J. Phys. Chem. C 2008, 112, 11059–11062
11059
Low-Temperature Dehydrogenation of LiBH4 through Destabilization with TiO2 X. B. Yu,*,†,‡ D. M. Grant,§ and G. S. Walker*,§ Department of Materials Science, Fudan UniVersity, Shanghai 200433, China, Energy Science and Technology Laboratory, Shanghai Institute of Microsystem and Information Technology, Chinese Academy of Sciences, Shanghai 200050, China, and School of Mechanical, Materials and Manufacturing Engineering, UniVersity of Nottingham, UniVersity Park, Nottingham NG7 2RD, U.K. ReceiVed: January 22, 2008; ReVised Manuscript ReceiVed: April 9, 2008
The hydrogen desorption properties of lithium borohydride, LiBH4, ball milled with titania, TiO2, in varying proportions have been investigated. The multicomponent LiBH4-TiO2 materials have been found to dehydrogenate at much lower temperatures. For example, the onset of dehydrogenation was 150 °C for a LiBH4-TiO2 mixture with a mass ratio of 1:4 and the majority of the hydrogen could be released below 220 °C. XRD and in situ neutron diffraction results revealed that the destabilization of LiBH4 by the oxide resulted from the formation of lithium titanate. Introduction The safe and compact storage of hydrogen is one of the most challenging issues in complementing fuel cell technology for automotive use.1 The U.S. Department of Energy (DOE) has set a target of 6 wt % by 2010 for hydrogen storage capacity for on-board applications, at moderate temperatures and pressures.2 However, none of the currently available candidate storage materials are able to meet this DOE target. Recently, a number of new hydrogen storage systems have been studied based on destabilized complex hydrides, such as alanates (AlH4-), amides (NH2-), and borohydrides (BH4-), because of their high storage capacities.3–9 Among these complex hydrides lithium borohydride, LiBH4, has promising prospects for onboard applications as the full dehydrogenation of LiBH4 yields 18.3 wt % hydrogen,. However, to release all of the hydrogen requires a reaction temperature of up to 900 °C.10 Previous results have demonstrated that the onset hydrogen release temperature of LiBH4 can be reduced to below 400 °C by ball milling with MgH211,12 and combination with LiNH2 formed a new borohydride-amide salt with an onset of decomposition at 250 °C.13 Zu¨ttel et al.10 reported that hydrogen released from a LiBH4-SiO2 mixture (25 wt % LiBH4 and 75 wt % SiO2) started at 200 °C, but the actual hydrogen storage capacity for the mixture was about 2.3 wt %. Here we report a dramatic reduction in the dehydrogenation temperature of LiBH4 through destabilization by titania, TiO2. Experimental Procedures The source materials, namely, LiBH4 (95%, Fisher), 7Li11BD4 (99%, Katchem Ltd., Czech Republic), and TiO2 (99%, Fisher), were obtained commercially. LiBH4 and the isotopically enriched LiBD4 were used without further purification. To remove adsorbed water from TiO2, it was heated to 450 °C for 2 h and then promptly transferred to a glovebox and allowed to cool under an argon atmosphere. The dried TiO2 was shown to be anatase by powder X-ray diffraction (XRD). Mixtures of * Corresponding authors. E-mail:
[email protected] (X.B.Y.);
[email protected] (G.S.W.). † Fudan University. ‡ Chinese Academy of Sciences. § University of Nottingham.
LiBH4-TiO2 with weight ratios of 4:1, 1:1, and 1:4 and LiBD4-TiO2 with a weight ratio of 1:4 were ball milled (EPX 800 Spex shaker ball mill) for different times 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 of 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. In situ neutron diffraction (ND) measurements were undertaken using the GEM diffractometer at ISIS (the pulsed neutron facility at the Rutherford Appleton Laboratory, U.K.). The LiBD4-TiO2 sample (mass ratio of 1:4) was loaded in a quartz reactor tube and kept under vacuum. The sample was heated at a rate of 2 °C min-1, while diffraction data were continuously collected (each pattern was an average of data collected over 5 min). Results and Discussion Hydrogen evolution and weight loss for LiBH4 and a range of LiBH4-TiO2 samples, ball milled for 1.5 h, are shown in Figure 1. Two hydrogen evolutions at 440 and 560 °C, respectively, were observed for LiBH4 (S1) with a total weight loss of 8.3 wt % (Figure 1b), indicating that less than half the hydrogen was released from the LiBH4 by 600 °C, in good agreement with the data in the literature.10,13 In the case of the 4:1 LiBH4-TiO2 sample there were three main hydrogen evolutions at 325, 405, and 525 °C. The total weight loss of the S2 sample was 11.9 wt % (Figure 1b). For the 1:1 LiBH4-TiO2 sample, it was found that the onset of dehydrogenation decreased to 150 °C with the hydrogen being evolved in three steps, having peak temperatures of 245, 390, and 465 °C. The total weight loss was 9.0 wt %. In contrast to the other samples, only one dehydrogenation peak was observed for the 1:4 LiBH4-TiO2 sample at 245 °C, but this peak did have a shoulder at 180 °C. The total weight loss was 3.65 wt %. Figure
10.1021/jp800602d CCC: $40.75 2008 American Chemical Society Published on Web 07/01/2008
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Figure 1. MS (a) and TGA (b) results for the evolution of H2 from LiBH4 (S1) and LiBH4-TiO2 mixtures with mass ratios of 4:1 (S2), 1:1 (S3), and 1:4 (S4) after 1.5 h ball milling. Heating rate was 10 °C min-1.
Figure 2. MS results for gas evolution from the LiBH4-TiO2 (mass ratio, 1:4) mixture, heated at a rate of 10 °C min-1. The rough peak characteristics compared with S4 sample in Figure 1 is due to a slower sampling rate as a full mass spectrum (m/z 1-50 amu) was recorded for each scan.
2 shows the MS results for the gas evolution from a hand milled LiBH4-TiO2 (mass ratio, 1:4) mixture. Neither BH3 nor B2H6 was detected during the decomposition of the borohydride. This is encouraging as these species are poisons for fuel cells, although they are normally associated with the decomposition of transition metal borohydrides. There was a very small increase in H2O intensity around 260-280 °C (enlarged region in Figure 2), but there was also a corresponding decrease in the O2 intensity. This suggests that the very small amount of water detected was merely from a gas phase reaction between evolved hydrogen and a small oxygen impurity in the argon carrier gas. Note that the peak temperature in Figure 2 is 50 °C higher than that of S4 in Figure 1, suggesting that the more energetic ball milling (in contrast to just hand milling) also contributed to decreasing the decomposition temperature of LiBH4-TiO2, presumably a result of more intimate mixing of the two phases. The fact that no H2O, BH3, and B2H6 were detected by MS for the LiBH4-TiO2 samples suggests that the total weight loss can be attributed solely to hydrogen release. Hence, 81.3% of the available hydrogen was released from the 4:1 LiBH4-TiO2 sample by 600 °C, and the weight loss from 1:1 and 1:4 samples correspond to 98.4% and 99.7% of the available hydrogen, respectively, a significant improvement compared to ball milled LiBH4 with no TiO2, which released only half of its available hydrogen. These results show that increasing the TiO2 content was effective in improving the hydrogen release from LiBH4 at low temperatures. TGA results revealed that ball milling for 1.5 h did not result in any significant loss of capacity. 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. This is illustrated in Figure 3, which displays the TGA-MS result for a 12 h ball milled LiBH4/TiO2 (mass,
Figure 3. TGA-MS results for evolution of H2 from LiBH4-TiO2 mixtures with mass ratio of 1:4 after 12 h ball milling. Heating rate was 10 °C min-1.
Figure 4. Dehydrogenation kinetic properties for 1.5 h milled LiBH4-TiO2 (mass ratio 1:4) sample at 180 and 220 °C.
1:4) mixture. Note that although the onset of hydrogen release decreased to around 100 °C, the total weight loss for this materials is about 1.7 wt % compared with 3.7 wt % for the sample milled for 1.5 h, suggesting that more than half the available hydrogen was released during the ball milling. The onset of dehydrogenation for the 1:4 sample ball milled for 1.5 h was at 150 °C, showing the potential for the system to operate at temperatures below 200 °C. Figure 4 shows the percentage dehydrogenation with time for the 1:4 sample at 180 and 220 °C. This sample exhibited relatively fast kinetic properties at 180 °C, releasing 68% of the available hydrogen in 70 min. Furthermore, the desorption capacity increased to 90% of the available hydrogen at 220 °C in the same time period. The phases present in the starting materials and end products were investigated by XRD; see Figures 5 and 6. The as-prepared 1:4 sample gave only weak diffuse TiO2 anatase reflections. No LiBH4 phase was identified because of the line broadening
Low-Temperature Dehydrogenation of LiBH4
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Figure 5. XRD results for 1.5 h milled LiBH4-TiO2 before heating. (a) Mass ratio 1:1; (b) mass ratio 1:4.
Figure 7. In situ neutron diffraction patterns for 1.5 h milled LiBD4-TiO2 (mass ratio 1:4) sample. The heating rate was 2 °C min-1.
Figure 6. XRD results for 1.5 h milled LiBH4-TiO2 after heating to 600 °C. (a) Mass ratio 1:1; (b) mass ratio 1:4.
expected after ball milling (Figure 5b). After heating to 600 °C, the XRD pattern for the end products corresponded to LiTiO2 (Figure 6b). This 1:4 wt % sample had an approximately equimolar ratio of LiBH4 and TiO2 (the mole ratio was 1:1.1); hence the reaction for this mixture was
LiBH4 + TiO2 f LiTiO2 + B + 2H2
(1)
The fact that no boron-containing phases were detected by XRD suggests that the element was in an amorphous state. The XRD patterns for the as-prepared 1:1 LiBH4-TiO2 sample was found to consist of weak diffuse patterns for both the TiO2 and LiBH4 phases. XRD was able to detect the latter phase as there was a larger proportion of the borohydride in this sample compared to the 1:4 sample (Figure 5a). After the sample was heated to 600 °C, the product was found to contain TiB2 and Li2O (Figure 6a), suggesting an alternative reaction mechanism. The TGA-MS results in Figure 1 show three dehydrogenation peaks; the first at 245 °C most probably results from the formation of a lithium titanate phase as for the 1:4 sample. As no lithium titanate is present in the end products after heating to 600 °C, the intermediate phase must react with the remaining LiBH4. Considering that the LiBH4 to TiO2 mole ratio for this sample was ca. 4:1 (the actual mole ratio was 3.7:1), the predominant reaction by 600 °C for the mixture was
TiO2 + 4LiBH4 f 2Li2O + TiB2 + 2B + 8H2
(2)
To investigate the reaction path for the rapid low-temperature dehydrogenation of the 1:4 mass ratio sample, in situ ND was used with LiBD4-TiO2 to identify any intermediate phases. Figure 7 gives the ND patterns for a sample heated at 2 °C min-1 under dynamic vacuum. The as-prepared sample gave ND patterns of anatase and the low-temperature (orthorhombic) form of LiBD4. At 200 °C, the LiBD4 phase had transformed into the high-temperature (hexagonal) phase; this phase trans-
Figure 8. XRD for sol-gel prepared TiO2 calcined at different temperatures for 2 h.
formation occurs at 108 °C.13 The LiBD4 peak started to decrease in intensity as the temperature increased from 200 °C. Concomitantly, the TiO2 phase transformed gradually to LiTiO2 and the reaction was complete by 260 °C, slightly lower than the end temperature for the TGA-MS results in Figure 1, which is most likely a result of the lower heating rate for the neutron experiment. No other intermediate phases were identified; hence the destabilized reaction was a result of the direct formation of LiTiO2. The formation of LiTiO2 raises an interesting comparison with Li insertion in the TiO2 anode for rechargeable Li ion batteries. The application of TiO2 as the electrode material in Li ion batteries is related to the large number of Li ions that can be reversibly intercalated into the TiO2 structure.14 Anatase is generally considered to be the most electroactive Li insertion host as far as TiO2 is concerned, while Li insertion into rutile is usually reported to be negligible.15 Lithium insertion results in a phase transition from the original tetragonal (space group I41/amd) anatase structure toward the orthorhombic LixTiO2 phase, which we refer to as the lithium titanate phase (space group Imma).16 During intercalation of microsized crystal particles, the two-phase equilibrium between these two phases is maintained by a continuous Li ion exchange.17 The Li ion storage capacity and the rate with which they can be inserted and extracted depend strongly on the dimensions of the TiO2 crystallites. This suggests that the structure of TiO2 may also have an effect on the hydrogen release of LiBH4. Figure 8 shows the XRD for sol-gel prepared TiO2 calcined for 2 h at 200, 450, and 600 °C, respectively. The sample calcined at 200 °C shows a single anatase phase. With the calcining temperature increase to 450 °C, a small amount of rutile TiO2 phase
11062 J. Phys. Chem. C, Vol. 112, No. 29, 2008
Yu et al. ponent systems. Of particular importance is the fact that the 1:4 sample can liberate the majority of its hydrogen below 200 °C at acceptably fast evolution rates. Further improvements in the capacity of LiBH4-based multicomponent systems are predicted via the use of oxides of either lighter metals and/or metals with more favorable redox behavior. Conclusions
Figure 9. MS results for the evolution of H2 from LiBH4-TiO2 mixtures with a mass ratio of 1:1 after 5 min hand milling. Heating rate was 10 °C min-1.
appeared. For the sample calcined at 600 °C a mixture of anatase and rutile TiO2 was observed. The MS data in Figure 9 show the hydrogen evolution from the LiBH4/TiO2 samples (mass ratio of 1:1), which were hand milled for about 5 min. Two main hydrogen release peaks at 280 and 356 °C were observed for the sample with TiO2 calcined at 200 °C. In the case of the sample with TiO2 calcined at 450 °C, the first hydrogen peak is also at 280 °C, but the second peak is at a higher temperature of 377 °C. Furthermore, for the sample with TiO2 calcined at 600 °C, both hydrogen evolution peaks are broader and shifted to higher temperature, 297 and 392 °C, respectively. This indicates that anatase TiO2 forms LiTiO2 with LiBH4 more readily than TiO2 with a rutile phase, exhibiting a similar trend to that for the Li insertion in TiO2. However, unlike Li insertion in TiO2, in which Li ion can be reversibly intercalated into the anatase structure, preliminary calorimetric studies indicate that the dehydriding reaction is exothermic, implying that the reverse reaction is not thermodynamically favored. Attempts to directly rehydride the decomposition product by heating under 400 °C and 10 MPa H2 pressure have not achieved significant hydrogenation; hence an alternative hydrogenation route would be required to re-form the LiBH4. These results show that TiO2 has improved the hydrogen release from LiBH4 through a redox reaction, with the titanium cation reducing from Ti4+ to Ti3+ in reaction 1, while in reaction 2 this is an intermediate step to finally forming TiB2. The titanium has been reduced again in forming TiB2; however, the precise oxidation state is under debate as modeling work shows that in addition to metallic Ti-Ti bonding there should also be some charge transfer between the Ti and B yielding some covalent and ionic Ti-B bond character.18 The identified reactions suggest that other metal oxides with variable oxidation states should also be able to play a similar role, being reduced by the LiBH4. This strategy has led to a significant reduction in the dehydrogenation temperature for LiBH4-based multicom-
In conclusion, ball milling LiBH4 with TiO2 reduced the onset temperature for dehydrogenation to 150 °C. Destabilization of the borohydride was via a redox reaction with the TiO2 forming LiTiO2 and liberated all the available hydrogen from LiBH4. Furthermore, lithium titanate can also accelerate the decomposition of LiBH4, forming Li2O and TiB2, resulting in a total hydrogen release of 9.0 wt %. These new reaction mechanisms show the possibility of expanding the rationale for destabilizing LiBH4 to other metal oxides. Acknowledgment. The authors would like to acknowledge the support of the EPSRC and ISIS (RB520467), who funded this work, and to Dr. W Kockelmann for his help with the ND experiments. X.B.Y. also acknowledges the 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.; Zuttle, A. Nature 2001, 414, 353–358. (2) Li, Y.; Yang, R. T. J. Am. Chem. Soc. 2006, 128, 8163. (3) Orimo, S.; Nakamori, Y.; Zuttel, A. Mater. Sci. Eng., B 2004, 108, 51. (4) Xiong, Z.; Wu, G.; Hu, J.; Chen, P. AdV. Mater. 2004, 16, 1522. (5) Wang, P.; Jensen, C. M. J. Phys. Chem. B 2004, 108, 15827. (6) Bogdanovic´, B.; Felderhoff, M.; Kaskel, S.; Pommerin, A.; Schlichte, K.; Schu¨th, F. AdV. Mater. 2003, 15, 1012. (7) Jensen, C. M.; Gross, K. J. Appl. Phys. A 2001, 72, 213. (8) Seayad, A. M.; Antonelli, D. M. AdV. Mater. 2004, 16, 765. (9) Chen, P.; Xiong, Z.; Lou, J.; Lin, J.; Tan, K. L. Nature 2002, 420, 302. (10) Zuttel, A.; Wenger, P.; Rentsch, S.; Sudan, P.; Mauron, Ph.; Emmenegger, Ch. J. Power Sources 2003, 118, 1. (11) Vajo, J. J.; Skeith, S. L.; Mertens, F. J. Phys. Chem. B 2005, 109, 3719. (12) Yu, X. B.; Grant, D. M.; Walker, G. S. Chem. Commun. 2006, 37, 3906. (13) Pinkerton, F. E.; Meisner, G. P.; Meyer, M. S.; Balogh, M. P.; Kundrat, M. D. J. Phys. Chem. B 2005, 109, 6. (14) Bechinger, C.; Ferrer, S.; Zaban, A.; Sprague, J.; Gregg, B. A. Nature 1996, 383, 608. (15) Hu, Y. S.; Kienle, L.; Guo, Y. G.; Maier, J. AdV. Mater. 2004, 16, 765. Hu, Y. S.; Kienle, L.; Guo, Y. G.; Maier, J. AdV. Mater. 2006, 18, 1421. (16) Cava, R. J.; Murphy, D. W.; Zahurak, S.; Santoro, A.; Roth, R. S. J. Solid State Chem. 1984, 53, 64. (17) Wagemaker, M.; Kentgens, A. P. M.; Mulder, F. M. Nature 2002, 418, 397. (18) Mouffok, B.; Feraoun, H.; Aourag, H. Mater. Lett. 2006, 12, 1433.
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