J. Phys. Chem. C 2009, 113, 17945–17949
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Dehydrogenation of LiBH4 Destabilized with Various Oxides X. B. Yu,*,† D. M. Grant,‡ and G. S. Walker*,‡ Department of Materials Science, Fudan UniVersity, Shanghai 200433, China, and Fuels and Power Technology Research DiVision, Faculty of Engineering, UniVersity of Nottingham, UniVersity Park, Nottingham NG7 2RD, U.K. ReceiVed: July 10, 2009; ReVised Manuscript ReceiVed: August 22, 2009
Lithium borohydride is a promising candidate for hydrogen storage and fuel cell application due to its high hydrogen content. In the present work, the effect of various oxides on the dehydrogenation of LiBH4 was investigated. The MS-TG results showed that the LiBH4/oxide mixtures were able to dehydrogenate at much lower temperatures; for example, the onset of dehydrogenation was below 100 °C for a LiBH4-Fe2O3 mixture with a mass ratio of 1:2, and the majority of the hydrogen (∼6 wt %) could be released after heating to 200 °C. The order of destabilization effect for LiBH4 for the studied oxides was Fe2O3 > V2O5 > Nb2O5 > TiO2 > SiO2. XRD results revealed that the destabilization of LiBH4 by the oxides resulted from a redox reaction of LiBH4 + MOx f LiMOx + B + 2H2. Investigation of mixed metal oxides showed that a TiO2/SiO2 sample produced an even greater effect, decreasing the temperature of hydrogen release from LiBH4 by more than either TiO2 and SiO2 alone. Introduction The prospect of hydrogen fuel cell for automotive application has stimulated intense interest in developing highly efficient onboard hydrogen storage systems. However, none of the currently available candidate solid state storage materials is adquate for automobile fueling applications due to the limitations of either unfavorable thermodynamics and kinetics or low hydrogen capacity.1,2 Recently, complex hydrides, such as alanates (AlH4-),3-5 amides (NH2-),6,7 and borohydrides (BH4-),8-11 have been investigated intensively because of their high storage capacities compared to conventional metal hydrides. Among these promising candidates, borohydrides have the highest hydrogen capacity. For LiBH4, the decomposition to Li + B + 2H2 yields 18.5 wt % hydrogen. However, the decomposition temperature is too high for practical applications, starting to release hydrogen at around 400 °C, and reversibility has been observed only under high pressure and temperature conditions (350 atm H2 and 600 °C).12 In spite of these drawbacks, substantial improvements in the dehydrogenation kinetics and/ or reversibility have been achieved through particle size reduction,13 doping (i.e., C,13 Al,14 and oxides,15,17), and reactive hydride mixtures such as LiBH4/MgH2,16,17 LiBH4/LiNH2,18 and LiBH4/LiNH2/MgH2.19 Our recent results have demonstrated that the destabilization of LiBH4 by TiO2 resulted from the formation of lithium titanate.20 It predicts the posibility of further improvements in the capacity of LiBH4-based multicomponent systems via the use of oxides of either lighter metals and/or metals with more favorable redox behavior. Here we report a systematic study in dehydrogenation of LiBH4 through reacting with various oxides. Experimental Procedures The source materials of LiBH4 (95%) and high-purity oxides of Fe2O3, V2O5, and Nb2O5 (>99%) were obtained commercially, * Corresponding author. E-mail:
[email protected]; gavin.walker@ nottingham.ac.uk. † Fudan University. ‡ University of Nottingham.
and LiBH4 was used without further purification. To remove adsorbed water, all the oxides were heated to 450 °C for 2 h and then promptly transferred to a glovebox and allowed to cool under an argon atmosphere. TiO2, SiO2, and SiO2-TiO2 mixtures were prepared by an alkoxide sol-gel technique. The oxide products were calcined at 450 °C for 2 h to remove any water. Mixtures of LiBH4/oxides with various weight ratios were hand milled for 5 min or ball milled (EPX 800 Spex shaker ball mill) for up to 4 h under an inert gas (Ar). Hydrogen release measurements were performed by thermogravimeteric analysis (TGA, TA Instruments STD 600) connected to a mass spectrometer (MS, Hiden HPR20) using a heating rate of 10 °C min-1 under 1 atm argon and a carrier flow rate of 200 cm3 min-1. Typical sample quantities were 5-10 mg. Powder XRD (Bruker D8, Cu KR source) measurements were conducted to confirm the crystalline phase. Samples were mounted on a Si single crystal. Samples were mounted in a glovebox, and an amorphous polymer tape was used to cover the surface of the powder to avoid oxidation during the XRD measurement. The surface morphology of the oxides was examined by a scanning electron microscope (SEM, JSM6700F). Results and Discussion Hydrogen evolution and weight loss for LiBH4 with several typical oxides, TiO2, Nb2O5, Fe2O3, and V2O5, with a mass ratio of 1:1, which were hand milled, are shown in Figure 1, compared to LiBH4. For LiBH4, two hydrogen evolutions at 440 and 560 °C, respectively, were observed with a total weight loss of 8.3 wt % (Figure 1b) by heating to 600 °C. In the case of the LiBH4-TiO2 sample, there were three main hydrogen evolutions at 291, 385, and 481 °C, which are higher than that of the ball milled LiBH4-TiO2 sample.20 This may result from the fact that the hand milled sample cannot exhibit an adequate mixing compared to the ball milled sample. For the LiBH4-Nb2O5 sample, two main evolution peaks at 273 and 446 °C were observed, while two other small peaks were located at 340 and 370 °C. The LiBH4-V2O5 sample showed a similar
10.1021/jp906519p CCC: $40.75 2009 American Chemical Society Published on Web 09/18/2009
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Figure 1. MS (a) and TGA (b) results for the evolution of H2 from 1 h ball milled LiBH4 and 5 min hand milled LiBH4/oxide mixtures with a mass ratio of 1:1.
Figure 2. MS (a) and TGA (b) results for the evolution of H2 from LiBH4-TiO2 (mass ratios of 1:4), LiBH4-Nb2O5 (mass ratios of 1:4), LiBH4-Fe2O3 (mass ratios of 1:2), and LiBH4-V2O5 (mass ratios of 1:2) mixtures after 5 min hand milling.
hydrogen evolution process with the LiBH4-Nb2O5 sample, but the temperature of the two main peaks decreased to 253 and 397 °C. In contrast to the other samples, only two dehydrogenation peaks at 250 and 363 °C were observed for the LiBH4-Fe2O3 sample, in which the main hydrogen evolution happened at the first step, indicating its favorable dehydrogenation compared to other LiBH4/oxides samples. Obviously, the onset temperature of dehydrogenation for all the LiBH4/oxides samples is below 300 °C, a significant reduction compared with neat LiBH4, in which dehydrogenation starts at around 400 °C. The TG results in Figure 1b revealed that the weight loss for the investigated LiBH4/oxide mixtures was around 9 wt %, which is comparable with the theoretical capacity of LiBH4 in a 1:1 (w/w) mixture, indicating that all the hydrogen was released from the LiBH4/oxide mixtures below 600 °C. However, only 8.3 wt % hydrogen, corresponding to 44.3% of hydrogen capacity, was released from the neat LiBH4 when it was heated to 600 °C. The above results indicate that addition of these oxides is effective in improving the dehydrogenation of LiBH4. Although the onset of dehydrogenation was decreased, full decomposition of LiBH4 in the above LiBH4/oxides mixtures (with a mass ratio of 1:1) still required a temperature of about 600 °C to completely remove all the hydrogen. The multiple dehydrogenation peaks is an indication of insufficient oxide in the mixtures, thus in an attempt to decompose the LiBH4 in one step, the effect of increasing the oxide content on the dehydrogenation of the LiBH4/oxides mixtures was investigated. The ratios of LiBH4 to oxides were calculated from the proportion of hydrogen released in the first step as shown in Figure 1. As a result, the calculated mass ratios in LiBH4-TiO2, LiBH4-Nb2O5, LiBH4-Fe2O3, and LiBH4-V2O5 for full decomposition in one step were 1:4, 1:4, 1:2, and 1:2, respectively. Figure 2 shows the MS and TG results for the LiBH4/oxide mixtures with the increased oxide addition. Only one main
Figure 3. MS results for the evolution of H2 from the LiBH4-Fe2O3 (mass ratios of 1:1) mixture with various ball milling times.
hydrogen peak was observed for all the four samples. The temperature of the hydrogen release peaks of LiBH4/TiO2 (mass ratio, 1:4), LiBH4/Nb2O5 (mass ratio, 1:4), LiBH4/Fe2O3 (mass ratio, 1:2), and LiBH4/V2O5 (mass ratio, 1:2) samples located at 299, 270, 265, and 250 °C, respectively, are comparable to that of the first step reaction in Figure 1. Especially, the LiBH4-Fe2O3 and LiBH4-V2O5 samples exhibit superior hydrogen desorption properties, releasing the majority of hydrogen below 300 °C (Figure 2b). Furthermore, ball milling was employed to improve the dehydrogenation of LiBH4oxides. As an example, Figure 3 shows the effect of ball milling time on the dehydrogenation of the LiBH4-Fe2O3 mixture with a mass ratio of 1:1. After ball milling for 2 h, the first peak temperature from 251 °C of the 0.5 h milled sample decreased to 200 °C. However, no further improvement was observed for the sample after prolonging the ball milling time to 4 h suggesting that ball milling for 2 h is an optimum processing time for the LiBH4-Fe2O3 sample. Figure 4 presents the MS and TG results for LiBH4-TiO2, LiBH4-Nb2O5, LiBH4-Fe2O3, and LiBH4-V2O5 samples with their optimized ball milling time. Significant improvement was observed for these samples. The dehydrogenation onset temperature was decreased to about
Dehydrogenation of LiBH4 Destabilized with Oxides
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Figure 4. MS (a) and TGA (b) results for the evolution of H2 from ball milled LiBH4-TiO2 (mass ratios of 1:4), LiBH4-Nb2O5 (mass ratios of 1:4), LiBH4-Fe2O3 (mass ratios of 1:2), and LiBH4-V2O5 (mass ratios of 1:2) mixtures.
Figure 5. XRD results for the LiBH4-V2O5 (mass ratios of 1:2) sample after ball milling for 10 and 30 min and the 10 min ball milled sample after dehydrogenation.
Figure 6. XRD results for the 5 min hand milled LiBH4/oxide mixtures before dehydrogenation. (a) LiBH4/TiO2 (mass ratios of 1:4); (b) LiBH4/ Nb2O5 (mass ratios of 1:4); (c) LiBH4/Fe2O3 (mass ratios of 1:2); (d) LiBH4/V2O5 (mass ratios of 1:2).
100 °C, and the terminal temperature for the four samples was decreased to 300, 270, 250, and 220 °C, respectively, with similar weight loss to their hand milled samples. However, it was found that ball milling for prolonged periods may lead to a decrease in capacity due to decomposition during the ball milling process. Figure 5 displays the XRD result for a LiBH4-V2O5 (mass, 1:2) mixture after ball milling for 10 and 30 min and the 10 min ball milled sample after dehydrogenation. Unlike the 10 min ball milled sample, which showed a mixture of V2O5 and LiBH4, the phase structure of the 30 min milled sample had transformed to an HCP structure, similar to the dehydrogenated LiBH4-V2O5 sample. It suggests that the majority of available hydrogen was released during the ball milling. To understand the reaction mechanism, the phase compositions of LiBH4-TiO2, LiBH4-V2O5, LiBH4-Nb2O5, and LiBH4-Fe2O3 samples before and after dehydrogenation were investigated by XRD. No new phases were formed after ball milling for any of the samples, showing mixtures of LiBH4 and corresponding oxides, before dehydrogenation (Figure 6). After full dehydrogenation, all the samples showed no evidence of
Figure 7. XRD results for the 5 min hand milled LiBH4/oxide mixtures after dehydrogenation to 600 °C. (a) LiBH4-Nb2O5 (mass ratios of 1:4); (b) LiBH4-TiO2 (mass ratios of 1:4); (c) LiBH4-V2O5 (mass ratios of 1:2); (d) LiBH4-Fe2O3 (mass ratios of 1:2).
Figure 8. MS results for the evolution of H2 from LiBH4-TiO2 (mass ratios of 1:1), LiBH4-SiO2 (mass ratios of 1:1), and LiBH4-TiO2/ SiO2 (mass ratios of 1:1) mixtures after 5 min hand milling.
Figure 9. XRD for TiO2, SiO2, and TiO2/SiO2 calcined at 450 °C for 2 h.
the starting phases, but the product phase had an FCC diffraction pattern (Figure 7), suggesting similar reactions were occurring for all the LiBH4-MOx samples. Table 1 listed the d-spacings for each sample after dehydrogenation. In the case of LiBH4/ TiO2 sample, the XRD pattern corresponds to LiTiO2. As the LiBH4-TiO2 (mass, 1:4) mixture has a mole ratio of LiBH4 to
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Figure 10. SEM for (a) TiO2 and (b) TiO2/SiO2 calcined at 450 °C for 2 h.
TABLE 1: Comparison of d-Spacings for LiBH4-Nb2O5 (Mass Ratios of 1:4), LiBH4-TiO2 (Mass Ratios of 1:4), LiBH4-V2O5 (Mass Ratios of 1:2), and LiBH4-Fe2O3 (Mass Ratios of 1:2) d-spacings (Å) sample
(111)
(200)
(220)
(311)
(222)
LiBH4-Nb2O5 LiBH4-TiO2 LiBH4-V2O5 LiBH4-Fe2O3
2.4405 2.3942 2.3817 2.3434
2.1155 2.0614 2.0557 2.0262
1.4960 1.4675 1.4542 1.4356
1.2752 1.2506 / /
1.2206 1.1987 / /
TiO2 of approximately 1:1, the main reaction process would appear to be
LiBH4 + TiO2 f LiTiO2 + B + 2H2 as reported previously.20 The reason that no boron was detected in the XRD patterns is most likely due to it forming an amorphous phase as found for the dehydrogenation of neat LiBH4 and some multicomponent systems based on LiBH4.17,20 In the case of LiBH4/V2O5 (mass ratio, 1:2), LiBH4/Nb2O5 (mass ratio, 1:4), and LiBH4/Fe2O3 (mass ratio, 1:2) samples, there are no relevant LixMyOz (MyOz ) V2O5, Nb2O5, and Fe2O3) in the XRD database. However, according to the mole ratio of LiBH4 to oxides in the mixtures, these products after dehydrogenation could be Li4.52V2O5, Li3.76Fe2O3, and Li3.13Nb2O5, suggesting that these oxides could react with LiBH4 to form nonstoichiometric LixMyOz. On the basis of the above analysis, the possible dehydrogenation mechanism in the LiBH4/oxide system might be speculated as xLiBH4 + MyOz f LixMyOz + xB + 2xH2. A similar thermodynamic modeling for the equilibrium reactions of SiO2, Al2O3, and ZrO2 with LiBH4 has been described by Opalka et al.21 However, it is difficult to prove the above mechanism using the limited data presented here. More experimental evidence is needed, especially for determining the exact composition of the reaction products. The above results show that oxides have improved the hydrogen release from LiBH4 through a redox reaction. We are now exploring the effect of different mixtures of oxides, and some encouraging results were achieved. For example, a TiO2/ SiO2 mixture (mass ratio, 2:1) showed an even lower temperature of hydrogen release from LiBH4 than either TiO2 or SiO2 on their own. Figure 8 shows the MS results for LiBH4-TiO2, LiBH4-SiO2, and LiBH4-TiO2/SiO2 samples with a mass ratio of 1:1. In the case of LiBH4-TiO2, there are three desorption peaks located at 291, 385, and 481 °C. The first dehydrogenation peak in LiBH4-SiO2 is located about 310 °C. With further heating, a broader peak at 320-370 °C was observed. Clearly,
full dehydrogenation in the two samples required temperatures above 450 °C. However, in the case of the LiBH4-TiO2/SiO2 sample, two dehydrogenation peaks were found at 264 and 307 °C, and the terminal dehydrogenation temperature was around 350 °C, which was lower than both of the LiBH4-TiO2 and LiBH4-SiO2 samples. Figure 9 presents the XRD for TiO2, SiO2, and TiO2/SiO2 samples. The TiO2/SiO2 exhibited the same amorphous structure as SiO2. SEM showed the mixed oxide had a small particle size, indicating a high surface area, as shown in Figure 10. This increased surface area for the TiO2 in the TiO2/SiO2 mixture might be one of the reasons for the improved effect on the dehydrogenation of LiBH4. For pure LiBH4, its rehydrogenation only can be achieved at extreme conditions of 350 atm H2 and 600 °C.12 Attempts to rehydrogenate these dehydrogenated products at 100 atm H2 and 400 °C were tried but were unsuccessful. Thermodynamic results showed that equilibrium reactions of various oxides with LiBH4 could favorably lead to the formation of very stable Libearing oxide phases. While the formation of these phases could effectively promote LiBH4 dehydrogenation through a destabilization reaction, the reverse reaction in the LiBH4/oxide system is not thermodynamically favored as indicated by the exothermic nature of the dehydrogenation reactions. Conclusions The dehydrogenation of LiBH4 hand milled or ball milled with oxides of Fe2O3, V2O5, Nb2O5, TiO2, SiO2, and TiO2/SiO2 mixtures in varying proportions was investigated. The main conclusions are: (1) All the studied oxides were able to destabilize the decomposition of LiBH4 to lower temperature. The order of destabilization of the LiBH4 for the studied oxides was Fe2O3 > V2O5 > Nb2O5 > TiO2 > SiO2. The 2 h milled LiBH4-Fe2O3 sample (mass ratio of 1:2) exhibited the optimum hydrogen desorption properties, releasing 6 wt % hydrogen below 200 °C. (2) XRD results showed that, after dehydrogenation, all the studied LiBH4/oxide samples showed a HCP structure. It was speculated that the improved dehydrogenation in LiBH4/oxides is due to a redox reaction. (3) The TiO2/SiO2 mixture exhibits an improved effect on destabilizing the decomposition of the LiBH4 phase compared to TiO2 or SiO2 alone. The terminal dehydrogenation temperature from >450 °C of the individual TiO2 and SiO2 decreased to around 350 °C. XRD and SEM results suggest that this was due to the increased surface area of the TiO2 in the TiO2/SiO2 mixture. Acknowledgment. The authors would like to acknowledge the support of the EPSRC. X.B. Yu also acknowledges the
Dehydrogenation of LiBH4 Destabilized with Oxides support of the Shanghai Leading Academic Discipline Project (B113), the Shanghai Rising-Star Program (05QMX1463), and the Hi-Tech Research and Development Program of China (2007AA05Z107). References and Notes (1) Schlapbach, L.; Zu¨ttel, A. Nature 2001, 414, 353. (2) Grochala, W.; Edwards, P. P. Chem. ReV. 2004, 104, 1283. (3) Vegge, T. Phys. Chem. Chem. Phys. 2006, 8, 4853. (4) Bogdanovic, B.; Felderhoff, M.; Pommerin, A.; Schuth, T.; Spielkamp, N. AdV. Mater. 2006, 18, 1198. (5) Balde, C. P.; Hereijgers, B. P. C.; Bitter, J. H.; de Jong, K. P. Angew. Chem., Int. Ed. 2006, 45, 3501. (6) Chen, P.; Xiong, Z.; Lou, J.; Lin, J.; Tan, K. L. Nature 2002, 420, 302. (7) Xiong, Z. T.; Wu, G. T.; Hu, H. J.; Chen, P. AdV. Mater. 2004, 16, 1522. (8) Orimo, S. I.; Nakamori, Y.; Ohba, N.; Miwa, K.; Aoki, M.; Towata, S.; Zuttel, A. Appl. Phys. Lett. 2006, 89, 021920. (9) Lodziana, Z.; Vegge, T. Phys. ReV. Lett. 2004, 93, 145501. (10) Filinchuk, Y. E.; Yvon, K.; Meisner, G. P.; Pinkerton, F. E.; Balogh, M. P. Inorg. Chem. 2006, 45, 1433.
J. Phys. Chem. C, Vol. 113, No. 41, 2009 17949 (11) Barkhordarian, G.; Jensen, T. R.; Doppiu, S.; Bo¨senberg, U.; Borgschulte, A.; Gremaud, R.; Cerenius, Y.; Dornheim, M.; Klassen, T.; Bormann, R. J. Phys. Chem. B 2008, 112, 2743. (12) Orimo, S.; Nakamori, Y.; Kitahara, G.; Miwa, K.; Ohba, N.; Towata, S.; Zuttel, A. J. Alloy Comp. 2005, 404, 427. (13) Yu, X. B.; Wu, Z.; Chen, Q. R.; Li, Z. L.; Weng, B. C.; Huang, T. S. Appl. Phys. Lett. 2007, 90, 034106. (14) Kang, X. D.; Wang, P.; Ma, L. P.; Cheng, H. M. Appl. Phys. A: Mater. Sci. Process. 2007, 89, 963. (15) Zu¨ttel, A.; Wenger, P.; Rentsch, S.; Sudan, P.; Mauron, Ph.; Emmenegger, Ch. J. Power Sources 2003, 118, 1. (16) Vajo, J. J.; Skeith, S. L.; Mertens, F. J. Phys. Chem. B 2005, 109, 3719. (17) Yu, X. B.; Grant, D. M.; Walker, G. S. Chem. Commun. 2006, 37, 3906. (18) Pinkerton, F. E.; Meisner, G. P.; Meyer, M. S.; Balogh, M. P.; Kundrat, M. D. J. Phys. Chem. B 2005, 109, 6. (19) Yang, J.; Sudik, A.; Siegel, D. J.; Halliday, D.; Drews, A.; Carter, R. O., III; Wolverton, C.; Lewis, G. J.; Sachtler, J. W. A.; Low, J. J.; Faheem, S. A.; Lesch, D. A.; Ozolins, V. Angew. Chem., Int. Ed. 2008, 47, 882. (20) Yu, X. B.; Grant, D. M.; Walker, G. S. J. Phys. Chem. C 2008, 112, 11059. (21) Opalka, S. M.; Tang, X.; Laube, B. L.; Vanderspurt, T. H. Nanotechnology 2009, 20, 204024.
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