O2 Batteries with LiI ... - ACS Publications

Subscriber access provided by READING UNIV

Article 2

Electrolyte Composition in Li/O Batteries with LiI Redox Mediators: Solvation Effects on Redox Potentials and Implications for Redox Shuttling Azusa Nakanishi, Morgan L. Thomas, Hoi-Min Kwon, Yuki Kobayashi, Ryoichi Tatara, Kazuhide Ueno, Kaoru Dokko, and Masayoshi Watanabe J. Phys. Chem. C, Just Accepted Manuscript • DOI: 10.1021/acs.jpcc.7b11859 • Publication Date (Web): 09 Jan 2018 Downloaded from http://pubs.acs.org on January 9, 2018

Just Accepted “Just Accepted” manuscripts have been peer-reviewed and accepted for publication. They are posted online prior to technical editing, formatting for publication and author proofing. The American Chemical Society provides “Just Accepted” as a free service to the research community to expedite the dissemination of scientific material as soon as possible after acceptance. “Just Accepted” manuscripts appear in full in PDF format accompanied by an HTML abstract. “Just Accepted” manuscripts have been fully peer reviewed, but should not be considered the official version of record. They are accessible to all readers and citable by the Digital Object Identifier (DOI®). “Just Accepted” is an optional service offered to authors. Therefore, the “Just Accepted” Web site may not include all articles that will be published in the journal. After a manuscript is technically edited and formatted, it will be removed from the “Just Accepted” Web site and published as an ASAP article. Note that technical editing may introduce minor changes to the manuscript text and/or graphics which could affect content, and all legal disclaimers and ethical guidelines that apply to the journal pertain. ACS cannot be held responsible for errors or consequences arising from the use of information contained in these “Just Accepted” manuscripts.

The Journal of Physical Chemistry C is published by the American Chemical Society. 1155 Sixteenth Street N.W., Washington, DC 20036 Published by American Chemical Society. Copyright © American Chemical Society. However, no copyright claim is made to original U.S. Government works, or works produced by employees of any Commonwealth realm Crown government in the course of their duties.

Page 1 of 48 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

The Journal of Physical Chemistry

Electrolyte Composition in Li/O2 Batteries with LiI Redox Mediators: Solvation Effects on Redox Potentials and Implications for Redox Shuttling Azusa Nakanishi, Morgan L. Thomas, Hoi-Min Kwon, Yuki Kobayashi, Ryoichi Tatara, Kazuhide Ueno, Kaoru Dokko and Masayoshi Watanabe*

Department of Chemistry and Biotechnology, Yokohama National University, 79-5 Tokiwadai, Hodogaya-ku, Yokohama 240-8501, Japan

CORRESPONDING AUTHOR FOOTNOTE: To whom correspondence should be addressed. Telephone/Fax: +81-45-339-3955. E-mail: [email protected] (M. W.)

1 ACS Paragon Plus Environment

The Journal of Physical Chemistry 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 2 of 48

ABSTRACT The use of LiI as redox mediator for the charge reaction in non-aqueous Li/O2 cells has been widely studied recently, as a possible means to fulfill the great promise of the Li/O2 system as a high energy density “beyond Li-ion” battery. In this work, we highlight the importance of considering the redox potential for both the I−/I3− and I3−/I2 redox couples, and how the electrolyte solvent (here tetraglyme (G4) and dimethylsulfoxide (DMSO)), and concentration (here 1.0 M and 2.8 M), has a profound influence on these potentials. Through a combination of galvanostatic cycling, electrochemical mass spectrometry and cyclic voltammetry, we thus consider the influence of solvent and electrolyte concentration on both the redox mediation and redox shuttle processes, and suggest that this important aspect must be considered for further studies with mediators in Li/O2 and related systems. We demonstrate that in our system, 100 mM LiI in 1.0 M Li[TFSA]/DMSO provides the most effective redox mediation amongst the electrolytes we have studied, but conversely exhibits the highest degree of redox shuttling (in the absence of O2). The balance between effective limitation of redox shuttle and ease of mediator diffusion to discharge products is of great importance and should be considered in any future cell design utilizing a mediator.


Page 3 of 48 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


INTRODUCTION Rechargeable Li/O2 batteries have recently attracted much attention as next generation energy storage devices due to their much higher theoretical energy density than prevalent lithium-ion batteries.1,2 Through diverse research efforts by many researchers worldwide, many severe problems which impede the practical application of lithium-air batteries have been revealed, such as chemical instability of cell components (notably organic electrolytes3–11 and carbon based cathodes12–14), extremely high overpotentials (especially during charge1,15) and so on. The high overpotential during charge is thought to be due to the low electronic conductivity of Li2O2 and side products, such as Li2CO3, and this overpotential can promote unwanted electrolyte decomposition leading to further side reactions.16 Moreover, it is generally known that high voltage hysteresis between charge and discharge is a major contributor to lowering energy efficiency.17 Thus, for stable, efficient and reversible operation of Li/O2 batteries, it is necessary to inhibit the generation of high charge overpotential. The use of soluble redox mediators (RMs), typically introduced into the electrolyte as reductants, is a promising strategy to lower the overpotential. Ideally, during the first discharge, the RM should not undergo any electrochemical reactions, and thus the main 3 ACS Paragon Plus Environment


Page 4 of 48

discharge product, Li2O2, would be formed directly. In turn, during charge, the RM should be electrochemically oxidized at the cathode, and then the oxidized form would chemically react with Li2O2, yielding O2 and the reduced form of the RM. In this way, since the charge voltage is determined by the redox potential of the RM, the charge voltage can be greatly suppressed, provided a RM with an appropriate redox potential is chosen. Lim et al.18 and Pande et al.19 discussed the importance of considering the redox potential in selecting suitable RM candidates for Li-O2 batteries. So far, many RMs have been reported, such as tetrathiafulvalene

(TTF)20–22,

2,2,6,6,-tetramethylpiperidinyloxyl

(TEMPO)23,24,

tris[4-(diethylamino)phenyl]amine (TDPA)25, iron phthalocyanine (FePc)26, LiBr27–29, cobalt(Ⅱ) bis(terpyridine) (Co(Terp)2)30, and so on with a significant lowering of the charge voltage observed in many cases. However, among these RMs, LiI has been studied the most extensively.27,31–51 Earlier reports on LiI-based RMs for the Li/O2 battery revealed substantially lowered charge potentials, and improved cycle lifetimes.31,32,36 However, more recent studies have revealed a number of possible limitations, and have highlighted (i) the importance of the redox shuttle effect38,39, (ii) the necessity for on-line measurement of gas evolution during charge to evaluate the mediator performance19,52, and (iii) the important role played by non-negligible quantities of protic species (water, alcohols) in the presence 4 ACS Paragon Plus Environment

Page 5 of 48 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


of iodine/iodide species in creating an alternative charge reaction pathway.27,39,40,51,53 Moreover, for much of the earlier literature reported for the LiI RM in Li/O2 batteries (see Table S1 in the Supporting Information), thorough quantitative consideration regarding the possible contributions of the RM redox couple to the capacity is limited due to lack of data on either the electrolyte volume (necessary in order to calculate the molar quantity of iodide) or the cathode mass (or area, necessary in order to calculate the actual coulombs of charge passed from the reported gravimetric (or areal) capacities). In the present work, the role of the electrolyte composition in determining the redox potentials of the LiI redox couples, and the implications for both the mediation reaction (i.e. cycling performance of the Li/O2 battery) and the redox shuttle are explored. In consideration of recent studies (vide supra), electrochemical mass spectrometry (ECMS)11 is employed as an analytical tool to evaluate the performance of the RM, and the water content of the electrolytes has been strictly limited. In addition, we consider the possible contributions of the RM redox couples to the observed capacity. Iodide undergoes two redox reactions, namely I−/I3− and I3−/I2 (see (1) and (2) below), whose redox potentials are higher than that of the O2/Li2O2 redox couple. Thus, the oxidants of each of the iodine redox couples, namely triiodide (I3−) and iodine (I2) are both 5 ACS Paragon Plus Environment


Page 6 of 48

capable of reacting with Li2O2 as follows ((3) and (4)). 3 I− → I3− + 2 e−

(1)

2 I3− → 3 I2 + 2e−

(2)

Li2O2 + I3− → 2 Li+ + O2 + 6 I− Li2O2 + 3 I2 → 2 Li+ + O2 + 2 I3−

(3) (4)

Recently, Zhang et al.40 and Zhu et al.41 have reported no reactivity of triiodide toward Li2O2 in aprotic solutions from the viewpoint of kinetics. Although there has been some discussion in the literature over the possibility of I2 generation in Li/O2 cells containing LiI RMs (for example Wang and co-workers37,41, Aurbach and co-workers54, McCloskey and co-workers39, and Zhou and co-workers51), there are few reports that explore which of the two redox couples (I−/I3− or I3−/I2) is active as a RM from the viewpoint of thermodynamics. Here, it should be noted that by variation of the electrolyte solvent structure and salt concentration, it is possible to tune the redox potential of redox active species. Thus, the selection of suitable RMs must also involve consideration of the electrolyte. Glymes (such as diglyme, triglyme (G3, triethylene glycol dimethylether) and tetraglyme (G4, tetraethylene glycol dimethylether)) and dimethyl sulfoxide (DMSO) have been widely studied as electrolyte solvents for Li/O2 batteries. Glymes have been studied as they offer 6 ACS Paragon Plus Environment

Page 7 of 48 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


improved stability against side-reactions in Li/O2 batteries compared to conventional Li-ion battery electrolytes55, and more recently DMSO has emerged as a promising alternative.56 It is therefore unsurprising that for most existing studies on iodide-based RMs for the Li/O2 battery, either glyme or DMSO were selected as the electrolyte solvents, see Table S1 in the Supporting Information. In this work, Li[TFSA] solutions of 1.0 M and 2.8 M in either G4 or DMSO were explored as the electrolytes. It should be noted that in several earlier reports36,39, G1 (i.e. dimethoxyethane, DME) was employed as the glyme solvent. However, the high volatility of this solvent (see Table S2) poses an obstacle to maintaining a fixed electrolyte salt concentration during experiments, particularly in ECMS measurements (performed under a continuous helium carrier gas flow) for which we have already reported the slow evaporation of G3 (notably less volatile than DME, see Table S2).11 Thus in this study, G4 (with lower volatility) was selected. The physicochemical properties of the electrolytes employed in this study are summarized in Table 1. It should be further noted that through judicious choice of the concentrations, the 1.0 M Li[TFSA] electrolytes correspond neatly to a solvent/salt molar ratio of 4:1 and 12:1 for G4 and DMSO, respectively. In the same manner, the highly concentrated 2.8 M Li[TFSA] electrolytes correspond to solvent/salt molar ratios of 1:1 and 3:1 for G4 and DMSO respectively. These 7 ACS Paragon Plus Environment


Page 8 of 48

2.8 M electrolyte solutions represent examples of the recently intensively studied “salt-in-solvent” (or “superconcentrated”) electrolytes.57–59 The former, G4/Li[TFSA] with molar ratio 1:1, is a solvate ionic liquid (i.e. [Li(G4)][TFSA]) as we have previously reported in detail.60–62 The latter DMSO solution was also studied in our investigation of concentration effects on the Li/O2 battery in Li[TFSA]/DMSO electrolytes of various concentrations.56

Table 1. Physicochemical properties of Li[TFSA]-based electrolytes at 30 °C. Li+ concentration Solvent

G4

/ mol dm−3

csolvent:cLi[TFSA]

η

d

σ

/ mPa s

/ g cm−3

/ mS cm−1

0.98

4:1

9.7

1.155

3.22

2.75

1:1

106

1.400

1.60

1.00

12:1

3.9

1.214

10.71

2.76

3:1

86

1.439

2.12

a

DMSOb

a = Ref. 63, b = Ref. 56

Although many recent works have sought to shed light onto the mechanism of the iodide reactions in nominally aprotic electrolytes, in contrast we focus here on gaining further understanding regarding the interplay between the redox mediation reactions and the redox 8 ACS Paragon Plus Environment

Page 9 of 48 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


shuttle mechanism, and how selection of the electrolyte solvent and salt concentration are critical in determining this, and in particular in relation to the redox potentials. In the following discussion, we first describe the phenomena observed during typical cell cycling conditions, further clarify these observations using ECMS measurements, and then rationalize the findings using CV and “redox shuttle” (i.e. cycling under Ar) measurements.

EXPERIMENTAL Electrolytes. Purified G4 (kindly supplied by Nippon Nyukazai Co., Ltd.), and DMSO (Kishida Chemical, Japan) were dried with 4 Ⅱ molecular sieves. Lithium bis(trifluoromethanesulfonyl)amide (Li[TFSA], kindly supplied by Solvay, Japan) and anhydrous LiI (Sigma-Aldrich, 99.99%) were used as received. All chemicals were stored in an Ar-filled glove box (VAC, [H2O] ≤ 0.5 ppm) without exposure to the air. For electrolytes without a RM, two Li[TFSA] concentrations (1.0 M and 2.8 M) G4-based and DMSO-based electrolytes were prepared and the water content of these electrolytes, as determined by Karl Fischer titration, were less than 30 ppm. For electrolytes with a LiI RM, 100 mM or 200 mM LiI were added into G4-based and DMSO-based electrolytes. Cathode Preparation. A mixture of 80 wt% Ketjenblack (KB) powder (Lion, Japan) and 9 ACS Paragon Plus Environment


20

wt%

poly(tetrafluoroethylene)

(PTFE)

binder

(Aldrich)

Page 10 of 48

was

dispersed

in

N-methylpyrrolidinone (NMP) with an automixer (ARE-320, Thinky) to obtain a slurry. This slurry was coated onto carbon paper (TGP-H-060, Toray, Japan). After evaporating the NMP at 80 °C in air for 12 hours, electrodes of 16 mm diameter were punched and then dried under vacuum at 120 °C for over 12 hours and directly transferred into an Ar-filled glove box without exposure to the air. Battery Test and Electrochemical Mass Spectrometry (ECMS) of Li-O2 Cell. An ECC-Air electrochemical test cell (EL-Cell, Germany) was used for Li-O2 cell tests. The cell was assembled in an Ar-filled glove box using a Li metal foil anode (16 mm diameter), a porous glass separator (GA-55, Advantec, Japan, 17 mm diameter) wetted with 80 µL of the electrolyte, and a KB/PTFE on carbon paper cathode (16 mm diameter). Before battery tests, the atmosphere inside the test cell was substituted to pure O2. Battery tests were measured with a charge/discharge measurement instrument (HJ1001SD8, Hokuto Denko, Japan) at a constant current of 50 µA cm−2 at 30 °C. Typical KB mass in the cathode was about 0.4 mg cm−2, giving an overall mass for the KB/PTFE/carbon paper cathode of about 9 mg cm−2. For the ECMS measurement, the pre-discharged cell was connected to a GCMS-QP2010 10 ACS Paragon Plus Environment

Page 11 of 48 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Ultra (Shimadzu, Japan). The details of the ECMS measurement were described previously.11 The iodine intensity (m/z = 254) was normalized with the calibration result of O2 (m/z = 32). Cyclic Voltammetry (CV) For CV measurements, the three electrode setup consisted of a working electrode, a Li/Li+ reference electrode, and a Pt coil counter electrode. For the working electrodes, a Pt disk electrode (φ = 3 mm) and a glassy carbon (GC) electrode (φ = 3 mm) were employed for measuring redox potential of I−/I3− and O2/Li2O2 respectively. For the reference electrode, Li metal was immersed into 1 M Li[TFSA]/G4 solution separated from the CV cell solution with a liquid junction (Vycor glass). XRD measurements After discharging Li-O2 cells using LiI containing electrolytes to a capacity of 1.5 mAh cm−2, they were initially purged with dry Ar. Subsequently they were transferred, without exposure of the internal components to air, to an Ar-filled glove box, and disassembled. Discharged cathodes were washed with anhydrous dimethoxyethane (DME), dried in a vacuum chamber at ambient temperature, and then transferred into an air-tight sample holder with beryllium windows. The measurement was performed with a Rigaku Ultima IV 11 ACS Paragon Plus Environment


Page 12 of 48

using Cu Kα (λ = 0.154 nm) radiation at a scan rate of 1 ° min−1.

RESULTS AND DISCUSSION Discharge and Charge Profiles in G4

Figure 1. Initial 2 cycles for Li/O2 cells, containing (a) 1.0 M Li[TFSA] in G4, (c) 1.0 M Li[TFSA] in G4 with 100 mM LiI, (e) 1.0 M Li[TFSA] in G4 with 200 mM LiI, and 12 ACS Paragon Plus Environment

Page 13 of 48 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


(b,d,f) ECMS measurement of 1st charge, limited to 4.0 V for clarity, for Li/O2 cells containing the same electrolytes. Vertical red dotted lines in (c) and (e) represents theoretical capacity for the I−/I3− redox couple. Horizontal orange dotted lines in (b), (d) and (f) represent theoretical oxygen evolution rate based on 2e−/O2.

Figure 1 shows the discharge and charge profiles for the initial 2 cycles of Li/O2 cells using 1.0 M Li[TFSA] in G4 in the absence of LiI (a), and in the presence of 100 mM LiI (c) and 200 mM LiI (e) (extended cycling behavior of these cells is provided in the Supporting Information, Figure S1). The charge voltage gradually increased and reached ~4.1 V in the cell using pristine 1.0 M Li[TFSA]/G4 electrolyte. The overpotential during charging was greatly suppressed upon addition of LiI, in broad agreement with earlier reports.,32,33,35 As can be further observed in Figure 1(c), for the cell with 100 mM LiI, there is only one plateau in the ORR during the 1st discharge. In turn, during charge, two plateaus around 3.0 V and 3.2-3.3 V appeared, which are tentatively assigned to 3I− → I3− + 2e− and 2I3− → 3I2 + 2e−

respectively (contribution from the direct electro-oxidation of

Li2O2 may occur simultaneously as will be further elaborated based on ECMS evaluation, vide infra). It should be noted here that in order to focus on these two plateaus, we used a relatively high concentration of LiI in the electrolyte and a relatively small capacity limit of 0.25 mAh cm−2. For the 100 mM LiI containing electrolyte, the theoretical capacity for the 13 ACS Paragon Plus Environment


Page 14 of 48

I−/I3− redox couple is indicated by the vertical dotted red line in Figure 1(c) (0.07 mAh cm−2, calculated from the LiI concentration and electrolyte volume). An efficient mediation reaction via the I−/I3− redox couple (Li2O2 + I3− → 2Li+ + O2 + 3I−) should extend the first oxidation plateau, at around 3.0 V, beyond that capacity until all Li2O2 formed during discharge is completely oxidized. The observed capacity for this plateau is lower, suggesting that the I−/I3− redox couple does not operate as an effective RM. The theoretical capacity for the I3−/I2 redox couple should be half of that of the I−/I3− redox couple (based on stoichiometry). However, the length of second plateau (tentatively assigned to I3−/I2) is greater than that of the first plateau. This suggests that the I3−/I2 redox couple operates as an active RM and a mediation reaction (Li2O2 + 3I2 → 2Li+ + O2 + 2I3−) proceeds, in addition to the direct oxidation of Li2O2. This is further supported by the appearance of a plateau at around 3.0 V during the second discharge, tentatively assigned to the reduction of I3− (remaining in the electrolyte after I2 reaction with Li2O2 during charge) to I−. Burke et al.39 reported a similar two plateau appearance for the discharge in the presence of significant water contamination. In that work, DME with 0.25 M Li[TFSI] and 0.2 M LiI was used as an electrolyte, which has much lower viscosity than the electrolytes used in the present study (see Table S2). We assume here that the origin of these two plateaus was not water 14 ACS Paragon Plus Environment

Page 15 of 48 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


contamination, but suppression of the redox shuttle by relatively high viscosity of the electrolyte as will be further elaborated (vide infra). In order to confirm the active species of the mediation reaction of LiI in G4, electrochemical mass spectrometry (ECMS) analysis was also performed. Figure 1(d) shows the charge curve and online gas evolution profiles for a Li/O2 cell containing 1.0 M Li[TFSA] with 100 mM LiI, which was initially discharged to 0.25 mAh cm−2 (corresponding ECMS analysis in the absence of LiI for this electrolyte, is provided in Figure 1(b)). It should be noted here that the cell used for ECMS measurements has a slightly different configuration, see Experimental, which may to some extent rationalize the small differences in the shape of the charge curves between comparable data in Figure 1(a), (c) and (e) and Figure 1(b), (d) and (f) respectively. Namely, as will be further elaborated upon considering the redox shuttle (vide infra), the observable cycling results may also be a function of the cell type and configuration used for the measurement. As observed in the cycling tests (Figure 1(c)), two plateaus around 3.0 V and 3.2-3.3 V were observed. Oxygen was detected from the beginning of the charge. The oxygen evolution was much smaller than the theoretical value until 0.05 mAh cm−2. Since the voltage of the first plateau was above 3.0 V, both the direct oxidation of Li2O2 and the 15 ACS Paragon Plus Environment


Page 16 of 48

oxidation of I− to I3− might occur simultaneously in the initial stage of discharge, considering the theoretical capacity of I− oxidation (0.07 mAh cm−2). And then, after the first plateau, once the oxidation of I− to I3− is complete, the mediation reaction between I2 and Li2O2 along with the subsequent re-oxidation of I3− should be occurring, resulting in oxygen evolution, since there is no plateau around 3.2-3.3 V in the cell without LiI (Figure 1(b)). However, regarding the possibility for simultaneous direct electro-oxidation of Li2O2, it is noteworthy that although the potential for this reaction (second plateau in 1st cycle, Figure 1(c)) is considerably lower than the potential observed for the majority of the charge process in the absence of LiI (Figure 1(a)), in the corresponding ECMS analysis, there is non-negligible oxygen evolution below 3.3 V. This is an indication that direct electro-oxidation of Li2O2 might be occurring along with the mediation reaction at such potentials. Additionally, lower potential for direct electro-oxidation may be a result of differences in the purity/morphology of Li2O2 produced in the absence/presence of LiI64-67, some minor contribution of other iodide-related species to the oxidation, or some other unknown process. It is further noted here that similar behavior, i.e. a two-plateau charge, and two plateau discharge (excluding the first cycle) was observed for a cell cycled with a higher 16 ACS Paragon Plus Environment

Page 17 of 48 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


concentration of LiI (200 mM, see Figure 1(e)). The length (i.e. capacity) of the first charge plateau was greater than that of the 100 mM LiI system, and indeed extends to almost the full capacity. However, a plateau at around 3.0 V is again observed during the second discharge, and a lower capacity, corresponding to the I−/I3− redox couple, is observed in the second charge. This supports the suggestion that that oxidation of I− to I3− occurred during the first plateau and the mediation reaction via I−/I3− redox couple did not occur efficiently. The relatively high capacity for the I−/I3− plateau during the first charge may indicate some contribution of redox mediation, but also redox shuttle, or low overpotential direct electrooxidation of Li2O2 (vide supra). This is further supported by the corresponding ECMS measurement (Figure 1(f)), where a more distinct two-plateau profile was observed. The initial plateau around 3.0 V provided only a small amount of oxygen evolution, whilst the later plateau (3.1-3.3 V) provides relatively high oxygen evolution. Taken together, these results suggest that in this medium, I3− (at these concentrations) does not react with Li2O2 and thus the I−/I3− redox couple could not operate as an active RM in the 1.0 M Li[TFSA]/G4 electrolyte system. We note here however that a recent report by Zhou and co-workers indicates significant 17 ACS Paragon Plus Environment


Page 18 of 48

oxygen evolution for a 1 M LiI/G4 electrolyte (i.e. no additional supporting electrolyte) containing additional water from the beginning of the charge process, and conversely no oxygen evolution for a nominally dry electrolyte (where LiOH formation is proposed51). As we have taken great care to ensure our electrolytes and cells are dry, we assume that the contrasting behavior is likely caused by differences in (i) supporting electrolyte and (ii) concentration of RM (and particularly, the higher LiI concentrations used in the previously reported study may promote LiOH formation). Indeed, in our measurements we note that the potential where oxygen evolution is observed is considerably higher. It is also noteworthy that Zhou and co-workers reported similar charge curves to those observed in our study when lower LiI concentrations are employed in the presence of Li[TFSA]. In 1.0 M Li[TFSA]/G4 electrolyte, having lower viscosity, it is reasonable to expect that the RM would more easily diffuse, and hence would exhibit a more facile reaction with the Li2O2. Thus, further experiments investigated higher concentration.


Page 19 of 48 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Figure 2. Initial 2 cycles for Li/O2 cells, containing (a) 2.8 M Li[TFSA] in G4, (c) 2.8 M Li[TFSA] in G4 with 100 mM LiI, (e) 2.8 M Li[TFSA] in G4 with 200 mM LiI, and (b,d,f) ECMS measurement of 1st charge, limited to 4.0 V for clarity, for Li/O2 cells containing the same electrolytes. Vertical red dotted lines in (c) and (e) represents theoretical capacity for the I−/I3− redox couple. Horizontal orange dotted lines in (b), (d) and (f) represent theoretical oxygen evolution rate based on 2e−/O2.



Page 20 of 48

Figure 2 shows the discharge and charge profiles for the inital 2 cycles of Li/O2 cells using 2.8 M Li[TFSA] in G4 in the absence of LiI (a), and in the presence of 100 mM LiI (c) and 200 mM LiI (e) (extended cycling behavior of these cells is provided in the Supporting Information, Figure S2). As observed for 1.0 M Li[TFSA]/G4, the charge voltage gradually increased and reached 4.5 V in the cell using pristine 2.8 M Li[TFSA]/G4 electrolyte. The charge voltage was greatly suppressed, below 3.5 V, in the cell with the LiI RM. Broadly similar behavior was observed for the discharge in the first two cycles (one plateau for first discharge, two plateaus for second discharge). Also, there are two plateaus around 3.0 V and 3.2-3.3 V during charge in ECMS, Figure 2(d,f). With 100 mM LiI, during the first plateau around 3.0 V, which was assigned to the oxidation of iodide to triiodide (vide supra), there was no oxygen evolution (similar behavior was also observed for the longer plateau at higher concentration of 200 mM LiI, see Figure 2(f)). This result strongly supports the suggestion that I3− could not react with Li2O2 and the I−/I3− redox couple was not active as a RM. On the other hand, oxygen evolution was clearly observed during the second plateau, which was assigned here to the oxidation of I3− to I2, due to the observation of I2 in the ECMS analysis. The relatively high 20 ACS Paragon Plus Environment

Page 21 of 48 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


capacity contributed by this plateau further suggests that I2 reacted with Li2O2 and the I3−/I2 redox couple operated as a RM. And then, when the oxygen evolution decreased, suggesting that the amount of Li2O2 available had decreased, the I2 was no longer consumed by reaction and was detected by ECMS. This is also evidence for release of I2 from the system at a later stage of charge, and might be one of the reasons why the mediation reaction decayed during cycling (see Figure S2(b)). The lack of observation of I2 in the ECMS analysis for 1.0 M Li[TFSA]/G4 may reflect the relatively lower quantity produced, but also the relative ease of dissolving solid I2 into the electrolyte (we speculate that in the more highly concentrated electrolyte, the solid I2 on the surface of the carbon electrode may sublime giving rise to the MS signal). It is assumed here that in addition to the I2 observed by ECMS, there is also some proportion of I3− in the electrolyte at the end of the first charge, leading to the plateau at around 2.9 V (vide supra) in the second discharge. Of course, we cannot rule out the possibility of a small contribution of direct electro-oxidation of Li2O2, considering the non-negligible oxygen evolution at the early (low potential) stage of charge in the absence of LiI (Figure 2(b)), and for other factors, as already outlined for 1.0 M Li[TFSA]/G4. However, the extension of the oxygen evolution across the entire capacity range (and below 3.5 V), combined with the observation of I2, is 21 ACS Paragon Plus Environment


Page 22 of 48

strong evidence for a major contribution of the I3−/I2 redox couple as RM in this electrolyte. These results suggest that in these Li[TFSA]/G4 media (at both 1.0 M and 2.8 M concentration), I3− does not react with Li2O2 and thus the I−/I3− redox couple could not operate as an active RM.

Discharge and Charge Profiles in DMSO


Page 23 of 48 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Figure 3. Initial 2 cycles for Li/O2 cells, containing (a) 1.0 M Li[TFSA] in DMSO, (c) 1.0 M Li[TFSA] in DMSO with 100 mM LiI, (e) 1.0 M Li[TFSA] in DMSO with 200 mM LiI, and (b,d,f) ECMS measurement of 1st charge, limited to 4.0 V for clarity, for Li/O2 cells containing the same electrolytes. Vertical red dotted lines in (c) and (e) represents theoretical capacity for the I−/I3− redox couple. Horizontal orange dotted lines in (b), (d) and (f) represent theoretical oxygen evolution rate based on 2e−/O2.



Page 24 of 48

In order to investigate the effect of electrolyte solvent, we conducted the discharge and charge test using 1.0 M Li[TFSA]/DMSO in the absence of LiI (Figure 3(a)), and in the presence of 100 mM LiI (Figure 3(c)) and 200mM LiI (Figure 3(e)) (extended cycling behavior of these cells is provided in the Supporting Information, Figure S3). In the system without a RM, the charge voltage gradually increased. On the other hand, with 100 mM LiI RM (Figure 3(c,d)), entirely different discharge and charge profiles were observed. During charging, only one plateau around 3.3 V was observed and this plateau can be assigned to the oxidation of I− to I3−. Since the length of this plateau exceeds the theoretical capacity for the I−/I3− redox couple, it is suggested that I3− reacts with Li2O2 in the DMSO-based electrolyte, unlike in the G4-based electrolytes. Since it was assumed that there was no formation of molecular iodine, mediation reactions were more stable in this DMSO system than the corresponding G4 system. The online gas evolution profile during charge of a cell pre-discharged to 0.25 mAh cm−2 using 1.0 M Li[TFSA] in DMSO containing 100 mM LiI electrolyte is shown in Figure 3(d). Unlike the G4-based electrolyte systems, oxygen was immediately generated during charging in the DMSO-based electrolyte system. When the concentration of iodide was increased to 200 mM (Figure 3(e,f)), this behavior was similar. Although it should be noted 24 ACS Paragon Plus Environment

Page 25 of 48 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


that we cannot rule out the possibility of a slight contribution from direct electro-oxidation, at low potential (Figure 3(b)), as was described for the G4-based electrolytes, these results strongly suggest that I3− reacts with Li2O2 and the mediation reaction via the I−/I3− redox couple proceeds efficiently in this case.



Page 26 of 48

Figure 4. Initial 2 cycles for Li/O2 cells, containing (a) 2.8 M Li[TFSA] in DMSO, (c) 2.8 M Li[TFSA] in DMSO with 100 mM LiI, (e) 2.8 M Li[TFSA] in DMSO with 200 mM LiI, and (b,d,f) ECMS measurement of 1st charge, limited to 4.0 V for clarity, for Li/O2 cells containing the same electrolytes. Vertical red dotted lines in (c) and (e) represents theoretical capacity for the I−/I3− redox couple. Horizontal orange dotted lines in (b), (d) and (f) represent theoretical oxygen evolution rate based on 2e−/O2.

We further conducted the discharge and charge test using 2.8 M Li[TFSA]/DMSO. Figure 4 shows the discharge and charge profiles for the initial 2 cycles of Li/O2 cells using 2.8 M Li[TFSA] in G4 in the absence of LiI (a), and in the presence of 100 mM LiI (c) and 200 mM LiI (e) (extended cycling behavior of these cells is provided in the Supporting Information, Figure S4). Also in this case, the overpotential during charge was greatly suppressed with the addition of LiI. During the first discharge in the presence of LiI, there was only one plateau in the ORR. Subsequently, during the first charge, there were two plateaus around 3.1 V and 3.3 V, which were again tentatively assigned to the oxidation of I− to I3− and I3− to I2. Since the length of the first plateau was shorter than the theoretical capacity for I−/I3−, I3− could not react with Li2O2 and it is thus clear that, as observed for the corresponding G4 system, the I−/I3− redox couple was not active as a RM. During the second discharge, there was another plateau around 3.0 V which is notably different from the ORR potential observed in the first discharge, and was again assigned to the reduction 26 ACS Paragon Plus Environment

Page 27 of 48 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


of I3− to I−, as was seen for the G4 systems. Again, these assignments are supported by the ECMS measurement. In Figure 4(d), during the first part of charging, oxygen evolution was much smaller than the theoretical value until around 0.05 mAh cm−2. This behavior was also discussed for 1.0 M Li[TFSA]/G4 (Figure 1(d)), and in keeping with that system it is further noted that, upon increasing the I− to 200 mM in this 2.8 M Li[TFSA]/DMSO system (Figure 4(f)), the region where oxygen evolution was much smaller than the theoretical value was extended. It is also noted that since the value of oxygen evolution in this region is slightly higher than that in 2.8 M Li[TFSA]/G4 with 200 mM LiI (Figure 2(f)), I3− might react with Li2O2. Again, we cannot rule out the possibility of a slight contribution from direct electro-oxidation, at low potential (Figure 4(b)), as was described for the G4-based electrolytes. However, the reactivity of I3− towards Li2O2 in 2.8 M Li[TFSA]/DMSO was much lower than that in 1.0 M Li[TFSA]/DMSO due to both thermodynamic and kinetic reasons as remarked below. It has been widely recognized that the main product for the conversion reaction during discharge in a non-aqueous Li/O2 cell is lithium peroxide (Li2O2).68–70 Conversely, some reports have suggested that not only Li2O2 but also LiOH precipitates onto the cathode, when LiI is added to the electrolyte.27,35,39,41,50,51 Especially, Liu et al. indicated with solid 27 ACS Paragon Plus Environment


Page 28 of 48

state NMR data that only LiOH was present in the discharge product, despite efforts to eliminate water contamination.36 Moreover, LiOH formation was also reported in the work of Zhou and co-workers (vide supra).51 In this work, in order to identify discharge products on the cathode, XRD measurements were performed (Figure S5) for cathodes discharged to 1.5 mAh cm−2 (the capacity being larger than that used for the earlier measurements to ensure a sufficient quantity of discharge product for analysis). From the XRD measurements, only Li2O2 was detected for all of the electrolytes considered in this study i.e. both 1.0 M / 2.8 M Li[TFSA] in G4 or DMSO, with and without added LiI. The absence of clear LiOH formation is thought to be a result of great care taken in preparing nominally anhydrous electrolytes, careful drying of the cell components and so on, in order to minimize the water content of the assembled cells.

Shift of Redox Potentials of I−/I3− and O2/Li2O2 (OER) Figure 5 shows the results of the cyclic voltammetry using a three electrode cell with 10 mM LiI in (a) 1.0 M Li[TFSA]/G4 and 1.0 M Li[TFSA]/DMSO and (b) 2.8 M Li[TFSA]/G4 and 2.8 M Li[TFSA]/DMSO electrolytes on a platinum electrode measured at a 10 mV s−1 scan rate under Ar atmosphere. For both electrolytes, similar overall 28 ACS Paragon Plus Environment

Page 29 of 48 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


appearance of two pairs of cathodic and anodic peaks was observed, which were assigned to the redox of I−/I3− and I3−/I2 at lower and higher potential respectively. Bentley et al.71 reported that the formal potential of I−/I3− (E°′(I−/I3−)) varies with solvent. The electrochemical step of oxidation of I− to I3− can be considered as oxidation of I− to I2. Since iodine is a strong Lewis acid and iodide is a Lewis base, the resulting I2 then reacts with another I− to form the triiodide (I− + I2 → I3−). Thus, the redox potential of I−/I3− is affected by the interaction of solvents towards both I− and I2. In a solvent which has a high dielectric constant, iodide is more stabilized than neutral molecular iodine. In the case of DMSO and G4, since the dielectric constant of DMSO (εr = 48.0) is much higher than that of G4 (εr = 7.79)72, the redox potential of I−/I3− in DMSO should shift to a more positive potential than that in G4 due to stronger stabilization of iodide in DMSO. Moreover, DMSO has a larger Gutmann acceptor number (AN = 19.3) than that of G4 (AN = 10.5)72, which causes stronger stabilization of iodide in DMSO due to stronger interaction between DMSO and iodide. Indeed, the redox peaks of I−/I3− in DMSO shifted to a more positive potential than in G4, in both the more dilute (1.0 M) and more concentrated (2.8 M) electrolytes, indicating that iodide was more difficult to oxidize in DMSO. It is also noteworthy that the observed peak currents in 1.0 or 2.8 M Li[TFSA]/G4 were 29 ACS Paragon Plus Environment


Page 30 of 48

lower than those observed for 1.0 or 2.8 M Li[TFSA]/DMSO. This is tentatively assigned to the difference in viscosity of the electrolytes (see Table 1).

Figure 5. Cyclic voltammograms on Pt disk electrode at a scan rate of 10 mV s−1 in 1.0 M Li[TFSA] in G4 or DMSO containing 10 mM LiI; (b) 2.8 M Li[TFSA] in G4 or DMSO containing 10 mM LiI under Ar atmosphere.

It is of course also necessary to consider the potential of the OER (Li2O2 → 2Li+ + O2 + 2e−). Figure 6 shows the results of the cyclic voltammetry using a three electrode cell with (a) 1.0 M Li[TFSA] in G4 or DMSO electrolytes and (b) 2.8 M Li[TFSA] in G4 or DMSO electrolytes on a GC electrode with 10 mV s−1 scan rate under O2 atmosphere. We used pristine electrolytes, i.e. LiI free electrolytes, in order to exclude redox of iodide. The potential of anodic peak (OER) in DMSO shifted to a more negative potential than that in


Page 31 of 48 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


G4 in both 1.0 M and 2.8 M Li[TFSA] electrolytes. The OER is described as follows, considering the effect of a solvent (sol): Li2O2 + x sol → 2[Li(sol)x]+ + O2 +2e− And the electrode potential of the O2/Li2O2 redox pair (E (O2/Li2O2)) is described as follows: (O2 /Li2 O2 )=

2∆f G(Lisolx + - ∆f G(Li2 O2 ) - x∆f G(sol) 2F

Since the standard Gibbs energy of formation of Li2O2 is constant, E(O2/Li2O2) is determined by the difference between the magnitude of the standard Gibbs energy of formation of a solvated lithium cation ([Li(sol)x]+) and a free solvent, which indicates the stabilization energy of lithium cation upon solvation. DMSO has a relatively high Gutmann donor number (DN = 29.8), compared with G4 (DN = 16.6).72 The higher the DN of a solvent, the stronger the interaction with Li+. Since Li+ in DMSO is stabilized more +

strongly than that in G4, it is assumed that the ∆f G([Lisolx ] ) in DMSO is lower than that in G4, thus shifting E(O2/Li2O2) in DMSO to a more negative potential.



Page 32 of 48

Figure 6. Cyclic voltammograms on GC disk electrode at a scan rate of 10 mV s−1 in 1.0 M Li[TFSA] in G4 or DMSO; (b) 2.8 M Li[TFSA] in G4 or DMSO under O2 atmosphere.

The reaction Gibbs energy of a mediation reaction (∆r G) via the I−/I3− redox couple (Li2O2 + I3− → 2Li+ + O2 + 3I−) is described as follows: ∆r G = 2 ×F ×EO2 ⁄Li2 O2 - EI- /I3 - The reaction Gibbs energy is thus determined by the difference between E(O2/Li2O2) and E(I−/I3−). The shift of formal potentials of E(I−/I3−) and E(O2/Li2O2) by changing the electrolyte solvent was investigated by using CV measurement, as remarked above. In DMSO, higher and lower potentials of E(I−/I3−) and E(O2/Li2O2) were observed respectively, resulting in a larger magnitude of ∆r G than that in G4. The reaction rate becomes higher with an increasing magnitude of reaction Gibbs energy until reaching the Marcus inverted region. In Table 2, the difference between anodic peak potential of 32 ACS Paragon Plus Environment

Page 33 of 48 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


O2/Li2O2 and I−/I3− was calculated. The absolute value of [Eap(O2/Li2O2) − Eap(I−/I3−)] is not meaningful, since the working electrodes for measuring Eap(O2/Li2O2) and Eap(I−/I3−) were different, however consideration of the relative value is worthwhile. The value of [Eap(O2/Li2O2) − Eap(I−/I3−)] follows the order 2.8 M Li[TFSA]/G4 > 1.0 M Li[TFSA]/G4 > 2.8 M Li[TFSA]/DMSO > 1.0 M Li[TFSA]/DMSO. Thus, we can estimate that the reaction of I3− with Li2O2 was most facile in 1.0 M Li[TFSA]/DMSO and most difficult in 2.8 M Li[TFSA]/G4, which provides a strong justification for the cell test/ECMS observation that I3− in 1.0 M Li[TFSA]/DMSO could react with Li2O2, but the I−/I3− redox couple in the other electrolytes conducted in this work could not work effectively as an active RM.

Table 2. Anodic peak potentials of O2/Li2O2 (Eap(O2/Li2O2)) on GC disk electrode and I−/I3−(Eap(I−/I3−)) Pt disk electrode. Eap(O2/Li2O2)

Eap(I−/I3−)

Eap(O2/Li2O2) − Eap(I−/I3−)

[V]

[V]

[V]

on GC

on Pt

without LiI

with 10 mM LiI

1.0 M Li[TFSA]/DMSO

2.93

3.33

−0.40

2.8 M Li[TFSA]/DMSO

3.29

3.40

−0.11

1.0 M Li[TFSA]/G4

3.26

3.24

+0.02

2.8 M Li[TFSA]/G4

3.43

3.37

+0.06

Base electrolyte



Page 34 of 48

Redox Shuttle and Viscosity When employing a RM, the redox shuttle of the RM due to diffusion from the cathode to the Li anode must be considered. Using FTIR analysis, Zhang et al.38 confirmed LiI precipitation on the surface of the Li anode in 0.5 M LiClO4/DMSO containing 50 mM LiI electrolyte due to diffusion to the anode and subsequent reduction of the I3−, formed at the cathode during charging. Burke et al.39 also suggested that the extremely low charge overpotential (2.95-3.10 V) in 0.25 M Li[TFSI]/DME containing 0.2 M LiI electrolyte related in part to the redox shuttle of I−/I3− to the Li anode. So far, almost all studies employing RMs have preferably used 0.2-1.0 M lithium salts electrolyte due to their high conductivity, low viscosity, and ease of diffusion to Li2O2 in proximity to the site of iodide reduction on the cathode. However, in such electrolytes, it is assumed that RMs also can easily diffuse to Li anode and thus result in redox shuttle of the RM. Figure 7 shows the results of charge (without prior discharge) under Ar atmosphere using cells with the LiI containing electrolytes. In this situation, in the absence of the redox shuttle, the oxidation of I− to I3− should start at first and then, the oxidation of I3− to I2 should occur when the oxidation of I− cannot maintain the current. If the redox shuttle occurs, the oxidation of I− is expected to prolong due to regeneration of I− from the reduction of I3− at the Li anode. 34 ACS Paragon Plus Environment

Page 35 of 48 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


In the case of electrolytes which had relatively high viscosity (see Table 1), 1.0 M Li[TFSA]/G4 (η = 9.7 mPa s), 2.8 M Li[TFSA]/G4 (η = 106 mPa s), 2.8 M Li[TFSA]/DMSO (η = 86 mPa s), the redox shuttle of iodide was suppressed due to sluggish diffusion of I−, and subsequent oxidation of the I3− could also be observed. It should be noted that this effect may also be observed in the poor efficiency of the mediation reaction with Li2O2. On the other hand, in the case of an electrolyte which had relatively low viscosity, 1.0 M Li[TFSA]/DMSO (η = 3.9 mPa s), the redox shuttle of I− occurred and the oxidation of I− to I3− prolonged due to the ease of RM diffusion. In a Li-O2 battery, since Li2O2 must be present at the cathode, the influence of the redox shuttle might be limited if the reaction between the RM oxidant and Li2O2 is significantly faster than the diffusion of the RM. This may support the observation that in the case of 100 mM LiI in 1.0 M Li[TFSA]/DMSO electrolyte, mediation reaction by triiodide was observed although some redox shuttle might also occur. Thus, the balance between the difficulty of redox shuttle and the ease of RM diffusing to discharge products is important, and should be considered in any future cell design utilizing a RM.



Page 36 of 48

Figure 7. Charge curves (under Ar) of Li/O2 cells using (a) 1.0 M Li[TFSA] in G4; (b) 1.0 M Li[TFSA] in DMSO; (c) 2.8 M Li[TFSA] in G4; (d) 2.8 M Li[TFSA] in DMSO containing 100 mM LiI before discharging.

CONCLUSIONS In this work, we explored the relationship between the redox potential of the LiI RM and Li2O2, and reactivity of I3− towards Li2O2. We first established the phenomena observed in typically galvanostatic cell tests, then rationalized these findings with ECMS, CV and redox shuttle measurements. The redox potential of O2/Li2O2 and I−/I3− were tuned by the electrolyte solvents, which influenced the reactivity of I3− towards Li2O2. For selecting a 36 ACS Paragon Plus Environment

Page 37 of 48 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


RM, both the choice of solvents and the concentration of Li salts play a critical role from the viewpoints of redox potential and diffusivity of the RM, and indicate that great care should be taken in comparing data sets with different solvents, salts and cell configurations. With this caveat in mind, we have reported our finding that the use of 100 mM LiI in 1.0 M Li[TFSA]/DMSO provides the most effective redox mediation amongst the electrolytes we have studied, but conversely exhibits the most redox shuttling (in the absence of O2). We suggest that, on the basis of our results, further studies on this (or related systems) should take into account not only the three aspects elaborated in the Introduction i.e. (i) the redox shuttle effect (ii) the necessity for on-line measurement of gas evolution during charge and (iii) the role played by protic species creating alternative charge reaction pathways, but also (iv) the possible contribution of the mediator redox couple to the cell capacity and (v) solvation/electrolyte concentration effects on the redox potential of the RM, and its interplay with the above-mentioned considerations. Moreover, further quantitative study of the efficiency of Li2O2 production/decomposition, along with the proportions of iodine-containing species, as a function of electrolyte solvent type, electrolyte salt concentration, and state of charge, is an attractive target for future study. The use of iodometric titration, as previously reported, would be beneficial, although 37 ACS Paragon Plus Environment


Page 38 of 48

limitations due to possible interferences of I2 in the titration should be further established.39 Furthermore, variations in superoxide anion stability/solubility, which may influence Li2O2 precipitation on the separator rather than cathode,73 should also be considered; as a function of electrolyte salt concentration,11,56 and in the presence of varying mediator concentrations.

ASSOCIATED CONTENT Supporting Information Summary of data from relevant literature (Table S1), structures and boiling points of DMSO and glymes (Table S2), charge/discharge curves (all cycles) for cells with G4- or DMSO-based electrolytes in the absence of LiI and presence of 100 mM LiI and 200 mM LiI (Figure S1-4), and XRD patterns of discharged cathodes (Figure S5)

AUTHOR INFORMATION Corresponding Author [email protected] (M. W.)

Notes


Page 39 of 48 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


The authors declare no competing financial interest.

ACKNOWLEDGMENTS This study was supported in part by the “Research and Development Initiative for Scientific Innovation of New Generation Battery (RISING project)” of the New Energy and Industrial Technology Development Organization (NEDO) of Japan and the JSPS KAKENHI (No. 15H03874) from the Japan Society for the Promotion of Science (JSPS).

REFERENCES (1)

Bruce, P. G.; Freunberger, S. A.; Hardwick, L. J.; Tarascon, J.-M. Li–O2 and Li–S Batteries with High Energy Storage. Nat. Mater. 2012, 11, 19–30.

(2)

Luntz, A. C.; McCloskey, B. D. Nonaqueous Li–Air Batteries: A Status Report. Chem. Rev. 2014, 114, 11721–11750.

(3)

Freunberger, S. A.; Chen, Y.; Peng, Z.; Griffin, J. M.; Hardwick, L. J.; Bardé, F.; Novák, P.; Bruce, P. G. Reactions in the Rechargeable Lithium–O2 Battery with Alkyl Carbonate Electrolytes. J. Am. Chem. Soc. 2011, 133, 8040–8047.

(4)

Xu, W.; Xiao, J.; Wang, D.; Zhang, J.; Zhang, J.-G. Effects of Nonaqueous Electrolytes on the Performance of Lithium/Air Batteries. J. Electrochem. Soc. 2010, 157, A219.

(5)

Xu, W.; Xu, K.; Viswanathan, V. V.; Towne, S. A.; Hardy, J. S.; Xiao, J.; Nie, Z.; Hu, D.; Wang, D.; Zhang, J. G. Reaction Mechanisms for the Limited Reversibility 39 ACS Paragon Plus Environment


Page 40 of 48

of Li–O2 Chemistry in Organic Carbonate Electrolytes. J. Power Sources 2011, 196, 9631–9639. (6)

McCloskey, B. D.; Scheffler, R.; Speidel, A.; Bethune, D. S.; Shelby, R. M.; Luntz, A. C. On the Efficacy of Electrocatalysis in Nonaqueous Li–O2 Batteries. J. Am. Chem. Soc. 2011, 133, 18038–18041.

(7)

Freunberger, S. A.; Chen, Y.; Drewett, N. E.; Hardwick, L. J.; Bardé, F.; Bruce, P. G. The Lithium–Oxygen Battery with Ether-Based Electrolytes. Angew. Chem., Int. Ed. 2011, 50, 8609–8613.

(8)

Bardé, F.; Chen, Y.; Johnson, L.; Schaltin, S.; Fransaer, J.; Bruce, P. G. Sulfone-Based Electrolytes for Nonaqueous Li–O2 Batteries. J. Phys. Chem. C 2014, 118, 18892–18898.

(9)

Sharon, D.; Afri, M.; Noked, M.; Garsuch, A.; Frimer, A. A.; Aurbach, D. Oxidation of Dimethyl Sulfoxide Solutions by Electrochemical Reduction of Oxygen. J. Phys. Chem. Lett. 2013, 4, 3115–3119.

(10)

Trahan, M. J.; Mukerjee, S.; Plichta, E. J.; Hendrickson, M. A.; Abraham, K. M. Studies of Li-Air Cells Utilizing Dimethyl Sulfoxide-Based Electrolyte. J. Electrochem. Soc. 2012, 160, A259–A267.

(11)

Kwon, H.-M.; Thomas, M. L.; Tatara, R.; Oda, Y.; Kobayashi, Y.; Nakanishi, A.; Ueno, K.; Dokko, K.; Watanabe, M. Stability of Glyme Solvate Ionic Liquid as an Electrolyte for Rechargeable Li−O2 Batteries. ACS Appl. Mater. Interfaces 2017, 9, 6014–6021.

(12)

Thotiyl, M. M. O.; Freunberger, S. A.; Peng, Z.; Bruce, P. G. The Carbon Electrode in Nonaqueous Li−O2 Cells. J. Am. Chem. Soc. 2013, 135, 494–500.

(13)

McCloskey, B. D.; Speidel, A.; Scheffler, R.; Miller, D. C.; Viswanathan, V.; Hummelshøj, J. S.; Nørskov, J. K.; Luntz, A. C. Twin Problems of Interfacial Carbonate Formation in Nonaqueous Li−O2 Batteries. J. Phys. Chem. Lett. 2012, 3, 997–1001. 40 ACS Paragon Plus Environment

Page 41 of 48 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


(14)

Itkis, D. M.; Semenenko, D. A.; Kataev, E. Y.; Belova, A. I.; Neudachina, V. S.; Sirotina, A. P.; Hävecker, M.; Teschner, D.; Knop-Gericke, A.; Dudin, P.; Barinov, A.; Goodilin, E. A.; Shao-Horn, Y.; Yashina, L. V. Reactivity of Carbon in Lithium−Oxygen Battery Positive Electrodes. Nano Lett. 2013, 13, 4697–4701.

(15)

McCloskey, B. D.; Burke, C. M.; Nichols, J. E.; Renfrew, S. E. Mechanistic Insights for the Development of Li–O2 Battery Materials: Addressing Li2O2 Conductivity Limitations and Electrolyte and Cathode Instabilities. Chem. Commun. 2015, 51, 12701–12715.

(16)

Radin, M. D.; Siegel, D. J. Charge Transport in Lithium Peroxide: Relevance for Rechargeable Metal–Air Batteries. Energy Environ. Sci. 2013, 6, 2370–2379.

(17)

Meister, P.; Jia, H.; Li, J.; Kloepsch, R.; Winter, M.; Placke, T. Best Practice: Performance and Cost Evaluation of Lithium Ion Battery Active Materials with Special Emphasis on Energy Efficiency. Chem. Mater. 2016, 28, 7203–7217.

(18)

Lim, H.-D.; Lee, B.; Zheng, Y.; Hong, J.; Kim, J.; Gwon, H.; Ko, Y.; Lee, M.; Cho, K.; Kang, K. Rational Design of Redox Mediators for Advanced Li–O2 Batteries. Nat. Energy 2016, 1, 16066.

(19)

Pande, V.; Viswanathan, V. Criteria and Considerations for the Selection of Redox Mediators in Nonaqueous Li–O2 Batteries. ACS Energy Lett. 2017, 2, 60–63.

(20)

Chen, Y.; Freunberger, S. A; Peng, Z.; Fontaine, O.; Bruce, P. G. Charging a Li–OⅡ Battery Using a Redox Mediator. Nat. Chem. 2013, 5, 489–494.

(21)

Schaltin, S.; Vanhoutte, G.; Wu, M.; Bardé, F.; Fransaer, J. A QCM Study of ORR-OER and an In Situ Study of a Redox Mediator in DMSO for Li–O2 Batteries. Phys. Chem. Chem. Phys. 2015, 17, 12575–12586.

(22)

Qiao, Y.; Ye, S. Spectroscopic Investigation for Oxygen Reduction and Evolution Reactions with Tetrathiafulvalene as a Redox Mediator in Li–O2 Battery. J. Phys. Chem. C 2016, 120, 15830–15845.



(23)

Page 42 of 48

Bergner, B. J.; Schürmann, A.; Peppler, K.; Garsuch, A.; Janek, J. TEMPO: A Mobile Catalyst for Rechargeable Li-O2 Batteries. J. Am. Chem. Soc. 2014, 136, 15054–15064.

(24)

Bergner, B. J.; Hofmann, C.; Schürmann, A.; Schröder, D.; Peppler, K.; Schreiner, P. R.; Janek, J. Understanding the Fundamentals of Redox Mediators in Li–O2 Batteries: A Case Study on Nitroxides. Phys. Chem. Chem. Phys. 2015, 17, 31769– 31779.

(25)

Kundu, D.; Black, R.; Adams, B.; Nazar, L. F. A Highly Active Low Voltage Redox Mediator for Enhanced Rechargeability of Lithium–Oxygen Batteries. ACS Cent. Sci. 2015, 1, 510–515.

(26)

Sun, D.; Shen, Y.; Zhang, W.; Yu, L.; Yi, Z.; Yin, W.; Wang, D.; Huang, Y.; Wang, J.; Wang, D.; Goodenough, J. B. A Solution-Phase Bifunctional Catalyst for Lithium−Oxygen Batteries. J. Am. Chem. Soc. 2014, 136, 8941–8946.

(27)

Kwak, W.-J.; Hirshberg, D.; Sharon, D.; Afri, M.; Frimer, A. A.; Jung, H.-G.; Aurbach, D.; Sun, Y.-K. Li−O2 Cells with LiBr as an Electrolyte and Redox Mediator. Energy Environ. Sci. 2016, 9, 2334–2345.

(28)

Liang, Z.; Lu, Y. C. Critical Role of Redox Mediator in Suppressing Charging Instabilities of Lithium−Oxygen Batteries. J. Am. Chem. Soc. 2016, 138, 7574–7583.

(29)

Xin, X.; Ito, K.; Kubo, Y. Highly Efficient Br–/NO3– Dual-Anion Electrolyte for Suppressing Charging Instabilities of Li−O2 Batteries. ACS Appl. Mater. Interfaces 2017, 9, 25976–25984.

(30)

Yao, K. P. C.; Frith, J. T.; Sayed, S. Y.; Bardé, F.; Owen, J. R.; Shao-Horn, Y.; Garcia-Araez, N. Utilization of Cobalt Bis(terpyridine) Metal Complex as Soluble Redox Mediator in Li−O2 Batteries. J. Phys. Chem. C 2016, 120, 16290–16297.

(31)

Chase, G. V.; Zecevic, S.; Walker, W.; Uddin, J.; Sasaki, K. A.; Giordani, V.; Bryantsev, V.; Blanco, M.; Addison, D. D. Soluble Oxygen Evolving Catalysts for Rechargeable Metal-Air Batteries. US Patent 2012, US 20120028137A1. 42 ACS Paragon Plus Environment

Page 43 of 48 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


(32)

Lim, H. D.; Song, H.; Kim, J.; Gwon, H.; Bae, Y.; Park, K. Y.; Hong, J.; Kim, H.; Kim, T.; Kim, Y. H.; Leprõ, X.; Ovalle-Robles, R.; Baughman, R. H.; Kang, K. Superior Rechargeability and Efficiency of Lithium−Oxygen Batteries: Hierarchical Air Electrode Architecture Combined with a Soluble Catalyst. Angew. Chem., Int. Ed. 2014, 53, 3926–3931.

(33) Kim, D. S.; Park, Y. J. Effect of Multi-Catalysts on Rechargeable Li–Air Batteries. J. Alloys Compd. 2014, 591, 164–169. (34)

Yu, M.; Ren, X.; Ma, L.; Wu, Y. Integrating a Redox-Coupled Dye-Sensitized Photoelectrode into a Lithium-Oxygen Battery for Photoassisted Charging. Nat. Commun. 2014, 5, 5111.

(35)

Kwak, W.-J.; Hirshberg, D.; Sharon, D.; Shin, H.-J.; Afri, M.; Park, J.-B.; Garsuch, A.; Chesneau, F. F.; Frimer, A. A.; Aurbach, D.; Sun, Y.-K. Understanding the Behavior of Li–Oxygen Cells Containing LiI. J. Mater. Chem. A 2015, 3, 8855– 8864.

(36)

Liu, T.; Leskes, M.; Yu, W.; Moore, A. J.; Zhou, L.; Bayley, P. M.; Kim, G.; Grey, C. P. Cycling Li-O2 Batteries via LiOH Formation and Decomposition. Science 2015, 350, 530–533.

(37)

Zhu, Y. G.; Jia, C.; Yang, J.; Pan, F.; Huang, Q.; Wang, Q. Dual Redox Catalysts for Oxygen Reduction and Evolution Reactions: Towards a Redox Flow Li–O2 Battery. Chem. Commun. 2015, 51, 9451–9454.

(38)

Zhang, T.; Liao, K.; He, P.; Zhou, H. A Self-Defense Redox Mediator for Efficient Lithium–O2 Batteries. Energy Environ. Sci. 2015, 9, 1024–1030.

(39)

Burke, C. M.; Black, R.; Kochetkov, I. R.; Giordani, V.; Addison, D.; Nazar, L. F.; McCloskey, B. D. Implications of 4e– Oxygen Reduction via Iodide Redox Mediation in Li–O2 Batteries. ACS Energy Lett. 2016, 747–756.



(40)

Page 44 of 48

Zhang, W.; Shen, Y.; Sun, D.; Huang, Z.; Zhou, J.; Yan, H.; Huang, Y. Promoting Li2O2 Oxidation via Solvent-Assisted Redox Shuttle Process for Low Overpotential Li-O2 Battery. Nano Energy 2016, 30, 43–51.

(41)

Zhu, Y. G.; Liu, Q.; Rong, Y.; Chen, H.; Yang, J.; Jia, C.; Yu, L.-J.; Karton, A.; Ren, Y.; Xu, X.; Adams, S.; Wang, Q. Proton Enhanced Dynamic Battery Chemistry for Aprotic Lithium–Oxygen Batteries. Nat. Commun. 2017, 8, 14308.

(42)

Landa-Medrano, I.; Olivares-Marín, M.; Pinedo, R.; Ruiz de Larramendi, I.; Rojo, T.; Tonti, D. Operando UV-Visible Spectroscopy Evidence of the Reactions of Iodide as Redox Mediator in Li–O2 Batteries. Electrochem. Commun. 2015, 59, 24– 27.

(43)

Shen, Y.; Zhang, W.; Chou, S.-L.; Dou, S.-X. Comment on “Cycling Li-O2 Batteries via LiOH Formation and Decomposition.” Science 2016, 352, 667a.

(44)

Liu, T.; Kim, G.; Carretero-Gonzalez, J.; Castillo-Martinez, E.; Grey, C. P. Response to Comment on “Cycling Li-O2 Batteries via LiOH Formation and Decomposition.” Science 2016, 352, 667b.

(45)

Viswanathan, V.; Pande, V.; Abraham, K. M.; Luntz, A. C.; McCloskey, B. D.; Addison, D. Comment on “Cycling Li-O2 Batteries via LiOH Formation and Decomposition.” Science 2016, 352, 667c.

(46)

Liu, T.; Kim, G.; Carretero-González, J.; Castillo-Martínez, E.; Bayley, P. M.; Liu, Z.; Grey, C. P. Response to Comment on “Cycling Li-O2 Batteries via LiOH Formation and Decomposition.” Science 2016, 352, 667d.

(47)

Zeng, X.; Leng, L.; Liu, F.; Wang, G.; Dong, Y.; Du, L.; Liu, L.; Liao, S. Enhanced Li-O2 Battery Performance, Using Graphene-like Nori-Derived Carbon as the Cathode and Adding LiI in the Electrolyte as a Promoter. Electrochim. Acta 2016, 200, 231–238.

(48)

Landa-Medrano, I.; Olivares-Marín, M.; Bergner, B.; Pinedo, R.; Sorrentino, A.; Pereiro, E.; Ruiz De Larramendi, I.; Janek, J.; Rojo, T.; Tonti, D. Potassium Salts as 44 ACS Paragon Plus Environment

Page 45 of 48 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Electrolyte Additives in Lithium–Oxygen Batteries. J. Phys. Chem. C 2017, 121, 3822–3829. (49)

Yoon, S. H.; Park, Y. J. Polyimide-Coated Carbon Electrodes Combined with Redox Mediators for Superior Li-O2 Cells with Excellent Cycling Performance and Decreased Overpotential. Sci. Rep. 2017, 7, 42617.

(50)

Tułodziecki, M.; Leverick, G. M.; Amanchukwu, C. V.; Katayama, Y.; Kwabi, D. G.; Bardé, F.; Hammond, P. T.; Shao-Horn, Y. The Role of Iodide on the Formation of Lithium Hydroxide in Lithium–Oxygen Batteries. Energy Environ. Sci. 2017, 10, 1828–1842.

(51)

Qiao, Y.; Wu, S.; Sun, Y.; Guo, S.; Yi, J.; He, P.; Zhou, H. Unraveling the Complex Role of Iodide Additives in Li–O2 Batteries. ACS Energy Lett. 2017, 2, 1869–1878.

(52)

McCloskey, B. D.; Addison, D. A Viewpoint on Heterogeneous Electrocatalysis and Redox Mediation in Nonaqueous Li-O2 Batteries. ACS Catal. 2017, 7, 772–778.

(53)

Li, Z.; Ganapathy, S.; Xu, Y.; Heringa, J. R.; Zhu, Q.; Chen, W.; Wagemaker, M. Understanding the Electrochemical Formation and Decomposition of Li2O2 and LiOH with Operando X-Ray Diffraction. Chem. Mater. 2017, 29, 1577–1586.

(54)

Kwak, W.-J.; Hirshberg, D.; Sharon, D.; Shin, H.-J.; Afri, M.; Park, J.-B.; Garsuch, A.; Chesneau, F. F.; Frimer, A. A.; Aurbach, D.; Sun, Y.-K. Understanding the Behavior of Li–Oxygen Cells Containing LiI. J. Mater. Chem. A 2015, 3, 8855– 8864.

(55)

Peng, Z.; Freunberger, S. A.; Hardwick, L. J.; Chen, Y.; Giordani, V.; Bardé, F.; Novák, P.; Graham, D.; Tarascon, J. M.; Bruce, P. G. Oxygen Reactions in a Non-Aqueous Li+ Electrolyte. Angew. Chem., Int. Ed. 2011, 50, 6351–6355.

(56)

Tatara, R.; Kwabi, D. G.; Batcho, T. P.; Tulodziecki, M.; Watanabe, K.; Kwon, H.-M.; Thomas, M. L.; Ueno, K.; Thompson, C. V.; Dokko, K.; Shao-Horn, Y.; Watanabe, M. Oxygen Reduction Reaction in Highly Concentrated Electrolyte



Page 46 of 48

Solutions of Lithium Bis(trifluoromethanesulfonyl)amide/Dimethyl Sulfoxide. J. Phys. Chem. C 2017, 9162–9172. (57)

Suo, L.; Hu, Y.-S.; Li, H.; Armand, M.; Chen, L. A New Class of Solvent-in-Salt Electrolyte for High-Energy Rechargeable Metallic Lithium Batteries. Nat. Commun. 2013, 4, 1481.

(58)

Yamada, Y.; Yamada, A. Review—Superconcentrated Electrolytes for Lithium Batteries. J. Electrochem. Soc. 2015, 162, A2406–A2423.

(59)

Watanabe, M.; Thomas, M. L.; Zhang, S.; Ueno, K.; Yasuda, T.; Dokko, K. Application of Ionic Liquids to Energy Storage and Conversion Materials and Devices. Chem. Rev. 2017, 117, 7190–7239.

(60)

Tamura, T.; Yoshida, K.; Hachida, T.; Tsuchiya, M.; Nakamura, M.; Kazue, Y.; Tachikawa, N.; Dokko, K.; Watanabe, M. Physicochemical Properties of Glyme–Li Salt Complexes as a New Family of Room-Temperature Ionic Liquids. Chem. Lett. 2010, 39, 753–755.

(61)

Yoshida, K.; Tsuchiya, M.; Tachikawa, N.; Dokko, K.; Watanabe, M. Change from Glyme Solutions to Quasi-Ionic Liquids for Binary Mixtures Consisting of Lithium Bis(trifluoromethanesulfonyl)amide and Glymes. J. Phys. Chem. C 2011, 115, 18384–18394.

(62) Yoshida, K.; Nakamura, M.; Kazue, Y.; Tachikawa, N.; Tsuzuki, S.; Seki, S.; Dokko, K.; Watanabe, M. Oxidative-Stability Enhancement and Charge Transport Mechanism in Glyme-Lithium Salt Equimolar Complexes. J. Am. Chem. Soc. 2011, 133, 13121–13129. (63)

Zhang, C.; Ueno, K.; Yamazaki, A.; Yoshida, K.; Moon, H.; Mandai, T.; Umebayashi, Y.; Dokko, K.; Watanabe, M. Chelate Effects in Glyme/Lithium Bis(trifluoromethanesulfonyl)amide Solvate Ionic Liquids. I. Stability of Solvate Cations and Correlation with Electrolyte Properties. J. Phys. Chem. B 2014, 118, 5144-5153.


Page 47 of 48 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


(64) Gallant, B. M.; Kwabi, D. G.; Mitchell, R. R.; Zhou, J.; Thompson, C. V; Shao-Horn, Y. Influence of Li2O2 Morphology on Oxygen Reduction and Evolution Kinetics in Li–O2 Batteries. Energy Environ. Sci. 2013, 6, 2518–2528. (65) Radin, M. D.; Monroe, C. W.; Siegel, D. J. How Dopants Can Enhance Charge Transport in Li2O2. Chem. Mater. 2015, 27, 839–847. (66) Sharon, D.; Hirsberg, D.; Salama, M.; Afri, M.; Frimer, A. A.; Noked, M.; Kwak, W.; Sun, Y. K.; Aurbach, D. Mechanistic Role of Li+ Dissociation Level in Aprotic Li-O2 Battery. ACS Appl. Mater. Interfaces 2016, 8, 5300–5307. (67) Matsuda, S.; Kubo, Y.; Uosaki, K.; Hashimoto, K.; Nakanishi, S. Improved Energy Capacity of Aprotic Li–O2 Batteries by Forming Cl-Incorporated Li2O2 as the Discharge Product. J. Phys. Chem. C 2016, 120, 13360–13365. (68)

Freunberger, S. A.; Chen, Y.; Peng, Z.; Griffin, J. M.; Hardwick, L. J.; Bardé, F.; Novák, P.; Bruce, P. G. Reactions in the Rechargeable Lithium–O2 Battery with Alkyl Carbonate Electrolytes. J. Am. Chem. Soc. 2011, 133, 8040–8047.

(69)

McCloskey, B. D.; Bethune, D. S.; Shelby, R. M.; Girishkumar, G.; Luntz, A. C. Solvents’ Critical Role in Nonaqueous Lithium–Oxygen Battery. J. Phys. Chem. Lett. 2011, 2, 1161–1166.

(70)

Gallant, B. M.; Mitchell, R. R.; Kwabi, D. G.; Zhou, J.; Zuin, L.; Thompson, C. V; Shao-Horn, Y. Chemical and Morphological Changes of Li–O2 Battery Electrodes upon Cycling. J. Phys. Chem. C 2012, 116, 20800–20805.

(71)

Bentley, C. L.; Bond, A. M.; Hollenkamp, A. F.; Mahon, P. J.; Zhang, J. Voltammetric Determination of the Iodide/Iodine Formal Potential and Triiodide Stability Constant in Conventional and Ionic Liquid Media. J. Phys. Chem. C 2015, 119, 22392–22403.

(72)

Laoire, C. O.; Mukerjee, S.; Abraham, K. M.; Plichta, E. J.; Hendrickson, M. A. Influence of Nonaqueous Solvents on the Electrochemistry of Oxygen in the Rechargeable Lithium–Air Battery. J. Phys. Chem. C 2010, 114, 9178–9186. 47 ACS Paragon Plus Environment


(73)

Page 48 of 48

Kwon, H.-M.; Thomas, M. L.; Tatara, R.; Nakanishi, A.; Dokko, K.; Watanabe, M. Effect of Anion in Glyme-based Electrolyte for Li-O2 Batteries: Stability/Solubility of Discharge Intermediate. Chem. Lett. 2017, 46, 573-576.

TOC Graphic


O2 Batteries with LiI ... - ACS Publications

Recommend Documents