The Enhanced Electrochemical Performance of Nanocrystalline Li[Li0.26Ni0.11Mn0.63]O2 Synthesized by the Molten-Salt Method for Li-ion batteries

Electrochimica Acta 117 (2014) 285–291

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The Enhanced Electrochemical Performance of Nanocrystalline Li[Li0.26 Ni0.11 Mn0.63 ]O2 Synthesized by the Molten-Salt Method for Li-ion batteries Wang ZhenYao, Li Biao, Ma Jin, Xia DingGuo ∗ Department of Energy and Resources Engineering, College of Engineering, PekingUniversity, Beijing100871, PR China

a r t i c l e

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Article history: Received 3 July 2013 Received in revised form 18 November 2013 Accepted 20 November 2013 Available online 4 December 2013 Keywords: Lithium rich materials xLi2 MnO3 -(1-x)LiMO2 Molten-salt synthesis Nanocrystalline Li-ion batteries

a b s t r a c t Nanocrystalline Li[Li0.26Ni0.11Mn0.63]O2 were easily prepared by using Ni0.15Mn0.85(OH)2 and Li2 CO3 as precursors and KCl as melt-salt for the high capacity materials of Li-ion storage. The obtained nanoparticles showed same morphology of polygonal shape and the particle size distribution increased with increasing sinter temperature. The Li[Li0.26Ni0.11Mn0.63]O2 electrode sintered at 800 ◦ C for 12 h exhibits a reversible capacity of more than 300 mAh g−1 at 0.1 C rate between 2 V and 4.8 V and the capacity retention remains 86% and 90% after 90 cycles at the rate of 0.5 C and 1 C, respectively. These superior electrochemical performances are discussed in detail and ascribed to the low dimension and well-crystallized particles. The low dimension provides a short diffusion path and fast transport channels for the lithium ion insertion/extraction reactions and the well-crystallized structure restrains the elimination of oxide ion vacancies and metal ions rearrangement during charge–discharge cycling. © 2013 Elsevier Ltd. All rights reserved.

1. Introduction Recently, lithium-rich metal oxides, such as xLi2 MnO3 -(1x)LiMO2 are receiving significant interest as the cathode materials for Li-ion batteries. These materials can provide a larger capacity of over 200 mAh g−1 when they are initially charged to 4.5 V or higher voltage [1–5]. However, some barriers make it difficult to be used as the high-energy cathodes for Li-ion batteries, especially for the electric vehicle application, such as the high first cycle irreversibility, the poor rate capability and the significant decrease in the discharge voltage plateau with the successive cycling, thereby reducing the energy density of the cell [6–13]. To improve the electrochemical performances of the layered xLi2 MnO3 -(1-x)LiMO2 materials, a number of research groups have recently adopted different synthetic methods to optimize the kinetic suitable morphologies and sizes of these materials [14,15]. Usually, the reduction of the particle size can considerably enhance the rate capability of these materials by shortening the lithium diffusion pathways [16–19]. Another method is to fabricate nanostructures, such as nanowires and nanoplates [13,15]. It has been reported that the nanowires of the layered xLi2 MnO3 (1-x)LiMO2 electrodes can deliver a capacity of 200 mAh g−1

∗ Corresponding author. Tel.: +86 10 62767962; fax: +86 10 62768316. E-mail address: [email protected] (X. DingGuo). 0013-4686/$ – see front matter © 2013 Elsevier Ltd. All rights reserved. http://dx.doi.org/10.1016/j.electacta.2013.11.124

at a rate of 5 C [20]. Wei et al. have recently reported that Li[Li0.17 Ni0.25 Mn0.58 ]O2 nanoplates exhibited greater rate capabilities at a 6 C rate comparing with a bulk counterpart prepared via a conventional co-precipitation method [21]. It is well established that battery performances are affected significantly by the electrode materials and the structure of the electrode materials are directly determined by synthesis routes. Consequently, many synthetic methods havebeen developed to prepare the cathode materials, including solid state reactions [22], co-precipitation methods [23], sol-gel methods [24] and others [25]. However, these methods can be still difficult to synthesize cathodes materials consisting of relatively small primary particles with a better reproducible electrochemical performance. The molten-salt synthesis method using the molten salt as solvents or reactants has been regarded as one of the simplest and relatively low cost methods to prepare the cathode materials. This method favors the formation of nanocrystallines and has been recently used to prepare LiMn2 O4 [26], Li(Ni0.5 Mn1.5 )O4 [27] and LiCoO2 [28]. However, there are only a few reports on the synthesis of Lithium-rich materials by using molten-salt methods [29,30]. In this study, a simple melt-salt assisted method is employed to prepare nanocrystalline Li[Li0.26 Ni0.11 Mn0.63 ]O2 powders, using M(OH)2 and Li2 CO3 as starting materials and KCl as the melt salt via a sintering method at 800-900 ◦ C. The obtained powders were investigated by XRD, SEM, TEM, and exhibit a high reversible capacity and good cycling performance.

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2. Experimental 2.1. Sample preparation The transition metal (TM) hydroxides precursor (Ni0.15Mn0.85(OH)2 ) was prepared by a co-precipitation method. In detail, an aqueous solution of 2 mol L−1 TM sulfates (0.3 mol L−1 NiSO4 , and 1.7 mol L−1 MnSO4 ) and an aqueous solution of 2 mol L−1 NaOH (including 0.36 mol L−1 NH4 OH as a chelating agent) were titrated simultaneously into a 1 L continuous stirred tank reactor (CSTR) at 60 ◦ C. The reaction was conducted under an inert atmosphere by bubbling N2 into the CSTR to mitigate the oxidation of Mn2+ and Ni2+ . After all the solutions added, the reaction solution was continuously stirred for 5 h to form TM hydroxide (Ni0.15Mn0.85(OH)2 ) precipitation. Theprecipitation was then centrifuged, washed, filtered, and dried in an oven at 100 ◦ C for 10 h in vacuum. Li[Li0.26 Ni0.11 Mn0.63 ]O2 was synthesized by a molten-salt assisted method. The dried TM precursors (Ni0.15Mn0.85(OH)2 ) were mixed with 5 wt.% excess of Li2 CO3 , and then 6 times weight of KCl was added as the molten salt. This mixture was grounded for 30 min to ensure adequate mixing and then calcined at 500 ◦ C for 4 h and further calcined at 800, 850, 900 ◦ C for 12 h in air. Samples were brought back to room temperature by furnace cooling and then washed thoroughly by deionized water to remove the melt salt KCl. The prepared samples were referred to as “KC800”, “KC850”, and “KC900”, respectively. 2.2. Characterizations The morphologies of the cathode materials were characterized with cold field emission scanning electron microscopy (SEM, Hitachi S-4800) and transmission electron microscopy (TEM, FEI, TECNAI F20). The precursor was first gently crushed with a pestle and then uniformly suspended in ethanol with the help of ultrasonic dispersion to prepare the TEM specimen. The suspension was dropped onto a carbon-coated copper grid for TEM characterization. The structures of samples was analyzed by X-ray diffraction (XRD) using a D8 X diffractometer (Bruker,Germany) employing Cu K␣ radiation operated at 40 kV and 40 mA. The scan data were collectedin a 2␪ range of 10–90◦ , with a step size of 0.02◦ and a counting time of 0.5 s. The chemical compositions of the samples were analyzed by inductively coupled plasma atomic emission spectrometry (ICP-AES).

potentials throughout the paper are in reference to Li/Li+ couple. Electrochemical impedance spectra (EIS) of all the samples were conducted at open-circuit voltage in the frequency range of 100 kHz to 10 mHz with an AC voltage amplitude of 5 mV using a potentiostat/galvanostat SP-240 frequency response analyzer (Bio-logic SAS, France). 3. Results and discussion 3.1. Structure analysis The XRD patterns of the prepared samples at different temperatures for 12 h are shown in Fig. 1. All XRD patterns are similar to those of ␣-NaFeO2 structural type. Most of the peaks can be indexed based on the R-3 m structure except for the peaks between 20◦ and 30◦ , which can only be identified as the monoclinic (C2/m) structure. On this basis, the material can be characterized as a combination of two phases, namely, a trigonal R-3 m phase (Li[Ni0.5 Mn0.5 ]O2 ) and a monoclinic C2/m phase (Li2 MnO3 ). It is can be seen from the patterns, the splitting of doublets (006)/(102) and (108)/(110) is apparent, suggesting a well-defined layered structure formed in the lattice [31,32]. In general, the integrated intensity ratio (R) of the (003) to (104) lines in the XRD patterns can be used to denote the degree of cation mixing in the Li-layers of these materials. If the R value is greater than 1.2, the cation mixing could be considered to be negligible [33–35]. In the XRD pattern of the samples, the relative intensity ratios of the (003) to (104) lines are 1.53, 1.31, and 2.03, respectively, which are all above 1.2, suggesting that the disordered arrangement of Ni and Li ions in the Li-layers could be ignored for all samples. Fig. 1b shows the carefully examining XRD patterns between 20◦ and 30◦ . The peaks at 2␪= 20.8◦ , 21.7◦ , 24.2◦ are caused by the superlattice ordering of Li+ , Ni2+ and Mn4+ in transition metal layer [4]. It can be seen that the intensity of superstructure reflections was slightly enhanced with increasing sintering temperature, demonstrating an increased content of layered-type structure. The molar ratio Li:Ni:Mn determined by ICP-AES was 1.245: 0.117: 0.624, which is in excellent agreement with the target ratio. The residual K is less than 0.1wt% of the total mass and can be negligible. The XRD and ICP results demonstrate that a Li[Li0.26 Ni0.11 Mn0.63 ]O2 cathode material with a good layered structure was successfully synthesized by the molten-salt synthesis method. 3.2. Morphology and particle-size of the particles

2.3. Electrochemical measurements Electrochemical measurements were carried out using CR2032 coin-type cells. The synthesized cathode materials were mixed with acetylene black and polyvinylidene difluoride binder with the weight ratio of 80:10:10 and coated onto an aluminum foil. Lithium cells were assembled with Li[Li0.26 Ni0.11 Mn0.63 ]O2 cathode and lithium metal as the counter electrode inside an argon-filled glove box (H2 O level of <1 ppm). The electrolyte was 1 mol L−1 LiPF6 dissolved in a mixture of ethylene carbonate (EC) and ethyl methyl carbonate (EMC) (1:1 v/v). The cells were galvanostatically charged and discharged on battery testers (NEWARE CT-3008 W, Neware Technology Co., Ltd., P.R. China) between 2.0 and 4.8 V (vs. Li/Li+ ) at room temperature. Capacities were calculated by only considering the active mass of the electrodes. The fabricated cells were tested at various rates (0.1, 0.2, 0.5, 1.0, 2.0 and 5.0 C) to investigate the rate capability of the prepared materials. 1 C rate is equivalent to a current density of 200 mA g−1 in our definition. The cells were cycled at current density of 0.5 C (100 mA g−1 ) and 1 C (200 mA g−1 ) to compare cycling performances of the materials prepared by different methods. The

Morphologies and particle sizes of the samples were examined by SEM. Fig. 2 shows the SEM photographs of Li[Li0.26 Ni0.11 Mn0.63 ]O2 powders heated at different temperatures for 12 h. All samples show a nearly same morphology of polygonal shape, but the particles size distribution increase with increasing sinter temperature. Among all samples, KC800 has the smallest particle size of around 50 -100 nm, and lots of particles show certain agglomeration. Sample KC850 shows a larger particle-size distribution in the 75–150 nm range and some particles grow to about 300 nm. When the heating temperature further increased to 900 ◦ C, the particle size of KC900 average at 100-200 nm and some particles even grow to larger than 500 nm. Moreover, the particles for KC850 and KC900 disperse more individually and exhibit more clearly polygonal shape. Therefore, it can be concluded that molten-salt sinter methods lead to polygonal-shape particles and increasing the sintering temperature could enlarge particle size. The TEM characterization of sample KC800 is displayed in Fig. 3. Fig. 3a shows a good crystallinity of the bulk Li[Li0.26 Ni0.11 Mn0.63 ]O2 material, which is in good agreement with SEM results. The layer distance of

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Fig. 1. XRD patterns of the as-prepared Li[Li0.26 Ni0.11 Mn0.63 ]O2 samples.

two adjacent lattice fringes in the enlarged picture (Fig. 3b) is ∼0.47 nm, corresponding to the interplanar distance of the (003) planes of the R-3 m structure, which agrees very well with the results obtained from XRD analysis. 3.3. Electrochemical studies 3.3.1. Charge/discharge cycling Fig. 4 (a, b, and c) shows the charge and discharge cures of sample KC800, KC850, KC900, respectively, at the rate of 0.1 C (20 mA g−1 ) between 2 and 4.8 V. During the first charge process, a slope and plateau has been assigned to the oxidation of Ni2+ to Ni4+ and the removal of the superfluous lithium from the transition metal layer along with the simultaneous oxygen evolution, respectively.

During the oxygen loss plateau, the Li2 MnO3 component is activated to form MnO2 component and thus the material could deliver high discharge capacity during the subsequent discharge process. Notably, the voltage of plateau for all samples is located at about 4.7 V, which is higher than the usually platform located between 4.5 V and 4.6 V. It may be attributed to the high crystalline and the polygonal shape of the particles, which makes the activation energy of “Li2 O” removing from the structure higher. The first charge and discharge capacity of KC800 was measured as 300 mAh g−1 and 255 mAh g−1 , respectively, with a coulombic efficiency of 85%. The first charge/discharge capacities of KC850 and KC900 were 262.5/206.8 mAh g−1 and 80/64 mAh g−1 , with coulombic efficiency of 79% and 80%, respectively. In the initial several cycles, the charge and discharge capacities

Fig. 2. SEM of the as-prepared samples Li[Li0.26 Ni0.11 Mn0.63 ]O2 prepared at different temperate (a) KC800, (b) The magnification of KC800, (c) KC850, and (d) KC900.

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Fig. 3. TEM images of the sample KC800.

increased highly with the cycles caused by the activation process of the electrodes. After five cycles, the discharge capacity of sample KC800, KC850 and KC900 increased to 317, 258 and 238 mAh g−1 , respectively. For such abnormal high discharge capacity of these cathode materials, there are several different interpretations. Tsutomu Ohzukuet. al. [36] found that Li[Ni0.20 Li0.20 Mn0.60 ]O2 electrodes showed a high capacity of more than 300 mAh g−1 at high temperate (55 ◦ C) or low current (4.8 mA g−1 ) between 2 and 5 V. They assumed that one possibility is high-valent manganese ions (Mn5+ , Mn6+ ) may participate in the reaction. Another possibility is an anion redox reaction (2O2- /O2 2− ) in the solid matrix. Armstrong et al. [2] had demonstrated directly that O2 evolution associated with the 4.5 V plateau on charging by in situ differential electrochemical mass spectrometry (DEMS). They reported that the oxide ion vacancies formed during the first charge processwere eliminated with cationic and anionic rearrangements at

the end of the first charge. This also eliminated a corresponding number of lithium sites, i.e. as every two oxide ion vacancies eliminated, one cation site in the lithium plane and one cation site in the transition metal plane were also eliminated. However, Kim [15] assumed that Li[Ni2+ 0.25 Li0.15 Mn4+ 0.6 ]O2 was fully charged to [][Ni4+ 0.25 Li0.15 Mn4+ 0.6 ]O1.78 in the first charge and these oxygen deficient materials hereafter reacted reversibly with lithium, showing a discharge capacity of about 300 mAh g−1 . A. Manthiram et al. [37,38] also suggested that part of the oxide ion vacancies may be retained in the layered lattice after the first charge as the observed irreversible capacities were less than that the expected on the basis of the complete elimination of oxide ion vacancies. Moreover, they indicated that higher vacancies of oxide ion were retained in the layered lattice by the surface modification, resulting in a decrease in irreversible capacity and an increase of the discharge capacity.

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in the layered lattice after the first charge [37,38]. Accordingly, we prefer to suggest that parts of the oxide ion vacancies and lithium ion vacancies may be retained in the layered lattice after the first charge, which contributes to the high reversible capacity, though it is difficult to confirm the existence of oxygen ion vacancies at present. And the gradual deintercalation of lithium ions result in the rearrangements of the structure step by step with cycles, which is in favor of the remains of ion vacancies and stabiles of layer structures. The similar results can be found in the surface modified materials [37,38] and step-charging process [39]. KC800 showed a discharge capacity of 255 and 320 mAh g−1 at the first and 3rd cycle, which correspond to intercalation of about 0.78 and 1 mol of Li+ , respectively. Based on the above discussions, the following reactions can be proposed for KC800: reaction (1) as the slope region, reaction (2) inthe plateau region of first charge, (3) as rearrangement, reaction (4) forthe firstdischarge, reaction (5) and (6) for the third charge and discharge:

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Li[Li0.26 Ni0.11 Mn0.63 ]O2 → 0.22Li+ + Li0.78 [Li0.26 Ni4+ 0.11 Mn4+ 0.63 ]O2 + 0.22e− (1)72mAhg−1

In this study, the reversible capacity of Li[Li0.26 Ni0.11 Mn0.63 ]O2 calculated based on metal valence change (Ni2+/4+ , Mn3+/4+ ) is 243 mAh g−1 . From the charge profile and capacity, it can be seen that not all lithium ions deintercalate from the lattices in the first cycle, but deintercalate gradually during the initial several cycles. Especially for sample KC900, the platform is very short in the first charge profile, which means only small parts of lithium ions deintercalate from the lattices. Then the charge capacity increase obviously with cycles, responding to more and more lithium ions deintercalate from the lattices, and reach to 246 mAh g−1 at 5th cycle. For sample KC800, the observed discharge capacity (255 and 320 mAh g−1 for the first and 3rd discharge capacity, respectively) and first coulombic efficiency (85%) are both much higher than the theoretically calculated data based on metal valence change (243 mAh g−1 and 59%). This result will imply the availability of more number of lithium sites for lithium insertion/extraction after the first charge, for example, if parts of the oxide ion vacancies are retained

Li0.78 [Li0.26 Ni4+ 0.11 Mn4+ 0.63 ]O2 → 0.7Li+ + Li0.08 [Li0.26 Ni4+ 0.11 Mn4+ 0.63 ]O1.65 0.35 + 0.7e+ 0.175O2 (2)229mAhg−1

(2)

Li0.08 [Li0.26 Ni4+ 0.11 Mn4+ 0.63 ]O1.65 0.35 → Li0.20 [Li0.13 0.13 Ni4+ 0.11 Mn4+ 0.63 ]O1.65 0.35 → Li0.20 [Li0.13 0.11 Ni4+ 0.11 Mn4+ 0.63 ]O1.65 0.31

→ Li0.98 [Li0.13 0.11 Ni2+ 0.11 Mn3.1+ 0.63 ]O1.65 0.31 255mAhg−1 (4)

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Fig. 6. (a) cyclic performance at different rates, and electrochemical charge/discharge profiles of (b) KC800, (c) KC850, and (d) KC900.

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. . .. . .. . . Li0.98 [Li0.13 0.11 Ni2+ 0.11 Mn3.1+ 0.63 ]O1.65 0.31 → 0.98Li+ + 0.98e- + 0.05O2 + [Li0.13 0.11 Ni4+ 0.11 Mn4+ 0.63 ]O1.55 0.41 320mAhg−1

(5)

[Li0.13 0.11 Ni4+ 0.11 Mn4+ 0.63 ]O1.55 0.41 +0.98Li+ + 0.98e→ Li0.98 [Li0.13 0.11 Ni2+ 0.11 Mn2.8+ 0.63 ]O1.55 0.41 320mAhg−1 (6)

3.3.2. Rate and cycling performances Fig. 5 shows the rate performance of the samples. It can be seen that after activation by several initial cycles the electrodes reach to a stable discharge capacity of 317, 258 and 238 mAh g−1 for KC800, KC850 and KC900 at 0.1 C, respectively. The rate performance of KC800 is much higher than that of the other two samples, and shows discharge capacities of 229 and 128 mAh g−1 at 1 and 5 C rate, respectively. However, KC850 and KC900 show much lower rate performances. The discharge capacity of KC850 is only 163 and 40.6 mAh g−1 at 1 and 5 C rate. The rate performance of KC900 is even slight lower than KC850, being 148 and 31.7 mAh g−1 , respectively. It can be seen that the smaller particle size show higher capacity and better rate performances, because low-dimension of particle size can provide short diffusion paths and increased active surface area for Li+ transportation in the cathode materials. Cycling performances of the samples at different current densities after initial activation are shown in Fig. 6a. At both current densities, KC800 shows the most pronounced cycle stability and its capacity retention remains 86% and 90% after 90 cycles at 0.5 and 1 C rate, respectively. The capacity retention of KC 850 and KC900 are 76% and 83% after 90 cycles at 0.5 C rate, which are slightly lower than that of KC800. The corresponding electrochemical charge/discharge profiles for cycles 1, 5, 20, 50 and 90 are shown in Fig. 6b, c and d. Further, it can be seen that the capacity decrease is accompanied by a slight voltage drop during the discharge process,which can be ascribed to the spinel transformation, lattice break down, and possible deteriorated electrode/electrolyte interface that hampers the reversible lithium ion intercalation/deintercalation [40]. The cycling performances attribute to well crystalline as well as small particle size. The wellcrystallized structure may restrain the elimination of oxide ion vacancies and metal ions rearrangement during charge–discharge cycling. Moreover, small particle size may hamper phase transformation and lattice break. However, higher sinter temperature will result in higher crystallinity and bigger particle size, which may hamper Li ion migration and cause big stress during the electrode cycling, leading torelatively poor cycling performances of the cathode materials. 3.3.3. Electrochemical impedance spectra EIS is a powerful technique for studying the electrode kinetics of lithium ions because Li intercalation and de-intercalation from electrodes can be normally modeled as a multi-step process that involves different time scales. EIS spectra of the samples are employed to explain the origin of the improvement in rate capability of the samples. Fig. 7 shows the Nyquist plots for the electrodes of KC800, KC850 and KC900. EIS of all the samples are tested prior to electrochemical oxidation and show standard Nyquist plots comprising of a small interrupt and a small semicircle in the high frequency region, and a slope in the low frequency region. The small interrupt in the high frequency corresponds to the solution resistance (Re). The small semicircle in the high frequency is assigned to the resistance (Rf) of lithium-ion diffusion in the surface layer

Fig. 7. The Nyquist plots of KC800, KC850 and KC900 samples.

(including the SEI film and the surface modification layer). The slope in the low frequency region refers to lithium ion diffusion in the bulk materials [41,42]. From Fig. 7, it can be seen that all three samples show negligible ohmic resistance while exhibiting different diffusion resistance. The lithium-ion diffusion resistance (Rf) of samples increases in the following order: KC800 < KC850
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[12]

[13]

[14]

[15]

[16] [17]

[18] [19] [20]

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[22]

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