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Evaluation of Stability of Charged Lithium-rich Layer-structured Cathode Material at Elevated Temperature
Electrochimica Acta 169 (2015) 310–316
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Electrochimica Acta journal homepage: www.elsevier.com/locate/electacta
Evaluation of Stability of Charged Lithium-rich Layer-structured Cathode Material at Elevated Temperature Hiroaki Konishi * , Tatsumi Hirano, Daiko Takamatsu, Akira Gunji, Xiaoliang Feng, Sho Furutsuki Hitachi Research Laboratory, Hitachi Ltd. 7-1-1 Omika-cho, Hitachi, Ibaraki 319-1292, Japan
A R T I C L E I N F O
A B S T R A C T
Article history: Received 3 December 2014 Received in revised form 7 March 2015 Accepted 31 March 2015 Available online 11 April 2015
The stability of charged Li1.2Ni0.13Mn0.54Co0.13O2 at elevated temperature was investigated by using thermal desorption spectrometry-mass spectrometry (TDS-MS), X-ray diffraction (XRD), and X-ray absorption fine structure (XAFS) measurements. The TDS-MS and XRD spectra indicated that the crystal structure of the delithiated Li1.2Ni0.13Mn0.54Co0.13O2 changed from layer to spinel, and oxygen was released at high temperature. However, the amount of oxygen released from Li1.2Ni0.13Mn0.54Co0.13O2 in a full-charge state was much less than that released from LiNi0.8Mn0.1Co0.1O2 in a full-charge state. The XAFS spectra indicated that the high oxidation state of manganese was more stable than that of nickel and cobalt. The high stability of charged Li1.2Ni0.13Mn0.54Co0.13O2 at high temperature was, therefore, attributed to the high manganese content of the transition metal. ã 2015 Elsevier Ltd. All rights reserved.
Keywords: Lithium-ion battery Lithium-rich layer-structured cathode Thermal decomposition Crystal structure Oxidation state
1. Introduction Lithium-ion batteries have recently been applied to not only small devices such as cellular phones but also automotive products such as hybrid electric vehicles (HEVs), plug-in hybrid electric vehicles (PHEVs), and electric vehicles (EVs). These applications require both high energy density and safety. These properties are closely related to those of materials used for the electrodes in batteries. Ni-based layer-structured cathode materials, LiNi1-aMaO2 (M: metal), are promising in that regard due to their high capacity [1–4]. However, the stability of Li1-yNi1-aMaO2 at elevated temperature is poor, especially in a highly delithiated state. The crystal structure of Li1-yNi1-aMaO2 changes from layer to spinel, and then rocksalt, and oxygen is released at high temperature [5–15]. The oxygen-release from the cathode can lead to thermal runway by reacting with electrolyte, and cause fire and explosion of lithium-ion battery. Therefore, the evaluation of thermal decomposition of the cathode is important. We elucidated the stability of an Ni-based cathode, LiNi0.8Mn0.1Co0.1O2, in a previous report, and its low stability due to heating was attributed to the instability of Ni3+ and Ni4+ in a high charge state [12,15]. Its poor stability has prevented it from being applied
* Corresponding author. Tel.: +81 294 52 5111x6089; fax: +81 294 52 7636. E-mail address: [email protected] (H. Konishi). http://dx.doi.org/10.1016/j.electacta.2015.03.217 0013-4686/ ã 2015 Elsevier Ltd. All rights reserved.
to large-scale batteries; therefore, new cathode materials that satisfy both high capacity and high stability criteria are required. It has been reported that lithium-rich layer-structured materials, Li[Li1-a-b-cNiaMnbMc]O2 (M: metal), exhibit high capacity around 250 Ah kg1 [16–20]. The crystal structure of Li1.2-xNi0.15Mn0.55Co0.1O2 changed from layer to spinel at high temperature [21]. Therefore, the oxygen-release from charged Li-rich layer-structured cathode at high temperature might be lower than that from charged Ni-based layer-structured cathode. However, there are no results compared by the identical condition. Furthermore, the cause of the difference in structure change has yet to be clarified. The stability of Li1.2Ni0.13Mn0.54Co0.13O2 in a charge state at high temperature was evaluated in the present study, and compared with that of LiNi0.8Mn0.1Co0.1O2 in a full-charge state. The oxygen release, crystal structure change, and oxidation state of each transition metal were investigated by using thermal desorption spectrometry-mass spectrometry (TDS-MS), X-ray diffraction (XRD), and X-ray absorption fine structure (XAFS) analyses. 2. Experimental Li1.2Ni0.13Mn0.54Co0.13O2 sample was synthesized by reacting stoichiometric mixtures of lithium acetate, nickel acetate, manganese acetate, and cobalt acetate. The mixture was stirred in distilled water for 1 h, and then dried. It was pre-heated at 500 C for 12 h in air, and then heated at 850 C for 12 h in air.
H. Konishi et al. / Electrochimica Acta 169 (2015) 310–316
3. Results and discussion 3.1. Crystal structure of Li1.2Ni0.13Mn0.54Co0.13O2 and LiNi0.8Mn0.1Co0.1O2 Fig. 1 shows the XRD patterns of Li1.2Ni0.13Mn0.54Co0.13O2 and LiNi0.8Mn0.1Co0.1O2 powder. Fig. 1(a) indicates that all the main
(018) / (110) (113)
(104)
(107)
(b)
(015)
(101) (006) / (102)
Intensity (a.u.)
(003)
Superlattice
(a) 10
20
30
40 2 /°
50
60
70
Fig. 1. XRD patterns of Li1.2Ni0.13Mn0.54Co0.13O2 and LiNi0.8Mn0.1Co0.1O2 powder.
peaks of Li1.2Ni0.13Mn0.54Co0.13O2 were assigned to a hexagonal a-NaFeO2 structure with space group R-3m. Small peaks also appeared between 2u = 20 and 25 , indicating the presence of a superlattice structure such as LiMn6 and LiMn5Ni in the transition metal layer [17,20]. Fig. 1(b) indicates that all the peaks of LiNi0.8Mn0.1Co0.1O2 were indexed to a hexagonal a-NaFeO2 structure. 3.2. Electrochemical properties of Li1.2Ni0.13Mn0.54Co0.13O2 and LiNi0.8Mn0.1Co0.1O2 The charge-discharge capacities of Li1.2Ni0.13Mn0.54Co0.13O2 and LiNi0.8Mn0.1Co0.1O2 were measured within a potential range of 2.5–4.6 V for the former and 2.5–4.3 V for the latter (vs. Li/Li+). Fig. 2 shows the initial charge and discharge curves of Li1.2Ni0.13Mn0.54Co0.13O2 and LiNi0.8Mn0.1Co0.1O2. The initial charge and discharge capacities of Li1.2Ni0.13Mn0.54Co0.13O2 were 327 and 251 Ah kg1, respectively. Even though 87% of lithium ions were de-intercalated in the charge process, only 67% of them were
+
/ V (vs. Li/Li )
4.5
Potential
LiNi0.8Mn0.1Co0.1O2 was synthesized by reacting lithium hydroxide, nickel oxide, manganese oxide, and cobalt oxide with 3% excess of lithium hydroxide. The mixture was pre-heated at 600 C for 10 h in oxygen, and then heated at 850 C for 10 h in oxygen. The crystal structures of the Li1.2Ni0.13Mn0.54Co0.13O2 and LiNi0.8Mn0.1Co0.1O2 samples were evaluated by XRD (Rigaku, Rint-2200). The diffraction data were recorded at a 0.02 step width over a 2u range from 10 to 70 using a Rigaku diffractometer with Cu Ka radiation (l = 1.54 Å). The electrochemical properties of Li1.2Ni0.13Mn0.54Co0.13O2 and LiNi0.8Mn0.1Co0.1O2 samples were measured by using a threeelectrode electrochemical cell containing Li1.2Ni0.13Mn0.54Co0.13O2 or LiNi0.8Mn0.1Co0.1O2 electrodes as a cathode, lithium foil as an anode, and a reference. The electrode was composed of the prepared sample, carbon, and binder (85: 10: 5 wt%). The electrolyte was a 1 M LiPF6 solution containing a mixture of ethylene carbonate (EC), ethyl methyl carbonate (EMC), and dimethyl carbonate (DMC) (1: 2: 2 vol%) solvents. The charge– discharge capacities of the samples were measured within a potential range of 2.5–4.6 V or 2.5–4.3 V (vs. Li/Li+). The cells were charged at 0.05 C (1 C = 260 A kg1) to 4.6 or 4.3 V, and then held at 4.6 or 4.3 V until the current became less than 0.005 C. After that, the cells were discharged at 0.05 C to 2.5 V. The lithium content after charging was measured using inductively coupled plasma atomic emission spectroscopy (ICP-AES, PerkinElmer, OPTIMA-3300XL). The oxygen released from the cathode by heating was evaluated by using TDS-MS (ESCO Ltd. EMD-WA1000). The samples for TDS-MS were prepared with the following procedure. The Li1.2Ni0.13Mn0.54Co0.13O2 and LiNi0.8Mn0.1Co0.1O2 electrodes were charged to the prescribed composition, Li1.2(1-x)Ni0.13Mn0.54Co0.13O2 (x = 0.6, 0.7, 0.8, 0.87) and Li1-yNi0.8Mn0.1Co0.1O2 (y = 0.8), at a rate of 0.05 C. After that, the electrochemical cells were decomposed, and the electrodes were washed with DMC. The delithiated samples were placed in the TDS-MS chamber and heated to 400 C in a vacuum at a rate of 5 C/min. The crystal structure of the Li1.2Ni0.13Mn0.54Co0.13O2 and LiNi0.8Mn0.1Co0.1O2 samples in a full-charge state at high temperature was evaluated by XRD. The samples for XRD were prepared with the following procedure. The Li1.2Ni0.13Mn0.54Co0.13O2 and LiNi0.8Mn0.1Co0.1O2 electrodes were charged up to a full-charge state, such as Li1.2(1-x)Ni0.13Mn0.54Co0.13O2 (x = 0.87) and Li1yNi0.8Mn0.1Co0.1O2 (y = 0.8). After that, the electrochemical cells were decomposed, and the electrodes were washed with DMC. The charged electrodes were heated at 150–400 C in an argon atmosphere for 1 h to expose them, and then cooled to room temperature (RT). The oxidation state of each transition metal was investigated by XAFS at the BL-9C of the Photon Factory (PF) for High Energy Accelerator Research Organization (KEK) in Japan. The samples for XAFS were prepared with the same procedure as that used to prepare the XRD samples. The oxidation state of each transition metal was mainly evaluated by X-ray absorption near-edge structure (XANES) analysis. The Ni, Co, and Mn K-edge XANES spectra of the prepared samples were measured in transmission mode by using an Si(111) double-crystal monochromator. The XANES spectra were compared with those of the reference samples such as lithium transition metal oxide and transition metal oxide.
311
4.0
3.5
3.0 Li1.2Ni0.13Mn0.54Co0.13O2 LiNi0.8Mn0.1Co0.1O2
2.5 0
100 Capacity
200 / Ah kg
300 -1
Fig. 2. Initial charge and discharge curves of Li1.2Ni0.13Mn0.54Co0.13O2 and LiNi0.8Mn0.1Co0.1O2 within potential range of 2.5–4.6 V, and 2.5–4.3 V vs. Li/Li+.
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Table 1 Lithium content calculated from charge capacity and that measured using ICP analysis. Cathodes
Calculated values
x = 0.87 Li1.2(1-x)Ni0.13Mn0.54Co0.13O2 Li1-yNi0.8Mn0.1Co0.1O2 y = 0.8 Li1.2(1-x)Ni0.13Mn0.54Co0.13O2: Ni + Mn + Co = 0.8 Li1-yNi0.8Mn0.1Co0.1O2: Ni + Mn + Co = 1.0
Experiment values x = 0.87 y = 0.78
intercalated in the discharge process. The initial charge and discharge capacities of LiNi0.8Mn0.1Co0.1O2 were 222 and 198 Ah kg1, respectively. This indicates that 80% of lithium ions were extracted in the charge process, and 72% of them were inserted in the discharge process. The irreversible capacity of Li1.2Ni0.13Mn0.54Co0.13O2 was larger than that of LiNi0.8Mn0.1Co0.1O2. The large irreversible capacity of the former was attributed to the irreversible loss of oxygen from the lattice [17–20].
(a)
=0.87 =0.8 =0.7 =0.6
200 Temperature
300 / °C
100
3.4. Change in oxygen release from Li1.2Ni0.13Mn0.54Co0.13O2 in a fullcharge state during charge-discharge cycle
400
(b)
100 150
200 Temperature
Table 1 lists the lithium content calculated from the charge capacity, and that measured using inductively coupled plasma (ICP) analysis. Table 1 indicates that there were no large differences in the calculated values from charge capacity and experimental values from ICP analysis. Fig. 3 shows the TDS-MS spectra of Li1.2(1-x)Ni0.13Mn0.54Co0.13O2 (x = 0.6, 0.7, 0.8, and 0.87) and Li1-yNi0.8Mn0.1Co0.1O2 (y = 0.8). Fig. 3(a) indicates that as delithiated content x increased, the amount of oxygen release increased and the oxygen-release onset temperature decreased. Fig. 3(b) indicates that the amount of oxygen release from Li1.2(1-x)Ni0.13Mn0.54Co0.13O2 (x = 0.8 and 0.87) was much less than that from Li1-yNi0.8Mn0.1Co0.1O2 (y = 0.8). In contrast, the oxygen-release onset temperature of the former was lower than that of the latter.
1st 2nd 3rd Intensity (a.u.)
Intensity (a.u.)
Li1.2(1- )Ni0.13Mn0.54Co0.13O2 ( =0.87) Li1.2(1- )Ni0.13Mn0.54Co0.13O2 ( =0.8) Li1- Ni0.8Mn0.1Co0.1O2 ( =0.8)
50
3.3. Comparison of stability of Li1.2(1-x)Ni0.13Mn0.54Co0.13O2 (x = 0.6, 0.7, 0.8, and 0.87) and Li1-yNi0.8Mn0.1Co0.1O2 (y = 0.8) at high temperature
Fig. 4 shows the TDS-MS spectra of Li1.2Ni0.13Mn0.54Co0.13O2 in a full-charge state during initial, second, and third charge process. Fig. 4 indicates that two oxygen release peaks were observed in the initial cycle. As cycle number increased, the oxygen release at low temperature was reduced, and only one peak was observed in the third cycle. The oxygen release at low temperature might be related to the charge reaction process of this material. It has been reported that oxygen was extracted in lithium-rich layer-structured material, and transition metal migrated from the surface to the bulk in the last state of charge reaction even at RT [22,23]. Since the migration of transition metal might be slow at RT, this reaction might not fully progress at RT. The migration of transition metal might have been prompted, and oxygen was released due to heating; therefore,
Intensity (a.u.)
100
Li1.2(1-x)Ni0.13Mn0.54Co0.13O2 (x = 0.87) and Li1-yNi0.8Mn0.1Co0.1O2 (y = 0.8) were defined as the full-charge state of each material.
300 / °C
400
Fig. 3. TDS-MS spectra of (a) Li1.2(1-x)Ni0.13Mn0.54Co0.13O2 (x = 0.6, 0.7, 0.8, 0.87), and (b) Li1.2(1-x)Ni0.13Mn0.54Co0.13O2 (x = 0.8, 0.87) and Li1-yNi0.8Mn0.1Co0.1O2 (y = 0.8).
100
200 Temperature
300 / °C
400
Fig. 4. TDS-MS spectra of Li1.2Ni0.13Mn0.54Co0.13O2 in a full-charge state during initial, second, and third charge process.
H. Konishi et al. / Electrochimica Acta 169 (2015) 310–316
(220) spinel
(a)
400°C
300°C
Intensity (a.u.)
3.5. Change in structure, oxidation state, and local structure of Li1.2Ni0.13Mn0.54Co0.13O2 and LiNi0.8Mn0.1Co0.1O2 in initial full-chargestate at high temperature The change in structure, oxidation state, and local structure of Li1.2Ni0.13Mn0.54Co0.13O2 in a full-charge state at high temperature was compared with that of LiNi0.8Mn0.1Co0.1O2 in a full-charge state reported in our previous paper [15]. Fig. 5 shows the XRD patterns of Li1-yNi0.8Mn0.1Co0.1O2 (y = 0.8) and Li1.2(1-x)Ni0.13Mn0.54Co0.13O2 (x = 0.87) at 25 C and heated at 150–400 C.
350°C
Fig. 5(a) indicates that the layer structure was maintained from 25 to 150 C. A new peak at 2u = 31 emerged at 200 C, indicating that some of the transition metal ions occupied tetrahedral sites [11,24,25]. Furthermore, the (018) and (110) peaks at 2u = 65 merged at 200 C. These features are characteristics of a spinel structure [5–15]. However, the peak at 2u = 31 vanished at 300 C. Furthermore, the peak at 2u = 18 decreased, and the peaks at 2u = 38, 44, and 64 shifted to smaller angles. This indicates that the layer structure was changed to spinel, and then rocksalt, accompanying oxygen release, according to the following reactions [26,27].
250°C
200°C
150°C
25°C
Li1-yMO2 ! (2-y)/3[Li3(1-y)/(2-y)M(2 O2
20
30
40 50 2 /°
60
350°C
Intensity (a.u.)
300°C
250°C
150°C
25°C
40 50 2 /°
60
! [Li1-yM]O2-y + [(2-y)/6]O2 + [(2y-1)/3]O2
(2)
The relative intensity of each peak changed from 250 to 300 C. A new peak at 2u = 31 emerged at 300 C, and this was clearly observed even at 400 C. This indicates that some of the transition metal ions occupied tetrahedral sites from 300 to 400 C [11,24,25]. Furthermore, the (018) and (110) peaks around 2u = 65 merged at 300 C. However, the peaks did not change much from 300 to 400 C.
200°C
30
(1)
Fig. 5(b) indicates that the relative intensity ratio of the (003) peak at 2u = 18 and the (018) peak at 2u = 64 was low, compared with that of the pristine state, as shown in Fig. 1. This indicates that there were two phases, which have different lattice parameter c, at the end of the charge process [23,28]. Even though the layer structure was maintained from 25 to 250 C, the relative intensity of the (003) and (018) peaks increased. This might be attributed to the two phases being changed into one phase due to heating. This change might be accompanied by the migration of transition metal and oxygen release, therefore, oxygen release was observed in this temperature range, as shown in Fig. 3(b).
400°C
20
y-1)/(2-y)][M]2O4 + [(2y-1)/3]
70
(220) spinel
(b)
313
70
Fig. 5. XRD patterns of (a) Li1-yNi0.8Mn0.1Co0.1O2 (y = 0.8) and (b) Li1.2(1 x)Ni0.13Mn0.54Co0.13O2 (x = 0.87) at 25 C and heated at 150–400 C.
oxygen release was observed at low temperature in the highly delithiated Li1.2Ni0.13Mn0.54Co0.13O2. As the cycle number increased, the migration of transition metal progressed; therefore, oxygen release at low temperature might be reduced. In contrast, oxygen release at high temperature was clearly observed after a few cycles, then it might be related to the crystal structure change.
A reaction similar to formula (1) was also observed for delithiated Li1.2Ni0.13Mn0.54Co0.13O2. However, no structural change like that in formula (2) was observed. We concluded that the suppressed reaction in formula (2) could be attributed to the fact that the amount of oxygen release from Li1.2Ni0.13Mn0.54Co0.13O2 in a full-charge state was lower than that from LiNi0.8Mn0.1Co0.1O2 in a full-charge state. Fig. 6 plots Ni, Co, and Mn K-edge XANES spectra in Li1-yNi0.8Mn0.1Co0.1O2 (y = 0.8) at 25 C and heated at 200, 300, 400 C, together with those of the reference samples such as lithium transition metal oxide and transition metal oxide. All spectra exhibit small pre-edge and large peaks. The former is related to dipole forbidden 1s ! 3d electronic transition. However, this transition is partly allowed due to the hybridization of the 3d-4p orbital caused by structural distortion around the central transition metal. The latter is associated with dipole allowed 1s ! 4p electronic transition. The energy of the spectra is related to the oxidation state of transition metal [29–36], and the pre-edge
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2.0
2.0 (a)
(a) 1.5
1.0 8328 8332 8336 LiNi0.8Co0.15Al0.15O2 NiO 400°C 300°C 200°C 25°C
0.5
0.0 8330
8335
8340 8345 Energy / eV
8350
Normalized (a.u.)
Normalized (a.u.)
1.5
1.0 8328 8332 8336 LiNi0.8Co0.15Al0.15O2 NiO 400°C 300°C 200°C 25°C
0.5
0.0 8330
8355
8335
8340 8345 Energy / eV
8350
2.0
2.0
(b)
(b)
1.5
1.0 7704 7708 7712 LiCoO2 Co3O4 CoO 400°C 300°C 200°C 25°C
0.5
0.0 7705
7710
7715 7720 Energy / eV
7725
Normalized (a.u.)
Normalized (a.u.)
1.5
1.0 7704 7708 7712 LiCoO2 Co3O4 CoO 400°C 300°C 200°C 25°C
0.5
0.0 7705
7730
7710
7715 7720 Energy / eV
7725
7730
2.0
2.0
(c)
(c)
1.0 6536 6540 6544 MnO2 Mn2O3 400°C 300°C 200°C 25°C
0.5
0.0 6540
6545
6550 6555 Energy / eV
6560
6565
Fig. 6. (a) Ni, (b) Co, and (c) Mn K-edge XANES spectra of Li1-yNi0.8Mn0.1Co0.1O2 (y = 0.8) at 25 C and heated at 200, 300, 400 C, together with those of reference samples such as lithium transition metal oxide and transition metal oxide.
Normalized (a.u.)
1.5
1.5 Normalized (a.u.)
8355
1.0 6536 6540 6544 MnO2 Mn2O3 400°C 300°C 200°C 25°C
0.5
0.0 6540
6545
6550 6555 Energy / eV
6560
6565
Fig. 7. (a) Ni, (b) Co, and (c) Mn K-edge XANES spectra of Li1.2(1-x)Ni0.13Mn0.54Co0.13O2 (x = 0.87) at 25 C and heated at 200, 300, 400 C, together with those of reference samples such as lithium transition metal oxide and transition metal oxide.
H. Konishi et al. / Electrochimica Acta 169 (2015) 310–316
intensity is related to the occupation sites of the transition metal ions. This is stronger when transition metal ions occupy tetrahedral sites than when they occupy octahedral sites [13,14,29–32]. As we can see from the XRD results in Fig. 5(a) and (b), some of the transition metal ions migrate from octahedral to tetrahedral sites due to heating. Therefore, the change in pre-edge intensity was used to elucidate what kinds of transition metal ions migrated. Fig. 6(a) indicates that Ni K-edge spectrum at 25 C was on an energy side higher than that of LiNi0.8Co0.15Al0.05O2 (Ni3+). The spectra shifted to a lower energy side from 25 to 300 C, and they approached that of NiO (Ni2+), indicating that nickel was reduced due to heating [13,14,33,34]. In contrast, the pre-edge intensity decreased due to heating. This suggests that nickel ions did not demonstrate the symptoms of occupation of tetrahedral sites at high temperature. Fig. 6(b) indicates that the energy of Co K-edge spectrum at 25 C was close to that of LiCoO2 (Co3+). The spectra shifted to a lower energy side with increasing temperature. In contrast, the pre-edge intensity increased from 25 to 200 C, and it decreased from 200 to 400 C. This indicates that cobalt ions moved from octahedral to tetrahedral sites at 25–200 C, and then octahedral sites at 200–400 C. This coincides with the XRD results in Fig. 5(a). Fig. 6(c) indicates that Mn K-edge spectra shifted to the lower energy side with increasing temperature. However, the change was less than that of Ni and Co. Furthermore, the pre-edge intensity indicates that manganese ions did not exhibit the symptoms of occupation of tetrahedral sites at high temperature. Fig. 7 plots the Ni, Co, and Mn K-edge XANES spectra of Li1.2(1 x)Ni0.13Mn0.54Co0.13O2 (x = 0.87) at 25 C and heated at 200, 300, 400 C, together with those of the reference samples. Fig. 7(a) indicates that the Ni K-edge spectra at 25 C were similar to those of LiNi0.8Co0.15Al0.05O2 (Ni3+). Nickel might be oxidized from 2+ to 4+ during the charge process; however, the energy of Ni K-edge spectra of Li1.2Ni0.13Mn0.54Co0.13O2 in a full-charge state at 25 C was almost the same as that of LiNi0.8Co0.15Al0.05O2 (Ni3+). This might be related to the following process. Nickel ions migrated from the transition metal layer to the lithium layer and oxygen was released; therefore, the energy of the spectra shifted to the lower energy side at the end of the charge process [37,38]. The spectra significantly shifted to the lower energy side when the temperature increased from 25 to 300 C, and finally approached that of NiO (Ni2+). This indicates that nickel was reduced from 25 to 300 C, and approached Ni2+. In contrast, the pre-edge intensity decreased due to heating, indicating that nickel ions did not demonstrate the symptoms of migration from octahedral to tetrahedral sites. Fig. 7(b) indicates that Co K-edge spectra shifted to the lower energy side with increasing temperature. Furthermore, some inflection points were observed in the spectrum at 300 C, and the shape of the spectrum at 300 C was similar to that of Co3O4. The shape of spectrum was maintained even at 400 C, and it did not change into CoO-like spectrum. In contrast, the energy above 300 C was lower than that of Co3O4. This indicates that cobalt might be reduced below 2.67+. In contrast, the pre-edge intensity increased from 200 to 400 C, indicating that cobalt ions might migrate from octahedral to tetrahedral sites due to heating, and occupy tetrahedral sites even at 400 C. The pre-edge intensity and the shape of the main spectra indicate that some of the cobalt ions occupied tetrahedral sites, and formed Co3O4-like structures with increasing temperature. This coincided with the XRD results in Fig. 5(b). Fig. 7(c) indicates that Mn K-edge spectra shifted to the lower energy side with increasing temperature. However, the change was less than that of Ni and Co K-edge spectra in Fig. 7(a) and (b). This indicates that the oxidation of manganese was less
315
changeable than that of nickel and cobalt. Furthermore, the pre-edge intensity of Mn K-edge spectra decreased due to heating. This indicates that manganese ions did not exhibit the symptoms of migration from octahedral to tetrahedral sites. The migration of transition metal observed in the XRD results in Fig. 5(b) was attributed to the migration of cobalt ions from the results of the pre-edges of Ni, Co, and Mn K-edge XANES spectra due to heating. 3.6. Difference between stability of charged lithium-rich layerstructured cathode and charged nickel-based layer-structured cathode at high temperature As shown in Fig. 3(b), the amount of oxygen released from Li1.2(1-x)Ni0.13Mn0.54Co0.13O2 (x = 0.8) was much less than that from Li1-yNi0.8Mn0.1Co0.1O2 (y = 0.8). Furthermore, even if 87% of lithium ions had been extracted from Li1.2Ni0.13Mn0.54Co0.13O2 in a full-charge state, the amount of oxygen released from Li1.2(1-x)Ni0.13Mn0.54Co0.13O2 (x = 0.87) would also be much less than that from Li1-yNi0.8Mn0.1Co0.1O2 (y=0.8) in a full-charge state. The high stability of Li1.2Ni0.13Mn0.54Co0.13O2 is considered to be as follows. The high oxidation state of nickel, such as Ni3+ and Ni4+, is unstable, so it is easily reduced to stable Ni2+ with the release of oxygen. Furthermore, nickel ions unstably occupy tetrahedral sites, so they easily move from octahedral sites in the transition metal layer to octahedral sites in the lithium layer through tetrahedral sites [13,14,39]. This movement changes the crystal structure from layer to spinel and then rocksalt, and a large amount of oxygen is released in Li1-yNi0.8Mn0.1Co0.1O2 (y = 0.8). Even though cobalt ions move from octahedral sites in the transition metal layer to tetrahedral sites, they stably occupy tetrahedral sites at high temperature in Li1.2(1-x)Ni0.13Mn0.54Co0.13O2 (x = 0.87). However, cobalt ions migrate from octahedral to tetrahedral, and then octahedral sites due to heating in Li1-yNi0.8Mn0.1Co0.1O2 (y = 0.8). Therefore, the effect of cobalt on stability at high temperature is slight, and there might be another reason for high stability in Li1.2Ni0.13Mn0.54Co0.13O2. As the high oxidation state of manganese, such as Mn4+, is more stable than that of Mn2+ and Mn3+, the oxidation state of manganese is less changeable than that of nickel and cobalt. Therefore, the suppression of oxygen release from delithiated Li1.2Ni0.13Mn0.54Co0.13O2 can be attributed to the high stability of manganese in a high oxidation state. The percentage of nickel to metal in the transition metal layer is only 13% for Li[Li0.2Ni0.13Mn0.54Co0.13]O2, and that of manganese exceeds 50%. The high stability of the lithium-rich layer-structured cathode, compared with the nickel-based cathode, can be attributed to the high manganese content of the transition metal. The former cathode is thus expected to provide both high capacity and high stability at elevated temperature. 4. Conclusions Li1.2Ni0.13Mn0.54Co0.13O2 sample was synthesized and found to exhibit high capacity of around 250 Ah kg1. The stability of the delithiated Li1.2Ni0.13Mn0.54Co0.13O2 at elevated temperature was evaluated by TDS-MS, XRD, and XAFS. According to the results from evaluations, the amount of oxygen released by heating increased with increasing delithiated content. Moreover, the amount of oxygen released from Li1.2Ni0.13Mn0.54Co0.13O2 in a full-charge state was lower than that released from LiNi0.8Mn0.1Co0.1O2 in a full-charge state. In contrast, the oxygen-release onset temperature in the former was lower than that in the latter. However, oxygen release at low temperature was reduced after a few cycles. The oxygen release at low temperature might be related to the structural change at the end of the charge process.
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In contrast, the suppression of oxygen release was related to the change in crystal structure at high temperature. The crystal structure of delithiated LiNi0.8Mn0.1Co0.1O2 changed from layer to spinel and then to rocksalt, and a large amount of oxygen was released during this structural change. In contrast, as the structure of delithiated Li1.2Ni0.13Mn0.54Co0.13O2 only changed in the first-step (i.e., layer to spinel), the amount of oxygen that was released was low. The XAFS spectra indicated that the oxidation state of manganese is less changeable than that of nickel or cobalt; therefore, the high stability of charged lithium-rich layerstructured cathode at high temperature was attributed to the high manganese content in the transition metal layer. Acknowledgements This work was supported by the New Energy and Industrial Technology Development Organization (NEDO) of Japan, and was carried out under the approval of the Photon Factory Program Advisory Committee (Proposal Nos. 2012C211 and 2013C211). References [1] J.R. Dahn, U. von. Sacken, M.W. Juzkow, H. Al-Janaby, J. Electrochem. Soc. 138 (1991) 2207. [2] T. Ohzuku, A. Ueda, M. Nagayama, J. Electrochem. Soc. 140 (1993) 1862. [3] K.S. Lee, S.T. Myung, K. Amine, H. Yashiro, Y.K. Sun, J. Electrochem. Soc. 154 (2007) A971. [4] H.J. Noh, S. Youn, C.S. Yoon, Y.K. Sun, J. Power Sources 233 (2013) 121. [5] M. Guilmard, L. Croguennec, C. Delmas, Chem. Mater. 15 (2003) 4484. [6] W.S. Yoon, M. Balasubramanian, X.Q. Yang, J. McBreen, J. Hanson, Electrochem. Solid State Lett. 8 (2005) A83. [7] W.S. Yoon, J. Hanson, J. McBreen, X.Q. Yang, Electrochem. Commun. 8 (2006) 859. [8] W.S. Yoon, K.Y. Chung, M. Balasubramanian, J. Hanson, J. McBreen, X.Q. Yang, J. Power Sources 163 (2006) 219. [9] I. Belharouak, D. Vissers, K. Amine, J. Electrochem. Soc. 153 (2006) A2030. [10] K.W. Nam, W.S. Yoon, X.Q. Yang, J. Power Sources 189 (2009) 515. [11] N. Yabuuchi, Y.T. Kim, H.H. Li, Y. Shao-Horn, Chem. Mater. 20 (2008) 4936. [12] H. Konishi, T. Yuasa, M. Yoshikawa, J. Power Sources 196 (2011) 6884.
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Electrochimica Acta journal homepage: www.elsevier.com/locate/electacta
Evaluation of Stability of Charged Lithium-rich Layer-structured Cathode Material at Elevated Temperature Hiroaki Konishi * , Tatsumi Hirano, Daiko Takamatsu, Akira Gunji, Xiaoliang Feng, Sho Furutsuki Hitachi Research Laboratory, Hitachi Ltd. 7-1-1 Omika-cho, Hitachi, Ibaraki 319-1292, Japan
A R T I C L E I N F O
A B S T R A C T
Article history: Received 3 December 2014 Received in revised form 7 March 2015 Accepted 31 March 2015 Available online 11 April 2015
The stability of charged Li1.2Ni0.13Mn0.54Co0.13O2 at elevated temperature was investigated by using thermal desorption spectrometry-mass spectrometry (TDS-MS), X-ray diffraction (XRD), and X-ray absorption fine structure (XAFS) measurements. The TDS-MS and XRD spectra indicated that the crystal structure of the delithiated Li1.2Ni0.13Mn0.54Co0.13O2 changed from layer to spinel, and oxygen was released at high temperature. However, the amount of oxygen released from Li1.2Ni0.13Mn0.54Co0.13O2 in a full-charge state was much less than that released from LiNi0.8Mn0.1Co0.1O2 in a full-charge state. The XAFS spectra indicated that the high oxidation state of manganese was more stable than that of nickel and cobalt. The high stability of charged Li1.2Ni0.13Mn0.54Co0.13O2 at high temperature was, therefore, attributed to the high manganese content of the transition metal. ã 2015 Elsevier Ltd. All rights reserved.
Keywords: Lithium-ion battery Lithium-rich layer-structured cathode Thermal decomposition Crystal structure Oxidation state
1. Introduction Lithium-ion batteries have recently been applied to not only small devices such as cellular phones but also automotive products such as hybrid electric vehicles (HEVs), plug-in hybrid electric vehicles (PHEVs), and electric vehicles (EVs). These applications require both high energy density and safety. These properties are closely related to those of materials used for the electrodes in batteries. Ni-based layer-structured cathode materials, LiNi1-aMaO2 (M: metal), are promising in that regard due to their high capacity [1–4]. However, the stability of Li1-yNi1-aMaO2 at elevated temperature is poor, especially in a highly delithiated state. The crystal structure of Li1-yNi1-aMaO2 changes from layer to spinel, and then rocksalt, and oxygen is released at high temperature [5–15]. The oxygen-release from the cathode can lead to thermal runway by reacting with electrolyte, and cause fire and explosion of lithium-ion battery. Therefore, the evaluation of thermal decomposition of the cathode is important. We elucidated the stability of an Ni-based cathode, LiNi0.8Mn0.1Co0.1O2, in a previous report, and its low stability due to heating was attributed to the instability of Ni3+ and Ni4+ in a high charge state [12,15]. Its poor stability has prevented it from being applied
* Corresponding author. Tel.: +81 294 52 5111x6089; fax: +81 294 52 7636. E-mail address: [email protected] (H. Konishi). http://dx.doi.org/10.1016/j.electacta.2015.03.217 0013-4686/ ã 2015 Elsevier Ltd. All rights reserved.
to large-scale batteries; therefore, new cathode materials that satisfy both high capacity and high stability criteria are required. It has been reported that lithium-rich layer-structured materials, Li[Li1-a-b-cNiaMnbMc]O2 (M: metal), exhibit high capacity around 250 Ah kg1 [16–20]. The crystal structure of Li1.2-xNi0.15Mn0.55Co0.1O2 changed from layer to spinel at high temperature [21]. Therefore, the oxygen-release from charged Li-rich layer-structured cathode at high temperature might be lower than that from charged Ni-based layer-structured cathode. However, there are no results compared by the identical condition. Furthermore, the cause of the difference in structure change has yet to be clarified. The stability of Li1.2Ni0.13Mn0.54Co0.13O2 in a charge state at high temperature was evaluated in the present study, and compared with that of LiNi0.8Mn0.1Co0.1O2 in a full-charge state. The oxygen release, crystal structure change, and oxidation state of each transition metal were investigated by using thermal desorption spectrometry-mass spectrometry (TDS-MS), X-ray diffraction (XRD), and X-ray absorption fine structure (XAFS) analyses. 2. Experimental Li1.2Ni0.13Mn0.54Co0.13O2 sample was synthesized by reacting stoichiometric mixtures of lithium acetate, nickel acetate, manganese acetate, and cobalt acetate. The mixture was stirred in distilled water for 1 h, and then dried. It was pre-heated at 500 C for 12 h in air, and then heated at 850 C for 12 h in air.
H. Konishi et al. / Electrochimica Acta 169 (2015) 310–316
3. Results and discussion 3.1. Crystal structure of Li1.2Ni0.13Mn0.54Co0.13O2 and LiNi0.8Mn0.1Co0.1O2 Fig. 1 shows the XRD patterns of Li1.2Ni0.13Mn0.54Co0.13O2 and LiNi0.8Mn0.1Co0.1O2 powder. Fig. 1(a) indicates that all the main
(018) / (110) (113)
(104)
(107)
(b)
(015)
(101) (006) / (102)
Intensity (a.u.)
(003)
Superlattice
(a) 10
20
30
40 2 /°
50
60
70
Fig. 1. XRD patterns of Li1.2Ni0.13Mn0.54Co0.13O2 and LiNi0.8Mn0.1Co0.1O2 powder.
peaks of Li1.2Ni0.13Mn0.54Co0.13O2 were assigned to a hexagonal a-NaFeO2 structure with space group R-3m. Small peaks also appeared between 2u = 20 and 25 , indicating the presence of a superlattice structure such as LiMn6 and LiMn5Ni in the transition metal layer [17,20]. Fig. 1(b) indicates that all the peaks of LiNi0.8Mn0.1Co0.1O2 were indexed to a hexagonal a-NaFeO2 structure. 3.2. Electrochemical properties of Li1.2Ni0.13Mn0.54Co0.13O2 and LiNi0.8Mn0.1Co0.1O2 The charge-discharge capacities of Li1.2Ni0.13Mn0.54Co0.13O2 and LiNi0.8Mn0.1Co0.1O2 were measured within a potential range of 2.5–4.6 V for the former and 2.5–4.3 V for the latter (vs. Li/Li+). Fig. 2 shows the initial charge and discharge curves of Li1.2Ni0.13Mn0.54Co0.13O2 and LiNi0.8Mn0.1Co0.1O2. The initial charge and discharge capacities of Li1.2Ni0.13Mn0.54Co0.13O2 were 327 and 251 Ah kg1, respectively. Even though 87% of lithium ions were de-intercalated in the charge process, only 67% of them were
+
/ V (vs. Li/Li )
4.5
Potential
LiNi0.8Mn0.1Co0.1O2 was synthesized by reacting lithium hydroxide, nickel oxide, manganese oxide, and cobalt oxide with 3% excess of lithium hydroxide. The mixture was pre-heated at 600 C for 10 h in oxygen, and then heated at 850 C for 10 h in oxygen. The crystal structures of the Li1.2Ni0.13Mn0.54Co0.13O2 and LiNi0.8Mn0.1Co0.1O2 samples were evaluated by XRD (Rigaku, Rint-2200). The diffraction data were recorded at a 0.02 step width over a 2u range from 10 to 70 using a Rigaku diffractometer with Cu Ka radiation (l = 1.54 Å). The electrochemical properties of Li1.2Ni0.13Mn0.54Co0.13O2 and LiNi0.8Mn0.1Co0.1O2 samples were measured by using a threeelectrode electrochemical cell containing Li1.2Ni0.13Mn0.54Co0.13O2 or LiNi0.8Mn0.1Co0.1O2 electrodes as a cathode, lithium foil as an anode, and a reference. The electrode was composed of the prepared sample, carbon, and binder (85: 10: 5 wt%). The electrolyte was a 1 M LiPF6 solution containing a mixture of ethylene carbonate (EC), ethyl methyl carbonate (EMC), and dimethyl carbonate (DMC) (1: 2: 2 vol%) solvents. The charge– discharge capacities of the samples were measured within a potential range of 2.5–4.6 V or 2.5–4.3 V (vs. Li/Li+). The cells were charged at 0.05 C (1 C = 260 A kg1) to 4.6 or 4.3 V, and then held at 4.6 or 4.3 V until the current became less than 0.005 C. After that, the cells were discharged at 0.05 C to 2.5 V. The lithium content after charging was measured using inductively coupled plasma atomic emission spectroscopy (ICP-AES, PerkinElmer, OPTIMA-3300XL). The oxygen released from the cathode by heating was evaluated by using TDS-MS (ESCO Ltd. EMD-WA1000). The samples for TDS-MS were prepared with the following procedure. The Li1.2Ni0.13Mn0.54Co0.13O2 and LiNi0.8Mn0.1Co0.1O2 electrodes were charged to the prescribed composition, Li1.2(1-x)Ni0.13Mn0.54Co0.13O2 (x = 0.6, 0.7, 0.8, 0.87) and Li1-yNi0.8Mn0.1Co0.1O2 (y = 0.8), at a rate of 0.05 C. After that, the electrochemical cells were decomposed, and the electrodes were washed with DMC. The delithiated samples were placed in the TDS-MS chamber and heated to 400 C in a vacuum at a rate of 5 C/min. The crystal structure of the Li1.2Ni0.13Mn0.54Co0.13O2 and LiNi0.8Mn0.1Co0.1O2 samples in a full-charge state at high temperature was evaluated by XRD. The samples for XRD were prepared with the following procedure. The Li1.2Ni0.13Mn0.54Co0.13O2 and LiNi0.8Mn0.1Co0.1O2 electrodes were charged up to a full-charge state, such as Li1.2(1-x)Ni0.13Mn0.54Co0.13O2 (x = 0.87) and Li1yNi0.8Mn0.1Co0.1O2 (y = 0.8). After that, the electrochemical cells were decomposed, and the electrodes were washed with DMC. The charged electrodes were heated at 150–400 C in an argon atmosphere for 1 h to expose them, and then cooled to room temperature (RT). The oxidation state of each transition metal was investigated by XAFS at the BL-9C of the Photon Factory (PF) for High Energy Accelerator Research Organization (KEK) in Japan. The samples for XAFS were prepared with the same procedure as that used to prepare the XRD samples. The oxidation state of each transition metal was mainly evaluated by X-ray absorption near-edge structure (XANES) analysis. The Ni, Co, and Mn K-edge XANES spectra of the prepared samples were measured in transmission mode by using an Si(111) double-crystal monochromator. The XANES spectra were compared with those of the reference samples such as lithium transition metal oxide and transition metal oxide.
311
4.0
3.5
3.0 Li1.2Ni0.13Mn0.54Co0.13O2 LiNi0.8Mn0.1Co0.1O2
2.5 0
100 Capacity
200 / Ah kg
300 -1
Fig. 2. Initial charge and discharge curves of Li1.2Ni0.13Mn0.54Co0.13O2 and LiNi0.8Mn0.1Co0.1O2 within potential range of 2.5–4.6 V, and 2.5–4.3 V vs. Li/Li+.
312
H. Konishi et al. / Electrochimica Acta 169 (2015) 310–316
Table 1 Lithium content calculated from charge capacity and that measured using ICP analysis. Cathodes
Calculated values
x = 0.87 Li1.2(1-x)Ni0.13Mn0.54Co0.13O2 Li1-yNi0.8Mn0.1Co0.1O2 y = 0.8 Li1.2(1-x)Ni0.13Mn0.54Co0.13O2: Ni + Mn + Co = 0.8 Li1-yNi0.8Mn0.1Co0.1O2: Ni + Mn + Co = 1.0
Experiment values x = 0.87 y = 0.78
intercalated in the discharge process. The initial charge and discharge capacities of LiNi0.8Mn0.1Co0.1O2 were 222 and 198 Ah kg1, respectively. This indicates that 80% of lithium ions were extracted in the charge process, and 72% of them were inserted in the discharge process. The irreversible capacity of Li1.2Ni0.13Mn0.54Co0.13O2 was larger than that of LiNi0.8Mn0.1Co0.1O2. The large irreversible capacity of the former was attributed to the irreversible loss of oxygen from the lattice [17–20].
(a)
=0.87 =0.8 =0.7 =0.6
200 Temperature
300 / °C
100
3.4. Change in oxygen release from Li1.2Ni0.13Mn0.54Co0.13O2 in a fullcharge state during charge-discharge cycle
400
(b)
100 150
200 Temperature
Table 1 lists the lithium content calculated from the charge capacity, and that measured using inductively coupled plasma (ICP) analysis. Table 1 indicates that there were no large differences in the calculated values from charge capacity and experimental values from ICP analysis. Fig. 3 shows the TDS-MS spectra of Li1.2(1-x)Ni0.13Mn0.54Co0.13O2 (x = 0.6, 0.7, 0.8, and 0.87) and Li1-yNi0.8Mn0.1Co0.1O2 (y = 0.8). Fig. 3(a) indicates that as delithiated content x increased, the amount of oxygen release increased and the oxygen-release onset temperature decreased. Fig. 3(b) indicates that the amount of oxygen release from Li1.2(1-x)Ni0.13Mn0.54Co0.13O2 (x = 0.8 and 0.87) was much less than that from Li1-yNi0.8Mn0.1Co0.1O2 (y = 0.8). In contrast, the oxygen-release onset temperature of the former was lower than that of the latter.
1st 2nd 3rd Intensity (a.u.)
Intensity (a.u.)
Li1.2(1- )Ni0.13Mn0.54Co0.13O2 ( =0.87) Li1.2(1- )Ni0.13Mn0.54Co0.13O2 ( =0.8) Li1- Ni0.8Mn0.1Co0.1O2 ( =0.8)
50
3.3. Comparison of stability of Li1.2(1-x)Ni0.13Mn0.54Co0.13O2 (x = 0.6, 0.7, 0.8, and 0.87) and Li1-yNi0.8Mn0.1Co0.1O2 (y = 0.8) at high temperature
Fig. 4 shows the TDS-MS spectra of Li1.2Ni0.13Mn0.54Co0.13O2 in a full-charge state during initial, second, and third charge process. Fig. 4 indicates that two oxygen release peaks were observed in the initial cycle. As cycle number increased, the oxygen release at low temperature was reduced, and only one peak was observed in the third cycle. The oxygen release at low temperature might be related to the charge reaction process of this material. It has been reported that oxygen was extracted in lithium-rich layer-structured material, and transition metal migrated from the surface to the bulk in the last state of charge reaction even at RT [22,23]. Since the migration of transition metal might be slow at RT, this reaction might not fully progress at RT. The migration of transition metal might have been prompted, and oxygen was released due to heating; therefore,
Intensity (a.u.)
100
Li1.2(1-x)Ni0.13Mn0.54Co0.13O2 (x = 0.87) and Li1-yNi0.8Mn0.1Co0.1O2 (y = 0.8) were defined as the full-charge state of each material.
300 / °C
400
Fig. 3. TDS-MS spectra of (a) Li1.2(1-x)Ni0.13Mn0.54Co0.13O2 (x = 0.6, 0.7, 0.8, 0.87), and (b) Li1.2(1-x)Ni0.13Mn0.54Co0.13O2 (x = 0.8, 0.87) and Li1-yNi0.8Mn0.1Co0.1O2 (y = 0.8).
100
200 Temperature
300 / °C
400
Fig. 4. TDS-MS spectra of Li1.2Ni0.13Mn0.54Co0.13O2 in a full-charge state during initial, second, and third charge process.
H. Konishi et al. / Electrochimica Acta 169 (2015) 310–316
(220) spinel
(a)
400°C
300°C
Intensity (a.u.)
3.5. Change in structure, oxidation state, and local structure of Li1.2Ni0.13Mn0.54Co0.13O2 and LiNi0.8Mn0.1Co0.1O2 in initial full-chargestate at high temperature The change in structure, oxidation state, and local structure of Li1.2Ni0.13Mn0.54Co0.13O2 in a full-charge state at high temperature was compared with that of LiNi0.8Mn0.1Co0.1O2 in a full-charge state reported in our previous paper [15]. Fig. 5 shows the XRD patterns of Li1-yNi0.8Mn0.1Co0.1O2 (y = 0.8) and Li1.2(1-x)Ni0.13Mn0.54Co0.13O2 (x = 0.87) at 25 C and heated at 150–400 C.
350°C
Fig. 5(a) indicates that the layer structure was maintained from 25 to 150 C. A new peak at 2u = 31 emerged at 200 C, indicating that some of the transition metal ions occupied tetrahedral sites [11,24,25]. Furthermore, the (018) and (110) peaks at 2u = 65 merged at 200 C. These features are characteristics of a spinel structure [5–15]. However, the peak at 2u = 31 vanished at 300 C. Furthermore, the peak at 2u = 18 decreased, and the peaks at 2u = 38, 44, and 64 shifted to smaller angles. This indicates that the layer structure was changed to spinel, and then rocksalt, accompanying oxygen release, according to the following reactions [26,27].
250°C
200°C
150°C
25°C
Li1-yMO2 ! (2-y)/3[Li3(1-y)/(2-y)M(2 O2
20
30
40 50 2 /°
60
350°C
Intensity (a.u.)
300°C
250°C
150°C
25°C
40 50 2 /°
60
! [Li1-yM]O2-y + [(2-y)/6]O2 + [(2y-1)/3]O2
(2)
The relative intensity of each peak changed from 250 to 300 C. A new peak at 2u = 31 emerged at 300 C, and this was clearly observed even at 400 C. This indicates that some of the transition metal ions occupied tetrahedral sites from 300 to 400 C [11,24,25]. Furthermore, the (018) and (110) peaks around 2u = 65 merged at 300 C. However, the peaks did not change much from 300 to 400 C.
200°C
30
(1)
Fig. 5(b) indicates that the relative intensity ratio of the (003) peak at 2u = 18 and the (018) peak at 2u = 64 was low, compared with that of the pristine state, as shown in Fig. 1. This indicates that there were two phases, which have different lattice parameter c, at the end of the charge process [23,28]. Even though the layer structure was maintained from 25 to 250 C, the relative intensity of the (003) and (018) peaks increased. This might be attributed to the two phases being changed into one phase due to heating. This change might be accompanied by the migration of transition metal and oxygen release, therefore, oxygen release was observed in this temperature range, as shown in Fig. 3(b).
400°C
20
y-1)/(2-y)][M]2O4 + [(2y-1)/3]
70
(220) spinel
(b)
313
70
Fig. 5. XRD patterns of (a) Li1-yNi0.8Mn0.1Co0.1O2 (y = 0.8) and (b) Li1.2(1 x)Ni0.13Mn0.54Co0.13O2 (x = 0.87) at 25 C and heated at 150–400 C.
oxygen release was observed at low temperature in the highly delithiated Li1.2Ni0.13Mn0.54Co0.13O2. As the cycle number increased, the migration of transition metal progressed; therefore, oxygen release at low temperature might be reduced. In contrast, oxygen release at high temperature was clearly observed after a few cycles, then it might be related to the crystal structure change.
A reaction similar to formula (1) was also observed for delithiated Li1.2Ni0.13Mn0.54Co0.13O2. However, no structural change like that in formula (2) was observed. We concluded that the suppressed reaction in formula (2) could be attributed to the fact that the amount of oxygen release from Li1.2Ni0.13Mn0.54Co0.13O2 in a full-charge state was lower than that from LiNi0.8Mn0.1Co0.1O2 in a full-charge state. Fig. 6 plots Ni, Co, and Mn K-edge XANES spectra in Li1-yNi0.8Mn0.1Co0.1O2 (y = 0.8) at 25 C and heated at 200, 300, 400 C, together with those of the reference samples such as lithium transition metal oxide and transition metal oxide. All spectra exhibit small pre-edge and large peaks. The former is related to dipole forbidden 1s ! 3d electronic transition. However, this transition is partly allowed due to the hybridization of the 3d-4p orbital caused by structural distortion around the central transition metal. The latter is associated with dipole allowed 1s ! 4p electronic transition. The energy of the spectra is related to the oxidation state of transition metal [29–36], and the pre-edge
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2.0
2.0 (a)
(a) 1.5
1.0 8328 8332 8336 LiNi0.8Co0.15Al0.15O2 NiO 400°C 300°C 200°C 25°C
0.5
0.0 8330
8335
8340 8345 Energy / eV
8350
Normalized (a.u.)
Normalized (a.u.)
1.5
1.0 8328 8332 8336 LiNi0.8Co0.15Al0.15O2 NiO 400°C 300°C 200°C 25°C
0.5
0.0 8330
8355
8335
8340 8345 Energy / eV
8350
2.0
2.0
(b)
(b)
1.5
1.0 7704 7708 7712 LiCoO2 Co3O4 CoO 400°C 300°C 200°C 25°C
0.5
0.0 7705
7710
7715 7720 Energy / eV
7725
Normalized (a.u.)
Normalized (a.u.)
1.5
1.0 7704 7708 7712 LiCoO2 Co3O4 CoO 400°C 300°C 200°C 25°C
0.5
0.0 7705
7730
7710
7715 7720 Energy / eV
7725
7730
2.0
2.0
(c)
(c)
1.0 6536 6540 6544 MnO2 Mn2O3 400°C 300°C 200°C 25°C
0.5
0.0 6540
6545
6550 6555 Energy / eV
6560
6565
Fig. 6. (a) Ni, (b) Co, and (c) Mn K-edge XANES spectra of Li1-yNi0.8Mn0.1Co0.1O2 (y = 0.8) at 25 C and heated at 200, 300, 400 C, together with those of reference samples such as lithium transition metal oxide and transition metal oxide.
Normalized (a.u.)
1.5
1.5 Normalized (a.u.)
8355
1.0 6536 6540 6544 MnO2 Mn2O3 400°C 300°C 200°C 25°C
0.5
0.0 6540
6545
6550 6555 Energy / eV
6560
6565
Fig. 7. (a) Ni, (b) Co, and (c) Mn K-edge XANES spectra of Li1.2(1-x)Ni0.13Mn0.54Co0.13O2 (x = 0.87) at 25 C and heated at 200, 300, 400 C, together with those of reference samples such as lithium transition metal oxide and transition metal oxide.
H. Konishi et al. / Electrochimica Acta 169 (2015) 310–316
intensity is related to the occupation sites of the transition metal ions. This is stronger when transition metal ions occupy tetrahedral sites than when they occupy octahedral sites [13,14,29–32]. As we can see from the XRD results in Fig. 5(a) and (b), some of the transition metal ions migrate from octahedral to tetrahedral sites due to heating. Therefore, the change in pre-edge intensity was used to elucidate what kinds of transition metal ions migrated. Fig. 6(a) indicates that Ni K-edge spectrum at 25 C was on an energy side higher than that of LiNi0.8Co0.15Al0.05O2 (Ni3+). The spectra shifted to a lower energy side from 25 to 300 C, and they approached that of NiO (Ni2+), indicating that nickel was reduced due to heating [13,14,33,34]. In contrast, the pre-edge intensity decreased due to heating. This suggests that nickel ions did not demonstrate the symptoms of occupation of tetrahedral sites at high temperature. Fig. 6(b) indicates that the energy of Co K-edge spectrum at 25 C was close to that of LiCoO2 (Co3+). The spectra shifted to a lower energy side with increasing temperature. In contrast, the pre-edge intensity increased from 25 to 200 C, and it decreased from 200 to 400 C. This indicates that cobalt ions moved from octahedral to tetrahedral sites at 25–200 C, and then octahedral sites at 200–400 C. This coincides with the XRD results in Fig. 5(a). Fig. 6(c) indicates that Mn K-edge spectra shifted to the lower energy side with increasing temperature. However, the change was less than that of Ni and Co. Furthermore, the pre-edge intensity indicates that manganese ions did not exhibit the symptoms of occupation of tetrahedral sites at high temperature. Fig. 7 plots the Ni, Co, and Mn K-edge XANES spectra of Li1.2(1 x)Ni0.13Mn0.54Co0.13O2 (x = 0.87) at 25 C and heated at 200, 300, 400 C, together with those of the reference samples. Fig. 7(a) indicates that the Ni K-edge spectra at 25 C were similar to those of LiNi0.8Co0.15Al0.05O2 (Ni3+). Nickel might be oxidized from 2+ to 4+ during the charge process; however, the energy of Ni K-edge spectra of Li1.2Ni0.13Mn0.54Co0.13O2 in a full-charge state at 25 C was almost the same as that of LiNi0.8Co0.15Al0.05O2 (Ni3+). This might be related to the following process. Nickel ions migrated from the transition metal layer to the lithium layer and oxygen was released; therefore, the energy of the spectra shifted to the lower energy side at the end of the charge process [37,38]. The spectra significantly shifted to the lower energy side when the temperature increased from 25 to 300 C, and finally approached that of NiO (Ni2+). This indicates that nickel was reduced from 25 to 300 C, and approached Ni2+. In contrast, the pre-edge intensity decreased due to heating, indicating that nickel ions did not demonstrate the symptoms of migration from octahedral to tetrahedral sites. Fig. 7(b) indicates that Co K-edge spectra shifted to the lower energy side with increasing temperature. Furthermore, some inflection points were observed in the spectrum at 300 C, and the shape of the spectrum at 300 C was similar to that of Co3O4. The shape of spectrum was maintained even at 400 C, and it did not change into CoO-like spectrum. In contrast, the energy above 300 C was lower than that of Co3O4. This indicates that cobalt might be reduced below 2.67+. In contrast, the pre-edge intensity increased from 200 to 400 C, indicating that cobalt ions might migrate from octahedral to tetrahedral sites due to heating, and occupy tetrahedral sites even at 400 C. The pre-edge intensity and the shape of the main spectra indicate that some of the cobalt ions occupied tetrahedral sites, and formed Co3O4-like structures with increasing temperature. This coincided with the XRD results in Fig. 5(b). Fig. 7(c) indicates that Mn K-edge spectra shifted to the lower energy side with increasing temperature. However, the change was less than that of Ni and Co K-edge spectra in Fig. 7(a) and (b). This indicates that the oxidation of manganese was less
315
changeable than that of nickel and cobalt. Furthermore, the pre-edge intensity of Mn K-edge spectra decreased due to heating. This indicates that manganese ions did not exhibit the symptoms of migration from octahedral to tetrahedral sites. The migration of transition metal observed in the XRD results in Fig. 5(b) was attributed to the migration of cobalt ions from the results of the pre-edges of Ni, Co, and Mn K-edge XANES spectra due to heating. 3.6. Difference between stability of charged lithium-rich layerstructured cathode and charged nickel-based layer-structured cathode at high temperature As shown in Fig. 3(b), the amount of oxygen released from Li1.2(1-x)Ni0.13Mn0.54Co0.13O2 (x = 0.8) was much less than that from Li1-yNi0.8Mn0.1Co0.1O2 (y = 0.8). Furthermore, even if 87% of lithium ions had been extracted from Li1.2Ni0.13Mn0.54Co0.13O2 in a full-charge state, the amount of oxygen released from Li1.2(1-x)Ni0.13Mn0.54Co0.13O2 (x = 0.87) would also be much less than that from Li1-yNi0.8Mn0.1Co0.1O2 (y=0.8) in a full-charge state. The high stability of Li1.2Ni0.13Mn0.54Co0.13O2 is considered to be as follows. The high oxidation state of nickel, such as Ni3+ and Ni4+, is unstable, so it is easily reduced to stable Ni2+ with the release of oxygen. Furthermore, nickel ions unstably occupy tetrahedral sites, so they easily move from octahedral sites in the transition metal layer to octahedral sites in the lithium layer through tetrahedral sites [13,14,39]. This movement changes the crystal structure from layer to spinel and then rocksalt, and a large amount of oxygen is released in Li1-yNi0.8Mn0.1Co0.1O2 (y = 0.8). Even though cobalt ions move from octahedral sites in the transition metal layer to tetrahedral sites, they stably occupy tetrahedral sites at high temperature in Li1.2(1-x)Ni0.13Mn0.54Co0.13O2 (x = 0.87). However, cobalt ions migrate from octahedral to tetrahedral, and then octahedral sites due to heating in Li1-yNi0.8Mn0.1Co0.1O2 (y = 0.8). Therefore, the effect of cobalt on stability at high temperature is slight, and there might be another reason for high stability in Li1.2Ni0.13Mn0.54Co0.13O2. As the high oxidation state of manganese, such as Mn4+, is more stable than that of Mn2+ and Mn3+, the oxidation state of manganese is less changeable than that of nickel and cobalt. Therefore, the suppression of oxygen release from delithiated Li1.2Ni0.13Mn0.54Co0.13O2 can be attributed to the high stability of manganese in a high oxidation state. The percentage of nickel to metal in the transition metal layer is only 13% for Li[Li0.2Ni0.13Mn0.54Co0.13]O2, and that of manganese exceeds 50%. The high stability of the lithium-rich layer-structured cathode, compared with the nickel-based cathode, can be attributed to the high manganese content of the transition metal. The former cathode is thus expected to provide both high capacity and high stability at elevated temperature. 4. Conclusions Li1.2Ni0.13Mn0.54Co0.13O2 sample was synthesized and found to exhibit high capacity of around 250 Ah kg1. The stability of the delithiated Li1.2Ni0.13Mn0.54Co0.13O2 at elevated temperature was evaluated by TDS-MS, XRD, and XAFS. According to the results from evaluations, the amount of oxygen released by heating increased with increasing delithiated content. Moreover, the amount of oxygen released from Li1.2Ni0.13Mn0.54Co0.13O2 in a full-charge state was lower than that released from LiNi0.8Mn0.1Co0.1O2 in a full-charge state. In contrast, the oxygen-release onset temperature in the former was lower than that in the latter. However, oxygen release at low temperature was reduced after a few cycles. The oxygen release at low temperature might be related to the structural change at the end of the charge process.
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In contrast, the suppression of oxygen release was related to the change in crystal structure at high temperature. The crystal structure of delithiated LiNi0.8Mn0.1Co0.1O2 changed from layer to spinel and then to rocksalt, and a large amount of oxygen was released during this structural change. In contrast, as the structure of delithiated Li1.2Ni0.13Mn0.54Co0.13O2 only changed in the first-step (i.e., layer to spinel), the amount of oxygen that was released was low. The XAFS spectra indicated that the oxidation state of manganese is less changeable than that of nickel or cobalt; therefore, the high stability of charged lithium-rich layerstructured cathode at high temperature was attributed to the high manganese content in the transition metal layer. Acknowledgements This work was supported by the New Energy and Industrial Technology Development Organization (NEDO) of Japan, and was carried out under the approval of the Photon Factory Program Advisory Committee (Proposal Nos. 2012C211 and 2013C211). References [1] J.R. Dahn, U. von. Sacken, M.W. Juzkow, H. Al-Janaby, J. Electrochem. Soc. 138 (1991) 2207. [2] T. Ohzuku, A. Ueda, M. Nagayama, J. Electrochem. Soc. 140 (1993) 1862. [3] K.S. Lee, S.T. Myung, K. Amine, H. Yashiro, Y.K. Sun, J. Electrochem. Soc. 154 (2007) A971. [4] H.J. Noh, S. Youn, C.S. Yoon, Y.K. Sun, J. Power Sources 233 (2013) 121. [5] M. Guilmard, L. Croguennec, C. Delmas, Chem. Mater. 15 (2003) 4484. [6] W.S. Yoon, M. Balasubramanian, X.Q. Yang, J. McBreen, J. Hanson, Electrochem. Solid State Lett. 8 (2005) A83. [7] W.S. Yoon, J. Hanson, J. McBreen, X.Q. Yang, Electrochem. Commun. 8 (2006) 859. [8] W.S. Yoon, K.Y. Chung, M. Balasubramanian, J. Hanson, J. McBreen, X.Q. Yang, J. Power Sources 163 (2006) 219. [9] I. Belharouak, D. Vissers, K. Amine, J. Electrochem. Soc. 153 (2006) A2030. [10] K.W. Nam, W.S. Yoon, X.Q. Yang, J. Power Sources 189 (2009) 515. [11] N. Yabuuchi, Y.T. Kim, H.H. Li, Y. Shao-Horn, Chem. Mater. 20 (2008) 4936. [12] H. Konishi, T. Yuasa, M. Yoshikawa, J. Power Sources 196 (2011) 6884.
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