16171
2008, 112, 16171–16173 Published on Web 10/01/2008
Modeling the Hydrogen Storage Materials with Exposed M2+ Coordination Sites Monica Kosa,*,† Matthias Krack,‡ Anthony K. Cheetham,§ and Michele Parrinello† Computational Science, Department of Chemistry, and Applied Biosciences, ETH Zurich, c/o USI Campus, Via Giuseppe Buffi 13, CH-6900 Lugano, Switzerland, Paul Scherrer Institut, CH-5232 Villigen PSI, Switzerland, and Department of Materials Science and Metallurgy, UniVersity of Cambridge, Pembroke Street, Cambridge CB2 3QZ U.K. ReceiVed: July 20, 2008; ReVised Manuscript ReceiVed: September 12, 2008
Recent experimental work has shown that exposed coordination sites of metals can enhance the binding of hydrogen in nanoporous inorganic and hybrid inorganic-organic framework materials. First principles calculations have been carried out on molecular analogs of such sites in order to compare the binding of dihydrogen molecules to Ni2+ and Mg2+. The binding to Mg2+ is found to be 1 kJ/mol or less, whereas that to Ni2+ varies between 6 and 23 kJ/mol, depending on the geometrical arrangement of the ligands. Analysis of the computational results shows that the preferred binding to Ni2+ is consistent with the general binding mode of H2 to metal centers. The implications of these results for the design of superior materials for hydrogen storage are discussed. The safe and efficient storage of hydrogen is one of the major challenges relating to the future use of fuel cells for vehicular transportation.1 There are two broad classes of hydrogen storage media. Metal hydrides and chemical hydrides store hydrogen by a chemisorptive mechanism that involves heats of adsorption that are typically greater than 25 kJ/mol.2 Porous materials, such as activated carbons, zeolites, clathrates, and metal organic frameworks (MOFs), bind hydrogen by a physisorptive mechanism with a much smaller heat of adsorption in the range of 5 kJ/mol or less.3 However, a few examples of zeolites with extra framework Mg2+ ions have shown enhanced adsorption enthalpies as a consequence of the enhanced charge-quadrupole interactions between hydrogen and the low coordination number cations.4 Metal hydrides and chemical hydrides often suffer from poor reversibility and kinetics, while the low heats of adsorption of porous media imply that they can only store useful amounts of hydrogen at low temperatures. However, a series of recent papers on both inorganic molecular sieves and porous metal organic frameworks have shown that the heat of adsorption in porous materials can be enhanced when under-coordinated metal sites are created within the framework.5 Specifically, exposed Ni2+ sites in the nickel phosphates, VSB-5,6 and the hybrid NaNi3(OH)(5-sulfoisophthalate)2 (Ni-SIP),7 as well as Mn2+ and Cu2+ sites8,9 in other MOFs, bind hydrogen with the heat of adsorption of 10 kJ/mol or greater. In all of these systems, the experimental data show that the adsorption occurs without the dissociation of the dihydrogen bond, similarly to the manner in which hydrogen is bound in the well known Kubas complexes.10 Nevertheless, these new porous materials still do not meet the DOE target of >6 wt % at room temperature. The replacement of the divalent metal ions (e.g., Mn2+, Ni2+, Cu2+, Ni2+) by * To whom correspondence should be addressed. E-mail: monica.kosa@ phys.chem.ethz.ch. † ETH Zurich, c/o USI Campus. ‡ Paul Scherrer Institut. § University of Cambridge.
10.1021/jp806394g CCC: $40.75
SCHEME 1: Schematic Representations of the Calculated Clusters
lighter Mg2+ may offer a solution, though attempts to make such materials have not yet yielded stable systems with exposed Mg2+ sites. In order to shed light on the viability of MOFs with exposed magnesium framework sites for hydrogen storage, we have performed first principles calculations on a model cluster that is based upon the available experimental structures.11 The exposed M2+ coordinated sites were modeled by a neutral square pyramidal cluster, ML3L′2, with L ) CH3OCH3 and L′ ) OCH3-. The precise starting structures were varied by placing the two OCH3- groups in cis, trans, and axial positions, Scheme 1. Each structure was minimized with the Gaussian 03 package12 prior to introducing the dihydrogen molecule, always keeping the exposed face planar. The resulting clusters were further minimized in the presence of dihydrogen, again keeping the face planar. All the structures were calculated at the BLYP/6-311G(d) level of theory. In order to test the credibility of our approach, calculations were performed using several methods and varying size basis sets.13 The basis set superposition error (BSSE) was calculated with the counterpoise method, suggested by Boys and Bernardi,14a and implemented in G03,14b together with geometry optimization. The range of the hydrogen binding energies for the magnesium complexes, calculated on different levels of theory, is ∼10 kJ/mol, which we shall show is less than the difference between the Mg and the corresponding Ni clusters; full details can be found in the Supporting Information. 2008 American Chemical Society
16172 J. Phys. Chem. C, Vol. 112, No. 42, 2008
Letters
Figure 1. Calculated complexes. Color assignment: blue, nickel; green, magnesium; red, oxygen; cyan, methyl group; gray, hydrogen.
TABLE 1: Binding Energies, kJ/mol, and NPA Charges18 of the Three Possible Isomers of M(OCH3)2(O(CH3)2)3, M ) Mg, Ni, Calculated at BLYP/6-311G(d) ∆E
isomer Mg-cis Mg-trans Mg-axial Ni-cis Ni-trans Ni-axial a
∆E
CP
CP a
ΝPA charge of the metal center
0.08 1.13 0.33 -10.3 -6.36 -23.78
1.74 1.75 1.73 1.03 1.07 1.03
is calculated according to eq 1.
The calculated species are shown in Figure 1 and the calculated binding energies, estimated according to eq 1, are given in Table 1.
∆ECP ) ECP[M(MeO)2(Me2O)3 · H2] {E[H2] + E[M(MeO)2(Me2O)3]}(1) All Mg complexes show very weak binding to the H2 molecule. In the trans and axial isomers, the H2 lies very far from Mg and the binding is very weak. In the free H2 (calculated at the same level of theory, BLYP/6-311G(d)) the H-H bond length is 0.740 Å, and it remains practically unchanged (0.746 and 0.747 Å) when H2 is weakly bound to the complex (in the trans and axial isomers, respectively). In the cis isomer, the H2 resides away from Mg, but closer to one of the oxygens. Again, the interaction, as indicated by the interaction energy and the bond length change for H2, is very weak (H-H ) 0.750 Å). The “true H2 complex” is characterized by an H-H distance in the range 0.8-1.0 Å.15 Thus we conclude that in the models considered in this study, magnesium does not form stable dihydrogen complexes. In the Ni complexes, the H-H bond distances in the cis, trans, and axial isomers are 0.749, 0.746, and 0.818 Å, respectively. This trend is also reflected in the binding energies, which are 10.3, 6.36, and 23.78 kJ/mol. It must be noted, however, that both in the cis and trans isomers, the H2 reside closer to one of the oxygens than to the nickel atom. In the more tightly bound axial isomer, the H2 resides in a side-on mode with respect to the nickel. The calculated natural population analysis (NPA) charges indicate that in the nickel complexes more electron density is
transferred from the ligands to the metal atom compared to magnesium complexes. Thus, in all three isomers, the nickel atom charge decreases from the formal charge of +2 to approximately +1.04 while the magnesium atom charge decreases from the formal charge of +2 to only +1.74. This general difference between nickel and magnesium complexes is due to the fact that in magensium complexes the extra charge (around. 0.20 el. in all three complexes) is almost fully transferred only to the 3s valence orbital while in the nickel complexes the charge is spread among the low lying 3d orbitals and the higher energy 4s orbital (around 8.94 el. in the 3d and 4s orbitals). The better binding ability of dihydrogen to Ni-axial compared to other nickel isomers and magnesium complexes can be understood following the general binding mode of H2 to metal centers. The dihydrogen binds to transition metals via two interactions. The first one is through donation of σ(H-H) electrons to some empty orbital on the metal. The second interaction is via back-donation of the electron density from the metal to the σ*(H-H) orbital. This general rule is maintained in the case of Ni-axial. According to natural bond orbital (NBO) analysis,16 the σ(H-H) bond is depopulated to 1.87 el. and interacts with the hybrid orbital of 4s and 3dz2 of the Ni atom. This hybrid orbital is contributed by the nickel atom to the antibonding orbital of the nickel and the axial oxygen atoms, σ*(Ni-O). The σ*(Ni-O) resides mainly on the nickel atom, rather than on the corresponding oxygen atom and it is populated by 0.27 el. The metal to ligand back-donation in the Ni-axial complex occurs via interaction of the nickel 3dxz orbital with the antibonding dihydrogen orbital, σ*(H-H). Thus, the 3dxz orbital is depopulated to 1.82 el while the σ*(H-H) orbital is populated by 0.18 el. In summary, our study leads to the following interesting conclusions. (i) Exposed Ni2+ sites bind H2 much more effectively than Mg2+; (ii) The binding energy of H2 to the nickel complexes is sensitive to the positions of charged ligands in our model. This is important because it suggests that the heat of adsorption of hydrogen into nickel-containing nanoporous solids is likely to be strongly dependent on the fine details of the crystal structure. The binding energy of H2 to Mg complexes is very small and there is no basis to assume that the small differences between various isomers arise from their different binding abilities (method error could give rise to such differ-
Letters ences); (iii) Our study shows that even in the presence of a transition metal such as Ni2+, H2 can migrate to the neighboring oxygen atoms, as was demonstrated in isomers Ni-cis and Nitrans.17 This is consistent with temperature-programmed desorption and inelastic neutron scattering data on such materials, which show multiple binding sites for hydrogen. (iv) The electronic structure analysis revealed that although all the nickel complexes bear practically the same charge, only in the Niaxial isomer the H2 binding occurs via specific orbital interactions. Acknowledgment. This work was supported by the EERE Program of the U.S. Department of Energy under Award No. DE-FC36-05GO15004. The CSCS is gratefully acknowledged for providing the computational resources. Supporting Information Available: Binding energies of the clusters calculated at high level ab initio and DFT levels of theory. Surface model schemes. This material is available free of charge via the Internet at http://pubs.acs.org. References and Notes (1) (a) Schlapbach, L.; Zu¨ttel, A. Nature 2001, 414, 353. (b) Von Helmolt, R.; Eberle, U. J. Power Sources 2007, 165, 833. (c) van den Berg, A. W. C.; Otero Area´n, C. Chem. Commun. 2008, n/a, 668. (2) (a) Grochala, W.; Edwards, P. P. Chem. ReV. 2004, 104, 1283. (b) Orimo, S.; Nakamori, Y.; Eliseo, J. R.; Züttel, A.; Jensen, C. M. Chem. ReV. 2007, 107, 4111. (3) Bordiga, S.; Vitillo, J. G.; Ricchiardi, G.; Regli, L.; Cocina, D.; Zecchina, A.; Arstad, B.; Bjørgen, M.; Hafizovic, J.; Lillerud, K. P. J. Phys. Chem. B 2005, 109, 18237. (4) For example see: (a) Turnes Palomino, G.; Llop Carayol, M. R.; Otero Area´n, C. J. Mater. Chem. 2006, 16, 2884. (b) Otero Area´n, C.; Turnes Palomino, G.; Llop Carayol, M. R. Appl. Surf. Sci. 2007, 253, 5701. (5) (a) Rowsell, J. L. C.; Yaghi, O. M. Angew. Chem., Int. Ed. 2005, 44, 4670. (b) Kaye, S. S.; Long, J. R. Chem. Commun. 2007, n/a, 4486. (6) Forster, P. M.; Eckert, J.; Chang, J.-S.; Park, S.-E.; Ferey, G.; Cheetham, A. K. J. Am. Chem. Soc. 2003, 125, 1309. (7) Forster, P. M.; Eckert, J.; Heiken, B. D.; Parise, J. B.; Yoon, J. W.; Jhung, S. H.; Chang, J.-S.; Cheetham, A. K. J. Am. Chem. Soc. 2006, 128, 16846. (8) Dincã, M.; Dailly, A.; Liu, Y.; Brown, C. M.; Neumann, D. A.; Long, J. R. J. Am. Chem. Soc. 2006, 128, 16876.
J. Phys. Chem. C, Vol. 112, No. 42, 2008 16173 (9) (a) Peterson, V. K.; Liu, Y.; Borwn, C. M.; Kepert, C. J. J. Am. Chem. Soc. 2006, 128, 15578. (b) Liu, Y.; Brown, C. M.; Neumann, D. A.; Peterson, V. K.; Kepert, C. J. J. Alloys Compd. 2007, 446-447, 385. (10) Kubas, G. J.; Ryan, R. R.; Swanson, B. I.; Vergamini, P. J.; Wasserman, H. J. J. Am. Chem. Soc. 1984, 106, 451. (11) (a) Dinca˜, M.; Long, J. R. J. Am. Chem. Soc. 2005, 127, 9376. (b) Hulvey, Z; Cheetham, A. K. Solid State Sci. 2007, 9, 137. (12) Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Scuseria, G. E.; Robb, M. A.; Cheeseman, J. R.; Montgomery, J. A., Jr.; Vreven, T.; Kudin, K. N.; Burant, J. C.; Millam, J. M.; Iyengar, S. S.; Tomasi, J.; Barone, V.; Mennucci, B.; Cossi, M.; Scalmani, G.; Rega, N.; Petersson, G. A.; Nakatsuji, H.; Hada, M.; Ehara, M.; Toyota, K.; Fukuda, R.; Hasegawa, J.; Ishida, M.; Nakajima, T.; Honda, Y.; Kitao, O.; Nakai, H.; Klene, M.; Li, X.; Knox, J. E.; Hratchian, H. P.; Cross, J. B.; Bakken, V.; Adamo, C.; Jaramillo, J.; Gomperts, R.; Stratmann, R. E.; Yazyev, O.; Austin, A. J.; Cammi, R.; Pomelli, C.; Ochterski, J. W.; Ayala, P. Y.; Morokuma, K.; Voth, G. A.; Salvador, P.; Dannenberg, J. J.; Zakrzewski, V. G.; Dapprich, S.; Daniels, A. D.; Strain, M. C.; Farkas, O.; Malick, D. K.; Rabuck, A. D.; Raghavachari, K.; Foresman, J. B.; Ortiz, J. V.; Cui, Q.; Baboul, A. G.; Clifford, S.; Cioslowski, J.; Stefanov, B. B.; Liu, G.; Liashenko, A.; Piskorz, P.; Komaromi, I.; Martin, R. L.; Fox, D. J.; Keith, T.; Al-Laham, M. A.; Peng, C. Y.; Nanayakkara, A.; Challacombe, M.; Gill, P. M. W.; Johnson, B.; Chen, W.; Wong, M. W.; Gonzalez, C.; Pople, J. A. Gaussian 03, revision D.01; Gaussian, Inc.: Wallingford, CT, 2004. (13) The credibility of the estimation of the geometries and the small binding energies of H2 with organic and metal moieties by DFT was demonstrated also in: Negri, F.; Saendig, N. Theor. Chem. Acc. 2007, 118, 129. (14) (a) Boys, S. F.; Bernardi, F. Mol. Phys. 1970, 19, 553. (b) Simon, S.; Duran, M.; Dannenberg, J. J. J. Chem. Phys. 1996, 105, 11024. (15) Kubas, G. J. Proc. Natl. Acad. Sci. U.S.A. 2007, 104, 6901. (16) For a comprehensive review of NBO methods see: Weinhold, F. Encyclopedia of Computational Chemistry; Schleyer, P. v. R., Allinger, N. L., Clark, T., Gasteiger, J., Kollman, P. A., Schaefer, H. F., III, Schreiner, P. R., Eds.; John Wiley & Sons: Chichester, UK, 1998; Vol. 3, pp 17921811. (17) The interaction of the dihydrogen molecule with nonmetal sites is not specific, and we suspect that other energetically similar isomers can be found where the dihydrogen molecule lies closer to one of the three remaining facial oxygens. Our preliminary studies on the surface models of the corresponding materials have confirmed this assumption. See Supporting Information for details. (18) (a) Reed, A. E.; Weinhold, F. J. Chem. Phys. 1983, 78, 4066. (b) Reed, A. E.; Weinstock, R. B.; Weinhold, F. J. Chem. Phys. 1985, 83, 735.
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