J. Phys. Chem. C 2008, 112, 131-138
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Comparison of the Calculated and Experimental Scenarios for Solid-State Reactions Involving Ca(AlH4)2 Nobuko Hanada,* Wiebke Lohstroh, and Maximilian Fichtner Institut fu¨r Nanotechnologie, Forschungszentrum Karlsruhe, Postfach 3640, D-76021 Karlsruhe, Germany ReceiVed: June 12, 2007; In Final Form: September 21, 2007
Solid-state reactions of Ca(AlH4)2 and various additives were investigated experimentally for hydrogen desorption. The Ca(AlH4)2 + Si, Ca(AlH4)2 + 2MgH2, Ca(AlH4)2 + 2LiH, and Ca(AlH4)2 + 2LiNH2 systems were chosen among reactions proposed theoretically1 to study their hydrogen storage capacity and an appropriate reaction enthalpy. For all systems investigated, the reversible reactions proposed should have more than 6.5 mass % hydrogen capacity and reaction enthalpies in the range of 30-55 kJ/mol H2. However, most of the experimentally observed reactions do not conform to theoretical propositions because different final products were obtained in all cases but one. Two tendencies were observed in the experiments. One is that the reaction comprises several distinct steps at different temperatures as observed in the Ca(AlH4)2 + Si and Ca(AlH4)2 + 2 MgH2 systems. In this case, Ca(AlH4)2 decomposes first and is followed by a reaction of the remaining compounds. The same kinetic reaction barriers are encountered here as in pure Ca(AlH4)2. Second, LiH and LiNH2 additions yield exothermic reactions of Ca(AlH4)2 and the other reactant as early as in the ball milling or annealing steps, leading to final products not considered in the reactions calculated.
3/ H ). 2 2
1. Introduction Hydrogen storage in solids is currently considered one of the most important options for on-board hydrogen storage because the familiar physical limits to storing a compressed hydrogen gas or a cryocooled liquid prevent targets from being met with respect to gravimetric and volumetric storage capacities.2 This makes the development of hydrogen storage materials of high hydrogen capacity and rapid kinetics at low decomposition temperature an important task in support of budding hydrogen fuel cell technologies. Complex aluminum hydrides, such as NaAlH4, have been considered promising materials for reversible hydrogen storage because transition-metal dopants were found to decrease the kinetic barriers considerably for both hydrogen absorption and desorption.3 Thus, reversible hydrogen uptake and release is possible at moderate temperatures and pressures (e.g., 150 °C, 100 bar H2). Other alanates with high hydrogen contents have been investigated for their hydrogen sorption properties, for example, Mg(AlH4)2 with 7.0 mass % hydrogen (decomposition to MgH2 and 2 Al) or Ca(AlH4)2 with 5.9 mass % H2 (decomposition to CaH2 and 2 Al).4-8 The thermodynamic stability of Mg(AlH4)2 is too low for reversible hydrogen storage (ref 7). For practical application, the storage material should contain a large amount of reversibly stored hydrogen (gravimetric density above 6 mass %) with a decomposition enthalpy in the range of 30-40 kJ/ mol H2. The latter requirement ensures that the working temperature of the storage material is adapted to PEM fuel cells in mobile applications. So far, singlecomponent systems have not been able to meet the requirements. Especially, most complex hydrides of high storage capacity are quite stable, for example, LiBH4 with 13.9 mass % and 67 kJ /mol H29 (decomposition into the binary hydride, LiH + B + * Corresponding author. E-mail:
[email protected]. Tel: +49-7247-82-6377. Fax: +49-7247-82-6368.
Several reactions applying a destabilization concept to complex hydrides have been studied recently. Destabilization means that thermodynamically more stable products are generated in the dehydrogenated state by having the complex hydride react with a second component. In this way, the reaction enthalpy can be reduced compared to the decomposition enthalpy of the pure material. The pioneering work by Reilly and Wiswall on hydrogen storage provides the first example of destabilization.10 In their work, MgH2 is destabilized by Cu addition, with Mg2Cu forming upon dehydrogenation. Recently, Vajo et al. used Si as an additive to LiH and MgH2, thus producing stable compounds, such as Li4Si and Mg2Si, during dehydrogenation of the respective hydride.11 Moreover, adding MgH2 to LiBH4 leads to the production of MgB2 and LiH in the dehydrogenated state. The reaction enthalpy of the latter system is lower than the decomposition enthalpy of pure LiBH4. Reversible hydrogenation and dehydrogenation of LiBH4 is observed at lower temperatures in the mixed system.9 The solidstate reactions of LiH with LiNH2, Mg(NH2)2, and Ca(NH2)2 were studied as well.12-19 The destabilization reactions of LiNH2 with LiAlH4 or Li3AlH6 were explored: the reaction of LiNH2 with Li3AlH6 featured hydrogen reversibility.20-22 These results suggest that there are indeed opportunities for lowering the working temperature in solid-state reactions of hydrides. Alapati et al. performed first-principles density functional theory (DFT) calculations for a large number of potential materials1,23,24 to approximate the enthalpy of formation (at 0 K in the absence of zero-point energy corrections, ∆Hf ≈ ∆Uf (T ) 0 K), where ∆Uf is the change in total electronic energy). From these data, they obtained the reaction enthalpy, ∆Hr , for a number of potential solid-state reactions. Several promising reactions involving Ca(AlH4)2 are shown in their paper in terms of hydrogen content and reaction enthalpies. Mamatha et al. studied the desorption properties of pure Ca(AlH4)2 by DSC and in situ XRD measurements.7,8 Decomposi-
10.1021/jp074534z CCC: $40.75 © 2008 American Chemical Society Published on Web 12/13/2007
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Hanada et al.
Figure 1. DSC measurements of Ca(AlH4)2 (+ 2NaCl) in 3 bar of H2 (solid line) and 3 bar of He (dashed line) (heating rate 5 K/min).
Figure 2. DSC measurement of Ca(AlH4)2 (+ 2NaCl) + Si ball milled for 2 h in 3 bar of He.
tion comprised the following three reactions: Ca(AlH4)2 f CaAlH5 + Al + 3/2H2 f CaH2 + 2Al + 3H2 f Al2Ca + 4H2. According to the DSC results, the reaction enthalpies of the first and second steps are -7 kJ/mol H2 and +32 kJ/mol H2, respectively.7 This result indicates a possibility of reversible rehydrogenation in the second step. However, this has not yet been achieved experimentally. The total enthalpy change amounts to 12 kJ/mol H2 for the decomposition to CaH2 and Al.1 Alapati et al. suggest various reactions involving Ca(AlH4)2 with reaction enthalpies slightly higher than the decomposition reaction.1 They argued that these reactions could be observed in experiments if the decomposition in the pure system is limited by the high kinetic barrier. The authors studied four different combinations of materials out of these proposals: Ca(AlH4)2 + Si, Ca(AlH4)2 + 2MgH2, Ca(AlH4)2 + 2LiH and Ca(AlH4)2 + 2LiNH2 in solid-state reactions, see Table 1. The results of hydrogen desorption in the reactions proposed and the effectiveness of theoretical calculations in predicting suitable solid-state reactions will be discussed.
The vial was sealed in an argon-filled glovebox and ball milled for 2 h at 400 rpm in a Fritsch P6 planetary mill. For structural characterization, X-ray powder diffraction was performed in a Philips X’PERT diffractometer (Cu KR radiation). The powder was spread on a silicon single crystal and sealed in the glovebox with an airtight hood of Kapton sheet. To examine the hydrogen desorption properties, a HP-DSC 204 Netzsch high-pressure differential scanning calorimeter (DSC) was used at 3 bar H2 or 3 bar He static pressure at a heating rate of 5 K/min. In addition, thermogravimetry (TG) and mass spectrometry of the gas phase (TG-MS) were performed in a He flow at a heating rate of 5 K/min in a Netzsch STA 409C analyzer equipped with a Balzers quadrupole mass spectrometer for analysis of the evolved gas produced. All powders were handled in an argon-filled glovebox with partial pressures of oxygen and water below 0.1 ppm.
2. Experimental Procedures Calcium alanate Ca(AlH4)2 was prepared by ball milling CaCl2 with 2NaAlH4, as described in ref 6. The product of the composition Ca(AlH4)2 + 2NaCl was used without further purification or removal of NaCl. Subsequently, Ca(AlH4)2 (+ 2NaCl) was mixed by ball milling with other materials, such as Si (powder produced from a Si wafer by ball milling for 2 h), MgH2 (Aldrich, 90% purity), LiH (Aldrich, 30 mesh powder, 95% purity), and LiNH2 (Merck, 98% purity). The molar ratios of each combination are shown in Table 1. For each of combination studied, 1 g of the mixture was put into an 80 mL silicon nitride vial with 5 silicon nitride balls (1 ball 20 mm in diameter and 4 balls 10 mm in diameter).
3. Results For comparison with combined Ca(AlH4)2 systems, the DSC measurements of Ca(AlH4)2 (+ 2NaCl) under 3 bar of H2 and 3 bar of He are shown in Figure 1. The signal obtained under 3 bar of He is similar to that reported by Mamatha et al.8 The exothermic peak at around 150 °C is due to the decomposition of Ca(AlH4)2 to CaAlH5 + Al + 3/2H2. The endothermic peak around 250 °C is the decomposition of CaAlH5 to CaH2 + Al + 3/2H2. The second and third endothermic peaks at around 300-400 °C are due to the reactions of CaH2 and Al to 1/ Al Ca + 1/ CaH and to Al Ca and hydrogen, respectively. 2 4 2 2 2 These last two reactions are suppressed in a hydrogen atmosphere of 3 bar as shown in Figure 1. NaCl is still present in the material and does not take part in the reactions. 3.1. Reaction of Ca(AlH4)2 + Si. Figure 2 shows the DSC data obtained at 3 bar of He for Ca(AlH4)2 (+ 2NaCl) + Si, ball milled for 2 h. The DSC trace is similar to that obtained
TABLE 1: Proposed Reactions of Ca(AlH4)2 and Additives, Theoretical Amount of Stored H2 [mass % H2], and Calculated Reaction Enthalpy, ∆Hr [kJ/mol H2], See Reference 1a
a
calculated reactions involving Ca(AlH4)2
mass % H2
∆H
experimental composition
Ca(AlH4)2 + Si f 2Al + CaSi + 4H2 (1) Ca(AlH4)2 + 2MgH2 f CaMg2 + 2Al + 6H2 (2) 6Ca(AlH4)2 + 17MgH2 f Al12Mg17 + 6CaH2 + 35H2 (3) Ca(AlH4)2 + 2LiH f 2AlLi + CaH2 + 4H2 (4) Ca(AlH4)2 + 2LiNH2 f CaH2 + 1/2LiN3 + 3/2AlLi + 1/2AlN + 5H2 (5)
6.30 7.83 6.67 6.85 6.82
30.7 51.6 35.4 33.3 43.4
Ca(AlH4)2 + Si Ca(AlH4)2 + 2MgH2
The last column shows the experimental compositions tested for this paper.
Ca(AlH4)2 + 2LiH Ca(AlH4)2 + 2LiNH2
Solid-State Reactions Involving Ca(AlH4)2
J. Phys. Chem. C, Vol. 112, No. 1, 2008 133
Figure 3. XRD profiles of Ca(AlH4)2 (+ 2NaCl) + Si ball milled for 2 h and heated to 450 °C in 3 bar He. The peak positions of NaCl are shown by the dashed lines. The peak positions of the ICDD powder diffraction files and Ca(AlH4)2 obtained from the structural data of ref 25 are included for comparison.
Figure 4. DSC measurement of (a) Ca(AlH4)2 (+ 2NaCl) + 2MgH2 after ball milling for 2 h in H2 at 3 bar (solid line) and He at 3 bar (dashed line), and (b) Ca(AlH4)2 (+ 2NaCl) + 2(MgH2 + 1 mol % V2O5) after ball milling for 2 h under H2 at 3 bar (solid line) and He at 3 bar (dashed line).
with pure Ca(AlH4)2 (+ 2NaCl) at 3 bar of He shown in Figure 1. The XRD profiles in Figure 3 seem to indicate that, after ball milling, the sample contains only the starting materials, Ca(AlH4)2, NaCl, and Si. After heating to 450 °C (i.e., after DSC measurement), Al2Ca, a small amount of Al4Ca, and Si are observed. Under the experimental conditions chosen, Si does not react with CaH2 or Ca-related phases. After the two-step decomposition of Ca(AlH4)2 to CaH2 and Al, CaH2 reacts with Al instead of reacting with Si. The DFT-calculated reaction enthalpies for
CaH2 + Si f CaSi + H2
(6)
CaH2 + 2Al f Al2Ca + H2
(7)
and
Figure 5. XRD profiles of Ca(AlH4)2 (+ 2NaCl) + 2MgH2 after ball milling for 2 h and heating to 200, 270, and 400 °C under He at 3 bar (i.e., after the DSC measurements). The peak positions of the ICDD powder diffraction files, Ca(AlH4)2 and CaAlH5, are included for comparison.
are 69 and 70 kJ/mol H2, respectively, calculated from the enthalpies of formation at 0 K supplied in ref 1. For the two reactions (6 and 7), the reaction enthalpies are almost identical. However, only the latter is observed, which may be an indication of better reaction kinetics of reaction 7. In fact, CaSi can be obtained from ball milled CaH2 + Si after heating to 600 °C for 5 h under a vacuum.26 There is no indication that the calculated reaction for the mixture of Ca(AlH4)2 + Si (see Table 1, eq 1) actually occurs under our experimental conditions. 3.2. Reaction of Ca(AlH4)2 + 2MgH2. Figure 4a shows the DSC data of measurements performed under 3 bar of He and H2 for Ca(AlH4)2 (+ 2NaCl) + 2MgH2 ball milled for 2 h. Up to 270 °C, the same reactions occur as in pure Ca(AlH4)2 (+ 2NaCl) in both the He and H2 atmospheres. After milling, the sample consists of a mixture of Ca(AlH4)2, NaCl, and MgH2. The subsequent reactions can be traced by the XRD profiles taken after annealing the sample to various temperatures in the DSC (3 bar He, see Figure 5). After heating to 200 °C, CaAlH5 and Al phases appear from the decomposition of Ca(AlH4)2. At 270 °C, CaAlH5 decomposes completely to CaH2 and Al, while MgH2 still remains in the sample. Above that temperature, behavior in He and H2 atmospheres varies: At around 270350 °C, there is an endothermic peak in a He atmosphere, which shifts to a higher temperature when measured in a H2 atmosphere. As shown in Figure 5, the sample heated to 400 °C in He contains CaH2 and small amounts of Al and Al12Mg17 (with a peak position slightly shifted to lower angles compared to the ICDD data of Al12Mg17 no. 01-1128). Apparently, MgH2 is decomposed in that temperature range and reacts with Al to form Al12Mg17 at around 270-350 °C. Other reactions may occur at the same time because two unknown peaks at 35° and 38° are observed in the XRD profile. The reaction of MgH2 + Al is shifted to higher temperatures when measured in a hydrogen atmosphere. The reactions observed experimentally in Ca(AlH4)2 + 2MgH2 thus can be summarized as follows:
Ca(AlH4)2 + 2MgH2 f CaAlH5 + Al + 2MgH2 + 3/2H2 140 °C (8)
134 J. Phys. Chem. C, Vol. 112, No. 1, 2008
f CaH2 + 2Al + 2MgH2 + 3H2 250-270 °C
Hanada et al.
(9)
f CaH2 + 2/17Al12Mg17 + 10/17Al + 5H2 270-350 °C (10) The two decomposition steps of Ca(AlH4)2 and the formation of Al12Mg17 (together with the decomposition of MgH2) occur at distinct temperatures. Because Al is already available from the decomposition of Ca(AlH4)2 at low temperatures, the production of Al12Mg17 could depend on the reactivity of MgH2 because uncatalyzed MgH2 is known to have slow hydrogen desorption kinetics. To improve the hydrogen desorption kinetics for MgH2, MgH2 catalyzed with 1 mol % V2O5 (99.9+% metal) was used next. V2O5 is known to be one of the catalysts improving the hydrogen sorption kinetics of MgH2.28,29 MgH2 + 1 mol % V2O5 was ball milled for 20 h and subsequently mixed with Ca(AlH4)2 (+ 2NaCl). As shown in Figure 6, the mixture ball milled for 2 h contains only Ca(AlH4)2 and MgH2 phases. The results obtained for decomposition of the catalyzed sample in a 3 bar He atmosphere will be discussed below; the corresponding DSC measurement is plotted in Figure 4b. The DSC profile shows an endothermic peak at around 250300 °C. At this stage, CaH2 and Al12Mg17 phases are found in the XRD pattern, see Figure 6. This means that the addition of a catalyst yields Al12Mg17 below 300 °C. Because the corresponding endothermic peak in the DSC profile exhibits a shoulder at 270 °C, presumably two endothermic events are superimposed upon each other. It is likely that the reaction of MgH2 + Al is shifted to lower temperatures by the effect of the catalyst on MgH2. Because the XRD profile also shows pure Al and Mg phases, the reaction to Al12Mg17 was not complete. Another small endothermic peak appears at around 330 °C. After heating to 400 °C (Figure 6), the XRD data show a CaH2 phase disappearing and a diminishing Al peak intensity, while unknown peaks appear at the same time. These were also observed in the sample without a catalyst after heating to 400 °C. Moreover, the peak position of Al12Mg17 shifts to lower angles, as in the sample without a catalyst. It is concluded from these results that CaH2, Al, and Al12Mg17 react to form an unknown phase at temperatures above 300 °C in a He atmosphere. TG-MS results show a weight loss of 6.1 mass % (this value is recalculated by using the weight without 2NaCl) until 400 °C with just hydrogen desorption occurring in the same temperature range as the DSC signals. Decomposition under 3 bar of static hydrogen pressure produces slightly different results: The DSC profile (Figure 4b) shows that the endothermic reaction, starting at around 250 °C, extends to ∼350 °C. Even if the kinetics of MgH2 is improved by the catalyst, the reaction temperature depends on hydrogen pressure because of the thermodynamics of MgH2. However, the decomposition temperatures of Ca(AlH4)2 and CaAlH5 are the same in He at 3 bar and 3 bar static H2 pressure. Hence, these temperatures mainly reflect the kinetic barriers to decomposition rather than the thermodynamics. In the XRD profile, after heating to 370 °C and cooling under H2 (not shown), CaH2, Al3Mg2, and MgH2 phases can be assigned. Moreover, an exothermic peak at around 300 °C is observed in the DSC measurement (not shown) during cooling in a H2 atmosphere. Apparently, Al12Mg17 or Mg phases are rehydrogenated to Al3Mg2 and MgH2 during cooling, as reported in refs 30-32. Even under 100 bar of hydrogen at 300 °C, complete reversibility and reformation of Ca(AlH4)2 or CaAlH5 is not observed in this Ca-Mg-Al-H system.
Figure 6. XRD profiles of Ca(AlH4)2 (+ 2NaCl) + 2(MgH2 + 1 mol % V2O5) after ball milling for 2 h and heating to 310 and 400 °C under He at 3 bar by using DSC.
To summarize, two-step decomposition of Ca(AlH4)2 is followed by an independent reaction of MgH2 with Al. Starting from CaH2, MgH2, and Al, these reactions are to occur
(i) 17MgH2 + 12Al f Al12Mg17 + 17H2
(11)
2MgH2 + 3Al f Al3Mg2 + 2H2
(12)
(ii) CaH2 + 2MgH2 f CaMg2 + 3H2
(13)
(iii) CaH2 + 2Al f Al2Ca + H2
(14)
with calculated enthalpies of 57, 89, and 70 kJ/mol H2 for reactions 11, 13, and 14, respectively (using enthalpies of formation at 0 K supplied in ref 1). For reaction 12, an enthalpy of +63.4 kJ/mol H2 is reported from hydrogenation experiments of Al3Mg2.33 Among these, reaction 11 has the smallest enthalpy change, and this reaction is also observed experimentally after the two-step decomposition of Ca(AlH4)2. Accordingly, the final products are in agreement with the calculated reaction in Table 1, eq 3. It is not the second reaction in the table that was anticipated from the stoichiometric ratio of the starting materials. However, the reaction temperature in the experiment is too high to reflect the calculated enthalpy change of 35.4 kJ/mol H2 for which, thermodynamically, hydrogen desorption can be expected at around 0 °C under 1 bar of H2. One possible reason is that the same kinetic barrier is effective in the combined system, which hinders CaAlH5 decomposition at its expected decomposition temperature from thermodynamics (an enthalpy change of 29.8 kJ/mol H234 decomposition is expected of at around -40 °C under 1 bar H2). Kinetic enhancement of MgH2 by means of a catalyst only decreases the temperature of the MgH2 + Al reaction and has no influence on the two decomposition steps of Ca(AlH4)2. Thus, there is no beneficial kinetic effect on the proposed reaction compared to pure Ca(AlH4)2. This means that, although the overall reaction enthalpy, considering the starting materials and the final products, could be on the order of the calculated value of 32 kJ/ mol H2, (value calculated according to the experimental composition, see reaction 10), it may be of limited practical use because each of the reaction steps takes place at a different temperature. In our linear heating experiments, the energy gain from exothermic step 1 has already
Solid-State Reactions Involving Ca(AlH4)2
Figure 7. DSC measurement under He at 3 bar for Ca(AlH4)2 (+ 2NaCl) + 2LiH ball milled for 2 h.
J. Phys. Chem. C, Vol. 112, No. 1, 2008 135 Al. The lattice parameters of these two phases are similar and cannot be distinguished in our XRD data. Furthermore, Al has high solubility for Li metal (at temperatures
