N-Alkylpyrrolidine·Alane Compounds for Energy Applications - The

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N‑Alkylpyrrolidine·Alane Compounds for Energy Applications Chengbao Ni,†,* Liu Yang,‡ James T. Muckerman,‡ and Jason Graetz† †

Sustainable Energy Technologies Department, ‡Chemistry Department, Brookhaven National Laboratory, Upton, New York 11973-5000, United States S Supporting Information *

ABSTRACT: Aluminum hydride (AlH3) has high gravimetric and volumetric hydrogen densities and is thus a promising hydrogen storage material. As part of our effort to develop a cost-effective regeneration pathway for AlH3, we report the synthesis and characterization of N-alkylpyrrolidine·alane complexes (NMPy)2·AlH3, NMPy·AlH3, and NEPy·AlH3 (NMPy = N-methylpyrrolidine, NEPy = N-ethylpyrrolidine); the reversible formation of (NMPy)2·AlH3 from titaniumdoped aluminum metal (denoted as Al*), H2 gas, and NMPy; and thermal decomposition studies of these alane adducts. Depending on the stoichiometric ratios of NMPy to AlH3, both the 2:1 complex (NMPy)2·AlH3 and the 1:1 complex NMPy·AlH3 can be selectively synthesized by direct reactions of NMPy with AlH3, whereas NEPy gives only the 1:1 adduct NEPy·AlH3 regardless of the ratio of NEPy to AlH3. In addition, the reversible formation of (NMPy)2·AlH3 from titanium-doped aluminum powder (Al*) and NMPy under a H2 pressure of 1000 psi was observed, while no alane formation was detected with NEPy. Theoretical calculations of the molecular geometries and absolute free energies are in good agreement with experimental observations, and indicate that the dramatic differences between NMPy and NEPy are caused by the steric effects imposed by the alkyl group (methyl or ethyl) on the pyrrolidine ring. Finally, we established the role of LiH in the decomposition of amine·alane adducts, and showed that the hydrogenation of Al* with NMPy and the decomposition of N-alkylpyrrolidine·alane adducts in the presence of LiH could be combined to generate donor-free LiAlH4.

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INTRODUCTION Among traditional metal hydrides and their associated compounds,1−3 aluminum hydride (or alane, AlH3) is a particularly promising material for hydrogen storage because of its high volumetric (1.48 g/mL) and gravimetric (∼10%) hydrogen capacities and its relatively low decomposition temperature (∼100 °C). These features, combined with the fast H2 release rates, make AlH3 ideal for a number of applications, including low-temperature fuel cell applications.4,5 Unfortunately, the conventional synthesis route6 for AlH3 is costly because it involves reactions of aluminum trichloride with alkali alanates (LiAlH4 or NaAlH4) or alkali metal hydrides (LiH or NaH). In addition, the direct hydrogenation of Al metal to obtain AlH3 is infeasible under practical conditions because of the extremely high equilibrium pressure of H2 gas (∼7 kbar at 298 K).7 Using triethylenediamine (TEDA) as a complexing agent, Ashby reported in 1964 the direct synthesis of TEDA·AlH3 from Al metal, H2, and TEDA at lower pressures (∼5000 psi).8 Since then, several other tertiary amines, including trimethylamine (TMA), dimethylethylamine (DMEA), quinuclidine, and hexamethylenetetramine (HMTA), have also been shown to form their corresponding amine·alane adducts via hydrogenation.2,9 A common feature of these amines is their very strong basicity.10,11 While TMA and DMEA are the least bulky and thus the strongest acyclic tertiary amines, quinuclidine, TEDA, and HMTA are even stronger because the three alkyl © 2013 American Chemical Society

groups are tethered in such a way that the steric hindrance is significantly reduced. This consideration is supported by highlevel theoretical calculations on the thermodynamics and stabilities of AlH3 adducts with various Lewis bases.12 The formation of amine·alane adducts by direct hydrogenation may be an important first step in alane regeneration, provided that the corresponding adduct can be easily separated. In this case, the ideal stabilizing ligand (usually a Lewis base) forms a bond with Al (e.g., Al−N) that is not too strong or too weak. Unfortunately, the amine·alane adducts that can be formed via direct hydrogenation all have strong Al−N bonds and thus cannot be easily decomposed to give AlH3 in a single step.13,14 On the other hand, the amine·alane complexes (e.g., TEA·AlH3, TEA = triethylamine) that have relatively weak Al− N bonds and can be separated to recover AlH3 cannot be formed by direct hydrogenation.15,16 Because of such a “paradox”, there is currently much interest in identifying a “Goldilocks” amine (or other Lewis base) that is strong enough to stabilize alane, but weak enough for the Al−N bond to be broken under mild temperature and pressure conditions. As part of our continuing effort to develop a cost-effective regeneration pathway for AlH3, we sought to tune the bulkiness and basicity of tertiary amines for hydrogen storage Received: November 2, 2012 Revised: January 22, 2013 Published: January 22, 2013 2628

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applications. N-Alkylpyrrolidine (Figure 1), with two alkyl groups fused into a five-membered ring, has a basicity between

sealed in a 300 mL stainless Parr reactor, which was purged with H2 gas twice before a final pressure of ca. 1000 psi was applied. Upon mechanical stirring (∼300 rpm), initial dissolution of H 2 into the organic mixture occurred immediately. The mixture was stirred for 20 h, by which time the reactor pressure dropped significantly, suggesting consumption of H2 by the reaction mixture. Excess H2 was released, and the reactor was opened in the glovebox. The reaction mixture was filtered; the colorless filtrate was evaporated to give a white crystalline solid with a yield of 1.04 g (60.7%). The yield determined from the consumption of H2 was slightly higher, indicating that partial decomposition of (NMPy)2·AlH3 occurred during the workup process. 1H NMR (400.14 MHz, toluene-d8): δ 1.65−1.68 (m, NCH2CH2, 8H), 2.27 (s, NCH3, 6H), 2.54−2.58 (m, NCH2, 8H), 3.61 (s, br, AlH, 3H). 13C NMR (100.62 MHz, toluene-d8): δ 24.6 (NCH2CH2), 45.0 (NCH3), 57.1 (NCH2CH2). The proton chemical shifts and peak widths (especially Al-H) depend on the concentration of the sample. Al−H (cm−1): 1709 (both in a CaF2 cell and in a KBr pellet). (NMPy)2·AlH3 via γ-AlH3. NMPy (0.358 g, 0.348 mL, 4.2 mmol) was added dropwise at room temperature to a suspension of freshly prepared γ-AlH3 (0.060 g, 2.0 mmol) in Et2O (10 mL). The mixture was stirred for 30 min. The resulting colorless solution was filtered. The filtrate was evaporated to give a white crystalline solid with a yield of 0.394 g (98.3% based on AlH3). The compound was confirmed to be (NMPy)2·AlH3 by NMR, IR, and XRD studies. NMPy·AlH3 via γ-AlH3. NMPy (0.213 g, 0.217 mL, 2.5 mmol) was added dropwise at room temperature to a suspension of γ-AlH3 (0.090 g, 3.0 mmol) in Et2O (8 mL). The mixture was stirred for 2 h, by which time a clear solution with some white precipitate was obtained. The mixture was then filtered; the filtrate was evaporated to give a colorless, viscous liquid with a yield of 0.288 g (95.7% based on NMPy). 1 H NMR (400.14 MHz, toluene-d8): δ 1.39 (br, NCH2CH2, 4H), 2.00 (s, NCH3, 3H), 2.35 (br, NCH2CH2, 4H), 4.07 (s, br, Al-H, 3H). 13C NMR (100.62 MHz, toluene-d8): δ 23.7 (NCH2CH2), 45.6 (NCH3), 58.2 (NCH2CH2). Al−H (cm−1): 1773 (in a CaF2 cell). NMPy·AlH 3 via (NMPy) 2 ·AlH 3 . A solution of (NMPy)2·AlH3 (0.200 g, 1.0 mmol) in toluene (2 mL) was held under vacuum at room temperature while being stirred. Once the solvent was evaporated, another 2 mL of toluene was added and evaporation was continued. After a few cycles, a colorless, viscous liquid was obtained, which was identified to be NMPy·AlH3 by NMR and IR spectroscopy. NEPy·AlH 3 . A similar procedure to that used for NMPy·AlH3, in which γ-AlH3 (0.090 g, 3.0 mmol) and NEPy (0.248 g, 2.5 mmol) were employed, produced a viscous liquid with a yield of 0.310 g (96.2% based on NEPy). 1H NMR (400.14 MHz, toluene-d8): δ 0.99 (t, 3J = 8.0 Hz, 3H, CH3), 1.46 (br, 4H, NCH2CH2), 2.27 (q, 3J = 8.0 Hz, 2H, NCH2CH3), 2.42 (br, 4H, NCH2CH2), 4.08 (s, br, Al-H, 3H). 13C NMR (100.62 MHz, toluene-d8): δ 11.8 (CH3), 23.7 (NCH2CH2), 54.4 (NCH2CH3), 56.1 (NCH2CH2). Al−H (cm −1): 1773 (in a CaF2 cell). Attempts to prepare (NEPy)2·AlH3 with more than 2 equiv of NEPy were unsuccessful. Thermal Decomposition in the Presence of LiH. Thermal decomposition studies were conducted in a Buchi vacuum furnace inside a glovebox. In a typical experiment, an alane·amine compound (1.0 mmol, 1.0 equiv) and LiH (0.008

Figure 1. The molecular structures of (a) NMPy and NEPy; (b) commonly observed amine·alane complexes. R3N is a tertiary amine. Type I represents the structure of a 1:1 monomeric complex, Type II that for a 2:1 complex, and Type III that for a 1:1 dimeric complex.

the acyclic amines (e.g., DMEA) and the bicyclic amines (e.g., quinuclidine and TEDA), and has been identified by theoretical studies as a potential candidate.12 In addition, previous studies that focused on the synthesis of NMPy alane complexes and the hydrogen storage capacities in and H2 release from these adducts have shown favorable properties of these NMPy alane complexes for energy applications.17−20 Thus, we wished to explore N-alkylpyrrolidine alane complexes for the generation of nonsolvated metal hydride species (e.g., AlH3 or LiAlH4) for hydrogen storage purposes. In the present paper, we report the synthesis and characterization of alane complexes of NMPy and NEPy, the reversible formation of (NMPy)2·AlH3 from Al* (titanium-doped aluminum powder), and the thermal decomposition studies of these adducts for the recovery of metal hydrides. In addition, we performed computations of the molecular geometries, binding free energies, and the Al−H vibrational frequencies of several commonly observed amine·alane structural types (Figure 1) to better understand the hydrogenation process and the dramatic differences between NMPy and NEPy. Interactions between AlH3 and Et2O were also included for comparison. The computational results are generally in good agreement with experimental observations.

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EXPERIMENTAL SECTION General. All manipulations were carried out in a vacuum atmosphere glovebox filled with argon. H2 (99.95%) was purchased from Praxair corporation and used as received. Diethyl ether (99.7%), N-methylpyrrolidine (98.5%), Netheylpyrrolidine (99.0%), and toluene (99.8%) were purchased from Sigma-Aldrich. Deuterated toluene (99.5%, toluene-d8) was obtained from Cambridge Isotope Laboratories. Al* and γ-AlH3 were prepared according to published procedures.13,21 Liquid chemicals were dried over 3 Å molecular sieves prior to use. NMR spectra were collected with use of a Bruker 400 spectrometer in toluene-d8. 1H (or 13 C) NMR resonances are reported in ppm relative to solvent resonances, taken as 7.13 (or 137.91) for toluene-d8. Infrared spectra were recorded on a Perkin-Elmer SpectrumOne spectrometer either in a 0.1 mm CaF2 liquid cell or in KBr pellets. X-ray powder diffraction (XRD) patterns were collected on a Phillips Cu Kα diffractometer. Yields indicate the amount of isolated compound and reactions are unoptimized. (NMPy)2·AlH3 via Direct Hydrogenation. The synthesis is similar to that of (DMEA)2·AlH3.15 NMPy (20 mL, 241 mmol), Al* (0.230 g, 8.5 mmol), and Et2O (80 mL) were 2629

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the generation of Li3AlH6 as a byproduct rendered a low overall yield for this reaction. To overcome the above disadvantages, we employed the reaction of NMPy or NEPy with γ-AlH3 (eqs 1 and 2) in our studies. Although we chose γ-AlH3 in our synthesis, it should be noted that other AlH3 polymorphs, such as α-AlH3, could also be used. Because of the high purity of γ-AlH36 and the absence of undesired reactions, this method is very clean, giving the desired products in high yield and purity. Depending on the stoichiometric ratios, both (NMPy)2·AlH3 and NMPy·AlH3 can be selectively prepared; however, NEPy forms only NEPy·AlH3, regardless of the amount of NEPy.

g, 1.0 mmol, 1.0 equiv) were added to Et2O (2 mL). The mixture was heated to 50 °C while being held under reduced pressures. After 2 to 3 h at 50 °C, a white powder was obtained, which was identified to be LiAlH4 by IR and XRD analyses. With less than 1.0 equiv of LiH, a mixture of LiAlH4 and Al metal was formed. Theoretical Calculations. Geometry optimizations of the isolated molecules and alane complexes of NMPy, NEPy, and Et2O were carried out at the MP222−26/6-31+G(d,p)27−29 level of theory in the gas phase. Three different types of molecular structures (Figure 1) were examined. A vibrational frequency calculation on each species at the same level of theory was performed to confirm the stationary point. The gas-phase zeropoint energy and thermal correction (ΔG0ZPE&Th(g)) to the electronic energy were obtained as the difference between the gas-phase free energy, G0(g), and the optimized electronic energy, E0el,(g). A more accurate electronic energy of each species in Et2O (E0el,(S)) was then obtained by a single-point electronic energy calculation at the gas-phase optimized geometry carried out at the MP2/6-311+G(2d,2p)30−33 level of theory with a CPCM34−36 treatment of the Et2O solvent, using UAHF radii. The absolute free energy of each complex in Et2O is then calculated as G∗(S) = G0(S) + ΔG0→* = E0el,(S) + ΔG0ZPE&Th(g) + ΔG0→*. The standard state correction (ΔG0→* = 1.894 kcal·mol−1), the work required to compress an ideal gas from a molar volume of 24.47 to 1 L, is included to account for the change in standard state from the gas phase (1 atm), the quantum chemical standard state for all species, to the experimental standard state in solution (1 mol/L). It should * values do not account for be noted that the reported G(S) differences in the thermal correction between the gas phase and aqueous phase, i.e., no vibrational frequency analysis was carried out with the continuum solvation model. For all cases of intermolecular bonding, a counterpoise calculation was also carried out to determine the basis set superposition error (BSSE). To calculate the absolute free energy of an Et2O molecule in Et2O, we need to calculate the “self-solvation free energy”, which is the free energy associated with moving the gas phase molecule at 1 mol/L to the pure liquid phase. The “selfsolvation energy” of Et2O is −3.440 kcal/mol as calculated by using the relation ΔG∗self = −2.303 RT log[(Mliq/Mo)/(pvapor/ Po)].37 Here Mo is equal to 1 mol/L, Po is the pressure (24.47 atm) of an ideal gas at 1 M concentration and 298.15 K, and Mliq is 9.619 mol/L, which is the molarity of Et2O in its pure liquid form. It can be calculated by using the relation Mliq = ρliq/MWliq. The absolute free energy of Et2O in its liquid phase is then equal to G∗(liq) = G0(g) + ΔG0→* + ΔG∗self.

γ − AlH3 + n NMPy ⎯⎯⎯⎯⎯⎯⎯⎯⎯→ (NMPy)n ·AlH3

(1)

γ − AlH3 + NEPy → NEPy·AlH3

(2)

n = 1 or 2

(NMPy)2·AlH3 is a white crystalline solid at room temperature, while NMPy·AlH3 and NEPy·AlH3 are liquid. Under moderate vacuum, (NMPy)2·AlH3 readily loses one NMPy molecule to form NMPy·AlH3; thus, proper conditions have to be applied in order to isolate (NMPy)2·AlH3 in high purity. (NMPy)2 ·AlH 3 is slightly more thermally stable than NMPy·AlH3 and NEPy·AlH3. However, these complexes are all unstable at room temperature as formation of Al metal is often observed after a few days. Despite this instability, these complexes can be stored below 0 °C for several months without any discernible decomposition. Calculations of Binding Free Energy and Molecular Geometries. To shed further light on the differences among these alane adducts, we examined the interactions in all three structural types (Figure 1) between AlH3 and NMPy, NEPy, and Et 2 O. The optimized geometries of Et 2 O·AlH 3 , (Et2O)2·AlH3, and isolated pyrrolidine molecules are provided in the Supporting Information (SI). The calculated binding Gibbs free energies with BSSE correction of different binding configurations are summarized in Table 1. The optimized Table 1. Binding Free Energy (with respect to 1 mol AlH3) of Different Configurations of Alane Adducts at 298.15 K in Et2O at the MP2 Level of Theory with BSSE Correction binding free energy in Et2O with BSSE correction (kcal·mol−1)

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alane adduct

Type I

Type II

Type III

NMPy NEPy Et2O

−19.91 −14.72 −13.24

−22.11 −19.49 −15.92

−18.57 −15.94 −13.77

structures of all possible geometries of amine·alane complexes in Et2O are shown in Figure 2. The Cartesian coordinates of their corresponding geometries are listed in the SI. Although Et2O can coordinate to AlH3, the Al−O bond is relatively weak (−15.92 kcal·mol−1) compared to the Al−N bond (N from an amine). Thus, formation of Et2O·AlH3 is negligible in the presence of NMPy or NEPy. In an Et2O/ NMPy solution, AlH3 prefers the coordination of one NMPy molecule as indicated by the calculated binding free energies. However, both the 1:1 and 2:1 complexes are thermally accessible. The small free energy difference suggests possible conversions between these two configurations, and is further supported by our experimental observations that (NMPy)2·AlH3 readily loses one NMPy to give NMPy·AlH3 under reduced pressures. The calculated binding free energies

RESULTS AND DISCUSSION Synthesis. The synthesis of (NMPy) 2 ·AlH 3 and NMPy·AlH3 has been previously reported.17,20 The most commonly used method involves the reaction of LiAlH4 with AlCl3 in the presence of a certain amount of NMPy.18 However, the isolated products are usually a mixture of (NMPy) 2 ·AlH 3 and NMPy·AlH 3 , possibly due to the coordination of NMPy to AlCl3, AlH3, and LiCl.18,20,38 In addition, this synthetic route is sensitive to the stoichiometric ratios of the starting materials; a slight alternation in ratio could introduce undesired chloride impurities (e.g., AlH2Cl or AlHCl2). Alternatively, (NMPy)2·AlH3 has been synthesized by a direct reaction of LiAlH4 with excess NMPy.17 However, 2630

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centered at 1709 cm−1, which is also observed in (TMA)2·AlH3 (1710 cm−1)16 and (DMEA)2·AlH3 (1710 cm−1)15 and other related alane adducts. The 1:1 adducts NMPy·AlH3 and NEPy·AlH3 (see SI) show Al−H absorptions at 1773 cm−1, which is typically associated with the monomeric structure.16 The similar Al−H absorptions of these 1:1 compounds may reflect similar structures in solution. The Al−H stretching frequencies of these complexes in the gas phase were predicted at the MP2/6-31+G(d,p) level of theory. As shown in Table 2, the predicted Al−H stretch Table 2. Calculated Al−H Stretching Frequencies (cm−1) at the MP2/6-31+G(d,p) Level of Theorya

Figure 2. Optimized geometries of alane adducts of NMPy (a-I, a-II, and a-III) and NEPy (b-I, b-II, and b-III) in Et2O.

of the NEPy alane adducts suggest that the 1:1 monomer (Type II) is the dominant species in Et2O, a finding that is also consistent with experimental results. The Type III dimeric structure with two bridging hydrides is a common solid-state structural unit for several alane adducts when the complexes are closely packed.39,40 However, the weak binding free energy of the dimer structure indicates that the dimer dissociates into monomers in the solution phase. Characterization. The alane complexes were characterized by NMR and IR spectroscopy. For (NMPy)2·AlH3, X-ray powder diffraction (XRD) patterns were also studied. To understand the effects of amine coordination on the hydrides, we performed 1H NMR studies. The Al−H signal of (NMPy)2·AlH3 is upfield by ca. 0.4 ppm compared to those of the 1:1 adducts,19,41 suggesting more shielded hydrides in (NMPy)2·AlH3. Thus, the Al−H chemical shift is an indicator of the coordination environment of AlH3. In addition, because of coupling to the 27Al nucleus,42 the Al−H proton signals were usually very broad. With increasing concentrations, the Al−H signals became sharper, and the positions of all proton signals shifted slightly. These features are consistent with previous observations.20,42 IR spectroscopy is a useful tool for studies of alane adducts because the Al−H stretches can be easily observed in the IR region. As shown in Figure 3, the Al−H absorption of (NMPy)2·AlH3 (in both a CaF2 cell and a KBr pellet) is

alane adduct

Type I

Type II

NMPy

1805 (1700)

1900 (1789)

NEPy

1805 (1700)

1900 (1789)

Type III 1695 (1598), 1902 (1793) 1697(1598), 1904 (1793)

experiment 1709, 1773 1773

a

Values in parentheses are corrected frequencies by a factor of 0.9418.44

frequencies of the Type II complexes of both NMPy and NEPy are 1789 cm−1, and those of Type I are 1700 cm−1. These values are close to experimental frequencies. In addition, calculations showed splits of the Al−H absorptions in Type III structures. The frequency at 1598 cm−1 corresponds to the stretching of the bridging hydrides (Al−H−Al), and that at 1793 cm−1, with stronger intensity, is due to the stretching of the terminal Al−H groups. These results are consistent with previous studies of the H3N·AlH3 dimer.43 We note that the absence of such a split of Al−H stretches in our experiments may originate from the different structures in solution and the gas phase. Despite numerous attempts, single crystals of (NMPy)2·AlH3 have not yet been obtained. Thus, X-ray powder diffraction was employed. As shown in Figure 4, (NMPy)2·AlH3 exhibits two

Figure 4. XRD patterns of independently synthesized (NMPy)2·AlH3 and the hydrogenation product.

relatively strong reflections at lower angles (2θ = 16.1° and 18.1°) and weak reflections at higher angles. Because of the low crystallinity of (NMPy)2·AlH3 and the lack of literature information, no crystal structural information could be derived from the current data set. Nevertheless, the XRD patterns display features similar to those of the related amine·alane adducts13,14,41 and provide useful information for comparisons.

Figure 3. The IR spectra of independently synthesized (NMPy)2·AlH3 and NMPy·AlH3, and the hydrogenation product. 2631

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Hydrogenation. As shown in Figure 5a, the initial fast drop in pressure was due to H2 dissolution; this process usually

CaF2 cell and a KBr pellet) to those of the independently synthesized sample, suggesting the formation of (NMPy)2·AlH3 in the reactor. As expected, the XRD patterns (Figure 4) and 1 H NMR chemical shifts also confirmed the formation of (NMPy)2·AlH3. The formation of (NMPy)2·AlH3 by direct hydrogenation is reversible. As illustrated in Figure 5b, heating the reaction mixture (after hydrogenation) to 100 °C followed by cooling to 20 °C led to an increase of pressure by 47.7 psi in the first cycle, indicating that H2 adsorption did occur in the reactor. Although only two cycles are shown in Figure 5b, complete decomposition of (NMPy)2·AlH3 formed in the reactor might require more cycles with extended reaction time.13,20 After complete decomposition, rehydrogenation of decomposed products occurred immediately after H2 dissolution with a yield comparable to that of the initial hydrogenation. Compared to other known tertiary amines2,9 that facilitate the hydrogenation of Al*, NMPy has several advantages. The gaseous nature of TMA at room temperature complicates the setup and recovery processes (of TMA),41 and the solid nature of TEDA and quinuclidine requires additional steps to purify the hydrogenation products from the amines.13,14 However, NMPy is a liquid and the hydrogenation product, (NMPy)2·AlH3, can be easily separated and purified. Additionally, (DMEA)2·AlH3 is “too” unstable,17 and TEDA·AlH3 and (quinuclidine)2·AlH3 are “too” stable13,45 for further manipulations or transformations. In contrast, (NMPy)2·AlH3 has an intermediate stability, and is thus more favorable for energy storage applications. Hydrogenation reactions of Al* with NEPy were also attempted; however, no consumption of H2 was observed even under elevated pressures (up to ∼2400 psi, see SI). Accordingly, no alane was detected by IR spectroscopy. This can be attributed to the significantly reduced stability of NEPy·AlH3 because of the distorted geometry of the ethyl group in NEPy·AlH3 compared to the unstrained structure of NEPy (see SI). Similar effects were also observed in other amine·alane molecules with ethyl groups (such as TEA·AlH3).12,15 On the other hand, during the hydrogenation process, AlH3 formed on the Al surface binds strongly to the surface, hindering further hydrogenation by blocking the access of H2 to Al atoms.46 To “harvest” the surface adsorbed AlH3, the binding energy between the stabilizing reagent (usually a tertiary amine) and AlH3 has to be higher than that between AlH3 and the Al surface. In the case of NEPy, the interactions between NEPy and AlH3 were not strong enough to pull off the surface adsorbed AlH3. Consequently, no NEPy·AlH3 was formed in solution by hydrogenation. Thermal Decomposition. For NMPy alane adducts, previous decomposition studies focused on the reaction kinetics and the release of H2 gas from these complexes directly.20 In our studies, however, we wished to use these alane adducts as intermediates to recover donor-free metal hydrides (e.g., AlH3 or LiAlH4) as hydrogen storage materials. In this regard, we conducted studies by heating a mixture of an alane complex and a catalytic amount of metal hydrides (e.g., LiH or LiBH4) in toluene or Et2O under reduced pressures. Unfortunately, no AlH3 was identified under a variety of conditions, including different hydride catalysts, solvents, temperatures, and pressures. In addition, we noticed that the decomposition process was complicated by evaporation and/or sublimation of the alane complexes.

Figure 5. (a) Reactor pressure during the hydrogenation reactions at 1 and 20 °C (final pressures listed on the right). Experiment setting: 1.200 g of Al*, 20 mL of NMPy, and 80 mL of Et2O in a 300 mL reactor. (b) Reactor pressure during the thermal decomposition of (NMPy) 2·AlH3 generated in situ. In the first and second cycles, the reactor pressure increased by 47.7 and 7.4 psi, respectively.

reached equilibrium in a few minutes. The slow pressure drop afterward was caused by a reaction of Al* with H2. Although no induction period was present in this case, previous studies on similar materials occasionally showed an induction period,3 which we assumed was due to the aluminum oxide coating on Al*. Depending on the reaction rate, the amount of time required to reach the hydrogenation equilibrium ranges from a few hours to several days. In addition, previous studies showed that cooling the reaction mixture could enhance the reaction rate.3,14,15 To study this effect, we performed hydrogenation reactions at 1 and 20 °C. In either case, H2 adsorption occurred immediately after dissolution. However, the reaction at 1 °C was faster than that at 20 °C as illustrated by the faster drop in reactor pressure at 1 °C. Consequently, the reaction at 1 °C reached equilibrium in 15 h with a total pressure drop of 68.7 psi, while that at 20 °C resulted in a pressure drop of only 52.6 psi. Despite the initial differences, stirring the reaction mixture at 20 °C for 2 to 3 days eventually caused a total pressure drop close to 68.7 psi. Thus, lowering the temperature increases the reaction rate, but has a negligible effect on the final yield. Once hydrogenation reached equilibrium, the reaction mixture was brought to atmospheric pressure. After standard workup, a white solid was obtained in ca. 61% yield. We noted that the isolated yield was usually slightly lower than that based on the pressure drop because partial decomposition of the hydrogenation product occurred during the workup process. As shown in Figure 3, the product isolated from hydrogenation exhibited similar IR spectra and Al−H absorptions (in both a 2632

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Figure 6. (a) XRD patterns of LiAlH4, Al metal, and the decomposition products of (NMPy)2·AlH3 with varying amounts of LiH; (b) IR spectra of (NMPy)2·AlH3, AlH3, LiAlH4, and the decomposition product in KBr pellets.

presence of LiH could be combined to generate nonsolvated LiAlH4.

Although Al metal was the major product in our studies, we observed a small amount of LiAlH4 in the product when a catalytic amount of LiH was used (as catalyst). To understand the role of LiH in the reaction, we attempted decomposition studies with varying amounts of LiH. As illustrated in Figure 6a, the decomposition products varied from pure Al metal in the case of no LiH, through a mixture of Al metal and LiAlH4 when less than 1.0 equiv of LiH (with respect to 1.0 equiv of (NMPy)2·AlH3) was used, to nearly pure LiAlH4 when 1.0 equiv of LiH was present. Accordingly, the IR spectra in Figure 6b clearly show the disappearance of Al−H absorption from (NMPy)2·AlH3 and the appearance of Al−H absorptions from LiAlH4 during the decomposition process. Considering the fact that (NMPy)2·AlH3 can be synthesized by a direct reaction of LiAlH4 with NMPy,17 the “reverse” process discussed above is particularly interesting and can be applied to other amine·alane adducts. Decomposition of NMPy·AlH3 and NEPy·AlH3 yielded similar results: without LiH, both alane adducts gave only Al metal; with 1.0 equiv of LiH, LiAlH4 was obtained in nearly quantitative yields in both cases. Thus, LiH plays an important role in the decomposition pathway of amine·alane adducts. Although these N-alkylpyrrolidine alane adducts do not decompose to give AlH3 directly, they can be coupled with LiH for the generation of nonsolvated LiAlH4,3,47 a useful reducing agent in a variety of organic transformations. Et 2O

2NMPy + Al* + 1.5H 2 ⎯⎯⎯⎯→ (NMPy)2 ·AlH3

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The IR spectrum of NEPy·AlH3, hydrogenation studies of Al* with NEPy, 1H and 13C NMR spectra of the alane complexes, and the optimized geometries for AlH3, Et2O, NMPy, NEPy, and the related alane complexes. This material is available free of charge via the Internet at http://pubs.acs.org.

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50 ° C

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Tel: 631-344-3511. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS We thank Yusuf Celebi, Weimin Zhou, James Wegrzyn, James Reilly, and John Johnson for assistance. C.N. acknowledges financial support from the Goldhaber Distinguished Fellowship Program at Brookhaven National Laboratory (BNL). We thank the BNL Chemistry Department for access to the NMR spectrometer. This work was carried out at BNL under Contract No. DE-AC02-98CH10886 with the U.S. Department of Energy and supported by its Office of Basic Energy Sciences. Calculations were carried out in large part using the Computational Cluster at the BNL Center for Functional Nanomaterials (CFN) under a user proposal by J.T.M.

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vacuum

(NMPy)2 · AlH3 + LiH ⎯⎯⎯⎯⎯⎯→ LiAlH4 + 2NMPy↑

ASSOCIATED CONTENT

S Supporting Information *

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REFERENCES

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CONCLUSIONS We have shown the synthesis and characterization of several alane complexes and the reversible formation of (NMPy)2·AlH3 via direct hydrogenation of Al*. Compared to other amine·alane adducts that can be formed via hydrogenation of Al*, (NMPy)2·AlH3 exhibits more favorable properties for energy storage and may permit further transformation/ manipulation (such as transamination) for the recovery of donor-free AlH3. In addition, through thermal decomposition studies of these N-alkylpyrrolidine·alane adducts, we established the role of LiH in the decomposition reaction, and showed that the hydrogenation of Al* with NMPy and the decomposition of N-alkylpyrrolidine·alane adducts in the 2633

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