Understanding the Formation and Evolution of Lithium-Ion Battery Solid Electrolyte Interphase

Lithium ion batteries are widely used as energy storage devices in a variety of products such as smartphones, tablets, laptops and other portable electronics. Thanks to their high energy density and cyclability, they are currently being used by and developed for electric vehicles. There is a growing need for cost reduction; increase in energy density; wider operating temperature range; and improved safety characteristics of the batteries. Organic carbonates are the primary solvents used in lithium-ion battery electrolytes along with electrolyte additives. The reversibility of current lithium-ion batteries is dependent upon the electrolyte used in the batteries. During the initial charging cycles of the cell, a solid electrolyte interface (SEI) is formed by reduction of organic carbonates, electrolyte salts and/or electrolyte additives on the surface of the graphitic anode in lithium-ion batteries. The generation of a stable anode SEI prevents continuous electrolyte reduction on the surface of the anode. The SEI functions as a Li ion conductor but an electrical insulator. The reduction reactions of the electrolytes on the graphitic anode surface have been investigated for many years and it been proposed to contain a complicated mixture of products including lithium oxalate, lithium alkoxides, and lithium oxide from the carbonate solvents and LiF and lithium fluorophosphates from the reduction of LiPF6. Similar ambiguity exists about the components of SEI formed from electrolyte additives and other electrolyte salts. Despite the extensive investigations, the structure, formation mechanisms and evolution of the SEI are poorly understood. Understanding the mechanisms of the reduction reactions of organic carbonates, electrolyte salts and electrolyte additives along with the products of the reactions which result in the generation of the SEI is essential for the development of safer lithium-ion batteries with wider operating temperature range. Lithium naphthalenide has been investigated as a one electron reducing agent for organic carbonates solvents, some of the most robust additives and salts used in lithium ion battery electrolytes. The reaction precipitates have been analyzed by IRATR, XPS and solution NMR spectroscopy. The evolved gases and the volatile components have been analyzed by GC-MS. The reduction products of ethylene carbonate and propylene carbonate are lithium ethylene dicarbonate (LEDC) and ethylene and lithium propylene dicarbonate (LPDC) and propylene, respectively. The reduction products of diethyl and dimethyl carbonate are lithium ethyl carbonate (LEC) and ethane and lithium methyl carbonate(LMC) and methane, respectively. Electrolyte additives, FEC and VC reductively decompose to HCO2Li, Li2C2O4, Li2CO3, and polymerized VC. All the fluorine containing salts generate LiF upon reduction. In addition to LiF, LiBF4 generates LixByFz species; LiBOB and LiDFOB generate lithium oxalate and boron-oxalatoesters; LiPF6 yields LiPF2 species and LiTFSI produces lithium bis[N-(trifluoromethylsulfonylimino)] trifluoromethanesulfonate. The poor thermal stability of the SEI layer has been attributed to exothermic reactions between lithium alkyl carbonates and LiPF6. While the relationship between capacity fade and SEI instability is clear, and there have been some investigations of SEI component evolution, the mechanism of SEI component decomposition is complicated by the presence of many different components. The thermal stability of Li2CO3, LMC, and LEDC in the presence of LiPF6 in dimethyl carbonate (DMC), a common salt and solvent, respectively, in lithium ion battery electrolytes, has been investigated to afford a better understanding of the evolution of the SEI. The residual solids from the reaction mixtures have been characterized by a combination of X-ray photoelectron spectroscopy (XPS) and infrared spectroscopy with attenuated total reflectance (IR-ATR), while the solution and evolved gases have been investigated by nuclear magnetic resonance (NMR) spectroscopy and gas chromatography with mass selective detection (GC-MS). The thermal decomposition of Li2CO3 and LiPF6 in DMC yields CO2, LiF, and F2PO2Li. The thermal decomposition of LMC and LEDC with LiPF6 in DMC results in the generation of a complicated mixture including CO2, LiF, ethers, phosphates, and fluorophosphates.

STATEMENT OF THE PROBLEM AND PROPOSED SOLUTIONS
Organic carbonates are the primary solvents used in lithium-ion battery electrolytes along with electrolyte additives. The reversibility of current lithium-ion batteries is dependent upon the electrolyte used in the batteries. 4 During the initial charging cycles of the cell a solid electrolyte interface (SEI) is formed by reduction of organic carbonates, electrolyte salts and/or electrolyte additives on the surface of the graphitic anode in lithium-ion batteries. The generation of a stable anode SEI prevents continuous electrolyte reduction on the surface of the anode. The SEI functions as a Li ion conductor but an electrical insulator. 3 The reduction reactions of carbonates on the graphitic anode surface have been investigated for many years. Initially a single two electron reduction mechanism of propylene carbonate to generate Li2CO3 and propylene was proposed, 5 later Aurbach and co-workers proposed two sequential one electron reduction reactions of cyclic carbonates to generate lithium alkyl carbonates and alkenes. 6 Numerous other researchers have investigated the composition of the SEI on graphitic anodes in lithium ion batteries. [7][8][9][10][11][12][13][14][15][16][17] In addition to lithium alkyl carbonates and lithium carbonate, the SEI has been proposed to contain a complicated mixture of products including lithium oxalate, lithium alkoxides, and lithium oxide from the carbonate solvents and LiF and lithium fluorophosphates from the reduction of LiPF6. [7][8][9][10][11][12][13][14][15][16][17] Similar ambiguity exists about the components of SEI formed from electrolyte additives and other electrolyte salts. Capacity fade at elevated temperature is connected to the exothermic reactions between lithium alkyl carbonates and LiPF6. 16,18 While the relationship between capacity fade and SEI instability is clear, 16 The results provide significant insights into the formation and decomposition mechanism of the anode SEI.

INTRODUCTION
Organic carbonates are the primary solvents used in lithium-ion battery electrolytes. The reversibility of current lithium-ion batteries is dependent upon the electrolyte used in the batteries. 1 During the initial charging cycles of the cell a solid electrolyte interface (SEI) is formed by reduction of organic carbonates on the surface of the graphitic anode in lithium-ion batteries. The generation of a stable anode SEI prevents continuous electrolyte reduction on the surface of the anode. The SEI functions as a Li ion conductor but an electrical insulator. 2 Understanding the mechanisms of the reduction reactions of organic carbonates along with the products of the reactions which result in the generation of the SEI is essential for the development of better lithium-ion batteries.
The reduction reactions of carbonates on the graphitic anode surface have been investigated for many years. Initially a single two electron reduction mechanism of propylene carbonate to generate Li2CO3 and propylene was proposed, 3 later Aurbach and co-workers proposed two sequential one electron reduction reactions of cyclic carbonates to generate lithium alkyl carbonates and alkenes. 4 Numerous other researchers have investigated the composition of the SEI on graphitic anodes in lithium ion batteries. [5][6][7][8][9][10][11][12][13][14][15] In addition to lithium alkyl carbonates and lithium carbonate, the SEI has been proposed to contain a complicated mixture of products including lithium oxalate, lithium alkoxides, and lithium oxide from the carbonate solvents and LiF and lithium fluorophosphates from the reduction of LiPF6. [5][6][7][8][9][10][11][12][13][14][15] Despite the extensive investigations, the structure and formation mechanisms of the SEI are poorly understood.
A detailed analysis of binder free graphitic anodes cycled in simplified electrolytes composed of a single carbonate solvent and LiPF6 has been reported. 16,17 These investigations suggest that the initial reduction reaction of the carbonates generate lithium alkyl carbonates and LiF as the predominant components of the anode SEI. As an expansion of these investigations, lithium naphthalenide, a well-known one electron

NMR Spectra Of Precipitates
The molecular structures of the precipitates formed in the reaction between lithium naphthalenide and various carbonates are analyzed via a combination of 1 H, 13 C NMR and IR-ATR spectroscopy. As previously reported, the 1 H NMR spectrum of the reduction product of EC contains a singlet at 3.6 ppm characteristic of (-OCH2CH2O-) and peaks at 62.5 and 161.1 ppm in the 13 C NMR spectrum, characteristic of (-CH2O-) and a C=O, respectively (Figure 2.1). 16 The resonances match those previously reported for lithium ethylene dicarbonate (LEDC). 7 The reaction produces LEDC in high yield ~ 95 %, and no other products are observed. The 1 H and 13 C NMR spectra of the precipitate formed by the reaction of lithium naphthalenide with PC is provided in  In addition to analysis of the reduction products of cyclic carbonates, the reduction of dialkyl carbonates has been investigated. The 1 H NMR spectra of the reaction product of the Lithium naphthalenide reduction of DMC is provided in Figure   2  to the peaks characteristic of lithium alkyl carbonates, a weak absorption is observed at ~1450 cm -1 characteristic of Li2CO3. While there appears to be some Li2CO3 present in all spectra, lithium alkyl carbonates decompose to form Li2CO3, as discussed above.
Thus, the thermal, Lewis acid, or Lewis Base catalyzed decomposition of lithium alkyl carbonates is the most likely source of Li2CO3 since Li2CO3 is not observed in the 13 C NMR spectra of fresh reduction products. In addition, all of the samples have increases in the intensity of the absorption of Li2CO3 upon exposure to air.

GC-MS ANALYSIS OF GASES
Upon reaction of lithium naphthalenide with carbonates. significant gas evolution is observed. The gases evolved during reaction were analyzed by GC-MS.
Analysis of the evolved gas during the reduction of EC confirms that the gas produced is ethylene. Reaction of PC results in similar gas evolution and the gaseous product is propylene. The two cyclic carbonates have very similar reaction mechanisms which are consistent with those originally proposed by Aurbach and co-workers. 4 The reduction follows two sequential single electron transfer reactions (Scheme 2.1). The first electron generates a radical anion, the second electron generates lithium alkylene dicarbonates and an alkene. Aurbach and co-workers. 4 There is no evidence for the generation of either Li2CO3 or CO2 during the initial reduction reactions. It is well known that the anode SEI changes upon additional cycling and thermal abuse. 14

02881, United States
The following is published in the journal, Chemistry of Materials, and is presented here in manuscript format.

ABSTRACT
We have synthesized the products of fluoroethylene carbonate (FEC) and vinylene carbonate (VC) via lithium naphthalenide reduction. By analyzing the resulting solid precipitates and gas evolution, our results confirm that both FEC and VC decomposition products include HCO2Li, Li2C2O4, Li2CO3, and polymerized VC. For FEC, our experimental data supports a reduction mechanism where FEC reduces to form VC and LiF, followed by subsequent VC reduction. In the FEC reduction product, HCO2Li, Li2C2O4, and Li2CO3 were found in smaller quantities than in the VC reduction product, with no additional fluorine environments being detected by solid-state nuclear magnetic resonance or X-ray photoelectron spectroscopy analysis. With these additives being practically used in higher (FEC) and lower (VC) concentrations in the base electrolytes of lithium-ion batteries, our results suggest that the different relative ratios of the inorganic and organic reduction products formed by their decomposition may be relevant to the chemical composition and morphology of the solid electrolyte interphase formed in their presence.

INTRODUCTION
Additives are widely used to improve performance of Li-ion batteries, offering an economically viable method of performance enhancement compatible with existing manufacturing infrastructure. 1 Generally, the function of additives is sacrificial: they are reduced at different voltage potentials compared to the base electrolytes to which they are added, forming decomposition products that are incorporated into a protective layer on electrodes. 1−5 This protective layer is called the solid electrolyte interphase (SEI). 6,7 The formation of a stable SEI is essential for all Li-ion batteries, preventing further electrolyte decomposition, thereby underlying capacity retention. 1,5−8 The SEI also represents an electronically insulating barrier between the electrodes and electrolyte, with its composition, thickness, and structure influencing the lithium transport across the interphase. 9 Practically, the SEI layer is extremely air-sensitive. 30,31 Moreover, with a thickness of less than 100 nm, it is very difficult to study experimentally. Here, our strategy is to synthesize the reduction products of FEC and VC in order to experimentally confirm their reduction products. Lithium naphthalenide (Li-Nap), a well-known reducing agent, is known to react with solvents in a similar manner to those which may occur on lithiated anodes. 32 It is used here to reduce FEC and VC, modeling a reduction process in a similar manner to that which may occur in a lithium-ion battery.
Solid products are analyzed with X-ray photoelectron spectroscopy (XPS), solid-state NMR (ssNMR), and Fourier transform infrared spectroscopy (FTIR). Gas evolution is monitored using gas chromatography mass spectrometry (GC-MS). The many techniques provide chemical signatures for future work. Viable reactions to form the detected decomposition products are proposed. For FEC, we propose a reduction scheme where FEC reduces to form LiF and VC, followed by further reduction of VC to polymerized VC (poly(VC)). The poly(VC) contains repeating EC units joined by cross-linking sites; our analysis shows no evidence for F−C bonds in the polymer.
HCO2Li, Li2C2O4, and Li2CO3 are also found in small quantities. For VC, we detect lithium environments of HCO2Li, Li2C2O4, and Li2CO3, in addition to poly(VC).

Synthesis.
All reagents were used as obtained, without further purification. Batterygrade VC and FEC were obtained from BASF. Naphthalene (99 + %, Scintillation grade) and THF (Anhydrous, 99.9%) were purchased from Acros organics. Lithium discs were XPS. XPS spectra of the dried precipitates were acquired using a thermo scientific Kalpha XPS instrument. Samples were made into circular pellets with a press and transferred from the glovebox to the XPS chamber using a vacuum transfer module without exposure to air. C 1s, O 1s, and Li 1s spectra were obtained from the VC precipitate, whereas C 1s, O 1s, Li 1s, and F 1s spectra were acquired from the FEC precipitate. An Argon flood gun was used to avoid surface charge accumulation during sample analysis. The binding energy was corrected on the basis of the C 1s of hydrocarbon at 284.8 eV. The data was processed and analyzed using the Thermo Avantage, XPS Peak 4.1 and the Origin software.
ssNMR. Multinuclear ssNMR spectra were obtained on 16.4 T Bruker Avance III 700 MHz and 11.7 T Bruker Avance III 500 MHz spectrometers. Samples were packed in an Ar glovebox (typically O2 and H2O < 0.1 ppm), avoiding any exposure to ambient air, into rotors of 1.3, 3.2, and 4 mm outer diameters. Magic-angle spinning (MAS) frequencies ranged from 10 to 60 kHz, spinning under N2. 1 H and 13 C chemical shifts were externally referenced to adamantane ( 1 H 1.9 ppm, 13 C 38.5 ppm, CH2) and 7 Li and 19 F to LiF ( 7 Li −1 ppm, 19 F−204 ppm). 33,34 The data were processed using the Bruker TOPSPIN software and analyzed using the dmfit software. 35  Each of the FEC and VC precipitates were investigated using 1 H, 7 Li, and 19 F ssNMR, using 1.3 mm rotors and 60 kHz spinning frequency. 13 C ssNMR experiments were performed using larger 3.2 and 4 mm rotors and spinning frequencies ranging from 10 to 12 kHz. The larger samples provided greater sensitivity. The 13 C spectra were acquired using sweptfrequency two-pulse phase modulation (swfTPPM) 36 1 H decoupling at 80−100 kHz. Direct excitation 13 C experiments provided quantitative information. 1 H− 13 C, 7 Li− 13 C, and 19 F− 13 C correlation experiments were used to probe spatial proximity of these nuclei by transferring magnetization from 1 H, 7 Li, and 19 F nuclei by cross-polarization to C nuclei. Dipolar dephasing (interrupted decoupling) 1 H− 13 C crosspolarization experiments allowed differentiation between protonated and nonprotonated environments. 34,37 Further ssNMR experimental details are given in the Supporting Information.
FTIR. FTIR analysis was performed on each of the precipitates prepared with nondeuterated naphthalene. FTIR spectra of the dried precipitates were acquired on a Bruker Tensor 27 spectrometer, equipped with germanium crystal, in attenuated total reflectance (IR-ATR) mode. Samples were transferred using airtight vials, and the spectrometer was operated inside a nitrogen filled glovebox to avoid sample exposure to ambient air. Each spectrum was acquired with 128 scans from 700 to 4000 cm −1 at the spectral resolution of 4 cm −1 . The data was processed and analyzed using the OPUS and Origin software. Helium was used as carrier gas at a flow rate of 24 mL/min. The initial column temperature was 40 °C, and the temperature was ramped at 10 °C/min to 200 °C and held at that temperature for 2 min with the total run time of 18 min. The mass spectra obtained were compared to the NIST library to determine their molecular structures.

GC-MS
Computational Methods. Chemical shifts were calculated using density functional theory (DFT) using Gaussian 09 38 and estimated using ChemNMR implemented in ChemBioDraw 13.0; see Table S1. ChemNMR approximates 13  Li2CO3. 21,44 The peak at 288.5 eV is assigned to CO2 environments contained in HCO2Li and/or Li2C2O4, while the peaks at lower binding energies of 286.8 and 284.8 eV indicate C−O and C−C bonds, respectively. Li2CO3 has a larger contribution in the VC precipitate compared with the FEC precipitate, also seen by 7 Li and 13 C ssNMR later.
In addition, the relative signal intensity of the HCO2Li/Li2C2O4 peak is larger in the VC precipitate, these carboxylate environments being confirmed by 13 50 We note that residual THF may contribute to the signals of 1.5 and 3.6 ppm. In the 7   to this C environment. A 7 Li− 13 C experiment was also attempted, but no signal was detected: the null result is in agreement with the 7 Li ssNMR assignment (Figure 3. 2) showing that very little Li2CO3 (or similar environment resonating near 0 ppm in the 7 Li spectrum) is present in the FEC precipitate. 1 H− 13 C cross-polarization experiments, Figure 3.3a-iii, were performed, further confirming the 13 C assignments. The signal intensity in these experiments depend on the dynamics of the functional groups and the molecules in the SEI, and spatial proximity of 1 H and 13 C nuclei. Following a similar strategy used in our previous paper to assign different carbon local environments, 52 a delay time is introduced following the cross-polarization step in the experiment to perform a dipolar dephasing (interrupted decoupling) experiment, the experimental details being described in the Supporting Information. When the delay times are varied in the experiment, different C functional groups can be identified on the basis of their attenuation rates. Protonated C is attenuated more rapidly than non-protonated C. Also, the signal intensities for rigid CH/CH2 environments attenuate more rapidly than signal from mobile species such as rotating CH3 methyl groups (due to a reduced dipolar coupling The minor peak I (13 ppm) is characteristic of CH3R environments; the resonance likely has some contribution from residual diethyl ether (CH3CH2−O−CH2CH3) used to rinse the precipitates during synthesis. However, reactions forming these environments in minor quantities may also contribute to the signal.
The 7 Li− 13 C cross-polarization experiment ( Figure 3.3b-ii) indicates the Li + coordination environments by the carboxyl and carbonate groups. The broad resonance of A is consistent with the HCO2Li assignment. The majority of the signal contributing to the asymmetric peak at B is assigned to Li2CO3, the small shoulder being assigned to Li2C2O4. The 7 Li ssNMR spectra are consistent with these assignments (Figure 3.2).
In the dipolar dephasing experiment (Figure 3.3b- (Figure 3.3b-iii), which is not consistent with its assignment solely to HCO2Li. This spectrum was collected at a higher magnetic field strength than the spectra shown in Figure 3.3b-i,-ii, and there is now a severe overlap with the now much more intense D-Napth. spinning sideband (labeled with an asterisk), this signal not being attenuated in the dephasing experiment. Similarly, no attenuation is expected for an acetate resonance.
FTIR. FTIR spectra of the precipitates obtained on reduction of FEC and VC are displayed in Figure 3.4, confirming chemically bonded groups assigned in our XPS and ssNMR spectra. Our assignment here is based on comparison of the spectra to related studies. 18,32,44,47 The FEC and VC reduction products have similar FTIR signatures, with some relative intensity differences at approximately 1300, 1400−1500, and 1750 cm −1 .
In addition to the VC/FEC reduction products, some residual naphthalene is seen (788, 3064 cm −1 ). In each of the samples, the previously assigned Li2CO3 is again observed (878, 1449, and 1488 cm −1 ). As also seen by 13 C ssNMR (Figure 3.3b-ii) and the C 1s XPS spectra (Figure 3.1), the Li2CO3 is more prevalent in the VC sample.
In the FEC reduction product, peaks for carbonate C=O (1795 cm −1 ) and C−O (1080, 1171 cm −1 ) bonds are seen. These peaks are assigned to bonds contained in ROCO2R environments, resembling those assigned to a poly(VC) product in our previous study. 18  To gain further insight into the decomposition mechanisms, an additional experiment was performed with a half molar equivalent of Li-Nap, providing FTIR spectra comparable to that of Figure 3.4 (see Supporting Information). The spectra show much weaker intensities for the peaks assigned to Li2CO3, Li2C2O4, and HCO2Li, relative to the peaks assigned to ROCO2R environments assigned to poly(VC), revealing a Li concentration dependence in the formation of these inorganics.
GC-MS Analysis of Gases. GC-MS analysis was performed, providing more insight for the reduction mechanisms of FEC and VC, resulting in the solid precipitates. For FEC, the reduction with Li-Nap yields a mixture of CO and CO2. The ratio of CO to CO2 peak areas is 1:4.4. For VC, the reduction with Li-Nap yielded carbon monoxide as the only gaseous product (i.e., no CO2 was detected), CO2 having been detected previously. 45,47 The absence of CO2 detection was attributed to its consumption in further reactions, the experiment being performed in a closed system, with an abundance of Li.
Proposed Reduction Products. On the basis of the above analysis, we propose that the reduction product poly(VC) is present in the precipitates of both FEC and VC, as well as Li2C2O4, Li2CO3, and HCO2Li, Figure 3.5. The relative ratio of these products differs for VC and FEC.
The ROCO2R environment observed in the XPS, ssNMR ( group and contribute to the 13 C ssNMR signal of resonance H (Table 3.2). 13 C ssNMR resonance D, assigned to protonated C environments adjacent to two OR groups, indicates the possibility of a cross-linking site for poly(VC) (seen in Figure 3.5). The cross-linking site may also contain C environments adjacent to one O, contributing to resonance E. The signals from the distribution of carbons not adjacent to O (Table 3. Definitive assignments for C-groups and fluorinated-species contained in the reduction products were aided by the large chemical shift dispersion of the 13 Figure 3.5). The carboxylate and carbonate assignments in the 13  With the reduction products detected being nearly identical to FEC, the difference being the LiF product, similar reduction mechanisms are expected. One possible mechanism for the reaction of FEC is nearly identical to the reaction of VC, except that the first step of the reaction involves the reduction of FEC to generate VC, LiF, and 1/2 H2, Scheme 3.1-ii. We note that LiF was generated nearly quantitatively via the Li-Nap reduction of FEC in our experiments. While we were unable to observe H2 generation, as the mass of H2 is below the detection limit of our GC-MS, the detection of H2 during the reduction of FEC was previously reported as part of the 4 electron reduction mechanism of FEC by Jung et al. 54 The rapid polymerization of the VC generated from We note that Shkrob et al. 53  and lithium divinylene dicarbonate (LDVD), these SEI decomposition products being suggested by prior theoretical investigations (see Supporting Information for our estimated NMR shifts of these predicted products). 20 They observed similar 13 C NMR peaks at ∼154 and ∼70−80 ppm in 13 C liquid NMR spectra of the SEI formed on the graphite electrode, dissolved in DMSO-d6, which they assigned to an oligomer of VC and/or poly(VC), the poly(VC) assignment in agreement with the ssNMR results here (see Table 3.2, peaks C and E); they also observed the distinctive 13 C ssNMR resonance at ∼100 ppm, seen in this study (see Table 3.2, peak D) which they assigned to an oligomer of VC. Here, we have assigned the 100 ppm resonance to a cross-linking site of poly(VC) (see Scheme 3.1-i, Figure 3.5). We have also observed broad peaks at 36 and 40 ppm, indicating a distribution of RCH2R′ environments (see Table 3.2, peaks F and G), assigned to the cross-linking site. Finally, in our previous study of the SEI composition on Si anodes formed in the presence of FEC and VC additives, 18 we have observed an FTIR adsorption peak at ∼1800 cm −1 increasing with additive concentration, the adsorption peak being assigned to poly(VC), as in this study.
The absence of the production of CO2 during the reduction of VC, seen by GC-MS, is in contrast to the literature. For example, the study by Ota et al., 47 which used pure VC as an electrolyte solvent, observed CO2 as the major gaseous product and a small amount of CO. Similarly, CO2 has been reported as the major gaseous product, when VC is used as an electrolyte additive. 45,49 The discrepancy is likely due to the reduction of CO2 by excess Li napthalenide to generate CO, Li2CO3, and Li2C2O4 (see Scheme 3.1iii). 55           inorganic components, and the organic phases were utilized for the analyses. Helium was used as carrier gas at a flow rate of 24 mL/min. The initial column temperature was 40°C and the temperature was ramped at 10°C/min to 200°C and held at that temperature for 2 minutes with the total run time of 18 minutes. The mass spectra obtained were compared to the NIST library to determine their molecular structures.
THF, Et2O (solvents) and naphthalene (starting material) and were the only volatile components present in the reaction mixtures. The gas analyses were performed by sampling the head spaces of the reaction mixtures in RB flasks with a 10 µL GC syringe.
Helium was used as the carrier gas at a flow rate of 1.5 mL/min. The initial column temperature was set to 40 °C, and the temperature was ramped at 1°C/min to 43°C and held at that temperature for 2 min with the total run time of 5 min. The mass spectra obtained were compared to the NIST library to determine their molecular structures.
IR-ATR spectra of the dried solid residues were acquired on a Bruker Tensor 27 spectrometer equipped with a germanium crystal in attenuated total reflectance (IR-ATR) mode. Samples were transferred using air-tight vials and the spectrometer was operated inside a nitrogen filled glovebox to avoid air exposure. Each spectrum was acquired with 128 scans from 700 cm -1 to 4000 cm -1 at the spectral resolution of 4 cm -1 . The data were processed and analyzed using the OPUS and Originlab software.
NMR spectra of the samples were collected with a Bruker Avance III 300 MHz NMR spectrometer at room temperature. The solids were dissolved in D2O in the nitrogen filled glovebox and 19 F, 31 P, 11 B, & 13 C NMR spectra of the solutions were acquired. The spectra were processed and analyzed using MestReNova 10.0.2.
XPS spectra of the dried precipitates were acquired using a Thermo Scientific K-alpha XPS. Samples were made into circular pellets with a press or stuck on a conductive carbon tape as a thin layer and transferred from the glovebox to the XPS chamber using a vacuum transfer module without exposure to air. An argon flood gun was used to avoid surface charge accumulation during sample analysis. The binding energy was corrected based on the C 1s of hydrocarbon at 284.8 eV. The data were processed and analyzed using the Thermo Avantage, XPS Peak 4.1 and the Originlab software.

Reduction of electrolyte salts
The

NMR analysis of the solids
The residual organic solvent insoluble solids have been analyzed via a combination of solution NMR spectroscopy in D2O, Infrared spectroscopy with attenuated total reflectance (IR-ATR), and X-ray photo electron spectroscopy. The residual solids have been dissolved in D2O for NMR analysis. While most of the residual solids dissolve in D2O, some of the solid does not readily dissolve. Some of the reduction products may react with water to generate subsequent hydrolysis products.
The dissolved solids were analyzed via a combination of 11 B, 13 C, 19 F, and 31 P NMR spectroscopy. Representative NMR spectra of the solids are provided in

FTIR analysis of the solids
In an effort to further understand the composition of the solids obtained from reduction, the reduction products of salts have been analyzed with IR-ATR. The IR-ATR spectra of the solids generated from the reduction of LiBOB and LiDFOB are provided in Figure 4.3. IR-ATR spectra of the residual solids for the other salts were also acquired, but the spectra were dominated by residual solvent and naphthalene since the decomposition products do not contain any functional groups which strongly absorb IR radiation, consistent with the observation of LiF, in NMR analysis.

INTRODUCTION
Lithium-ion batteries (LIB) are widely used as energy storage devices in portable electronics 1 and increasingly in electric vehicles due to their high energy density. However, LIB exhibit poor capacity retention at moderately elevated temperatures, 2 which is undesirable for many of the intended applications. Impedance growth and loss of cyclable lithium are reported to be the main contributors to capacity fade. [3][4][5] Lithium-ion batteries typically contain a graphite negative electrode, a lithiated transition metal oxide positive electrode, and an electrolyte composed of LiPF6 dissolved in a mixture of organic carbonate solvents. 6 The SEI (solid electrolyte interphase) is formed on the surface of the anode from the electrochemical reduction of the electrolyte and plays a crucial role in the long-term cyclability of LIB. 7 While the SEI has been reported to be a complex mixture of compounds, the initially formed components of the SEI are dominated by LiF, Li2CO3, lithium ethylene dicarbonate ((CH2OCO2Li)2, LEDC) and lithium alkyl carbonates (ROCO2Li). [7][8][9][10][11] The poor thermal stability of the SEI layer has been attributed to exothermal reactions between lithium alkyl carbonates and LiPF6. 12,13 While the relationship between capacity fade and SEI instability is clear, [12][13][14] and there have been some investigations of SEI component evolution 15,16 the mechanism of SEI component decomposition is complicated by the presence of many different components. The limited understanding of the evolution of the SEI components over time has significantly limited efforts to understand the mechanism of ion transport through the SEI via computational modeling. [17][18][19][20][21][22][23][24][25] A more comprehensive understanding of the decomposition reactions will aid computational scientists to develop a better physical understanding of the evolution of ion conducting mechanisms in the SEI and will help to improve the calendar life of lithium-ion batteries.
The SEI components lithium ethylene dicarbonate (LEDC), and lithium methyl carbonate (LMC) were independently synthesized by reduction of EC and DMC with lithium naphthalenide. 10 The decomposition reactions of Li2CO3, LEDC and LMC in the presence of LiPF6 have been investigated. The decomposition products were analyzed via a combination of nuclear resonance spectroscopy (NMR), infrared spectroscopy with attenuated total reflectance (IR-ATR), X-ray photoelectron spectroscopy (XPS) and gas chromatography with mass selective detection (GC-MS).

Experimental
Battery-grade DMC and LiPF6 were obtained from BASF. Li2CO3 was purchased from Sigma-Aldrich. LMC and LEDC were synthesized and purified as previously described. 26 All the reagents were stored in nitrogen filled glove box at room temperature and used without further purification.
The concentrations of LiPF6 and lithium carbonates were fixed at 0.65 mmol/mL in DMC. Samples were prepared inside the nitrogen filled glovebox. The samples were added to dry NMR tubes with DMSO-d6 capillaries, sealed with rubber septa, transferred out of the glovebox, flame sealed without air exposure, and analyzed by NMR spectroscopy. The samples were then stored at 55°C for 48 hours in an oil bath followed by analysis with NMR and GC-MS. Comparable samples were prepared on larger scale in glass ampules for the analysis of the solid resides. The solid residues were washed with DMC three times, dried overnight at room temperature, and analyzed with IR-ATR and XPS. Comparable samples were prepared in a Schlenk tube for overhead gas analyses by GC-MS. Comparable samples were prepared in stainless steel coin cells with comparable results to confirm that the glass containers play no role in the observed reactions.
NMR spectra of the samples were collected with a Bruker Avance III 300 MHz NMR spectrometer at room temperature before and after the high-temperature storage with and without proton decoupling. 19  GC-MS analyses were conducted with Agilent 6890-5973N GC equipped with an Agilent G973N mass selective detector. Liquid samples were diluted with dichloromethane, quenched with water to remove inorganic components, and the organic phase was utilized for the analysis. Helium was used as carrier gas at a flow rate of 24 mL/min. The initial column temperature was 40°C and the temperature was ramped at 10°C/min to 200°C and held at that temperature for 2 minutes with the total run time of 18 minutes. The gas analyses were performed by sampling the head spaces of a Schlenk tubes with a 10 µL GC syringe. Helium was used as carrier gas at a flow rate of 1.5 mL/min. The initial column temperature was set to 40 °C, and the temperature was ramped at 1 °C/min to 43 °C and held at that temperature for 2 min with the total run time of 5 min. The mass spectra obtained were compared to the NIST library to determine their molecular structures.
FTIR spectra of the dried solid residues were acquired on a Bruker Tensor 27 spectrometer equipped with a germanium crystal in attenuated total reflectance (IR-ATR) mode. Samples were transferred using air-tight vials and the spectrometer was operated inside a nitrogen filled glovebox to avoid air exposure. Each spectrum was acquired with 128 scans from 700 cm -1 to 4000 cm -1 at the spectral resolution of 4 cm -1 . The data were processed and analyzed using the OPUS and Originlab software.
XPS spectra of the dried precipitates were acquired using a Thermo Scientific K-alpha XPS. Samples were made into circular pellets with a press or stuck on a conductive carbon tape as a thin layer and transferred from the glovebox to the XPS chamber using a vacuum transfer module without exposure to air. An argon flood gun was used to avoid surface charge accumulation during sample analysis. The binding energy was corrected based on the C 1s of hydrocarbon at 284.8 eV. The data were processed and analyzed using the Thermo Avantage, XPS Peak 4.1 and the Originlab software.

Reactivity of lithium carbonates with LiPF6 in DMC.
In an effort to better understand the stability and decomposition products of components of the anode SEI with the electrolyte, the reactions of three different lithium carbonates, Li2CO3, LMC, and LEDC with LiPF6 have been investigated in DMC.
LEDC and LMC have been independently prepared via the chemical reduction by lithium naphthalenide while Li2CO3 is commercially available. [9][10][11] The stability of Li2CO3 has been investigated in DMC with and without added

NMR Spectroscopy of the solutions
All of the samples were analyzed by 19 F and 31 P NMR spectroscopy before and after storage at elevated temperature. The NMR spectra after storage at 55 o C are provided in Figure 5.1 and the spectral data listed in

FTIR Spectroscopy of the solid residues
The IR-ATR spectra of Li2CO3, LEDC, and LMC before and the residue after reaction of the different lithium carbonates with LiPF6 in DMC are provided in Figure   5.2. The IR spectrum of Li2CO3 contains two strong peaks centered at 1490 and 1450 cm -1 and a weak peak at 858 cm -1 . After storage of Li2CO3 in the presence of DMC for 2 days at 55°C, the insoluble residue was isolated. The IR-ATR spectrum exhibits essentially the same absorbance patterns as of pure Li2CO3 starting material, indicating little reactivity between lithium carbonate and DMC under the storage conditions. However, after storage of Li2CO3 in the presence of 0.65 M LiPF6 in DMC under similar conditions, the residue does not exhibit any peaks associated with Li2CO3 and instead contain several weak absorbances at 1300, 1162, and 758 cm -1 . The structure of the compound associated with these IR absorbances is unclear, but the absence of Li2CO3 in the residual solid is very clear.
Similarly, the characteristic peak of lithium alkyl carbonates, corresponding to C=O bonds, is observed at 1650 cm -1 . The residues obtained from samples containing LMC or LEDC in DMC exhibit IR absorptions similar to the starting material, whereas IR-ATR spectra of the residues obtained from samples containing LMC or LEDC and LiPF6 in DMC contain no peaks associated with LMC or LEDC, respectively. Absorptions are observed in the 720-740 cm -1 region which remain unidentified, but the absence of the lithium alkyl carbonates is clear. The absence of LMC or LEDC in the precipitate suggests that the majority of the lithium carbonates react with LiPF6 in DMC during storage at 55 o C.

X-ray photoelectron spectroscopy of the residues
The XPS spectra of the residues after reaction of the different lithium carbonates with LiPF6 in DMC are provided in Figure 5.3. The XPS spectrum of the residue from the reaction of Li2CO3 with LiPF6 has a very high concentration of Li and F, 38 and 42 %, respectively, and very low concentrations of C, O, and P 8, 6, and 6 % respectively.
The F 1s spectrum is dominated by a peak at 685 eV and the Li 1s spectrum is dominated by a peak at 56.4 eV coupled with the ~ 1:1 ratio of F to Li suggest that the residue is predominantly LiF. The F1s spectrum also contains a small shoulder at 687.5 eV along with related O 1s and P 2p peaks at 533 and 136 eV, respectively, consistent with the presence of a low concentration of LixPFyOz. The C 1s spectrum is dominated by a peak at 285 eV which likely results primarily from residual naphthalene or hydrocarbon contamination. There is no evidence for any residual Li2CO3 at ~290 eV in the C1s XPS spectra, consistent with the IR-ATR spectra.
The The reaction of Li2CO3 with LiPF6 results in the quantitative decomposition of the Li2CO3. A single gas, CO2, is observed by GC-MS. The residual solid from the reaction is predominantly LiF, as supported by XPS and IR-ATR. The solution phase contains a single decomposition product, F2PO2Li consistent with previous reports. 30 The low concentration of P in the residual solid is consistent with the generation of a soluble P containing species, F2PO2Li. The generation of soluble P containing species is the likely reason that the composition of the SEI typically has a much higher ratio of F:P than the 6:1 expected for LiPF6. When a 1:1 stoichiometry of Li2CO3 to LiPF6 is used, ~50 % of the LiPF6 is converted to F2PO2Li and LiF. This reaction is consistent with the equation 1.


Pitzer thermodynamic modeling study on solid – liquid equilibria of the quinary system LiCl – NaCl – CaCl 2 – SrCl 2 – H 2 O at 298.15 K

The Pitzer thermodynamic model for solid – liquid equilibria in the quinary system LiCl – NaCl – CaCl 2 – SrCl 2 – H 2 O at 298.15 K was constructed by combining the proper parameters and solubility product constants of the equilibrium solids for the subsystems in the literature. The solubility data of the systems CaCl 2 – SrCl 2 – H 2 O and LiCl – CaCl 2 – H 2 O were used to evaluate the model. The comparisons between the experimental and calculated solubilities indicate that the model is reliable. The solubilities of the quaternary system NaCl – CaCl 2 – SrCl 2 – H 2 O and the quinary system LiCl – NaCl – CaCl 2 – SrCl 2 – H 2 O were calculated, which can be used as a theoretical reference for comprehensive exploitation and utilization of this type of brine resource.

Introduction
The oileld brines in Nanyishan Section in the Qaidam Basin of the Qinghai-Tibet Plateau, which belong to the CaCl 2 -type, have high contents of lithium, potassium, calcium, and strontium accompanied with sodium, bromine, boron, and many other useful components. The concentration of lithium, calcium, and strontium in the brines is up to 0.98 g L À1 , 69.10 g L À1 , and 4.45 g L À1 , respectively, which is much higher than those in the salt lake brines in Qinghai-Tibet Plateau. 1 Aer a multistep exploitation of boron, potassium, and bromine, the brine largely consists of the complex system LiCl-NaCl-CaCl 2 -SrCl 2 -H 2 O. The phase equilibria and phase diagrams (solubility data) of the brine systems are theoretical foundations for the exploitation of the brine resources and describe the geochemical behavior of the brine and mineral systems. 2 Therefore, studies of the phase diagrams and thermodynamics of the brine systems containing lithium and strontium are necessary and urgent to extract natural resources. 2,3 In order to effectively exploit the oileld brine containing lithium and strontium, many systems containing lithium and strontium such as LiCl-NaCl-KCl-SrCl 2 -H 2 O at 298.15 K and SrCl 2 -KCl-NaCl-H 2 O at 298. 15 and 323 K had been investigated previously. [4][5][6][7] However, there are few reports in the literature for the systems containing both lithium and strontium with calcium. Experimental work and thermodynamic modeling are usually combined to complete the description of the phase equilibria for the salt-water systems. Global thermodynamic models, even though they are largely empirical, provide convenient representations of the thermodynamic properties for practical applications and for further research. 8 The Pitzer and Harvie-Weare (HW) chemical model, [9][10][11][12] which combines the Pitzer parameters and the solubility product constants of the equilibrium solids, have been widely used in the solubility predictions and the calculations of thermodynamic properties.
A number of experimental and theoretical studies on the subsystems of the complex system LiCl-NaCl-CaCl 2 -SrCl 2 -H 2 O at 298.15 K have been reported in recent decades. The solubilities of the systems LiCl-CaCl 2 -H 2 O, LiCl-SrCl 2 -H 2 O, NaCl-SrCl 2 -H 2 O, and CaCl 2 -SrCl 2 -H 2 O, which are the important subsystems of the oileld brine system, have been reported at 298.15 K, 13-18 but the phase diagram for the quinary system at 298.15 K is still lacking. The Pitzer parameters and the solubility product constants for the species in the quinary system have also been reported many times in the literature, 4,19-22 but it is difficult to select a consistent set of parameters out of these results. Therefore, the Pitzer thermodynamic model for the system LiCl-NaCl-CaCl 2 -SrCl 2 -H 2 O has not been constructed before. In this paper, the Pitzer thermodynamic model for the solid-liquid equilibria in the quinary system was constructed by choosing the appropriate parameters and solubility product constants for the species. Then, the solubilities of the quinary system were predicted.

Model approach
Pitzer developed an ion-interaction model and published a series of papers, 9,10 which provided a set of expressions for the osmotic coefficients of a solution and the mean activity coefficients of the electrolytes in the solution. On the basis of Pitzer's semi-empirical equations, Harvie and Weare developed the chemical equilibrium model, which is more convenient to use in the solubility calculations. 11,12 Since these equations are based on the excess free energy, all of the activity expressions are consistent and suitable for the application on different types of data (e.g. osmotic coefficients, activity coefficients, water activity, and solubility data) in the parameter regression and the calculation of other thermodynamic functions. Model validation involves the comparison of model predictions with data not used in the parameter evaluation process. The solubility data were calculated in this paper to affirm the model accuracy. The compositions of the solution and coexisting solid minerals can be identied with these equations. The equations are the main expressions for the model, shown as follows: In expressions (1) to (3), c, c 0 , and M represent cations and a, a 0 , and X represent anions. In addition, g i and m i are the activity coefficient and molality (mol kg À1 ) of the ions, respectively, z i is the valence state of the ions, and ø is the osmotic coefficient. Other symbols in eqn (1) to (3) are all described in the ref. 9-12. The solubility product constant (K sp ) of a hydrated salt M n 1 X n 2 $n 0 H 2 O at a stated temperature and pressure is shown in eqn (5).
In eqn (5), m represents the saturated concentration (mol kg À1 ) of the ions. In eqn (6), a w and M w represent the water activity and molar mass of water (kg mol À1 ), respectively, and the sum contains all solute species.

Model parameterization
The Pitzer model of the systems (Li + Na + K + Mg + Cl + SO 4 + H 2 O) 21 and (LiCl-NaCl-KCl-SrCl 2 -H 2 O) 4,5 at 298.15 K were successfully constructed. Therefore, the binary Pitzer parameters for LiCl, NaCl, and SrCl 2 and the solubility product constants for LiCl$H 2 O, NaCl, SrCl 2 $6H 2 O, and SrCl 2 $2H 2 O used in this research for the quinary system were taken from the literature. 4,21 It should be pointed out that the parameters and standard chemical potentials for LiCl were tted again using the osmotic coefficients, activity coefficients or the solubility data to suit the high brine concentration in Chinese salt lake brines (up to 20 mol kg À1 ) by Song et al. 21 The binary Pitzer parameters for CaCl 2 used in the research were acquired from the literature, 20 which can be used for concentrations up to 6 mol kg À1 . The mixing parameters for the quinary system, which are evaluated in the literature, 4,21 were also used in this study. All three types of parameters required in the quinary system: Pitzer binary parameters, mixing parameters, and solubility product constants, are listed in Tables 1-3, respectively.

Evaluation of parameters in the LiCl-CaCl 2 -H 2 O system
The solubilities of the system LiCl-CaCl 2 -H 2 O were calculated by Christov et al. 19 The binary parameters for LiCl and CaCl 2 are different from those in our study. In the literature, the binary parameters for LiCl and CaCl 2 were only used when their concentration did not exceed 19 mol kg À1 and 2.5 mol kg À1 , respectively. 19 However, the binary parameters used in this study can be used for a higher concentration. Therefore, the binary parameters for LiCl and CaCl 2 are more accurate. The mixing parameters q Li,Ca and J Li,Ca,Cl used in this study were the same as those from Christov et al. 19 With different parameters, the solubility product constants of CaCl 2 $6H 2 O, CaCl 2 -$4H 2 O and LiCl$CaCl 2 $5H 2 O were obtained again with the activity product constants, which differ from those in the literature. 19 The reference solubility data for this ternary system were used to evaluate the model. 13,14 The solubility data reported in these two references, 13,14 particularly the invariant point data, are different and shown in Fig. 1. The calculated data in this study and from Christov et al. 19 are also shown in Fig. 1. The calculated solubility curves saturated with CaCl 2 $6H 2 O and LiCl$H 2 O are nearly the same as those obtained in the study from Christov. 19 The calculated data in the curves saturated with CaCl 2 $4H 2 O and LiCl$CaCl 2 $5H 2 O in this study agree with the experimental data, but are smaller than those obtained in Table 1 Pitzer binary parameters of the quinary system LiCl-NaCl-

Evaluation of parameters in the CaCl 2 -SrCl 2 -H 2 O system
The solubility data were calculated with the PSC model, 22 but the calculation with the Pitzer model for the system was still lacking. The phase equilibrium of the ternary system was reported in detail by Bi et al. 18 The experimental diagram comprised of two invariant points, which were saturated with The solid solution (Ca,Sr) 2 $6H 2 O was found in the system. From the literature, 21 the interaction between the CaCl 2 and SrCl 2 salts is quite weak and the binary model parameters can represent the properties (component activities) and be used to evaluate the ternary system. The equilibrated solid phase is probably the ideal solid solution in the entire concentration range of the ternary system at T ¼ 298.15 K, rather than the single pure solid phase. Therefore, the mixing parameters q Ca,Sr and J Ca,Sr,Cl for the Pitzer model in this study were considered as zero. Herein, we also assumed that the solid solution was an ideal solution and predicted its solubility isotherm, which is the same as that in the literature. 22 The calculated solubilities and the experimental results are shown in Fig. 2

Solubility prediction
Phase equilibria and phase diagrams are the theoretical foundation for the exploitation of brine resources. The LiCl concentration is very small at the beginning of the evaporation for the mother liquor of the oileld brine, which was acquired from Nanyishan district in the Qaidam Basin. The quaternary

NaCl-CaCl 2 -SrCl 2 -H 2 O system
The solubility data for the quaternary system were calculated using the parameters in Tables 1-3. The experimental solubility data for the quaternary system at 298.15 K are not reported in the literature. With the calculated data, the dry-salt phase diagram was plotted using the J€ anecke indices of CaCl 2 and SrCl 2 with the unit g/100 g dry salt (m NaCl + m CaCl 2 + m SrCl 2 ), as shown in Fig. 3a. The phase diagram of the system consists of two crystallization elds: NaCl and (Ca,Sr)Cl 2 $6H 2 O. There is only one solubility curve AB co-saturated with NaCl and (Ca,Sr) Cl 2 $6H 2 O in the phase diagram. The phase diagram with the molalities of CaCl 2 and SrCl 2 as X-axis and Y-axis was also drawn in Fig. 3b. The pattern of the curve is likely the same as that in the ternary system CaCl 2 -SrCl 2 -H 2 O in Fig. 2. 18 No invariant points for the quaternary system were found.    Fig. 4. The probable saturated equilibrium solids for the invariant points of the quinary system are listed in Table 4. There could be three invariant points: (F, G, and H), (F, I, and J) or (J, K, and L). By combining the parameters in Tables 1-3 and the probable saturated solids in Table 4, the solubilities of the quinary system were predicted. There are some errors in the data of the point I, and the solubilities of the point K cannot be calculated.  Fig. 4. The J€ anecke indices of CaCl 2 and SrCl 2 , whose units are g/100 g dry salt (m LiCl + m CaCl 2 + m SrCl 2 ), are used as X-axis and Y-axis, respectively. There are six univariant solubility curves saturated with two salts (Fig. 4). The crystallization areas decrease in the sequence (Ca,Sr) 2 $6H 2 O, SrCl 2 $2H 2 O, LiCl$H 2 O, CaCl 2 $4H 2 O, and LiCl$CaCl 2 $5H 2 O. The concentrations of NaCl, CaCl 2 , and SrCl 2 are very small when the LiCl concentration is high in the solution, which shows that LiCl has a strong salting-out effect on other salts.

Conclusion
The Pitzer thermodynamic model for solid-liquid equilibria in the quinary system LiCl-NaCl-CaCl 2 -SrCl 2 -H 2 O at 298.15 K was constructed by selecting the appropriate parameters from the literature. The solubility data of the systems LiCl-CaCl 2 -H 2 O and CaCl 2 -SrCl 2 -H 2 O, not used in the parameterization process, were used to evaluate the model. Good agreement between the experimental and calculated solubilities shows that the model is reliable. By combining the Pitzer parameters and the solubility equilibrium constant equations of the equilibrium solids, the solubilities of the NaCl-CaCl 2 -SrCl 2 -H 2 O system were calculated. The invariant points of the quinary system LiCl-NaCl-CaCl 2 -SrCl 2 -H 2 O were affirmed and the solubilities of the quinary system were predicted. The thermodynamic model obtained in this study is essential for the development of universal thermodynamic models for brine systems containing calcium chloride and strontium chloride.

Conflicts of interest
There are no conicts to declare.


Adsorption properties of NO molecules on the hexagonal LaCoO 3 (0 0 1) surface: a density functional theory study †

Six types of adsorption con ﬁ gurations, together with two di ﬀ erent adsorption sites for NO adsorption on LaCoO 3 , were investigated via density functional theory. There were two types of terminal con ﬁ gurations for the LaCoO 3 (0 0 1) surface: CoO 2 - and LaO-terminated. We used the calculated adsorption energy, Mulliken and Hirshfeld charge analysis, the electron density di ﬀ erence and projected density of states analysis methods to determine the electronic and chemical properties of NO adsorbed on LaCoO 3 . We found that the Co-d and N-p orbitals played primary roles during electron transfer for NO adsorption on CoO 2 -terminated LaCoO 3 , and the O-p and N-p orbitals during the electron transfer for NO adsorption on LaO-terminated LaCoO 3 . Therefore, we assert that the optimal adsorption con ﬁ guration is Co – NO (N-end of NO on adsorption CoO 2 – LaCoO 3 ).

Introduction
With the rapid development of the science and technology of materials research, perovskites have received extensive attention 1 for their unique properties, such as stable structure, 2 electromagnetic characteristics, 3 thermal stability, 4 moderate price, widespread sources and gas sensitivity. In addition to the worsening environmental problems, 5-12 the air pollution due to particles is of concern. 13 Perovskites are sensitive to gas and act as gas sensing materials [14][15][16] that can adsorb and detect a very large number of poisonous gases, e.g., CO, NO x , SO 2 and VOCs. 17 These gases have received a wave of attention from research communities studying environmental catalytic materials, 18 and it can be seen that there is a lot of work to be done regarding the research on materials for environmental protection.
Studies have approved perovskites as catalytic materials for gas sensors 19,20 or locomotive tail gas processors to detect and adsorb the various harmful gases. Nunzio Russo et al. have expounded that the Li-substituted chromite catalyst, whose chemical formula is La 0.8 Cr 0.9 Li 0.1 O 3 , an alkali-metal-substituted perovskite compound, exhibits the highest activity in adsorbing poisonous gases, and they found that the catalyst had a great amount of weakly chemisorbed O À species that played key roles in the soot oxidation state. 21 Y. Teraoka et al. also revealed that the La-K-Mn-O perovskites showed activity in getting rid of NO x and soot, and thought that the potassium content was crucial since the activity of the catalyst varied with potassium content. 22 The nano-structured PrCrO 3 on CeO 2 was prefabricated by Debora Fino and Vito Specchia et al., who used the combustion synthesis technique, and then they adopted typical engine vibrations to test the catalytic activity. 23 All of the above cases involved research into the treatment of the exhaust gases by perovskites; however, the research on gas sensitive materials is also quite extensive. Changmin Shi et al. used density functional theory (DFT) to investigate LaMnO 3 , which can be used to detect CO, and calculate the adsorption energy for the CO molecule adsorption on LaMnO 3 . 24 Yongjia Zhang et al. also used DFT to study SrFeO 3 , but they selected the NO molecule to study the probable potential. 25 LaCoO 3 is a classical compound that has been widely studied and applied in various elds. Lihui Sun et al. used LaCoO 3 as the gas sensing material for detecting the CO molecule. 26 Beatriz Rivas-Murias et al. developed ferromagnetic insulating barriers of LaCoO 3 for devising spin-ltering tunnels in spintronic applications, according to the magnetic properties of LaCoO 3 . 27 NO gas has been, in recent years, recognized as a detriment to the atmospheric environment. In this study, we investigate the properties of LaCoO 3 as a catalytic agent for wiping out NO. Studying the electronic structures of the adsorption congurations can provide theoretical guidance for relevant experimental research.

Theoretical methods and computational details
CASTEP 28-33 package of Materials Studio (MS) 8.0. We did some optimization tests for the sake of determining the best optimized structure of the compound . We chose 14 types of orbital  cutoff values (300, 350, 400, 450, 500, 550, 600, 650, 700, 750, 800, 850, 900 and 950 eV) in computer simulation, which are shown in Fig. S1 in the ESI † and tried to gure out the most suitable energy cutoff value. Fig. S1 † shows that the orbital cutoff value of 600 eV is the best value. Based on this, the Kpoints were tested and the data are marked in Fig. S2. † Furthermore, in the process of geometry optimization, the functions were optimized under the LDA and GGA approximations to determine the best one. For these tests, we identied the orbital cutoff value as 600, K-points as 2 Â 2 Â 2, and the function as GGA with the PBEsol method to calculate the exchange and correlation energy. The constringency standards of optimal geometry, with respect to the maximum force, energy and maximum displacement, were 0.05 eVÅ À1 , 2.0 Â 10 À5 eV per atom and 0.002Å, respectively, with ultraso pseudopotentials. 34,35 The dipole correction was used to deal with the electric eld gradients in three-dimensional periodic supercells, which can affect electronic densities and binding energies.
In order to simulate the adsorption behavior of NO, the periodic supercell of 2 Â 2 Â 1 was constructed. This supercell for the CoO 2 -terminated surface of LaCoO 3 (0 0 1) contains 120 atoms, and 180 atoms for the LaO-terminated surface. Based on the literature, the length of the N-O bond was set as l N-O ¼ 1.17 A. 36 In our study, a NO molecule was attracted to the LaCoO 3 surface and the adsorption energy E ad was dened as follows: where E total is the total energy of NO adsorbed by the LaCoO 3 system, and E LaCoO 3 and E NO represent the energy of the LaCoO 3 and the free state of the NO molecule, respectively. There are two types of terminations in LaCoO 3 : LaO and CoO 2 terminations. The periodic boundary qualication was employed in the investigation with the reduplicated slab model. For the LaO-terminated surface, the slab consisted of nine La layers and nine Co layers, and for the CoO 2 -terminated surface, the slab consisted of six Co layers and six La layers, and both are shown in Fig. 1. Numerous early studies have authenticated that the (0 0 1) surface is the best in the ABO 3 , 37,38 and hence the LaOterminated and the CoO 2 -terminated (0 0 1) surfaces of LaCoO 3 are the most stable interfaces. The density functional theory calculations were used to study the adsorbing effect of NO on the LaO and CoO 2 -terminated LaCoO 3 (0 0 1) surface. The data could then guide the NO reaction in subsequent studies.

Results and discussion
3.1 Select functional, test orbital cutoff and K-points LDA and GGA are two types of general exchange-correlation energy approximation methods used during the process of geometry optimization. The structural optimization of LaCoO 3 using the 2 Â 2 Â 1 supercell was conducted by CA-PZ under LDA, and PW91, PBE, WC, RPBE, PBEsol under GGA. The specic computational conditions are shown in Table 1. By comparing the obtained results with experimental values, we found that LDA approximation methods were more applicable for calculating bonds lengths 39 and lattice parameters. 40 The Co-O bond length was 1.925Å and the lattice parameter was 5.446Å as calculated by CA-PZ, which approached the experimental values (1.929Å and 5.444Å); however the electron distributions were strongly restricted by LDA approximation, and the absolute value of the forbidden band width and the system binding energy were not suitable for analysis by LDA. Generally, the metal band gap calculated by GGA was closer to the experimental value than that calculated by LDA.
Relative deviations (%) of the optimized crystal parameters for hexagonal LaCoO 3 were as follows: Taking the experimental values of the lattice constant and energy band gap into consideration, we chose the GGA  approximation PBEsol functional to execute the subsequent study. There were two evident variation trends in Fig. S1: † when the cutoff energy ranged from 300 eV to 350 eV, the curve declined; when the cutoff energy ranged from 350 eV to 450 eV, the curve remained at. However, there was a signicant decline between 450 eV and 600 eV; hence, 350 eV was not the best choice. Aer 600 eV, there was a more stable trend in the image, making it clear that it converged very well in the structure optimization; therefore, we chose 600 eV as the value of orbital cutoff energy. Fig. S2 † clearly shows that it reached convergence for the Kpoints test using the GGA-PBEsol approximation when the value of the cutoff energy was 600 eV. It is clear that when the K-points value was 2 Â 2 Â 2, the image became smooth and did not change much; therefore, we chose 2 Â 2 Â 2 as the K-points value.
3.2 LaO-terminated and CoO 2 -terminated LaCoO 3 (0 0 1) surfaces Fig. 1 shows that the LaO and CoO 2 -terminated LaCoO 3 (0 0 1) surfaces have been adequately relaxed. We used the above functions and parameters to optimize them, and we could nd no drastic change for the LaCoO 3 (0 0 1) surface. The O atoms moved downward while the Co atoms moved upward on the CoO 2 -terminated LaCoO 3 (0 0 1) surface, which is in agreement with Sun's results. 41 During this process, we checked the length of the Co-O bond for the CoO 2 -terminated LaCoO 3 (0 0 1) surface. The length of the Co-O bond was always 1.937Å in every layer, which revealed that relaxation did not occur in this structure. This result is consistent with the slight relaxation phenomenon mentioned by Sun. 41 Nevertheless, we found that the La atoms moved downward and the O atoms shied upward for the LaO-terminated LaCoO 3 (0 0 1) surface. The length of the Co-O bond increased from 1.947Å to 1.953Å. This phenomenon showed that relaxation occurred in this structure.

NO adsorption on LaCoO 3
There were six adsorption congurations to choose from. When NO was adsorbed on different surfaces, including the CoO 2terminated and LaO-terminated LaCoO 3 (0 0 1) surfaces, all possible sites were investigated and only six congurations were suitable, which are listed in Fig. 2-4. We focused on adsorbing NO on the top site and bridge site of CoO 2 -terminated and LaOterminated LaCoO 3 (0 0 1) surfaces.
3.3.1 NO adsorption on the CoO 2 -terminated LaCoO 3 (0 0 1) surface. In this section, we focus on various phenomena and properties of NO adsorbed on the CoO 2 -terminated LaCoO 3 (0 0 1) surface, including adsorption energy, Mulliken charge and Hirshfeld charge analysis, the electron density difference and projected density of states (PDOS).
A. Adsorption energy. First, the clean CoO 2 -terminated LaCoO 3 (0 0 1) surface was fully optimized. Then, we put NO on the support and optimized this adsorption structure. In order to understand the optimal adsorption site of the NO molecules adsorbed on the CoO 2 -terminated LaCoO 3 (0 0 1) surface, we considered several different adsorption congurations. However, we found that only four types of adsorption congurations were not repetitive structures aer optimization, shown in Fig. 2    of these three O atoms is very unusual and they make the surface look like a hole. These O atoms are all connected to the same Co atom and they are also connected to three other Co atoms. In Fig. 4(B), the O atom of the NO molecule is close to the O atoms of the CoO 2 -terminated surface (NO O-end CoO 2 -LaCoO 3 bridge); these two adsorption sites belong to the bridge site.
Adsorption energies of all congurations were calculated on the basis of eqn (1); it can be clearly seen from Table 2 that all the values of adsorption energy were negative, implying that the adsorption structures were stable, and the adsorption processes were exothermic. At the top site (Co-NO and Co-ON), the values of the adsorption energies were À3.59 and À2.165 eV, respectively, and the values of the corresponding bond lengths were 1.649 and 1.732Å, respectively. At the bridge sites (NO N-end and NO O-end of the CoO 2 -LaCoO 3 bridge), the adsorption energies were À1.152 and À1.406 eV, respectively, and the corresponding bond lengths were 2.307 and 2.725Å, respectively.
Based on the above description, we will discuss the four adsorption congurations in this study. Since we will focus on discussing the NO reaction on the LaCoO 3 surface in another report, we will not consider the other factors. The Co atom formed a chemical bond with the NO molecule on the top site. Shi 42 pointed out that when CO was adsorbed on LaMnO 3, forming the covalent bond, the C-O bond length of the adsorbed CO molecule would be greater than that of the free CO molecule, which is in accordance with our calculated results. As mentioned above, the adsorption energy at the top site is relatively large and the values are within the range of chemisorption energy. However, the adsorption energy of the bridge site is much smaller than that of the top site and we found that NO molecules adsorbed on LaCoO 3 formed covalent bond interactions; therefore, we considered that NO molecules adsorbed on the bridge site were not stable chemisorptions. The most outstanding state conguration was the Co-NO conguration. The NO molecule is inclined to be adsorbed on the CoO 2terminated LaCoO 3 (0 0 1) surface, and the results are consistent with the literature, in which Sun 43 obtained the most favorable adsorption structure with the Fe-NO conguration (the N atom of NO molecule connected to the Fe site). Zhang 44 gured out in the adsorption of NO on the TiO 2 -terminated SrTiO 3 (0 0 1) that the N atom of the NO molecule was downward, which was more benecial for adsorption. Yongjia Zhang 45 reported that the optimal conguration was the Fe-NO structure for the FeO-terminated SrTiO 3 (0 0 1) surface adsorbed NO. This indicated that the B element in the ABO 3 perovskite was the key element for adsorbing gases, based on the above conclusions.
B. Mulliken charge and hirshfeld charge analysis. In order to present detailed results, we calculated the Mulliken charge and Hirshfeld charge for NO adsorption on the CoO 2 -terminated LaCoO 3 (0 0 1) surface in Table 2. Aer NO molecules were adsorbed on the CoO 2 -terminated LaCoO 3 (0 0 1) surface, the equilibrium of the initial charge distribution was destroyed. Table 2 shows whether the Mulliken charges of the O atoms of NO were negative for NO adsorption at the top site or bridge site on the CoO 2 -terminated LaCoO 3 (0 0 1) surface. The Mulliken charge of the N atoms of NO was negative on the top site; however, the charge of N atoms on the bridge site was positive. The total charge of O and N atoms (Q O+N Mulliken (e)) was also   (0 0 1). On the other hand, the stable chemical bond did not form at the bridge site, the changes of the values were extremely small, and the electron transfer occurred with NO.
C. The electron density difference analysis. Aer the adsorption of NO on LaCoO 3 , the charge on LaCoO 3 was equably distributed in the chemical bond, instead of collecting on a certain atom, revealing a strong covalent bond. In order to be more straightforward in explaining the impact of the charge density for NO adsorption on LaCoO 3 , we have calculated and drawn the electron density difference for NO adsorption on CoO 2 -terminated LaCoO 3 (0 0 1), as shown in Fig. 5. We dened the electron density difference (Dr) as follows: where r NO-LaCoO 3 is the total electron density of the adsorbed system, r LaCoO 3 is the electron density of the pure LaCoO 3 and r NO was the electron density of the free state of the NO molecule.
As can be seen from Fig. 5(A), the pronounced charge accumulation occurs on the bond between the N atom of NO and Co atom of CoO 2 -terminated LaCoO 3 , indicating that a large electronic accumulation was the key factor for bonding, and the formative chemical bond was very strong. However, Fig. 5(B) shows a similar phenomenon in comparison to Fig. 5(A), between the O atom of NO and the Co atom of CoO 2 -terminated LaCoO 3 , which also forms a stable chemical bond; however, the accumulated charge in Fig. 5(A) is more abundant than that in Fig. 5(B), suggesting that the formative chemical bond for the Co-NO conguration is more stable than that for the Co-ON conguration. This result proves that the adsorption energy of the Co-NO conguration is larger, compared to the Co-ON conguration. Both the O atom and N atom of the NO molecule obtained electrons for NO adsorption on the CoO 2 -terminated LaCoO 3 , and these results agree well with those mentioned above in the Mulliken and Hirshfeld charge analysis.
D. Projected density of states (PDOS) analysis. To obtain a further insight into the contributions of the Co, N and O atoms, we investigated the PDOS for NO adsorption on the CoO 2 -terminated LaCoO 3 (0 0 1) surface, as shown in Fig. 7-10. Fig. 10 shows the PDOS, including the N-p orbital of NO and the O-p orbital of LaCoO 3 , and NO before and aer adsorption. There was an evident change in the O-p PDOS near the Fermi level, and the overlapping peaks aer adsorption were signicantly increased, compared to those before adsorption, indicating that the electrons were transferred from LaCoO 3 to NO and there was a certain interaction between NO and the surface. It is evident from Fig. 9 that the PDOS of the O-p orbital of LaCoO 3 before and aer adsorption had scarcely any change. However, the PDOS of the N-p orbital of NO had some change, suggesting that there was no stronger bonding interaction between NO and LaCoO 3 . Both of the above results were consistent with the Mulliken charge and Hirshfeld charge analysis. We can conclude that less transferred charges meant weaker interactions between NO and the surface.   Compared to the before adsorption, the aer adsorption had a notable change. From the PDOS of N and Co for NO adsorption on the CoO 2 -terminated LaCoO 3 (see Fig. 7 aer), it is clear that the Co-d states offered tremendous contributions to bands in the scope of about À5 eV to 0 eV, with a negligible component of the N-p states. Because the surface lost a lot of electrons, and electrons were transferred to NO, the Co-N bond was formed. A strong overlap of Co-d and N-p states in the range of about À8 eV to 1 eV denoted the appearance of some hybridization between the d orbital of Co and the p orbital of N, indicating that an intense reaction occurred between the adsorbed NO and the Co of the surface. These results were also consistent with Mulliken charge analysis, Hirshfeld charge analysis and the electron density difference analysis. Therefore, we claim that the p orbitals of NO and the d orbitals of Co are the critical factors leading to bonding with adsorbed NO.  Table 2), it was clear that the charges of the O atoms of NO were negative, the charges of the N atoms were 0.160e, 0.120e, 0.120e and 0.05e, respectively, and the charges of N atoms were positive for NO (N downward Fig. 7 The projected density of states (PDOS) of the NO N-p orbital, and Co-d orbital on the CoO 2 -terminated LaCoO 3 (0 0 1) surface at the top site. "Before" indicates that NO was not adsorbed on the LaCoO 3 surface, and NO and CoO 2 -terminated LaCoO 3 were calculated in a single crystal structure (free-NO and free-CoO 2 -terminated LaCoO 3 ); "after" indicates that NO was adsorbed on the LaCoO 3 surface and the N-end (N atom downward) on the CoO 2 -terminated LaCoO 3 . Fig. 8 The projected density of states (PDOS) of the NO O-p orbital, and Co-d orbital on the CoO 2 -terminated LaCoO 3 (0 0 1) surface at the top site. "Before" indicates that NO was not adsorbed on the LaCoO 3 surface, and NO and CoO 2 -terminated LaCoO 3 were calculated in a single crystal structure (free-NO and free-CoO 2 -terminated LaCoO 3 ); "after" indicates that NO was adsorbed on LaCoO 3 surface and O-end (O atom downward) on CoO 2 -terminated LaCoO 3 . Fig. 9 The projected density of states (PDOS) of the NO N-p orbital, and O-p orbital on the CoO 2 -terminated LaCoO 3 (0 0 1) surface at the bridge site. "Before" indicates that NO was not adsorbed on the LaCoO 3 surface, and the NO and CoO 2 -terminated LaCoO 3 were calculated in a single crystal structure (free-NO and free-CoO 2terminated LaCoO 3 ); "after" indicates that NO was adsorbed on the LaCoO 3 surface and N-end (N atom downward) on CoO 2 -terminated LaCoO 3 . or O-NO) adsorption on the LaO-terminated surface. The total charges (Mulliken and Hirshfeld charges) of N and O atoms (Q O+N ) were negative for the N-end of NO adsorption on the LaOterminated surface, while they were positive for the O-end of NO on the LaO-terminated surface. These results revealed that the O atom gained electrons, the N atom lost electrons, and the NO (N-end) acted as an acceptor, due to the NO (N-end) forming stable chemical bonds with the surface for the N-end of NO adsorption on the LaO-terminated surface. Nevertheless, for the O-end of NO adsorption on the LaO-terminated surface, the NO (O-end) acted as a charge donor, forming bonds with the surface. The calculated adsorption energy for the N-end of NO adsorption on the LaO-terminated surface was À3.213 eV, which is much larger than that of the O-end of NO adsorbed on the LaO-terminated surface because the O-NO could form bonds and the O-ON could not form bonds. Fig. 6 shows the electron density difference for NO adsorption on the LaO-terminated surface. It was unmistakable that the charges accumulated between O and NO (Fig. 6A), in which the blue indicated losing electrons and the red indicated obtaining charges. From different electron density maps, we noticed that the result of the population analysis was relative to the electron density, which proved that our calculation was reasonable.
To ascertain the particular atomic orbital in the charge transfer during NO adsorption on the LaO-terminated surface, we investigated the PDOS of N and O, which are shown in Fig. 11 and 12, respectively, in the ESI. † By comparison of the PDOS before and aer adsorption, it was evident that the peaks of the held states shied to lower energy for NO adsorption on the LaO-terminated surface. A strong overlap of N-p and O-p states from À8 eV to 0 eV signied that there was some hybridization between the p states of O and p states of N. This result is in good agreement with Mulliken charge, Hirshfeld charge and the electron density difference. We also considered that the p orbitals of N and O played an important role in explaining the bonding mechanism.

Conclusions
In conclusion, we have conducted an overall inquiry into the adsorption properties of the NO molecule on the LaCoO 3 surface with different adsorption congurations and sites. Analysis of the adsorption energy, Mulliken charge and Hirshfeld charge analysis, electron density difference and projected Fig. 10 The projected density of states (PDOS) of the NO O-p orbital, and O-p orbital on the CoO 2 -terminated LaCoO 3 (0 0 1) surface at the bridge site. "Before" indicates that NO was not adsorbed on the LaCoO 3 surface, and NO and CoO 2 -terminated LaCoO 3 were calculated in a single crystal structure (free-NO and free-CoO 2 -terminated LaCoO 3 ); "after" indicates that NO was adsorbed on the LaCoO 3 surface and the O-end (O atom downward) on CoO 2 -terminated LaCoO 3 . Fig. 11 The projected density of states (PDOS) of the NO N-p orbital and O-p orbital on the LaO-terminated LaCoO 3 (0 0 1) surface at the top site. "Before" indicates that NO was not adsorbed on the LaCoO 3 surface, and NO and LaO-terminated LaCoO 3 were calculated in a single crystal structure (free-NO and free-LaO-terminated LaCoO 3 ); "after" indicates that NO was adsorbed on the LaCoO 3 surface and the N-end (N atom downward) on the LaO-terminated LaCoO 3 . Fig. 12 The projected density of states (PDOS) of the NO O-p orbital, and O-p orbital on the LaO-terminated LaCoO 3 (0 0 1) surface at the top site. "Before" indicates that NO was not adsorbed on the LaCoO 3 surface, and NO and LaO-terminated LaCoO 3 were calculated in a single crystal structure (free-NO and free-LaO-terminated LaCoO 3 ); "after" indicates that NO was adsorbed on the LaCoO 3 surface and the O-end (O atom downward) on LaO-terminated LaCoO 3 . density of states revealed the primary role of the Co-d orbital and N-p orbital during the electron transfer for NO adsorption on the CoO 2 -terminated LaCoO 3 and the pivotal role of the O-p orbital and N-p orbital during the electron transfer for NO adsorption on LaO-terminated LaCoO 3 . All results reected the fact that NO was adsorbed on the LaCoO 3 surface to form bonds such as Co-O, Co-N and O-N. I. Twagirashema et al. 46 and M. Pirez-Engelmann et al. 47 used the experimental method and also observed that LaCoO 3 removal of NO has a good effect.
Our calculations elucidated the predominant role of LaCoO 3 in adsorbing NO exhaust gas. If we allow the CoO 2 -terminated LaCoO 3 to act as the adsorption surface, the absorption efficiency is better. These results provide theoretical support for studying the reaction of NO and exhaust gas on LaCoO 3 , as well as excellent guidance for NO exhaust gas treatment.