Role of Cu-Ion Doping in Cu-α-MnO2 Nanowire Electrocatalysts for

Article pubs.acs.org/JPCC

Role of Cu-Ion Doping in Cu-α-MnO2 Nanowire Electrocatalysts for the Oxygen Reduction Reaction Danae J. Davis,† Timothy N. Lambert,*,† Julian A. Vigil,† Mark A. Rodriguez,‡ Michael T. Brumbach,‡ Eric N. Coker,§ and Steven J. Limmer∥ †

Department of Materials, Devices & Energy Technologies, ‡Materials Characterization & Performance, §Advanced Materials Laboratory, and ∥Physics Based Microsystems, Sandia National Laboratories, Albuquerque, New Mexico 87185, United States S Supporting Information *

ABSTRACT: The role of Cu-ion doping in α-MnO2 electrocatalysts for the oxygen reduction reaction in alkaline electrolyte was investigated. Cu-doped α-MnO2 nanowires (Cu-α-MnO2) were prepared with varying amounts (up to ∼3%) of Cu2+ using a hydrothermal method. The electrocatalytic data indicate that Cu-α-MnO2 nanowires have up to 74% higher terminal current densities, 2.5 times enhanced kinetic rate constants, and 66% lower charge transfer resistances that trend with Cu content, exceeding values attained by α-MnO2 alone. The observed improvement in catalytic behavior correlates with an increase in Mn3+ content at the surface of the Cu-α-MnO2 nanowires. The Mn3+/Mn4+ couple is the mediator for the rate-limiting redox-driven O2/OH− exchange. O2 adsorbs via an axial site (the eg orbital on the Mn3+ d4 ion) at the surface or at edge defects of the nanowire, and the increase in covalent nature of the nanowire with Cu-ion doping leads to stabilization of O2 adsorbates and faster rates of reduction. A smaller crystallite size (roughly half) for Cu-α-MnO2 leading to a higher density of (catalytic) edge defect sites was also observed. This work is applicable to other manganese oxide electrocatalysts and shows for the first time there is a correlation for manganese oxides between electrocatalytic activity for the oxygen reduction reaction (ORR) in alkaline electrolyte and an increase in Mn3+ character at the surface of the oxide.

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INTRODUCTION Efficient catalysis of the oxygen reduction reaction (ORR) and/ or oxygen evolution reaction (OER) is required for the effective development of renewable energy technologies, including the operation of fuel cells, metal-air batteries, and electrolysis cells. The performance of such electrochemical storage and electricitygenerating devices relies upon the efficiencies and kinetics of the ORR (e.g., during discharging) and the OER (e.g., during charging) occurring at the cathode. The OER is also the anode reaction employed in electrolysis cells. Interest in high capacity primary metal air batteries1 and fuel cells2 requires developing ORR catalysts3−5 with emphasis on the ORR in alkaline aqueous electrolytes.3,6 It has long been established that platinum (Pt) and Pt-based alloys are leading catalysts for the ORR, and hence they are commonly used as the research and development benchmark for gauging catalytic performance. However, the widespread implementation of Pt-containing catalysts in current and future technologies is not feasible due to the prohibitive cost and rarity of Pt. Pt and Pt-based catalysts also suffer from reaction poisoning and poor electrochemical selectivity in alkaline fuel cells, such as in carbon monoxide poisoning and methanol crossover, respectively. Hence, alternative catalysts that function with high electroselectivity, are cost-effective and stable, and exhibit low overpotential and fast kinetics are still needed. A number of catalysts are being developed, including but not limited to ceramics, nanoscale carbons, and blends thereof.7−10 Manganese oxide (MnOx) has been utilized extensively as a catalyst in alkaline electrolyte for the ORR. First investigated in © 2014 American Chemical Society

the early 1970s, commercially available crystalline manganese oxide (MnO2) was demonstrated to improve the performance of carbon-air electrodes when added in small amounts, generally less than 10 wt %.11 Since then, numerous MnOx catalysts have demonstrated promising catalytic behavior.12−14 In addition to favorable catalytic performance, MnOx meets the requisite criteria of being environmentally benign, earth-abundant, and overall potentially cost-effective, making their use in widespread electrochemical energy applications promising. Manganese can exist in numerous valences (+2, +3, +4, +6, and +7), a property that has been hypothesized to be a key factor in the catalytic activity of MnOx. MnOx can also occur in a variety of stoichiometries, crystalline phases, and morphologies, factors that are known to have a significant impact on the catalytic performance. For example, the order of ORR catalytic activity for phases of manganese oxide is as follows: γ-MnOOH > Mn2O3 > Mn3O4 > Mn5O8.15 For MnO2 of differing phases, the catalytic activities were demonstrated to be α-MnO2 ≈ δ-MnO2 > γMnO2 > λ-MnO2 > β-MnO2.16 The α-MnO2 crystalline lattice possesses larger (2 × 2) pore tunnels of approximately 4.6 Å diameter, as compared to the (1 × 1) and (1 × 2) tunnels in the β and γ phases with 1.89 and 2.3 Å diameters, respectively. However, another study found α-MnO2 > β-MnO2 > γ-MnO2, citing the enhanced electrical conductivity of the β-phase as Received: April 23, 2014 Revised: July 1, 2014 Published: July 9, 2014 17342

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We demonstrate for the first time that Cu-ion doping leads to smaller crystallite domains [i.e., a higher density of edge defect sites] and an increase in the amount of Mn3+ at the surface of the Cu-α-MnO2 nanowire and rationalize how this improves the catalytic activity for the ORR. We also demonstrate that the increased activity of Cu-α-MnO2 relative to α-MnO2 is a result of an increased covalent nature for Mn−O bonds of the nanowire, which results in stabilization of O2 adsorbates and faster kinetic rates.

compared to the γ-species.12 Cheng et al. further demonstrated the effect that the particle size and the morphology of the catalyst have on activity.12 Nanostructures were found to be superior to microstructures, and 1-dimensional α-MnO2 nanowires were more catalytically active than 3-dimensional nanospheres of the same crystalline phase with similar surface area (32.9 and 40.1 m2 g−1, respectively).12 In general agreement with Cheng et al., a recent study that examined the α- and δ-MnO2 phases determined that α-MnO2 nanorods and nanotubes are more catalytically active than mixed microsphere/nanosheet δMnO2.13 Methods to further increase the catalytic activity of MnOx include inducing surface defects or oxygen vacancies in the crystalline lattice, increasing conductivity via metal ion doping, and/or altering the average valence of Mn. For example, the performance of β-MnO2 was improved by introducing oxygen defects to the surface.17 Roche et al. prepared MnOx/carbon composites, with a largely amorphous, birnessite-like phase of MnOx,18 doped with Ni or Mg cations and found the doped catalysts to yield higher n-values, indicating an improvement in the ORR mechanism toward a more apparent four-electron process.19 The Ni-doped species was also shown to result in a significant decrease in the production of peroxide intermediate; it was suggested that the Ni dopant was stabilizing the Mn3+ intermediate obtained from the electro-reduction of the ceramic in the ORR pathway.19 The activity of spinel (Mn3O4) catalysts has also been improved upon the introduction of metal ion dopants (e.g., CuxMn3−xO4).20 We recently reported on Ni- and Cu-doped α-MnO2 nanowire catalysts and demonstrated their activity was superior to the already highly active α-MnO2 nanowire (73.6 m2 g−1), despite having lower surface areas (54.1 and 43.7 m2 g−1, respectively).3 In addition to attaining higher current densities, the Ni- and Cudoped species demonstrated substantially improved kinetics. When simply blended with a graphene-like carbon (GLC), the 20% Cu-α-MnO2/80% GLC composite catalyst compared favorably with the commercial E-Tek 20% Pt/C benchmark catalyst. Combining MnOx with graphene-based carbons is an attractive extrinsic approach to boost activity, via increased conductivity, and in some cases additional, synergistic activity with the graphene component.3,21−25 While it was shown with Xray diffraction and elemental mapping from energy-dispersive Xray spectroscopy (EDS) that the Cu ions in the Cu-α-MnO2 were doped substitutionally throughout the nanowire,3 the exact role of Cu ions on the activity of Cu-α-MnO2 for the ORR remained unanswered. Given their promising activity, further studies into the origin of the improved electrocatalytic activity for Cu-αMnO2 were warranted. In general, cation doping of MnOx has previously been suggested to increase the Mn3+ content, resulting in increased electrocatalytic activity;19,20,26 however, to the best of our knowledge, until now there was no experimental evidence to support this. A recent study on Cu ions in α-MnO2 focused on changes in the crystalline structure and rodlike morphology, with a lithium-ion battery cathode as the material’s application space.27 Numerous papers have provided information on the generally proposed mechanisms/pathways of ORR with MnOx, and we refer the readers to those for more details.12,13,16,19 Here we present the results of our investigation of the role of Cu-ion doping in Cu-α-MnO2 nanowires by examining Cu-αMnO2 with varying amounts of Cu2+ dopant. This study examines how Cu2+ alters the nanowire, leading to improvements in the electroactivity for the ORR catalysis in alkaline electrolyte.

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EXPERIMENTAL SECTION Materials Synthesis. All chemicals were obtained commercially and used as received, without further purification. MnSO4· H2O and Cu(NO3)2·3H2O were purchased from Alfa Aesar and Acros, respectively. KMnO4, isopropanol, and Nafion solution (5 wt % solution in water and lower aliphatic alcohols) were used as received from Sigma-Aldrich. The synthesis of Cu-α-MnO2 nanowires was performed via a hydrothermal reaction in an acid digestion bomb (Teflon liner, 45 mL capacity, Parr). KMnO4 (0.287 g, 1.8 mmol) was added to a solution of MnSO4·H2O (0.102 g, 0.6 mmol) and specified amounts of Cu(NO3)2·3H2O [(0.145 g, 0.6 mmol), (0.0725 g, 0.3 mmol), or (0.0362 g, 0.15 mmol)] in deionized water (18 mL), and the resulting reaction mixture was heated in an acid digestion bomb at 140 °C for 120 h to ensure nanowire morphology.3,12 The Cu(NO3)2·3H2O was added in molar ratios of 1:1, 1:0.5, and 1:0.25 (Mn:Cu) to achieve Cu-α-MnO2 with different Cu doping levels. Undoped α-MnO2 was obtained by a hydrothermal reaction that did not include the Cu precursor.3,12 Upon completion, all reactions were cooled to room temperature. The solids were washed with deionized H2O (4 × 10 mL) and ethanol (4 × 10 mL), isolated via centrifugation, and finally dried in a vacuum oven at 65 °C overnight. The Parr bombs were cleaned in between use by heating them at 140 °C for 24 h with aqueous EDTA (3 wt %, 20 mL) and then deionized water (20 mL). Materials Characterization. Powder X-ray Diffraction (PXRD). Catalyst powders were mounted directly onto a zero background holder. Samples were scanned at a rate of 0.02°/2 s in the 2θ range of 10−80° on a Bruker D8 Advance Diffractometer in Bragg−Brentano geometry with Cu Kα radiation and a diffracted beam graphite monochromator; longer scans (t = 12 h) were performed at a rate of 0.024°/15 s in the range of 10−80°. Phase identification was achieved through use of the Jade 9.0 software suite and the Joint Committee on Powder Diffraction Standards (JCPDS). Diffraction peak positions were determined by profile fitting of peak profiles within the program Jade 9.0 using a pseudo-Voigt profile function. Using these peak positions, lattice parameters were then derived using a least-squares refinement procedure embedded within the Jade 9.0 software. Initial values of the unit cell were input based on the α-MnO2 structure (JCPDS file 00-44-0141). Thermogravimetric Analysis: Differential Scanning Calorimetry with Mass Spectrometry (TGA/DSC/MS). Simultaneous TGA/DSC with analysis of off-gas via MS was conducted using an STA 449 F3 Jupiter TGA/DSC (Netzsch Instruments, Selb, Germany) coupled to an HPR-20 MS (Hiden Analytical, Warrington, UK). The specimen was loaded into a platinum crucible and was allowed to degas and stabilize in the TGA/DSC under flowing argon overnight prior to initiating the run. The MS was connected via a “T” to the vent from the TGA/DSC furnace and analyzed a continuous stream of gas during the TGA/DSC 17343

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not to identify absolute concentrations. Mn 3s peak fitting was performed with a linear background and two unconstrained Gaussian/Lorentzian (70/30) peaks. X-ray Fluorescence Analysis. X-ray fluorescence analysis was performed using a Bruker M4 micro-XRF system (Madison, WI). The beam source was a microfocus X-ray source (Rh anode). Energy spectrum data were collected using a Silicon Drift Detector (SDD). XRF spectra were collected in point mode using a polycapillary optic on the microfocus X-ray source to define a small sampling location (∼30 μm diameter spot size) within the cathode region of the specimen. Analysis of the energy spectrum for chemical species identification was performed within the M4 software provided with the system. XRF analysis conducted on an electrode after RDE analysis confirmed no Pt contamination on the working from the counter occurred. Detected species (above sodium) were S, K, Mn, and Cu. No Pt was detected (see Supporting Information). Elemental mapping of the entire cathode suggested a relatively even distribution of the S, K, Mn, and Cu concentrations and no indication of significant Pt presence. Atomic Absorption Spectrophotometric Analysis. Atomic absorption spectrophotometry (ASS) was performed on a PerkinElmer 900T Atomic Absorption spectrophotometer in flame mode. Specimens (∼2 mg) were digested in 10 mL of conc. HCl and then diluted with deionized water to the required concentration range for analysis. All blanks were matrix-matched to standards and sample solutions. Electrocatalytic Measurements. Oxygen reduction reaction studies were performed using a rotating disk electrode (RDE) and rotating ring disk electrode (RRDE) apparatus. Catalyst inks consisting of the catalyst material (5.0 mg), isopropanol (200 μL), and Nafion solution (300 μL, 5 wt %) in a 4 mL glass vial were mixed using bath sonication for approximately 20 min. For RDE studies, a 5 μL aliquot was drop cast on a glassy carbon RDE electrode [Bioanalytical Systems, Inc. (BASi), active area = 0.0788 cm2] and allowed to dry giving 0.63 mg cm−2 of active catalyst. For RRDE studies, 10 μL of the prepared ink was drop cast onto the glassy carbon disk, surrounded by a platinum ring electrode (Pine, disk area = 0.2475 cm2, ring area = 0.1866 cm2) giving 0.40 mg cm−2. Prior to each drop-coating process, the RDE or RRDE electrode surface was polished with alumina slurry (0.05 μm) and rinsed with deionized water and ethanol and allowed to air-dry. Linear scanning voltammograms (LSVs) were executed in a potential window from 0.2 to −0.6 V (V vs Ag/AgCl) at a scan rate of 1 mV/s. For RDE studies, LSVs were run at rotation rates of 500, 900, 1600, 2500, and 3600 rpm. For RRDE studies, LSVs were obtained at 500 and 2500 rpm. The collection efficiency of the ring electrode was determined to be 37% by using a solution of K3Fe(CN)6.28,29 All experiments were run in 0.1 M KOH electrolyte that had been thoroughly purged with UHP oxygen or nitrogen (for subtraction of background) prior to each run. The appropriate gaseous atmosphere was maintained throughout the experiment by gaseous blanketing. RDE experiments were performed in a three-electrode cell, manufactured by Bioanalytical Systems, Inc., connected to a Versastat 4 potentiostat operated by VersaStudio software. Electrodes (from BASi) included the glassy carbon working electrode, prepared as described, a Pt coil counter electrode, and a Ag/AgCl reference electrode (3 M NaCl, potential = 0.209 V vs NHE). RRDE experiments were performed in a three-electrode cell, manufactured by Gamry Instruments connected to a Gamry Series G 750 Test System Bipotentiostat. Electrodes included the

run. Due to the close coupling of the MS inlet to the TGA/DSC furnace, lag time between gas evolution from the specimen and its detection by MS was considered negligible on the time frame of the experiment and was ignored. The specimen was heated in the TGA/DSC at a rate of 5 °C min−1 from 25 to 750 °C under a flow of 100 mL min−1 ultrahigh purity (UHP) argon. The evolution of oxygen (m/z = 32) was monitored continuously by MS during the course of the TGA/DSC run. Due to the small sample size we estimate that TGA/DSC errors are ∼3% on DSC peak areas and 2% (relative) on mass changes. Scanning Electron Microscopy (SEM). The samples were dispersed onto carbon tape and imaged using a Zeiss Supra 55VP field emitter gun scanning electron microscope (FEGSEM). A Noran EDS detector and Noran System Six software were used for the acquisition of the EDS spectra. EDS data were replotted with Kaleidagraph 4 for presentation purposes. Samples were sputter coated with gold−palladium prior to analysis. Brunauer−Emmett−Teller (BET) Surface Area Analysis. Gas sorption experiments were performed on a Micrometrics Tristar 3000 sorptometer or an Autosorb iQ2-Chemi instrument (Quantachrome Instruments, Boynton Beach, FL, USA) at 77 K using UHP nitrogen as adsorbate. Specimens were degassed at 140 °C for 2−3 h (the exact degas time was determined automatically based upon periodic evaluation of rate of gas evolution from each specimen). The specimen cells were quartz glass with a 9 mm outer diameter, and cell volumes were determined prior to each run, using UHP helium. Isotherms were typically measured using 40 adsorption points and 40 desorption points, with an optional 30 additional adsorption points in the micropore range. Surface areas were calculated using the Brunauer−Emmett−Teller (BET) method using five adsorption points in the range P/P0 = 0.1−0.2. Total pore volumes were calculated based on the total gas adsorbed at P/P0 values of 0.993. Average pore sizes were calculated from the measured total pore volume and surface area values. The errors in the sorption data are dominated by the errors in weighing the small samples. Error estimates are ±0.6 m2 g−1 on each surface area, ± 0.1 nm on pore size, and ±0.02 cm3 g−1 on each pore volume. X-ray Photoelectron Spectroscopy. X-ray photoelectron spectroscopy (XPS) was performed with a Kratos Axis Ultra DLD instrument with base pressures less than 5 × 10−9 Torr. Powder samples were pressed into carbon tape for analyses. Charge neutralization was not required; however, the neutralizer was used for all samples to ensure that differential charging did not influence the spectral line shapes. XPS was performed with a monochromatic Al Kα (1486.7 eV) source operated at 150 W. The analysis spot size was elliptical: 300 × 700 μm. Several locations from each sample were analyzed. Survey spectra were taken with pass energy of 80 eV, 600 meV step, and 100 ms dwell times. With Al Kα X-rays the Mn LMM Auger transition overlaps with the Cu 2p core-level peaks, so XPS was also performed with a Mg X-ray source (1253.6 eV) operated at 180 W. The energy resolution is lower with the Mg source, so step sizes were 800− 1000 meV for survey spectra. XPS with the Mg source was helpful for resolving the Cu 2p peaks on a relatively flat background. High-resolution spectra were recorded with a 20 eV pass energy, 200 ms dwell time, and step sizes ranging from 30 to 50 meV (100 meV for Mg X-ray XPS). Data processing was performed with CasaXPS. Quantifications were performed using the built-in relative sensitivity factors, which are Scofield cross sections. Only Cu, Mn, and O were included in the quantifications to avoid possible interferences from the carbon tape. The primary objective was to observe trends in changing surface composition, 17344

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Values of the Cu content of each sample were determined using atomic absorption spectrophotometry (AAS), after acid digestion of the samples (Table 1). The sample synthesized at

Gamry RDE710 rotating working electrode prepared as described, Pt foil as the counter electrode, and a Hg/HgO reference electrode (0.1 M KOH, potential = 0.174 V vs NHE) at a scan rate of 1 mV/s. RDE and RRDE experiments were not corrected for IR drop. n values (indicating the reaction order) were determined following published protocols.3,12 The onset potential, typically defined as the potential at which the current first becomes negative, was alternatively defined in this study. The onset potential was determined from the averaged LSVs obtained at 2500 rpm and was approximated from the intersection of a tangential line from 0 V and a tangential line from the half-wave potential. This approach likely underestimates the onset as compared to literature values but was found to be more reliable for relative comparisons here.

Table 1. Average Atomic Percent Composition for Cu-αMnO2 Synthesized at Varying Reactant Ratiosa

a

initial reactant ratio (Mn:Cu)

Cu %

Mn %

1:1 1:0.5 1:0.25

2.92 ± 0.047 2.39 ± 0.0329 1.30 ± 0.0279

31.15 ± 0.465 31.11 ± 0.687 31.44 ± 0.507

As determined using acid digestion/AAS.

the initial reactant ratio of 1:1 (Mn:Cu) contained 2.92% atomic percent Cu; the sample prepared at the reactant ratio of 1:0.5 (Mn:Cu) contained 2.39% Cu; and the sample prepared at the reactant ratio of 1:0.25 (Mn:Cu) contained 1.30% Cu, indicating that an increased amount of Cu precursor in the reaction correlates to an increased amount of Cu in the sample. For purposes of discussion these samples will be referred to as Cu2.9, Cu-2.4, and Cu-1.3. Powder X-ray diffraction (PXRD) was used to confirm that the crystalline phase with α-MnO2 (JCPDS file 00-44-0141) was obtained in all cases (Figure 1d−g and SI Figure S2). These results are consistent with the substitutional doping of Cu ions into the α-MnO2 lattice, as previously determined by TEM.3 The PXRD data were further analyzed to determine the effects of the Cu dopant on the unit cell parameters of the α-MnO2 lattice (Table 2). For the control sample of undoped α-MnO2, the

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RESULTS AND DISCUSSION The synthesis of Cu-α-MnO2 nanowires was achieved by adapting a hydrothermal synthesis method for α-MnO2.3,12 Briefly, KMnO4 was added to a solution of MnSO4·H2O and Cu(NO3)2·3H2O in deionized water. Molar ratios of Mn to Cu in the initial reaction mixtures were 1:1, 1:0.5, or 1:0.25 (Mn:Cu). The reaction mixture was then heated in an acid digestion bomb at 140 °C for 120 h. Contrary to previous reports,12 nanowires were not completely obtained in 12 h in our hands, hence the longer reaction time was employed. The scanning electron microscopy (SEM) images shown in Figure 1a−c demonstrate that for all reactions nanowires were

Table 2. Effects of Cu Dopants on the Unit Cell Parameters of the α-MnO2 Lattice sample

Cu %

a (Å)

c (Å)

volume (Å3)

JCPDS file 00-44-0141 α-MnO2 Cu-α-MnO2 Cu-α-MnO2 Cu-α-MnO2

0 0 1.30 2.39 2.92

9.785 9.836 (3) 9.838 (3) 9.84 (1) 9.866 (5)

2.863 2.858 (1) 2.859 (1) 2.857 (2) 2.857 (1)

274.1 276.45 (3) 276.74 276.5 278.12

measures of the a-axis (9.836 Å) and unit cell volume (276.45 Å3) were found to be greater than the measures documented in JCPDS file 00-44-0141 (9.785 and 274.1 Å3, respectively), while the c-axis was found to be smaller (2.858 Å) in the nanowire (vs 2.863 Å for α-MnO2). The a-axis (and volumes) expanded further from 9.836 Å (276.45 Å3) for α-MnO2 to 9.866 Å (278.12 Å3) for Cu-2.9, the Cu-α-MnO2 nanowire sample with the highest Cu content (Figure 2). The substitution of larger, sixcoordinate Cu2+ ions (0.73 Å) in place of smaller Mn3+ (0.645 Å)

Figure 1. (a−c) SEM images and (d−f) XRD patterns for Cu-α-MnO2 nanowires synthesized at Mn:Cu ratios of 1:1, 1:0.5, and 1:0.25, respectively. (g) JCPDS 00-44-0141 index.

obtained. While exact aspect ratios are hard to determine due to the nanowire aggregation, the wires are ∼100 nm in width and can achieve lengths of several micrometers, giving aspect ratios that can range from ∼10 to 30. EDS obtained via SEM provides evidence of Cu in the nanowire samples (SI, Figure S1a−c).

Figure 2. Tetragonal (I4/m) structure: MnO6 octahedra are shown, along with the a- and c-axis. Analyses of PXRD data revealed expansion of the unit cell along the a-axis upon Cu doping. Note a = b.31 17345

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or Mn4+ ions (0.530 Å) is consistent with an overall expansion in the lattice.30 Our results contradict a previous report where Cudoped α-MnO2 nanowires exhibited a decrease along the a-axis and corresponding decreases in cell volume.27 However, this contradiction can be rationalized in terms of the total water content of the nanowire, whereby higher water content leads to larger pore volumes.27 TGA/DSC/MS analysis (vide infra) clearly indicates higher water content in Cu-doped α-MnO2 versus α-MnO2, as prepared here. The growth of the α-MnO2 nanowires has been shown to occur along the c-axis direction,3 which, as illustrated in Figure 2, is the coincident propagation direction of the large (2 × 2) channel through the structural lattice. The peak width appears to have an (hkl) dependence, showing sharp peak widths for strongly c-axis dependent peaks, e.g., the (002), but considerably larger peak widths for a-axis dependent reflections. This results from the very significant aspect ratio of the α-MnO2 fiber morphology as well as the crystallographically dependent relationship of the c-axis coincident with the axis of the fiber. Surface areas, pore size, and pore volume were obtained by N2 adsorption−desorption measurements, as presented in Table 3

Figure 3. Averaged LSVs with increasing current obtained at 500, 900, 1600, 2500, and 3600 rpm for: (a) Cu-2.9, (b) Cu-2.4, and (c) Cu-1.3. (d) Overlay of LSVs obtained at 2500 rpm for Cu-1.3 (red triangles), Cu-2.4 (green squares), Cu-2.9 (blue circles), and α-MnO2 (orange diamonds). All LSVs are the average of three independent trials. All experiments at 1 mV s−1 scan rate, in O2-saturated 0.1 M KOH with N2 background subtracted.

Table 3. Surface Areas, Pore Sizes, and Pore Volumes for Cuα-MnO2 Samples Cu %

surface area (m2 g−1)

average pore diameter (nm)

pore volume (cm3 g−1)

0 1.30 2.39 2.92

73.6 59.1 80.7 83.8

13.4 9.4 9.2 11.6

0.31 0.14 0.19 0.24

Ag/AgCl) only ∼10 mV more positive than Cu-1.3 (at −0.107 V vs Ag/AgCl). The average onset potential for a benchmark catalyst, 20% Pt on Vulcan XC-72 Carbon (E-Tek) (data not shown), was found to be ∼50 mV more positive (−0.051 V vs Ag/AgCl). A ∼20 mV positive shift in the half wave potential with increasing Cu content was also observed, with half-wave values ranging from approximately −0.290 to −0.312 V. The onset for the Cu-doped α-MnO2 catalysts can be further improved by combining with GLC.3 The kinetic rate constants, extracted from Koutecky−Levich plots shown in Figure 4a, indicate an increase in reaction rate with increasing Cu content. Cu-2.9 was shown to have an average kinetic rate constant of 1.9 × 10−2 cm s−1, while Cu-2.4 and Cu1.3 gave average rate constants of 9.7 × 10−3 and 7.8 × 10−3 cm s−1, respectively. For comparison, the kinetic rate constant obtained under identical conditions for α-MnO2 and Pt/C was 7.5 × 10−3 and 5.84 × 10−2 cm s−1, respectively. Surprisingly, the average number of electrons (n) involved in the ORR process, calculated using the Koutecky−Levich equation, revealed little difference between samples of varying Cu content; n-values for all three catalysts were n = 3.2−3.3. The two factors of n-value and reaction rate, when considered together, are suggestive that, while the ORR mechanism is not highly influenced by the Cu content of the α-MnO2 catalyst, the reaction kinetics become more rapid as more Cu is added. RRDE experiments were used to determine the amount of peroxide formation. The fraction of evolved peroxide was calculated by XH2O2 = 2(IR/N)/[(IR/N) + (ID)] where IR is the ring current, N the collection efficiency of the ring electrode, and ID the disk current, and the percentage of evolved peroxide can obtained as %XH2O2 = XH2O2 × 100%.28,29 Figure 4b provides the data from RRDE experiments at 500 rpm and demonstrates that the fraction of evolved peroxide is low (