Thermodynamics of the α, β and γ polymorphs of AlH3

Article abstract of DOI:10.1016/j.jallcom.2005.11.086

The thermodynamics of the α, β and γ polymorphs of AlH₃ were determined using differential scanning calorimetry and ex situ X-ray diffraction. These results demonstrate that at around 100 ° C the decomposition of the β and γ polymorphs occurs by an initial phase transition to the α polymorph followed by decomposition of the α phase. The total heat evolved during the β → α transition is 1.5 ± 0.4 kJ/mol AlH₃ and 2.8 ± 0.4 kJ/mol AlH₃ during the γ → α transition. The transformation to the α phase is exothermic and is therefore likely to occur spontaneously at room temperature. A formation enthalpy of approximately - 10 kJ/mol AlH₃ was measured for α-AlH₃, which is in good agreement with previous experimental and calculated results.

Full text of DOI:10.1016/j.jallcom.2005.11.086

Journal of Alloys and Compounds 424 (2006) 262–265
Thermodynamics of the ␣, ␤ and ␥ polymorphs of AlH3
∗
Jason Graetz , James J. Reilly
Department of Energy Sciences and Technology, Brookhaven National Laboratory, Building 815, Upton, NY 11973, USA
Received 24 October 2005; accepted 2 November 2005
Available online 8 February 2006
Abstract
The thermodynamics of the ␣, ␤ and ␥ polymorphs of AlH
3
were determined using differential scanning calorimetry and ex situ X-ray diffraction.
These results demonstrate that at around 100 C the decomposition of the ␤ and ␥ polymorphs occurs by an initial phase transition to the ␣ polymorph
followed by decomposition of the ␣ phase. The total heat evolved during the ␤→␣ transition is 1.5 ± 0.4 kJ/mol AlH and 2.8 ± 0.4 kJ/mol AlH
◦
3
3
during the ␥→␣ transition. The transformation to the ␣ phase is exothermic and is therefore likely to occur spontaneously at room temperature.
A formation enthalpy of approximately −10 kJ/mol AlH
3 3
was measured for ␣-AlH , which is in good agreement with previous experimental and
calculated results.
©
2005 Elsevier B.V. All rights reserved.
Keywords: Energy storage materials; Hydrogen storage materials; Thermal analysis; Calorimetry
1
. Introduction
to determine the reaction mechanism and enthalpy. In addition,
the transition enthalpy of the conversion of the ␤ and ␥ phases
to the ␣ phase was also delineated.
Aluminum hydride (AlH3) is a covalently bonded, metastable
solid at room temperature with a large gravimetric and volumet-
3
ric hydrogen capacity (10.1 wt.% and 149 kg/m , respectively).
2. Experimental
The high capacity and rapid kinetics [1–3] has generated consid-
erable interest in using AlH3 as an H2 source in automotive fuel
cells. AlH3 was originally synthesized as a nonsolvated solid
using an organometallic synthesis route by Brower et al. [4]. In
addition to the ␣ phase, Brower et al. successfully prepared six
Differential scanning calorimetry was performed using a Mettler-
Toledo DSC822 . This instrument uses the Boersma, or heat-flux con-
e
figuration whereby energy released or absorbed is determined by mea-
suring the heat flow between the sample and a reference crucible.
The reaction enthalpies were determined by measuring the heat ab-
ꢀ
other non-solvated AlH3 polymorphs, i.e. ␣, ␣ , ␤, ␥, ␦, ⑀ and
◦
◦
sorbed/released during a temperature ramp between 35 C and 300 C
␨. Over the past 30 years the ␣ phase has been thoroughly in-
◦
at a rate of 10 C/min. In this study it is assumed that the magnitude
vestigated [5–13,1,2]. However, little is known about the other
solid alane polymorphs. In this letter we present new informa-
tion on the thermal stability of freshly synthesized ␣, ␤ and ␥
polymorphs of AlH3.
The thermal decomposition reaction of the subject poly-
morphs is quite straight forward as shown below:
of the decomposition enthalpy is equivalent to the formation enthalpy.
Although measurements were performed during the decomposition re-
action, the values are reported as formation enthalpies (with a sign
change) to be consistent with similar studies in the literature. Ex situ X-
ray diffraction measurements were performed at room temperature on
a Philips diffractometer using Cu K␣ radiation. X-ray samples received
the same thermal treatment as the DSC samples, but were quenched in a
room temperature water bath after reaching a given point in the thermal
ramp.
AlH3 → Al + 3/2H2
(1)
This reaction was studied using differential scanning
calorimetry (DSC) in conjunction with ex situ X-ray diffraction
The synthesis of the ␣, ␤ and ␥ polymorphs were previously
described in detail; for convenience a brief summary is presented
here [3]. AlH
with LiAlH to produce an etherated species of aluminum hydride,
AlH O] [4]. The removal of the associated ether complex
·0.3[(C
was accomplished by heating the solvated AlH in the presence of a
3 3
was synthesized via an ethereal reaction of AlCl
4
∗
Corresponding author. Tel.: +1 631 344 3242; fax: +1 631 344 2359.
3
2 5 2
H )
E-mail address: graetz@bnl.gov (J. Graetz).
3
0
925-8388/$ – see front matter © 2005 Elsevier B.V. All rights reserved.
doi:10.1016/j.jallcom.2005.11.086

J. Graetz, J.J. Reilly / Journal of Alloys and Compounds 424 (2006) 262–265
263
◦
◦
Fig. 1. (a) Differential scanning calorimetry plot of ␣-AlH3 in the temperature range of 35–300 C ramped at a rate of 10 C/min; (b) ex situ diffraction patterns
◦
acquired at room temperature before thermal treatment (I) and after a temperature ramp to 300 C (II).
complex metal hydride (LiAlH
is extremely sensitive to the desolvating conditions (e.g. temperature
and time) and small alterations can lead to the precipitation of a differ-
4
, LiBH
4
) under vacuum. The synthesis
predicted by Claudy et al. [13], but considerably higher than the
values measured by Baranowski and Tkacz. [7].
The DSC plot of ␤-AlH3 is shown in Fig. 2a. In addition to the
ent AlH
while ␤-AlH
In both cases, a slightly higher temperature (∼ 70 C) and/or a longer
heating time will lead to the formation of ␣-AlH
3
polymorph. ␥-AlH
3
forms in the presence of excess LiAlH
4
,
◦
large endothermic peak observed around 170 C, there is a sig-
3
forms in the presence of excess LiAlH
4
and LiBH [4,3].
4
◦
nificant exothermic peak that appears at approximately 100 C.
◦
Diffraction patterns were acquired prior tothe thermaltreatment,
3
.
◦
◦
after a ramp to 130 C, and after a ramp to 290 C and are shown
in Fig. 2b. Prior to thermal treatment the sample is predominately
composed of the ␤ phase with a small amount of a ␥ phase im-
purity as shown in pattern I. Subsequent to the temperature ramp
3
. Results and discussion
◦
Fig. 1a shows a DSC plot of ␣-AlH3 over a temperature range
up to 130 C the sample is almost entirely composed of ␣-AlH3
◦
of 35–300 C. A large endothermic peak is observed around
1
with a small amount of Al metal and residual ␤-AlH3 as shown
in pattern II. As the temperature is further increased the ␣ phase
has decomposes to Al metal (and H2) as shown in pattern III. It
is clear that the ␤→␣ transition is the preferred decomposition
◦
70 C and is attributed to the decomposition of AlH3 (reac-
tion (1)). Fig. 1b displays the corresponding diffraction pattern
◦
◦
at 35 C and after ramping to 290 C. As expected, the diffrac-
tionpattern demonstrates that ␣-AlH3 is completelytransformed
to Al metal (and H2) during the temperature ramp. The mea-
sured formation enthalpy is ꢀH1 = −9.9 ± 0.4 kJ/mol AlH3
as shown in Table 1. This is similar to the enthalpy measured
by Sinke et al. (−11.4 ± 0.8 kJ/mol AlH3) using calorime-
try and calculated values by Wolverton et al. (−6.95 kJ/mol
AlH3) [14] and Ke et al. (−12.35 kJ/mol AlH3) [15]. The Gibbs
free energy of formation at room temperature (ꢀG298 K) was
calculated using the measured formation enthalpy and an en-
tropy change (ꢀS) of 130.7 J/mol K, which is the entropy as-
sociated with the transformation of hydrogen from a disor-
dered gas to an ordered solid. The Gibbs formation energy is
◦
pathway at around 100 C. However, it should be noted that at
lower temperatures or faster heating rates the ␤ phase may de-
compose more readily to Al + 3/2H2. The transition enthalpy
(ꢀH0), heat of formation for Al + 3/2H2 →␣-AlH3 (ꢀH1), to-
tal formation enthalpy (ꢀH
= ꢀH0 + ꢀH1) and the Gibbs
total
free energy for ␥-AlH3 are shown in Table 1. The measured
transition enthalpy may be slightly higher than the actual value
due to the contribution from the ␥ phase impurity. Therefore,
the magnitudes for ꢀHtotal and ꢀG298 K may underestimate the
actual values.
The DSC plot of ␥-AlH3, shown in Fig. 3a, also exhibits a
◦
large exothermic peak around 100 C. The ex situ diffraction
ꢀ
G298 K = 48.5 ± 0.4 kJ/mol AlH3 as shown in Table 1. It is
patterns acquired prior to the thermal treatment, after a ramp to
133 C, andafterarampto290 CareshowninFig. 3b. Thestart-
ing material is composed of pure ␥-AlH3, which transforms to
◦
◦
interesting to note that this value yields an H2 equilibrium pres-
sure of 5 × 10 bar at 298 K, which is comparable to the value
5
Table 1
Formation energies and temperature onsets for ␣, ␤ and ␥-AlH3 showing transition enthalpy (ꢀH0) and temperature onset (T0), heat of formation for Al + 3/2H2 →␣-
AlH3 (ꢀH1) and temperature onset (T1), total formation enthalpy (ꢀHtotal = ꢀH0 + ꢀH1) and the Gibbs free energy (ꢀG298 K = ꢀHtotal − TꢀS where ꢀS =
1
30.7 J/mol K)
◦
◦
Polymorph
ꢀH0 (kJ/mol AlH3)
T0 ( C)
ꢀH1 (kJ/mol AlH3)
T1 ( C)
ꢀHtotal (kJ/mol AlH3)
ꢀG298 K (kJ/mol AlH3)
␣
-AlH3
-AlH3
–
–
−9.9 ± 0.6
−9.5 ± 0.6
−9.9 ± 0.6
162.1
157.1
158.3
−9.9 ± 0.6
−8.0 ± 1.0
−7.1 ± 1.0
48.5 ± 0.6
50.5 ± 1.0
51.4 ± 1.0
␤
1.5 ± 0.4
91.0
100.7
␥-AlH3
2.8 ± 0.4

2
64
J. Graetz, J.J. Reilly / Journal of Alloys and Compounds 424 (2006) 262–265
◦
◦
Fig. 2. (a) Differential scanning calorimetry plot of ␤-AlH3 in the temperature range of 35–300 C ramped at a rate of 10 C/min; (b) ex situ diffraction patterns
◦
◦
acquired at room temperature before thermal treatment (I), after a temperature ramp to 130 C (II) and after a temperature ramp to 300 C (III).
◦
◦
Fig. 3. (a) Differential scanning calorimetry plot of ␥-AlH3 in the temperature range of 35–300 C ramped at a rate of 10 C/min; (b) ex situ diffraction patterns
◦
◦
acquired at room temperature before thermal treatment (I), after a temperature ramp to 130 (II) and after a temperature ramp to 300 C (III).
◦
␣
-AlH3 upon a temperature ramp to 133 C. Additional heating
for the ␥→␣ transition. These reactions are exothermic and are
therefore likely to occur spontaneously even at room tempera-
ture. The measured formation enthalpy for ␣-AlH3 is approxi-
mately −10 kJ/mol AlH3, which is consistent with other exper-
imental and calculated results.
◦
to 290 C exhibits the expected decomposition of the ␣ phase to
Al metal (and H2). These results demonstrate that, similar to the
␤
phase, the decomposition of ␥-AlH3 occurs by an initial ␥→␣
transition followed by decomposition of the ␣ phase. Similar ob-
servations of the ␥→␣ transition were observed by Claudy et
◦
al. [16]. Despite the complete ␥→␣ transition at 100 C, it is
Acknowledgements
possible that the ␥ phase decomposes directly to the elements
at other temperatures. The transition enthalpy, heat of forma-
tion for Al + 3/2H2 →␣-AlH3 and the Gibbs free energy for
TheauthorsgratefullyacknowledgeGarySandrockforhisin-
sights and his encouragement to investigate aluminum hydride.
This work was supported by the Department of Energy’s Office
of Energy Efficiency and Renewable Energy. This manuscript
has been authored by Brookhaven Science Associates, LLC un-
der Contract No. DE-AC02-98CH1-886 with the U.S. Depart-
ment of Energy.
␥
-AlH3 are shown in Table 1.
4
. Conclusion
Ex situ X-ray diffraction of ␤ and ␥-AlH3 performed during
thermal treatment has demonstrated that the decomposition oc-
◦
curs via a transition to the more stable ␣ phase at around 100 C
References
followed by the decomposition of the ␣ phase to Al and H2.
The heats evolved during the polymorph transformations are
[
1] G. Sandrock, J. Reilly, J. Graetz, W.M. Zhou, J. Johnson, J. Wegrzyn, Appl.
Phys. A 80 (2005) 687.
1
.5 kJ/mol AlH3 for the ␤→␣ transition and 2.8 kJ/mol AlH3

J. Graetz, J.J. Reilly / Journal of Alloys and Compounds 424 (2006) 262–265
265
[
2] G. Sandrock, J. Reilly, J. Graetz, W.-M. Zhou, J. Johnson, J. Wegrzyn, J. [9] P.J. Herley, O. Christofferson, R. Irwin, J. Phys. Chem. 85 (1981) 1874.
Alloys Comp. (2006), in press.
[10] P.J. Herley, O. Christofferson, J. Phys. Chem. 85 (1981) 1882.
[11] P.J. Herley, O. Christofferson, J. Phys. Chem. 85 (1981) 1887.
[12] P.J. Herley, O. Christofferson, J.A. Todd, J. Solid State Chem. 35 (1980)
391.
[13] P. Claudy, B. Bonnetot, J. Etienne, G. Turck, J. Therm. Anal. 8 (1975)
255.
[
[
3] J. Graetz, J. Reilly, J. Phys. Chem. 109 (2006) 22181.
4] F.M. Brower, N.E. Matzek, P.F. Reigler, H.W. Rinn, C.B. Roberts, D.L.
Schmidt, J.A. Snover, K. Terada, J. Am. Chem. Soc. 98 (1976) 2450.
5] J.W. Turley, H.W. Rinn, Inorg. Chem. 8 (1969) 18.
[
[
6] G.C. Sinke, L.C. Walker, F.L. Oetting, D.R. Stull, J. Chem. Phys. 47 (1967)
2
759.
[14] C. Wolverton, V. Ozolins, M. Asta, Phys. Rev. B 69 (2004) 144109.
[15] X. Ke, A. Kuwabara, I. Tanaka, Phys. Rev. B 71 (2005) 184107.
[16] P. Claudy, B. Bonnetot, J.M. Letoffe, J. Therm. Anal. 15 (1979) 129.
[
[
7] B. Baranowski, M. Tkacz, Z. Phys. Chem. 135 (1983) 27.
8] P.J. Herley, R.H. Irwin, J. Phys. Chem. Solids 39 (1978) 1013.