Article pubs.acs.org/cm
Instability at the Electrode/Electrolyte Interface Induced by Hard Cation Chelation and Nucleophilic Attack Yi Yu,†,‡,□ Artem Baskin,§,∥,□ Carlos Valero-Vidal,†,∥ Nathan T. Hahn,∥,⊥ Qiang Liu,†,# Kevin R. Zavadil,∥,⊥ Bryan W. Eichhorn,‡ David Prendergast,*,§,∥ and Ethan J. Crumlin*,†,∥ †
Advanced Light Source, Lawrence Berkeley National Laboratory, One Cyclotron Road, Berkeley, California 94720, United States Department of Chemistry and Biochemistry, University of Maryland, College Park, Maryland 20742, United States § Molecular Foundry, Materials Science Division, Lawrence Berkeley National Laboratory, Berkeley, California 94720, United States ∥ Joint Center for Energy Storage Research, Lawrence Berkeley National Laboratory, One Cyclotron Road, Berkeley, California 94720, United States ⊥ Sandia National Laboratories, Albuquerque, New Mexico 87185, United States # State Key Laboratory of Functional Materials for Informatics, Shanghai Institute of Microsystem and Information Technology, Chinese Academy of Sciences, Shanghai 200050, People’s Republic of China ‡
S Supporting Information *
ABSTRACT: Electrochemistry is necessarily a science of interfacial processes, and understanding electrode/electrolyte interfaces is essential to controlling electrochemical performance and stability. Undesirable interfacial interactions hinder discovery and development of rational materials combinations. By example, we examine an electrolyte, magnesium(II) bis(trifluoromethanesulfonyl)imide (Mg(TFSI)2) dissolved in diglyme, next to the Mg metal anode, which is purported to have a wide window of electrochemical stability. However, even in the absence of any bias, using in situ tender X-ray photoelectron spectroscopy, we discovered an intrinsic interfacial chemical instability of both the solvent and salt, further explained using first-principles calculations as driven by Mg2+ dication chelation and nucleophilic attack by hydroxide ions. The proposed mechanism appears general to the chemistry near or on metal surfaces in hygroscopic environments with chelation of hard cations and indicates possible synthetic strategies to overcome chemical instability within this class of electrolytes.
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INTRODUCTION While the driving force for electrochemical processes is thermodynamic in nature, the performance of a given materials system may be determined almost entirely by the kinetics of charge and mass transport at materials interfaces. For example, the specific chemistry of the solid electrode/liquid electrolyte interface can dictate the efficiency of active ion desolvation from the electrolyte and deposition or insertion at the electrode. And, if the associated free-energy barriers are insurmountable within accessible potentials or temperatures, then the system may exhibit little useful electrochemical function, with respect to accessible capacity and desirable rates of charge or discharge. Key advances in energy storage and conversion may rely on active chemical control of these interfaces. However, to do this, it is essential to be able to probe and understand interfacial chemistry specifically. Experimentally, we have largely relied on ex situ characterization or indirect inferences from standard electrochemical methods, both of which struggle to unambiguously isolate and identify © 2017 American Chemical Society
chemical and electrochemical processes across functioning electrode/electrolyte interfaces.1−3 Similarly, theoretical methodologies4−6 have commonly derived electrolyte stability at electrodes in terms of bulk characteristics of the electrolyte in idealized conditions with respect to a pristine electrode.7−11 Reevaluating the assumptions regarding interfacial chemistry under operating conditions would seem timely, making the best use of cutting-edge techniques for in situ characterization and first-principles simulations to provide much-needed molecular-scale insight. Our test case is the electrolyte comprising magnesium(II) bis(trifluoromethanesulfonyl) imide (Mg(TFSI)2) dissolved in diglyme and the Mg metal electrode.12 This shall serve as a model for the increased interest in exploiting the high capacities provided by metal anodes and the exploration of multivalent Received: August 11, 2017 Revised: September 19, 2017 Published: September 21, 2017 8504
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Figure 1. Diglyme instability against Mg surfaces. (a) Scheme of the three-electrode electrochemistry setup in the APXPS end-station after the dip and pull procedure in diglyme solvent. (b) APXPS data of core level Mg 1s, O 1s, and C 1s collected on the Mg surface before and after the introduction of diglyme solvent: pristine Mg under vacuum, vapor phase diglyme (1−2 Torr), Mg surface with condensed diglyme, Mg covered by a thin layer of diglyme, and pure liquid diglyme (a droplet). (Main components are labeled. Further information about peak assignments can be found in Table S1 and Figure S1.) (c) Schematics of diglyme decomposition on MgO and Mg(OH)2 surfaces with defects and water impurities. (i) Diglyme chelating the under-coordinated Mg at the MgO kink and H2O dissociating near the active O site. (ii) Free OH− deprotonating the deformed diglyme and the induced bond breaking. (iii) Decomposition products and the regeneration of water. (iv) Diglyme chelating the open Mg site at the Mg(OH)2 kink and the hopping of native surface OH−. (v) Deprotonation of the strained diglyme by OH− and the induced bond cleavage. (vi) Decomposition products and the regeneration of water. 8505
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Figure 2. A plausible mechanism of diglyme decomposition and initial decomposition products. Chelation of the Mg2+ ion with diglyme introduces strain in its backbone and exposes a proton for abstraction. A proton extraction by nucleophilic OH− induces diglyme cleavage and the formation of water. Any under-coordinated and sterically unhindered Mg2+ ion should suffice, including those at kinks on MgO or Mg(OH)2 surfaces.
ions as cheaper or safer alternatives to lithium. At first glance, this would seem a promising combination. The diglyme solvent is nominally inert with respect to Mg metal,13 with a large window of electrochemical stability (up to 8 V,14 with an anodic limit ∼3.9 V versus Mg/Mg2+ on Pt15) and high ionic conductivity resulting from a large solvent donor number15 that provides efficient solvation of metal dications.16 Furthermore, Mg metal is expected to form a stable surface passivation in the presence of oxidative impurities without the propensity for dendrite growth infamous in the case of Li metal and Li-ion anodes. Interestingly, despite these attractive characteristics, in practice the system delivers dismal Coulombic efficiency even after electrolyte conditioning and distillation (≤87%) and a significant overpotential for Mg deposition and stripping.17 The origin of its poor electrochemical performance is still debated due to conflicting experimental observations and theoretical explanations of the Mg anode passivation mechanisms as well as the role of chemical impurities such as water, oxygen, and HTFSI.15,18−21 What is needed is a probe of interfacial activity under the conditions of the working cell. We employed an in situ characterization technique, ambient pressure tender X-ray photoelectron spectroscopy (APXPS), combined with the “dip and pull” method22−24 (Figure 1a) and interpret the results using first-principles quantum chemistry calculations employing density functional theory. Our experimental observations indicate a previously unexplored chemical instability of the electrolyte. In stark contrast to the perceived inertness of diglyme with respect to Mg metal, APXPS results undoubtedly indicate solvent instability under open circuit conditions. Our computational modeling reveals that the nominally benign passivation of the Mg anode via oxidation is the origin of the diglyme instability. In the presence of trace quantities of water, either in the hygroscopic solvent or adsorbed on the introduced anode, surface defects can act as reactive sites to produce hydroxide ions. Under-coordinated Mg sites at the surface become chelated by diglyme, rendering the solvent vulnerable to decomposition following deprotonation initiated by the hydroxide ions. Furthermore, at the interface, contact-ion-pair formation is the major cause of chemical decomposition of TFSI−, also due to reaction with hydroxide ions. This work clarifies the performance limitations of the Mg(TFSI)2−diglyme electrolyte at the Mg electrode surface and provides atomistic insight for selection or design of stable electrolytes or electrode surfaces for Mg, other multivalent ions, or reactive metal anodes more generally. It also offers an illustrative example of the nontrivial role of
interfacial chemistry in the destabilization of a nominally stable electrolyte.
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RESULTS AND DISCUSSION Diglyme Instability at the Mg Electrode Surface. To investigate the interactions between diglyme and the Mg electrode surface, in situ APXPS measurements were performed systematically on the Mg metal electrode, the Mg(TFSI)2− diglyme electrolyte, and increasingly realistic models of their working interface. The experimental setup is depicted in Figure 1a (see Methods for more details). The electrode was first studied alone under vacuum condition in the chamber. When a beaker of the solvent was introduced into the analysis chamber, the diglyme vapor spectra and the condensed diglyme features on the Mg surface were determined. Dipping the Mg electrode into the diglyme solvent and then removing them allowed us to access spectra of both the bulk solvent and the electrode/ electrolyte interface by placing the electrode in front of analyzer in the respective locations. Under vacuum conditions, due to the known reactive nature of Mg metal in the presence of trace impurities, a film was found covering the metal surface, which included magnesium oxide, -hydroxide, and -carbonate components (Figure 1b, Vacuum). When diglyme is introduced and its vapor reaches equilibrium with the Mg surface (Figure 1b, Diglyme Condensed), photoemission features of diglyme vapor (532.825,26 and 286.5 eV;27−32 Figure 1b, Vapor) and liquid diglyme (531.625 and 285.8 eV;33 Figure 1b, Diglyme Bulk) are observed. But the most important findings in this Diglyme Condensed condition are the unexpected features: 533.6 eV and 532.5 eV in O 1s and 289.2 and 287.5 eV in C 1s spectra (Figure 1b, Diglyme Condensed). These features are not attributable to the aforementioned Mg surface components, nor diglyme vapor or liquid components. The signals at 533.6, 532.5, and 289.2 eV indicate (OC)−O moieties,26,34,35 and that at 287.5 eV signals the presence of carbonyl, C O.27,31,32,36 If they originate from the solvent, these signals imply bond-breaking along the diglyme backbone, which is comprised exclusively of ether moieties (C−O−C). The liquid diglyme features are reaffirmed upon performing the dip and pull procedure (Figure 1b, Diglyme Layer and Diglyme Bulk). The results in Figure 1b clearly indicate that new chemical species are formed at the interface when diglyme is in contact with the Mg electrode surface. A plausible mechanism for the observed diglyme decomposition as well as its initial decomposition products is depicted in Figure 2. Though one can assume other concomitant (cascade) reactions leading 8506
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Figure 3. Mg decompositions at the electrode/electrolyte interface. (a) Core level APXPS data of Mg 1s, O 1s, C 1s, S 1s, F 1s, and N 1s collected on the Mg electrode with the presence of 0.8 M Mg(TFSI)2−diglyme thin layer solution under a pressure of 1−2 Torr: Mg electrode covered by a thin layer of electrolyte solution after the dip and pull procedure before electrochemical CV (Before CV), Mg electrode covered by a thin layer of electrolyte solution after CV (After CV), and pure liquid electrolyte solution (Electrolyte Bulk) (main components are labeled in the figure for simplicity; a summary of XPS peak assignments can be found in Table S1 and Figure S1). (b) Simulations showing the weakening of the C−S bond in the [Mg(TFSI)]+ ion pair in the liquid phase as a result of nucleophilic attack by OH− (diglyme molecules not directly coordinating Mg are not shown).
Density functional theory (DFT) calculations indicate that isolated diglyme molecules do not undergo spontaneous decomposition upon single-electron reduction or oxidation (Figure S2). When the diglyme molecule is placed near the pristine hexagonal close-packed (0001) Mg metal surface (Figure S3), a strong physisorption (ca. 1.4 eV per molecule w/
eventually to the products containing carbonyl, carboxyl, or ester groups, here, we limited ourselves to determining the initial step in diglyme decomposition. Quantum mechanical calculations, Figure 1c, are provided to support that Mg2+ chelation and trace water impurities21 play key roles in this process. 8507
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Simulations show that low Miller index, pristine surfaces of magnesium oxide or hydroxide, both confirmed to be present on the mechanically abraded Mg surface (Figure 1b), appear to have no appreciable reactivity with diglyme (Figure S16). However, as it is shown in the previous studies,40−44 water may undergo dissociative adsorption on well-ordered MgO surfaces and, depending on the amount of the content, can promote the formation of Mg(OH)2 and other aggregates containing Mg(OH) along with hydroxide ions. In our study, we showed that MgO (and Mg-suboxide) surfaces with under-coordinated (e.g., threefold) O2− sites can spontaneously dissociate water producing surface-bound and free OH− ions (Figures S14, S15, and S17).40,42,43 Moreover, under-coordinated Mg sites at MgO or Mg(OH)2 kinks can effectively act as coordination centers for diglyme molecules with sufficient space for chelation, inducing the destabilizing strain mentioned above (Figure S11). It should be emphasized here that both the water dissociation process and the diglyme decomposition are self-limiting processes that depend on the surface activation (for water dissociation) and the availability of surface defects to induce the chelation. Therefore, we propose the following intersection of surface chemistries: (1) diglyme molecules readily chelate under-coordinated Mg2+ sites on the oxidized anode surface; (2) OH− is present at the surface, produced either from trace water dissociation or released from Mg(OH)2 by intersite hopping; and (3) without direct access to the Mg2+ itself, OH− could readily deprotonate the deformed diglyme molecules and induce their spontaneous cleavage (Figures S8). Though other mechanisms of water dissociation are possible (e.g., on the bare Mg surface, yet with a high potential barrier ∼2.5 eV/per water molecule45), the general scheme indicated in Figure 2 seems most plausible for a freshly introduced anode with surface oxidation (before any electrochemical deposition) and consistent with initiating the formation of moieties resulting from diglyme cleavage, evident in APXPS and highlighted schematically in Figure 1c. Mg(TFSI)2 Decomposition near the Open Circuit Equilibrated Mg Electrode Surface. Next, we address chemical instability of the electrolyte salt component Mg(TFSI)2. Characterizations of the interface between the Mg electrode and the electrolyte containing 0.8 M Mg(TFSI)2 solution in diglyme were performed systematically. The APXPS measurements were obtained at three different conditions: (1) before cyclic voltammetry (CV), prior to any externally applied potential (Figure 3a, Before CV, see Experimental Section for more details); (2) after reimmersion into the electrolyte solution for cycling (Figure 3a, After CV); and (3) in the bulk electrolyte solution far from the electrode surface in (Figure 3a, Electrolyte Bulk). APXPS results clearly indicate that TFSI components chemically decompose at the electrode/electrolyte interface prior to CV. Upon introduction of the Mg(TFSI)2 salt, S 1s, F 1s, and N 1s photoemissions are detected. In Figure 3a, Electrolyte Bulk, the bulk electrolyte solution presents a single sulfur and fluorine feature (2476.446,47 and 687.6 eV13,27−30,33,34), which are associated with the TFSI’s trifluoromethylsulfone (−SO2CF3) component. However, when the electrode/electrolyte interface (before any electrochemistry) is probed, additional sulfur (2474.2 and 2470.2 eV) and fluorine (684.5 eV) features are observed (Figure 3a, Before CV). These peaks can be assigned to oxidized sulfur (−SOx),46,47 sulfide (S2−),46,47 and fluoride (F−)13,27−31,33,34,48 species, respectively. The observation of interfacial sulfur and fluorine species different
r to the vacuum) is observed; however, diglyme remains reductively stable, and no spontaneous decomposition occurs. In previous work,37 the adsorption energy of monoglyme on metallic magnesium was estimated as little as ∼0.1 eV per molecule with the bulk solvent as the reference state. The discrepancy originates not only from the choice of the reference state but also from the complications associated with the procedure for evaluating the free energy difference of the solvent molecule in the adsorbed state and in the bulk, which is not a trivial task. Here, we used the simplified approach that shows the strong physisorption cannot explain the decomposition of diglyme on the metallic Mg surface. The weaker adsorption would favor the stability of diglyme even more. In Figure S4, we show the results of the analysis for an interfacial layer of diglyme in vicinity to the Mg metal with the assumption of surface neutrality or extra negative charge localized on the top layer. Ab initio molecular dynamics (AIMD) simulations show the stability of the interfacial layer and no spontaneous decomposition of the solvent. On the other hand, the electronic projected density of states (PDOS) of a single diglyme molecule on the Mg(0001) surface shows that the frontier molecular orbitals (highest occupied or lowest unoccupied molecular orbitals, HOMO and LUMO, respectively) of diglyme reside ∼3 eV away from the Fermi level for the neutral system. Only subtle changes of their relative energies are observed upon charging of the Mg surface (Figure S5). Our calculations show that diglyme decomposition is not induced by direct interactions with the pristine or defective Mg metal surface (Figure S6), and we see no evidence for unimolecular reactions leading to diglyme cleavage. Although diglyme is known to possess great stability in basic environments, it may decompose in the presence of Lewis acids and strong nucleophiles.38 In this study, the Mg2+ cation and OH− anion could serve as a sufficiently strong Lewis acid and nucleophile, respectively (Figure S7). However, simulations of Mg2+(diglyme)1 complexes suggest that for decomposition to occur, an intermediate step, the deprotonation of diglyme by OH−, cannot be reached without significant deformation of the diglyme molecule (Figures S9 and S10). Once diglyme is deprotonated, its backbone becomes vulnerable to bond scissions. In isolation and in the neat liquid, diglyme molecules tend to remain linear or extended to reduce electrostatic dipole interactions between the three ether moieties. However, diglyme’s ability to solvate Mg2+ dications is driven by the strong chelation, where multiple neighboring ether oxygens are involved in coordination (Figure S12). In this molecular conformation, deprotonation by nucleophilic OH− results in spontaneous cleavage of diglyme, as described in Figure 2. In fact, analysis of the most stable configurations involving different simple metal cations (Figure S13) shows that the largest backbone strain is induced by Li+ and Mg2+ due to their small ionic radii and short bonds to oxygen.39 This might explain the known stability of diglyme molecules in common basic solutions (e.g., NaOH or KOH) from the proposed instability upon chelation of Mg2+ (or even Li+ by the same logic). However, we note that deprotonation by hydroxide ions may be unlikely (or undetectable) in the bulk electrolyte, where we expect very low concentrations of water and even less of its decomposition products. And, without constraints, in the liquid phase, it is more likely that hydroxide ions will compete with diglyme to coordinate Mg2+ directly. Instead, we expect that diglyme decomposition resulting from deprotonation following Mg2+ chelation (Figure 2) is prevalent at the electrode surface. 8508
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Chemistry of Materials from the −SO2CF3 structure indicate that, at the electrode/ electrolyte interface, the Mg(TFSI)2 decomposes through a chemical redox process.21,49 The lower core-level binding energy (BE) of sulfur in −SOx compared to that of −SO2CF3 can be explained by the loss of the −CF3 group. This is consistent with theoretical predictions that cleavage of the S−C bond is the first reductive decomposition step of TFSI components in Mg(TFSI)2−diglyme systems,4,19 although we do not invoke electrochemical reduction in this context. The interfacial sulfide S2−, which has a much lower BE, must come from further reduction of the −SOx group, while interfacial fluoride F− results from −CF3 decomposition. This is most likely driven thermodynamically by reactions with Mg metal to form stable sulfides and fluorides of magnesium. The APXPS measurements in Figure 3a, After CV show no obvious differences at the interface after cycling, further indicating that the observed decomposition/reduction reactions are not due to electrochemical modifications associated with cathodic polarizations on the electrode. Instead, such reactions are chemical interactions between the electrode and the electrolyte solution. Our simulation studies suggest that such processes could be assisted by electrode surface defects (see below). We also observe from electrochemical APXPS measurements that the decomposition products disappear during electrochemical Mg dissolution under an anodic bias (Figure S18). Once the applied potential returns to zero, the reaction products soon recover at the interface. This provides evidence for the self-limiting property of these chemical processes at the electrode/electrolyte interface and indicates a mechanism for possible interphase formation due to the accumulation of (assumed to be electronically insulating) reaction products at the interface; further decomposition/ reduction of the electrolyte salts is inhibited. Our spectroscopic findings contrast with previous experimental and theoretical studies of TFSI-based ionic liquids or electrolytes with Li or Mg cations, which customarily consider the TFSI decomposition as an electrochemical process driven by electronic charge transfer across the interface.10,20,50 Experimental studies indicate that the reduction potential of the TFSI− anion in ionic liquids is well above the reduction potentials of Mg or Li and that the decomposition is dependent on the electrode substrate and water content.50 Theoretical investigations are divided in their assessment of TFSI stability next to anode surfaces. A combined classical molecular dynamics/DFT study of the methyl, propyl-pyrrolidinium bis(trifluoromethylsulfonyl)imide) (Py1,3TFSI, P13/TFSI) ionic liquid claims the reductive stability of TFSI− next to the clean Li surface,10 implying its stability with respect to the Mg surface as well. On the other hand, the explicit modeling of Li metal interfaces with P13/TFSI ionic liquid or various LiTFSI electrolyte solutions shows the intrinsic (reductive) instability of the TFSI− anion, which decomposes through C−S bond cleavage and is associated with a substantial charge transfer.2,7,8 In conjunction with the previous studies, our computational studies suggest that Mg(TFSI)2 in diglyme is chemically stable in its solvated form as well as against the pristine Mg metal surface. Quantum mechanical calculations and AIMD studies of the diglyme-solvated reduced anion (TFSI2−) indicate no spontaneous decomposition (Figure S20). Although the electron affinity (EA) is significantly enhanced when the [Mg(TFSI)]+ contact ion-pair is formed (Table S2),51 further solvent coordination of Mg2+ in the contact ion-pair (Figure S21) stabilizes the charge state (reducing its EA) so that ion-
pair reduction is nonspontaneous at open circuit potential. AIMD simulations of the TFSI− anion and [Mg(TFSI)]+ ionpair next to the bare Mg surface (Figure S22) show no spontaneous reduction or chemical instability (Figure S23) as well. However, inspired by our findings for diglyme, further investigations into the role of nucleophiles in this system shed light on possible reaction pathways. Although the TFSI− ion is stable in basic environments,52 we find that the cationic [Mg(TFSI)]+ ion-pair readily reacts with OH− ions. A key finding is that the nucleophilic attack of OH− results in a metastable state with an overcoordinated sulfur atom and a significantly weakened C−S bond (see Figure 3b). As discussed in the previous section, OH− ions could be produced at the MgO interface from water impurities or simply by releasing OH− ions from the Mg(OH)2 surface. This implies that the [Mg(TFSI)]+ decomposition could become spontaneous at the electrode surface where OH− ions are in appreciable concentration. Further decomposition could occur near the surface Mg- or O-defect sites to produce Mg-CF3 and sulfites, which might further decompose next to a large source of magnesium to MgF2 and MgS,53 respectively. These findings from theoretical simulations support our APXPS interpretation that TFSI decomposition at this electrode/electrolyte interface is a chemical process, occurring prior to any electrochemically driven processes. The fact that defect sites on the electrode surface and the trace amount of water impurities are responsible for the decomposition reactions agrees well with our observation that electrolyte decomposition occurs only at the electrode/electrolyte interface. Once the products accumulate at the interface and passivate active sites of the electrode surface, further decomposition of the electrolyte solution is inhibited. Hence, no evidence for decomposition is observable in the bulk solution (Figure 3a, Electrolyte Bulk).
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CONCLUSION/OUTLOOK We performed in situ characterizations of the Mg electrode/ diglyme solvent interface as well as the Mg electrode/0.8 M Mg(TFSI)2−diglyme electrolyte interface using in situ APXPS combined with the dip and pull technique. The introduction of diglyme solvent and electrolyte solution on the Mg electrode surface and comparative analysis of APXPS results among various conditions allowed us to identify the instabilities of diglyme solvent and TFSI electrolyte species at the Mg electrode surface. Theoretical simulations suggest that defects on the oxidized surface, water impurities, and/or free OH− ions released from Mg(OH)2 surfaces conspire to induce the chemical decomposition of diglyme and TFSI at the electrode/ electrolyte interface before any electrochemistry. The fact that diglyme and Mg(TFSI)2 are potentially chemically unstable for the same reason, namely decomposition by nucleophilic OH− attack, highlights the generality of insight gained from this APXPS study as well as its value in revealing the importance of trace impurities at evolving solid/liquid interfaces. The importance of this finding becomes more pronounced if one considers the variance of water content in commercially available solvents and intrinsic reactivity of metal electrode surfaces. We discovered a direct correlation between conformational strain in diglyme induced by chelation of hard cations and its subsequent susceptibility to chemical decomposition via deprotonation. This suggests that enhanced solvent stability vs Mg2+ or Li+ could be achieved by appropriate 8509
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implicit solvent model (IEF-PCM), and the solute/solvent interface was constructed using the van der Waals surface defined from the universal force field60 radii scaled by 1.1. The dielectric constant for the diglyme solvent was set as 7. e. Effective Screening Medium (ESM) Calculations. For the isolated diglyme, the TFSI− anion and the [Mg(TFSI))+ ion pairs with different configurations of the anion and cation with respect to the Mg-surface and the mutual alignment of Mg state with respect to the Mg-slab Fermi energy were evaluated within the ESM model61 implemented in the SIESTA package62 (PBE functional and DZP basis set). The slab was modeled by a unit cell (10 × 10 × 3, 360 atoms, a = b = 3.208 Å). The distance between the exposed surface of the slab and the fictitious electrode was set as 17.5 Å. The cutoff energy of 400 Ry was used. The Brillouin zones of the unit cells were sampled by the Monkhorst−Pack grid with the k-point grid 3 × 3 × 1. Geometry optimization was carried out until all the forces were less than 0.04 eV/Å and the stress in the periodic direction was smaller than 0.01 GPa. f. Ab Initio Molecular Dynamics Simulations. AIMD simulations were conducted for solvent molecules, the TFSI anion, and [Mg(TFSI)]+ ion pairs at the PBE-D3 level with the TZV2PX (DZVP for Mg) basis set as implemented in the CP2K package.63 For all AIMD simulations, the energy grid cutoff was set as 320 Ry, and the total energy was sampled at the Γ-point. Using a sampling of 0.5 fs, we performed simulations of ca. 20 ps within canonical ensemble (NVT) at 300 K with the Nosé−Hoover thermostat. For simulations of the molecular species next to the metal slab a box of 33.35305 × 50.0 × 32.094 Å3 was used. For non-neutral systems, the universal background compensating charge scheme and the Poisson wavelet solver with two-dimensional (x−z) periodic boundary conditions were used. We used the Fermi−Dirac smearing (electronic temperature 300 K) and the Broyden mixing (fraction of new density set as 0.1, denominator parameter for Kerker damping set as 1.5).
chemical modifications of the ethers or through the introduction of a scavenger mechanism for OH− removal. On the other hand, the vulnerability of the contact ion pair [Mg(TFSI)]+ to nucleophilic OH− attack could be alleviated by seeking anion substitutes with reduced (single-point) coordination of Mg2+. Additionally, strategies to tailor the working Mg electrode surface to prevent OH− release, water dissociation, or to reduce the population of under-coordinated Mg2+ should be explored. What is clear is that direct in situ observations of the electrode/electrolyte interface are key to inspiring firstprinciples molecular-scale mechanistic investigation of chemical processes that limit electrochemical performance. Further studies of this kind, which self-consistently combine experimental observation and theoretical insight, will expand our ability to understand and control interfacial processes relevant to electrical energy storage.
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EXPERIMENTAL SECTION
a. Electrodes and Solution Preparation. The Mg foils (99.9%, Alfa Aesar) were scraped clean with a razor under nitrogen atmosphere and used as the WE, RE, and CE. The 0.8 M Mg(TFSI)2 solution was prepared by adding predetermined amounts of Mg(TFSI)2 (Alfa Aesar) to diglyme (99.5%, Sigma-Aldrich) in an argon filled glovebox. The chemicals were used as received. b. APXPS Measurement and Analysis. APXPS experiments were performed at Beamline 9.3.1 of the Advanced Light Source. The X-ray incident angle was 75°, and the electron emission angle was 0° with respect to the sample normal. The photon energy was 4.0 keV. A 0.3 mm diameter aperture cone was used. The analyzer (R4000 Hipp2, Scienta) pass energy was 200 eV with step size of 100 meV and a dwell time of 0.2 s. Pressure ranged from vacuum (
