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Hexafluorophosphate-based Solutions for Mg Batteries and the Importance of Chlorides Ivgeni Shterenberg, Michael Salama, Yosef Gofer, and Doron Aurbach Langmuir, Just Accepted Manuscript • DOI: 10.1021/acs.langmuir.7b01609 • Publication Date (Web): 26 Jun 2017 Downloaded from http://pubs.acs.org on July 2, 2017

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Hexafluorophosphate-based Solutions for Mg Batteries and the Importance of Chlorides Ivgeni Shterenberg*, Michael Salama, Yosef Gofer, and Doron Aurbach

Department of Chemistry, Bar-Ilan University, Ramat Gan 5290002, Israel

ACS Paragon Plus Environment

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Abstract The selection of viable conventional magnesium salts in electrolyte solutions for Mg secondary batteries is very limited. Reversible magnesium deposition was only demonstrated with MgTFSI2, in ethereal solutions. A recent report has suggested that Mg can be reversibly deposited from a solution of Mg(PF6)2 in THF and acetonitrile. In this paper, we dispute that claim and show that PF6- anions passivate Mg anodes and completely inhibit any Mg deposition/dissolution process. We show that addition of chlorides suppresses the passivation phenomena and allows for reversible Mg deposition/dissolution processes to commence. The Mg deposits have been examined with elemental analysis, SEM and XRD measurements, depicting a highly oriented, preferential Mg growth. This study evaluates the feasibility of employing PF6-based electrolytes for Mg batteries, and exemplifies the aptitude of chlorides for suppressing passivation phenomena.


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Introduction Magnesium-based secondary batteries have been receiving a lot of attention lately, as reflected by the influx of recent publications.1-10 It is only natural, as Li-ion technology is reaching its limit and new energetic alternatives are sought to replace it. Although an operational model of a secondary battery has been demonstrated more than 15 years ago,11 Mg secondary cells have yet to reach a satisfactory performance level to become operational. Unlike the well-established Li systems, Mg currently has a very limited selection of cell components, i.e. salts, solvents and cathode materials.12-13 The limited selection of salts and solvents stems from the precondition of the Mg surface to be passivation-free. So far, only ethers were demonstrated to be sufficiently stable with Mg to enable reversible electro-deposition/dissolution of Mg ions. There is however a limited selection of Mg salts/compounds that are soluble in ethers. The most notable electrolytic solutions for Mg batteries are based on the organo-haloaluminate complexes,14 along with MACC solutions15-17 and Mg-carborane solutions.18 An interesting venture is to use Mg salts that have an equivalent in Li-based systems. Of those, only MgTFSI2 was demonstrated to reversibly deposit/dissolve Mg in pure ethers.19-21 Recently, a report on the performance of a new electrolytic solution, based on magnesium hexafluorophosphate, Mg(PF6)2,

in a 1:1 solution of THF and acetonitrile (AN), was

presented by Keyzer et al..22 There are several serious issues, unaddressed by the researchers, which put great doubt on whether metallic Mg can actually be electrochemically deposited or dissolved from these solutions. The most glaring issue is the presence of AN in the solution. AN is known to decompose at around -2.5V vs. SHE, very close to the reduction potential of Mg, making deposition of Mg from solutions containing


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AN highly implausible. There have been no reliable reports that suggest reversible Mg deposition or dissolution from AN-based solutions. Previously, Tran et al. presented a thorough report on AN-based solutions for Mg secondary batteries, displaying crucial and important results that cannot be ignored. They concluded that AN decomposes at -0.2V vs. Mg, and that researchers must be aware that this decomposition process can mimic Mg plating processes.23 Unfortunately, Keyzer et al. did not cite this important paper, and hardly addressed these concerns. In fact, they did not provide clear-cut evidence for Mg electro-deposition. Providing such evidence is straightforward; for instance one can perform XRD after electrodepositing Mg on an electrode that contains no Mg prior to deposition, thus verifying crystalline growth of Mg. The evidence provided was insufficient to prove, unequivocally, that metallic Mg plating did indeed ensue. This is particularly important as the experiments were performed in a solvent that is incompatible with Mg electrochemistry. Any evaluation of Mg(PF6)2, or of any other salt, must be performed in solvents that were proven to be stable with Mg, namely in ethers. In this study we assess PF6--based solutions in a system that had been proven to be completely chemically and electrochemically stable with Mg. The most obvious way of running such a study would be to use pure Mg(PF6)2 in ethereal solutions. Unfortunately pure Mg(PF6)2 is not available commercially, and its reported synthesis produces Mg that is tenaciously coordinated by six molecules of AN,22 which must be avoided. Therefore, we opted for a model system that simulates the behavior of Mg anodes in PF6-/ether conditions without using plain Mg(PF6)2 salt directly. For this purpose we made use of a solution of highly pure LiPF6 in pure and dry dimethoxyethane (DME). LiPF6 was chosen as it can be easily obtained in very high purity and dryness, and its electrochemical performance is well-established.24 Also, lithium’s redox potential is much lower than the operational


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voltages for the Mg deposition/dissolution processes. This ensures the avoidance of misinterpreting the electrochemical measurements. Experimental Solution preparation and handling. LiPF6 (99.99%), AgPF6 (99.99%), LiTFSI (99.95%), LiCl (99.99%), MgCl2 (99.99%), and dimethoxyethane (DME) (99.5%) were purchased from Sigma Aldrich. All solutions were prepared as follows: DME was added to predetermined amounts of LiPF6/LiTFSI with or without MgCl2/LiCl and stirred for 6h at 600C. The solutions were then filtered. All solution and sample preparations and electrochemical measurements were performed in a glovebox filled with Ar (99.999%). Electrochemical measurements and cell configuration. All electrochemical measurements were performed in sealed three-electrode flooded glass cells. Mg strips and foils were used as the counter and reference electrodes for all measurements. Platinum and stainless steel (SS) were used as the working electrodes. All measurements were performed with a PAR 273A potentiostat. Macro-reversibility measurements were performed as follows: a 1.5 µm layer (calculated, 1 C/cm2) of Mg was electrodeposited on an SS electrode at 0.5 mA/cm2. Next, 20% of the initial deposited material was electrodeposited/dissolved for 100 cycles. Finally, the remaining Mg was dissolved electrochemically. The reversibility was calculated with the following formula:

‫ܧ‬% =

݊ܳ − ܳ௜௡ + ܳ௘௫ ∗ 100 ݊ܳ

Where Q is the charge cycled, n is the number of cycles, Qin is charge of the initial Mg electrodeposition, and Qout is the residual Mg charge at the end of the experiment.


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Elemental analysis. The Mg:Li ratios in the electrodeposits were determined by a PerkinElmer atomic absorption spectrometer, AAnalyst400. Samples were dissolved in nitric acid and diluted with doubly distilled deionized water. Morphological and structural analysis. Surface images and elemental analysis were carried out with Magellan 400L and Helios 600 scanning electron microscope systems. XRD measurements were performed using a Rigaku TTRAX-III X-ray diffractometer (18 kW), operating with a Cu anode.


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Results and discussion In the following set of experiments we aimed to register, among the regular electrochemical responses, the potentials developing on the Mg counter electrode (CE) during the electrodeposition of Li on an inert working electrode (WE). This was done in a conventional three-electrode cell (Mg reference). In theory, under passivation-free conditions, the process on the CE during Li deposition should be the dissolution of Mg. Dissolution of Mg is expected to ensue at approx. 0V - 1V vs. Mg. No other species in the LiPF6/DME solution are oxidizable at this range, and therefore a compliance voltage around these values (0V - 1V vs. Mg) would strongly advocate the process of electrochemical Mg dissolution. During the galvanostatic deposition of Li on a Pt electrode from a LiPF6 (0.18M)/DME solution, using Mg as the RE and CE, the steady-state voltage on the Mg CE (vs. RE/WE) was measured to be 50V (±1V) vs. Mg. Such a high compliance voltage, suggesting very high interfacial resistance, is caused, most certainly, due to strong passivation of the Mg CE. For comparison, we performed a similar experiment with a LiTFSI/DME solution. With this solution, the steady-state voltage on the Mg CE did not exceed 0.5V vs. Mg during Li deposition. This value is in accordance with the voltage expected for the Mg oxidation process in TFSI-based solutions on electrodes that were not passivated.19 From this simple experiment, by observing the potential developed on the Mg CE, we can definitively state that Mg is strongly passivated in DME/PF6- solutions. Moreover, by comparing LiPF6 and LiTFSI solutions it can be deduced that the anion, PF6-, is responsible for this passivation. Next, in order to see if we can influence the passivation phenomenon, we added chlorides to the LiPF6/DME solutions. Chlorides are a very important component in most electrolytic solutions for Mg batteries. They are hypothesized to inhibit passivation of Mg, and have shown to improve the electrochemical performance in most cases.19, 25 We have added chlorides in the form of ultra-pure and dry MgCl2, which, of its own accord, is insoluble in


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DME. Nonetheless, in the presence of LiPF6, up to a Li:Mg ratio of 4:1, MgCl2 is soluble, akin to the solubility levels of MgCl2 in LiTFSI/DME solutions.19 A voltage profile for the galvanostatic deposition of Li/Mg from a LiPF6 (0.18M) + MgCl2 (0.045M) DME solution is presented in Figure 1a. The first point to note is that the steady-state voltage measured on the Mg CE during the process never exceeded 0.2V vs. Mg. This indicates that the anodic process in this case is electrochemical Mg dissolution, rather than solution decomposition. Remarkably, the addition of chlorides prevented the strong passivation that was observed in pure LiPF6/DME solutions. We have also performed similar experiments with higher LiPF6- concentrations (up to 0.5M) and obtained similar results. We have to note here that we do not have direct and solid evidence elucidating the mechanism by which chlorides ions suppress the Mg passivation. We suggest however, that chloride ions prevent the approach of PF6- anions to the Mg surface (and/or driving away their decomposition products) by preferentially and persistently adsorbing onto its surface and creating a barrier that wanes the reaction between the Mg metal and the PF6- anions. Obviously, this hypothesized barrier is not inhibiting magnesium ions to approach to the electrode’s surface. During the galvanostatic process, the deposition overvoltage on the Pt WE drops slowly from -1.3V to -0.4V vs. Mg. Li deposition is usually observed at -1.2V – (-1V) vs. Mg,19 in non-aqueous solutions. Hence, a voltage of -0.4V strongly advocates Mg deposition processes. Cyclic voltammetry (CV) that was performed on a fresh Pt electrode after the galvanostatic procedure, revealed a characteristic reversible performance of the Mg deposition/dissolution processes (Fig 1b) with about 90% micro-cycling efficiency. At this point we would like to note that, as will be elaborated in the next section, the electrodeposition of metallic Mg has been verified via multiple techniques. The initial galvanostatic reduction process, probably associated with Li and/or Mg deposition, as well as the reduction of contaminants, has enabled reversible deposition/dissolution processes


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of Mg with relatively high columbic efficiencies. A CV measurement performed prior to the galvanostatic processes showed only the deposition/dissolution of Li and no Mg related processes (Fig 1b inset). This behavior is not uncommon, and can be attributed to “solution conditioning”, a phenomenon that has been demonstrated previously in non-Grignard solutions.13, 16 The conditioning process consists of passing faradaic currents, during CV, galvanostatic cycling or, in this case, constant current experiments, which results in a significant improvement of the solution electrochemical performance. We still hold firmly the notion that the conditioning process simply “cleans” the solution from various detrimental contaminants. It is possible that the conditioning process may involve a different mechanism than what we proposed, since in different solutions formulations, different conditioning mechanisms may be dominant, while others are of minor influence. Several reports have suggested other explanations for the conditioning process.16-17, 26 They were however all related to MACC solutions and Mg:Al speciation. Hence, they are less relevant to the present study. This is the first time, to our knowledge, that reversible Mg deposition/dissolution processes on an inert electrode from a PF6--based solution were truly demonstrated.


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Figure 1. (a) Voltage profile of Li/Mg deposition on a Pt WE from a LiPF6 (0.18M) + MgCl2 (0.045M) DME solution. The average voltage measured on the CE during the deposition process was 0.2V vs. Mg. (b) CV performed on a fresh Pt WE after deposition at 25 mV/sec. Inset shows the CV prior to the galvanostatic process. Mg foil was used as the CE and RE for both measurements.


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The ratio between Mg and Li in the solution after the galvanostatic conditioning process was determined by elemental analysis (atomic absorption). For that, 70C of charge was passed during the galvanostatic procedure: the theoretical amount needed to consume all of the Li in the solution and replace it with Mg. Surprisingly, the concentration of Mg did not change during the galvanostatic conditioning process and remained 0.045M. The concentration of Li dropped only slightly, from 0.18M to 0.16M. In order to gain a better understanding of the Mg:Li speciation during the galvanostatic procedure, we created analogous experiments with solutions containing only Li, using LiCl instead of MgCl2 to formulate a LiPF6 (0.18M) + LiCl (0.09M) solution. The voltage measured on the Mg CE (vs. WE) during the galvanostatic deposition of Li from this Li-based solution (Fig 2a) was 0.2V, exemplifying once again the passivation-free conditions afforded by the addition of chlorides. The initially high overvoltage, of around -1V, drops slowly, until it stabilizes at a steady-state voltage of ca. -0.4V. Since initially no Mg ions were present in the solution, this voltage drop signifies the replacement of Li ions for Mg.

Atomic absorption (AA)

measurements were performed on the LiPF6 (0.18M) + LiCl (0.09M) solution prior to the galvanostatic process, after transferring 70C of charge (enough to consume two thirds of the present Li), and after transferring an additional 70C. Prior to the galvanostatic process, the Li concentration was measured to be 0.27M, as expected. After transferring 70C, the measured concentrations were 0.2M for Li and 0.023M for Mg. Considering that 2 mols of Li are necessary to generate 1 mol of Mg, there is 0.02 mol of Li that is unaccounted for, similar to the elemental analysis results obtained with the previous solution containing the LiPF6 (0.18M) + MgCl2 (0.045M) initial formulation. This slight loss of Li is most likely due to parasitic reactions occurring near the Mg counter electrode during the deposition process, such as Mg corrosion and possibly the consumption of contaminants in the solution. This would be in line with our firm belief that the conditioning process is first and foremost the


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elimination of reactive contaminants from the electrolytic solution. After passing an additional 70C (140C in total), elemental analysis showed that both Li and Mg maintained their previous concentrations of 0.2M and 0.023M respectively. At this stage, Li is no longer being converted to Mg, as Mg deposition and dissolution dominate the cathodic and anodic processes with high faradaic efficiency. There is also no loss of either Li or Mg, indicating the end of the loss of charge to parasitic reactions. This newly formulated LiPF6 (0.2M) + MgCl2 (0.023M) solution exhibits good electrochemical performance, depicted by the CV experiment in Figure 2b. In this case we have managed to prepare a PF6--based solution that can reversibly deposit/dissolve Mg from a solution containing only Li, by partial electrochemical replacement of Li ions with Mg. Unfortunately, it is not possible to replace all the Li and formulate a Mg-only solution under these conditions, since the redox potential for Li is higher than for Mg. We have tried the same method with an AgPF6/MgCl2 solution, unfortunately MgCl2 was found to be insoluble in AgPF6/ether solutions.


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Figure 2. (a) Voltage profile of Li/Mg deposition on a Pt WE from a LiPF6 (0.18M) + LiCl (0.09M) DME solution. The measured average voltage on the CE during the deposition process was 0.2V vs. Mg. (b) CV performed on a fresh Pt WE after deposition at 25 mV/sec. Mg foil was used as the CE and RE for both measurements.


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For a deeper insight into the electrochemical performance of Mg in PF6--based solutions, macro-cycling experiments were performed on conditioned LiPF6 (0.16M) + MgCl2 (0.045) solutions. In these experiments, a charge density of 1 C/cm2 at 0.5 mA/cm2 worth of Mg (equivalent to a 1.5 µm thick layer) was deposited onto pure SS electrodes, followed by a galvanostatic cycling of 20% depth of discharge (DOD) at 0.5 mA/cm2 for 100 cycles (Fig 3). The Mg deposition/dissolution processes in these solutions were complete for only 60-70 cycles, as indicated by the large oxidation over-voltage spikes. The calculated reversibility value (see experimental) for the first 68 reversible cycles is 93.8%. The steady-state potential for Mg deposition is retained at approx. -0.28V vs. Mg throughout the entire 100 cycles, excluding the initial deposition spikes at the beginning of each cycle. The potential for Mg dissolution was also constant, at 0.05V vs. Mg. These results differ from the ones obtained for TFSI-based solutions that showed ever increasing deposition/dissolution potentials throughout cycling. From these experiments we get a clear indication for the lack of passivation or surface film formation that deteriorates the reversibility of the Mg deposition/dissolution processes. However, the low cycling efficiency obtained suggests that either electrochemical parasitic reactions take place, which consume part of the charge, or chemical, corrosion-like reactions lead to material loss. Alternatively, poor adherence of the newly deposited magnesium may lead to material loss by mechanical means. The second possibility is less probable since no precipitates are observed in these experiments.


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Figure 3. (a) Galvanostatic cycling, 20% DOD, at 0.5 mA/cm2 using a conditioned LiPF6 (0.16M) + MgCl2 (0.045M) DME solution. (b) Enlarged image of the first seven cycles. (C) Enlarged image of cycles 61-75. Mg foil was used as the CE and RE. SS was used as the WE.


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SEM micrographs of a nominally 1.5 µm thick (calculated based on charge capacity and Mg density) electrodeposited Mg from a LiPF6 + MgCl2 solution, are presented in Figure 4. The electro-deposition was performed on SS at current densities of 0.05 mA/cm2 (Fig 4a-b), 0.5 mA/cm2 (Fig 4c-d), and 1 mA/cm2 (Fig 4e-s). AA measurements made on the deposits unequivocally proved that the vast majority of the initial deposit is indeed metallic Mg. A very small Li concentration was detected in the metallic deposit, at a Mg:Li 66:1 ratio. The SEM micrographs show a very uniform crystalline morphology with no clear signs of surface films that would indicate a massive reaction between the Mg surface and the solution components. Yet, very thin and compact passivation films would be, obviously, impossible to detect with SEM. At lower current densities, the deposits obtained were more closely packed with less exposed edges, but otherwise the morphology looks very similar at all current densities up to 1 mA/cm2 and charge density of 1 C/cm2. Interestingly, the Mg deposits are very uniformly distributed across all submerged surface areas, including the back of the electrodes, showing that the throwing power at 0.5 mA/cm2 and 1 mA/cm2 current densities is very high. At 0.05 mA/cm2, the nucleation seems much more selective and only partial coverage was observed. The morphology of the deposited Mg does not seem to resemble the typical Mg hexagonal crystallites as obtained in many previous studies.27-29 It appears that the Mg, in this case, deposits as micro-crystals with preferential orientation resembling layered formation. To test whether the specific growth orientation develops slowly during the deposition processes or is initiated since its inception, we examined thinner films of Mg deposits. SEM image of 0.01 C/cm2 worth of Mg (equivalent to a 15 nm thick Mg layer), deposited on a SS electrode, shows uniform nucleation of Mg across the surface, with no observed preferential orientation (Fig 4g). At a 150 nm thickness, 0.1 C/cm2 (Fig 4h), the morphology begins to resemble the layered morphology from Figures 4a-f.


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Figure 4. SEM images of Mg, deposited on pure SS electrodes from a LiPF6 (0.16M) + MgCl2 (0.045M) (DME) solution at: (a-b) 0.05 mA/cm2 1 C/cm2, (c-d) 0.5 mA/cm2 1 C/cm2, and (e-f) 1 mA/cm2 1 C/cm2. (g) 0.5 mA/cm2 0.01 C/cm2. (h) 0.5 mA/cm2 0.1 C/cm2.

To explore the morphology further, XRD measurements were performed on the Mg deposits and are presented in Figure 5. The XRD diffraction pattern of the deposited Mg is compared to hexagonal Mg pattern to reveal several key differences, most notably, the absence of the (002) plane at diffraction angle 34 2Θ. The absence of this plane


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corroborates the SEM observation that a specific orientated growth of Mg is preferable in these electrolytic solutions, under these specific electrodeposition conditions. As for the extent of this study we cannot give definitive answers explaining the preferential Mg growth in these electrolytic solutions. There may be numerous causes for this phenomenon. The most glaring is the minor presence of Li in the deposits. It may be that some, diluted Mg:Li alloy formation may initiate specific Mg growth. To properly address these questions, a thorough in-depth study on the crystallography of the Mg and Mg:Li alloy deposits from these solutions is required, Such study is beyond the scope of this work.

Figure 5. XRD pattern of deposited Mg from a LiPF6 (0.16M) + MgCl2 (0.045M) (DME), compared to a hexagonal Mg pattern, PDF Card No.: 00-035-0821. Radiation source wavelength 1.54056 Å.

Next, we examined the Mg deposit’s morphology after macro-cycling at 10% and 20% DOD (see above). A stark change in the morphology of the deposits is observed after 10 cycles (Fig 6a,b). The deposited Mg surface shows crevices and pits (marked with red arrows) which resembles corrosion morphologies. Although Mg deposits evenly during the first


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electrodeposition

cycle,

continuous

cycling

forms

an

ever

increasing

uneven

distribution/packing of the Mg deposits. After 100 cycles (ending with oxidation), a decent amount of Mg deposits is still clearly visible on the surface of the electrode even though at this stage the Mg is no longer electrochemically oxidizable (Fig 6c,d). When a single full cycle of Mg deposition/dissolution at similar conditions (1 C/cm2, 0.5 mA/cm2) is performed, some Mg still remains on the electrode (Fig 6e,f), although at a much smaller quantities compared to the long term experiment (100 cycles). Some magnesium, at least, therefore becomes “trapped”, electronically isolated as a function of cycling, leading to low cycling efficiency. EDAX measurements were performed on all samples and revealed similar results: only Mg and oxygen are present in any meaningful amounts. F is present at 1-4% across all samples. The difference in the F concentration does not follow any distinguishable trend. It is important to mention however that even 1% of fluorine on the magnesium surface is potentially enough to fully passivate the Mg surface, or form electronically isolated islands of metallic Mg, in the extreme case where MgF2 is the sole or dominant species and is formed in a compact, pinhole free manner. Most likely, a compact layer of MgF2 does not form in the current case since the Mg electrode isn’t passivated. The presence of F can also be attributed to remaining PF6- anions or their derivatives.


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Figure 6. SEM images of Mg, deposited on pure SS electrodes from a LiPF6 (0.16M) + MgCl2 (0.045M) (DME) solution. (a-b) After 10 cycles at 20% DOD (red arrows indicate crevices). (c-d) After 100 cycles at 20% DOD. (e) After one full deposition/dissolution cycle of 1C/cm2. (f) Mg EDAX mapping of (e).

Summary and Conclusions In this work, we have evaluated the feasibility of using PF6--based electrolytic solutions for Mg batteries. Since Mg(PF6)2 is not available for purchase, and is very difficult to obtain as pure material by synthesis, we used a model LiPF6/DME electrolytic solution. By observing the potential on the Mg CE, which reaches up to 50V during the Li deposition process, we can safely ascertain that the Mg anode is fully passivated in PF6--based solutions. We


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discovered that by adding chlorides, through either MgCl2 or LiCl, the passivation on Mg no longer develops and reversible Mg deposition/dissolution processes commence. This was clearly demonstrated with CV experiments performed on inert Pt electrodes. We have verified, beyond doubt, through AA, EDAX, HRSEM, and XRD, that practically pure Mg is being deposited in uniform, crystalline, pure form with strong preferential orientation. This work shows that pure PF6--based solutions fully passivate Mg. As evidenced from the SEM micrographs, the main mechanisms leading to the low cycling efficiency rely on corrosion of the metallic magnesium with the PF6 anion. It leads to both electronically isolated metallic magnesium and to some direct metal loss by corrosion. This study also exemplifies the critical role of chlorides in electrolyte solutions for Mg batteries. Previously, chlorides were shown to enhance the electrochemical performance. In our case, the addition of chlorides completely overhauls the electrochemical behavior of Mg anodes, transforming them from fully passivated to reversible. It is perhaps the strongest case presented yet for the ability of chlorides to remove and inhibit passivation of Mg anodes.


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References 1. Salama, M.; Shterenberg, I.; Gizbar, H.; Eliaz, N. N.; Kosa, M.; Keinan-Adamsky, K.; Afri, M.; Shimon, L. J. W.; Gottlieb, H. E.; Major, D. T.; Gofer, Y.; Aurbach, D. Unique Behavior of Dimethoxyethane (DME)/Mg(N(SO2CF3)2)2 Solutions. J. Phys. Chem. C 2016, 120 (35), 1958619594. 2. Wetzel, D. J.; Malone, M. A.; Gewirth, A. A.; Nuzzo, R. G. Anisotropic Mg Electrodeposition and Alloying with Ag-based Anodes from Non-Coordinating Mixed-Metal Borohydride Electrolytes for Mg Hybrid Batteries. Electrochimica Acta 2017, 229, 112-120. 3. Truong, Q. D.; Kempaiah Devaraju, M.; Nguyen, D. N.; Gambe, Y.; Nayuki, K.; Sasaki, Y.; Tran, P. D.; Honma, I. Disulfide-Bridged (Mo3S11) Cluster Polymer: Molecular Dynamics and Application as Electrode Material for a Rechargeable Magnesium Battery. Nano Letters 2016, 16 (9), 5829-5835. 4. Yu, Y.; Baskin, A.; Valero-Vidal, C.; Hahn, N.; Zavadil, K. R.; Eichhorn, B. W.; Prendergast, D.; Crumlin, E. J. In Fundamental investigations of magnesium electrode/electrolyte interface using operando ambient pressure x-ray photoelectron spectroscopy, 2017; American Chemical Society, pp ENFL-283. 5. Mandai, T.; Akita, Y.; Yagi, S.; Egashira, M.; Munakata, H.; Kanamura, K. A key concept of utilization of both non-Grignard magnesium chloride and imide salts for rechargeable Mg battery electrolytes. J. Mater. Chem. A 2017, 5 (7), 3152-3156. 6. Hou, T.; Monroe, C. W. In Synthesis of a magnesium battery electrolyte without need for conditioning or promoters, 2017; American Chemical Society, pp ENFL-282. 7. Sun, X.; Bonnick, P.; Nazar, L. F. Layered TiS2 Positive Electrode for Mg Batteries. ACS Energy Lett. 2016, 1 (1), 297-301. 8. Cho, J.-H.; Ha, J. H.; Lee, S.-H.; Cho, B. W.; Chung, K. Y.; Na, B.-K.; Kim, K.-B.; Oh, S. H. Effect of 1-allyl-1-methylpyrrolidinium chloride addition to ethylmagnesium bromide electrolyte on a rechargeable magnesium battery. Electrochimica Acta 2017, 231, 379-385. 9. Tian, H.; Gao, T.; Li, X.; Wang, X.; Luo, C.; Fan, X.; Yang, C.; Suo, L.; Ma, Z.; Han, W.; Wang, C. High power rechargeable magnesium/iodine battery chemistry. Nature Communications 2017, 8, 14083. 10. Ma, Z.; Kar, M.; Xiao, C.; Forsyth, M.; MacFarlane, D. R. Electrochemical cycling of Mg in Mg[TFSI]2/tetraglyme electrolytes. Electrochemistry Communications 2017, 78, 29-32. 11. Aurbach, D.; Lu, Z.; Schechter, A.; Gofer, Y.; Gizbar, H.; Turgeman, R.; Cohen, Y.; Moshkovich, M.; Levi, E. Prototype systems for rechargeable magnesium batteries. Nature 2000, 407 (6805), 724-727. 12. Canepa, P.; Sai Gautam, G.; Hannah, D. C.; Malik, R.; Liu, M.; Gallagher, K. G.; Persson, K. A.; Ceder, G. Odyssey of Multivalent Cathode Materials: Open Questions and Future Challenges. Chemical Reviews 2017, 117 (5), 4287-4341. 13. Shterenberg, I.; Salama, M.; Gofer, Y.; Levi, E.; Aurbach, D. The challenge of developing rechargeable magnesium batteries. MRS Bull. 2014, 39 (5), 453-460. 14. 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 (8), 2265-2279. 15. Doe, R. E.; Han, R.; Hwang, J.; Gmitter, A. J.; Shterenberg, I.; Yoo, H. D.; Pour, N.; Aurbach, D. Novel, electrolyte solutions comprising fully inorganic salts with high anodic stability for rechargeable magnesium batteries. Chem Commun (Camb) 2014, 50 (2), 243-5.


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