Article pubs.acs.org/JPCC
Limited Stability of Ether-Based Solvents in Lithium−Oxygen Batteries Kate R. Ryan, Lynn Trahey, Brian J. Ingram, and Anthony K. Burrell* Electrochemical Energy Storage Department, Chemical Sciences and Engineering Division, Argonne National Laboratory, Argonne, Illinois, Untied States S Supporting Information *
ABSTRACT: Li−O2 batteries offer the tantalizing promise of a specific energy much greater than current Li ion technologies; however, many challenges remain before the development of commercial energy storage applications based on the lithium−oxygen couple can be realized. One of the most apparent limitations is electrolyte stability. Without an electrolyte that is resistant to attack by reduced oxygen species, optimizing other aspects of the redox performance is challenging. Thus, identifying electrolyte decomposition processes that occur early in the redox process will accelerate the discovery process. In this study, ATR−FTIR was used to examine various reported Li−O2 electrolytes taken directly from the cell separators of cycled electrochemical cells. Specifically, we examined, 1 M LiPF6 in propylene carbonate (PC), 1 M LiCF3SO3 in tetraethyleneglycoldimethylether (TEGDME), and 1 M LiCF3SO3 in a siloxane ether (1NM3) and looked for soluble decomposition products. Each electrolyte was tested using a regular Li−O2 cathode with no catalyst and either an O2 atmosphere or an Ar atmosphere and a Li metal anode as well as in a Li−Li symmetric cell. The 1NM3 electrolyte was found to form soluble decomposition products under all cell conditions tested, and a decomposition pathway has been proposed. It was also found that 1NM3 and TEGDME were consumed as part of the charging process in a working Li−O2 cell, even at moderate voltages in the absence of O2.
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for Li ion batteries,6,7 they are decomposed by the Li−O2 discharge products8 to irreversibly form lithium carbonate and other related compounds.8,9 Consequently, there is now a strong drive to develop new, more stable electrolytes for Li−O2 batteries. Ethers are the main solvents under investigation, in particular, various glymes10,11 and dimethoxyethane (DME). Glymes, although more stable than the carbonates, still decompose on discharge with carbonate formation8 and appear to be consumed on charging.11 Conversely, although stable on discharge, DME oxidizes and consumes peroxide at high voltages during charging.12,13 There is also an increasing amount of interest in functionalized silanes, such as 2-[2-[2-[2-methoxy]ethoxy]ethoxy]ethoxytrimethyl silane (1NM3) and related compounds. These silanes have been reported to offer several advantages over the ethers and organic carbonates, including higher chemical stability and lower flammability, and are more environmentally benign.14 In their 2011 paper, Zhang et al.15 claimed 1NM3 was stable with respect to the Li−O2 discharge products and did not form lithium carbonate. They also attributed the observed reduction in charge overpotential for 1NM3-containing cells to favorable peroxide/oxide formation on discharge. All of the decomposition products mentioned in these studies were solids found on the cathode. Although the
INTRODUCTION The Li−O2 battery has the potential to lead the next generation of battery technology. Projections place its specific energy at around four times that of the best Li ion batteries, ∼800 Wh/ kg.1,2 However the use of the term battery in this system has led many research groups to oversimplify the chemistry that is required to work at 100% conversion to actually act as a rechargeable battery. The generally accepted reactions in a rechargeable lithium oxygen cell are: Li ↔ Li+ + e−
2Li+ + 2e− + O2 ↔ Li 2O2
(E 0 = 2.96 V vs Li/Li+)
This simple representation of these reactions is deceptive. In reality, they are highly complex systems with a prodigious number of parameters to optimize and side reactions to minimize. The oxygen reduction reaction presumably progresses in two one-electron steps with the disproportionation reaction of the one electron superoxide competing against the further reduction to peroxide.3 In terms of minimizing the number of competing side reactions (e.g. the formation of LiCO3), electrolyte stability is one of the most important parameters. The nonaqueous electrolytes required to give the highest energy density4 are particularly subject to attack from the reduced oxygen within the electrochemical environment, reactive radical species, and even the Li metal itself.5 Until recently, organic carbonates (e.g., propylene and ethylene carbonate) were the most commonly used electrolytes for Li−O2 batteries; however, while these carbonates work well © 2012 American Chemical Society
Received: July 9, 2012 Revised: August 10, 2012 Published: August 23, 2012 19724
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under constant current at a slow rate, typically 50−125 mA/g of carbon. ATR-FTIR measurements of the samples were obtained on a Perkin-Elmer Spectrum 100 using a Ge crystal, housed in an Ar glovebox. The tested cells were disassembled in the glovebox, and the harvested separators were placed directly onto the Ge crystal. Reference spectra were collected by moistening a separator with a few drops of pristine electrolyte. All spectra were normalized, and difference plots were obtained by subtracting the normalized cell spectra from the reference spectra.
formation of lithium carbonates is a facile process from organic carbonates, the formation of similar species in the ether-based electrolytes requires multiple electron steps, and solid carbonate or acetates are not necessarily the initial products. Detection of carbon dioxide or carbonates may not necessarily indicate solvent decomposition, until significant decomposition has occurred, and may arise from other carbon sources, for example, the carbon cathode. To fully understand the decomposition process occurring within the Li−O2 cell, it is important to characterize any species that may remain soluble in the electrolyte. Unfortunately, extracting the electrolyte from a disassembled cell for characterization with conventional liquid analytical techniques, for example, GC/MS, is very difficult. Only small quantities of electrolyte are used to make a cell which becomes fully absorbed into the separator during cycling. Under these conditions, attenuated total reflectance (ATR)FTIR is an ideal technique for analyzing the electrolytesaturated separators.16 Another advantage of FTIR is its sensitivity, with detection limits as low as a few parts per million.17 NMR would be a powerful technique for identifying the exact structures of the impurities; however, NMR cannot readily identify compounds present at CH2) bend.20 The C−O stretch is likely to be located at ∼1250 cm−1. Second, they failed to note significant changes in their IR spectra which indicate the electrolyte had altered. Two bands originating from the C−O stretch at 1250 cm−1 and a C−C stretch at 840 cm−1 disappeared. Furthermore, the appearance of a new band at 1622 cm−1 in the CC stretching region was not noted in the text. The peak is seen in the spectra of all three of their cell components, albeit very faintly in their anode spectra. Examination of Figure 2 indicates a reduction in the bands at 840 cm−1 and the emergence of the band at 1622 cm−1. 19726
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decomposition products under the cell conditions tested. However, the electrochemical stability of the ether solvents may not be as great as previously thought. 1NM3 and TEGDME are consumed as part of the charging process in working lithium oxygen cells, even at moderate voltages in the absence of O2. In the presence of reduced oxygen species, ether electrolytes decompose to intermediates containing alkenes even before carbonate is produced.
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ASSOCIATED CONTENT
S Supporting Information *
FTIR spectra of separators from disassembled Li−Li cells run with PC and TEGDME electrolytes. FTIR spectra of separators from disassembled carbon cathode cells run with PC and TEGDME in O2 and Ar chambers. This material is available free of charge via the Internet at http://pubs.acs.org.
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AUTHOR INFORMATION
Corresponding Author
*Phone: 1-630-252-2629. E-mail:
[email protected]. Author Contributions
The manuscript was written through contributions of all authors. All authors have given approval to the final version of the manuscript.
Figure 3. Graph of the cycling data from the three cells run in an Ar environment; 5 h charge and discharge @ 0.05 mA (4.5 V cut off).
Notes
The authors declare no competing financial interest.
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interest that the PC does not oxidize at similar low potentials, consistent with its being more stable to electrochemical oxidation than the ethers 1NM3 and TEGDME. FTIR spectra collected from the soaked separators from the disassembled cells are shown in Figure 2 and the SI. The 1NM3 cell (Figure 2) :charged: to 3.5 V and did show the same spectral changes as described in the previous section. Not surprisingly, no changes were observed in the PC spectrum (SI) because the cell did not reach a plateau region. However, despite reaching a plateau of 4.25 V and the TEGDME spectrum (SI) was also unchanged indicating its oxidation products are insoluble in the electrolyte. These results indicate that 1NM3 and TEGDME are consumed as part of the “charging” process in lithium oxygen cells, even at moderate voltages. This has important implications for Li−O2 research and data interpretation of 1NM3 and TEGDME as a component of a “stable” electrolyte,10 as fully reversible cells are impossible if electrolyte is being consumed during the charge process, and from an electrochemical point of view, the observed voltages and current capacities are the sum of several different reactions. It is therefore extremely difficult to know the contribution of the desired reactions vs the electrolyte decomposition, which may give the illusion of a reversible cell by giving artificially high charge capacities.
ACKNOWLEDGMENTS Financial support from the U.S. Department of Energy is gratefully acknowledged. The submitted manuscript has been created by UChicago Argonne, LLC, Operator of Argonne National Laboratory (“Argonne”). Argonne, a U.S. Department of Energy Office of Science laboratory, is operated under Contract No. DE-AC02-06CH11357. Argonne National Laboratory assisted in meeting the publication costs of this article.
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REFERENCES
(1) Bruce, P. G.; Freunberger, S. A.; Hardwick, L. J.; Tarascon, J. M. Nat. Mater. 2012, 11 (1), 19−29. (2) Christensen, J.; Albertus, P.; Sanchez-Carrera, R. S.; Lohmann, T.; Kozinsky, B.; Liedtke, R.; Ahmed, J.; Kojic, A. J. Electrochem. Soc. 2012, 159 (2), R1−R30. (3) Laoire, C. O.; Mukerjee, S.; Abraham, K. M.; Plichta, E. J.; Hendrickson, M. A. J. Phys. Chem. C 2009, 113 (46), 20127−20134. (4) Zheng, J. P.; Liang, R. Y.; Hendrickson, M.; Plichta, E. J. J. Electrochem. Soc. 2008, 155 (6), A432−A437. (5) Shao, Y.; Park, S.; Xiao, J.; Zhang, J.-G.; Wang, Y.; Liu, J. ACS Catal. 2012, 844−857. (6) Balbuena, P. B.; Wang, Y., Lithium-Ion Batteries: Solid-Electrolyte Interphase; Imperial College Press: London, 2004. (7) Aurbach, D.; Talyosef, Y.; Markovsky, B.; Markevich, E.; Zinigrad, E.; Asraf, L.; Gnanaraj, J. S.; Kim, H. J. Electrochim. Acta 2004, 50 (2−3), 247−254. (8) Freunberger, S. A.; Chen, Y.; Drewett, N. E.; Hardwick, L. J.; Barde, F.; Bruce, P. G. Angew. Chem. Int Ed. 2011, 50 (37), 8609− 8613. (9) Mizuno, F.; Nakanishi, S.; Kotani, Y.; Yokoishi, S.; Iba, H. Electrochemistry 2010, 78 (5), 403−405. (10) Jung, H. G.; Hassoun, J.; Park, J. B.; Sun, Y. K.; Scrosati, B. Nat. Chem. 2012, 2, 579. (11) Ó Laoire, C.; Mukerjee, S.; Plichta, E. J.; Hendrickson, M.; Abraham, K. M. J. Electrochem. Soc. 2011, 158 (3), A302−A308.
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CONCLUSIONS The following conclusions can be drawn from the combined electrochemical and spectroscopic study of the Li−O2 electrolytes 1 M LiPF6 in propylene carbonate, 1 M LiCF3SO3 in tetraethylene glycol dimethyl ether, and 1 M LiCF3SO3 in 2-[2[2-[2-methoxy]ethoxy]ethoxy] ethoxy trimethyl silane. The 1NM3 electrolyte was found to form soluble decomposition products under all cell conditions tested. In contrast, the TEGDME and PC electrolytes did not form any soluble 19727
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(12) McCloskey, B. D.; Bethune, D. S.; Shelby, R. M.; Girishkumar, G.; Luntz, A. C. J. Phys. Chem. Lett. 2011, 2 (10), 1161−1166. (13) Wang, H.; Xie, K. Electrochim. Acta 2012, 64, 29−34. (14) Dong, J.; Zhang, Z.; Kusachi, Y.; Amine, K. J. Power Sources 2011, 196 (4), 2255−2259. (15) 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 (51), 25535−25542. (16) FT-IR SpectroscopyAttenuated Total Reflectance (ATR). http://shop.perkinelmer.com/content/TechnicalInfo/TCH_ FTIRATR.pdf, 2005. (17) Acha, V.; Meurens, M.; Naveau, H.; Agathos, S. N. Biotechnol. Bioeng. 2000, 68 (5), 473−487. (18) Burrell, A. K.; Del Sesto, R. E.; Baker, S. N.; McCleskey, T. M.; Baker, G. A. Green Chem. 2007, 9 (5), 449−454. (19) Clayden, J. Organolithiums: Selectivity for Synthesis, 1st ed.; Pergamon: Amsterdam, Boston; 2002; Vol. 23, p 400. (20) Coates, J. Interpreation of Infrared Spectra, A Practical Approach. In Encycolpedia of Analytical Chemistry; Meyers, R. A., Ed.; John Wiley and Sons Ltd: Chichester, 2000; pp 10815−10837. (21) Oberhammer, H.; Boggs, J. E. J. Am. Chem. Soc. 1980, 102 (24), 7241−7244. (22) Gembus, V.; Marsais, F.; Levacher, V. Synlett 2008, 10, 1463− 1466. (23) Fitt, J. J.; Gschwend, H. W. J. Org. Chem. 1984, 49 (1), 209− 210.
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