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2009, 113, 16568–16570 Published on Web 08/31/2009
Electrostatic Cooperativity of Hydroxyl Groups at Metal Oxide Surfaces Jean-Franc¸ois Boily*,† and Roberto D. Lins‡ Department of Chemistry, Umeå UniVersity, SE 901 87 Umeå, Sweden, and Pacific Northwest National Laboratory, P.O. Box 999, Richland Washington 99352 ReceiVed: June 30, 2009; ReVised Manuscript ReceiVed: August 21, 2009
The O-H bond distribution of hydroxyl groups at the {110} goethite (R-FeOOH) surface was investigated by molecular dynamics. This distribution was strongly affected by electrostatic interactions with neighboring oxo and hydroxo groups. The effects of proton surface loading, simulated by emplacing two protons at different distances of separation, were diverse and generated several sets of O-H bond distributions. DFT calculations of a representative molecular cluster were also carried out to demonstrate the impact of these effects on the orientation of oxygen lone pairs in neighboring oxo groups. These effects should have strong repercussions on O-H stretching vibrations of metal oxide surfaces. Metal oxide surfaces play a commanding role in a variety of environmental and technological processes.1,2 Many of their catalytic capabilities are attributed to reactive hydroxyl functional groups that control the spatial distributions and reactivities of adsorbed molecules. Although these groups are certainly affected by the number of underlying metal atoms to which they are coordinated, hydrogen bonds with neighboring (hydr)oxo groups are also of paramount importance. Extensive networks of hydrogen bonds can have a considerable impact on surfacecatalyzed reactions. The structures, dynamics, and stabilities of these networks are however not well resolved. The environmentally and technologically important {110} surface of goethite (R-FeOOH) (Figure 1) exhibits a number of well-controlled possibilities for networks of hydrogen bonds that are amenable to both experimental and theoretical studies.2-8 Hydrogen bonding can notably occur between µ3-O(H) and -O(H) groups and between adjacent -O(H) groups of this surface (Figure 1). Interactions among other groups are, on the other hand, presumed to be negligible due to steric hindrances. They are, however, likely to be considerably more varied on stepped/kinked surfaces, such as the ones studied in the laboratory. In a recent study,7 the gas-phase O-H stretching vibrations of these groups were studied on synthetic goethite particles. The bands varied from discrete-like (10-25 cm-1 half width maxima) to nearly unresolvable broad features that undergo important changes in position and intensity with changes in proton loading. Nonetheless, over 94% of the variance of the spectral changes could be explained by only two independent spectral components. Each component consisted of a mixture of correlated O-H bands but with positions spanning over 250 cm-1 units. Although such spreads in band positions are usually indicative of different surface species,9-11 they are, in truth, more likely to have arisen from slight variations in O-H distances of otherwise structurally indistinguishable hydroxyls. Theoretical * E-mail:
[email protected]. † Umeå University. ‡ Pacific Northwest National Laboratory.
10.1021/jp906124a CCC: $40.75
Figure 1. The {110} surface of goethite (Pbnm convention). Left: map view. Right: along the c axis and showing the µ3-OH · · · Ohydrogen bond in dashed lines at the surface and in the bulk (large open circles, surface O; small open circles, bulk O; gray circles, Fe; dark circles, H).
studies have in fact uncovered band shifts of the order of 100-175 cm-1 per pm change in O-H distance.8,12-14 Such changes can be caused by variations in metal-oxygen and/or hydrogen bond strengths. The latter, in particular, should involve cooperating networks of neighboring (hydr)oxo groups which are expected to operate on at least two levels: (1) the formation/ rupture of hydrogen bonds involving interconnected neighboring groups15-20 and (2) electrostatic attractive/repulsive forces brought by noncoordinated neighbors.21,22 Recognition of these effects is of special importance in reactions where molecules disrupt hydrogen bonds by binding to, or displacing, hydroxyl groups. In this study, we use molecular dynamics (MD) to illustrate the impact of electrostatic cooperativity effects on O-H bond strengths brought by noncoordinated neighboring hydroxyl groups. We focus on a nearly proton-depleted {110} goethite surface dominated by -O, µ-O, and µ3-OI,II,III sites (Figure 1) to isolate these effects and to avoid networks of interconnected hydroxyls. Details of the calculations can be found in the Supporting Information. A set of preliminary calculations confirmed that singly coordinated oxygen atoms (-O) possess the strongest affinity for protons (Figure 2, position 1 in 1OH). This group is a particularly reactive center, compared to the other sites, given its low metal atom coordination number. Five different configurations of surface protonation were investigated to test these electrostatic effects (Figure 2): (a) one isolated -OH 2009 American Chemical Society
Letters
Figure 2. Configurations of surface hydroxyls along rows of -O and µ3-O on the {110} surface. Hydrogen bonds are shown as dashed lines.
(1OH), (b) two adjacent -OH (2OH, 0.301 nm separation), (c) two -OH separated by one -O (2OH, 0.602 nm separation), (d) two -OH separated by two -O (2OH, 1.204 nm separation), and (e) a full row of -OH (100% OH). Each simulation (300 ps equilibration, followed by 200 ps production runs using a 1as time step) gave rise to symmetric distributions of bulk and surface O-H distances. Each distribution consists of one or several normal-like components whose median and area were resolved by fitting Gaussian functions. The most stable configuration for 1OH involves three different hydrogen bonds with three adjacent µ3-O acceptors. This scheme gives rise to a bimodal distribution of O-H distances. The longest O-H bond length, centered at 0.104 nm, results from a hydrogen bond with the closest adjacent µ3-O site (position 1). The shortest O-H distance (0.096 nm) corresponds to, on the other hand, a hydrogen bond with either two µ3-O acceptors on either side of the closest µ3-O site (position 3). The time spent on these two acceptors is moreover identical to that with the closest µ3-O site, as confirmed by the area of the respective distributions. Overall, this situation alone gives rise to O-H distances ranging from 0.090 to 0.110 nm. The 1OH bimodal distribution is retained when a proton is added onto an adjacent -O (2OH, 0.301 nm). Although this configuration should not be entropically viable, it remains nonetheless an instructive one to evaluate electrostatic coopera-
J. Phys. Chem. C, Vol. 113, No. 38, 2009 16569 tive effects. Hydroxyls are, in this case, strongly and continuously repelled from one another (median H-H distance of 0.352 nm vs O-O of 0.304 nm) without, however, restricting their migration between three adjacent µ3-O acceptors. The broad distribution of H-H distances (Figure 3b) moreover underscores a considerable degree of decoupled migrations between the various hydrogen bond acceptors. The range of O-H bonds is consequently extended by up to 0.005 nm, resulting in a broadening of 0.01 nm in the bond length distribution. Separating these two hydroxyls by one -O group (2OH, 0.602 nm) results in a drastic reduction in this electrostatic repulsion. The resulting median of H-H distances of 0.629 nm (shown as the 0.315 nm value in Figure 3b to facilitate comparison) nonetheless indicates that repulsive forces are still mutually exerted on neighboring hydroxyls. The narrower distribution of H-H distances underscores an increased coupling in the migrations of OH groups. The distribution of associated O-H distances is comparable to those of 1OH but consists of a third normal distribution with a median at 0.100 nm. This distribution arises from electrostatic repulsive effects on bond lengths when one hydroxyl is in position 1 and the second in position 3 or 3′ (Figure 2). The neighboring OH groups never simultaneously form a hydrogen bond in position 3. A separation of 1.204 nm causes further changes in the distribution of O-H distances. In this case, H-H repulsion is at its lowest with a median at 1.212 nm (i.e., 0.303 nm for normalized H-H distances). H-H distances should consequently be affected by atoms located within a 1.204 nm radius of any adsorption center. The associated O-H distribution is, on the other hand, the most convoluted of all configurations considered for this study and consists of five overlapping components. The first two (0.096 and 1.030 nm) correspond to coupled motions of protons at positions 1 and 3′, respectively (Figure 2). The third distribution (0.100 nm) arises from both protons at position 1, as in the 2OH 0.602 nm case. The fourth and fifth distributions (0.089 and 0.110 nm) correspond to hydroxyl groups hydrogen-bonded to two adjacent µ3-O groups (positions 3′ and 2). Finally, a surface containing a fully protonated row of -OH groups (100% OH) was used to investigate the effects of site saturation. This scenario corresponds to the highest proton loading possible without forming interconnected networks of hydrogen bonds. It is characterized by a normal distribution of O-H distances with a median at 0.100 nm and an associated median H-H distance of 0.301 nm. Proton migration patterns
Figure 3. Distributions of O-H (a) and H-H (b) distances of adjacent hydroxyls from the 200 ps production run. H-H distances are rescaled to the distance of separation between two adjacent -O groups along the c axis (0.301 nm) to facilitate comparison.
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Figure 4. Laplacian of the electron density of O valence electrons of the {110} surface of diaspore (R-AlOOH). Lone pairs are represented as peaks/maxima. The dashed line indicates the -OH · · · O-µ3 hydrogen bond. The arrows denote the oriented lone pairs toward the site of the proton.
are however more decoupled given the comparatively large breath of the H-H distribution. These simulations have consequently demonstrated how mere changes in proton surface loading affect distributions of O-H and H-H distances. Because lone pairs are involved in the stabilization of hydrogen bonds, they may also be affected by electrostatic cooperativity effects. To illustrate this point further, a topological analysis of the electron density23 was carried out from density functional theory calculations of a gas-phase cluster exhibiting the 1 OH configuration (Figure 2). The Laplacian of the electron density of this cluster reveals the loci and orientations of lone pairs in the cluster, seen as maxima in Figure 4. In this diagram, the lone pairs of the two -O and all three µ3-O groups are oriented toward the proton of -OH (position “1”). Another set of calculations however showed that these lone pairs can be readily reoriented if a hydrogen bond is formed with an adjacent µ3-O. As such, time-averaged distributions of hydrogen bonds and lone pairs should be coupled and dictated by the configuration of the hydrogen bonding network (e.g., 1 OH vs 2 OH). These concerted effects could moreover affect interactions between lone pairs and adsorbing molecules. For example, adsorption geometries of hydrogen-bonded species could be constrained by such effects. Elucidation of such possibilities should consequently warrant further studies along these lines. The results of this study are currently being used in the context of our experimental vibration spectroscopic studies of metal (hydr)oxide surfaces. They are most notably providing alternative interpretations to subtle frequency shifts and/or band splitting in spectra measured in the laboratory. These observations, linked to the high sensitivity of O-H bond distances with frequency (100-175 cm-1 per pm) and further computational efforts aimed at unraveling cooperativity effects in interconnected networks (i.e., larger proton loadings), will improve our knowledge of the molecular controls of adsorption reactions on metal oxide surfaces. Acknowledgment. This work was supported by the Kempe and Wallenberg Foundations (Sweden) and by the U.S. Department of Energy, Office of Basic Energy Sciences (Geosciences) Research Program. The MD calculations were performed in the
Letters Environmental Molecular Sciences Laboratory, a national scientific user facility sponsored by the Department of Energy’s Office of Biological and Environmental Research and located at Pacific Northwest National Laboratory. Pacific Northwest National Laboratory is operated by Battelle Memorial Institute for the U.S. Department of Energy under contract DE-AC0676RLO 1830. The quantum chemical calculations were carried out at the High Performance Computing Center North (HPC2N) cluster of Umeå University. Supporting Information Available: Computational details. This material is available free of charge via the Internet at http:// pubs.acs.org. References and Notes (1) Ertl, G.; Kno¨zinger, H.; Weitkamp, J. EnVironmental Catalysis; Wiley VCH: 1999. (2) Schwertmann, U.; Cornell, R. M. The Iron Oxides; Wiley-VCH: 2003. (3) de Leeuw, N. H.; Cooper, T. G. Geochim. Cosmochim. Acta 2007, 71, 1655. (4) Rustad, J. R.; Felmy, A. R. Geochim. Cosmochim. Acta 2005, 69, 1405. (5) Boily, J.-F.; Szany, J.; Felmy, A. R. Geochim. Cosmochim. Acta 2006, 70, 3613. (6) Kerisit, S.; Cooke, D. J.; Marmier, A.; Parker, S. C. Chem. Commun 2005, 24, 3027. (7) Boily, J.-F.; Felmy, A. R. Geochim. Cosmochim. Acta 2008, 72, 3338. (8) Rustad, J. R.; Boily, J.-F. Am. Mineral., in press. (9) Russell, J. D.; Parfitt, R. L.; Fraser, A. R.; Farmer, V. C. Nature 1974, 248, 220. (10) Kno¨zinger, H.; Ratnasmany, P. Catal. ReV.sSci. Eng. 1978, 17, 31. (11) Morterra, C.; Magnacca, G. Catal. Today 1996, 27, 497. (12) Dyan, A.; Cenedese, P.; Dubot, P. J. Phys. Chem. B 2006, 110, 10041. (13) Digne, M.; Sautet, P.; Raybaud, P; Euzen, P.; Toulhoat, H. J. Catal. 2002, 211, 1. (14) Ohno, K.; Okimura, M.; Akai, N.; Katsumoto, Y. Phys. Chem. Chem. Phys. 2005, 7, 3005. (15) Chen, B.; Ivanov, I.; Klein, M. L.; Parrinello, M. Phys. ReV. Lett. 2003, 91, 215503. (16) Barnes, P.; Finney, J. L.; Nicholas, J. D.; Quinn, J. E. Nature 1979, 282, 459. (17) Ruckenstein, E.; Shulgin, I. L.; Shulgin, L. I. J. Phys. Chem. B 2007, 111, 7114. (18) Znamenskiy, V. S.; Green, M. E. J. Chem. Theor. Comput. 2007, 3, 103. (19) de la Paz, M. L.; Gonzalez, C.; Vincent, C. Chem. Commun. 2000, 5, 411. (20) Carcabal, P.; Jockusch, R. A.; Hunig, I.; Snoek, L. C.; Kroemer, R. T.; Davis, B. G.; Gamblin, D. P.; Compagnon, I.; Oomens, J.; Simons, J. P. J. Am. Chem. Soc. 2005, 127, 11414. (21) Bordiga, S.; Roggero, I.; Ugliengo, P.; Zecchina, A.; Bolis, V.; Artioli, G.; Buzzoni, R.; Marra, G.; Rivetti, F.; Spano, G.; Lamberti, C. J. Chem. Soc., Dalton Trans. 2000, 21, 3921. (22) Gomez, M. A.; Pratt, L. R.; Kress, J. D.; Asthagiri, D. Surf. Sci. 2007, 601, 1608. (23) Bader, R. F. W. Atoms In Molecules: A Quantum Theory; Clarendon Press: Oxford, U.K., 1993.
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