Li–O2 Battery Degradation by Lithium Peroxide (Li2O2): A Model

Dec 11, 2012 - First, cathodes including carbon, α-MnO2 catalyst, and Kynar binder (PVdF-HFP) were exposed to Li2O2 and LiClO4 in propylene carbonate...
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Li−O2 Battery Degradation by Lithium Peroxide (Li2O2): A Model Study Reza Younesi,*,† Maria Hahlin,† Fredrik Björefors,† Patrik Johansson,‡,§ and Kristina Edström*,†,§ †

Department of Chemistry-Ångström Laboratory, Uppsala University, Box 538, SE-75121 Uppsala, Sweden Department of Applied Physics, Chalmers University of Technology, SE-41296 Göteborg, Sweden § Alistore, European Research Institute, 33 Rue Saint-Leu, 80039 Amiens Cedex, France ‡

S Supporting Information *

ABSTRACT: The chemical stability of the Li−O2 battery components (cathode and electrolyte) in contact with lithium peroxide (Li2O2) was investigated using X-ray photoelectron spectroscopy (XPS). XPS is a versatile method to detect amorphous as well as crystalline decomposition products of both salts and solvents. Two strategies were employed. First, cathodes including carbon, α-MnO2 catalyst, and Kynar binder (PVdF-HFP) were exposed to Li2O2 and LiClO4 in propylene carbonate (PC) or tetraethylene glycol dimethyl ether (TEGDME) electrolytes. The results indicated that Li2O2 degrades TEGDME to carboxylate containing species and that the decomposition products, in turn, degraded the Kynar binder. The α-MnO2 catalyst was unaffected. Second, Li2O2 model surfaces were kept in contact with different electrolytes to investigate the chemical stability and also the resulting surface layer on Li2O2. Further, the XPS experiments revealed that the Li salts such as LiPF6, LiBF4, and LiClO4 decomposed to form LiF or LiCl together with P−O or B−O bond containing compounds when exposed to Li2O2. PC decomposed to carbonate and ether based species. The degradation of the electrolytes increased from short to long exposure time indicating that the surface layer on Li2O2 became thicker by increasing time. Overall, it was shown that a mixture of ethylene carbonate and diethyl carbonate (EC/DEC) is more robust in contact with Li2O2 compared to PC. KEYWORDS: lithium−air battery, chemical decomposition, Li2O2, X-ray photoelectron spectroscopy, lithium−oxygen, lithium peroxide, XPS

1. INTRODUCTION The Li−O2 battery has recently attracted attention due to its potential to be used in, for example, future electric vehicles (EV) since it is predicted to provide a 2−3 times higher practical specific energy compared to the conventional Li-ion batteries.1 The high specific energy (gravimetric energy density) of the Li−O2 battery originates from the use of lithium metal as negative electrode and consumption of gaseous oxygen as “fuel” at the positive electrode. The lithium anode is based on the lightest metal resulting in the record specific capacity (3861 mAh/g), well beyond that of the common graphite anodes used commercially in Li-ion batteries.2 The consumption of oxygen is limited by the cathode’s intrinsic material properties (e.g., surface area, porosity, etc.); however, such cathodes can store more Li compared to intercalation cathodes.3 The anode and cathode are separated by either aqueous, nonaqueous, or solid state electrolyte, or by a combination of them.4,5 Considering the three parts (anode, cathode, and electrolyte) of the Li−O2 battery, they all suffer from several obstacles including the instability of the electrolyte solvents and salts, synthesis of an efficient catalyst, inefficient cathode formulation, lithium dendrite formation, etc.1−4 Among the aprotic solvents, © 2012 American Chemical Society

carbonate based ones, such as propylene carbonate (PC), have been widely used as a base for electrolytes for Li−O2 batteries.6−9 However, carbonate based solvents decompose during the discharge of the battery to form lithium carbonate, lithium alkyl carbonates, water, etc.6−9 Mass spectroscopy (MS) revealed that CO2 forms during charging7,8 and an X-ray photoelectron spectroscopy (XPS) study indicated that the decomposition products formed on the surface of the carbon cathode during discharge would disappear predominantly during charging,9 in line with the MS results. The failures in employing carbonate based electrolytes have led to suggestions to instead use ether based electrolytes,10 although they have also been demonstrated to decompose.11−13 In addition, also the Li salts in the electrolyte are degrading, which has added extra complexity to the issue of electrolyte stability. Li salts such as LiPF6, Li[B(CN)4], lithium bis(oxalato)borate (LiBOB), and lithium bis(trifluoromethanesulfonyl)imide (LiTFSI) all decompose during battery cycling,9,13−16 and therefore, finding Received: October 5, 2012 Revised: December 10, 2012 Published: December 11, 2012 77

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2. EXPERIMENTAL SECTION

new stable electrolyte solvents and salts is an urgent and important challenge. The degradation of the Li−O2 battery during operation has mainly been referred to be caused by attacks from reduced oxygen species, such as the super oxide radical, or by the Li anode suffering from an unstable solid electrolyte interface (SEI) in presence of oxygen.17 However, the chemical reactivity of Li2O2, which is the desired discharge product of the Li−O2 battery, was neglected in the early investigations of Li−O2 cells. Recently, this issue has been highlighted in both theoretical and experimental studies.18−21 The formation of Li2O2 as the discharge product has indeed been confirmed by different research groups.10,22−25 Similar to other alkali metal peroxides, Li2O2 is a very strong oxidizing agent,26,27 and therefore, we find that it is equally important to clarify how chemical decomposition of electrolytes can be a result of Li2O2 reactivity to ascertain that an electrolyte will be stable not only against reduced oxygen species but also against Li2O2. It is also important to investigate the stability toward Li2O2 for other components, including binders, catalysts, current collectors, separators, etc. A recent computational study suggested that Li2O2 acts as a degradation agent to decompose PC to alkyl carbonates by abstraction of hydrogen and by a nucleophilic attack followed by PC ring-opening.18 The authors further suggested that Li2O2 may form during discharge of the battery, but it immediately converts due to the reaction with PC.18 Similar results have recently been reported regarding Li2O2 decomposition of dimethoxyethane (DME), rendering formation of lithium alkyl carbonates.19 It has also been concluded that the carbon cathode reacts with Li2O2 to form lithium carbonates. Oswald et al. proposed that LiPF6 and LiBOB salts may decompose in contact with bulk Li2O2 to form LiF and B2O3, respectively.20 In contrast, it has recently been reported that LiBOB, LiBF4, and LiTFSI salts and PC, EC, and DME solvents are stable toward Li2O2, while LiPF6 is not.21 Here, we present a detailed study on how different electrolytes react when exposed to Li2O2. We also investigate the stability of the different components of a cathode. XPS, being a powerful analytical technique to analyze surface compositions, was used to investigate the decomposition products. The choice of XPS allowed us to study both crystalline and amorphous degradations products. First, we present the results regarding the stability of carbon cathode components. The most common cathode formulation for Li− O2 batteries was used in which catalyst, binder, and carbon are mixed.6−9,11−14 The carbon cathodes were exposed to mixtures of Li2O2 and electrolyte and then analyzed by XPS. This experimental situation is similar to the discharge of a Li−O2 battery when Li2O2 and electrolyte meet at the cathode surface. Second, XPS experiments designed to investigate reactions between Li2O2 and an electrolyte were performed; layers of Li2O2 powder was placed on an aluminum substrate and exposed to several electrolytes, and then, it was brought to the XPS instrument. This is a simple design in which no binder, carbon, or catalyst were used; therefore, the decomposition of electrolyte solvents and salts could be investigated separately, since there was no overlap between the XPS peaks of electrolytes and the cathode components. This design allowed us to determine the surface layer formed on the Li2O2 as a result of the electrolyte decomposition. Carbonate and ether based electrolytes were chosen together with three common Li salts: LiPF6, LiClO4, and LiBF4.

2.1. Sample Preparation. Two types of substrates were used in this study: (i) the porous carbon cathode of a Li−O2 battery and (ii) Li2O2 powder placed on an aluminum plate. 2.1.1. Porous Carbon Cathode. The porous cathodes were made of carbon Super P (Lithium battery grade, Erachem Comilog), Kynar 2801 (PVdF, Arkema) as binder, and α-MnO2 nanowire as catalyst in a weight ratio of 25:33:42. Propylene carbonate (PC) and acetone were used as plasticizer and solvent to prepare a slurry of cathode components. The slurries were kept under magnetic stirring for about 4 h and then cast onto an aluminum mesh. The cathodes were dried at 120 °C overnight in a vacuum furnace placed within an argon-filled glovebox (H2O, O2 < 2 ppm). The cathodes were thereafter soaked in mixtures of 0.5 mL electrolytes and 20 mg Li2O2 powder (SigmaAldrich) and kept in a glass bottle in the glovebox for 24 h. 2.1.2. Li2O2 Samples. Binder free samples were prepared by placing Li2O2 powder on an aluminum plate. Since Li2O2 powder has very low conductivity, which induces differential charging effects during XPS measurements, a thin layer of Li2O2 powder was carefully laminated on to an aluminum substrate. To ensure that a fully covering Li2O2 layer was obtained, with negligible spectroscopic contribution from the substrate, the absence of characteristic substrate element (Al) was confirmed. The Li2O2 layers were then exposed to different electrolytes in the glovebox for 10 min and 48 h, respectively. 2.1.3. Electrolytes. Several carbonate and ether based electrolytes were investigated. PC (Ferro, Purolyte), EC/DEC (2:1) (Ferro, Purolyte), DME (Ferro, Purolyte), and TEGDME (Tetraethylene glycol dimethyl ether) (Sigma Aldrich) solvents, and LiPF6 (Ferro), LiClO4 (GFS), and LiBF4 (Tomiyama) salts were used for the electrolyte preparation. All the electrolytes used in this study are listed in Table 1. Prior to use, the salts were dried in a vacuum furnace

Table 1. List of Electrolytes Used symbol

electrolyte

I II III IV V VI VII VIII

0.1 M LiClO4 in TEGDME 1 M LiClO4 in PC 0.1 M LiClO4 in DME 1 M LiPF6 in PC 1 M LiPF6 in EC/DEC (2:1) 0.8 M LiBF4 in EC/DEC (2:1) 0.1 M LiPF6 in PC 0.1 M LiClO4 in TEGDME

first part of the study

second part of the study

placed within the glovebox (LiPF6 and LiBF4 at 60 °C and LiClO4 at 110 °C). The water content of the electrolytes was ≤10 ppm (measured by Karl Fisher titration or indicated by the manufacturer). 2.2. XPS Measurements. Prior to the XPS measurements, the carbon cathode or Li2O2 samples were washed with dimethyl carbonate (DMC) within the glovebox to remove the remaining Li2O2 powder and electrolytes from the samples. To avoid contact with air, the electrodes were transferred from the glovebox to the XPS analysis chamber using an airtight argon filled module. The XPS measurements were performed on a commercial PHI 5500 spectrometer, using monochromatic Al Kα radiation (1.4 keV) and an electron emission angle of 45°. All spectra were energy calibrated by using the hydrocarbon peak at binding energy 285.0 eV. To reduce the effect of differential charging on Li2O2 samples, a flood gun was also used during the measurements. However, at relatively low binding energies a minor charging effect was observed in all spectra.

3. RESULTS AND DISCUSSION This section is divided into two parts. In the first section, the results regarding the exposure of carbon cathodes to mixtures of Li2O2 and electrolytes are presented. In the second section, detailed surface characterizations of Li2O2 samples after exposure to different electrolyte combinations are discussed. 78

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3.1. Surface Characterization of Carbon Cathodes. In a nonaqueous Li−O2 battery, Li2O2 forms at the carbon cathode during discharge. Li2O2 is a highly reactive compound, and therefore, all cathode components need to be stable toward this discharge product otherwise degradation will lead to consumption of battery material and/or to passivation of the carbon cathode surface. Degradation of cathode components may also clog the pores in the cathode, which are needed to obtain a high discharge capacity.28,29 The chemical stability of the Kynar binder and the α-MnO2 catalyst was studied by exposing carbon cathodes to mixtures of Li2O2 and electrolytes for 24 h, and subsequently analyzing the surfaces by XPS. A commonly used formulation for the porous cathodes, in which carbon, Kynar binder (PVdF-HFP), and catalyst are mixed, was used.6−9,11−14 To investigate the Kynar binder degradation, we used a nonfluorinated salt, LiClO4, with TEGDME as an ether based solvent and PC as a carbonate based solvent (both have been demonstrated earlier in rechargeable Li−O2 batteries). The F 1s spectra of the samples are presented in Figure 1. The top spectrum belongs to a cathode after exposure to

Table 2. Summary of XPS Peak Assignments

F 1s

O 1s

C 1s

Cl 2p B 1s

P 2p

Mn 2p

assignment

chem. bond

binding energy, eV

Kynar (PVdF-HFP) LiPF6/LixPOyFz LiBF4/LixBOyFz LiF ethers carboxylates/carbonates Li2O2 carbonates carboxylates ethers hydrocarbons LiClO4 LiCl LiBF4 decomposition product of LiBF4 LiPF6 decomposition product of LiPF6 MnO2

CF PF BF LiF CO CO LiOOLi CO3 OCO CO CH ClO ClLi BF BO

688 687.5 687.2 685 533.6 532.2 531.5 289.7 289.1 287 285 207.4 197.5 195.4 193.3

PF PO

136.6 133.8

OMn4+O

642.3 653.9

peak reveals formation of LiF on the surface, and since the binder is the only source of F atoms in the system, we can conclude that LiF originated from decomposition of the Kynar binder. In contrast, only traces of LiF were observed when a cathode was exposed to a mixture of Li2O2 and electrolyte II. The trace amount was equal to the quantity observed in the reference sample (a cathode exposed to only electrolyte II). The two bottom F 1s spectra show results from cathodes after being exposed to electrolyte II without and with presence of Li2O2, respectively. These two spectra show mainly one peak at 688 eV, representing the binder of the cathode. It should be emphasized that the absence of the LiF peak in the PC based samples is not due to covering of the LiF by other decomposition products such as carbon based compounds. This is concluded based on the following: first, the O 1s and C 1s spectra of PC based samples are more or less similar in the absence and presence of Li2O2 (Figure S1 in Supporting Information), indicating that no carbon or oxygen containing decomposition products are present on the surface of the electrodes. Second, MnO2 peaks are clearly visible in the Mn 2p spectra with no changes in the intensities of the peaks. Therefore, the absence of the LiF peak in the F 1s spectra of the PC based samples together with the presence of the MnO2 peaks in the Mn 2p spectra reveal that no LiF was formed on the surface of the carbon cathode when exposed to Li2O2 and electrolyte II. The binder decomposition in the presence of Li2O2 is thus noticeable when a cathode is exposed to a TEGDME based electrolyte (I), while it is not when a cathode was exposed to a PC based electrolyte (II). To investigate possible binder decomposition with other ether based electrolytes, a similar experiment was also performed using electrolyte III (DME based). Also in this case, LiF formed (Figure S2 in the Supporting Information), which again indicates that the Kynar binder degrades in contact with a mixture of Li2O2 and ether based electrolytes.

Figure 1. F 1s and Mn 2p spectra of cathodes after being exposed to electrolyte I (purple spectra), electrolyte I + Li2O2 (green spectra), electrolyte II (blue spectra), and electrolyte II + Li2O2 (red spectra).

electrolyte I (see Table 1). It shows one main peak at a binding energy 688 eV, which represents the Kynar binder (a summary of all peak assignments using references30−34 is presented in Table 2). However, when a cathode was exposed to a mixture of Li2O2 and electrolyte I (the green spectrum), a peak at a lower binding energy (685 eV) increased dramatically. This 79

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Figure 2. F 1s, O 1s, C 1s, and P 2p spectra of Li2O2 after being in contact with electrolytes IV or V for short (10 min) or long (48 h) time.

Figure 3. (a) Relative element surface composition of Li2O2 samples after being exposed to electrolytes IV or V for 10 min or 48 h, respectively. (b) The assigned compounds to the P, F, and C elements and their relative amounts.

oxygen and form hydroperoxides and/or dialkyl peroxide.35−37 These products are chemically very reactive and can react with the battery components, including the Kynar binder. The oxygen atoms in Li2O2 may similarly abstract hydrogen from ethers, resulting in formation of hydroperoxides. It is worth mentioning that we have recently reported that the Kynar binder decomposes during the cycling of a Li−O2 battery using an ether based electrolyte, and consequently, LiF was observed on the cathode surface.13 However, the reason for Kynar binder

Based on these observations, we propose that Kynar may decompose either due to reactions with Li2O2 or due to reactions with degradation products formed by reactions between the electrolyte and Li2O2. However, direct decomposition of Kynar by Li2O2 is less likely since no binder decomposition was observed using electrolyte II. Therefore, the indication is that the Kynar binder decomposes in contact with ether electrolyte decomposition products. Via a process called “autooxidation”, ethers oxidize slowly in presence of excess 80

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Figure 4. F 1s, C 1s, and B 1s spectra of Li2O2 after being in contact with electrolyte VI for short (10 min) or long (48 h) time.

The compounds containing P−O bond form due to LiPF6 salt decomposition.9 It is also observed that the relative intensity of the LiF increases (almost twice) comparing the 10 min to 48 h exposure time, see Figures 2 and 3b. This shows that more of the LiPF6 salt decomposes by longer exposure time. The corresponding increase in the peak at 133.8 eV (P−O bond) supports this finding. Figure 3b shows that even after 10 min exposure about 14−17% of the surface composition of both samples is made up of LiF. The decomposition in such a short time indicates that LiPF6 is very unstable in contact with Li2O2. Figure 3a shows that the relative amounts of F and P increase by increasing the exposure time while the relative amounts of O and Li decrease. This suggests that a layer built up of decomposed salt is covering the surface of Li2O2. The C 1s spectra of all samples show three contributions at the binding energies of 285, 287, and 289.7 eV, representing hydrocarbons, ethers/alkoxides, and carbonates, respectively. The hydrocarbon peak originates from the solvents and also from carbon contamination (always present in ex-situ prepared samples). In all PC, EC, and DEC molecular structures, two C atoms are bonded to one O atom (ether) while one C atom is bonded to three O atoms (carbonate). Therefore, if assuming that the solvents (not decomposed solvents) remained on the surface, it is expected that ethers and carbonates would appear in the relative amount of 2:1 (ethers/carbonates). However, in the case of PC-based samples the relative amount of ethers to carbonates is ∼1.1:1 after 10 min exposure time. This ratio increases to ∼4.1:1 after 48 h. The relatively small value of the ratio after 10 min implies that some carbonate species formed on the surface. This is in line with a recent computational study suggesting that PC in contact with Li2O2 undergoes ringopening decomposition to form lithium alkyl carbonates (ROCO2Li),18 in which the ratio is 1:1. The 4-fold increase in the ratio by increasing the exposure time indicates that PC decomposes to ether compounds, in addition to alkyl carbonates. It has been shown that cyclic carbonate solvents such as PC may produce oligomer chains of poly(ethylene oxide) (CH2CH2O)n (PEO) and/or ROLi during the ring-opening decomposition due to reaction with Lewis acids such as PF5.39,40 As concluded above from the F 1s and P 2p spectra, the relative amount of LiPF6 decomposition products increase by increasing exposure time. Thus, the increase in the ether peak in C 1s spectra could be related to the presence of a larger amount of PF5. However, PC solvents can also degrade directly by Li2O2, which is, similar to PF5, a very strong oxidizing agent. This is discussed further in Section

degradation was ambiguous. It has been shown that the Kynar binder decomposes chemically in contact with lithium super oxide (LiO2).38 On the other hand, here, we report that an ether-based solvent in contact with Li2O2 produces decomposition products, which in turn will decompose the Kynar binder. The Mn 2p spectra are similar for all four samples (Figure 1), with peaks characteristic for the α-MnO2 catalyst, which thus is stable in contact with Li2O2. 3.2. Decomposition of Electrolyte Solvents and Salts in Contact with Li2O2. To investigate possible reactions between Li2O2 and the electrolytes single-handedly, Li2O2 surfaces were exposed to the electrolytes listed in Table 1 (electrolytes IV−VIII), and the results are summarized in electrolyte groups. 3.2.1. Surface Analysis of Exposure to LiPF6 in PC or EC/ DEC. In the first set of experiments, Li2O2 were exposed to two carbonate-based electrolytes, IV and V, which have commonly been used for Li−O2 and Li ion batteries. Two different exposure times, 10 min and 48 h, were used. The results from the shorter exposure time, 10 min, suggest that the electrolytes react as soon as Li2O2 forms during the discharge. The results from the longer exposure time, 48 h, about the length of 1−3 cycles in a Li−O2 cell, indicate the condition of electrolytes in contact with Li2O2 after a few cycles. Taken together the variation in reaction products in short and long time contact can be monitored. The survey scans of Li2O2 samples after being exposed to electrolytes IV and V showed that F, O, C, P, and Li containing compounds are present on the surface. The deconvoluted F 1s, O 1s, C 1s, and P 2p spectra are shown in Figure 2, while the relative surface compositions and the assigned compounds to F, P, and C elements are presented in Figure 3. It should be mentioned that a small shoulder at relatively low binding energies was present in all the spectra, caused by the poor conductivity of Li2O2. A flood gun was used during the XPS measurements to reduce the effect of differential charging, however, a minor charging effect was still observed in the spectra. The F 1s spectra consist of two peaks at 685 and 687.4 eV, representing LiF and LixPFy/LixPOyFz. The dominating contribution to the F 1s spectra, the LiF peak, forms due to the partial decomposition of LiPF6 by Li2O2. The two spin− orbit split peaks in the P 2p spectra further confirm the LiPF6 decomposition. The peaks at 133.8 and 136.6 eV in the P 2p spectra (Figure 2) represent P−O and P−F bonds, respectively. 81

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Figure 5. (a) Relative element surface composition of Li2O2 samples after being exposed to electrolyte VI for 10 min or 48 h. (b) Assigned compounds to the B, F, and C elements and their relative amounts.

Figure 6. F 1s, O 1s, C 1s, and P 2p spectra of Li2O2 samples after being in contact with PC solvent, electrolytes VII (0.1 M) and IV (1 M) for 48 h.

and ∼2:1 after 10 min and 48 h exposure times, respectively. As these values are very close to the 2:1 ratio in the structure of EC and DEC solvents, this suggests that the EC/DEC solvent mixture is more stable than PC in contact with Li2O2. 3.2.2. Surface Analysis of Exposure to LiBF4 in EC/DEC. To further investigate the stability of the EC/DEC solvent mixture to Li2O2, similar experiments were performed using LiBF4 instead of LiPF6. The survey scans of Li2O2 samples after being exposed to electrolyte VI showed that F, O, C, B, and Li containing compounds are present. The F 1s, C 1s, and B 1s spectra after 10 min and 48 h exposure time are presented in Figure 4. The C 1s spectrum for the sample exposed for 10 min shows a similar amount of ethers and carbonates indicating formation of carbonate species due to decomposition of EC/ DEC. Although this ratio increases to ∼2.6:1 after 48 h the increase is relatively small compared to the increase in the case of LiPF6 in PC electrolyte.

3.2.3, where attempts are made to discriminate between decomposition by Li2O2 and PF5. The ether compounds formed contribute to the C 1s ether peak but are also observable in the O 1s spectra, the peak at higher binding at 533.6 eV. Thus, the relative increase of this peak by increasing the exposure time supports the findings from the C 1s spectra. The other peak in the O 1s spectra at 531.5 eV matches the Li2O2 reference (Figure S2 in the Supporting Information). The carbonate compounds also appear at almost the same binding energy and overlap with the O peak of Li2O2. However, since the relative amount of carbonate compounds was small (concluded from C 1s spectra), the main contribution to the peak at 531.5 eV is from Li2O2. Therefore, the surface layer on Li2O2, which is made of the decomposed compounds such as LiF and ethers species, is not thicker than 5 nm.41 For the EC/DEC based samples, the relative amount of ethers/carbonates (as calculated from C 1s spectra) is ∼2.2:1 82

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Figure 7. O 1s, C 1s, and Cl 2p spectra of a Li2O2 sample after being in contact with electrolyte VIII for 48 h.

ing the charging process in the Li−O2 battery where Li2O2 needs to be decomposed. The higher amount of F and P may hamper decomposition of Li2O2 and reduce the rechargeability of the cell. Furthermore, the presence of a surface layer on Li2O2 would increase the charging overpotential.19 These results suggest that the Li−O2 battery performance suffers when using higher concentration of LiPF6 salt due to (i) a larger degree of PC decomposition and (ii) a larger amount of decomposition products of PC and LiPF6 on the surface of Li2O2 formed in a battery. 3.2.4. Surface Analysis of Exposure to LiClO4 in TEGDME. Since LiPF6 seems to be critical for electrolytes decomposition and since ether-based solvents have been suggested for the Li− O2 battery, we exposed a Li2O2 sample to electrolyte VIII for 48 h. The survey scans showed the presence of O, C, Cl, and Li on the surface of Li2O2. The O 1s, C 1s, and Cl 2p spectra are presented in Figure 7. The C 1s spectrum presents three peaks at binding energies of 285, 287, and 289 eV representing hydrocarbons, ethers, and carboxylates, respectively. The carboxylate peak with a relatively high contribution to the C 1s spectrum clearly indicates that TEGDME decomposes in contact with Li2O2 and forms COO containing compounds. It is known that ethers in the presence of excess of O2 oxidize to produce hydroperoxides,35,36 which contain a COO bond. Thus, the autoxidation of ethers offers a plausible explanation. The presence of COO containing compounds is also confirmed in the O 1s spectrum at 532.1 eV. The other two peaks in the O 1s spectrum at 531.5 and 533.3 eV represent Li2O2 and ethers. Also, the ether peaks in both C 1s and O 1s spectra originate from decomposed rather than remaining solvent since TEGDME evaporates in the XPS vacuum chamber. The Cl 2p spectrum shows two peaks at 197.5 and 207.4 eV representing LiCl and LixClOy, respectively. Thus, also the LiClO4 salt decomposes in contact with Li2O2.43

Taken together, the results obtained using LiPF6 and LiBF4 salts suggest that EC/DEC (2:1) is more robust compared to PC in contact with Li2O2. Based on the previously mentioned hydrogen abstraction in a computational study,18 EC/DEC (2:1) can be more stable compared to PC because it contains fewer accessible methyl group. The F 1s spectra in Figure 4 show two peaks at binding energies of 685 and 687.2 eV representing LiF and LiBF4/ LixBOyFz, respectively.32 Thus, LiBF4 also decomposes in contact with Li2O2, supported by the B 1s spectrum; the peaks at 193.3 and 195.4 eV belong to B−O and B−F, respectively.39 The relative surface compositions of Li2O2 samples exposed to LiBF4 based electrolytes and the assigned compounds to F, P, and C elements are presented in Figure 5. The relative amounts of both F and B increase by increasing the exposure time while the relative amount of O decreases. The relative amount of LiF increases from about 4 to 26 at% between 10 min and 48 h exposure time, respectively. This indicates that more LiBF4 salt decomposes on Li2O2 electrode from shorter to longer exposure time, similar to what we observed in the LiPF6 based samples. 3.2.3. LiPF6 Salt Concentration in PC Electrolytes. As mentioned earlier, both Li2O2 and PF5 (formed due to decomposition of LiPF6 by Li2O2) may contribute to the decomposition of PC. To investigate this decomposition in detail, Li2O2 samples were exposed to both PC solvent (without LiPF6) and to 0.1 LiPF6 in PC, which are compared to 1 M LiPF6 in PC. The results obtained after 48 h exposure are presented in Figure 6. The C 1s spectrum upon PC exposure shows a ratio of ∼1.5:1 carbonates to ethers. Thus, PC decomposes in contact with Li2O2 even in the absence of any Li salt, in agreement with the computational study.18 The C 1s spectra also show that the relative intensity of the ether peak, at 287 eV, increases with salt concentration, 0→0.1→1 M, and is thus mainly due to the formed PF5 reacting with PC.42 Moving to the F 1s and P 2p spectra, they both indicate that LiPF6 salt decomposes regardless of concentration. Both F and P contents are approximately twice as high when using the higher concentration. In addition, the relative amount of O on the surface layer is almost half for the more concentrated electrolyte. These findings are particularly important consider-

4. CONCLUSIONS XPS as a surface sensitive technique, capable to determine both crystalline and amorphous compounds, was used in this study to investigate the reactions between Li 2 O 2 and cell components. Primarily, the three common Li salts, LiPF6, LiBF4, and LiClO4, were all verified to be unstable and decompose in contact with Li2O2. These salts form products 83

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such as LiF or LiCl together with P−O or B−O containing species on the surface of Li2O2. Second, also the electrolyte solvents were found to decompose as follows: PC to carbonate species, while EC/DEC (2:1) is more stable although it also decomposes by Li2O2, and finally, TEGDME solvent degrades and forms carboxylates to a large extent when exposed to Li2O2. Importantly, we here highlight and exemplify that in a Li−O2 cell, the decomposition of one cell component may cause degradation of other parts of the cell. For example, degradation of ether based solvents by Li2O2 leads in turn to decomposition of the Kynar binder. The decomposition of Kynar binder consequently results in the formation of LiF on the surface of carbon cathode, leading to passivation of the cathode. We also show that the decomposition of LiPF6 by Li2O2 increases the degradation of the PC solvent. Taken together, the results show that the chemical stability of electrolyte solvents and salts, and also other cell components, in contact with Li2O2 should be carefully investigated to find a stable chemistry for the Li−O2 battery.

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ASSOCIATED CONTENT

S Supporting Information *

XPS spectra of carbon cathodes after being exposed to 0.1 M LiClO4 in TEGDME or to 1 M LiClO4 in PC in the presence or absence of Li2O2. XPS spectra of a carbon cathode after being exposed to 0.1 M LiClO4 in DME+Li2O2. XPS spectra of Li2O2 powder placed on an aluminum plate. This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected] (R.Y.); kristina.edstrom@ kemi.uu.se (K.E.). Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS Support from the Swedish Energy Agency (STEM), the Swedish Research Council (VR), StandUp for Energy, and the Storage project within KIC EIT Innoenergy are hereby acknowledged.

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REFERENCES

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dx.doi.org/10.1021/cm303226g | Chem. Mater. 2013, 25, 77−84