Concentration Effects on the Entropy of Electrochemical Lithium

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Concentration Effects on the Entropy of Electrochemical Lithium Deposition: Implications for Li+ Solvation Matthias J. Schmid,† Junjie Xu,† Jeannette Lindner,† Petr Novák,‡ and Rolf Schuster*,† †

Institut für Physikalische Chemie, Karlsruhe Institute of Technology, Fritz-Haber-Weg 2, D-76131 Karlsruhe, Germany Electrochemistry Laboratory, Paul Scherrer Institut, CH-5232 Villigen PSI, Switzerland

‡

ABSTRACT: The solvation behavior of Li+ in ethylene carbonate and dimethylcarbonate upon dilution has been investigated by electrochemical microcalorimetry. We measured the heat effects at a Li electrode upon electrochemical Li deposition and dissolution from Li+ solutions of varying concentration. The exchanged heat is correlated to the entropy of lithium deposition and therefore reveals information about the solvation of Li+. Lithium deposition from electrolytes with lower concentrations showed less entropy gain than deposition from electrolytes with higher concentrations. This can be fully explained by the entropy of dilution of the Li+ ions. From our data we further concluded that the inner coordination shell of the Li+ ions does not significantly change between 0.01 and 1 M Li+ concentration. The results also suggest that contact ion pairs, which are probably present in the electrolyte, show a similar solvation behavior as regular lithium ions.

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INTRODUCTION Lithium-ion batteries have become the standard for energy storage in many fields, e.g., in portable electronic devices.1 Ongoing research efforts aim to extend their use in new fields such as electromobility. Li-ion batteries operate with an electrolyte consisting of a lithium salt in organic solvents. Usually a mixture of cyclic and acyclic carbonates is employed in order to combine advantageous properties such as high dielectric constant, low viscosity, low melting point, high boiling point, and high conductivity.2,3 Particularly the viscosity and conductivity of the electrolyte are a function of the concentration of the lithium salt. At high concentrations lithium and its counterion may form neutral ion contact pairs, which do not contribute to the conductivity.4 Additionally the coordination number of the lithium ion may change with both, concentration and electrolyte composition.2,4−7 Information on the behavior of solvated lithium and the molecular structure of lithium-solvent aggregates can be helpful to study cointercalation,8 and the formation of layers covering the electrode (also called solid electrolyte interphase, SEI),9 and to design electrolytes.10 The solvation of lithium ions in battery electrolytes has therefore been studied by several methods, including Raman- and IR-spectroscopy,4 DFT- and MDcalculations,6,8 electro spray ionization mass spectroscopy,11 NMR-spectroscopy,5,7,9 and conductivity measurements.10 In this study we measured heat effects during the electrochemical deposition of lithium from a solution of LiPF 6 in a mixture of ethylene carbonate (EC) and dimethylcarbonate (DMC) using electrochemical microcalorimetry (see refs 12−16 and references therein). In electrochemical reactions the reversibly exchanged heat in a half cell reaction, also called Peltier heat, corresponds to the entropy change of the process.12 The entropy change in turn contains contributions from all species involved in the reaction, © XXXX American Chemical Society

including the solvation shell. Previously we studied the deposition of lithium from a 1 M solution to investigate the solvation of lithium and to determine the number of solvating molecules.17 Here we extend the scope to lower concentrations in order to gain information on the solvation behavior of Li+ upon dilution, possible changes in coordination, and the influence of contact ion pair formation.

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METHOD AND EXPERIMENTAL PROCEDURE In this work we used an electrochemical microcalorimeter which can be operated both in a glovebox as well as under ambient conditions. With this calorimeter it is possible to measure heat effects of electrochemical reactions with submonolayer conversions, i.e., close to the thermal equilibrium. The idea of the method is to study the heat generation of a half cell reaction, in our case the oxidation of lithium and the reduction of lithium ions, in order to gain information about the entropy of the process. We used a thin gold-coated polyvinylidene fluoride (PVDF) foil as temperature sensor. Close contact between the working electrode and the PVDF foil enabled the quantitative detection of the temperature change generated by electrochemical reactions on the working electrode. A similar concept was introduced by Borroni-Bird and King, who measured the heat of adsorption by the temperature change of thin samples in UHV.18 Lew et al. and Stuckless et al. used a PVDF foil as a pyroelectric temperature sensor.19,20 Further details on our experimental setup can be found in refs 21 and 22. Information on the theoretical background is summarized in refs 12, 13, 15, and 16 In this study we used a thin polished nickel sheet (50 μm thickness) Received: August 7, 2015 Revised: September 28, 2015

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DOI: 10.1021/acs.jpcb.5b07670 J. Phys. Chem. B XXXX, XXX, XXX−XXX

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Figure 1. Typical transients of potential (top, blue), current (middle, black), and temperature (bottom, red) during potential pulses with negative (left) and positive (right) amplitude. The potential pulses of −140 or +140 mV started at 10 ms and were switched off at 20 ms. Lithium deposition (left) led to cooling, lithium dissolution (right) to warming of the electrode.

with an active electrode area of ca. 0.2 cm2 as working electrode. The counter and reference electrodes consisted of either platinum or lithium, as shown below. To minimize the exchange of heat with the surrounding, i.e. the solution and the mount of the temperature sensor, we applied only short potential or current pulses with a duration of 10 ms to drive the electrochemical reaction. On this time scale the overall heat loss is very small while the electrode-sensor assembly is in thermal equilibrium. Under the above conditions the temperature change of the electrode during the pulse is proportional to the heat generated by the electrochemical reaction.21 Furthermore, the concentration gradients in the bulk electrolyte can be neglected under such conditions. The experiment consisted of calorimetric measurements using a series of different electrolytes. We used a 1 M solution of LiPF6 in a 1:1 mixture of EC and DMC, commercially available as LP30 (battery grade, BASF). The diluted solutions were prepared from the 1 M solution and a 1:1 mixture of EC (99%, Alfa Aesar) and DMC (99%, Alfa Aesar). The solvent mixture was stored over a molecular sieve, and the diluted solutions were freshly prepared before the experiment. For the calibration of the experiment we used an aqueous 0.1 M [Fe(CN)6]4−/0.1 M [Fe(CN)6]3− solution made from K4[Fe(CN)6] and K3[Fe(CN)6] (p.a., Merck) and ultrapure water (Arium, Sartorius). The oxidation and reduction of this redox couple is well-studied and the reversible molar heat, i.e. the molar Peltier heat, is available from literature.14 The calibration was conducted before the lithium deposition experiments since lithium irreversibly changes the electrode surface due to SEI formation, as generally known. For the calibration in the aqueous electrolyte we used a Pt (pseudo) reference and counter electrode. After the calibration, the cell was rinsed with acetone (p.a., Merck) and ultrapure water. Subsequently the setup was flange-mounted to a glovebox, evacuated with a turbomolecular pump to ca. 10−5 mbar, and flushed with argon. Inside the glovebox LiPF6 of the respective concentration in EC and DMC was used as electrolyte, and the reference and counter electrodes were changed to Li electrodes. In our previous experiments17 we used platinum electrodes as counter electrode and pseudo reference electrode for all parts of the experiment. Using lithium electrodes, however, provides a better-defined reference potential. In addition the reaction at the lithium counter electrode becomes reversible. We found however no differences between results obtained with lithium and platinum reference and counter electrodes; in particular we measured the same heat effects with both sets of electrodes. Prior to the calorimetric measurement of lithium deposition and dissolution a cyclic voltammogram was recorded to assess

the purity of the electrolyte in the cell and to form the SEI on the nickel current collector. Then the nickel working electrode was covered with about 104 monolayers of lithium by electrochemical lithium deposition to obtain a lithium bulkmetal working electrode.

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RESULTS We measured the exchanged heat during short pulses of lithium deposition or dissolution. Figure 1 shows the transients of potential, current, and temperature during typical pulses. The starting potential was the Li+/Li equilibrium potential. Potential pulses with amplitudes of −140 mV (Figure 1, left) or +140 mV (Figure 1, right) were applied between 10 and 20 ms. At the end of the potential pulse the flow of the external cell current was electronically interrupted, i.e. the working electrode was switched to open cell potential. Following this cutoff the potential jumps back to the Li+/Li equilibrium potential. The charge, i.e., the amount of converted lithium ions, can be calculated from the current transient. Negative current corresponds to lithium deposition and positive current to lithium dissolution. For the pulses shown in Figure 1 the charge was −6.5 μC respectively 6.6 μC, which corresponds to about 20% of a monolayer of lithium, assuming a surface density of 1015 atoms·cm−2. The temperature change started immediately with the beginning of the pulse. After the end of the pulse the increase respectively decrease continued only for a short time (∼1 ms) before the temperature signal started to relax to its equilibrium value. The temperature change during the potential pulse directly reflects the exchanged heat during the process. By calibrating our setup (see Experimental Section and refs 15 and 21) we were able to quantitatively determine the exchanged heat during the reaction. As previously reported,17 lithium deposition (left) led to a cooling and lithium dissolution (right) to a warming of the electrode. This symmetrical behavior shows the reversibility of the reaction. In order to obtain the reversibly exchanged heat, we performed calorimetric pulse measurements with various amplitudes, i.e., overpotentials, for both lithium deposition and dissolution. The exchanged heat was calculated from the difference between the temperature signal at the beginning and at the end of the pulse. The exchanged heat divided by the charge which flowed during the respective pulse will be referred to as molar heat of the reaction. In Figure 2 the molar heat for the deposition and dissolution of lithium from a 1 M solution is plotted versus the overpotential, i.e., the amplitude of the corresponding pulse. The error of the measurement will be discussed later. Analyzing the molar heat, two main contributions have to be considered. In one contribution the B

DOI: 10.1021/acs.jpcb.5b07670 J. Phys. Chem. B XXXX, XXX, XXX−XXX

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within one set of pulse measurements can be derived from the linear extrapolation and is usually in the order of 1%, which shows that a linear fit is a good approximation. Typical deviations of the extrapolated molar Peltier heat between different measurements within one experiment, i.e. without dismantling the cell, were of the order of 1 kJ·mol−1. For the comparison of data from different experiments the error is of the order of 7 kJ·mol−1, because in this case an additional error due to calibration has to be considered. The slope of the linear fit in Figure 2 represents the increasing contribution of irreversible heat with increasing overpotential, as follows from eq 1. Ideally it should match the Faraday constant. Real measurements show deviations due to electrolyte resistance, which creates a potential difference between the working electrode and the reference electrode. Hence the effective overpotential at the electrode/electrolyte interface becomes lower than the potential pulse amplitude and the slope of the line in Figure 2 becomes less than the Faraday constant. It should be noted that this effect does not change the resulting molar Peltier heat because the value for zero overpotential is independent of the slope. Considering the electrolyte resistance and the current density for the data in Figure 2 the deviation from F in the slope can be explained by an ohmic drop along about 3.6 mm of the current path in the electrolyte. This seems reasonable considering the geometry of our cell and the size and position of the reference electrode. In this study we determined the molar Peltier heat of lithium deposition from electrolytes with different concentrations to gain information on the behavior of lithium ions upon dilution. Figure 3 shows the exchanged molar heat for lithium deposition

Figure 2. Exchanged molar heat for the deposition (negative potential pulse) and dissolution (positive potential pulse) of lithium from a 1 M solution vs the potential pulse amplitude. The reversibly exchanged molar heat, i.e., the Peltier coefficient of 48.2 kJ·mol−1, was obtained by extrapolating the data to zero potential pulse amplitude. The error bars were calculated from the signal-to-noise ratio of the temperature transient.

heat flow is reversed upon reversal of the current flow. This reversible contribution is due to the electrochemical reaction.12 The other contribution is irreversible, i.e., it causes warming for both directions of the current flow. In our case it is mainly due to the overpotential during the electrochemical reaction, causing a polarization of the electrode with respect to its equilibrium potential.15 Joule heat, another common contribution to irreversible heat caused by current flow through the electrolyte, can be neglected for short pulsed experiments and electrolytes with rather high conductivity, as in our case.21 The overall heat change δq can therefore be written as δq = T ΔS dξ − |η|neF dξ

(1)

with temperature T, molar entropy change ΔS, reaction variable of the oxidation reaction ξ, overpotential η, stoichiometric number of electrons ne, and Faraday constant F. By convention δq is negative for heat transferred from the system to the surrounding. The reversibly exchanged heat, i.e., the heat in the limit of vanishing overpotential, is called Peltier heat and the corresponding molar value TΔS is called Peltier coefficient Π:

Π = T ΔS

(2)

The molar entropy change ΔS upon the electrochemical reaction includes the molar entropy of the main electrochemical reaction as well as the entropy of possible side reactions. In addition the transport of ions and electrons contributes to the molar entropy change.12 However, in our case this contribution is negligible due to the large Stokes radius of the lithium ion.17 In order to extract the reversible contribution to the molar heat, i.e., the Peltier coeffient, from our experiment, the data from all measurements with different potential pulse amplitudes, i.e., overpotentials, was extrapolated to an overpotential of zero by linear regression. For the data of the 1 M solution of LiPF6 in EC/DMC in Figure 2 a Peltier coefficient of 48.2 kJ·mol−1 mol was estimated. In Figure 2 the error bars of the single potential pulse measurements are also included. The error is calculated from the signal-to-noise ratio of the temperature transient. Understandably it is large for pulses with small overpotentials, i.e., small conversions. The statistical error of the molar Peltier heat

Figure 3. Exchanged molar heat for the deposition and dissolution of lithium vs the potential pulse amplitude for different Li+ concentrations. The sequence of the measurements was: 1 M (red, crosses), 0.1 M (black, triangles), 1 M (red, squares), and 0.01 M (blue, circles). For each concentration two sets of data were recorded, both marked with the same symbol. Each data set was separately extrapolated to zero pulse amplitude by a linear regression. All data were recorded within one experiment without dismantling the cell.

from 1, 0.1, and 0.01 M solutions of Li+ in EC/DMC for different potential pulse amplitudes, which were determined within one experiment without changing the cell configuration. We measured two sets of potential pulse measurements for each solution. The solutions were measured in the following order: 1, 0.1, 1, and 0.01 M. When changing between different concentrations the cell was thoroughly rinsed using the next solution. The transients of all pulses qualitatively matched the transients in Figure 1. As can be seen from the extrapolation of C

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concentration, with stronger deviations for higher concentrations. Therefore, considering activities instead of concentrations will lower the effect of dilution on the molar Peltier heat. Indeed, employing the Davies equation, an extension of the Debye−Hückel equation for higher concentrations,26 yielded a reduction of the Peltier coefficient by 5.4 kJ·mol−1 for dilution by a factor of 10. Hence, the entropy of dilution provides a sufficient explanation for our experimental results. Nonetheless it is important to discuss other possible contributions. As a second contribution to the changes in the solvation of Li+ ions we consider the effect of changes in the coordination of Li+ with electrolyte concentration. It has been found by various methods like Raman-, IR-, and NMR-spectroscopy as well as by DFT- and MD-calculations that the coordination number of Li+ increases upon dilution, e.g., in solutions of LiBF4 in EC/DMC, LiClO4 in EC, and LiBF4 in MEC (methoxymethyl ethylene carbonate).4−6,27 However, in those studies contributions from inner and outer solvation shells are not discriminated. Our previous results showed that the entropy of lithium deposition is mainly due to the entropy gain of the solvent molecules when they are released from the inner solvation shell upon deposition of the Li+ ion.17 The contribution to the entropy change by the liberation of solvent molecules from the inner solvation shell upon lithium deposition was estimated to be about 42 J·(mol· K)−1 for 1 M Li+ in EC:DMC, referred to the moles of liberated solvent molecules. In other words, an increase of the coordination number in the inner solvation shell by about one upon dilution of the electrolyte would enlarge the molar Peltier heat by about 12 kJ·mol−1. From our data, however, we determined a decrease of the molar Peltier heat upon dilution, which can be fully accounted for by the entropy of dilution. Note that activity variations with concentration cannot compensate for changes in the inner coordination shell, since the effect of activity also tends to lower the differences of the Peltier coefficients upon dilution, as discussed above. We therefore conclude that the inner shell coordination of Li+ ions is not changing considerably upon dilution and that the reported changes of the coordination number of Li+ with concentration rather reflect variations in the coordination beyond the first coordination shell. Finally as a third contribution to the changes in the solvation of Li+ ions with concentration we consider the effect of the formation of ion contact pairs. Ion association is favored in electrolytes with high concentration and/or small permittivity. Contact pairs between lithium and its counterion have been identified in various salts and solvents using Raman- and IRspectroscopy,4,27,28 Hartree−Fock, MD, and DFT calculations,6,8,29 and conductivity measurements.10 The extent of ion contact pair formation also has been discussed in literature. Qiao et al. studied the ion contact pair formation of LiBF4 in 4methoxymethyl-ethylene carbonate (MEC) at different temperatures.27 They report a Raman band corresponding to ion contact pair interactions even for their measurement with the lowest concentration (0.26 mol·L−1). The intensity of the band increased with concentration. Tasaki et al. calculated the solvation of lithium in a 1 M electrolyte with propylene carbonate (PC) and reported a partial solvation number for Li+ by PF6ions of 0.35, which can be interpreted as a measure for the contact ion pair concentration.8 Salomon and Plichta measured the conductivity of various Li-electrolytes and derived the activity of the ions and the thermodynamic association constant using the Fuoss-Hsia equation.10 For LiClO4 in PC

the single sets of potential pulses toward zero amplitude the molar Peltier heat decreases with concentration. The values for the molar Peltier heat were 48.1 kJ·mol−1 for 1 M, 41.7 kJ· mol−1 for 0.1 M, and 38.3 kJ·mol−1 for 0.01 M. The reproducibility was confirmed by repeating the measurement of each solution after 10 min. The average difference between the two measurements of the molar Peltier heat of the same solution is less than 0.5 kJ·mol−1. The difference between the molar Peltier heat for the 1 M solution and the diluted solutions is 6.4 kJ·mol−1 for a dilution by a factor of 10 respectively 9.8 kJ·mol−1 for dilution by a factor of 100. Any time dependent changes or aging effects can be ruled out because the first and the repeated measurement of the 1 M solution showed the same result. It is noteworthy that the slope of the linear regression becomes slightly smaller with decreasing electrolyte concentration. This is in line with the increasing electrolyte resistance and larger ohmic drop at the lower concentrations, as discussed above. As mentioned before, we repeated the whole experiment 5 times. Some of the experiments showed a slightly smaller difference of the molar Peltier heat values. However, all of the differences for the molar Peltier heat values were in the order of 5 ± 1.5 kJ·mol−1 for a dilution by a factor of 10 respectively 10 ± 2 kJ·mol−1 for dilution by a factor of 100. In some experiments the molar Peltier heat for the very first lithium deposition from a 1 M solution was smaller than that of all following measurements of lithium deposition from a 1 M solution in the same experiment. This is logical because the first lithium deposition is coupled with the formation (or further growth and/or rearrangement) of the SEI due to the change of the micromorphology of the electrode surface. We attribute this heat effect therefore to side reactions due to the building of a covering SEI layer, which was not completed at that time. For more information on the surface layer on lithium metal, normally called solid electrolyte interphase (SEI), see, e.g., refs 23−25.

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DISCUSSION The reaction pathways for electrochemical lithium deposition from the different concentrations share the same thermodynamic final state but differ in the starting one. Hence, differences in the molar Peltier heat of the lithium deposition, i.e., in the entropy of the process, can be traced back to changes in the entropy of the lithium ion in the solution for different concentrations. Upon dilution of the electrolyte three effects have to be considered: The entropy of dilution for Li+, changes in coordination of Li+, and the formation of ion contact pairs between Li+ and its counterion. The effect of dilution on the entropy of Li deposition can be readily calculated. The entropy of the Li+ ion SLi+ changes with its activity aLi+ according to SLi+ = SLi ° + − R ln(aLi+) − RT((∂ln(aLi+))(∂T)).13 Here S°Li+ is the standard molar entropy of Li+. Neglecting the temperature dependence of the activity, which is usually small, and using concentrations instead of activities, the entropy change of the Li+ due to dilution becomes ΔSdil ≈ − R ln(c2/c1). Correspondingly the Peltier coefficient of the Li deposition process will vary by ΔΠ ≈ − RT ln(c2/c1). A dilution by a factor of 10 will lower the molar Peltier heat by 5.7 kJ·mol−1, a dilution by a factor of 100 will yield a reduction by 11.4 kJ/mol. These values already come close to the experimental observations of 5 ± 1.5 and 10 ± 2 kJ·mol−1. In a real electrolyte the activity of Li+ will deviate from the concentration. We expect that in our experimental concentration interval the activity is generally lower than the D

DOI: 10.1021/acs.jpcb.5b07670 J. Phys. Chem. B XXXX, XXX, XXX−XXX

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The Journal of Physical Chemistry B they found an association constant Ka of 1.3 L·mol−1 which corresponds to 31% contact ion pairs at 1 M. It is difficult to predict the solvation of ion contact pairs. Because they possess only a dipole moment but no net charge, the extent of solvation might initially be expected to be reduced compared to the single ions. Therefore, deposition of lithium from contact ion pairs should liberate less solvent molecules than deposition from isolated solvated Li+ ions. The formation of a free anion upon lithium deposition can be neglected in this context because it is generally very weakly solvated.30 Since the fraction of contact ion pairs should decrease upon dilution, their effect on the Peltier coefficient should oppose the influence of dilution entropy, discussed above. Therefore, at first glance, the differences between the Peltier coefficients for the different concentrations should be lowered by the formation of contact ion pairs. From our data, however, we see no such effect within our error margins. This suggests that the solvation entropy of ion contact pairs is similar to the solvation entropy of free lithium ions. Indeed, on the molecular level only one side of the Li+ ion is shielded by its counterion. On the unshielded side the Li+ ion will still strongly polarize the neighboring solvent molecules. This reasoning is supported by DFT-calculations of Tasaki et al., who report an almost unchanged energy for the desolvation of lithium in ion contact pairs compared to solvated lithium at 1 M concentration.8

compensate the variations due to the dilution of the electrolyte. Hence, the trend for the variation of the Peltier heat with Li+ concentration, as observed in Figure 3, may eventually reverse its sign for superconcentrated Li electrolytes.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The authors gratefully acknowledge technical support by D. Waltz and his colleagues from the mechanical workshop. This work benefitted from experimental improvements funded by the Deutsche Forschungsgemeinschaft (Grant No. SCHU 958/ 7-1).

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REFERENCES

(1) Armand, M.; Tarascon, J. M. Building Better Batteries. Nature 2008, 451, 652−657. (2) Xu, K. Electrolytes and Interphases in Li-Ion Batteries and Beyond. Chem. Rev. 2014, 114, 11503−11618. (3) Aurbach, D.; Talyosef, Y.; Markovsky, B.; Markevich, E.; Zinigrad, E.; Asraf, L.; Gnanaraj, J. S.; Kim, H.-J. Design of Electrolyte Solutions for Li and Li-Ion Batteries: A Review. Electrochim. Acta 2004, 50, 247−254. (4) Wang, Z.; Huang, B.; Xue, R.; Chen, L.; Huang, X. Ion Association and Solvation Studies of LiClO4/Ethylene Carbonate Electrolyte by Raman and Infrared Spectroscopy. J. Electrochem. Soc. 1998, 145, 3346−3350. (5) Castriota, M.; Cazzanelli, E.; Nicotera, I.; Coppola, L.; Oliviero, C.; Ranieri, G. A. Temperature Dependence of Lithium Ion Solvation in Ethylene Carbonate−LiClO4 Solutions. J. Chem. Phys. 2003, 118, 5537−5541. (6) Postupna, O. O.; Kolesnik, Y. V.; Kalugin, O. N.; Prezhdo, O. V. Microscopic Structure and Dynamics of LiBF4 Solutions in Cyclic and Linear Carbonates. J. Phys. Chem. B 2011, 115, 14563−14571. (7) Bogle, X.; Vazquez, R.; Greenbaum, S.; Cresce, A. v. W.; Xu, K. Understanding Li+ -Solvent Interaction in Nonaqueous Carbonate Electrolytes with 17O Nmr. J. Phys. Chem. Lett. 2013, 4, 1664−1668. (8) Tasaki, K.; Kanda, K.; Nakamura, S.; Ue, M. Decomposition of LiPF6 and Stability of PF5 in Li-Ion Battery Electrolytes: Density Functional Theory and Molecular Dynamics Studies. J. Electrochem. Soc. 2003, 150, A1628−A1636. (9) Seo, D. M.; Reininger, S.; Kutcher, M.; Redmond, K.; Euler, W. B.; Lucht, B. L. Role of Mixed Solvation and Ion Pairing in the Solution Structure of Lithium Ion Battery Electrolytes. J. Phys. Chem. C 2015, 119, 14038−14046. (10) Salomon, M.; Plichta, E. Conductivities and Ion Association of 1:1 Electrolytes in Mixed Aprotic Solvents. Electrochim. Acta 1983, 28, 1681−1686. (11) Cresce, A. V.; Xu, K. Preferential Solvation of Li+ Directs Formation of Interphase on Graphitic Anode. Electrochem. Solid-State Lett. 2011, 14, A154−A156. (12) Agar, J. N., Thermogalvanic Cells. In Advances in Electrochemistry and Electrochemical Engineering; Delehay, P., Ed.; Interscience Publishers: London, 1963; Vol. 3, pp 31−118. (13) Ozeki, T.; Ogawa, N.; Aikawa, K.; Watanabe, I.; Ikeda, S. Thermal Analysis of Electrochemical Reactions: Influence of Electrolytes on Peltier Heat for Cu(0)/Cu(II) and Ag(0)/Ag(I) Redox Systems. J. Electroanal. Chem. Interfacial Electrochem. 1983, 145, 53− 65. (14) Boudeville, P. Thermometric Determination of Electrochemical Peltier Heat (Thermal Effect Associated with Electron Transfer) of Some Redox Couples. Inorg. Chim. Acta 1994, 226, 69−78.

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CONCLUSION We determined the molar Peltier heat for electrochemical Li deposition from solutions of LiPF6 in DMC/EC with varying Li+ concentrations. Since entropy of transport can be neglected in such electrolyte solutions, the molar Peltier heat for the Li deposition directly reflects the entropy change upon the electrochemical Li deposition reaction. The molar Peltier heat decreased by 5 ± 1.5 kJ·mol −1 by reducing the Li + concentration from 1 to 0.1 M and by 10 ± 2 kJ·mol−1 comparing 1 and 0.01 M solutions. Assuming an ideal solution, we approximated a decrease of the molar Peltier heat by 5.7 kJ· mol−1 upon dilution of the electrolyte concentration by a factor of 10. Hence the measured variations of the molar Peltier heat can be well explained by the change of the entropy of the solvated Li+ ion upon dilution. Corrections which account for the deviation of the solution from ideality, the change of the inner shell coordination number of Li+, or the formation of ion pairs should all reduce the differences of the molar Peltier heat among the different concentrations. Hence, we neither expect strong changes of the activity coefficient nor of the inner shell coordination number of the Li+ ions among the investigated Li+ concentrations. Furthermore, although the formation of ion contact pairs has been reported for electrolytes similar to the one used in this study, we find no evidence of changes in the solvation entropy for lithium ions. This suggests that ion contact pairs are similarly solvated as lithium ions at Li+ concentrations of 1 M and below. It is worth to note that recently superconcentrated Li+ based electrolytes with Li+ concentrations exceeding 2 M were investigated for applications, e.g., in fast-charging Li ion batteries.31 At such high Li+ concentrations the solvation shells of the Li+ ions are interpenetrating and often also the electrochemical properties of the electrolyte deviate considerably from those at lower concentrations.32 We expect that in our system, for Li+ concentrations far beyond 1 M, the effects on the Peltier heat by the changes of the inner coordination shell and the formation of ion pairs may eventually overE

DOI: 10.1021/acs.jpcb.5b07670 J. Phys. Chem. B XXXX, XXX, XXX−XXX

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DOI: 10.1021/acs.jpcb.5b07670 J. Phys. Chem. B XXXX, XXX, XXX−XXX