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      Magnesium Borohydride: From Hydrogen Storage to Magnesium Battery**

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          Since Bogdanović and Schwickardi illustrated the possibility of reversibly storing hydrogen in sodium alanate, [1] extensive research efforts have been dedicated to investigating the hydrogen storage potential of complex metal hydrides. In particular, borohydrides have attracted great interest because of their superior gravimetric hydrogen content. [2] Of these, magnesium borohydride Mg(BH4)2, first reported in 1950 [3] and more recently studied for hydrogen storage, has attracted attention because of its relatively low hydrogen-release temperature and reversibility. [2a] , [4] Furthermore, borohydrides are strong reducing agents that are widely used in organic and inorganic syntheses. This reducing power translates to high stability against electrochemical reduction; this stability could be exploited in highly reductive environments, such as battery anodes. Therefore, for the first time, we have conducted research towards harnessing this property of borohydrides for their use in rechargeable batteries. In particular, we have been focusing on utilizing a Mg(BH4)2 based electrolyte in a rechargeable magnesium battery. Recently, magnesium batteries have received increased attention as alternatives to the lithium-based battery because of the high volumetric capacity (3832 mA h cm−3), improved safety (nondendritic), and abundance of Mg metal. [5] Despite the potential of Mg batteries, several key challenges need to be overcome for this technology to become viable. For instance, current state-of-the-art electrolytes use organomagnesium salts and complexes as they are the only ones known to be compatible with the Mg anode that allow for reversible electrochemical Mg deposition and stripping. [5b] , [6] Although some of these electrolytes have shown impressive stability against electrochemical oxidation, they were also found to be corrosive. [6] This property was attributed to the presence of chlorides in either/both their cations and anions. [6] Conventional inorganic and ionic salts such as Mg(ClO4)2 were found to be incompatible with the Mg anode as a result of the formation of an ion-blocking layer formed by their electrochemical reduction. [6] Hence, the discovery of halide-free electrolytes with high reductive stabilities is crucial for realizing a practical rechargeable Mg battery system. Herein, we propose a new class of electrolytes based on Mg(BH4)2 for a Mg battery. We show the first example of electrochemical reversible Mg deposition/stripping in a halide-free inorganic salt in both tetrahydrofuran (THF) and dimethoxyethane (DME) solvents. An increase of several orders of magnitude in the current densities, and high coulombic efficiencies of up to 94 % are observed in DME when LiBH4 is used as an additive. Furthermore, we use this electrolyte in a rechargeable Mg battery, thus giving the first example of a borohydride electrolyte in a battery system. This work also illustrates the unique properties of borohydrides and opens the door for designing a whole new class of electrolytes for Mg batteries. Mg deposition/stripping was studied for Mg(BH4)2 in ether solvents. Figure 1 a shows the cyclic voltammogram obtained for 0.5 M Mg(BH4)2/THF where a reversible reduction–oxidation process took place with onsets at −0.6 V/0.2 V and a 40 % coulombic efficiency (Figure 1 a, inset), thus indicating reversible Mg deposition and stripping. X-ray diffraction (XRD) confirmed that the deposited product from the galvanostatic reduction of the above solution (Figure 1 b) was hexagonal Mg, hereby establishing the compatibility of Mg(BH4)2 with Mg metal. The electrochemical oxidative stabilities measured on platinum, stainless steel, and glassy carbon electrodes were 1.7, 2.2, and 2.3 V, respectively (Figure S7). These results showed that for the first time: 1) Mg(BH4)2 is electrochemically active in THF, that is, ionic conduction is possible, and 2) reversible magnesium deposition/stripping from an inorganic, relatively ionic (Mg Bader charge is +1.67) [7] and halide-free salt is feasible. Although these results are promising, to make this electrolyte more practical for use in batteries the electrochemical performance needs to be improved by lowering the overpotentials, and achieving higher current density and coulombic efficiency. In addition, the demonstration of this performance in less-volatile solvents would make Mg(BH4)2 based electrolytes even more practical. Therefore, DME was selected (its boiling temperature is 19 °C higher than that of THF) for further investigations. The cyclic voltammogram obtained for 0.1 M Mg(BH4)2/DME is shown in Figure 1 c where a substantial improvement in the electrochemical performance compared to Mg(BH4)2/THF was evident from: 1) a 10-fold increase in the current density, 2) a reduction in the overpotentials (deposition/stripping onsets at −0.34 V/0.03 V versus −0.6 V/0.2 V in THF), and 3) a higher coulombic efficiency of 67 % (40 % in THF). These findings suggested that the Mg electroactive species was present in higher concentration and had increased mobility in DME despite the lower solubility of Mg(BH4)2 in DME versus THF. Figure 1 For 0.5 M Mg(BH4)2/THF: a) Cyclic voltammogram (8 cycles), inset shows deposition/stripping charge balance (third cycle), and b) XRD results following galvanostatic deposition of Mg on a Pt working electrode. c) Cyclic voltammogram for 0.1 M Mg(BH4)2/DME compared to 0.5 M Mg(BH4)2/THF. Inset shows deposition/stripping charge balance for Mg(BH4)2/DME. All experiments used Pt working electrode and Mg reference/counter electrodes. These results demonstrated that for the Mg(BH4)2 electrolyte, the electrochemical performance in DME is higher than that in THF. In contrast, organomagnesium electrolytes show an optimum electrochemical performance in THF.5b To further improve the electrochemical performance, it was pertinent to characterize the electroactive species in Mg(BH4)2 solutions. Therefore, IR and NMR spectroscopic analyses (Figure 2) were conducted for 0.5 M Mg(BH4)2/THF and 0.1 M Mg(BH4)2/DME. The IR B–H stretching region (2000–2500 cm−1) showed two strong widely separated bands (Mg(BH4)2/THF: 2379 cm−1, 2176 cm−1 and Mg(BH4)2/DME: 2372 cm−1, 2175 cm−1); note that the spectra for 0.1 and 0.5 M of Mg(BH4)2 in THF are similar (Figure S2). These IR spectra are similar to those of covalent borohydrides [8] and those of Mg(BH4)2 solvates from THF and diethyl ether [9] where two hydrogen atoms in BH4 − are forming a bridge to one metal atom (μ bonding). Therefore, we assigned the bands at the higher and lower B–H frequencies to terminal and bridging B–H vibrations (B–Ht and B–Hb), respectively. The band and shoulder at 2304 and 2240 cm−1 were assigned to asymmetric B–Ht and B–Hb vibrations, respectively. As complete dissociation of Mg(BH4)2 into discreet ions is unlikely (as other borohydrides are in ethers), [10] we propose that Mg(BH4)2 is present as the contact ion pair Mg[(μ-H)2BH2]2, which partially dissociates into [Mg{(μ-H)2BH2}]+ and BH4 − as in [Eq. (1)]; since the different B–H bands most likely overlap, it is not possible to distinguish all the species. 1 Figure 2 For Mg(BH4)2 in THF (red line) and in DME (black line): a) IR spectra, b) 11B NMR spectra, and c) 1H NMR spectra. Where [Mg{(μ-H)2BH2}]+ may further dissociate: 2 For the spectrum of Mg(BH4)2/DME, although the main features present in the spectrum of Mg(BH4)2/THF were retained, the νB–Ht band is broader and shifted to a lower value and the νB–Hb intensity is relatively weaker. Although νB–Ht band broadening suggests a pronounced presence of a species similar to that found in THF, the shift in the band maximum indicates a more-ionic B–H bond (the νB–Ht shift is similar to those resulting from BH4 − ions that have enhanced ionic character, such as in stabilized covalent borohydrides). [8] In addition, the relative weakening in νB–Hb intensity suggests that there is more free BH4 −. The NMR spectrum of BH4 − in DME (Figure 2 b and c) indicates that there is increased boron shielding as the associated signal is shifted by about 0.5 ppm (quintet in 11B NMR spectrum), and slightly reduced proton shielding (0.01 ppm, quartet in 1H NMR spectrum); these results are consistent with B–H bonds that have a higher ionic character than those in BH4 − in THF (distinguishing B–Ht from B–Hb is not possible likely because of rapid hydrogen exchange). These findings are evidence of weaker interactions between Mg2+ and BH4 − within the ion pair and an enhanced dissociation in DME [Eq. (1) and (2)]. So despite the fact that DME has a slightly lower dielectric constant (7.2) compared to THF (7.4), its chelation properties (owing to the presence of two oxygen sites per molecule) [11] resulted in an enhanced dissociation and thus an improved electrochemical performance. Based on the understanding gained of the nature of Mg(BH4)2 in solution, we hypothesized that electrochemical performance would be enhanced when the association within the ion pair is weakened. To achieve this, an additive that has an acidic cation with the following characteristics is desirable: 1) reductive stability comparable to Mg(BH4)2, 2) nonreactive, 3) halide free, and 4) soluble in DME. Hence, LiBH4 was selected as it fulfils all of the above criteria. Mg deposition and stripping was studied in DME using different molar ratios of LiBH4 to Mg(BH4)2; Figure 3 a shows the cyclic voltammogram obtained for 3.3:1 molar LiBH4 to Mg(BH4)2 (Figure S8a and S8b show the cyclic voltammograms for different concentrations). The use of LiBH4 resulted in an increase of two orders of magnitude in the current density (i.e. oxidation peak current Jp=26 mA cm−2), and in a higher coulombic efficiency of up to 94 %. We attribute the deposition/stripping currents solely to Mg because of the absence of Li after galvanostatic deposition (Figure 3 b), and also the lack of electrochemical activity in a LiBH4/DME solution (Figure S8a). The ionic character of BH4 − was enhanced, as evidenced by lower νB–Ht and higher νB–Hb bands in the IR spectrum (Figure 3 c), thus implying that LiBH4 has a role in increasing Mg(BH4)2 dissociation (the B–H bands for LiBH4/DME occur at lower values, Figure S9). The coulombic efficiency was proportional to the molar ratios of LiBH4/Mg(BH4)2 (Figure S10). A rechargeable Mg battery with a Chevrel phase Mo6S8 cathode, an Mg metal anode, and this optimized electrolyte (Figure 4) demonstrated reversible cycling capabilities at a 128.8 mA g−1 rate (capacity retention and cathode magnesiation are shown in Figure S11 and Figure S12). We are currently investigating the sources of the overcharge and capacity fade. Figure 3 For LiBH4 (0.6 M)/Mg(BH4)2 (0.18 M) in DME: a) Cyclic voltammogram (inset shows deposition/stripping charge balance). b) XRD results following galvanostatic deposition of Mg on a Pt disk. c) IR spectra (red | indicates band maxima for Mg(BH4)2/DME). Figure 4 Charge/discharge profiles with Mg anode/Chevrel phase cathode for 3.3 molar LiBH4/Mg(BH4)2 in DME. Cycle 1 (blue), cycle 2 (red), cycle 20 (black), cycle 40 (green). In summary, unprecedented reversible Mg deposition and stripping from an inorganic and relatively ionic salt was obtained in THF and DME. Higher current density and lower overpotentials were achieved in DME compared to those in THF. Substantial enhancement in the coulombic efficiency and the current density was accomplished by the addition of LiBH4. Battery performance was demonstrated using a Chevrel phase cathode. Although the oxidative stability (1.7 V vs. Mg on platinum) currently limits Mg(BH4)2 utilization with high voltage cathodes, reversibility in the absence of halides and THF makes this salt extremely unique and these findings very important for designing a whole new class of Mg(BH4)2 based electrolytes. Currently, we are investigating improving the oxidative stability of Mg(BH4)2. In addition, the exact nature of the electroactive species in the presence and the absence of the additive is being studied to guide the design of Mg(BH4)2 based electrolytes. This work provides a stepping stone for extending the applications of Mg(BH4)2 and underscores the beauty and versatility of the chemistry of borohydrides. Experimental Section Magnesium borohydride (Mg(BH4)2, 95 %) lithium borohydride (LiBH4, 90 %), anhydrous tetrahydrofuran (THF), and dimethoxyethane (DME) were purchased from Sigma–Aldrich. Cyclic voltammetry was conducted in a three-electrode cell with Mg wire/ribbon as reference/counter electrodes. The electrochemical testing was conducted in an argon filled glovebox with O2 and H2O amounts kept below 0.1 ppm. Details of the analyses and battery testing conducted are described in the Supporting Information.

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          Building better batteries.

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            Prototype systems for rechargeable magnesium batteries.

             R Turgeman,  Z. Lu,  G Aurbach (2000)
            The thermodynamic properties of magnesium make it a natural choice for use as an anode material in rechargeable batteries, because it may provide a considerably higher energy density than the commonly used lead-acid and nickel-cadmium systems. Moreover, in contrast to lead and cadmium, magnesium is inexpensive, environmentally friendly and safe to handle. But the development of Mg batteries has been hindered by two problems. First, owing to the chemical activity of Mg, only solutions that neither donate nor accept protons are suitable as electrolytes; but most of these solutions allow the growth of passivating surface films, which inhibit any electrochemical reaction. Second, the choice of cathode materials has been limited by the difficulty of intercalating Mg ions in many hosts. Following previous studies of the electrochemistry of Mg electrodes in various non-aqueous solutions, and of a variety of intercalation electrodes, we have now developed rechargeable Mg battery systems that show promise for applications. The systems comprise electrolyte solutions based on Mg organohaloaluminate salts, and Mg(x)Mo3S4 cathodes, into which Mg ions can be intercalated reversibly, and with relatively fast kinetics. We expect that further improvements in the energy density will make these batteries a viable alternative to existing systems.
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              Structure and compatibility of a magnesium electrolyte with a sulphur cathode

              Energy diversification is necessary for global sustainability and the minimization of industrial and automotive pollution. Low-cost, high energy density batteries utilizing environmentally friendly elements used for storing intermittent energy from renewable resources such as solar and wind are bound to increase in demand. High energy density, rechargeable batteries will have a large role in powering market-competitive electric vehicles, where the available space to mount the battery packs dictates the volumetric energy density is more important than gravimetric energy density. Attractive choices are alkaline/alkaline earth metal anodes, which provide some of the highest theoretical volumetric capacities of any anode material: the volumetric capacity of lithium, sodium, calcium and magnesium are 2,062, 1,128, 2,073 and 3,832 mAh cm−3, respectively. For comparison, current graphite anodes for lithium-ion batteries have a volumetric capacity of 777 mAh cm−3. Additionally, metallic anodes do not require solid-state diffusion of ions to transfer material from the charged to the discharged state, but merely the successful deposition/dissolution of the ions onto/from the surface of the metal. Magnesium metal has a high negative reduction potential (−2.356 V versus NHE) and the highest volumetric capacity of the practical choices from group I and II metals (beryllium metal is not a practical choice because of its high cost of $7,480 per kg), which make it a superior alternative as an anode material for high energy density batteries1. Furthermore, Mg is not plagued by dendrite formation, which is a significant safety issue that has dissuaded the commercialization of rechargeable batteries utilizing a lithium metal anode2 3 4. The first rechargeable batteries with Mg metal anodes were demonstrated in 2000. These Mg batteries showed impressive cycle life (>3,500 cycle measured), low capacity fading over prolonged cycling, negligible self-discharge and wide temperature operating range5. However, these batteries were then considered only as replacements for nickel–cadmium or lead–acid batteries because the high formula weight of the Chevrel phase MgxMo3S4 cathode lowered the overall energy density. Further research on alternative high-energy Mg battery systems has also been hindered by the surface chemistries of Mg, which greatly limits the choice of available electrolytes and cathodes6 7. With regard to high-energy systems, one of the ideal materials to couple Mg with is sulphur 8, which has a high theoretical capacity (1,671 mAh g−1 or 3,459 mAh cm−3). The combination of a magnesium anode and a sulphur cathode is of great interest because the theoretical energy density of this battery is estimated to be over 4,000 Wh l−1, which is approximately twice that of a Li ion battery composed of a graphite anode and a cobalt oxide cathode. Unfortunately, Mg electrolytes reported so far6 9 10, while having high coulombic efficiencies, are nucleophilic, and, therefore, preclude the use of electrophilic cathodes such as sulphur. Consequently, the feasibility and performance of a Mg/S battery is completely unknown, because there is no electrolyte compatible with both Mg and S. To couple the two electrodes, an electrolyte able to transport Mg2+ ions between the anode and cathode is essential. In general, the prerequisites for battery electrolytes include electrochemical/chemical stability, ionic conduction and electronic insulation4. Magnesium organohaloaluminate electrolytes, generated in situ from the reaction between a Lewis acid and a Lewis base, are nucleophilic. For example, a 2:1 mixture of phenylmagnesium chloride and aluminum trichloride (AlCl3) in tetrahydrofuran (THF) is incompatible with an electrophilic sulphur cathode. Gas chromatography–mass spectroscopy analysis confirmed that this electrolyte directly reacts with sulphur to form phenyl disulphide and biphenyl sulphide. Consequently, our synthetic strategy was to avoid a direct reaction with sulphur by focusing on utilizing non-nucleophilic bases. The fact that potassium hexamethyldisilazide (KN(SiMe3)2) is a non-nucleophilic base suggests that hexamethyldisilazide magnesium chloride (HMDSMgCl) is an excellent candidate, because it has been reported to be capable of reversible Mg deposition10. Unfortunately, the coulombic efficiency, voltage stability and current density of the HMDSMgCl electrolyte are far inferior to magnesium organohaloaluminate electrolytes, and, as a result, it has not received much attention. Here we report the performance enhancement of HMDSMgCl, through the addition of a Lewis acid AlCl3. Crystallization of the electrochemically active species resulted in a dramatic improvement in the potential stability and coulombic efficiency and, furthermore, it is the critical step in synthesizing a non-nucleophilic electrolyte that is chemically compatible with an electrophilic sulphur cathode. Although the dissolution of sulphur and polysulphides plagues the Mg/S system with rapid fade, we demonstrate a proof of concept for the first rechargeable Mg/S battery. Results Electrochemical analysis of the electrolyte To enhance the electrochemical performance of the HMDSMgCl electrolyte, we investigated its reactivity with a Lewis acid, AlCl3. By varying the ratio of acid to base and the reaction time, we found the optimum electrochemical performance of the electrolyte to be when the ratio of HMDSMgCl:AlCl3 was 3:1 and the reaction time was 24 h. It is noteworthy to state that organomagnesium chemistry is acutely sensitive to moisture and air, and, therefore, all reactions and electrochemical experiments were performed in a glovebox under an argon atmosphere. The current density for Mg deposition is increased by almost a factor of seven by the addition of AlCl3, as shown by the green and blue lines in Figure 1a. Unfortunately, the voltage stability of the HMDSMgCl electrolyte was not improved (Fig. 1b). To clarify the product from the reaction between HMDSMgCl and AlCl3, a crystal was obtained by slow diffusion of hexane. The crystal structure [Mg2(μ-Cl)3·6THF][HMDSAlCl3], 1, was determined by single-crystal X-ray diffraction (Fig. 2), which was solved to reveal a cation consisting of two octahedrally coordinated Mg centres bridged by three chlorine atoms. The three remaining sites on each Mg are occupied by THF molecules coordinated through the oxygen. The counter anion is an aluminum atom tetrahedrally coordinated by one HMDS group and three chlorine atoms. The [Mg2(μ-Cl)3·6THF]+ cation has been previously isolated in the solid state from a THF-based Grignard reagent solution11, from the selective synthesis of a 1,3,4-triphospholide anion12, and from a catalytic metathesis reaction between ZnCl2 and tBuMgCl (ref. 13). Figure 1a compares the electrochemical activity of the crystal redissolved in THF (red) to that of the electrolyte generated in situ from the reaction between HMDSMgCl and AlCl3 (blue). It is impressive that the voltage stability of the electrolyte increased by almost 0.8 V after crystallization. We propose that the purification step removes any unreacted HMDSMgCl, that starts to electrochemically oxidize and decompose around 2.5 V. This is supported by spiking the redissolved, crystallized electrolyte with a HMDSMgCl solution, which results in a decrease of voltage stability from 3.2 to 2.5 V. In addition, crystallization of the electrolyte increases the coulombic efficiency from 95 to 100%, as shown in Figure 1a, inset. The superior electrochemical performance of the crystal 1, is fascinating, considering its structural similarity to the allegedly inactive crystal, [Mg2(μ-Cl)3·6THF][C2H5AlCl3]. This crystal was reported to be obtained from an active electrolyte composed of di-n-butylmagnesium ([CH3(CH2)3]2Mg, Bu2Mg) and ethylaluminum dichloride (C2H5AlCl2, EtAlCl2), but electrochemically inactive when dissolved in THF 5. This contradiction prompted us to reevaluate the electrochemical performance of the crystals obtained from the reaction product of Bu2 Mg and EtAlCl2. We found the crystal was indeed electrochemically active, although the coulombic efficiency was reduced from 100 to 90% (Fig. 1c, and inset). Chemical analysis of the electrolyte The 1H NMR spectrum of the white crystalline product obtained from the reaction mixture of HMDSMgCl and AlCl3 (3:1) dissolved in d8-THF displayed 2 singlets at 0.12 and −0.01 p.p.m.. The 13C {1H} NMR spectrum shows 3 resonances, two high frequency peaks assigned to the protic THF atoms and the peak at 6.1 p.p.m. assigned to the -CH3 groups of the HMDS ligand. The 27Al and the 25Mg NMR both displayed a broad singlet peak at 103.9 and 5.02 p.p.m., respectively. The mass spectroscopy analysis showed an exact mass and isotope pattern consistent with both an HMDSAlCl3 − anion and an HMDS2AlCl2 − anion. Based on the fact that the HMDSAlCl3 − anion will be shifted further downfield than the HMDS2AlCl2 − anion, we assign the peak at 0.12 p.p.m. to the methyl protons of 1, and the peak at −0.01 p.p.m. to the same group in [Mg2(μ-Cl)3·6THF][HMDS2AlCl2], 1. The integration of the two peaks clearly shows that the ratio of 1 to 1 is 97:3. Unfortunately, we could not observe the [HMDS2AlCl2]− co-product by 13C NMR owing to the low yield. However, we feel the presence is sufficiently evidenced through mass spectroscopy and 1H NMR analysis. Equations (1,2,3,4) show the proposed key step in the formation of 1 as the transmetallation of the HMDS group to yield the HMDSAlCl3 − anion [equation (1)]. The resulting MgCl+ cation is not observed, and one plausible scenario to account for the cation in 1 is that the MgCl+ rapidly reacts with MgCl2 [equation (3)]. Some MgCl2 is present from the Schlenk equilibrium [equation (2)], and depletion of the MgCl2 drives the equilibrium to the right. Evidence for the formation of HMDS2Mg was confirmed through 1H NMR by spiking the original reaction mixture with a commercial sample of HMDS2Mg. A proposed pathway to account for the minor product, 1, involves the reaction of 1 with HMDSMgCl [equation (5)]. Evidence for this reaction was found by increasing the concentration of HMDSMgCl and the reaction time, which consequently amplified the percentage of 1 formed. As affirmation of the compatibility, we investigated the reactivity of the electrolyte with elemental sulphur through NMR. We observed no change in the 33S NMR of elemental sulphur in the presence of the electrolyte, even after one week. The 1H, 13C and 27Al NMR of the electrolyte also remains unchanged in the presence of elemental sulphur. Based on the NMR studies, the non-reactive nature of the electrolyte with sulphur clearly demonstrates that the electrolyte is indeed non-nucleophilic and therefore chemically compatible with an electrophilic sulphur cathode. 1 2 3 4 5 X-ray photoelectron spectroscopy of the sulphur cathode Our non-nucleophilic electrolyte is the key that opens the door to examining the performance of a Mg/S battery. Coin cells with a metallic magnesium anode, separator, and a sulphur cathode consisted of elemental sulphur dispersed in carbon black and a polymeric binder, were assembled to test the feasibility of our electrolyte. As Figure 3a shows, a typical Mg/S coin cell displayed an excellent capacity of 1,200 mAh g−1 (2,484 mAh cm−3, based on the mass of sulphur) for the first discharge. The starting potential is 0.55 V and slowly increases up to 0.89 V during the discharge process. The rise is most likely due to the fracturing of the resistive surface layer on the magnesium anode, evidenced by the higher voltage at the start of the second discharge, when the surface layer is no longer present. The second discharge capacity dramatically decreased to 394 mAh g−1 (816 mAh cm−3), which stimulated our investigations into the reasons for the large capacity fade. To examine the components after undergoing a discharge/charge cycle, we dismantled a Mg/S coin cell for visual inspection and chemical analysis. The apparent yellow discolouration of the separator indicated that the main cause of capacity fade is polysulphide or sulphur dissolution14, which also explains the overcharging behaviour observed in Figure 3a. We note that our preliminary coin cells are not optimized to prevent sulphide dissolution, and therefore display typical problems of a S cathode. Polysulphide dissolution, a rapid capacity fade, and overcharging are well-known problems in Li/S battery system15. Polysulphides can be charged to various chain lengths and migrate between the anode and cathode via a sulphide shuttle mechanism. During overcharge of Li/S battery, soluble polysulphides are formed on the cathode, diffuse to the anode to be reduced there, and, then, diffuse back to the cathode to be reoxidized. Consequently, the current can continue to flow without actually oxidizing Li2S to S, or 'recharging' the battery. One of the most direct ways to gain a deeper understanding of the battery chemistry is X-ray photoelectron spectroscopy (XPS), because it can determine the change in the oxidation state of sulphur on the surface of the cathode as the result of battery cycling. Figure 3b compares high-resolution (HRES) S 2p spectra obtained from sulphur cathodes (as-prepared, after a first discharge, and after a complete discharge/charge cycle), to standard samples of S and magnesium sulphide (MgS) powder. The as-prepared sulphur cathode was composed of elemental sulphur and is evidenced by the S 2p3/2 peak located at 164.0 eV. After discharging the Mg/S coin-cell, the oxidation state of the sulphur in the cathode changed dramatically. Curve-fit analysis of the S 2p peak showed that three oxidation states of sulphur were present. As expected from the battery chemistry, the majority of the sulphur was reduced to lower oxidation states indicated by the lower binding energies. The 2p3/2 peak positioned at 160.9 eV confirmed the conversion to MgS. The extra peaks observed between S and MgS are most reasonably assigned to magnesium polysulphides (MgSx, 1 99.0%, Fluka). Freshly prepared ethyl magnesium chloride (EtMgCl, 70 ml, 140 mmol in THF) is added slowly via the addition funnel and the mixture is refluxed until the EtMgCl is completely consumed, as indicated by phenanthroline. The reaction is refluxed for 8 h. Once all EtMgCl is consumed extra EtMgCl is added in 0.5 ml portions with refluxing until an excess persists. This excess EtMgCl is then carefully back-titrated in a similar manner with further hexamethyldisilazane until the end point is achieved. The final concentration is determined by titration with diphenylacetic acid. Preparation of (Mg2(μ-Cl)3·6THF)(HMDSnAlCl4−n) (n=1, 2) In an argon-filled glovebox, AlCl3 (0.5 M solution in THF, 3 ml, 1.5 mmol) was mixed with 3 equivalents of freshly prepared HMDSMgCl (1.44 M solution in THF, 3.125 ml, 4.5 mmol) in a 20 ml screw capped vial. The vial was immediately capped and vigorously stirred for 24 h. The crystals were formed by slow diffusion of anhydrous hexane (Sigma-Aldrich, 6 ml). The resulting crystals were washed with hexane and dried under vacuum to furnish ~800 mg of white crystalline product. (60% yield) 1H-NMR (d8-THF, 500 MHz): δ 0.12 p.p.m. (s, CH3 groups on 1), −0.01 (s, CH3 groups on 1). 13C-NMR (d8-THF, 500 MHz): δ 68.4, 26.4, 6.1. 27Al-NMR (d8-THF, 500 MHz): δ 103.9. 25Mg-NMR (d8-THF, 300 MHz): δ 5.00. Crystallography Crystal data for C34H74AlCl6Mg2NO7Si2; M r=953.42; Monoclinic; space group P21/c; a=11.6517(7) Å; b=13.6722(8) Å; c=32.490(2) Å; α=90°; β=93.9030(10)°; γ=90°; V=5163.8(5) Å3; Z=4; T=200(2) K; λ(Mo-Kα)=0.71073 Å; μ(Mo-Kα)=0.459 mm−1; d calc=1.226g cm−3; 52,464 reflections collected; 8,792 unique (R int=0.0386); giving R 1=0.0846, wR 2=0.2347 for 5610 data with [I>2σ(I)] and R 1=0.1228, wR 2=0.2713 for all 8,792 data. Residual electron density (e−.Å−3) max/min: 1.245/−0.493. An arbitrary sphere of data were collected on a colourless rod-like crystal, having approximate dimensions of 0.50 mm×0.27 mm×0.20 mm, on a Bruker Kappa X8-APEX-II diffractometer using a combination of ω- and ϕ-scans of 0.5° (ref. 23). Data were corrected for absorption and polarization effects and analysed for space group determination24. The structure was solved by direct methods and expanded routinely. The model was refined by full-matrix least-squares analysis of F2 against all reflections. All non-hydrogen atoms were refined with anisotropic thermal displacement parameters. Unless otherwise noted, hydrogen atoms were included in calculated positions. Thermal parameters for the hydrogens were tied to the isotropic thermal parameter of the atom to which they are bonded (1.5× for methyl, 1.2× for all others). Methyl groups on the ((Me)3Si)2AlNCl3 anion were found to be disordered. The methyl carbon positions were located at the ellipse locii that described the elongated thermal envelope. These half-occupancy carbon atoms were then refined with isotropic thermal parameters. Large thermal motion, bordering on disorder, was also observed in all THF moieties. It was decided that refinement, as an anisotropic model, led to an overall satisfactory structure. The disorder results in a lower-than-desired maximum angle for observed data as well as unusual thermal parameter relationships. Carbon atoms in the THF and methyl groups typically have thermal parameter envelopes greater than their neighbouring atoms, whereas the heavier atoms are more well-located and have smaller thermal displacement ellipsoids than their neighbouring carbon atoms. This pseudo-disorder results in a larger number of alerts, beyond those immediately addressed. However, the conclusions drawn by the structure are not affected by the thermal motion of peripheral atoms. Hydrogen atoms were all included in geometrically calculated positions. Electrochemistry All solutions were diluted to 0.4 M with regard to Mg. Cyclic voltammograms were obtained using a Solartron 1,287 potentiostat in a conventional 3-electrode cell with a Pt disk working electrode, a Mg wire reference electrode, and a Mg ribbon counter electrode at a scan rate of 0.025 V s−1. All Measurements were performed in an argon-filled glovebox. Battery Coin cells were built in an argon-filled glovebox using standard parts for 2032-type cells. Anodes comprised 100-μm-thick Mg foil (ESPI Metals). Cathodes were prepared by coating sulphur–carbon composite ink (61% elemental sulphur, 35% carbon black, 4% poly(tetrafluoroethylene) binder) on a porous carbon substrate. Cathode analysis The XPS analyses were carried out with a PHI5802 Multitechnique instrument using a monochromatic Al Kα source (1486.6 eV). Air-sensitive samples were transferred into the instrument under inert gas environments. Non-linear least squares curve-fitting was applied to selected HRES spectra (χ 2<1). NMR spectroscopy All NMR experiments were performed at magnetic field strengths of 11.7 (1H, 13C, 27Al) and 7.05 (25Mg, 33S) T corresponding to 1H resonance frequencies of 499.89 and 299.89 MHz, respectively, and at ambient temperature (~21 °C) using Varian Inova and UnityPlus spectrometers. The Varian Inova spectrometer was equipped with a 5 mm broadband probe to measure one-dimensional (1D) 1H, 13C, and 27Al spectra. The Varian UnityPlus spectrometer was equipped with a 10 mm broadband probe to measure 1D 25Mg and 33S spectra. Usually 20 mg of sample were dissolved in 0.6 ml of THF-d8. Chemical shift values δ are given in p.p.m.. 1H and 13C spectra were referenced to residual solvent signals; δ=1.73 for 1H and δ=25.4 for 13C. 25Mg, 27Al, and 33S spectra were referenced indirectly using external reference standards (δ=0) MgCl2, Al(NO3)3, and saturated ammonium sulphate, respectively, in D2O. Author contributions J.M. devised the original concept, developed the experimental design and was responsible for the organometallic synthesis. H.S.K. performed the electrochemical experiments and supported the synthesis. G.D.A. synthesized HMDSMgCl. J.Z. and J.M. collected and analysed all the NMR data. A.O. collected and analysed the crystallographic data. J.G.N., H.S.K., A.E.R. and T.A. analysed the XPS spectra. W.B. and J.M. collected and analysed the mass spectroscopy data. J.M., H.S.K. and T.A. wrote the first draft of the manuscript; and all authors participated in manuscript revision. Additional information How to cite this article: Kim, H. S. et al. Structure and compatibility of a magnesium electrolyte with a sulphur cathode. Nat. Commun. 2:427 doi: 10.1038/ncomms1435 (2011).

                Author and article information

                Angew Chem Int Ed Engl
                Angew. Chem. Int. Ed. Engl
                Angewandte Chemie (International Ed. in English)
                WILEY-VCH Verlag (Weinheim )
                24 September 2012
                21 August 2012
                : 51
                : 39
                : 9780-9783
                Dr. R. Mohtadi, Dr. M. Matsui, Dr. T. S. Arthur Materials Research Department Toyota Research Institute of North America Ann Arbor, MI 48105 (USA)
                Division of Chemistry and Chemical Engineering California Institute of Technology (USA)
                Author notes

                The authors would like to thank Dr. T. Matsunaga and Dr. C. Bucur for the discussions. Funding was provided by Toyota Motor Engineering and Manufacturing North America Inc.

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                borohydrides, hydrogen, magnesium, electrochemistry


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