Aprotic Li–O2 Battery: Influence of Complexing Agents on Oxygen

Jan 29, 2014 - School of Chemistry, University of St. Andrews, The Purdie Building, ... UMR CNRS 6007, 33 rue Saint-Leu, 80039 Amiens Cedex, France...
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Aprotic Li−O2 Battery: Influence of Complexing Agents on Oxygen Reduction in an Aprotic Solvent Chunmei Li,†,‡ Olivier Fontaine,§ Stefan A. Freunberger,∥ Lee Johnson,† Sylvie Grugeon,‡ Stéphane Laruelle,‡ Peter G. Bruce,*,† and Michel Armand‡ †

School of Chemistry, University of St. Andrews, The Purdie Building, North Haugh, St. Andrew, Fife KY16 9ST, United Kingdom Laboratoire de Réactivité et Chimie des Solides  UMR CNRS 6007, 33 rue Saint-Leu, 80039 Amiens Cedex, France § Institut Charles Gerhardt Montpellier, UMR 5253, Equipe Chimie Moléculaire et Organisation du Solide; Université Montpellier 2, CC 1700, Place Eugène Bataillon, 34095 Montpellier Cedex 5, France ∥ Christian Doppler Laboratory for Lithium Batteries, and Institute for Chemistry and Technology of Materials, Graz University of Technology, Stremayrgasse 9, 8010 Graz, Austria ‡

S Supporting Information *

ABSTRACT: Several problems arise at the O2 (positive) electrode in the Li-air battery, including solvent/electrode decomposition and electrode passivation by insulating Li2O2. Progress partially depends on exploring the basic electrochemistry of O2 reduction. Here we describe the effect of complexing-cations on the electrochemical reduction of O2 in DMSO in the presence and absence of a Li salt. The solubility of alkaline peroxides in DMSO is enhanced by the complexing-cations, consistent with their strong interaction with reduced O2. The complexing-cations also increase the rate of the 1-electron O2 reduction to O2•− by up to six-fold (k° = 2.4 ×10−3 to 1.5 × 10−2 cm s−1) whether or not Li+ ions are present. In the absence of Li+, the complexing-cations also promote the reduction of O2•− to O22−. In the presence of Li+ and complexing-cations, and despite the interaction of the reduced O2 with the latter, SERS confirms that the product is still Li2O2.

1. INTRODUCTION The confluence of a looming energy crisis and great strides in technological progress makes the harnessing of the second most electronegative element, oxygen, a nemesis for scientists. Li-air batteries have been a topic of considerable interest for almost a decade.1 Progress in Li° plating or its protection by an impervious solid electrolyte would still leave three fundamental problems at the positive electrode: (i) the insulating nature of Li2O2 and its poor solubility in aprotic solvents; (ii) the reactivity of the intermediate one-electron reduction products, O2•− or LiO2, which are strong nucleophiles and destructive to most solvents; and (iii) a large voltage gap between the reduction process O2 + 2e− + 2Li+ → Li2O22 and the reverse reoxidation.3−5 To date, a number of strategies have been reported that use additives to overcoming the poor solubility of Li2O2. For example, strong Lewis acids (C6F5)3B or simple sp3 borate esters have been added to the system to dissolve the insulating product.6,7 A conceptually elegant anion carrier selective to O22− versus O2•−, due to the multiplicity of spatially arranged H bonds, was shown to act as transition-metal-free superoxide dismutase.8 Other approaches have shown that additives can be used to facilitate the oxidation of the insulating Li2O2 using a redox-mediating molecule that is oxidized directly at the electrode and then in turn oxidizes the Li2O2.9−16 Given the above interest in additives for the Li-air battery, we wished to explore further the effect of additives on the ORR. © 2014 American Chemical Society

We made the hypothesis that onium salts with delocalization and a “π” system were likely to interact with the orbital in the O2x− molecule promoting solubility, provided the system was stable to the high basicity of, in particular, O22−. For this purpose, we mainly chose an imidazolium salt with the C2 position methylated. Two or more of these cations were further conjectured to create relatively hydrophobic and highly charged “pockets” in the solvent (here dimethyl sulfoxide, DMSO) favorable to both solubility of alkaline peroxides and faster electron injection to neutralize the high density of cationic charges. We begin by discussing the effect of the complexingcations on the solubility of alkaline peroxides, as this is a strong indicator of their interaction with O22−. We then go on to describe how the kinetics of reduction of O2 to O2•− and O22− in the absence and presence of Li+ is effected by complexation to the cations.

2. EXPERIMENTAL SECTION The solubilities of Li2O2, Na2O2, NaF, and KF in DMSO were measured by atomic absorption spectrometry (AAS). Solutions of the 0.1 M complexing-cations (see Table 1) in DMSO with an excess amount of Li2O2, Na2O2, NaF, or KF were stirred overnight at room temperature. After removing the undissolved Received: September 19, 2013 Revised: January 24, 2014 Published: January 29, 2014 3393

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Table 1. List of Chemical Structures and Their Abbreviations for the Synthesized Complexing-Cations

inlet and outlet. Polycrystalline Au disks (BAS Inc.) (diameter 1.6 mm or 100 μm for the ultramicroelectrode) were used as the working electrodes. Prior to use, the working electrodes were polished with 0.05 μm alumina slurry and rinsed with copious amounts of distilled water, followed by drying under vacuum. The surface area of the Au ultramicroelectrode (UME) was quantified by steady-state polarization in 0.5 M TBATFSI in acetonitrile containing 1.0 mM ferrocene (Fc) using the known diffusion coefficient, D = 1.7 × 10−5 cm2 s−1,17 for Fc. A platinum wire served as the counter electrode. A silver wire immersed in a glass tube containing the supporting electrolyte, separated from the main solution by a porous glass frit, was used as the reference electrode, which was calibrated against the Fc+/Fc couple after each experiment. It should be noted that

solids by centrifugation, the clear supernatant solutions were collected and diluted with distilled water for AAS. DMSO was distilled with the addition of NaNH2, then further dried for several days over freshly activated molecular sieves (type 4 Å), resulting in a final water content of ≤4 ppm (determined using a Mettler-Toledo Karl Fischer titration apparatus). Electrochemical-grade tetrabutylammonium bis(trifluoro(methanesulfonyl)imide) (TBATFSI) and batterygrade lithium bis(trifluoro(methanesulfonyl)imide) (LiTFSI) were used for preparing the electrolytes. Prior to use, both TBATFSI and LiTFSI were dried by heating under vacuum at 150 and 80 °C, respectively, for 24 h. Electrochemical measurements were performed within a glovebox in a multinecked, airtight glass cell with valves to control the gas 3394

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date)) (Di-ImTFSI). Di-ImTFSI was prepared from α,α′dibromo-o-xylene (10 mmol, 2.64 g) and 1,2-dimethylimidazole (20 mmol, 1.92 g) by the same procedure as compound C8TFSI, and the yield of final product was ∼87.5%. mp 89.7 °C; δH (300 MHz; DMSO-d6) 2.60 (6H, s), 3.86 (6H, s), 5.56 (4H, s), 6.91 (2H, q), 7.41 (2H, q), 7.53 (2H, d), 7.74 (2H, d); δC (75 MHz; DMSO-d6) 9.8 (CH3), 35.3 (NCH3), 48.4 (CH2), 113.5 (CF3), 117.8 (CF3), 122.0 (CF3), 126.3 (CF3), 121.7 (CH), 123.4 (CH), 127.3 (CH), 129.3 (CH), 132.5 (C), 145.8 (N2C). 2.1.5. Preparation of m-Xylene-α,α′-diylbis(1,2-dimethyl1H-imidazol-3-ium) Di(bis-(trifluoromethanesulfonylmidate)) (1,3-Di-ImTFSI). This compound was synthesized from starting chemicals α,α′-dibromo-m-xylene (10 mmol, 2.64 g) and 1,2-dimethylimidazole (20 mmol, 1.92 g) by the same procedure as C8TFSI. δH (300 MHz; DMSO-d6) 2.60 (6H, s), 3.79 (6H, s), 5.42 (4H, s), 7.29 (2H, d), 7.35 (1H, s), 7.46 (1H, t), 7.68 (4H, d); δC (75 MHz; DMSO-d6) 9.8 (CH3), 35.3 (NCH3), 50.8 (CH2), 113.5 (CF3), 117.8 (CF3), 122.0 (CF3), 126.3 (CF3), 121.7 (CH), 123.2 (CH), 127.6 (CH), 128.1 (CH), 130.2 (C), 135.7 (C), 145.2 (N2C). 2.1.6. Preparation of o-Xylene-α,α′-diylbis(1-methyl-1Hpyrazole-2-ium) Di(bis-(trifluoromethanesulfonylmidate)) (Di-PyTFSI). Compound Di-PyTFSI was prepared from α,α′dibromo-o-xylene (10 mmol 2.64 g) and 1-methylpyrazole (20 mmol, 1.64 g) by the same procedure as C8TFSI. mp 132.7 °C; δH (300 MHz; DMSO-d6) 4.07 (6H, s), 5.89 (4H, s), 6.92 (2H, t), 6.96 (2H, q), 7.51 (2H, q), 8.33 (2H, d), 8.60 (2H, d); δC (75 MHz; DMSO-d6) 36.5 (NCH3), 49.6 (CH2), 107.5 (CH), 113.5 (CF3), 117.8 (CF3), 122.0 (CF3), 126.3 (CF3), 128.5 (CH), 130.0 (CH), 130.6 (C), 137.7 (CH), 139.0 (CH). 2.1.7. Preparation of 2,2′,2″-Tris((1,2-dimethyl-1H-imidazol-3-ium)-ethyl)amine Tri(bis-(trifluoromethanesulfonylmidate)) (Tri-ImTFSI). 1-Methylpyrazole (12 mmol, 1.15 g) was added to a solution of tris(chloroethyl)amine hydrochloride (3 mmol, 0.72 g) in absolute ethanol (20 mL). The mixture was stirred for 2 days at room temperature. After evaporating the solvent, the white solid was collected, dried, and dissolved in distilled water (10 mL), then added to a solution of LiTFSI (9 mmol, 2.58 g) in distilled water (10 mL). The resulting white solid was recrystallized from boiling water to obtain the title compound. mp 71.8 °C; δH (300 MHz; DMSO-d6) 2.59 (9H, s), 2.97 (6H, t), 3.79 (9H, s), 4.12 (6H, t), 7.55 (3H, d), 7.63 (3H, d); δC (75 MHz; DMSO-d6) 9.8 (CH3), 35.2 (NCH3), 45.4 (CH2), 52.0 (CH2), 113.5 (CF3), 117.8 (CF3), 122.0 (CF3), 126.3 (CF3), 121.5 (CH), 122.7 (CH), 145.1 (N2C).

potentials are reported versus Li+/Li. Electrochemical measurements were carried out at room temperature using an Autolab PG30 electrochemical workstation. In situ surface-enhanced Raman spectroscopy (SERS) was carried out in a three-electrode electrochemical cell. A previously electrochemically roughened Au working electrode (see reference for experimental procedure)18 was placed behind a 1 mm thick sapphire window. Raman spectra were recorded using a customized Renishaw Raman system with an excitation wavelength of 632.8 nm. A Leica inverted microscope with a 50× objective lens was the collection optic. The power of the laser beam delivered to the electrode surface was estimated to be 2.5 mW, and the spectrum acquisition time was 10 s. 2.1. Materials. The following chemicals were used in this study: 1,2-dimethylimidazole (Alfa Aesar), diethyl sulfate (Fluka), LiTFSI (Solvionic), 1,4-dibromobutane, á,á′-dibromo-o-xylene, á,á′-dibromo-m-xylene, 1-methylpyrazole, 1,8dibromooctane, 1,1,1-tris(hydroxymethyl)ethane, methanesulfonic anhydride, tris(chloroethyl)amine hydrochloride, absolute ethanol, acetic acid, and acetonitrile. Except where stated otherwise, the chemicals were supplied by Aldrich. 2.1.1. Preparation of 3-Ethyl-1,2-dimethyl-1H-imidazol-3ium Bis-(trifluoromethanesulfonylmidate) (MonoTFSI). A solution of 1,2-dimethylimidazole (20 mmol, 1.92 g) in acetonitrile was added to diethyl sulfate (20 mmol, 3.08 g). The mixture was stirred for 2 h at room temperature. After the solvent, acetonitrile, was extracted from the mixture, the remaining white solid was dissolved in distilled water (20 mL) and added to a solution of LiTFSI (20 mmol, 5.74 g) in distilled water (20 mL). The ionic liquid was obtained after evaporation of the water. δH (300 MHz; DMSO-d6) 1.39 (3H, t), 2.61 (3H, s), 3.79 (3H, s), 4.17 (2H, quin), 7.59 (1H, d), 7.63 (1H, d); δC (75 MHz; DMSO-d6) 9.2 (CH3), 14.9 (CH3), 34.8 (NCH3), 43.3 (CH2), 113.5 (CF3), 117.8 (CF3), 122.0 (CF3), 126.3 (CF3), 120.6 (CH), 122.7 (CH), 144.4 (N2C). 2.1.2. Preparation of Octane-1,8-diylbis(1,2-dimethyl-1Himidazol-3-ium) Di(bis-(trifluoromethanesulfonylmidate)) (C8TFSI). A mixture of 1,8-dibromooctane (10 mmol, 2.72 g), 1,2-dimethylimidazole (20 mmol, 1.92 g), and acetonitrile (20 mL) was heated to 60 °C with stirring for 1 week. The resulting white precipitate was collected by centrifugation, washed with large amount of acetonitrile, and dried. The white powder was dissolved in distilled water (20 mL) and added to a solution of LiTFSI (20 mmol, 5.74 g) in distilled water (20 mL). The resulting white solid was recrystallized from boiling water to obtain the title compound. mp 77.8 °C; δH(300 MHz; DMSOd6) 1.31 (8H, m), 1.72 (4H, quin), 2.59 (6H, s), 3.77 (6H, s), 4.11 (4H, t), 7.63 (4H, d); δC(75 MHz; DMSO-d6) 9.5 (CH3), 26.0 (NCH3), 28.8 (CH2), 29.5 (CH2), 35.0 (CH2), 47.9 (CH2), 113.5 (CF3), 117.8 (CF3), 122.0 (CF3), 126.3 (CF3), 121.2 (CH), 122.7 (CH), 144.6 (N2C).19 2.1.3. Preparation of Butane-1,4-diyl-bis-(4,4′(dimethylamino)pyridine-ium Di(bis-(trifluoromethanesulfonylmidate)) (DMAPTFSI). DMAPTFSI was synthesized from starting compounds 1,4-dibromobutane (10 mmol, 2.16 g) and 4-(dimethylamino) pyridine (20 mmol, 2.44 g) by the same procedure as C8TFSI. mp 157.9 °C; δH (300 MHz; DMSO-d6) 1.76 (4H, quin), 3.20 (12H, s), 4.21 (4H, t), 7.04 (4H, d), 8.28 (4H, d); δC (75 MHz; DMSO-d6) 27.2 (NCH3), 39.8 (CH2), 56.4 (CH2), 108.2 (CH), 113.5 (CF3), 117.8 (CF3), 122.0 (CF3), 126.3 (CF3), 142.2 (CH), 156.2 (C). 2.1.4. Preparation of o-Xylene-α,α′-diylbis(1,2-dimethyl1H-imidazol-3-ium) Di(bis-(trifluoromethanesulfonylmi-

3. RESULTS AND DISCUSSION 3.1. Solubility Enhancement of Alkaline Peroxides and Fluorides. The solubilities of Li2O2, Na2O2, NaF, and KF were tested with complexing-cations in DMSO. Figure 1a,b shows the poor solubility of both Li2O2 and Na2O2 in DMSO. When the MonoTFSI was added to the solution, the solubility increased from 0.79 to 2.08 mM for Li2O2 and from 2.41 to 58.90 mM for Na2O2, as a result of the formation of an ion pair between the MonoTFSI and O22−. The much higher solubility of Na2O2 compared with Li2O2 is due, at least in part, to the lower lattice energy of Na2O2 (2309 kJ mol−1 vs Li2O2 2592 kJ mol−1).20 Upon linking these cations in various ways to form complexing di- and trications, the solubility increased further and in the order DMAPTFSI < Di-PyTFSI < 1,3-Di-ImTFSI < C8TFSI < Di-ImTFSI. Di-ImTFSI dissolved 3.38 mM of Li2O2 and 89.96 mM of Na2O2, while Tri-ImTFSI promoted a four3395

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strong interaction with reduced O2. In the following, we discuss the effect of different complexing-cations on the electrochemistry of O2 reduction, starting with the first electron reduction of O2, followed by the second reduction and finally exploring the effect of adding Li+ to the system. 3.2.1. O2 Reduction to Superoxide. Voltammograms were recorded at an Au electrode in O2-saturated DMSO with the various complexing-cations (Figure 2a). The complexing-

Figure 2. Cyclic voltammograms recorded in O2 saturated DMSO with 0.1 M of various complexing-cations (a) at a 1.6 mm diameter Au disk electrode at 100 mV s−1 and (b) at a 100 μm diameter Au disk electrode at 20 mV s−1.

cations based on pyridinium (DMAPTFSI) and pyrazolium (Di-PyTFSI) cations are not stable against superoxide, demonstrated by the absence of an anodic current in the CVs for O2 reduction in the presence of these complexingcations (Figure S10 in the Supporting Information), which is consistent with an irreversible following chemical reaction. These complexing-cations (DMAPTFSI and Di-PyTFSI) were therefore not investigated further in the electrochemical experiments. In the case of all imidazolium complexing-cations, a cathodic peak corresponding to the reduction of O2 to O2•− and an anodic peak corresponding to its oxidation to O2 were observed, confirming that the species O2•− was stable in the presence of the complexing-cations. This was further corroborated by overlapping multiple CV scans (Figure S9 in the Supporting Information). Notably, the structure and charge of the cations had an effect on the formal potential, E0′, which, in general, shifted to more positive potentials in the presence of the di- and trications (Figure 2a). However, it should be noted that the relatively small magnitude of the shift (Table 2) was close to the error obtained when correcting the potential to the internal ferrocene reference, making the trend difficult to quantify. Nonetheless, this trend mirrored that seen in the solubility enhancement of the alkaline peroxides, indicating that both effects are due to the strength of interaction between the complexing-cation and reduced O2 species. We suggest that both the di- and trication can provide a positively charged

Figure 1. Solubility of (a) Li2O2, (b) Na2O2 in DMSO, and (c) NaF (black markers) and KF (red markers) with and without 0.1 M complexing-cations.

fold increase in Li2O2 solubility and a forty-fold increase in Na2O2. The trication is more effective because of the higher density of positive charges. Di-ImTFSI appears to be the best among dications because of the geometry of the structure, as the rigid benzene ring forces the two imidazolium rings together in a way that favors complexation of O22−. To confirm the effect of complexing-cations, we also measured the solubility of NaF and KF in DMSO. As expected, Di-ImTFSI increased the concentration of both NaF and KF, two-fold to 5.22 and 4.64 mM, respectively, as shown in Figure 1c. The stability of all complexing-cations toward both peroxides was confirmed, as shown in the Supporting Information. Overall, the solubility studies corroborate the validity of using complexing-cations formed by a delocalized π-system to complex O22−. 3.2. Influence of Complexing-Cations on the O2 Reduction Reaction. The previously described results are encouraging and indicate that the complexing-cations have a 3396

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Table 2. E0′ of the O2•−/O2 Couple and k° for the 1-Electron Reduction of O2 in DMSO in the Presence of Different Complexing-Cations (0.1 M) electrolytes

E0′/V vs Li+/Li

103 k°/cm s−1

TBATFSI MonoTFSI C8TFSI 1,3-Di-ImTFSI Di-ImTFSI Tri-ImTFSI

2.37 2.36 2.37 2.37 2.43 2.46

2.4 5.0 8.0 8.4 8.6 15

pocket that can complex O2•−, and this enhanced stability results in the shift of E0′ to positive potentials. As well as a change in the energetics of O2 reduction, due to complexation of the O2•−, it is clear from the voltammograms in Figure 2a that the peak-to-peak separation decreased as the cation charge increased, indicating an enhancement in the rate of electron transfer. To explore this, we obtained the standard heterogeneous rate constants, k°, for O2 in the presence of the various complexing-cations using the method of Nicholson and Shain.21,22 (See the Supporting Information.) The resulting rate constants are shown in Table 2, and the highest k° of 1.5 × 10−2 cm s−1 was observed in the presence of Tri-ImTFSI, which also gave the highest solubility enhancement of Li2O2 and Na2O2. This value is ∼six-fold higher than the k° of 2.4 × 10−3 cm s−1 obtained in a solution of TBATFSI in DMSO. The other complexing-cations also showed an enhancement of the k° for O2 reduction. It is important to note that simply increasing the concentration of TBATFSI does not result in a similar effect on either k° or E0′, excluding the possibility that these data are due to an increase in electrolyte strength alone. Voltammograms of O2 reduction were also recorded at an UME, and these are shown in Figure 2b. When using TBATFSI as the supporting electrolyte, a typical steady-state CV was observed. However, in the solutions that contained the larger complexing-cations, the bulky complexing-cation/O2•− ion-pair diffusion was sufficiently slow that a peak appeared in the reverse sweep.23−27 Importantly, this is good evidence of the interaction of O2•− and the complexing-cations. The slopes of the voltammograms increase with the complexing strength of the cations consistent with the increase in the standard heterogeneous rate constants for the O2 redox process found previously. In summary, these data show that the introduction of a complexing-cation enhances the kinetics for O2 reduction in an organic electrolyte and also appears to have some effect on the formal potential. 3.2.2. Superoxide Reduction to Peroxide. Complexingcations not only increase the ko for the reduction of O2 to superoxide, they also promote the reduction of superoxide to peroxide in DMSO. Figure 3a shows voltammograms for O2 reduction in the presence of various concentrations of DiImTFSI. In the solution containing only TBATFSI, there is no reduction wave from O2•− to O22− until 1.5 V. However, in the presence of Di-ImTFSI, the data show a second reduction wave that shifts to more positive potentials as the cation concentration increases. Thus, the data indicate that Di-ImTFSI has a marked effect on both the k° and E0′ for the redox couple O2•−/O22−. Although the 1-electron reduction of O2 to superoxide has a classical shape, the superoxide reduction to peroxide is different; considerable hysteresis is observed during the reverse sweep, and this is likely due to fouling of the electrode by reduced O2 species. A similar response was

Figure 3. Cyclic voltammograms recorded at a 100 μm diameter Au disk electrode at 20 mV s−1 in O2-saturated DMSO with (a) various concentrations of Di-ImTFSI and (b) 0.1 M of various complexingcations.

observed when performing O2 reduction in the presence of the other complexing-cations (Figure 3b). Notably, Tri-ImTFSI appeared to provide the largest enhancement to the kinetics for the reduction of superoxide to peroxide, as indicated by the positive shift of voltammograms shown in Figure 3b. These results can be explained by the pocket of positive charge formed by the complexing di- and trications stabilizing O22− effectively. 3.2.3. O2 Reduction in the Presence of Di-ImTFSI and Li+. Having shown the effects of complexing-cations on O2 reduction without Li+, we now turn to the behavior in the presence of Li+. Voltammograms were collected for different concentration ratios of Di-ImTFSI and LiTFSI in DMSO (The total electrolyte concentration was 100 mM). We selected DiImTFSI to study because it has shown the best complexation effect on superoxide among the complexing dications. The reduction of O2 to O2•− in the presence of Li+, when restricted to a relatively positive cutoff potential on reduction, shows a well-defined oxidation peak, consistent with a 1electron redox process (Figure 4a). When performing the reverse scan to 4 V, no peak was observed for the oxidation of Li2O2 (Figure S8 in the Supporting Information). As the ratio of Di-ImTFSI concentration was increased, the O2 reduction peak shifted to more positive potentials, indicating that the kinetics of O2 reduction to superoxide were enhanced. Because of the restricted potential limit, it is not possible to apply the method of Nicholson and Shain to obtain values for k° for O2 reduction in the presence of Li+. Here we used the simulation package DigiSim (Supporting Information for details) to estimate the relative effect of Di-ImTFSI on O2 reduction kinetics in the presence of Li+. The resulting fits to experimental voltammograms with and without Di-ImTFSI are shown in Figure 5d,c. For completeness, we also include fits to O2 reduction voltammograms recorded with TBATFSI 3397

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Table 3. k° for the 1-Electron Reduction of O2 in DMSO in the Presence of Different Complexing-Cations (0.1 M) Obtained from the Fits Shown in Figure 5 electrolytes

103 × k°/cm s−1

TBATFSI LiTFSI Di-ImTFSI 50 mM LiTFSI + 50 mM Di-ImTFSI

2.9 2.5 6.9 6.0

previously (Table 2). A k° value of 2.5 × 10−3 cm s−1 was obtained for the 1-electron reduction of O2 in the presence of Li+, and this value is close to that obtained using TBATFSI (k° = 2.9 × 10−3 cm s−1), indicating that Li+ has little effect on this first electron transfer. Upon the addition of Di-ImTFSI, an increase in k° occurs to 6.0 × 10−3 cm s−1. This is consistent with the data obtained with Di-ImTFSI alone (k° = 6.9 × 10−3 cm s−1) and the hypothesis that the complexing effect of DiImTFSI enhances O2 reduction. It may seem surprising that DiImTFSI has a larger effect on the voltammetry than the Li+ it is replacing, as Li+ is a strong electrophile and hence expected to complex O2•− more effectively than Di-ImTFSI. However, it has been shown that the donor number of DMSO is sufficiently high such that Li+ is strongly solvated and, as a result, behaves as a relatively soft Lewis acid,28 sufficiently so that the DiImTFSI can compete for the attention of the O2•− and thus dominate the electrochemistry. In the solution containing LiTFSI only, extending the electrochemistry to lower potentials reveals the second reduction process to form peroxide (Figure 4b), as noted by others for solutions containing Li+ ions.29,30 As the LiTFSI was replaced by Di-ImTFSI, the second reduction peak shifted to negative potentials as the concentration of Di-ImTFSI was increased. This is unsurprising because the data in Figure 3a show that complexed O2•− reduction occurs at ∼2 V.

Figure 4. Cyclic voltammograms recorded at a 1.6 mm diameter Au disk electrode at 100 mV s−1 (a) at higher cutoff and (b) at lower cutoff, in O2 saturated DMSO with various concentrations of DiImTFSI in the presence of LiTFSI. The total electrolyte concentration was 100 mM in all cases, and the numbers on the plot indicate the concentration of Di-ImTFSI, where the remaining concentrations are LiTFSI.

(Figure 5a) and Di-ImTFSI (Figure 5b) only. For the latter two, k° values (Table 3) were consistent with those obtained

Figure 5. (Black line) Cyclic voltammograms recorded at a 1.6 mm diameter Au disk electrode in O2-saturated DMSO at 100 mV s−1 with (a) 0.1 M TBATFSI, (b) 0.1 M Di-ImTFSI, (c) 0.1 M LiTFSI, and (d) 50 mM LiTFSI + 50 mM Di-ImTFSI. (Red line) Fits to the experimental data. 3398

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Therefore, the shift of the second reduction peak as the ratio of Di-ImTFSI and LiTFSI is changed is a result of the changing equilibrium between complexed O2•− and LiO2 (and possibly the products Li2O2 and complexed O22−) and their different standard reduction potentials. Clearly it would be advantageous to identify the species involved in this system, and below we explore this using in situ SERS. 3.3. Spectroelectrochemical Analysis of O2 Reduction in the Presence of Di-ImTFSI and Li+. To confirm the reduction products of O2 in DMSO in the presence of DiImTFSI in our electrochemical studies, in situ SERS (Figure 6a) was carried out at various potentials on reduction and

Di‐ImTFSI‐O2·− + e− → Di‐ImTFSI‐O2 2 −

(2)

O2 + Di‐ImTFSI + e− → Di‐ImTFSI‐O2·−

(1)

Di‐ImTFSI‐O2·− + e− → Di‐ImTFSI‐O2 2 −

(2)

or

If reaction 1 does occur in two steps 1a and 1b, then 1b is likely to occur very quickly after 1a. In situ SERS obtained during O2 reduction in the presence of Li+ and Di-ImTFSI in DMSO is shown in Figure 6b, and the corresponding voltammograms are shown in Figure 4b (50 mM LiTFSI + 50 mM Di-ImTFSI). At 2.5 V, peaks appeared at 1160 and 1216 cm−1; the same peaks were observed in Figure 6a and belong to the complexed O2•− and rotation of the C−N bond, respectively. Notably, we did not observe a peak for LiO2 at any potential, and this may be because the peak of LiO2 (at ∼1137 cm−1)32,33 is covered by the broad peak at 1160 cm−1, or it may be that LiO2 is in solution and therefore not detectable. At 2.3 V, there is a new peak at ∼800 cm −1 corresponding to Li2O2,34 which may be formed from the following routes: 2LiO2 → Li 2O2 + O2

(3)

LiO2 + Li+ + e− → Li 2O2

(4)

Di‐ImTFSI‐O2·− + 2Li+ + e− → Li 2O2 + Di‐ImTFSI (5) Figure 6. In situ SERS during O2 reduction and reoxidation on a roughened Au electrode in O2-saturated (a) 100 mM Di-ImTFSI in DMSO and (b) 50 mM Di-ImTFSI + 50 mM LiTFSI in DMSO. Spectra were collected at reduction potentials of 2.5, 2.3, and 2.0 V versus Li+/Li, followed by an oxidation potential of 3.6 V. (1) C−N bond at 1216 cm−1, (2) O−O stretch of complexed O2•− at 1160 cm−1, (3) O−O stretch of complexed O22− at 832 cm−1, and (4) O−O stretch of Li2O2 at ∼800 cm−1.

Significantly, the peak associated with the O−O stretch of complexed O22− at 832 cm−1 was not observed, indicating that in the presence of Di-ImTFSI and Li+ the product of O2 reduction at 2.0 V is Li2O2.

4. CONCLUSIONS Several complexing-cations have been synthesized, consisting of two to three cations tethered together to form a positively charged pocket. The resulting cations improve the solubility of Li2O2 and Na2O2 in DMSO considerably, with the trication, 2,2′,2″-tris((1,2-dimethyl-1H-imidazol-3-ium)-ethyl)amine tri(bis-(trifluoromethanesulfonylmidate)) (Tri-ImTFSI), having the strongest effect and providing a four-fold increase in Li2O2 solubility and a forty-fold increase in Na2O2 solubility. The trend in complexing strength toward reduced O2 species as observed during alkaline peroxide solubility is mirrored by the effect observed during electrochemical reduction of O2 to O2•−, where the kinetics and equilibrium potential were affected by the presence of the complexing-cations and the rate constant increased by up to six times. Notably, this is the first demonstration of an appreciable enhancement of O2 reduction kinetics in an organic electrolyte in the absence of heterogeneous catalysis. The complexing-cations also promote the further reduction of complexed O2•− to complexed O22−, which benefits from the “positively charged pocket” provided by the di- and trications that stabilize the reduced O2 product. When O2 reduction is performed in the presence of both complexing-cations and Li+ in DMSO, a similar enhancement of the 1-electron reduction rate of O2 to O2•− occurs; furthermore, the product of the second reduction remains Li2O2. We believe that exploring the effect of similar compounds on O2 reduction is an appealing target for future research that may also benefit Li-air technology.

oxidation, corresponding to the voltammograms in Figure 3a (100 mM Di-ImTFSI). At the open circuit potential (OCP), a background spectrum was collected. DMSO shows bands due to C−S stretching at 669 and 699 cm−1 and at 1058 cm−1 from SO stretching and 953 cm−1 from a CH3 rock.31 On O2 reduction at 2.5 V, a peak at 1160 cm−1 appears that corresponds to O−O stretching of complexed O2•−. We propose that in the presence of O2•−, the two imidazolium rings form a “cation pocket” that complexes the O2•−. This may lead to rotation of the C−N bond resulting in the peak at 1216 cm−1. Such changes in the spectrum confirm the solvation of superoxide by Di-ImTFSI. A peak at 832 cm−1 associated with O−O stretch of complexed O22− appeared when the complexed O2•− was further reduced at 2.0 V. The O2•− stretching peak at 1160 cm−1 still remains, which indicates that complexed O2•− is only partially reduced to complexed O22− at 2 V. On oxidation, when the potential reached 3.6 V, complexed O2•− and O22− were oxidized to O2, which results in the disappearance of the peaks at 832, 1160, and 1216 cm−1. These results give directly spectroscopic evidence that ORR in the presence of Di-ImTFSI in DMSO are

O2 + e− → O2·−

(1a)

O2·− + Di‐ImTFSI → Di‐ImTFSI‐O2·−

(1b) 3399

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ASSOCIATED CONTENT

S Supporting Information *

Calculation of diffusion coefficient and concentration of O2, calculation of heterogeneous electron transfer rate constant, calculation of heterogeneous electron transfer rate constant using DigiSim, wider range of Figure 4a, stability of complexing-cations in the presence of superoxide and peroxide, and photograph of solutions prepared to test the solubility of Na2O2 by AAS. 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]. Phone: +44-1334463825. Fax: +44-1334-463808. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS P.G.B. acknowledges ALISTORE, the European Network of Excellence on Lithium Batteries, and the EPSRC, including the SUPERGEN program. S.A.F. acknowledges financial support by the Austrian Federal Ministry of Economy, Family and Youth and the Austrian National Foundation for Research, Technology and Development.

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