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Synthesis mechanisms of lithium cobalt oxide prepared by hydrothermal–electrochemical method
Journal of Alloys and Compounds 430 (2007) 222–225
Synthesis mechanisms of lithium cobalt oxide prepared by hydrothermal–electrochemical method Ying Tao a,∗ , Baojun Zhu b , Zhenhua Chen b a
Institute of Materials Science and Engineering, Central South University, Changsha 410083, China b Institute of Materials Science and Engineering, Hunan University, Changsha 410082, China Received 5 February 2006; received in revised form 28 April 2006; accepted 28 April 2006 Available online 27 June 2006
Abstract Thermodynamic calculation method was adopted to predict the reaction mechanism of LiCoO2 prepared by hydrothermal–electrochemical process. It was found that in the Co–LiOH–H2 O system, Co was oxidized to HCoO2 − , Co(OH)2 (100 ◦ C) or CoO (150 ◦ C), CoOOH in sequence with the increase of electrode potential, then the ion-exchange reaction of CoOOH and lithium ion occurred and LiCoO2 came into being. The optimum synthesis parameter was obtained through thermodynamic calculation and it was validated experimentally by cyclic voltammogram method. © 2006 Elsevier B.V. All rights reserved. Keywords: Oxide materials; Chemical synthesis; Thermodynamic; Phase transitions
1. Introduction LiCoO2 has been commercially used as cathode material in lithium ion battery during the past 10 years for its high potential, excellent reversibility, and easy preparation on an industrial scale [1–3]. Significant research has been done on the formation of LiCoO2 via solid-state reaction and sol–gel methods. These techniques require a high energy and material consumption [4–6]. So an attempt to develop an alternative synthetic route, such as hydrothermal method, electrochemical–hydrothermal method, to produce the desired cathode materials in an economical and environmentally friendly way is in progress [7,8]. Hydrothermal treatment of cobalt metal plates leads directly to the formation of LiCoO2 films [9]. By using electrochemical–hydrothermal approach, crystallized and electrochemically active LiCoO2 films can be effectively prepared in a single synthetic step from cobalt metal plates [10–13]. In this paper, a series of thermodynamic calculation was used to predict the optimum synthesis conditions of LiCoO2 powders and films in lithium hydroxide solution. The optimum synthesis parameter obtained was validated experimentally by cyclic voltammogram method. A possible formation mechanics ∗
Corresponding author. E-mail address: [email protected] (Y. Tao).
0925-8388/$ – see front matter © 2006 Elsevier B.V. All rights reserved. doi:10.1016/j.jallcom.2006.04.074
of hexagonal LiCoO2 in the Co–LiOH–H2 O system was proposed. 2. Theoretical concept of potential–pOH diagram drawing system The details of the construction of potential–pOH diagrams were described in the previous paper [14]. The thermodynamic data of chemical species required for drawing the potential–pOH diagrams in Co–H2 O–LiOH systems are summarized in Table 1 [15–19]. The reaction equations of chemical species with OH− in Co–LiOH systems are shown in Table 2. From the thermodynamic data in Table 1, the changes of Gibbs energy at various temperatures for each reaction equation in Table 2 were calculated by using a personal computer. Base on these values, the relationship between potential and pOH was evaluated and the potential–pOH diagrams in Co–LiOH systems were drawn at 100 and 150 ◦ C. In addition, the solubilities of chemical species were calculated at given temperature. Base on this information, the thermodynamic stabilities of Co were evaluated in 4 M LiOH solution at various temperatures. 3. Results and discussion Figs. 1 and 2 show the potential–pOH diagrams in Co–H2 O system at 100 and 150 ◦ C, respectively. As could be seen from
Y. Tao et al. / Journal of Alloys and Compounds 430 (2007) 222–225
223
Table 1 Thermodynamic properties of components at 298.15 K Species
H2 O H2 O2 H+ OH− Li+ Co Co(OH)2 CoO Co3 O4 Co2 O3 CoOOH CoO2 Co2+ HCoO2 − LiCoO2
State
l g g aq aq aq c c c c c c c aq aq c
G0f,298.15 (kJ mol−1 )
−237.18 0 0 0 −157.29 −293.8 0 −454.4 −214.2 −774.0 −558.1 −349.5 −216.9 −54.4 −347.2 −561.4
0 S298.15 (abs) (J mol−1 K−1 )
69.91 130.58 205.03 −5 10.2 9.2 30.07 79.5 52.97 102.5 43.77 66.98 53.10 −155 62.8 34.91
Fig. 1. Potential–pOH diagram of the Co–H2 O system at 100 ◦ C.
Cp (J mol−1 K−1 ) A
B
C
20.335 27.28 29.96
1.092×10−1
2.033×106 5.0×104 −1.67×105
19.94 82.39 55.1 136.65 81.55 172.24 128.22
68.87
3.264×10−3 4.14×10−3 Cation Anion Cation 1.656×10−2 4.846×10−2 −4.6×10−3 2.753×10−2 3.213×10−2 −2.369×10−1 −1.35×10−1 Cation Acid oxyanion 3.617×10-2
−1.647×106
these figures, the numbers on the lines represent the reaction equations given in Table 2. Dotted lines (a) and (b) showed the equilibrium potentials for hydrogen and oxygen evolution reactions in Table 2, respectively. The small numbers in parenthesis showed the logarithms of the activity of ionic species. The pOH values of 4 M LiOH solution, which was used as an electrolyte, were −0.88 at 100 ◦ C and −0.77 at 150 ◦ C. According to Fig. 1, it could be seen that at 100 ◦ C, the equilibrium potentials for Co/CoO and Co/Co(OH)2 couples were more positive than that of the hydrogen evolution reaction at pOH −0.88. While the equilibrium potential for Co/HCoO2 − couple (Eq. (3)) at an activity of 10−4.7 M was equal to that of hydrogen evolution reaction. Therefore, at 100 ◦ C and pOH −0.88 the first product of Co was HCoO2 − ion, whose saturated concentration was 10−4.7 . At the force of electric field, HCoO2 − ions moved to anode. Then the concentration of HCoO2 − ion was above its saturated concentration at the surrounding of anode. HCoO2 − ion further transformed to Co(OH)2 which was thermodynamically more stable than CoO at that condition. The equilibrium potential for Co(OH)2 /CoOOH couple (Eq. (7)) at 100 ◦ C and pOH −0.88 Table 2 The reactions for Co in LiOH solution
Fig. 2. Potential–pOH diagram of the Co–H2 O system at 150 ◦ C.
−6.0×103 2.1×104 −1.67×105 −2.289×106 −2.122×106 7.11×105 −1.795×106
H2 O + e → OH− + H(g) O2 + 2H2 O + 4e → 4OH− Co + 2OH− → Co(OH)2 + 2e Co + 2OH− → CoO + H2 O + 2e Co + 3OH− → HCoO2 − + H2 O + 2e CoO + H2 O → Co(OH)2 Co(OH)2 + OH− → HCoO2 − + H2 O CoO + OH− → HCoO2 − Co(OH)2 + OH− → CoOOH + H2 O + e CoO + OH− → CoOOH + e HCoO2 − → CoOOH + e CoOOH + OH− → CoO2 + H2 O + e CoOOH + Li+ + OH− → LiCoO2 + H2 O Co + Li+ + 4OH− → LiCoO2 + 2H2 O + 3e
(a) (b) (1) (2) (3) (4) (5) (6) (7) (8) (9) (10) (11) (12)
224
Y. Tao et al. / Journal of Alloys and Compounds 430 (2007) 222–225
Fig. 4. Cyclic voltammogram of Co electrode in a 4 M LiOH solution at a scan rate of 3 mV s−1 . Fig. 3. Schematic diagram for tracing out an area of formation LiCoO2 .
was 0.15 V versus NHE. At 100 ◦ C and pOH −0.88, the reaction proceeded in the following order: Co, HCoO2 − , Co(OH)2 , and CoOOH. At 150 ◦ C and pOH −0.77, as shown in Fig. 2, because the equilibrium potential for H2 O/H couple equaled to the equilibrium potentials for Co/HCoO2 − couple at an activity of 10−3.6 M, metal Co was easily oxidized to HCoO2 − ion, whose saturated concentration was 10−3.6 M. At the force of electric field, HCoO2 − ions moved to anode. Then the concentration of HCoO2 − ion was above its saturated concentration at the surrounding of anode, so HCoO2 − ion would transform to CoO which was thermodynamically more stable than Co(OH)2 at that condition. The equilibrium potential for the CoO/CoOOH couple (Eq. (6)) at 150 ◦ C and pOH −0.77 was 0.10 V versus NHE. At 150 ◦ C and pOH −0.77, the reaction proceeded in the following order: Co, HCoO2 − , CoO, and CoOOH. The potential to form CoOOH phase shifted toward the negative direction with increasing temperature. Without the influence of other ions, Co was finally oxidized to CoO2 . The equilibrium potential for the CoOOH/CoO2 couple (Eq. (10)) is shown in Figs. 1 and 2. At 100 ◦ C and pOH −0.88, the potential was 0.436 V. At 150 ◦ C and pOH −0.77, the potential was 0.396 V. In a 4 M LiOH solution, there was a possibility of a Li+ /H+ exchange to obtain LiCoO2 from CoOOH [20]. Because its Gibbs energy was less than zero, as shown in Fig. 3, Eq. (11) could spontaneously proceed in the left area of line 11. LiCoO2 could be synthesized in the area enclosed by Y-axis, line 10, line 11, and line 7 (100 ◦ C) or line 8 (150 ◦ C), as shown in Fig. 3. In 4 M LiOH solution, LiCoO2 could be prepared in the potential range of 0.1–0.396 V versus NHE at 150 ◦ C and 0.15–0.436 V versus NHE at 100 ◦ C, respectively. The cyclic voltammogram (CV) of LiCoO2 film electrode posited on cobalt substrate in 4 M LiOH is shown in Fig. 4. The CV measures were made in a conventional three-electrode cell using cobalt as work electrode, Ag/AgCl as reference electrode, and platinum slice as counter electrode, and the reaction temperatures were 100 and 150 ◦ C, respectively. The cycling was performed between the voltages −1.0 and 0.5 V at a scanning
rate of 3 mV s−1 . The experiments were performed by means of a Solartron Analytical-1287 Potentiostat. It could be seen that there were two clear anodic peaks in Fig. 4. The first peak was associated with the oxidation of Co/HCoO2 − . When the concentration of HCoO2 − ion was above its saturated concentration, HCoO2 − ion would transform to Co(OH)2 and CoO at 100 and 150 ◦ C, respectively. The second one was corresponds to the reaction of Co(OH)2 /CoOOH or CoO/CoOOH. No cathodic peak is found in Fig. 4, which mean no CoOOH and CoO2 phase existed at the end of the positive scan. Based on the analysis, a conclusion could be drawn that LiCoO2 was prepared by the following H+ /Li+ ion exchange reaction. Li+ + CoOOH → LiCoO2 + H+
(13)
The result of cyclic voltammetry measures in Fig. 4 was concordance with that of thermodynamic analysis. The X-ray diffraction pattern of LiCoO2 film prepared by the hydrothermal electrochemical method is shown in Fig. 5. As was evident, the diffraction patterns of film showed not only the reflections of the film itself, but also those of cobalt substrate. After peaks corresponding to cobalt substrate were excluded, the left diffraction pattern can be indexed to a hexagonal unit ¯ space group which corresponds to LiCoO2 film. cell with R3m
Fig. 5. XRD pattern of LiCoO2 film obtained after hydrothermal electrochemical treatment of Co substrate.
Y. Tao et al. / Journal of Alloys and Compounds 430 (2007) 222–225
4. Conclusion Potential–pOH diagrams in Co–H2 O systems at 100 ◦ C and 150 ◦ C were drawn and equilibrium potential and solubility of cobalt species were calculated. In 4 M LiOH, the reaction proceeded in the following order: Co, HCoO2 − , Co(OH)2 (100 ◦ C) or CoO (150 ◦ C), CoOOH, LiCoO2 . LiCoO2 was synthesized in the potential range of 0.15–0.436 V versus NHE at 100 ◦ C and 0.1–0.396 V versus NHE at 150 ◦ C. The thermodynamic result was in concordance with the cyclic voltammogram analysis. LiCoO2 was synthesized via H+ /Li+ ion-exchange reaction in the CoOOH precursors. Acknowledgement The authors would like to thank for the support from the Natural Science Foundation of China (Grant no. 50401011). References [1] [2] [3] [4]
A. Ermete, F. Maurizio, J. Solid State Chem. 117 (1995) 1. S. Nieto-Ramos, M.S. Tomar, Int. J. Hydrogen Energy 26 (2001) 159. A.G. Ritchie, J. Power Sources 96 (2001) 1. C. Delmas, J.J. Braconnier, P. Hagenmuller, Mater. Res. Bull. 17 (1982) 117.
225
[5] A. Kalyan, T. Masao, K. Shinichi, Mater. Res. Bull. 33 (1998) 1845. [6] E. Antolini, J. Eur. Ceram. Soc. 18 (1998) 1405. [7] J.M. Fernandez-Rodriguez, J. Morales, J.L. Tirado, React. Solids 4 (1987) 163. [8] A.G. Ritchie, J. Power Sources 96 (2001) 1. [9] J.M. Fernadez-Rodriguez, L. Hernan, J. Morales, Mater. Res. Bull. 23 (1988) 899. [10] K.S. Han, S.W. Song, H. Fujita, M. Yoshimura, Solid State Ionics 135 (2000) 273. [11] T. Fujiwara, M. Yoshimura, J. Electroanal. Chem. 559 (2003) 63. [12] K.S. Han, S.W. Song, S.S. Tsurimoto, Solid State Ionics 151 (2002) 11. [13] Y. Tao, Z.H. Chen, B.J. Zhu, W.Z. Huang, Solid State Ionics 161 (2003) 187. [14] Y. Tao, Z.H. Chen, B.J. Zhu, Physica B 362 (2005) 76. [15] C.M. Criss, J.W. Cobble, J. Am. Chem. Soc. 86 (1964) 5390. [16] C.M. Chen, K. Aral, Corrosion-Nace. 38 (1982) 183. [17] M.W. Chase Jr., C.A. Davies, J.R. Downey Jr., D.J. Frurip (Eds.), JANAF Thermochemical Tables, third ed., American Institute of Physics, New York, 1986. [18] O. Knacke, O. Kubaschewski, K. Hesselmann (Eds.), Thermochemical Properties of Inorganic Substance, Springer-Verlag Verlag Stahleisen, 1991. [19] Y.C. Zhang, H. Tagawa, S. Asakura, J. Electrochem. Soc. 144 (1997) 4345. [20] D. Larcher, M.R. Palacin, G.G. Amatucci, J.M. Tarascon, J. Electrochem. Soc. 144 (1997) 408.
Synthesis mechanisms of lithium cobalt oxide prepared by hydrothermal–electrochemical method Ying Tao a,∗ , Baojun Zhu b , Zhenhua Chen b a
Institute of Materials Science and Engineering, Central South University, Changsha 410083, China b Institute of Materials Science and Engineering, Hunan University, Changsha 410082, China Received 5 February 2006; received in revised form 28 April 2006; accepted 28 April 2006 Available online 27 June 2006
Abstract Thermodynamic calculation method was adopted to predict the reaction mechanism of LiCoO2 prepared by hydrothermal–electrochemical process. It was found that in the Co–LiOH–H2 O system, Co was oxidized to HCoO2 − , Co(OH)2 (100 ◦ C) or CoO (150 ◦ C), CoOOH in sequence with the increase of electrode potential, then the ion-exchange reaction of CoOOH and lithium ion occurred and LiCoO2 came into being. The optimum synthesis parameter was obtained through thermodynamic calculation and it was validated experimentally by cyclic voltammogram method. © 2006 Elsevier B.V. All rights reserved. Keywords: Oxide materials; Chemical synthesis; Thermodynamic; Phase transitions
1. Introduction LiCoO2 has been commercially used as cathode material in lithium ion battery during the past 10 years for its high potential, excellent reversibility, and easy preparation on an industrial scale [1–3]. Significant research has been done on the formation of LiCoO2 via solid-state reaction and sol–gel methods. These techniques require a high energy and material consumption [4–6]. So an attempt to develop an alternative synthetic route, such as hydrothermal method, electrochemical–hydrothermal method, to produce the desired cathode materials in an economical and environmentally friendly way is in progress [7,8]. Hydrothermal treatment of cobalt metal plates leads directly to the formation of LiCoO2 films [9]. By using electrochemical–hydrothermal approach, crystallized and electrochemically active LiCoO2 films can be effectively prepared in a single synthetic step from cobalt metal plates [10–13]. In this paper, a series of thermodynamic calculation was used to predict the optimum synthesis conditions of LiCoO2 powders and films in lithium hydroxide solution. The optimum synthesis parameter obtained was validated experimentally by cyclic voltammogram method. A possible formation mechanics ∗
Corresponding author. E-mail address: [email protected] (Y. Tao).
0925-8388/$ – see front matter © 2006 Elsevier B.V. All rights reserved. doi:10.1016/j.jallcom.2006.04.074
of hexagonal LiCoO2 in the Co–LiOH–H2 O system was proposed. 2. Theoretical concept of potential–pOH diagram drawing system The details of the construction of potential–pOH diagrams were described in the previous paper [14]. The thermodynamic data of chemical species required for drawing the potential–pOH diagrams in Co–H2 O–LiOH systems are summarized in Table 1 [15–19]. The reaction equations of chemical species with OH− in Co–LiOH systems are shown in Table 2. From the thermodynamic data in Table 1, the changes of Gibbs energy at various temperatures for each reaction equation in Table 2 were calculated by using a personal computer. Base on these values, the relationship between potential and pOH was evaluated and the potential–pOH diagrams in Co–LiOH systems were drawn at 100 and 150 ◦ C. In addition, the solubilities of chemical species were calculated at given temperature. Base on this information, the thermodynamic stabilities of Co were evaluated in 4 M LiOH solution at various temperatures. 3. Results and discussion Figs. 1 and 2 show the potential–pOH diagrams in Co–H2 O system at 100 and 150 ◦ C, respectively. As could be seen from
Y. Tao et al. / Journal of Alloys and Compounds 430 (2007) 222–225
223
Table 1 Thermodynamic properties of components at 298.15 K Species
H2 O H2 O2 H+ OH− Li+ Co Co(OH)2 CoO Co3 O4 Co2 O3 CoOOH CoO2 Co2+ HCoO2 − LiCoO2
State
l g g aq aq aq c c c c c c c aq aq c
G0f,298.15 (kJ mol−1 )
−237.18 0 0 0 −157.29 −293.8 0 −454.4 −214.2 −774.0 −558.1 −349.5 −216.9 −54.4 −347.2 −561.4
0 S298.15 (abs) (J mol−1 K−1 )
69.91 130.58 205.03 −5 10.2 9.2 30.07 79.5 52.97 102.5 43.77 66.98 53.10 −155 62.8 34.91
Fig. 1. Potential–pOH diagram of the Co–H2 O system at 100 ◦ C.
Cp (J mol−1 K−1 ) A
B
C
20.335 27.28 29.96
1.092×10−1
2.033×106 5.0×104 −1.67×105
19.94 82.39 55.1 136.65 81.55 172.24 128.22
68.87
3.264×10−3 4.14×10−3 Cation Anion Cation 1.656×10−2 4.846×10−2 −4.6×10−3 2.753×10−2 3.213×10−2 −2.369×10−1 −1.35×10−1 Cation Acid oxyanion 3.617×10-2
−1.647×106
these figures, the numbers on the lines represent the reaction equations given in Table 2. Dotted lines (a) and (b) showed the equilibrium potentials for hydrogen and oxygen evolution reactions in Table 2, respectively. The small numbers in parenthesis showed the logarithms of the activity of ionic species. The pOH values of 4 M LiOH solution, which was used as an electrolyte, were −0.88 at 100 ◦ C and −0.77 at 150 ◦ C. According to Fig. 1, it could be seen that at 100 ◦ C, the equilibrium potentials for Co/CoO and Co/Co(OH)2 couples were more positive than that of the hydrogen evolution reaction at pOH −0.88. While the equilibrium potential for Co/HCoO2 − couple (Eq. (3)) at an activity of 10−4.7 M was equal to that of hydrogen evolution reaction. Therefore, at 100 ◦ C and pOH −0.88 the first product of Co was HCoO2 − ion, whose saturated concentration was 10−4.7 . At the force of electric field, HCoO2 − ions moved to anode. Then the concentration of HCoO2 − ion was above its saturated concentration at the surrounding of anode. HCoO2 − ion further transformed to Co(OH)2 which was thermodynamically more stable than CoO at that condition. The equilibrium potential for Co(OH)2 /CoOOH couple (Eq. (7)) at 100 ◦ C and pOH −0.88 Table 2 The reactions for Co in LiOH solution
Fig. 2. Potential–pOH diagram of the Co–H2 O system at 150 ◦ C.
−6.0×103 2.1×104 −1.67×105 −2.289×106 −2.122×106 7.11×105 −1.795×106
H2 O + e → OH− + H(g) O2 + 2H2 O + 4e → 4OH− Co + 2OH− → Co(OH)2 + 2e Co + 2OH− → CoO + H2 O + 2e Co + 3OH− → HCoO2 − + H2 O + 2e CoO + H2 O → Co(OH)2 Co(OH)2 + OH− → HCoO2 − + H2 O CoO + OH− → HCoO2 − Co(OH)2 + OH− → CoOOH + H2 O + e CoO + OH− → CoOOH + e HCoO2 − → CoOOH + e CoOOH + OH− → CoO2 + H2 O + e CoOOH + Li+ + OH− → LiCoO2 + H2 O Co + Li+ + 4OH− → LiCoO2 + 2H2 O + 3e
(a) (b) (1) (2) (3) (4) (5) (6) (7) (8) (9) (10) (11) (12)
224
Y. Tao et al. / Journal of Alloys and Compounds 430 (2007) 222–225
Fig. 4. Cyclic voltammogram of Co electrode in a 4 M LiOH solution at a scan rate of 3 mV s−1 . Fig. 3. Schematic diagram for tracing out an area of formation LiCoO2 .
was 0.15 V versus NHE. At 100 ◦ C and pOH −0.88, the reaction proceeded in the following order: Co, HCoO2 − , Co(OH)2 , and CoOOH. At 150 ◦ C and pOH −0.77, as shown in Fig. 2, because the equilibrium potential for H2 O/H couple equaled to the equilibrium potentials for Co/HCoO2 − couple at an activity of 10−3.6 M, metal Co was easily oxidized to HCoO2 − ion, whose saturated concentration was 10−3.6 M. At the force of electric field, HCoO2 − ions moved to anode. Then the concentration of HCoO2 − ion was above its saturated concentration at the surrounding of anode, so HCoO2 − ion would transform to CoO which was thermodynamically more stable than Co(OH)2 at that condition. The equilibrium potential for the CoO/CoOOH couple (Eq. (6)) at 150 ◦ C and pOH −0.77 was 0.10 V versus NHE. At 150 ◦ C and pOH −0.77, the reaction proceeded in the following order: Co, HCoO2 − , CoO, and CoOOH. The potential to form CoOOH phase shifted toward the negative direction with increasing temperature. Without the influence of other ions, Co was finally oxidized to CoO2 . The equilibrium potential for the CoOOH/CoO2 couple (Eq. (10)) is shown in Figs. 1 and 2. At 100 ◦ C and pOH −0.88, the potential was 0.436 V. At 150 ◦ C and pOH −0.77, the potential was 0.396 V. In a 4 M LiOH solution, there was a possibility of a Li+ /H+ exchange to obtain LiCoO2 from CoOOH [20]. Because its Gibbs energy was less than zero, as shown in Fig. 3, Eq. (11) could spontaneously proceed in the left area of line 11. LiCoO2 could be synthesized in the area enclosed by Y-axis, line 10, line 11, and line 7 (100 ◦ C) or line 8 (150 ◦ C), as shown in Fig. 3. In 4 M LiOH solution, LiCoO2 could be prepared in the potential range of 0.1–0.396 V versus NHE at 150 ◦ C and 0.15–0.436 V versus NHE at 100 ◦ C, respectively. The cyclic voltammogram (CV) of LiCoO2 film electrode posited on cobalt substrate in 4 M LiOH is shown in Fig. 4. The CV measures were made in a conventional three-electrode cell using cobalt as work electrode, Ag/AgCl as reference electrode, and platinum slice as counter electrode, and the reaction temperatures were 100 and 150 ◦ C, respectively. The cycling was performed between the voltages −1.0 and 0.5 V at a scanning
rate of 3 mV s−1 . The experiments were performed by means of a Solartron Analytical-1287 Potentiostat. It could be seen that there were two clear anodic peaks in Fig. 4. The first peak was associated with the oxidation of Co/HCoO2 − . When the concentration of HCoO2 − ion was above its saturated concentration, HCoO2 − ion would transform to Co(OH)2 and CoO at 100 and 150 ◦ C, respectively. The second one was corresponds to the reaction of Co(OH)2 /CoOOH or CoO/CoOOH. No cathodic peak is found in Fig. 4, which mean no CoOOH and CoO2 phase existed at the end of the positive scan. Based on the analysis, a conclusion could be drawn that LiCoO2 was prepared by the following H+ /Li+ ion exchange reaction. Li+ + CoOOH → LiCoO2 + H+
(13)
The result of cyclic voltammetry measures in Fig. 4 was concordance with that of thermodynamic analysis. The X-ray diffraction pattern of LiCoO2 film prepared by the hydrothermal electrochemical method is shown in Fig. 5. As was evident, the diffraction patterns of film showed not only the reflections of the film itself, but also those of cobalt substrate. After peaks corresponding to cobalt substrate were excluded, the left diffraction pattern can be indexed to a hexagonal unit ¯ space group which corresponds to LiCoO2 film. cell with R3m
Fig. 5. XRD pattern of LiCoO2 film obtained after hydrothermal electrochemical treatment of Co substrate.
Y. Tao et al. / Journal of Alloys and Compounds 430 (2007) 222–225
4. Conclusion Potential–pOH diagrams in Co–H2 O systems at 100 ◦ C and 150 ◦ C were drawn and equilibrium potential and solubility of cobalt species were calculated. In 4 M LiOH, the reaction proceeded in the following order: Co, HCoO2 − , Co(OH)2 (100 ◦ C) or CoO (150 ◦ C), CoOOH, LiCoO2 . LiCoO2 was synthesized in the potential range of 0.15–0.436 V versus NHE at 100 ◦ C and 0.1–0.396 V versus NHE at 150 ◦ C. The thermodynamic result was in concordance with the cyclic voltammogram analysis. LiCoO2 was synthesized via H+ /Li+ ion-exchange reaction in the CoOOH precursors. Acknowledgement The authors would like to thank for the support from the Natural Science Foundation of China (Grant no. 50401011). References [1] [2] [3] [4]
A. Ermete, F. Maurizio, J. Solid State Chem. 117 (1995) 1. S. Nieto-Ramos, M.S. Tomar, Int. J. Hydrogen Energy 26 (2001) 159. A.G. Ritchie, J. Power Sources 96 (2001) 1. C. Delmas, J.J. Braconnier, P. Hagenmuller, Mater. Res. Bull. 17 (1982) 117.
225
[5] A. Kalyan, T. Masao, K. Shinichi, Mater. Res. Bull. 33 (1998) 1845. [6] E. Antolini, J. Eur. Ceram. Soc. 18 (1998) 1405. [7] J.M. Fernandez-Rodriguez, J. Morales, J.L. Tirado, React. Solids 4 (1987) 163. [8] A.G. Ritchie, J. Power Sources 96 (2001) 1. [9] J.M. Fernadez-Rodriguez, L. Hernan, J. Morales, Mater. Res. Bull. 23 (1988) 899. [10] K.S. Han, S.W. Song, H. Fujita, M. Yoshimura, Solid State Ionics 135 (2000) 273. [11] T. Fujiwara, M. Yoshimura, J. Electroanal. Chem. 559 (2003) 63. [12] K.S. Han, S.W. Song, S.S. Tsurimoto, Solid State Ionics 151 (2002) 11. [13] Y. Tao, Z.H. Chen, B.J. Zhu, W.Z. Huang, Solid State Ionics 161 (2003) 187. [14] Y. Tao, Z.H. Chen, B.J. Zhu, Physica B 362 (2005) 76. [15] C.M. Criss, J.W. Cobble, J. Am. Chem. Soc. 86 (1964) 5390. [16] C.M. Chen, K. Aral, Corrosion-Nace. 38 (1982) 183. [17] M.W. Chase Jr., C.A. Davies, J.R. Downey Jr., D.J. Frurip (Eds.), JANAF Thermochemical Tables, third ed., American Institute of Physics, New York, 1986. [18] O. Knacke, O. Kubaschewski, K. Hesselmann (Eds.), Thermochemical Properties of Inorganic Substance, Springer-Verlag Verlag Stahleisen, 1991. [19] Y.C. Zhang, H. Tagawa, S. Asakura, J. Electrochem. Soc. 144 (1997) 4345. [20] D. Larcher, M.R. Palacin, G.G. Amatucci, J.M. Tarascon, J. Electrochem. Soc. 144 (1997) 408.