Potential-Dependent Generation of O2– and LiO2 ... - ACS Publications

Jan 28, 2016 - Discharging of the aprotic Li–O2 battery relies on the oxygen reduction reaction (ORR) producing Li2O2 in the positive electrode, whi...
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Potential-Dependent Generation of O2− and LiO2 and Their Critical Roles in O2 Reduction to Li2O2 in Aprotic Li−O2 Batteries Yelong Zhang,†,§ Xinmin Zhang,† Jiawei Wang,† William C. McKee,‡ Ye Xu,‡ and Zhangquan Peng*,† †

State Key Laboratory of Electroanalytical Chemistry, Changchun Institute of Applied Chemistry, Chinese Academy of Sciences, Changchun, Jilin 130022, P.R. China § University of Chinese Academy of Sciences, Beijing 100039, P.R. China ‡ Department of Chemical Engineering, Louisiana State University, Baton Rouge, Louisiana 70803, United States S Supporting Information *

ABSTRACT: Discharging of the aprotic Li−O2 battery relies on the oxygen reduction reaction (ORR) producing Li2O2 in the positive electrode, which remains incompletely understood. Here, we report a mechanistic study of the Li-ORR on a model system, i.e., an Au electrode in a Li+ dimethyl sulfoxide (DMSO) electrolyte. By spectroscopic identification of the reaction intermediates coupled with density functional theory calculations, we conclude that the formation of O2− and LiO2 in the Li-ORR critically depends on electrode potentials and determines the Li2O2 formation mechanism. At low overpotentials (> 2.0 V vs Li/ Li+) O2− is identified to be the first surface intermediate, which diffuses into the bulk electrolyte and forms Li2O2 therein via a solution-mediated disproportionation mechanism. At high overpotentials (ca. 2.0−1.6 V vs Li/Li+) LiO2 has been observed, which can rapidly transform to Li2O2 by further electro-reduction, suggesting a surface-mediated mechanism. The solution-mediated Li2O2 formation that can account for the widely observed toroid-shaped discharged Li2O2 particles has also been thoroughly examined. Thus, O2− formation controls the overall reaction onset potential, and LiO2 formation demarcates the change from a solution- to surface-mediated reaction mechanism. The new findings and improved understandings of the Li-ORR in DMSO will contribute to the further development of aprotic Li−O2 batteries.

1. INTRODUCTION Molecular oxygen (O2), a natural product of photosynthesis under the sun, has drastically transformed our planet.1−3 For instance, the great oxygenation2 and the evolution of aerobic organisms3 are two notorious events that were triggered by O2 as it started to appear in the atmosphere some 109 years ago. O2 also plays a vital role in human beings’ daily activities from respiration to transportation.4−6 A common feature of utilizing O2 to sustain life resides in its unique oxidation ability to generate energy when reacting with organic compounds in food chains.4,5 As a strong oxidizer, O2 has also been considered to react electrochemically with a spectrum of inorganics including H2,7 Zn,8 and more recent Li9 to build efficient energy conversion and storage devices (H2−O2 fuel cell, Zn−O2 and Li−O2 batteries, respectively) that have the potential to tackle climate change by mitigating CO2 emissions.10,11 The Li−O2 battery, particularly the aprotic version, has attracted a great deal of interest because of its high theoretical energy density that is far beyond what the Li-ion cells can achieve, and has been considered as a potentially transformational technology.10−14 Discharging of the aprotic Li−O2 battery relies on the O2 reduction reaction (ORR) producing Li2O2 in the positive electrode. Therefore, toward developing © 2016 American Chemical Society

high performance Li−O2 batteries, it is crucial to have a fundamental understanding of the ORR at the positive electrode/aprotic Li+ electrolyte interface. Although O2 reduction to Li2O2 presents a seemingly simple case of ORR because no O−O bond scission is involved, the mechanisms including the first charge transfer step, the stability of O2− in Li+ electrolyte and the role of LiO2 intermediate, however, remain disputed so far, due to a lack of spectroscopic evidence of the ORR intermediates,12 and the formation of solid products that can easily obscure the interpretation of experimental results.12−14 Here, we report a mechanistic study of a model system, i.e., ORR on an Au electrode in a Li+ dimethyl sulfoxide (DMSO) electrolyte. By spectroscopic identification of the ORR products and intermediates coupled with detailed density functional theory (DFT) calculations, it is concluded that the generation of O2− and LiO2 intermediates depends critically on the electrode potentials and determines the Li2O2 formation mechanism. At high potentials (low overpotentials), O2−* Received: December 16, 2015 Revised: January 19, 2016 Published: January 28, 2016 3690

DOI: 10.1021/acs.jpcc.5b12338 J. Phys. Chem. C 2016, 120, 3690−3698

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The Journal of Physical Chemistry C (νO−O of 1110 cm−1; * denotes the adsorbed state) has been identified as the first intermediate, which can diffuse into the electrolyte and transform to Li2O2 therein via a solutionmediated disproportionation mechanism. At low potentials (high overpotentials), molecular LiO2* forms and agglomerates (νO−O of 1137 cm−1), which can transform to Li2O2 by further electro-reduction, suggesting a surface-mediated mechanism. For the solution-mediated Li2O2 formation that can account for the toroid-shaped Li2O2 particles widely observed in discharged Li−O2 cells, the kinetics and mechanism of the reaction of O2− and Li+ in solution phase have been examined in detail, both experimentally and theoretically. The new findings and clarified understandings of the Li-ORR reported here will contribute to the future development of Li−O2 batteries.

to fine powders prior to transferring to a 500 mL round-bottom flask with 20 g of 3 mm-diameter glass beads. The roundbottom flask with the dry solid mixture was transferred from the glovebox to the rotary evaporator and rotated under vacuum for 3 days. The solid reaction proceeds according to the reaction of [(Me4N)OH]·H2O+3KO2 = (Me4N)O2 + 3/ 2O2 + 3KOH. Finally, (Me4N)O2, i.e., TMA+O2−, was separated from KOH and superfluous KO2 by repeated extraction (three times) with liquid NH3 (∼50 mL for each time) using a specially designed Soxhlet extractor. After the ammonia was completely removed, the slightly yellow solid TMA+O2− (2.30 g) was obtained and its purity has been determined to be 98% by iodometric titration.14 Further characterization by Raman spectroscopy and magnetic susceptibility is shown in Figure S4, which are consistent with those reported by Dietzel et al.20 2.2. Physical Characterization. Electrochemical measurements were performed with either a VMP3 electrochemical workstation (BioLogic) or a CHI 700 D bipotentiostat (CHI Inc.). A multinecked, water-jacketed and airtight glass cell equipped with inlet and outlet valves for gas control was used throughout. The working electrode used for CV and spectroelectrochemistry was a 2 mm-diameter polycrystalline Au disk electrode, and for hydrodynamic voltammetry a rotating ring disk electrode (RRDE) with Au disk-Au ring configuration was used (Pine Research Instrumentation, part no. AFE7R8AUAU, collection efficiency 0.22). Prior to use, the working electrodes were polished with 0.05 μm alumina slurry to obtain a mirror-like surface. A platinum wire served as the counter electrode. The reference electrode was a partially delithiated LiFePO4, which was constructed by mixing active material with Super P (Timcal) and PTFE in the ratio of 8:1:1 (m/m). The prepared electrodes were vacuum-dried at 200 °C for 24 h and then preoxidized (10% of total capacity) to LixFePO4 (x = 0.9). The LiFePO4 reference electrode has a stable potential of 3.45 V vs Li/Li+. The assembly of the cell and the injection of electrolyte were conducted in an Ar-filled glovebox. After transferring the assembled electrochemical cell out of the glovebox, the electrolyte was bubbled with O2 for more than 10 min to make a saturated solution, and a continuous flow of O2 was maintained over the solutions during the electrochemical measurements. In situ surface enhanced Raman spectroscopy (SERS) was carried out with an airtight three-compartment spectroelectrochemical cell. An electrochemically roughened Au working electrode was placed behind a 1 mm thick sapphire window, the Au roughening procedure and the cell design can be found in ref 21. Raman spectra were recorded with a customized LabRAM HR800 confocal Raman microscope (Horiba Jobin Yvon). The spectrometer was equipped with an 18 mW He:Ne 633 nm laser source for excitation, a 1800 or 600 lines/mm grating to disperse the scattering light with different resolutions, and a long working distance objective lens (Nikon 50 × 0.45 NA) to focus laser beam on and collect scattering light from the Au electrode surface. The details of differential electrochemical mass spectrometry (DEMS) for chemical analysis have been reported elsewhere.15 Briefly, it is based on a commercial quadrupole mass spectrometer (Thermo Fischer) with turbo molecular pump (Pfeiffer Vacuum) that is backed by a dry scroll pump (Edwards), and a leak inlet which samples from the purge gas stream. A glass vial (15 mL) equipped with inlet and outlet valves was incorporated into the purging system of the DEMS.

2. EXPERIMENTAL AND COMPUTATIONAL METHODS 2.1. Materials. Commercially available anhydrous DMSO (Sigma-Aldrich, > 99.9%) was first dried over freshly activated molecular sieves (type 4 Å) for 3 days, then distilled under vacuum with added sodium amide resulting in a final water content of 4 ppm (determined by Karl Fischer titration apparatus). Electrochemical grade tetrabutylammonium perchlorate (TBAClO4, Fluka, > 99.0%) and battery grade lithium perchlorate (LiClO4, Aldrich, > 99.99%) were dried by heating under vacuum at 80 and 160 °C, respectively, for 24 h. Commercial lithium peroxide (Li2O2, Aldrich, 90%), potassium superoxide (KO2, Alfa Aesar, 96%), tetramethylammonium hydroxide pentahydrate ([(Me4N)OH]·5H2O, Acros, 99%) and liquid ammonia (NH3, Dalian Special Gases Co., LTD, 99.999%) were used as received. DMSO has been used as the electrolyte solvent for the study of ORR due to its proved stability against reduced O2 species,15,16 despite its limited stability toward the Li anode and high voltage.17 The purity and stability of the distilled DMSO solvent were further examined prior to its use in the ORR study. The reason to establish an appropriate solvent for the fundamental study is that the reduced O2 species involved in ORR are very sensitive to the impurities, particularly protic ones such as residual water and organic acids.18 These protoncontaining species can compete with Li+ ions to protonate the reduced O2 species (forming HO2 and H2O2) and interfere severely with the spectroelectrochemical signals of the main reactions (formation/decomposition of Li2O2) making the interpretation of experimental results difficult.19 The purity of the DMSO solvent was examined by GC-MS and no detectable impurities have been identified (Figure S1). The short-term stability of the DMSO against O2− was proved by the reversible redox behavior of O2/O2− couple in cyclic voltammetry (CV) with a scan rate range of 0.01−0.1 V/s (Ia/Ic = 1, where I is the current on the anodic (a) and cathodic (c) sweeps; see Figure S2); while the long-term stability was verified by UV−vis spectroscopy of DMSO containing dissolved tetramethylammonium superoxide20 (TMA+O2−, a stable superoxide compound at room temperature synthesized for this study; see following paragraph for details), in which essentially no change in the absorbance of O2− was observed for 1 week aging (Figure S3). TMA+O2− was prepared from metathesis reaction of KO2 and [(Me4N)OH]·H2O (obtained by vacuum drying of [(Me4N)OH]·5H2O), according to a modified synthetic procedure of ref 20. Briefly, in an Ar-filled glovebox, [(Me4N)OH]·H2O (10.90 g, 0.10 mol) and a 5-fold excess of KO2 (35.00 g, 0.50 mol) were ground separately in mortars 3691

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Li+ concentration is taken to be 1 M in DMSO. A different concentration would shift the free energy of Li+ by + kT ln([Li+]) assuming ideal solution behavior. Therefore, at, e.g., [Li+] = 0.1 M each Li atom incorporated into a surface species would be destabilized by 0.06 eV. Lower-coordination Au sites stabilize O2− by 0.1 eV and molecular LiO2 and Li2O2 by 0.3− 0.4 eV compared to Au(111).19,36

For the detection of the gases evolved, Ar carrier gas was used as an internal tracer with a flow rate of 2.0 mL/min, which was precisely controlled with a digital mass flow controller (Bronkhorst). GC-MS was performed using a MD800 benchtop quadrupole mass spectrometer coupled to a GC8000 gas chromatograph (Fisons Instruments, Manchester, UK). UV−vis spectra were obtained on an UV−visible spectrophotometer (UV-2550, SHIMADZU, Japan). Powder X-ray diffraction (PXRD) was carried out using a STOE STADI/P diffractometer operating in transmission mode with a primary beam monochromator and position sensitive detector. Fe Kα1 radiation (λ = 1.936 Å) was employed. Magnetic susceptibility was recorded with a MPMS7 Quantum Design SQUID magnetometer at magnetic fields of 0.1 T in the temperature range 2−300 K. A SX20 stopped-flow spectrometer was used to measure the kinetics of the reaction of Li+ and O2− in DMSO at 25 °C. 2.3. Computational Methods. DFT calculations for the reaction of O2 species in the solution phase were performed using Gaussian 09 at the B3LYP/6-311++G** level.22 Solvent effects were accounted for via the SMD continuum solvation model of Cramer and Truhlar.23 The free energy of each optimized species was calculated from its solution phase partition functions24 using the rigid rotor-harmonic oscillator approximation, and included zero-point energy correction. Harmonic vibrational frequencies for each species were computed analytically without scaling. For the LiO2 and Li2O2 dimers, the singlet−triplet splitting was evaluated and the lowest energy spin state for each was used in the analysis. Several possible structures for the LiO 2 dimer were investigated, while the Li2O2 dimer structure that we used was that reported to be the lowest energy configuration by Lau et al.25 Charged species (i.e., O2−, Li+) were modeled by complexing each with an explicit DMSO molecule in addition to the implicit solvent. Periodic DFT calculations for O2 species on Au(111) were performed in the generalized gradient approximation (GGARPBE26) as implemented in the VASP (version 5.3).27,28 The valence electrons were expanded in Kohn−Sham one-electron orbitals up to 700 eV, and the core electrons were described using the projector-augmented wave method.29 The RPBE equilibrium lattice constant of Au (4.198 Å) is in good agreement with the experimental value (4.08 Å).30 The Au(111) surface was modeled by a 2√3 × 2√3 surface unit cell and a slab of four metal layers. Neighboring slabs were separated by vacuum equivalent to seven metal layers (ca. 17 Å) with electrostatic decoupling in the z direction. The surface Brillouin zone was sampled on a Γ-centered 7 × 7 × 1 Monkhorst−Pack k-point grid. The DMSO/Au(111) interfacial environment was modeled as a layer of two DMSO molecules per 2√3 × 2√3 surface unit cell, for a coverage of 1/6 ML. Further details can be found in ref 19. The Fermi level of the Au(111) surface was verified to be located between the HOMO and LUMO of the DMSO layer.31 To avoid the error in the GGA total energy of gas-phase O2, the tabulated standard enthalpy and entropy for gas-phase H2O formation32 was used in calculating the free energy of O2.33 The top two layers of Au and all adsorbate and solvent molecules thereon were fully relaxed until the maximum force in all degrees of freedom decreased to below 0.01 eV/Å. The definition of the computational Li electrode34−36 and the procedure to account for the potential of the surface O2− state19 can be found in the respective references. The reference

3. RESULTS AND DISCUSSION 3.1. ORR Intermediates on Au in DMSO Electrolytes. The ORR in DMSO electrolytes was examined using CV with relatively high potential scan rates, typically 20 V/s. The reason for using high scan rates lies in that the ORR product of solid Li2O2 is insulating and can easily form a blocking film on electrode surface, which will change the electrode’s active surface area and obscure the interpretation of voltammetric results.13 It shall be noted here that a constant electrode surface area was assumed when the relevant equations were derived on stationary and rotating electrodes, such as Randles-Sevcik and Levich equations.37 High scan rates can partially alleviate this dilemma, albeit maybe at the expense of the quality of the voltammograms, as fewer amounts of reaction products are formed (could be less than a monolayer but well within the effective detection range of SERS38 used in this study), and therefore the electrode reaction under study is less inhibited. First we examined the O2 electrochemistry in a 0.1 M TBAClO4 DMSO electrolytes (without Li+) at a scan rate of 20 V/s (CVs with lower scan rates can be found in Figure S2). Figure 1a (black curve) shows the result, in which a cathodic peak at 1.82 V and a corresponding anodic peak at 2.93 V were observed with E0 of 2.38 V consistent with previous results in DMSO.39 The large peak separation of 1.11 V at high scan rate indicates a kinetically sluggish process.37 It has been established that in DMSO electrolyte containing tetraalkylammonium cations, O2 can be reversibly reduced to O2−.40 To confirm O2− formation on reduction of O2 in our experiments, in situ SERS measurement was performed. SERS spectrum was first collected at open circuit potential (OCP) prior to the ORR, and only the signals associated with DMSO electrolyte were observed (Figure 1b, spectrum no. OCP). However, SERS spectrum (Figure 1b, spectrum no. 1) collected at the end of the potential scan to 1.10 V (marked as 1 in Figure 1a) demonstrated the existence of O2− with a 1110 cm−1 band, despite a relatively small amount of charge (123 μC/cm2) that was passed to generate O2−. This 1110 cm−1 band is consistent with the vibration of O2− generated on Au by ORR in other solvents including acetonitrile21 and ionic liquid of 1-butyl-1methylpyrrolidinium bis(trifluoromethylsulfonyl)imide,41 and with the SERS spectrum of TMA+O2− dissolved in DMSO (see Figure S5). When a 0.1 M LiClO4 DMSO electrolyte was used for the O2 electrochemistry, remarkably different CV was obtained (Figure 1a, orange curve). For the ORR part of the potential scan (from OCP to 1.1 V), two reduction peaks, distorted due to high potential scan rate, were observed at 2.0 and 1.6 V, respectively. At lower scan rates these two peaks are better resolved,42,43 see also Figure S6. The other two features associated with the CV of ORR are (i) the onset potential of the ORR is positively shifted to 2.51 V compared to 2.40 V obtained in TBAClO4 DMSO electrolyte and (ii) the peak height of ORR is considerably increased. As the potential scan was reversed from 1.1 to 4.6 V for O2 evolution reaction (OER), no peak at 2.93 V corresponding to the oxidation of O2− was observed. 3692

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ance of O2−* (Figure S7). Based on these observations, we propose the following Li2O2 formation mechanism at potentials close to the typical discharge plateau voltage of aprotic Li−O2 batteries. Here “(g)” and “(sol)” indicate gas-phase and solvated species, respectively: O2 (g) + * + e− = O2−*

(1)

O2−* = * + O2−(sol)

(2)

2O2−(sol) + 2Li+(sol) = Li 2O2 (sol) + O2 (g)

(3) −

In this mechanism, the adsorbed superoxide (O2 *) is proposed as the first ORR intermediate, which can desorb from the Au surface and diffuse into the bulk electrolyte. The O2− in the bulk solution then reacts with Li+ ions therein producing Li2O2 that precipitates back onto the Au surface and is detected by SERS at long reaction time (Figure S7). A detailed study of the reaction of O2− and Li+ in solution is in section 3.2. Similar observation of solely O2−* and no other O2 species has been found in the potential range of OCP down to 2.2 V; see Figure S8. A recently work by Dilimon and co-workers45 suggested the possibility of Li2O formation at this modest overpotential via a chemical reaction of Li2O2 and LiO2, which in principle could occur. However, its kinetics shall be limited because the scission of the O−O bond of Li2O2 merely by LiO2 without any assistance of catalysts (either homogeneous or heterogeneous) would be difficult so that Li2O did not manifest itself in our SERS study (Figure S7). As the electrode potential was scanned from OCP to 1.9 V (marked as 2 in Figure 1a and the charge passed was 76 μC/ cm2) just beyond the first reduction peak, in situ SERS spectrum (Figure 1b spectrum no. 2) showed both the 1110 cm−1 band and two new bands located at 1137 and 790 cm−1, respectively. The 1110 cm−1 band is again assigned to O2−*. The 1137 cm−1 band is assigned to bulk LiO2 based on experimental results in acetonitrile21,43 and theoretical investigation of possible bulk LiO2 structures.25 The 790 cm−1 band is assigned to bulk Li2O2.21 When the potential scan was reversed for OER, two oxidation peaks were observed. The peak centered at 2.93 V was consistent with the oxidation peak obtained in TBAClO4 DMSO electrolyte and was attributed to the oxidation of O2−. The broad peak between 3.1 and 4.0 V was attributed to the oxidation of a mixture of LiO2 and Li2O2 based on SERS results. It is of significance to identify three reduced O2 species at this relatively low potential because these spectroscopic results provide invaluable information for the formulation of the reaction pathways at this stage. The presence of O2−* suggests that the solution-mediated pathway (eqs 1−3) may still be operating. However, the rapid formation of LiO2 and Li2O2 indicated that some other pathways for the Li2O2 formation exist, in addition to the solution-mediated mechanism. Based on the spectroscopic evidence of LiO2 and Li2O2, we proposed that at this potential, the adsorbed O2−* (via eq 1) can chemically bind a Li+ ion forming the other key intermediate, molecular LiO2*, which can further transform to molecular Li2O2* on Au surfaces via a surface-mediated mechanism:

Figure 1. (a) CVs on Au in O2-saturated 0.1 M TBAClO4 DMSO (black), and in O2-saturated 0.1 M LiClO4 DMSO with different cathodic cutoff potentials of 1.9 (red), 1.6 (green), and 1.1 V (orange) at a scan rate of 20 V/s. (b) SERS spectra collected at OCP and OSP, at the end of various cathodic cutoff potentials for ORR marked with numbers from 1 to 4, and at the end of two anodic potentials marked with 5 and 6 in (a). Hollow rectangles highlight the Raman band of the reaction products and intermediates.

Instead a broad oxidation peak between 3.0 and 4.1 V was obtained. All of the above observations suggest a complex electrode reaction, in which more than one electron was transferred and multiple steps were involved.37 For more insight into the ORR in LiClO4 DMSO electrolyte, CVs with various cathodic cutoff potentials and the corresponding SERS spectra were collected with the aim to identify the respective ORR products and intermediates. A SERS spectrum was first collected at 2.50 V, i.e., the onset potential (OSP) of ORR (marked as OSP in Figure 1a), in which a Raman band at 1110 cm−1 was identified (Figure 1b spectrum no. OSP). This 1110 cm−1 band was consistent with the vibration of O2− observed in TBAClO4 DMSO electrolyte during ORR (Figure 1b spectrum no. 1) and was assigned to the adsorbed O2−, which is ionically paired with, instead of being chemically perturbed by,20 solvated Li+ because essentially no band shift of O−O− is observed in the presence of high concentration Li+ ions. The observation of O2−* provides the direct evidence that the first step of ORR in DMSO containing Li+ ions is the formation of O2−, rather than the concerted Li+ and e− transfer forming LiO2, as proposed by other authors.13,14,39,44 Holding the electrode potential at 2.51 V for an extended period of time (several minutes) led to the appearance of bulk Li2O2 (790 cm−1) and gradual disappear3693

O2−* + Li+(sol) = LiO2 *

(4a)

O2 (g) + * + e− + Li+(sol) = LiO2 *

(4b)

DOI: 10.1021/acs.jpcc.5b12338 J. Phys. Chem. C 2016, 120, 3690−3698

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(5)

2LiO2 * = Li 2O2 * + O2 (g)

(6)

soluble LiO2 and Li2O2 at low potentials. Therefore, it would be interesting to examine the “running away” of O2− during ORR and study its reactions in bulk electrolytes. This “running away” issue was first examined using RRDE voltammetry, in which the potential applied to the disk electrode was scanned from OCP to 1.8 V at 25 mV/s to produce reduced O2 species, and the ring electrode potential was set to 2.7 V, a value that can exclusively oxidize O2− to O2. By using this configuration, O2− species generated at the disk electrode was indeed detected at the ring electrode (Figure 2 solid black curves), highlighting the

The molecular LiO2* and Li2O2* can agglomerate and form the respective solid phases, or more likely, a mixture of both at low potentials. As the electrode potential was scanned from OCP to 1.6 V within the second reduction peak (marked as 3 in Figure 1a and 130 μC/cm2 charge passed), the SERS spectrum (Figure 1b spectrum no. 3) showed that the ORR products and intermediates, based on the above assignments, were a mixture of bulk LiO2 (1137 cm−1) and Li2O2 (790 cm−1), but the O2−* species was absent. When the potential scan was reversed from 1.6 to 4.6 V for OER, two oxidation peaks were once again observed. The one at 2.93 V, attributed to the oxidation of O2−, decreased compared to that of the red curve in Figure 1a (with 1.90 V cutoff potential). The decrease in peak height suggests that much of the O2− had reacted away. Here we did not expect the disproportionation reaction (eq 6) to contribute significantly to the decrease in the oxidation peak at 2.93 V, because its reaction rates, estimated by cyclic voltammetry,39,43 are considerably slow, thus in the time scale of tens of milliseconds (variation of potential from 1.9 to 1.6 V at a scan rate of 20 V/ s) its contribution shall be minor. It is intriguing that CVs clearly indicate the existence of O2−, which, however, does not manifest itself in the SERS spectrum (Figure 1b spectrum no. 3). We attribute this observation to the diminished O2− concentration in the system due to surface electro-reduction (eqs 4−5) and due to displacement of O2− from the electrode surface by solid LiO2 and Li2O2, which is unfavorable for its Raman signal enhancement38 but does not preclude the electrooxidation of O2− during the reverse scan. We will further address the issue of O2− loss into bulk electrolyte during ORR by RRDE voltammetry and model its disproportionation reaction in section 3.2. The broad oxidation peak between 3.0 and 4.0 V increased in peak height and was due to the oxidation of a mixture of LiO2 and Li2O2 based on our SERS results. SERS spectrum (Figure 1b spectrum no. 4) was also collected at the end of the potential scan down to 1.1 V (marked as 4 in Figure 1a, 283 μC/cm2 charge passed and the products deposited were estimated to be about one monolayer46 and within the range of effective SERS enhancement38), and only the signal of Li2O2 (790 cm−1) was identified. Reverse scan from 1.1 to 4.6 V showed only one oxidation peak corresponding to Li2O2 oxidation, which was further corroborated by SERS (Figure 1b, spectra nos. 5 and 6) collected at the end of different anodic cutoff potentials during oxidation (marked as 5 and 6 in Figure 1a). This observation means that complete reduction of O 2 − * and LiO 2 * intermediates to Li2O2* occurs more rapidly at more cathodic potentials and highlights that Li2O2 formation at low potentials is dominated by surface-mediated mechanism. In addition, the electrochemical oxidation of Li2O2 to O2 does not produce stable LiO2 (or O2−) intermediate possibly due to the relatively high potentials applied, consistent with the previously results based on DEMS studies.21,47 3.2. Reaction of Li+ and O2− in Bulk Electrolyte. In the previous section, we spectroscopically identified three surface adsorbed O2 species of O2−, LiO2, and Li2O2, and noticed that O2−, i.e., the primary intermediate, can diffuse into the bulk electrolyte at high potentials, and be displaced by the less

Figure 2. RRDE voltammograms for ORR in 0.1 M LiClO4 DMSO electrolyte, in which a potential scan at 25 mV/s was applied to the Au disk electrode, and the Au ring electrode potential was set either at 2.7 V (black solid) or 4.0 V (red dot) to collect the reduced oxygen species swept from the disk.

importance of understanding the reactions of O2− and Li+ in bulk electrolytes. For comparison, a ring electrode potential of 4.0 V was also used to oxidize all the ORR products that were swept from the disk electrode, and an increased ring current was observed (Figure 2, red dashed curve). The increase in the ring current can be attributed to the additional oxidation of LiO2 and/or Li2O2 transformed from the “running away” O2−. The “running away” O2− has recently been noticed by others, but its fate in the bulk electrolyte remains debated.43,48−50 It is worth noting the shape of i-E curve obtained on the disk electrode (Figure 2, lower panel). In the potential range between OCP and 2.0 V, the ORR was not impeded and even a steady state current plateau was achieved within the potentials of 2.0−2.3 V. Below 2.0 V, the current decreased drastically indicating the formation of blocking films on the disk electrode. This U-shaped current response also highlights the electrode potential’s critical role in the Li2O2 formation mechanisms during ORR in the aprotic Li−O2 batteries. To study the reaction of Li+ and O2− in bulk solution, TMA+O2− having excellent solubility in organic solvents has been used as the O2− source to react with Li+ ions in DMSO. In a typical reaction, 0.240 mmol of TMA+O2− dissolved in 8 mL of DMSO and 0.260 mmol of LiClO4 in 2 mL of DMSO were mixed in a glass vial. The vial was equipped with inlet and outlet valves and was incorporated into the purging system of a quantitative DEMS.15 Upon mixing, gas bubbles formed in the vial vigorously, and the gas was identified to be solely O2, see Figure 3a. The amount of the gas was quantified to be 2.593 mL (or 0.1158 mmol), according to a published procedure.15 After reaction, white deposits were found at the bottom of the 3694

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The Journal of Physical Chemistry C

Figure 4. Reaction free energy profile at 298 K for the solution-phase disproportionation reaction leading to bulk Li2O2 formation. Zero on the y axis represents solvated O2− and Li+. “Sol”, “g”, and “s” indicate solution phase, gas phase, and bulk phase. Energies of all solvated species are calculated at infinite dilution. Snapshots of optimized lithiated O2 species are included; white = Li; red = O.

for the gas phase,51 which involves the rearrangement of a lower energy C2v LiO2 dimer to a structure similar to the C1 species presented here prior to O2 dissociation. However, the C2v LiO2 dimer reported in ref 51 is higher in energy than the present C1 LiO2 dimer in DMSO, and hence is not expected to lie on the minimum energy pathway for disproportionation. Upon the expelling of O2, further agglomeration and eventually precipitation of Li2O2 becomes the driving force of the overall reaction because the formation of bulk Li2O2 from molecular Li2O2 is highly exergonic. This eventually results in the oversaturation of O2 in the DMSO solvent and evolution of O2 from it. The endergonicity of 0.31 eV of eq 7c (value base on infinite dilute O2 concentration) is thus the main reaction barrier for this mechanism, the size of which is small enough to be efficiently overcome at ambient temperature at sufficient reactant concentrations. It should be kept in mind that the stability of the solution-phase reaction intermediates and products is not subject to the electrode potential as surface species are (discussed in section 3.3) and the reaction rate is not a direct function of the electrode potential, as long as the molecules are in no contact with the electrode or in contact with a part of the electrode that is not electrically conductive. It would be interesting to measure the kinetics of the reaction of Li+ and O2− in solution because it is the key chemical step that follows the charge transfer at potentials close to the discharge plateau of the aprotic Li−O2 batteries.43,44,47−52 To do so, a stopped-flow UV−vis spectrophotometer was employed to measure the overall kinetics of the reaction, eq 7a-7c, in which a 1 mM TMA+O2− DMSO was rapidly mixed with excess of LiClO4 in DMSO and the decay of absorbance of O2− at 300 nm was recorded, as seen in Figure 3d. It is evident that the spontaneous reaction follows the pseudo-first-order reaction (ln(At/εl) = −k1t + ln(A0/εl), At = εl[O2−]t, k1 = k2[Li+], ε300 nm = 9.63 × 102 L/mol cm, where k1 and k2 are the pseudo-first-order and second order rate constants, respectively) because a plot of log absorbance against time was well linear with the slope equal to k1 (or k2[Li+]). Variation of Li+ concentration confirms that the rate for the disproportionation process is first order in [Li+], see Figure 3d inset, and the second-order rate constant of the reaction of Li+ and O2− was measured to be 24.6 M−1 s−1. For the rate constant of the reaction of Li+ and O2−, Johnson et al.43 have reported a value of 0.07 s−1 for k1 (transformed to k2 = 0.7 M−1 s−1), based on

Figure 3. Reaction of TMA+O2− and LiClO4 in DMSO studied with (a) DEMS, (b) PXRD, (c) Raman spectroscopy, and (d) stopped-flow UV−vis spectrophotometry.

vial, which were identified to be Li2O2 by PXRD (Figure 3b) and Raman spectroscopy (Figure 3c). Combined, the above results indicate that Li+ and O2− can react readily at ambient conditions to produce Li2O2 and O2, possibly via the following solution-phase mechanism that involves the disproportionation of LiO2 to Li2O2 and O2 (eqs 7a-7d,e). Here “(s)” indicates a bulk phase: O2−(sol) + Li+(sol) = LiO2 (sol)

(7a)

2LiO2 (sol) = {LiO2 }2 (sol)

(7b)

{LiO2 }2 (sol) = Li 2O2 (sol) + O2 (sol)

(7c)

2Li 2O2 (sol) = {Li 2O2 }2 (sol)···Li 2O2 (s)

(7d,e)

We have computationally investigated this reaction mechanism. The reaction free energy (at 298 K) profile is shown in Figure 4 (structures and properties of the various O2 species as solvated in DMSO are given in Figure S9 and Table S1, respectively). The key disproportionation step, eq 7c, is calculated to be endergonic by 0.31 eV without any additional kinetic barrier. It proceeds via the direct dissociation of O2 from a C1 triplet LiO2 dimer. The present solution-mediated disproportionation step differs from the one recently proposed 3695

DOI: 10.1021/acs.jpcc.5b12338 J. Phys. Chem. C 2016, 120, 3690−3698

Article

The Journal of Physical Chemistry C

beyond O2−*, and Li2O2 formation is dominated by the solution-phase mechanism. The νO−O of O2−* are previously calculated to be 1028 cm−1,19 in line with the 1100 cm−1 band observed in SERS, which may be due to a higher equilibrium coverage of O2−* than directly modeled (0.08 ML). Consistent with the expectation that Li2O2 can form in the solution and precipitate on to the Au electrode, we observed the gradual increase in Raman signal of Li2O2 and the concomitant decrease and finally disappearance of O2−* even at the OSP at long reaction time (Figure S7). When the electrode potential falls below 1.90 V, LiO2* becomes stable with respect to O2(g), and eqs 4 and 5 become operative, with molecular LiO2* and Li2O2* now being stable species on Au (see Table S2 for their properties). In this intermediate overpotential regime (experimentally, ca. 2.0−1.6 V) O2(g) can be reduced to O2−* first, or directly to LiO2* (eq 4b). In the high end of this regime, we hypothesize that the further electro-reduction of LiO2* is slow, so it can agglomerate on the Au surface and form solid LiO2, most likely in a mixture with Li2O2. Lau et al. have reported that structures in and near the global minimum configuration for bulk Li2O2 (the Föppl structure) is calculated to have peak O−O vibrational frequencies in the 799−865 cm−1 range,54 which is consistent with the broad peak centered at 790 cm−1 in SERS (Figure 1b, spectra nos. 2 and 3). The existence of solid LiO2 in the Li-ORR system, on the other hand, has been disputed on thermodynamic grounds.25,55 Lau et al.25 and Kang et al.55 have both reported the Pnnm structure to be the ground state of bulk LiO2, and calculated its equilibrium potential to be 2.61 or 2.68 V vs Li/Li + respectively, which is less than bulk Li2O2 and Li2O, but reached different conclusions on whether bulk LiO2 should appear in the ORR. The calculated peak O−O frequency for Pnnm differed (1103 {ref 25} vs 996 {ref 55} cm−1), and Kang et al. suggested that the Raman signal attributed to LiO2 may be due to the P63/mmc structure of LiO2 instead (1089 cm−1). Our experimental and theoretical findings suggest that the appearance of bulk LiO2 is a dynamic phenomenon as its SERS signal appears on Au only when molecular LiO2* as a key intermediate forms on Au at below 2 V. Although bulk LiO2 is less stable than bulk Li2O2, its further reduction to Li2O2 may need to overcome additional thermochemical barriers owing to the solid-state transformation involved. Thus, bulk LiO2 would have an appreciable lifetime to be detected in SERS, although it eventually disappears if the system is held in potentiostatic conditions for a prolonged period of time.21 The theoretically predicted stabilities of the reduced O2 intermediates lend further support to the dual-mechanism interpretation of the Li-ORR on Au in DMSO, and reveal the critical roles that O2−* plays in the solution-mediated mechanism in the low overpotential regime (> 2.0 V) and that LiO2* plays in the surface-mediated mechanism in the intermediate (ca. 2.0−1.6 V) and high (