suementary inrmatin · 2014-11-20 · 1 eastchem, school of chemistry, university of st andrews,...

NATURE CHEMISTRY | www.nature.com/naturechemistry 1

SUPPLEMENTARY INFORMATIONDOI: 10.1038/NCHEM.2101

1

A unified mechanism of O2 reduction in aprotic Li+-‐ electrolytes and its consequences for Li-‐O2 batteries

Lee Johnson1,6‡, Chunmei Li1,2‡, Zheng Liu1,6, Yuhui Chen1,6, Stefan A. Freunberger3, Praveen C.

Ashok4, Bavishna B. Praveen4, Kishan Dholakia4, Jean-Marie Tarascon5, Peter G. Bruce6*

1 EastChem, School of Chemistry, University of St Andrews, North Haugh, St Andrews,

Fife, KY16 9ST, UK

2 Laboratoire de Réactivité et Chimie des Solides — UMR CNRS 6007, 33 rue Saint-Leu, 80039 Amiens Cedex,

France

3 Christian Doppler Laboratory for Lithium Batteries, and Institute for Chemistry and Technology of Materials,

Graz University of Technology, Stremayrgasse 9, 8010 Graz, Austria

4 SUPA, School of Physics & Astronomy, University of St Andrews, North Haugh, St Andrews, Fife, KY16 9SS, UK

5 Collège de France, 11 place Marcelin Berthelot, 75231 Paris Cedex, France

6 Departments of Materials and Chemistry, University of Oxford, UK

Table of Contents

Materials and methods .......................................................................................................................................... 2 Stability of electrolyte solutions ............................................................................................................................. 3 Determination of donor number using NMR ......................................................................................................... 3 Figure S1 ................................................................................................................................................................. 4 Table S1 .................................................................................................................................................................. 4 Simulation of voltammograms of O2 reduction obtained in DMSO and Me-‐Im ..................................................... 4 Figure S2 ................................................................................................................................................................. 5 Table S2 .................................................................................................................................................................. 6 Figure S3 ................................................................................................................................................................. 6 Shifts in the E0 for O2/O2

-‐ on replacing TBA+ with Li+ in Me-‐Im and DMSO ............................................................ 6 UV-‐vis analysis of O2

-‐ disproportionation ............................................................................................................... 7 Figure S4 ................................................................................................................................................................. 7 Figure S5 ................................................................................................................................................................. 8 Figure S6 ................................................................................................................................................................. 8 Hydrodynamic voltammograms for O2 reduction in the presence of Li

+ ................................................................ 8 Figure S7 ................................................................................................................................................................. 9 Estimating the free energy of equation (1) ............................................................................................................ 9 Calculation of capacity.......................................................................................................................................... 10 Influence of the solvent on the morphologies of Li2O2 at high and low voltages ................................................ 10 References ............................................................................................................................................................ 12

1

The role of LiO2 solubility in O2 reduction in aprotic solvents

and its consequences for Li–O2 batteries Lee Johnson1,6‡, Chunmei Li1,2‡, Zheng Liu1,6, Yuhui Chen1,6, Stefan A. Freunberger3, Praveen C.



Fife, KY16 9ST, UK


France






Table of Contents

Materials and methods .......................................................................................................................................... 2 Stability of electrolyte solutions ............................................................................................................................. 3 Determination of donor number using NMR ......................................................................................................... 3 Figure S1 ................................................................................................................................................................ . 4 Table S1 .................................................................................................................................................................. 4 Simulation of voltammograms of O2 reduction obtained in DMSO and Me-‐Im ..................................................... 4 Figure S2 ................................................................................................................................................................ . 5 Table S2 .................................................................................................................................................................. 6 Figure S3 ................................................................................................................................................................ . 6 Shifts in the E0 for O2/O2


-‐ disproportionation ............................................................................................................... 7 Figure S4 ................................................................................................................................................................ . 7 Figure S5 ................................................................................................................................................................ . 8 Figure S6 ................................................................................................................................................................ . 8 Hydrodynamic voltammograms for O2 reduction in the presence of Li

+ ................................................................ 8 Figure S7 ................................................................................................................................................................ . 9 Estimating the free energy of equation (1) ............................................................................................................ 9 Calculation of capacity.......................................................................................................................................... 10 Influence of the solvent on the morphologies of Li2O2 at high and low voltages ................................................ 10 References ............................................................................................................................................................ 12

1

The role of LiO2 solubility in O2 reduction in aprotic solvents

and its consequences for Li–O2 batteries Lee Johnson1,6‡, Chunmei Li1,2‡, Zheng Liu1,6, Yuhui Chen1,6, Stefan A. Freunberger3, Praveen C.



Fife, KY16 9ST, UK


France






Table of Contents

Materials and methods .......................................................................................................................................... 2 Stability of electrolyte solutions ............................................................................................................................. 3 Determination of donor number using NMR ......................................................................................................... 3 Figure S1 ................................................................................................................................................................ . 4 Table S1 .................................................................................................................................................................. 4 Simulation of voltammograms of O2 reduction obtained in DMSO and Me-‐Im ..................................................... 4 Figure S2 ................................................................................................................................................................ . 5 Table S2 .................................................................................................................................................................. 6 Figure S3 ................................................................................................................................................................ . 6 Shifts in the E0 for O2/O2


-‐ disproportionation ............................................................................................................... 7 Figure S4 ................................................................................................................................................................ . 7 Figure S5 ................................................................................................................................................................ . 8 Figure S6 ................................................................................................................................................................ . 8 Hydrodynamic voltammograms for O2 reduction in the presence of Li

+ ................................................................ 8 Figure S7 ................................................................................................................................................................ . 9 Estimating the free energy of equation (1) ............................................................................................................ 9 Calculation of capacity.......................................................................................................................................... 10 Influence of the solvent on the morphologies of Li2O2 at high and low voltages ................................................ 10 References ............................................................................................................................................................ 12

© 2014 Macmillan Publishers Limited. All rights reserved.



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Materials and methods

Me-‐Im was distilled under vacuum. DMSO was distilled under vacuum over NaNH2. CH3CN and DME were distilled under argon over CaH2. All solvents were further dried for several days over freshly activated molecular sieves (type 4 Å). The final water content was ≤ 4 ppm (determined by Karl Fischer titration). Battery grade tetrabutylammonium perchlorate (TBAClO4) and battery grade lithium perchlorate (LiClO4) were dried as described previously1. High purity N5 O2 was obtained from BOC industrial gasses and was used in all measurements. All materials were stored in an argon glove box.

Electrochemical measurements were performed using a VMP3 electrochemical workstation (Biologic) and a multi-‐necked, air-‐tight glass cell within a glove box. Electrochemical measurements were carried out at room temperature and IR correction was used. 2 mm diameter polycrystalline Au disks (BAS Inc.) were employed as the working electrodes. Prior to use the working electrodes were polished in a glove box with 0.05 μm alumina slurry in ethanol and rinsed with copious amounts of ethanol followed by drying under vacuum. A platinum wire served as the counter electrode. Measurements were performed using a reference electrode based on LiFePO4. It was first pre-‐oxidized (20% of total capacity) to LixFePO4, which has a constant, potential. This is then placed in 0.1 M LiClO4 in DMSO separated from the bulk electrolyte with a porous frit, thus providing a fixed potential. The LixFePO4 reference was corrected to Li/Li+ by subtracting 3.45 V. Electrolytes were saturated with O2. Rotating ring/disk electrode, RRDE, measurements were performed using a MSR rotator and a removable disk rotating ring/disk electrode containing an Au disk (5 mm diameter) and a glassy carbon ring (Pine instruments). Measurements were performed in a glove box in a round bottom flask and all necks other than the neck containing the shaft were sealed. The electrode was polished as described above between measurements. The collection efficiency at the ring, N, in each experiment was determined using an ideal redox couple. For the polarization measurements in Fig. 6, discharge at a planar 2 mm Au working electrode was performed at a constant current of 60 µA cm-‐2 in O2 saturated 100 mM LiClO4 solutions of the various aprotic solvent.

In situ surface enhanced Raman spectroscopy (SERS) was carried out in an airtight 3-‐electrode electrochemical cell. A previously electrochemically roughened Au working electrode (see reference for experimental procedure)2 was placed behind a 1 mm thick sapphire window. Raman spectra were recorded using a custom-‐built confocal Raman microscope with a 50x 0.45 NA objective (Nikon). The Raman system was equipped with a 785 nm diode laser (Sacher Lasertechnik) and a spectrometer (Shamrock SR-‐303i, Andor Technology) which employed a 400 lines/mm grating, blazed at 850 nm and was equipped with a deep depletion, back-‐illuminated and thermoelectrically cooled CCD camera (Newton, Andor Technology) for the detection of Raman signal.




3

Porous composite cathodes for the study of solvent dependent Li2O2 morphologies were prepared using carbon (Super P, Timcal Ltd) and 10 %w.t. PTFE, cast onto a stainless steel mesh. Discharge of the cathodes was carried out at constant voltage in a three electrode cell. LixFePO4 was again used as the reference electrode. The morphology of the pristine carbon electrode and discharged samples were observed by FE-‐SEM using a JSM-‐6700F (JEOL).

Stability of electrolyte solutions

Parasitic side-‐reactions at the electrolyte/cathode interface in Li-‐O2 cells have been reported widely. In parallel, many authors have reported successful fundamental electrochemical studies in the same or similar electrolyte solutions3,4. Some of the side-‐reactions are attributable to decomposition of the carbon cathode, used in most studies of Li-‐O2 cells, but not in the fundamental studies reported here, which use gold. Furthermore, the short timescales of electrochemical studies (e.g. CV) compared with investigations of side-‐reactions on discharge/charge in Li-‐O2 cells, means that the quantity of side reaction products is much less in the electrochemical studies. This is reinforced by the fact that the electrochemical data presented in the paper are well fitted by O2 reduction without any corrections for side-‐reactions and no evidence of side-‐reaction products were detected in SERS.

Determination of donor number using NMR

As the DN of Me-‐Im was unknown, 23Na NMR was used to estimate this value. Here we exploit the linear trend in DN against 23Na NMR shift and extract the DN of Me-‐Im from this trend, Supplementary Fig. S15. 23Na NMR was carried out in a 10 mM solution of NaTFSI in a range of aprotic solvents (listed in Table S1). Using these data the DN of Me-‐Im was estimated to be 47.




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Figure S1 | Plot of 23Na NMR shift versus solvent DN of NaTFSI dissolved in various aprotic solvents. The solvents used are listed in Table S1.

Table S1 | Solvent DN and the corresponding 23Na NMR shift in each solvent.

DN 23Na NMR shift CH3CN 14 -‐6.92 DME 20 -‐5.89 DMSO 30 -‐0.14 Pyridine 33 1.24 Me-‐Im 47 8.84

1,2-‐diaminoethane 55 12.53

Simulation of voltammograms of O2 reduction obtained in DMSO and Me-‐Im

BASi DigiSim Simulation Software was used to generate fits to the experimental voltammograms of O2 reduction in DMSO and Me-‐Im. CVs obtained at 100 mV s-‐1 were fit using a single E step and CVs obtained at 5 mV s-‐1 were fit using an EC mechanism, where the C step is irreversible. The resulting fits to experimental data for 100 mM TBA+ and 100 mM Li+ are shown in Supplementary Fig. S2 and the parameters obtained are given in Table S2. As E01 varies with the concentration of LiClO4 (Fig. 2), this parameter was allowed to change during simulation.




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Fits to experimental CVs obtained at 5 mV s-‐1 were carried out in order to observe the disproportionation of O2

-‐:

2O2-‐(sol) + 2Li+(sol) → Li2O2(solid) + O2(sol) (S1)

The standard heterogeneous rate constant, k0, and the O2 concentration, C, were maintained at the values found at higher scan rates. The CVs are shown in Supplementary Fig. S3 and the first order rate constants, k, for LiO2 disproportionation, equation (S1), in Me-‐Im and DMSO with 100 mM LiClO4 are 0.015 s-‐1 and 0.030 s-‐1, respectively.

Figure S2 | (circles) CVs and (red lines) fits for O2 reduction in Me-‐Im and DMSO at a Au electrode in 100 mM TBAClO4 and 100 mM LiClO4. CVs obtained at 100 mV s-‐1 and fits based on a single E step. The resulting variables obtained from the fits are given in Table S2.




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Table S2 | Variables obtained from fits to the CVs shown in Supplementary Fig. S2. D – diffusion coefficient and k0 –standard heterogeneous electron transfer rate constant for O2/O2

-‐. During fitting, the O2 concentration, C, was constant at 1.34 mM and the mean α used were 0.48 and 0.44 for DMSO and Me-‐Im, respectively.

Electrolyte 105 x D / cm2 s-‐1 103 x k0 / cm s-‐1 Me-‐Im

TBAClO4 9.49 2.72 LiClO4 10.4 2.75

DMSO TBAClO4 8.6 2.34

LiClO4 8.9 2.34

Figure S3 | (circles) CVs and (red lines) fits for O2 reduction in Me-‐Im and DMSO at a Au electrode in 100 mM LiClO4. CVs obtained at 5 mV s-‐1 and the fits are based on an EC mechanism.

Shifts in the E0 for O2/O2-‐ on replacing TBA+ with Li+ in Me-‐Im and DMSO, Fig 2

The shifts in the standard potential for O2/O2-‐, E01, on replacing TBA+ with Li+ in Me-‐Im and

DMSO, extracted from the CVs, are plotted in Fig. 2 in the main paper. For Me-‐Im the variation is small and linear, whereas for DMSO it is larger and curves upwards, indicating that the potential increases more rapidly at higher Li+ concentrations. Recognizing that the main difference between Me-‐Im and DMSO is that the latter has a lower DN, we interpret the different trends as indicating that in Me-‐Im (high DN), O2

-‐ interacts relatively weakly with Li+ cations, e.g. possibly via solvent separated ion pairs (Li+-‐S-‐O2

-‐); explaining why there is little difference in the interactions between O2

-‐ and TBA+ and those between O2-‐ and Li+




7

(very similar CVs from TBA+ to Li+, Fig. 2). Whereas in DMSO, solvent separated ion pairs at low Li+ concentrations increasingly give way to contact ion pairs (Li+-‐O2

-‐) or larger aggregates (clusters) as the concentration increases, resulting in a greater positive shift in potential. The existence of ion pairing and the trend towards contact ion pairs and aggregates with increasing salt concentration is well known in organic electrolytes6,7.

UV-‐vis analysis of O2-‐ disproportionation

In order to determine the rate of O2-‐ disproportionation, equation (S1), in aprotic solvents

containing Li+, the reaction was followed by UV-‐vis spectroscopy using a Lambda 35 spectrometer (PerkinElmer Instrument). Excess KO2 was added to the selected solvent and stirred overnight. This was then centrifuged and the clear KO2 saturated solution was decanted in an Ar filled glove box. 2.7 ml of KO2 solution was added to a sealed cuvette. 0.3 ml of 1 M LiClO4 solution was quickly injected into the cuvette to initiate the reaction and precipitate Li2O2; the absorbance, A, at 300 nm wavelength was recorded as a function of time. The analysis was performed in Me-‐Im and DMSO. However, due to an overlap between the adsorption peaks of O2

-‐ and Me-‐Im at 300 nm a rate constant could not be obtained in this solvent. The resulting plot of ln(A) against time is shown in Supplementary Fig. S4, consistent with a first-‐order reaction. Using these data the first order rate constant for the disproportionation of O2

-‐ in DMSO was calculated to be 0.07 s-‐1.

Figure S4 | UV-‐vis spectroscopy analysis of O2-‐ disproportionation in DMSO. UV-‐vis

adsorption by O2-‐ at 300 nm recorded using a saturated solution of KO2 in DMSO after

addition of 100 mM LiClO4. Circles – experimental points, solid line – linear least squares fit.




8

Figure S5 | SERS of superoxide during O2 reduction. In DMSO on O2 reduction the peak for adsorbed O2

-‐ in the 100 mM TBA+ solution is unchanged on replacing TBA+ by Li+. The Raman shifts for O2-‐ and LiO2 are 1110 cm-‐1 and 1132 cm-‐1, respectively.

Figure S6 | CVs obtained in O2 saturated Me-‐Im, DMSO and CH3CN at a Au electrode in 0.5 M LiClO4. The scan rate was 100 mV s-‐1. The 2nd reduction peaks have moved to higher potentials with higher Li+ concentration. The CVs collected with the higher voltage limit (black curves) show that O2

-‐ is still formed in solution even at these high Li+ concentrations in high DN solvents.

Hydrodynamic voltammograms for O2 reduction in the presence of Li+

RRDE measurements were performed at a polycrystalline Au disk in Me-‐Im, DMSO, DME and CH3CN and the resulting polarization curves for O2 reduction in 0.1 M LiClO4 solutions are shown in Fig. 4. Here we scan the potential of the disk and record the current while rotating the electrode to enhance mass transport. The ring is held at a constant potential such that any O2

-‐ present in solution is oxidized and the recorded current is proportional to the

amount of O2-‐(sol) formed during O2 reduction at the disk. As the 2nd reduction process forms

solid Li2O2 on or near the disk surface its oxidation cannot be observed at the ring. In the data shown, the ring current is corrected to account for the ratio of product formed at the disk to that collected at the ring using an ideal one electron redox mediator8. The black lines in Fig. 4 show the current recorded at the disk, iD and the green lines correspond to the




9

Figure S7 | Polarization curves obtained in O2 saturated Me-‐Im, DMSO, and CH3CN containing 0.5 M LiClO4. The RRDE was a 5 mm diameter Au disk with a GC ring and the rotation rate was 2000 RPM.

current detected at the ring iR (presented as a negative current for easy comparison to the disk current). Finally, the red lines correspond to the difference (iD -‐ iR), i.e., that which does not lead to O2

-‐ in solution; we assign this current to the surface reduction to form Li2O2, which is confirmed by in situ SERS. The RRDE results in both Me-‐Im and DMSO show no red area i.e. (iD – iR) at high potentials, and the peak current was dependent on the rotation rate of the RRDE. Both of these results are consistent with the reduction of O2 to O2

-‐ in solution and not with formation of Li2O2 on the disk surface. At lower voltages, the current ratio between the ring and disk began to drop below one (red area grows) and at slightly more negative potentials the current at the ring began to decrease, consistent with the onset of reduction of O2

-‐ to Li2O2* at the disk. This process also initiated a decrease in the current at the disk due to the Li2O2 blocking the electrode surface. It should be noted that RRDE is a steady state technique and that this decrease in current is not due to any diffusion limitation. In the lowest DN solvent, CH3CN, there is almost no ring current and the red area dominates indicating almost no O2

-‐ in solution. The disk current was relatively insensitive to rotation rate and the peak current was considerably lower than in the absence of Li+. These results indicate a surface confined process consistent with the formation of a Li2O2 film on the electrode. Finally, in DME properties of both the high and low DN systems were observed. The magnitude of the red area is between the high and low DN solvents, there is a significant ring current indicating the presence of O2

-‐ in solution, similar to high DN solvents, but at the same time the ring current never equals the disk current, even at high potentials. In addition, the current was noticeably lower than in the absence of Li+, confirming a surface process occurred at all potentials.

Estimating the free energy of equation (1)

Equation (2) is the sum of the two redox couples, O2/O2-‐ and Li/Li+, therefore, summing the

two E0 values for these couples and recalling ΔG0 = -‐nFE0, gives the Gibbs free energy for




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equation (2). The E0 values for the O2/O2-‐ couple were measured in each of the four solvents

in the presence of the TBA+ salt (i.e.TBAClO4 solutions). The E0 values for the Li/Li+ couple were measured in each of the four solvents in the presence of ClO4

-‐ ions (i.e. in LiClO4 solutions). The reference electrode used in each case is described in the Methods section and is the same as that used throughout this work. LixFePO4 was used instead of Li metal, because Li is not stable in every solvent. 3.45 V was subtracted to give the value for Li/Li+. The free energy of equation (3) was obtained from the pressure of O2(g) in the head space above the solution and the concentration of O2(sol) in the solvent, which are in equilibrium. Combining equations (2) and (3) with (4), for which the free energy is known (-‐221 kJ mol-‐1)9, gives an estimate for the free energy of equation (1). The approximations inherent in the estimation of ΔG0 for equation (1) and why, despite such approximations, they still predict the correct trend from soluble LiO2 in high DN to surface LiO2* in low DN solvents, is discussed in the paper.

The ΔG0 for DME is 25 kJ mol-‐1 so the equilibrium, equation (1), lies towards surface species, however a concentration of 0.4 mM for O2

-‐ is predicted from ΔG0 = -‐RTln[Li+(sol)][O2-‐(sol)]),

Li+ 0.1 M.

Calculation of capacity

We consider an electrode based on carbon composed of hexagonally close packed cylindrical pores, which we optimize for deposition by a surface and separately a solution mechanism. An optimized cathode structure for a surface mechanism (low DN solvents) has the highest possible surface area commensurate with pores of sufficient size to permit mass transport in the electrolyte within the pores. The maximum thickness of a Li2O2 layer formed on the surface has been reported as ∼7 nm10. Taking a pore diameter of 50 nm as a minimum requirement for mass transport and a wall thickness of 3 nm, this corresponds to a BET surface area of 170 m2 g-‐1. The resulting capacity for a 7 nm Li2O2 film is 2500 mAh g-‐1(C). An electrode optimized for the solution growth of Li2O2 (high DN solvent) requires the largest possible pore volume. Taking a pore diameter of 500 nm for the cylindrical pores and assuming 80 % of the porous volume can be filled with Li2O2 during discharge (the remaining 20 % contains the electrolyte), we obtain a specific capacity of 8500 mAh g-‐1(C), which is more than a threefold increase in capacity.

Influence of the solvent on the morphologies of Li2O2 at high and low voltages

For each of the solvents studied here the morphology of Li2O2 growth has been examined as a function of potential, Fig. 7. For such studies a cathode with a relatively large surface area and porosity was required, so that both solution and surface pathways can contribute and to allow sufficient Li2O2 to form so that the morphology of Li2O2 may be readily imaged by




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SEM. Recognizing that the vast majority of morphological studies have employed carbon, we used electrodes formed from a composite of Super P carbon and PTFE (see Methods in SI), which fulfill the criteria of a relatively large surface area and porosity. For the highest DN solvent (Me-‐Im), and at high potentials ≥ 2.5 V (high V in Fig. 7 b), large agglomerates of Li2O2 particles (∼200-‐300 nm) form between the carbon particles. At low potentials ∼2.2 V (low V in Fig. 7 c), i.e. corresponding to the second reduction peak in the CVs for Me-‐Im, there is little evidence of Li2O2 particles. These morphological observations are in accord with the mechanism we have proposed, as well as with the SERS and electrochemical data, i.e. particles at high V and a film on the electrode at low V (see e.g. SERS, Fig. 3, for Li2O2 film growth on the electrode surface). A similar behavior is observed in DMSO. We do not know the detailed mechanism of how dissolved LiO2, forms a Li2O2 precipitate in high DN solvents at high voltages. Presumably there is desolvation, formation of an activated complex, disproportionation and nucleation and growth. Computational studies have given valuable insight into this process in the gas phase, although of course conditions in solution are somewhat different11. Some of the O2

-‐ near the electrode surface will disproportionate and form Li2O2 likely using the electrode or existing Li2O2 surfaces as a convenient nucleation site. Most will diffuse away but in practical, porous, electrodes, will precipitate in the pores, because the rate of disproportionation is quite fast. The morphology in CH3CN is very different from high DN solvents at high voltage. There is no evidence of Li2O2 particles at any potential, Fig. 7 h and i, but SERS (Fig. 3) and the electrochemistry confirms Li2O2 is present at high and low potentials. In the case of DME, Li2O2 particles are observed in the pores at high potentials. Our electrochemical analysis and SERS, Fig. 3, indicates the simultaneous formation of a Li2O2 film at these potentials. Below this potential no Li2O2 particles are evident and SERS (Fig. 3) confirms a Li2O2 film on the electrode surface.

The variation in Li2O2 morphology with solvent type is in excellent agreement with the change in the pathway of O2 reduction from one dominated by solution soluble LiO2 in high DN solvents, (giving rise to Li2O2 particles in the pores of the electrode at high potentials) to the surface mechanism in low DN solvents (Li2O2 film on the electrode surface). These results are also in good accord with the polarization measurements shown in Fig. 6. The intermediate nature of the DN in DME, results in both solution and surface mechanisms of oxygen reduction occurring at the same potential, the result being the simultaneous formation of particles that fill the pores in the electrode as well as a film that coats the surface.




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suementary inrmatin · 2014-11-20 · 1 eastchem, school of chemistry, university of st andrews,...

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