Hydrogen Storage of Li2NH Prepared by Reacting Li with NH3

Nov 22, 2005 - at 280 °C. The obtained Li2NH particles reversibly absorb hydrogen and have ... development of a hydrogen economy for transportation.1...
5 downloads 0 Views 98KB Size
Download PDF
182

Ind. Eng. Chem. Res. 2006, 45, 182-186

Hydrogen Storage of Li2NH Prepared by Reacting Li with NH3 Yun Hang Hu* and Eli Ruckenstein Department of Chemical Engineering, State UniVersity of New York at Buffalo, Amherst, New York 14260

In this paper, Li2NH was prepared by reacting Li particles with NH3 at 200 °C, followed by dehydrogenation at 280 °C. The obtained Li2NH particles reversibly absorb hydrogen and have slow kinetics during the first hydrogenation and much faster kinetics during the subsequent rehydrogenations. Furthermore, their hydrogen capacity increases with the cycle number. After 15 cycles, the reversible hydrogen capacity increases to 3.1 wt % from the initial value of about 2 wt %. A larger number of cycles are expected to increase the hydrogen capacity. 1. Introduction A low-cost hydrogen storage technology with high storage capacity and fast kinetics constitutes a critical factor in the development of a hydrogen economy for transportation.1 In recent years, the attention was focused on solid storage materials. Hydrogen can form hydrides with numerous metals and alloys. However, only some of them can provide reversible hydrogenation-dehydrogenation cycles.2 Although the reaction of hydrides (such as LiH and CaH2) with water has been widely employed to generate hydrogen for meteorological balloons,3 this reaction could not be reversed in any efficient manner. The most famous hydrogen storage material, Mg, exhibits reversible hydrogen storage, but the hydrogenation of magnesium to MgH2 occurs only under severe conditions (high temperatures, above 350 °C, and high pressures, above 50 atm) and very slowly and incompletely. In addition, the rate of dehydrogenation of the MgH2 hydride is too low.4 However, the kinetics of hydrogenation has been enhanced through the development of nanocrystalline Mg, Mg alloys, and Mg composites.5-9 High-surfacearea Mg materials (often containing catalysts or other hydriding additives) are now available, which can be rapidly hydrided to >6 wt % hydrogen, even at temperatures close to room temperature.5 Nevertheless, the dehydrogenation of the hydrogenated nanocrystalline Mg materials still has thermodynamic limitations, which make it hard to recover the H2 at temperatures below 100 °C. The low-temperature reversible hydrides, such as LaNi5H6 and TiFeH2, exhibit suitable dehydrogenation kinetics at very low temperatures;10 they have, however, very low hydrogen storage capacities (1.5 wt % for LaNi5H6 and 1.8 wt % for TiFeH2). The complex hydrides of light metals (Li, Na, and Al), such as LiAlH4 (10.5 wt % H2) and NaAlH4 (7.4 wt % H2), have relatively high hydrogen storage capacities, but they are nonreversible.11 Bogdanovic et al.12-14 demonstrated recently that, by doping these light metal hydrides with titanium compounds, the dehydrogenation of NaAlH4/Na3AlH6/Na2LiAlH6 could be facilitated and rendered reversible under moderate conditions. This breakthrough was followed by progress in the development of catalysts for the reversible dehydrogenation of NaAlH4.15-22 However, a reversible hydrogen storage capacity greater than 6 wt % with good cyclability is still a challenge for these catalyst-doped complex hydrides. Borohydride complexes with suitable alkali or alkaline earth metals also constitute * To whom correspondence should be addressed. Phone: 7166452911 ext 2253. Fax: 716-6453822. E-mail: [email protected].

a promising class of compounds for hydrogen storage. Their hydrogen capacity can be as high as 18 wt %. However, their dehydrogenation is very difficult. To reach a 9 wt % hydrogen desorption, a temperature of 600 °C was required.23-25 It was found that SiO2 mixed with LiBH4 (25:75 by weight) lowered the desorption temperature.26 However, our experiments showed that, although SiO2 can reduce the dehydrogenation temperature of LiBH4, the dehydrogenated LiBH4 could reabsorb only about 1 wt % hydrogen at about 250 °C.27 Other promising materials for hydrogen storage are the nanostructured composites. Hydrogen storage in carbon nanotubes attracted tremendous experimental and theoretical interest.28-36 However, their hydrogen storage was not as efficient as it was expected to be.31-34 For this reason, recent investigations have shifted away from them. Although other types of nanotubes, such as boron nitride,37 MoS2,38 and TiS239 nanotubes, could be used for hydrogen storage, their hydrogen capacities were lower than 3 wt % even at high pressures. Recently, Yaghi et al. reported about hydrogen storage in microporous metal-organic frameworks.40 Although this observation led to an interesting direction of research, the current microporous metal-organic frameworks require a temperature as low as 78 K to yield a 4.5 wt % hydrogen storage capacity. At room temperature and 20 atm, their hydrogen storage capacity is 1 wt %. Dafert and Miklauz41 were the first to examine the hydrogenation and dehydrogenation of Li3N in 1910. They reported that the reaction between Li3N and H2 generates Li3NH4,

Li3N + 2H2 ) Li3NH4 that Li3N absorbs 10.4 wt % hydrogen, and that Li3NH4 can be decomposed to release H2.41 Furthermore, Ruff and Goeres42 noted that Li3NH4 is a mixture of LiNH2 and 2LiH. Although Li3N can be a useful storage material, it did not attract attention for nearly a century. Recently, interest in the hydrogen storage of N-based Li materials has been renewed.43-50 Although the high capacity of hydrogen storage (9.0 wt %) of Li3N was reproduced,41,43 a critical issue regarding Li3N is that its reversible hydrogen capacity is only about 5.5 wt %.45,47 This occurs because LiNH2/2LiH, which are the products of Li3N hydrogenation, dehydrogenate in two steps:43,45,49 LiH + LiNH2 ) Li2NH + H2 and LiH + Li2NH ) Li3N + H2. The first step, which provides about 5.5 wt % hydrogen capacity, takes place easily even at temperatures below 200 °C, whereas the second step requires high temperatures (>400 °C). Furthermore, it was found that although a high dehydrogenation temperature can

10.1021/ie050690l CCC: $33.50 © 2006 American Chemical Society Published on Web 11/22/2005

Ind. Eng. Chem. Res., Vol. 45, No. 1, 2006 183

increase the dehydrogenation, the resulting solid has almost completely lost the ability to reabsorb hydrogen.45 The mole ratio LiNH2/LiH of the hydrogenated Li3N is 0.5, and consequently, only half of the LiH releases hydrogen during the first reaction step, at acceptable temperatures. In contrast, if the initial material is lithium imide (Li2NH) instead of Li3N, its hydrogenation products are LiNH2/LiH (1:1). This means that all the LiH in the mixture (LiNH2/LiH) can release hydrogen during the first dehydrogenation step.45-47,49 As a result, the amount of reversible hydrogen can increase above 6 wt %. However, even though doping Li2NH with Mg or Ca may reduce the dehydrogenation temperature of hydrogenated Li2NH,50 the operation temperatures of Li2NH for hydrogen storage are still higher than the Department of Energy (DOE) target. A fine Li2NH powder, which provides a large interface to hydrogen, has good kinetics for hydrogen absorption. However, for on-board hydrogen storage in fuel cell vehicles, a fine powder is not a good choice, because the vibration of the vehicle results in a packing tightening, which reduces the hydrogen flow. In contrast, larger particles can keep a larger space between them, which allows the H2 gas to flow more easily, thus reducing the local H2 pressure and increasing the hydrogen desorption. For this reason, we examine the hydrogen absorption and desorption of relatively large Li2NH particles. In this paper, Li2NH particles were prepared by reacting Li particles (2 mm) with NH3, followed by dehydrogenation. Furthermore, the produced Li2NH particles were investigated regarding their hydrogen absorption and desorption as well as phase changes.

Figure 1. XRD patterns: (a) Li treated with NH3 at 200 °C; (b) Li treated with NH3 at 200 °C, followed by dehydrogenation at 280 °C for 14 h; (c) Li treated with NH3, followed by dehydrogenation at 280 °C for 14 h and hydrogenation at 230 °C for 24 h; (d) Li treated with NH3, followed by dehydrogenation at 280 °C for 14 h and hydrogenation at 230 °C and 14 other cycles of dehydrogenation-rehydrogenation at 230 °C.

2. Experimental Section

3. Results and Discussions

2.1. Materials. Large lithium particles (2 mm) were bought from Aldrich Chemical Co. The dehydrogenated NH3-treated Li was obtained via the reaction between Li and NH3, followed by dehydrogenation (described in the next section). X-ray powder diffraction (XRD) showed that the resulting product was lithium imide, Li2NH (described in the Results and Discussions section). 2.2. Volumetric Test of Hydrogen Storage. A volumetric method, described in a previous paper,47a was employed to accurately determine the hydrogen absorption of Li2NH formed via the in situ NH3 treatment of Li followed by dehydrogenation: An amount of 0.25 g of Li was loaded into a reactor located inside an electrical tubular furnace. First, a NH3 flow (100 mL/min) was passed over the Li at 200 °C for 4 h, followed by dehydrogenation at 280 °C in a vacuum for 14 h. Then, the material was subjected to the first hydrogen absorption at an initial pressure of 7 atm. Before any reabsorption of hydrogen, the sample was subjected to a vacuum (p < 10-5 Torr). The change in the gas-phase pressure of H2 during absorption was followed using a digital pressure gauge. To investigate the rehydrogenation, the hydrogenated sample was exposed to a vacuum to desorb the hydrogen at a selected temperature, followed by reabsorption. It should be noted that the temperature was measured outside the reactor. Therefore, the reaction temperature does not account for the hot spots generated during reaction. The hydrogen capacity is defined as the percentage of hydrogen absorbed based on the weight of the lithium imide (Li2NH). 2.3. X-ray Powder Diffraction (XRD). The X-ray powder diffraction of NH3 treated Li and of its dehydrogenated and hydrogenated samples was determined using a Siemens D500 X-ray diffraction instrument, equipped with a Cu KR source, at 40 kV and 30 mA.

XRD was employed to identify the products after the Li particles were treated with NH3, followed by dehydrogenation. As shown in Figure 1a, the Li specimen treated with NH3 at 200 °C for 4 h exhibits three peaks, located at 2θ ) 30.7, 35.6, and 51.2, which can be assigned to both LiNH2 and Li2NH. Because the other two main peaks at 17.4° and 19.6° belong only to LiNH2, one can conclude that the NH3-treated Li surely contains LiNH2 and may also contain Li2NH. The two peaks at 33.7° and 56.5° can be assigned to Li2O and occur because of the impurities (H2O and O2) present in NH3. After the NH3treated Li was subjected to dehydrogenation at 280 °C, the two main peaks at 17.4° and 19.6°, which can be assigned to LiNH2, disappeared from the XRD patterns, but the other three peaks at 30.7 °, 35.6 °, and 51.2 ° remained (Figure 1b). This means that the entire LiNH2 specimen was transformed into Li2NH during dehydrogenation. Hence, lithium imide (Li2NH) particles are generated by reacting Li particles with NH3, followed by dehydrogenation. The reversible hydrogen absorption by the Li2NH particles was determined by the volumetric method described above. As shown in Figure 2, the first hydrogen absorption of the sample was slow; the hydrogen capacity increased slowly with the reaction time to only about 0.7 wt % (based on Li2NH weight) in 10 min and required 500 min to reach about 2 wt %. To investigate the effect of the number of cycles on the reversible hydrogen capacity, a dehydrogenation was carried out at 230 °C for 14 h before the next rehydrogenation. Figure 2 shows that, during the second cycle, the hydrogen absorption was much faster than that during the first cycle; the hydrogen capacity reached 1.5 wt % in just 10 min, and then, it increased slowly with the reaction time. The absorption kinetics was fast during the subsequent recycles, and the hydrogen absorption capacity increased with increasing recycle number. During the 14th rehydrogenation cycle, the reversible hydrogen absorption

184

Ind. Eng. Chem. Res., Vol. 45, No. 1, 2006

Figure 3. Hydrogen absorption by Li2NH prepared from Li after 15 hydrogenation-dehydrogenation cycles. Absorption conditions: initial hydrogen pressure ) 7 atm; final pressure ≈ 4 atm; absorption temperature ) (a) 230 °C, (b) 200 °C, (c) 180 °C, or (d) 150 °C.

Figure 2. Hydrogen absorption by Li2NH prepared from Li. Absorption conditions: initial hydrogen pressure ) 7 atm; final pressure ≈ 4 atm; absorption temperature ) 230 °C. The sample was subjected to dehydrogenation at 230 °C for 14 h before rehydrogenation.

capacity could reach about 2.8 wt % (based on Li2NH weight) in only 10 min and 3.0 wt % in 60 min, much higher than during the first cycle. However, the absorption of hydrogen was not much different during the 15th rehydrogenation cycle from that during the 14th one. Furthermore, the highest reversible hydrogen capacity of Li2NH was only about 3.1 wt %, about half of its theoretical value of 6.8 wt %. A larger number of cycles is needed to achieve a larger reversible hydrogen capacity. The difference in the kinetic behavior between the first and second cycle most likely occurs because the fresh sample is partially covered with Li2O, which does not absorb hydrogen. Li2O was generated during the treatment with NH3, which contains traces of H2O and O2. The Li2O surface layer was, however, dispersed into the bulk of the sample during the hydrogenation and dehydrogenation. As a result, after the first cycle, the sample had a fresh surface layer free of Li2O, which provided faster initial kinetics than the fresh one. All of the absorption curves after the first cycle have a similar shape. The increasing hydrogen capacity with the increasing recycle number might be explained as follows. The reaction between NH3 and Li is exothermic,51

Li + NH3 ) LiNH2 + 1/2H2

∆H ) -130 kJ/mol (1)

and the high reaction heat can result in the sintering of LiNH2. For this reason, the Li2NH generated via dehydrogenation has a low surface area. Furthermore, the melting point of Li metal (179 °C) is lower than the reaction temperature (200 °C) and this provides another factor that leads to a low surface area of Li2NH. The changes with the cycle number could be a result of the surface area increase because the dehydrogenation and rehydrogenation increase the number of pores. The larger the surface, the higher the amount of Li2NH that reacts with hydrogen. Therefore, the best configuration of Li2NH for onboard hydrogen storage for fuel cell vehicles might be as large particles of porous Li2NH, because large particles reduce the local H2 pressure and the pores increase the hydrogen absorption due to better contact with hydrogen.

Figure 4. Hydrogen desorption of hydrogenated Li2NH prepared from Li. The desorption temperature was 230 °C, and the amount of desorbed hydrogen was calculated based on the weight of Li2NH.

As shown in Figure 3, the absorption kinetics of Li2NH, which was previously subjected to 15 rehydrogenationdehydrogenation cycles at 230 °C, remained fast even at 180 °C; the hydrogen capacity reached about 2.8 wt % in 10 min. However, when the absorption temperature was decreased to 150 °C, the absorption became slow, reaching 0.7 wt % in 10 min and 1.6 wt % after about 240 min. We also examined the desorption of hydrogenated Li2NH prepared from Li, which was previously subjected to 15 rehydrogenation-dehydrogenation cycles at 230 °C. Figure 4 shows that 65% of the total reversible hydrogen capacity of Li2NH could be desorbed in 30 min at 230 °C, nearly 80%, in 60 min, and nearly 90%, in 120 min. Consequently, the prepared material exhibited relatively rapid dehydrogenation kinetics. XRD was employed to identify the phases formed during the various treatments of the specimens. After Li2NH (formed via NH3 treatment followed by dehydrogenation) was hydrogenated, the two main peaks at 17.4° and 19.6° reappeared (Figure 1c), indicating that LiNH2 was formed. In addition, two new peaks at 38.4° and 44.5° were present, which can be assigned to LiH. Therefore, the hydrogen absorption can be expressed as follows:

Li2NH + H2 ) LiNH2 + LiH

(2)

Because this reaction is reversible,43-49 Li2NH can reversibly absorb hydrogen. This explains why the dehydrogenated NH3treated Li can reversibly store hydrogen. Furthermore, Table 1 shows that the intensity ratios of the peaks at 44.5° to one at 19.8° (representing, respectively, LiH and LiNH2) are almost the same (0.558 and 0.570 for the samples subjected to the 1st and the 15th rehydrogenation, respectively; see Figure 1c and d). Because the XRD intensity ratio of the two peaks should be proportional to the molar ratio LiH/LiNH2, the molar ratios LiH/

Ind. Eng. Chem. Res., Vol. 45, No. 1, 2006 185 Table 1. Intensity Ratio of Peaks in XRD XRD intensity ratios of the peaks sample

at 44.5° and 19.9°

at 44.5° and 51.2°

after 1st hydrogenation after 15th hydrogenation

0.558 0.570

0.118 0.312

Table 2. Molar Ratio of Li2O/LiNH2 in the Rehydrogenated Li2NH after 14 Cycles

components

XRD intensity ratio of the peaks at 33.6° (Li2O) and 19.9° (LiNH2)

molar ratioa

Li2O/LiNH2

4.89

0.233

a

The ratio of the XRD response factors, which were obtained from a standard sample measurement, is 20.97 for Li2O (peak at 33.6°) to LiNH2 (peak at 19.9°).

LiNH2 should be the same for both specimens. This is not surprising, because, for each mole of LiNH2 formed through Li2NH hydrogenation, a mole of LiH is also formed (see eq 2) and their ratio is, therefore, equal to unity. The intensity ratios of the peaks at 44.5° and at 51.2° (representing LiH and LiNH2 + Li2NH, respectively) are, however, different for the samples subjected to the 1st and the 15th hydrogenation, being 0.118 for the former and 0.312 for the latter case. As a result, the molar ratio LiH/(LiNH2 + Li2NH) is much smaller for the first hydrogenation than for the 15th one and the ratio Li2NH/LiNH2 is larger for the former than for the latter hydrogenation. Consequently, less Li2NH reacted with hydrogen in the former case than in the latter, in agreement with the above hydrogen absorption results. Furthermore, we also calculated the composition of Li2NH after the 15th hydrogenation from curve d in Figure 1. The highest reversible hydrogen capacity of Li2NH was 3.1 wt % (Figure 2), which is 44.9% of its theoretical value (6.9 wt % hydrogen). This indicates that only 44.9% Li2NH was hydrogenated to LiNH2 and LiH. After rehydrogenation, the material consists of Li2O, LiNH2, LiH, and unhydrogenated Li2NH. According to eq 2, LiNH2 and LiH should have equal molar concentrations. According to curve d in Figure 1, our calculations indicate that the molar ratio of Li2O/LiNH2 is 0.233 (see Table 2). This means that 10.4% Li2NH transformed into Li2O (due to H2O and O2 impurities in NH3). Therefore, after the 15th hydrogenation, 44.9% of Li2NH transformed into LiNH2 and LiH and 10.4% transformed into Li2O. In other words, 44.7% of the original Li2NH remained unchanged. Consequently, the composition of the material after the 15th hydrogenation was 7.2 mol % Li2O, 31.0 mol % LiNH2, 31.0 mol % LiH, and 30.8 mol % Li2NH. 4. Conclusions In conclusion, the reaction between Li particles and NH3 at 200 °C, followed by dehydrogenation at 280 °C, transformed Li into Li2NH, which can reversibly absorb hydrogen. The produced Li2NH has slow kinetics during the first hydrogenation and faster kinetics during the subsequent rehydrogenations. Furthermore, its hydrogen capacity increases with the cycle number. However, even after 14 cycles, there is 30.8 mol % Li2NH that is nonhydrogenated during the 15th hydrogenation. Literature Cited (1) Song, C. S. Fuel processing for low-temperature and high-temperature fuel cells;challenges, and opportunities for sustainable development in the 21st century. Catal. Today 2002, 77, 17.

(2) Sandrock, G. Hydride storage. In Handbook of Fuel Cell; Vielstich, W., Lamm, A., Gasteiger, H. A., Eds.; John Wiley & Sons Ltd: New York, 2003; p 101. (3) Gibb, T. R. P., Jr. J. Chem. Educ. 1948, 25, 577. (4) Genossar, J.; Rudman, P. S. Catalytic role of Mg2Cu in the hydriding and dehydriding of Mg. Z. Phys. Chem. (Muenchen, Ger.) 1979, 116, 215. (5) Huot, J.; Liang, G.; Shultz, R. Mechanically alloyed metal hydride systems. Appl. Phys. A 2001, 72, 187. (6) Orimo, S.; Fujii, H. Materials science of Mg-Ni-based hydrides. Appl. Phys. A 2001, 72, 167. (7) Zaluska, A.; Zaluski, L.; Strom-Olsen, J. O. Structure, catalysis and atomic reactions on the nanoscale: a systematic approach to metal hydrides for hydrogen storage. Appl. Phys. A 2001, 72, 157. (8) Ivanov, E.; Konstanchuk, I.; Bokhonov, B.; Boldyrev, V. Comparative study of first hydriding of Mg-NaF and Mg-NaCl mechanical alloys. J. Alloys Compd. 2003, 360, 256. (9) Tran, N. E.; Imam, M. A.; Feng, C. R. Evaluation of hydrogen storage characteristics of magnesium-misch metal alloys. J. Alloys Compd. 2003, 359, 225. (10) Buchner, H. Energiespeichrung in Metallhydriden; SpringerVerlag: Wein, 1982. (11) Wiswall, R. In Hydrogen in Metals II; Alefeld, G., Vo¨lkl, J., Eds.; Springer-Verlag: New York, 1978; p 201. (12) Bogdanovic, B.; Schwickardi, M. Ti-doped alkali metal aluminium hydrides as potential novel reversible hydrogen storage materials. J. Alloys Compd. 1997, 253, 1. (13) Bogdanovic, B.; Sandrock, G. Catalyzed complex metal hydrides. MRS Bull. 2002, 27, 712. (14) Bogdanovic, B.; Brand, R. A.; Marjanovic, A.; Schwickardi, M.; Tolle, J. Metal-doped sodium aluminium hydrides as potential new hydrogen storage materials. J. Alloys Compd. 2000, 302, 36. (15) Jensen, C. M.; Zidan, R.; Mariels, N.; Hee, A.; Hagen, C. Advanced titanium doped of sodium aluminum hydride: segue to a practical hydrogen storage material? Int. J. Hydrogen Energy 1999, 24, 461. (16) Zidan, R. A.; Takara, S.; Hee, A. G.; Jensen, C. M. Hydrogen cycling behavior of zirconium and titanium-zirconium hydride doped aluminum. J. Alloys Compd. 1999, 285, 119. (17) Jensen, C. M.; Gross, K. J. Development of catalytically enhanced sodium aluminum as a hydrogen-storage material. Appl. Phys. A 2001, 72, 213. (18) Anton, D. L. Hydrogen desorption kinetics in transition metal modified NaAlH4. J. Alloys Compd. 2003, 356, 400. (19) Majzoub, E. H.; Gross, K. J. Titanium-halide catalyst-precursors in sodium aluminum hydrides. J. Alloys Compd. 2003, 356, 363. (20) Balogh, M. P.; Tibbetts, G. G.; Pinkerton, F. E.; Meisner, G. P.; Olk, C. H. Phase changes and hydrogen release during decomposition of sodium alanates. J. Alloys Compd. 2003, 350, 136. (21) Sandrock, G.; Gross, K.; Thomas, G.; Jensen, C.; Meeker, D.; Takara, S. Engineering considerations in the use of catalyzed sodium alanates for hydrogen storage. J. Alloys Compd. 2002, 330, 696. (22) Gross, K. J.; Thomas, G. J.; Jensen, C. M. Catalyzed alanates for hydrogen storage. J. Alloys Compd. 2002, 330-332, 683. (23) M. Seayad, A.; Antonelli, D. M. Recent advances in hydrogen storage in metal-containing inorganic nanostructures and related materials. AdV. Mater. 2004, 16, 765. (24) Fedneva, E. M.; Alpatova, V. L.; Mikheeva, V. I. Russ. J. Inorg. Chem. 1964, 9, 826. (25) Zu¨ttel, A.; Rentsch, S.; Fischer, P.; Wenger, P.; Sudan, P.; Mauron, P.; Emmenegger, C. Hydrogen storage properties of LiBH4. J. Alloys Compd. 2003, 356-357, 515. (26) Zu¨ttel, A.; Wenger, P.; Rentsch, S.; Sudan, P.; Mauron, P.; Emmenegger, C. LiBH4 a new hydrogen storage. J. Power Sources 2003, 118, 1. (27) Hu, Y. H.; Ruckenstein, E. Unpublished experimental results. (28) Dillon, A. C.; Jones, K. M.; Bekkedahl, T. A.; Kiang, C. H.; Bethuune D. S.; Heben, M. J. Storage of hydrogen in single-walled carbon nanotubes. Nature 1997, 386, 377. (29) Ye, Y.; Ahn, C. C.; Witham, C.; Fultz, B.; Liu, J.; Rinzler, A. G.; Colbert, D.; Smith, K. A.; Smalley, R. E. Hydrogen adsorption and cohesive energy of single-walled carbon nanotubes. Appl. Phys. Lett. 1999, 74, 2307. (30) Liu, C.; Fan, Y. Y.; Cheng, H. M.; Dresselhaus, M. S. Hydrogen storage in single-walled carbon nanotubes at room temperature. Science 1999, 286, 1127. (31) Dagani, R. Tempest in a tiny tube. Chem. Eng. News 2002, 80 (2), 25. (32) Yang, R. T. Hydrogen storage by alkali-doped carbon nanotubesrevisited. Carbon 2000, 38, 623.

186

Ind. Eng. Chem. Res., Vol. 45, No. 1, 2006

(33) Hirscher, M.; Becher, M.; Haluska, M.; Dethlaff-Weglikowska, U.; Quintel, A.; Duesberg, G. S. Hydrogen storage in sonicated carbon materials. Appl. Phys. A 2001, 72, 129. (34) Cheng, H.; Pez, G. P.; Cooper, A. C. Electrochemical hydrogen storage in MoS2 nanotubes. J. Am. Chem. Soc. 2001, 123, 5845. (35) Wang, Q.; Johnson, J. K. Computer Simulations of Hydrogen Adsorption on Graphite Nanofibers. J. Phys. Chem. B 1999, 103, 277. (36) Dresselhaus, M. S.; Williams, K. A.; Eklund, P. C. Hydrogen adsorption in carbon materials. MRS Bull. 1999, 24, 45. (37) Ma, R.; Bando, Y.; Zhu, H.; Sato, T.; Xu, C.; Wu, D. Hydrogen Uptake in Boron Nitride Nanotubes at Room Temperature. J. Am. Chem. Soc. 2002, 124, 7672. (38) Chen, J.; Kuriyama, N.; Yuan, H.; Takeshita, H. T.; Sakai, T. Electrochemical Hydrogen Storage in MoS2 Nanotubes. J. Am. Chem. Soc. 2001, 123, 11813. (39) Chen, J.; Li, S. L.; Tao, Z. L.; Shen, Y. T.; Cui, C. X. Titanium Disulfide Nanotubes as Hydrogen-Storage Materials. J. Am. Chem. Soc. 2003, 125, 5284. (40) Rosi, N. L.; Eckert, J.; Eddaoudi, M.; Vodak, D. T.; Kim, J.; O’Keefe, M.; Yaghi, O. M. Hydrogen storage in microporous metal-organic frameworks. Science 2003, 300, 1127. (41) Dafert, F. W.; Miklauz, R. New compounds of nitrogen and hydrogen with lithium. Monatsh. Chem. 1910, 31, 981. (42) Ruff, O.; Goeres, H. Li imide and some compounds of N, H and Li. Berichte 1910, 44, 502. (43) Chen, P.; Xiong, Z.; Luo, J.; Lin, J.; Tan, K. L. Interaction of hydrogen with metal nitride and imides. Nature 2002, 420, 302. (44) Hu, Y. H.; Ruckenstein, E. Ultrafast reaction between LiH and NH3 during H2 storage in Li3N. J. Phys. Chem. A 2003, 107, 9737. (45) Hu, Y. H.; Ruckenstein, E. H2 storage in Li3N. Temperatureprogrammed hydrogenation and dehydrogenation. Ind. Eng. Chem. Res. 2003, 42, 5135.

(46) Ichikawa, T.; Isobe, S.; Hanada, N.; Fujii, H. Lithium nitride for reversible hydrogen storage. J. Alloys Compd. 2004, 365, 271. (47) (a) Hu, Y. H.; Ruckenstein, E. Highly Effective Li2O/Li3N with Ultrafast Kinetics for H2 Storage. Ind. Eng. Chem. Res. 2004, 43, 2464. (b) Hu, Y. H.; Ruckenstein, E. High reversible hydrogen capacity of LiNH2/ Li3N. Ind. Eng. Chem. Res. 2005, 44, 1510. (c) Hu, Y. H.; Yu, N. Y.; Ruckenstein, E. Hydrogen storage in Li3N: Deactivation caused by a high dehydrogenation temperature. Ind. Eng. Chem. Res. 2005, 44, 4304. (d) Hu, Y. H.; Yu, N. Y.; Ruckenstein, E. Effect of the heat pretreatment of Li3N on its H2 storage performance. Ind. Eng. Chem. Res. 2004, 43, 4174. (e) Hu, Y. H.; Yu, N. Y.; Ruckenstein, E. High reversible-hydrogen storage capacity with ultrafast kinetics of LiNH2/Li3N. Ind. Eng. Chem. Res. 2005, 44, 1510. (48) Kojima, Y.; Yasuaki, K. Hydrogen storage of metal nitride by a mechanochemical reaction. Chem. Commun. 2004, 2210. (49) (a) Leng, H. Y.; Ichikawa, T.; Hino, S.; Hanada, N. New metalN-H system composed of Mg(NH2)2 and LiH for hydrogen storage. J. Phys. Chem. B 2004, 108, 8763. (b) Ichikawa, T.; Hanada, N.; Isobe, S.; Leng, H.; Fujii, H. Mechanism of novel reaction from LiNH2 and LiH to Li2NH and H2 as a promising hydrogen storage system. J. Phys. Chem. B 2004, 108, 7887. (50) Luo, W. (LiNH2-MgH2): a viable hydrogen storage system. J. Alloys Compd. 2004, 381, 284. (51) Brauer, G. Hand book of preparatiVe inorganic chemistry; Academic press: New York, London, 1963; pp 463-464.

ReceiVed for reView June 13, 2005 ReVised manuscript receiVed October 28, 2005 Accepted October 31, 2005 IE050690L