Robust Benzimidazole-Based Electrolyte Overcomes High-Voltage

Article pubs.acs.org/cm

Robust Benzimidazole-Based Electrolyte Overcomes High-Voltage and High-Temperature Applications in 5 V Class Lithium Ion Batteries Fu-Ming Wang,*,†,‡ Sylvia Ayu Pradanawati,† Nan-Hung Yeh,† Shih-Chang Chang,† Ya-Tang Yang,† Shih-Han Huang,† Ping-Ling Lin,† Jyh-Fu Lee,§ Hwo-Shuenn Sheu,§ Meng-Lin Lu,§ Chung-Kai Chang,§ Alagar Ramar,† and Chia-Hung Su∥ †

Graduate Institute of Applied Science and Technology, National Taiwan University of Science and Technology, Taipei 106, Taiwan Sustainable Energy Center, National Taiwan University of Science and Technology, Taipei 106, Taiwan § National Synchrotron Radiation Research Center, Hsin-Chu 300, Taiwan ∥ Department of Chemical Engineering, Ming-Chi University of Technology, Taipei 243, Taiwan ‡

S Supporting Information *

ABSTRACT: Electric vehicles (EVs) are poised to dominate the next generation of transportation, but meeting the power requirements of EVs with lithium ion batteries is challenging because electrolytes containing LiPF6 and carbonates do not perform well at high temperatures and voltages. However, lithium benzimidazole salt is a promising electrolyte additive that can stabilize LiPF6 through a Lewis acid−base reaction. The imidazole ring is not eligible for high-voltage applications owing to its resonance structure, but in this research, electronwithdrawing (−CF3) and electron-donating (−CH3) substitutions on imidazole rings were investigated. According to the calculation results, the CF3 substitution facilitates a high electron cloud density on imidazole ring structures to resist the electron releases from bezimidazole in oxidation reactions. In addition, through CF3 substitution, electrons are accepted from the lattice oxygen (O2−) in lithium-rich layer material and O− is converted by an electron released. The O− is then adsorbed with the ethylene carbonate and catalyzed to alkyl carbonate by Ni2+. The −CF3 substituted benzimidazole triggers a further reaction with alkyl carbonate and forms a new polyionic liquid solid electrolyte interphase on the cathode’s surface. Furthermore, the cycle performance tested at 60 °C and 4.8 V showed that the CF3 substitution maintains the battery retention effectively and exhibits almost no fading compared with both the blank electrolyte and the CH3 substitution.

1. INTRODUCTION

electrolyte solvents make these Li-rich materials ineligible for executing high anodic electrochemical reactions such as with ethylene carbonate (EC), propylene carbonate, and dimethyl carbonate (DMC). These carbonate solvents cannot withstand drastic oxidation because of the gas evolution and polymer formation caused by the breaking of chemical bonds.7,8 Recently, researchers have attempted to develop electrolyte additives9 and new electrolyte solvents, including ionic liquid,10,11 to improve the interfacial properties between highvoltage cathode surfaces and electrolytes. Li and colleagues used tris(trimethylsilyl)borate (TMSB) as an electrolyte additive, reporting that although the capacity retention of Li[Li0.2Mn0.54Ni0.13Co0.13]O2 was only 19% in the blank

Lithium ion batteries, the most popular energy storage source in the world, face two substantial limitations in electric vehicle (EV) and hybrid electric vehicle (HEV) applications.1,2 These are caused by high temperatures (≥60 °C), along with lithium salt decomposition and substantial side reactions on the electrode’s surface from PF5 formation.3,4 In addition, HEV and EV require lithium ion batteries with high power and high energy density. Innovation is required to meet these stringent standards and lower the cost. In the literature, high-voltage (>4.8 V; 5 V class) Li-rich layer compounds (Li[NixLi(1−2x)/3Mn(2−x)/3]O2) have been reported as a promising new candidate for cathode materials instead of the conventional LiCoO2 or LiNixMnyCozO2 where increased power and energy density are required. These Li-rich layer materials have reversible capacities over 270 mAh g−1, which is almost twice as high as that of LiCoO2.5,6 However, the current © 2017 American Chemical Society

Received: February 27, 2017 Revised: June 9, 2017 Published: June 11, 2017 5537

DOI: 10.1021/acs.chemmater.7b00824 Chem. Mater. 2017, 29, 5537−5549

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Chemistry of Materials

omethyl benzimidazole (LITFB) were synthesized using a deprotonation method between LiOH (Sigma-Aldrich) and a mixture of methyl benzimidazole (Sigma-Aldrich) and trifluoromethyl benzimidazole (Alfa Caesar). To prepare the LIMB/LITFB, the mixture of methyl benzimidazole (0.2 M) and trifluoromethyl benzimidazole (0.2 M) was stirred with LiOH (0.4 M) and then placed into the solvent tetrahydrofuran (which had been dried prior to use) for 24 h in a three-necked round-bottomed flask. After stirring, the slurry was filtered and placed in a vacuum at 80 °C overnight. The resulting Li benzimidazole salts were characterized using nuclear magnetic resonance (NMR) spectroscopy, mass spectrometry, and differential scanning calorimetry (DSC). The high-voltage cathode constituted 90 wt % Li1.2Mn0.6Ni0.2O2 (LLNMO), 5 wt % Super-P as a conductive carbon, and 5 wt % PVDF as a binder. The graphite anode comprised 93 wt % mesocarbon microbeads (MCMBs), 3 wt % vapor-grown carbon fibers as a conductive carbon, and 4 wt % PVDF as a binder. 2.2. Electrolyte Preparation. Four electrolytes were prepared for the experiments. Sample A was 1 M lithium hexafluorophosphate (LiPF6) in EC and diethyl carbonate (DEC) (volume ratio of 1:2, Battery grade, water content less than 20 ppm), purchasing from Unionward company in Taiwan. The remaining samples were identical to Sample A, except for the addition of 0.1 wt % lithium benzimidazole (LIB) to Sample B, 0.1 wt % LIMB to Sample C, and 0.1 wt % LITFB to Sample D. The battery fabrication and electrolyte preparation were performed inside a glovebox in an Ar gas atmosphere to avoid the influence of moisture. 2.3. Instrumentation. The lithium salts were characterized using NMR spectroscopy, mass spectrometry, and DSC. For the NMR measurement, the lithium salts were dissolved in d-DMSO or dacetone solvent and then being placed in an NMR tube. The thermal stabilities of the lithium salts were studied using DSC. PerkinElmer DSC-4000 was used to conduct the experiment from 25−250 °C at a scan rate of 10 °C min−1. The samples were pressed using gold pans and placed in ambient nitrogen to avoid moisture. Ionic conductivities were measured using electrochemical impedance spectroscopy (EIS) from 100.0 K−0.1 Hz with an AC amplitude of 5 mV, at temperatures of 25−90 °C. The equipment was controlled using EC-Lab electrochemical software (Biologic, Inc.). Temperaturedependent conductivities were obtained by placing the electrochemical cell in an oven. The impedance spectra exhibited depressed semicircles at high frequencies and slanted lines at low frequencies; the resistance was determined based on an extrapolation of the intercept of the slanted line and the real axis of each impedance spectrum plot (imaginary Z″ versus real Z′). The specific ionic conductivity σ was obtained from σ = l/AR, where the distance between the two stainless electrodes is l = 0.5 cm. For the stainless electrode, the measured area is A = 1 cm2, and R is the measured resistance (Ω). The compatibility of the electrolytes with the stainless electrode was determined using EIS analysis under open-circuit conditions at various temperatures. Ionic conductivity measurements at various temperatures were used to determine the activation energy of the electrolyte from an Arrhenius plot.

electrolyte, it improved to 74% when 0.5% TMSB was applied.12 In addition, they demonstrated that a novel electrolyte additive, phenyl trifluoromethyl sulfide (PTS), could be used to improve the interfacial stability of highvoltage lithium nickel manganese oxide (LiMn1.5Ni0.5O4). At a rate of 1C, the addition of 0.5% PTS improves the capacity retention of LiMn1.5Ni0.5O4 from 65% to 84% after 450 cycles at room temperature, and from 64% to 95% after 100 cycles at evaluated temperature.13 Some reports on the additive of borates for spinel14 and layer15 high voltage materials are also investigated. The design of high-voltage electrolytes is focused on the stabilization of the carbonate backbone structure and interfacial properties on the cathode’s surface; however, adopting hightemperature electrolytes requires varying criteria for the lithium salt as well as interfacial properties on the anode’s surface. LiPF6 is a common salt in lithium ion batteries because of its excellent solid electrolyte interphase (SEI) formation,4 excellent solubility in carbonate solvents,14 wide electrochemical operation window,16 and high ionic conductivity.16 However, LiPF6 suffers thermal stability problems at the temperature above 60 °C, leading an important side reaction in eq I:3,17−21 LiPF6 → LiF + PF5

(I)

The side product, PF5 (a Lewis acid) is usually reacting with the SEI in the presence of moisture to form HF and LiF. The following reactions (eqs II to V) demonstrate the formations of insulators such as LiF, POF3, and RCOF, thereby increasing the interfacial impedance and corroding the current collectors.22−24 PF5 + H 2O → POF3 + 2HF

(II)

RCO2 Li + PF5 → RCOF + LiF + POF3

(III)

Li 2CO3 + PF5 → POF3 + 2LiF + CO2

(IV)

ROCO2 Li + PF5 → RF + LiF + CO2 + POF3

(V)

Wang and colleagues proposed a new Lewis base compound, benzimidazole-based Li salt additive, to neutralize PF5 at high temperatures in an electrolyte containing LiPF6. This result revealed a pentafluorophosphate benzimidazole anion synthesized in situ by a Lewis acid−base reaction between the benzimidazole anion and PF5. The new compound, pentafluorophosphate benzimidazole anion, inhibits the decomposition of LiPF6 by inhibiting PF5 side reactions, which leads to a well-maintained battery performance at 60 °C.3 The benzimidazole and imidazole salts had been reported to their potential applications in lithium ion battery. In accordance with the calculation results, high electron-withdrawing groups such as cyano and trifluoro substitution on structure greatly enhance the anodic stability of electrolyte.25,26 In this research, two lithium benzimidazole-based salt additives within electron-withdrawing (−CF3) and electrondonating (−CH3) functional group substitutions were synthesized in accordance with a high-voltage and high-temperature mutual function. By combining previous observations3,24,27 within this study, the substitution functional group effects on both electrodes’ surfaces at high voltages and high temperatures were revealed.

ln σ = −

Ea + ln A RT

slope = −

Ea R

The anodic stability of the electrolytes in three-electrode Teflon cell28 was measured through cyclic voltammetry (CV; Biologic VMP3) from open-circuit potential (OCP) to 5.6 V at a scan rate of 0.2 mV s−1 within a platinum as a working electrode and Li metal foil as counter and reference electrodes. The electrolyte fills the space between the working and the counter electrodes. The anodic electrochemical stability of the electrolytes in battery was measured through CV from OCP to 4.90 V at a scan rate of 0.1 mV s−1 using a coin cell (CR2032) within a high-voltage cathode (LLNMO) as a working electrode and Li metal foil as a counter electrode (1.0 cm2) with the electrolyte filling the space between the

2. EXPERIMENTAL SECTION 2.1. Materials and the Preparation of Lithium Benzimidazole. Lithium methyl benzimidazole (LIMB) and lithium trifluor5538


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Chemistry of Materials working and the counter electrodes. EIS was performed using a Biologic VMP3 from 106 to 0.01 Hz with an AC amplitude of 5 mV at 25 °C in a 100% state of discharge. The charge−discharge test in cathode and anode half cells was conducted using the constant current−constant voltage (CC−CV) mode with a voltage range from OCP to 4.9 and 0.005 V to OCP at 0.1 C/0.1 C at 60 °C, measured using a U-bic battery tester. Scanning electron microscopy (SEM) and energy dispersive spectrometry (EDS) were used to observe the MCMB electrode’s morphology (after Pt-coating the electrode in a glovebox) at an accelerating voltage of 15 kV using an LEO-1530 microscope. Before any observations were made, the specimens were disassembled and washed with DMC in a dry room, then dried overnight in a vacuum. The samples were placed into a custom-built high vacuum stainless steel holder to transfer the electrodes from the dry room to the SEM instruments. The SEI was not influenced by the treatment before assembling and drying the cell. All geometry optimizations and frequency determinations were performed using the Gaussian 09 software package.29 Density functional theory (DFT) was selected with the B3LYP method, which is a hybrid function containing Becke’s three-parameter exact exchange function (B3) in conjunction with the nonlocal gradientcorrected correlation function of Lee−Yang−Parr (LYP).29 B3LYP was applied with the triple split valence basis set 6-311G and a set of p, d polarization functions on heavy atoms and hydrogen atoms.29 Spinunrestricted calculations were used to allow for any possible bond cleavage during the geometric optimization of the radical species. The LLNMO structures of all the samples were observed with a high-resolution transmission electron microscope (HR-TEM; HITACHI H-9000UHR III) and analyzed through selected area electron diffraction (SAED). The HR-TEM was operated at 300 kV. The Synchrotron XRD analysis was conducted by using the beamline 01C2 facility at the National Synchrotron Radiation Research Center in Taiwan. The wavelength of the incident X-ray was 1.0223 Å. Each XRD spectrum was obtained with an acquisition time of 180 s. The data was collected using a mar345 imaging plate detector. By using FIT2D software, the 2D Debye−Scherrer rings were integrated to the equivalent of 2 scan and were corrected for polarization and tangent geometry. The scattering angle was calibrated by standard sample, which is a mixture of silver behenate and Si powders. X-ray absorption spectra (XAS) were obtained at the National Synchrotron Radiation Research Center (NSRRC), Hsinchu, Taiwan, where the storage ring of the electronic accelerator can supply 1.5 GeV with an operating current of 360 mA. Hard X-ray absorption data for the Ni and Mn K-edges were collected at beamline BL17C. The in situ experiment was carried out in operando, at C/50, in a modified coin cell battery, the only difference in the electrochemical cell being the electrolyte used. A Si(111) double-crystal monochromatic was used to perform the energy scan, in which the parallelism was adjusted to eliminate high-order harmonics. All of the spectra were recorded in transmission mode. Ionization chambers were used as detectors to monitor the intensity of the incident and transmitted beams through the specimen, thereby allowing the absorption coefficient to be calculated from the logarithm of the intensity ratio of the incident and transmitted beams. Ni and Mn reference foils were positioned in front of the window of the third ionization chamber to allow for simultaneous measurement to act as a standard, that is, calibration energy, for each scan.

Figure 1. 1H NMR spectrum of (a) Methyl-Benzimidazole and (b) Lithium Methyl-Benzimidazole (LIMB).

indicating that the synthesis was successful. The specific peaks of the methyl benzimidazole (Figure 1a) correlated with δ2.5 (quintet, d-DMSO); δ3.3; δ12.1 (singlet, N−H); δ2.6 (triplet, −CH3); δ7.1 (singlet, 1H); δ7.4 (multiplet, 2H); and δ7.5 (multiplet, 3H). In addition, the specific peaks of LIMB (Figure 1b) correlated with δ2.5 (quintet, d-DMSO); δ2.6 (triplet, −CH3); δ8.8 (singlet, 1H); and δ7.4 (multiplet, 2H/3H). In addition, the mass spectrum was detected (not provided in this manuscript), MS(EI) m/z 132 (M+, 100). Figure 2 shows the 1H NMR spectra of the precursor and LITFB. A bonding absorbance difference of 3.4 and 13.9 ppm

Figure 2. 1H NMR spectrum of (a) trifluoro methyl-benzimidazole and (b) lithium trifluoro methyl-benzimidazole (LiTFB).

was observed, whereas the reacted functional group did not show a −N−H group after the deprotonation reaction, indicating that the synthesis was successful. The specific peaks of the trifluoromethyl benzimidazole (Figure 2a) correlated with δ2.5 (quintet, d-DMSO); δ3.4; δ13.9 (singlet, N−H); δ7.4 (singlet, 1H); and δ7.7 (multiplet, 2H/3H). In addition, the specific peaks of LITFB (Figure 1b) correlated with δ4.8 (singlet, D2O); δ7.7 (singlet, 1H); and δ7.3 (multiplet, 2H/3H). The mass spectrum was also detected (not provided in this paper), MS(EI) m/z 162 (M+, 100). 3.2. Analysis of DFT Calculation. According to qualitative molecular orbital theory,16 the highest occupied molecular orbital (HOMO) indicates the characteristic of a nucleophilic component. Hence, a low-energy HOMO indicates that electrons are not easily donated to the lowest unoccupied molecular orbital for a chemical bonding reaction. Table 1 shows the HOMO values for three lithium bezimidazole salts, indicating the anion oxidation stability. The DFT calculations estimate that the HOMO of LITFB salt is predicted to be −5.957 eV, which is lower than that of LIB (−5.304 eV) and LIMB (−5.114 eV). High electro-withdrawing functional group (CF3) substitution on the imidazole ring results in an

3. RESULTS AND DISCUSSION 3.1. LIMB and LITFB Salts Synthesis. LIMB and LITFB were synthesized using a deprotonation method.3,30 Benzimidazole with a −N−H bond on an imidazole ring was replaced by tuning the relative ratios of the −N−H bond and LiOH. Figure 1 shows the 1H NMR spectra of the precursor and LIMB. A bonding absorbance difference of 3.3 and 12.1 ppm is displayed in Figure 1, whereas the reacted functional group did not show a −N−H group after the deprotonation reaction,

Table 1. DFT Analysis of Lithium Benzimidazole-Based Salts

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salt

HOMO (eV)

Li-anion (Å)

Eox (V)

ΔEd (kcal mol−1)

LIB LIMB LITFB

−5.304 −5.114 −5.957

1.795 1.786 1.840

4.96 (BTB)29 4.73 (BTMB)29 5.05 (BTTB)29

115.2 (BTB)29 116.5 (BTMB)29 113.2 (BTTB)29


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Chemistry of Materials improvement of anion oxidation potential and weakening of the corresponding ion pair. The CF3 substitution facilitates a high electron cloud density on the structure to resist the electron releases from bezimidazole in the oxidation (anodic) reaction. In addition, the high electron cloud density also causes two adjacent (C−N) bond elongations because of the steric repulsion. This repulsion effect produces lower ion pair dissociation in accordance with weaker coordination because of extensive charge delocalization.31,32 The bond length of an ion pair in LITFB salt (1.840 Å) is therefore longer than those of LIB (1.795 Å) and LIMB (1.786 Å). Figure S1 and Table S1 show the ionic conductivity and activation energy of all the electrolytes. According to the results, Sample D has the lowest activation energy and the highest ionic conductivity. By contrast, CH3 substitution in the imidazole structure shows an opposite result. This CH3 electro-donating functional group substitution causes less electron negativity on the imidazole ring compared with CF3. Jiang and colleagues demonstrated the effects of substituents in their bis(trufluoroborane)−benzimidazole salts through DFT calculation (Table 1). As with this study, their results reflected that an electron-withdrawing group substitution (bis(trifluoroborane)-2-(trifluoromethyl)benzimidazolide, BTTB anion) is used to enhance the anion oxidation stability (5.05 V) and lower the ion pair dissociation energy (133.2 kcal mol−1), which is suitable for application in high-voltage lithium ion batteries.29 However, their calculation does not contain and discuss about the electrochemical effects on the electrode’s surface. Therefore, some electrochemical tests such as CV, EIS, and morphology are shown in this study. 3.3. Analysis of Thermal, NMR, and Electrochemical Characteristics of the Electrolyte. Benzimidazole has an imidazole ring and aromatic resonance structure that enhances the thermal stability of the electrolyte and the imidazole ring bonds of PF5 through the Lewis acid−base reaction.3 The DSC result (Figure 3) showed that the commercial electrolyte contains three Li benzimidazole salt additives that demonstrate extremely different thermal behaviors. In this study, Sample B displayed almost no exothermic and endothermic reactions below 210 °C. The two endothermic reactions at temperatures of 110−114 °C and 220−223 °C were assigned to the boiling point of DEC and the transesterification reaction between DEC and PF5. Figure S2 shows an enlarged area of DSC results in Figure 3, displaying the sensitive boiling reaction of DEC on all samples. The heat trend on Samples B and C displays continuous endothermic reaction after 120 °C rather than the Sample D, indicating the gaseous DSC joins the transesterification reaction with PF5 at further temperature range. Interestingly, Sample C indicated that the electron-donating group (−CH3) on the imidazole ring strengthened the transesterification reaction and lowered the main endothermic temperature to 170−172 °C. In comparison with three Li benzimidazole salts, the Sample A shows the dissolution of residual LiPF6 at 50−70 °C in the carbonates, the DEC boiling point at 105−110 °C the acid-catalyzed transesterification reaction of DEC at 130−160 °C, and the electrolyte decomposition above 210 °C, respectively.3 This behavior can be explained by the highest HOMO value of the LIMB salt. However, the increment tendency of the endothermic behavior on the DSC curve proceeded to the end of the experiment, indicating that the LIMB salt was ineligible for accepting PF5 and that the PF5 continuously reacted with other Lewis bases in the electrolyte. Sample D demonstrated an opposite result compared with Sample C. The electro-

Figure 3. DSC results of benzimidazole based electrolytes at the scanning rate of 5 °C min−1.

withdrawing group (−CF3) on the imidazole ring terminated the transesterification reaction and the main endothermic temperature was not be found until 250 °C. In addition, the tendency of the endothermic behavior on the DSC curve could no longer be observed and the reaction turned to slightly exothermic behavior, indicating that the PF5 was almost neutralized by the LITFB salt. According to this result, a specific exothermic reaction was measured in the region of 187−201 °C, which can be assigned to the carbonate decomposition.33,34 This result concurs with the DFT calculation, in which the LITFB salt has the lowest HOMO value. To analyze the side products after the aging process, four electrolyte systems were stored at 60 °C for 12 h. Sample A clearly displayed two peaks (−153.61 and −153.57 ppm; Figure 4a) and one peak (−85.55 ppm; Figure 4b), which can be assigned to HF and POF3 formation after the aging process,35−37 indicating that the reaction in eq II occurred.

Figure 4. 19F NMR spectrum of electrolytes of (a) HF and (b) POF3 evolution after the electrolytes were aged at 60 °C for 12 h. 5540


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a higher energy band gap) compared with the other salt additives, indicating that the strong electron-withdrawing group is able to provide free electron pairs against the high anodic reaction on the imidazole ring. In Figure S4, the Samples D shows an outstanding performance at 0.2C/0.2C as well as 60 °C compare with other Samples. The higher rate testing does not affect the SEI formation on cathode’s surface. 3.5. Analysis of Electrochemical Impedance Spectroscopy. Figure 6 shows the EIS spectra of all the batteries. The

However, Samples B and C demonstrated two peaks (−152.98 and −152.95 ppm; Figure 4a) and one peak (−85.55 ppm: Figure 4b), with the chemical shifts resulting from the electron negativity effect of the Benzimidazole salts, although the intensity of all the peaks became lower. Thus, benzimidazolebased salt additives can be used to eliminate the formation of PF5 through the Lewis acid−base neutralization reaction. Interestingly, Sample D showed of almost no HF and POF3 formation after the aging process. This result demonstrated that the electron-withdrawing group (−CF3) strengthens the Lewis acid−base neutralization reaction more than the electrodonating group (−CH3) and can be used to terminate all the side reactions in the electrolyte. Figure S3 shows the CV measurements on Sample A and Sample D. According to the CV, the Sample D containing LiTFB shows significantly stable when the testing potential arises to more than 5.5 V. In comparison with the Sample D, Sample A displays a clear decomposition reaction above potential 5.2 V, indicating that the LiTFB maintains the electrolyte. 3.4. Analysis of Cycle Performance. Figure 5 shows the cycle performance of the LLNMO cathode half-cell at 60 °C.

Figure 6. EIS spectra of Sample A, B, C, and D in LLNMO cathode half cells after the 20th cycle at 60 °C. Inset table represents the simulation results of equivalent circuit of four electrolytes.

resistance of the electrolyte (Re) simulated by the circuit model was markedly similar, indicating that the bezimidazole salt additives do not substantially change the ionic conductivity of the electrolyte. However, the values of Rsei and Rct varied greatly with the addition of bezimidazole salt additives. In this study, Rsei represented the passivation layer resistance and Rct was the interfacial charge transfer resistance on the LLNMO surface. According to the table in Figure 6, Sample D exhibited an Rsei value of 16.93 Ω, which was less than that of Sample A (80.43 Ω; 78.9% decrease) and Sample C (71.78 Ω; 76.4% decrease). Furthermore, Sample D revealed an Rct value of 118.5 Ω, which was less than that of Sample A (372 Ω; 68.1% decrease) and Sample C (258.2 Ω; 54.1% decrease). These EIS results demonstrated that Sample D with its strong electron-withdrawing group causes a unique SEI formation on the LLNMO cathode, in which the high diffusivity compounds fabricate the SEI and are used to eliminate the interfacial impedance. 3.6. Analysis of SEM and EDS. Figure 7 shows the LLNMO morphologies of the four electrolyte systems after 20 cycles at 60 °C, visualized using SEM (the scale bars in Figure 7 are 1 μm). The four photographs show that the morphologies of the LLNMO are distinct. Figure 7a indicates that the LLNMO particle distribution was uniform in the electrode used with Sample A and that the secondary particles were separate, with a diameter distribution of 300−400 nm. In terms of the EDS results in Table 2, the electrode used with Sample A had the highest fluorine content (14.72 at. %) compared with the other bezimidazole samples, implying that undesirable compounds, mostly LiF was disclosed in the cathode’s surface because of the high-temperature effects on LiPF6 decomposition. In addition, Table 2 also shows that the Mn/Ni atomic ratio (2.79) on the LLNMO electrode used with Sample

Figure 5. Capacity retention of Sample A, B, C, and D in LLNMO cathode half cells at 60 °C.

Compared with the other samples, Sample D had the best capacity retention, showing almost no fading after 20 cycles. In comparison with Sample A, the electrolyte containing lithium salt only, LiPF6 decomposed to LiF and PF5 easily at high temperature.3 The high temperature triggered side reactions, including LiF interrupting the ionic pathway on the electrode surface and PF5 reacting with the SEI. These reactions indicated that the ionic transfer on the electrode surface eventually terminates. Sample B improved the battery performance at high temperatures because of the self-reaction between LIB salt and PF5 (i.e., the Lewis acid−base reaction).3 Figure 5 also reveals that the cycle retention of the batteries is followed by the sequence of Sample D (LITFB), Sample B (LIB), Sample C (LITMB), and Sample A (blank). The results indicated that a strong electron-withdrawing group on the structure strengthens the Lewis acid−base reaction. Although Samples B and C confirmed the battery performance at high temperatures, neither of them are suited to application at high voltages (approximately 4.8 V). The intrinsic problem for Samples B and C was the imidazole ring. The DFT calculation in Table 1 demonstrates that Sample D has the lowest HOMO value (with 5541


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Table 2. EDS Analysis of LLNMO Electrodes with the Electrolytes of Sample A, B, C, and D at 60 °C after 20 Cycles name element

Sample A

Sample B

Sample C

Sample D

C O F P Mn/Ni

27.63 40.36 14.72 0.57 2.79

22.79 58.85 4.41 0.43 2.98

20.46 48.35 7.48 0.52 3.01

19.84 65.09 3.89 0.56 3.02

A did not conform to the original design (atomic ratio 3.0), indicating that part of the Mn had been dissolved through HF corrosion (eqs I and II). Compared with Sample A, the electrodes used with Samples B and D showed interesting morphologies in that the SEI sensitively adhered to the secondary particles and aggregated them, yielding SEI-coated LLNMO particles. These photographs confirm that the effects of the electron-withdrawing group substitution cause the SEI formation on the LLNMO’s surface. 3.7. Analysis of in Situ Gas Chromatography−Mass Spectrometry. Figure 8 shows the in situ gas evolution observed in the Sample A and D systems. In Figure 8a, two potentials, 4.21 and 4.6 V, are used as watersheds to distinguish the stages of gas evolution. Lower than 4.21 V, no gas evolution occurred, indicating a normal metal transition (Ni2+/Ni4+) reaction on the LLNMO electrode. Higher than 4.21 V, CO was generated, and CO2 was released at 4.5 V. Both the CO

Figure 7. SEM micrograph of LLNMO electrodes with the electrolytes on (a) Sample A, (b) Sample B, (c) Sample C, and (d) Sample D after the 20th cycle at 60 °C. All the scale bars are 1 μm.

Figure 8. In situ GC/MS measurement of LLNMO electrodes with the electrolytes on (a) Sample A and (b) Sample D. 5542


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Sample D triggered the ring opening reaction of EC by the −CF3 electron-withdrawing group. However, a new unknown peak was found at 285.7 eV, which did not appear for Sample A. We suspect that the addition of Sample D may have triggered a new reaction beyond 4.4 V, causing new SEI formation in the system. The XPS results confirm the SEM analysis that O− greatly catalyzed a new SEI formation onto the LLNMO. Figure 9c depicts the O 1s spectra of the LLNMO electrode used with Sample A, showing that a multicomponent lode with an oxygen atom at 528−536 eV was deconvoluted and revealing peaks for the CO of organic SEI at 533 eV and Li2CO3 at 532 eV.38−45 Highly abundant Li2CO3 and lower amounts of organic SEI were found, which is consistent with the conclusions of Figure 9a,b. Interestingly, a third new unknown peak in Figure 9d is found at 529.9 eV, which did not appear for Sample A. This new compound may correlate with the triggering of new SEI formation by O− from the addition of Sample D. Figure 9e shows the Li 1s spectra of the LLNMO electrode used with Sample A, which consists of a main compound, Li2CO3, with a peak at 54.8 eV.38−45 Figure 9f indicates that the electrode used with Sample D had a low natural abundance of Li2CO3 formation on the surface. This result is in agreement with those of Figure 9a−d. In addition, a peak located at 49.1 eV demonstrated the new SEI formation. Figure 9g shows the N 1s spectra of of the LLNMO electrode used with Sample A, which does not reveal and discover any N related compounds on surface. However, Figure 9h displays an exciting result that a new compound consists of −C−N− bonding at 400.2 eV52 is created on the LLNMO’s surface, which can be demonstrated to the new SEI formation. 3.9. Analysis of Cyclic Voltammogram. Figure 10 shows the CV results of all of the electrolyte systems, which display similar electrochemical behavior consisting of two main redox reactions and one irreversible reaction in the first cycle. These reactions can be assigned to the Ni2+/Ni4+ at 3.88 V, O2−/O− at 4.21 V, and oxygen activation after 4.4 V, O−/O2− at 4.32 V, and Ni4+/Ni2+ at 3.77 V.5,47,53 The four samples exhibited similar redox behavior except for the reaction’s current intensity. For the reaction of Ni2+/Ni4+, Sample C shows the lowest current intensity (1.43 mA cm−2) compared with the others owing to the electro-donating group (−CH3) limiting the electron cloud density on the electrode’s surface and leading the electrochemical polarization increment of the Ni2+/ Ni4+ reaction. In addition, Sample D showed the highest current intensity (1.73 mA cm−2) of the O2−/O− reaction, indicating that the electron-withdrawing group (−CF3) effectively attracted electrons from the lattice oxygen ions (O2−), which were converted to O− in great quantity. Regarding this strong electron-withdrawing effect, the higher current density of O− conversion in Sample D was increased and accelerated after the potential of 4.39 V compared with that of Sample A (4.42 V). Furthermore, the in situ gas chromatography−mass spectrometry (GC−MS) result in Figure 9 reveals that the oxygen evolution of Sample D was weak, implying that the O− catalyzed the new SEI formation in Figure 7d at this potential, not the reaction of oxygen activation (O−/O2). Hwang and colleagues argued that the oxygen activation leads to the possibility of Mn participation where some Mn4+ ions are reduced to Mn3+ after the first discharge, which can participate in subsequent cycles.5 Although this Mn4+/Mn3+ reaction is reversible, the activation of oxygen

and CO2 evolutions were independent of the lattice oxygen of LLNMO, indicating that the synthesis of CO and CO2 in the battery must correlate with the carbonate oxidation catalyzed by O−. The possible reaction mechanism after 4.21 V is as follows. O2 − → O− + e−

O− + C (from carbonate electrolyte) → CO

O− + CO → CO2 The second gas evolution occurred after 4.5 V. This time, the CO and CO2 were accompanied by the oxygen evolution from bulk LLNMO, demonstrating that some O− directly oxidized to O2 and synthesized CO and CO2. The reaction mechanism after 4.5 V is as follows. O2 − → O− + e−

2O− → O2 + e− 1/2O2 + C (from carbonate electrolyte) → CO 1/2O2 + CO → CO2

Compared with the result of Sample A (LLNMO), Figure 8b shows that Sample D delivers improved gas evolution. 3.8. X-ray Photoelectron Spectroscopy Analysis. Figure 9 displays the X-ray photoelectron spectroscopy (XPS) spectra of the LLNMO electrode surface where Samples A and D were used; the spectra were detected after finishing the 20th cycle at 60 °C in Figure 5. Figure 9a shows the C 1s spectra of the LLNMO electrode used with Sample A. The carbonate group was assigned to 288−292 eV and the superimposed group was assigned to 283−287 eV. These two groups contained the C−C (graphite) peak at 284.4 eV, the organic SEI (ROCO2Li) peak at 289−290 eV, the Li2CO3 peak at 290 eV, the PVdF peak at 286.2 eV, and the Li2C2 peak at 283 eV.38−45 Figure 9b shows an entirely separate result. The peak intensity of Li2CO3 was substantially decreased by adding Sample D, indicating that the SEI on the electrode used with Sample D had less abundant Li2CO3. Hwang and colleagues reported that LLNMO demonstrates an O2−/O− redox couple at the potential range of 4.4−4.8 V. In situ Raman and other observations revealed that O− reacted with Li+ and formed Li2 O. Li 2O and CO2 further formed Li2CO3.5,46 Our modification of the reaction mechanisms, based on our XPS results, is as follows. O2 − → O− + e− 2O− → O2 + 2e−

1/2O2 + 2Li+ + 2e− → Li 2O Li 2O + CO2 → Li 2CO3

Some of the O2− ions are continuously reduced to oxygen gas, which triggers a phase exchange from the layer to spinel structure of LLNMO.5,47−50 The oxygen gas is then reduced with lithium ions and electrons through an oxygen reduction reaction (ORR) to form Li2O. According to the above analysis, this research proposes as a reaction mechanism that Sample D may eliminate the formations of CO2 and Li2O. Figure 9b shows high amounts of polycarbonate at around 288.8 eV compared with Figure 9a.51 We assumed that the addition of 5543


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from the initiation of oxygen evolution, which was in agreement with the XPS results of Figure 8. Compared with the results of Sample D, Sample C (with the strong electron-donating group) did not trigger the catalytic performance on LLNMO, indicating that the reaction of the oxygen evolution behavior was weaker than those of the other samples. 3.10. Analysis of Transmission Electron Microscopy. Figure 11 shows the SAED patterns and HR-TEM images of

Figure 11. SAED patterns of LLNMO electrodes of (a) Sample A, (b) Sample D and the HR-TEM microscopies of (c) Sample A, (d) Sample D electrodes after 20th cycle at 60 °C in Figure 5.

Figure 9. XPS spectra of LLNMO electrodes with the electrolytes after the 20th cycle at 60 °C. C 1s spectra of (a) Sample A and (b) Sample D. O 1s spectra of (c) Sample A and (d) Sample D. Li 1s spectra of (e) Sample A and (f) Sample D. N 1s spectra of (g) Sample A and (h) Sample D.

the Sample A and Sample D electrodes after cycle testing (Figure 5). The lattice points of the LLNMO structure on Sample A become unclear, according to the results, because the crystalline structure was demolished by the oxygen activation during the first charge (Figure 11a). This result explains the poor cycle performance in Figure 5. However, Figure 11b shows direct insights into the crystal structure microregions on the Sample D electrode, which indicate clear lattice points. In addition, the contrast of the (111), (220), and (311) facets are striking compared with Figure 11a, demonstrating that the new SEI formed from LITFB dramatically enhances the stability of the LLNMO structure. This result validates the gas evolution observation in Figure 9b, which exhibits less CO2, CO, and O2 formation during the first charge. Furthermore, the layer and spinel structures of the Sample A electrode are displayed in Figure 11c; they indicate that the reaction of O2−/O− at the potential of 4.21 V during the first charge signficantly triggers the irreversible phase change reaction from layer to spinel. Interestingly, the irreversible phase change was not observed in the Sample D electrode (Figure 11d); the image shows only the layer structure comprising the (003) facet.56 In addition, an amorphous compound was found at the edge of the surface in Figure 11d, indicating that the new SEI significantly affected the performance of LLNMO. In contrast with Figure 11d, Figure 11c exhibits no SEI formation. 3.11. Analysis of in Situ Synchrotron Measurements. Figure 12 shows the in situ XRD patterns of the Sample A and

Figure 10. First cycle of cyclic voltammogram of Sample A, B, C, and D in LLNMO electrode at the scanning rate of 0.1 mV s−1.

leaves the LLNMO highly reactive and unstable (i.e., similar to a catalyst or reactive species).5,54−56 This can explain why Sample A had the poorest battery cycle performance (Figure 5). According to these results, the ORR reaction of Li2O leads to the Li2CO3 formation at the potential of 4.42 V on Sample A 5544


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Figure 13. In situ Ni K-edge XANES spectra of (a) Sample A, (b) Sample D and Mn K-edge XANES spectra of (c) Sample A, (d) Sample D at the first charge.

when charged to 4.4 V, (3) Ni3+ oxidizes to close to Ni4+ when charged to 4.6 V, and (4) the Ni4+ forms in the fully charging material.60−63 However, the Sample D shows extremely different reaction sequences compare with Sample A. The reaction sequences can be described in terms of Figure 13b that (1) Ni2+ exists in the uncharging material, (2) Ni2+ is almost not oxidized when charged to 4.4 V, (3) Ni2+ oxidizes to between Ni3+ and Ni4+ when charged to 4.6 V, and (4) the Ni4+ forms in the fully charging material. These observations suggest that a new reaction arises and is initialized by the Ni2+ between 4.2 and 4.4 V. According to the CV results in Figure 10, the in situ XANES experiments confirm that the Sample D showed the highest current intensity at 4.21 V, implying the formation of this new reaction is included in the reaction of O2−/O−. Figure 13c,d shows the in situ Mn K-edge XANES of the Sample A and Sample D at 6558 eV is the valence state of Mn4+ due to the electric dipole transfers from a 1s to a 4p state.61 The peaks in Figure 13c (Sample A) shift dramatically at inflection points of 6554 eV to higher energy from OCP to 4.8 V. It is addressed that the peaks shift to higher energy with increasing potential, indicating the Mn4+ ions are partly reduced to Mn3+.61 On the contrary, the peaks in Figure 13d (Sample D) almost no shift at inflection points of 6554 eV compare with the Sample A, implying that the Mn4+ ions remain exist and stable in the lattice by the LiTFB salt effects. The pre-edge absorptions at 6539 and 6541 eV have been assigned to quadrupole-allowed transitions from 1s to 3deg and 3dt2g states for manganese oxides with MnO6 octahedra for LiMn2O4 spinel compounds.60 No peak shifts were observed on both samples in charged state; however, a change in intensity in the pre-edge peaks above 4.4 V was discovered. This behavior indicates that the local distortion of the MnO6 octahedral structure occurred, which implies extraordinary oxygen states.60 Unfortunately, our testing at 4.8 V on Sample D suffers a sudden power shutdown in the Synchrotron center and cannot obtain any results for comparison. The changes on Sample A and Sample D in the Ni K-edge EXAFS spectra during the first charge are presented in Figure 14a,b, respectively. Figure 15a shows the Ni−O bond length in Sample A is 2.013 Å (Ni2+), which decreases to 1.995 Å (Ni3+). When charged to 4.8 V, the length of Ni−O dramatically decreases to 1.869 Å (Ni4+) owing to the lithium ions are deintercalated from the Li2MnO3 with charged compensation

Figure 12. In situ XRD patterns in 18−19° range of (a) Sample A, (b) Sample D and in 20−22° range of (c) Sample A, (d) Sample D.

Sample D during the first charge. Two regions (18−19° and 20−22°) are selected in order to evaluate the changes of dspacing of the lithium layers and the transition metal (TM) layers.6,47,57−59 According to the results in Figure 12a,b, the (003) peak on both samples shift obviously from higher to lower angle, indicating the d-spacing of the lithium layers expands owing to the increment of the c-parameter by the deintercalation of lithium ions and causes the strong repulsion to the neighboring oxygen layers. Compared with the (003) peak observation, the intensities of (020) and (110) peaks between Sample A and Sample D show dramatically different behaviors. The lithium ions in the TM layers will be deintercalated during the high voltage operation at the first cycle, causing the demolishment of the LiTM6 ordering and the decay of the (020) and (110) facets.6,47,57−59 Therefore, the preservation of these peaks provides strong evidence to the fact that the LiTM6 ordering in the TM layers of Sample D with the addition of LiTFB salt is stable and the lithium ions are remain existed. However, the Sample A shows extremely unstable, indicating the changing of the honeycomb ordering of TM layers is already performed. Figure 13a,b shows the in situ Ni K-edge XANES of the Sample A and Sample D during the first charge. In Figure 13a, it is clear to observe that the valence state of Ni changes to the high energy from OCP to fully charged (4.8 V) on Sample A, indicating the reaction sequences were followed by (1) Ni2+ exists in the uncharging material, (2) Ni2+ oxidizes to Ni3+ 5545


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Scheme 1. Proposed Reaction Mechanism of Poly Ionic Liquid SEI Formation on LLNMO Electrode with Sample D

Figure 14. k3-Weighted Fourier transforms for the Ni absorber of (a) Sample A, (b) Sample D and Mn absorber of (c) Sample A, (d) Sample D at the first charge.

migration and oxygen activation. In the case of Sample D, the strong electron-withdrawing group (−CF3) on benzimidazole activated a new reaction that accompanies O2−/O− at the potential of 4.21 V on the surface of LLNMO in which the EC is then adsorbed with O−.65 In addition, some reports have discussed that the Ni2+ is used to catalyzes the ring opening polymerization66,67 and the cathodic decomposition of EC.68 The proposed reaction mechanism in this research suggests that the Ni2+ catalyzes the O− adsorbed EC into alkyl carbonate and O−. The benzimidazole ring provides one lone pair on each N atom, indicating that a polymerization69,70 takes place with alkyl carbonate to form polyionic liquid SEI on LLNMO’s surface. The XPS measurements in Figure 8 were used to identify the chemical composition of polyionic liquid SEI, namely the unknown peaks in the C 1s, O 1s, Li 1s, and N 1s spectra. Because the polyionic liquid SEI provided lithium ions for its infrastructure, part of the lithium lattice no longer deintercalated during the charge−discharge process. The irreversible reactions of Ni4+ migration and oxygen activation during the first cycle were able to eradicate and the Ni2+ substantially maintain the layer structure of LLNMO. The in situ GC−MS results in Figure 9, the TEM images in Figure 11, and the in situ Synchrotron measurements in Figures 12−15 confirm that the electrolyte additive performance of LITFB with LLNMO is beneficial.

Figure 15. (a) Ni−O and (b) Mn−O bond length changes estimated from in situ EXAFS measurements for the first charge.

on the oxygen ions.60 Interestingly, the Sample D effectively shortens the reduction of the Ni−O bond length, especially at the potential of 4.6 V (1.961 Å) compares with the Sample A (1.887 Å). This information confirms the in situ GC−MS result in Figure 9 that the effects on the LiTFB salt addition can significantly decrease the oxygen evolution from the material structure. The changes on Sample A and Sample D in the Mn K-edge EXAFS spectra during the first charge are presented in Figure 14c,d, respectively. In Figure 15b, the bond length of Mn−O on both samples varies slightly compares with the huge changes on Ni−O. However, it still can be found that the peak intensity of the Mn−O in Sample A decreases sensitively after the charging potential achieves to 4.4 V and more, which is dedicated by release of the oxygen atoms and mainly to the increase of the spinel-like structure.64 In comparison with Sample D, the peak intensity of the Mn−O almost no change when charged to 4.4 V and further 4.6 V, indicating the less formation of spinel structure. These results are compatible to the observations of TEM in Figure 11, Sample D maintains the preservation of the layer structure after electrochemical reaction. On the other hand, the peak intensities of MnMetal (Mn-M) bond on both samples are independent to the charging process from 4.4 to 4.8 V. 3.12. Proposed Reaction Mechanism Underlying New Solid Electrolyte Interphase Formation. The results of this study confirm that Sample D greatly reacted with the alkyl carbonate and therefore less Li2CO3 formation on electrode’s surface was found. A possible reaction mechanism is shown in Scheme 1. At the potential of 4.5 V in Sample A, the reaction plateau indicated that most of the lithium ions deintercalated from the structure lattice of LLNMO through a phase change (layer to spinel), which caused a mutual reaction of Ni4+

4. CONCLUSION The strong electron-withdrawing group (−CF3) branched on benzimidazole was used to make a lithium ion battery that was robust at high temperatures and voltages. This study developed a polyionic liquid SEI formation on an LLNMO electrode. This new SEI eliminated gas evolutions, maintained the original cathode structure, and resulted in the high efficiency electrochemical performance of LLNMO. This new electrolyte additive delivered a dramatic impact that can improve and extend Li ion battery life for EV and HEV applications.

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.chemmater.7b00824. The ionic conductivity measurements; the activation energy calculations; the DSC results; the anodic 5546


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electrochemical stability measurement and the cycle performance at 60 °C (PDF)

AUTHOR INFORMATION

Corresponding Author

*F.-M. Wang. E-mail: [email protected]. ORCID

Fu-Ming Wang: 0000-0003-4407-3554 Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The authors are grateful for the financial support from the Ministry of Science and Technology (MOST) of Taiwan, R.O.C, under grant numbers 102-2221-E-011-016-MY3, 1043113-E-011-002, 104-2745-8-011-001, 105-3113-E-011-002, 105-2628-E-011-005-MY3, 105-2811-E-011-021, 106-3113-E011-001, 106-2923-E-036-002-MY3, and the facilitates support from the National Synchrotron Radiation Research Center (NSRRC).

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ABBREVIATIONS DEC, diethyl carbonate EC, ethylene carbonate EV, electric vehicle HEV, hybrid electric vehicle SEI, solid electrolyte interphase

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REFERENCES

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