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May 21, 2015 - Systematic variations of the cosalt reveal two broad classes of electrolytes, namely, the group 13 electrolytes, which require electrol...
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Exploring Salt and Solvent Effects in Chloride-Based Electrolytes for Magnesium Electrodeposition and Dissolution Christopher J. Barile,†,‡ Ralph G. Nuzzo,†,‡ and Andrew A. Gewirth*,†,‡ †

Joint Center for Energy Storage Research, Argonne National Laboratory, Argonne, Illinois 60439, United States Department of Chemistry, University of Illinois at Urbana−Champaign, Urbana, Illinois 61801, United States

‡

S Supporting Information *

ABSTRACT: We describe in this work Mg electrodeposition and dissolution from a wide range of inorganic ethereal electrolytes consisting of MgCl2 and a second chloride salt. Systematic variations of the cosalt reveal two broad classes of electrolytes, namely, the group 13 electrolytes, which require electrolytic cycling to improve their performance, and electrolytes based on heavy p-block chlorides, which exhibit Mg intermetallic formation. Results from electrospray ionization mass spectrometry demonstrate that Mg deposition and stripping only occur in electrolytes containing Mg multimers. We also explore the role of solvent in determining the electrochemical performance of chloride-based electrolytes. Our analysis establishes thermodynamic parameters that dictate the ability of a solvent to support Mg electrochemistry in the MgCl2−AlCl3 system. In their totality, these results illustrate important electrolyte design guidelines for future Mg-ion batteries.

1. INTRODUCTION Mg-ion batteries have the potential to outperform currently used Li-ion batteries due to their high theoretical volumetric and gravimetric energy densities.1 One of the largest barriers preventing the successful implementation of Mg-ion batteries is the lack of a suitable electrolyte that supports reversible Mg electrodeposition and dissolution at the Mg metal anode.2 The carbonate-based solvents traditionally utilized in Li-ion batteries do not support reversible Mg deposition due to the difficulty of desolvating Mg2+ and the growth of passivation layers on the anode surface.3,4 Gregory et al. first reported that Mg can reversibly deposit from Grignard reagents in ethers and suggested that the electron-donating nature of the organic moiety facilitates Mg reduction.5 Because of their high reactivity, however, Grignard-based electrolytes suffer from poor anodic stabilities.6,7 Later work by Aurbach and coworkers established that the reactivity of Grignard reagents could be modulated by reacting them with Al species to form Mg organohaloaluminates.8 By manipulating the organic substituents of these electrolytes, systems with good coulombic efficiencies and electrochemical windows up to ∼2.5 V have been reported.9 A handful of studies also suggest electrolytes based on Mg alkoxides and Mg bis(trifluoromethanesulfonyl)imide may be promising.10−13 In recent work, Aurbach et al. and Liu et al. demonstrated reversible Mg deposition and stripping from all inorganic electrolytes made by combining AlCl3 and MgCl2 in ethereal solvents.14,15 These electrolytes, which contain the so-called magnesium aluminum chloride complex (MACC), exhibit electrochemical windows greater than 3 V. Their high anodic stabilities make them promising candidates for coupling with as of yet undiscovered high-voltage Mg-ion cathodes. However, © XXXX American Chemical Society

studies in organohaloaluminates have determined that the chloride anion is incompatible with common battery current collectors such as those made from stainless steel and aluminum, which may limit the practicality of chloride-based electrolytes such as those containing the MACC (MgCl2− AlCl3).16,17 We recently investigated MgCl2−AlCl3 electrolytes and determined that the performance of these electrolytes varies dramatically with cycle number.18 In fact, reversible deposition can only be attained after extensive cycling in a process known as electrolytic conditioning. These studies further show that during the conditioning process, the speciation of Mg cations in solution shifts to Mg multimers, which in turn facilitates facile and reversible Mg deposition. We found that the performances of the MgCl2−AlCl3 electrolytes in tetrahydrofuran (THF) and 1,2-dimethoxyethane (DME) increase dramatically as they cycle. Once fully conditioned, the MgCl2−AlCl3 electrolyte in THF possesses a coulombic efficiency of 100% and an overpotential for Mg deposition of ∼200 mV. By comparison, the conditioned MgCl2 −AlCl 3 electrolyte in DME exhibits an inferior coulombic efficiency of 92% but an improved Mg deposition overpotential of ∼150 mV. We hypothesized that the differences in performance of the two electrolytes stem from differences in the structures of Mg−Al complexes present. Despite these studies, the exact nature of the species responsible for Mg deposition in the MgCl2−AlCl3 electrolyte remains unknown. The lack of fundamental understanding of Received: April 11, 2015 Revised: May 21, 2015

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DOI: 10.1021/acs.jpcc.5b03508 J. Phys. Chem. C XXXX, XXX, XXX−XXX

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Figure 1. Cyclic voltammograms of cycle 1 (black), cycle 10 (red), cycle 100 (blue), and cycle 200 (green) of MgCl2−BCl3 in DME at 5 mV/s on a Pt electrode (A). Mg deposition potential and coulombic efficiency as a function of cycle number (B).

To synthesize completely inorganic-based electrolytes, the desired solvent was added dropwise to a metal chloride (30 mM) at ∼0 °C over the course of 10 min and stirred at room temperature for ∼30 min. MgCl2 (60 mM) was then added to an equal volume of the same solvent at room temperature and stirred for ∼30 min. The two solutions were then combined and stirred for at least 10 h. Before use in electrochemical experiments, any insoluble solids were filtered from the electrolyte solution. For electrolytes containing BCl3, the requisite amount of BCl3 (30 mM) was added dropwise at ∼0 °C to DME or THF containing MgCl2 (60 mM) and tetrabutylammonium (TBA) perchlorate (100 mM) from a 1.0 M solution of the gas dissolved in hexanes, resulting in a final solvent composition of 97% DME or THF and 3% hexanes. 2.3. Mass Spectrometry Experiments. For mass spectrometry (MS) experiments, the electrolytes were diluted with additional DME, removed from the glovebox in a sealed vial, and directly injected into a Quattro Ultima (Waters) electrospray ionization (ESI) mass spectrometer using a purged, Ar-filled syringe. MS was performed in negative ionization mode. 2.4. Scanning Electron Microscopy. Imaging studies of the electrodes by scanning electron microscopy (SEM) were performed using a Hitachi S-4700 Cold FE-SEM (Hitachi High Technologies) with an acceleration voltage of 20 kV. Quantitative energy-dispersive X-ray spectroscopy (EDS) was performed with an Oxford Instruments ISIS EDS X-ray Microanalysis System using Cu foil as a standard. The surfaces were rinsed thoroughly with DME inside the glovebox before transferring them for SEM-EDS analysis.

how chloride-based electrolytes support Mg deposition hampers further progress in electrolyte development. In this work, we systematically alter the components of the MgCl2− AlCl3 electrolyte to gain insight into the parameters necessary for reversible Mg deposition. In particular, we explore the effects of changing the solvent of the electrolyte and the identity of the cosalt added to MgCl2, finding within these data important rate−structure correlations that underpin efficient electrochemical conversion between Mg and Mg2+.

2. EXPERIMENTAL DETAILS 2.1. General Procedures. Chemicals were purchased from Sigma-Aldrich unless otherwise stated. All materials were handled in an Ar-filled glovebox, which contained BCl3 > ZnCl2 > AlCl3 > SbCl3.33 The finding that BCl3, AlCl3, and SbCl3 support quasi-reversible Mg electrochemistry, whereas FeCl3 and ZnCl2 do not, suggests that the Lewis acidity of MClx is unimportant in determining its efficacy as a cosalt. Since the dominant Lewis base in solution is the ethereal solvent, we next sought to determine the effect of the solvent on the performance of the MgCl2−AlCl3 system. 3.7. Solvent Effects in the MgCl2−AlCl3 Electrolyte. As discussed in the introduction, we previously compared the Mg electrochemistry of the MgCl2−AlCl3 electrolyte in THF and DME.18 We now analyze the electrochemical behavior of this system in a wide range of ethereal solvents. Figure 11 shows

dimers or trimers may have to do with some of the properties of Ga that deviate from ordinary periodic trends. For example, normal periodic trends dictate that moving down a family of the elements results in decreasing ionization energies. For the group 13 elements, B, Al, and In follow this canonical trend, but Ga and Tl are exceptions due to the effect of d-block and fblock contractions, respectively.20 Preliminary experiments indicate that, like Ga, Tl electrolytes do not support Mg deposition or stripping, but difficulties in preparing anhydrous TlCl3 have stymied us from presenting stronger conclusions about this system. Table 2 lists the anodic decomposition voltages of the cycled MgCl2−MClx electrolytes in DME that support Mg deposition Table 2. Approximate Anodic Decomposition Voltages for Various Cycled MgCl2−MClx Electrolytes in DME and the Electronegativity of M (Pauling Scale) electrolyte (MgCl2−MClx)

anodic decomposition voltage (V vs Mg/Mg2+)

electronegativity of M (χ)

MgCl2−AlCl3 MgCl2−InCl3 MgCl2−BCl3 MgCl2−SnCl2 MgCl2−BiCl3 MgCl2−SbCl3

3.2 2.0 1.8 1.5 1.5 1.2

1.6 1.8 2.0 2.0 2.0 2.1

and stripping alongside the electronegativity of the element used in the chloride cosalt.20 In general, as the chloride cosalt becomes more electronegative, the anodic stability of the MgCl2−MClx electrolyte decreases. This finding suggests that MClx catalyzes the oxidative decomposition of DME. As MClx becomes more electronegative, it can more easily withdraw electrons from DME, resulting in a lower oxidative stability of the electrolyte. However, since the anodic decomposition voltage of each electrolyte changes with cycling, the determined correlation must only be considered approximate. Table 3 lists 30 anhydrous chlorides from across the periodic table that we screened for their ability to support reversible Mg

Figure 11. Cyclic voltammograms of the conditioned MgCl2−AlCl3 electrolyte in 2-MeTHF (black) and diglyme (red) at 5 mV/s on a Pt electrode.

cyclic voltammograms of the conditioned MgCl 2−AlCl3 systems in 2-methyltetrahydrofuran (2-MeTHF) and diglyme. Both of these solvents support Mg deposition and stripping, but the onset potentials for Mg deposition and the coulombic efficiencies are inferior to those obtained in THF. Figures 12A and 13A display cyclic voltammograms of the MgCl2−AlCl3 electrolytes in 2-MeTHF and diglyme, respectively, as the electrolytes cycle. In both cases over the first 50 cycles of the conditioning process, the performances of the electrolytes markedly improve. This improvement is similar to what we observed previously for the MgCl2−AlCl3 systems in THF and DME.18 For the case of the electrolyte in 2-MeTHF, the Mg deposition potential and coulombic efficiency vary erratically during the remainder of the conditioning process (Figure 12B). In contrast, the coulombic efficiency of the diglyme electrolyte reaches a steady value of ∼94% before declining after ∼200 cycles (Figure 13B). During this decline in coulombic efficiency, however, the Mg deposition potential remains constant at ca. −0.38 V. Supporting Information, Figures S8 and S9 show cyclic voltammograms of the MgCl2−AlCl3 electrolyte in 2,5dimethyltetrahydrofuran (2,5-Me2THF) and 2,2,5,5-tetramethyltetrahydrofuran (2,2,5,5-Me4THF), respectively. Even after 75 cycles, these electrolytes display no significant signs of Mg

Table 3. List of the 30 Chlorides Screened in MgCl2−MClx in DME Electrolytes chloridesa

group alkali metals alkaline earth metals transition metals group 13 group 14 pnictogens other

LiCl, KCl CaCl2, BaCl2 YCl3, VCl3, NbCl5, TaCl5, WCl6, MnCl2, ReCl3, FeCl3, CoCl2, NiCl2, PdCl2, PtCl4, CuCl2, AgCl, ZnCl2, HgCl2 BCl3, AlCl3, GaCl3, InCl3 SnCl2, PbCl2 SbCl3, BiCl3 SeCl4, TBACl

a

Underlined and bolded salts generated electrolytes that display Mg electrodeposition and stripping.

deposition in the MgCl2-MClx electrolyte in DME. Of the chlorides studied, only BCl3, AlCl3, InCl3, SnCl2, SbCl3, and BiCl3 exhibit Mg deposition and stripping to any appreciable extent. The ability of the MgCl2−AlCl3 system to support reversible Mg deposition and stripping has been attributed previously in part to the strong Lewis acidity of AlCl3.14 In this framework, G

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Figure 12. Cyclic voltammograms of cycle 1 (black), cycle 10 (red), cycle 30 (blue), cycle 100 (green), and cycle 150 (purple) of the MgCl2−AlCl3 electrolyte in 2-MeTHF at 5 mV/s on a Pt electrode (A). Mg deposition potential and coulombic efficiency as a function of cycle number (B).

Figure 13. Cyclic voltammograms of cycle 1 (black), cycle 50 (red), cycle 100 (blue), cycle 150 (green), and cycle 200 (purple) of the MgCl2−AlCl3 electrolyte in diglyme at 5 mV/s on a Pt electrode (A). Mg deposition potential and coulombic efficiency as a function of cycle number (B).

deposition and stripping. The finding that 2,5-Me2THF and 2,2,5,5-Me4THF do not support Mg electrochemistry, whereas 2-MeTHF and THF do, suggests that too much steric crowding around the oxygen atom of ethers prevents Mg deposition and stripping from taking place in the MgCl2−AlCl3 electrolyte. Supporting Information, Figures S10 and S11 display cyclic voltammograms of the MgCl2−AlCl3 electrolyte in triglyme and tetraglyme, respectively. These solvents also do not support appreciable Mg deposition and stripping even after 75 cycles. The observation that DME (monoglyme) and diglyme support Mg electrochemistry, but triglyme and tetraglyme do not, suggests that if a solvent is too chelating, it will prevent Mg deposition and stripping from occurring in the MgCl2−AlCl3 electrolyte. Several other solvents were evaluated for their ability to support Mg deposition and stripping with the MgCl2−AlCl3 electrolyte. These include the ethers 3-methyltetrahydrofuran, oxetane, tetrahydropyran, 1,3-dioxane, 1,2-diethoxyethane, and dibutyl ether and the related solvents γ-butyrolactone, furan, phthalan, and tetrahydrothiophene. In all of these cases, the MgCl 2 −AlCl 3 system did not display appreciable Mg deposition or stripping even after 75 cycles at 5 mV/s. These solvents possess a wide range of steric and chelating properties. Their inability to support Mg deposition and stripping demonstrates that the speciation of the MgCl2−AlCl3 electrolyte is very sensitive to solvent. 3.8. Solvent Effects in the APC Electrolyte. To better understand the effect of solvent in Mg electrolytes, we also tested the PhMgCl−AlCl3 electrolyte, or “all phenyl complex” (APC) in different ethers. Figure 14 shows representative cyclic voltammograms of the APC electrolyte in 2-MeTHF (black) and 2,5-Me2THF (red). In both of these solvents, the APC electrolyte possesses ∼98% coulombic efficiency. However, the

Figure 14. Cyclic voltammograms of cycle 20 of APC in 2-MeTHF (black) and 2,5-Me2THF (red) at 5 mV/s on a Pt electrode.

current densities achieved in 2,5-Me2THF are significantly less than in 2-MeTHF probably because the unfavorable steric bulkiness of 2,5-Me2THF hinders the formation of Mg species that facilitate Mg deposition. Nonetheless, the finding that the APC electrolyte supports Mg electrochemistry in 2,5-Me2THF is in stark contrast to the lack of significant Mg deposition and stripping observed in the MgCl2−AlCl3 system. These results suggest that some property of the APC electrolyte allows for fairly reversible Mg deposition and dissolution despite the adverse steric bulkiness of 2,5-Me2THF. Like the MgCl2−AlCl3 electrolyte, since the APC electrolyte does not exhibit Mg deposition and stripping in 2,2,5,5-Me4THF, we conclude that the extreme steric bulkiness of 2,2,5,5-Me4THF effectively thwarts any Mg electrochemistry (Supporting Information, Figure S12). Additionally, the current densities for Mg deposition and stripping obtained from the APC in 2MeTHF and 2,5-Me2THF are less than we and others observe from the APC in THF.9 This decrease in current density is due both to the lower solubility of the APC in 2-MeTHF and 2,5H

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(DME), diglyme, triglyme, and tetraglyme in order of increasing chelation. We hypothesize that entropic considerations play a large role in dictating the ability of a solvent to facilitate Mg deposition and stripping and consider the expected trends in entropic changes for the glyme series. Table 5 lists the plausible

Me2THF and inhibited Mg deposition kinetics that arise from the steric bulkiness of these THF derivatives. Figure 15 shows representative cyclic voltammograms of the APC electrolyte in diglyme, triglyme, and tetraglyme. In all

Table 5. Coordination Structures Considered for Six-, Five-, and Four-Coordinate MgCl2 and Mg2Cl3+ Species for the Glyme Solvent Series Mg coordination number

solvent

MgCl2 structure 2

six

DME

MgCl2(κ -DME)2

five

DME

Figure 15. Cyclic voltammograms of cycle 20 of APC in diglyme (black), triglyme (red), tetraglyme (blue) at 5 mV/s on a Pt electrode.

four

DME

MgCl2(κ2-DME) (κ1-DME) MgCl2(κ2-DME)

six

diglyme

MgCl2(κ2-diglyme)2

three solvents, the APC electrolyte exhibits coulombic efficiencies of 98% to 99%. Since the conditioned MgCl2− AlCl3 system does not display significant Mg deposition and stripping in triglyme and tetraglyme, but the APC system does, these results again demonstrate that the APC electrolyte is less sensitive to changes in solvent than the MgCl 2 −AlCl 3 electrolyte. 3.9. Analysis of Solvent Effects. Table 4 compares the deposition potentials and coulombic efficiencies of the MgCl2− AlCl3 and APC electrolytes in eight different ethereal solvents. The solvents studied can be divided into two categories based upon their structural similaritiesa series consisting of THF, 2MeTHF, 2,5-Me2THF, and 2,2,5,5-Me4THF in order of increasing steric bulkiness and a series consisting of monoglyme

five

diglyme

MgCl2(κ3-diglyme)

four

diglyme

MgCl2(κ2-diglyme)

six

triglyme

MgCl2(κ4-triglyme)

five

triglyme

MgCl2(κ3-triglyme)

four

triglyme

MgCl2(κ2-triglyme)

six

tetraglyme

five

tetraglyme

four

tetraglyme

MgCl2(κ4tetraglyme) MgCl2(κ3tetraglyme) MgCl2(κ2tetraglyme)

solvent THF

2-MeTHF

2,5-Me2THF

2,2,5,5Me4THF monoglyme (DME) diglyme

triglyme

tetraglyme

electrolyte MgCl2− AlCl3 APC MgCl2− AlCl3 APC MgCl2− AlCl3 APC MgCl2− AlCl3 APC MgCl2− AlCl3 MgCl2− AlCl3 APC MgCl2− AlCl3 APC MgCl2− AlCl3 APC

highest coulombic efficiency

−0.2 V

100%

−0.2 V −0.27 V

100% 97%

−0.16 V

98% 0%

−0.2 V

98% 0%

−0.15 V

0% 92%

−0.38 V

94%

−0.24 V

99% 0%

−0.24 V

98% 0%

−0.32 V

99%

Mg2(μ-Cl)3(κ2DME)3+ Mg2(μ-Cl)3(κ2DME)2+ Mg2(μ-Cl)3(κ2DME)+ Mg2(μ-Cl)3(κ3diglyme)2+ Mg2(μ-Cl)3(κ2diglyme)2+ Mg2(μ-Cl)3(κ2diglyme)+ Mg2(μ-Cl)3(κ3triglyme)2+ Mg2(μ-Cl)3(κ4triglyme)+ Mg2(μ-Cl)3(κ2triglyme)+ Mg2(μ-Cl)3(κ3tetraglyme)2 Mg2(μ-Cl)3(κ4tetraglyme) Mg2(μ-Cl)3(κ2tetraglyme)+

structures of MgCl2 and Mg2Cl3+ in each of the glyme solvents with coordination numbers of four to six. The exact coordination number of Mg in the MgCl2−AlCl3 electrolyte is a matter of debate in the literature. In the solid state, Mg in the form of MgCl2 and Mg2Cl3+ complexes with ethers predominantly in a six-coordinate octahedral geometry with bridging chloride ligands in the case of the dimer.15,34 In situ Xray studies also suggest that a six-coordinate Mg dimer with bridging chloride ligands exists near the interface of the working electrode during Mg deposition and stripping from a Mg organohaloaluminate.35 Recent computational studies, however, argue that at least in bulk solution, MgCl2 and Mg2Cl3+ prefer four- or five-coordinate geometries.36,37 We consider how the coordination number in the glyme series affects the expected entropic changes in the nominal reactions that occur in the synthesis of the MgCl2−AlCl3 and APC electrolytes (Scheme 1). Supporting Information, Table S1 lists the change in the number of solvent molecules bound to Mg that is expected to occur when Mg2Cl3+ forms from two MgCl2 species. Those reactions that result in a larger net loss of solvent molecules are expected to be more entropically favored

Table 4. Comparison of the Most Positive Mg Deposition Potentials and Highest Coulombic Efficiencies Obtained during the Cycling of MgCl2−AlCl3 and APC Electrolytes in Various Solvents most positive deposition potential (vs Mg/Mg2+)

Mg2Cl3+ structure

Scheme 1. Nominal Reactions Occurring in the Syntheses of MgCl2−AlCl3 (1) and APC (2) Electrolytes and Their Heat of Reactions Calculated Using Bond-Dissociation Energies

I

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effect of the solvent and chloride cosalt added to MgCl2 was investigated. We find that MgCl2−MClx electrolytes can be classified into two broad categoriesthe group 13 electrolytes, which require electrolytic conditioning, and the heavy p-block electrolytes, which exhibit electrochemistry controlled by Mg intermetallic formation. ESI-MS results indicate that electrolytes only exhibit appreciable Mg deposition and stripping if they contain multimeric Mg species. Furthermore, we suggest that Mg deposition and stripping occur in a wider range of solvents for the APC electrolyte than for the MgCl2−AlCl3 electrolyte because of the greater enthalpic favorability of Mg multimers in the APC electrolyte. On the basis of entropic arguments, we also propose that six-coordinate Mg dimers facilitate Mg electrodeposition and dissolution in the MgCl2− AlCl3 electrolyte. Taken together, these findings serve as a guide for future electrolyte development efforts for Mg batteries.

since free solvent molecules have more disorder than bound ones. We analyze the nine possible permutations of four-, five-, and six-coordinate species in the dimerization reaction for each solvent (Table S1). The analysis reveals that the dimerization of two six-coordinate MgCl2 species to give a six-coordinate Mg2Cl3+ species is entropically favorable for DME and diglyme but not for triglyme and tetraglyme. This finding matches our experimental results in which Mg deposition and stripping occur in DME and diglyme but not in triglyme and tetraglyme. We conclude that the formation of six-coordinate Mg dimers required to facilitate reversible Mg deposition is not entropically favorable in triglyme and tetraglyme, thus explaining their inability to support Mg electrochemistry in the MgCl2−AlCl3 electrolyte. Strikingly, we find that the all six-coordinate dimerization reaction is the only set of coordination numbers that matches our results. These findings suggest that as was proposed for a Mg organohaloaluminate system,35 a sixcoordinate Mg2Cl3+ species in the MgCl2−AlCl3 electrolyte participates in Mg deposition at the electrode interface. In a broad sense, the results from the solvent effect studies demonstrate that the APC electrolyte can support Mg deposition and stripping in a much wider range of solvents than can the MgCl2−AlCl3 electrolyte. We propose that these differences are explained by specific thermodynamic parameters of Mg and Al speciation present in each electrolyte system. Scheme 1 lists the central reactions that occur during the initial synthesis of the MgCl2−AlCl3 and APC electrolytes.9,38 Both reactions generate dimeric Mg cations, which, depending on the Mg−Al stoichiometry in the electrolyte, may continue to react with additional Mg species to form higher-order Mg multimers. It is these multimeric Mg species that are thought to be active with respect to Mg electrodeposition and stripping.18 Through the use of tabulated gas-phase homolytic bond dissociation enthalpies,20,39 we estimated the enthalpy changes for both of these reactions. The results show that the reaction occurring in the synthesis of the APC electrolyte (reaction 2) is much more enthalpically favorable than the reaction that occurs upon making the MgCl2−AlCl3 electrolyte (reaction 1). These differences highlight that the exchange of Mg−C bonds for Al− C bonds in reaction 2 is much more thermodynamically favored than the exchange of Mg−Cl bonds for Al−Cl bonds in reaction 1. The changes in entropy for both of these reactions are negative since in both cases, two more-ordered species form from three less-ordered species. The magnitude of this entropic change will vary widely depending on the nature of the solvent. If the entropic change in a given solvent is too negative, it may outweigh the enthalpic driving force and render the overall reaction thermodynamically unfavorable. Since the reaction describing the formation of Mg multimers in the APC electrolyte is more enthalpically favored than the analogous reaction for the MgCl2−AlCl3 electrolyte, the APC electrolyte can tolerate solvents that give more negative entropic changes for the formation of Mg multimers before the reaction becomes thermodynamically unfavorable. In other words, these thermodynamic considerations explain why the APC electrolyte can support Mg deposition and stripping in a much wider range of solvents than can the MgCl2−AlCl3 system.

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ASSOCIATED CONTENT

S Supporting Information *

Cyclic voltammograms, SEM-EDS illustrations, ESI-MS spectra, tabulated data indicating change in number of solvent molecules bound upon formation of Mg2Cl3+. The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jpcc.5b03508.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Phone: +1 (217) 333-8329. Fax: +1 (217) 244-3186. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was supported as part of the Joint Center for Energy Storage Research, an Energy Innovation Hub funded by the U.S. Department of Energy, Office of Science, Basic Energy Sciences. C.J.B. acknowledges a National Science Foundation Graduate Research Fellowship (No. NSF DGE-1144245), a Springborn Fellowship, and insightful discussions with Dr. E. Chenard. This work was carried out in part in the Frederick Seitz Materials Research Laboratory Central Facilities, Univ. of Illinois at Urbana−Champaign.

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REFERENCES

(1) Yoo, H. D.; Shterenberg, I.; Gofer, Y.; Gershinsky, G.; Pour, N.; Aurbach, D. Mg Rechargeable Batteries: An On-going Challenge. Energy Environ. Sci. 2013, 6, 2265−2279. (2) Saha, P.; Datta, M. K.; Velikokhatnyi, O. I.; Manivannan, A.; Alman, D.; Kumta, P. N. Rechargeable Magnesium Battery: Current Status and Key Challenges for the Future. Prog. Mater. Sci. 2014, 66, 1−86. (3) Aurbach, D.; Gofer, Y.; Lu, Z.; Schechter, A.; Chusid, O.; Gizbar, H.; Cohen, Y.; Ashkenazi, V.; Moshkovich, M.; Turgeman, R.; et al. A Short Review on the Comparison Between Li Battery Systems and Rechargeable Magnesium Battery Technology. J. Power Sources 2001, 97−98, 28−32. (4) Vardar, G.; Sleightholme, A. E. S.; Naruse, J.; Hiramatsu, H.; Siegel, D. J.; Monroe, C. W. Electrochemistry of Magnesium Electrolytes in Ionic Liquids for Secondary Batteries. ACS Appl. Mater. Interfaces 2014, 6, 18033−18039. (5) Gregory, T. D.; Hoffman, R. J.; Winterton, R. C. Nonaqueous Electrochemistry of Magnesium: Applications to Energy Storage. J. Electrochem. Soc. 1990, 137, 775−780.

4. CONCLUSIONS In conclusion, we have analyzed the ability of a wide series of electrolytes comprised of completely inorganic chlorides to support Mg electrodeposition and dissolution. In particular, the J

DOI: 10.1021/acs.jpcc.5b03508 J. Phys. Chem. C XXXX, XXX, XXX−XXX

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DOI: 10.1021/acs.jpcc.5b03508 J. Phys. Chem. C XXXX, XXX, XXX−XXX