Layered Molybdenum (Oxy)Pyrophosphate as Cathode for Lithium-Ion

Aug 13, 2013 - Brookhaven National Laboratory, Upton, New York 11973, United States .... Fredrick Omenya , Peter G. Khalifah , and M. Stanley Whitting...
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Layered Molybdenum (Oxy)Pyrophosphate as Cathode for LithiumIon Batteries Bohua Wen,† Natasha A. Chernova,† Ruibo Zhang,† Qi Wang,†,‡ Fredrick Omenya,§ Jin Fang,† and M. Stanley Whittingham*,†,§,∥ †

Institute for Materials Research and §Department of Chemistry, State University of New York at Binghamton, Binghamton, New York 13902-6000, United States ‡ Brookhaven National Laboratory, Upton, New York 11973, United States ∥ Northeastern Center for Chemical Energy Storage, Department of Chemistry, Stony Brook University, Stony Brook, New York 11794-3400, United States ABSTRACT: The layered structure of molybdenum (oxy)pyrophosphate (δ-(MoO2)2P2O7) was synthesized by heating MoO2HPO4·H2O precursor at 560 °C. The synthesis temperature was selected using in situ high-temperature Xray diffraction (XRD) depicting phase transformations of the precursor from room temperature up to 800 °C. Electrochemical evaluation reveals that up to four Li ions per formula unit can be intercalated into δ-(MoO2)2P2O7 upon discharge to 2 V. Three voltage plateaus are observed at 3.2, 2.6, and 2.1 V, lower than the theoretical predictions. The first plateau corresponds to the intercalation of 1.2 Li forming δLi1.2(MoO2)2P2O7, the same structure formed upon chemical lithiation with LiI. In-situ XRD indicates two-phase reaction upon the first lithium insertion and expansion of the lithiated phase unit cell in the a direction. Intercalation of the second lithium results in a different lithiated structure, which is also reversible, giving the capacity of about 110 mAh/g between 2.3 and 4 V. More lithium-ion intercalation leads to loss of crystallinity and structural reversibility. The Mo reduction upon lithiation is consistent with the amount of Li intercalated as confirmed by the Xray absorption fine structure. KEYWORDS: Li-ion batteries, cathodes, molybdenum, (oxy)pyrophosphate

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INTRODUCTION The first commercial lithium-ion batteries (LIBs) were based on the transition metal oxide cathodes, typically with layered structure, such as LiCoO2 appearing in the market in the early 1990s.1 Introduction of olivine lithium iron phosphate (LiFePO4) with enhanced safety and improved rate capability has led to a strong focus of the battery community on polyanionic compounds and especially on phosphates.2,3 However, batteries based on polyanionic compounds have much lower volumetric energy densities than those based on oxides. To increase the energy density one strategy is to consider more than one-electron transfer per redox center.4 Recent results of ab initio calculations reported by Hautier et al. systematically evaluate the properties of most redox transition metal couples, such as voltages, specific capacity, and so forth, based on thousands of known and virtual phosphate compounds.5 An inspection of their data suggests that molybdenum (Mo3+/4+, Mo4+/5+, Mo5+/6+) and vanadium (V2+/3+,V3+/4+, V4+/5+) are the only two multiple-valent elements, which can possibly enable two or more electron transfers within the acceptable voltage range (3−4.5 V) in phosphates, as illustrated in Figure 1. Iron pyrophosphates have © 2013 American Chemical Society

Figure 1. Computational prediction of the cell voltages for a wide range of transition metal phosphates. Two couples are identified in the red ovals within the electrolyte stability limit of 4.5 V (red line), modified from reference.5

Received: June 15, 2013 Revised: August 9, 2013 Published: August 13, 2013 3513

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been investigated by us and others,6,7 but the second Li falls outside the stability range of today’s electrolytes. To date, twoor more electron transfers have been demonstrated in a few vanadium phosphate structures; however, difficulties arise from the separated redox couples, which may fall outside the applicable voltage window.8−12 This inspired us to conduct research toward harnessing the three close redox couples of molybdenum as a potential cathode for LIBs. Because of the varied oxidation states of molybdenum, a myriad of molybdenum-phosphates have been reported. They consist of varied phosphate groups, including monophosphate (PO43−), diphosphate (pyrophosphate P2O74−), metaphosphate ions (PO3−), oxyphosphates, and even their mixture, for instance MoOPO4,13−16 MoP3O9,17 MoP2O7, and Mo2P4O15.18 Most structures are based on the MoO6 unit which can be isolated, bi-, trioctahedral or infinitely connected.19 Furthermore, molybdenophosphates such as A-Mo-P-O (A = Na, K, Rb, Cs) have been isolated, and their original frameworks involve either Mo3+, Mo4+, Mo5+, Mo6+ or mixed valencies.20 This group includes compounds such as K2Mo6+O2P2O7,21 A2(Mo5+O2)2P2O7(A = K, Rb),22,23 Li3Mo5+/6+3O5(PO4)3,24 Na3(Mo5+O)2(PO4)3,25 Li2Na(Mo5+O)2(PO4)3,26 AMo5+OP2O7(A = Li, Cs),27,28 A3(Mo5+O)4(PO4)5(A = K, Na),19,29 A(Mo6+O2)2(PO4)2 (A = Ba, Pb),30 AMo4+2(PO4)3(A = K, Rb), and AMo3+P2O7(A = Li, K, Na, Rb),31−33 with the most work focused on their structure. Considering the high atomic weight of Mo, only a few molybdenophosphates are viable as candidates for cathodes with more than one electron transfer. γ-(MoO2)2P2O7 with a three-dimensional (3D) structure appears to be the only molybdenophosphate investigated for lithium-ion and sodiumion batteries. Initially, this condensed structure can insert 4 lithium ions or 3 sodium ions, but it turns amorphous when charged to 4.2 V and loses capacity.34 Here we investigate a layered structure δ-(MoO2)2P2O7 (Figure 2)15 with the goal to

Article

EXPERIMENTAL SECTION

The MoO2HPO4·H2O precursor was prepared, as reported in the literature.15,16,35 First, 15 g of MoO3 (Alfa Aesar) were dissolved in 45 mL of 85% phosphoric acid (J. T. Baker) at around 180 °C. When a green solution formed, it was cooled down and refluxed with 400 mL of 16 M HNO3 (Fisher Scientific) at 140 °C. The resulting white solid was washed with deionized water and acetone and then dried in air. Blue powder of δ-(MoO2)2P2O7 was synthesized by heating the precursor MoO2HPO4·H2O to 560 °C at a rate of 0.33 °C/min, keeping it at 560 °C for 6 min, and then cooling down naturally.15 A chemically lithiated sample was synthesized by stirring δ(MoO2)2P2O7 and LiI (molar ratio 6:1) in acetonitrile for 7 days. A dark gray sample was obtained after filtering the mixture while washing with acetonitrile and then drying in a nitrogen glovebox. The structure of the samples was characterized by powder X-ray synchrotron diffraction at the National Synchrotron Light Source (NSLS) beamlines X14A and X7B, wavelengths 0.7777 and 0.3196 Å, respectively. The Rietveld refinement of the X-ray diffraction (XRD) patterns was done using the GSAS/EXPGUI package.36,37 In situ hightemperature XRD analysis was performed at beamline X7B, wavelength 0.3196 Å, with a heating rate of 5 °C/min under O2 flow. In situ XRD upon cell discharge was taken at beamline X14A, wavelengths 0.7777 Å, using 2325 coin cells with 3 mm X-ray windows sealed with Kapton tape. A current density of 22 mA/g was applied upon discharge from open circuit voltage (OCV) to 2 V, and 9 min XRD scans were taken continuously. X-ray absorption spectroscopy (XAS) experiments were performed at beamline X18A at NSLS, Brookhaven National Laboratory. A double-crystal Si(111) monochromator was used to scan X-ray energy from −200 to +1000 eV relative to Mo K edge (20,000 eV). The electrode samples loaded with 10−15 mg of active material (δ-(MoO2)2P2O7) were discharged at various conditions. The cells were disassembled in a helium glovebox, and 12 mm diameter electrode samples on an aluminum current collector were washed, dried, and press-sealed between Kapton tape. The samples were stored in the glovebox prior to being subjected to XAS detection. Fine powders of reference compounds (e.g., MoO2, MoO3) were brushed uniformly onto the Kapton tape which was then folded several times to achieve a suitable total thickness for the measurement. Transmission XAS measurements were carried out with the pure Mo metal foil measured in reference mode simultaneously for X-ray energy calibration and data alignment. X-ray absorption near edge structure (XANES) data were processed by using the Athena program of the IFEFFIT package.38 The morphology and particle size were characterized by Scanning Electron Microscopy (SEM), on a ZeissSupra field emission scanning electron microscope operating at 8.0 kV. The Fourier transform infrared spectroscopy (FTIR) were recorded using KBr pellets between 400 cm−1 and 4000 cm−1 (Perkin-Elmer Model 1600). An Inductively Coupled Plasma (ICP) test was conducted on a Varian Vista-MPX Axial ICP-OES instrument. The instrument was calibrated with NIST traceable standards. Cathodes were prepared by mixing δ-(MoO2)2P2O7 with carbon black and polyvinylidene fluoride (PVDF) in a weight ratio of 80:10:10 using 1-methyl-2-pyrrolidinone as solvent. The slurry formed was then cast onto an aluminum current collector and dried under vacuum. The dried electrodes, of area 1.2 cm2, containing 5−6 mg of active material were placed in 2325-type coin cells in a He-filled glovebox with pure lithium foil (Aldrich, thickness 0.38 mm) as the counter and reference electrodes. The electrolyte was 1 M LiPF6 (lithium hexafluorophosphate) dissolved in a mixture solution of ethylene carbonate (EC) and dimethyl carbonate (DMC) in a volume ratio of 1:1; a Celgard 2400 separator (Hoechst Celaese) was used. The electrochemical properties of δ-(MoO2)2P2O7 were evaluated using a VMP multichannel potentiostat (Biologic). The cyclic voltammetry (CV) test was done at a slow scanning rate of 0.01 mV/s while the galvanostatic charge and discharge experiments were performed at current densities of 0.04 mA/cm2, which corresponds to approximately 9 mA/g or C/40 (374 mAh/g/40 h = 9.35 mA/g).

Figure 2. Layered structure of δ-(MoO2)2P2O7. MoO6 octahedra are red and PO4 tetrahedra are blue.

examine its Mo redox activity, structural changes upon lithium cycling, and to systematically evaluate its electrochemical lithium intercalation behavior. This compound is layered along a axis, and the layers are built of chains of corner-shared MoO6 octahedra in the b-direction connected by disordered P2O7 groups.15 3514

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RESULTS AND DISCUSSION Phase Transformations of the MoO2HPO4·H2O Precursor. A high resolution synchrotron XRD pattern of MoO2HPO4·H2O is shown in Figure 3, together with the

Figure 4. SEM images of (upper) MoO2HPO4·H2O and (lower) δ(MoO2)2P2O7.

Figure 3. Synchrotron XRD pattern of MoO2HPO4·H2O precursor (wavelength 0.3196 Å) with the Rietveld refinement.

Rietveld refinement, which shows excellent fit and no impurity peaks. Because of the chain structure39 and one-dimensional morphology of the precursor, some preferred orientation was observed. The refinement was done in the P21/m space group, and the results, shown in Table 1, are close to those previously reported.16 The morphology of both the precursor MoO2HPO4·H2O and δ-(MoO2)2P2O7 are quite similar (Figure 4). Both have large particles size ∼10 μm in length and ∼2 μm in width. Higher magnification reveals that these particles consist of long, thin sticks ∼ 100 nm in diameter, which justifies the preferred orientation planes processed in Rietveld refinement. ICP results reveal the molar ratios of Mo to P of 1.03:1 for the precursor and 1.07:1 for the pristine sample. We investigated the thermal stability of MoO2HPO4·H2O precursor in oxygen atmosphere at temperatures up to 800 °C with a heating rate of 5 °C/min to find optimum conditions for the δ-(MoO2)2P2O7 formation. As illustrated in Figure 5, MoO2HPO4·H2O starts to transform to β-MoOPO4 around 250 °C. This phase shows broad peaks and may contain some amorphous component. δ-(MoO2)2P2O7 starts to form above 400 °C, and is stable up to 750 °C. Above 750 °C, γ(MoO2)2P2O7 starts to form and becomes the main phase at 800 °C. Compared with the two-dimensional film representation of powder diffraction data upon precursor heating that was reported,15 our in situ data indicates a broader stability region for the δ-(MoO2)2P2O7 with the upper limit increased to 750 °C compared to 500 °C previously reported. This discrepancy may be attributed to the purity and morphology of the samples, along with differences in high-temperature diffraction experiments. Based on the in situ XRD results, 560 °C was chosen as the optimum temperature to form the δ-(MoO2)2P2O7 phase.

Figure 5. In-situ X-ray synchrotron data (wavelength 0.3196 Å) while heating precursor MoO2HPO4·H2O under oxygen atmospheric conditions up to 800 °C with step 5 °C/min.

Figure 6 displays the XRD data of the product together with the Rietveld refinement using the structural model previously

Figure 6. Synchrotron XRD pattern of δ-(MoO2)2P2O7 (wavelength 0.7777 Å) and its Rietveld refinement.

reported.15 The XRD pattern indicates a pure phase, and similarly to the precursor, because of the layered structure and

Table 1. Lattice Parameters of MoO2HPO4·H2O and δ-(MoO2)2P2O7 at Different States of Charge compounds MoO2HPO4·H2O δ-(MoO2)2P2O7 A (2.75 V discharged) D (4 V charged) δ-Li1.2(MoO2)2P2O7 a

a

a (Å)

b (Å)

c (Å)

V (Å3)

wRp (%)

Rp (%)

6.7388(9) 16.189(2) 17.438(3) 16.239(0) 17.340(9)

6.3298(1) 3.8826(3) 3.7333(0) 3.8758(1) 3.7086(9)

7.0402(1) 6.2641(2) 6.4260(0) 6.2878(6) 6.3989(1)

281.8(9) 393.7(4) 418.3(5) 395.7(3) 411.5(2)

2.02 6.76 2.39 3.76 6.14

1.58 4.79 1.62 2.79 4.58

Note: β = 101.171°. 3515

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anodic peaks decreases dramatically, and the CV shows no obvious redox peaks beyond the second cycle confirming structural irreversibility within the 2.0 to 4.0 V range. When the lower cutoff voltage is raised from 2.0 to 2.3 V and the same current density is applied, the initial capacity achieves 115 mAh/g, corresponding to almost 1.9 lithium ions (per formula unit) being intercalated, which reduces Mo6+ to Mo5+ (Figure 8). The subsequent 20 cycles demonstrate shorter

elongated plate morphology (Figure 4), some preferred orientation was included in the refinement. The resulting lattice parameters are included in Table 1. Electrochemical Performance of δ-(MoO2)2P2O7. The electrochemical performance of δ-(MoO2)2P2O7 was evaluated in lithium cells. When δ-(MoO2)2P2O7 is cycled at a current density of 0.04 mA/cm2 (∼C/40 rate) in the voltage range of 2.0−4.0 V (vs Li+/Li) (Figure7a), the capacity reaches 240

Figure 8. Galvanostatic charge-discharge curves of δ-(MoO2)2P2O7 in the 1st, 5th, 10th, and 20th cycle; inset shows cycling performance of δ-(MoO2)2P2O7 cycled between 2.3−4.0 V at 0.04 mA/cm2.

discharge plateaus at 3.3 V, and long sloping profiles in both lithium-ion insertion and removal. Eighty percent of the discharge capacity is retained after 20 cycles (inset in Figure 8), implying this intercalation process is not completely reversible, which will be discussed later. On increasing the lower cutoff voltage to 2.75 V, the structure can insert 1.2 lithium ions initially (per formula unit), and the reversible capacity is over 70 mAh/g for 30 cycles (Figure 9a). The initial discharge plateau voltage is observed at 3.18 V. Interestingly this is slightly lower than the plateau voltage of subsequent cycles, 3.3 V. This may be due to some irreversible structural changes during the initial Li insertion or higher polarization in the initial discharge. The shape of the curves and the capacity are well maintained after the fifth cycle. A stable cycling capacity of 72 mAh/g is retained for the first 30 cycles as shown in Figure 9b. The rate capability test upon charging and discharging at current densities from C/20 to 2C between 2.75 to 4.0 V indicates good capacity retention up to C/5, while the capacities at 1C and 2C rates drop to 40 and 30 mAh/g, respectively. The capacity at C/20 after the rate test is reversed back to 60 mAh/g (Figure 9c). These results imply the reversible insertion of 1.2 lithium ions and structural stability of this layered compound between 2.75 and 4.0 V. CV curves between 2.75−4.0 V and 2.3−4.0 V are plotted in Figure 10. When the cells are cycled in the wider voltage range of 2.3−4.0 V, both cathodic and anodic processes reveal more than one redox peak with the positions of the peaks shifting and the intensity decreasing in subsequent cycles. This is consistent with the capacity fade and the decrease in the plateau capacity for the 20 cycles described in Figure 8, confirming slight structural change. However, in the 2.75−4.0 V voltage window, one main redox pair at 3.2 and 3.4 V, with small humps at 3.0 and 3.1 V, appears. Except for the initial scan with different peak position, the subsequent CV curves show very stable shape and peak positions, confirming the reversibility of reaction and structural stability for cycling between 2.75−4.0 V. This is in good agreement with the charge-discharge data. The difference between the oxidation and reduction peak positions is about 0.2 V, exhibiting low polarization of δ-

Figure 7. (a) Galvanostatic charge-discharge curves of δ(MoO2)2P2O7 at 0.04 mA/cm2; (b) CV curves of δ-(MoO2)2P2O7 at 0.01 mV/s cycled over a voltage range of 2.0−4.0 V.

mAh/g in the initial discharge process, corresponding to 3.9 lithium ions insertion per formula unit. The theoretical capacity is 62.35 mAh/g per lithium inserted into δ-(MoO2)2P2O7, and with a maximum of 6 lithium ions inserted, assuming Mo can be reduced to 3+, it adds up to 374 mAh/g. Experimental data indicates that Mo is reduced from Mo6+ to Mo4+, working as a two-electron transfer center in the investigated voltage range. Three plateaus are observed at 3.2, 2.6, and 2.1 V; however, when the cell is charged to 4 V, a sloping profile is observed, which may be due to capacitive current or change in the reaction mechanism from two-phase to single-phase because of structural change which we will discuss in detail later. The redox potentials observed in δ-(MoO2)2P2O7 are lower than the calculated ones5 and are close to those reported for γ(MoO2)2P2O7, where the first Li intercalates at just above 3 V, and the voltage decreases gradually to 2 V upon intercalation of up to 4 Li per formula unit. Both (MoO2)2P2O7 polymorphs show increased potentials in comparison with layered α-MoO3, where average reaction voltage is about 2.5 V,40 because of the inductive effect of phosphate groups. Cyclic voltammetry in a 2.0−4.0 V range shown in Figure 7b reveals consistent results with the galvanostatic chargedischarge curves. At a 0.01 mV/s scanning rate, three highintensity cathodic peaks appear during negative scanning, without clear corresponding anodic peaks in positive scanning. In the second and third scan cycles, the intensity of these 3516

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previously15 with 1.2 lithium ions inserted per formula unit as evidenced by ICP analysis. The synchrotron XRD data with Rietveld refinement of this lithiated phase referred to as δLi1.2(MoO2)2P2O7 is shown in Figure 11. The refined lattice

Figure 11. X-ray synchrotron diffraction pattern (λ = 0.7777 Å) and Rietveld refinement of δ-Li1.2(MoO2)2P2O7; inset shows the SEM image.

parameters presented in Table 1 are close to the reported values.15 Thus, with 1.2 lithium ions inserted, the a lattice parameter increases most significantly, by about 1.2 Å, indicative of the unit cell expansion in the a direction. Further increase of the molar ratio of LiI to δ-(MoO2)2P2O7 in chemical lithiation does not increase the amount of Li inserted, probably because of the small potential difference between I−/ I2 and Mo6+/Mo5+.4 The SEM image of the chemically lithiated sample (Figure 11 inset) shows similar morphology with the precursor MoO2HPO4·H2O and δ-(MoO2)2P2O7 (Figure 4); the secondary bulk particles are aggregates of one-dimensional long crystals ∼100 nm in diameter. Furthermore, we have investigated the electrochemically lithiated phases of δ-(MoO2)2P2O7 at different states of charge and discharge after disassembling the charged coin cells in a glovebox. We collected high energy synchrotron XRD (λ = 0.3196 Å) patterns at five points of the initial charge-discharge process as illustrated in Figure 12. The pattern of the material discharged to 2.75 V (point A) fit well with the reported

Figure 9. (a) Galvanostatic charge-discharge curves of δ(MoO2)2P2O7 in the 1st, 5th, 10th, and 30th cycles; (b) cycling performance and (c) rate performance of δ-(MoO2)2P2O7 cycled between 2.75 and 4.0 V at 0.04 mA/cm2.

Figure 10. CV curves of δ-(MoO2)2P2O7 in the voltage window of (a) 2.3−4.0 V and (b) 2.75−4.0 V at 0.01 mV/s.

(MoO2)2P2O7 which is less than upon cycling between 2.3 and 4.0 V as shown in Figure 10. Structural Changes in δ-Lix(MoO2)2P2O7 upon Insertion of One Li Ion. Since the electrochemical performance of δ-(MoO2)2P2O7 depends strongly upon the amount of lithium inserted, the exploration of structural changes upon chemical and electrochemical lithiation can help to understand these electrochemical phenomena. Chemical lithiation of δ-(MoO 2 ) 2 P 2 O7 using LiI in acetonitrile gives a pure lithiated phase similar to that reported

Figure 12. (top) Electrochemistry curve and (bottom) ex-situ XRD patterns (λ = 0.3196 Å) of δ-(MoO2)2P2O7 during the 1st cycle between 2.75 and 4.0 V. The upper ticks indicate XRD peaks of the starting material δ-(MoO2)2P2O7 and the lower ticks belong to δLi1.2(MoO2)2P2O7. 3517

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diffraction patterns upon discharge of δ-(MoO2)2P2O7 from OCV to 2 V and subsequent charge to 4 V (Figure 13b) together with the pattern of the material discharged to 2.3 V and then charged to 4 V (Figure 13c). The corresponding electrochemical states of each cell are indicated in Figure 13a. Obviously, 2.3 V discharged phase (E) with two lithium ions inserted is different from 2.75 V discharged material (C), which intercalated 1.2 lithium per formula unit. As highlighted in frames, peaks of 3.06 V discharged material (pattern B) appear as a mixture of pristine δ-(MoO2)2P2O7 (A) and 2.75 V discharged material (C) δ-Li1.2(MoO2)2P2O7, consistent with Figure 12. Moreover, 2.5 V (D) and 2.3 V (E) discharged patterns display a mixture of 2.75 V (C) and 2.10 V (F) discharged states. These mixtures of phases confirm the twophase behavior in the electrochemical cycling. Ex-situ patterns of the 2.10 V (F) and 2 V (G) discharged material show significantly broader peaks mainly consistent with those of 2.3 V discharged phase (E). However, the intensities of the main peaks decrease significantly, and the XRD pattern does not recover to that of pristine δ-(MoO2)2P2O7 upon charge (H, I and J), indicating the irreversible structure change. Note that upon discharge to 2.3 V and subsequent charge back to 4 V, the structure returns to that of pristine δ-(MoO2)2P2O7, indicating the reversible structural transformation in this voltage range (Figure 13c). The starred box highlights two small diffraction peaks between 15° and 16° consistently showing up in all electrochemical states except for the pristine materials. These stable peaks of currently unknown origin may suggest an irreversible phase transformation contributing to the irreversible capacity of δ-(MoO2)2P2O7. Figure 14 illustrates the in situ diffraction patterns of δ(MoO2)2P2O7 in the initial discharge process form OCV to 2 V. In-situ measurements give information about the phase during transformation, while the ex-situ results represent equilibrated material. Frames in Figure 14 highlight the gradual transformation of the main peaks, which confirms the two-phase behavior. The main peak positions of δ-(MoO2)2P2O7 maintain well, while the new peaks of δ-Lix(MoO2)2P2O7 shift. Furthermore, new peaks show up when discharged below 2.40 V, in agreement with the ex-situ data in Figure 13. We have determined the unit cell parameters of both δ(MoO2)2P2O7 and δ-Lix(MoO2)2P2O7 from the positions of

structure of the lithiated phase δ-Li1.2(MoO2)2P2O7, while the material charged to 4 V (point D) fits well with the pristine δ(MoO2)2P2O7 structure.15 XRD patterns at points B, C, and E exhibit a lithiated and pristine phase mixture, suggesting twophase behavior. Also, the XRD patterns at points C and E are similar, confirming the reversibility of the structure during cycling. Table 1 summarizes the lattice parameters of δLix(MoO2)2P2O7; the reversible expansion of the a lattice constant during lithiation is clear. Structural Changes upon Two or More Lithium Ion Insertions. Figure 13 shows a series of ex-situ synchrotron

Figure 13. (a) Electrochemistry curve and (b) ex-situ synchrotron XRD patterns of δ-(MoO2)2P2O7 discharged to 2 V and charged to 4 V (λ = 0.7777 Å). Sharp peaks at 19° belong to Al current collector. The upper ticks indicate XRD peaks of the starting material δ(MoO2)2P2O7 and the lower ticks belong to δ-Li1.2(MoO2)2P2O7. (c) Ex-situ synchrotron XRD patterns of pristine δ-(MoO2)2P2O7 and the sample discharged to 2.3 V and charged to 4 V (λ = 0.7777 Å).

Figure 14. In-situ synchrotron XRD pattern of δ-(MoO2)2P2O7 discharged to 2 V (λ = 0.7777 Å). 3518

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Figure 13b), but the transformation is not complete under the in situ cycling conditions. FTIR analysis was performed to further examine the change of bonding in MoO2HPO4·H2O and δ-(MoO2)2P2O7 upon lithiation to δ-Li1.2(MoO2)2P2O7 (Figure 16). In the precursor,

several of most intense peaks in the in situ diffraction patterns from OCV to 2.69 V (Figure 15); the a lattice parameter was

Figure 16. FTIR spectra of precursor MoO2 HPO4·H2 O, δ(MoO2)2P2O7, and chemically lithiated δ-Li1.2(MoO2)2P2O7.

peaks at 3160, 3470, and 3538 cm−1 can be attributed to the vibrations of O-H bonds, while the 1610 cm−1 band is assigned to Mo-O-H bonds. The vibrations of P-O bonds in PO4 are observed for all these compounds in the range of 850−990 cm−1 for the symmetric vibration (νs) and in the 990−1400 cm−1 range for the asymmetric vibration (νas). P-O-P νs bonds in δ-(MoO2)2P2O7 are clearly seen at 741 cm−1, and νas at 938 cm−1. Moreover, O-P-O bonds in the range of 400−645 cm−1 are observed in MoO2HPO4·H2O and δ-(MoO2)2P2O7.41 Stretching vibrations of the oxygen atoms linked to two molybdenum atoms (Mo-O-Mo) are found at 858 cm−1 in δ(MoO2)2P2O7, and the low wavenumber peak is attributed to Mo-O bonds.42 FTIR spectra of the lithiated phase show broader bands of lower intensity, which may be attributed to lower crystallinity or increased structural disorder, consistent with the diffraction results. X-ray Absorption Near-Edge Structure. XANES was studied to further evaluate the electronic structure of the Mo in δ-(MoO2)2P2O7 compound and its lithiated forms. The XANES spectra of pristine δ-(MoO2)2P2O7 as well as the electrodes discharged to 2.75 V, 2.30 V, and 2.0 V are depicted in Figure 17a; also shown are the spectra from reference compounds Mo6+O3 and Mo4+O2. Two prominent features (labeled “A” and “B”) characterizing the valence state and local geometry of Mo are highlighted in the spectra and will be discussed in this section. The pre-edge structure labeled with “A” arises mainly from the electronic excitation of Mo 1s→4d, which shares the similar spectroscopic properties of 1s→3d transitions for 3d elements. Such features are formally forbidden while gaining the intensity by p-d hybridization in the final states as a result of noncentrosymmetric environment around the central metal.43 A distinct pre-edge peak is observed in pristine δ-(MoO2)2P2O7 while the feature becomes “silent” upon 4 Li intercalation. The peak can be observed at 1.2 Li and 1.9 Li insertion, though with much decreased intensities. Such variation of the pre-edge feature in the molybdenum oxyphopyrophosphate series resembles the trend demonstrated in the molybdenum oxide references (Figure 17a). It has been recognized that the octahedral oxygen coordination is greatly distorted in the MoO3 compound: the four shorter Mo-O bonds (1.64 Å−1.96 Å) form a tetrahedron-like substructure, while the other two bonds are much longer (2.24 Å and 2.31 Å). This significantly

Figure 15. Lattice parameters of δ-(MoO2)2P2O7 (black squares) and δ-Lix(MoO2)2P2O7 (red circles) from in situ XRD upon discharge from OCV to 2.69 V. Open symbols in the top plot show a lattice parameters calculated from (200) and (400) peak positions.

also independently determined from the (200) and (400) reflections. It is clear that the lattice parameters of δ(MoO2)2P2O7 do not change much, while the lattice parameter a of δ-Lix(MoO2)2P2O7 increases significantly as more lithium is intercalated. Such lattice parameter change during two-phase reaction is, possibly, an artifact of the in situ measurement, where a nonequilibrium transient state of newly formed structure is observed. Upon further lithiation, a significant shift of the (400) peak toward lower angle is observed, indicating a further increase of interlayer spacing. The new peaks appearing at the 2.4 V state of discharge are consistent with those observed in the ex-situ experiment (point E in 3519

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accuracy of the Rietveld refinement complicated by the preferred orientation and pyrophosphate disorder.15 Also, the presence of amorphous phases, which can only be observed in the X-ray absorption data, can contribute to this discrepancy. The energy (position) of the Mo K-edge (label “B” in Figure 17a), a parameter sensitive to the oxidation states, was inspected very closely as well. The pristine compound exhibits an absorption edge at a similar energy to that of MoO3, suggesting the Mo is approximately hexavalent in δ(MoO2)2P2O7 as expected from the chemical composition. Mo K-edge absorption position shifts toward lower energies when lithium is intercalated into the structure. Specifically, the peak position for the samples discharged to 2.75 and 2.30 V resides between those of MoO2 and MoO3; and shifts further to an energy lower than that in MoO2 for the sample discharged to 2.00 V. One can conclude that the Mo6+ in δ-(MoO2)2P2O7 is electrochemically reduced as discharging voltage is applied and may reach an oxidation state lower than 4+ as ∼3.9 lithium is inserted. An approximate linear correlation has previously been established for the Mo K absorption edges and chemical charges in the series of Mo reference compounds.46,47 In the current study, the Mo chemical valences at various stages of lithium intercalation were estimated by adopting this approach with the reference compounds including Mo0 metal, Mo4+O2, and Mo6+O3. The method and reference choices were considered sufficiently validated, as 6 oxygen atoms construct the first coordination shell around Mo in the δ-(MoO2)2P2O7 structure and its Li-inserted forms. Good linear correlation was derived between valence states of reference compounds and the corresponding edge energies, for which the first derivative peak of main absorptions were used. From this linear dependence the chemical valences were determined to be approximately 5.8+ for the pristine δ-(MoO2)2P2O7; and 5.1+, 4.8+, 3.7+ for the samples discharged to 2.75 V, 2.30 V, and 2.00 V respectively (Figure 17b). The results are consistent with the electrochemical analysis which reveals the amount of intercalated lithium to be ∼1.2, 1.9, and 3.9, respectively. The Mo oxidation state was also calculated for the chemically lithiated sample and is found to be 5.1+, comparable with that of the sample electrochemically discharged to 2.75 V, which intercalates 1.2 lithium in the structure. The structural and chemical equivalence of δ-Li1.2(MoO2)2P2O7 produced by chemical lithiation and 2.75 V electrochemically discharged is further supported by the observed similarity in Mo K-edge XANES spectra (data not shown). This analysis confirms that the electrochemical lithium intercalation in δ-(MoO2)2P2O7 within acceptable potentials involves a ∼2-electron redox process, and up to 4 lithiums can be inserted, while the chemical reaction saturates upon insertion of only 1.2 lithium.

Figure 17. (a) Mo K-edge XANES spectra of pristine and electrochemically lithiated (2.75 V, 2.30 V, 2.0 V) δ-(MoO2)2P2O7 together with reference compounds MoO3 and MoO2; (b) linear correlation between average Mo valence states of reference compounds and the corresponding edge energies together with the edge energies of pristine and electrochemically lithiated δ(MoO2)2P2O7 and their calculated valence states.

enhances the Mo s-d excitation, thus resulting in the strong preedge feature in Mo K-edge XANES. Such a pre-edge peak is not observed in MoO2 in contrast, for the reason that the rutile-like structure of this crystal consists of lower distortion in its MoO6 octahedral coordinations.44 In the case of δ-(MoO2)2P2O7, as evidenced by the crystallographic data,15 the pristine material features a highly distorted MoO6 octahedron in its first coordination shell because of the significant differences among the six Mo-O bonds: four equatorial Mo-O bonds are between 1.93 Å and 1.95 Å, forming a square-planar-like structure, while the two axial bonds are 1.67 Å and 2.08 Å, resulting in strongly distorted octahedral coordination. Such local structural geometry explains well the presence of the pre-edge peak in the pristine δ-(MoO2)2P2O7. While the observed decrease in pre-edge intensity of XAS suggests that on average the noncentrosymmetry of MoO6 is alleviated upon lithium intercalation, the crystallographic refinement in reference 15 indicates otherwise. In comparison to δ-(MoO2)2P2O7, besides the change of axial Mo-O bonds to 2.14 Å and 1.71 Å, the equatorial Mo-O bonds in the Li2(MoO2)2P2O7 contract and elongate collinearly. The two pairs of 1.87 Å and 2.07 Å bonds deform the equatorial quasi-“square-planar” structure present in δ-(MoO2)2P2O7 to a parallelogram in its chemically lithiated counterpart, which adds further distortion to the MoO6 octahedron in the structure. On the contrary, the evolution of local lattice upon δ-(MoO2)2P2O7 lithiation suggests elongation of the average Mo-O bond distance, which may induce the decrease of pre-edge peak intensity, since the wave functions are exponentially decaying with distance.45 The discrepancy between the long-range (XRD) and local (X-ray absorption) structural data may be attributed to the limited

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CONCLUSIONS The layered structure of δ-(MoO2)2P2O7 was synthesized by heating the MoO2HPO4·H2O precursor at 560 °C. δ(MoO2)2P2O7 can reversibly intercalate 1.2 lithium per formula in a 2.75−4.0 V voltage window with a capacity of 70 mAh/g, and 1.9 lithium (110 mAh/g) between 2.3 and 4.0 V with little capacity fading over 50 cycles. The electrochemical tests reveal better reversibility when only one lithium ion is intercalated, though XRD shows structural reversibility for both. Synchrotron XRD demonstrates two different patterns for the 2.3 and 2.75 V discharged samples indicating that two different lithiated structures are formed. When more than two lithium ions are inserted upon discharge to 2 V, an irreversible crystallinity loss 3520

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(12) Ellis, B. L.; Ramesh, T. N.; Davis, L. J. M.; Goward, G. R.; Nazar, L. F. Chem. Mater. 2011, 23, 5138−5148. (13) Kierkegaard, P.; Westerlund, M. Acta Chem. Scand. 1964, 18, 2217−2225. (14) Kierkegaard, P.; Longo, J. M. Acta Chem. Scand. 1970, 24, 427− 432. (15) Lister, S. E.; Rixom, V. J.; Evans, J. S. O. Chem. Mater. 2010, 22, 5279−5289. (16) Rangan, K. K.; Gopalakrishnan, J. Inorg. Chem. 1996, 35, 6080− 6085. (17) Watson, I. M.; Borel, M. M.; Chardon, J.; Leclaire, A. J. Solid State Chem. 1994, 111, 253−256. (18) Lister, S. E.; Evans, I. R.; Evans., J. S. Inorg. Chem. 2009, 48, 9271−9281. (19) Hoareau, T.; Leclaire, A.; Borel, M. M.; Grandin, A.; Raveau, B. J. Solid State Chem. 1995, 114, 61−65. (20) Leclaire, A.; Borel, M. M.; Chardon, J.; Raveau, B. J. Solid State Chem. 1995, 116, 364−368. (21) Zid, M. F.; Driss, A.; Jouini, T. Acta Crystallogr., Sect. E: Struct. Rep. Online 2003, 59, I65−I67. (22) Gueno, C.; Borel, M. M.; Grandin, A.; Leclaire, A.; Raveau, B. J. Solid State Chem. 1993, 104, 202−208. (23) Guesdon, A.; Leclaire, A.; Borel, M. M.; Grandin, A.; Raveau, B. Acta Crystallogr., Sect. C: Cryst. Struct. Commun. 1994, 50, 1852−1854. (24) Ledain, S.; Leclaire, A.; Borel, M. M.; Provost, J.; Raveau, B. J. Solid State Chem. 1997, 133, 391−399. (25) Ledain, S.; Leclaire, A.; Borel, M. M.; Raveau, B. J. Solid State Chem. 1997, 132, 249−256. (26) Ledain, S.; Leclaire, A.; Borel, M. M.; Raveau, B. J. Solid State Chem. 1997, 129, 298−302. (27) Ledain, S.; Borel, M. M.; Leclaire, A.; Provost, J.; Raveau, B. J. Solid State Chem. 1995, 120, 260−267. (28) Guesdon, A.; Borel, M. M.; Leclaire, A.; Grandin, A.; Raveau, B. J. Solid State Chem. 1993, 108, 46−50. (29) Leclaire, A.; Hoareau, T.; Borel, M. M.; Grandin, A.; Raveau, B. J. Solid State Chem. 1995, 114, 543−549. (30) Masse, R.; Averbuch-Pouchot, M. T.; Durif, A. J. Solid State Chem. 1985, 58, 157−163. (31) Ledain, S.; Leclaire, A.; Borel, M. M.; Raveau, B. Acta Crystallogr., Sect. C: Cryst. Struct. Commun. 1996, C52, 1593−1594. (32) Leclaire, A.; Borel, M. M.; Grandin, A.; Raveau, B. J. Solid State Chem. 1988, 76, 131−135. (33) Riou, D.; Leclaire, A.; Grandin, A.; Raveau, B. Acta Crystallogr., Sect. C: Cryst. Struct. Commun. 1989, 45, 989−991. (34) Uebou, Y.; Okada, S.; Yamaki, J.-i. J. Power Sources 2003, 115, 119−124. (35) Kierkegaard, P. Acta Chem. Scand. 1958, 12, 1701−1714. (36) Toby, B. H. J. Appl. Crystallogr. 2001, 34, 210−213. (37) Larson, A. C.; Von Dreele, R. B. General Structure Analysis System (GSAS), Los Alamos National Laboratory Report LAUR 86748; Los Alamos National Laboratory: Los Alamo, NM, 2000. (38) Ravel, B.; Newville, M. J. Synchrotron Radiat. 2005, 12, 537− 541. (39) Gilson, T. R.; Weller, M. T. Inorg. Chem. 1989, 28, 4059−4061. (40) Whittingham, M. S. Prog. Solid State Chem. 1978, 12, 41−99. (41) Kim, H.; Lee, S.; Park, Y. U.; Kim, H.; Kim, J.; Jeon, S.; Kang, K. Chem. Mater. 2011, 23, 3930−3937. (42) Mai, L.; Hu, B.; Chen, W.; Qi, Y.; Lao, C.; Yang, R.; Dai, Y.; Wang, Z. L. Adv. Mater. 2007, 19, 3712−3716. (43) Yamamoto, T. X-Ray Spectrom. 2008, 37, 572−584. (44) Lützenkirchen-Hecht, D.; Frahm, R. J. Phys. Chem. B 2001, 105, 9988−9993. (45) Bunker, G. Introduction to XAFS: A practical Guide to X-ray Aborption Fine Structure Spectroscopy; Cambridge University Press: Cambridge, U.K., 2010. (46) Cramer, S. P.; Eccles, T. K.; Kutzler, F. W.; Hodgson, K. O. J. Am. Chem. Soc. 1976, 98, 1287−1288. (47) Ressler, T.; Wienold, J.; Jentoft, R. E.; Neisius, T. J. Catal. 2002, 210, 67−83.

occurs. In-situ XRD indicates two-phase reaction upon insertion of the first Li ion, accompanied by expansion of δLi1.2(MoO2)2P2O7 unit cell in the a direction. The X-ray absorption fine structure study reveals that molybdenum in assynthesized δ-(MoO2)2P2O7 is nearly hexavalent (5.8+). Insertion of two lithium ions reduces Mo in the structure to 5+, while the average Mo oxidation state can be reduced to lower than 4+ when as many as four lithium ions are intercalated. These results confirm that molybdenum-based phosphate compounds can provide a multielectron redox center; however, the redox potential drops fast to below 3 V upon Mo reduction to 5+, and the structure of δ-(MoO2)2P2O7 does not tolerate more than two lithium ions per formula unit becoming amorphous. For a successful use of both Mo6+/5+ and Mo5+/4+ redox pairs a different structure has to be found, with (1) the maximized inductive effect of the PO4 groups and (2) stability toward reversible intercalation of more than 1 Li per Mo. Meeting both conditions seems to be challenging, as the inductive effect is maximized in the structures with edge-sharing polyhedra, as in LiFePO4, which is not typical for Mo phosphates. At the same time, open structures, capable of providing sites for multiple Li insertion tend to provide lower voltages because of the smaller free energy difference between the lithiated and delithiated states.

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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 This research is supported as part of the Northeastern Center for Chemical Energy Storage, an Energy Frontier Research Center funded by the U.S. Department of Energy, Office of Science, Basic Energy Sciences under Award Number DESC0001294. Use of the National Synchrotron Light Source at Brookhaven National Laboratory is supported by the U.S. Department of Energy, Office of Science, Basic Energy Sciences, under Contract No. DE-AC02-98CH10886.

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REFERENCES

(1) Whittingham, M. S. Chem. Rev. 2004, 104, 4271−4301. (2) Whittingham, M. S.; Song, Y.; Lutta, S.; Zavalij, P. Y.; Chernova, N. A. J. Mater. Chem. 2005, 15, 3362−3379. (3) Padhi, A. K.; Nanjundaswamy, K. S.; Goodenough, J. B. J. Electrochem. Soc. 1997, 144, 1188−1194. (4) Murphy, D. W.; Christian, P. A. Science 1979, 205, 651−656. (5) Hautier, G.; Jain, A.; Ong, S. P.; Kang, B.; Moore, C.; Doe, R.; Ceder, G. Chem. Mater. 2011, 23, 3495−3508. (6) Zhou, H.; Upreti, S.; Chernova, N. A.; Hautier, G.; Ceder, G.; Whittingham, M. S. Chem. Mater. 2011, 23, 293−300. (7) Nishimura, S.; Nakamura, M.; Natsui, R.; Yamada, A. J. Am. Chem. Soc. 2010, 132, 13596−13597. (8) Gaubicher, J.; Wurm, C.; Goward, G.; Masquelier, C.; Nazar, L. Chem. Mater. 2000, 12, 3240−3242. (9) Yin, S. C.; Grondey, H.; Strobel, P.; Anne, M.; Nazar, L. F. J. Am. Chem. Soc. 2003, 125, 10402−10411. (10) Patoux, S.; Wurm, C.; Morcrette, M.; Rousse, G.; Masquelier, C. J. Power Sources 2003, 119−121, 278−284. (11) (a) Song, Y.; Zavalij, P. Y.; Whittingham, M. S. J. Electrochem. Soc. 2005, 152, A721−A728. (b) Chen, Z.; Chen, Q.; Chen, L.; Zhang, R.; Zhou, H.; Chernova, N. A.; Whittingham, M. S. J. Electrochem. Soc. 2013, 160, A1777−A1780. 3521

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