ARTICLE pubs.acs.org/JPCC
Experimental and Theoretical Investigations of Dimethylacetamide (DMAc) as Electrolyte Stabilizing Additive for Lithium Ion Batteries Mengqing Xu,*,†,‡ Liansheng Hao,†,‡ Yanlin Liu,†,‡ Weishan Li,*,†,‡ Lidan Xing,†,‡ and Bin Li†,‡ † ‡
School of Chemistry & Environment, South China Normal University, Guangzhou 510006, China Key Laboratory of Technology on Electrochemical Energy Storage and Power Generation in Higher Education Guangdong Institutes, South China Normal University, Guangzhou 510006, China ABSTRACT: Dimethylacetamide (DMAc) is used as an electrolyte stabilizing additive for lithium ion battery. The effects of DMAc on the enhancements of electrolyte thermal stability and the solid electrolyte interphases (SEIs) on graphite anode and LiFePO4 cathode were investigated via a combination of electrochemical methods, nuclear magnetic resonance (NMR), Fourier transform infrared-attenuated total reflectance (FTIRATR), as well as X-ray photoelectron spectroscopy (XPS). It was found that 1.0 M LiPF6 EC/DMC/DEC (1/1/1,weight ratio) electrolyte with 1% DMAc incorporation can be stable at 85 °C for over 6 months without precipitation and color change. In addition, the addition of 1% dimethylacetamide (DMAc) can significantly improve the cyclic performance of a LiFePO4/graphite cell at elevated temperature. These improved performances are ascribed to the enhancement of the thermal stability of the electrolyte and the modification of SEI components on graphite anode and LiFePO4 cathode. The explicit working mechanism of DMAc stabilizing LiPF6-based electrolyte is also discussed by the density functional theory (DFT) calculations.
1. INTRODUCTION Lithium-ion batteries (LIBs) have been widely used for portable consumer electronics for their high energy density, high voltage, and long cycle life.1 However, several problems including limited operating temperature and loss of power and capacity upon storage or prolonged use limit the application for large scale power sources, such as hybrid electronic vehicles (HEV) or plugin hybrid electronic vehicles (PHEV).2-4 There are several factors that limit the thermal stability of LIB, and it is generally acknowledged that the thermal stability of electrolyte at elevated temperature and the reactions of the electrolyte with the surface of the electrode materials are frequently reported to be the most important.5-7 A typical lithium-ion battery consists of a graphite anode, a transition metal oxide (such as LiMn2O4, LiCoO2, LiFePO4, etc.) cathode, and a nonaqueous organic electrolyte, which acts as an ionic conductor between electrodes and separates the two electrode materials. The electrolytes used in commercial lithiumion batteries are prepared by dissolving LiPF6 into binary or ternary organic carbonates or ethers, including dimethyl carbonate (DMC), diethyl carbonate (DEC), ethylene carbonate (EC), as well as 1,3-dioxane.8,9 However, the LiPF6-based electrolytes have poor thermal stability due to the decomposition of LiPF6 at moderate temperatures (>55 °C). Previous studies have investigated the mechanisms of thermal decomposition of LiPF6-based electrolytes.10-12 It is generally accepted that PF5 will be generated during the thermal decomposition of LiPF6, and PF5 reacts rapidly with trace protic impurities in the electrolyte, such as ROH or H2O, to form HF r 2011 American Chemical Society
and OPF3, which then initiates an autocatalytic decomposition of the electrolyte. The olivine structure LiFePO4 has been widely investigated as a promising cathode material to substitute LiCoO2 for its low price, rich source, and high cyclic stability. These advantages make it attractive for developing large-scale lithium ion batteries for hybrid electronic vehicles (HEV) or electronic vehicles (EV). However, LiFePO4 is not stable in LiPF6-based electrolyte because iron trends to dissolve in the electrolyte due to the presence of HF. Therefore, it is essential to develop stabilizing additives to inhibit the detrimental thermal reactions of the electrolyte and the dissolving of iron in the electrolyte. Lewis base such as dimethylacetamide (DMAc) was first reported to be used as a stabilizing additive by Lucht’s research group.13,14 The thermal reactions of electrolyte with electrode materials were significantly inhibited through introduction of DMAc into the electrolyte. Basically, they acknowledged that DMAc can capture Lewis acidic PF5 generated during the thermal dissociation of LiPF6. However, some specific questions are not clear yet. For example, why does Lewis acidic PF5 prefer to complex with DMAc rather than attack the carbonate solvents although it is at a much lower concentration in the electrolyte? And to our knowledge, the applications of stabilizing additive in LiFePO4-based cells to improve the elevated temperature are rarely reported. In this work, the effects of DMAc on the thermal Received: October 5, 2010 Revised: January 10, 2011 Published: March 04, 2011 6085
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stability of carbonate electrolyte and the explicit working mechanism of DMAc complex with PF5, as well as thermal reactions of electrolyte containing DMAc with electrode materials are investigated via combination of nuclear magnetic resonances (NMR), X-ray photoelectron spectroscopy (XPS), Fourier transform infrared-attenuated total reflectance (FTIR-ATR), as well as density functional theory (DFT).
2. EXPERIMENTAL SECTION Battery-grade carbonates were purchased from Tinci Company (Guangzhou Tinci Materials Technology Co. Ltd., China) and used without further purification. Battery-grade lithium hexafluorophosphate (LiPF6) was obtained from Hashimoto Chemical Corporation and used without further purification. Dimethylactemide (DMAc) was purchased from Acros Organic Chemical Co. and used without further purification. The composition of the electrolyte used in all of the cells was a 1.0 M LiPF6 solution in EC/DMC/DEC (1/1/1, weight ration) (STD) with and without DMAc additive (1% weight ratio). Water and free acid (HF) contents in the electrolyte were controlled to be less than 20 and 50 ppm, respectively, determined by Karl Fischer 831 Coulometer (Metrohm) and Karl Fischer 798 GPT Titrino (Metrohm). Samples for NMR spectroscopy were prepared in a glovebox filled with high purity argon followed by flame sealing under reduced pressure. The sealed samples were heated in a silicon oil bath of 85 °C. Samples were weighed before and after storage to confirm seal. NMR analyses were conducted on a JEOL 400 MHz NMR spectrometer. Chemical shifts are referenced to ethylene carbonate (EC) at 4.51 ppm, and LiPF6 at 65.0 ppm, for 1H, and 19F resonances, respectively. Pouch cells were fabricated with LiFePO4 cathodes and graphite anodes as active materials and Celgard 2325 separators. The anode electrodes contained 95% graphite, 1% acetylene carbon black, 1.5% carboxymethyl cellulose (CMC) and 2.5% styrene butadiene rubber (SBR); cathode electrodes contained 90% LiFePO4, 5% acetylene carbon black, and 5% binder. The cells were cycled with a constant-current-constant-voltage charge and a constant-current discharge of 1 C rate between 3.85 and 2.0 V using BS-9300R type battery cycle (Guangzhou, China). The graphite anodes and LiFePO4 cathodes were washed with anhydrous DMC solvent 3 times to remove residual EC and LiPF6 salt followed by vacuum drying overnight at room temperature. The XPS spectra were carried out with ESCALAB 250 (Thermo Fisher Scientific), using a focused monochromatized Al KR radiation (hυ = 1486.6 eV). Lithium was not monitored due to its low inherent sensitivity and small change of binding energy. Calibration of XPS peak position was made by recording XPS spectra for reference compounds, which would be presented on the electrode surfaces: LiF, Li2CO3, LixPOyFz and lithium alkyl-carbonate. The hydrocarbon contamination peak at 284.8 eV was used as a reference for the final adjustment of the energy scale in the spectra. The spectra obtained were fitted using XPS peak software (version 4.1). Line syntheses of elemental spectra were conducted using a Guassian-Lorentzain (80:20) curve fit with Shirley background subtraction. The element concentration was calculated based on the equation Cx = (Ix/ Sx)/(∑Ii/Si), where Ix is the intensity of the relative element and Si is the sensitivity number of the element. Fourier transform infrared-attenuated total reflectance (FTIR-ATR) analysis of the anode and cathode electrodes was carried out with a
Figure 1. Linear sweeps of glass carbon electrode in 1.0 M LiPF6 EC/ DMC/DEC (1/1/1) with 2% DMAc electrolyte, scan rate 5 mV/s.
Bruker Tensor 27 spectrometer, with germanium crystal and 4 cm-1 resolution and 64 scans. All calculations are performed on Guassian 03 package.15 The equilibrium state structures are optimized by B3LYP method16 at 6-311þþG (d,p) basis set.17 To investigate the role of solvent effects, the equilibrium state structures are optimized by using the polarized continuum model (PCM).18 To confirm each optimized stationary point and make zero-point energy (ZPE) corrections, frequency analyses are done with the same basis set. Charge distribution is analyzed by the natural bond orbital (NBO) theory.
3. RESULTS AND DISCUSSION 3.1. Electrochemical Window. Figure 1 shows the electrochemical window of 1.0 M LiPF6/EC þ DMC þ DEC (1:1:1, weight ratio) with 2% DMAc electrolyte obtained by measuring the electrochemical reduction and oxidation of electrolyte on glassy carbon (GC) electrode, respectively. With respect to the cathodic sweep, the electrolyte decomposes at 0.6 V (vs Li/Liþ) corresponding to the reduction of electrolyte, a typical reduction reaction of EC-based electrolyte. As the potential becomes more negative, the reduction current increases, which corresponds to an increase of decomposition of electrolyte. With respect to the anodic sweep, the oxidative current was observed at 4.7 V (vs Li/Liþ), which is consistent with the decomposition of EC-based electrolyte on an inert electrode at high potential. This suggests that DMAc is stable at GC electrode between 0.6 to 4.7 V (vs Li/Liþ), showing a wide electrochemical window. 3.2. Cycling Performance. Pouch cells were constructed containing STD and STD with 1% DMAc electrolytes. Figure 2 shows cyclic performances of LiFePO4/graphite cells with STD and STD with 1% DMAc electrolytes at 60 °C. Cells with STD electrolyte show a slightly higher initial discharge capacity than the cells with 1% DMAc-containing electrolyte. This slight difference is most likely related to the discrepancy of expected capacity of the cells since DMAc displays a wide electrochemical window as discussed above. However, cells containing STD electrolyte experience much greater fast capacity fading than the cells with 1% DMAc addition electrolyte. Cells with STD electrolyte experience a discharge capacity drop from 448.0 mAh 6086
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Figure 2. Cycling performance of LiFePO4/graphite cells with and without DMAc (1%, weight ration) electrolyte at 60 °C.
on the first cycle to 246.4 mAh on the 50th cycle at 60 °C. The capacity retention is about 55.0% after 50 cycles. When STD with 1% DMAc electrolyte is used the capacity drops from 427.9 mAh on the first cycle to 373.2 mAh at 50th cycle at 60 °C, or about 87.2% after 50 cycles. This suggests that addition of 1% DMAc can significantly improve the cycling stability at elevated temperature. Thermal stability of the electrolytes with and without DMAc, as well as surface analyses of electrodes extracted from cells with and without DMAc, as described below, provides insight into the sources of the difference in capacity retention at elevated temperature. 3.3. Thermal Stability of LiPF6/Carbonate Electrolyte Containing DMAc. The thermal stability of LiPF6 in carbonate solvents has been reported in previous literatures.10-12 Storage of LiPF6-base carbonate electrolytes at elevated temperatures results in the thermal dissociation of LiPF6 to LiF and PF5. The strong Lewis acidic PF5 reacts rapidly with trace protic impurities in the electrolyte, such as ROH or H2O, to form OPF3, which then initiates an autocatalytic decomposition of the electrolyte. Previous investigations of LiPF6/carbonate electrolytes indicate nearly quantitative decomposition upon storage for several days at 85 °C, forming a number of byproducts, including CO2, alkyl ethers (R2O), alkyl fluorides, phosphorus oxyfluoride (OPF3), and fluorophosphates [OPF2OR, OPF(OR)2].10-12,19 However, upon incorporation of 1% DMAc to LiPF6/carbonate electrolyte the thermal decomposition is dramatically inhibited. Figure 3 shows the 1H and 19F NMR spectra of 1.0 M LiPF6 EC/DMC/DEC with 1% DMAc before and after thermal storage. Besides the resonances of carbonate solvents, three small peaks were also presented in Figure 3a, consistent with the resonances of protons in DMAc molecule. As respect to the 19F spectrum before thermal storage, besides the resonances of PF6-, additional resonances (δ = 54.4 ppm, d, JP-F = 960) were presented in Figure 3b. The additional resonances can be assigned to the OPF2 (OLi) due to the trace amount of OPF3, which is inevitably formed by PF5 reacting with trace protic impurities in the electrolyte. 1H and 19F NMR spectra after thermal storage at 85 °C for 2 months are shown in Figure 3, parts c and d, respectively. Analysis of 1H NMR spectrum suggests that no decomposition of carbonate solvents took place after thermal storage at 85 °C for 2 months, consist with the 13C
Figure 3. 1H and 19F NMR spectra of carbonate solvents with DMAccontaining electrolyte before thermal storage (a, b) and after thermal storage (c, d) at 85 °C for 2 months.
NMR spectrum and GC-MS results. As respect to the 19F NMR spectrum, only small set of resonances characteristic of OPF3 (19F, δ = 49.6 ppm, d, JP-F = 1116), LiF/FH (19F, ∼-10 ppm) are present after 2 months storage at 85 °C. Continued storage at 85 °C for over 6 months only results in small changes to the spectra suggesting that the thermal decomposition reactions are dramatically inhibited. 3.4. XPS Analysis of Graphite Anodes. In order to better understand the sources of superior capacity retention of cells containing 1% DMAc electrolyte, the cells after 50 cycles at 60 °C were disassembled in an argon-filled glovebox and the surfaces of the electrodes were analyzed by XPS. Analyses of the anodes reveal significant differences in the surface species (Figure 4, Table 1). The concentration of C is dramatically decreased for both cycled anodes while the concentrations of O, F, and P are increased compared to the fresh anode. This suggests the active electrode material is covered by a SEI film. The concentrations of C and O are higher for the anode cycled with 1% DMAc-containing electrolyte than those for the anode with STD electrolyte. However, the concentrations of F and P are much lower for the anode cycled with 1% DMAc-containing electrolyte than those for the anode with STD electrolyte. The C 1s spectrum of the fresh anode consists of three peaks. The first one at 284.3 eV is attributed to graphite.20,21 The second one at 285.0 eV is assigned to SBR binder and also to hydrocarbon contamination, while the peak observed at 286.7 eV is attributed to the C-O like carbon atoms in the CMC binder.22 After 50 cycles at 60 °C, the C 1s spectrum of the anode cycled with STD electrolyte contains three peaks characteristic of graphite (284.3 eV), C-O (286.3 eV), CdO bonds in lithium 6087
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Figure 4. C 1s, O 1s, F 1s, and P 2p XPS spectra of fresh graphite anode (top) and cycled with STD (middle) and STD þ 1% DMAc (weight ration) electrolyte (bottom) at 60 °C after 50 cycles.
alkyl carbonates and polycarbonates (289 eV). The peak at 285 eV corresponding to SBR binder has disappeared, which suggests the electrode surface has been covered by a thickness of SEI film. The C 1s spectrum of the anode cycled with DMAc-containing electrolyte contains peaks at 286-288 eV, consistent with the presence of C-O bonds in esters and carbonates, CdO bonds in lithium alkyl carbonates and polycarbonates (290 eV), as well as graphite (284.3 eV). The O 1s spectrum of the fresh anode consists of two main peaks assigned to the oxygen atoms of the CMC binder. For the anode cycled with STD electrolyte and DMAc-containing electrolyte at 60 °C, we can observe that the enhancement of the O 1s components of the CMC binder, suggesting the deposition of new oxygenated species, also supported by the new C 1s characterization. The O 1s spectrum of the anode cycled with the STD electrolyte contains a broad peak characteristic of C-O and CdO containing components at 533.7 and 532.5 eV, respectively. This suggests the deposition of carbonate salt, lithium alky carbonates ROCO2Li, or polycarbonates on the surface of the electrode, resulting from the decomposition of the solvents and salt.23 Several formation mechanisms of the carbonate species can be found in the literature.24,25In addition, a new peak at 531.5 eV, characteristic of Li2CO3, was also observed in the O 1s spectrum, suggesting the graphite anode was covered by inorganic component. The O 1s spectrum of the anode cycled containing STD with 1% DMAc electrolyte is slightly different from that of the anode containing STD electrolyte. Except the C-O and CdO peaks, the Li2CO3 peak was not detected in O 1s spectrum, which suggests that the graphite anode cycled with DMAc-containing electrolyte covered by much more organic components than the electrode cycled with STD electrolyte at elevated temperature.
Table 1. Surface Concentrations of Different Elements on Fresh Anode and Anode from Cycled Cells with STD Electrolyte and with DMAc-Containing Electrolyte sample
C 1s (%)
O 1s (%)
F 1s (%)
P 2p (%)
fresh
82.8
14.7
2.5
STD electrolyte
32.4
29.8
35.8
2.0
STD þ 1% DMAc
38.9
56.4
3.8
1.0
The F 1s spectrum of the anode cycled with STD electrolyte at 60 °C shows three main peaks. The peak at 684.5 and 685.9 eV, attributed to LiF and LixPOyFz, respectively, decomposition products of LiPF6 commonly observed at electrode/ electrolyte interface at elevated temperature. An additional peak in F 1s at 686.8 eV is attributed to the remaining salt LiPF 6 , despite washing the electrode with DMC before XPS analysis. The F 1s spectrum of the anode extracted from a cell containing STD with 1% DMAc is similar but the intensity of LiF is much lower than that of the anode cycled with the STD electrolyte, which is consistent with the increase in concentration of F for STD cells as showed in Table 1. In addition, the peak at 686.8 eV characteristic of LiPF6 is not present in the F 1s spectrum. The P 2p spectrum of the graphite anode cycled in the STD electrolyte has two main peaks at about 133.7 and 136.9 eV, the peak at 133.7 eV is characteristic of LixPOyFz, and the peak at 136.9 eV could be assigned to LiPF6 precipitated on the electrode surface. However, as respect to P 2p spectrum of the electrode cycled with STD and 1% DMAc-containing electrolyte, only one peak at 133.6 eV was observed, characteristic of LixPOyFz. The results obtained from P 2p spectra are consistent well with the F 1s spectra. 6088
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Figure 5. C 1s, O 1s, F 1s, and P 2p XPS spectra of fresh LiFePO4 cathode (top) and cycled with STD (middle) and STD þ 1% DMAc (weight ratio) electrolyte (bottom) at 60 °C after 50 cycles.
3.5. XPS Analysis of LiFePO4 Cathodes. The fresh LiFePO4 cathode is characterized by peaks corresponding to the metal oxide, 531.4 eV in O 1s spectrum, 133.5 eV in P 2p spectrum, as well as Fe 2p spectrum. The Fe 2p spectra are not shown here since not too much difference was observed whether the cathode cycled with STD electrolyte or with 1% DMAc-containing electrolyte. The Fe 2p spectrum is split into two parts due to spin-orbit coupling (Fe 2p3/2 and Fe 2p1/2) with an intensity ration of about 2/1.23 The peak at 284.3 eV in C 1s is attributed to the black carbon. An additional peak at ∼285 eV in C 1s of the fresh LiFePO4 is also present, characteristic of the binder. Analysis of the cathode surface reveals significant differences between the cells cycled with STD electrolyte and DMAc-containing electrolyte after 50 cycles at 60 °C (Figure 5, Table 2). Compared to the fresh cathode electrode, the concentrations of C and Fe are dramatically decreased for both cycled cathodes, while the concentrations of O, F and P are significantly increased. This suggests that the LiFePO4 cathode surface is also covered by a SEI layer. The C 1s spectra show slight differences between the cathode cycled with STD electrolyte and STD with 1% DMAc-containing electrolyte. The C 1s spectrum of the electrode cycled in STD electrolyte shows three main peaks, dominated by hydrocarbon contamination (284.8 eV), C-O bonds in ethers and carbonate salt (286-287 eV), as well as Li2CO3 (290 eV). However, as respective to the electrode cycled with DMAc-containing electrolyte, the C 1s spectrum contains three peaks and dominated by hydrocarbon contamination (284.8 eV), C-O bonds in ethers and carbonate salt (286.4 eV), as well as CdO bonds in lithium alkyl carbonates (R-CH2OCO2-Li) or polycarbonates (288.7 eV) . The intensity of CdO and C-O observed on the surface of
Table 2. Surface concentrations of different elements on fresh LiFePO4 cathode and cathode from cycled cells with STD electrolyte and with DMAc-containing electrolyte sample
C 1s (%) O 1s (%) F 1s (%) P 2p (%) Fe 2p (%)
fresh
60.1
22.3
0.9
7.9
8.8
STD electrolyte
38.5
56.1
39.4
1.0
0.5
STD þ 1% DMAc
43.5
23.4
29.2
1.9
2.0
the LiFePO4 electrode cycled with STD and DMAc-containing electrolyte is much stronger than that of the cathode cycled with the STD electrolyte (Figure 5, Table 2). In addition, the peak associated with inorganic decomposition product of Li2CO3 is not present in the C 1s spectrum of the electrode cycled in DMAc-containing electrolyte. This suggests that the LiFePO4 cathode electrode cycled with STD electrolyte was covered with more inorganic degradation products, such as Li2CO3 or LiF, than that of the electrode cycled with DMAc-containing electrolyte. The O 1s spectra have some similar peak presences, containing the lithium metal oxide, LiFePO4 (531.4 eV) and C-O bonds in polycarbonates and carbonate salts (533.5 eV), as well as CdO bonds in ethers and carbonate salts (532.5 eV), which are the products from electrolyte degradation. However, the intensity of LiFePO4 (531.4 eV) of the electrode cycled in DMAc-containing electrolyte is slightly lower than that of the electrode cycled in STD electrolyte, which suggests a thicker SEI covered on the LiFePO4 electrode in DMAc-containing electrolyte. This is consistent with the stronger intensity of C-O (532.1 eV) species observed in the O 1s spectra of the electrode cycled with DMAc-containing electrolyte. 6089
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Figure 6. FTIR-ATR spectra of the graphite electrodes cycled with and without DMAc electrolytes at 60 °C.
Figure 7. FTIR-ATR spectra of the LiFePO4 electrodes cycled with and without DMAc electrolytes at 60 °C.
The F 1s spectra are similar and contained two peaks: LiF (684.5 eV) and LixPOyFz (686.2 eV). However, the intensity of LiF (684.5 eV) of the LiFePO4 electrode extracted from the cell cycled with STD electrolyte is slightly stronger than that of the electrode cycled with DMAc-containing electrolyte, suggesting more inorganic species covered on the surface with STD electrolyte, consistent with the C 1s spectra as discussed above. The P 2p spectra are dominated by LiFePO4 (133.5 eV) and the deposition of a small amount of LiPF6 (137 eV). 3.6. FTIR Analysis of Electrodes. Analysis of the graphite anode (Figure 6) and LiFePO4 cathode (Figure 7) by FTIR-ATR
spectra provides additional insight into the differences of electrode surface in the cells with and without DMAc electrolyte after 50 cycles at 60 °C. The IR spectra of the anodes extracted from cells containing different electrolytes are similar and are dominated by the absorptions characteristic of polycarbonate (1775 cm-1), Li2CO3 (1427 cm-1), LixPOyFz (1040 cm-1, P-O), and LixPFy (847 cm-1, P-F), as well as some organic species (1230 cm-1), C-H (1617 cm-1).26-29 However, the intensities of LixPOyFz (1040 cm-1) and LixPFy (847 cm-1) for the anode cycled in DMAc-free electrolyte are slightly stronger than that of the anode cycled with DMAc-containing electrolyte, 6090
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Scheme 1. Reduction Mechanisms of EC and Linear Carbonates
Scheme 2. Thermal Dissociation Mechanisms of LiPF6Based Electrolytes with and without DMAc
which suggests that the anode surfaces are covered by more inorganic decomposition species in the DMAc free electrolyte. The intensity at 1617 cm-1 (δ C-H) is much stronger for the anode cycled with DMAc-containing electrolyte than that in DMAc free electrolyte, which also supports that much more organic SEI components are generated with the DMAc-containing electrolyte. The IR data acquired from the anode electrodes cycled at 60 °C are consistent with the XPS data discussed above. The electrodes are covered by lithium alkyl carbonate, Li2CO3, polycarbonate, as well as LixPOyFz, generated by the reductive decomposition of electrolyte at anode surface and thermal dissociation of LiPF6 at elevated temperature. Figure 7 displays the FTIR spectra of the LiFePO4 cathodes for the pristine electrode, after 50 cycles at 60 °C with and without 1% DMAc addition, respectively. FTIR spectrum for pristine electrode dominates between 932 and 1136 cm-1, which can be assigned to P-O vibrations of the PO43- polyanion.30-32 The absorption at 1136 cm-1 originates from symmetric and antisymmetric stretching vibration of O-P-O. The absorption at 1035 cm-1 originates from symmetric stretching vibration of P-O. The absorption at 932 cm-1 is attributed to the P-O stretching vibration. The absorptions at 633, and 646 cm-1 may be attributed to the symmetric stretching vibration of P-O and intramolecular symmetric stretching vibration of the Fe-O, respectively.33 After 50 cycles, new absorptions were observed of the LiFePO4 electrodes due to the surface reactions of electrode/electrolyte, and thermal dissociation of LiPF6-based electrolyte. The absorption at 1752 cm-1 can be attributed to CdO stretching mode in polycarbonate or lithium alkyl carbonate, consistent with the XPS data as discussed above. The peak at 1451 cm-1 is assigned to Li2CO3, which generally exists in the electrode surface due to the decomposition of carbonate electrolyte. The absorptions at 1236 and 842 cm-1 are related to LixPOyFz and LiF,34 originating from the thermal dissociation of LiPF6 at elevated temperature, which is also observed in the XPS spectra discussed above. On the basis of the surface chemistry on electrode/electrolyte layers discussed above, the main reduction mechanisms for the alkyl carbonate solvents (ethylene carbonate and linear carbonate solvents) are proposed in Scheme 1. This singleelectron reduction pathway is well proposed by Aurbach’s group.35,36 Scheme 2 suggests the reactions related to the
thermal dissociation of LiPF6 with and without the stabilizing additive, direction reduction of salt anions, as well as possible effects of PF5 and HF, similar decomposition mechanisms suggested by Marom et al.29 The reaction mechanisms suggested in Scheme 1 and 2 explain well the XPS and FTIR spectra of anode and cathode surface ingredients with and without DMAc. The standard electrolyte, 1.0 M LiPF6 EC/ DMC/DEC, experiences a significant thermal decomposition at elevated temperature, resulting in OPF3, HF, LiF, and LixPOyFz, et al. However, the thermal stability of electrolyte with DMAc is significantly enhanced due to the formation of the DMAc-PF5 complex. This explains that why so many inorganic ingredients are present on the anode and cathode surface with the DMAc-free electrolyte at elevated temperature besides the reduction products of carbonate solvents. In addition, the iron can dissolve in the electrolyte without DMAc at elevated temperature due to the high concentration of HF generated by the thermal dissociation of LiPF6, as suggest in Scheme 2, also supported by Aurbach et al.37,38 This difference may also account for the cycling difference of LiPePO4/ MCMB cell at elevated temperature. 3.7. DFT Calculations. The optimized geometry and the atomic charges of LiPF6 and PF5 are listed in Table 3, obtained from B3LYP/6-311þþ G (d,p) calculations at gas-phase and solvent-phase. In the optimized structure of LiPF6, the Li atom coordinates with three F atoms, almost the equidistance, 1.906, 1.906 and1.904 Å, respectively. This is consistent with Kasaki’s previous report,39 forming a stable tripod structure; however, the distance between Li and F presented here is slightly longer than that of Kasaki’s research work, 1.882 Å. In order for the dissociation of LiPF6 to take place, Li has to move to the extension of one of the P-F bonds with the P-F--Li angle being 180 o, the same direction of the P-F antibonding orbital. This coordination allowed the maximum overlap of atomic orbitals between the Li and F atom, facilitating the reaction of Li and F taking place, producing LiF and PF5. 6091
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Table 3. Optimized Geometries and Atomic Charges for Solutes in Various Solventsa gas phase
DEC
DMC
mixb
EC
1.761
1.769
Table 4. Thermodynamic Properties for Decomposition of LiPF6 in Various Solvents (kJ/mol) a
LiF r LiF
1.583
q Li
0.960
0.962
0.963
0.961
0.961
qF LiPF6
-0.960
-0.962
-0.963
-0.961
-0.961
r Li-F6
1.906
1.938
1.953
2.057
2.069
r Li-F2
1.904
1.938
1.951
2.057
2.069
r P-Li
2.479
2.832
2.845
2.931
2.942
r P-F3
1.591
1.608
1.610
1.619
1.620
r P-F2
1.714
1.705
1.702
1.673
1.671
q Li
0.964
0.977
0.977
0.978
0.978
q F2 qP
-0.625 2.539
-0.645 2.550
-0.644 2.550
-0.635 2.550
-0.634 2.551
1.568
1.674
1.682
PF5 r P-F2,4,5
1.570
1.569
1.569
1.568
r P-F3,6
1.604
1.608
1.608
1.611
1.611
qP
2.629
2.629
2.629
2.629
2.629
q F2,4,5
-0.511
-0.511
-0.511
-0.510
-0.511
q F3,6
-0.548
-0.548
-0.548
-0.549
-0.549
a
r and q refer to the distance in angstroms and the atomic charge, respectively. All results are obtained from B3LYP/6-311þþG(d,p) and PCM-B3LYP/6-311þþG(d,p) calculations in gas phase and in solvents. b mix refers to the mixture solvents, EC/DMC/DEC (1/1/1, weight ratio).
Next, we take into account the solvents effects. The optimized geometries and the atomic charges are summarized in Table 3, obtained at the PCM-B3LYP/6-311þþG(d,p) level in solvent phase and natural bond orbital (NBO) theory for charge distribution. The optimized geometries and atomic charges in gas phase are also presented in Table 3. As we can see from Table 3, the discrepancy of the geometry and the atomic charge becomes greater as the polarity of the solvent increases. For instance, the distance between Li and F of LiF increases in going from DEC 1.674 Å to EC 1.769 Å, the distance between Li and P of LiPF6 increases from DEC 2.832 Å to EC 2.942 Å, suggesting qualitatively the experimental trend of a larger dissociation of LiPF6 in a more polar solvent.40 The geometry and the atomic charge of PF5 change little over a wide range of solvents. This is due to its molecular symmetry. Table 4 summarizes the thermodynamic properties (ΔE, ΔE þ ΔZPE, ΔH, ΔG) obtained from B3LYP/6-311þþG(d,p) in gas phase and PCM-B3LYP/6-311þþG(d,p) in various solvents represented by their dielectric constants (ε). Frequency analysis was performed at B3LYP/6-311þþG(d,p) and PCM-B3LYP/ 6-311þþG(d,p) levels to obtain the thermal corrections to the enthalpy and the entropy. According to the PCM-DFT calculations, the decomposition is largely endothermic in these solvent and shifts to lower with increasing polarity of the solvent, ΔG = 110.56 kJ/mol for DEC and ΔG = 50.1 kJ/mol for EC, which suggests that the dissociation of LiPF6 is getting easier with increasing polarity of the solvents. However, we know that LiF is insoluble in the solvents used in this study; therefore, the thermodynamic properties (ΔE, ΔE þ ΔZPE, ΔH, ΔG) obtained in Table 4 is not accurate and cannot
gas
DEC
DMC
mix
EC
ε
1
2.8
3.1
31.9
89.8
ΔE ΔE þ ΔZPE
177.29 150.74
152.48 146.30
151.94 145.98
97.5 90.7
97.4 90.6
ΔH
174.24
149.22
148.70
93.2
93.1
ΔG
128.37
110.56
106.12
50.4
50.1
Calculated is the free energy difference for the decomposition: ΔG = G(LiF) þ G(PF5) - G(LiPF6).
a
be directly compared to the experimental data. Furthermore, PF5 is a very strong Lewis acid and tends to react with the solvents in the electrolyte. Therefore, the interactions between the solvents with PF5 should be considered. In this calculation, a model was constructed to consider the interactions between the solvents with PF5 by immersing solvent complex of PF5 in each solvent continuum. The structures of the complexes were constructed from the optimized of PF5 with the P atom coordinating with carbonyl oxygen of the solvent, as well as DMAc in the solvents. The complexes with the P atom pointing at the ether oxygen of the solvents were also investigated. It was found that the P atom pointing at the carbonyl oxygen structure is much more stable than that pointing at the ether oxygen of all solvents. Herein, we only show the stable structure results. The optimized geometry obtained from PCM-B3LYP/6-311þþG(d,p) calculations in EC/DMC/DEC solution is displayed in Figure 8. Table 5 presents the relative energy (ΔE) of each PF5solvent complex calculated in EC/DMC/DEC solution, obtained to the same calculations in Table 4, for the reaction PF5 þ S f PF5-S, S refers to the solvent and stabilizing additive. As we can see from Table 5, the relative energy of PF5-DEC, PF5-DMC, PF5-EC, and PF5-DMAc is -15.2, -14.6, -35.0, and -88.2 kJ/mol, respectively, suggesting the complex of PF5DMAc is much more stable than that of PF5-DEC, PF5-DMC, and PF5-DMAc in EC/DMC/DEC solvents. This indicates that PF5 prefers to attack DMAc than other electrolyte components, EC, DMC, and DEC, forming a stable complex of PF5-DMAc. This is also supported by the evidence that the bond distance of P-(Cd)O in PF5-DMAc complex (1.799 Å) is much shorter than that of PF5-DEC (1.863 Å), PF5-DMC (1.893 Å), and PF5-EC (1.879 Å), suggesting a strong interaction between PF5 and DMAc. This calculation results are in good agreement with the NMR experimental results. The atomic charges of PF5-solvents and PF5-stabilizing additive complexes were analyzed by the natural bond orbital (NBO) theory at PCM-B3LYP/6-311þþG(d,p) level in the EC/DMC/DEC (1/1/1) solvents, shown in Table 5. The atomic charge of P for PF5-DMC, PF5-DEC, PF5-EC, and PF5-DMAc complex is 2.544, 2.599, 2.546, and 2,541, respectively. This is also suggests that the PF5-DMAc complex is much more stable than the PF5-DMC, PF5-DEC, PF5-EC complexes. However, the atomic charges of carbonyl oxygen of PF5DMAc is slightly larger than that of PF5-EC, -0.689 and 0.622, respectively. This may be caused by the cyclic structure of EC. On the basis of the calculation results above, it is found that a strong interaction existed between PF5 with carbonate solvents, and stabilizing additive. However the stabilizing additive DMAc prefers to attack PF5, resulting in a much more stable complex in 6092
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Figure 8. Geometry of the PF5-solvent complex optimized at PCM-B3LYP/6-311þþG (d,p) level in EC/DMC/DEC (1/1/1, weight ration) solvents.
Table 5. Relative Energy (kJ/mol), Atomic Charges, and Bond Distances (Å) of PF5-Solvent Complexes Optimized at PCM-B3LYP/6-311þþG(d,p) in EC/DMC/DEC (1/1/1, Weight Ratio) Solution ΔE qP
PF5-DEC
PF5-DMC
PF5-EC
PF5-DMAc
-15.2 2.554
-14.6 2.559
-35.0 2.546
-88.2 2.541
qO
-0.733
-0.723
-0.622
-0.689
r P-O
1.863
1.893
1.879
1.799
solution than that of the carbonate solvents, DMC, DEC, and EC. This calculation results are in good agreement with the NMR experimental observation that addition of DMAc into the carbonate-based electrolyte can suppress the thermal decomposition of carbonate solvents and significantly improve the thermal stability of the carbonate electrolyte.
4. CONCLUSIONS Addition of 1% dimethylacetamide (DMAc) to 1.0 M LiPF6 EC/DMC/DEC (1/1/1) electrolyte can significantly improve the thermal stability of the electrolyte and the cyclic performance of LiFePO4/graphite cells at 60 °C. The electrolyte containing 1% DMAc can be stable at 85 °C for over 6 months. The discharge capacity retention increased from 55.0% to 87.2%. This can be ascribed to the enhancement of the thermal stability of the electrolyte, as well as the modification of SEI components of anodes and cathodes in the presence of DMAc. The mechanisms of enhancements of thermal stability of the electrolyte and interface layers have been investigated by NMR, FTIR, and XPS. It was found that the addition of DMAc to LiPF6-based electrolytes inhibits the autocatalytic thermal decomposition of the electrolyte due to the formation of PF5-DMAccomplex. DFT calculation results suggest that the stabilizing additive DMAc prefers to attack PF5, resulting in a much more stable complex in solution than that of the carbonate solvents, DMC, DEC, and EC. ’ AUTHOR INFORMATION Corresponding Author
*E-mail:
[email protected] (M.X.);
[email protected] (W.L.).
’ ACKNOWLEDGMENT This work is supported by the National Natural Science Foundation of China (No.NSFC21003054), Natural Science Foundation
of Guangdong Province (No. 10351063101000001), Specialized Research Fund for the Doctoral Program of Higher Education (Grant No. 20104407120008) and Project of Guangdong Province (Grant No. 2009B050700039).
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