cathode: Spatiotemporal Growth of Li

2

O

2

During Redox Reaction

References

Chapter 9: Metal–Air Battery:

In Situ

Spectroelectrochemical Techniques

9.1 Raman Spectroscopy

9.2 Infrared Spectroscopy

9.3 UV/Visible Spectroscopic Studies

9.4 Electron Spin Resonance

9.5 Summary and Outlook

References

Chapter 10: Zn–Air Batteries

10.1 Introduction

10.2 Zinc Electrode

10.3 Electrolyte

10.4 Separator

10.5 Air Electrode

10.6 Conclusions and Outlook

References

Chapter 11: Experimental and Computational Investigation of Nonaqueous Mg/O

2

Batteries

11.1 Introduction

11.2 Experimental Studies of Magnesium/Air Batteries and Electrolytes

11.3 Computational Studies of Mg/O

2

Batteries

11.4 Concluding Remarks

References

Chapter 12: Novel Methodologies to Model Charge Transport in Metal–Air Batteries

12.1 Introduction

12.2 Modeling Electrochemical Systems with GPAW

12.3 Second Principles for Material Modeling

Acknowledgments

References

Chapter 13: Flexible Metal–Air Batteries

13.1 Introduction

13.2 Flexible Electrolytes

13.3 Flexible Anodes

13.4 Flexible Cathodes

13.5 Prototype Devices

13.6 Summary

References

Chapter 14: Perspectives on the Development of Metal–Air Batteries

14.1 Li–O

2

Battery

14.2 Na–O

2

Battery

14.3 Zn–air Battery

References

Index

End User License Agreement

List of Tables

Chapter 04

Table 4.1 Comparison of O

2

solubility (

c

(O

2

)) and diffusivity (

D

(O

2

)) and liquid viscosity in a selection of ILs and molecular solvent‐based electrolytes, where the stated accuracy concurs with the provided reference.

Chapter 05

Table 5.1 Desired chemistries for nonaqueous Li–oxygen batteries.

Chapter 07

Table 7.1 Oxygen solubility

and TEMPO diffusion coefficient (

D

) of a 0.1 M LiTFSI–diglyme solution with different TEMPO concentrations.

Table 7.2 Reported charge mediators and their electrochemical, chemical, and diffusion properties.

Table 7.3 Electrochemical properties of the oxoammonium cation/nitroxide redox couples.

Table 7.4 Electrochemical performance of the charge mediators.

Table 7.5 A comparison of the voltage profile shape and the experiment conditions using LiI or TTF as a charge mediator.

Table 7.6 Reported discharge mediators and their electrochemical and chemical properties.

Chapter 09

Table 9.1 Raman bands for ORR discharge products in cm

−1

.

Chapter 10

Table 10.1 Strategies for improving zinc electrode performance.

Table 10.2 Important measurement and kinetic parameters for transition metal oxide electrocatalysts for OER.

Chapter 11

Table 11.1 Calculated limiting potentials, thermodynamic overpotentials, and efficiencies associated with various discharge and charge reactions in an Mg/O

2

cell.

Table 11.2 Formation energies and migration barriers for various defects in MgO and MgO

2

.

Table 11.3 Calculated diffusion coefficient, mobility, and conductivity of dominant defects in MgO and MgO

2

.

Chapter 12

Table 12.1 Charge transfer parameters for electron polaron in

‐sulfur.

Chapter 13

Table 13.1 Composition and performance of different flexible electrolytes for Li–air batteries.

List of Illustrations

Chapter 01

Figure 1.1 Schematic illustration of Li–O

2

battery based on (a) aqueous electrolyte, (b) aprotic electrolyte, (c) solid‐state electrolyte, and (d) hybrid electrolyte.

Chapter 02

Figure 2.1 Schematic illustration of the configuration of a nonaqueous Li–O

2

battery.

Figure 2.2 (a,b) SEM images of top view (a) and cross section (b) of CPL coated on Li metal surface. (c) Cyclic voltammetry curve for the electrolyte of 1 M LiClO

4

‐TEGDME with 0.05 M TEMPO additive at the carbon electrode at a scan rate of 10 mV s

−1

under O

2

gas. (d) Discharge–charge curves between cells without and with TEMPO additive. (e,f) Impedances of the Li/Li symmetric cells with bare Li (e) and with the CPL‐coated Li (f). (g,h) Discharge–charge profiles of 0.05 M TEMPO‐containing Li–O

2

cell using bare Li electrodes (g) and CPL‐coated Li electrodes (h). (i) Corresponding cycling stabilities of above cells at a discharge–charge depth of 1000 mAh g

−1

.

Figure 2.3 (a,b) SEM images of the top view (a) and cross‐sectional view (b) of Li metal coated with CPL (inset: photograph of CPL‐coated Li metal). (c,d) Discharge–charge profiles of Li–O

2

cells using Li metal (c) without CPL and (d) with CPL at 0.1 mA cm

−2

. (e,f) SEM images of the Li electrode (e) without CPL and (f) with CPL after 80 cycles (inset: photograph of cycled Li metal). (g,h) XPS spectra of (g) C 1 s and (h) S 2p of the cycled Li metal in Li–O

2

cells.

Figure 2.4 (a–c) Cycling performance of Li–O

2

cells with (a) PLM, (b) TLM, and (c) F‐TLM anodes under limited capacity of 1000 mAh g

−1

at a current density of 300 mA g

−1

. (d–f) XRD patterns obtained from (d) PLM, (e) TLM, and (f) F‐TLM anode surface at multiple cycles. (g–l) SEM images of (g) PLM, (h) TLM, (i) F‐TLM before cycles, and (j) PLM, (k) TLM, (l) F‐TLM after 60 cycles.

Figure 2.5 (a) Voltage profiles of Li/Li symmetrical cells with different electrolytes (LiTFSI/TEGDME, LiFSI/TEGDME, and LiFSI/TEGDME‐DX) at a current density of 0.25 mA cm

−2

(8 h for each cycle). (b) Extended cycling of Li/Li symmetrical cell with LiFSI/TEGDME‐DX electrolyte for 800 h and enlarged voltage curves selected from area 1, 2, 3, 4 in their above voltage profiles. (c) Nyquist plots of Li/Li cells with LiFSI/TEGDME‐DX and LiFSI/TEGDME electrolytes as a function of storage time after 5 cycles. (d) Galvanostatic discharge–charge voltage profiles of Li–O

2

cells with LiFSI/TEGDME‐DX electrolyte at a current density of 200 mA g

−1

to the fixed capacity protocol of 800 mAh g

−1

.

Figure 2.6 (a) LSV curves on carbon paper electrodes with a scan rate of 0.05 mV s

−1

under 1 atm O

2

gas. Inset: galvanostatic charging of Li–O

2

cells at 0.1 mA cm

−2

in 0.5 M LiTFSI/DMA. (b) Galvanostatic cycling in three‐electrode cell including a Super P carbon cathode, a Li metal anode, a Li reference electrode, and 1 M LiNO

3

/DMA electrolyte cycled at 0.1 mA cm

−2

(where the discharge was limited to 1 mAh cm

−2

, and the charge was limited to 3.8 V). (c) Cycling performance of Li–O

2

cell (Li anode/Super C carbon air electrode) cycled at 0.1 mA cm

−2

(in which the discharge was limited to 10 h, and the charge was limited to 4.2 V). Inset: plot of charge–discharge capacity as a function of cycle number.

Figure 2.7 (a–c) Discharge–charge voltage profiles of Li–O

2

cells with different concentration LiTFSI‐DMSO electrolytes (LiTFSI‐3DMSO (a), LiTFSI‐4DMSO (b), LiTFSI‐12DMSO (c)). (d) Their corresponding cycling performance. (e,f) The voltage profiles of Li/Li symmetric cells cycled to 1 mAh cm

−2

at 1.0 mA cm

−2

with (e) 1 M LiTFSI‐12DMSO electrolyte and (f) concentrated LiTFSI‐3DMSO electrolyte. (g–i) Cross‐sectional SEM images of the Li metal anodes after 90 cycles with (g) LiTFSI‐3DMSO electrolyte, (h) LiTFSI‐4DMSO electrolyte, and (i) LiTFSI‐12DMSO electrolyte. (j–l) Schematic illustrations of the respective components electrolytes: (j) pure DMSO solvent, (k) LiTFSI‐12DMSO dilute electrolyte, (l) LiTFSI‐4DMSO concentrated electrolyte, and (m) LiTFSI‐3DMSO highly concentrated electrolyte.

Figure 2.8 (a–c) Voltage profiles of the Li–O

2

cells with (a) 1 M, (b) 2 M, and (c) 3 M LiTFSI/DME electrolytes at 0.1 mA cm

−2

using a capacity‐limited protocol (1000 mAh g

−1

). (d) Their corresponding cycling performance.

Figure 2.9 (a) The schematic illustration of the synthesis of TBF and PF on Li metal surface. (b,c) Optical images of Li before and after exposure to air without TBF (b) and with TBF protective films (c). (d,e) The corresponding FTIR characterizations on Li metal surfaces without (d) and with TBF (e). (f,g) The cycle life of Li–O

2

cells at a current density of 200 mA g

−1

and a capacity protocol of 500 mAh g

−1

(f) and at a current density of 300 mA g

−1

and capacity protocol of 1000 mAh g

−1

(g). (h) A schematic representation of operation mechanism of TBF. (i) XRD patterns of side reaction products on Li metal anodes without and with TBF.

Figure 2.10 (a) Schematic of the Li–O

2

cell used in the work of Ref. [70]. (b) Galvanostatic discharge–charge profiles of the Li–O

2

cells with GF or AAO‐GF separators. 1 M LiNO

3

in DMAc is served as electrolyte. (c,d) In situ DEMS characterizations of Li–O

2

cell with two kinds of separators: oxygen evolution (c) and carbon dioxide evolution (d). (e) Cyclic performance of Li–O

2

cells without and with AAO separator cycled with 0.5 mA for 1000 s (blue: AAO/GF; red: GF). (f) Li CE of the Li–O

2

cells as a function of capacity. (g,h) SEM images of Li metal surface after 15 cycles without AAO (g) and with AAO (h).

Figure 2.11 (a) The schematic of addition of an interlayer between a water‐stable solid electrolyte and Li metal to efficiently protect Li (within LMP: Li metal phosphate). (b) Schematic of a Li–air battery based on a PLE technology.

Figure 2.12 (a) Schematic illustration of a Li‐ion–oxygen battery with Super P carbon as air electrode and lithiated silicon/carbon as anode. (b) Voltage profile of the Li

x

Si–O

2

full cell with carbon air electrode and lithiated Si–C anode for first cycle under a capacity protocol of 1000 mAh g

−1

. (c–f) XRD patterns of lithiated Si–C anode after discharge (c) and recharge (e) and Super P carbon electrode after discharge (d) and recharge (f). (g) Corresponding voltage profiles and (h) cycling performance of full cell cycled at 200 mA g

−1

.

Chapter 03

Figure 3.1 Schematic cell configurations for the four types of Li–O

2

batteries.

Figure 3.2 Discharge–charge profile of an aprotic Li–O

2

cell with a cutoff capacity of 1000 mAh g

−1

based on a pure Li metal anode and Super P carbon cathode. The electrolyte is 1 M lithium trifluoromethanesulfonate (LiCF

3

SO

3

) in tetraethylene glycol dimethyl ether (TEGDME), and the applied current is 100 mA g

−1

.

Figure 3.3 (a) Crystalline structure of Li

2

O

2

. (b) The experimental X‐ray diffraction (XRD) of commercial Li

2

O

2

(Sigma‐Aldrich) powders and the simulated XRD of the Föppl structure based on a wavelength of Cu Kα, 1.5406 nm [10]. (c) Different charge transport pathways predicted in amorphous Li

2

O

2

and crystalline Li

2

O

2

.

Figure 3.4 Scanning electron microscopy (SEM) images of morphology of Li

2

O

2

formed in the

Ketjenblack

(

KB

) carbon cathodes at different discharge capacities: (a) pristine cathode, (b) 250 mAh g

−1

, (c) 1300 mAh g

−1

, (d) 4000 mAh g

−1

, (e) 5200 mAh g

−1

, and (f) 7500 mAh g

−1

. The electrolyte is 0.5 M LiCF

3

SO

3

in diethylene glycol dimethyl ether (DEGDME). The applied current is 0.1 mA cm

−2

.

Figure 3.5 (a,b) TEM images of the particles discharged at 50 mA g

c

−1

with a cutoff capacity of 1000 mAh g

c

−1

and 90 mA g

c

−1

with a cutoff capacity of 13 000 mAh g

c

−1

, respectively. (c) Schematic diagram of the structure evolution of toroidal Li

2

O

2

. (d) Electron diffraction of Li

2

O

2

with simulated zone axis (red and blue dots). (e) Side‐view and top‐view schematics of a stack of crystallite plates, which compose the disc and toroid particles. The electrolyte is 0.1 M LiClO

4

in dimethoxyethane (DME) in this work.

Figure 3.6 SEM images of Li

2

O

2

obtained in (a) DMSO at high potentials (low overpotentials), (b) DMSO at low potentials (high overpotentials), (c) CH

3

CN at high potentials (low overpotentials), and (d) CH

3

CN at low potentials (high overpotentials) of 100 mM LiClO

4

. The applied current is 60 μA cm

−2

.

Figure 3.7 Crystal structure of LiO

2

. (a) The marcasite unit cell (marked with black dashed lines). Each Li is bonded with six O atoms. (b) A 3 × 2 × 3 supercell model with unit cells outlined by green lines.

Figure 3.8 (a) Typical discharge–charge profile of the Li–O

2

cell with an Ir–rGO cathode. (b) SEM image and TEM image (inset) of the discharge product on Ir–rGO cathode. Both the main panel and inset are from the discharged electrode with a current density of 100 mA g

−1

and a cutoff capacity of 1000 mAh g

−1

. (c) HE‐XRD patterns of discharge product on Ir–rGO cathode as a function of aging time. (d) Voltage plots for the Ir–rGO cathode first discharged in O

2

to a capacity of 1000 mAh g

−1

at 100 mA g

−1

and then discharged in Ar during which time it attained a capacity 956 mAh g

−1

.

Figure 3.9 Schematic diagram of the evolution of the carbon surface during charge. In the voltage profile (the lower panel), at charge, the first step is forming Li

2

CO

3

on the surface of carbon cathode and then covered by Li

2

O

2

with carbonate species dispersed inside (upper right). More carbonate species are generated on the outer surface of the undischarged Li

2

O

2

(upper middle) during charge. CO

2

is released at the end of charge when the voltage goes up to 4.2 V (upper left).

Figure 3.10 SEM images of (a) pristine carbon fiber electrode and electrodes discharged in electrolyte with (b) trace amount of H

2

O (<20 ppm), (c) 200 ppm H

2

O, and (d) 1000 ppm H

2

O. The electrolyte is 0.5 M LiTFSI in DEGDME and the applied current is 1 μA. (e) X‐ray diffraction patterns of pristine carbon fiber electrode and electrodes discharged in electrolytes with trace amount of H

2

O (<20 ppm), 1% H

2

O water, 1000 ppm H

2

O, and 200 mM HClO

4

. The straight lines represent standard patterns of Li

2

O

2

(PDF 09‐0355), LiOH (PDF 32‐0564), and LiOH·H

2

O (PDF 24‐0619). The inset shows a zoom in the diffractogram from 14.4° to 16.8° where all three compounds have their main reflections.

Figure 3.11 (a) Discharge–charge profiles of the hybrid air cell in 1 M LiClO

4

in EC‐DEC/LATP/saturated LiOH with 10 M LiCl/KB aqueous solution at 0.64 mA cm

−2

. (b) XRD patterns of the LiOH·H

2

O reference and air electrode after discharged at 0.88 mA cm

−2

for 240 h.

Figure 3.12 (a) Galvanostatic discharge profiles of Li cells discharged under three atmospheres: pure CO

2

, pure O

2

, and a 10 : 90 CO

2

 : O

2

mixture. The cathode is XC72‐based electrodes and the current is 0.47 mA cm

−2

. (b) FTIR of cathodes discharged (0.9 mA cm

−2

, 4.7 mAh cm

−2

) under pure O

2

and a 10 : 90 CO

2

 : O

2

mixture. In the pure O

2

spectra, peaks at 1064, 1138, 1202, and 1340 cm

−1

can all be attributed to the electrolyte salt (LiTFSI). P50 carbon was used as the cathode. The electrolyte is 1 M LiTFSI in DME.

Chapter 04

Figure 4.1 Comparison of solvent stability against nucleophilic attack by proton abstraction (

y

‐axis) and the tendency of a solvent to promote the solution‐based mechanism by dissolution of the adsorbed LiO

2

*

discharge intermediate (

y

‐axis). The legends provided in the corners of each quadrant describe the properties along the relative scales, consequently highlighting the bottom‐left quadrant as the most desirable position. MeOH, methanol; DMSO, dimethyl sulfoxide; DMA,

N,N

‐dimethylacetamide; NMP,

N

‐methyl‐2‐pyrrolidone; DMF,

N,N

‐dimethylformamide; MeCN, acetonitrile; DME, dimethoxyethane.

Figure 4.2 (a) First cycle discharge–charge profiles and (b) cycling stability of Li–air cells with different lithium salts in TEGDME. Cells were constructed with a lithium‐metal anode and a porous carbon cathode and were discharged at 0.05 mA cm

−2

between 2 and 4.5 V voltage limits for discharge and charge, respectively.

Figure 4.3 (a) Illustration of catalytic properties of the [NO

2

]

−

anion formed by reductive decomposition of the [NO

3

]

−

electrolyte anion. (b) Voltage profiles of carbon paper electrodes in diethylene glycol dimethyl ether (DEGDME)‐based electrolyte solvent with 1 mol dm

−3

Li[NO

3

] (black trace) or 1 mol dm

−3

Li[TFSI] (red trace).

Figure 4.4 Proposed mechanisms of superoxide radical‐mediated decomposition of ether hydroperoxides.

Figure 4.5 (a) First cycle and (b) voltage versus time profiles for galvanostatic discharge–charge cycling of Li–O

2

cells with DME (black lines) and DMDMB (blue lines) electrolytes. (c) CO

2

evolution as a function of charging capacity and (d) discharge–charge capacities for the DME (black circles) and DMDMB (blue circles) electrolytes. Li[TFSI] is used as the electrolyte salt in each case to form complexes [(DMDMB)

2

Li][TFSI] and [(DME)

2

Li][TFSI].

Figure 4.6 Comparison of the CO

2

evolution of two Li–air cells containing different electrolytes – 0.5 mol dm

−3

lithium perchlorate (Li[ClO

4

]) in DMSO (red) and 0.5 mol dm

−3

Li[PF

6

] in tetraglyme (green). After being removed from the cells, cathodes were subsequently treated with acid and Fenton's reagent to decompose Li

2

[CO

3

] (to give [12]CO

2

, a) and lithium carboxylates (to give Org [12]CO

2

, b), respectively. The numbers 2 and 5 on the

x

‐axis correspond to the CO

2

evolution at the end of the second and fifth discharge cycles, respectively.

Figure 4.7 (a) Cyclic voltammetry (1 V s

−1

) at a Au electrode in O

2

‐saturated MeCN with 0.1 mol dm

−3

[Et

4

N][ClO

4

] salt containing different concentrations of Li[ClO

4

] as shown in the legend. (b)

In situ

SERS at the Au electrode in MeCN with 0.1 mol dm

−3

Li[ClO

4

] collected at the potentials/times shown. Peaks 1 and 4 were assigned to solvent and [ClO

4

]

−

anion stretches, and peaks 2 and 3 show the O–O stretch of adsorbed LiO

2

(1137 cm

−1

) and the O–O stretch of Li

2

O

2

(808 cm

−1

).

Figure 4.8 Illustration of novel separation of reaction and storage processes in a Li–air cell, where the cathode electrolyte is an IL with no added Li salt and the cross section of a cell based on this concept.

Chapter 05

Figure 5.1 Oxygen reduction mechanisms in Li–oxygen batteries.

Figure 5.2 Decomposition mechanisms of organic electrolytes in Li–oxygen batteries triggered by reactive oxygen species.

Figure 5.3 Nucleophilic attack of ester bonds by superoxide species.

Figure 5.4 (a) The mechanism of autoxidation in DME. (b,c) Methylation as an approach to reduce the autoxidation.

Figure 5.5 (a) Li

2

O

2

abstracting protons from DMSO molecules generating LiOH. (b) Two possible routes of Hoffman β‐H elimination in PYR

14

+

cations by superoxides.

Figure 5.6 A proposed mechanism of DMA decomposition by Li metal. No proper insoluble product can be formed to serve as SEI layers.

Figure 5.7 Quantitative analysis of carbon corrosion in the discharge process (green arrow).

Figure 5.8 The chemistry of carbon corrosion during the ORR process generating carbonates and epoxy groups.

Figure 5.9 Better cathode stability can be achieved through (1) the physical protection effect by isolating carbon from Li

2

O

2

and (2) complete Li

2

O

2

decomposition by uniform OER catalyst loading.

Figure 5.10 Catalyzed decomposition of DME by the introduction of catalyst including Pt, MnO

2

, and Au. (a) Electrochemical profile of the discharge and recharge processes. (b) Oxygen evolution rate quantified by DEMS. (c) CO

2

generation rate quantified by DEMS.

Figure 5.11 Redox mediator facilitated electrochemical decomposition of Li

2

O

2

.

Figure 5.12 Decomposition of PVDF by lithium peroxide.

Figure 5.13 Surface SEI formation on Li metal after the reaction with electrolytes. Dendritic growth of Li metal may appear upon the breakage of SEI.

Figure 5.14 The corrosion of Li‐metal anode during the cycling of Li–oxygen batteries.

Figure 5.15 The composition and morphology of SEI on Li metal could be changed by the involvement of O

2

.

Figure 5.16 New Li

2

O

2

decomposition pathway introduced by the addition of H

2

O.

Chapter 06

Figure 6.1 SEM images of the CNT fibril cathode.

Figure 6.2 (a) SEM images of the functionalized graphene sheets cathode. (b) TEM image of the hierarchically porous carbon cathode.

Figure 6.3 Optimized geometries of the Li

5

O

6

cluster and (a) rGO and (b) B‐rGO. Corresponding binding energy values are also listed. (c) Schematic picture of B‐rGO as the substrate gaining electrons from the Li

2

O

2

.

Figure 6.4 High magnification TEM images of PNT‐LSM.

Figure 6.5 (a) TEM image of nanoporous N‐doped graphene nanosheets with uniformly dispersed encapsulated RuO

2

nanoparticles with the inset being high‐resolution TEM of the RuO

2

nanoparticles surrounded by 2–3 layers of N‐doped graphene. (b) High‐resolution TEM of the RuO

2

surrounded by 2–3 layers of N‐doped graphene at the charge state after 50 cycles. (c) High‐resolution TEM image of the unencapsulated RuO

2

nanoparticles before test. (d) TEM images of the unencapsulated RuO

2

nanoparticles after 62th cycles.

Figure 6.6 (a) Proposed mechanism for Li

2

O

2

nucleation on MNP@CNTs for MNP@CNTs, the surface electron density is strengthened, where the nucleation and growth of Li

2

O

2

is promoted on the entire CNTs surface. (b) RuNPs@CNTs cathode after 10th cycle discharged to 250 mAh g

−1

. (c) Proposed mechanism for Li

2

O

2

nucleation on MNP–CNTs for MNP–CNTs, the regional enrichment of electron density around noble metal NPs leads to localized distribution of Li

2

O

2

aggregation on the exposed noble metal NPs. (d) RuNPs–CNTs cathode after 10th cycle discharged to 250 mAh g

−1

.

Figure 6.7 The rate capability of the Li–O

2

cells with each type of cathode at current densities of 100 mA g

−1

and 2 A g

−1

.

Chapter 07

Figure 7.1 The working principle of redox mediator in Li–O

2

batteries.

Figure 7.2 The working mechanism of shuttle additive for overcharge protection of Li‐ion battery [44]. This publisher falls under STM opt out.

Figure 7.3 Mechanism of redox targeting of insulating LiFePO

4

by a molecular redox shuttle S.

Figure 7.4 A schematic illustration of a typical cyclic voltammogram of a reversible or quasi‐reversible redox reaction.

Figure 7.5 Molecular orbital energies of RMs and solvents. (a) The HOMO/LUMO energies of five representative solvents used in the Li–O

2

battery. (b) HOMO and LUMO energies of original RMs (black line) and first oxidized RMs (red line) in DMSO electrolyte using DFT calculations. (c) HOMO/LUMO energies of original RMs (black bar) and first oxidized RMs (red bar) in TEGDME electrolyte based on DFT calculation.

Figure 7.6 CV of (a) TTF, (b) FC, and (c) TMPD in Ar and O

2

to probe their stability in O

2

environment.

Figure 7.7 A schematic illustration of catalytic shuttling and parasitic shuttling of redox mediator.

Figure 7.8 CV of LiI in a 1 M LiTFSI–TEGDME electrolyte.

Scheme 7.1 Electrochemical reactions of LiI as a redox mediator in Li–O

2

batteries.

Figure 7.9 Gas evolution during charge of a cell using LiI as a redox mediator. (a) Charging without LiI. (b) Charging with LiI.

Scheme 7.2 Electrochemical reactions of LiBr as a redox mediator in Li–O

2

batteries.

Figure 7.10 CV of LiBr in 1 M LiTFSI–diglyme electrolyte.

Figure 7.11 Voltage and gas evolution profiles during charging (a) without and (b) with 10 mM LiBr as a mediator. (c) Statistics of the gas evolution at different LiBr concentration based on at least three replications.

Figure 7.12 XRD patterns of carbon cathodes discharged to 2 V (a) using diglyme/LiBr and (b) diglyme‐/LiI‐based electrolyte solutions.

Scheme 7.3 Electrochemical reactions of nitroxide as a redox mediator in Li–O

2

batteries.

Scheme 7.4 Electrochemical reactions of TEMPO as a redox mediator in Li–O

2

batteries.

Figure 7.13 CV of 10 mM TEMPO in 1 M LiTFSI–diglyme electrolyte.

Figure 7.14 Gas evolution during charge with TEMPO.

Figure 7.15 Potentiostatic measurements of 1 M LiTFSI–diglyme with various TEMPO concentration using a glassy carbon working electrode, a LiFePO

4

counter electrode, and a lithium reference electrode; compare Figure 7.1; (a) current–time profiles for different TEMPO concentrations

c

0

using

E

 = 3.95 V vs Li

+

/Li and diffusion distance

d

 = 200 mm; (b) stationary shuttle current

j

ss

at different

c

0

, derived from the final data points of (a); (c) stationary shuttle current

j

ss

at different

d

, obtained at

c

0

 = 100 mM and

E

 = 3.95 V vs Li

+

/Li; and (d) stationary shuttle current

j

ss

at different

E

, obtained at

c

0

 = 100 mM and

d

 = 200 mm.

Figure 7.16 Chemical structures of the investigated nitroxides. AZADO, 2‐azaadamantane

N

‐oxyl; TMAO, 1,1,3,3‐tetramethyl‐2,3‐dihydro‐2‐azaphenalene‐2‐yloxyl.

Figure 7.17 Cyclic voltammograms (CVs) of 10 mM TEMPO

1

, 10 mM 4‐methoxy‐TEMPO

2

, 10 mM AZADO

3

, 10 mM 1‐Me‐AZADO

4

,

and 10 mM TMAO

5

in 1 M LiTFSI–diglyme electrolyte.

Scheme 7.5 Electrochemical reactions of TTF as a redox mediator in Li–O

2

batteries.

Figure 7.18 CV of TTF in 0.1 M TBAClO

4

–DMSO.

Figure 7.19 Constant current discharge/charge with and without TTF [80].

Figure 7.20 Li

2

CO

3

Raman peak intensity evolution during constant current cycling in TTF‐free and TTF‐containing cells.

Figure 7.21 CV of TDPA in 0.1 M LiTFSI–TEGDME under argon (black) and oxygen (red) atmosphere.

Scheme 7.6 Electrochemical reactions of TDPA as a redox mediator in Li–O

2

batteries. Source: Kundu et al. 2015 [82]. Reprinted (adapted) with permission of American Chemical Society.

Figure 7.22 Charge–voltage curve (red) and the corresponding O

2

(black) and CO

2

(blue) evolution profile for the Li−O

2

cell (a) with and (b) without 50 mM TDPA in the 0.5 M LiTFSI–TEGDME electrolyte.

Scheme 7.7 Electrochemical reactions of FePc as a redox mediator in Li–O

2

batteries.

Figure 7.23 CV of FePc in 0.1 M LiTFSI–DMSO electrolyte.

Figure 7.24 Cycling voltage profile of cells without and with FePc.

Figure 7.25 SEM images of the discharged cathodes (a) without and (b) with FePc.

Scheme 7.8 Reactions of DBBQ as a redox mediator in Li–O

2

batteries.

Figure 7.26 CV of DBBQ in 1 M LiTFSI–TEGDME.

Figure 7.27 SEM images of cells (a) fully discharged without DBBQ, (b) half discharged with DBBQ, and (c) fully discharged with DBBQ.

Chapter 08

Figure 8.1 The experimental setup. (a) Operando μ‐XRD study of a Li–O

2

cell under cycling condition. (b) Cell's schematic design. (c) SEM image of a cross section of the Li–O

2

cell. The yellow bar represents the actual X‐ray beam dimension compared with different section of the cell.

Figure 8.2 Discharge–charge voltage profiles of Li–O

2

cells. (a) A new cell. (b) A

rebuilt

cell using new cathode–separator but used lithium anode taken from the cell used in (a) at the end of the cycling. Cycle numbers are marked by charging voltage curves.

Figure 8.3 Photographic images of used anodes. (a) The top view (separator–anode contact side) and (b) bottom view (anode current collector contact side) of Li anode taken from a nonaqueous Li–O

2

cell after initial failure. (c) Li anode after multiuses until no capacity remained. The diameter of the anode is 12 mm.

Figure 8.4 Monitoring the change of anodic Li during discharge–charge cycling by a microfocused synchrotron X‐ray. (a) Discharge–charge voltage profile of the

operando

Li–O

2

cell as the function of cycling time. The points at which XRD data sets were collected are marked as

a, b, c

… on the curve. (b) Four representative XRD sets taken at the cycling stage

a, c, f

, and

i

, and each set are consisted of seven selected XRD patterns collected at various anode–separator interfacial depths (marked as 1, 2, …, 7 on the left of each set). (c) An expanded XRD data of set

i

covering the whole anode through adjacent separator region. The inset is a representative XRD pattern taken from the position marked by a blue line that includes references of Li and LiOH.

Figure 8.5 Quantitative analysis of anode composition changes. (a) The change in the amount of LiOH and Li metal contents at the anode–separator interfacial region as a function of cycling time with the contents individually normalized to the highest points of respective components. The cycling time is corresponding to Figure 8.3a, and data collecting positions (

a

,

b

, ..., 

i

) during cycling were same as those marked in Figure 8.3a. (b) Changes in the amount of LiOH at different interfacial depths (as marked by the scan number in Figure 8.3b) as the function of cycling time. The amounts at all depths are normalized to the highest intensity (depth 4 at 16th hour). (c) Changes in the amount of Li at different interfacial depths as the function of cycling time. The amounts at all depths are normalized to the highest intensity (depth 5 at beginning of cycling).

Figure 8.6 A full set of μ‐XRD spectra for a cell. (a) Microfocused X‐ray diffraction data of the cell after 14 cycles with a current of 0.2 mA and cutoff voltage of 2.2–4.5 V for each discharge–charge. This cell was also used for generating 3D microtomography image shown in Figure 8.6 and the video clip. (b,c) Phase identification at LiOH dominating layer (II) and Li dominating region (III), respectively.

Figure 8.7 X‐ray tomography images of the used anode. (a) Planar cross section of LiOH layer on the anode. (b,c) Vertical cross sections through the anode, including a small portion of separator.

Figure 8.8 Electrochemical performance of Li–O

2

batteries. (a) A gradual decay in capacity during cycling of a fresh battery and a rebuilt battery with a used separator. (b) The discharge–charge curves of 3rd and 23rd cycle of a fresh battery and 2nd cycle of a rebuilt battery.

Figure 8.9

Ex situ

X‐ray diffraction study on a failed Li–O

2

battery. (a) Progression of a series of diffraction patterns from the anode interface (bottom) to the cathode region (top) at 20 μm interval. (b) Comparison of diffraction pattern taken at the midsection of separator with references. (c) Intensity distribution of accumulated Li

2

CO

3

as function of distance from the cathode (left) to the anode (right). The origin of the distance is assigned at the interface between the cathode and the separator.

Figure 8.10 SEM images of (a) fresh separator and (b–i) separator slices from a failed battery sampled at equal distance from cathode to anode. The scale bar is 2 μm.

Figure 8.11 (a) Discharge–charge voltage profiles for the 1st and 11th cycles of a Li–O

2

cell. (b) X‐ray diffraction patterns collected from a cathode position near separator at different cycling time as marked by * in (a). (c) Analysis of Li

2

O

2

relative amount (integrated intensity of the peak (101) and normalized to the largest peak at time of 83.4 h) versus cycling time. Solid squares represent the position of the cathode close to the separator, and the empty circles represent the position closer to GDL.

Figure 8.12 (a) Voltage profile of a prolonged discharge of the Li–O

2

cell. The markers on the

x

‐axis indicate the moments when the diffractions were taken. (b) The increase of diffraction intensities of the three strongest Li

2

O

2

peaks taken at the same layer position on the cathode, where the vertical arrow indicates the direction of the discharge depth progresses. (c) Changes of the primary crystalline size and relative concentration of Li

2

O

2

as the function of discharge capacity obtained by fitting the peak (101). The peak intensities were normalized to the largest peak at the end of discharge.

Figure 8.13 (a) X‐ray diffraction of Li

2

O

2

(100) and (101) collected from the GDL to the cathode–separator interface at a cell discharge capacity of 584 mAh g

−1

. (b) Li

2

O

2

integrated intensity of peak (101) and crystallite size distributions at different cell depths. The peak intensities were normalized to the largest peak at position of 80 μm. The

x

‐axis origin is defined at the GDL–cathode layer contact.

Figure 8.14 Evolutions of (a) integrated peak intensity and (b) primary crystalline size, obtained from the diffraction of (110), as the functions of cathode layer position and the discharge capacity. The two‐dimensional data were acquired through a series of scans at different layer thickness positions and were taken at any given discharge capacity. For (a), the peak intensities were normalized to the largest peak at the end of discharge.

Figure 8.15 Fittings of the experimentally observed crystallite size (diamond) and concentration (triangle) growths as the function of discharge capacity.

Chapter 09

Figure 9.1 Comparison of Raman scattering processes.

Figure 9.2 (a) Cycling procedures for roughening an Au working electrode in argon purged 0.1 M KCl solution at 23 °C. Potential versus the Ag/AgCl reference electrode. (b) Voltammetric oxidation/reduction cycling of a gold electrode in argon purged 0.1 M KCl. (c) Electrodes before (i) and after cycling (ii) showing tan color of roughened electrode.

Figure 9.3 Schematic of an example

in situ

electrochemical Raman cell.

Figure 9.4 Potential dependence of the superoxide Raman signals in O

2

‐saturated 0.1 M TBAClO

4

in DMSO.

Figure 9.5

In situ

SERS of O

2

reduction on a roughened gold electrode in 0.1 M LiClO

4

in MeCN. The spectra were collected along the discharge reduction curve showing the formation of LiO

2

(1137 cm

−1

) and Li

2

O

2

(808 cm

−1

).

Figure 9.6

In situ

SERS of O

2

‐saturated 0.1 M NaOTf in (a) DMSO, (b) DMA, (c) 1 M NaOTF in

diethylene glycol dimethyl ether

(

DEGDME

), and (d) 0.1 M NaOTf in MeCN on a roughened Au working disk electrode at 23 °C, 0.1 V s

−1

at varying potentials vs Na

+

/Na.

Figure 9.7

In situ

SERS of a roughened gold electrode in O

2

‐saturated 0.1 M TBAClO

4

in DMSO and 20 mM phenol. This shows the cathodic sweep from 3.1 to 2.2 V vs Li+/Li and observes the formation of H

2

O

2

at the interface.

Figure 9.8

In situ

Raman of O

2

reduction in O

2

‐saturated 0.1 M TBAClO

4

on (a) smooth, (b) roughened, and (c) SHINs drop cast on the surface of polycrystalline gold. All potentials vs Li

+

/Li (SHINERS and O

2

are not to scale).

Figure 9.9

In situ

SERS collected at

open‐circuit potential

(

OCP

) (black) and then at the end of the cathodic reaction (red). This is then followed by spectra showing further cathodic reactions over a time period of 5 h discharging under 16O

2

.

Figure 9.10 Schematic showing typical

in situ

setups used for IR reflection–absorption spectroscopy (a) as well as internal reflection SEIRAS spectroscopy (b).

Figure 9.11

In situ

IR spectra showing the formation of dimethyl sulfone at positive potentials on an Au working electrode arranged in external reflection configuration. The spectra were taken in an O

2

‐saturated electrolyte composed of 0.1 M LiPF

6

in DMSO.

Figure 9.12 Spectroelectrochemical cell used by Calvo and coworkers for their

in situ

infrared spectroscopy studies.

Figure 9.13 Schematic of the modified ATR‐SEIRAS setup used by Adzic and coworkers.

Figure 9.14

In situ

ATR SEIRA spectra showing no major changes in bands in deoxygenated electrolyte (a), the formation of propylene carbonate ring‐opened species (ROCO

2

Li) in an O

2

‐saturated electrolyte containing Li salt and propylene carbonate. (b)

Figure 9.15 Schematics of UV/Vis cells.

Figure 9.16 Time‐resolved UV/Vis spectra of O

2

(a) reduction and (b) evolution in O

2

−

saturated 0.1 M TBAClO

4

in DMSO showing the formation and disappearance of

. Reproduced with permission of American Chemical Society.

Figure 9.17 (a) Voltammetric analysis of O

2

‐saturated 1 M LiOTf and 0.1 M KI dissolved in

tetraethylene glycol dimethyl ether

(

TEGDME

), 5 mV s

−1

. (b) UV/Vis spectra at different stages of the voltammetry displayed. Reproduced with permission of Elsevier.

Figure 9.18 Structure of tetrathiafulvalene.

Figure 9.19 Zeeman splitting. The energy gap (Δ

E

) between spin‐up (+½) and spin‐down (−½) states is directly proportional to the magnetic field strength (

B

0

).

Figure 9.20

In operando

EPR cell design used by Wandt et al. [62] for the detection of singlet O

2

in an aprotic Li–O

2

battery [62]. Top left: cell housing with (1) lid containing a connection for gas purging and three feedthrough wires for contacting of working, counter, and reference electrodes and (2) EPR tube containing the electrochemical cell. Center: tubular electrochemical cell with (3)

poly(tetrafluoroethylene)

(

PTFE

) spacer. Bottom right: cut‐through electrochemical cell, (4) Vulcan working electrode coated on Celgard separator, (5) reference electrode, (6) LiFePO

4

counter electrode coated on Al wire (10), (7) glass fiber separator, (8) Al wire (0.1 mm diameter) as working electrode current collector, (9) PTFE tube, and (10) Al wire (2.0 mm diameter) as counter electrode current collector.

Chapter 10

Figure 10.1 Schematic configuration of a zinc–air battery with alkaline electrolyte.

Figure 10.2 (A) A schematic image of the solid‐state, flexible, and rechargeable zinc–air battery. (Left inset: a photo showing the flexibility of the cell (a); the cross‐sectional SEM image of the laminated structure of the battery (b).) The optical pictures of (B) the freestanding zinc electrode film, (C) the bifunctional catalytic air electrode film using LaNiO

3

/NCNT, and (D) the porous gelled PVA electrolyte membrane. The corresponding SEM images of (E) the freestanding zinc film, (F) the bifunctional catalytic air electrode (the inset of TEM image illustrates the core–corona structured LaNiO

3

nanoparticle and intertwined nitrogen‐doped carbon nanotubes), and (G) the porous gelled PVA electrolyte membrane.

Figure 10.3 Scheme of modified electrode using polysulfonium 1.

Figure 10.4 Model structure of air electrode with metal mesh/foam‐supported

oxygen‐permeable membrane

(

OPM

, for filtering out H

2

O or CO

2

), gas diffusion layer (GDL), and catalyst layer (CL). Suggested porous structure of the catalyst layer with hierarchical macro‐/meso‐/interparticle pores for both efficient electrolyte wetting and gas transportation.

Figure 10.5 Representation of the steps for the ORR/OER in ILs.

Figure 10.6 Activity map for the ORR obtained for different classes of Pt‐based materials. Improvement factors are given on the basis of activities compared with the values for polycrystalline Pt and the state‐of‐the‐art Pt/C catalyst established by RDE measurements in 0.1 M HClO

4

at 0.95 V.

Figure 10.7 Synthesis scheme and electrochemical oxygen reduction performance of Ag

–Co surface alloy materials. (a) Schematic of Ag

3

[Co(CN)

6

] precursor formation and rapid reduction to form Ag

–Co surface alloys. Ag, blue; Co, violet; C, green; N, black. (b) Leaching CV in Ar‐purged 0.001 M H

2

SO

4

 + 0.1 M Na

2

SO

4

, which shows dissolution of Co on the first cycle, followed by featureless scans afterward. (c) Rotating disk electrode I

–V polarization curves (potentials relative to the RHE) taken in O

2

‐saturated 0.1 M NaOH at 25 °C and 900 rpm for the Ag

–Co surface alloy, as well as a leached alloy, a segregated Ag/Co sample and pure Ag and Co. (d) Mass transport corrected kinetic current Tafel plots for the samples. The inset shows Pb‐stripping voltammograms used to measure the electrochemical surface area at 10 mV s

−1

in Ar‐purged electrolyte with 125 μM Pb(NO

3

)

2

added after the activity measurements. (e) Summary of kinetic current densities for oxygen reduction on each sample, measured at 0.8 and 0.85 V

RHE

. Error bars represent the standard deviation for three or more separately synthesized samples (except commercial Pt, which varied mainly because of a higher electrolyte sensitivity).

Figure 10.8 (a) Crystal structure of RuO

2

. (b) OER performance of Ru‐ and Ir‐based electrocatalysts in acid electrolyte from different groups.

Figure 10.9 Oxygen electrode activities of the nanostructured Mn oxide thin film, nanoparticles of Pt, Ir, and Ru, and the GC substrate. The Mn oxide thin film shows excellent activity for both the ORR and the OER.

Electrochemical application of nanocrystalline CoMnO‐B and CoMnO‐P as ORR and OER electrocatalysts. (a) Voltammograms of the ORR using catalyst‐modified RDEs in O

2

‐saturated alkaline electrolyte. (b) K–L plots for the ORR. (c) Galvanostatic discharge curves of the prototype zinc–air battery made with a nanoparticle catalyst layer. (d) Voltammetry curves of the OER measured at different catalyst‐modified electrodes.

Structural characterization of exfoliated LDH nanosheets. (a) Optical images of colloidal solutions of exfoliated nanosheets. Tyndall effect was visible when irradiated with a laser beam. (b) TEM images of bulk NiCo LDH nanoparticles. Inset is the corresponding

selected area electron diffraction

(

SAED

) pattern. Scale bar, 400 nm. (c) TEM images of exfoliated single‐layer nanosheet of NiCo LDH. Scale bar, 100 nm. Inset is the corresponding SAED pattern where the spots can be indexed to crystalline plane family of {110}. (d) AFM image of exfoliated single‐layer nanosheets of NiCo LDH. (e) Height profile of the exfoliated monolayer of NiCo LDH. The profiles 1, 2, 3, and 4 correspond to the numbered lines in Figure 10.3d, respectively. The average thickness is about 0.8 nm.

Design of molecular catalysts for water oxidation. (a) Redox flexibility arising from a multinuclear core. Multielectron transfer to afford several oxidation states and electron rearrangement among valence tautomers enables the accumulation of positive charges required for water oxidation. (b) Adjacent water‐activation sites to promote intramolecular O

–O bond formation. (c) Ball‐and‐stick representations of the molecular structure (left) and the Fe

5

O core structure (right) of [Fe

II

4

Fe

III

(μ

3

‐O)(μ‐L)

6

]

3+

[1]; the chemical structure of LH is also shown (bottom right). Three pentacoordinated iron centers are bridged by an oxygen atom in μ

3

‐fashion to form a triangle structure, and two hexacoordinated iron centers are connected to the triangle structure by six Ls.

SEM images (a) and TEM image (b) of the prepared TPPy, BET results, (c) and FTIR spectra (d) of the PPy. SEM images of the as‐prepared TPPy‐supported air electrodes (e). Contact angles of the nonaqueous electrolyte on AB (inset) and PPy; the electrolyte is LiTFSI in DME (f).

(a) Scheme of the synthesis of Co

4

N/CNW/CC. (b–d) Low‐ and high‐magnification SEM images of Co

4

N/CNW/CC. (e) XRD patterns of PNW/CC, ZIF‐67/PNW/CC, and Co

4

N/CNW/ CC. (f) TEM image of Co

4

N/CNW/CC (inset HRTEM image of Co

4

N). (g) SEM EDS element mapping of C, Co, and N for Co

4

N/CNW/CC.

(a) A schematic of a zinc–air battery. (b) Polarization and (c, d) discharge curves at 200 and 250 mA cm

−2

, respectively, of zinc–air full cells with different air electrodes: amorphous MnO

x

nanowires on Ketjenblack composites and 20% Pt on Vulcan XC‐72 (E‐tek). A commercial air electrode (Meet) was used for comparison.

Detection of molecular oxygen during electrocatalytic water oxidation by Et‐Fl

+

at different potentials. (a) Cyclic voltammograms of 0.1 M TBAP in acetonitrile in the presence of varying concentrations of pH = 2 water: 0, 30, 60, 120, 240, 360, and 480 mM. (b) Cyclic voltammograms obtained under the same conditions as in (a) but in the presence of 3 mM Et‐Fl

+

. The scan direction was −1.2∼ + 2.7∼ − 1.2 V versus NHE, sweep rate was 100 mV s21, and electrodes were a carbon working electrode, platinum counter electrode, and nonaqueous Ag/Ag

+

reference electrode. (c–e), Oxygen evolution during electrolysis of 0.1 M aqueous phosphate buffer solution at pH = 2 in the presence (red line) and absence (blue line) of 0.4 mM Et‐Fl

+

(

n

Et‐Fl+

 = 40 mM) at a constant potential of +1.8 (c), +2.2 (d), and +3.0 V (e) versus NHE. The red line in panel (d) represents the anticipated O

2

evolution for a process with 30% faradic efficiency. The electrolysis experiments were carried out in two solvent systems: [1] an acetonitrile/water mixture (2  :  1) and [2] water only. The two experiments gave similar results, so here we present only the results of aqueous electrolysis.

Preparation of the N and P codoped porous carbon (NPMC) electrocatalysts. a, Schematic illustration of the preparation process for the NPMC foams. An aniline (i) phytic acid (ii) complex (iii) is formed (for clarity, only one of the complexed anilines is shown for an individual phytic acid), followed by oxidative polymerization into a three‐dimensional PANi hydrogel crosslinked with phytic acids. As each phytic acid molecule can complex with up to six aniline monomers, phytic acid can be used as the crosslinker and protonic dopant to directly form the three‐dimensional PANi hydrogel network; for clarity, only a piece of the two‐dimensional network building block is shown in the enlarged view under the three‐dimensional PANi hydrogel. The PANi hydrogel is freeze‐dried into an aerogel and pyrolyzed in Ar to produce the NPMC (for clarity, only a piece of the two‐dimensional NPMC network building block is shown in the enlarged view under the three‐dimensional NPMC). (b,c) SEM images of PANi aerogel (b) and NPMC‐1000 (c). Inset in (c): digital photo images of PANi aerogel before (left) and after (right) pyrolysis at 1000 °C. (d,e) High‐resolution TEM image (d) and TEM image (e, left), with corresponding element mapping images of NPMC‐1000 (e). The TEM image shows a piece of interconnected network‐like scaffold. The element mapping images for C, N, and P show a uniform distribution of the elements.

Chapter 11

Figure 11.1 Theoretical volumetric and gravimetric energy densities for metal/O

2

and Li ion battery chemistries. For metal/O

2

chemistries, the compound in parenthesis indicates the assumed discharge product.

Figure 11.2 Cyclic voltammograms of neat BMIM‐Tf

2

N (light gray) and 10 mM Mg(Tf

2

N)

2

in BMIM‐Tf

2

N (black) on a 50 μm diameter Pt working electrode at room temperature; scan rate 100 mV s

−1

.

Figure 11.3

Pair distribution function

(

PDF

) for the series of A(TFSI)

x

–diglyme solutions, A = Li

+

, Na

+

, K

+

, Mg

2+

, Ca

2+

, Zn

2+

, highlighting (a) the common TFSI–TFSI and TFSI–diglyme atom–atom correlations as determined through principal component analysis, (b) the as‐measured total PDFs, and (c) the differential PDFs corresponding to the A

+/2+

solvation environment. Data have been offset for clarity.

Figure 11.4 Cyclic voltammograms of neat DEME‐BF

4

(light gray) and 100 mM Mg(BH

4

)

2

in DEME‐BF

4

(black). Cyclic voltammetry scan limits are constrained to −1 to 1 V vs Mg/Mg

2+

. Fifty microns diameter Pt‐disk working electrode; 100 mV s

−1

scan rate; room temperature.

Figure 11.5 Successive cyclic voltammograms for 0.5 M Mg(BH

4

)

2

/MPEG

7

PyrTf

2N

at 25 mV s

−1

.

Figure 11.6 Cell potentials and half‐reaction potentials for several metal/O

2

battery chemistries. The dashed red line corresponds to the potential at which O

2

reduces to superoxide.

Figure 11.7 Discharge/recharge cycles for a room temperature Mg/O

2

cell using 4 : 1 PhMgCl:Al(OPh)

3

/THF at 5 μA cm

−2

(superficial). Curves are labeled with the corresponding cycle numbers.

Figure 11.8 SEM images of the positive electrode surface on the side closest to the O

2

gas inlet for Mg/O

2

cell using 4 : 1 PhMgCl:Al(OPh)

3

/THF electrolyte. The dashed circles represent boundaries of the regions that were directly exposed to O

2

through perforations in the Pt‐coated current collector. (a) An electrode after first discharge. (b) Higher magnification of the first‐discharge product with an inset image of a control electrode exposed to O

2

in a cell held at open circuit. (c) An electrode at the end of first recharge. (d) Higher magnification of the residual product after first recharge.

Figure 11.9 Discharge–charge curves of the nonaqueous Mg/O

2

battery with iodine at 60C. The black, pink, green, and red lines correspond to the first, second, third, and fourth cycles, respectively. The blue line represents a discharging–charging profile in the absence of iodine.

Figure 11.10 Cell voltage vs capacity for a Mg/O

2

cell using the MACC electrolyte discharged at current densities ranging from 0.02 to 1 mA cm

−2

(superficial). The inset shows a typical discharge/charge cycle at 0.02 mA cm

−2

.

Figure 11.11 (a) Capacity per geometric electrode area achieved at the 0.6 V cutoff potential as a function of discharge rate. Capacity falls as a power law with respect to discharge rate at rates higher than 0.02 mAh cm

−2

. (b) Cell voltage at 50% depth of discharge as a function of discharge rate.

Figure 11.12 (a) Nyquist plots for a Mg/Mg cell containing MACC/DME electrolyte under Ar at various times during an

open circuit

(

OC

) hold (black) and after passing 0.075 mAh cm

−2

through the cell at 0.05 mA cm

−2

(blue). (b) Nyquist plots for a Mg/Mg cell using MACC/DME before O

2

exposure (black), under O

2

at various times after O

2

exposure during an OC hold (red), and after passing 0.075 mAh cm

−2

(blue). (c) Equivalent‐circuit model fit values of

R

SEI

during OC holds for Mg/Mg cells under Ar (black circles) and under O

2

(red squares).

Figure 11.13 Variation of

R

bulk

with time for Mg/Mg cells exposed to O

2

using MACC/DME (blue) and modified Grignard (green) electrolytes.

Figure 11.14 Electrochemical impedance spectra for a Mg/O

2

cell using MACC/DME electrolyte before O

2

exposure (black), after OC hold under O

2

(red), middischarge (light blue), and after discharge (blue). The inset indicates the stage of (pre)discharge at which the EIS data was measured. The ECM used to model the spectra is shown above.

Figure 11.15 The rock salt crystal structure of MgO (a) and the pyrite crystal structure of MgO

2

(b). Most stable surface terminations of MgO and MgO

2

(c–f). Red indicates oxygen atoms, and yellow indicates magnesium atoms.

Figure 11.16 Calculated free energy diagram for the discharge and charging of an Mg/O

2

cell on a MgO (100) surface. The black line refers to the single‐step pathway, and the blue line refers to the multistep pathway (with peroxide intermediates). The identity of the surface adsorbed species is indicated with text. Energies are plotted assuming the application of a potential,

U

, equal to the theoretical cell potential,

U

 = 

U

0

.

Figure 11.17 Calculated free energy diagram for discharge and charge of an Mg/O

2

cell for single‐step reactions occurring on the oxygen‐rich MgO

2

(111)

Orich‐1

(black curve) and MgO

2

(100)

Orich‐3

(blue curve) surfaces. The identity of the surface adsorbed species is indicated with text. Energies are plotted assuming the application of a potential,

U

, equal to the theoretical cell potential,

U

 = 

U

0

. (Note that the maximum value for the ordinate (3.5 eV) used in this plot is half the value used in Figure 11.16.)

Figure 11.18 Formation energies of intrinsic defects in (a) MgO and (b) MgO

2

calculated using the HSEα functional. Magnesium vacancies (V

Mg

) are depicted using blue lines, magnesium interstitials (Mg

I

) with blue‐dashed lines, oxygen vacancies (V

O

) with red lines, oxygen interstitials (O

I

) with red‐dashed lines, and oxygen divacancies (

) with green lines. Hole and electron polarons are shown in black. The slope of each line corresponds to its respective charge state; values of −2, −1, 0, +1, and +2 were considered (positive slopes correspond to defects with positive charges). The dashed line indicates the position of the Fermi level.

Figure 11.19 Migration energy barriers of dominant defects in MgO and MgO

2

calculated using the NEB method. (a) Hole polaron and (c) magnesium vacancy (V

Mg

2−

) in MgO. (b) Hole polaron and (d) electron polaron in MgO

2

.

Figure 11.20 Arrhenius plot of the conductivity of MgO, illustrating the three Arrhenius branches reported in experiments: (red)

high temperature

(

HT

) branch, (blue)

low temperature

(

LT

) branch, and (gray) 1 eV branch.

Chapter 12

Figure 12.1 Schematic view of the Marcus theory. The green and red lines define the ground and excited adiabatic states, respectively, whereas the dotted lines present the initial and final diabatic states.

Figure 12.2 Electron polaron (red isosurface) in

‐sulfur located between two sulfur atoms. The blue arrow shows the intra‐ring electron hop, while the black arrow is used to denote the intra‐ring hop.

Chapter 13

Figure 13.1 (a) Conductivity as a function of KOH content at 25 °C with 0, 5, 10, and 15 g of acrylic acid. (b) Conductivity curve of a polymer alkaline gel electrolyte with 6 wt% PAA at 25 °C. The inset is a photograph of the colorless, transparent, and elastic electrolyte gelatin. (c) SEM image for the PVA/PAA polymer membrane samples of top surface. (d) SEM image for the PVA/PAA polymer membrane samples of cross section. (e) The AC impedance spectra for the alkaline PVA/PAA polymer electrolyte membranes of various compositions at 25 °C. (f) The discharge curves of the Zn–air batteries with the different composition ratios of PVA/PAA polymer electrolyte membranes at C/5 discharge rate, both at 25 °C. (g) Photographs of the cross‐linked hydrogel electrolyte under bending. Scale bar: 1 cm. (h) Discharge curves of fiber‐shaped Al–air batteries at current densities of 0.5, 0.75, and 1.0 mA cm

−2

.

Figure 13.2 (a) Discharge curves of 1 M LiTf/triglyme‐based Li–O

2

battery and the

solid polymer electrolyte

(

SPE

)‐based Li–O

2

battery at different current densities. (b) Cycling stability of the SPE‐based Li–O

2

battery at a current density of 0.2 mA cm

−1

. (c) Voltage profiles of the galvanostatic cycling test of a Li–O

2

polymer battery operating at controlled capacity regime by applying a 100 mA g

−1

current with a limited capacity of 500 mAh g

−1

. (d) Voltage profiles of the galvanostatic cycling test of a Li–O

2

polymer battery operating at controlled capacity regime by applying a 100 mA g

−1

current with a limited capacity of 500 mAh g

−1

. (e) Cycling performance of the fiber‐shaped Li–air battery at current density of 1400 mA g

−1

in air. (f) Charge–discharge curves of flexible Li–O

2

battery and (g) the corresponding cycling performance. (h) Discharge–charge curves of a preliminary Li–O

2

battery based on the 1 M LiTFSI/P13TFSI/PVDF‐HFP polymer electrolyte membrane at a current density of 0.05 mA cm

−1

. (i) Discharge curves of lithium–air batteries with different electrolytes in ambient atmosphere at the discharge current density of 0.02 mA cm

−1

.

Figure 13.9 Common ionic liquid cations (

1‐ethyl‐3‐methylimidazolium

(

EMI

),

N

‐butyl‐

N

‐methylpyrrolidinium

(

PYR

),

1‐propyl‐2,3‐dimethylimidazolium

(

PMMI

),

N

‐methyl‐

N

‐propylpiperidinium

(

PP13

)) and anion (TFSI) used in Li–air battery research.

Figure 13.3 Flexible lithium–air battery based on lithium foil anode under (a) initial, (b) bending, and (c) twisting state. (a–c) (d) Freestanding zinc–carbon nanotube network composite electrode film. (e) Zn–air battery with 90° bending angle. (f) SEM image of (d). (g) Synthesis of Li–rGO composite electrode film. (h, i) Photographs of wound and twisted flexible Li–rGO composite electrode film, respectively.

Figure 13.4 SEM images of (a) SS mesh current collector prior to the growth and (b) densely coated Co

3

O

4

NW array. (c) Photograph of the Co

3

O

4

/NCNT/SS electrode. (d) SEM image (inset: photograph of the PCN–CFP) of P‐g‐C

3

N

4

nanosheets grown on CFP. (e) LSV curves of PCN–CFP under folding and rolling up. (f) Schematic illustration to the aligned cross‐stacked MWCNT sheet served as air cathode. (g) SEM image of CNT sheets with cross‐stacking angle of 90°.

Figure 13.5 (a) Schematic illustration of the solid‐state, flexible, and rechargeable zinc–air battery. The left two images correspond to the photograph showing the flexibility of the battery at the top and the cross‐sectional SEM image of the laminated structure of the battery at the bottom. (b) A demonstration of wearable prototype integrated with a tandem device in series to power an LED under bending. (c) Schematic illustration of the all‐solid‐state rechargeable Zn–air battery. (d) Photograph of the all‐solid‐state rechargeable Zn–air battery. (e) Galvanostatic discharge–charge cycling curve at 2 mA cm

−2

for the all‐solid‐state rechargeable Zn–air battery with NCNF‐1000 as catalyst, applying bending strain every 2 h. (f) Photograph of a blue LED powered by three all‐solid‐state Zn–air batteries in series. (g) Photographs of a blue LED powered by four all‐solid‐state Zn–air micro batteries in series.

Figure 13.6 (a) Schematic illustration to the battery composed of TiO

2

NAs/CT (cathode), glass fiber (separator), and lithium foil (anode). (b, c) Photographs of the Li–O

2

battery bent at 360° and the cycle performance under bending, respectively. (d) Schematic of the fabrication and working mechanism of a foldable Li–O

2

battery pack. (e) Multilayered structure of the stretchable Li–air battery. (f) Discharge curves of the stretchable Li–air battery under increasing strains. The inserted photographs display a red LED lit up by the Li–air battery under increasing strains.

Figure 13.7 (a) A prototype fiber‐shaped Zn–air battery. (b) Cross‐sectional image of the battery. (c) Discharge curves of the battery under periodic bending deformations (every 20 min) at a discharge current density of 0.1 mA cm

−2

. (d) Fabrication of the fiber‐shaped Al–air battery. (e) A commercial LED watch powered by two fiber‐shaped Al–air batteries connected in series and woven into a textile.

Figure 13.8 (a) Schematic illustration to the preparation of the fiber‐shaped Li–O

2

battery. (b) An LED display screen powered by the battery at different bending and twisting conditions. (c) Discharge curves of the battery corresponding to various deformations. (d) Charge–discharge curves of the battery after bending for thousands of cycles.

Guide

Cover

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