Article pubs.acs.org/Langmuir
Electrochemical and Spectroscopic Analysis of Mg2+ Intercalation into Thin Film Electrodes of Layered Oxides: V2O5 and MoO3 Gregory Gershinsky,†,‡ Hyun Deog Yoo,†,‡ Yosef Gofer,† and Doron Aurbach*,† †
Department of Chemistry and Bar-Ilan Institute of Nanotechnology and Advanced Materials, Bar-Ilan University, Ramat-Gan, Israel 52900 S Supporting Information *
ABSTRACT: Electrochemical, surface, and structural studies related to rechargeable Mg batteries were carried out with monolithic thin-film cathodes comprising layered V2O5 and MoO3. The reversible intercalation reactions of these electrodes with Mg ion in nonaqueous Mg salt solutions were explored using a variety of analytical tools. These included slow-scan rate cyclic voltammetry (SSCV), chrono-potentiometry (galvanostatic cycling), Raman and photoelectron spectroscopies, high-resolution microscopy, and XRD. The V2O5 electrodes exhibited reversible Mg-ion intercalation at capacities around 150−180 mAh g−1 with 100% efficiency. A capacity of 220 mAh g−1 at >95% efficiency was obtained with MoO3 electrodes. By applying the electrochemical driving force sufficiently slowly it was possible to measure the electrodes at equilibrium conditions and verify by spectroscopy, microscopy, and diffractometry that these electrodes undergo fully reversible structural changes upon Mg-ion insertion/deinsertion cycling.
1. INTRODUCTION R&D of rechargeable Li-ion batteries can be considered as one of the most impressive and important successes of modern electrochemistry. During the last two decades, Li and Li-ion batteries are conquering the markets of electrochemical power sources. Most of the mobile equipment today is powered by Liion batteries. Their success pushes Li-ion battery technology to the challenge of electromobility. Despite their great commercial success, lithium-based batteries have some drawbacks associated with safety issues, high cost, and resource scarcity. For long-term support of portable energy needs, load-leveling applications (e.g., storage and conversion of sustainable energy), and even electrochemical propulsion, there is a strong interest to develop power sources based on other, more abundant elements than lithium. Indeed, post-lithium systems are these days under serious scrutiny by an increasingly large number of research groups and companies. In spite of its inferiority vs Li anode in terms of specific (mass) capacity and higher redox potential (i.e., a drawback for high energy density) magnesium, as a bivalent metal, is expected to possess superior volumetric capacity (3833 mAh cc−1) compared to lithium (2046 mAh cc−1). Mg is also the fifth most abundant element on the earth’s crust. However, development of magnesium rechargeable batteries has been hampered by a variety of intrinsic problems related to the use of magnesium metal anodes and magnesium-ion intercalation cathodes.1−3 Passivation of magnesium metal in all conventional nonaqueous solutions is a serious problem. In contrast to Li electrodes, where surface films are Li-ion conductors, when surface films comprising ionic compounds are formed on Mg metal electrodes, they do not allow Mg-ion migration through © 2013 American Chemical Society
them. Hence, reversible Mg electrodes exist only in passivation-free situations.4,5 Consequently, Mg electrodes are reversible only in unique, complex solutions that were described well in the literature.6 Unfortunately, many of these complex solutions have a limited anodic stability and poor compatibility with many potential cathode materials based on transition metal oxides. Finally, in most of the transition metal oxides and sulfides which can serve as excellent Li-ion insertion electrodes, Mg-ion intercalation is very slow or even impossible. Notwithstanding these difficulties, prototype magnesium rechargeable batteries were successfully developed in 2000, utilizing solutions comprising Mg organo−halo−aluminate complexes and ether solvents (THF, glymes) that are compatible with both magnesium metal anodes and Mo6S8 cathodes.7 Since then, more profound understanding has been achieved of several types of complex solutions in which Mg electrodes are reversible.8 Several types of Chevrel phase (MgxMo6T8, T = S, Se) cathodes for rechargeable Mg batteries have been thoroughly explored.9−11 Other potential cathode materials for these systems are being intensively explored,12−20 and new concepts and strategies are being developed for advanced magnesium batteries.21−25 Various candidates were suggested in recent years for magnesium battery cathode materials with capacity and redox voltage higher than those of Chevrel phases (120 mAh/g; 1.1 V vs Mg).1,2,15,19,26−28 Among them, layered transition metal oxides are promising because they may possess the structural Received: June 25, 2013 Revised: July 25, 2013 Published: July 26, 2013 10964
dx.doi.org/10.1021/la402391f | Langmuir 2013, 29, 10964−10972
Langmuir
Article
flexibility29 needed in order to accommodate the expected severe structural deformations upon insertion of bivalent ions such as Mg2+ of high charge density. Specifically, divanadium pentoxide (V2O5) is known as a high-voltage/capacity cathode material that may be able to insert Mg ion reversibly, and it has been studied very intensively as a cathode material for Li-ion batteries.30 V2O5 electrodes were also explored in connection with rechargeable Mg batteries.2 These studies indicated the possibility that this compound may insert Mg ion in conventional nonaqueous Mg salt solutions with some degree of reversibility. However, none of these studies demonstrated performance that would make it possible to consider V2O5 as a cathode material for rechargeable Mg batteries. Too low capacity and slow rates of Mg-ion insertion into this host material could be attributed to strong Coulombic interactions of the bivalent cations within the host lattice. Such interactions can be mitigated by utilizing appropriate complexing ligands, tenaciously bonded to the divalent Mg ion, thus acting as partial charge screening. An example is a water molecule that is known to strongly hydrate Mg ion. It was indeed found that the presence of a trace amount water in nonaqueous Mg salt solution increases the rate and capacity of Mg-ion insertion into V2O5 electrodes.31 For the same reason, V2O5 aerogels exhibited the highest rate and capacity of Mg-ion insertion found until now, as the layers are intrinsically hydrated.32−36 Unfortunately, this approach has strong drawbacks: it leads to a large effective radius of the intercalated moieties, which may lead to excessive structural deformations of the host. Also, wet solutions are not compatible with magnesium anodes. A different and well-studied approach for mitigating slow solid-state diffusion difficulties is to decrease the total diffusion path by nanosizing the cathode material.37 V2O5 aerogels exhibit relatively high rate capability of Mg-ion insertion due also to the shortened diffusion path in their nanostructure (in addition to the above mention shielding effect due to the presence of water molecules). These previous studies left many unanswered questions because it was impossible to separate shielding and size effects. An excellent platform for studying the mechanism of Mg intercalation into V2O5 and MoO3 materials is the use of highly pure thin films of the said oxides on an inert (Pt) substrate. Using such electrodes allows one to investigate the intrinsic behavior of the host materials by eliminating many complexities in the electrochemical and spectral response of Mg insertion electrodes associated with composite electrode formulations and variations (carbon, binder, substrate, etc.). In this work we prepared highly pure thin-film V2O5 and MoO3 electrodes of nanoscale thickness and particle size. Their interactions with Mg ions were studied by electrochemical, spectroscopic, microscopic, and XRD techniques. Electrochemical analysis was carried out under near-equilibrium conditions. The corresponding phase changes were studied with spectroscopic tools.
Surface morphological changes and elemental analysis were carried out by HR-SEM with EDS (Magellan 400L, FEI Co.). All electrodes were washed twice with AN and dried under Ar gas after electrochemical cycling. Phase changes were monitored by X-ray diffraction (XRD, AXS D8 with Cu Kα, Bruker Co.) and Raman spectroscopy (inVia Raman microscope, Renishaw Co.) for electrodes at different stages of intercalation and cycling. This structural information was correlated with surface analysis by ex-situ X-ray photoelectron spectroscopy (XPS, axis HS, Kratos Analytical Co.). These measurements included also depth profiling by Ar+ sputtering and enabled as to follow changes in the binding energies of the various atoms as a function of the Mg intercalation level. Either 0.1 M magnesium bis(trifluoromethane sulfonyl)imide (MgTFSI2, Strem Co., 97%) or 0.5 M magnesium perchlorate (Mg(ClO4)2, Aldrich Co., ACS reagent grade) in acetonitrile (AN, CH3CN, Aldrich Co., anhydrous, 99.8%) was used as the Mg2+containing electrolyte solution. Salts were used after vacuum drying at 180 °C. The water contents of the electrolyte solutions were 32 and 212 ppm for 0.1 M MgTFSI2/AN and 0.5 M Mg(ClO4)2/AN, respectively (monitored by Karl Fischer, 652 KF coulometer, Metrohm). Electrochemical cycling of the Mg2+ solutions was done either galvanostatically or potentiodynamically by a combination of slow-scan rate voltage sweeping and constant-voltage steps (VMP-2 potentiostat, Bio-Logic Co.). For potentiodynamic analysis, postmeasurement charge integration was used for calculation of the attained intercalation level and specific capacity. For all electrochemical measurements, a sealed flooded cell under argon gas was used to observe the visual changes during electrochemical measurements. Quasi-reference electrodes were activated carbon (AC) cloth attached to a Pt wire because the electrode potential of such electrodes is known to be stable in AN-based electrolytes.38 The immersion potential of the AC electrodes was measured to be 0.120 V vs ferrocene/ferrocenium−Pt electrodes, calculated as 2.38 V vs Mg2+/ Mg and 3.05 V vs Li+/Li. Mg electrodes cannot behave reversibly in any conventional nonaqueous solutions due to passivation phenomena that totally block them, and therefore, they cannot serve as counter electrodes.34 Ion adsorption/desorption processes in the electric double layer (EDL) are usually fully reversible. It was found possible to prepare counter electrodes for the present studies based on high surface area, AC cloth electrodes whose total surface area is high enough to allow a full charge balance for the Mg intercalation process that we explored (Figure S1, Supporting Information).
3. RESULTS AND DISCUSSION 3.1. Mg2+ Ion Intercalation into V2O5 Thin-Film Electrode (1 e− + 0.5 Mg2+ per V2O5). Prior to the study of Mg intercalation into V2O5, the well-documented electrochemical behavior of the thin-film V2O5 electrodes with lithium served as benchmarking. In fact, the electrochemical response of the V2O5 film electrodes with Li+ ions was also used as an indicator for the correct deposition procedure as well as for mass load verification. The annealed films behaved according to expectations based on literature data, with two redox peaks at exactly the expected potentials (3.2 and 3.4 V vs Li+/Li) and high reversibility (Figure S2, Supporting Information). Since the V2O5 films are very thin, the electrochemical response is very fast; thus, we could measure cyclic voltammograms (CV) at scan rates higher than 0.5 mV s−1 with no peak widening, no voltage shifting, and full charge reversibility. As can be seen in Figure S3, Supporting Information, cyclic voltammograms of V2O5 thin-film electrodes exhibit highly reversible (about 100% of Coulombic efficiency) electrochemical insertion and deinsertion of magnesium ions at 0.1 mV s−1. Insertion takes place at 2.3 V vs Mg2+/Mg (about 3.0 V vs Li+/Li). The half-wave potential for Mg2+ insertion is about 2.6 V vs Mg2+/Mg (3.3 V vs Li+/Li, Figure S3, Supporting
2. EXPERIMENTAL METHODS V2O5 films were vacuum deposited thermally from prefused V2O5 powder (Sigma Aldrich). Film thickness was usually around 200 nm. Electrodes were then annealed at 415 °C for 4.5 h in air to obtain the desired crystalline phase. MoO3 films were electrodeposited using 50 mM molybdic acid aqueous solution, followed by annealing at 350 °C for 3 h (the films thickness was around 100 nm). Qualitative analysis was performed by XRD and Raman spectroscopy, which have unique and unambiguous response for the various crystalline and amorphous phases. 10965
dx.doi.org/10.1021/la402391f | Langmuir 2013, 29, 10964−10972
Langmuir
Article
Figure 1. HR-SEM images of the V2O5 thin-film electrodes: (a) pristine, (b) magnesiated (Mg0.5V2O5), and (c) demagnesiated Mg(ClO4)2/AN solution. EDS patterns are presented in the images (overlapped), and visual photographs of the electrodes are presented in the insets.
Information), which is very close to the potentials of Li+ insertion (3.2 and 3.4 V vs Li+/Li, Figure S2, Supporting Information). In Figure 1 a−c SEM imaging shows the morphological effect of the electrochemical process on these electrodes. Each image includes an inset that shows the electrode’s color. The relevant EDS spectrum is printed at the bottom of each image. Figure 1a shows a SEM micrograph of deposited and annealed V2O5 thin film. Primary particles of about 20−50 nm are uniformly deposited over the entire surface. The EDS spectrum indicates the presence of V and O elements (only qualitatively because of superposition of V(L1) and O(K) peaks). Figure 1b shows a SEM micrograph of an electrode that underwent Mg2+ insertion process: 1 e− and 0.5 Mg2+ per unit V2O5 magnesiation. Only minor morphological changes were observed. The presence of Mg is clearly identified by the EDS spectrum; however, due to the limitations of this technique it is impossible to distinguish between surface and bulk origin of the peaks (the analytical information in EDS can be obtained from depths as high as 1 μm and not carrying direct depth information). The change of the electrode color upon the electrochemical process from yellow to faint green also indicates the change in the oxidation state of V (see the inset to Figure 1a and 1b). After Mg deinsertion, with 98% Coulombic efficiency of the cyclic process, the morphology of the electrode resembles that of the pristine one (Figure 1c). A more rigorous observation (see the images in Figure S4, Supporting Information) reveals that the cyclic process forms cracks. As seen in Figure 1c, after a full cycle, no Mg can be detected by EDS and the electrode’s color returns to yellow. The electrochemical, visual, morphological, and spectral information provided in Figures S1, Supporting Information, and 1, related to a full cycle of these thin V2O5 electrodes, indicates that they undergo fully reversible, bulk Mg-ion insertion/deinsertion processes. By galvanostatic cathodic polarization of these electrodes in the Mg salt solution at low constant current density (0.5 μA cm−2) down to 2.2 V vs Mg (the cathodic potential limit for these experiments), a capacity of 150 mAh g−1 was achieved, corresponding to 1 e− and 0.5 Mg2+ per V2O5 unit (Figure 2). At these conditions, complete deintercalation of Mg ions occurs at about 2.4 V vs Mg2+/Mg with 97% Coulombic efficiency. The rate-determining step of magnesium insertion is the ion transport within the host (i.e., solid-state diffusion of Mg2+).34 In the layered transition metal oxides, the layers are weakly bound (e.g., by van der Waals forces), where guest ions can be
Figure 2. Typical galvanostatic titration curve of a V2O5 thin-film electrode in 0.1 M MgTFSI2 in AN (current density = 0.5 μA cm−2).
intercalated in between the oxygen-terminated slabs, while the transition metal ions provide electroneutrality by changing their oxidation state. The intercalation goes on as long as the layered oxide can sustain the structural expansion and the electronic structural changes. The high charge density Mg2+ ions induce slow insertion kinetics due to the strong electrostatic interactions with the host and the significant difficulty in maintaining local electroneutrality in the vicinity of the doubly charged, small ions. A simple, comparative way to learn about the kinetics of intercalation/deintercalation reactions is a comparison of intercalation or deintercalation overpotentials in the cyclic voltammogram. We compared the intercalation response of Li and Mg insertion with almost identical setups: the difference was only the cation in the solution, LiClO4 for Li and Mg(ClO4)2 for Mg (both in AN). We can assume that there are no electronic conductivity issues for the thin-film V2O5 electrodes because of their semiconducting nature39 and their being used as thin films on metallic (Pt) current collectors. In such experiments the slower the process the larger is the observed overpotential. In the Mg insertion into the thin film of V2O5, an overpotential of about 700 mV ([peak potential] − [middle potential between the peaks]) was observed at a scan rate of 100 μV s−1 (Figure S3, Supporting Information). For comparison, the overpotential for Li insertion into similar electrodes at 100 μV s−1 was around 25 mV (Figure S2, Supporting Information). The larger overpotential for Mg2+ insertion compared to that for Li+ is an indication of the slower kinetics of Mg2+ insertion processes. In order to understand these charge−discharge processes, we computed the differential capacity (dQ/dV) from the 10966
dx.doi.org/10.1021/la402391f | Langmuir 2013, 29, 10964−10972
Langmuir
Article
galvanostatic profile (Figure 3). Several (at least four) peaks at 2.35, 2.32, 2.30, and 2.29 V vs Mg2+/Mg are resolved during the
transport. We believe that due to the lower activation energies differences for the demagnesiation processes, the very low current galvanostatic process shows very weak patterns in the dQ/dV curve. (3) At both high and low current densities, higher overpotential is required to complete magnesiation compared with demagnesiation. Differentiation of the low-current galvanostatic curves proved to be a very sensitive tool to distinguish between multiple processes and gain insight to their comparative kinetic properties. dQ/dV analysis provided distinction between multiple processes and insight to their comparative kinetic properties. Apparently, magnesiation of V2O5 undergoes multiple insertion processes with different thermodynamic and kinetic properties. The different thermodynamics and kinetics may be attributed to different phase formations. Ten different phases have been postulated in a Mg−V2O5 phase diagram.2 However, the definite processes and their association with different phases or kinetic parameters are beyond the scope of this manuscript. As the magnesiation step was found to be slower, we modified the mode of cycling to combined cyclic voltammetry and constant potential steps at the cutoff potentials (for 2 h). Operating by this voltammetric−amperometric protocol, the obtained capacity drops initially from 180 mAh g−1, stabilizing at 150 mAh g−1 over the first 15 cycles, with ∼100% of Coulombic efficiency for the following cycles (Figure 4a). The constant potential step was essential for completion of the magnesiation process. During the potential step, the current decreased from 120 to virtually 0 mA g−1. On the other hand, demagnesiation is completed via a current decrease from 40 to
Figure 3. Typical differential capacity (dQ/dV) plot of the V2O5 thinfilm electrodes at a current density of (a) 0.5 and (b) 1 μA cm−2.
magnesiation curve in Figure 3a, while demagnesiation occurred as virtually a single process, with a very sharp peak at 2.42 V vs Mg2+/Mg. A closer look at this peak reveals some fine structure, with a medium step at 2.35 V, a deflection point at about 2.4 V, and a flat top peak (Figure 3a). When the current density doubles, the difference in the magnesiation/ demagnesiation potentials becomes larger (from 0.11 to 0.26 V) (Figure 3b). While the magnesiation process is still a multiple-step process, the demagnesiation process is now observed in a clear manner as also a multiple-step process with 4 new peaks appearing at 2.49, 2.54, 2.60, and 2.86 V vs Mg2+/ Mg. From this behavior, the following can be concluded. (1) There are at least four different stages for magnesiation and demagnesiation of V2O5 corresponding to the four resolved dQ/dV peaks. At an exceedingly slow rate, the faster processes, namely, the demagnesiation, proceeds as a virtually single continuous process, washing out most of the fine details. (2) At sufficiently low current density (0.5 μA cm−2), both processes are close to their thermodynamic equilibrium value. The processes take place between −0.4 and 0.0 V, and we observe strong response to the sluggish kinetics for both processes by increasing the rate of galvanostatic titration (1 μA cm−2), as the window of these electrochemical reactions widens drastically to from −0.4 to 0.2 V. It is also clearly observed that each process has its unique overpotential, especially identifiable during faster demagnesiation, presumably due to the different activation energy for the solid-state ionic
Figure 4. (a) Presentation of the cycle life (capacity and cycling efficiency vs cycle number) of the V2O5 thin-film electrodes by potentiodynamic mode of charge−discharge in 0.5 M Mg(ClO4)2 in AN (scan rate = 0.15 mV s−1), followed by potential steps for 2 h after reaching the cutoff potentials. (b) Presentation of the repeated corresponding cyclic voltammograms, related to the same experiment. 10967
dx.doi.org/10.1021/la402391f | Langmuir 2013, 29, 10964−10972
Langmuir
Article
0 mA g−1, as expected for the faster process. The CV curve (in Figure 4b) testifies again that the magnesiation process is appreciably slower than demagnesiation. Note that the two demagnesiation peaks are resolved even on the CV at −0.05 and 0.3 V vs Pt quasi reference (E[Pt quasi] = 2.5 V vs Mg2+/ Mg). Reversible structural changes in the V2O5 lattice during magnesiation and demagnesiation were monitored with XRD and Raman. The magnesiated and demagnesiated samples were prepared by the first cycles of magnesiation and demagnesiation, respectively. Surface chemistry changes were analyzed by XPS measurements, which revealed that as expected the vanadium oxidation state decreases upon magnesiation (V 2p peaks shift to lower BE) and returns to its original state after demagnesiation (Figure S5, Supporting Information). Additionally, it can be seen that some surface species accumulate on the V2O5 layer, e.g., ClO4−, and nitrogen-containing molecules, probably as a result of minor electrolyte decomposition (Figure S6, Supporting Information). The XRD pattern of the pristine annealed V2O5 films corresponds well to the database for an orthorhombic phase with preferentially oriented (010) plane (Figure 5a).40 For the Mg0.5V2O5 phase, the peaks’ intensity was attenuated and the main (010) peak was shifted by 0.1− 0.9° to lower angle and broadened (fwhm increased to 0.9° from, initially, 0.3°). This peak shift is similar to an earlier report on Li intercalation in thin-film vanadium pentoxide electrodes.41 The shift of the (010) peak can be attributed to the enlarged interlayer distance value, because the interlayers of
V2O5 are stacked in the (010) direction.2,42 The emergence of new peaks for magnesiated V2O5 at 2θ = 29.5° and 37.2°, unseen in the Li+ intercalation case,41 suggests formation of a new phase. After demagnesiation the patterns return to that of the pristine material. Raman spectroscopic results show similar trends (Figure 5b). The Raman signal of 1e− per unit magnesiated V2O5 shows a single, broadened peak at about 900 cm−1. The absence of other peaks may indicate that the bonding became severely disordered, which results in increased randomization of bond lengths and angles, thus creating a continuum of vibration modes and energies, and unresolved peaks. Indication of disordering was also obtained in the XRD measurements in the broadening of the main peak (Figure 5a). One of the pitfalls in Raman spectroscopy is the possibility of sample damage due to laser heating. Hence, in order to decrease the possibility of damage we carried out the measurement at 5% laser intensity. When the laser intensity increased to 50% and 100% clear indications of sample damage were observed, with evolution of a new and different phase, similar to the pristine material. This kind of laser-induced sample damage was also reported for δLi1V2O5.43 The structural reversibility of the material during intercalation/deintercalation cycles was substantiated by the reversion of the Raman spectra and XRD patterns. 3.2. Mg2+ Ion Intercalation into MoO3 Thin-Film Electrode. The same analytical approach as described above was also utilized to study the intercalation of magnesium ion into another layered compound, namely, MoO3. Orthorhombic α-MoO3 is a layered oxide material that has been shown to intercalate Mg2+ ion.2 Mg intercalation into this material is known to be plagued with extremely slow kinetics and large volume changes that often lead to severe capacity fading.44 The electrodeposited and annealed MoO3 thin films showed reversible color and EDS spectrum changes upon magnesiation and demagnesiation (Figure 6). Several reports showed that for reversible galvanostatic cycling of MoO3 electrodes an electrochemical conditioning must be pursued as a vital pretreatment; otherwise, the material suffers from severe irreversible structural change upon the first insertion of the guest ions (Figure S7, Supporting Information).44,45 Electrochemical conditioning is realized by initial intercalation with small galvanostatic current until the charge reaches a predetermined limiting value. As can be seen in Figure 7, after such conditioning, the voltage profile during slow galvanostatic titration (0.3 μA cm−2) shows about 220 mAh g−1 reversible magnesiation capacity at 1.7−2.8 V vs Mg2+/Mg, with 95% Coulombic reversibility. In the dQ/dV plot (inset of Figure 7), the magnesiation process shows at least two stages as sharp peaks at 1.80 and 1.74 V vs Mg2+/Mg. Demagnesiation proceeds as a single oxidation peak, centered at 2.15 V vs Mg2+/Mg, followed by a sluggish process at 2.2− 2.8 V vs Mg2+/Mg. Reversible specific capacity of 210 mAh g−1 was retained for the first 10 cycles with an average of 95% Coulombic efficiency (Figure S8, Supporting Information). The voltage difference between the magnesiation and the demagnesiation processes was found to be 0.2−0.4 V (Figure 7), 3−6 times larger than that for the V2O5 film electrodes. This indicates greater kinetic limitations. This is probably due both to the larger diameter of MoO3 disc-shaped flakes, ca. 500 nm (while the thickness of the flakes is ca. 100 nm), as evaluated from the SEM measurements, and to an intrinsically lower Mgion mobility within the MoO3 lattice. As Mg ions are inserted along the (020) layers and the MoO3 thin-film deposits are preferentially oriented along the b direction, the diameter of the
Figure 5. (a) XRD patterns and (b) Raman spectra of pristine, magnesiated (Mg0.5V2O5), and demagnesiated V2O5 thin-film electrodes. For the magnesiated electrode, Raman spectra were obtained by changing the laser intensity (5%, 50%, and 100%) at the same spot. 10968
dx.doi.org/10.1021/la402391f | Langmuir 2013, 29, 10964−10972
Langmuir
Article
Figure 6. Typical HR-SEM images of MoO3 thin-film electrodes: (a) pristine, (b) magnesiated (Mg0.5MoO3), and (c) demagnesiated. EDS patterns are presented (overlapped), and the visual electrodes’ photographs are presented in the insets.
Figure 7. Typical galvanostatic titration curve of a MoO3 thin-film electrode in 0.1 M MgTFSI2/AN solutions (current density = 0.3 μA cm−2). Corresponding dQ/dV plot is provided in the inset.
MoO3 discs (i.e., along the ac plane) is a critical parameter for the intercalation kinetics. On the other hand, the demagnesiation overvoltage is increased by the degree of demagnesiation. It may be due to the voltage-induced damage (i.e., structural change) upon magnesiation/demagnesiation.46 The white color of the pristine phase, in comparison with the blue for the magnesiated phase, is consistent with lower electronic conductivity for the demagnesiated MoO3.47,48 We suggest that Mg2+ intercalation is assisted by the increasing electronic conductivity of the material during magnesiation, realizing electron paths throughout the material as the degree of intercalation increases. During the deintercalation, formation of low-conductivity domains may slow down the process and even trap isolated partially magnesiated islands. The XRD patterns of the pristine MoO3 film electrodes correspond to the database of orthorhombic phase with preferential orientation of the (010) direction (Figure 8a, (020) and (040) planes at 2θ = 12.8° and 25.7°, respectively).45 The XRD pattern of the 1 e− per unit magnesiated phase (Mg0.5MoO3) exhibits new peaks at 2θ = 12.3°, 24.8°, and 37.5°. This may be proof for a two-phase reaction, as shown in the detailed view of the 11−14° region (Figure 8b). For the magnesiated MoO3, two peaks are observed at 12.3° and 12.7°. The peak at 12.7° may be associated with a pristine MoO3 phase, and we associate the peak at 12.3° with a Mg0.5MoO3 phase with increased interlayer distance, similar to the Li+ intercalation compound.49 The shift of the (020) peak can be directly associated with the interlayer distance, because the interlayers of MoO3 are stacked in the (010) direction.
Figure 8. (a) XRD patterns of pristine, magnesiated (Mg0.5MoO3), and demagnesiated MoO3 thin-film electrodes. (b) Detailed view at 2θ = 11−14°.
Certainly, more detailed structural analysis is needed to unambiguously prove this point. Upon demagnesiation, the peaks are restored to the original positions. The Raman spectrum for Mg0.5MoO3 shows completely different bands compared to the pristine film (Figure 9). New peaks are observed at 1020, 960, and 450 cm−1. The pristine MoO3 peaks can hardly be seen in the magnesiated, Mg0.5MoO3, spectrum. Upon demagnesiation, the pristine phase pattern is largely restored but some weak peaks at 1010 and 450 cm−1 that correspond to the magnesiated phase remain. This may be due to remaining islands of magnesiated MoO3, isolated by loss of electronic paths, because of the low electronic conductivity of the pure MoO3 phase. Alternatively, it may be due to the materials’ mechanical degradation (cracking) because of the large volumetric expansion/ 10969
dx.doi.org/10.1021/la402391f | Langmuir 2013, 29, 10964−10972
Langmuir
Article
as reported in the literature. The potentials we observed are well correlated with these literature values.2 In the case of V2O5, only one broad peak is observed for Mg2+ intercalation (Figure 10b) while two sharp peaks are clearly resolved for the Li+ intercalation (Figure 10a). Considering that the intercalation charge is the same for Li+ and Mg2+ insertion (∼150 mAh g−1), the broad peak for Mg2+ insertion can be attributed to the merging of several (two or more) peaks, as indicated in the dQ/ dV analysis. In the case MoO3, the deintercalation of Mg2+ shows two different peaks at 1.75 and 2.25 V vs Mg2+/Mg, and most of the deintercalation relates to the latter peak (Figure 10d). On the other hand, electrochemically reversible intercalation/deintercalation peaks are observed for the Li+ intercalation into MoO3 after an irreversible phase change at 2.8 V vs Li+/Li (i.e., 2.1 V vs Mg2+/Mg) upon the first lithiation (Figure 10c). For both materials, higher overpotentials are needed for the intercalation/deintercalation of Mg2+ ion than for Li+ ion.
Figure 9. Raman spectra of pristine, magnesiated (Mg0.5MoO3), and demagnesiated MoO3 thin-film electrodes.
contraction during cycling that inhibits complete demagnesiation. Considering the relative surface sensitivity of Raman spectroscopy compared to that of XRD, the finding may indicate that the Mg0.5MoO3 phase is preferably developed on the surface of the MoO3 particles while a portion of unreacted MoO3 phase remains in the core of the particles. Figure 10 shows the cyclic voltammograms of V2O5 and MoO3 thin-film electrodes upon Li+ or Mg2+ intercalation. For both Li+ and Mg2+ intercalation the average redox potential of V2O5 (ca. 2.6 V vs Mg2+/Mg) is higher than that of MoO3 (ca. 1.7 V vs Mg2+/Mg). Generally, the redox potential of a transition metal oxide is determined by the redox potential of the transition metal and its associated structural effect. The redox potential of Mo in MoO3 is lower than that of V in V2O5,
4. CONCLUSION This study of Mg-ion intercalation into thin-film electrodes comprising highly pure V2O5 and MoO3 showed that highly reversible Mg insertion/deinsertion is possible with these systems. V2O5 thin-film electrodes can be cycled over a potential range of 2.2−3.0 V vs Mg2+/Mg with a specific capacity of 150 mAh g−1 corresponding to 1e− transfer per unit. These electrodes can be cycled at 100% Coulombic efficiency and demonstrate very stable capacity upon cycling. Full chemical and structural reversibility was ascertained by Raman spectroscopy, XRD analysis, visual inspection, and EDS analysis. The differential capacity plots (dQ/dV vs V) for
Figure 10. Comparison of cyclic voltammograms for Li+ and Mg2+ intercalation processes into the V2O5 and MoO3 thin-film electrodes (v = 0.15 and 0.10 mV s−1 for Li+ and Mg2+ insertion, respectively): (a) Li+/V2O5, (b) Mg2+/V2O5, (c) Li+/MoO3, and (d) Mg2+/MoO3. 10970
dx.doi.org/10.1021/la402391f | Langmuir 2013, 29, 10964−10972
Langmuir
Article
(9) Aurbach, D.; Suresh, G. S.; Levi, E.; Mitelman, A.; Mizrahi, O.; Chusid, O.; Brunelli, M. Progress in Rechargeable Magnesium Battery Technology. Adv. Mater. 2007, 19, 4260−4267. (10) Mitelman, A.; Levi, E.; Lancri, E.; Aurbach, D. On the Mg trapping mechanism in Electrodes Comprising Chevrel Phase. ECS Trans. 2007, 3, 109−115. (11) Gershinsky, G.; Haik, O.; Salitra, G.; Grinblat, J.; Levi, E.; Nessim, G. D.; Zinigrad, E.; Aurbach, D. Ultra fast elemental synthesis of high yield copper Chevrel phase with high electrochemical performance. J. Solid State Chem. 2012, 188, 50−58. (12) Li, X. L.; Li, Y. D. MoS2 nanostructures: Synthesis and electrochemical Mg2+ intercalation. J. Phys. Chem. B 2004, 108, 13893−13900. (13) Nuli, Y.; Guo, Z. P.; Liu, H. K.; Yang, J. A new class of cathode materials for rechargeable magnesium batteries: Organosulfur compounds based on sulfur-sulfur bonds. Electrochem. Commun. 2007, 9, 1913−1917. (14) Nuli, Y. N.; Yang, J.; Wang, J. L.; Li, Y. Electrochemical Intercalation of Mg2+ in Magnesium Manganese Silicate and Its Application as High-Energy Rechargeable Magnesium Battery Cathode. J. Phys. Chem. C 2009, 113, 12594−12597. (15) NuLi, Y.; Yang, J.; Li, Y.; Wang, J. Mesoporous magnesium manganese silicate as cathode materials for rechargeable magnesium batteries. Chem. Commun. 2010, 46, 3794−3796. (16) Li, Y.; Nuli, Y. N.; Yang, J.; Yilinuer, T.; Wang, J. L. MgFeSiO4 prepared via a molten salt method as a new cathode material for rechargeable magnesium batteries. Chin. Sci. Bull. 2011, 56, 386−390. (17) NuLi, Y.; Zheng, Y.; Wang, Y.; Yang, J.; Wang, J. Electrochemical intercalation of Mg2+ in 3D hierarchically porous magnesium cobalt silicate and its application as an advanced cathode material in rechargeable magnesium batteries. J. Mater. Chem. 2011, 21, 12437− 12443. (18) NuLi, Y. N.; Zheng, Y. P.; Wang, F.; Yang, J.; Minett, A. I.; Wang, J. L.; Chen, J. MWNT/C/Mg1.03Mn0.97SiO4 hierarchical nanostructure for superior reversible magnesium ion storage. Electrochem. Commun. 2011, 13, 1143−1146. (19) Liang, Y.; Feng, R.; Yang, S.; Ma, H.; Liang, J.; Chen, J. Rechargeable Mg Batteries with Graphene-like MoS2 Cathode and Ultrasmall Mg Nanoparticle Anode. Adv. Mater. 2011, 23, 640−643. (20) Zheng, Y.; NuLi, Y.; Chen, Q.; Wang, Y.; Yang, J.; Wang, J. Magnesium cobalt silicate materials for reversible magnesium ion storage. Electrochim. Acta 2012, 66, 75−81. (21) Vatsala Rani, J.; Bhavana Rushi, S.; Kanakaiah, V.; Palaniappan, S. Green Fluorination of Natural Graphite and its Application in Rechargeable Magnesium Ion Transfer Battery. J. Electrochem. Soc. 2011, 158, A1031−A1035. (22) Kim, H. S.; Arthur, T. S.; Allred, G. D.; Zajicek, J.; Newman, J. G.; Rodnyansky, A. E.; Oliver, A. G.; Boggess, W. C.; Muldoon, J. Structure and compatibility of a magnesium electrolyte with a sulphur cathode. Nat. Commun. 2011, 2, 427. (23) Bradwell, D. J.; Kim, H.; Sirk, A. H. C.; Sadoway, D. R. Magnesium−Antimony Liquid Metal Battery for Stationary Energy Storage. J. Am. Chem. Soc. 2012, 134, 1895−1897. (24) Arthur, T. S.; Singh, N.; Matsui, M. Electrodeposited Bi, Sb and Bi1‑xSbx alloys as anodes for Mg-ion batteries. Electrochem. Commun. 2012, 16, 103−106. (25) Singh, N.; Arthur, T. S.; Ling, C.; Matsui, M.; Mizuno, F. A high energy-density tin anode for rechargeable magnesium-ion batteries. Chem. Commun. 2013, 49, 149−151. (26) Doe, R. E.; Mueller, T. K.; Ceder, G.; Barker, J. Electrode materials for magnesium batteries. WO/2011/150093, 2011. (27) Ling, C.; Banerjee, D.; Song, W.; Zhang, M.; Matsui, M. Firstprinciples study of the magnesiation of olivines: redox reaction mechanism, electrochemical and thermodynamic properties. J. Mater. Chem. 2012, 22, 13517−13523. (28) Zhang, R.; Yu, X.; Nam, K.-W.; Ling, C.; Arthur, T. S.; Song, W.; Knapp, A. M.; Ehrlich, S. N.; Yang, X.-Q.; Matsui, M. α-MnO2 as a cathode material for rechargeable Mg batteries. Electrochem. Commun. 2012, 23, 110−113.
the low current density galvanostatic cycling revealed that for crystalline V2O5 there are at least four different Mg2+ insertion stages or processes with different thermodynamic and kinetic characteristics. MoO3 thin-film electrodes were reversibly cycled at a relatively high capacity: 220 mAh g−1 at potential around 1.8 V vs Mg2+/Mg. Much larger overpotentials were observed for the reversible Mg intercalation processes into thin MoO3 electrodes than for V2O5 film electrodes. Raman spectroscopy and XRD measurements revealed that a portion of the magnesiated phase remains unreacted after electrochemical demagnesiation to high potentials. This may be due to a reduction in the electronic conductivity or to partial disintegration and isolation resulting from the large volume changes during cycling. XRD and Raman spectra clearly showed that a two-phase reaction occurs during the magnesiation of MoO3. For both V2O5 and MoO3 electrodes, higher overpotentials are needed for intercalation/deintercalation of Mg2+ ion than for Li+ ion.
■
ASSOCIATED CONTENT
S Supporting Information *
Detailed scheme and results of electrochemical and spectroscopic analyses. This material is available free of charge via the Internet at http://pubs.acs.org.
■
AUTHOR INFORMATION
Corresponding Author
*E-mail:
[email protected]. Author Contributions ‡
These authors contributed equally.
Notes
The authors declare no competing financial interest.
■
ACKNOWLEDGMENTS This research was supported by Pellion Technologies and ARPA-E award DE-AR0000062 and the BSF, Israel−US Binational Science Foundation.
■
REFERENCES
(1) Gregory, T. D.; Hoffman, R. J.; Winterton, R. C. Nonaqueous electrochemistry of magnesium. Applications to energy storage. J. Electrochem. Soc. 1990, 137, 775−780. (2) Novak, P.; Imhof, R.; Haas, O. Magnesium insertion electrodes for rechargeable nonaqueous batteries - a competitive alternative to lithium? Electrochim. Acta 1999, 45, 351−367. (3) Gofer, Y.; Chusid, O.; Aurbach, D. Magnesium Batteries (Secondary and Primary). In Encyclopedia of Electrochemical Power Sources; Jürgen, G., Ed.; Elsevier: Amsterdam, 2009; pp 285−301. (4) Aurbach, D. The electrochemical behavior of active metal electrodes in nonaqueous solutions. In Nonaqueous Electrochemistry; Aurbach, D., Ed.; Marcel Dekker: New York, 1999; pp 289−409. (5) Lu, Z.; Schechter, A.; Moshkovich, M.; Aurbach, D. On the electrochemical behavior of magnesium electrodes in polar aprotic electrolyte solutions. J. Electroanal. Chem. 1999, 466, 203−217. (6) Muldoon, J.; Bucur, C. B.; Oliver, A. G.; Sugimoto, T.; Matsui, M.; Kim, H. S.; Allred, G. D.; Zajicek, J.; Kotani, Y. Electrolyte roadblocks to a magnesium rechargeable battery. Energy Environ. Sci. 2012, 5, 5941−5950. (7) 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, 724−727. (8) Yoo, H. D.; Shterenberg, I.; Gofer, Y.; Gershinsky, G.; Pour, N.; Aurbach, D. Mg rechargeable batteries: an on-going challenge. Energy Environ. Sci. 2013, 6, 2265−2279. 10971
dx.doi.org/10.1021/la402391f | Langmuir 2013, 29, 10964−10972
Langmuir
Article
(29) Bhandavat, R.; David, L.; Singh, G. Synthesis of SurfaceFunctionalized WS2 Nanosheets and Performance as Li-Ion Battery Anodes. J. Phys. Chem. Lett. 2012, 3, 1523−1530. (30) Chernova, N. A.; Roppolo, M.; Dillon, A. C.; Whittingham, M. S. Layered vanadium and molybdenum oxides: batteries and electrochromics. J. Mater. Chem. 2009, 19, 2526−2552. (31) Novák, P.; Desilvestro, J. Electrochemical Insertion of Magnesium in Metal Oxides and Sulfides from Aprotic Electrolytes. J. Electrochem. Soc. 1993, 140, 140−144. (32) Novák, P.; Scheifele, W.; Joho, F.; Haas, O. Electrochemical Insertion of Magnesium into Hydrated Vanadium Bronzes. J. Electrochem. Soc. 1995, 142, 2544−2550. (33) Le, D. B.; Passerini, S.; Coustier, F.; Guo, J.; Soderstrom, T.; Owens, B. B.; Smyrl, W. H. Intercalation of Polyvalent Cations into V2O5 Aerogels. Chem. Mater. 1998, 10, 682−684. (34) Amatucci, G. G.; Badway, F.; Singhal, A.; Beaudoin, B.; Skandan, G.; Bowmer, T.; Plitz, I.; Pereira, N.; Chapman, T.; Jaworski, R. Investigation of Yttrium and Polyvalent Ion Intercalation into Nanocrystalline Vanadium Oxide. J. Electrochem. Soc. 2001, 148, A940−A950. (35) Imamura, D.; Miyayama, M. Characterization of magnesiumintercalated V2O5/carbon composites. Solid State Ionics 2003, 161, 173−180. (36) Imamura, D.; Miyayama, M.; Hibino, M.; Kudo, T. Mg Intercalation Properties into V2O5 gel/Carbon Composites under High-Rate Condition. J. Electrochem. Soc. 2003, 150, A753−A758. (37) Yang, S.; Gong, Y.; Liu, Z.; Zhan, L.; Hashim, D. P.; Ma, L.; Vajtai, R.; Ajayan, P. M. Bottom-up Approach toward SingleCrystalline VO2-Graphene Ribbons as Cathodes for Ultrafast Lithium Storage. Nano Lett. 2013, 13, 1596−1601. (38) Ruch, P. W.; Cericola, D.; Hahn, M.; Kotz, R.; Wokaun, A. On the use of activated carbon as a quasi-reference electrode in nonaqueous electrolyte solutions. J. Electroanal. Chem. 2009, 636, 128− 131. (39) zad, A. I.; Kavei, G.; Tabar, M. R. R.; Allaei, S. M. V. Height fluctuations and intermittency of V2O5 films by atomic force microscopy. J. Phys.: Condens. Matter 2003, 15, 1889. (40) Baddour-Hadjean, R.; Raekelboom, E.; Pereira-Ramos, J. P. New Structural Characterization of the LixV2O5 System Provided by Raman Spectroscopy. Chem. Mater. 2006, 18, 3548−3556. (41) Liu, Y.; Li, J.; Zhang, Q.; Zhou, N.; Uchaker, E.; Cao, G. Porous nanostructured V2O5 film electrode with excellent Li-ion intercalation properties. Electrochem. Commun. 2011, 13, 1276−1279. (42) Haber, J.; Witko, M.; Tokarz, R. Vanadium pentoxide I. Structures and properties. Appl. Catal. A: Gen. 1997, 157, 3−22. (43) Baddour-Hadjean, R.; Pereira-Ramos, J. P.; Navone, C.; Smirnov, M. Raman Microspectrometry Study of Electrochemical Lithium Intercalation into Sputtered Crystalline V2O5 Thin Films. Chem. Mater. 2008, 20, 1916−1923. (44) Spahr, M. E.; Novak, P.; Haas, O.; Nesper, R. Electrochemical insertion of lithium, sodium and magnesium in molybdenum(VI) oxide. J. Power Sources 1995, 54, 346−351. (45) Hashem, A. M.; Askar, M. H.; Winter, M.; Albering, J. H.; Besenhard, J. O. Two-phase reaction mechanism during chemical lithium insertion into α-MoO3. Ionics 2007, 13, 3−8. (46) Bullard, J. W.; Smith, R. L. Structural evolution of the MoO3(010) surface during lithium intercalation. Solid State Ionics 2003, 160, 335−349. (47) Julien, C.; Mohammad Hussain, O.; El-Farh, L.; Balkanski, M., Electrochemical studies of lithium insertion in MoO3 films. Solid State Ionics 1992, 53−56, Part 1, 400-404. (48) Julien, C.; Nazri, G. A. Transport properties of lithiumintercalated MoO3. Solid State Ionics 1994, 68, 111−116. (49) Tsumura, T.; Inagaki, M. Lithium insertion/extraction reaction on crystalline MoO3. Solid State Ionics 1997, 104, 183−189.
was included. The corrected version of the paper was reposted on August 12, 2013.
■
NOTE ADDED AFTER ASAP PUBLICATION This paper was published on the Web on August 7, 2013. Because of a production error, an incorrect version of Figure 6 10972
dx.doi.org/10.1021/la402391f | Langmuir 2013, 29, 10964−10972
