Formation Reactions and the Thermodynamics and Kinetics of

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Formation Reactions and the Thermodynamics and Kinetics of Dehydrogenation Reaction of Mixed Alanate Na2LiAlH6 Fenghuai Wang, Yongfeng Liu,* Mingxia Gao, Kun Luo, Hongge Pan,* and Qidong Wang Department of Materials Science and Engineering, Zhejiang UniVersity, Hangzhou 310027, People’s Republic of China ReceiVed: February 9, 2009; ReVised Manuscript ReceiVed: March 14, 2009

Na2LiAlH6 was synthesized by ball-milling a mixture of NaH and LiAlH4 at a molar ratio of 2:1. NaH and LiAlH4 were readily converted in the initial ball-milling process to LiH and NaAlH4, which subsequently reacted with the remaining excessive NaH to form Na2LiAlH6. The thermodynamic and kinetic mechanisms of dehydrogenation of Na2LiAlH6 were systematically elucidated. Approximately 6.7 wt % of hydrogen was found stored reversibly in Na2LiAlH6 through a few sequential reactions. An enthalpy change of 63.8 kJ/ mol-H2 and an apparent activation energy of about 173 kJ/mol were determined for the first-step hydrogen storage reaction, indicating the reaction was thermodynamically relatively stable with a high kinetic barrier for the decomposition of Na2LiAlH6. In depth kinetic investigations showed that the first-step dehydrogenation reaction of Na2LiAlH6 could be well interpreted with a nucleation and growth model, and its reaction rate was controlled by the diffusion of substance. The dehydrogenation mechanism developed in this work can be helpful for further efforts on the improvement of the hydrogenation/dehydrogenation performances of Na2LiAlH6. 1. Introduction One serious barrier that prevents the wide adoption of hydrogen as an energy carrier for onboard applications is the lack of an efficient and safe hydrogen storage system.1 Since Bogdanovic´ and Schwickardi demonstrated revolutionarily that Ti-doped NaAlH4 could store H2 reversibly under mild conditions in 1997,2 metal alanates have been widely regarded as potential hydrogen storage materials.3-16 Numerous studies on alanates, both experimental and theoretical, have been made to improve hydrogen storage capability and to understand the catalytic mechanism in the hydrogenation/dehydrogenation process.6-16 Unfortunately, the targets of 5-6 wt % hydrogen storage capacity and low operating temperature around 80-100 °C have not been achieved until now.17 Thus, mixed alanates formulated as MM′(AlH4)n+m or M3-nM′nAlH6 (M and M′ ) alkali metals or alkaline earth metals) with their good performances achieved are attracting more and more attention of investigations, and improved thermodynamic hydrogen absorption/desorption properties are expected on mixing two or more alkali or alkaline earth metals in alanates.18-25 Several mixed alanates have been successfully synthesized and characterized, e.g., Na2LiAlH6, K2LiAlH6, K2NaAlH6, and LiMg(AlH4)3.2,18-26 Among them, Na2LiAlH6 was proved to be reversible for hydrogen storage through the following reaction:24,25

3 Na2LiAlH6 T 2NaH + LiH + Al + H2 2

(1)

Na2LiAlH6, which possesses a cubic close-packed structure of AlH6 with all Li atoms in the octahedral sites and all Na atoms in the tetrahedral sites, was usually prepared either by the reaction of LiAlH4 with 2NaH in toluene or by a solid-state reaction at elevated temperatures and high H2 pressure (50 kbar * To whom correspondence should be addressed. Phone: +86 571 87952615. Fax: +86 571 87952615. E-mail: [email protected] (Y.F.L); [email protected] (H.G.P.).

and 400 °C).18,26 Recently, Huot et al. developed a facile preparation method by ball-milling a mixture of NaH, LiH, and NaAlH4.23 Specimens of the same composition were also obtained by ball-milling a mixture of NaH and LiAlH4 at a molar ratio of 2:1.25 Furthermore, hydrogen storage performances of Na2LiAlH6 were improved by the introduction of some catalysts.2,24,25 Bogdanovic´ and his co-workers reported that Na2LiAlH6 with TiCl3 as the additive had one plateau in P-C isotherms with a reversible hydrogen capacity larger than 3 wt % at 210 °C.2 Thermodynamic investigations on TiF3-catalyzed Na2LiAlH6 revealed that its dissociation enthalpy was about 56.4 kJ/mol-H2, slightly higher than 47 kJ/mol-H2 for Na3AlH6.24 Ma et al. investigated the hydrogenation/dehydrogenation properties of Na2LiAlH6 catalyzed by metal oxides and halides, and their results revealed that a higher reversible capacity and reasonable kinetics were both achieved by using CeO2 as the catalyst.25 Although many efforts have been directed toward synthesis methods, structure characterization, and kinetics improvement of Na2LiAlH6, fundamental mechanisms have not been clearly understood, including the mechanochemical reaction mechanism for its formation in the ball-milling process, the mechanism for the thermal decomposition process, and the kinetic mechanism for the dehydrogenation process, etc. In this work, Na2LiAlH6 was first synthesized by ball-milling a mixture of NaH and LiAlH4 at a molar ratio of 2:1. The mechanochemical reaction mechanism during ball-milling was investigated systematically by examining the phase changes in samples at different milling stages with X-ray diffraction (XRD). The thermal decomposition process of Na2LiAlH6 was then studied by using a temperatureprogrammed decomposition (TPD) and a quantitative volumetric method. For the dehydrogenation kinetic mechanism, isothermal hydrogen desorption measurements were conducted on the firststep decomposition of Na2LiAlH6, and the curves attained were analyzed by the Johnson-Mehl-Avrami (JMA) equation. The dehydrogenation process of Na2LiAlH6 was found to be a diffusion-controlled reaction. The desorption enthalpy change

10.1021/jp9011697 CCC: $40.75  2009 American Chemical Society Published on Web 04/07/2009

Dehydrogenation Reaction of Mixed Alanate Na2LiAlH6

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and the apparent activation energy were further calculated and discussed. The information obtained in this work provides us with insights into the dehydrogenation mechanisms of the mixed alanate Na2LiAlH6 and provides ideas for lowering the high hydrogen desorption temperature. 2. Experimental Section Sample Preparation. The starting chemicals LiAlH4 (purity 95%) and NaH (purity 95%) were purchased from SigmaAldrich and used without any further purification. Na2LiAlH6 was synthesized by ball-milling a mixture of NaH and LiAlH4 at a molar ratio of 2:1 under 1 atm of pure argon atmosphere on a planetary ball mill (QM-3SP4, Nanjing) at 500 rpm. Typically, ∼1.2 g (0.05 mol) of NaH and ∼0.95 g (0.025mol) of LiAlH4 were loaded into a milling jar equipped with a gas valve. The ball-to-powder weight ratio was about 60:1. The handling of samples was carried out in a glovebox (MBRAUN) filled with argon, in which both H2O and O2 concentrations were kept below 1 ppm. Structural Characterization. Phase structures were characterized by means of X-ray diffraction (XRD) on powder samples. Data were acquired at room temperature on a Rikagu D/Max-RA X-ray diffractometer with Cu KR radiation at 40 kV and 30 mA from 10° to 90° (2θ) with step increments of 0.05°. Samples were protected under an argon atmosphere in a homemade container. Property Measurements. Temperature-programmed desorption (TPD) measurement of the sample was performed on a homemade system with a gas chromatograph (GC) attached. In this experiment, approximately 25 mg of sample was loaded and tested with continuous flow of pure Ar as the carrier gas passing through the sample, which was heated from 25 to 620 °C at a constant rate of 2 °C/min. Quantitative temperature-hydrogen absorption/desorption curves were determined by volumetric methods with a homemade Sieverts-type apparatus. For each experiment, ∼200 mg of powder sample was loaded into a stainless-steel tube reactor, which was connected to a Sieverts-type apparatus. Both isothermal and nonisothermal measurements were made. In the nonisothermal experiment, the sample was gradually heated to a desired temperature at an average ramp of 2 °C/min (initially in vacuum) for desorption and at 1 °C/min (initial H2 pressure being 135 bar) for absorption, respectively. Temperatures and pressures of the sample and gas reservoirs were automatically monitored and recorded. For isothermal desorption kinetic measurements the sample was quickly heated to and then kept at a given temperature. The pressure change in the reactor with time was recorded over the temperature range of 200-210 °C. Quantities of hydrogen absorbed/desorbed were determined by the pressure changes in the reactor by means of the ideal gas law calculation. The desorption heat effect was obtained by differential scanning calorimetry (DSC) on a Netzsch DSC 200 3F unit. Each time, 3-4 mg of a sample was used and heated in an aluminum crucible under pure argon with a heating rate of 10 °C/min. 3. Results and Discussion 3.1. Synthesis Mechanism of Na2LiAlH6. Na2LiAlH6 was synthesized by ball-milling a mixture of NaH and LiAlH4 with a molar ratio of 2:1. No hydrogen released from the mixture during ball-milling was detected by measuring the gaseous pressure change within the milling jar up to 24 h. For the study of the mechanochemical reaction mechanism, samples at dif-

Figure 1. XRD patterns of the mixture of 2NaH-1LiAlH4 hand-mixed for 5 min and ball-milled for 0.2-24 h.

ferent ball-milling stages were taken out from the milling jar for XRD analyses. Figure 1 shows the XRD patterns of samples milled for 0.2-24 h. For comparison, the XRD pattern of the hand-mixed 2NaH-LiAlH4 sample using an agate mortar and pestle milling for 5 min was also presented in Figure 1. It is interesting to see that two new phases, NaAlH4 and LiH, were observed in the sample just milled for 0.2 h with weakened intensity of diffraction peaks of LiAlH4. As the milling time increased to 0.4 h, the typical diffraction peaks of LiAlH4 disappeared completely and the sample was composed of NaAlH4, LiH, and NaH. This result indicated that a metathesis reaction between LiAlH4 and NaH occurred during the energetic ball-milling treatment as described below:

NaH + LiAlH4 f NaAlH4 + LiH

(2)

The values of standard enthalpies of formation for LiAlH4, NaH, NaAlH4, and LiH are -116.3, -56.3, -115.5, and -90.5 kJ/mol,27 respectively. With these figures, the reaction enthalpy change calculated for eq 2 is -33.4 kJ/mol, indicating a reaction of exothermic nature. We believe therefore that the reaction can be initialized by energetic ball-milling and then carried on spontaneously since enough energy is supplied with the generation of high pressure in the order of GPa in the solid by colliding balls.28 After ball-milling for 1 h, a new set of peaks at 20.6°, 34.2°, 49.1°, and 61.3° appeared although the NaAlH4, NaH, and LiH phases were still dominating in the XRD profile. Matched with the PDF-2 database of JCPDS-ICDD, this set of peaks indicated that a cubic Na2LiAlH6 phase was formed. With increasing ballmilling time, the Na2LiAlH6 phase gradually became the dominating one with the intensities of peaks of NaAlH4, NaH, and LiH all weakened. As the milling time reached 8 h, Na2LiAlH6 was the only phase detected by XRD along with the complete disappearance of the diffraction peaks of NaAlH4, NaH, and LiH. The following reaction is therefore suggested to be carrying out in the milling period of 0.4-8 h:

NaH + NaAlH4 + LiH f Na2LiAlH6

(3)

With further prolongation of ball-milling time to 24 h, the diffraction peaks of Na2LiAlH6 broaden slightly, signifying a decrease in the crystal grain size. 3.2. Hydrogen Desorption Performance and Mechanism. The hydrogen desorption character of the as-milled sample was first studied by means of TPD as shown in Figure 2. Three peaks of hydrogen evolution are obviously shown in the temperature

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Figure 2. TPD curve of the 2NaH-1LiAlH4 sample after being milled for 24 h.

Wang et al.

Figure 4. XRD patterns of the post-24 h milled 2NaH-1LiAlH4 samples after dehydrogenation at different temperatures.

in the XRD profile. At 230 °C, the intensities of diffraction peaks of the three newly developed phases of NaH, LiH, and Al are all distinctly increased along with the weakening of peaks of Na2LiAlH6. After heating to 250 °C, all reflection peaks for Na2LiAlH6 phase disappeared, and only the well-crystallized NaH, LiH, and Al phases could be observed. These results indicate that Na2LiAlH6 is completely decomposed to yield NaH, LiH, and Al with hydrogen released in the temperature range of 190-250 °C according to the following reaction as proposed previously.2

3 Na2LiAlH6 f 2NaH + LiH + Al + H2 2

Figure 3. Hydrogen desorption (a) and absorption (b) curves (H/MT) for the 2NaH-LiAlH4 sample.

range of 200-430 °C with peak temperatures at 255, 346, and 395 °C, respectively. It indicates that the dehydrogenation process of Na2LiAlH6 is a multistep reaction. The quantity of hydrogen desorbed from each step reaction was further measured by the volumes of hydrogen gathered starting from vacuum. Figure 3a shows the amount of hydrogen desorbed as a function of temperature for the 2NaH-LiAlH4 sample milled for 24 h (symbol 9). Hydrogen desorption starts at about 190 °C with three evident desorption steps, comparing well with the TPD results. In the temperature ranges of 190-250, 320-380, and 380-480 °C, three dehydrogenation steps take place one after the other to desorb 3.35, 2.28, and 1.10 wt % hydrogen as 3, 2, and 1 mol of H atoms, respectively. In total, ∼6.73 wt % hydrogen, equivalent to 6 mol of H atoms, is released from the 2NaH-LiAlH4 sample milled for 24 h, indicating all hydrogen in the Na2LiAlH6 sample is evolved when it is heated to 480 °C. For understanding well the chemical reactions occurring in the dehydrogenation process, samples at different dehydrogenation stages were collected for XRD examination. The results are shown in Figure 4. As mentioned above (see Figure 1), the sample milled for 24 h exhibits typical diffraction peaks of Na2LiAlH6. After dehydrogenated at 215 °C, the NaH, LiH, and Al phases all become discernible with very low intensity of diffraction peaks, while the Na2LiAlH6 phase is still dominating

(4)

This reaction is consistent with the amount of hydrogen desorbed as shown by filled squares in Figure 3a. In the meantime, it should be mentioned that a small amount of Na3AlH6 can still be detected at 215 and 230 °C, and it becomes undetectable at 250 °C, implying that Na3AlH6 is first formed and then consumed. A similar phenomenon has been reported by Mamatha et al.29 As described above, hydrogen is gradually released from Na2LiAlH6 during its transformation into NaH, LiH, and Al when heated to above 190 °C. At 215-230 °C, partial newly developed NaH reacts possibly with Na2LiAlH6 to form Na3AlH6 and LiH. With further increase of the temperature to 250 °C, Na3AlH6 decomposes into NaH, Al, and gaseous hydrogen. The reaction process proceeds as below.

Na2LiAlH6 + NaH f Na3AlH6 + LiH

(5)

Na3AlH6 f 3NaH + Al + 3H2

(6)

On the basis of XRD analyses, we believe that reaction 1 should be the primary reaction and reaction 5 is just a competitive by-reaction, which is responsible for the very low intensity of the Na3AlH6 phase in Figure 4. After being heated to 350 °C, the diffraction peaks of the LiH and Al phases still persist in the XRD profile. Metallic Na is clearly identified from its peaks at 29.3°, 42.0°, and 52.2°, with lowering intensity of peaks assignable to NaH, indicating the newly formed NaH is decomposed into metallic Na and H2 at the higher temperatures. Consequently, the second step dehydrogenation in the temperature range of 320-380 °C can be expressed by the following reaction:


1 NaH f Na + H2 2


(7)

According to the above formula, 2 mol of NaH decomposed to desorb 1 mol of hydrogen gas, which is in good agreement with the experimental result (Figure 3a). As the temperature was further elevated to 450 °C, metallic Na and LiAl are the only detectable phases along with the disappearance of the NaH, LiH, and Al phases. It can be concluded therefore that the chemical reaction between LiH and Al gives rise to the third-step dehydrogenation to result in LiAl and H2.

1 LiH + Al f LiAl + H2 2

(8)

Here, we point out that the previous opinion on the last-step dehydrogenation of Na2LiAlH6 as the self-decomposition of LiH should be corrected as no metallic Li was detected in the dehydrogenation product at 450 °C in the present study.29 In addition, investigations of Dilts et al. on LiAlH4 had also revealed that H in LiH was liberated by reacting with Al in the last-step dehydrogenation rather than its self-decomposition.30 Thus, it can be ascertained that the last 1 mol of H atom in the dehydrogenation process of Na2LiAlH6 comes from reaction 8. According to the above discussions, the dehydrogenation reaction of Na2LiAlH6 at different temperature ranges should be expressed by the following three-step reaction: 190 - 250◦C 3 Na2LiAlH6 98 2NaH + LiH + Al + H2 2 320 - 380◦C 5 98 2Na + LiH + Al + H2 2

(9)

380 - 480◦C

98 2Na + LiAl + 3H2 The above reaction releases theoretically 6.97 wt % of hydrogen, which agrees well with the experimental measurements as shown in Figure 3a. 3.3. Hydrogen Storage Reversibility of Na2LiAlH6. In principle, a thermodynamically reversible hydrogen storage process requires an endothermic nature for hydrogen desorption reaction.31 DSC measurement is well-known as a suitable and quick method for the determination of desorption heat.32 Figure 5 shows the DSC curve of a 2NaH-LiAlH4 sample milled for 24 h. Obviously, the DSC curve exhibits endothermic nature for hydrogen desorption, implying that the hydrogen desorption/ absorption of Na2LiAlH6 should be thermodynamically reversible. Three endothermic peaks were observed corresponding to the three-step dehydrogenation of reaction 9, consisting well with the results of TPD and volumetric release. The discrepancy in the peak temperatures originates from the different temperature ramping rates in the TPD (2 °C/min) and the DSC (10 °C/min) experiments. Hydrogenation/redehydrogenation experiments were then performed on the fully dehydrogenated sample up to 530 °C. The results are also presented in Figure 3. Figure 3b shows the hydrogen uptake curve of the dehydrogenated sample with increasing temperature at 135 atm of hydrogen pressure (symbol b). It can be seen that hydrogen uptake starts sluggishly around 100 °C. With a long enough holding time, the sample absorbs 6.6 wt % of hydrogen at 285 °C, very close to the amount of hydrogen evolved during its first-time dehydrogenation, which is obviously higher than those reported previously which focused only on the first-step hydrogen storage reaction of Na2LiAlH6.2,24,25

Figure 5. DSC curve of the 2NaH-LiAlH4 sample milled for 24 h.

The follow-up dehydrogenation measurement further displays that all hydrogen charged into the sample is redesorbed as shown by open circles in Figure 3a. These results confirm that the hydrogen storage of Na2LiAlH6 is completely reversible. In the meantime, it should be noticed that the operating temperature used for redehydrogenation of the hydrogenated sample is slightly higher than that required for the as-milled sample, especially for the first-step dehydrogenation reaction. We attribute the higher temperature required to the increased particle size of the hydrogenated sample. 3.4. Thermodynamics and Kinetics of the First-Step Dehydrogenation. It can be seen in Figure 3a that the operating temperatures of the second and third steps for the dehydrogenation of Na2LiAlH6 are both too high for practical applications. Therefore, detailed thermodynamic and kinetic investigations were presently focused only on the first-step dehydrogenation reaction. It is well-known that the thermodynamic parameters, the desorption enthalpy change (∆H), and entropy change (∆S) of a reversible hydrogen storage system can be calculated from the temperature-dependent equilibrium pressure (P) in the van’t Hoff equation:33

ln P )

1 -∆H ∆S + T R R

(

)

(10)

where T is the absolute temperature and R is the gas constant. Figure 6 is the van’t Hoff plot of the first-step dehydrogenation process of Na2LiAlH6 prepared by ball-milling a mixture of 2NaH-LiAlH4 for 24 h. The plot of ln P versus 1/T is a straight line, indicating a good van’t Hoff relationship between these two parameters. The resultant van’t Hoff equation is expressed numerically as

ln P )

-7685.3 + 18.3 T

(11)

Comparing eq 10 and eq 11, we calculated the enthalpy change (∆H) to be 63.8 kJ/mol-H2 and the entropy change (∆S) to be 152.1 J/K · mol-H2 for the first-step dehydrogenation reaction of Na2LiAlH6, indicating a relatively high thermal stability of the reaction, which should be the thermodynamic reason for the high operating temperature of dehydrogenation of Na2LiAlH6. Here we also noticed that the value of the enthalpy change is slightly larger than 62.8 kJ/mol-H2 reported by Claudy et al.,26 which is likely due to the difference in the measurement methods. By extrapolating the linear curve of Figure 6, an operating temperature of 145 °C was determined

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Wang et al. a straight line as shown in Figure 7b. From the slope of the straight line, the apparent activation energy Ea was determined to be approximately 172.6 kJ/mol for the first-step dehydrogenation of Na2LiAlH6, much higher than that of the other alanates as reported previously.9 For further elucidating the details and mechanism of the reaction kinetics, isothermal dehydrogenation experiments were carried out at 200-210 °C, and the rate-determining step was determined by analyzing the isothermal hydrogen desorption curves with the Johnson-Mehl-Avrami (JMA) equation shown below,35 which has been widely used in the study for NaAlH4, LiAlH4, and AlH39,36,37

R(t) ) 1 - exp[-(kt)n]

Figure 6. van’t Hoff plot for the first-step dehydrogenation of Na2LiAlH6.

for the equilibrium hydrogen pressure of 1 atm. However, hydrogen desorption was detected experimentally only at temperatures above 190 °C. The big difference of desorption temperatures indicates the presence of a relatively high kinetic barrier in dehydrogenation reaction. Quantitative estimation of kinetic barrier was then carried out by the determination of the apparent activation energy (Ea) with Kissinger’s method34

( )

ln

Ea β )2 RT Tm

(12)

where β is the heating rate in °C/min, Tm is the absolute temperature for the maximum desorption rate, and R is the gas constant. In the present study, Tm was extracted from the kinetic dehydrogenation curves at variable heating rates as shown in Figure 7a. Hydrogen desorption curves were shifted to higher temperatures with increasing the heating rate from 0.1 to 5 °C/ min as expected. Differentiating the hydrogen desorption curves, we obtained the Tm values to be 199, 210, 222, 234, and 242 °C corresponding to the heating rate of 0.1, 0.4, 1, 3, and 5 °C/min, respectively. The plot of ln(β/Tm2) versus 1/Tm is also

(13)

in which R(t) is the fraction already reacted at time t, k is the rate constant, and n is the Avrami exponent which reflects generally the nucleation and growth morphology, and the magnitude of which provides us with insights into the mechanism of dehydrogenation kinetics. In general, eq 13 is rearranged in a form more convenient to use as

ln[-ln(1 - R(t))] ) n ln(t) + n ln(k)

(14)

From eq 14, the plot of ln[-ln(1 - R(t))] versus ln t is expected to be a straight line at a given temperature, and the slope and intercept of the line can be used to calculate out n and k, respectively. Figure 8a shows the isothermal dehydrogenation curves of Na2LiAlH6 at 200, 203, 207, and 210 °C. All isothermal dehydrogenation curves exhibit the typical sigmoidal shape with a short induction period, followed by an acceleration period for dehydrogenation, and finally a decaying period. As the operating temperature increased, the dehydrogenation rate is distinctly speeded up and the duration for full dehydrogenation shortens. The first-step dehydrogenation completes in 60 min at 210 °C while it takes 140 min at 200 °C. To obtain the Avrami exponent n and the rate constant k, the isothermal dehydrogenation data were plotted as ln[-ln(1 R(t))] versus ln t as shown in Figure 8b. It can be clearly seen that ln[-ln(1 - R)] varies linearly with ln t over the range of 0.15 e R e 0.85 with a linearity constant of R2 > 0.997, which means that the JMA equation can well be used to describe the first-step dehydrogenation process of Na2LiAlH6. The values

Figure 7. Hydrogen desorption curves (a) of the first-step dehydrogenation of Na2LiAlH6 at various heating rates and the Kissinger’s plot (b).



Figure 8. Isothermal hydrogen desorption curves (a) of Na2LiAlH6 at 200-210 °C and the JMA plots (b).

TABLE 1: Kinetic Data for the Hydrogen Desorption of Na2LiAlH6 T (°C) 200 203 207 210

slope (n)

-ln k

k (×10-4 s-1)

2.36 2.14 2.33 2.11

8.39 8.27 7.93 7.57

2.27 2.56 3.59 5.17

of n obtained are in the range of 2.11-2.36 as shown in Table 1, suggesting the dehydrogenation process belongs to a diffusion-controlled reaction.38 As a consequence, the rate of hydrogen desorption from Na2LiAlH6 is limited by the diffusion of substance here rather than interface reaction. This result is in good agreement with recent investigations on other alanates.14 The rate constant k determined by the intercept of the straight line in Figure 8b is also listed in Table 1. As expected, the value of k is increased with temperature, indicating faster dehydrogenation rates at higher operating temperatures. The temperature dependence of k of chemical reactions can be described by the Arrhenius equation38

( )

-Ea k ) A exp RT

(15)

where A is the pre-exponential factor, Ea is the apparent activation energy, R is the gas constant, and T is absolute temperature. The natural logarithm of eq 15 gives a first-order linear equation on 1/T. By plotting ln k versus 1/T, Ea and A can be extracted from the linear slope and intercept. Figure 9 is the Arrhenius plot for the first-step dehydrogenation. A good linearity between ln k and 1/T is obtained with R2 ) 0.992. The apparent active energy (Ea) calculated is about 173.2 kJ/ mol, and the corresponding pre-exponential factor (A) is 2.37 × 1015 s-1. As a result, the rate equation of the first-step dehydrogenation of Na2LiAlH6 can be expressed as

k ) 2.37 × 1015e-173200⁄RT

(16)

Here, it is interesting to note that the value of Ea (173.2 kJ/ mol) derived from the Arrhenius equation is almost identical with that determined by Kissinger’s method (172.6 kJ/mol), a rather undesirably large value. It indicates that a relatively high kinetic barrier needs to be surmounted for hydrogen desorption

Figure 9. Arrhenius plot for the first-step dehydrogenation of Na2LiAlH6.

from Na2LiAlH6. We believe therefore that a reduction in the particle size,4,13,39-43 an increase in the crystal defects (such as dislocations, vacancies, and increased number of grain boundaries),14,39-43 and a suitable catalyst doping4,40-43 will be certainly helpful for the improvement of the hydrogen absorption/ desorption performances of Na2LiAlH6. 4. Conclusion Na2LiAlH6 was synthesized by ball-milling a mixture of 2NaH-LiAlH4. Reactions in the ball-milling process were first investigated. XRD examinations showed that a metathesis reaction first took place between NaH and LiAlH4 to convert to LiH and NaAlH4 in the initial ball-milling stage, and then the mixture of LiH, LiAlH4, and NaH reacted together to yield Na2LiAlH6. The thermodynamic and kinetic mechanisms for the hydrogen storage process of the synthesized Na2LiAlH6 were then systematically studied. Approximately 6.73 wt % of hydrogen could be released from the as-milled sample in the temperature range of 190-480 °C with a three-step reaction corresponding to the ordinal decomposition of Na2LiAlH6, the decomposition of NaH, and the reaction between LiH and Al.

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Hydrogen desorption exhibited endothermic nature with an enthalpy change of 63.8 kJ/mol-H2 for the first-step reaction. The hydrogen uptake experiment revealed that all of hydrogen could be recharged at 135 atm and 285 °C. Kinetic investigations indicated that a relatively high kinetic barrier needed to be surmounted for hydrogen desorption from Na2LiAlH6 with an apparent activation energy of about 173 kJ/mol. Diffusioncontrolled kinetic mechanism was determined for the first-step dehydrogenation of Na2LiAlH6. Consequently, reduction in the particle size, increase in the crystal defects, and catalyst doping can be used to improve the hydrogen aborption/desorption performances of Na2LiAlH6. Acknowledgment. We acknowledge financial support from the National Natural Foundation of China (Grant Nos. 50701040 and 50631020), from the National High-Technology Research and Development Plan (863 Program, Grant No. 2006AA05Z127), from Qianjiang Talent Project of Zhejiang Province (Grant No. QJD0702005), and from the Scientific Research Foundation of the State Education Ministry for Returned Overseas Chinese Scholars. We thank the reviewers for their insightful comments and suggestions. References and Notes (1) Schlapbach, L.; Zu¨ttel, A. Nature (London) 2001, 414, 353–358. (2) Bogdanovic´, B.; Schwickardi, M. J. Alloys Compd. 1997, 253254, 1–9. (3) Grochala, W.; Edwards, P. P. Chem. ReV. 2004, 104, 1283–1315. (4) Bogdanovic´, B.; Eberle, U.; Felderhoff, M.; Schu¨th, F. Scr. Mater. 2007, 56, 813–816. (5) Orimo, S.; Nakamori, Y.; Eliseo, J. R.; Zu¨ttel, A.; Jensen, C. M Chem. ReV. 2007, 107, 4111–4132. (6) Jensen, C. M.; Zidan, R.; Mariels, N.; Hee, A. G.; Hagen, C. Int. J. Hydrogen Energy 1999, 23, 461–465. (7) Chen, J.; Kuriyama, N.; Xu, Q.; Takeshita, H. T.; Sakai, T. J. Phys. Chem. B 2001, 105, 11214–11220. (8) Wang, J.; Ebner, A. D.; Ritter, J. A. J. Am. Chem. Soc. 2006, 128, 5949–5954. (9) Kircher, O.; Fichtner, M. J. Alloys Compd. 2005, 404-406, 339– 342. (10) Bogdanovic´, B.; Felderhoff, M.; Pommerin, A.; Schu¨th, F.; Spielkamp, N. AdV. Mater. 2006, 18, 1198–1201. (11) Wang, J.; Ebner, A. D.; Ritter, J. A. J. Phys. Chem. C 2007, 111, 14917–14924. (12) Zheng, S. Y.; Fang, F.; Zhou, G. Y.; Chen, G. R.; Ouyang, L. Z.; Zhu, M.; Sun, D. L. Chem. Mater. 2008, 20, 3954–3958. (13) Balde´, C. P.; Hereijgers, B. P. C.; Bitter, J. H.; de Jong, K. P. J. Am. Chem. Soc. 2008, 130, 6761–6765.

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