Influence of Temperature on LithiumâOxygen Battery Behavior

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Influence of the temperature on the lithium-oxygen battery behavior Jin-Bum Park, Jusef Hassoun, Hun-Gi Jung, Hee-Soo Kim, Chong Seung Yoon, I. Oh, Bruno Scrosati, and Yang-Kook Sun Nano Lett., Just Accepted Manuscript • DOI: 10.1021/nl401439b • Publication Date (Web): 16 May 2013 Downloaded from http://pubs.acs.org on May 19, 2013

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Influence of the temperature on the lithium-oxygen battery behavior Jin-Bum Parka,+, Jusef Hassounb,+,*, Hun-Gi Junga, Hee-Soo Kimc, Chong Seung Yoonc, In-Hwan Ohd, Bruno Scrosatib,* and Yang-Kook Suna,e,* a Department of Energy Engineering and c Department of Materials Science and Engineering, Hanyang University, Seoul 133-791, South Korea b Department of Chemistry, University of Rome Sapienza, 00185, Rome, Italy d Fuel Cell Center, Korea Institute of Science and Technology, Seoul 136-791, South Korea e Department of Chemistry, King Abdulaziz University, P. O. Box 80203, Jeddah 21589, Saudi Arabia +

These authors equally contributed to this work

*Corresponding authors: [email protected]; [email protected]; [email protected] key words: lithium-air cell; temperature; electrochemical process Abstract In this paper we report an electrochemical and morphological study of the response of lithiumoxygen cells cycled at various temperatures, i.e. ranging from -10°C to 70 °C. The results show that the electrochemical process of the cells is thermally influenced in an opposite way, i.e., by a rate decrease, due to a reduced diffusion of the lithium ions from the electrolyte to the electrode interface, at low temperature and a rate enhancement, due to the decreased electrolyte viscosity and consequent increased oxygen mobility, at high temperature. In addition, we show that the temperature also influences the crystallinity of lithium peroxide, namely of the product formed during cell discharge.

Introduction Due to its very high value of energy density (1,2), the lithium-oxygen battery has received considerable attention worldwide. Recent results have demonstrated that the battery, if operated under proper conditions, may provide acceptable values of cycle life and capacity (3-5). In its most classical configuration, the Li/O2 battery is formed by a lithium metal anode, a liquid organic electrolyte and a carbon-supported (with or without catalyst) air electrode (6-9). Recently, a new ACS Paragon Plus Environment

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configuration, where the lithium metal is replaced by a lithium alloy silicon anode, has also been reported (10). Key parameters in assuring proper Li/O2 battery behavior are i): the optimization of the positive electrode structure, in terms of use of an adequate gas diffusion layer and of effective catalysts (11,12) and ii): the choice of an electrolyte stable to superoxide attack (3,5,13,14). It is in fact well known that the basic electrochemical cell process, leading to the reversible formation and dissolution of lithium peroxide, involves an intermediate oxygen anion radical O2-. (8,9,14), namely a highly reactive base that readily attacks and decomposes conventional electrolytes, such as organic carbonate solutions (15-17). Di-methoxy ethane, DME-based (18) and ionic liquid-based (19,20) solutions have been proposed as alternative electrolyte media, however with scarce success. Recent works demonstrated that best results in terms of Li/O2 battery stability and cycling, may be obtained with the use of long chain, ether-based glymes, such as tetra ethylene glycol dimethyl ether (TEGDME) electrolyte solutions (3,5,7,10,13). In a previous paper we have reported a detailed Transmission Electron Microscopy, TEM study showing that Li/O2 batteries based on the TEGDME electrolyte indeed show a very promising behavior at room temperature (21). In this paper we extend the study by investigating the role of temperature on influencing the response of Li/TEGDME/O2 cells.

Results and discussion The lithium-oxygen cells studied in this work are based on a specially developed Gas Diffusion Layer, GDL oxygen electrode and on a tetraethylene glycol dimethyl ether-lithium triflate, TEGDME-LiCF3SO3 electrolyte. The gas diffusion layer was coated with Super-P carbon as matrix to host the lithium oxygen reaction products that are mainly constituted by Li2O2 nanospheres and hollow nanospheres formed at the interface with the tetraglyme-based solution. Work in our (3,21) as well as in other (5,13,14) laboratories has demonstrated that this electrolyte is

ACS Paragon Plus Environment

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stable in the Li/O2 cell environment allowing long charge-discharge cycling life with high specific capacity and good energy efficiency. FIGURE 1 Figure 1 shows the initial cycling response of a Li/O2 cell at various temperatures, reported in terms of voltage profile versus capacity. All the cycles are run under a fixed capacity regime of 1,000 mAh g-1carbon. As expected, the cell polarization, particularly at the charge, OER, process, decreases as the temperature increases. For instance, moving from 25°C to 50°C, the charging maximum voltage passes from about 4.2V to about 3.8V, compare figures 1a and 1b. Even more remarkable decrease is observed at 70°C where the maximum voltage is reduced to about 3.4V, see figures 1c. This temperature-related trend may be associated with an enhancement of the electrode kinetics, in turn promoted by an increasing amount of oxygen dissolution in the TEGDME-based electrolyte. It is in fact the oxygen transported by the electrolyte that mostly contributes to the ongoing of the cell reaction. The increase of the oxygen transport at the higher temperature is most likely associated with the decrease of the electrolyte viscosity (7), this inducing a favorable solid/liquid interphase: the higher mobility of the oxygen reduces the polarization, supposedly for both charge and discharge processes, although a prevalence is expected for the former since it is the one affected by the greater kinetic limits (6,7,14). The oxygen mobility and its diffusion in the electrolyte is influenced by its viscosity according to the well known Stokes-Einstein relation, as already demonstrated by the work of J. Read using ether-based electrolyte (22), reporting that once a certain level of oxygen solubility is reached, the decreased viscosity of the electrolyte becomes the determining factor in improving the rate capability and the cell performances. Furthermore, the almost constancy of the oxygen solubility within the considered temperature range in this class of solvents (23) indicates that the oxygen amount transported to the cathode side is determined by the electrolyte viscosity.


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In line with this scenario, the polarization strongly increases as the temperature is lowered. For example, at 0°C and at -10 °C the charge voltage rises above 4.2V, see Figures 1d, and 1c. We have plotted in Figure 1f the dependence of cell charge and discharge voltages as function of temperature. The Figure reflects the above considerations by showing that the charge-discharge gap decreases with increasing temperature. It is however to be remarked that, even at a temperature as low as -10°C, the Li/O2 cell is able to operate for 10 cycles at a capacity of 1,000 mAhg-1carbon . At the best of our knowledge, there have not been previous reports demonstrating a similar stable cell response in a comparable wide temperature range. We believe that this unique behavior may only be achieved if the cell benefits by a very stable electrolyte as in fact is the here adopted TEGDMELiCF3SO3 solution. FIGURE 2 To confirm the above conclusion, we have performed an ex-situ NMR analysis of the oxygen electrode after two galvanostatic discharge/charge tests, run at 25°C and at 70°C, respectively, see figures 2a and 2b. The NMR has been considered as a well suitable and sensitive technique for the determination of the electrolyte decomposition products, as demonstrated by S.A. Freunberger et. al. (24) using the same tetraglyeme solvent with a different, more aggressive salt, i.e.LiPF6, in a lithium air cell. In our electrolyte, only the signals associated with TEGDME are detected with no evidence of possible reaction decomposition products, such carbonates, alcohols or other secondary compounds (15-17). The NMR data give then a very convincing validation of the stability of the TEGDME-based electrolyte, even when tested in the wide range of temperatures. The TEM analysis, combined with selected area diffraction patterns (SADP), of the electrode at the discharged state reported in figures 2c and 2d evidences some differences in terms of Li2O2 crystallinity formed at the two temperatures. Although the morphology is basically the same, i.e., of hollow-sphere-like type, the Li2O2 formed on the electrode discharged at 25 oC is partially crystalline (Li2O2, JCPDS#-73-1640 as by XRD, figure S1 in supplementary information section),


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see Figure 2c, while those operated at 70°C and at 50 oC show mostly amorphous Li2O2, see Figure 2d and figure S2 in supplementary information section, respectively. The reason of this difference is not yet totally clear; one, however, may reasonably assume that the temperature variation promotes favorable changes in the electrode/electrolyte interface. FIGURE 3 In the attempt to shed more light on this point, we have investigated the kinetics of the oxygen electrode process by cyclic voltammetry, CV. Figures 3a and 3b show the response at decreasing and increasing temperatures, respectively. The cycling voltammetry is performed by an “ad-hoc’ experimental set up using a low scan rate to avoid an excessive cell polarization, this in order to well determine the process reversibility at the various tested temperatures by comparing the voltammetric peak magnitude. Accordingly, the height of the oxidation peak at 3.5V is plotted versus temperature in Figure 3c; the trend is very complex and its simulation required the use of a polynomial equation of the 4th order. We tentatively assume that the peak height, i.e. the magnitude of the cell charge process, is ruled by four different parameters, namely: i) the oxygen transport and the TEGDME-based electrolyte viscosity (7,22); ii) the lithium diffusion in the electrolyte bulk (25); iii) the interphase charge transfer process (26) and iv) the kinetics of the Li-O2 reaction (27). Although all these parameters are sensitive to the temperature variation, we believe that the most critical effects to rule the electrode kinetics are mainly two, i.e., i): limits in lithium ions diffusion in the electrolyte in the low temperature range and ii): increase of the oxygen transport in the high temperature range induced by the decrease of the electrolyte viscosity (7,22), see Figure S3 in Supplementary Information, with consequent positive reflections in the Li/O2 cell response in terms of both kinetics and capacity (6,22,26) and effects on the electrode morphology, see also additional TEM images in Figure S4 in Supplementary Information section.

Conclusion


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The issue of the electrolyte stability in the lithium-air battery is still under debate and various alternatives are proposed worldwide. One of them is the tetraethylene glycol dimethyl etherlithium triflate, TEGDME-based solution. There have been evidences in many laboratories that this solution is indeed a suitable electrolyte for assuring a stable performance of the battery (3,5,13,14,16,21) and the results reported in this work confirm the validity of the choice. Another aspect highly investigated is the study of the parameters that may influence the kinetics of the electrochemical process of the lithium-air battery. In this work we show that the temperature of operation plays an important role and, to the best of our knowledge for the first time, that a lithiumair battery, fabricated under optimized electrode and electrolyte conditions, may successfully operate in a temperature range extending from -10°C to 70°C.

Experimental 1. Li-O2 cell preparation and electrochemical measurements A carbon material (Super P) and a polyvinylidene fluoride binder (PVDF) were intimately mixed with a weight ratio of 8:2 in a N-methyl-2-pyrrolidone (NMP) solvent. Then, the resulting slurry was coated on a gas-diffusion layer (TGP-H-030 carbon paper, Torray) with carbon loading density of 1.0 ± 0.1 mgcarbon cm-2. The prepared electrode was dried for 12 hours at 100 °C under vacuum to remove the residual solvent. R2032 coin type Li-O2 cells were assembled in an argonfilled glove box (MBRAUN, H2O < 0.1 ppm, O2 < 0.1 ppm). The positive top cover was machinedrilled to evenly distributed 21 × 1.0 mm diameter holes to provide the oxygen flow. The cell consisted of a metallic lithium foil anode (400 µm thick), the SP carbon-GDL cathode and glass filter (Whatman®) separator. A solution of LiCF3SO3 (Aldrich) in a tetra (ethylene glycol) dimethyl ether (TEGDME) solvent with a molar ratio of 1:4, respectively, was used as the electrolyte. The LiCF3SO3 (Aldrich) salt was vacuum-dried at 100 °C for 24 hours prior to use. All solvents were also dried for several days over activated, dry, molecular sieves in order to completely remove


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traces of water. The prepared electrolytes had a moisture contents of < 10 ppm determined using a Mettler-Toledo Karl Fischer instrument. The Li-O2 cells were electrochemically tested by charge-discharge galvanostatic cycling under oxygen atmosphere carried out with a VMP3 Biologic-instrument on the time-controlled mode at a current density of 100 mA gcarbon-1 for 10 hours (1000 mAh gcarbon-1 capacity limit). For the NMR and TEM study the cells were cycled on the time-controlled mode at a current density of 500 mA gcarbon-1 for 20 hours (10,000 mAh gcarbon-1) . 2. Characterization of the oxygen electrode after cycling A high-resolution transmission electron microscopy (HR-TEM, model JEM-2010, JEOL) was employed to observe the morphology of cycled air electrode samples. The samples were protected from exposure to air during the transfer to the TEM chamber by conductive tape and carbon grid applied in the argon-filled glove box. The 1H NMR analysis was carried out using a portion of electrolyte extracted from the tested cells collected from the top of the GDL-electrode, disassembled in argon filled globe box, in DCCl3 (chloroform-d6) as solvent, and tetramethylsilane (TMS) as a reference on a VNMRS 600 MHz instrument. The eventual Li-Air decomposition products are here studied indirectly by NMR, since lithium peroxide and lithium carbonate are not soluble in DCCl3 or DMSO NMR. Mainly, these decomposition products can be seen in their hydrated form, i.e. H2O, HCO2Li, or as CO2 as already reported by literature papers on Li-Air batteries cycled in ether based electrolyte (24). Furthermore, NMR can reveal organic materials such as bifuctional(ether and ester) that are soluble in organic NMR solvents and produced by the eventual electrolyte decomposition.

Acknowledgments This work was supported by the Human Resources Development program (No. 20124010203310) of the Korea Institute of Energy Technology Evaluation and Planning (KETEP)


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grant funded by the Korea government Ministry of Knowledge Economy and by Regione Lazio project, Italy.

List of Figures Figure 1. Cycling behavior of TEGDME-based electrolyte lithium-oxygen cells at 25 oC (a), 50 oC (b), 70 oC (c), 0 oC (d), -10 oC (e). The figure also shows the cell voltage vs. temperature trend (f). Capacity limited to1000 mAh g-1Carbon. Current: 100 mA g-1Carbon. All the tests were performed using TEGDME4LiCF3SO3 electrolyte.. Figure 2. Ex-situ NMR analyses of the lithium-air cathode after galvanostatic discharge and charge at 25 oC (a) and at 70 oC (b). TEM images of the lithium-air cathode after galvanostatic discharge at 25 oC (c) and at 70 oC (d). Capacity limited to10,000 mAh g-1Carbon. Current: 500 mA g-1Carbon. All the tests were performed using TEGDME4LiCF3SO3 electrolyte. Figure 3. Cycling voltammetry profiles of the lithium-oxygen cell performed at -10°C, 0°C and 25 o

C (a), 50 oC and 70 oC (b). 3.5V-peak oxidation current vs. temperature curve (c). All the tests

were performed using TEGDME4LiCF3SO3 electrolyte. Scan rate: 0.1 mV s-1.

References 1. Scrosati, B. Hassoun J. & Sun, Y-K., Lithium ion batteries. Energy & Environmental Science, 4, 3287–3295 (2011) 2. Bruce, P.G., Freunberger, S. A., Hardwick, L. J., Tarascon, J.M., Nature Materials, 11, 19-29 (2012) 3. Jung, H.-G., Hassoun, J., Park, J.-B., Sun, Y.-K. & Scrosati, B., Nature Chemistry 4, 579–585 (2012) 4. Peng, Z., Freunberger, S.A., Chen, Y., Bruce, P.G., Science 337 (2012) , pp. 563-566


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5. Oh, S. H., Black, R., Pomerantseva, E., Lee, J.-H. & Nazar, L. F., Nature Chemistry 4, 1004– 1010 (2012) 6. A. Débart, A.J. Paterson, J. Bao and P.G. Bruce, Angew. Chem. Int. Ed. 2008, 47, 4521-4524. 7. C. O. Laoire, S. Mukerjee, K. M. Abraham, E. J. Plichta, M. A. Hendrickson, J. Phys. Chem. C 2010, 114, 9178–9186. 8. Hassoun, J., Croce, F., Armand, M. & Scrosati, B., Angew. Chem. Int. Ed. 2011, 50, 2999 – 3002 9. Shao, Y., Ding, F., Xiao, J., Zhang, J., Xu, W., Park, S., Zhang, J.-G., Wang, Y., Liu, J., Adv. Funct. Mater, 2013, 23, 987–1004 10. Hassoun, J., Jung, H.-G., Lee, D.-J., Park, J.-B., Amine, K., Sun, Y.-K., & Scrosati, B., Nano Lett., 2012, 12 (11), pp 5775–5779 11. Harding, J. R., Lu, Y.-C., Tsukada, Y. & Shao-Horn, Y., Phys. Chem. Chem. Phys., 2012, 14, 10540–10546 12. Cao, Y., Wei, Z., He, J., Zang, J., Zhang, Q., Zheng, M., & Dong, Q., Energy Environ. Sci., 2012, 5, 9765 13. Laoire, C. Ó, Mukerjee, S., Plichta, E. J., Hendrickson, M. A., & Abraham, K. M., Journal of The Electrochemical Society, 158 (3) A302-A308 (2011) 14. Laoire, C. O., Mukerjee, S., & Abraham, K. M., J. Phys. Chem. C 2010, 114, 9178–9186 15. Freunberger, S. A., Chen, Y., Peng, Z., Griffin, J. M., Hardwick, L. J., Barde, F., Novak, P., & Bruce, P. G., J. Am. Chem. Soc. 2011, 133, 8040–8047 16. Peng, Z., Freunberger, S. A., Hardwick, L. J., Chen, Y., Giordani, V., Barde, F., Novak, P., Graham, D., Tarascon, J.-M., & Bruce, P. G., Angew. Chem. Int. Ed. 2011, 50, 6351 –6355 17. Xu, W., Xu, K., Viswanathan, V. V., Towne, S. A., Hardy, J. S., Xiao, J., Nie, Z., Hu, D., Wang, D., Zhang, J.-G., Journal of Power Sources 196 (2011) 9631– 9639


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18. McCloskey, B. D., Bethune, D. S., Shelby, R. M., Girishkumar, G., & Luntz. A. C., Phys. Chem. Lett. 2011, 2, 1161–1166 19. Cecchetto, L., Salomon, M., Scrosati, B., Croce, F., Journal of Power Sources 213 (2012) 233238 20. Das, S.K., Xu, S., Emuas, A.H., Lu, Y.Y., Srivastava, S., Archer, L.A., Energy Environ. Sci., 2012, 5, 8927-8931 21. Jung, H.-G., Kim, H.-S., Park, J.-B., Oh, I.-H., Hassoun, J., Yoon, C. S., Scrosati, B. & Sun, Y.K., Nano Lett., 2012, 12 (8), 4333–4335 22. Read, J., J. Electrochem. Soc., 2006, 153, A96-A100 23. Fischer, K., Wilken, M., Chem. Thermodynamics, 2001, 33, 1285–1308 24. S. A. Freunberger, Y. Chen, N. E. Drewett, L. J. Hardwick, F. Bard, P. G. Bruce, Angew. Chem. Int. Ed. 2011, 50, 8609 –8613 25. D. Aurbach, J. Power Sources, 2000, 89, 206–218 26. Lu, Y.-C. , Xu, Z., Gasteiger, H.A., Chen, S., Hamad-Schifferli, K., Shao-Horn, Y., J. Am. Chem. Soc., 132, 12170–12171 (2010) 27. M. Park, X. Zhang, M. Chung, G.B. Less, A.M. Sastry, J. Power Sources, 2010, 195, 7904– 7929


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Fit Data1_B to y0+A1e^(-x/t1): Chi^2 0.00137 R^2 0.99719 --------------------------------------------y0 3.21403 0.19958 A1 1.32293 0.20978 t1 59.15837 13.92565 ---------------------------------------------

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Parameter Value Error -----------------------------------------------------------A -0.95861 0 B1 0.01146 0 B2 -2.81618E-4 0 B3 -3.30131E-6 0 B4 1.17231E-7 0 -----------------------------------------------------------R-Square(COD) SD N P -----------------------------------------------------------1 0 5

Influence of Temperature on LithiumâOxygen Battery Behavior

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Influence of Temperature on LithiumâOxygen Battery Behavior