Investigation of MnO₂ and Ordered Mesoporous Carbon Composites as Electrocatalysts for Li-O₂ Battery Applications.

The electrocatalytic activities of the MnO₂/C composites are examined in Li-O₂ cells as the cathode catalysts. Hierarchically mesoporous carbon-supported manganese oxide (MnO₂/C) composites are prepared using a combination of soft template and hydrothermal methods. The composites are characterized by X-ray powder diffraction, scanning electron microscopy, transmission electron microscopy, small angle X-ray scattering, The Brunauer-Emmett-Teller (BET) measurements, galvanostatic charge-discharge methods, and rotating ring-disk electrode (RRDE) measurements. The electrochemical tests indicate that the MnO₂/C composites have excellent catalytic activity towards oxygen reduction reactions (ORRs) due to the larger surface area of ordered mesoporous carbon and higher catalytic activity of MnO₂. The O₂ solubility, diffusion rates of O₂ and O₂•- coefficients (DO₂ and DO-₂), the rate constant (kf) for producing O₂•-, and the propylene carbonate (PC)-electrolyte decomposition rate constant (k) of the MnO₂/C material were measured by RRDE experiments in the 0.1 M TBAPF₆/PC electrolyte. The values of kf and k for MnO₂/C are 4.29 × 10-2 cm·s-1 and 2.6 s-1, respectively. The results indicate that the MnO₂/C cathode catalyst has higher electrocatalytic activity for the first step of ORR to produce O₂•- and achieves a faster PC-electrolyte decomposition rate.

1. Introduction

Energy storage devices with high energy and power densities <span class="Chemical">arn>e being developed for use as power sources for electric vehicles (EV) and hybrid electric vehicles (HEV) [1,2,3]. Over the past few decades, the vast majority of relevant rese<span class="Chemical">arch has focused on upgrading the performance of conventional <span class="Chemical">lithium-ion batteries for EV or HEV applications; however, their energy densities and specific charge capacities still fail to satisfy commercial requirements such as long-range driving, low cost, and fast charging [1,2,4]. In recent years, rechargeable nonaqueous Li-air batteries have attracted much interest owing to their low cost, environmental friendliness, and high theoretical energy density (~3500 Wh·kg−1), which is nearly equivalent to a nine-fold increase over conventional Li-ion batteries (~400 Wh·kg−1) [4,5,6,7]. Despite these favorable characteristics, their practical applications are still hampered by several serious challenges including limited rate capability, poor cycling stability due to the instability of the electrode and electrolyte, and low round-trip efficiency induced by the rather large polarization, resulting in a wide charge–discharge voltage gap [3,8,9,10,11,12,13,14,15]. These critical problems are highly attributable to the O2 cathode.

A typical rech<span class="Chemical">argeable <span class="Chemical">Li-O2 battery is constituted by a porous <span class="Chemical">oxygen diffusion cathode, a lithium metal anode, and an Li+-conducting electrolyte. In general, the O2 cathode is an oxygen catalyst loaded with porous carbon material, which enables both Li2O2 deposition (oxygen reduction reactions, ORRs) and decomposition (oxygen evolution reactions, OERs) reactions to occur upon battery discharge and charge, respectively. Many reports [1,4,5,6,8,9,11,15,16,17,18] have pointed out that the electrochemical performance of Li-O2 batteries depends on many factors such as: the nature and microstructure of the O2 electrode, electrolyte formula (especially, the composition of solvent), O2 partial pressure, possible presence of reactive contaminants (e.g., trace water), and the choice of catalysts. In order to enhance the properties of rechargeable Li-O2 batteries, several strategies have been followed over the years to explore the electrolyte formula, choice, and microstructure design of the O2 electrode and optimization of the operating parameters [1,3,5,8,9,10,11].

<span class="Chemical">Carbon materials with v<span class="Chemical">arious nanostructures have been developed and used as <span class="Chemical">O2 cathodes in Li-O2 batteries [4,6,10,19]. It has been well demonstrated that the performance of Li-O2 batteries is related to the properties of carbon, such as the morphology, surface area, porous structure, and conductivity [6,9,20]. The design of porous carbon cathodes requires larger intraparticulate voids and open frameworks in their architecture structure to accommodate the insoluble discharge products. These voids and frameworks should help improve discharge capacity and cycling performance [19,20,21]. Obviously, designing an optimum pore structure for carbon materials can effectively improve the electrochemical performance of Li-O2 batteries. Although various porous carbon structures have been explored, some studies have demonstrated that hierarchically porous honeycomb-like carbon cathodes with mesoporous/macroporous pore size can increase the specific capacity of Li-O2 batteries [4,5,6,15,19,20,21,22,23,24,25,26]. Moreover, it is well known that an ideal cathode catalyst can facilitate the complete reversibility of ORRs and OERs with low polarization in Li-O2 batteries [21]. Several potential catalysts have recently been proposed to promote ORRs and OERs, including nitrogen-doped carbon, metal oxides, metal nitrides, precious and nonprecious metals, etc. [1,3,8,13,15,19,27,28,29]. Among metal oxides, MnO2 is a catalyst material of great interest because of its low cost, environmental friendliness, abundance, and electrocatalytic activity for ORRs in Li-O2 batteries [13,28,30,31,32]. This study of Li-O2 batteries focuses on MnO2-based catalysts.

In the first p<span class="Chemical">art of this work, we created a hier<span class="Chemical">archically <span class="Chemical">mesoporous carbon-supported β-manganese oxide (MnO2/C) as an O2 cathode material. We present a detailed study of the Li-O2 electrochemistry of the MnO2/C material using an electrolyte of 1 M LiPF6 in a propylene carbonate (PC, which was used in many of the initial works on Li-O2 batteries) solvent. Although there have been many studies of MnO2/C materials for Li-O2 battery applications, few studies have examined the poor stability of the electrolyte due to its reaction with the superoxide radical (O2•−) produced upon the discharge at the MnO2/C electrode. In this work, the stability of the electrolyte against the O2•− of the MnO2/C electrode was first explored by the RRDE technique. The RRDE was developed about 50 years ago and has been verified to be a powerful tool for the study of electrochemical reactions. RRDE consists of two concentric electrodes (disk and ring electrodes) in a cylindrical holder with both of the electrodes facing downward into the solution. Products generated at the disk reaction are swept outward by the convection caused by rotation, and can be detected electrochemically at the ring by fixing the potential on the ring electrode. In this study, a disk electrode coated with MnO2/C materials and a Pt ring electrode was fixed at an O2•−/O2 oxidation potential to collect the O2•− ions in electrolytes. Therefore, in the second part, we emphasize aspects of the PC-based electrolyte reaction against O2•− and the related kinetic information of O2•− in the MnO2/C electrode by studying rotating ring disk electrode (RRDE) experiments and using a lithium-free non-aqueous electrolyte due to the stability of the intermediate O2•−. In addition, the oxygen solubility in the electrolyte and the oxygen diffusion velocity throughout the whole O2 electrode play key roles in determining battery performance, especially at high current densities [33]. In this work, the O2 solubility, diffusion rates of O2 and superoxide radical (O2•−) coefficients ( and ), rate constant (k) for producing O2•−, and PC-electrolyte decomposition rate constant (k) of the MnO2/C electrode were quantified.

2. Experimental Methods

<span class="Chemical">MnO2/C composites were prepared by supramolecul<span class="Chemical">ar self-assembly methods followed by a hydrothermal process. A modification of the mesoporous metal oxides and carbon nanocomposites procedure of Huang et al. [34] was applied to synthesize the MnO2/C composites. The first step was to synthesize a 20 wt. % resol ethanolic solution according to an established method [34,35]. A solution was prepared by dissolving 1.5 g of triblock copolymer Pluronic F127 (OH(CH2CH2O)-(CH2CH(CH3)O)-(CH2CH2O)H, EO106PO70EO106, Sigma Aldrich, St. Louis, MO, USA) in 10 g of anhydrous ethanol, then 5 g 20 wt. % resol ethanolic solution and 0.28 g MnCl2·4H2O (J.T. Baker, 99.8%) were added into the above solution slowly under stirring for 30 min at an ambient temperature. The homogeneous mixture was then transferred into a Petri dish at an ambient temperature for 24 h. After being dried, the films were heated at 100 °C for another 24 h to form orange transparent membranes. The as-made products were scraped from the Petri dish and ground into powders and then calcinated at 400 °C for 5 h under an Ar atmosphere with a heating rate of 1 °C·min−1 to yield Mn/C powders. To obtain MnO2/C composites, the as-prepared Mn/C powders were subjected to a hydrothermal process at 180 °C for 12 h with 0.22 g KMnO4 (J.T. Baker) and 30 mL of deionized water in a Teflon-lined stainless steel autoclave.

A Rigaku-D/MaX-2550 diffractometer (Rigaku, Tokyo, Japan) with Cu Kα radiation (λ = 1.54 Å) was used to obtain X-ray diffraction (XRD) patterns for the samples. Small angle X-ray scattering (SAXS) measurements were taken on a Nanost<n class="Chemical">span class="Chemical">ar U small-angle X-ray scattering system (Bruker, K<spn>an class="Chemical">arlsruhe, Germany) using Cu Kα radiation (40 kV, 35 mA). The morphology of the sample was observed using a scanning electron microscope (SEM, Hitachi S-3400 (Hitachi Limited, Tokyo, Japan)) and transmission electron microscope (TEM, JEOL JEM-3010 (JEOL, Tokyo, Japan)). Selected area electron diffraction (SAED) was applied to examine samples’ crystallinity. The Brunauer–Emmett–Teller (BET) method was used to measure the specific surface area of the powders (ASAP2020). The residual carbon content of the samples was measured by an automatic elemental analyzer (EA, Elementar vario, EL III (Elementar Analysensysteme GmbH, Hanau, Germany)).

For electrochemical evaluation, the <span class="Chemical">MnO2n>/C electrodes were prep<span class="Chemical">ared by wet coating, and were made from as-prep<span class="Chemical">ared MnO2/C composites with super P and a poly(vinylidene difluoride) (PVDF) binder (MKB-212C, Atofina, Serquigny, France) in a weight ratio of 64:16:20. The MnO2/C composites and super-P were first added to a solution of PVDF in N-methyl-2-pyrrolidone (NMP, Riedel-deHaen, Seelze, Germany). To make a slurry with an appropriate viscosity, the mixture was stirred for 20 min at room temperature using a magnetic bar, and then for 5 min using a turbine at 2000 rpm. The resulting slurry was coated onto a piece of separator (Celgard 2400, Charlotte, NC, USA) and dried at 60 °C under vacuum for 12 h. The coating had a thickness of ~100 μm with an active material mass loading of 8 ± 1 mg·cm−2. The quantity of active materials on the electrodes was kept constant. Electrodes were dried overnight at 100 °C under a vacuum before being transferred into an argon-filled glove box for cell assembly. The Li-O2 test cell (EQ-STC-LI-AIR, MTI Corporation, Richmond, CA, USA) was constructed with lithium metal as the negative electrode and the MnO2/C electrode as the positive electrode. A solution of 1 M LiPF6 in a PC solvent was used as the electrolyte in all cells. After assembly, the test cell was taken away from the Ar-filled glove box and attached to a gas pipe that was constantly purged with dry O2. Electrochemical tests were carried out after the cell was flushed with O2 for 6 h. The cells were cycled galvanostatically with a BAT-750B (Acu Tech System, Taipei, Taiwan) at a constant current of 100 mA/g with a voltage region of 2.0–4.3 V vs. Li/Li+ at room temperature.

For the RRDE experiments, the RRDE system (AFMT134DCPTT, Pine Rese<n class="Chemical">span class="Chemical">arch Instrumention, Durham, NC, USA) with interchangeable disk consisted of a 5 mm diameter glassy <spn>an class="Chemical">carbon electrode and a Pt ring electrode (1 mm width) with a 0.5 mm gap between them. The collection efficiency with this geometry is 0.24. The rotating ring-disk assembly was operated on a Pine AFMSRX rotator and CH705 Bipotentiostat (CH Instruments, Austin, TX, USA) with a computerized interface. Experiments were conducted using a three-electrode cell containing 10 mL of the electrolyte of interest and assembled in a dry Ar-filled AtmosBag (Sigma-Aldrich Z108450, St. Louis, MO, USA). Figure 1 shows the schematic of a four-neck, jacketed glass cell with the RRDE system. The counter electrode was a Li foil connected to a Ni wire, which was isolated by a layer of Celgard 2400 separator to prevent convective oxygen transport to the electrode. The Ag/Ag+ reference electrode consisted of an Ag wire immersed into 0.1 M AgNO3 in CH3CN and sealed with a vycor frit at its tip. All potentials in this study were referenced to the Li/Li+ potential scale (volts vs. Li+/Li or VLi+), obtained by calibration of the reference electrode against a fresh lithium wire before the experiments (0 VLi = −3.46 ± 0.01 V vs. Ag/Ag+). The working electrode consisted of a catalyst-covered glassy carbon disk and was immersed into the Ar or O2-purged electrolyte for 30 min before each experiment. Prior to the RRDE measurements, Alternating current (AC) impedance measurements were carried out to determine the uncompensated ohmic electrolyte drop between working and reference electrodes by applying a 10 mV perturbation (0.1 MHz to 10 mHz) at the open circuit. IR (drop) correction to remove ohmic losses was performed by considering a total cell resistance of ~293 Ω measured by AC impedance. The capacitive-corrected ORR currents were calculated by subtracting the current measurement under Ar from that obtained in pure O2 under identical scan rates, rotation speeds, and catalyst loadings.

Figure 1

Schematic of a four-neck, jacketed glass cell with a rotating ring-disk electrode (RRDE) system.

3. Results and Discussion

The phase composition and structure of the prepared <span class="Chemical">MnO2/C composites were examined by the wide-angle XRD and SAXS patterns given in Figure 2a,b. As shown in Figure 2a, all peaks can be identified as a pure and well-crystallized β-<span class="Chemical">MnO2 phase (JCPDS 24-0735) with an ordered tetragonal structure indexed to the P42/mnm space group. Moreover, the XRD curves did not show any evidence of the formation of crystalline or amorphous carbon. It appears that when using resol/Pluronic F127 templates as a carbon source, the final product is most likely to remain amorphous or in a low crystalline carbon state. The appearance of the scattering peak in the SAXS pattern, as shown in Figure 2b, indicates the long-range regularity and highly ordered nature of the mesoporous structures of the prepared MnO2/C composite.

Figure 2

(a) Wide-angle X-ray diffraction (XRD) patterns; (b) small angle X-ray scattering (SAXS) patterns of MnO2/C composites.

The morphology of the prepared <span class="Chemical">MnO2/C composite was observed using SEM and TEM, as shown in Figure 3a–f. From the SEM images of the <span class="Chemical">MnO2/C composite (Figure 3a,b), it is clear that the oriented tetragonal MnO2 nanorods are arranged on the surface of the carbon matrix. The prepared β-MnO2 nanorods, typically 2–3 μm in length, have a square cross-section with an edge length in the range of 200–300 nm. Figure 3c,d show the TEM images of the MnO2/C composite at different magnifications. Large domains of highly ordered stripe-like 1D channels are clearly observed. Figure 3e displays a TEM image of a typical nanorod with a smooth surface, and a SAED pattern based on a single nanorod (Figure 3f), indicating single-crystalline nature. The SEM and TEM analysis suggest that hollow MnO2 nanorods grow homogeneously on the ordered mesoporous carbon frameworks to form the structure of the hierarchically mesoporous MnO2/C composite.

Figure 3

Scanning electron microscope (SEM) images (a) MnO2/C composites; (b) high magnification of the region marked with a square in (a); and transmission electron microscope (TEM) images (c,e) MnO2/C composites; (d) high magnification of the region marked with a square in (c); and (f) selected area electron diffraction (SAED) pattern of the region marked with a square in (e).

(a) Wide-angle X-ray diffraction (XRD) patterns; (b) small angle X-ray scattering (SAXS) patterns of <span class="Chemical">MnO2/C composites.

Scanning electron microscope (SEM) images (a) <span class="Chemical">MnO2n>/C composites; (b) high magnification of the region m<span class="Chemical">arked with a squ<span class="Chemical">are in (a); and transmission electron microscope (TEM) images (c,e) MnO2/C composites; (d) high magnification of the region marked with a square in (c); and (f) selected area electron diffraction (SAED) pattern of the region marked with a square in (e).

The pore structure of the <span class="Chemical">mesoporous <span class="Chemical">MnO2/C composite was determined by <span class="Chemical">nitrogen adsorption-desorption isothermal measurements. As shown in Figure 4, the adsorption isothermal curve of the MnO2/C composite has a well-defined step as in typical IV classification with a H1-type hysteretic loop in the p/po range of 0.40–1.0, indicating mesoporous material character. These findings suggest that the MnO2/C composite sample does not contain framework-confined pores but is rather made up of individual nanorods. This is in agreement with the results from the SEM and TEM images. The Barrett–Joyner–Halenda (BJH) pore size distribution for the mesoporous MnO2/C composite, shown in the insert of Figure 4, reveals peaks centering at 4.8 and 35 nm. This result confirms that most of the pore channels in the ordered mesoporous carbon are not blocked by the loading of MnO2 nanorods. The nanoarchitecture of ordered mesoporous channels is maintained, which is desirable for the O2 electrode in Li-O2 batteries. Moreover, the measured BET surface area of the MnO2/C composite is relatively high, at about 424 m2·g−1. The hierarchical microstructure of the MnO2/C composite results in a large specific surface area. This is important for enhancing the electrochemical properties of an O2 cathode material.

Figure 4

Nitrogen sorption isotherms of MnO2/C composites. The insert is the Barrett–Joyner–Halenda (BJH) desorption pore size distribution.

<span class="Chemical">Nitrogen sorption isotherms of <span class="Chemical">MnO2/C composites. The insert is the Barrett–Joyner–Halenda (BJH) desorption pore size distribution.

<span class="Chemical">MnO2n> has been known as a highly active ORR catalyst for some time [30,32,36] and has recently been applied as an <span class="Chemical">O2 cathode catalyst in <span class="Chemical">Li-O2 batteries. Due to the studies of the electrocatalytic activity of MnO2, the following discussions regarding electrochemical tests make comparisons between Super-P carbon (SP) and MnO2/C materials. To better study the catalytic activity of the electrodes, cyclic voltammetry (CV) and charge-discharge voltage measurements were carried out. At first, CV was carried out in the Ar-purged electrolyte and subsequently in the same solution saturated with O2. The capacitively-corrected CV curves derived from both measurements are shown in Figure 5a. The CV plots of the O2 electrodes prepared from MnO2/C and SP cycled between 1.5 and 4.5 V with 2 mV·s−1 and the O2-saturated 1 M LiPF6/PC electrolyte are shown in Figure 5a. From the CV curves, the reduction peak voltage is shifted toward positive voltage, exhibiting electrocatalytic activity in the ORR of both samples. However, the MnO2/C offers more positive onset reduction peak potential and a larger peak current, which clearly indicate the superior electrocatalytic activity of MnO2/C compared to SP. Furthermore, the onset oxidation peaks appearing in the CV curves are about 2.7 and 2.9 V for MnO2/C and SP, respectively. This demonstrates that MnO2/C, with its lower onset oxidation peak, is more efficient for Li2O2 decomposition and has higher catalytic activity for the OER. The initial charge–discharge voltage profiles for both samples are shown in Figure 5b. The charge–discharge profiles of the MnO2/C electrode exhibit much lower charge overpotential than do those of the SP electrode, although the reduction of the total overpotential is only about 25%. The round-trip efficiencies of the Li-O2 batteries with a MnO2/C electrode were lower than those with the SP electrode. These results indicate that the MnO2/C composite can facilitate the complete reversibility of ORR and OER with low polarization for a Li-O2 battery. This finding is in good agreement with the CV measurement. The initial discharge capacities of the MnO2/C and SP electrodes were 612 mAh·g−1 and 589 mAh·g−1, respectively. The good electrochemical performance of the MnO2/C electrode may be due to the hierarchical mesostructure and large specific surface area, and the catalytic activity of the MnO2/C composite.

Figure 5

(a) CV curves were recorded at a scanning rate of 2 mV·s−1 for MnO2/C and Super-P carbon samples; (b) initial charge–discharge profiles for MnO2/C and Super P samples at a current density of 0.2 mA·cm−2.

(a) CV curves were recorded at a scanning rate of 2 mV·s−1 for <span class="Chemical">MnO2n>/C and <span class="Chemical">Super-P <span class="Chemical">carbon samples; (b) initial charge–discharge profiles for MnO2/C and Super P samples at a current density of 0.2 mA·cm−2.

The rotating ring disk electrode (RRDE) technique was also used to investigate the kinetics of ORR since the ORR current is strongly relevant to hydrodynamic conditions [31]. Here, we used a glassy <span class="Chemical">carbonn> (GC) electrode and an as-prep<span class="Chemical">ared <span class="Chemical">MnO2/C composite coated on the GC (MnO2/C-GC) electrode as the working electrodes to study the stability of the electrolyte at the MnO2/C electrode. Many reports [1,5,9] have shown the O2•− produced in the first step of the ORR upon battery discharge:

The reaction between the <span class="Chemical">O2•− and the electrolyte is the critical problem that causes poor <span class="Chemical">Li-O2 battery cyclability. In the PC-based electrolyte, the ethereal <span class="Chemical">carbon atom in PC suffers from nucleophilic attacks by O2•−, yielding carbonate, acetate, and formate species (among others), according to Equation (2) [37,38]:

Here, we applied rotating disk electrode (RDE) voltammetry to measure the rate constant (k) when reducing <span class="Chemical">O2n> to <span class="Chemical">O2•− for Equation (1) in the 0.1 M <span class="Chemical">TBAPF6/PC electrolyte. The reaction rate constant, k, can be evaluated via the Koutecky–Levich (K–L) equation for a first-order reaction as follows [31]: where i and i represent kinetics and diffusion limiting current density (A·m−2), respectively; n is the number of electrons exchanged in the electrochemical reaction; F is Faraday’s constant (96,485 C·mol−1); k is the rate constant for Equation (1); is the diffusion coefficient of O2 in the solution; ν is the kinematic viscosity; ω is the angular frequency of the rotation; and is the saturation concentration of O2 in the solution. Additionally, knowing the values of ν and for an electrolyte, one can obtain the concentration of oxygen () and rate constant (k) for Equation (1) by linearly fitting the K-L plots of i−1 vs. ω−0.5, as follows

Prior to estimating the value of k from the K–L equation, the kinematic viscosity (ν) of the electrolyte and the diffusion coefficients of <span class="Chemical">O2n> and <span class="Chemical">O2•− ( and ), need to be quantified. The value of ν for PC with 0.1 M <span class="Chemical">TBAPF6 is 2.59 × 10−2 cm2·s−1 at 25 °C (ρ = 1.2 g·mL−1 and η = 3.13 mPa·s) and was measured by a Rheometer (Malvern Gemini, Malvern Instruments Ltd., Malvern, UK). For a known viscosity, the diffusion coefficients can be directly determined from the transit-time (T) measurement by the RRDE technique, as reported previously [37,39]. Figure 6a shows an example of T measurement in O2-saturated solutions of 0.1 M TBAPF6 in PC at ω = 100 rpm; T, the origin of which is taken at time = 2 s (the time at which the disk is conducted cathodic potential at 1.85 VLi), is measured graphically from the intercept of the base steady ring current and the fast attenuate ring current line. Figure 6b,c show measurements of steady ring currents at the rotation rates (ω) of different electrodes, yielding T values for O2 and O2•−. Then, the obtained T is related to the ω and the ratio of ν and the diffusion coefficient (D), according to Equation (7) [37,39]: where K is proportionally constant depending on the RRDE’s geometry; K = 43.1[log(r2/r1)]2/3 (for T, reported in s and ω in rpm). For the RRDE used here, with r1 = 0.25 cm and r2 = 0.325 cm, the value of K is 10.1 rpm·s. Table 1 shows the estimated values of the diffusion coefficients of O2 and O2•− calculated from Equation (7) based on the slopes of T ω−1 obtained from Figure 6d. Figure 7a shows that well-defined O2 diffusion-limited currents are obtained for the ORR on a GC electrode in an O2-saturated 0.1 M TBAPF6/PC solution. The K–L plot for the disk current values at 1.50 VLi reveals the expected linear relation between the inverse of the limiting current and ω−0.5 (see Equation (6)). As shown in Table 1, the concentration of oxygen () on the GC electrode was estimated from the slope of the K–L plot using the prior measured values of ν and , where n = 1 (according to the reaction of Equation (1)). The value of is 6.1 M, which is higher than the finding of a previous report (4.8 M) [37]. This can be attributed to the larger O2 flow rate in this experiment. The estimated value of was also applied in the following calculations of the MnO2/C-GC electrode since the same operation parameters (i.e., O2 flow rate, electrolyte composition, and amount) were used, as listed in Table 1. The rate constant for producing O2•−, k for GC and the MnO2/C-GC electrodes can be obtained by linearly fitting the K–L plots of i−1 vs. ω−0.5 (see Equation (6)), as shown in Figure 7a,b. The values of k for GC and the MnO2/C-GC electrodes are 1.92 × 10−2 cm·s−1 and 4.29 × 10−2 cm·s−1, respectively. This result indicates that the MnO2/C cathode catalyst exhibits a larger k value, resulting from higher electrocatalytic activity for the first step of the ORR (see Equation (1)) which produces a higher concentration of O2•−.

Figure 6

(a) Example of determination of the superoxide radical (O2•−) transit-time (T) in O2-saturated solutions of 0.1 M TBAPF6 in propylene carbonate (PC) at ω = 100 rpm, Edisk = 1.85 V and Ering = 2.6 V. Transit time (Ts) values at different rotation rates for the diffusion of (b) O2 and (c) O2•−; (d) relation between the inverse of the rotation speed and the transient time for O2 and O2•−.

Table 1

Summary of the electrolyte properties estimated with the proposed RRDE-based methodology and comparison with findings reported in the literature.

Disk Material/Electrolyte	ν (cm2·s−1)	D_O_2 (cm2·s−1)	D_O2•− (cm2·s−1)	C_O_2 (mM)	Reference
GC/0.1 M TBAPF6, PC	2.6 × 10−2	1.9 × 10−5	8.6 × 10−6	6.1	This work
MnO2/C-GC/0.1 M TBAPF6, PC	2.6 × 10−2	1.9 × 10−5	1.8 × 10−6	6.1	This work
GC/0.2M TBATFSI, PC	2.6 × 10−2	2.5 × 10−5	6.8 × 10−6	4.8	[37]

Figure 7

(a) Steady-state CV curves of a glassy carbon rotating disk electrode (RDE) in an O2-saturated 0.1 M TBAPF6/PC solution at a scan rate of 50 mV/s between 1.5 and 2.8 VLi with different rotation rates. The insert is the Koutecky–Levich plot derived from the disc current values at 1.50 VLi; (b) steady-state CV curves of a MnO2/C RDE in an O2-saturated 0.1 M TBAPF6/PC solution at a scan rate of 50 mV/s between 1.2 and 2.8 VLi with different rotation rates.

(a) Example of determination of the <span class="Chemical">superoxide radical (<span class="Chemical">O2•−) transit-time (T) in <span class="Chemical">O2-saturated solutions of 0.1 M TBAPF6 in propylene carbonate (PC) at ω = 100 rpm, Edisk = 1.85 V and Ering = 2.6 V. Transit time (Ts) values at different rotation rates for the diffusion of (b) O2 and (c) O2•−; (d) relation between the inverse of the rotation speed and the transient time for O2 and O2•−.

Summ<span class="Chemical">ary of the electrolyte properties estimated with the proposed RRDE-based methodology and comp<span class="Chemical">arison with findings reported in the literature.

(a) Steady-state CV curves of a glassy <span class="Chemical">carbon rotating disk electrode (RDE) in an <span class="Chemical">O2-saturated 0.1 M <span class="Chemical">TBAPF6/PC solution at a scan rate of 50 mV/s between 1.5 and 2.8 VLi with different rotation rates. The insert is the Koutecky–Levich plot derived from the disc current values at 1.50 VLi; (b) steady-state CV curves of a MnO2/C RDE in an O2-saturated 0.1 M TBAPF6/PC solution at a scan rate of 50 mV/s between 1.2 and 2.8 VLi with different rotation rates.

Recently, Herranz et al. [37] used RRDE voltammetry to quantify the stability of an electrolyte against <span class="Chemical">O2n>•− by the rate constant (k) for Equation (2) According to their methods, the <span class="Chemical">O2•− produced at the disk electrode in Equation (1) and the amount of <span class="Chemical">O2•− were quantified at the ring electrode. The amount of O2•− consumed depends on the effective transient time, T, between the disk and the ring and the rate constant, k, for Equation (2). Longer T and larger k values cause increasing consumption of O2•− due to its reaction with the electrolyte, resulting in a lower O2•− oxidation current at the ring. Therefore, the collection efficiency, N, for O2•− at the ring electrode decreases with increasing transient time, which, in turn, depends on the geometry of ring and disk electrode, the diffusion coefficient of O2•− in the electrolyte, , and the electrode rotation speed, ω. The correlation with the collection efficiency is the absolute ratio of ring and disk currents and can be characterized by the following equation [37,40]: where A1 = 1.288, A2 = 0.643 ν1/6 1/3, β = 3ln(r3/r2), U = k−1tanh(A1k) and T2 = 0.718ln(r2/r1), whereby r1–r3 refer to the radius of the disk and internal and external ring radii, respectively; ν is the kinematic viscosity; ω is the rotation rate; k is the rate constant for Equation (2); and is the diffusion coefficient of O2•−. N is the geometrical collection efficiency of the RRDE corresponding to the fraction of a species electrochemically generated at the disk. This species is detected at the ring due to the lack of side-reactions with the electrolyte. Equation (8) shows the variation of N where the rotation rate and the rate constant (k) can be calculated at higher rotation rates, which show that the N is close to a constant value. Figure 8a shows the RRDE profiles of the MnO2/C sample coating on the disk electrode. The disk and ring currents are recorded in an O2-saturated 0.1 M TBAPF6/PC solution at rotation rates between 300 and 2100 rpm, with continuous holding of the Pt ring at 2.85 VLi. The ring current increases with the rotation rates because the shorter transient time at higher rotation rates reduces the reaction time between O2•− and the PC electrolyte so that a higher concentration of superoxide radical can be oxidized at the ring. Also, the N increases with rotation rates (ω) and is close to a constant value (0.14) at ω = 2100 rpm, as shown in Figure 8b. The PC-electrolyte decomposition rate constant (k) can be calculated by Equation (8) using the N value at a rotation speed of 2100 rpm with the kinematic viscosity (ν) and listed in Table 1. Table 2 shows the rate constant for producing O2•−, k, and the PC-electrolyte decomposition rate constant, k, on the GC and MnO2/C-GC electrodes. The value of k (1.5 s−1) on the GC electrode is close to that of a previously reported measurement (k = 1.3 s−1) [37]. Obviously, the k value on the MnO2/C-GC electrode of 2.6 s−1 is larger than that on the GC electrode. This result shows that MnO2/C is more active for the first step of the ORR (larger rate constant; k), producing a higher concentration of O2•− and leading to faster PC-electrolyte decomposition due to the attack by a large amount of O2•−. Therefore, it is important to choose an appropriate electrolyte to avoid decomposition by O2•− attack for highly active catalyst applications on the cathode materials in Li-O2 batteries. More detailed RRDE experiments and analysis will be carried out to estimate the decomposition rates of various electrolytes with different active catalysts.

Figure 8

(a) RRDE profiles of MnO2/C recorded at 50 mV·s−1 in an O2-saturated 0.1 M TBAPF6/PC solution, at rotation rates between 300 and 2100 rpm with continuous holding of the Pt ring at 2.85 VLi; (b) evolution of the absolute ratio between the ring and disk current (N) and the electrode rotation speed (ω).

Table 2

The rate constant for producing O2•−, k, and the PC-electrolyte decomposition rate constant, k, on the GC and MnO2/C-GC electrodes.

Disk Material/Electrolyte	kf (cm·s−1)	k (s−1)	Reference
GC/0.1 M TBAPF6, PC	1.9 × 10−2	1.5	This work
MnO2/C-GC/0.1 M TBAPF6, PC	4.3 × 10−2	2.6	This work
GC/0.2M TBATFSI, PC		1.3	[37]

The rate constant for producing <span class="Chemical">O2•−, k, and the PC-electrolyte decomposition rate constant, k, on the GC and <span class="Chemical">MnO2/C-GC electrodes.

(a) RRDE profiles of <span class="Chemical">MnO2/C recorded at 50 mV·s−1 in an <span class="Chemical">O2-saturated 0.1 M <span class="Chemical">TBAPF6/PC solution, at rotation rates between 300 and 2100 rpm with continuous holding of the Pt ring at 2.85 VLi; (b) evolution of the absolute ratio between the ring and disk current (N) and the electrode rotation speed (ω).

4. Conclusions

A hier<span class="Chemical">archically <span class="Chemical">mesoporous <span class="Chemical">carbon-supported manganese oxide (MnO2/C) has been synthesized by a combination of soft template and hydrothermal methods. SEM and TEM analysis confirmed that hollow MnO2 nanorods grow homogeneously on ordered mesoporous carbon frameworks to form a hierarchically mesoporous MnO2/C composite structure. The CV and galvanostatic charge–discharge tests indicate that MnO2/C composites have excellent catalytic activity towards ORR due to the larger surface area of ordered mesoporous carbon and higher catalytic activity of MnO2.

The <span class="Chemical">O2 solubility, the diffusion rates of <span class="Chemical">O2 and <span class="Chemical">O2•− coefficients, the rate constant for producing O2•− (k), and the PC-electrolyte decomposition rate constant (k) of the MnO2/C composites have been measured by RRDE experiments and analysis in the 0.1 M TBAPF6/PC electrolyte. The results indicate that MnO2/C is more active for the first step of the ORR (larger rate constant; k), produces a higher concentration of O2•−, and leads to faster PC-electrolyte decomposition due to the attack by a large amount of O2•−. The stability of the electrolyte is very important when applying an active catalyst on a cathode material in Li-O2 batteries. More detailed RRDE experiments and analysis will be carried out to estimate the decomposition rates of various electrolytes. These results seem to be interesting for the design of advanced Li-O2 batteries with high electrochemical performance.