(DMDMP): An Aprotic Solvent Designed for Stability in Li-O2 cells

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2,4-Dimethoxy-2,4-dimethylpentan-3-one (DMDMP): An Aprotic Solvent Designed for Stability in Li-O2 cells Daniel Sharon, Pessia Sharon, Daniel Hirshberg, Michael Salama, Michal Afri, Linda JW Shimon, Won-Jin Kwak, Yang-Kook Sun, Aryeh A. Frimer, and Doron Aurbach J. Am. Chem. Soc., Just Accepted Manuscript • DOI: 10.1021/jacs.7b06414 • Publication Date (Web): 08 Aug 2017 Downloaded from http://pubs.acs.org on August 8, 2017

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Journal of the American Chemical Society

2,4-Dimethoxy-2,4-dimethylpentan-3-one (DMDMP): An Aprotic Solvent Designed for Stability in Li-O2 cells

Daniel Sharon a, Pessia Sharon a, Daniel Hirshberg a, Michael Salama a, Michal Afri a, Linda J.W Shimon b, Won-Jin Kwak c, Yang-Kook Sun c, Aryeh A. Frimer a and Doron Aurbach a a.

b. c.

Department of Chemistry and BINA (BIU Institute for Nano-Technology and Advanced Materials) , Bar Ilan University, Ramat-Gan, 5290002, Israel Department for Chemical Research Support, Weizmann Institute of Science, Rehovot 76100, Israel Department of Energy Engineering, Hanyang University, Seoul, 04763, South Korea

Supporting Information Placeholder ABSTRACT: In this study we present a new aprotic solvent, 2,4-dimethoxy-2,4-dimethylpentan-3-one (DMDMP), which is designed to resist nucleophilic attack and hydrogen abstraction by reduced oxygen species. Li-O2 cells using DMDMP solutions were successfully cycled. By various analytical measurements we showed that even after prolonged cycling only a negligible amount of DMDMP was degraded. We suggest that the observed capacity fading of the Li-O2 DMDMP-based cells was due to instability of the lithium anode during cycling. The stability towards oxygen species makes DMDMP an excellent solvent candidate for many kinds of electrochemical systems which involve oxygen reduction and assorted evaluation reactions.

The search for suitable aprotic solvents for Li-O2 cells has been an ongoing pursuit for almost a decade. Once researchers recognized that the standard carbonate-based media are reactive towards reduced oxygen species1, studies on other candidates began in earnest. The first family of solvents that was suggested were the generally inert polyether solvents.2 Surprisingly, these solvents were found to degrade during ORR and possess a low oxidation potential.3 Focus turned next to a study of the compatibility of amides and sulfoxides (such as DMF, DMA and DMSO) with LiO2 cells.4,5 In addition to noticeable degradation during the ORR, these solvents proved reactive towards lithium metal anodes.6,7 The conclusion from these studies was that neither amides nor sulfoxides were completely suitable for Li-O2 cells. After a thorough but disappointing screening process of most available solvents, our research has continued on two fronts. The first approach is to optimize the current available solvents by using additives, changing the cell setup, or simply by highlighting the selection process of the solvent. The second approach is to

synthesize new solvents modified to resist the harsh conditions existing in Li-O2 cells. Synthesizing new solvents for batteries is a very challenging task. Many factors such as purity, electrochemical window and solventelectrodes interface must be taken into account. In the particular case of Li-O2 batteries, there is another important feature that must be mentioned. During the ORR, a variety of reduced oxygen species such as O2-, O2-2, OH- and HO2- are formed. These species are very reactive and therefore can degrade the electrolyte solution. Hence, candidates proposed for synthesis need to be inert to such attack. Because of these strict requirements, there are not many examples of such studies in the field of metal-oxygen batteries. Some researchers have concentrated on chemical modification of existing aprotic solvents. Zhang et al. introduced a newly modified polyether solvents.8 The backbone of triglyme was modified by substituting one terminal methyl group with a bulky trimethylsilane moiety to form 2,2-dimethyl-3,6,9,12-tetroxa-2-silatridecane (1NM3). Another polyether modification was proffered by Adams et al.9 By substituting the internal protons of the dimethoxyethane (DME) backbone with methyl groups, they formed 2,3dimethyl-2,3-dimethyoxybutane (DMDMB). These authors proposed that by replacing the reactive backbone protons, assorted reactions of the solvent with superoxide radical would be avoided. Indeed, DMDMB showed improved cycling performance as compared to that of DME. Nonetheless, cycled DMDMB in 1NM3 cells exhibited the formation of side-products that are typical of ethereal solution decomposition. In the present work, we report the synthesis of a novel polar aprotic solvent, 2,4-dimethoxy-2,4-dimethylpentan-3-one (DMDMP) according to the steps outlined in Scheme 1. DMDMP combines

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the polarity of a carbonyl and the complexing ability of ether groups.

Scheme 1. Synthetic route to 2,4-dimethoxy-2,4-dimethyl-3pentanone (DMDMP, 3) CH3 CH3

O C

C C

Br

Br

CH3 CH3

K2CO3 / H2O

CH3

reflux, 4 days CH3

1 CH3 1. NaH, THF, 0.5hrs. 2. CH3I, 24 hrs.

CH3

O C C

C

OH

OH

2

O C C

C

CH3 CH3

CH3

CH3 OCH 3 OCH3 3

As noted in the introduction, our goal in synthesizing DMDMP was to prepare a solvent which would resist degradation by reduced oxygen species. Two modes of attack by these reactive species are feasible. One would be removal of acidic protons in a basic/E2 elimination process, while the other would involve a nucleophilic attack on the carbonyl carbon. If we compare DMDMP with acetone (CH3COCH3), we see that both these processes seem to be prevented by the fact that in DMDMP bears no acidic hydrogens alpha to the carbonyl. We designed the DMDMP in such way that none of the hydrogens are activated by strongly electron withdrawing groups. The pKa of the hydrogens in DMDMP should be around 50-56 .10 The replacement of the alpha hydrogens with methyl or methoxy groups in DMDMP also means that the carbonyl is sterically protected from attack by a bulky methyl or methoxy groups. Thus DMDMP is structurally similar to di-t-butylketone.

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Figure 1. (a) The FTIR spectrum carbonyl region of pure DMDMP and solutions containing different LiTFSI concentrations (b) structure for DMDMP(2)·LiTFSI from SCXRD analysis. In the current study we choose to assemble Li-O2 cells with electrolyte solution that consists of 0.2 M LiTFSI in DMDMP which had an ionic conductivity of 0.644 mS/cm at 25 °C. To charaterize the DMDMP-Li+-TFSI- solvated complex, we measured the FTIR spectra of pure DMDMP solvent and solutions containing different LiTFSI concentrations. The carbonyl region is shown in Figure 1a. The C=O stretching mode for pure DMDMP was observed at 1715 cm-1. When 0.2 M LiTFSI was introduced to the solvent, the carbonyl peak broadened and shifted slightly to 1714 cm-1. The broadening and shifting of the C=O stretching mode is the result of Li+-carbonyl coordination. When we increased the concentration to 0.6 M LiTFSI, the C=O stretch was further shifted to 1713 cm-1. In addition a new shoulder appeared at 1706 cm-1. The separation of the carbonyl bending into two peaks is sufficient to distinguish between the coordinated and uncoordinated DMDMP. We used Single Crystal X-ray Diffraction (SCXRD) to associate the vibrational changes with the spatial structure of the solution complex. The single crystals were grown from a saturated solution of LiTFSI in DMDMP. The SCXRD analysis generated the structure shown in Figure 1b. From the SCXRD we can observe that the Li+ cation is penta-coordinated. The Li+ is solvated by two DMDMP molecules. Each DMDMP is coordinated to the Li+ via the carbonyl and one methoxide oxygens. A sulfone oxygen from the TFSI - is also coordinated to the Li+. Therefore, we can describe the LiTFSI-(DMDMP)2 as a contact ion pair. The low conductivity of the electrolyte solution can also indicate the formation of a soluble contact ion-pair in the solution.

Although the DMDMP ketone moiety might undergo nucleophilic attack by either O2- or O2-2, nevertheless in the absence of a suitable leaving group, the attack will reverse without solvent decomposition. The ethereal terminal groups are also very poor leaving groups and, hence, protected from nucleophilic attack.

Figure 2. (a) Cyclic voltammetry (1 mV/s) on Au working electrodes versus lithium reference in argon and oxygen; (b) Voltage profiles in oxygen and argon of the first cycles of Li-O2 cells; (c) Voltage profiles during prolong limited capacity cycling; and (d) The obtained cycling performance. All measurements were conducted with 0.2 M LiTFSI DMDMP electrolyte solutions. Figure 2a exhibits the voltammetric response of 0.2 M LiTFSI DMDMP in argon and oxygen atmosphere on gold working electrode and lithium counter. Sweeping the voltage anodically to 2 V in argon does not produce a noticeable current response. However, during the cathodic sweep around 4.25 V, a strong oxidation peak appears. This response represents the stability limits of the solvent.

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When sweeping anodically in oxygen atmosphere, a broad reduction peak around 2.4 V is observed, which is correlated to a typical ORR in aprotic solvents. During the subsequent cathodic sweep, we can observe three different peaks, around 3.3, 3.8 and 4.3 V. The oxidation process occurring between 3 and 4.2 V can be correlated to the OER, while the process above these values can originate from the electrolyte solution oxidation. The galvanostatic discharge and charge process profile of 0.2 M LiTFSI DMDMP cells are presented in Figure 2b. The galvanostatic response in oxygen resembles the voltammetric response in Figure 1. The voltage slope between 2.5 and 2.3 V represents the broad ORR peak in Figure 1. The charge process was set to allow full coulombic reversibility. The multiple voltage steps in the galvanostatic mode during the OER can be correlated to the peaks that were observed in the voltammetric mode. It was clear that we obtain fully reversible reduction of oxygen to Li2O2 and reoxidation to O2, as was reflected by XRD measurements of the electrodes and in-situ pressure measurements in Figure S1 and S2 (SI). The cycling performances of Li-O2 cells with 0.2 M LiTFSI DMDMP are presented in Figure 2c. The discharge capacity of the cell was restricted to 0.64 mAh/cm2. The capacity of the cell starts to fade after 40 cycles (Figure 2d). The fading process suggests that the ORR and OER might involve parasitic reactions with the cell components such as the electrolyte solution, carbon cathode and lithium anode. To identify side-products from the electrolyte solution decomposition, we performed 1H and 13C NMR measurements on cells containing 0.2 M LiTFSI DMDMP in different charging states. The spectra of the samples are presented in Figure S3 (SI). Samples were taken from both residue electrolyte solution and from extractions of the cathodes and separators, by washing them with CDCl3 or D2O. The samples that were run after the first discharge and charge process did not exhibit additional peaks to those of the pristine DMDMP solvent. The absence of side products was verified by GC-MS (Figure S3d) of the NMR sample. This outstanding behavior proves that our solvent is extremely durable to attack by active oxygen species. In the extracted sample of the Li-O2 cell that was cycled until failure (after 48 cycles), we could identify very small amounts of additional side products: 2,4-dimethyl-3pentanone (4; diisopropyl ketone) and 2-methoxy-2,4-dimethyl-3pentanone (5). A proposed mechanism for the formation of these products is shown in Scheme S1.

Figure 3. (a) Photos of lithium metal anodes after prolonged cycling in Li-O2 cell containing 0.2 M LiTFSI DMDMP; and (b) after ageing in DMDMP. (c) Voltage profile; and (d) coulombic

efficiency of lithium deposition/dissolution at 1 µA cm-2 on flat copper electrodes using 0.2 M LiTFSI DMDMP electrolyte solutions. The negligible amount of side products found in the electrolyte solution and cathode cannot, by themselves, explain the failure mechanism. The image of lithium metal anodes from Li-O2 cells that were cycled in 0.2 M LiTFSI DMDMP are presented in Figure 3a. We observe that black dots appeared on the lithium surface, which can be associated to the formation of dendrites. We also observe that the morphology and color of the lithium metal is not homogeneous. The lithium disk that was aged in DMDMP did not show any color or morphological changes (Figure 3b). Thus, we can assume that the observable changes over the metal anode are more related to the dissolution and deposition process of lithium ions - and not to the possible inherent chemical reactivity of the solvent with the lithium metal. To further investigate the lithium dissolution/deposition efficiency in DMDMP based solutions, we assembled cells with working Li metal and counter Cu electrodes. The voltage profiles of cells with 0.2 M LiTFSI DMDMP at argon atmosphere are presented in Figure 3c. The reduction process (negative potentials) corresponds to the deposition of Li on top of the Cu substrate, while the charge process corresponds to the dissolution of Li from the Cu electrode. During the first discharge process, the voltage crosses the 0 V (versus Li) potential quickly from the starting OCV value of 2.4 V. Presumably, negative overpotentials are required for the deposition of the lithium, due to the nucleation process occurring on the Cu surface.11 However, the sharp drop of the voltage curve to -0.6 V and absence of stable voltage plateaus implies that the traditional nucleation and growth mechanism of the lithium metal layer is not obtained. During the next cycles, the reduction overpotential increases until it reaches the set voltage cutoff of -1.2 V. After 15 cycles, a voltage plateau builds up near the cutoff voltage. Because of the high reduction potential, we assume that the deposition of lithium is mixed with parasitic reactions, such as electrolyte solution reduction. The high overpotential during the charging process implies that the dissolution of Li from Cu is also inefficient. Figure 3d presents the columbic efficiency of lithium dissolution/deposition on the Cu electrode. After efficiency increase during the first cycles, due to irreversible formation reactions11, the efficiency stabilized around 50%. These results indicate that solutions containing DMDMP solvent do not allow reasonable Li metal deposition and dissolution. For comparison, we present in Figure S4 (SI) the behavior of lithium deposition/dissolution on Cu electrodes in another popular lithium battery solvents - EC-DMC (ethylene carbonate - dimethyl carbonate), DMSO (dimethyl sulfoxide) and diglyme. These solvents exhibit better lithium deposition processes even at a higher current density of 10 µA cm-2. Nevertheless, we note that except of the EC-DMC based solution which is not relevant to Li-O2 applications, neither diglyme nor DMSO based solutions reached 100 % columbic efficiency. In light of these results, we suggest that the instability of the lithium anode is the primary reason for the Li-O2 cell failure in cells containing DMDMP solvent. It is likely that during the first cycles of the Li-O2 DMDMP based cells, the excess source of lithium within the metal anode can enable reasonably "stable" ORR and OER on the cathode. However, due to inefficient dissolution/deposition of lithium on the anode surface, the capacity

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drops. The deactivation of the lithium metal can result from two parallel processes. First, because of the high overpotential used in the deposition of lithium, the electrolyte solution is in danger of being reduced. This can cause the consumption of the solution and passivation of the lithium surface by side products. Second, due to the uncontrolled manner in which the lithium is deposited on the anode, dendritic growth occurs, which leads to local shorts and finally to cell failure.

Experimental procedures of all measurements and methods, synthesis of 2,4-dihydroxy-2,4-dimethyl-3-pentanone, synthesis of 2,4-dimethoxy-2,4-dimethyl-3-pentanone, XRD pattern of cycled cathodes, cycling behavior of Li-O2 cells containing LiBr or LiNO3, proposed degradation mechanism, behavior of lithium deposition/dissolution on Cu electrodes in different lithium battery solvents. This material is available free of charge via the Internet at http://pubs.acs.org

Our attempts to add to the DMDMP solutions LiNO3 as a lithium passivation agent or LiBr as a redox-mediator are described in the supporting information (Figure S6).12,13,14 Neither showed any improvements to the Li-O2 cell performance. We suspect that this intrinsic challenge requires more extreme measures. Among all the possible directions, we suggest that the replacement of the lithium metal by alternative anode material, such as alloys or carbonaceous materials, might be the best chance for solvents such as DMDMP.15, 16,17 It is also too early to rule out the option of using pretreated lithium metals by methods such as ALD18 or chemical treatment.17 In order to eliminate the detrimental effect of the Li anodes, thus concentrating in optimization of electrolyte solutions for most efficient ORR and OER in Li-O2 cells, we developed recently a new methodology of bi-compartments Li-O2 cells. In these cells the Li anode is cycled in fluorinated ethylene carbonate (FEC) based solutions while being fully isolated by a selective Li ions conducting membrane from the cathode side. In several FEC based solutions Li metal anodes behave fully reversibly. This methodology is an ideal tool to explore exclusively the cathode side in Li-oxygen cells.19 Work with DMDMP solutions using these cells were initiated (beyond the scope of this communication)

AUTHOR INFORMATION

The aim of this work was to demonstrate a new strategy in advancing Li-oxygen battery technology by developing a new aprotic solvent which is resistant to the reduced active oxygen species that are formed during the Li-O2 cell operation. By sensitive analytical methods we have shown that only a minor amount of solvent decomposes during prolonged cycling of Li-O2 cells containing DMDMP LiTFSI solutions. The enhanced stability of the solvent contributed to the protection of sensitive functional groups - such as the ketone and terminal glycols of the DMDMP solvent. We also propose a rational mechanism for the unavoidable side reactions that are occurring. After we established the solvent stability on the cathode side, we investigated its suitability for the counter lithium anode. We found that DMDMP based solutions do not allow efficient lithium deposition and dissolution during cycling. We suggested that the inefficient lithium deposition/dissolution is the primary reason for the capacity fading and cell failure.

Corresponding Author

*[email protected]

ACKNOWLEDGMENT Acknowledgements: A.A.F. thanks the Israel Science Foundation (ISF; Grant No. 1469/13) as well as the Ethel and David Resnick Chair in Active Oxygen Chemistry for their kind and generous support. A partial support for this study was also obtained by the Israeli Committee of High Education and the Prime-Minister office in the framework of the INREP project.

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(2) (3)

(4) (5) (6) (7) (8)

(9)

(10) (11) (12) (13)

(14)

The introduction of DMDMP is a step in the right direction. With that being said, a solution for one parameter of the battery oftentimes exposes the remaining challenges even more. Therefore, systematic work on the solvent-anode stability should be carried out in parallel with attempts performed on the solvent-cathode side in Li-O2 cells. Because it is hard to separate between the different challenges of the Li-O2 cell, the experimental isolation of every cell component is crucial for the development of these promising batteries.

ASSOCIATED CONTENT

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(16) (17) (18) (19)

Freunberger, S. A; Chen, Y.; Peng, Z.; Griffin, J. M.; Hardwick, L. J.; Bardé, F.; Novák, P.; Bruce, P. G. J. Am. Chem. Soc. 2011, 133, 8040. Read, J. J. Electrochem. Soc. 2006, 153, A96. Freunberger, S. A.; Chen, Y.; Drewett, N. E.; Hardwick, L. J.; Bardé, F.; Bruce, P. G. Angew. Chem. Int. Ed. Engl. 2011, 50, 8609. Trahan, M. J.; Mukerjee, S.; Plichta, E. J.; Hendrickson, M. A.; Abraham, K. M. J. Electrochem. Soc. 2012, 160, A259. Walker, W.; Giordani, V.; Uddin, J.; Bryantsev, V. S.; Chase, G. V; Addison, D. J. Am. Chem. Soc. 2013, 135, 2076. Chen, Y.; Freunberger, S. a; Peng, Z.; Bardé, F.; Bruce, P. G. J. Am. Chem. Soc. 2012, 134 (18), 7952. Sharon, D.; Hirsberg, D.; Afri, M.; Garsuch, A.; Frimer, A. A.; Aurbach, D. J. Phys. Chem. C 2014, 118, 15207. Zhang, Z.; Lu, J.; Assary, R. S.; Du, P.; Wang, H.-H.; Sun, Y.K.; Qin, Y.; Lau, K. C.; Greeley, J.; Redfern, P. C.; Iddir, H.; Curtiss, L. A.; Amine, K. J. Phys. Chem. C 2011, 115, 25535. Adams, B. D.; Black, R.; Williams, Z.; Fernandes, R.; Cuisinier, M.; Berg, E. J.; Novak, P.; Murphy, G. K.; Nazar, L. F. Adv. Energy Mater. 2015, 5, 1400867. Bordwell, F. G. Acc. Chem. Res. 1988, 2, 456. Bieker, G.; Winter, M.; Bieker, P. Phys. Chem. Chem. Phys. 2015, 17, 8670. Liang, Z.; Lu, Y. J. Am. Chem. Soc. 2016, 138, 7574. Sharon, D.; Hirsberg, D.; Afri, M.; Chesneau, F.; Lavi, R.; Frimer, A. a.; Sun, Y.-K.; Aurbach, D. ACS Appl. Mater. Interfaces 2015, 7, 16590. Kwak, W.-J.; Hirshberg, D.; Sharon, D.; Afri, M.; Frimer, A. A.; Jung, H.-G.; Aurbach, D.; Sun, Y.-K. Energy Environ. Sci. 2016, 9, 2334. Hirshberg, D.; Sharon, D.; De La Llave, E.; Afri, M.; Frimer, A. A.; Kwak, W.-J.; Sun, Y.-K.; Aurbach, D. ACS Appl. Mater. Interfaces 2017, 9, 4352. Wu, S.; Zhu, K.; Tang, J.; Liao, K.; Bai, S.; Yi, J.; Yamauchi, Y.; Ishida, M.; Zhou, H. Energy Environ. Sci. 2016, 9, 3262. Wu, H.; Cao, Y.; Geng, L.; Wang, C. Chem. Mater. 2017, 29, 3572. Wood, K. N.; Noked, M.; Dasgupta, N. P. ACS Energy Lett. 2017, 664. Kwak, W.-J.; Jung, H.-G.; Aurbach, D.; Sun, Y.-K. Adv. Energy Mater. 2017, in press DOI: 10.1002/aenm.201701232.

Supporting Information

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(DMDMP): An Aprotic Solvent Designed for Stability in Li-O2 cells

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