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Insights into Electrochemical Oxidation of NaO2 in Na−O2 Batteries via Rotating Ring Disk and Spectroscopic Measurements Robert Morasch,†,‡ David G. Kwabi,*,‡,§,△ Michal Tulodziecki,†,‡ Marcel Risch,†,‡,# Shiyu Zhang,∥ and Yang Shao-Horn*,†,‡,§,⊥ †
Research Laboratory of Electronics, ‡Electrochemical Energy Laboratory, §Department of Mechanical Engineering, ∥Department of Chemistry, and ⊥Department of Materials Science & Engineering, Massachusetts Institute of Technology, 77 Massachusetts Avenue, Cambridge, Massachusetts 02139, United States S Supporting Information *
ABSTRACT: O2 reduction in aprotic Na−O2 batteries results in the formation of NaO2, which can be oxidized at small overpotentials ( 45) like DMSO (0.18 eV), implying that NaO2 dissolution is thermodynamically unfavorable. In a glyme-based electrolyte with 0.1 M Na+, this results in a rather low expected O2− solubility of 4.5 × 10−22 mM. SEM analysis of Na−O2 cell GDL cathodes show that NaO2 was predominantly oxidized at its interface with the conductive GDL. Na−O2 cells were discharged at 200 μAcm−2 to 1000 μAhcm−2 and subsequently charged to 50%, 75% and 90% of the discharge capacity, as shown in Figure 2a. Raman spectroscopy revealed only NaO2 for all discharged electrodes (Figure S2). Discharged electrodes were found to have faceted
our findings, yet our RRDE results show that such a mechanism is negligible compared to a direct oxidation at the cube/ electrode interface. Our RRDE measurements are sensitive enough to detect concentrations of soluble species on par with ESR measurements. The RRDE ring charge measured during oxygen reduction came to ∼14.4 nAh which corresponds to 26.9 nmol/L (20 mL electrolyte) of measured species over the course of 40s. ESR is similarly sensitive, in the range of 1 × 10−9 to 1 × 10−8 mol/L.27 Also, the total collected charge during reduction (ring and disk) closely matched the oxidation charge (see Figure S1a), both of which were constant across a wide range of rotation rates (400−1600 rpm). We take this as evidence for sufficient accuracy of our RRDE setup in determining relative contributions of solid/soluble species regardless of mass transport characteristics for soluble O2−. It should also be mentioned here that Øpstad et al.28 observed the formation of DMSO radicals in the presence of OH−, assumed to occur because of a proton abstraction from DMSO. O2− species may likewise induce proton abstraction;29−31 however, such reactivity has been shown to require several hours/days,30 whereas the amount of time between O2− formation on the disk and detection on the ring is on the order of milliseconds.32 We thus believe that the influence of the chemical instability of DMSO on our RRDE measurements is negligible, and attribute C
DOI: 10.1021/acsami.6b08355 ACS Appl. Mater. Interfaces XXXX, XXX, XXX−XXX
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Figure 3. (a) FY and (b) TEY XANES O K-edge and (c) Valence band XPS spectra data of powderous NaO2 (light blue) and Li2O2 (orange), KO2 (dark blue), and Li2CO3 (purple). Similar spectra for electrochemically formed NaO2 and Li2O2 are shown in dashed lines, from CNT electrodes discharged at 25 mA/gC to 4000 mAh/gC in Li−O2 and Na−O2 cells. Markers indicate approximate positions of σ* and π* features. (c) Fermi level is marked as EF. KO2 was included in the comparison as it exhibited chemical similarity to NaO2, and similarly low charging overpotential39 to NaO2. (d) Molecular orbital diagram of the superoxide (O2−) and peroxide (O22−) anion.
description in the Experimental Methods section) and KO2 powders (Sigma-Aldrich), using XANES measurements at the O K-edge (Figure 3). Bulk-sensitive fluorescence yield (FY) measurements revealed a sharp feature at ∼528.7 eV in both powderous and discharged NaO2 (Figure 3a). Based on previous O K-edge studies of alkali metal−oxygen species that have explained spectral features in terms of transitions from the O 1s core level to 2p frontier orbitals (π* and σ* antibonding states) of O2− anions,34−37 we assign the peak at ∼528.7 eV to the O 1s → π* transition. We sought to complement XANES measurements of unoccupied states, with analysis of occupied states using X-ray Emission Spectroscopy (XES),38 however the high photon fluxes required to compensate for comparatively weaker photoemission cross sections resulted in severe beam damage of the samples. We thus turned to valence band XPS measurements (Figure 3c), which show prominent peaks at binding energies of ∼4.8 and 9.0 eV for NaO2. FY XANES and XPS measurements of KO2 powder show similar features compared to NaO2, providing direct evidence of the similarity in electronic structure between NaO2 and KO2, with valence orbitals of O2− comprised of partially empty π* states (Figure 3d).18 In contrast, the FY XANES spectra of both electrochemically formed and powderous Li2O2 (Figure 3a) consist of a broadened peak corresponding to the O 1s → σ* transition,37 lying ∼2 eV above the π* peaks of NaO2 and KO2, and centered around ∼530.8 eV. This difference between NaO2/KO2 and Li2O2 spectra originates from the fact that in Li2O2, π* states would be expected to be completely filled18 (Figure 3d). This results in a significant energetic gap between occupied π* and σ* states, while a lower energy transition would be expected from filled to unfilled states within the π* manifold. XPS measurements show that band edges for all three compounds lay roughly equidistant from the zero binding energy position,
NaO2 cubes with edge dimensions varying between 3.0 to 7.8 μm (Figure 2b). Figure S3 shows additional images of NaO2 cubes after discharge, especially such cubes which are in direct contact with the conductive GDL fibers. The charging overpotential at 200 μA cm−2 (∼180 mV) was comparable to that reported for GDL electrodes ( 90%), synthesized NaO2 (see synthesis D
DOI: 10.1021/acsami.6b08355 ACS Appl. Mater. Interfaces XXXX, XXX, XXX−XXX
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Figure 4. Proposed NaO2 formation and decomposition mechanism. The cubes are first formed from aggregation of soluble NaO2. Initial oxidation occurs at the NaO2/GDL interface with only little oxidation at the surface, limited to the edge sites. At later stages the oxidation on the edges increases until the product loses its cubic shape. More uniform oxidation then occurs until the cube is mostly decomposed and parasitic reactions become dominant.
originates from the O 1s → π* transition in the CO32− anion,37 with the O 1s → π* peak from the NaO2 bulk phase in particular greatly diminished. The formation of thin layers of carbonate on electrochemically discharged NaO2 potentially explains the brief high-overpotential spikes observed at the beginning of charge (Figure 2a) as removal of passivating carbonate layers, as previously suggested.10 For Ti metal, the escape depth of total electrons is ∼5 nm at 560 eV,43 assuming that the escape depth scales as 1/ρ per material, where ρ is density, results in an escape depth for NaO2 of ∼10 nm. Another possibility for these potential spikes could be the nucleation of kinetically active sites for NaO2 oxidation; however, further study is required to explore this hypothesis experimentally.
which we take to indicate the Fermi level, EF. Combining these results with the above XANES measurements, which clearly show that the lowest unoccupied states of NaO2 and KO2 lie below those of Li2O2, strongly points to the conclusion that the band gap of NaO2/KO2 is lower than that of Li2O2. This is in agreement with DFT calculations by Lee et al.24 which suggest a higher electronic conductivity for superoxide phases compared to peroxides. It is important to note that XANES and XPS measurements only indirectly probe the band gap, the determination of which requires detailed measurement of occupied states with respect to unoccupied states on the same energy scale.38 The hypothesis that NaO2 has a lower band gap than Li2O2 is, however, further supported by the following arguments. The band gap of Li2O2 lies between filled π* and empty σ* states (Figure 3d), and may thus be approximated by the energetic separation between π* and σ* states in KO2 and NaO2 XANES measurements (Figure 3a). This results in an expected band gap for Li2O2 of about 6.5 eV, which is strikingly consistent with DFT estimates (5.70−7.76 eV).40 This band gap is likely to be much bigger than the energetic splitting between filled and unfilled π* states, which define the band gap in NaO2. Indeed, a combined XANES-DFT study of the electronic structure of KO2 by Kang et al.36 found that Coulombic interactions and spin−orbit coupling caused a ∼2.5 eV split between occupied and unoccupied π* states, with unoccupied states only 2 eV above the Fermi level. Likewise, a study of temperature-dependent KO2 conductivity has reported it to be semiconducting, with an activation energy of 1.3 eV.41 In further support of the above, both KO2 and NaO2 are yellowish in color, in contrast to Li2O2 and Na2O2, which are white. Such visible coloration is consistent with a band gap not much greater than in the 2−3.3 eV range, i.e., in the visible region, and reminiscent of semiconductors like ZnS and CdS, the latter of which is similarly yellow in color has a band gap around 2.42 eV.42 Surface-sensitive total electron yield (TEY) measurements (Figure 3b) showed clear evidence of carbonate species on both NaO2 and KO2 as seen by the emergence of a peak at 533.7 eV (Figure 3b). Comparing NaO2 and KO2 spectra to a reference Li2CO3 powder (Figure 3b) spectrum confirmed that the peak
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CONCLUSION In conclusion, RRDE measurements show that no considerable amount of soluble species is formed during the oxidation of NaO2, dismissing any solution-mediated oxidation process as the dominant oxidation path. Additional SEM images show that said oxidation occurs predominantly via charge transfer from solid NaO2 to the current collector. Furthermore, XPS and XANES measurements show that valence orbitals of the O2− anion in NaO2 are comprised of partially empty π* states, and suggest a smaller band gap in NaO2 than Li2O2, which has fully occupied π* states. This further supports that NaO2 is more electronically conductive than Li2O2 and that NaO2 oxidation occurs by charge transfer from NaO2. We therefore propose a charging mechanism for NaO2 where it is oxidized at the NaO2/GDL interface first, as seen in Figure 4. Most of the oxidation initially happens at this interface, but also at the energetically favored edges of the macroscopically cubic discharge product. At later stages of the oxidation, the process is dominated by oxidation at the NaO2/GDL interface, causing NaO2 to significantly decrease in size and lose its cubic shape. At the last stage, where the discharge product is mostly decomposed, the oxidation of resistive parasitic products begins to increase, thus significantly increasing the overpotential. Because understanding the charging mechanism of NaO2 is a crucial step toward rational design of Na−O2 battery E
DOI: 10.1021/acsami.6b08355 ACS Appl. Mater. Interfaces XXXX, XXX, XXX−XXX
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air exposure samples were sealed and stored in argon before rapidly moving the holder into the SEM chamber. NaO2 Synthesis. NaO2 was synthesized using a modified literature procedure.44 In the glovebox, finely ground Na2O2 (2.400 g, 30.77 mmol) was placed in a glass container that fits in the Parr reactor. The reactor was assembled, sealed and taken out of the glovebox. The system was slowly filled with 1100 psi O2. The bomb was sealed and slowly heated to 344 °C, at which point the pressure of the reactor reached 1700 psi. After 9 days, the reaction was allowed to cool down and O2 was slowly vented. The system was vacuumed and refilled with N2 for four cycles before it was brought back into the glovebox. The bright yellow solid was collected to give 2.380 g product (70% yield). X-ray Spectroscopy Measurements. XANES spectra from discharged electrodes and reference powders were obtained at the Canadian Light Source (CLS) on the spherical grating monochromator (SGM) beamline, where all the samples were mounted in an argonfilled glovebox before being loaded into the vacuum chamber under N2 gas flow. XANES spectra were recorded in the surface-sensitive TEY mode using specimen current and bulk FY mode using four energyresolved silicon drift detectors. All O K-edge spectra were calibrated to the t2g peak in rutile TiO2 at 530.5 eV.45 All XPS spectra were collected using a PHI 5000 VersaProbe II (ULVAC-PHI, INC.) using a monochromatized Al Kα source and a charge neutralizer. The powder samples were transferred from the Ar glovebox to the XPS chamber of the spectrometer using a sample transfer vessel (ULVACPHI, INC.) avoiding exposure of the sample to air. All spectra were recorded with pass energy of 23.5 eV and calibrated with the C 1s photoemission peak of adventitious carbon at 284.8 eV. Both KO2 (Sigma-Aldrich) and Li2O2 powder (Alfa Aesar, > 90%, ball-milled) were obtained commercially, whereas NaO2 was synthesized (see above).
EXPERIMENTAL METHODS
Cyclic Voltammetry Measurements. Three-electrode cells consisted of a 5 mm-diameter GC disk as the working electrode (5 mm diameter), an Au ring electrode (6.5 mm inner diameter, 7.5 mm outer diameter), and a sodium foil counter and reference electrode for DEGDME. For the DMSO measurement an Ag+/Ag electrode was used as reference electrode and nickel foam was used as counter electrode. RRDE experiments were performed in a water-free argon glovebox (H2O < 0.1 ppm, O2 < 1%). NaClO4 (98%, Sigma-Aldrich) salt was vacuum-dried at 150 °C for 24 h prior to electrolyte mixing. Diethylene glycol dimethyl ether (anhydrous, 99.5% Sigma-Aldrich) and dimethyl sulfoxide (anhydrous, 99.9% Sigma-Aldrich) solvent was dried over molecular sieves for more than 10 days. The water content was measured regularly and was below 15 ppm, determined by C20 Karl Fisher coulometer (Mettler Toledo). RRDE measurements were performed in O2-saturated electrolyte by sweeping the disk between 1.8 and 4.0 V vs Na+/Na at 50 mV s−1 while the ring is held at 3.0 V vs Na+/Na to oxidize soluble O2− species convected to it from the disk under rotation up to 1600 rpm. We assume the ring oxidation charge reflects the detection of solution-based O2− species produced during ORR, and estimate the fraction of QORR charge that is comprised of these species by normalizing it to the theoretical collection efficiency η of the RRDE geometry: Qring/ηQORR, where in our case, η is 23.5%. η was calibrated using the reversible O2/O2− couple in DMSO, however the collection efficiency using the well-known ferrocene/ferrocenium (Fc/Fc+) couple yielded a similar value of 25.8%. The other component of QORR thus reflects surface-bound or insoluble species, represented by QOER/QORR. The standard potential of reduction of the Ag+/Ag electrode is 0.8 V vs SHE and that of Na+/Na is −2.71 V. The potential scale for the DMSO measurement was therefore adjusted by 3.5 V and the potential measured was between −1.8 and 0.6 V vs Ag+/ Ag which correlates to 1.7 and 4.1 V vs Na+/Na, respectively. Cell Assembly and Electrochemical Measurements. The fabricated Na−O2 cells consisted of a sodium metal anode and commercial gas diffusion layers (FuelCellsEtc, USA) were used as cathodes (circular, diameter ∼1.27 cm/0.5 in.). The electrodes were vacuum-dried at 100 °C for 8 h, and transferred to a glovebox (H2O < 0.1 ppm, O2 < 0.1 ppm, Mbraun, USA) without exposure to the ambient environment. Diethylene glycol dimethyl ether (anhydrous, 99.5% Sigma-Aldrich) was used as the electrolyte solvent and sodium perchlorate (NaClO4, 98%, Sigma-Aldrich) as the conducting salt. DEGDME and sodium perchlorate were dried over molecular sieves (3 Å) for 10 days and under vacuum at 150 °C for 24 h, respectively. The electrolyte solution of 0.1 M NaClO4 in DEGDME was prepared in a glovebox with a final water content of
