Effect of Ti-catalyst content on the reversible hydrogen storage properties of the sodium alanates

Journal of Alloys and Compounds 339 (2002) 299–308

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Effect of Ti-catalyst content on the reversible hydrogen storage properties of the sodium alanates G. Sandrock a

a,b ,

*, K. Gross a , G. Thomas a

Sandia National Laboratories, P.O. Box 969, MS 9402, Livermore, CA 94551, USA b SunaTech, Inc., 113 Kraft Pl., Ringwood, NJ 07456, USA

Received 27 March 2001; received in revised form 4 June 2001; accepted 29 November 2001

Abstract The reversible hydrogen storage properties of Ti-catalyzed NaAlH 4 (and associated Na 3 AlH 6 ) were studied as a function of Ti-content using a dry preparation technique consisting of the ball-milling of NaAlH 4 1TiCl 3 mixtures (0–9 mol.% TiCl 3 ). This process is believed to result in the in situ solid-state introduction of metallic Ti via the reduction of the TiCl 3 by the Na-component of NaAlH 4 (to form NaCl). Properties studied were hydriding and dehydriding rates and reversible gravimetric H-capacity as a function of starting TiCl 3 content. Detailed isothermal kinetic studies were done over a wide temperature range (20–225 8C) and treated by Arrhenius analysis. All kinetics were found to follow the Arrhenius equation and the changes of thermal activation energies and rate constants with TiCl 3 -content were observed that may give valuable insights into the as-yet unknown mechanistic factors that control H 2 desorption and absorption kinetics. Ti increases both dehydriding and hydriding kinetics (and associated practical engineering rates), but at the substantial expense of H-capacity. The finite room temperature decomposition of the catalyzed NaAlH 4 phase was reconfirmed and rates better quantified. Published by Elsevier Science B.V. Keywords: Hydrogen absorbing materials; Sodium alanates; NaAlH 4 ; Na 3 AlH 6 ; Ti-catalysis; Hydriding and dehydriding kinetics; Arrhenius analysis; Hydrogen storage capacity

1. Introduction The historical availability of reversible metal hydrides with high-capacity hydrogen storage has been hampered by thermodynamic and kinetic limitations. On the one hand, interstitial hydrides that are easily reversible around room temperature (e.g. those based on V or the AB, AB 2 and AB 5 intermetallic compounds) represent essentially metallic H-bonding and are limited to only about 1.5–2.5 wt.% reversible gravimetric H-capacity [1,2]. At the other extreme, we have reversible hydrides that exhibit strong covalent or ionic H-bonding (e.g. MgH 2 and LiH, respectively) and can provide good H-capacity (7–13 wt.%), but unfortunately require temperatures greater than 250 8C to release the bound H. Until 1996, there were essentially no practical choices between these extremes. Starting in 1996, Bogdanovic’ and co-workers reported that the Na-alanates, which were generally considered *Corresponding author. E-mail address: [email protected] (G. Sandrock). 0925-8388 / 02 / $ – see front matter Published by Elsevier Science B.V. PII: S0925-8388( 01 )02014-X

nonreversible, could be made reversible by doping with Ti catalysts [3,4]. Thus the practical possibility of the following two-stage reaction was demonstrated: NaAlH 4 ~1 / 3Na 3 AlH 6 1 2 / 3Al 1 H 2 ~NaH 1 Al 1 3 / 2H 2

(1)

Stoichiometrically, the first step consists of 3.7 wt.% H 2 release and the second step 1.9 wt.%, for a theoretical net reaction of 5.6 wt.% reversible gravimetric H-storage. Furthermore, the temperatures for atmospheric pressure desorption are about 33 8C for the first step (NaAlH 4 decomposition) and about 130 8C for the second step (Na 3 AlH 6 decomposition) [4–6]. In summary, this material offers a new and previously unoccupied thermodynamic realm for reversible hydrides. The remaining problem is the rather slow hydriding and dehydriding kinetics below about 150 8C. If rapid rates could be accomplished below 100 8C, the catalyzed alanate would be of significant practical value for high-capacity, on-board H 2 storage for low-temperature fuel-cell vehicles. In fact, there has been

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significant work done in the past 5 years to improve the kinetics of the catalyzed Na-alanates through particle size and catalyst control. This work has been summarized in two recent reviews [5,6].

2. Solid state Ti-doping with the catalyst precursor TiCl 3 The first (Bogdanovic’ et al.) work on the doping of NaAlH 4 with Ti used solution chemistry techniques whereby nonaqueous liquid solutions or suspensions of NaAlH 4 and either TiCl 3 or alkoxide Ti(OBu n) 4 catalyst precursors were decomposed to precipitate solid Ti-doped NaAlH 4 [3,4]. An alternate, nonsolution approach was taken by Jensen and Zidan at the University of Hawaii, whereby the liquid Ti(OBu n) 4 precursor was simply ball-milled with the solid NaAlH 4 [7,8]. They also added Zr(OPr i ) 4 to help stimulate the kinetics of the second step (Na 3 AlH 6 decomposition) of the reaction shown in Eq. (1). We started developing this ‘semi-dry’ ball-milling concept into practical engineering levels via a cooperative effort with these investigators. There is some practical merit in this approach because of ease of production and the potential for minimizing the handling of regulated solvents such as tetrahydrofuran (THF) and toluene. The problem we found was that the alkoxide precursors decomposed with ballmilling and subsequent service, thus contaminating the H 2 released with high levels of hydrocarbons [9]. In addition, H-capacities were only about half of the 5.6 wt.% reversible H-capacities predicted in Eq. (1). Why not turn to a completely dry, inorganic process? A preliminary ballmilling experiment with 2 mol.% solid TiCl 3 198 mol.% solid NaAlH 4 suggested we could substantially reduce both the impurity and H-capacity problems encountered with the alkoxides [9]. We have carried this dry inorganic approach further and report herein the effects of varying the level of the TiCl 3 precursor on the reversible hydrogen storage properties of the Ti-doped alanates (dehydriding and rehydriding kinetics, along with H-capacity). There are no earlier data that systematically cover the effects of TiCl 3 -catalyst precursor level. It is important to understand our use of the term ‘catalyst precursor’ when describing the starting chemical species TiCl 3 or Ti(OBu n) 4 . As shown by the solution approach, gaseous H 2 is evolved when these species are added to toluene suspensions of NaAlH 4 , with the resultant H 2 yields suggesting both the TiCl 3 and alkoxides decompose to form zero-valent (metallic) Ti [4]. A similar effect occurs in the semi-dry (alkoxide) or dry (TiCl 3 ) ball milling processes. For example, X-ray diffraction patterns of our ball-milled TiCl 3 –NaAlH 4 mixtures clearly show that the TiCl 3 is completely reduced by the Na in NaAlH 4 to form NaCl and most likely zero-valent (metallic) Ti [10]. The general solid state reaction can probably be written as follows:

(1 2 x)NaAlH 4 1 xTiCl 3 → (1 2 4x)NaAlH 4 1 3xNaCl 1 xTi 1 3xAl 1 6xH 2 ,

(2)

where x is the mole fraction of TiCl 3 added to the initial ball milling charge. Those familiar with Ti production metallurgy will immediately recognize similarities to the old Hunter process, whereby liquid TiCl 4 is metallothermically reduced by hot Na to form sponge Ti metal (closely related to the more common Kroll process where hot Mg is used) [11]. The main point is that it is the zero-valent Ti that must be the catalyst and TiCl 3 is only the precursor. We do not represent that Eq. (2) is precisely correct in the sense it is not necessarily elemental Ti that is the final catalyst. It could be TiH 2 , a Ti-subhydride, a Ti-alloy or an intermetallic compound (or hydride thereof) that is the true catalyst. For example, hydrides of Ti 3 Al have been suggested [12]. Whatever the catalyst is, it is either amorphous or too small to be identified by XRD. The fact is that the Ti species is created by the apparently fine in situ reduction of the initial TiCl 3 by the partial destruction of NaAlH 4 and the Ti so produced has a very powerful catalytic effect on the remaining NaAlH 4 , as we shall dramatically show in this paper.

3. Experimental procedures Crystalline NaAlH 4 was made by cryopumping THF from a 1.0 M solution of NaAlH 4 in THF (Aldrich Chemical no. 40,424-1), followed by vacuum drying with a mechanical and / or turbomolecular pump (down to 73 10 25 Pa dynamic pressure). The crystalline TiCl 3 was 99.999% pure (Aldrich Chemical no. 51,438-1). Mixtures of NaAlH 4 and TiCl 3 were weighed in a purified argon glovebox in the levels of 0.9, 2, 4, 6 and 9 mol.% TiCl 3 . These mixtures were then ball-milled in argon for 3 h, using a high-energy SPEX  mill. The balls and milling vial were WC. X-ray diffraction (using special airless sample holders) was done before and after milling; in all cases the complete conversion of TiCl 3 to NaCl (Eq. (2)) during milling was confirmed. Dehydriding and hydriding rates and capacities were obtained volumetrically with a carefully calibrated and instrumented Sieverts’ apparatus using a cylindrical 316 SS reactor (1.3 cm outer diameter by 0.12 cm wall thickness) containing about 1.5 g of catalyzed samples. This reactor also had 0.8 mm diameter internal and external thermocouples, was heated with external electrical heating tape and was cooled by ambient external air. Absorption pressure changes were quantified with a calibrated 20 MPa pressure transducer and desorption pressures with a calibrated 0.13 MPa (absolute) Baratron  capacitance manometer. Data were recorded via computer and hydride / dehydride (H / D) runs lasting from minutes to several days, depending on the TiCl 3 level and test pressure and temperature conditions. In order to reduce the

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ubiquitous trace levels of residual THF, all milled samples were additionally turbomolecular-pumped overnight before starting to obtain H / D data. Our synthesis of the starting NaAlH 4 is artificially considered the first absorption half-cycle, 1A. The first desorption half-cycle (1D) was performed by starting at room temperature and stepping the temperature up to 150 8C over the period of an hour or so, and finally holding at 150 8C until H 2 desorption was complete. Although this TPD-like initial desorption produces some useful qualitative kinetic data, it generally did not yield a reliable capacity value because it was obvious that some H 2 was lost during ball-milling (especially with the higher catalyst loadings). All of the ‘complete’ hydriding and dehydriding curves shown herein were measured during the second absorption and desorption half-cycles (2A and 2D) and were performed at a nominal 125 8C. During absorption, the applied H 2 pressure was generally in the 8–9 MPa range, well above the 3–4 MPa plateau pressure for NaAlH 4 at 125 8C [4–6]. For the desorption experiments, the backpressure during NaAlH 4 decomposition was kept below 0.1 MPa (absolute) and during Na 3 AlH 6 decomposition below 0.025 MPa (abs.), well below the Na 3 AlH 6 plateau pressure of about 0.2 MPa. Although the ‘complete’ hydriding and dehydriding curves described above represent excellent indications of catalytic effects and also provide very useful engineering data, they are not very appropriate for kinetic analysis because they are not very isothermal. Substantial exothermic and endothermic temperature excursions occur during ‘fast’ absorption and desorption, respectively. In fact, it is even possible to ‘self-heat’ the sample to greater than the 183 8C melting temperature of NaAlH 4 [6,9]. Therefore, we performed desorption Arrhenius analyses (isothermal desorption kinetics vs. temperature) using a different technique from the quasi-isothermal tests discussed above. We started with samples in the fully hydrided condition that had been cycled a few times. After the final hydriding half-cycle, the samples were cooled to room temperature and desorption begun. After getting the room-temperature kinetics from the pressure rise in a known volume, the temperature was periodically stepped higher, waiting after each step to collect enough data to precisely determine the slopes of the desorption curve. The pressure rises were always virtually linear with time. The stepping–wait procedure was continued up to 150 8C, holding at that temperature until the NaAlH 4 decomposition step (Eq. (1)) was finished. The sample was then quickly cooled to 40–60 8C and the Na 3 AlH 6 rates determined by the same stepwise procedure up to 180 8C (225 8C for an uncatalyzed sample). It is important to use an internal thermocouple in good thermal contact to record the isothermal bed temperatures precisely during each step. The resultant isothermal data was used to fit the wellknown Arrhenius equation rate 5 k exp(2Q /RT )

(3)

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to determine the preexponential rate constant k and the activation energy Q for each of the catalyzed samples, as well as for a reference uncatalyzed sample (NaAlH 4 ballmilled 3-h without TiCl 3 ). Graphical examples of the least squares Arrhenius fits are shown herein as log (rate) vs. 1 /T plots. Because we are interested in achieving the theoretical 5.6 wt.% H 2 content of NaAlH 4 , most of the rate and gravimetric H-capacity data are presented in terms of only the alanate weight of the mixture before ball-milling (i.e. only the NaAlH 4 component of the left side of Eq. (2)). The engineering wt.% values (based on alanate and inert species) will also be given in appropriate context.

4. Results and discussion

4.1. Engineering properties—effects of Ti-content on hydriding /dehydriding rates Comparative 125 8C hydriding and dehydriding curves for all the catalyzed samples studied (0.9–9 mol.% TiCl 3 precursor) are shown in Figs. 1 and 2, respectively. For each series, the test conditions, initial temperature and H 2 pressure were held as closely as possible. Both the hydriding and dehydriding show the same trends, increasing rates and decreasing H-capacities with increasing TiCl 3 content. On the scales shown in Figs. 1 and 2, the uncatalyzed sample (0% TiCl 3 ) showed virtually no signal, so the effect of Ti-catalyst doping is very strong even at its lowest level. This will be discussed in more detail in Section 4.3. The hydriding curves (Fig. 1) show noteworthy differences from the dehydriding curves (Fig. 2). Except for the 0.9% TiCl 3 level, all samples in Fig. 1 were essentially saturated within the 2-h time frame. The hydriding curves rise smoothly to H-saturation. It is likely that both the Na 3 AlH 6 and NaAlH 4 form simultaneously in a nonequilibrium manner during rapid charging. On the other hand, the dehydriding curves of Fig. 2 show two distinct stages, a rapid first stage followed by a much slower second stage. The break between the two stages is rather sharp. Referring to Eq. (1), we interpret the break as the end of the NaAlH 4 →Na 3 AlH 6 stage [13]. In other words, the desorption process seems to be closer to equilibrium conditions in following the stepwise right reaction sequence of Eq. (1) [14]. Unfortunately, from a practical point of view, the rates of the Na 3 AlH 6 →NaH second stage are too slow for most engineering purposes. Even though there were positive effects of Ti-catalysis, none of the samples were completely desorbed within the 5-h time frame of Fig. 2. The dehydriding measurements were carried out to longer times (50–70 h). The 0.9% TiCl 3 sample was not completely desorbed even after 70 h; the 2 and 4% samples took about 50 h for full desorption. To summarize the practical engineering consequences of

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Fig. 1. H 2 absorption curves starting from NaH1Al (dehydrided NaAlH 4 ) as a function of added TiCl 3 catalyst precursor (expressed in mol.%). T i 5125 8C. Applied PH 2 58.8–7.9 MPa.

the 125 8C rate data, full charging of 2–9% Ti-catalyzed alanates can be accomplished in less than an hour at 8–9 MPa applied H 2 pressure. Similarly, discharging of the NaAlH 4 →Na 3 AlH 6 first stage can also be accomplished within an hour or so for the 2–9% Ti-levels. The second stage Na 3 AlH 6 →NaH desorption is the problem and will especially need further work for practical utilization.

4.2. Engineering properties—effects of Ti-content on Hcapacity For engineering purposes we are interested in the true capacity of the catalyzed alanate bed, i.e. H 2 stored per total weight. Firstly, simply mixing of the catalyst precursor TiCl 3 dilutes the sample. Secondly, the reaction of

Fig. 2. H 2 desorption curves starting from NaAlH 4 as a function of added TiCl 3 catalyst precursor (expressed in mol.%). T i 5125 8C.

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the TiCl 3 with NaAlH 4 (Eq. (2)) eliminates more Hcapacity (i.e. results in NaCl, Al and Ti, all of which do not store H 2 per se). Thus in an engineering sense, the very positive kinetic benefits of Ti-catalysis are seriously offset by losses in reversible H-capacity. The measured H-capacities as a function of TiCl 3 addition are shown in Fig. 3 (with the inclusion of data from a 225 8C bakeout of similarly ball-milled but uncatalyzed NaAlH 4 ). The data are plotted in two ways, with a linear fit applied to each. The higher (solid) curve represents the H-capacity as based only on the weight of the starting NaAlH 4 on the left side of Eq. (2) (i.e. what we used in Figs. 1 and 2). The lower (dashed) curve represents the true (engineering) H-capacity and is based on the weight of the starting NaAlH 4 1TiCl 3 , i.e. the entire left side of Eq. (2). The negative effect of the TiCl 3 loading on H-capacity is obvious. In addition, our uncatalyzed (0% TiCl 3 ) showed only 5.1 wt.% H 2 , significantly below the 5.6 wt.% expected from Eq. (1). How does the measured engineering H-capacities compare with what might be expected from Eq. (1) (5.6 reversible wt.% H 2 with textbook NaAlH 4 ) and Eq. (2) (reaction losses)? This is shown in Fig. 4. Here we take our measured true (engineering) H-capacities and backcalculate the NaAlH 4 -content assuming that has the theoretical capacity of 5.6 wt.% H (data points and linear-fit lower solid line). This data is compared to the theoretical expected curves using both the left side (solid, as mixed) and right side (dashed, after reaction) of Eq. (2) (also assuming 5.6 wt.% H for pure NaAlH 4 ). Comparison of the actual line with the ‘after reaction’ line shows slightly different slopes and negative offsets ranging from 9 to 18 percentage points. The roughly similar slopes imply Eq. (2) is basically correct (less the aforementioned possibility if and how the Ti and Al alloy). The offset suggests

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questions relative to the quality of our starting NaAlH 4 and / or possible damage that might have occurred during ball-milling, e.g. leaks of the milling vial. In summary, the losses of capacity with TiCl 3 addition are serious from a practical storage point of view. Our future work will focus on that problem, both toward reducing the capacity offset and reducing the deleterious effects of the basic Eq. (2) TiCl 3 reduction reaction to introduce the catalyst.

4.3. Arrhenius analyses Isothermal kinetic studies, including Arrhenius analysis, add very useful information on the basic effects of Ticatalyst level, not to mention providing useful engineering data as a function of temperature. We have done such desorption analyses on both the NaAlH 4 and Na 3 AlH 6 decomposition reactions for the 0, 0.9, 2, 4, and 6 mol.% TiCl 3 samples. An Arrhenius plot consists of plotting ln (rate) vs. 1 /T, with the slope of the plot representing the activation energy Q and the 1 /T 5 0 intercept representing the rate constant k (see Eq. (3)). Examples of the resultant Arrhenius plots for 0 and 4% TiCl 3 are shown in Fig. 5. All of the data can be represented quite well with an exponential equation of the form Eq. (3). (Actually, careful examination of the NaAlH 4 data shows a slight concaveupward curvature, but that curvature is small for our purposes and we shall ignore it for now.) The differences between the two sets of Arrhenius lines in Fig. 5 are dramatic. Note the multiple order of magnitude rate shifts for both the NaAlH 4 and Na 3 AlH 6 lines from 0 to 4% TiCl 3 (or equivalently the substantial shifts in temperatures for given rates). Note the NaAlH 4 and Na 3 AlH 6 lines are relatively close for the uncatalyzed sample, and have similar activation energies (slopes). The 4% data show the

Fig. 3. Effect of TiCl 3 addition level on H-capacity of ball-milled NaAlH 4 1TiCl 3 .

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Fig. 4. Estimated actual NaAlH 4 content after ball-milling as a function of TiCl 3 -loading, in comparison with theoretical levels expected from both sides of Eq. (2).

NaAlH 4 and Na 3 AlH 6 lines are further apart, have significantly different, and lower, activation energies. Let us explore these remarkable differences by examining the Arrhenius data for all the samples evaluated. The Arrhenius parameters are tabulated in Table 1 and the values of activation energy Q and rate constant k are plotted versus TiCl 3 level in Figs. 6 and 7, respectively. As

can be seen in Fig. 6, the Qs are nearly equal for the decompositions of both the NaAlH 4 and Na 3 AlH 6 with 0% TiCl 3 , but drop to much lower and different levels when even a small amount of catalyst is added (0.9%). This indicates a substantial easing of the thermally activated mechanism. Q then remains virtually constant with TiCl 3 level above 0.9%, suggesting no further change in fun-

Fig. 5. Arrhenius lines for NaAlH 4 and Na 3 AlH 6 decompositions for 0 and 4 mol.% added catalyst precursor TiCl 3 .

G. Sandrock et al. / Journal of Alloys and Compounds 339 (2002) 299 – 308 Table 1 Experimentally derived parameters for the Arrhenius equation, rate5k exp(2Q /RT ). Included are the true H-capacities of each reacted sample (from Fig. 3) TiCl 3 level mol.% 0 0.9 2 4 6

True Hcapacity wt.% H 2 5.12 4.94 4.25 3.85 2.91

Rate constant k wt.% H 2 / h NaAlH 4

Activation energy Q kJ / mol H 2

Na 3 AlH 6 13

1.98310 2.06310 9 7.19310 10 1.81310 11 1.63310 11

13

1.43310 1.51310 11 5.33310 11 1.85310 12 1.85310 12

NaAlH 4

Na 3 AlH 6

118.1 72.8 79.5 80.0 78.5

120.7 97.1 97.1 97.5 98.2

damental mechanism after the first small Ti-addition. Fig. 7 shows similar decreases of k, including NaAlH 4 – Na 3 AlH 6 differentiation starting with the first small Tiaddition. Unlike Q, however, k increases monotonically with further increases in TiCl 3 addition, and is therefore alone responsible for all the further increases in kinetics above the first 0.9% Ti-addition. In summary, a minimal addition of Ti-catalyst is necessary to change the thermal activation mechanism for H 2 desorption and reabsorption. In addition, decomposition rates are also proportional to the quantity of catalyst added. Although the exact identity of the catalyst and the mechanism of catalysis are as yet unknown, we can assume it initially works to aid the breaking and reforming of covalent Al:H bonds in the [AlH 4 ] 21 and [AlH 6 ] 23 complexes by lowering the activation energy Q for that process. Further increases in kinetics resulting from increased catalyst levels are manifested through increases in the rate constant k only and may be a consequence of

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reduced particle sizes and / or diffusion distances. It is also possible that the otherwise inert NaCl serves as an antisintering agent, thus maintaining small particles and short diffusion distances. Mechanistic studies continue. The values of the Arrhenius parameters in Table 1 have important engineering value by conveniently allowing one to calculate true isothermal kinetics at any temperature and also model non-isothermal operation of H 2 -storage tanks or other commercial processes. The 125 8C desorption rates were independently calculated from Table 1 and Eq. (3) and are shown in Fig. 8. (The 9% TiCl 3 points were not calculated but rather taken directly from the data of Fig. 2.) In a rather practical way, these data reiterate the valuable effects of Ti-catalysis. The improvements in kinetics with increasing Ti-catalyst-level must be weighed against the loss of H-capacity. However, a reasonable balance of kinetics and capacity can be made with a 4 mol.% TiCl 3 blend. We have chosen that level to further explore larger-bed engineering studies of temperature, variable charging pressure and extended H / D cycling.

4.4. Room temperature desorption The ability to desorb H 2 from the decomposition of the NaAlH 4 phase at room temperature may have practical use in low-demand devices, as we pointed out in our earlier publication [9]. We explored this speculation by making an 800-h (33-day) desorption run on a 6% TiCl 3 sample at room temperature (20–25 8C). These data are shown in Fig. 9. Over an 800-h period the sample desorbed 0.44 wt.% H 2 . During the first 20–40 h, the rate drops off somewhat, but even after 800 h desorption was continuing

Fig. 6. Activation energies Q for NaAlH 4 and Na 3 AlH 6 decompositions as a function of added TiCl 3 .

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Fig. 7. Rate constants k for NaAlH 4 and Na 3 AlH 6 decompositions as a function of added TiCl 3 .

in a more or less linear manner. At that point the desorption rate is somewhat lower than what would be calculated from Eq. (3) using the parameters of Table 1 for

the 6% TiCl 3 addition level. The run used for Fig. 9 was half-cycle 9D. Kinetics may have been influenced by the several previous cycles, including heating to 180 8C (par-

Fig. 8. Desorption rates for NaAlH 4 and Na 3 AlH 6 as a function of added TiCl 3 , as calculated for 125 8C from the Arrhenius equation (Eq. (3)) using the experimental parameters listed in Table 1.

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Fig. 9. Slow room temperature desorption of H 2 from NaAlH 4 with 6 mol.% TiCl 3 catalyst precursor added. The gaps in the data represent datalogger malfunctions.

tial NaH and / or Al sintering?) during the latter stage of the Arrhenius run (cycle 4D). We show the data only to demonstrate conclusively that long-term H 2 generation from NaAlH 4 is possible at room temperature.

3.

5. Conclusions This represents the first systematic study of the effect of the TiCl 3 catalyst precursor level on the reversible hydrogen properties of the system NaAlH 4 –Na 3 AlH 6 –NaH. It focuses especially on hydriding / dehydriding kinetics and H-capacity. It is the first to make extensive use of the dry NaAlH 4 –TiCl 3 ball-milling preparation method. It is also the first to make use of isothermal Arrhenius analysis over a wide range of temperature (20–225 8C), and includes both decomposition steps (Eq. (1)). The main conclusions can be summarized as follows:

4.

5. 1. TiCl 3 decomposes during ball-milling to form NaCl and probably metallic Ti, the latter of which in some form serves as a powerful catalyst for enhancing both dehydriding and hydriding kinetics, even at the lowest level tried (0.9 mol.%). The final species of the Ticatalyst formed in situ has not yet been precisely defined. 2. Systematic Arrhenius analysis of isothermal desorption kinetics demonstrates that the first small addition of Ti-catalyst significantly and discontinuously lowers the

6.

thermal activation energies Q for both decomposition steps, apparently producing a new mechanism for enhanced desorption kinetics. Doping with higher concentrations of TiCl 3 beyond 0.9% further enhances the kinetics, but only via increasing the pre-exponential rate constant k. After the initial decreases with 0.9% TiCl 3 , the activation energy Q, on the other hand, remains constant with increasing TiCl 3 level. From a practical engineering point of view, Ti-catalyzed alanates have reasonably good charging rates at 125 8C. Discharge from NaAlH 4 to Na 3 AlH 6 is also fast enough for practical applications at 125 8C, but the decomposition of the Na 3 AlH 6 to NaH is relatively slow. Improvements in the catalysis of this second desorption step are especially needed. From an overall practical point of view, improved catalysts will be needed for practical charging and discharging kinetics below 100 8C. Improvements in kinetics gained via increased concentrations of the Ti-catalysts are achieved at the significant expense of H-capacity. In addition to the decrease in gravimetric capacity due to the additional weight of the catalyst precursor, significant additional capacity loss comes as a result of the TiCl 3 →NaCl reaction, which destroys some of the starting NaAlH 4 (Eq. (2)). The starting uncatalyzed NaAlH 4 has about 10% lower desorption capacity than it should (5.1 wt.% H 2 vs. 5.6 wt.%).

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7. The slow decomposition of the NaAlH 4 phase at room temperature has been confirmed with an 800-h run.

Acknowledgements We thank Don Meeker for NaAlH 4 preparation, blending with TiCl 3 and ball-milling, i.e. all steps used to make the samples studied. We are also grateful to Eric Majzoub for helpful comments on the manuscript. We also wish to acknowledge the pioneering work of Boris Bogdanovic’ in the area of catalyzed alanates and his positive comments on and encouragement toward our earlier reported alanate work along these lines [9]. We thank Craig Jensen and Ragaiy Zidan for their valuable cooperative interactions with Sandia and for their important early work in developing the ball-mill catalyzing approach for NaAlH 4 [7,8]. This work was performed under support from the Hydrogen Program of the US Department of Energy.

References [1] G. Sandrock, J. Alloys Comp. 293–295 (1999) 877. [2] G. Sandrock, G. Thomas, IEA / DOE / SNL Hydride Databases, http: / / hydpark.ca.sandia.gov

[3] B. Bogdanovic’, M.J. Schwickardi, J. Alloys Comp. 253 (1997) 1. [4] B. Bogdanovic’, R.A. Brand, A. Marjanovic’, M. Schwikardi, J. ¨ Tolle, J. Alloys Comp. 302 (2000) 36. [5] C.M. Jensen, K.J. Gross, Appl. Phys. A 72 (2001) 213. [6] K.J. Gross, G.J. Thomas, C.M. Jensen, J. Alloys Comp. 330–332 (2002) 683. [7] C.M. Jensen, R.A. Zidan, N. Mariels, A.G. Hee, C. Hagen, Int. J. Hydrogen Energy 24 (1999) 461. [8] R.A. Zidan, S. Takara, A.G. Hee, C.M. Jensen, J. Alloys Comp. 285 (1999) 119. [9] G. Sandrock, K. Gross, G. Thomas, C. Jensen, D. Meeker, S. Takara, J. Alloys Comp. 330–332 (2002) 696. [10] K.J. Gross, G. Sandrock, G. Thomas, J. Alloys Comp. 330–332 (2002) 691. [11] S.R. Seagle, in: Titanium and Titanium Alloys, Kirk–Othmer Encyclopedia of Chemical Technology, Vol. 24, John Wiley, New York, 1997, pp. 186–224. ˚ [12] A.J. Maeland, B.C. Hauback, H. Fjellvag, M. Sørby, Int. J. Hydrogen Energy 24 (1999) 163. [13] G.J. Thomas, S.E. Guthrie, K. Gross, Hydride development for hydrogen storage, in: Proceedings US DOE Hydrogen Program Review, Report NREL / CP-570-26938, 1999, http: / / www.eren.doe.gov / hydrogen / docs / 26938toc.html. [14] K.J. Gross, S.E. Guthrie, S. Takara, G.J. Thomas, J. Alloys Comp. 297 (2000) 270.