12881
2007, 111, 12881-12885 Published on Web 08/14/2007
Phase Boundaries and Reversibility of LiBH4/MgH2 Hydrogen Storage Material Frederick E. Pinkerton,*,† Martin S. Meyer,† Gregory P. Meisner,† Michael P. Balogh,‡ and John J. Vajo§ Materials and Processes Laboratory and Chemical and EnVironmental Sciences Laboratory, General Motors Research and DeVelopment Center, 30500 Mound Road, Warren, Michigan 48090-9055, and HRL Laboratories, LLC, 3011 Malibu Canyon Road, Malibu, California 90265-4797 ReceiVed: June 3, 2007; In Final Form: July 19, 2007
The coupled system LiBH4 + 1/2MgH2 T LiH + 1/2MgB2 + 2H2 demonstrates improved hydrogen cycling thermodynamics compared to either LiBH4 or MgH2 alone; in effect, formation of MgB2 “destabilizes” the decomposition of LiBH4. Here we establish the thermodynamically and kinetically stable region of the H2 pressure-temperature phase diagram for reversible hydrogen storage in TiCl3-catalyzed LiBH4 + 1/2MgH2. Although MgB2 formation was thermodynamically favored at elevated temperature, it was kinetically more favorable for MgH2 and LiBH4 to decompose independently in a two-step dehydrogenation starting with MgH2 T Mg + H2. At high temperature and low H2 pressure, direct LiBH4 decomposition is both thermodynamically allowed and kinetically favored; thus, the second dehydrogenation step from LiBH4 produced LiH and amorphous boron along with the Mg metal from the first step. From this state, recombination of LiH with amorphous boron had very poor kinetics and the system did not fully rehydrogenate. Applying an H2 gas overpressure of at least 3 bar during dehydrogenation, however, suppressed direct decomposition of LiBH4 and reaction of Mg with LiBH4 produced LiH and MgB2, which was fully reversible.
High-performance on-board hydrogen storage systems are a critical element in the development of practical fuel-cellpowered vehicles. Hydrogen storage in solid hydrides is an attractive option because many such materials offer volumetric hydrogen densities substantially greater than that of compressed gas and comparable to or exceeding that of liquid hydrogen,1 without very-high-pressure containment vessels or cryogenic tanks. Conventional transition-metal hydrides, however, cannot meet the gravimetric density requirements,2 stimulating considerable recent research on light-element complex hydrides such as alanates (AlH4-),3-9 amides (NH2-),10-17 borohydrides (BH4-),18,19 and mixed anion quaternary hydrides such as Li4BN3H10.20-26 Complex hydrides, however, frequently have strongly covalently or ionically bound hydrogen and thus are too thermodynamically stable to operate at practical temperatures and H2 pressures (1-5 bar); the equilibrium temperature T(1 bar), defined as the minimum thermodynamically allowed temperature for dehydrogenation into 1 bar of H2 gas, is undesirably high. The LiBH4/MgH2 system reported recently by Vajo et al.27 is one example of “destabilized” complex hydrides; by reacting the complex hydride with another compound (in this case a binary hydride) to form a mixed compound in the dehydrogenated state, high hydrogen capacity can be retained in a system with reduced thermodynamic stability.28-30 Although in detail * Corresponding author. E-mail:
[email protected]. † Materials and Processes Laboratory, General Motors Research and Development Center. ‡ Chemical and Environmental Sciences Laboratory, General Motors Research and Development Center. § HRL Laboratories.
10.1021/jp0742867 CCC: $37.00
hydrogen loss may occur in more than one step,18,31 the overall thermal decomposition of pure LiBH4 is commonly approximated by the reaction
LiBH4 f LiH + B + 3/2H2
(1)
theoretically yielding 13.9 wt % H2. Unfortunately, it has been shown experimentally18 that LiBH4 is too stable for practical hydrogen storage; temperatures exceeding 400 °C are required for decomposition even in the absence of H2, and reaction 1 is very difficult to reverse. Similarly, MgH2 releases 7.7 wt % hydrogen according to the reaction
MgH2 T Mg + H2
(2)
with ∆H ) 74.5 kJ/mol H2 and T(1 bar) ) 279 °C.32 When combined, however, a new reaction pathway is created through the formation of MgB2:
LiBH4 + 1/2MgH2 f LiBH4 + 1/2Mg + 1/2H2 f LiH + 1/2MgB2 + 2H2 (3a) LiH + 1/2MgB2 + 2H2 f LiBH4 + 1/2MgH2
(3b)
Reaction 3a reflects the observation that dehydrogenation occurs in two distinct steps, whereas rehydrogenation, reaction 3b, occurs in a single step.27 The measured enthalpy change is ∆H ) 40.5 kJ/mol H2 [T(1 bar) ) 225 °C], that is, at a given temperature the equilibrium H2 pressure is higher than that for either LiBH4 or MgH2 alone.27 MgB2 is thermodynamically © 2007 American Chemical Society
12882 J. Phys. Chem. C, Vol. 111, No. 35, 2007
Figure 1. Volumetric measurements of the first and second hydrogenation for LiH + 1/2MgB2 + 0.03TiCl3, Samples A and B. Hydrogen uptake has been normalized to the original LiH + 1/2MgB2 content to compensate for the added weight of TiCl3. Hydrogenation was performed in a 60 cm3 sample cell initially at 100 bar H2.
more stable than either Mg or B, effectively “destabilizing” the LiBH4 + 1/2MgH2 system. The theoretical capacity of reaction 3 is 11.7 wt % hydrogen, and reversible hydrogen storage exceeding 9 wt % has been demonstrated in this system when catalyzed by adding 2-3 mol % TiCl3.27 Indeed, Barkhordarian et al. have shown that combinations of MgB2 with binary hydrides can be hydrogenated to form LiBH4, NaBH4, and Ca(BH4)2.33 Even when catalyzed, however, reaction 3 has slow kinetics, and several hours at elevated temperature are generally required to drive the reaction to completion.27 Vajo et al.27 observed that dehydrogenation according to reaction 3a required an H2 gas overpressure of several bar; dehydrogenation under vacuum led to the formation of Mg instead of MgB2 and resulted in loss of capacity during rehydrogenation. Recent calorimetric measurements have also reported the formation of Mg rather than MgB2 unless H2 gas was present.34 Here we examine the impact of temperature and H2 pressure during dehydrogenation of LiBH4 + 1/2MgH2 on the decomposition pathway and products. By examining the dehydrogenation products and hydrogen cycling behavior, we establish the thermodynamically and kinetically stable region of the H2 pressure-temperature phase diagram for reversibility of reaction 3. Formation of MgB2 during dehydrogenation, and hence fully reversible hydrogen storage, required application of sufficient H2 gas pressure to suppress direct decomposition of LiBH4 via reaction 1. TiCl3-catalyzed LiBH4 + 1/2MgH2 required an H2 gas pressure equal to or exceeding about 3 bar for dehydrogenation at 425 °C. Samples for volumetric measurements were prepared in the dehydrogenated state by ball-milling TiCl3-doped mixtures of LiH + 1/2MgH2. Figure 1 shows the first hydrogenation at 100 bar for two samples distinguished only by sample size, designated A (1.018 g) and B (0.409 g). As expected, the initial hydrogenation is essentially independent of sample mass, as shown by the curves labeled “A-first” (solid line) and “B-first” (dashed line) in Figure 1. The hydrogen uptake values displayed in Figure 1 have been normalized to the LiH and MgB2 content of the sample, thereby compensating for the added TiCl3 content. After hydrogenation was complete, each sample was dehydrogenated starting in vacuum by heating to 450 °C at 2 °C/ min, as shown in Figure 2. Sample A (1.018 g) was dehydrogenated using a small sample-cell volume (220 cm3), whereas Sample B (0.409 g) was dehydrogenated using a large sample-
Letters
Figure 2. Volumetric measurements of the dehydrogenation following the first hydrogenation of LiH + 1/2MgB2 + 0.03TiCl3, samples A-C, and LiH + 1/2MgB2 + 0.03TiH2, sample D. Samples A and D were Measured using large sample sizes in a small sample volume starting from vacuum. Samples B and C were measured using small sample sizes in a large sample volume, Sample B starting from vacuum and Sample C starting with a 4.2 bar H2 gas overpressure. Hydrogen desorption has been normalized to the original LiH + 1/2MgB2 content to compensate for the added weight of TiCl3 or TiH2.
cell volume (725 cm3). The temperature dependence of dehydrogenation for Sample A (solid line) and Sample B (dashed line) were again nearly identical and yielded similar hydrogen capacity. It is important to note, however, that because of the different combinations of sample size and cell volume the absolute pressure increase in the sample cell was 4.6 bar for Sample A, but nearly an order of magnitude smaller at only 0.55 bar for Sample B. A major difference was observed when the resulting dehydrogenated materials were again rehydrogenated under 100 bar H2. The second hydrogenation of Sample A, labeled “A-second” (dash-dot curve) in Figure 1, followed the pattern reported previously by Vajo et al.27 After being cycled once, the second hydrogenation occurred relatively rapidly at 300 °C, and showed essentially unchanged hydrogen capacity. Sample B, however, began to hydrogenate at a temperature consistent with hydrogen absorption by Mg metal, and saturated at only about 3.3 wt % hydrogen uptake after long exposure to high-pressure H2 at 350 °C, as shown by “Bsecond” (dash-dot-dot curve) in Figure 1. This is about the uptake predicted if the only effect was formation of MgH2 according to reaction 2. Underlying the different rehydrogenation behaviors of Samples A and B are profound differences in the dehydrogenation products evident in the X-ray diffraction (XRD) patterns shown in Figure 3. Sample A after dehydrogenation reformed LiH and MgB2 according to reaction 3. The XRD pattern of dehydrogenated Sample B, however, shows a combination of Mg metal and LiH, with traces of residual MgH2. Samples B and A clearly followed different reaction pathways, in accord with previous observations.27,34 No diffraction lines are observed corresponding to boron or boron-containing compounds; we presume that the boron is present as amorphous boron and thus does not contribute significant X-ray intensity. The dissimilar behaviors of Samples A and B arise from the interplay between thermodynamic equilibrium and the kinetics of dehydrogenation. This can be illustrated by reference to Figure 4, showing the pressure-temperature diagram for the LiBH4/MgH2 system. The solid lines represent thermodynamic equilibrium boundaries for two relevant reactions. The leftmost
Letters
Figure 3. X-ray diffraction patterns for Samples A-C after dehydrogenation.
Figure 4. Temperature-pressure phase diagram for the Li-Mg-B-H system. Solid lines represent thermodynamic stability boundaries for LiBH4 + 1/2MgH2 f LiH + 1/2MgB2 + 2H2 (left) and MgH2 f Mg + H2 (right). The solid square symbol is the thermodynamic equilibrium T(1 bar) value estimated for decomposition of liquid LiBH4 via reaction 1 as described in the text. Dash-dotted curves represent the temperature-pressure trajectories of samples A-D during dehydrogenation. Open circles are the onsets of dehydrogenation for pure LiBH4 measured by DSC at a heating rate of 10 °C/min; the open triangle is a DSC measurement performed at an H2 pressure of 1.1 bar with a heating rate of 2 °C/min. Filled circles are the temperatures at which the sample foamed out of the sample bucket in TGA experiments at 1.3 bar H2. The dashed line is a guide to the eye through the combined DSC and TGA points representing the approximate kinetic boundary for LiBH4 decomposition.
line shows the equilibrium pressure as a function of temperature for reaction 3, calculated using the reaction enthalpy ∆H ) -40.5 kJ/mol H2 and entropy ∆S ) -0.0813 kJ/(K mol H2) measured by Vajo et al.27 Decomposition of LiBH4 + 1/2MgH2
J. Phys. Chem. C, Vol. 111, No. 35, 2007 12883 to LiH + 1/2MgB2 + 2H2 is thermodynamically favored throughout the region to the right of this line. We note that replacing MgH2 by Mg metal as in reaction 3a yields an even lower T(1 bar); LiH, MgB2, and H2 are thermodynamically stable with respect to LiBH4 + Mg everywhere in the temperature-pressure regime of Figure 4 (i.e., the thermodynamic boundary lies below 250 °C at 5.5 bar). The second line represents reaction 2 for independent dehydrogenation of MgH2 calculated from the ∆H and ∆S values for MgH2 provided in the Sandia metal hydrides database.32 The thermodynamic boundary for LiBH4 decomposition, reaction 1, is more problematical because specific heat measurements do not exist for LiBH4 in this temperature regime. Extrapolating the values of ∆H ) 67 kJ/mol H2 and ∆S ) 97 J/(K mole H2) for LiBH4 at 30 °C in the Outokumpu HSC Chemistry software35 and assuming LiH, crystalline B, and H2 as decomposition products yields T(1 bar) ) 408 °C.36 A similar value is obtained from the NIST-JANAF thermodynamic tables.37 Although often used in the literature, these values do not include the orthorhombic-to-hexagonal structural transformation38 at 113 °C or the melting transition at 288 °C, and thus are inappropriate at high temperature. Recently, Zu¨ttel et al. have included the enthalpies of the phase transition and of melting to estimate ∆H ) 61 kJ/mol H2 for reaction 1 starting from liquid LiBH4.39 Similar calculation of ∆S using entropies from HSC Chemistry and including the entropy changes ∆Ss ) ∆Hs/Ts and ∆Sm ) ∆Hm/Tm for the structural transformation and melting, respectively, gives ∆S ) 84 J/(K mole H2) and T(1 bar) ) 457 °C. The latter value is indicated on Figure 4 by the filled square symbol. We performed DSC measurements on pure LiBH4 under H2 partial pressures from 0.5 to 5 bar to try to further elucidate the appropriate boundary for LiBH4 decomposition. Because the working gas was a 50% H2/50% N2 mixture (to prevent thermal saturation of the instrument), the DSC response was likely dominated by a large contribution to the heat flow due to the change in thermal conductivity of the gas as hydrogen was released from the sample. Consequently, the thermal signature was substantially larger than the heat of reaction alone. It does, however, produce a valid measurement of hydrogen release. To obtain high sensitivity, a heating rate of 10 °C/min was used for most scans. At each H2 partial pressure, we obtained the onset of hydrogen release by looking for the temperature at which an excess thermal signature was first detectable. These points are represented by the open circles in Figure 4. For comparison, a DSC scan at 2 °C/min, equivalent to the heating rate used in the volumetric measurements, was performed at an H2 partial pressure of 1.1 bar; the shift in the observed onset temperatures due to DSC heating rate (∼16 °C) is comparable to the scatter in the data (∼11 °C). Although the scatter is significant, and the selection of onset temperatures is somewhat subjective, it is clear that the DSC data lie substantially to the left of the liquid LiBH4 T(1 bar) (filled square symbol) calculated from reaction 1. The inverse is also true, that is, at the equilibrium temperatures predicted by the ∆H and ∆S values derived above, the DSC data already have substantial hydrogen signatures (see the Supporting Information). This observation received additional support from an unusual behavior noted in thermogravimetric measurements of LiBH4 + 1/2MgH2 + 0.02TiCl3 performed under 1.3 bar pure H2 at a heating rate of 2 °C/min. After partial dehydrogenation corresponding to decomposition of MgH2, the sample foamed out of the sample bucket, making contact with the surrounding support structure (after which we could no longer measure sample mass). Two
12884 J. Phys. Chem. C, Vol. 111, No. 35, 2007 such experiments were attempted, and the temperatures at which the samples lost confinement are shown as solid circles in Figure 4. One of the experiments was nevertheless continued through a 3 h soak at 450 °C, after which XRD of the recoverable material showed the same Mg and LiH products as in the H2free case (except for a small quantity of retained MgH2). We infer that these instability points represent evolution of H2 gas from the decomposition of LiBH4, and their location is consistent with the values obtained from DSC. We include the dashed line in Figure 4 as a guide to the eye representing the aggregate DSC and TGA data for the observed earliest onset of liquid LiBH4 decomposition. This almost certainly represents a kinetic rather than a thermodynamic boundary (hence, the line is dashed); the equilibrium thermodynamic boundary must be to the left of the observed decomposition. Our data are inconsistent with the estimated T(1 bar) value for liquid LiBH4 decomposition (filled square). We conclude that the derived T(1 bar) does not adequately describe the actual decomposition behavior of pure LiBH4, probably because reaction 1 oversimplifies the actual chemical reaction(s) for LiBH4 decomposition;31 intermediate phases may be forming during thermal decomposition18,19,31 via reactions characterized by lower equilibrium temperatures. The temperature-pressure trajectories of Samples A and B during dehydrogenation are superimposed on Figure 4 as dashdotted curves. Each sample began at room temperature to the left, and moved to the right as the sample was heated according to the temperature profile shown in Figure 2. As Sample A dehydrogenated, the pressure in the sample cell increased as shown by the trajectory labeled “A”. Because this was a large sample (1.018 g) in a small cell volume (220 cm3), the pressure increased substantially as hydrogen was evolved, reaching 4.6 bar at the end of dehydrogenation. In contrast, Sample B was a small sample (0.409 g) in a large cell volume (725 cm3); thus, the pressure increase due to hydrogen evolution was about an order of magnitude smaller, 0.55 bar. In both cases, about the same weight percent of hydrogen was produced from the sample, as indicated in Figure 2. The trajectory of Sample A is clearly in the regime of thermodynamic stability of the decomposition products LiH + 1/ MgB + 2H at all temperatures shown in Figure 4. 2 2 2 Temperatures significantly higher than the thermodynamic boundary are required to initiate decomposition as a consequence of the limited reaction kinetics at low temperature. Furthermore, the two discrete decomposition steps of reaction 3a are evident in Figure 2. The hydrogen loss in the first step, between 280 °C and 340 °C, corresponds to dehydrogenation of the MgH2 component; in situ XRD measurements as a function of temperature confirm the formation of metallic Mg in this temperature range. The MgH2 and LiBH4 are kinetically decoupled, such that decomposition of MgH2 is kinetically favored in this temperature range, whereas the reaction between the MgH2 and LiBH4, although thermodynamically favored, remains kinetically blocked. Further decomposition occurs at higher temperature by the reaction of Mg with LiBH4 to form the MgB2 product. The important point here is that because the pressure increases rapidly with dehydrogenation for Sample A the sample is always in a temperature-pressure regime to the left of the LiBH4 decomposition boundary; thus, direct decomposition of the LiBH4 component is suppressed, as was also observed by Nakagawa et al.34 Moreover, these results demonstrate that the presence of Mg clearly assists the decomposition of LiBH4, which occurs as much as 25 °C below the DSC onsets for pure LiBH4. The maximum dehydrogenation rate occurs at ∼410 °C and H2 pressures at or above 3 bar, whereas
Letters the maximum decomposition rate for LiBH4 alone estimated from the DSC measurements (after adjusting for the higher heating rate in the DSC) occurs at least 70 °C higher at those pressures. A similar result was obtained by Yu et al. in a Mgrich LiBH4/(Mg or MgH2) system, where the presence of Mg was found to facilitate hydrogen desorption from LiBH4.40 The trajectory for Sample B also falls to the right of the boundary representing reaction 3, but now the low H2 pressure places the sample beyond the LiBH4 decomposition boundary as well. Once again, dehydrogenation began with the decomposition of MgH2. In this case, however, the H2 pressure was insufficient to suppress direct LiBH4 decomposition. Dehydrogenation of LiBH4 to LiH and amorphous boron, as observed in XRD, was kinetically preferred over the reaction with Mg, presumably because the former does not require mass transport. The amorphous boron in Sample B was kinetically very resistant to further reaction; thus, in Sample B only the metallic Mg rehydrogenated. In contrast, Sample A retained the boron in the form of MgB2, keeping it available for the reverse, hydrogenation, reaction. This interpretation is consistent with the recent suggestion of Barkhordarian et al.33 that the presence of boron as MgB2 rather than elemental boron substantially lowers the activation barrier for the formation of the [BH4]complexes during hydrogenation. It is interesting to note that, although samples at the LiBH4 + 1/2MgH2 stoichiometry of reaction 3 do not rehydrogenate to form LiBH4 unless MgB2 is present, Yu et al.40 have reported successfully regenerating LiBH4 during rehydrogenation of a heavily Mg-enriched composition, LiBH4 + 3.3MgH2 (containing more that 6 times the initial MgH2 content as our samples), even though MgB2 did not form during dehydrogenation. To verify our interpretation of dehydrogenation behavior in terms of Figure 4, Sample C was prepared in the same manner as Samples A and B, and initially hydrogenated by the same process as in Figure 1 (see the Supporting Information). Like Sample B, Sample C was dehydrogenated using a small sample size (0.399 g) placed within a large sample volume (725 cm3) to limit the pressure rise during dehydrogenation. In this case, however, the sample cell was initially charged with an H2 pressure of 4.2 bar. Its dehydrogenation behavior, shown in Figure 2, is similar to that of Samples A and B except that the MgH2 decomposition is shifted to higher temperature by the H2 overpressure, as expected. Its XRD pattern after dehydrogenation, seen in the lower panel of Figure 3, shows LiH and MgB2, as expected for reaction 3a. This can be understood from its trajectory in Figure 4; here we see that the pressure change during dehydrogenation is similar to that of Sample B but that the H2 overpressure now keeps the trajectory entirely to the left of the LiBH4 stability line. Sample C was rehydrogenated successfully in the same manner as Sample A (see the Supporting Information). Gravimetric measurements of dehydrogenation from LiBH4 + 1/2MgH2 doped with 2 mol % TiCl3 further confirmed these results. Samples run at 3 or 5 bar H2 pressure formed MgB2 and LiH, whereas a sample run in He gas with no H2 present formed metallic Mg and LiH (and presumably amorphous boron). We have already noted the unusual behavior at 1.3 bar H2 overpressure in which the sample foamed from the sample bucket, attributed to evolution of H2 gas from the direct decomposition of LiBH4. TGA experiments with flowing inert gas on Mg-rich LiBH4 + 3.3MgH2 by Yu et al.40 similarly found that LiBH4 dehydrogenated near 405 °C without the formation of MgB2 (although MgB2 did subsequently form at temperatures above 500 °C).
Letters Finally, the experiment was repeated with Sample D, comprised of LiBH4 + 1/2MgH2 + 0.03TiH2. TiH2 proved lesseffective as a catalyst for accelerating the dehydrogenation of LiBH4 + 1/2MgH2, as shown by the much slower kinetics of Sample D in Figure 2. Its trajectory is shown in Figure 4 during heating to 450 °C and soaking at that temperature (the nearly vertical rise at 450 °C). Like Sample B, the dehydrogenated state of this sample was characterized by LiH and Mg metal plus, we assume, amorphous boron. Similar to Sample B, Sample D could only be partially rehydrogenated to about 2.6 wt %. Its behavior is thus consistent with crossing the LiBH4 decomposition boundary into the region of LiBH4 instability. We emphasize that as a kinetic boundary, the dashed line is not an exact demarcation but rather reflects the point where the kinetics of LiBH4 decomposition begin to compete significantly with the kinetics of reaction 3a; to the left of the line reaction 3a is strongly dominant (as for Samples A and C), whereas beyond the line the kinetics of reaction 1 become increasingly dominant (as for Samples B and D). Our results on the LiBH4/MgH2 system emphasize the importance of fully understanding the interaction between the equilibrium thermodynamic phase diagram and the reaction kinetics in destabilized complex hydrides. Only by doing so can the most favorable conditions for destabilization be chosen. The best case is to tune the reaction kinetics, for example by finding appropriate catalysts, so that the transformation occurs within the desired thermodynamically stable region. Even when optimal kinetics are difficult to obtain, however, it is still possible, by tuning a suitable thermodynamic variable such as H2 pressure, to practically navigate the phase diagram and avoid detrimental reaction regions. In the present example, an applied H2 pressure of at least 3 bar was required to ensure that dehydrogenation of TiCl3-catalyzed LiBH4 + 1/2MgH2 always occurred within the region of LiBH4 stability, thus forming MgB2 and promoting reversible hydrogen storage. Acknowledgment. We thank Sky Skeith for technical assistance, Jan Herbst for many useful discussions, and Mark Verbrugge and Jim Spearot for their support of this research. Supporting Information Available: Synthesis procedures; details of volumetric, gravimetric, X-ray diffraction (XRD), and differential scanning calorimetry (DSC) measurements; hydrogenation data for samples C and D; gravimetric weight loss data; XRD results after gravimetric measurements; DSC scans. This material is available free of charge via the Internet at http:// pubs.acs.org. References and Notes (1) Zu¨ttel, A. Mater. Today 2003, September, 24-33. (2) U.S. Department of Energy - Energy Efficiency and Renewable Energy: Hydrogen, Fuel Cells and Infrastructure Technologies Program, Hydrogen Storage System Requirements, available at: http://www. eere.energy.gov/hydrogenandfuelcells/storage/current_technology.html. (3) Bogdanovic´, B.; Schwickardi, M. J. Alloys Compd. 1997, 253254, 1.
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