Structural and Electrochemical Characterization of Nanocrystalline Li

compounds of Li[Li(1-2x)/3NixMn(2-x)/3]O2 seem to be a very promising candidate as a cathodic material for the new genera- tion of Li-ion batteries du...
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J. Phys. Chem. B 2005, 109, 1148-1154

Structural and Electrochemical Characterization of Nanocrystalline Li[Li0.12Ni0.32Mn0.56]O2 Synthesized by a Polymer-Pyrolysis Route Lihong Yu, Hanxi Yang, Xinping Ai, and Yuliang Cao* Department of Chemistry, Wuhan UniVersity, China, 430072 ReceiVed: August 7, 2004; In Final Form: September 30, 2004

A nanocrystalline Li[Li0.12Ni0.32Mn0.56]O2 powder was synthesized via the pyrolysis of polyacrylate salt precursors prepared by in situ polymerization of the metal salts and acrylate acid. The pyrolysis mechanism of the polymeric precursor is discussed by use of thermal analysis. Layered crystalline structure of the resulting compound was characterized by powder X-ray diffraction and high-resolution transmission electron (TEM) micrographs. The results revealed that the Li[Li0.12Ni0.32Mn0.56]O2 as prepared has nanosized morphology and good crystallinity even if calcined at 900 °C, and the electrodes made from these nanoparticles exhibited high-rate characteristics and can deliver stable discharge capacity (163 mA‚h/g) with excellent capacity retention at appropriate charge-discharge voltage interval. Since this preparation method is simple and particularly suitable for preparation of mixed metal oxides, it can be used for industrial production of highly homogeneous metal oxide cathode materials of the lithium-ion batteries.

Introduction Layered LiMnO2 have been actively investigated in recent years as an alternative cathodic material for rechargeable Liion batteries because of their possible higher capacity and lower cost compared with currently used LiCoO2. The major drawback of layered LiMnO2 is the capacity decay due to the phase transformation of the layered structure to a spinel structure by Jahn-Teller distortion during charge-discharge cycles.1-7 To solve this problem, a number of research groups have tried to stabilize the layered structure by introducing a LiMO2 (M ) Ni, Co, Cr) phase into the layered Li2MnO3 to form a solid solution of mixed transition-metal oxides.8-14 Like layered LiMnO2, Li2MnO3 or Li[Li1/3Mn2/3]O2 has a perfect layered structure, which is composed of a pure lithium layer, a mixed metal layer ([Li1/3Mn2/3]), and a pure oxygen layer.15 Due to the existence of Mn4+ in Li2MnO3 structure, layered Li2MnO3 shows remarkable structure stability without Jahn-Teller distortion. However, because Mn4+ cannot be further oxidized to a higher oxidation state, Li2MnO3 is electrochemically inactive and cannot be used as a cathodic material. Nevertheless, in the LiMO2-Li2MnO3 solid solution, M can take part in oxidationreduction reaction during charge-discharge, whereas Mn4+ ions only act for stabilizing the layered structure. These chemical and structural properties enable the solid solution not only to deliver high capacity but also to have good cycling stability.16 Among the solid solutions being developed, the layered compounds of Li[Li(1-2x)/3NixMn(2-x)/3]O2 seem to be a very promising candidate as a cathodic material for the new generation of Li-ion batteries due to their thermal stability and good electrochemical performances. In previous studies, there are mainly two methods reported for preparation of the Li[Li(1-2x)/3NixMn(2-x)/3]O2 compounds. The most-used one was the mixed hydroxide method17 and the other was sol-gel method.18,19 The resulting compounds prepared from these methods have better crystal homogeneity and electrochemical performances, but the synthetic procedures * Corresponding author: Tel +86-027-68754526; fax +86-027-87884476; e-mail [email protected].

are very time-consuming, involving complicated pretreatment processes. Thus, it is of practical importance to seek for a simplified synthetic process for preparation of the layered LiNi-Mn oxides. In this work, we report a new and simple method for the synthesis of nanocrystalline Li[Li0.12Ni0.32Mn0.56]O2 compounds by pyrolysis of Li-Ni-Mn polyacrylates precursors prepared via in situ polymerization and describe the structural and electrochemical characteristics of the Li[Li0.12Ni0.32Mn0.56]O2 compounds. Experimental Section Li[Li(1-2x)/3NixMn(2-x)/3]O2 materials of various stoichiometry have been reported adequately in previous references.15,18,19 When the content of nickel is small such as x ) 0.25, the material shows high irreversible capacity and when the Ni content reaches x ) 0.5, the material shows low discharge capacity. Therefore, we chose Li[Li0.12Ni0.32Mn0.56]O2 material to demonstrate the novel polymer-pyrolysis method for preparation of inserting materials of mixed oxides. Nanocrystalline Li[Li0.12Ni0.32Mn0.56]O2 powders were prepared by a polymer-pyrolysis method using polyacrylates of Li, Ni and Mn as precursor compounds. The copolymeric precursor was made by solution polymerization of the mixed aqueous solution of acrylic acid in the presence of LiOH, Ni(NO3)2, and Mn(NO3)2, with (NH4)2S2O8 as initiator. The typical experimental procedure is first to dissolve 10.8 g of LiOH‚H2O (95%), 18.8 g of Ni(NO3)2‚6H2O (98%), and 40.2 g of Mn(NO3)2 (50% aqueous solution) in 63 g of acrylic acid aqueous solution (acrylic acid:H2O ) 70:30 wt %) under stirring. Then a small amount (2 g) of 5% (NH4)2S2O8 aqueous solution as initiator was added to the mixed acrylic acid solution to promote the polymerization. Under heating at ∼80 °C for 2 h, the mixed solution was dried to form the well-distributed polyacrylates of Li, Ni, and Mn. Afterward the resulting polyacrylates were dried at 120 °C for 24 h. The obtained copolymeric precursor was then decomposed at 450 °C for 5 h in air to get the powders of a layered Li-Ni-Mn oxide. To obtain homogeneous crystalline structure of the compounds, the powders were finally

10.1021/jp0464369 CCC: $30.25 © 2005 American Chemical Society Published on Web 01/04/2005

Nanocrystalline Li[Li0.12Ni0.32Mn0.56]O2

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Figure 1. FTIR spectra of (a, top panel) acrylic acid and (b. bottom panel) the copolymeric precursor of Li-Ni-Mn polyacrylates.

calcined at 450, 600, and 900 °C for 3 h in air, respectively, and then quenched to room temperature. The actual stoichiometric compositions of the as-prepared powders were determined by inductively coupled plasma (ICP) analysis. The individual polyacrylate of Li, Ni, or Mn was prepared by dissolving stoichiometric LiOH, Ni(NO3)2, or Mn(NO3)2 respectively, in acrylic acid and then polymerizing the precursor solution at 80 °C, followed by drying at 120 °C for 24 h. The structural feature of the polymeric Li-Ni-Mn-O precursor was determined by Fourier transform infrared spectroscopy (FTIR) and conducted on a Thermo Nicolet FTIR spectrometer (Model AVATAR 360 FT-IR). The IR samples were made by pressing the mixture of the polymeric precursor with KBr into a disk, and the spectra were obtained at a 2 cm-1 resolution and with the average of 20 scans, processed by an EZ OMNIC spectroscopy software package. The decomposition processes of the polymer precursor during heat treatment were characterized by thermogravimetric and differential thermal

analysis (TG/DTA) on a model WCT-1A thermobalance (Beijing Optical Instrument Factory) at temperature range of 30-950 °C with a heating rate of 10 °C/min in air. To reveal the crystalline structure of the Li[Li0.12Ni0.32Mn0.56]O2 powders, X-ray diffraction (XRD) and high-resolution transmission electron microscopy (TEM) were carried out on a Shimadzu XRD-6000 diffractometer with Cu KR and on a JEM-2010FEF TEM system, respectively. The cathode used in this work with a geometrical area of 2 cm2 consisted of 80% Li[Li0.12Ni0.32Mn0.56]O2 powders, 12% acetylene black, and 8% poly(tetrafluoroethylene) (PTFE) by weight. The electrode was prepared by mixing the Li[Li0.12Ni0.32Mn0.56]O2 powders, acetylene black, and PTFE emulsion with 2-propanol to form an electrode paste, then rolling the paste into a ca. 0.1-mm-thick film, and finally pressing the electrode film onto an aluminum foil. The galvanostatic charge-discharge measurements were performed in a three-electrode cell of a sandwiched design. The

1150 J. Phys. Chem. B, Vol. 109, No. 3, 2005

Figure 2. Schematic representation of polymeric chain for the copolymeric precursor of Li-Ni-Mn polyacrylates.

Li[Li0.12Ni0.32Mn0.56]O2 electrode and lithium electrode were separated by a microporous membrane (Celgard 2400) and pressed in parallel together by a pair of electrode holders. A small piece of lithium foil immersed in the electrolyte served as a reference electrode. The electrolyte was 1 mol L-1 LiPF6 dissolved in a mixed solution of ethylene carbonate (EC)dimethyl carbonate (DMC)-ethylene methyl carbonate (EMC) (1:1:1 w/w/w). All the cells were assembled in a drybox filled with argon gas. The galvanostatic charge-discharge test was conducted by a BTS-55 Neware battery testing system. Results and Discussion The polymer-pyrolysis method used in this work is to synthesize the copolymeric precursor of the mixed metallic ions and then to pyrolyze the precursor into the required oxide structure. The polymerization reaction of acrylate has been well studied and discussed in detail in previous papers.20,21 It is now well accepted that the reaction proceeds mainly through a free radical polymerization. Figure 1 compares the standard IR spectrum of acrylic acid (99%) with that of the precursor compound of copolymerized Li-Ni-Mn polyacrylates. In the standard spectrum of acrylic acid (Figure 1a), a strong vibration band appears at 1702 cm-1, characteristic of the CdO stretching mode of R,β-unsaturated carboxylic acid. All other vibrations, such as O-H stretching (3600-2500 cm-1), CdC stretching (∼1635 cm-1), (COO)- stretching (∼1616 and ∼1433 cm-1), C-O stretching (∼1297 and ∼1242 cm-1), and dC-H vibration (∼1045 and ∼984 cm-1), are all observed in the spectrum. In the spectrum of the copolymeric Li-Ni-Mn precursors (Figure 1b), the CdO stretching vibration (∼1702 cm-1) of R,βunsaturated carboxylic acid, the CdC stretching vibration (∼1640 cm-1) and the dC-H vibrations (∼1045 and ∼984 cm-1) of acrylic acid, characteristic of the vinylidine group in the acrylic monomers, are absent, suggesting that the polymerization of acrylates took place in the acrylates of Li, Mn, and Ni. Furthermore, the spectrum shows the appearance of the symmetric (∼1618 cm-1) and asymmetric (∼1451 cm-1) stretching bands of carboxylate salts. Compared with those in Figure 1a, these two bands are positively shifted, implying an associating interaction between the metallic ions and carboxylate ions. In addition, the IR bands for H2O (∼3428 cm-1) and NO3vibration (∼1384 and ∼839 cm-1) are also observable due to the residual H2O and NO3-. The copolymeric precursor compound is schematically represented in Figure 2. As it is shown, the metallic Li, Mn, or Ni

Yu et al. ions are bound by the strong ionic bonds between the metallic ions and carboxylate ions in a polymeric chain or between the polymeric chains. This uniform immobilization of metallic ions in the polymer chains favors the formation of uniformly distributed solid solution of the metallic oxides in the following pyrolysis process. To clarify the chemical reactions of the copolymeric precursor occurring in the pyrolysis process, we measured TG and DTA curves of the Li polyacrylate (LiPA), Ni polyacrylate (NiPA), Mn polyacrylate (MnPA), and the copolymeric Li-Ni-Mn precursor, separately, as shown in Figure 3. In the TG curve of LiPA (Figure 3a), there are two weight-losing processes present at 160∼530 and 530∼860 °C with total weight loss of 76.8%. The first step of weight loss is ca. 52.2% of total weight, corresponding to two exothermic peaks in the DTA curve, possibly due to the decomposition of LiPA into lithium carbonate (Li2CO3). In the second step of weight loss, the weight loss is ca. 24.8% and correspondingly, a wide endothermic peak in the DTA curve is observed in the temperature range, which is likely to arise from the decomposition of Li2CO3. In the thermal data of NiPA (Figure 3b), two exothermic peaks appear at 190∼335 and 335∼385 °C, respectively, corresponding to two steps of weight loss. The first step shows a weight loss of 42.1% and the latter has a weight loss of 23.6%, which may correspond to the decomposition of NiPA to form nickel carbonate (NiCO3) and the decomposition of NiCO3, respectively. Figure 3c shows the TG and DTA curves of MnPA. The main feature in the DTA curve is a single exothermic peak at 235 °C, corresponding to a weight loss reaction between 115 and 340 °C. The total weight loss is 55.8%, possibly due to the decomposition of MnPA to form MnO2. The TG and DTA curves of the copolymeric Li-Ni-Mn precursor are shown in Figure 3d. The total weight loss is 66.0% and it occurred in three steps up to 545 °C. The DTA curve could be regarded as an overlapped curve of the LiPA, NiPA, and MnPA in Figure 3a-c. In the first step of the weight loss (160∼245 °C), a 10.2% weight loss associated with an exothermic peak at 190 °C may be attributed to the decomposition of NiPA to form NiCO3. In the DTA curve, there is a peak at 295 °C, corresponding to the second step of the weight loss (245∼330 °C), due to the decomposition of MnPA. In the third step (330∼545 °C), the exothermic peaks at 385 and 425 °C may be related to the decomposition of NiCO3 and LiPA, respectively. In particular, a small weight loss can be seen in the temperature range 450∼545 °C, which may be attributed to the decomposition of Li2CO3, in agreement with the results shown in Figure 3a. The results in Figure 3d show that the thermal events detected in the copolymeric precursors are the same as those in Figure 3a-c. The observed shifts of the temperatures for the similar thermal events are due to the differences in the strength of the interaction between metallic cations and the polymeric chain. Table 1 gives the experimental TG data and the proposed reactions based on the theoretical analysis of the weight losses. Clearly, the experimental data from the TG measurements are in good agreement with the theoretical values expected from the proposed reaction schemes. Figure 4 shows the XRD patterns of the Li[Li0.12Ni0.32Mn0.56]O2 powders pyrolyzed from the copolymeric precursor and calcined at 450, 600, and 900 °C. Except for a very weak and broad band between 20° and 25°, all the XRD peaks in Figure 4 can be very well indexed from a hexagonal R-NaFeO2 structure (space group R3hm, 166),18 suggesting that the pyro-

Nanocrystalline Li[Li0.12Ni0.32Mn0.56]O2

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Figure 3. TG/DTA curves of (a) LiPA, (b) NiPA, (c) MnPA, and (d) the copolymeric precursor of Li-Ni-Mn polyacrylates.

TABLE 1: Thermogravimetric Analysis possible reaction LiPA

∆W1

LiAA 98 Li2CO3 ∆W2

Li2CO3 98 Li2O ∆W

LiAA 98 Li2O NiPA

∆W1

Ni(AA)2 98 NiCO3 ∆W2

NiCO3 98 NiO ∆W

Ni(AA)2 98 NiO MnPA

Mn(AA)2 98 MnO2

copolymeric precursor

[LiAA]1.12[Ni(AA)2]0.32[Mn(AA)2]0.56 98 Li1.12Ni0.32Mn0.56O2

∆W

∆W

∆W1

(1) Ni(AA)2 98 NiCO3 ∆W2

(2) Mn(AA)2 98 MnO2 NiCO3 f NiO (3) ∆W3 LiAA f Li2O

}

lyzed oxide has the same structure as layered LiCoO2 and Li2MnO3.22,23 Compared with the well-documented data in literature,24 the positions and relative intensities of the XRD lines in Figure 4 agree very well with the layered Li[Li0.12Ni0.32Mn0.56]O2 synthesized by sol-gel or mixed hydroxide methods; even the small feature of the weak band at ∼22° is also similar to those reported in refs 11 and 25. This small band has already been assigned to a short-range superlattice ordering of the Li,

theoretical weight loss (%)

experimental weight loss (%)

temperature range (°C)

52.6

52.2

160-530

28.2

24.8

530-860

80.9

76.8

160-860

40.9

42.1

190-335

21.9

23.6

335-385

62.8

65.5

190-385

55.9

55.8

115-340

65.9

66.0

160-545

10.0

10.2

160-245

23.5

23.3

245-330

32.4

330-545

5.4 27.0

Ni, and Mn atoms in the transition metal layers.18 As shown in Figure 4, all the XRD peaks of the Li[Li0.12Ni0.32Mn0.56]O2 powders become sharper and stronger with increasing calcining temperature from 450 to 900 °C, suggesting that the material has an improved crystallinity at higher calcination temperature. The lattice parameters of the Li[Li0.12Ni0.32Mn0.56]O2 powders calcined at 900 °C calculated from the XRD pattern based on the R-NaFeO2 structure model are a ) 2.871 Å and c ) 14.255

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Yu et al.

Figure 4. XRD patterns of the Li[Li0.12Ni0.32Mn0.56]O2 powders calcined for 3 h at (a) 450, (b) 600, and (c) 900 °C. The labeled Miller indices are based on R-NaFeO2 structure.

Å, in accordance with literature values.12,15,16,18,24,25,28-33 A noticeable feature in Figure 4 is a pair of peaks at 30∼32° for the Li[Li0.12Ni0.32Mn0.56]O2 powders calcinated at 450 °C (Figure 4a), indicative of the presence of Li2CO3. When the synthesis temperature increases to 600 °C, the XRD peaks of lithium carbonate disappear, indicating that the Li[Li0.12Ni0.32Mn0.56]O2 powders are free of lithium carbonate. These results are in good agreement with those of thermal analysis in Figure 3d. Figure 5a shows the TEM bright-field image of the Li[Li0.12Ni0.32Mn0.56]O2 powders calcined at 900 °C. It can be seen that the powders are composed of ultrafine particles with uniformly distributed size of ca. 70∼100 nm. No doubt, such nanosized particles would facilitate reducing the diffusion length of the lithium ions for intercalation and deintercalation. Figure 5b shows the high-resolution TEM (HRTEM) image and the magnified spot of the as-prepared powders with a [120] zone, which was calculated from the Fourier transform pattern of Figure 5b. This image clearly shows good crystallinity of the material with some stacking faults. These faults appear to arise from the cationic disorder due to occupation of Li atoms in the Ni, Mn sites as well as the stacking disorder.34 Also, the magnified image reveals a ∼0.47 nm modulation, which is equivalent to the spacing of (003) planes in rhombohedral structure and is in agreement with the results obtained in XRD analysis (Figure 4). In comparison with other preparation methods reported previously,17-19 the complex oxide synthesized from the polymerpyrolysis method shows a uniform morphology and narrow size distribution, and the preparation process is much simpler without coprecipitating, gelling, and grinding processes. However, it should be mentioned that the pretreatment process in the polymer-pyrolysis method can produce some oxides of nitrogen and CO2 due to the decomposition of the residual NO3- and polyacrylate salts in a similar way as the sol-gel method. But the waste gas could be treated during drying and pyrolysis process. Thus, the polymer-pyrolysis method is still a commercially worthy route compared with other methods. Figure 6 shows the charge-discharge curves of Li[Li0.12Ni0.32Mn0.56]O2 electrodes between 2.5 and 4.8 V for specific currents of 20, 50, 100, 200, and 400 mA/g, respectively. The Li[Li0.12Ni0.32Mn0.56]O2 electrodes at different specific currents show similar charge-discharge profiles. In the first charge process, the charge curve is composed of a sloping part before 4.4 V and a plateau around 4.5 V. According to the previous assignments in refs 12 and 15, the sloping part at the charge

Figure 5. (a, top panel) TEM bright-field image and (b, bottom panel) high-resolution TEM image in [120] zone of the Li[Li0.12Ni0.32Mn0.56]O2 powders calcinated at 900 °C.

Figure 6. Initial charge-discharge curves of the Li[Li0.12Ni0.32Mn0.56]O2 electrodes cycled between 2.5 and 4.8 V at current densities of (a) 20, (b) 50, (c) 100, (d) 200, and (e) 400 mA/g.

curve is due to the oxidation of Ni2+ to Ni4+ when Li+ is extracted. The voltage plateau around 4.5 V during the first charge process corresponds to the removal of the remaining lithium from the lithium layer along with the simultaneous ejection of oxygen.30 As shown in Figure 6, the Li[Li0.12Ni0.32Mn0.56]O2 electrode delivered an initial discharge capacity of 213 mA‚h/g at a low specific current of 20 mA/g and a

Nanocrystalline Li[Li0.12Ni0.32Mn0.56]O2

Figure 7. Capacity retention versus cycle number for Li/Li[Li0.12Ni0.32Mn0.56]O2 cell cycled between 2.5 and 4.8 V at current densities of (a) 20, (b) 50, (c) 100, (d) 200, and (e) 400 mA/g.

Figure 8. First two charge-discharge curves for Li/Li[Li0.12Ni0.32Mn0.56]O2 cell cycled at a current density of 50 mA/g at (a) 2.5 and 4.8 V; (b) 2.5 and 4.5 V; (c) first charged to 4.8 V and then cycled at 2.5-4.5 V.

reversible discharge capacity of 147 mA‚h/g at a high specific current of 400 mA/g, which are superior to those by the solgel method19,34 and by the mixed hydroxide method.15,18 The high-rate capability is probably due to the uniformly distributed nanosize of the Li[Li0.12Ni0.32Mn0.56]O2 powders prepared by the polymer-pyrolysis method. Since the solid-state diffusion of Li ions is the rate-determining step for the Li+ insertion reaction, the smaller particles would result in a larger diffusion rate. Figure 7 shows the capacity retention versus cycle number for the Li/Li[Li0.12Ni0.32Mn0.56]O2 cells cycled between 2.5 and 4.8 V at various rates, respectively. It can be seen that the capacity fading rates at low specific current (20 mA/g) and at high specific current (400 mA/g) are faster than those at the moderate specific currents. This is because the layered structure may be severely damaged at deep charge and discharge at low specific current and also cannot tolerate the high current due to the structure rearrangement during Li+ intercalation and deintercalation at high specific current. Thus, it is suitable for the Li[Li0.12Ni0.32Mn0.56]O2 electrode to cycle at moderate current rates to achieve a good cyclic performance. We found that the capacity behaviors of the Li[Li0.12Ni0.32Mn0.56]O2 electrode have a strong dependence on the voltage intervals set for charge and discharge. Figure 8 shows the first two charge-discharge curves for the Li/Li[Li0.12Ni0.32Mn0.56]-

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Figure 9. Capacity retention versus cycle number for Li/Li[Li0.12Ni0.32Mn0.56]O2 cell cycled at a current density of 50 mA/g at (a) 2.5 and 4.8 V; (b) 2.5 and 4.5 V; (c) first charged to 4.8 V and then cycled at 2.5-4.5 V.

O2 cells cycled in different voltage range at a constant current density of 50 mA/g. When the cells were cycled only at 2.5∼4.8 V (Figure 8a), the cells exhibited an initial discharge capacity of 189 mA‚h/g, whereas the initial discharge capacity was only 140 mA‚h/g when the cells were cycled at 2.5-4.5 V (Figure 8b). But, no significant change in the discharge capacity was observed in the second cycling curves, except an excess of discharge capacity is observed between 3.5 and 2.5 V when the cell is charged to 4.8 V (Figure 8a). This excess capacity has been considered to arise from the redox reaction of Mn below 3.5 V because of the ejection of oxygen at the voltage plateau around 4.5 V.30 However, when the cell was first charged to 4.8 V and then cycled at 2.5∼4.5 V (Figure 8c), it delivered a reversible capacity of 163 mA‚h/g in the second discharge. The second discharge curve of the cell below 3.5 V in Figure 8c is similar to that of the cell cycled only at 2.5∼4.8 V (Figure 8a) but different from that of the cell cycled only at 2.5∼4.5 V (Figure 8b), indicating that there is 18% higher discharge capacity than that of the cell cycled only at 2.5∼4.5 V. Figure 9 shows discharge capacity retention versus cycle number for the Li/Li[Li0.12Ni0.32Mn0.56]O2 cells cycled in different voltage range at a constant current density of 50 mA/ g. When the cell was cycled only at 2.5∼4.8 V, the initial discharge capacity of the Li[Li0.12Ni0.32Mn0.56]O2 electrode was 189 mA‚h/g, and the subsequent capacity decreased gradually and reduced to 133 mA‚h/g after 22 cycles. When the cell was cycled only at 2.5∼4.5 V, the discharge capacity of the Li[Li0.12Ni0.32Mn0.56]O2 electrode remained about 135 mA‚h/g during cycling. However, when the cell was first charged to 4.8 V and then cycled at 2.5∼4.5 V, it delivered a stable discharge capacity of about 160 mA‚h/g after the second cycle. The results indicate that the Li/Li[Li0.12Ni0.32Mn0.56]O2 cell exhibits high capacity and good cyclic stability when first charged to 4.8 and then cycled at 2.5∼4.5 V. Conclusions This paper describes a novel route of producing highly homogeneous nanocrystalline Li[Li0.12Ni0.32Mn0.56]O2 powders by the pyrolysis of a Li-Ni-Mn copolymeric precursor in situ polymerized by reaction of the metal salts and acrylic acid. The chemical reactions occurring during the pyrolysis and the following heat-treatment processes were discussed on the basis of the TG/DTA data. The results showed that the pyrolysis process of the Li-Ni-Mn copolymeric precursor is in good

1154 J. Phys. Chem. B, Vol. 109, No. 3, 2005 agreement with that of the individual metal (Li, Ni, Mn) polymeric precursor. The data from the XRD and TEM measurements revealed that the Li[Li0.12Ni0.32Mn0.56]O2 powders even calcined at 900 °C have good crystallinity with the layered R-NaFeO2 structure and a uniformly distributed nanosize of ca. 70∼100 nm. The charge-discharge experiments demonstrate that the Li[Li0.12Ni0.32Mn0.56]O2 electrode has a quite high capacity (163 mA‚h/g) and good cycling stability during defined cycling voltage intervals. In addition, the as-prepared Li[Li0.12Ni0.32Mn0.56]O2 electrode exhibited an excellent high-rate capability of 147 mA‚h/g at 400 mA/g at 2.5∼4.8 V. Since the polymer-pyrolysis method is simple and easy to control the structure and morphology of the reaction product, it could be used for industrial production of highly homogeneous nanocrystalline Li-Ni-Mn-O compounds. Acknowledgment. We acknowledge financial support by the 973 Program, China (Grant 2002CB211800) and technical assistance in TEM work by the Center for Electron Microscopy, Wuhan University. References and Notes (1) Ohzuku, T.; Ueda, A.; Nagayam, M. J. Electrochem. Soc. 1993, 140, 1862. (2) Armstrong, A. R.; Bruce, P. G. Nature 1996, 381, 499. (3) Gummow, R. J.; Thackeray, M. M. J. Electrochem. Soc. 1993, 140, 3365. (4) Shao-Horn, Y.; Hackeray, S. A. J. Electrochem. Soc. 1999, 146, 2404. (5) Armstrong, A. R.; Gitzendanner, R.; Robertson, A. D.; Bruce, P. G. Chem. Commun. 1988, 1833. (6) Jang, Y. I.; Huang, B.; Chiang, Y. M.; Sadoway, D. R. Electrochem. Solid-State Lett. 1998, 1, 13. (7) Ammundsen, B.; Desilvestro, J.; Groutso, T.; Hassell, D.; Metson, J. B.; Regan, E.; Steiner, R.; Pickering, P. J. J. Electrochem. Soc. 2000, 147, 4078. (8) Ammundsen, B.; Paulsen, J. AdV. Mater. 2001, 13, 943. (9) Ammundsen, B.; Desilvestro, J.; Steiner, R.; Pickering, P. In Proceedings of the 10th International Meeting on Lithium Batteries, Como, Italy, 28 May-2 June 2000, The Electrochem. Soc. Inc.: Pennington, NJ.

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