Facile in Situ Syntheses of Cathode Protective Electrolyte Additives for

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Facile In Situ Syntheses of Cathode Protective Electrolyte Additives for High Energy Density Li-Ion Cells Ilya A. Shkrob, Binghong Han, Ritu Sahore, Adam P. Tornheim, Lu Zhang, Daniel P. Abraham, Fulya Dogan, Zhengcheng Zhang, and Chen Liao Chem. Mater., Just Accepted Manuscript • DOI: 10.1021/acs.chemmater.8b05261 • Publication Date (Web): 18 Mar 2019 Downloaded from http://pubs.acs.org on March 19, 2019

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Chemistry of Materials

Facile In Situ Syntheses of Cathode Protective Electrolyte Additives for High Energy Density Li-Ion Cells Ilya A. Shkrob,a,b Binghong Han,a Ritu Sahore,c Adam P. Tornheim,a Lu Zhang,a,b Daniel P. Abraham,a Fulya Dogan,a Zhengcheng Zhang,a Chen Liao*,a,b aChemical

Sciences and Engineering Division, Argonne National Laboratory, 9700 South Cass

Avenue, Argonne, Illinois, U.S.A.

b

Joint Center for Energy Storage Research, Lemont, Illinois

60439, U.S.A. c Roll-to-Roll Manufacturing Group, Energy & Transportation Science Division, Oak Ridge National Laboratory, Oak Ridge, Tennessee, U.S.A. Abstract Increasing the energy densities of Li-ion batteries necessitates operation of layered lithiated oxide cathodes at potentials exceeding 4 V vs. Li/Li+. When continually exposed to such high potentials, these materials gradually deteriorate unless protected by sacrificial agents called electrolyte additives. During the formation cycles, these electrolyte additives decompose on the electrodes forming thin protective layers of insoluble products impeding further deleterious reactions. Some of these electrolyte additives spontaneously react when introduced into the electrolyte to yield specific surface-modifying products that alone protect the cathode; in other words, the nominal additive is the precursor and the secondary product is the protective agent. Guided by this insight, we used molecular engineering to obtain such surface-active secondary products in situ with 100% yield. Two of these electrolyte additives proved exceptional in that they delay both the impedance rise and capacity fade in the Li-ion cells. We demonstrate this protective action and scrutinize the activation of these additives in the electrolyte. By "taming" spontaneous reactions and regaining full control over the chemical structure, new avenues open to targeted synthesis of electrolyte additives extending the operation of high voltage Li-ion batteries.

* Corresponding author. E-mail: [email protected], Tel.: 630-2524597

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1.

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Introduction Due to their exceptional cycling stability and high energy, lithium-ion batteries (LIBs) with

layered transition metal (TM) oxide cathodes and graphite (Gr) anodes are ubiquitous in portable electronics and electric vehicles. During electrochemical cycling of such cells, lithium ions deintercalate from one electrode and intercalate into other, with an astounding efficiency and reversibility. Although LIBs have been commercialized since 1991, layered oxides such as LiMO2 (where M = Co, Ni, Mn) are still the cathode materials of choice in commercial batteries.1 Further increase in the energy density requires cathode materials that operate at voltages exceeding 4.0 V without performance degradation during aging of the cells. LiNixMnyCozO2 (x+y+z=1) cathodes are particularly promising in this regard as their properties can be fine-tuned to negotiate complex trade-offs between the energy density, safety, and cycling stability.2 The electrolyte plays a pivotal role in the protection of the electrodes, which is required for extended cycling. The molecules in the electrolyte react on energized electrodes and modify their surface, e.g., by forming protective layers. The classical example of this kind is the formation of solid electrolyte interphase (SEI) on lithiated graphite: the carbonate molecules in the electrolyte are chemically reduced forming a thin (1-5 nm) layer of insoluble ionic and polymeric products.3 This layer transports Li+ ions but hinders further reduction of electrolyte by inhibiting the access of solvent molecules to the reducing Gr surface and radicals towards the solvent. In this way, the electrochemical stability of electrolyte extends beyond the thermodynamic limit through the formation of a kinetic barrier. The SEI is formed during the first lithiation half-cycle and protects the electrolyte over many subsequent cycles. A similar process occurs on the cathode, but the protective layers are much thinner 4 and may consist of a single or multiple molecular layers. Still, these protective layers can significantly slow down surface modification, corrosion, and loss of TM ions and oxygen from the surface of delithiated cathode.5 Importantly, the cathode-electrolyte interface makes the greatest contribution to the cathode impedance, which in turn is the major component in a full-cell impedance.6 As the cell ages, the cathode and full-cell impedances gradually increase eventually reaching values at which the charge and discharge of the cell becomes prohibitively slow.7 This phenomenon is referred to as the impedance rise. To improve the surface protection and cell performance, electrolyte additives are frequently used. Empirically, a few weight percent of these compounds can significantly delay 2 ACS Paragon Plus Environment

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performance deterioration of the cell. Such additives proved to be efficient means for improving capacity retention and reducing the impedance rise in high-voltage Li-ion cells.8-14 Without these additives, electrolyte decomposition is difficult to avoid, and the electrochemical performance rapidly deteriorates.15 While that much is known, the details of this protective action remain unclear. In particular, it has been speculated that certain "primary" (nominal) additives become activated through their chemical transformation in the electrolyte, and the protective agents are, in fact, secondary products of such reactions. A recent example of this kind was given in our previous study,5 where we demonstrated that P(OSiMe3)3 additive spontaneously reacts with LiPF6 in the electrolyte to yield a secondary product, PF2OSiMe3, that served as the cathode protective agent in this electrolyte. Amongst all of the reaction products only PF2OSiMe3 contributed to the cathode protection, while most of P(OSiMe3)3 was consumed in side reactions not leading to the formation of products providing the cathode protection. That led us to an idea: instead of letting such reactions to occur by chance, can we carry them on purpose in a controlled environment, obtain the desired product in high yield, and use it directly as electrolyte additive without separation and purification? Furthermore: can we engineer a "primary" additive in such a way that it gives only the specific beneficial secondary product in situ quantitatively? Here we answer both of these questions positively. Specifically, we demonstrate that protective agent 2 shown in Figure 1 can be obtained in situ with ~100% yield by adding precursor 1 directly to the LiPF6 containing carbonate electrolyte as a "primary" additive.

In our

electrochemical tests, two products of this reaction (2a and 2b) outperformed several additives previously evaluated in our laboratory.16 In this case, protective agent 2 is the predominant secondary product, with no side products of consequence to its designated action. Previous studies of Lucht and co-workers

19,20

indicated that 2a (obtained through a different synthetic route, see

below) improves thermal stability of electrolyte; here we demonstrate that both 2a and 2b can counter impedance loss and capacity fade under normal operation conditions. We remind that compound 2a in Figure 1 has been obtained ex situ by heating of oxalic acid and LiPF6 with silicon tetrachloride,17,

18

and also via in situ reactions between lithium

bis(oxalato)borate and LiPF6 at 85 oC and lithium oxalate with PF5 19, 20. Here we demonstrate a novel synthetic route to 2a and 2b via facile cleavage of trimethylsilyl group from 3 ACS Paragon Plus Environment


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bis(trimethylsilyl) carboxylates 1a and 1b in a reaction of this precursor with LiPF6 that yields trimethylsilyl fluoride (TMSF) as byproduct, rapidly and without thermal activation (see Figure 1). Using this synthetic route, stable reagents 1a and 1b can be reacted in situ with a LiPF6containing carbonate electrolyte to yield 2a and 2b, while volatile TMSF (the byproduct of this reaction) can be easily removed by purging or outgassing of the solution. There is no need to separate and purify 2a and 2b as they are the only products of this reaction, and residual traces of TMSF do not affect electrolyte performance. Below we detail the synthetic procedures, examine the reaction products, and demonstrate their use as electrolyte additives in high-voltage LIBs. Our study suggests that "taming spontaneity" - that is, following the trail of chemical reactions spontaneously occurring in an electrolyte and then improving the yield and specificity of these reactions - can be a potent new approach to discovery of new electrolyte additives. To reduce the volume, some sections, figures, and tables are placed in the Supporting Information (SI). These are referred to in the text with the letter "S" next to them, as in Figure S1. 2. Methods All electrodes were fabricated at Argonne’s Cell Analysis, Modeling and Prototyping (CAMP) facility. The cathode was composed of 90 wt% Li1.03(Ni0.5Mn0.3Co0.2)0.97O2 (NMC532, TODA), 5.0 wt% polyvinylidene fluoride binder (PVdF, Solvay, 5130) and 5 wt% conductive C45 carbon black particles (Timcal) coated on an aluminum current collector. The loading density was 9.12 mg/cm2 and the porosity of the calendered electrode was 33.5%.

The Gr electrode was

composed of 92 wt% ConocoPhillips A12 graphite, 6 wt% PVdF binder (Kureha, 9300) and 2 wt% C45 carbon black particles (Timcal) coated on a copper current collector. The loading density was 6.07 mg/cm2 and the porosity was 35.6 %. The microporous separator (Celgard 2335) was 25 µm thick, 2.01 cm2 area, with 39% porosity. The total pore volume was ~ 7 µL, and the volume of added electrolyte was 25 µL, with 5 µL under the anode, 10 µL between the anode and separator, and 10 µL between the cathode and separator. See Table I in ref. 14 for more detail. The “baseline” Gen2 electrolyte was composed of 1.2 M LiPF6 in ethylene carbonate/ethyl methyl carbonate (EC/EMC) (3:7 w/w ratio) supplied by Tomiyama Chemical Industry, Japan. Six other electrolytes tested in this study were the baseline electrolyte containing 0.5, 1.0, or 2.0 wt%


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2a and 2b as additives. The cells containing these electrolytes are referred to as the 2a and 2b cells, respectively. Compounds 1a and 1b were synthesized according to refs. 21, 22; all other compounds were obtained from Sigma-Aldrich in their purest form and used as is. To obtain 2a and 2b, 1 equiv. LiPF6 and 2 equiv. 1a or 1b, respectively, were placed into a Nalgene vessel and stirred in dimethyl carbonate (DMC) for 48 h under the inert atmosphere. The solvent and the volatile reaction product (FSiMe3, TMSF in Figure 1) were removed by evaporation, and nuclear magnetic resonance (NMR) analyses of the solid residues were carried out as described below (see section S1 for spectroscopic parameters). To follow the progress of these reactions, NMR spectra of reaction mixtures were obtained using a Bruker Avance III 300 spectrometer (300 MHz, 1H; 282.3 MHz,

19F;

121.4 MHz,

31P).

Representative NMR spectra are shown in Figures 2 (for 2b) and Figures S1 and S2 (for 2a). The structural reasoning to identify the NMR spectra as originating specifically from products 2a and 2b (whose structure is shown in Figure 3) is scrutinized in section 3.4. Chemical shifts () are reported in parts per million (ppm) relative to tetramethylsilane. Following ref.

23,

fluorinated

ethylene propylene tube liners capped with polytetrafluoroethylene plugs (3 mm, Wilmad/VWR) were used to avoid direct contact between the corrosive electrolyte and glass NMR tubes. To lock the spectrometer, the gaps between the liner and the glass tubes were filled with acetonitrile-d3 or acetone-d6. To follow the reactions of 1 wt% 1a or 1b in Gen2 electrolyte, the solution was stirred for 2 days at 25 oC, and the aliquots were periodically collected and analyzed using NMR spectroscopy. Only 2a and 2b were found in these solutions. For the reactions between 1a and 2a with different LiPF6 ratio, see Figures S1 to S10 for representative

31P

and

19F

NMR spectra

obtained during and after the reaction. In our electrochemical tests, the electrodes were dried at 100°C prior to use, and 2032-type coin cells were assembled in Ar atmosphere. All coin cells were tested in triplicate. The figures in the text show average parameters for the three cells (the plots with the standard deviations are shown in the SI, see Figures S11 to S17). The performance metrics summarized in Table 1 were averaged over these three cells with the standard deviations given in the same table to illustrate variability in the data.



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The cathode, the Gr electrode, and separator were 14, 15 and 16 mm in diameter, respectively. The specific capacity of the cells is given per weight of the lithiated oxide material in the cathode, as initially the entire Li inventory was contained in the cathode. The currents are given as C-rates, with a rate of 1C corresponding to full charge in 1 h (~2.8 mA at 1C charge). All electrochemical data were collected using MACCOR cyclers at 30

oC.

Galvanostatic

charge/discharge cycling was conducted between 3.0 and 4.4 V using the protocol detailed in ref.24 Briefly, this protocol included five formation cycles at a C/10 rate (ensuring that the protective layers form on the electrodes) followed by a sequence of aging cycles at a C/3 rate which include a 3 h potentiostatic hold at 4.4 V between the cycles. Hybrid pulse power characterization (HPPC) tests using 10 s, 2C pulses were performed every 20 cycles throughout aging of the cells. Area specific impedance (ASI) was obtained as a charged cell was incrementally discharged at a C/10 rate, by comparison of voltages at the beginning and the end of the discharge pulse. A graphic description of impedance measurement is provided in Figure S18. The power density P (in W/cm2) was estimated using the ASI values at 90% state of charge and the known voltage at rest (~4.08 V) using equation P = 3 ×

(𝑉𝑟𝑒𝑠𝑡 ― 3) 𝐴𝑆𝐼

justified in ref 25. The HPPC sequence

was followed by two diagnostic cycles at the 1C and C/10 rates, respectively. To compare the effectiveness of different electrolyte additives, the figures of merit in energy (FOME) and in power (FOMP) for different cells can be used.16 These quantities are defined as the cycle number at which a performance metric reaches 80% of the initial value in the baseline cells. More detail of this ranking are given in section 3.2. Transmission Electron Microscopy (TEM) studies were performed using a JEOL 2100F microscope operating at 200 kV at the Center for Nanoscale Materials of Argonne National Laboratory. Electron energy loss spectra (EELS) with the energy resolution of 0.4 eV were obtained using the same instrument. The sample was prepared by removing the oxide material from the aged cathode and dispersing it in dimethyl carbonate by sonication. A small drop of this suspension was evaporated on the Cu grid. X-ray Photoelectron Spectroscopy (XPS) characterization was performed using a PHI 5000 VersaProbe II System (Physical Electronics) with a base pressure of 2×10-9 torr. The cathodes were harvested from the aged cells and lightly washed with dimethyl carbonate prior to these spectroscopic observations. The photoelectron spectra were obtained in the fixed analyzer 6 ACS Paragon Plus Environment

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transmission mode using an Al Kα radiation (hν = 1486.6 eV, 100 µm beam, 25 W) with Ar+ and electron beam sample neutralization. Peak fitting was performed using Shirley background correction and the Gaussian-Lorentzian curve synthesis available in CasaXPS suite. XPS spectra were aligned to the graphitic carbon at 284.7 eV, and manufacturer-reported relative sensitivity factors were used to normalize these spectra for elemental analyses. To compute molecular structures and reaction energetics, a density functional theory (DFT) method with the B3LYP functional26, 27 and 6-31+G(d,p) basis set from Gaussian 09 suite were used. The chemical shifts and spin-spin J coupling constants for magnetic nuclei were estimated using Gaussian 0928 for geometry optimized gas phase species.

3. Results and Discussion. 3.1. Synthesis of the cathode protective agents. The key intermediates in the synthesis of 2 are the bis-trimethylsilyl carboxylates 1 shown in Scheme 1. Using a modiﬁed procedure of Palomo et al.,29 the yield of these compounds was improved signiﬁcantly from below 30% to over 60%. 3-(Trimethylsilyl)-2-oxazolidinone provided the trimethylsilyl moiety, and the reaction was facilitated through the precipitation of 2-oxazolidinone. The ex situ syntheses discussed in section 3.4 were carried out to establish product speciation as a function of the mole ratio between the precursor (compound 1) and LiPF6, as the hexafluorophosphate anion potentially provides multiple reaction sites for the carboxylate moieties. The volatile product (TMSF) and unreacted precursor can be readily removed by evaporation or gas purging. Furthermore, product 2 can be generated in situ by adding 1-10 wt % of 1 to Gen2 electrolyte (which contains 1.2 M LiPF6). This reaction readily proceeds at room temperature (48 h) yielding almost exclusively 2. According to NMR spectroscopy (see section 3.4), addition of a few wt% of 1 to the Gen 2 electrolyte results in complete transformation of 1 to 2 after 48 h at room temperature. No significant difference in the performance of electrolytes containing 2 obtained ex situ and electrolytes containing 2 generated in situ was found in our electrochemical tests. Given these observations, only in situ generated 2a and 2b electrolytes will be discussed in the next section. 3.2. Electrochemical testing. 7 ACS Paragon Plus Environment


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Ideally, an electrolyte additive should protect both electrodes, which is difficult to achieve in practice given the considerable difference in chemistry between these two electrodes. This difficulty is compounded when the surface-modifying agent is a secondary product of a spontaneous reaction between the precursor and electrolyte, as such uncontrolled reactions yield numerous products, some of which can adversely affect stability of one or both of the electrodes. Our synthetic approach (section 3.1) eliminates the latter concern, potentially allowing to design solute molecules capable of protecting both of the electrodes. To achieve this goal, the cathode protective agent needs to behave in a certain way towards the lithiated Gr electrode. We remind that no organic molecule (including the solvent) is stable on the fully lithiated Gr electrode; if a solute molecule can access the highly reducing surface of this electrode, it will react and decompose. Stable SEI is known to form when the lithiated Gr reacts with ethylene carbonate in the electrolyte yielding specific decomposition products.30 The solute competes with the solvent for the electrons at the energized surface. Thus, there are two related concerns. First, electrochemical reduction of the solute adds to irreversible capacity loss through the formation of ionic compounds permanently trapping Li+ ions. Second, the modified SEI may not have the chemical and mechanical properties required for prolonged electrochemical cycling of the cell. A useful metric indicative of such problems is the fractional capacity loss in the first ("formation") half-cycle, when the SEI is formed. The capacity loss in this half-cycle is related to Li+ ion trapping in the anion products of solvent and/or solute decomposition. By examining data in Table 1 and capacity retention profiles shown in Figure 4, we can make the following observations. The first cycle discharge capacity for the 2a and 2b cells was always lower than in the baseline cell, suggesting that these additives become reduced and decomposed on the lithiated Gr electrode. Nevertheless, for the 0.5 wt% and 1.0 wt% 2b cells, the capacity loss during the first cycle was only slightly lower compared to the baseline cells. Furthermore, for the 0.5 wt% and 1.0 wt% 2a and 2b cells, while the initial (cycle-3) capacity is lower, the capacity retention after 119 cycles is improved and the mean loss per cycle is smaller, suggesting that this initial capacity loss soon becomes compensated by improved protection of the cathode. The latter is not surprising as the continuous loss of TM ions from unprotected oxide cathode (especially Mn2+ ions) is known to accelerate the capacity loss through their deposition in the SEI layer on the Gr electrode and catalyzing SEI decomposition.7 By protecting the oxide cathode, the additives indirectly protect the Gr electrode. 8 ACS Paragon Plus Environment

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The initial capacity loss observed in the 2a and 2b cells suggests that these additives are reduced at a somewhat higher potential than the solvent; in this case, as the Gr electrode becomes lithiated, these additives decompose first, yielding products that may not form high-quality SEI unlike the solvent itself. This scenario is also suggested by differential capacity (dQ/dV) profiles obtained during the first charge of the cells shown in Figure 5. In the baseline cell, there is a prominent peak at 3.0 V due to reduction of the ethylene carbonate at 0.7 V vs. Li/Li+. In the 2a and 2b cells, there are additional peaks observed at 2.0 and 2.6 V in the NMC532//Gr full cell (corresponding to +1.7 V and 1.1 V vs. Li/Li+, respectively, with the potential of the cathode estimated to be around 3.7 V vs. Li/Li+), which become more prominent as the concentration of the electrolyte additives increases.16 A lower redox potential for 2b is due to strong electron withdrawing effect of the fluorine in the malonate ring that was intentionally introduced to provide this effect. Having a harder-toreduce additive has paid off, as the initial capacity of the 2b cells was almost on the par with the baseline cells, unless so much of 2a was in the electrolyte that SEI became composed predominantly of the breakdown products of 2b on the Gr electrode, and SEI quality deteriorated. Thus, while compounds 2a and 2b do not actively protect the Gr electrode by forming a betterquality SEI than the solvent, they minimally interfere with the SEI formation. Furthermore, over time they improve the SEI quality by eliminating one of the leading causes for its gradual deterioration during aging of the cell, which is loss of TM ions from the cathode and their deposition into the SEI, as we discuss later.7 Turning to the cathode protection, one of the key functions of the electrode additive is to counter the gradual increase in cell impedance due to aging of the cathode and build up of decomposition products that slow down Li+ ion transport between the cathode and electrolyte. Figure 6 shows ASI during aging of the cell obtained every 20th cycle (see Table 1 for the initial and final ASI and power density at 4.08 V). The voltage dependence of ASI has the characteristic “bath tub” shape due to slower Li+ ion transport at either one of the electrodes becomes fully lithiated. While the initial ASI of the baseline cell is somewhat lower compared to the additivecontaining cells, the fractional impedance rise and loss of power density in the presence of 2a and 2b is considerably lower (Table 1), and it becomes lower still as the concentration of these additives increases. Thus, we can state that both 2a and 2b are indeed the cathode protective agents.



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This function can also be demonstrated by measuring parasitic currents flowing across the cell during a potentiostatic hold at a high voltage (4.5 V vs Li/Li+). As the cathode potential reaches and stays at the values at which the electrolyte can be oxidized, the associated leakage current characterizes the quality of electrolyte interface instability.31 Comparing kinetics for this leakage current in Figure 7 and Figure S19, we see that addition of 1 wt% 2a and 2b lowers the parasitic current as it settles after the initial decrease suggesting improved passivation of the cathode in the presence of these additives. This effect was observed at all concentrations of electrolyte additives, with the ASI rise being lower at higher concentration of these additives. Finally, we adopt the Figure of Merit (FOM) approach to rank the electrolytes in their ability to improve the electrochemical cycling performance of a cell (Table 1). 16 The FOMs are cycle numbers at which a given performance metric decreases to 80% of its original value for the baseline cell. To give a specific example, the initial energy density of the baseline cell cycled at a C/10 rate was 714 mWh g-1. Based on this definition, the energy FOM corresponds to the number of the C/10 cycle in which the energy density would decrease to 80% of this initial value (= 571 mWh g-1). To estimate this cycle number, the energy density was determined at a C/10 rate after each HPPC sequence every 20th cycle (giving the total of five such measurements over 120 cycles). A straight line was fitted interpolating between these points with the coefficient of determination R2>0.999, and it was used to extrapolate beyond. Through this linear extrapolation the cycle number corresponding to 20% loss in the initial energy density was estimated to be 170. The same strategy was adopted for the power FOM using a 10 s, 2C discharge pulse applied at 4.08 V to calculate and interpolate/extrapolate the power density. Figure 8 summarizes the two FOMs in a 2D chart. Each point in this plot represents an electrolyte composition, with the X-values representing the energy FOM and the Y-values representing the power FOM. The cross hairs in Figure 8 serve as the visual guide to differentiate the cell performance vis-à-vis the baseline cell. Electrolytes that have FOMs in the top right section of this chart give the best performance. According to these FOMs, the cells containing 2 wt% 2a and 1 wt% 2b gave the best overall performance. 3.3. Cathode surface characterization. Changes in the elemental composition are reflective of the chemical transformations in the cathode ocurring as the cell ages. Figure 9 shows X-ray photoelectron spectra (XPS) of the pristine 10 ACS Paragon Plus Environment

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and aged cathodes from the three 2a cells. The sharp C-C peak at 285 eV from the conductive carbon particles on the surface of the cathode can be used for qualitative characterization of deposits (these carbon particles are the integral part of the fabricated electrode, see section 2). As the cell ages and deposits grow thicker and/or denser, the amplitude of this peak becomes smaller due to stronger attenuation of outcoming photelectrons by the deposits. Clearly, in the presence of 2a this masking of the surface increases, and the degree of obstruction correlates with the concentration of this additive (the same was observed for 2b in Figure S20). Also seen in the photoelectron spectra are the 286-288 eV peak from the carbonyl carbons, the 531-536 eV peak (from the P-O and O-P-F speciess) from the oxygen, and 132-139 peak (from the same speciess) from the phosphorus, which all increase with the concentration of 2a and, therefore, can wholly or partially originate from decomposed 2a on the cathode surface. Overall, the features seen in the F1s, O1s, and P2p spectra correspond to the formation of oxofluorophosphates (LiPxOyFz) on the surface due to decomposition of 2a and/or increased decomposition of electrolyte in the presence of 2a. These features are already observed after the formation cycles (Figure S21), and they are also seen in the photoelectron spectra obtained from the 2b cells (Figure S20). Our conclusion is that the LiPxOyFz deposits form early during the cycling and subsequently protect the cathode by slowing oxidative decomposition of electrolyte and TM dissolution. While XPS characterizes chemical composition of the cathode surface, TEM images of the oxide particles in aged cells indicate morphological changes in the material. As seen from Figure S22, panels a to d, in the 0.5 wt%, 1 wt%, and 2 wt% 2a cells, the oxide surface maintains high crystallinity with smooth border lines. In comparison, in the oxide particles from the baseline cells (Figure S22, panel a), pits (indicated by red arrows in the upper image) and surface amorphization/distortion (shown in the lower image) are observed, which is likely due to TM ion loss when unprotected surface is exposed to hydrofluoric acid formed by LiPF6 hydrolysis near the cathode. The advanced state of damage is also seen through the negative shift in the L-edges for TM ions in the EEL spactra that are observed in the baseline cell, but not in the 2a cells (Figure S22, panels e to f). Such shifts indicate a lower oxidation state for TM ions on the surface, which is consistent with the loss of oxygen from the surface and subsurface regions of delithiated cathode.7 This change in the oxidation states for TM ions can also be seen through the decreased L3/L2 ratios for the surface TM ions relative to the bulk. The TM loss rationalizes the faster 11 ACS Paragon Plus Environment


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capacity fade observed in the baseline cells (Figure 4 and section 3.2). We surmise that compounds 2a and 2b decompose on the cathode forming a thin protective layer that slows down acid corrosion and oxygen loss from the cathode surface. All of these features are consistent with the cathode protective function demonstrated through electrochemical measurements in section 3.2. 3.4. Solution chemistry and idenitification of 2a and 2b. As was mentioned in section 2, our identification of 2a and 2b formed in reactions of 1a and 1b with LiPF6 is based on the NMR spectra alone. As one can potentially envision many ways in which the fluorines in the PF6- anion become substituted with the carboxylates, below we consider such reactions in a greater depth, as we have not yet demonstrated that 2a and 2b do have the structure ascribed to them. To facilitate this examination, hereafter we use ox2- for the oxalate and Z2- for cyclo-CFMe(CO2-)2 moieties. In bidentate substitution in the PF6- anion by such X2- dianions, one obtains PXnF6-2n- anions (n=1-3) corresponding to structures a, c, and d in Figure 10 in which these X2- anions become bidentate ligands (the analogous structures for Z2- are shown in Figure S23). According to our DFT calculations in Table 2, for 1a each such substitution would be exergonic, although the absolute reaction heat decreases from n=1 to n=3, while for 1b only the first such substitution would be exergonic. For this first substitution, each P-O bond formation contributes ≈50% towards the total reaction heat. That is, the first and the second P-O bond formation cost approximately the same in energy, so only entropic factors and repulsion between the two proximal PF6- anions cause the preference for the bidentate binding over the bis-monodentate binding shown in Figures 10b and Figure S23, panels b to d. Even more complex species can be imagined, and rich speciation chemistry was expected. Contrary to these expectations, the solution chemistry of 1a and 1b turned out to be remarkable simple. We begin with 1a. Table S1 compares the calculated NMR parameters for 1a with the ones obtained from the analyses of our NMR spectra. Representative 19F and 31P NMR spectra of aged solutions in DMC containing 1:1 equiv. LiPF6 and 1a are shown in Figures S1 and S2, respectively, and the resonance lines are classified in Table S2. After three days, the main reaction product is the 31P19F4 species shown in Figure 10a that is readily recognized through the dt patterns in the axial (ax) and equatorial (eq) fluorines (cf. Figure S1 and Tables S1 and S2). The corresponding tt pattern in the

31P

NMR spectrum is shown in Figure S2. This species accounts for 95% of the 12 ACS Paragon Plus Environment

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original PF6- anions, of which less than 5% remains after 3 days. In addition to this product, there is a minor product (1-3% yield) whose NMR spectra indicate it is a 31P19F2 species. Comparison with calculated NMR properties suggests that it is the P(ox)2F2- complex shown in Figure 10c (see Table S1). Apparently, the second bidentate substitution in the PF6- anion by the oxalate is inefficient, and we found no evidence that P(ox)3- substitution ever occurs. Having a molar excess of LiPF6 can shift the equilibria towards more extensive substitution, and in addition to this 1:1 equiv. solution, 2:1 and 3:1 equiv. solutions were examined (Tables S3 and S4, respectively). In these solutions, too, the parent PF6- anions were almost entirely consumed. When 1a:LiPF6 mole ratio is high (3:1), in addition to these previously identified products, traces of the third product were observed (Table S4). The yield of this product (< 1%) was too small to detect it using 31P NMR, but the single axial and two equatorial fluorines (Table S4) point to a 31P19F

3

species. As there are no complementary proton resonances from the TMS group from a

monosubstituted anion, we believe that this product is the symmetrical bis-substituted species shown in Figure 10b (which agrees well with our DFT calculations in Table S1). Neglecting this minor product, the ratio of P(ox)F4- and P(ox)2F2- species was ~4:1, so even a large molar excess of LiPF6 over 1a does not shift the equilibria towards the complete bidentate disubstitution. This strong preference for the bidentate monosubstitution accounts for the efficiency of the formation of 2a and high purity of this reaction product. As suggested by Table 2, for 1b the bidentate di- and tri- substitutions in PF6- are prohibitive energetically, so 2b shown Figure 3 prevails in the reaction mixture. In 10 wt% 1b in Gen2 electrolyte after two week aging (this concentrated mixture was examine to better understand in situ generation of 2b in the electrolyte), the prevalent species has a

31P19F

4

system with two

inequivalent axial and two equivalent equatorial fluorines that correspond either to a monomer or a dimer PZF4- anions shown in Figure 3 and Figure S23b, respectively (see Table S5). Importantly, the smallest coupling of ≈ 6 Hz (Table 3) cannot be from this 31P19F4 system; it can only originate from the ligand. In addition, there is a -156.1 ppm resonance from the CF*Me fluorine with a qdd pattern, which becomes a dd pattern when the protons are decoupled (so the quartet is from the methyl group). Our DFT calculation (Table 3) suggests that for this fluorine the coupling to 19Feq is negligible, whereas the coupling to

31P

and

19F

ax’

is estimated to be 10.4 Hz and 5.9 Hz,

respectively. The dd pattern in this fluorine corresponds to the J-couplings of 5.7 Hz and 2.2 Hz, respectively. The first one of these constants is close to ≈ 6 Hz in the phosphorus-31, which 13 ACS Paragon Plus Environment


Page 14 of 32

means that the 2.2 Hz coupling is from the axial fluorine. To further probe the structure of this product, we obtained

13C

NMR spectra with and without proton decoupling (Tables 3). Strong

J(13C-19F) coupling of 20-25 Hz was found between the carboxyl and methyl carbons and the CFMe fluorine (Table 3); in addition, the methyl carbon was coupled to a single fluorine (~8 Hz), which is clearly from the axial fluorine. The C*FMe carbon is also coupled to this axial fluorine (5.2 Hz); both of these spin-spin coupling in the bidentate ligand are consistent with our DFT calculations in Table 3. Importantly, these features are not be observed for the ring-shaped dimer anion shown in Figure S23b (see Table S6) since in this anion the CFMe group is removed further from the 31P19F4 subsystem than in the monomer anion shown in Figure 3 (Table 3). This completes our proof that the prevalent reaction product is indeed 2b and no other anion species. In this aged 10 wt% solution, 2b and PF6- are present in 1:2 mole ratio (Table S5). Shortly after the initiation of the reaction, resonance lines from another anion species were observed through its dd resonance in the 19Feq nuclei at -65.6 ppm (Table S5). As the solution ages, a dquin line from the 19Fax nuclei in this product is observed at -79.2 ppm, indicating a 31P19F5 species with four equivalent equatorial fluorines. At all stages of aging, the yield of this species was < 5 %. We identify this product with the monosubstituted anions shown in Figures S23a, as it is the first step towards the disubstituted anion shown in Figure 3, although the bis-substituted anion shown in Figure S23c also cannot be excluded (see our DFT calculations in Table S7). We consider the latter possibility unlikely. Indeed, in the 1H NMR (Table S7), a CFMe methyl is observed upfield of this proton in 2b (0.95 vs 1.09 ppm) with the similar J(1H-19F) coupling (21.6 Hz vs. 22.8 Hz) that is consistent with the monosubstituted anion, and there is a complementary resonance line from the TMS group upfield of 1b (-0.44 ppm vs. -0.55 ppm, see Table S7). We, therefore, suggest that the 31P19F

5 product is the postulated intermediate of monodentate substitution. Two other minor anions

were observed in dimethyl carbonate solutions containing stoichiometric excess of 1b over LiPF6; their attribution is discussed in section S2. To conclude, for multiple reasons the bidentate monosubstitution always prevails; the products of further substitution are either minor or not observed, and the monodentate substitution is almost entirely suppressed. This favorable chemistry underscores the success of the synthetic route shown in Figure 1.

4. Conclusion 14 ACS Paragon Plus Environment

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Here we report ex situ and in situ syntheses of two cathode protective electrolyte additives 2a and 2b through bidentate monoaddition of bis-trimethylsilyl diesters to the PF6- anion in carbonate solvents (Figure 1). These reactions proceed with great facility and almost 100% efficiency at room temperature, generating only traces of other products. Both of these additives are shown to be highly effective in slowing down the impedance rise and capacity fade in cells containing high-voltage NMC532 cathodes. Surface characterization of the aged cathode indicates no corrosive pitting of the oxide surface, reduced oxygen loss relative to the baseline electrolyte, and the formation of stable protective layers involving decomposed additives. The figure-of-merit comparison of the effectiveness of these additives in countering the loss of energy and power density during aging of the cells indicates that 2 wt% 2a and 1 wt% 2b compositions provide the best electrochemical performance. Our study exemplifies a new approach to searching for the cathode protective agents: rather than letting spontaneous chemical reactions in the electrolyte to transform the inactive “primary” additive into an active cathode-protective agent, we show how chemists can take matters in their own hands and use molecular design to obtain a specific product with the desired function, expanding the current art. Through this regained control over the chemical structure, a better stability of the cathode protection agent on the anode can be achieved, and possibly other useful features can be introduced.

ASSOCIATED CONTENT Supporting Information: A PDF file containing additional sections, figures, and tables. This material is available free of charge via the Internet at http://pubs.acs.org.

ACKNOWLEDGMENTS Support from the U.S. Department of Energy’s Vehicle Technologies Program (DOE-VTP), specifically from Peter Faguy and Dave Howell, is gratefully acknowledged. The electrodes and electrolytes used in this article are from Argonne’s Cell Analysis, Modeling and Prototyping (CAMP) Facility. Both facilities are supported within the core funding of the Applied Battery Research (ABR) for Transportation Program. TEM images were obtained at the Center for 15 ACS Paragon Plus Environment


Page 16 of 32

Nanoscale Materials at Argonne National Laboratory, which are supported by the U. S. Department of Energy, Office of Science, Office of Basic Energy Sciences, also under Contract No. DE-AC02-06CH11357. The submitted manuscript has been created by UChicago Argonne, LLC, Operator of Argonne National Laboratory (“Argonne”). Argonne, a U.S. Department of Energy Office of Science laboratory, is operated under Contract No. DE-AC02-06CH11357. The U.S. Government retains for itself, and others acting on its behalf, a paid-up nonexclusive, irrevocable worldwide license in said article to reproduce, prepare derivative works, distribute copies to the public, and perform publicly and display publicly, by or on behalf of the Government. REFERENCES 1.

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Page 20 of 32

Table 1. Summary of electrochemical performance metrics for the NMC532//Gr cells containing different electrolytes. Additive

Baseline

2a

2b

wt% Charge First-cycle specific Discharge capacity, a % Loss mAh/goxide Discharge Initial b specific Final b b capacity, % loss per mAh/goxide cycle b c ASI, Initial b 2 Ω/cm Final b % increase Power Initial b density, d Final b 2 mW/cm % Loss Figure of Energy merit, density full cycles Power e density

0.0 220±3 191.0±0.4 13.1

0.5 192.4±26.6 158.8±29.8 17.5

1.0 214±10 180.7±11.4 15.6

2.0 219.3±1.4 181.2±1.9 17.3

0.5 217.3±5.6 185.8±3.2 14.5

1.0 214.1±11.7 182.1±3.2 14.9

2.0 215.4±3.4 174.7±3.2 18.9

190.3±0.2 157.8±0.1 0.28

183.1±5.1 161.3±0.8 0.19

184.9±1.5 163.5±0.7 0.19

178.4±2.1 157.2±2.1 0.18

187.7±0.7 160.5±1.0 0.24

188.0±1.2 158.6±0.6 0.26

177.6±2.2 157.1±3.2 0.18

24.0±1.6 44.9±2.1 87.1 1358±0.5 71.7±2.1 47.1 170

31.2±1.4 45.8±2.2 46.8 106±5 73.3±3.7 30.9 201

26.2±1.9 36.2±1.7 26.7 128.4±9.1 93.2±4.1 35.2 217

25.1±0.5 31.3±1.9 24.7 133±2.5 105.9±4.9 20.4 188

27.4±4.2 44.1±1.1 58.8 124±18 74±1.1 40.3 180

23.8±1.5 36.8±1.3 54.6 141±8.7 89.2±3.1 36.7 167

30.6±3.2 35.1±2.2 14.7 109.2±11.7 91.8±5.9 15.9 172

23

8

48

87

21

54

10

a) Formation cycles at a C/10 rate; b) The 'initial' and 'final' correspond to cycles 3 and 118, respectively; c) full (cathode and anode) ASI as determined using 10 s, 2C discharge pulses applied at 4.08 V; d) defined as described in section 2; e) number of full cycles corresponding to 80% loss of a metric relative to the first cycle for the baseline cell under the identical cycling conditions. Average values and standard deviations for three independent measurements are given.


Page 21 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Table 2. Computed energetics for X/2F exchange in the gas phase PXnF6-2n- complexes. a PXnF6-2nanion n=1 n=2 n=3

Electron energy gain in X/2F exchange, eV X=ox X=Z 0.60 +0.126 0.39 -0.096 0.28 -0.164

a) PXn-1F8-2n- + X(TMS)2 → PXnF6-2n- + 2 TMSF

Table 3. Comparison of calculated and experimental NMR parameters for 2b anion. a

B3LYP/6-31+G(d,p) calculation nucleus δ, JP*, JF*, JH*, ppm Hz Hz Hz 31P -128.07 19F -45.06 855.3 80.1 (eq), 21.6 (ax') 2.1 ax 19F -59.75 876.8 82.5 (eq), 21.6 (ax) 0.6 ax' 19F -66.99 860.2 82.5 (ax), 80.1 (ax') 0.4 eq 19F -160.13 10.4 5.9 (ax'), 0.7 (eq) 15.8 C 1H 2.21 15.8 Me 13C 88.99 220, 3.8 (ax), 2.7 0.7 F (eq), 1.8 (ax') 0.2 13C 22.90 0.9 19.5, 7.8 (ax), 0.1 132.0 Me (ax'), 0.2 (eq) 13C 159.57 5.6 16.8, 3.9 (ax'), 1.4 3.1 O (ax), 1.1 (eq)

 ppm -151.85 -52.14 -58.19 -70.78 -156.07 87.28

experimental estimates JP*, JF*, Hz Hz 786.8 50.1 (x2) 768.4 51.1 (x2) 736.7 51.1 (x2) 5.7 22.8 189.9, 5.2

22.13

-

25.2, 8.0

132.5

165.78

-

21.22

-

a) The atoms are labeled following the conventions of Figure 3.


JH*, Hz 2.2 22.8 -


F F F P F F F

Li

Si O

O

Si O

O

1a

Li

Page 22 of 32

F O P F FO

F

O +

Si F

O

2a O

Si O

Si

F

O O 1b

Li

F O P F FO

F

O

F

+

Si F

O

2b

Figure 1. Reactions between LiPF6 and bistrimethylsilyl carboxylates in a carbonate solvent.


eq, x2

(a)

ax'

CF*Me

ax

-51 -52 -53

-58 -59 -70 -71 -156.0 19 ( F), ppm

-156.2

PF6

-

(b)

NMR signal

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


NMR signal, arb. u.

Page 23 of 32

-145

-150 31 ( P), ppm

-155

Figure 2. Partial (a) 19F NMR and (b) 31P NMR spectra of a 1:1 equiv. mixture of 1b and LiPF6 in dimethyl carbonate after three days of aging. In panel b, the resonance lines of the PF6- anion are highlighted in blue, and the resonance lines of 2b are indicated with open ovals.



Page 24 of 32

Figure 3. A structural model for 2b anion, with atoms labelled as discussed in the text.


Page 25 of 32

2 wt% 2b

2 wt% 2a Baseline

180

Specific Capacity/ mAh g-1

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


160 200

1 wt% 2a

1wt% 2b

0.5 wt% 2a

0.5 wt% 2b

180 160

180

160 0

50

100

Cycle Number

0

50

100

Cycle Number

Figure 4. Specific capacity per oxide weight for the NMC532//Gr cells containing 2a and 2b (red symbols) and the baseline electrolyte (gray symbols) obtained for cycling at a rate of C/3. The weight fractions of additives are given in the plot. The discontinuities in the data correspond to HPPC tests performed every 20th cycle. The scattered points represent 1C (lower points) and C/10 measurements (upper points) taken after each HPPC sequence.



Page 26 of 32

Figure 5. Differential electric charge (dQ/dV) profiles plotted vs. the cell voltage (red lines) for the first charge half cycle of the NMC532//Gr cells containing 2a (on the left) and 2b (on the right). The dQ/dV trace for the baseline electrolyte (gray line) is given in each panel for comparison. The weight fraction of electrolyte additives is also given in each panel.


Page 27 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Figure 6. Full-cell ASI measurements (filled circles) for the NMC532//Gr cells containing different additives in Gen2 electrolyte. The rainbow colors correspond to 1, 20, 40, 60, 80, 100, and 120 cycles (from bottom to the top), respectively. The baseline electrolyte is shown in grey in each panel for comparison. The ASI was estimated using 10 s, 2 C discharge pulses applied at different cell voltages during the incremental discharge of a fully charged Li-ion cell at a C/3 rate.


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Current/ A


Page 28 of 32

1 wt% 2a 1 wt% 2b Baseline

2x10-6 1x10-6 0

0

20

40

60

Time/ h Figure 7. Decay kinetics for leakage current flowing through the NMC532//Gr cells during a 60 h long potentiostatic hold at 4.5 V. The baseline electrolyte is indicated with the blue line and electrolytes containing 1 wt% 2a and 2b are indicated with the black and red lines, respectively.


Page 29 of 32

2wt% 2a

80

FOMP

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


1wt% 2b

40

1wt% 2a

Baseline 0.5wt% 2b

0

2wt% 2b

160

0.5wt% 2a

200 FOME

240

Figure 8. A two-dimensional representation for the figures of merit in power (FOMP) and in energy (FOME) for different electrolytes. Electrolytes containing 0.5 wt% (blue), 1 wt% (orange), and 2 wt% (green) 2a (circles) or 2b (squares) are indicated by colored symbols. The baseline electrolyte is represented by the red triangle in the cross hairs.



Counts

30 20

Pristine NCM532 Baseline 0.5 wt% 2a 1 wt% 2a 2 wt% 2a

8 O1s

C1s

4

10 0 296

0 288

280

536

528 P2p

F1s

Counts

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 30 of 32

8

0.4

0 696

0.0 688

680 144

Binding Energy/ eV

136

128

Binding Energy/ eV

Figure 9. Baseline corrected X-ray photoelectron spectra of the NMC532 cathode before (black lines) and after 119 cycles at a C/3 rate. The color coding is detailed in the inset.


Page 31 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Figure 10. Structural models for products of bidentate substitution by the oxalate (ox2-) in the PF6- anion.



Page 32 of 32

For Table of Content Only


Facile in Situ Syntheses of Cathode Protective Electrolyte Additives for

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