Identification and Implications of Lithium Superoxide in Li–O2

Identification and Implications of Lithium Superoxide in Li−O2 Batteries

L

can be formed by disproportionation of solid LiO2 or small clusters of LiO2 such as a dimer.

ithium oxides have recently been of much interest due to the large amount of research being done on Li−O2 batteries. Lithia (Li2O) and lithium peroxide (Li2O2) are stable in molecular and solid forms and have been wellcharacterized. Until recently, lithium superoxide (LiO2) had only been observed in low-temperature matrices (4 K)1 and the solid phase was believed to be unstable because of disproportionation2 to Li2O2 and O2, which is a thermodynamically favorable process.3 Recently, some studies of Li−O2 batteries have reported evidence for LiO2 in the discharge product, as discussed below, although the topic of lithium superoxide remains somewhat controversial.4 In this Viewpoint, we discuss the evidence for LiO2 (some examples of which are shown in Figure 1), the reason why

2LiO2(solid) → Li 2O2(solid) + O2(gas)

Most Li−O2 studies report Li2O2 as the major discharge product, although there have recently been numerous Li−O2 battery studies that have assigned a Raman peak at ∼1125 cm−1 to the O−O stretching of LiO2 in the discharge products.5−23 Whether LiO2 appears in the discharge product seems to depend on the cell configuration used, i.e., cathode, electrolyte, etc. In addition, some of these studies have also reported a Raman peak at around ∼1495− 1505 cm−1 that appears in conjunction with the ∼1125 cm−1 peak, which has been assigned to a LiO2−C mode where C is from the carbon-based cathode.5−16 Examples of Raman spectra from Li−O2 cells with these two peaks are shown in Figure 2.

Figure 1. Examples of characterization methods for a Li−O2 battery showing evidence of LiO2 formation (crystal structure shown in the center (unpublished results)), including selected area electron diffraction (SAED) giving evidence for some crystalline LiO2 (reprinted from ref 6), X-ray diffraction (XRD) pattern with some LiO2 peaks (reprinted from ref 6), differential electrochemical mass spectroscopy (DEMS) with evidence for one electron per O2 during charge (reproduced with permission from ref 8. Copyright 2016 Springer, Nature), and Raman spectra with two LiO2 peaks (unpublished results).

Figure 2. (a) Raman spectra of the toroids on the surface of the discharged AC cathode at the same discharge capacity of 1000 mAh/g with different current densities. The values of the peaks (in cm−1) are 250 (P1), 790 (P2), 1123 (S1), 1505 (S2), 1340 (D), and 1600 (G). S = superoxoide, P = peroxide, D and G = carbon peaks (reprinted from ref 5). (b) Raman spectra of discharge product on an Ir−rGO cathode for first and second discharges (reproduced with permission from ref 8. Copyright 2016 Springer, Nature.) (c) Raman spectra of Li−O2 cells with varying LiI contamination (reprinted from ref 12). (d) Raman spectra of the Fe7Ni3 cathode (reprinted from ref 11).

it can be stabilized, and the implications of its existence in Li−O2 batteries. In a Li−O2 battery, lithium superoxide can be formed by oneelectron reduction of O2 at the cathode surface and reaction with a Li cation, which can nucleate and grow as a solid. Li+(soln) + O2(surf) + e− → LiO2(soln) → LiO2(solid)

Received: March 9, 2018 Accepted: April 6, 2018

The competing process is a two-electron reduction of O2 and reaction with two Li cations to form Li2O2. Alternatively, Li2O2 © XXXX American Chemical Society

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DOI: 10.1021/acsenergylett.8b00385 ACS Energy Lett. 2018, 3, 1105−1109

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a dried PVDF binder (see also ref 8 Figure S8). In contrast, when PVDF is deliberately degraded in an alkaline solution and left to age in air, the intensities of the two peaks do not change significantly with time, as shown in Figure 3b. Second, if the discharge product is left in a vacuum, disproportionation of the LiO2 will occur and the intensity of the peaks should change. Figure 3c is an example of a case where the two Raman peaks decrease and a Raman peak from Li2O2 appears when left in a vacuum for 24 h.8 Third, the two Raman peaks in Li−O2 cells, which used a dried PVDF binder, should not disappear upon charge because PVDF does not decompose until >4.5 V.27 Figures 2d and 4 are examples

An alternative suggestion is that these peaks are from poly(vinylidene fluoride) (PVDF) binder that degraded during preparation of the cathode.4,24 The alternative assignment has been suggested because degradation of PVDF in alkaline solution results in Raman peaks25,26 at ∼1130 and ∼1525 cm−1, similar to the two reported for LiO2. The degradation of PVDF may occur from preparation of the PVDF binder at room temperature, instead of drying it at a higher temperature to drive off water, which could lead to an alkaline environment for degradation of the binder.24 If the preparation of the binder includes drying, then the PVDF binder will not degrade and peaks from it will not appear at ∼1130 and ∼1525 cm−1. Therefore, any peaks near these frequencies could be due to LiO2. For example, out of the 12 reports5−16 in the literature where the Raman peaks at ∼1125 and ∼1505 cm−1 were assigned to LiO2, seven5−8,11,13,15 reported that the PVDF binder was dried at 80−110 C, two12,14 used PTFE binder with drying and three9,10,16 used no binder. Thus, the two observed Raman peaks in these studies could not have come from degraded PVDF. There is further evidence, in addition to drying of the binder mentioned above, that the two Raman peaks are consistent with LiO2 and are not from degraded PVDF binder. First, the two Raman peaks, if they are from LiO2, will decrease significantly in intensity with aging of the discharge product because disproportionation of LiO2 to Li2O2 and O2 is favorable. In contrast, the peaks at ∼1130 and ∼1525 cm−1 from HF elimination25,26 of PVDF are not likely to disappear with time because the CC bonds from the degraded PVDF are stable during the aging process. This is confirmed by the decrease in intensity of the two LiO2 peaks after 70 h, shown in Figure 3a for a Li−O2 cell,6 which used

Figure 4. Galvanostatic discharge/charge curves of an activated carbon cathode. Three cells were first discharged with a current density of 0.1 mA/cm2, with a discharged capacity of 1000 mAh/g, and then were charged back with the same current density of 0.1 mA/cm2 until the voltage was up to (a) 3.5, (b) 3.8, and (c) 4.0 V. Raman spectra are shown for toroids on the surface of the cathodes for the three cells after three cycles with the upper cutoff voltage of (d) 3.5, (e) 3.8, and (f) 4.0 V. Reprinted from ref 5.

of the disappearance of the two LiO2 peaks upon charge due to electrochemical decomposition of LiO25 (see also ref 8 Figure S12). Finally, the FTIR spectrum of degraded PVDF has a new broad peak (compared to pristine PVDF)25 at 1530−1680 cm−1 (with a maximum at 1640 cm−1) that does not appear in the FTIR spectra (Figure S12a, ref 8) of a charged or discharged Li−O2 cathode used for studying LiO2 as a discharge product. Thus, there is significant evidence that the Raman peaks seen in Li−O2 batteries do actually correspond to LiO2 if the PVDF binder is prepared correctly, i.e., with heating to dry it. There is also considerable evidence from other techniques in addition to Raman that supports the identification of LiO2. For the battery based on lithium superoxide,8 other evidence was obtained from (i) XRD peaks that match LiO2, (ii) an EPR signal that matches molecular LiO2 observed in a low-temperature matrix, (iii) differential electrochemical mass spectrometry (DEMS) giving one electron per O2 during both discharge and charge, as expected for LiO2, and (iv) a Li−O2 discharge experiment to form Li2O2 from LiO2 in the presence of Ar (no O2) supports LiO2 formation. In a follow-up publication,28 titration evidence was presented. This evidence included titration by a TiOSO4 solution that showed that no Li2O2 was present in the discharge

Figure 3. (a) Raman spectra of the toroids on the surface of a discharged AC cathode for six different conditions. The values of the peaks (in cm−1) are 1123 (S1), 1505 (S2), 1340(D), and 1600 (G) (reprinted from ref 6). (b) Raman spectra of the PVDF binder degraded in alkaline solution to show little effect of aging on the 1125 and 1512 peaks (unpublished result). (c) Raman spectra of the discharge product from the Ir−rGO cathode after a 1000 mAh/g discharge (top) and after a 1000 mAh/g discharge followed by aging for 24 h in vacuum (bottom). The aging in vacuum results in decreased LiO2 and LiO2−C peaks and appearance of a Li2O2 peak at ∼780 cm−1 (Supporting Information in ref 8). Reproduced with permission from ref 8. Copyright 2016 Springer, Nature. 1106

DOI: 10.1021/acsenergylett.8b00385 ACS Energy Lett. 2018, 3, 1105−1109

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phase first. The DFT-computed activation barrier for disproportionation of the LiO2 dimer is ∼0.5 eV, which would indicate a short lifetime.35,36 However, the computed energy barrier for disproportionation of larger LiO2 clusters is significantly higher (∼1 eV) than the barrier in the LiO2 dimer.35 This means that the disproportionation rate would be slower for LiO2 clusters other than the dimer. As a result, the clusters may survive long enough to be incorporated into the growing discharge product. Because there is ample evidence for the existence of solid LiO2, as discussed above, either in a mixture with Li2O2 or by itself, it is of interest to consider O2 desorption from solid LiO2, a key step in disproportionation to Li2O2. This has been reported using DFT calculations for O2 desorption from crystalline LiO2. The initial step of O2 leaving the crystalline surface into a vacuum has a significant barrier of ∼0.9 eV8 for a low-energy LiO2 surface, as shown in Figure 6. The presence of an electrolyte in contact with

product, followed by a pH test consistent with the presence of LiO2. Several others, besides Lu et al.,8 have reported quantitative evidence for the existence of LiO2, although in these cases, LiO2 was one of several components of the discharge product in a Li−O2 battery. Bondue et al.29 used DEMS to show the amount of one-electron transfer that occurs to form LiO2 at certain overpotentials, which they used to support their hypothesis of a disproportionation reaction to form Li2O2. Mohzhukhina et al.30 used DEMS to show that superoxide has limited solubility in an ionic liquid and its presence in solution is impeded when the salt concentration is increased. In addition, Olivares-Mariń et al.22 used oxidation-state-sensitive full field transmission soft X-ray microscopy to quantify and localize with spatial resolution the distribution of the oxygen discharge products in a Li−O2 cell, which included significant amounts of superoxide-like phases (LiO2/Li2O2 ratio between 0.2 and 0.5). Numerous other papers have presented evidence, based on nonquantitative techniques, for LiO2 as one of the discharge products. In addition to the Raman studies5−23 already mentioned showing evidence for LiO2, other spectroscopic techniques have been used including EELS,31 XANES,22 and XPS.17 Other methods have also provided evidence for LiO2 including a quartz crystal microbalance,32 impedance measurements,33 electron diffraction,7,34 and STEM.31 For example, selected area electron diffraction (SAED)34 showed how LiO2 plays an important role in the formation mechanism of toroidal nanoparticles, as shown in Figure 5. We note that absolute identification of LiO2, as well as

Figure 6. DFT calculations of the energy barrier for desorption (ΔEact) of an O2 molecule from the (101)LiO2 surface in vacuum using the nudged elastic band (NEB). Reproduced with permission from ref 8. Copyright 2016 Springer, Nature.

the LiO2 should make O2 desorption more difficult and increase the barrier and thus the lifetime of LiO2 in a discharge product. The results of these computational studies of the reaction barrier to O2 desorption are consistent with the evidence for LiO2 in the discharge product. An implication of the existence of lithium superoxide in Li−O2 batteries is the possibility of its involvement in side reactions. The question of the reactivity of LiO2 in Li−O2 batteries is important because side reactions involving the discharge products can cause poor cycle life. In this discussion, we assume that the formation mechanism of LiO2 during discharge involves nucleation and growth of LiO2 from solution37 due to the solubility of LiO2 in nonaqueous solvents.38 In this mechanism, the LiO2 forms readily after reduction of O2 at a catalytic site on the cathode. The reactivity of the resulting solvated LiO2 species toward an ether solvent (proton or hydrogen abstraction) typically used in Li−O2 cells has been studied using DFT,39 and the ether was found to have a large barrier to decomposition; therefore, it is unlikely to occur. In this study, O2− was also found not likely to cause decomposition of ethers. Similarly, the charge process is likely to involve dissolution of LiO2 and migration of the solvated LiO2 to an oxygen evolution reaction (OER) site where the LiO2 is electrochemically decomposed to O2 and Li+. Thus, during discharge and charge, the LiO2 is not likely to cause decomposition of an ether electrolyte. Another possibility for side reactions is the reactivity of the surface of the discharge product with the electrolyte. If there are many defect sites on a LiO2 surface, as would likely occur in a toroid composed of both LiO2 and Li2O2 components, the surface will likely be reactive toward the electrolyte and cause side

Figure 5. In situ SAED analysis of the phase evolution and corresponding coupled reaction mechanisms. (a,b) Time-resolved SAED patterns illustrate the case of discharging, in which LiO2 is formed and subsequently evolves into Li2O2 and O2 through a disproportionation reaction and the O2 gas inflates the particle to a hollow structure (a), as illustrated schematically in (b). Reproduced with permission from ref 34. Copyright 2017 Springer, Nature.

other alkaline oxide species such as Li2O2, NaO2, etc., is very difficult and can only be done on the basis of evidence from various characterization techniques. However, when considered in totality, the evidence in the literature strongly supports the identification of LiO2 as one component of the discharge product under the right cell conditions. There have been several computational studies of the stability of LiO2 that have provided some explanation for the observation of it under some conditions. Even though thermodynamically LiO2 is unstable with respect to disproportionation (2LiO2(solid) → Li2O2(solid) + O2(gas)), kinetically it could exist for some time due to the reaction barrier for O2 elimination, if it forms in the solid 1107

DOI: 10.1021/acsenergylett.8b00385 ACS Energy Lett. 2018, 3, 1105−1109

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*E-mail: [email protected]. Tel.: +001 6302527380 (L.A.C.).

reactions. If the LiO2 surface is largely defect-free, it likely will not be reactive. It has also been suggested that singlet oxygen in the case of electrochemical decomposition of Li2O2 could be the cause of side reactions because of its potential reactivity.4 We note that, although singlet O2 could result from a charging mechanism involving electrochemical decomposition of solid Li2O2, the electrochemical decomposition of a solvated LiO2 molecule at a catalyst site will result in a triplet O2

ORCID

Kah Chun Lau: 0000-0002-4925-3397 Rajeev S. Assary: 0000-0002-9571-3307 Jun Lu: 0000-0003-0858-8577 Stefan Vajda: 0000-0002-1879-2099 Khalil Amine: 0000-0001-9206-3719 Larry A. Curtiss: 0000-0001-8855-8006

LiO2 (doublet)(adsorbed) → O2 (triplet)(soln) + Li+(soln) + e−

Notes

The triplet state results because oxidation of doublet LiO2 will lead to a thermodynamically preferred triplet O2 over the singlet oxygen. Note that oxidation of the LiO2 doublet to a LiO2 singlet cation species requires a higher (>4 V) charging potential. In terms of electronic conductivity, density functional calculations have shown that solid lithium superoxide is a half-metal.8 Relative to an insulating Li2O2 product, an advantage of having LiO2 as a component of the discharge product in a Li−O2 cell is that it will be electronically conductive and may lower the charge potential. This may be offset by instability involving disproportionation, unless methods are found to increase its lifetime and stabilize it. In summary, in this Viewpoint we have presented recent evidence from the literature using a variety of techniques that lithium superoxide can exist as a discharge product component in a Li−O2 cell under some conditions. Among these techniques, Raman spectroscopy has been widely used to identify LiO2 when formed on a carbon surface based on two peaks at ∼1123 and ∼1505 cm−1 from O−O stretching and LiO2 interaction with the carbon substrate, respectively. The disappearance of these peaks with aging or charging, as well as proper preparation of the binder used in the cathode, makes it unlikely that they are from some other source such as the binder. There have also been many other experimental techniques mentioned in the Viewpoint that provided additional evidence for the existence of LiO2. Identification of LiO2 is difficult and can only be done on the basis of evidence from various characterization techniques. When considered in totality, the evidence in the literature supports observation of LiO2 as one component of the discharge product that can occur when some cathode and electrolyte materials are used.

Views expressed in this Viewpoint are those of the authors and not necessarily the views of the ACS. The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was supported by the U.S. Department of Energy under Contract DE-AC02-06CH11357 by the Vehicle Technologies Office, Office of Energy Efficiency and Renewable Energy (H.-H.W., K.C.L, R.S.A, J.L., K.A., L.A.C), and by the Materials Sciences and Engineering Division, Basic Energy Sciences, Office of Science, (A.H., S.V.).

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Avik Halder† Hsien-Hau Wang† Kah Chun Lau‡ Rajeev S. Assary† Jun Lu§ Stefan Vajda† Khalil Amine*,§,∥,⊥ Larry A. Curtiss*,† †

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REFERENCES

Materials Science Division, Argonne National Laboratory, Argonne, Illinois 60439, United States ‡ Department of Physics and Astronomy, California State University, Northridge, California 91330, United States § Chemical Sciences and Engineering Division, Argonne National Laboratory, Argonne, Illinois 60439, United States ∥ IRMC, Imam Abdulrahman Bin Faisal University (IAU), Dammam 34212, Saudi Arabia ⊥ Material Science and Engineering, Stanford University, Stanford, California 94305, United States

AUTHOR INFORMATION

Corresponding Authors

*E-mail: [email protected]. Tel.: +001 6302523838 (K.A.). 1108

DOI: 10.1021/acsenergylett.8b00385 ACS Energy Lett. 2018, 3, 1105−1109

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DOI: 10.1021/acsenergylett.8b00385 ACS Energy Lett. 2018, 3, 1105−1109