Oxygen Reduction Properties of Bifunctional α-Manganese Oxide Electrocatalysts in Aqueous and...

Published: October 07, 2011

r 2011 American Chemical Society 22009 dx.doi.org/10.1021/jp2055443 | J. Phys. Chem. C 2011, 115, 22009–22017

ARTICLE

pubs.acs.org/JPCC

Oxygen Reduction Properties of Bifunctional α-Manganese OxideElectrocatalysts in Aqueous and Organic ElectrolytesE. M. Benbow, S. P. Kelly, L. Zhao, J. W. Reutenauer, and S. L. Suib*

Department of Chemistry, University of Connecticut, 55 North Eagleville Road Unit 3060, Storrs, Connecticut 06269, United States

’ INTRODUCTION

Environmentally friendly, high-energy dense power sources,such as metal�air batteries, are of extreme interest due to therapidly expanding market of mobile electronics and electricvehicles. Lithium is a promising material for metal�air batteriesdue to being extremely lightweight and due to its high theoreticalenergy density (13 kWh g�1). When electronically coupled tooxygen via a carbon-catalyzed air cathode, a high theoreticalspecific energy (5200Wh kg�1) and discharge capacities (∼4900mAh g�1) can be achieved, albeit at low current densities(0.1 mA cm�2).1,2 However, achieving a rechargeable Li�air cellhas proven to be problematic due to numerous issues, one being cellpolarization.3 Cell polarization for lithium�air batteries is aconsequence of the large gap between the charging and dischar-ging voltage. Reducing the cell polarization by lowering thecharge voltage via catalysts is therefore of great importance forincreasing efficiency and lifetimes of secondary Li�air cells.

In nonaqueous electrolytes, the discharge products of alithium�air battery accumulate within the pores of the aircathode, eventually restricting oxygen diffusion and preventingfurther discharge.1 In addition, the oxygen reduction mechanismin nonaqueous electrolytes was not fully understood with variousmechanisms proposed.4,5 Electrolyte decomposition in the pre-sence of oxygen reduction intermediates is an additional con-tributing factor further complicating matters.6,7 Recent in situstudies have identified two separate pathways for oxygen reduc-tion and oxygen evolution.8 An ECC mechanism [eqs 1�3] forthe oxygen reduction reaction (ORR) in nonaqueous electro-lytes is observed, while a direct pathway [eq 4] is observed for the

oxygen evolution reaction (OER).8 An ECC mechanism is anelectrochemical reduction [eq 1] step followed by two rapidchemical steps [eqs 2,3]. These chemical steps make properidentification of the reduction pathway by cyclic voltammetrydifficult.

O2ðgÞ þ e� f O2�ðsolÞ ð1Þ

O2�ðsolÞ þ LiþðsolÞ f LiO2ðsolÞ ð2Þ

2LiO2ðsolÞ f 2Li2O2ðsÞ ð3Þ

Li2O2ðsÞ f 2LiðsÞ þ O2ðgÞ ð4ÞCatalysts are extremely important for lithium�air batteries as

they have been shown to affect many performance characteristicssuch as discharge/charge potentials, capacity, and recharge-ability.1�3,9�12 The extent to which the catalysts reduce oxygenhas been shown to have great importance to the functionality ofthe Li-air battery. A good example of this is displayed with Pt andAu nanoparticle catalysts, which favor the four-electron reduc-tion product (OH�) and the two-electron reduction product(HO2

�), respectively.11 The Au catalysts have better dischargeperformance, whereas the Pt catalysts have better charge per-formance in Li�air cells.11 Synthesizing Au/Ptbifunctional

Received: June 13, 2011Revised: October 6, 2011

ABSTRACT: Several different preparation and morphologiesof α-manganese oxide catalysts were synthesized, and theirphysical properties were characterized. These catalysts werealso characterized by electrochemical means (RDV and CV) foroxygen reduction in both aqueous and organic electrolytes. Thesolvent-free preparation yielded catalysts with ideal physicalproperties: low average oxidation state (3.73) of manganese,small crystallite size (8.1 nm), small particle size (d = 10 nm,l = 30�100 nm), high surface area (144m2 g�1), and pore volume(0.470 cm3/g). This catalyst displayed the highest activity in both aqueous (iL = 3.66( 0.12 mA cm�2, i0 = 1.0� 10�5 mA cm�2)and organic electrolytes (iL = 2.25( 0.15 mA cm�2, αn = 0.51). Doping the solvent-free preparation with Ni2+ slightly improved itsoxygen reduction capabilities in aqueous (iL = 3.78( 0.15 mA cm�2, i0 = 2.3� 10�5 mA cm�2) and organic electrolytes (iL = 2.43(0.17 mA cm�2, αn = 0.47). A large increase in the ability of the catalysts to decompose hydrogen peroxide was observed (2.7( 0.26vs 0.53 ( 0.03 s�1 g�1), which is a good indicator of enhanced bifunctionality for application in secondary lithium air batteries.Stability tests in the presence of Li2O2 at elevated temperatures (100 �C) indicate a decrease in catalyst crystallinity as exposure timeis increased.

22010 dx.doi.org/10.1021/jp2055443 |J. Phys. Chem. C 2011, 115, 22009–22017

The Journal of Physical Chemistry C ARTICLE

nanoparticle catalysts combines the advantages of both the Auand the Pt catalysts.12 The Au/Pt catalysts likely possess anoxygen reduction pathway similar to those observed for manga-nese oxides resulting in the formation of two- and four-electronreduction products.

Oxygen reduction reaction mechanisms in aqueous electro-lytes on manganese oxide surfaces have been studied by Maoet al. and have identified a series pathway by which oxygen isreduced to the peroxide (two-electron reduction) [eq 5] andthen is subsequently reduced to the hydroxide (four-electronreduction) [eq 6] or undergoes a catalytic disproportionation ofperoxide intermediate [eq 7] to give a pseudo four-electronreduction of oxygen.13 Nickel doping of manganese oxides hasbeen shown to enhance the four-electron reduction of oxygen andenhance the decomposition reaction of hydrogen peroxide.14,15

Recent work by Giordaniet al. shows a strong correlation be-tween the ability of the catalysts to decompose hydrogenperoxide and the charging potential of Li�air cells.3 Therefore,nickel doping should increase the bifunctionality of manganeseoxide catalysts similar to that observed for Au/Pt catalysts at asignificant cost reduction.

O2ðgÞ þ e� f O2�ðaqÞ þ H2OðaqÞ þ e� f HO2

�ðaqÞ

þOH�ðaqÞ ð5Þ

HO2�ðaqÞ þ H2OðaqÞ þ 2e� f 3OH�

ðaqÞ ð6Þ

2HO2�ðaqÞ f 2OH�

ðaqÞ þ O2ðgÞ ð7ÞManganese oxides such as MnO2 (α, β, δ, EMD), Mn2O3,

Mn3O4, Mn5O8, and MnO(OH) have been widely used aselectrocatalysts for the reduction of oxygen.2,3,9,10,14�22 Thesehighly active, thermally stable catalysts can be mixed valent,possess high surface areas, and have environmental and econom-ic benefits. Of all of the manganese oxide polymorphs, the αphase appears to be the most active electrocatalyst for oxygenreduction, especially in Li�air applications.9,17,21 The higheractivity of the α phase has been reported to be related to thehigher number of defects and �OH groups present in thestructure, which facilitates O2 absorption and dissociation ofO�O bonds.23 It has been shown that different preparationmethods can yield vast differences in physical properties such asparticle size, exposed crystal facets morphology, surface area, andaverage metal oxidation state,17,21,24�26 all of which play animportant role in their catalytic efficacy.

Herein, we present a study of different preparations of α-MnO2 catalysts in an attempt to identify a simple and scalablepreparation method, which results in the most active catalyst forsecondary Li�air batteries. Such physical properties examinedinclude particle size, morphology, surface area, and averageoxidation state (AOS) in an effort to determine key physicalproperties for optimal electrocatalytic activity. Doping the man-ganese oxide framework with Ni2+ ions enhances the decom-position of hydrogen peroxide and has been shown to favor thefour-electron reduction pathway of oxygen; therefore, we studiedNi2+ doping of the most active preparation for enhancingbifunctionality.14,15 The catalysts have been characterized byX-ray diffraction, BET, SEM, EDAX, CV, RDV, and hydrogenperoxide rate constant decomposition studies. The electroche-mical studies were performed in both aqueous and organicelectrolytes.

’EXPERIMENTAL SECTION

Synthesis.The reflux and solvent-free preparation ofα-MnO2

catalysts were synthesized according to the literature.24�26 Themanganese oxide nanowires made at low temperature vide infrawere synthesized by the addition of K2SO4, K2S2O8, andMnSO4 3H2O to 200 mL of DDW in a 250 mL Erlenmeyer flaskat 50 �C. The solution was continuously stirred for 24 h in an oilbath at 50 �C. The manganese oxide dendrites were synthesizedunder similar conditions except 2 mL of concentrated H2SO4

was added to the 200 mL of DDW before addition of theprecursors. The black precipitate was then filtered and washedwith deionized water. The manganese oxide precipitate was thendried in an oven at 80 �C overnight.Characterization. Powder X-ray diffraction (XRD) patterns

were collected using a Scintag PDS 2000 diffractometer with aCu Kα X-ray radiation (λ = 0.15406 nm) source operating at45 kV and 30 mA. Field-emission scanning electron microscopy(FESEM) images were obtained using a Zeiss DSM 982 Geminifield-emission scanning electron microscope with a Schottkyemitter. Samples were suspended in ethanol and dispersed onAuPd-coated silicon chips mounted on stainless-steel sampleholders using double-sided carbon tape.Surface areas of the catalysts were measured via nitrogen

sorption at 77K using a Micrometrics ASAP 2010 acceleratedsurface area system. The samples were degassed at 150 �C for24 h. The average oxidation state (AOS) ofmanganese was deter-mined via a potentiometric titration.27

Cyclic voltammetry (CV) was conducted with a computer-controlled BAS 100B electrochemical workstation to test theelectrocatalytic activity of the manganese oxides. A gas-diffusionelectrode was prepared to check the ability to reduce oxygen. Thecell setup consisted of a working electrode with a slurry (10mg ofcatalyst, 10 mg of carbon (Vulcan XC-72), 5 drops of PTFE or3.5 mg of catalyst, 10 mg of carbon (Vulcan XC-72), 3 drops ofPTFE) cast on a pyrolytic graphite electrode (PG, AdvancedCeramics, A = 0.125 cm2), a saturated calomel reference elec-trode (SCE), and a Pt counter electrode in 1.0 M KOH solution.A similar cell setup was used for testing in organic electrolytes(1 M LiPF6 1:1:1 EC:DEC:DMC) except an Ag/AgCl electrodereference electrode was substituted for the SCE referenceelectrode. The working electrode was cycled repeatedly until areproducible signal was obtained. Data analysis was performedafter obtaining reproducible cyclic voltammetry scans usingmultiple working electrodes with similar preparations. Thesample closest to the mean was used for further data analysis.To investigate oxygen reduction in aqueous electrolytes,

rotating disk voltammetry (RDV) (Pine Instruments) wasperformed at 2500 rpm and a scan rate of 5 mV s�1. The cellsetup was the same as mentioned above. The PG electrode wasabraded with 400 grit SiC paper, and was then cast with 20 μL ofthe slurry, which was allowed to dry overnight. O2 was purged for30 min before each scan. All experiments were conducted atambient temperature (21 �C). All scans were backgroundcorrected under N2 flow, and analysis was performed afterreproducible scans were obtained. The sample closest to themean for used for further data analysis.Hydrogen peroxide decomposition studies were performed at

constant temperature (21 ( 1 �C) and volume under vigorousstirring. A typical experiment consisted of 10 mg of catalystssonicated in 10 mL of DDW to which 20 mL of stock hydrogenperoxide (30% v/v) was injected into a sealed 1 L round-bottom



flask submerged in a water bath. The rate constants werecalculated from the change in pressure (Torr), measured by aMKS pressure transducer with a sensitivity of 0.1 Torr, withrespect to time (s). A maximum pressure of 900 Torr was used asthe end point for the reaction, and measurements began 10 or15 s after injecting the hydrogen peroxide depending on the catalyst.

’RESULTS

Two-Dimensional Nanostructures. The X-ray diffractionpatterns (Figure 1) for all preparations can be indexed to thetetragonal cryptomelane structure also known as α-MnO2; thelattice parameters are summarized in Table 1. There is nosignificant difference in the lattice parameters with respect tothe different preparations. Cryptomelane is comprised of man-ganese oxide octahedra (MnO6

�) forming a 2 � 2 tunnelstructure. This manganese oxide framework is negativelycharged, which requires charge balancing by potassium cationswithin the tunnels (Figure 2). These manganese oxide nanoma-terials belong to a class of microporous transition metal oxidescalled octahedral molecular sieves (OMS). OMS materialspossess high surface area, are semiconducting, are mixed-valentframeworks with diverse tunnel structures (i.e., 1 � 1, 1 � 2,etc.), and have a wide variety of industrial applications.Upon examining the XRD patterns, one can already begin to

see differences in these materials according to the differentpreparations. The most significant difference can be observedfor the solvent-free preparation, which has very weak and broadpeaks. This is consistent with the literature as the solvent-freepreparation is known to produce very small particles (d = 10 nm,l = 30�100 nm).24 To compare differences in crystallite size, the

Scherrer equation was employed, which relates the broadening ofthe peaks to the crystallite size. According to the calculations ofcrystallite size from the [110] planes (Table 1), the differentpreparations rank in crystallite size nanorodSF < nanowireLT <nanorodR.FESEM images (Figure 3) of the synthesized manganese

oxide nanomaterials show the nanorod/nanowire morphologiesachieved by the different preparations. The reflux preparation(Figure 3a) results in a nanorod morphology consisting of verysmall and very large particles (d= 15�30 nm, l = 400�1000 nm).The low temperature (50 �C) preparation (Figure 3b and c)results in a nanowire morphology with a very uniform particlesize (d = 15 nm, l = 500 nm), which tends to aggregate in largeinterwoven bundles. The solvent-free preparation (Figure 3d)yields a nanorod morphology with very small particles, which canmake estimating particles size difficult using FESEM (d = 10 nm,l = 30�100 nm). The solvent-free preparation yields a particlesize very close to the ideal particle size (d = 10 nm, l = 200 nm)proposed by D�ebartet al.6

The mixed valency of α-MnO2 results in Mn3+, Mn4+, andpossible Mn2+ ions in the framework. The presence of chargebalancing tunnel ions in addition to oxygen vacancies is respon-sible for the mixed valency observed in OMS catalysts, whichis partly responsible for their high catalytic activity. The

Table 1. Summary of α-MnO2 (Cryptomelane) Catalysts’ Physical Properties

sample

lattice parameter

a = b (nm)

lattice parameter

c (nm)

crystallite size

(nm) [110]

surface area

(m2 g�1)

pore size

(nm)

pore volume

(cm3/g)

nanorodR 0.9810 0.2851 9.5 95 18.4 0.438

nanowireLT 0.9814 0.2850 8.9 92 18.7 0.428

dendriteLT 0.9803 0.2849 10.0 50 9.4 0.118

nanorodSF 0.9817 0.2853 8.1 144 13.0 0.470

Ni-nanorodSF 0.9814 0.2855 7.5 140 9.3 0.325

Figure 2. Structure of α-MnO2 (cryptomelane) octahedral molecularsieve catalysts with tunnel cations omitted.

Figure 1. X-ray powder diffraction patterns ofα-MnO2 (cryptomelane)catalysts. Ni-nanorodSF, nanorodSF, nanorodR, dendriteLT, and nano-wireLT, diffraction patterns correlate to their preparation from top tobottom, respectively.



solvent-free preparation has the lowest observed AOS (3.73)and increases for the nanowire (3.85) and the reflux (3.91)preparations.24,26 The low AOS of the solvent-free preparationcannot be fully explained by the charge balancing provided by thetunnel cation (potassium). The lowest achievable AOS providedsolely by potassium ions is 3.875 resulting from∼1/2 occupationof the tunnel site (K0.5Mn8O16) due to size constraints related tothe ionic radii of potassium. A high concentration of oxygendefects is likely present, further lowering the average oxidationstate.28

BET surface areas (Figure 4) of both the reflux (95 m2 g�1) andlow temperature (91 m2 g�1) preparation are nearly identical to thesolvent-free (144 m2 g�1) preparation possessing the highest surfacearea. All preparations yield materials with an average pore size within

the range designated for mesoporous materials (2�50 nm) and issummarized in Table 1. The solvent-free preparation possesses thesmallest pore size (13.0 nm) and increases for the reflux (18.4 nm)and nanowire (18.7 nm) preparations. The solvent-free preparationyields the largest pore volume of 0.470 cm3/g, with the reflux(0.438 cm3/g) and nanowire (0.428 cm3/g) preparations nearlyidentical.Three-Dimensional Nanostructure. To examine a potential

relationship between morphology and the catalytic activity, amanganese oxide catalyst with three-dimensional urchin-likemorphology was synthesized (Figure 3). These three-dimen-sional catalysts consist of nanorod crystallites branching from acentral core having diameters of 500�1000 nm. These manga-nese oxide dendrites have a BET surface area (Table 1, Figure 4)

Figure 3. EM images depicting morphology of α-MnO2 catalysts: (a) reflux, (b) nanowire, (c) dendrite, and (d) solvent-free.

Figure 4. BET isotherms comparing α-MnO2 catalysts (nanowireLTand nanorodR offset for clarity).

Figure 5. Background corrected oxygen reduction RDV scans of purecryptomelane catalysts in 1 M KOH vs SCE (v = 5 mV s�1).



of 50m2 g�1, an average pore size of 9.4 nm, and a pore volume of0.118 cm3/g. The XRD pattern (Figure 1) can also be indexed tothe cryptomelane structure with average crystallite size of 10 nmas determined by the Scherrer equation. Previously, thesemanganese oxide dendrites with three-dimensional morpholo-gies were hydrothermally synthesized using dichromates, a mildoxidizing agent.29,30 In this method, the dendritic structure wasachieved by a controlled oxidation of manganese(II) ions byperoxydisulfate (S2O8

2�), a strong oxidizing agent. This controlof the oxidation process was achieved by a small addition (2 mL)of concentrated sulfuric acid.Ni2+Doping.The solvent-free preparation appears to have the

best combination of physical properties of all of the catalystsprepared. Briefly, this preparation has the smallest crystallite size(8.1 nm), highest surface area (144 m2 g�1), smallest pore size(13.0 nm), and lowest AOS (3.73). Therefore, we selected thispreparation to examine framework doping with Ni2+, which hadbeen previously shown to enhance hydrogen peroxide decom-position and the four-electron reduction pathway.14,15 Dopinglevels of ∼1:40 Ni:Mn (Ni-nanorodSF) ratio were synthesizedand confirmed by EDAX. FESEM images show no noticeablechanges in morphology, and the XRD patterns (Figure 1) do notindicate the presence of any secondary phase. The measuredBET surface areas were nearly identical to the pure version with140 m2 g�1. However, doping with Ni2+ decreased the averagepore size (9.3 nm) and pore volume (0.325 cm3/g).RDV in 1M KOH.Rotating disk voltammetry (RDV) was used

to elucidate the differences in the ability of catalysts to reduceoxygen in an aqueous environment. The background corrected

results for the manganese oxide catalysts are shown (Figures 5, 6)and are summarized in Table 2. The limiting current densities forthe nanorodR (2.63 ( 0.09 mA cm�2) and dendriteLT (2.65 (0.13 mA cm�2) catalysts are significantly lower than those for thenanowireLT (3.59 ( 0.15 mA cm�2), nanorodSF (3.66 ( 0.12mA cm�2), and Ni-nanorodSF (3.78( 0.15 mA cm�2) catalysts.

log ik ¼ log io þ ð1� αnF=RTÞη ð8Þ

The Tafel equation [eq 8] was employed to investigate the linearrelationship between the rate of the electrochemical reaction andthe overpotential (η), where ik is the kinetic current, io is the exchangecurrent, R is the ideal gas constant (8.314 J mol�1 K�1), α is thetransfer coefficient, n is the number of electrons transferred, T is thetemperature in Kelvin, and F is the Faraday constant (9.6485� 104

Figure 6. Background corrected oxygen reduction RDV scans (top) ofnanorodSF and Ni-nanorodSF in 1 M KOH vs SCE (v = 5 mV s�1) andthe corresponding Tafel plots (bottom).

Table 2. Summary of α-MnO2 (Cryptomelane) Catalysts’Electrochemical Properties in 1 M KOH

sample

limiting current

(iL) (mA cm�2)

Tafel slope

(b) (mV dec�1)

exchange current

(io) (A cm�2)

nanorodR 2.63 ( 0.09 114 8.1 � 10�6

nanowireLT 3.59 ( 0.15 114 9.1 � 10�6

dendriteLT 2.65 ( 0.13 110 3.6 � 10�6

nanorodSF 3.66 ( 0.12 113 1.0 � 10�5

Ni-nanorodSF 3.78 ( 0.15 112 2.3 � 10�5

Figure 7. IR corrected oxygen reduction CV scans vs Ag/AgCl of pure(1:1 wt % catalyst:carbon) cryptomelane catalysts (top) and carbon,nanorodSF, and Ni-nanorodSF (bottom) in 1 M LiPF6 1:1:1 EC:DEC:DMC (v = 25 mV s�1).



Cmol�1). Two representative Tafel plots are shown in Figure 6 forthe nanorodSF and Ni-nanorodSF catalysts. The Tafel slopes (ω =2000 rpm) were calculated (Table 2) at high overpotentials, and allcatalysts show similar values near the ideal value of 120mV dec�1.31

To help further distinguish each catalyst, the exchange currentdensities (Table 2), which are measures of kinetics, were extra-polated from the Tafel plots and show the following trend:dendriteLT < nanorodR < nanowireLT < nanorodSF <Ni-nanorodSF.Because io,iL, the overpotential with respect to oxygen reduction isdue to the activation of charge transfer not mass transport.31

CV in 1 M LiPF6 1:1:1 EC:DEC:DMC. To test for potentialapplication in nonaqueous Li�air batteries, the oxygen reductionand oxygen evolution properties were electrochemically investi-gated in an organic electrolyte (1M LiPF6 1:1:1 EC:DEC:DMC).

CV scans are shown in Figure 7 (1:1 wt % catalyst:carbon) andFigure 8 (0.35wt% catalyst:carbon) comparing the catalytic abilityof the cryptomelane catalysts at different levels of catalyst loadingwith the results summarized in Table 3. The observed peak poten-tials appear within the same potential window previously observedby Ngalaet al.32 With respect to the limiting current density, theactivity of each catalysts at high catalyst loading for oxygenreduction is compared, nanorodR , dendriteLT < nanowireLT <nanorodSF < Ni-nanorodSF < carbon. At low catalyst loading, an∼50 mV increase in the peak potential is observed with acorresponding increase in the limiting current density with thefollowing trend: dendriteLT < nanorodR < nanowireLT < carbon <nanorodSF < Ni-nanorodSF. These results show a clear decrease inactivity due to catalyst overloading.To further elucidate the differences in the pure cryptomelane

catalysts, the reaction kinetics were investigated using transfer co-efficients (αn). The transfer coefficient was estimated (Table 3)using the difference between the peak potential and the half wavepotential [eq 9],31 where Ep is the peak potential, Ep2 is the halfwave potential, R is the ideal gas constant (8.314 J mol�1 K�1),T is the temperature in Kelvin, and F is the Faraday constant(9.6485 � 104 C mol�1). A comparison of the transfer coeffi-cients at high catalysts loading gives the following trend: nanor-odR < nanowireLT < nanorodSF < Ni-nanorodSF < dendriteLT <carbon. A significant improvement in the transfer coefficient isobserved when the catalyst loading is decreased with the catalystyielding values close to the ideal value of 0.5. A slight inflection inthe CV scans [Figures 7,8] is observed at∼600 mV in the regionwhere oxygen evolution would be expected to be observed. How-ever, the slight difference in the CV scans for each catalyst isnot obvious; therefore, no trend with respect to oxygen evolutioncan be established. This lack of observed oxygen evolution reac-tion may be attributed to reactions between the solvent andoxygen reduction intermediates in nonaqueous electrolytes pre-venting the generation of the desired lithium peroxide dischargeproduct.6�8

Ep � Ep2 ¼ 1:857RT=αnF ¼ 47:7 mV=αn ð9Þ

Hydrogen Peroxide Decomposition. Determination of therate constants for hydrogen peroxide decomposition was per-formed due to the lack of clear evidence distinguishing the abilityof different catalysts for the oxygen evolution reaction. The rateconstants for hydrogen peroxide decomposition are of interest asthey show a strong correlation to the charging voltage of the Li�airbattery.3 The results for these tests can be seen in Figure 9 andare summarized in Table 3. Comparison to the values presentedby Giordaniet al. is not possible due to unknown experimentalconditions.3 The solvent-free preparation yields the highest activity

Figure 8. IR corrected oxygen reduction CV scans vs Ag/AgCl of pure(0.35:1 wt % catalyst:carbon) cryptomelane catalysts (top) and carbon,nanorodSF, and Ni-nanorodSF (bottom) in 1 M LiPF6 1:1:1 EC:DEC:DMC (v = 25 mV s�1).

Table 3. Summary of α-MnO2 (Cryptomelane) Catalysts’ Electrochemical Properties in Organic Electrolyte and HydrogenPeroxide Rate Decomposition Studiesa

sample limiting current (iL)org (mA cm�2) HL/LL E (mv) vs Ag/AgCl transfer coefficient (αn) HL/LL k (s�1 g�1)

carbon 1.99 ( 0.07 �206 0.48 N/A

nanorodR 1.20 ( 0.13/1.61 ( 0.11 �382/�244 0.24/0.48 0.14 ( 0.03

nanowireLT 1.64 ( 0.08/1.80 ( 0.12 �297/�241 0.30/0.46 0.20 ( 0.03

dendriteLT 1.37 ( 0.09/1.48 ( 0.05 �254/�214 0.42/0.56 0.35 ( 0.03

nanorodSF 1.51 ( 0.04/2.25 ( 0.15 �313/�241 0.33/0.51 0.53 ( 0.03

Ni-nanorodSF 1.71 ( 0.09/2.43 ( 0.17 �308/�270 0.35/0.47 2.7 ( 0.26aHL = 1:1 catalyst:carbon, LL = 0.35:1 catalyst:carbon.



for hydrogen peroxide decomposition (0.53 ( 0.03 s�1 g�1)among the pure catalysts. A large increase in the rate constants wasobserved upon Ni2+ doping with Ni-nanorodSF, yielding thehighest value (2.7 ( 0.26 s�1 g�1). The standard deviation inthese measurements increased for the doped catalysts and is likelyrelated to a slight increase in gas temperature above the solutiondue to the extreme violence of the reaction. All tests wereperformed with the reaction vessel submerged in a water bath inan attempt to minimize the temperature fluctuations from theexothermic decomposition of hydrogen peroxide.After a long exposure (∼24 h) to the stock hydrogen peroxide

solution, almost all of the cryptomelane catalysts decomposed.This instability toward the peroxide ion may be a concern forsecondary Li�air batteries, because the catalysts would beexposed to lithium peroxide at high overpotentials (∼4 V).The loss of energy due to overpotential is released as thermalenergy; therefore, we exposed cryptomelane catalysts to lithiumperoxide at elevated temperatures (100 �C) in an effort tosimulate overpotential. The catalysts were ground together withlithium peroxide in an agate mortar and pestle and then subjectedto 24 h heat treatments (Figure 10). The catalysts where thenwashed/filtered, and their crystallinity was examined using X-raydiffraction. From the X-ray diffraction data, one can clearly see adecrease in peak height as exposure time is increased, indicating a

decrease in crystallinity of the cryptomelane catalyst. Theseresults would indicate a potential hindrance in using manganeseoxide catalysts in secondary Li�air batteries due to the repeatedexposure to lithium peroxide upon cycling.

’DISCUSSION

Of the two-dimensional pure manganese oxide catalysts, thenanorodSF preparation appears to have the best combination ofphysical and electrochemical properties. This preparation yieldsthe smallest crystallite size (8.1 nm), lowest AOS (3.73), highestsurface area (144m2 g�1), and pore volume (0.470 cm3/g) for allmaterials we studied. These properties produce a well-dispersedcatalyst that has a high degree of MnIII/MnIV couples ineasily accessible mesopores, resulting in high electrochemical activity.In 1MKOH, this catalysts possesses the highest limiting (3.66(0.12 mA cm�2) and exchange (1.0 � 10�5A cm�2) currentdensities. In the organic electrolyte at low catalyst loading, thehighest limiting current (2.25 ( 0.15 mA cm�2) was observedwith a transfer coefficient (0.51) near the ideal values of 0.5.

The most common preparation (nanorodR) of two-dimen-sional α-manganese oxide catalysts possesses the lowest limiting(2.63 ( 0.09 mA cm�2) and exchange (8.1 � 10�6A cm�2)current densities in 1 M KOH. The nanowireLT catalyst pos-sesses a higher limiting (3.59 ( 0.15 mA cm�2) and exchangecurrent density (9.1 � 10�6 A cm�2). In the organic electrolyteat low catalyst loading, the nanowireLT catalyst yields a higherlimiting current density (1.80 ( 0.12 vs 1.61 ( 0.11 mA cm�2)with similar transfer coefficients (0.46 vs 0.48), when comparedto the nanordR catalyst. These two catalysts have similar surfaceareas (96 vs 91m2 g�1) and pore volume (0.438 vs 0.428 cm3/g);therefore, these differences are likely related to the smallercrystallite size (8.9 vs 9.5 nm), lower AOS (3.85 vs 3.91), andmore uniform particle size distribution for the nanowireLTcatalyst. These slight differences can allow for a better dispersionof a more active catalysts resulting in decreased Ohmic resistanceof the electrode composite.

The dendriteLT catalyst has poor performance in 1M KOH asshown by the low limiting (2.65 ( 0.13 mA cm�2) exchange(3.6� 10�6A cm�2) current densities. In the organic electrolyte atlow loading, the dendriteLT catalyst did not observe a significantincrease in the limiting current density (1.48 ( 0.05 mA cm�2)as the other catalysts, yet still has a high transfer coefficient(0.56). These results are interesting considering this catalyst hasa large three-dimensional particle size (500�1000 nm), thelargest crystallite size (10.0 nm), the lowest BET surface area(50 m2 g�1), and pore volume (0.118 cm3/g). Comparing theresults at high and low catalyst loading may give insight intothis unique behavior. At high catalyst loading, all of the two-dimensional catalysts experienced sluggish kinetics evident bytheir low transfer coefficients, whereas the dendriteLT has a hightransfer coefficient (0.42). This is likely due to catalyst over-loading and therefore results in decreased electrode conductivityand accessibility of the pores of carbon, which with respect toperformance in organic electrolytes is critical. The 3-D nano-structure may create a more porous electrode and thereforeincreases accessibility of the interior of the electrodes. However,due to its large particle size, its catalytic performance is inhibited,such as with the nanorodR catalysts.

The Ni-nanorodSF catalysts show interesting electrochemicalproperties in both 1 M KOH and the organic electrolyte. In1 M KOH, the Ni-nanorodSF shows a limiting current density

Figure 9. Rate constant studies tracking the evolution of oxygen fromhydrogen peroxide using different catalysts. All experiments wereconducted at constant volume and temperature (20 �C).

Figure 10. X-ray diffraction patterns of reaction between cryptomelanecatalyst and lithium peroxide: (top) pure, (middle) 24 h, (bottom) 48 h.



(3.78(0.15mAcm�2) similar to that of thepure nanorodSF catalyst;however, there was an increase in the exchange current density(2.3 � 10�5 mA cm�2). Because they possess similar limitingcurrents, doping with Ni2+ maintains similar activity whiledisrupting the natural AOS of the nanorodSF, resulting in adecrease in the number ofMnIII/MnIV couples. However, dopingdid increase the exchange current density, which could be a resultof a decrease in the charge transfer resistance that is dominantwhen io,iL. This decrease in charge transfer resistance mayresult in decreased electrode polarization, which hinders second-ary Li�air batteries. In the organic electrolyte at low loading, theNi-nanorodSF catalyst possesses the highest limiting current(2.43 ( 0.17 mA cm�2) and a transfer coefficient of 0.47.

Examination of the rate constants for hydrogen peroxidedecomposition shows that for the pure catalysts the nanorodSFshows the greatest ability (0.53 ( 0.03 s�1 g�1) to decomposehydrogen peroxide. This value is significantly higher than thoseof the nanowireLT (0.20( 0.03 s�1 g�1), nanorodR (0.14( 0.03s�1 g�1), and dendriteLT (0.35( 0.03 s�1 g�1). The high surfacearea, small particle size, and low AOS all likely contribute to thehigh catalytic ability of nanorodSF to decompose hydrogenperoxide. Upon doping with Ni2+, a large increase in the rateconstant was observed for Ni-nanorodSF (2.7 ( 0.26 s�1 g�1).

’CONCLUSION

The pure nanorodSF catalyst shows the highest electrochemi-cal activity in both aqueous and organic electrolytes. Its uniquecombination of physical properties such as small crystallite size(8.1 nm), lowest AOS (3.73), high surface area (144 m2 g�1),and pore volume (0.470 cm3/g) all contribute to its highelectrochemical activity. The nanorodSF catalyst possessed thehighest limiting (3.66 ( 0.12 mA cm�2) and exchange current(1.0 � 10�5 A cm�2) densities in 1 M KOH. In organicelectrolyte at low catalyst loading, its transfer coefficient (0.51)is ideal with a high limiting current density (2.25( 0.15mA cm�2).This catalyst displayed the highest decomposition rate constant(0.53 ( 0.03 s�1 g�1) of all of the pure catalysts, which is a goodindicator of its bifunctionalilty. The data collected suggest that lowaverage oxidations states of manganese is a critical parameter forbifunctional catalysts.

Doping the nanorodSF catalyst with Ni2+ ions is shown toenhance the electrochemical activity in both aqueous and organicelectrolyte. This increase was observed despite decreases in poresize (13.0 vs 9.3 nm) and pore volume (0.470 vs 0.325 cm3/g).Upon doping, a significant increase in the exchange currentdensity (2.3 � 10�5 A cm�2, 1 M KOH) was observed, whichcould be indicative of decreased charge transfer resistance, whichmay result in decreased electrode polarization. In the organicelectrolyte, the limiting current density increased slightly (2.43(0.17 mA cm�2), while the transfer coefficient was slightly less(0.47), however still very close to the ideal value. A verysignificant increase (2.7 ( 0.26 s�1 g�1) in the decompositionof hydrogen peroxide was observed, indicating that doping ofmanganese oxide catalysts with Ni2+ ions enhances bifunction-ality and will be beneficial for decreasing the charging potentialfor Li�air batteries if suitable electrolytes are identified.

’AUTHOR INFORMATION

Corresponding Author*E-mail: [email protected].

’ACKNOWLEDGMENT

We would like to thank the Intelligence Community Post-doctoral Research Fellowship Program under the GeospatialIntelligence Agency program for support of this work. We wouldlike to thank Dr. Ahbay Vaze and Hui Huang for helpfuldiscussions. We would also like to thank J. F. Rusling for use ofthe Pine Instruments RDV apparatus.

’REFERENCES

(1) Abraham, K. M.; Jiang, Z. J. Electrochem. Soc. 1996, 143, 1–5.(2) Dobley, A.; Morein, C.; Roark, R. ECS Trans. 2008, 3,

83–88.(3) Giordani, V.; Freunberger, S. A.; Bruce, P. G.; Tarascon, J.-M.;

Larcher, D. Electrochem. Solid-State Lett. 2010, 13, A180–183.(4) Laoire, C. O.; Mukerjee, S.; Abraham, K. M.; Plichta, E. J.;

Hendrickson, M. A. J. Phys. Chem. C 2009, 113, 20127–20134.(5) Laoire, C. O.; Mukerjee, S.; Abraham, K. M.; Plichta, E. J.;

Hendrickson, M. A. J. Phys. Chem. C 2010, 114, 9178–9186.(6) Freunberger, S. A.; Chen, Y.; Drewett, N. E.; Hardwich, L. J.;

Barde, F.; Bruce, P. G. Angew. Chem., Int. Ed. 2011, 50, 8609–8613.(7) Freunberger, S. A.; Chen, Y.; Peng, Z.; Griffin, J. M.; Hardwick,

L. J.; Barde, F.; Novak, P.; Bruce, P. G. J. Am. Chem. Soc. 2011,133, 8040–8047.

(8) Peng, Z.; Freunberger, S. A.; Hardwick, L. J.; Chen, Y.; Giordani,V.; Barde, F.; Novak; Graham, D.; Tarascon, J.-M.; Bruce, P. G. Angew.Chem. 2011, 123, 6475–6479.

(9) Debart, A.; Paterson, A. J.; Bao, J.; Bruce, P. G. Angew. Chem., Int.Ed. 2008, 47, 4521–4524.

(10) Debart, A.; Bao, J.; Armstrong, G.; Bruce, P. G. J. Power Sources2007, 174, 1177–1182.

(11) Lu, Y.-C.; Gasteiger, H. A.; Parent, M. C.; Chiloyan, V.; Shao-Horn, Y. Electrochem. Solid-State Lett. 2010, 13, A69–72.

(12) Lu, Y.-C.; Xu, Z.; Gasteiger, H. A.; Chen, S.; Hamad-Schifferli,K.; Shao-Horn, Y. J. Am. Chem. Soc. 2010, 132, 12170–12171.

(13) Mao, L.; Zhang, D.; Sotomura, T.; Nakatsu, K.; Koshiba, N.;Ohsaka, T. Electrochim. Acta 2003, 48, 1015.

(14) Roche, I.; Chainet, E.; Chatenet, M.; Vondrak, J. J. Appl.Electrochem. 2008, 38, 1195–1201.

(15) Zhou, H.; Shen, Y. F.; Wang, J. Y.; Chen, X.; O’Young, C.-L.;Suib, S. L. J. Catal. 1998, 176, 321–328.

(16) Brock, S. L.; Duan, N.; Tian, Z. R.; Giraldo, O.; Zhou, H.; Suib,S. L. Chem. Mater. 1998, 10, 2619.

(17) Xiao, W.; Wang, D.; Lou, X. W. J. Phys. Chem. C 2010,114, 1694–1700.

(18) Lima, F. H. B.; Calegaro, M. L.; Ticianelli, E. A. J. Electroanal.Chem. 2006, 590, 152.

(19) Lima, F. H. B.; Calegaro, M. L.; Ticianelli, E. A. Electrochim.Acta 2007, 52, 3732.

(20) Cao, Y. L.; Yang, H. Z.; Ai, X. P.; Xiao, L. F. J. Electroanal. Chem.2003, 557, 127.

(21) Cheng, F.; Su, Y.; Liang, J.; Tao, Z.; Chen, J.Chem.Mater. 2010,22, 898–905.

(22) Crisostomo, V. M. B.; Ngala, J. K.; Alia, S.; Dobley, A.; Morein,C.; Chen, C. H.; Shen, X. F.; Suib, S. L. Chem. Mater. 2007, 19, 1832.

(23) Thackeray, M. M. Prog. Solid State Chem. 1997, 25, 1–71.(24) Ding, Y.; Shen, X.; Sithambaram, S.; Gomez, S.; Kumar, R.;

Crisostomo, V. M. B.; Suib, S. L.; Aindow, M. Chem. Mater. 2005,17, 5382–5389.

(25) Iyer, A.; Galindo, H.; Sithambaram, S.; King’ondu, C.; Chen,C.-H.; Suib, S. L. Appl. Catal., A 2010, 375, 295–302.

(26) DeGuzman, R. N.; Shen, Y.; Neth, E. J.; Suib, S. L.; O’Young,C.; Levine, S.; Newsam, J. M. Chem. Mater. 1994, 6, 815–821.

(27) Cai, J.; Liu, J.; Willis, W. S.; Suib, S. L. Chem. Mater. 2001,13, 2413–2422.

(28) Cheng, F.; Shen, J.; Peng, B.; Pan, Y.; Tao, Z.; Chen, J. Nat.Chem. 2011, 3, 79–84.



(29) Li, W.-N.; Yuan, J.; Shen, X.-F.; Gomez-Mower, S.; Xu, L.-P.;Sithambaram, S.; Aindow, M.; Suib, S. L. Adv. Funct. Mater. 2006,16, 1247.(30) Yuan, J.; Li, W.; Gomez, S.; Suib, S. L. J. Am. Chem. Soc. 2005,

127, 14184.(31) Bard, A. J.; Faulkner, L. R. Electrochemical Methods Fundamen-

tals and Applications, 2nd ed.; John Wiley & Sons: New York, 2001.(32) Ngala, J. K.; Alia, S.; Dobley, A.; Crisostomo, V. M. B.; Suib,

S. L. Chem. Mater. 2007, 19, 229–234.

Oxygen Reduction Properties of Bifunctional α-Manganese Oxide Electrocatalysts in Aqueous and...

Documents

Transcript of Oxygen Reduction Properties of Bifunctional α-Manganese Oxide Electrocatalysts in Aqueous and...