Electrochemical Performance of MnO2 Nanorods in Neutral Aqueous

Jul 14, 2009 - Dongchang Chen , Dong Ding , Xiaxi Li , Gordon Henry Waller , Xunhui Xiong , Mostafa A. El-Sayed , and Meilin Liu. Chemistry of Materia...
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Electrochemical Performance of MnO2 Nanorods in Neutral Aqueous Electrolytes as a Cathode for Asymmetric Supercapacitors Qunting Qu,† Peng Zhang,† Bin Wang,† Yuhui Chen,† Shu Tian,† Yuping Wu,*,† and Rudolf Holze*,‡ New Energy and Materials Laboratory (NEML), Department of Chemistry & Shanghai Key Laboratory of Molecular Catalysis and InnoVatiVe Materials, Fudan UniVersity, Shanghai 200433, China, and Technische UniVersita¨t Chemnitz, Institut fu¨r Chemie, D-09107 Chemnitz, Germany ReceiVed: December 22, 2008; ReVised Manuscript ReceiVed: May 20, 2009

The electrochemical performance of MnO2 nanorods prepared by a precipitation reaction was investigated in 0.5 mol/L Li2SO4, Na2SO4, and K2SO4 aqueous electrolyte solutions. Results show that at the slow scan rates, the nanorods show the largest capacitance (201 F/g) in Li2SO4 electrolyte since the reversible intercalation/ deintercalation of Li+ in the solid phase produces an additional capacitance besides the capacitance based on the absorption/desorption reaction. At fast scan rates they show the largest capacitance in the K2SO4 electrolyte due to the smallest hydration radius of K+, highest ionic conductivity, and lowest equivalent series resistance (ESR). An asymmetric activated carbon (AC)/K2SO4/MnO2 supercapacitor could be cycled reversibly between 0 and 1.8 V with an energy density of 17 Wh/kg at 2 kW/kg, much higher than those of the AC/K2SO4/AC supercapacitor and AC/Li2SO4/LiMn2O4 hybrid supercapacitor. Moreover, this supercapacitor exhibits excellent cycling behavior with no more than 6% capacitance loss after 23 000 cycles at 10C rate even when the dissolved oxygen is not removed. Introduction Electrochemical capacitors are energy storage devices which could provide higher power density compared to batteries and higher energy density than conventional electrostatic capacitors.1 They have promising applications as backup or auxiliary power sources in electric vehicles and other electronic devices for the purpose of power enhancement. Currently, much attention is paid to improvements of the energy density of electrochemical capacitors. As the energy density is proportional to the specific capacitance (F/g) and the square of the operating voltage, improvements of the energy can be achieved in two ways, i.e., the characteristics of the electrode materials and the selection of electrolytes. As far as the electrolytes are concerned, supercapacitors in aqueous electrolytes could deliver much higher power density than those in nonaqueous electrolytes due to their higher ionic conductivity. Moreover, aqueous electrolytes are environmentally friendly and do not require anhydrous atmosphere for the cell assembly.2,3 Although aqueous electrolytes are usually stable over a voltage window of 1.23 V, asymmetric aqueous supercapacitors, which make use of the different potential windows of two electrodes and the high oxygen and/or hydrogen overpotentials of the electrodes, could widen the maximum operating voltage of aqueous electrolytes.4 Asymmetric aqueous supercapacitors, such as activated carbon (AC)//LiMn2O4,5 AC//Ni(OH)2,6 AC//MnO2,7,8 and FeOOH// MnO2,9 were recently explored. MnO2, with the advantages of low cost, environmental friendliness, and the ability to charge-discharge rapidly, appears to be a promising material for electrochemical supercapacitors. Moreover, it can be used in neutral aqueous electrolytes, unlike * To whom correspondence should be addressed. Phone/fax: +86-2155664223. E-mail: [email protected] (Y.W.) and rudolf.holze@ chemie.tu-chemnitz.de (R.H.). † Fudan University. ‡ Technische Universita¨t Chemnitz, Institut fu¨r Chemie.

RuO2 · xH2O and NiOOH, which can only be used in strong acidic or alkaline electrolytes, thus causing environmental problems.10-16 In our previous study, we compared the electrochemical performances of AC in 0.5 mol/L Li2SO4, Na2SO4, and K2SO4 aqueous electrolytes, and it turns out that supercapacitors with K2SO4 electrolyte show the best rate behavior.17 In this work, we systematically investigated the electrochemical performance of MnO2 nanorods prepared by a precipitation reaction in the above-mentioned three aqueous electrolytes. The asymmetric supercapacitor AC//MnO2 also exhibited the best electrochemical performance in the K2SO4 electrolyte. Experimental Section MnO2 nanorods were synthesized by using a simple solution precipitation method. First, 200 mL of 0.1 mol/L MnSO4 solution was mixed with 200 mL of 0.1 mol/L K2S2O8 solution. Upon stirring this solution at room temperature, 100 mL of 1.2 mol/L NaOH was added dropwise, and a dark brown precipitate was immediately formed. At last, the precipitate was filtered, washed several times with deionized water, and dried at 60 °C for 24 h. The obtained MnO2 powder has a specific surface area of about 135 m2/g measured by the BET method. For comparison, δ-MnO2 microcrystals with a high crystallinity were prepared by thermal decomposition of KMnO4 at 800 °C, followed by washing with water to remove any soluble products. The crystalline structure of the products was characterized by X-ray diffraction (XRD), using a Rigaku D/MAX-Π A with Cu KR radiation source. Transmission electron micrographs (TEM) were obtained with a JEOL JEM-2010 transmission electron microscope. Scanning electron micrographs (SEM) were obtained on a Philip XL300 microscope. Elemental analysis was performed with a Thermo E. IRIS Duo inductively coupled plasma (ICP) instrument. Surface electronic states were

10.1021/jp8113094 CCC: $40.75  2009 American Chemical Society Published on Web 07/14/2009

MnO2 Nanorods in Neutral Aqueous Electrolytes investigated by X-ray photoelectron spectroscopy (XPS; PerkinElmer PHI 5000C ESCA, using Al KR radiation) and the binding energy values were calibrated by using C1s ) 284.6 eV as a reference. For electrochemical tests, the MnO2 electrode was prepared by the following methods. First, a thin film composed of a mushy mixture of MnO2, acetylene black, and poly(tetrafluoroethylene) (PTFE) in a weight ratio of 85:10:5 was prepared, and then punched into small disks with a diameter of 10 mm. At last, these disks were pressed onto Ni-grid at a pressure of 12 MPa and then dried at 70 °C for 5 h. Our preliminary investigations showed that the effect of electrolyte concentrations on the electrochemical behaviors of electrode materials is negligible when the electrolyte concentration exceeds 0.3 mol/ L;17 here we used 0.5 mol/L Li2SO4, Na2SO4, and K2SO4 solutions as the electrolytes. For comparison of the electrochemical properties of MnO2 electrode in the above three electrolytes, only electrode disks with the same mass were used, so as to lessen any possible errors since these disks possessed the same area and were pressed onto Ni-grid at the same pressure. The electrochemical tests of the individual MnO2 electrode were performed with use of a three-electrode cell, in which Ni-grid and saturated calomel electrode (SCE) were used as counter and reference electrodes, respectively. Cyclic voltammograms (CV) were collected between 0 and 1.0 V vs. SCE at different scan rates. From CV curves, the calculation of the specific discharge capacitance of the single MnO2 electrode is based on the following formula: Cs ) ∫ I dU/(2νm∆U), where I is the current (A), ∫I dU is the integration area for the CV curve of the MnO2 electrode, V is the scan rate (V/s), m is the mass (g) of the active material MnO2 in the composite electrode (not including the conductive and polymer additives), ∆U is the potential difference (V) during negative scan, and the factor 2 corrects the fact that the above integration area includes both the positive scan and the negative scan. The electrochemical impedance spectra were recorded from 105 to 0.01 Hz and the amplitude of the used perturbation was 10 mV. Activated carbon (AC) with a specific surface area of about 2800 m2/g measured by the BET method was purchased from Ningde Xinseng Chemical and Industrial Co., Ltd. and used as received without further treatment. The activated carbon electrode was prepared in the same way as the MnO2 electrode. A two-electrode cell consisting of the cathode and the anode with a distance of about 1 cm was used to test the electrochemical properties of the asymmetric supercapacitor AC//MnO2. From the charge and discharge curves, the specific capacitance of the asymmetric supercapacitor was calculated by using the formula C ) It/m∆U, where I is the current (A) used for charge/ discharge cycling, t is the discharge time (seconds), m is the total mass (g) of the two active electrode materials (including AC and MnO2), and ∆U is operating voltage window (V) of the capacitor during the discharge. Results and Discussion The crystalline structure and morphology of the as-prepared MnO2 are shown in Figure 1. The diffraction peaks at around 12.5° (001 planes) and 25.0° (002 planes) suggest that the obtained MnO2 belongs to a δ-type crystal with a layered structure.18,19 However, the crystallinity is lower than that prepared from sol-gel method since the diffraction peak for the 100 planes at around 37.6° and the other peaks could not be clearly identified.20 In the TEM micrographs it can be clearly seen that the prepared MnO2 consists of nanorods with a diameter of less than 10 nm. Elemental analysis by ICP shows

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Figure 1. (a) XRD pattern and TEM micrographs of the as-prepared MnO2 (b) at low magnification and (c) at high magnification.

that the as-prepared products contain a small portion of Na+ ions with a Na/Mn mole ratio of about 1/10, and the contents of other elements such as S and K are negligible. Figure 2 shows the cyclic voltammograms of MnO2 nanorods in 0.5 mol/L aqueous Li2SO4, Na2SO4, and K2SO4 electrolytes at scan rates of 1, 5, 10, 30, and 50 mV/s, respectively. At the slow scan rate of 1, 5, and 10 mV/s, Li2SO4-based electrolyte shows a roughly symmetric rectangular shape, typical of pseudocapacitive behavior. In addition, a small number of distinct redox peaks (at 0.37 and 0.64 V) is observed, which appear during the extended cycles and are also repeatable. These redox peaks can be regarded as the reversible intercalation/ deintercalation of lithium ions in the MnO2 solid phase, whereas in the other two electrolytes, no obvious redox peaks were observed. This can be ascribed to the smaller radius of the unsolvated Li+ (0.6 Å) compared with Na+ (0.95 Å) and K+ (1.33 Å), which makes the intercalation/deintercalation of Li+ ions in the MnO2 solid phase possible. This result is consistent

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Figure 2. Cyclic voltammograms of the MnO2 nanorods in 0.5 mol/L aqueous Li2SO4, Na2SO4, and K2SO4 electrolytes at various scan rates.

with that reported in the literature, where reversible Li+ intercalation/deintercalation into/from MnO2 can be found in 1 mol/L LiOH aqueous electrolyte, which is totally different from that in KOH electrolytes.21 However, these few redox peaks can only be observed at the slow scan rate. It can be seen that at the fast scan rates of 30 and 50 mV/s, there are no evident redox peaks corresponding to the intercalation/deintercalation of Li+ ions in the solid phase. Besides, all the CV curves at the fast scan rates deviate from the ideal capacitive behavior.22-24 Further studies of the different reaction activity of MnO2 electrode in the three electrolytes will be shown later in this report. The above phenomena can be explained more clearly from the schematics of the ionic motion (Figure 3). As for the charge storage mechanisms in MnO2, two mechanisms were proposed. One is based on adsorption/desorption of protons (H+) or alkaline cations (M+) on the MnO2 surface, accompanied by the valence conversion of Mn4+/Mn3+, which is likely to be

Figure 3. Schematics of the migration of hydrated alkaline ions during the potential scan.

predominant in amorphous MnO2. The other is based on intercalation/deintercalation of H+ or M+, also accompanied by faradic reaction, which may be predominant in crystalline MnO2. Since the crystallinity of the MnO2 nanorods we synthesized here is very low, the electrochemical reaction is expected to be mostly dominated by the former mechanism, which is verified

MnO2 Nanorods in Neutral Aqueous Electrolytes by the typical rectangular shape of the CV curves in Figure 2. However, it should be noted that these two mechanisms involve only the reaction of protons or cations on the electrode, anions are not involved in the electrode process. Therefore, here we concentrate on the motion of cations in the electrolyte. Since the only difference between the three neutral electrolytes is the alkaline cations, only the motion of hydrated Li+, Na+, and K+ cations is illustrated. When the potential is scanned negatively, these hydrated alkaline cations move to the surface of the MnO2 electrode from the electrolytes, and then desolvate and absorb on the MnO2 surface. On the reverse scan, these alkaline ions desorb from the MnO2 electrode and move back into the electrolytes. At the slow scan rates, all three kinds of hydrated alkaline ions have enough time to arrive at the surface of MnO2, and then desolvate, and the absorption reaction occurs favorably. Thus the nanorods show excellent pseudocapacitive behavior in all three electrolytes. In addition, the second charge storage mechanism of MnO2 involving intercalation/deintercalation of M+ cations into/from the solid MnO2 phase can be found in Li2SO4 electrolyte at the slow scan rates, which is suggested by a small number of redox peaks in CV curves. As a result, the MnO2 nanorods present the largest capacitance values (201 F/g at the scan rate of 1 mV/s) in the Li2SO4 electrolyte. At fast scan rates, some alkaline ions do not have enough time to absorb on the MnO2 surface due to the polarization of the desolvation process of the hydrated alkaline ions, and consequently the adsorption on the electrode surface was incomplete and all the CV curves deviate from the ideal rectangular shape. In the case of Li+, at the fast scan rate, the intercalation/ deintercalation of Li+ in the solid MnO2 phase is a rather tardy process with large polarization and Li+ does not have enough time to intercalate into the solid phase. In addition, since the charge density of Li+ ion is the largest among the three alkaline ions, its desolvation process would be the most difficult and the CV curve in the Li2SO4 deviates most seriously from the ideal capacitive behavior. XRD was used to investigate the crystalline changes of MnO2 electrode before cycling and after the first discharge in different electrolytes at the low current density of 200 mA/g and at the high current density of 2000 mA/g (Figure 4). Here the MnO2 electrodes were first charged and then discharged. The XRD pattern of the MnO2 electrode before cycling is similar to that of the original MnO2 powder, and the typical (001) and (002) peaks can still be observed. The diffraction peaks corresponding to the current collector Ni mesh appear for all the MnO2 electrodes. It can be seen that after the first discharge in Li2SO4 at the low current density of 200 mA/g, the crystallinity of MnO2 was maintained perfectly. However, after discharge in Na2SO4 and K2SO4 at the low current density, the crystallinity of MnO2 decreases significantly, which can be seen from the relative intensity of the (001) peak to the characteristic diffraction peak of Ni. When a large charge-discharge current density is used, the crystallinity of MnO2 decreases greatly after discharge in all three electrolytes. These phenomena can be explained as follows. As a small number of Na+ ions are incorporated into the interlayer space during the synthesis of MnO2 nanorods, Na+ ions may deintercalate from the interlayer during the first charge, and the deintercalation of Na+ may be accompanied by the intercalation of H2O.25 During the following discharge, cations in the electrolytes are supposed to intercalate into the interlayer. At the small current density, Li+ ions could intercalate into the interlayer space readily due to its small ionic radius, thus the layered structure of MnO2 was stabilized and the crystallinity

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Figure 4. XRD of the MnO2 electrodes before cycling and after the first discharge in different electrolytes (a) at the low current density of 200 mA/g and (b) at the high current density of 2000 mA/g. Inset: Magnification of the low-angle section.

of MnO2 was maintained. However, it is hard for Na+ and K+ ions to intercalate into the interlayer space because of their large size. As a result, there were not enough cations to stabilize the layered structure of MnO2 and MnO2 becomes more amorphous. At the large current density, Li+ ions do not have enough time to intercalate into the interlayer space, and consequently MnO2 becomes more amorphous, which is the same as that of MnO2 after discharge in Na2SO4 and K2SO4. In addition, after discharge at the large current density, the interlayer space of MnO2 increases slightly for all three electrolytes, as can be seen clearly from the shift of the (001) peak toward lower angles compared with those before cycling. This can be ascribed to the deintercalation of Na+ ions during the first charge,25 but no obvious cation intercalation occurs during the following discharge process, whereas after the discharge process at the small current density, the interlayer space of MnO2 was maintained most perfectly for the Li2SO4 electrolytes, which can be ascribed to the obvious intercalation of Li+ ions during the discharge process. At the MnO2 electrodes after discharge in Na2SO4 and K2SO4 at the low current density, a very small part of the Na+ or K+ ions intercalation also may be involved, but the change of the interlayer space is not obvious. XPS was used to investigate the surface electronic states of the MnO2 electrodes before cycling and after the first discharge in different electrolytes at the low current density of 200 mA/g and at the high current density of 2000 mA/g (Figure 5). The binding energy of O1s for the MnO2 before cycling (530.0 eV) is in agreement with that of the XPS standard database. After the first discharge in Li2SO4 at the low current density of 200 mA/g, the binding peak of O1s shifts significantly to a higher

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Figure 5. XPS of the MnO2 electrodes before cycling and after the first discharge in different electrolytes (a, b) at the low current density of 200 mA/g and (c, d) at the high current density of 2000 mA/g.

energy of 531.0 eV. At the same time, shifts of the binding peaks of O1s in the MnO2 electrodes after discharge in Na2SO4 and K2SO4 at the low current density are very slight. In addition, after the first discharge at the large current density of 2000 mA/ g, the binding energy peaks of O1s in the MnO2 electrodes exhibit a very slight shift for all three kinds of electrolytes. The above obvious shift of the binding energy of the O1s peak in the MnO2 electrodes after discharge in Li2SO4 at the low current density is consistent with the strange CV curve of the MnO2 electrode in the Li2SO4 solution at the slow scan rate, which was characterized by a small number of redox peaks. The positive shift of the binding energy peak of O1s can be ascribed to the intercalation of Li+ ions into the MnO2 lattice since Li-O coordination has a higher O1s binding energy of 531.3 eV. At a large current density, Li+ ions do not have enough time to intercalate into the solid lattice. As a result, the binding energy peak of O1s only shows a very slight shift. Because of the large sizes of Na+ and K+ ions, the binding energy peaks of O1s show no obvious shift at both small and large current densities. The binding energy peaks of Mn2p in the MnO2 electrode shift slightly toward lower position after discharge at the low current density in all three electrolytes, which could be due to the valence transformation of Mn4+ to Mn3+. At the large current density of 2000 mA/g, the binding energy peaks of Mn2p in the MnO2 electrodes show no obvious shifts after discharge in all three electrolytes. This is associated with the low capacity of the MnO2 electrode at the large current density. The specific capacitance values of the MnO2 nanorods were calculated by integrating the CV curves to obtain so-called voltammetric charges. The effect of the scan rate on the

Figure 6. Capacitance variations of the MnO2 nanorods at different scan rates in 0.5 mol/L aqueous Li2SO4, Na2SO4, and K2SO4 electrolytes.

capacitance of the MnO2 nanorods is shown in Figure 6. At the slow scan rate, the capacitance value in the three electrolytes increases in the order K2SO4 < Na2SO4 < Li2SO4. The main reason is that at slow scan rates, besides the absorption/ desorption reaction, the reversible intercalation/deintercalation of Li+ in the solid phase mentioned above produces an additional pseudocapacitance, thus leading to a capacitance higher than that obtained with Na2SO4 or K2SO4. This result is different from those reported where MnO2 showed no pseudocapacitive behavior in the LiCl-based electrolyte.26,27 This may be due to the different morphology of the prepared MnO2. In our experiment, the as-prepared MnO2 shows a nanorod morphology that makes the solvated Li+ reach the MnO2 surface easily, especially the inner surface, thus resulting in a high pseudocapacitance value. In the literature reports, the MnO2 particles were

MnO2 Nanorods in Neutral Aqueous Electrolytes

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Figure 7. The Nyquist plots of the MnO2 electrode after 20 cycles in 0.5 mol/L aqueous Li2SO4, Na2SO4, and K2SO4 electrolytes.

bigger, mostly of an average size of about 5 µm.27 As a result, the alkaline ions could not reach the MnO2 inner surface effectively, leading to noncapacitive behavior. When the scan rate is increased, the effective utilization for the redox reaction is limited to some extent, thus resulting in the decrease of the capacitance in all three electrolytes. As the hydrated ionic radius of K+ (3.31 Å) is the smallest and its ionic conductivity is the highest, its access to the inner surface of the MnO2 electrode is much easier and faster than that of Na+ (hydrated ionic radius: 3.58 Å) and Li+ (hydrated ionic radius: 3.82 Å) (see Figure 3).28,29 Besides, as the charge density of the K+ ion is the smallest, its polarization for the desolvation process is the smallest. As a consequence, in K2SO4-based electrolyte the MnO2 nanorods show the largest capacitance at the fast scan rates. The rate behavior of MnO2 in the three electrolytes is similar to that of activated carbon.17 The different electrochemical behavior of MnO2 in the three neutral electrolytes described above suggests that the alkaline cations are indeed involved in the reaction of the MnO2 electrode, which provides useful information for studies on the charge storage mechanism of MnO2. The Nyquist plots of the MnO2 nanorods electrode after 20 cycles in 0.5 mol/L aqueous Li2SO4, Na2SO4, and K2SO4 electrolytes are shown in Figure 7. On the whole, the Nyquist plots of the MnO2 electrode in the three electrolytes are similar to each other, consisting of a semicircle at midhigh frequency and a linear region at low frequency. The linear region of the plot exhibits an angle of about 60° relative to the real axis, indicating that the electrode process is not perfectly capacitive in nature but is also under diffusion control. The semicircles correspond to a parallel combination of charge-transfer resistance (Rct) and double-layer capacitance.30 It can be seen that the diameters of the semicircles decrease in the order of Li2SO4 > Na2SO4 > K2SO4, indicating that Rct is the smallest in the K2SO4 electrolyte, i.e., the electrochemical reaction on the electrode/electrolyte interface is the most facile. In addition, as can be seen from the value of the intercept at the real axis at high frequency, the electrolyte resistance decreases in the order of Li2SO4 > Na2SO4 > K2SO4. These results show that the equivalent series resistance (ESR) values for the three electrolytes decrease in the order of Li2SO4 > Na2SO4 > K2SO4. According to the equation Pmax (the maximum power) ) U02/ 4ESR, the MnO2 electrode is bound to possess the largest Pmax and exhibit the best rate behavior in the K2SO4 electrolyte. Figure 8 shows the changes of electrochemical impedance spectra of the MnO2 electrode before and after 20 cycles in the Li2SO4, Na2SO4, and K2SO4 electrolytes. For the purpose of

Figure 8. Changes of the electrochemical impedance spectra of the MnO2 nanorods before and after 20 cycles in 0.5 mol/L aqueous electrolytes: (a) Li2SO4, (b) Na2SO4, and (c) K2SO4.

observing the semicircles more clearly, here the Nyquist plots were presented in the frequency regions from 100 kHz to 1 or 0.31 Hz. It can be seen that the diameters of the semicircles decrease slightly after cycling for all three electrolytes, which is similar to that of activated carbon.17 It is believed that the initial several cycles are an activation process, which makes the electrolytes soak into the electrode surface completely. Furthermore, the MnO2 electrode is fully activated and the redox reaction on the entire surface will become easier. This activation mechanism is different from that of lithium ion batteries, which is always accompanied by the formation of a solid-electrolyteinterface (SEI) film.31 These results show that the MnO2 nanorods exhibit the best rate behavior in the K2SO4 electrolyte, which is similar to that of activated carbon.17 An asymmetric supercapacitor, AC//MnO2, was assembled by using 0.5 mol/L K2SO4 as the electrolyte. Figure 9 shows the potential-time curves of the individual electrode vs. SCE reference electrode and the voltage-time profile of the asymmetric AC//MnO2 aqueous supercapacitor

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Figure 9. Potential-time curves of the individual electrode vs. SCE reference electrode and the voltage-time profile of the asymmetric AC// MnO2 aqueous supercapacitor at a current rate of 2C in 0.5 mol/L K2SO4 electrolyte.

Figure 10. Cycling behavior of the asymmetric AC/0.5 mol/L K2SO4/ MnO2 aqueous supercapacitor at a current rate of 10C between 0 and 1.8 V.

at a current rate of 2C. The potential of the AC anode shows a typical linear relationship with time, characteristic of an electric double layer capacitance. The potential-time curves of the MnO2 cathode also exhibit an almost linear shape due to its fast pseudocapacitive nature. In the asymmetric AC//MnO2 aqueous supercapacitor, the weight ratio of AC to MnO2 was fixed at about 1:1, which was obtained from the specific capacitance of the AC electrode (200 F/g) and the MnO2 electrode (175 F/g) in 0.5 mol/L K2SO4. The asymmetric supercapacitor shows a sloping voltage profile from 0 to 1.8 V with excellent reversibility, and its specific discharge capacitance could be up to 53.7 F/g based on the total mass of the two active electrode materials. Figure 10 shows the cycling behavior of the asymmetric AC/ 0.5 mol/L K2SO4/MnO2 supercapacitor at a current rate of 10C between 0 and 1.8 V. Here we did not take any measures to

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Figure 12. Ragone plots of the asymmetric AC//MnO2 nanorods, AC// MnO2 microcrystals supercapacitor, and symmetric AC//AC supercapacitor in 0.5 mol/L K2SO4 electrolyte.

remove the dissolved oxygen from the electrolyte. This is quite different from previous reports, where oxygen is usually removed.7 The supercapacitor shows excellent cycling behavior with no more than 6% capacitance loss after 23 000 cycles. The excellent cycling performance of the supercapacitor can be ascribed to the following two aspects: (1) The AC anode stores energy through a nonfaradic reaction, which is well-known for its excellent cycle life, and (2) MnO2 stores energy through a fast absorption/desorption reaction that mainly occurs on the active surface due to the amorphous nanostructure of the prepared MnO2. The cycling behavior is much better than that reported by Brousse et al.20 In their report, there is a 24% decrease of energy density after 23 000 cycles in a narrower voltage range of 0 and 1.5 V. The main reason is presumably that the MnO2 material we prepared is more amorphous and has a unique nanorod structure. In addition, the SEM images of the MnO2 electrode (Figure 11) before and after 23 000 cycles show no obvious morphology changes after the long cycling. The nanorod morphology of MnO2 seems to be beneficial for the stability of the MnO2 electrode. Figure 12 shows the Ragone plots of the asymmetric supercapacitor AC//MnO2 nanorods, AC//MnO2 microcrystals, and symmetric supercapacitor AC//AC. All the data were calculated based on the total mass of the two active electrode materials. It can be seen clearly that the asymmetric supercapacitor has a much higher energy density than the symmetric supercapacitor and at the same time keeps a very high power density. The energy density of the asymmetric supercapacitor based on MnO2 nanorods is 28.4 Wh/kg at 150 W/kg and maintains 17 Wh/kg at 2 kW/kg, much higher than the asymmetric supercapacitor based on MnO2 microcrystals. The energy density of the AC/K2SO4/MnO2 nanorods supercapacitor at high power is also much higher than that of the most recently

Figure 11. SEM images of the MnO2 electrodes (a) before and (b) after 23 000 cycles.

MnO2 Nanorods in Neutral Aqueous Electrolytes described AC/Li2SO4/LiMn2O4 hybrid supercapacitors, which only have an energy density of 10 Wh/kg at a power density of 2 kW/kg.5 This can be ascribed to the fast pseudocapacitance reaction on MnO2 nanorods surface in contrast to the slow Li+ insertion/extraction reaction in the LiMn2O4 solid phase. In addition, in this asymmetric AC/K2SO4/MnO2 nanorods capacitor, the preparation of MnO2 nanorods is very easy and the K2SO4 electrolyte is cheaper and more easily available than Li2SO4. All these advantages strongly suggest its promising commercial applications in electric vehicles or other large power devices. Conclusions The electrochemical behavior of MnO2 nanorods prepared from a precipitation reaction in 0.5 mol/L Li2SO4, Na2SO4, and K2SO4 aqueous electrolytes was investigated. The MnO2 nanorods show the superior rate behavior in the K2SO4 electrolyte due to the smallest hydrated radius of K+, highest ionic conductivity, and smallest equivalent series resistance (ESR). Moreover, the initial cycles could be an effective way to activate the MnO2 electrode and make the electrolyte solution soak into the material surface completely. The assembled asymmetric AC/ K2SO4/MnO2 nanorods supercapacitor shows an excellent cycling behavior between 0 and 1.8 V and also exhibits a large energy density of 17 Wh/kg at a power density of 2 kW/kg. This supercapacitor has promising applications because of its advantages of low-cost, high energy and power density, and environmentally benign properties. Acknowledgment. Financial support from National Basic Research Program of China (973 Program No. 2007CB209702), Shanghai Committee of Science and Technology (09QH1400400), and Alexander von Humboldt (Institutional Academic Cooperation Program) is gratefully acknowledged. References and Notes (1) Conway, B. E. Electrochemical Supercapacitors Scientific Fundamentals and Technological Applications; Kluwer Academic/Plenum Press: New York, 1999. (2) Li, W.; Dahn, J. R.; Wainwright, D. Science 1994, 264, 1115– 1118. (3) Wang, G.; Fu, L.; Zhao, N.; Yang, L.; Wu, Y. P.; Wu, H. Angew. Chem., Int. Ed. 2007, 46, 295–297. (4) Hong, M. S.; Lee, S. H.; Kim, S. W. Electrochem. Solid-State Lett. 2002, 5, A227–A230.

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