J. Phys. Chem. B 2007, 111, 13301-13306
13301
Determination of the Heat of Hydride Formation/Decomposition by High-Pressure Differential Scanning Calorimetry (HP-DSC) Carine Rongeat,† Isabel Llamas-Jansa,† Stefania Doppiu,† Stefano Deledda,‡ Andreas Borgschulte,§ Ludwig Schultz,† and Oliver Gutfleisch*,† IFW Dresden, Institute for Metallic Materials, P.O. Box 270116, D-01171 Dresden, Germany, and Institute for Energy Technology, Department of Physics, P.O. Box 40, N-2027 Kjeller, Norway, EMPA, Section 138: Hydrogen and Energy, U ¨ berlandstrasse 129, CH-8600 Du¨bendorf, Switzerland ReceiVed: July 27, 2007; In Final Form: September 7, 2007
Among the thermodynamic properties of novel materials for solid-state hydrogen storage, the heat of formation/ decomposition of hydrides is the most important parameter to evaluate the stability of the compound and its temperature and pressure of operation. In this work, the desorption and absorption behaviors of three different classes of hydrides are investigated under different hydrogen pressures using high-pressure differential scanning calorimetry (HP-DSC). The HP-DSC technique is used to estimate the equilibrium pressures as a function of temperature, from which the heat of formation is derived. The relevance of this procedure is demonstrated for (i) magnesium-based compounds (Ni-doped MgH2), (ii) Mg-Co-based ternary hydrides (Mg-CoHx) and (iii) Alanate complex hydrides (Ti-doped NaAlH4). From these results, it can be concluded that HP-DSC is a powerful tool to obtain a good approximation of the thermodynamic properties of hydride compounds by a simple and fast study of desorption and absorption properties under different pressures.
1. Introduction Nowadays, there is an increasing awareness of the necessity to replace our energy sources based on fossil fuels. These fuels have a big impact on pollution and climate change and will only be available for a short period of time. Hydrogen could be an ideal new energy carrier on which a complete new energy infrastructure can be built. Regarding the use of hydrogen as fuel for the zero-emission vehicle, the main technological problem is the storage of hydrogen.1 Solid-state storage offers a safe alternative to storage in compressed or liquid form. Among solid-state storage materials, a large variety of metal hydrides or complex hydrides exists that have high storage capacities by volume or/and by weight. For mobile applications, hydrides have to fulfill various requirements. The required thermodynamic properties originate from the practical conditions necessary to use hydrogen storage materials. This means essentially an operation temperature (T) up to 100-150 °C and H2 pressures (p) around atmospheric pressure. The van’t Hoff equation
ln p ) -
∆Habs,des ∆S + RTabs,des R
the (nano-)structure of the material, the nucleation and growth mechanisms, the diffusion, and the catalytic properties. The experimental determination of ∆Habs,des and ∆S by volumetric ab/desorption experiments is difficult, time-consuming, and subject to errors (kinetics). The usual experimental method is pressure-composition isotherm (PCI) measurements, where the variation of the pressure with hydrogen content is determined at different temperatures. At sufficient kinetics and appropriate slow measurements, an equilibrium between hydrogen in the gas phase and in the (metal) hydride can be assumed, and thus the equilibrium pressure of hydride formation/ decomposition is obtained. Differential scanning calorimetry (DSC) measurements can provide thermodynamic data such as reaction enthalpy, miscibility gap width, and pressure-temperature relation, as well as kinetics data such as activation energy and hysteresis.2-4 DSC is a nonequilibrium measurement, and thus the directly exchanged heat does not resemble the heat of formation/decomposition. Nevertheless, in this work we show how approximate reaction enthalpies can be extracted from isobaric DSC measurements. 2. Experimental Methods
(1)
relates the thermodynamic conditions to the heat of hydride formation/decomposition ∆Habs,des and the corresponding entropy ∆S. R is the gas constant (8.314 J K-1 mol-1), and T is the absolute temperature. In addition, the kinetics for hydrogen release and absorption has to be sufficiently fast to fulfill the technological constraints. These properties are determined by * Corresponding author. E-mail:
[email protected]. † IFW Dresden, Institute for Metallic Materials, Dresden, Germany. ‡ IFE, Institute for Energy Technology, Kjeller, Norway. § EMPA, Section 138: Hydrogen and Energy, Du ¨ bendorf, Switzerland.
2.1. Material Synthesis. All of the materials used in this study were prepared by reactive ball milling under hydrogen atmosphere. The powders were handled in a glove box under Ar atmosphere. The Mg99Ni1 hydride was synthesized by mixing elemental Mg (Alfa Aesar, 99.8%, 325 mesh) and Ni (MaTeck GmbH, 99.99%, 100 mesh) powders in the nominal composition. The mixture was then milled under 90 bar H2 during 12 h using a planetary mill (Fritsch pulverisette P6). A dedicated highpressure milling vial was used. This vial (evico-magnetics) provides the possibility of monitoring the pressure and temperature during milling. The ball-to-powder ratio was 10:1, with 10-mm-diameter stainless-steel balls. The Mg-Co-H ternary compound was obtained by milling elemental Mg and Co with
10.1021/jp075954r CCC: $37.00 © 2007 American Chemical Society Published on Web 10/31/2007
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Rongeat et al.
Figure 1. Determination of equilibrium temperature: example of DSC traces for Na3AlH6 (5 K min-1).
a 2:1 ratio under 5.5 bar H2 and during 60 h. Finally, the doped sodium alanate was prepared in a one-step synthesis. NaH (Aldrich, 95%) and Al (MaTeck, -100 + 200 mesh) powders (1:1) were mixed with 4% mol TiCl3 powder (Aldrich, 99.999%). The mixture was milled during 5 h under 100 bar H2 with a ball-to-powder ratio of 50:1 (10-mm-diameter balls) using a planetary mill (Fritsch pulverizette P6). 2.2. Differential Scanning Calorimetry (DSC) Measurements. DSC measurements were performed using a highpressure DSC apparatus (HP-DSC) from Netzsch (DSC 204 HP Phoenix, range 1-150 bar) placed inside a dedicated glove box under Ar. The heating rate was 5 K min-1 for all the samples. Note that in a preliminary DSC study using different heating rates we observed only small differences in the onset temperature (not shown). The 5 K min-1 rate was considered to be a good compromise between the quality of the measured signal and the time necessary. The hydrogen flow was adapted during measurement to maintain a constant pressure (dynamic mode). For magnesium compounds, the temperature range was fixed between room temperature (RT) and 500 °C, for sodium alanate between RT and 300 °C. The onset temperatures were determined with the Netzsch Proteus software using the derivative of the curve or, in some cases, by taking the intersection between the background line and the slope of the peak. 3. Thermodynamic Considerations In a first approximation, the formation/decomposition of a hydride can be regarded as a simple exothermic/endothermic reaction, similar to, for example, the melting or the structural transformation of a compound. The distinction of such processes is difficult from the DSC traces obtained under constant pressure. As an example, the DSC traces of Ti-doped Na3AlH6 under 120 bar H2 are given in Figure 1. This figure shows peaks attributed to a structural transformation of Na3AlH6 (R to β phases) at 250 °C (Tstruc eq ) and the ab- and desorption of des hydrogen at 277 °C (Tabs eq ) and 288 °C (Teq ), respectively. Assuming quasi-equilibrium, the Gibbs free energy ∆G for this process is zero; that is
0 ) ∆G ) ∆H - Teq∆S
(2)
where ∆H and ∆S are the corresponding heat and entropy of reaction, respectively, and Teq is the equilibrium temperature, for example, the melting temperature. Note that Teq is the temperature at which a reaction will start considering only thermodynamics (in most cases an extrapolation). For DSC, such a temperature is assumed to be the onset temperature of the
peak. This temperature represents the starting of the reaction and the point where ∆G becomes different from zero (eq 2). The described situation is fulfilled in most phase transformations (melting and crystal structure transformations) as shown in Figure 1. In the DSC trace, the peak width corresponding to the structure transformation is less than two degrees, and no difference between Teq for heating and cooling is found. In this case, ∆H can be derived directly from the measured ∆Q. The situation is different for hydrogen sorption. The corresponding peak width is approximately 10°, and the extrapolated equilibrium temperatures for heating and cooling differ. The reason for these differences is the relatively slow kinetics for sorption process compared to structural transformations. Therefore, simple equilibrium conditions cannot be assumed and thus the measured heat exchange does not resemble ∆Q. In this paper, we present a way of using kinetic DSC measurements to estimate the equilibrium properties of hydrides. The idea is as follows: the DSC signal is used to determine the temperature at which the ab-/desorption starts. However, in many hydride systems the onset (at which a signal is visible) is determined by kinetics instead of by thermodynamics. A typical example is the desorption process of nanocrystalline MgH2. This material starts to desorb significant amounts of hydrogen at temperatures above 200 °C under dynamic vacuum,5 although the thermodynamic equilibrium temperature should be around 100 °C in these conditions.6 The desorption is limited by kinetic constraints, which can be overcome by using appropriate catalysts.5,7 In general, the experimental values of Teq (absorption des Tabs eq , desorption Teq ) obtained by an extrapolation of the DSC signal to zero, give the range in which the true Teq lies (Figure 1): des Tabs eq < Teq < Teq
(3)
Even in the case of PCI measurements, a hysteresis exists between absorption and desorption and the equilibrium conditions are never fully reached. The equilibrium plateau can be also considered to be located between the absorption and desorption values obtained from PCI curves.8 From HP-DSC experiments, it is possible to extract ∆H from the van’t Hoff plots using Tieq measured under various hydrogen pressures p. Moreover, by including hysteresis in the determination of Teq, the true ∆H value is expected to be between the slopes for absorption and desorption in the van’t Hoff plot:
slopeabs < ∆H < slopedes
(4)
4. Results and Discussions To confirm the validity of heat values obtained through the van’t Hoff plot derived from HP-DSC, we applied the same experimental procedure to different classes of hydrides: (i) magnesium-based hydride, (ii) Mg-Co-H ternary system, and (iii) sodium alanate complex hydride. For these compounds, the hydrogen bonding and reaction paths are fundamentally different. DSC experiments were performed under different hydrogen pressures depending on the compound. The results are summarized and compared with literature data (see Table 1). 4.1. Magnesium-Based Hydride. Mg99Ni1 was synthesized by high-pressure ball milling from a mixture of elemental Mg and Ni. XRD patterns measured after 12 h milling mainly displayed the presence of magnesium hydride β-MgH2.9 DSC experiments were done between room temperature and 500 °C under 10, 20, 30, and 40 bar hydrogen. The DSC traces obtained are shown in Figure 2. For each curve, the heating process gives
Heat of Hydride Formation/Decomposition
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TABLE 1: Summary of the Results Obtained in This Study for Different Reactionsa this study reactions Mg99Ni1Hz / Mg99Ni1
literature
slopeabs slopedes ∆Habs ∆Hdes (kJ/mol) (kJ/mol) (kJ/mol) (kJ/mol) ref 68.1
80.8
MgH2 / Mg + 2 wt % Ni
74 -68.2 -67
MgH2 / Mg
-74.5 -75 -75 -74.5
-78 70 -69.5 Mg2CoH5 / MgCo
NaAlH4 / Na3AlH6
Na3AlH6 / NaH a
72.5
-83.2 -79
11 13 4 12 15 7 6
76
13 12 11 8 14 26 8 26 14
75.4
?
28.4
37 -45.5 -36.7
35.2
51.9
47 -69.6 -89
Enthalpies from literature are also given for comparison.
Figure 2. DSC traces under different pressures for Mg99Ni1 prepared by ball-milling. Inset: Details of DSC trace under 40 bar H2. Arrows indicate the peaks related to metastable phases.
rise to a large endothermic peak ascribed to the desorption of MgH2.9 The onset temperature of this peak shifts to higher values when the hydrogen pressure increases from ∼380 °C under 10 bar H2 to ∼450 °C under 40 bar. The results are consistent with the value given by Borgschulte et al.4 for catalyzed MgH2 under 1 bar H2 (peak starting at ca. 320 °C). In the same way, only one exothermic peak is observed during cooling. This peak is related to the absorption of hydrogen in magnesium and a temperature shift is also observed from ca. 360 to 425 °C when increasing the applied pressure. Note that in this case the onset temperature for absorption is slightly lower than that for desorption (shift of ca. 20 °C). The difference points out the hysteresis existing for magnesium hydrides.4 Additionally for desorption, the peaks show small shoulders likely related to the decomposition of the metastable hydride phase γ-MgH2, which can be obtained during high-energy milling or under high pressures10 (see arrows and inset in Figure 2 for trace under 40 bar). The formation of γ-MgH2 during milling was confirmed by XRD.9 This phase is known to be less-stable, and the desorption peak arises before the peak related to the β phase.10 However, there is no indication confirming the rehydrogenation of this metastable γ phase during cooling. The shoulder observed for cooling before the main peak in Figure 2 may nevertheless be related to this phase. This shoulder becomes pronounced with increasing pressure and appears at a higher temperature than that for the β phase. To construct the van’t Hoff plot, we will
Figure 3. Van’t Hoff’s plots for absorption and desorption of Mg99Ni1 (squares) compared with the results from Stampfer et al.15 for MgH2 (circles). Closed symbols, desorption; open symbols, absorption.
take into consideration only the positions of the peaks corresponding to β-MgH2. The van’t Hoff plots for absorption and desorption are given in Figure 3. Data from literature for the decomposition of MgH2 are included in the figure. The results show that our data can be fitted by a linear regression. Using the less-square method, we found a slope of 80.8 ( 1.6 kJ mol-1 for desorption and 68.1 ( 4.1 kJ mol-1 for absorption, respectively. Although the hydride phase corresponding to Mg99Ni1 is probably not exactly MgH2, considering the small amount of Ni, the results can be compared with those values given for Mg/MgH2.11 From PCI equilibrium measurements,7,11-15 the enthalpies of MgH2 formation and decomposition are known to be around 75 kJ mol-1 (Table 1). These values are usually an average of the absorption and desorption results. In Figure 3, the PCI results15 are lying between those derived from DSC traces for absorption and desorption as expected from eq 4. The fact that these curves are in the same range of temperature and pressure confirms that the onset temperatures obtained from DSC measurements are close to the thermodynamic values. Therefore, the results indicate that the kinetics do not limit the reactions. The use of nickel as a catalyst helps to reach these properties because it is a known catalyst for magnesium compounds.9 Chen et al.13 also found a large hysteresis and gave enthalpies of -68.4 kJ mol-1 for absorption and -74.5 kJ mol-1 for decomposition. Therefore, our measurements with HP-DSC lead to a good approximation of the enthalpy values for decomposition and absorption. Differences between our results and literature can be explained by the different hydride phase (Mg99Ni1Hz) and by the delicate determination of the exact peak position due to the metastable hydride phase present after milling. As pointed out in Section 3, the “true” value of reaction enthalpy is predicted between the slopes found for absorption and desorption, that is, 68.1 and 80.8 kJ mol-1. This range includes the results of the different groups cited above. 4.2. Mg-Co-H Ternary System. Three hydride phases are known in the Mg-Co-H system: tetragonal β-Mg2CoH5, orthorhombic γ-Mg6Co2H11, and cubic δ-Mg2-xCoH.16 Previous studies13,17 described the synthesis of Mg2CoH5 by mechanical alloying from a 2Mg(H2) + Co mixture. Mg2CoH5 is a compound similar to Mg2FeH6, which was synthesized successfully from a mixture of 2MgH2 + Fe and then released 5 wt % of hydrogen.18 However, the rehydrogenation of Mg2FeH6 is difficult because of the preferential formation of MgH2 from elemental Mg instead of Mg2FeH6 because no intermetallic compound exists between Mg and Fe. On the contrary, a MgCo2 phase is predicted from the phase diagram and the formation
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Figure 4. DSC traces for Mg2CoH5 prepared by ball-milling (second cycles).
of Mg2-xCo was also reported.16,17 For this work, elemental magnesium and cobalt powders were milled with a 2:1 ratio under hydrogen atmosphere. The XRD pattern of the as-milled compound (not shown) displayed the formation of the ternary phase Mg2CoH5 after 60 h of milling. The DSC traces in Figure 4 are taken from the second heating/ cooling cycle of measurement for different hydrogen pressures. The first cycle is considered to be an activation cycle. In this first cycle (not shown), during heating, an endothermic peak was observed at ca. 375 °C under 3 bar H2. This peak corresponds to desorption of the hydride into cubic MgCo and probably elemental Mg. For the second cycles, there are two distinguishable endothermic events whose positions shift to higher temperature when increasing pressure. Considering the results of Chen et al.,13 the first endothermic peak is attributed to the desorption of MgH2, whereas the second peak is attributed to the desorption of Mg2CoH5. As explained above, the HPDSC curves display the second cycle of measurement. It is therefore possible that during the previous absorption (cooling stage for cycle 1) the cubic MgCo and Mg forms the ternary hydride Mg2CoH5 and some of the elemental Mg forms MgH2. In the cooling section of the DSC trace, one exothermic peak is observed and related to hydrogen absorption. However, it is very difficult to distinguish between the formation of Mg2CoH5 and MgH2. The onset temperature of the exothermic peak shifts to a higher value, from 360 to 460 °C, with increasing pressure. However, this temperature is always lower than the desorption peak of the ternary hydride and higher than that for the desorption of MgH2. This fact indicates that the exothermic peak may be related to the formation of Mg2CoH5 (compare with Figure 1). Alternatively, the formation of MgH2 must be either overlapped by the Mg2CoH5 peak or too broad to be observable (see tail to lower temperatures for absorption peaks during cooling). The corresponding van’t Hoff plots for Mg2CoH5 are given in Figure 5 together with data from literature. As reported above for magnesium hydride, the values can be fitted easily with a linear curve. The calculated slopes are -75.4 ( 0.1 and -72.5 ( 4.0 kJ mol-1 for desorption and absorption, respectively. From PCI measurements (see Table 1), Reiser et al.11 reported a value for the enthalpy of Mg2CoH5 formation of 76 kJ mol-1. Ivanov et al.12 obtained -79 ( 4 kJ mol-1 for desorption. Chen et al.13 measured ∆H ) -69.5 kJ mol-1 for absorption and ∆H ) -83.2 kJ mol-1 for desorption. Our values obtained from HP-DSC approach the “true” enthalpy values. Kinetics appears not to limit the reactions, and the values of temperature and pressure are in the same range as those obtained from PCI
Rongeat et al.
Figure 5. Van’t Hoff plot for absorption and desorption of Mg2CoH5 (squares) compared with the results of Ivanov et al.12 (circles). Closed symbols, desorption; open symbols, absorption.
measurements (see Figure 5). Thus, the onset temperatures are related to the thermodynamics and not the kinetics of the reactions. Moreover, the hysteresis between absorption and desorption is lower for the Mg2CoH5 ternary hydride compared with the results found for Mg99Ni1. This fact indicates that the measurement is very close to the equilibrium as depicted for structural transformations (see Section 3) with enthalpies between 72.5 and 75.4 kJ mol-1. 4.3. Sodium Alanate. Bogdanovic and Schwickardi19 demonstrated in 1997 the reversibility of sodium alanate. The decomposition and reformation occur in two steps:
NaAlH4 T 1/3Na3AlH6 + 2/3Al + H2
(5)
Na3AlH6 T 3NaH + Al + 3/2H2
(6)
Here, a sample prepared by one-step-synthesis is used. This sample was synthesized by high-pressure ball milling under 100 bar of hydrogen of a NaH and Al mixture with 4% mol TiCl3 added as a catalyst. XRD for the as-milled compound (not shown) displayed peaks corresponding to NaAlH4 as the main phase. The presence of Al and NaCl was also detected as minor byproducts resulting from the reaction between NaH (or NaAlH4) and TiCl3.20-23 Figure 6 shows the DSC traces measured for this sample between room temperature and 300 °C under various hydrogen pressures. The curve measured under 20 bar H2 was obtained from the as-milled powder, and the subsequent cycles were measured one after the other with the same powder. The dotted trace under 20 bar was performed after the complete sequence of pressures to compare with the initial state. The two curves measured under 20 bar are very similar, indicating the good reversibility of sodium alanate. Several phase transitions are observed during heating or cooling, depending on the hydrogen pressure. The peaks are numbered from the trace under 120 bar but represent the same transitions for each pressure. Endothermic peaks are observed during heating, whereas exothermic peaks are observed during cooling. On the basis of a previous study by Claudy et al.24 for pure NaAlH4 under Ar atmosphere, and the fact that the peak positions will shift with pressure only if hydrogen is exchanged during the transition, the different features are identified as follows: (1) decomposition of NaAlH4 (eq 5), which can overlap with melting if the peak temperature is close to 180 °C, (2) structural transformation of Na3AlH6, and (3) decomposition of Na3AlH6 (eq 6). Note that peaks 2 and 3 overlap in the DSC traces under low pressures. The formation after the desorption
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Figure 7. Van’t Hoff’s plots for both steps of absorption and desorption of TiCl3-doped NaAlH4 (squares for eq 6 and triangles for eq 5) compared with the results of Bogdanovic et al.8 (circles, eq 6; stars, eq 5). Closed symbols, desorption; open symbols, absorption.
Figure 6. DSC traces for TiCl3-doped NaAlH4 prepared by one-step synthesis. Vertical dashed lines indicate the temperatures of melting (Tm) and of structural transformation of Na3AlH6 (Tstruc).
steps of a mixture of NaH, Al, and some NaCl was confirmed by XRD (not shown). The phase transition of Na3AlH6 corresponds to the change from the monoclinic R-Na3AlH6 to the orthorhombic β-Na3AlH6, which is the high-temperature phase.25 Considering reversible reactions, the peaks observed during cooling are related to (4) formation of Na3AlH6, (5) structural transition of Na3AlH6, and (6) formation of NaAlH4, which can also overlap with the solidification peak at approximately 180 °C. Peaks 4 and 5 can also overlap at low pressures. Moreover, peak 6 appears clearly only for high pressures. This result is the consequence of the high equilibrium pressure of the hydride formation (around 60 bar at 150 °C8) and its slow kinetics at low temperatures. We then assume that the peak of NaAlH4 formation may be very broad and small and it is shadowed by the background of the curve. Additionally, a small exothermic “wave” often arises at temperatures lower than the one corresponding to peak 1. This feature can be related to hydrogen absorption during heating, when the hydride is not completely formed during cooling. An exception in the sequence of peaks described above is found for the trace measured under 30 bar H2. In this case, the differences are the consequence of the incomplete reabsorption in the previous measurement under 20 bar. Thus, under 30 bar, the first event corresponds to an exothermic peak ascribed to hydride formation and the second feature to an endothermic peak related to desorption. For higher H2 pressure, reabsorption of hydrogen is observed during cooling and then two desorption peaks arise in the following cycle. From the peaks’ sequence, it is possible to construct the van’t Hoff plot for absorption and desorption (Figure 7). These plots are compared with the results given by Bogdanovic et al.8 For the second desorption step, corresponding to the phase transition from Na3AlH6 to NaH (eq 6), we obtain a slope of 51.9 ( 1.2 kJ mol-1 and 35.2 ( 0.5 kJ mol-1 for the corresponding absorption (NaH to Na3AlH6). These values are in fact an average between the reactions steps from/to R- and β-Na3AlH6. From PCI measurements, Bogdanovic et al.8 reported ∆H ) 47 kJ mol-1 as an average between absorption and desorption values. Given that they used temperatures lower than 200 °C, their results involved only R-Na3AlH6 but they can nevertheless
be compared with our results. The enthalpy value of 47 kJ mol-1 appears between the HP-DSC results for absorption and desorption but the hysteresis is high because of the relatively slow kinetics of reaction. For the first desorption step from NaAlH4 to Na3AlH6 (eq 5), we find a slope of 28.4 ( 1.5 kJ mol-1. Bogdanovic et al.8 reported for this step ∆H ) 37 kJ mol-1 (average between absorption and desorption). Our value is not far from this enthalpy considering the fact that this reaction is much more difficult to study. Higher pressures are necessary, NaAlH4 is relatively instable, and kinetics is slow. Moreover, the corresponding peaks are often overlapped with melting or solidification peaks like for the trace under 120 bar H2. It is therefore difficult to determine the onset temperature. The results obtained by DSC measurements are on the order of those obtained by other authors using equilibrium measurements14,26 (see Table 1). Therefore, the method seems to give a reasonable approximation of the enthalpy of reaction for these materials. The equilibrium pressures have to be in the range of the HPDSC apparatus in order to apply our method. In the case of LiAlH4, the hydride is very unstable with predicted enthalpies around 10 kJ mol-1 for the two steps of desorption.27,28 Indeed, no shift in peak position is observable from 10 to 140 bar hydrogen (not shown). The low enthalpies imply a very high equilibrium pressure in the range of temperatures used in the DSC (.1000 bar). Therefore, the van’t Hoff plot for this compound does not correspond to equilibrium values. This remark also applies to PCI measurements. Nevertheless, because the target for hydrogen storage material is an enthalpy of reaction around 40 kJ mol-1, useful compounds can be studied with HP-DSC as demonstrated for NaAlH4. 5. Conclusions Determination of enthalpies of hydride formation/decomposition is a crucial point to evaluate the ability of a compound to absorb and desorb hydrogen. In this paper, we show that the HP-DSC technique is a powerful tool to obtain easily and rapidly reliable results. PCI experiments are the method of choice to get as close as possible to “equilibrium” measurements and to construct the van’t Hoff plots but the analyses are very delicate and laborious. The HP-DSC measurements can be applied to a large variety of hydrides by adapting the hydrogen pressures used. The main difficulty as in the case of complex hydrides is to understand correctly the sequence of peaks appearing during heating/cooling. Another important point is to obtain values for
13306 J. Phys. Chem. B, Vol. 111, No. 46, 2007 both absorption and desorption. Considering the difficulty of reaching equilibrium, the range given by the two curves is necessary to approach the true enthalpy value. The equilibrium temperature is assumed to be between the onset temperatures for absorption and desorption. It is worth noting that this applies equally to PCI measurement for which a hysteresis is also obtained. The equilibrium plateau pressure lies between the absorption and desorption curve. With this, HP-DSC measurements are a much faster and simpler method to observe absorption and desorption. Moreover, more subtle features can also be seen by DSC and not by PCI, for example, the presence of a metastable phase in the case of Mg99Ni1 or the phase transition of Na3AlH6. This furthers the complete understanding of the reactions. Of course to get as precise results as possible, sufficient kinetics of the reactions is necessary. This is the essential parameter to reduce the hysteresis. Nevertheless, even for compounds with slow kinetics, for example, NaAlH4, it is possible to obtain reasonable approximations of the thermodynamic properties. This procedure is very suitable for evaluating new compounds as promising reactive composites, for example, LiBH4-MgH2 for which the thermodynamic properties are yet to be determined.29 Acknowledgment. This work is financially supported by the Helmholtz Initiative FuncHy and the Marie-Curie Research Training network COSY (EU-RTN). Thanks are due to M. Herrich and B. Gebel for experimental assistance. References and Notes (1) Schlapbach, L.; Zu¨ttel, A. Nature 2001, 414, 353-358. (2) Bohmhammel, K.; Christ, B.; Wolf, G. Thermochim. Acta 1996, 271, 67-73. (3) Hagstro¨m, M. T.; Lund, P. D.; Vanhanen, J. P. Int. J. Hydrogen Energy 1995, 20, 897-909. (4) Borgschulte, A.; Boesenberg, U.; Barkhordarian, G.; Dornheim, M.; Borman, R. Catal. Today 2007, 120, 262-269. (5) Gutfleisch, O.; Schlorke-de Boer, N.; Ismail, N.; Herrich, M.; Walton, A.; Speight, J.; Harris, I. R.; Pratt, A.; Zu¨ttel, A. J. Alloys Compd. 2003, 356-357, 598-602. (6) Pedersen, A. S.; Kjoller, J.; Larsen, B.; Vigeholm, B. Int. J. Hydrogen Energy 1983, 8, 205-211.
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