In situ fabrication of porous-carbon-supported α-MnO2 nanorods at room temperature: application for...

Energy &Environmental Science

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aChemical Science and Engineering Divisio

South Cass Avenue, Lemont, IL 60439,

+1-630-252-3838bX-ray Science Division, Advanced Photon S

9700 South Cass Avenue, Lemont, IL 60439cElectron Microscopy Center, Material

Laboratory, 9700 South Cass Avenue, Lemo

† The submitted manuscript has beenOperator of Argonne National LaboraDepartment of Energy Office of Science la

Cite this: Energy Environ. Sci., 2013, 6,519

Received 26th September 2012Accepted 3rd December 2012

DOI: 10.1039/c2ee23621d

www.rsc.org/ees

This journal is ª The Royal Society of

In situ fabrication of porous-carbon-supported a-MnO2

nanorods at room temperature: application forrechargeable Li–O2 batteries†‡

Yan Qin,§a Jun Lu,§a Peng Du,a Zonghai Chen,a Yang Ren,b Tianpin Wu,a

Jeffrey T. Miller,a Jianguo Wen,c Dean J. Miller,c Zhengcheng Zhanga

and Khalil Amine*a

Lithium–O2 cells can be considered the “holy grail” of lithium batteries because they offer much superior

theoretical energy density to conventional lithium-ion systems. In this study, porous carbon-supported

MnO2 nanorods synthesized at room temperature were explored as an electrocatalyst for rechargeable

Li–O2 cells. Both high-energy X-ray diffraction and X-ray absorption fine-structure analyses showed that

the prepared MnO2 exhibited a tetragonal crystal structure (a-MnO2), which has proved to be one of

the most efficient catalysts to facilitate the charging of the Li–O2 cell. Under the current synthetic

approach, a-MnO2 was uniformly distributed onto the surface of a carbon support, without disrupting

the porous structure at the surface of the carbon cathode. As a result, the as-prepared catalysts

demonstrated good electrochemical behavior, with a capacity of �1400 mA h g�1 (carbon +

electrocatalyst) under a current density of 100 mA g�1 (carbon + electrocatalyst) during the initial

discharge. The charge potential was significantly reduced, to 3.5–3.7 V, compared with most of the

reported data, which are above 4.0 V. The mechanism of the capacity fade with cycling was also

investigated by analyzing the cathode at different states of discharge–charge by X-ray photoelectron

spectroscopy.

Broader context

Energy dense and environmentally friendly, the lithium–air battery is one of the most attractive emerging energy storage systems. A fully developed andoptimally packed Li–air battery could exceed a specic energy of 2000 W h kg�1, which is more than twice as high as other battery types, either primary orsecondary. One of the biggest hurdles for rechargeable lithium–O2 batteries is the large overpotential during charge and discharge, which depends on the natureof catalysts applied and their loading onto a high-surface-area cathode. Finding a catalyst that can facilitate the ORR–OER process, thereby improving the cycleefficiency of Li–air batteries, as reported in this paper based on systematic experimental results, is strongly required in the eld. Herein, we applied a wet-chemistry approach to synthesize porous carbon supported a-MnO2 nanorods at ambient temperature as an electrocatalyst for rechargeable Li–O2 cells. Bymeans of our unique synthetic approach, MnO2 can be uniformly dispersed onto the surface of a carbon support, while the porous structure and surface area ofcarbon are still well preserved. As a consequence, the as-prepared catalysts demonstrated good electrochemical behavior in cell tests.

1 Introduction

Recently, there has been an increasing demand for energystorage devices with high energy density for use in electric and/or hybrid vehicles. Metal–air batteries, including Zn–air, Al–air,

n, Argonne National Laboratory, 9700

USA. E-mail: [email protected]; Tel:

ources, Argonne National Laboratory,

, USA

Science Division, Argonne National

nt, IL 60439, USA

created by UChicago Argonne, LLC,tory (“Argonne”). Argonne, a U.S.boratory, is operated under contract

Chemistry 2013

and Li–air, have attracted considerable attention owing to theirextremely high specic capacity. In general, the metal–airbattery is a class of electrochemical cell comprising a metalanode, which is oxidized by releasing electrons to the externalcircuit during discharge, in combination with an oxygen

no. DE-AC02-06CH11357. The U.S. Government retains for itself, and othersacting on its behalf, a paid-up nonexclusive, irrevocable worldwide license insaid article to reproduce, prepare derivative works, distribute copies to thepublic, and perform publicly and display publicly, by or on behalf of theGovernment.

‡ Electronic supplementary information (ESI) available: Electrochemicalperformance of Sample 2 (50% MnO2 on carbon). See DOI: 10.1039/c2ee23621d

§ These authors contributed equally.

Energy Environ. Sci., 2013, 6, 519–531 | 519

http://dx.doi.org/10.1039/C2EE23621D

http://pubs.rsc.org/en/journals/journal/EE

http://pubs.rsc.org/en/journals/journal/EE?issueid=EE006002

Energy & Environmental Science Paper

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cathode, on which oxygen is obtained from ambient air andreduced at a catalytic surface during discharge of the metal–aircell.1

Lithium–O2 cells can be considered the “holy grail” oflithium batteries because they offer, in principle, a superiortheoretical energy density, as high as 5200 W h kg�1, as shownin Table 1. The inherent energy density of lithium (more than11 000W h kg�1) approaches that of gasoline (13 000W h kg�1),which makes Li–O2 batteries even more attractive for electricvehicle, stationary energy storage, and military applications.2–38

Although it provides signicantly higher energy density thanintercalation electrodes, much fundamental work is required tounderstand and optimize the performance of the O2 electrodefor Li–O2 batteries before it can be considered further forpractical applications.2,23,32

One of the biggest hurdles for rechargeable lithium–O2

batteries is the large overpotential during charge and discharge,even at very low current densities (0.01–0.05 mA cm�2), whichresults in very low round-trip efficiencies (<60%) and low powercapability.13 This performance is strongly believed to depend onthe nature of catalysts applied and their loading onto a high-surface-area cathode. In previous work by other researchgroups, metal complexes39 and metal oxides40–42 have beenexamined as the cathode catalysts for the oxygen reductionreaction (ORR), showing large differences in discharge capacityamong different catalysts. However, surprisingly, a nearlyidentical discharge voltage plateau of about 2.6 V was obtainedfor different catalysts.39–42 This effect probably indicates that theORR kinetics in a Li–O2 cathode is either not catalyticallysensitive or limited by the oxygen mass transport resistancetoward the catalyst. Yet, the charge voltage plateau of a singleLi–O2 cell strongly depends on the nature of the catalysts. Theoxygen oxidation plateau can be signicantly lowered whencatalysts are introduced to the cathode.39–42 Recently, Yang'sresearch group28 demonstrated that the catalyst can provide thepossibility of not only reducing the charge voltage but alsoraising the discharge voltage. In tests with a nano-structuredPtAu/C catalyst, the charge voltage was as low as 3.8 V, and thecycle efficiency was as high as 73% at a rate of 100 mA g�1.Yang's group proposed that the alloy nanoparticles serve asbifunctional catalysts. Gold sites enhance the oxygen reductionreaction that occurs during discharge, and platinum sitesfacilitate oxygen evolution during charging. Nevertheless,MnO2-based catalysts are considered as the most favorable

Table 1 Relative specific energies of Li–O2, lithium-ion, and gasoline system

System ReactionOCVa

(V)Theoretical specicenergy (W h kg�1)

Li–O2 2Li + O2 ¼ Li2O2 3.1 3623 (including O)4Li + O2 ¼ 2Li2O 2.9 5204 (including O)4Li + O2 ¼ 2Li2O 2.9 11 202 (excluding O)

Li-ion, e.g., C6/LiMO2

(M ¼ Mn, Ni, Co)LixC6 + Li1�xMO2

¼ C6 + LiMO2

3.6 �900

Gasoline (octane) C8H18 + O2 ¼8CO2 + 9H2O

— �13 000 (excluding O)

a OCV ¼ open circuit voltage.

520 | Energy Environ. Sci., 2013, 6, 519–531

trade-off between reactivity and cost. The present study is,therefore, focused on MnO2-based materials as a catalyst for theoxygen cathode in the Li–O2 cell.

In 2006 Bruce et al.35 demonstrated high-energy-density Li–O2

cells using a nonaqueous liquid electrolyte and a porous carboncathode with electrolytic manganese dioxide. Since that time,MnO2-based compounds have been widely investigated as thecathode catalyst to improve the performance of the oxygen elec-trode. In their work, MnO2 with different crystal structures andmorphologies was found to exhibit different catalytic reactivityfor theoxygenevolution reaction (OER).Ofparticular interest, theuseofa-MnO2nanowiresas thecatalyst for rechargeable lithium–

air batteries was found to deliver capacities of 3000 mA h g�1 ofcarbon, making it one of the most promising catalysts for theoxygen cathode.34 However, a problem has been encountered inmost previous studies10,28,34,35,43 because the catalysts are pre-synthesized and then loaded onto a carbon support by means ofmechanical milling. This process introduces not only thepotential of destroying the porous structure of the carbonsupport, if high-energy ball milling is applied, but also thepossibility of nonuniform dispersion of the catalysts on thecarbon support, whichmight severely degrade the discharge andcharge properties of the oxygen cathode, and thus the overallperformance of the Li–O2 cell.

The present study aims to explore a new method touniformly disperse the catalyst onto the carbon support whilestill preserving the original porous structure of carbon duringthe synthetic process. Both the high surface area and porousstructure of the carbon cathode are critical for the performanceof Li–O2 batteries during electrochemical reactions. Typically, alarger surface area provides more surface to uniformly dispersecatalytic particles and more active sites to aid the electro-chemical reactions. A porous structure with an appropriate poresize provides the space to store the discharge products. In thisstudy, MnO2 nanoparticles were incorporated into porouscarbon through a wet-chemistry approach at room temperature,which avoids any aggressive post-treatment of the carbonmaterials. Thus, the original structure of carbon was wellpreserved. As a result, the performance of the oxygen cathodeduring charge and discharge cycles is expected to improve.

2 Experimental section2.1 Preparation of the MnO2-catalyzed cathode andelectrochemical measurements

Unless otherwise noted, all reagent grade chemicals were usedas-received without any further purication. The startingmaterials, manganese sulfate (MnSO4) and potassiumpermanganate (KMnO4), were purchased from Sigma-AldrichChemical (Milwaukee, WI). Super P Li carbon, designated as“SPL carbon” hereaer, was purchased from Ultralife Batteries,Inc. This high-surface-area carbon with a porous structure wasused for the catalyst support as part of the cathode materials.

For in situ deposition of MnO2 onto carbon, we oxidized 5 gof SPL carbon by dispersing it in 100 ml of 1 M KMnO4 solutionby magnetic stirring for 12 h.44 The oxidized carbon was washedwith deionized water, ltered multiple times, and dried in an

This journal is ª The Royal Society of Chemistry 2013


Paper Energy & Environmental Science

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oven at 110 �C for 12 h. Then, the oxidized carbon was dispersedin 100 ml deionized water by magnetic stirring. Next, 0.92 gMnSO4 and 0.37 g KMnO4 were mixed, manually ground in amortar-and-pestle, and then transferred into the carbon solu-tion prepared above. Aer magnetic stirring of this solution for2 hours, 20 ml of 1 M H2SO4 solution was added in drops, andthe resulting solution was stirred overnight. Aer ltration, thenal solid product was washed repeatedly with deionized waterand ethanol, and dried in an oven at 110 �C for 12 h. Twosamples with different weight percentage of catalyst loadingwere synthesized by the above method: 10 wt% and 50 wt%MnO2 versus SPL carbon, referred to as Sample 1 and Sample 2,respectively. For comparison, a sample containing the samecarbon and catalyst (a-MnO2) was ball milled using acommercial SPEX mill for 4 hours.

Electrochemical characterizations were carried out in Swa-gelok-type cells composed of a lithium metal anode, an elec-trolyte of 1 M LiCF3SO3 in tetraethylene glycol dimethyl ether(TEGDME) impregnated into a glass ber separator, and aporous cathode (11 mm in diameter). The cathode was preparedby casting a mixture of the as-prepared MnO2/SPL carbon andKynar 2801 binder in a weight ratio of 50 : 50 on an aluminumgrid. The active material loading on the cathode laminate isabout 0.5–2mg cm�2. The cells were sealed except for the Al gridwindow that exposed the porous cathode to 1 bar O2 pressure toavoid any negative effects of humidity and CO2. The electro-chemical measurements were conducted with a MACCOR cyclerin the voltage range of 2.3–4.3 V at a constant current of 100 mAg�1. We normalized the observed capacity by the weight ofcarbon and catalyst for capacity comparison in this study.

2.2 Characterization

BET analysis. The porous structure and Brunauer–Emmett–Teller (BET) surface area of SPL carbon before and aer loadingthe MnO2 catalyst were characterized by use of a N2 adsorption–desorption isotherm (NOVA-2020 Micromeritics Co., USA). Thesamples were degassed at 220 �C prior to the BET measure-ments. The pore volume and pore diameter distribution werederived from desorption branches of the isotherms by theBarrett–Joyner–Halenda (BJH) model. The surface area wascalculated from the linear part of the BET plot.

High-energy X-ray diffraction. The structures of the as-prepared MnO2 were characterized by synchrotron X-raydiffraction (XRD), which was carried out at the 11-ID-C beam-line of the Advanced Photon Source (APS), Argonne NationalLaboratory. The X-ray wavelength was 0.10798 A. The sampleswere prepared by tapping the powder into a Kapton capillary,which was then sealed at both ends. The XRD patterns werecollected in the transmission mode using a Perkin Elmer largearea detector. The collected 2D patterns were then integratedinto conventional 1D patterns (intensity vs. 2q) for nal dataanalysis using the Fit2d soware.45 Based on XRD peak broad-ening, the crystallite size of the sample was estimated from theScherrer equation:

Dhkl ¼ 0:89l

bð2qÞcosðqÞ


where b(2q) is the full width at half-maximum (FWHM) of thepure diffraction peak in radians, l is the wavelength of theX-rays (0.10798 A), q is the diffraction angle, and Dhkl isthe average diameter of the crystallite.

X-ray adsorption ne structure analysis. Mn K-edge X-rayadsorption ne structure (XAFS) analysis was performed at theMaterials Research Collaborative Access Team (MRCAT) andComplex Materials Consortium (CMC) beamlines of the APS.The setup used was similar to that outlined by Castagnolaet al.46 The sample was loaded as a self-supporting waferwithout binder in the channels (i.d. ¼ 4 mm) of a stainless steelmulti-sample holder. The sample holder was then placed in thecenter of a quartz tube, which was equipped with gas andthermocouple ports and Kapton windows. The amount ofsample used was optimized for a step height of about 0.5. TheXAFS spectra were recorded in the transmission mode. Stan-dard procedures based on WINXAS 3.1 soware were used to tX-ray absorption spectroscopy data. The EXAFS coordinationparameters were obtained by a least-square t in the q- and r-space of the k2-weighted Fourier transform data for the isolatednearest neighbor.

Energy dispersive spectrometry and scanning and trans-mission electron microscopy. Electron microscopy was carriedout in electron microscopy center (EMC) at Argonne NationalLaboratory. Field-emission scanning electron microscopy (SEM,Hitachi S-4700) coupled with energy dispersive spectrometry(EDS) was employed to determine the morphology and estimatethe particle size of the samples. TEM was employed to obtainmore accurate estimates of the particle sizes and/or crystallitesizes. For the TEM analysis, we obtained a dilute suspensionfrom a sample that had been ultrasonically dispersed in ethanolfor 5 min, dropped the suspension onto a copper grid, and driedit. Samples were examined using both a Philips CM 30 operatedat 80 kV and a FEI Titan operated at 80 kV to allow minimume-beam damage when imaging.

X-ray photoelectron spectroscopy. Samples were analyzed byX-ray photoelectron spectroscopy (XPS) using a Kratos� AxisUltra DLD surface analysis instrument. The base pressure of theanalysis chamber during these experiments was 3 � 10�10 torr,with operating pressures around 1 � 10�9 torr. Spectra werecollected with a monochromatic Al Ka source (1486.7 eV) and a300 � 700 micron spot size. The Al source was operated at anemission current of 13 mA with the target anode set to 15 kV;the resulting power was 195 W. For survey spectra the data werecollected at a pass energy of 160 eV (xed analyzer transmissionmode), a step size of 1 eV, and a dwell time of 200 ms. High-resolution regional spectra were collected with a pass energy of20 eV (xed analyzer transmission mode), a step size of 0.1 eV,and a dwell time of 300 ms. For low signal-to-noise regions,multiple passes were made, and the results were averagedtogether.

Prior to introduction into the load-lock vacuum chamber ofthe XPS instrument, all air-sensitive samples were loaded intoan inert transfer module interfaced with the instrument.Samples were prepared for analysis in an Ar-lled glove box,with no more than 1 ppm O2 and 1 ppm H2O. Nonconductivesamples showed evidence of differential charging, resulting in




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peak shis and broadening. Photoelectron peak positions wereshied back toward their true values, and their peak widthswere minimized by ooding the samples with low-energy elec-trons and ions from the charge neutralizer system on theinstrument. Peak position correction was further corrected byreferencing the C 1s peak position of adventitious carbon for asample (284.8 eV, PHI Handbook of Photoelectron Spectroscopy),and shiing all other peaks in the spectrum accordingly.

Fitting was done by using the program CasaXPS. Each rele-vant spectrum was t to a Shirley type background to correct forthe rising edge of backscattered electrons that shis the base-line higher at high binding energies. Peaks were t as asym-metric Gaussian/Lorentzians, with 0–30% Lorentzian character.The FWHM of all sub-peaks was constrained to 0.7–2 eV, asdictated by instrumental parameters, lifetime broadeningfactors, and broadening due to sample charging. With thisnative resolution set, peaks were added, and the best t, using aleast-squares tting routine, was obtained while adhering to theconstraints mentioned above.

3 Results and discussion3.1 Synthesis and characterizations of carbon-supported a-MnO2 nanoparticles

The phase purity and crystal structure of the MnO2–Ccomposites were characterized by high-energy powder XRD, andthe results are presented in Fig. 1. The XRD patterns for theoriginal SPL carbon and the standard a-MnO2 calculated frominorganic crystal structure database (ICSD) are included forcomparison, as shown in the bottom two curves in Fig. 1. Thelarge peak at 2q ¼ 1.7 is from SPL carbon. All the other peakscan be readily indexed to a pure a-MnO2 phase with the latticeconstant a ¼ 9.82 A, b ¼ 9.82 A, and c ¼ 2.84 A, which are ingood agreement with the reported values found in the ICSD.Sample 2 can be seen to have higher MnO2 loading than Sample1 by comparing the peak intensity of these two samples. The

Fig. 1 High-resolution XRD patterns of the original SPL carbon and SPL carbonafter being loaded with the MnO2 catalyst. The corresponding standard a-MnO2

calculated from ICSD database is shown at the bottom of this figure forcomparison.


average grain size for Sample 2, estimated by the Scherrerequation, is about 20 nm. It is worth mentioning that, withoutadding carbon during the synthesis, a-MnO2 with much largerparticle size would be expected to form, as mentioned in ref. 47.With addition of carbon, the obtained a-MnO2 particle is alsorod-like (conrmed by TEM as discussed below) but muchsmaller, possibly because the carbon matrix limits the rod-likegrowth.

To further conrm the crystal structure of the sample, XAFSmeasurements were performed on the MnO2–C composites.Fig. 2a shows the Mn K-edge X-ray absorption near edge struc-ture (XANES) spectra of the prepared MnO2 catalyst (Sample 1)and that of a commercial manganese dioxide (g-MnO2). TheSample 1 XANES is located at the same energy position (6.562keV) as that of the commercial manganese dioxide, althoughthe shapes of the edge are a little different, indicating that theprepared MnO2 has the same chemical valence (+4) as that ofthe commercial manganese dioxide. Fig. 2b shows the magni-tude and the imaginary parts of the Fourier transform of the MnK-edge EXAFS for these two samples. The magnitude of theFourier transform for the rst shell coordination of Sample 1has two main peaks, one large peak at about 1.6 A and a smaller

Fig. 2 X-ray absorption data: (a) Mn K-edge spectra for the commercial g-MnO2

and the prepared MnO2–C composite and (b) Fourier transforms of the absorp-tion spectra of the magnitude part. Fitting results are provided in Table 2.




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peak at about 2.4 A. This spectrum is typical of the K-edgeEXAFS for a single 3d metal. The magnitude of the Fouriertransform of the commercial manganese dioxide has threemain peaks. Fitting both the magnitude and imaginary parts ofthe Fourier transform to a model that includes Mn enabled usto extract bond distances for Mn–Mn and Mn–O as well as thecoordination numbers, N. TheMn–O bond distance in Sample 1is 1.6 A, which is the same as that in the commercial g-MnO2.However, the Mn–Mn bond distance is quite different betweenSample 1 and the commercial g-MnO2. These intensity andbonding results provide further evidence that the preparedMnO2 has a different crystal structure from that of thecommercial g-MnO2.

The pore size and specic surface area of the carbon supportbefore and aer the reaction were characterized by nitrogenadsorption–desorption isotherm measurements, as shown inFig. 3. The results indicate that the original SPL carbon has aspecic surface area of about 70 m2 g�1 with a pore size ofaround 60 nm (top curve in Fig. 3a), which are similar to thoseof Sample 1 (bottom curve in Fig. 3a). The pore-size distributionof Sample 1 was further conrmed by BJH desorptionisotherms, as shown in Fig. 3b. These BET results suggest thatthe specic surface area and pore size of the original SPL carbonare well preserved aer being loaded with the catalyst.

Field-emission SEM was used to determine the particle sizesand morphologies of the original SPL carbon and Sample 1

Fig. 3 (a) N2 adsorption–desorption isotherms and (b) the corresponding BJHpore size distributions of the original SPL carbon and the as-prepared MnO2–Ccomposite (Sample 1).


and 2. As shown in Fig. 4a, the primary particles of the as-received SPL carbon powder are spherical, with an averageparticle size of 100 nm. These particles tend to agglomerate toform secondary particles and pores. Fig. 4b and c show the SEMimages of Sample 1 and Sample 2, which indicate the samemorphology and particle size as those of the original SPLcarbon. Comparing Fig. 4a–c reveals that the porous structure ofcarbon is well preserved during the process of synthesizing theMnO2 catalyst. Fig. 4d and e show the back scattering images ofthe same area of Fig. 4b and c, respectively, from which it isclearly shown that the MnO2 nanorods (white spots) areuniformly distributed onto the carbon matrix. Back scatteringSEM observation shows that Sample 2 has higher MnO2 loadingthan Sample 1.

TEM was used to further investigate the morphology andcrystalline structure of the samples. Fig. 5a and b present TEMimages for Sample 1 and Sample 2, respectively, clearly showingevidence that most of the prepared MnO2 particles areuniformly dispersed onto the carbonmatrix with rod-like shape.TEM observation also shows that Sample 2 has higher MnO2

loading than Sample 1, consistent with the XRD and SEMresults. Fig. 5c presents a TEM image of the original SPL carbonand the inset shows the morphology of graphitic layers insidecarbon particles. Carbon particles have an onion-like structure

Fig. 4 SEM images of (a) Pristine Super P Li carbon, (b) carbon loaded with 10%MnO2 and (c) carbon loaded with 50% MnO2. (d and e) Backscattered SEMimages for the same areas that are shown in (b) and (c), respectively. (f) Secondaryelectron and (g) backscattered electron images of a sample obtained from themixture of Super P Li carbon and 10% MnO2 after ball milling for 4 h.



Fig. 5 TEM images of (a) as-prepared MnO2–C composite (Sample 1) and (b) as-prepared MnO2–C composite (Sample 2); and (c) original SPL carbon; (d) and (e)high-resolution TEM image of Sample 1, and (f) EELS of Sample 1.


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and an average particle size of 50 nm. Fig. 5d shows a TEMimage of MnO2 nanorods (with dimensions of about 5 nm � 20nm) attached to carbon. Fig. 5e shows a high-resolution TEMimage of a MnO2 nanorod with a lattice spacing of 1.9 A, cor-responding to the (510) spacing of the bulk a-MnO2 (ICSD les,card number 20229). Fig. 5f shows the EELS spectrum from theMnO2 nanorod, the intensity ratio between L3 (large peak) andL2 (small peak) indicates that the valence of Mn is 4+, which isconsistent with the chemical composition of nanorodsmeasured by XRD and XANES.

Table 2 summarizes results derived from XRD, BET, XAFS,and TEM analyses for Samples 1 and 2. On the basis of all theseresults, we concluded that (1) the porous structure of the orig-inal SPL carbon is preserved during the synthesis of the MnO2

catalyst; (2) the as-prepared MnO2 is crystalline with a

Table 2 Physicochemical properties of the prepared MnO2/C samples from N2 sor

SampleSBET(m2 g�1)

Porevolume (cm3 g�1)

SPL carbon 70.26 0.356Sample 1 68.66 0.332Sample 2 67.47 0.309


tetragonal structure (a phase), and it shows a rod-like shapewith an average dimension of 5 � 20 nm; and (3) the preparedMnO2 nanorods are uniformly distributed on the carbonsurface. Because these as-prepared MnO2–C composites haveporous structures and high specic surface area, they canprovide more active sites for oxidation and reduction. Conse-quently, the catalytic activity for OER is expected to increase, asdemonstrated next in the tests with Li–O2 cells.

3.2 Electrochemical cycling performance of a lithium–

oxygen cell

Samples 1 and 2 were tested as cathodes in a Swagelok-type cellhaving a Li-foil anode. The cells were tested at 1 atm O2 using aMACCOR cycler. Sample 1 was chosen as an example todemonstrate the electrochemically catalytic property of theMnO2–C composites in the Li–O2 cell, since both samplesshowed similar electrochemical performance in the voltagewindow used in the current study (the electrochemical perfor-mance of Sample 2 is provided in the ESI‡). Fig. 6 shows voltageproles of the rst discharge–charge cycles obtained for the cellusing TEGDME/LiCF3SO3 as the electrolyte. The initialdischarge reaction occurred at approximately 2.62 V vs. Li+/Li(discharge rate is 100 mA g�1). A total capacity of �1400 mA hg�1 was delivered to 2.3 V in the rst discharge, most of whichwas delivered at the 2.62 V plateau. Severe polarization occurredduring the subsequent charging, which is a common feature ofcells in non-aqueous electrolytes. To ensure the full recovery ofthe discharged product, we employed an equal capacitycharging procedure, in which the cutoff charging condition isspecied as the charging capacity matches the previous dis-charging capacity. With this procedure, around 1400 mA h g�1

was recovered at the charge plateau of about 4.0 V, which couldbe attributed predominantly to oxidation of the dischargeproducts deposited on the cathode surface (i.e., OER). On thesecond discharge to 2.3 V, the capacity dropped dramatically (by600 mA h g�1), which is probably because some surface and/orpores in the oxygen cathode were blocked during the rst cycle.A capacity of about 800 mA h g�1 was maintained for the nextseveral cycles (<6) until the capacity nally faded (Fig. 6b),indicating that the cell charge–discharge is reasonably revers-ible at earlier cycles. A round-trip efficiency of �68% wascalculated based on the rst-cycle measurement. It should bepointed out that the charge potential increased to about 4.2 Vduring the subsequent cycles, suggesting that even more severepolarization occurs, probably because solid discharge productsform and ll the pores of the carbon support, as has beenobserved in earlier publications.35 Note that the intercalation

ption, XRD patterns, EXAFS and TEM images

Average poresize (nm)

MnO2 crystalstructure

Average MnO2

crystal size (nm)

84.2 N/A N/A80.4 Tetragonal (a) 20.278.1 Tetragonal (a) 25.5



Fig. 7 (a) Discharge and charge profiles of a Li–O2 cell containing 10% MnO2 inthe carbon matrix (Sample 1) using the TEGDME/LiCF3SO3 electrolyte between2.2 and 4.3 V versus Li/Li+ with a limited capacity of 500 mA h g�1 at a constantcurrent density of 100 mA g�1 (0.6 mg of carbon + catalyst); (b) cyclability of thesame cell shown in (a).

Fig. 6 Voltage profile of the first cycle for Sample 1 cathode containing 10%MnO2 in the carbon matrix cycling in 1 M TEGDME/LiCF3SO3 under a currentdensity of 100 mA g�1 (1.2 mg carbon + MnO2). (b) Cyclability of the same cellshown in (a).


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reaction of the a-MnO2 used in the cell is negligible, because theproportion used in the electrode is very small and contributes atmost �60 mA h g�1 (�0.2 mg a-MnO2 per air electrode themost) in total capacity.

Very recently, it has been reported that considerableimprovement of the cyclability can be achieved by controllingthe discharge and charge protocol of the Li–O2 cells, i.e.,limiting the depth of the discharge.48,49 Such limited discharge–charge protocols are analogous to the way that cathodes areused in commercial lithium ion batteries, including LiCoO2 andNCM (Li(NiMnCo)1/3O2) cathodes, because the use of morelithium involves phase transitions, which signicantly deterio-rate the cycle performance.50,51 Therefore, we also adopted thesame protocol in this study to investigate the cycle performanceof the Li–O2 cell when using the MnO2–C composite as theoxygen cathode. To avoid the electrolyte decomposition duringthe discharge and charge, we constrained the capacity to 500mA h g�1. Under this protocol, an energy density of 1350 W hkg�1 (500 mA h g�1 � 2.7 V) can be achieved, which is about twotimes the energy density of commercial cathodes in Li-ionbatteries (160 mA h g�1 � 4 V). The cell was discharged andcharged at a constant current of 100 mA g�1 with constantcurrent (CC) mode. Fig. 7 indicates that a considerable


improvement in cycle performance could be achieved using thecapacity limited discharge–charge protocol. Reversible cyclingcan be maintained for over 50 cycles using TEGDME/LiCF3SO3

as the electrolyte. In contrast, as shown in Fig. 6, an identicalcell operated without the capacity constraint can last for only 6cycles, although it delivered much higher capacity. Fig. 7bpresents the discharge capacity as a function of cycle number ata current of 100 mA g�1 for the same cell tested above,demonstrating that cycling can be sustained for more than 50cycles, aer which the capacity gradually dropped. This resultclearly indicated that the limited capacity protocol improved therechargeability of the Li–O2 cell by minimizing the possibledecomposition of the electrolyte, which could severely occur atthe deep discharge–charge condition in the long run otherwise.It should also be noted that the discharge potential wasincreased to about 2.9 V at the rst 20 cycles under the capacitylimited protocol, as shown in Fig. 7a. Based on the aboveresults, it is clearly indicated that the cyclability of a Li–O2 cellcan be improved by applying our MnO2–C composite as thecathode material along with the capacity limited protocolduring the discharge and charge process.

For comparison, the electrode prepared by the conventionalhigh-energy mechanical milling (SPEX Mill) with the same




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carbon and catalyst is also tested using the same discharge–charge protocol, and the results are shown in Fig. 8. It is clearlyindicated that only a few cycles can be sustained using themilled sample as the cathode material in the Li–O2 cell evenunder the capacity-limited protocol. As we pointed out earlier,high-energy ball milling will introduce not only the potential ofdestroying the porous structure of the carbon support, but alsothe possibility of nonuniform dispersion of the catalysts ontothe carbon support, as conrmed by the SEM images shown inFig. 4f and g. This clearly led to severe degradation of thedischarge and charge properties of the oxygen cathode, asshown in Fig. 8.

A different electrolyte, i.e., TEGDME/LiTFSI, was tested in aLi–O2 cell containing the as-prepared MnO2–C composites asthe active cathode materials. Fig. 9 presents voltage proles ofthe different discharge–charge cycles, indicating that goodcyclability (>40) can be achieved in such an electrolyte. A slightlyhigher discharge–charge rate, i.e. 150 mA g�1, was applied inorder to investigate the rate capability in a Li–O2 cell containingthe as-prepared MnO2–C composite as the cathode, and resultsare shown in Fig. 10. Good cyclability can still be achieved,although it is not as good as that from low rate (Fig. 7). More

Fig. 8 (a) Discharge and charge profiles of a Li–O2 cell containing 10% MnO2 inthe carbon matrix (Spex milled sample) using the TEGDME/LiCF3SO3 electrolytebetween 2.2 and 4.3 V versus Li/Li+ with a limited capacity of 500 mA h g�1 at aconstant current density of 100 mA g�1 (0.6 mg of carbon + catalyst); (b) cycla-bility of the same cell shown in (a).

Fig. 9 (a) Discharge and charge profiles of a Li–O2 cell containing 10% MnO2 inthe carbon matrix (Sample 1) using the TEGDME/LiTFSI electrolyte between 2.2and 4.3 V versus Li/Li+ with a limited capacity of 500 mA h g�1 at a constantcurrent density of 100 mA g�1 (0.5 mg of carbon + catalyst); (b) cyclability of thesame cell shown in (a).


importantly, it is also shown from Fig. 7 and 10 that the currentdensity does not affect the discharge and charge potentials.

The increased polarization during charge and discharge isthe cause of the capacity and energy efficiency loss in thecurrent lithium–O2 cell. This loss might be due to clogging ofthe pores of the carbon matrix by the solid discharge products,as pointed out earlier. However, the catalyst, a key componentin the cathode materials, may also contribute to the capacityfade. We therefore applied XPS to investigate the catalystsurface and identify the discharge and charge products duringthe electrochemical reactions.

3.3 XPS characterization of the oxygen electrodes aerdischarge and charge

X-ray photoelectron spectra (XPS) were used to investigate thecatalyst surface and identify the discharge and charge productsduring the electrochemical reactions, which were obtainedfrom the oxygen electrodes at the 1st and 10th cycle (both atdischarged and charged state under a time control protocol(500 mA h g�1)), respectively. The C 1s, Li 1s, O 1s and Mn 2pXPS spectra were recorded and tted, as shown in Fig. 11. XPSspectra of pristine carbon and commercial Li2O2 are also listedin Fig. 11 for comparison. The C 1s spectra (Fig. 11a) of allsamples showed typical C–O (�287 eV) and C]O (�288.7 eV)



Fig. 10 (a) Discharge and charge profiles of a Li–O2 cell containing 10%MnO2 inthe carbon matrix (Sample 1) using the TEGDME/LiCF3SO3 electrolyte between2.2 and 4.3 V versus Li/Li+ with a limited capacity of 500 mA h g�1 at a constantcurrent density of 150 mA g�1 (0.6 mg of carbon + catalyst); (b) cyclability of thesame cell shown in (a).


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functionalities, which were originated from the pristine carbonsurface. The presence of such functionalities is furtherconrmed by the O 1s peak as shown in Fig. 11c. O–C]O(�290.8 eV) functionality is also found in all the samplesincluding the pristine carbon. No change of the relative inten-sity of the O–C]O peak was observed, indicating that theelectrolyte decomposition is avoided even aer 10 cycles of thedischarge and charge, maybe due to the capacity-limitedprotocol used in the current study.

Fig. 11b (1st discharge) shows the Li 1s XPS spectrum for the1st discharge sample, from which it can be concluded thatlithium oxides appear to be formed as the dominant dischargeproduct of the electrochemical reactions. O 1s XPS shown inFig. 11c (1st discharge) for the same sample conrmed theformation of lithium oxides. However, it is difficult to distin-guish Li2O2 from Li2O solely based on the XPS results since thebinding energies of these two compounds are in the samerange.52 However, based on the previous results on the dis-charged sample using the same electrolyte48,53 and comparisonwith the standard Li2O2 (bottom spectra), these peaks (55.5 eVfor Li and 531.9 eV for O) can be assigned to that of lithiumperoxide (Li2O2). This is an encouraging result showing that the


desired discharge products (lithium peroxide) can be producedwhen a suitable catalyst, in this case the MnO2–C composite,was introduced to the cathode. Upon charging, the lithiumsignals from lithium peroxide completely disappeared (Fig. 11b,1st charge (Li 1s) and Fig. 11c (O 1s)), suggesting that thesedischarged products decompose during the charging process.This nding provides solid evidence that lithium peroxide canbe formed and decomposed reversibly in a TEGDME-based Li–O2 cell. Li 1s and O 1s spectra of the 10th cycle samples (Fig. 11band c) clearly indicated that such reversible reaction can besustained for many cycles.

Fig. 11d presents the Mn 2p core level XPS spectra ofthe prepared MnO2–C cathode for different electrochemicaltreatments (charge–discharge), as marked in the gure. Thespin–orbit splitting obtained was 11.47 eV, with an intensityratio of 0.5 between the Mn 2p1/2 and Mn 2p3/2. Analysis ofMn 2P spectra of the discharge electrodes (1st and 10th

discharge in Fig. 11d) showed no Mn signal, providing furtherevidence that the lithium-based products were sufficientlythicker to coat the MnO2 particle. In other words, the oxygenreduction reaction does take place on the surface of the catalystduring discharge of the cell.54 The Mn signal reemerged in thecharged samples (1st and 10th charge in Fig. 11d) since thedischarge products were decomposed and removed fromthe electrode surface.

We also performed XPS analysis on the samples obtainedfrom the 1st full discharge–charge cycle, and the results werepresented in Fig. 12. Fig. 12a shows the C 1s spectra of thesamples, from which no signicant change of the relativeintensity of the O–C]O peak was observed, indicating that theelectrolyte decomposition is avoided at the rst cycle, during afull discharge and charge protocol. Li 1s and O 1s spectra of thesamples shown in Fig. 12b and c clearly showed that lithiumperoxide can be formed and decomposed reversibly in aTEGDME-based Li–O2 cell. It was also suggested that the oxygenreduction reaction does take place on the surface of the catalystduring discharge of the cell by comparing the Mn 2p spectra ofdischarge and charge samples, as shown in Fig. 12d.

3.4 Electrochemical performance in a different electrolyte

It is worth pointing out that other factors, such as the electro-lyte, also play critical roles in the performance of the Li–O2

battery.6,7,19,20,55,56 For this reason, Sample 1 was tested as acathode in a Swagelok-type cell using LiPF6/PC as an electrolyte,which is a widely used electrolyte in Li–O2 batteries. The cellswere tested at 1 atm O2 using a MACCOR cycler, and the voltageproles of the rst three discharge–charge cycles for the cell areshown in Fig. 13a. The initial discharge reaction occurred atapproximately 2.7 V vs. Li+/Li (discharge rate is 100 mA g�1),which is consistent with reported data by Bruce et al.34,35 for asimilar cell containing electrochemical manganese dioxide asthe catalyst discharged in 1 atm of oxygen in a PC-based elec-trolyte. A total capacity of�1000 mA h g�1 was delivered to 2.0 Vin the rst discharge, most of which was delivered at the 2.7 Vplateau. Severe polarization occurred during the subsequentcharging, which is a common feature of cells when PC is used in



Fig. 11 XPS spectra of (a) C 1s, (b) Li 1s, (c) O 1s and (d) Mn 2p core peaks of the carbon electrodes at different charge–discharge status under a time-controlledprotocol (500 mA h g�1).

Fig. 12 XPS spectra of (a) C 1s, (b) Li 1s, (c) O 1s and (d) Mn 2p core peaks of the carbon electrodes at 1st discharge and charge status where a fully discharge (to 2.3 V)protocol was applied.

528 | Energy Environ. Sci., 2013, 6, 519–531 This journal is ª The Royal Society of Chemistry 2013


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Fig. 13 (a) Voltage profiles of the first 3 cycles for Sample 1 cathode containing10%MnO2 in the carbon matrix cycling in 1 M LiPF6/PC at 100 mA g�1 between 2and 4.4 V; and (b) discharge capacity as a function of cycle number for air elec-trodes containing 10% in carbon matrix cycling in 1 M LiPF6/PC at a constantcurrent density of 100 mA g�1 (1.4 mg of carbon + catalyst).

Fig. 14 XPS spectra of (a) C 1s and (B) Li 1s core peaks of the carbon electrodesat different charge–discharge status cycling in 1 M LiPF6/PC at a constant currentdensity of 100 mA g�1 (1.4 mg of carbon + catalyst).


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the electrolyte. To ensure the full recovery of the dischargedproduct, we employed an equal capacity charging procedure, inwhich the cutoff charging condition is specied as the chargingcapacity matches the previous discharging capacity. With thisprocedure, around 900 mA h g�1 was recovered at the chargeplateau of about 3.7 V, which could be attributed predominantlyto oxidation of the discharge products deposited on the cathodesurface (i.e., OER). On the second discharge to 2.0 V, thecapacity dropped dramatically (by 300 mA h g�1), which isprobably because some surface and/or pores in the oxygencathode were blocked during the rst cycle. A capacity of about600 mA h g�1 was maintained for the next several cycles untilthe capacity nally faded, indicating that the cell charge–discharge is reasonably reversible. A round-trip efficiency of�73% was calculated based on the rst-cycle measurement. Itshould be pointed out that the charge potential increased toabout 4.0 V during the subsequent cycles, suggesting that evenmore severe polarization occurs, probably because soliddischarge products form and ll the pores of the carbonsupport, as has been observed in earlier publications.35 Fig. 13bpresents the discharge capacity as a function of cycle number at100 mA g�1 for the same cell tested above, demonstrating thathigh capacity can be sustained for 10 cycles. Compared to theresults obtained from the LiCF3SO3/TEGDME electrolyte, it isclear that the cyclability of the cell tested in a PC-based


electrolyte is poor due to the stability of PC against the reducedspecies during the discharge, such as O2

�.In order to investigate the electrolyte decomposition, XPS

spectra were obtained from the carbon electrodes containingSample 1 (tested in PC-based electrolyte) in the 1st discharged,1st charged, 5th discharged and 10th charged states, respectively.The C 1s and Li 1s XPS spectra were recorded and tted, asshown in Fig. 14. The C 1s spectra (Fig. 14a) showed carbonate-type components in all of the samples, suggesting that thepropylene carbonate electrolyte decomposed during dischargeand charge of the cell, which is consistent with the previousreports for such electrolyte.19 Fig. 14b shows the Li 1s XPSspectrum for the 1st discharge sample, from which it is clearthat (1) a Li2CO3 type product is formed on the cathode duringdischarge, consistent with the C ls spectrum (Fig. 14a), and (2)more importantly, lithium peroxide appears to be the majordischarge products of the electrochemical reactions. Uponcharging, the lithium concentration is signicantly reduced(Fig. 14b, 1st charge), suggesting that the discharge productsdecompose during the charging process. The remaining lithiumspecies in the charged electrode can be attributed to Li2CO3 andLiF/LiPF6 from the lithium salt. However, only Li2CO3 is formedduring the 5th discharge, indicating that electrolyte decompo-sition is dominant in the later cycle performance.

In fact, Bruce et al.19 recently showed that side reactions withan organic carbonate-based electrolyte (e.g., propylene




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carbonate) are more severe than commonly anticipated,providing a rationale for the compromised reversibility in thesystem. Our recent work also provided strong evidence that thedischarge products are heavily dependent on the electrolytesused.52 Further research is necessary to better understand themechanism of the electrochemical reactions, which, in turn,would help us to identify, design, and develop more stableelectrolyte systems. In addition, it has also been reported thatthe binder used in making the cathode laminate can inuencethe stability of discharged intermediates.57 Despite the uncer-tainty of the stability of electrolytes and binder in the Li–O2

system, the current study is focused on developing a newapproach to uniformly disperse catalysts onto the carbonsupport under mild conditions and seeking a highly activeelectrocatalyst for the oxygen reduction reaction.

4 Conclusion

In this study, we applied a wet-chemistry approach to synthesizeporous carbon supported a-MnO2 nanoparticles at ambienttemperature as an electrocatalyst for rechargeable Li–O2 cells.Both high-energy XRD and XAFS analyses conrmed the crystalstructure of a-MnO2, which has proved to be one of the mostefficient catalysts to facilitate the charging of the Li–O2 cell. Bymeans of our unique synthetic approach, MnO2 can beuniformly dispersed onto the surface of the carbon support,while the porous structure and surface area of carbon are stillwell preserved. As a consequence, the as-prepared catalystsdemonstrated good electrochemical behavior in cell tests,which can deliver 1400 mA h g�1 of capacity (carbon + electro-catalyst) using a current density of 100 mA g�1. More signi-cantly, the charge potential has been lowered to 3.5 V using acapacity-limited discharge–charge protocol, which is one of thelowest charge potentials reported for the Li–O2 cell. Thereversible charge–discharge reaction can be sustained for morethan 50 cycles. XPS data of the cathodes demonstratedunequivocally that lithium oxide products partook in reversiblereactions in TEGDME-based electrolytes.

Acknowledgements

Research at Argonne National Laboratory was funded by U.S.Department of Energy, FreedomCAR and Vehicle TechnologiesOffice. J. Lu was supported by the Department of Energy (DOE)Office of Energy Efficiency and Renewable Energy (EERE)Postdoctoral Research Award under the EERE Vehicles Tech-nology Program administered by the Oak Ridge Institute forScience and Education (ORISE) for the DOE under DOEcontract number DE-AC05-06OR23100. The SEM and TEManalysis was performed at the Electron Microscopy Center(EMC), which is supported by the DOE Office of Science – BasicEnergy Sciences under contract DE-AC02-06CH11357. Use ofthe Advanced Photon Source was supported by the U.S.Department of Energy, Office of Basic Energy Sciences, undercontract no. DE-AC02-06CH11357. MRCAT operations aresupported by the Department of Energy and the MRCATmember institutions.


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In situ fabrication of porous-carbon-supported α-MnO2 nanorods at room temperature: application for...

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