Article pubs.acs.org/cm
LiFeO2-Incorporated Li2MoO3 as a Cathode Additive for Lithium-Ion Battery Safety Kyu-Sung Park,†,§,* Dongmin Im,† Anass Benayad,‡ Anthony Dylla,∥ Keith J. Stevenson,∥ and John B. Goodenough§ †
Battery group and ‡Analytical Engineering group, Samsung Advanced Institute of Technology (SAIT), PO Box 111, Suwon 440-600, South Korea § Texas Materials Institute and ∥Department of Chemistry and Biochemistry, University of Texas at Austin, Austin, Texas 78712, United States S Supporting Information *
ABSTRACT: Li2MoO3, with a Mo(IV)/Mo(VI) redox couple, has been tested as a cathode additive to make lithium-ion batteries safe under abnormal discharge conditions. Its high charging capacity and sloping discharge voltage below 3.4 V vs Li+/Li effectively prevents the Cu anode current collector from oxidative dissolution at the overdischarge condition. However, molybdenum dissolution from the electrochemically charged Li2MoO3 quickly deteriorates the battery performance at 45 °C. A solid solution of Li2MoO3 with LiFeO2 has stabilized the crystalline structure at the charged states and suppressed the Mo(VI) dissolution. The addition of 10 wt % 0.9Li2MoO3−0.1LiFeO2 to LiCoO2 electrode enables long-term, high-temperature cycling in the operating voltage range of 0.0−4.3 V vs graphite. KEYWORDS: lithium-ion battery, overdischarge, Li2MoO3, Mo dissolution, LiFeO2
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INTRODUCTION During the last 20 years, the lithium-ion battery (LIB) has proven to be a successful power source for portable electronic devices. Another major trend in battery development is the design and fabrication of large-scale LIBs for vehicle applications.1 The vehicles are typically classified according to the driving range provided by the batteries like electric vehicles (EVs) and hybrid electric vehicles (HEVs), and they are extensively studied in terms of developmental and applied research. The most distinctive difference between vehicle and portable electronics applications is the amount of stored energy, that is, the number of cells connected within a battery pack, which makes safety an even more critical issue. In the conventional lithium-ion battery system, the cathode is typically an intercalation compound; the cathode provides Li+ions to the anode and determines the cell capacity.1,2 Therefore, the basic design strategy for a full-cell operation would be balancing capacities and Coulombic efficiencies of the cathode and the anode as well as their voltages. Besides these strategies, another important design factor is battery safety. Battery components have been generally modified and changed to prevent possible safety hazards.3 For example, separators were modified to have a dimensional stability upon thermal exposure or coated with ceramic layers.4 However, modifying the cathode/anode cell design would be more effective to prevent any undesirable safety events in advance. This concept © 2012 American Chemical Society
is of more importance for EV and HEV batteries because, in a battery pack, hundreds of cells are integrated in series and in parallel with complicated control units and a cooling system. Different cell voltages, internal impedances, temperature distributions, and capacities will result in cell imbalance, which might potentially cause overcharge or overdischarge in a particular cell unless monitored and controlled properly. In this report, we show an effective method for stabilizing cell performance under abnormal voltage conditions. Discharge to 0 V constant current followed by a constant voltage is adopted as an extreme overdischarge characterization condition for testing the cycle stability both at room temperature and 45 °C. The most problematic phenomenon of the method in a full cell under the overdischarge condition is the anodic dissolution of the Cu current collector, which occurs above approximately 3.6 V vs Li+/Li, so the anode voltage must be managed to be less than the dissolution voltage.5 The voltage curves of the cathode are modified on introducing a cathode additive to control the voltage of the anode during discharge (Li+-ion extraction from the anode). Detailed material design rules for the cathode additive will be given after discussing the voltage behavior of an overdischarged Received: February 14, 2012 Revised: June 21, 2012 Published: June 25, 2012 2673
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Figure 1. (a) Half-cell voltage profiles of LiCoO2 and graphite at the first cycle. Graphite electrode was overcharged to show an anodic dissolution of Cu current collector. (b) Cell balancing of the LiCoO2 and graphite with the initial half-cell data. MPD diffractometer (40 kV, 40 mA; Cu Kα radiation). Crystal structure refinement was performed by the Rietveld method with the Fullprof software. The powder morphology was characterized by fieldemission scanning electron microscopy (FE-SEM, Hitachi S-4500). Time-of-flight secondary ion mass spectroscopy (TOF-SIMS) was performed on a TOF.SIMS 5 by ION-TOF GMBH with Bi3+2 ions accelerated at 30 kV as the analysis (primary) gun and O2 accelerated at 2 kV as the sputtering (secondary) gun. A randomly rastered 150 mm2 analysis area (128 × 128 pixels) inside a 300 mm2 sputtering area was used to acquire the 3-D ion-distribution maps. Secondary ions were detected in positive-ion mode. An electron flood gun was introduced between analysis and sputtering steps to avoid charging of the surface. Pelletized samples were loaded into the vacuum chamber and allowed to purge for a minimum of 12 h before introduction to the main analysis chamber. All analyses were performed with mainanalysis-chamber pressures between 5 and 9 × 10−9 mbar. Electrode XPS analysis was performed with a Φ Physical Electronics (Quantum 2000 Scanning ESCA Microprobe) spectrometer and a focused monochromatized Al Kα radiation (1486.6 eV). The residual pressure inside the analysis chamber was 7 × 10−9 Torr. Drastic precautions were taken in order to avoid any air contamination during the preparation or the transfer of the cell electrodes, including Li metal, from an Ar-filled glovebox to the XPS main chamber by using a specially designed air-proof transfer chamber. Powder XPS was carried out with a Kratos AXIS Ultra DLD system calibrated using the signals for C 1s at 284.7 eV. The electrochemical characterization was carried out by galvanostatic cycling with Li2MoO3 and/or LiCoO2 as a cathode and Li metal foil or graphite as an anode. The commercial LiCoO2 used has an
electrochemical cell. Li2MoO3 has been chosen as the cathode additive for that purpose in this work. James and Goodenough first reported the basic electrochemical properties of Li2MoO3 and its atomic structure in the transition-metal layer.6 We have tested Li2MoO3 as a potential cathode additive and identified its chemical instability, especially in the charged state, Li2−xMoO3 (0 < x < 2). Moreover, by making a solid solution of Li2MoO3 and LiFeO2, a strikingly stable high-temperature cycle performance could be achieved in the operating voltage range of 0.0−4.3 V vs graphite.
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EXPERIMENTAL SECTION
Li2MoO3 powder was prepared by solid-state reaction. A stoichiometric amount of Li2CO3 (Aldrich) and MoO3 (Aldrich) were thoroughly mixed and fired two times at 700 °C for 10 h under 5% H2 and 95% N2 flow to make crystalline Li2MoO3 powder without any impurity phases. Synthesizing a series of solid solutions between Li2MoO3 and LiFeO2 ((1−x)Li2MoO3−xLiFeO2; x = 0.1, 0.2, and 0.3) was also done by the solid-state reaction. A controlled amount of Li2CO3 and lab-made α-Fe2O3 were mixed with the synthesized Li2MoO3 and fired at 700 °C for 10 h under 5% H2 and 95% N2 atmosphere. The α-Fe2O3 was obtained by a thermal decomposition of iron(II) oxalate dehydrate (Aldrich, Fe(C2O4)·2H2O) at 600 °C in air. Synthesized 0.9Li2MoO3−0.1LiFeO2, 0.8Li2MoO3−0.2LiFeO2, and 0.7Li2MoO3−0.3LiFeO2 are referred to as M9F1, M8F2, and M7F3, respectively. The structural properties of Li2MoO3 and its solid solution with LiFeO2 have been studied by X-ray diffraction with a Philips X’Pert 2674
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Figure 2. (a) Initial charge, discharge voltage curves of LiCoO2 and its blend electrode with 10% Li2MoO3, which are done at the rate of C/10 between 1.5 and 4.3 V vs Li+/Li at room temperature. (b) Cell balancing of the cathodes and graphite with the initial half-cell data.
corresponding to 289 mAh g−1, which gives an initial Coulombic efficiency (I.C.E.) of 92.9% when discharged to 3.0 V vs Li+/Li. Abnormal discharging induces the electrochemical dissolution of the Cu current collector at higher voltage ranges, as is shown in Figure 1a. Because LiCoO2 has a higher ICE of 96.9%, overdischarge could potentially cause Cu dissolution as indicated with the dashed green circle in Figure 1b. Maintaining the ICE of LiCoO2 to less than that of graphite is a basic requirement for negating the possibility of too high a voltage increase of the anode during discharge. In this condition, the anode voltage remains in the low-voltage region versus lithium after the voltage of LiCoO2 fully drops. However, it has been reported that LiCoO2 cathodically decomposes at 1.636 V vs Li+/Li under equilibrium condition,8 so the voltage of the LiCoO2 electrode must be kept above that value. Hence, it is also required that the voltage of a Li1−xCoO2 cathode in the low-voltage region (i.e., below the Cu dissolution potential) should be sloping. To find a cathode additive to fulfill all these requirements, some design rules were considered. First, it should have high Li+-ion content that could be extractable as much as possible without O2 gas evolution during charging. Second, Li+-ion insertion (discharge) must be significantly restricted above the Cu-dissolution voltage, but it should have a sloping voltage curve below that. We have chosen Li2MoO3 as a possible cathode additive because the energy of
average particle size of ca. 20 μm. The electrode compositions for halfcell characterization were LiCoO2/carbon black/PVDF = 96:2:2 and Li2MoO3 or (1−x)Li2MoO3−xLiFeO2/carbon black/PVDF = 93:2:5 in weight. The cathode composition for the full-cell cathode-blend test was LiCoO2/Li2MoO3 or (1−x)Li2MoO3−xLiFeO2/carbon black/ PVDF = 86.4:9.6:2:2 in weight. The liquid electrolyte was 1.3 M LiPF6 in EC/DEC/EMC = 3/5/2 in volume. The electrochemical experiments were controlled with a TOYO system and performed at room temperature, 45 and 60 °C. Detailed charge/discharge test conditions with the exact composition of the blend electrodes will be given for each cell-test result.
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RESULTS 1. Specification for the Cathode Additive. LiCoO2 as a cathode material and graphite as an anode material are the most popular full-cell combination for the lithium-ion battery.1,7 We employed the standard cell chemistry to verify our cell design in this work, and different layered cathode materials considered for the vehicle application would bring the same results. In the combination, the graphite electrode typically has been designed to have a 10−20% larger charge capacity (Li+-insertion) than that of the cathode to avoid possible Li-metal plating on the anode. The graphite electrode used in this study was designed to have a charge capacity of ∼2 mAh cm−2 while the LiCoO2 cathode was designed to have ∼1.7 mAh cm−2. The cutoff condition for galvanostatic Li+-insertion into graphite in the half-cell configuration is a capacity limit of 1.7 mAh cm−2, 2675
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Figure 3. Voltage curves of full-cells employing graphite as an anode. The cells are cycled between 0 and 4.3 V vs graphite at the rate of C/2 with a constant-voltage cutoff condition at C/100 both at charge and discharge. Cathodes are (a) LiCoO2 and (b, c) blend electrodes of 90 wt % of LiCoO2 and 10 wt % of Li2MoO3. Cycling temperature is (a, b) room temperature and (c) 45 °C.
parameter was calculated with Fullprof software by assuming its crystallographic structure to be a trigonal unit cell (R3̅m): the unit-cell parameters were a = 2.8812(2) Å and c = 14.955(1) Å. From the unit cell information, the theoretical density of Li2MoO3 was calculated to be 7.31 g cc−1, which is higher than that of LiCoO2 (∼ 5 g cc−1). Therefore, Li2MoO3 can be blended into a LiCoO2 electrode without sacrificing electrode density.
Mo(VI)/Mo(IV) redox couple lies above the top of the O-2p bands, which allows extraction all the lithium without oxygen oxidation, and it has a low discharge voltage starting from ca. 3.4 V vs Li+/Li.6,9 Previous reports on Li2MoO3 were limited, because it has too low a discharge voltage to be a competitive cathode material and molybdenum is cost prohibitive. 2. Li2MoO3 as Additive. Single-phase Li2MoO3 was synthesized by solid-state reaction; its powder X-ray diffraction is shown in Figure S1 (Supporting Information). The lattice 2676
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Figure 4. Ex situ electrode XRD patterns of Li2MoO3 during charging to 4.4 V vs Li+/Li at the rate of C/20 and at 60 °C. The corresponding charging voltage curve is displayed with arrows indicating the state of charge.
Figure 5. Mo 3d XPS spectra on the counter electrode (Li metal) at various state of charge. Cathode is Li2MoO3 and charged to 4.4 V vs Li+/Li at the rate of C/20 and at 60 °C. Inset photograph shows the surface of Li metal at the end of the charge. A SEM image and the corresponding EDS spectrum were taken on the surface of separator at the end of the charge.
Li2MoO3 shows a charge capacity of 232 mAh g−1 (∼1.36 Li -extraction) and a discharge capacity of 228 mAh g−1 (∼1.34 Li+-insertion) between 4.4 and 1.5 V vs Li+/Li at room temperature, and the discharge starts below 3.4 V vs Li+/Li; see Figure S2 (Supporting Information). The charge/discharge feature of Li2MoO3 seems to be adequate for the cathode
additive because of the high charging capacity and no discharge capacity above 3.4 V vs Li+/Li. One more notable characteristic is the sloping discharge voltage curve below 3.4 V vs Li+/Li, which can modify the voltage curves of LiCoO2 electrode on blending with Li2MoO3. Gas evolution during charging of Li2MoO3 was also checked in parallel by making a pouch-cell,
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Figure 6. Powder XRD patterns of (a) Li2MoO3, (b) 0.9Li2MoO3−0.1LiFeO2, (c) 0.8Li2MoO3−0.2LiFeO2, and (d) 0.7Li2MoO3−0.3LiFeO2.
100 mAh g−1 steps were performed as shown in Figure 4. The charging temperature was set to 60 °C to accelerate the degradation. At 60 °C, the charging capacity was dramatically increased, but irreversibly, to 285 mAh g−1 (∼1.68 Li+extraction). Crystalline Li2MoO3 is gradually decomposed during charging and eventually becomes amorphous. Such a decomposition would pose a problem in the electrochemical test, but it is not clear whether the structural degradation is the origin of the cell degradation presented in Figure 3c. To analyze changes in cell components, the coin cell was disassembled after the high-temperature charging. The Li metal and separator were contaminated with a deposit as shown in Figure 5. SEM and EDS analyses of the separator showed that molybdenum-containing particles covered the porous polymeric separator surface. This observation shows there is a molybdenum dissolution from Li2MoO3. The color of the deposit varies from blue to brownish black depending on the time after the charging. The degree of reduction of molybdenum by Li metal from the anode may affect the color change. Ex situ XPS surface analysis was performed on the counter electrode to trace the molybdenum dissolution from Li2MoO3 during charging. The counter electrode (Li metal) was collected as soon as the Li2MoO3 in the coin cell reached a certain capacity, i.e., a different state of charge. In the glovebox filled with Ar gas, the Li metal was rinsed with DMC solvent and carefully transferred to the XPS sample chamber without exposing it to air; the results are presented in Figure 5. For the 0 mAh g−1 sample preparation, an assembled coin-cell was aged at room temperature for 1 day. There was no molybdenum on the anode surface before Li+ extraction from the cathode, but a Mo(VI) XPS signal was observed after even a low level of charging (100 mAh g−1): Mo 3d5/2 binding energy is 232.1 eV.10 The XPS peak intensity increased during the galvanostatic charging at 60 °C, indicating that Mo(VI) is continuously dissolved in the electrolyte and migrates to the anode surface. Moreover, a peak around 225 eV and broad noisy peaks present up to 230 eV indicate that molybdenum is further reduced to a metallic state together with some transitional oxidation states after charging over 200 mAh g−1,10 which would leave additional byproducts on the anode. Combining ex situ XRD and XPS results, it is evident that the
but no gas was detected during several charge/discharge cycles (data not included in this report). It is also noted that the OCV (open circuit voltage) after cell assembly is below 3.0 V, typically around 2.7 V vs Li+/Li. The modification of the discharge voltage curve of LiCoO2 was confirmed after blending 10 wt % of Li2MoO3 into LiCoO2; the results are illustrated in Figure 2. A 10 wt % substitution of Li2MoO3 made the specific charge capacity increase from 152.5 to 166.1 mAh g−1 ( 3.0 V vs Li+/Li; at the same time, it induced a sloping voltage curve at the end of discharge. The sloping Li+-insertion voltage curve delivers additional capacity and shows a discharge capacity of 160.8 mAh g−1 for V > 1.5 V vs Li+/Li. The importance of the Li2MoO3 addition could be seen in Figure 2b. With the aid of Li2MoO3, the discharge cutoff voltage of a graphite anode is reduced to an electrochemically stable region, which prevents the Cu-dissolution in the overdischarge condition. Therefore, it is technically possible to control precisely the voltage of an anode at 0 V vs graphite in a full-cell by controlling the blending ratio of the cathode composite. This design concept has been evaluated by making a coin full-cell, and the charge/discharge results are listed in Figure 3. A LiCoO2/graphite cell exhibits severe degradation when cycled between 4.3 and 0.0 V vs graphite (Figure 3a), while a (90 wt % LiCoO2 and 10 wt % Li2MoO3)/graphite cell cycles stably for 100 cycles (Figure 3b). Note that we used a harsh CV cutoff condition of C/100 even at the end of discharge, 0.0 V vs graphite. The sloping voltage curves, below 3.5 V vs graphite, are also reproducible during cycling, which indicates that the Li2MoO3 is reversibly charged/discharged at the cathode side and, more importantly, the discharge voltage of graphite is not increased above 3.0 V vs Li+/Li as designed. However, when the full-cell is exposed to high temperature, 45 °C, the cycle performance deteriorated immediately and the voltage signature of Li2MoO3 was not detected at discharge as shown in Figure 3c. LiCoO2 is thought to have structural and chemical stability at 45 °C, so Li2MoO3 was assumed to have caused the temperature-dependent degradation in the present case. To verify the structural and chemical instabilities of charged Li2MoO3, ex situ X-ray diffraction of Li2MoO3 electrodes with 2678
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chemical instability of charged Li2‑xMoO3 would be a main failure mechanism of a (90 wt % LiCoO2 + 10 wt % Li2MoO3)/ graphite cell during high-temperature cycling. It is known that molybdenum compounds are easily dissolved in water and other solvents, and there are numerous reports on soluble polyoxocomplexes of molybdenum.11 In this report, we show that a solid solution of Li2MoO3 with LiFeO2 greatly reduces the solubility of Mo−oxide species. 3. LiFeO2−Li2MoO3 as Additive. The evolution with LiFeO2 concentration of the structure of LiFeO2−Li2MoO3 solid solutions was monitored with powder XRD as presented in Figure 6. All compounds were assumed to have the layered structure with space group R3̅m, and the lattice parameters obey Vegard’s law well before having a notable antisite mixing of Fe to Li-layer (3b site) that happened at 0.7Li2−xMoO30.3LiFeO2 (see Figures S3 and S4 and Table S1, Supporting Information). The cation mixing induced an abrupt decrease in the lattice parameter c and c/a ratio, which also caused a notable shift of the (003) peak to higher Bragg angle. Refined total occupancies of cations (Li, Fe, and Mo) in the structure vary linearly, which represents well the starting composition. Therefore, it is suggested that Li2MoO3 forms a solid solution with LiFeO2, but it has an antisite mixing of Fe. In the refinement, the bond valence sum of Fe is calculated to be +2.793(3) and that of Mo is +3.937(5), which are close to their formal oxidation states in the composition and corresponds well to the XPS result (Figure S5, Supporting Information). Moreover, a peak at approximately 20° in the 0.9Li2MoO3− 0.1LiFeO2 sample was identified as the (212) peak of an Fe3O4 impurity. As illustrated in Figure S6 (Supporting Information), the particle shapes and size were also modified by the incorporation of LiFeO2; Li2CO3 and Fe2O3 additives to Li2MoO3 lower the sintering temperature, and the particle size increases with LiFeO2 concentration. Time-of-flight secondary-ion mass spectroscopy (TOFSIMS) was conducted in parallel to define the Fe distribution in the solid solution samples and to get three-dimensional elemental distribution information (Figure 7). In the images, the x−y plane (150 μm × 150 μm) in Figure 7abc represents the powder surface, and the x−z plane in Figure 7d−f shows the areal depth-profile of Fe at a certain y-position. It is observed that, with increasing Fe content, the distribution of Fe
becomes more homogeneous throughout the particle surface and depth. For example, in 0.9Li2MoO3−0.1LiFeO2, there is a higher image contrast between bright and dark areas on the surface, which represents higher and lower Fe content in each area, respectively (Figure 7a). The spatial distribution is also inhomogeneous along the x-direction and is more concentrated on the surface along the z-direction (Figure 7d,e with different y-positions). This result is consistent with the existence of an Fe3O4 impurity in the same sample as identified in the XRD pattern. It is suspected that a relatively small amount of Li2CO3 and α-Fe2O3 may not be able to make a homogeneous mixture with the Li2MoO3 agglomerate as seen in the Figure S6a (Supporting Information); but by 0.8Li2MoO3−0.2LiFeO2, the distribution of Fe is homogeneous throughout the bulk (Figure 7f). Before designing full-cells with the solid solution between LiFeO2 and Li2MoO3, their half-cell properties were characterized (see Figure S7, Supporting Information). Upon charging to 4.4 V vs Li+/Li, the capacities of 0.9Li2MoO3− 0.1LiFeO2 and 0.8Li2MoO3−0.2LiFeO2 were 235.3 and 221.7 mAh g−1, respectively, while Li2MoO3 delivered 244 mAh g−1. However, 0.7Li2MoO3−0.3LiFeO2 was found to be completely inactive during the charge/discharge. Longer constant-voltage charging region after the LiFeO2-incorporation is consistent with the larger cathode particles. On the other hand, there was a significant drop in the discharge capacities with LiFeO2 content, from 243 mAh g−1 for Li2MoO3 to 139.5 mAh g−1 with 0.1LiFeO2 and 89.9 mAh g−1 with 0.2LiFeO2. In the test, we did not fully investigate the origin of the change in the voltage curves both at charge and discharge, but collected cell parameters needed for the full-cell design. As identified in the high-temperature charging test, the most serious problem with Li2MoO3 is its chemical instability. Figures 8 and 9 display the 60 °C charge behavior of 0.9Li2MoO3−0.1LiFeO2 and 0.8Li2MoO3−0.2LiFeO2, respectively. The two materials deliver much higher capacity at this temperature compared to room-temperature: 281.4 mAh g−1 for 0.9Li 2 MoO 3 −0.1LiFeO 2 and 261.8 mAh g −1 for 0.8Li2MoO3−0.2LiFeO2 were achieved; a capacity increase of 46.1 and 40.1 mAh g−1 in each composition confirms the resistive nature of Li + extraction in the LiFeO 2 -rich compositions. Besides the cell performance, the structural stability of the charged states was evaluated with ex situ XRD analysis. The structural decomposition during the Li+ extraction was also observed in both compositions, but LiFeO2 clearly enhanced the stability of the layered structure even at the end of the charge. Another interesting phenomenon is a positive shift of the (003) peak with charge. In the layered transition metal oxide, Li+-extraction should increase the Coulombic repulsion between oxide ions in adjacent oxygen layers, which gives a positive lattice parameter increase in the c-axis, i.e., a negative shift of the (003) Bragg angle. However, a positive shift of the (003) peak in the results implies a decrease in the lattice parameter along the c-axis. The surface of the counter electrode (Li metal) was also investigated. Figure 8 shows photographs of the Li metal at different states of charge. Unlike the Li2MoO3 half-cell, the surface is remarkably clean without any deposit even at the end of the charge. In the 0.8Li2MoO3−0.2LiFeO2 half-cell, a Mo 3d XPS spectrum at the end of charge was obtained and compared to that of the Li2MoO3 half-cell (Figure 9). A significant decrease in Mo XPS signals could be seen, which is comparable
Figure 7. 2D Fe composition maps (150 μm × 150 μm) obtained by TOF-SIMS experiment. (a) 0.9Li 2 MoO 3 −0.1LiFeO 2 , (b) 0.8Li2MoO3−0.2LiFeO2, and (c) 0.7Li2MoO3−0.3LiFeO2. Depth distribution of Fe is also presented in (d,e) for 0.9Li2MoO3− 0.1LiFeO2 and (f) for 0.8Li2MoO3−0.2LiFeO2. 2679
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Figure 8. Ex situ electrode XRD patterns of 0.9Li2MoO3−0.1LiFeO2 during charging to 4.4 V vs Li+/Li at the rate of C/20 and at 60 °C. The corresponding charging voltage curve and the counter electrode (Li metal) photographs are displayed.
Figure 9. Ex situ electrode XRD patterns of 0.8Li2MoO3−0.2LiFeO2 (M8F2) during charging to 4.4 V vs Li+/Li at the rate of C/20 and at 60 °C. Mo 3d XPS spectrum on Li metal at the end of the charge are presented to compare with Mo dissolution when using Li2MoO3 as a cathode.
from Li2MoO3. There is another sloping voltage curve below 0.5 V vs graphite in the case of the LiCoO2/0.8Li2MoO3− 0.2LiFeO2 blend electrode. It is suspected that LiCoO2 begins to decompose at this voltage because of the reduced ICE and the abrupt voltage decrease below 2.5 V vs Li+/Li. The decomposition is quickly terminated by the abrupt voltage increase of the graphite anode in this case. The chemical stability of the LiFeO2-incorporated Li2MoO3 brings true benefits to the charge/discharge reversibility over a large voltage window from 0.0 to 4.3 V vs graphite as presented in Figure 11. The cells are cycled at the rate of C/2 with a constant-voltage cutoff condition at C/100 both at charge and discharge. In this condition, a LiCoO2 full cell is fully deteriorated within initial 10 cycles and the full cell with LiCoO2/Li2MoO3 blend electrode is only stable at room temperature. However, the full cell with LiCoO2/0.9Li2MoO3− 0.1LiFeO2 shows a remarkably stable cycle performance both at room temperature and 45 °C. For 100 cycles, the capacities only dropped from 153.8 to 141.2 mAh g−1 (91.8% retention) at room temperature and from 162.8 to 138.9 mAh g−1 (85.3% retention) at 45 °C. The 0.9Li2MoO3−0.1LiFeO2 additive is not only chemically stable even at high temperature but also
to the level of cobalt dissolution from LiCoO2 in the same test condition (Figure S8, Supporting Information). The incorporation of LiFeO2 clearly enhances the chemical stability of Li2−xMoO3. The cell properties of the blend electrodes having the solid solutions were characterized as presented in Figure 10. In halfcells, decreased initial Coulombic efficiencies above 3.0 V vs Li+/Li were observed at all samples, and the shape of the voltage curve below 3.0 V was modified as well. Below 3.0 V, the discharge voltage drops abruptly as the Fe-content is increased, so the LiCoO2/0.8Li2MoO3−0.2LiFeO2 blend electrode has an almost similar slope as that of the composite LiCoO2 as a result of the discharge characteristics of 0.8Li2MoO3−0.2LiFeO2 as discussed in the Figure S7 (Supporting Information). The half-cell test also confirmed the cathodic decomposition of LiCoO2 at a voltage below 1.2 V vs Li+/Li under a constant-current discharge condition, which gives a huge capacity of 878 mAh g−1 to the LiCoO2/ 0.9Li2MoO3−0.1LiFeO2 blend electrode.8 Therefore, a cathode, containing LiCoO2, should have its voltage kept higher than 1.2 V vs Li+/Li at deep discharge. In full-cells, as shown in Figure 3, all the cells show the signature sloping voltage curve 2680
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Figure 10. Initial charge, discharge voltage curves of LiCoO2 and its blend electrodes with 10% Li2MoO3, 0.9Li2MoO3−0.1LiFeO2 (M9F1), and 0.8Li2MoO3−0.2LiFeO2 (M8F2) at room temperature. (a) Half-cell test was done between 1.5 and 4.3 V vs Li+/Li except the LiCoO2/M9F1 blend electrode that was discharged to 0.5 V vs Li+/Li. (b) Full-cells are tested between 0 and 4.3 V vs graphite at the rate of C/10 with a constant-voltage cutoff condition at C/100 both at charge and discharge.
the Mo(IV) to maximize the bonding of each Mo(IV) to two neighboring Mo(IV) with 60° between the Mo−Mo bonds. This process would create a surface amorphous phase that grows as the delithiation process continues; the resistance to Li+ and Mo(VI) transport from the bulk would increase with the thickness of the amorphous layer. Incorporation of LiFeO2 introduces Fe3+ ions into the Li+ layer, thereby increasing the bonding between the transitionmetal layers to shorten the c-axis; but the Fe3+ in the Li+ layers impede the Li+ mobility and decrease the capacity of Li extraction/insertion. On the other hand, an Fe-rich surface together with the Fe3O4 impurity, as is evident in the 0.1LiFeO2 sample, decreases the surface Mo concentration, thus suppressing the surface disproportionation reaction and stabilizing the crystalline character of the surface layers.
electrochemically stable over a large voltage window. In the case of LiCoO2/0.8Li2MoO3−0.2LiFeO2 full-cell, stable cycling is not observed, presumably because of the side reaction at voltages below 0.5 V vs graphite, even though it is chemically stable in the charged states.
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DISCUSSION Discharged Li2MoO3 contains Li1/3Mo2/3O2 layers in which there is a random distribution of Mo3O13 clusters that are stabilized by Mo−Mo bonding across shared octahedral-site edges.6 On charging, oxidation of Mo(IV) to Mo(V) results, at least at the surface, in the disproportionation reaction 2Mo(V) = Mo(IV) + Mo(VI) with dissolution of the Mo(VI) into the electrolyte as evidenced by XPS analysis (Figure S5, Supporting Information). Higher temperature induces more rapid displacement of Mo(VI) to the electrolyte. Removal of Mo(VI) and Li+ from the solid would be accompanied by a rearrangement of 2681
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Figure 11. Full-cell cycle performance at (a) room temperature and (b) 45 °C. The cells are cycled between 0 and 4.3 V vs graphite at the rate of C/ 2 with a constant-voltage cutoff condition at C/100 both at charge and discharge. Cathodes are LiCoO2 and its blend electrodes with 10% Li2MoO3, 0.9Li2MoO3−0.1LiFeO2 (M9F1), and 0.8Li2MoO3−0.2LiFeO2 (M8F2).
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shows a strikingly stable cycle performance at 45 °C in the operating voltage range of 0.0 to 4.3 V vs graphite with 85.3% capacity retention after 100 cycles.
CONCLUSION
The evolution of large-scale lithium-ion batteries has brought a new set of materials problems whose solutions can answer the problems of the battery pack system. In this report, we show that Li2MoO3 can solve the cell voltage imbalance problem of series-connected lithium-ion batteries. It has a high charging capacity of 232 mAh g−1 at room temperature and a sloping discharge voltage curve that starts down from 3.4 V vs Li+/Li, which can modify the voltage curves of a cathode composite and therefore control the charge/discharge characteristics of a lithium-ion battery. A 10 wt % Li2MoO3−blended LiCoO2 cathode makes it possible to discharge down to 0.0 V vs graphite. However, a chemical instability in the charged states was encountered at higher temperature, which was associated with the molybdenum dissolution from Li2−xMoO3. With LiFeO2 incorporated into Li2MoO3 as a solid solution, we were able to stabilize the crystalline structure at the charged states and suppress the molybdenum dissolution. The full cell with 10 wt % 0.9Li2MoO3−0.1LiFeO2 blended into a LiCoO2 cathode
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ASSOCIATED CONTENT
S Supporting Information *
Refined XRD information, XPS spectra, powder SEM images, and electrochemical cell data. This material is available free of charge via the Internet at http://pubs.acs.org.
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AUTHOR INFORMATION
Corresponding Author
*E-mail:
[email protected]. Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS K.-S.P. thanks Dr. S. Larrégola for helpful discussion. The TOFSIMS experiment was supported as part of the program 2682
dx.doi.org/10.1021/cm300505y | Chem. Mater. 2012, 24, 2673−2683
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“Understanding Charge Separation and Transfer at Interfaces in Energy Materials (EFRC:CST)”, an Energy Frontier Research Center funded by the U.S. Department of Energy, Office of Science, Office of Basic Energy Sciences under Award Number DE-SC0001091. We also gratefully acknowledge the TOFSIMS facility (which includes a TOF.SIMS 5 instrument produced by ION-TOF GmbH (Germany) in 2010 and purchased through the NSF grant DMR-0923096), part of the Texas Materials Institute at University of Texas at Austin, for helping us with the TOF-SIMS data acquisition and interpretation.
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