Reactivity of Aluminum Cluster Anions with Water: Origins of Reactivity

J. Phys. Chem. A 2010, 114, 6071–6081

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Reactivity of Aluminum Cluster Anions with Water: Origins of Reactivity and Mechanisms for H2 Release Arthur C. Reber and Shiv N. Khanna* Physics Department, Virginia Commonwealth UniVersity, 701 W. Grace Street, Richmond, Virginia 23284-2000

Patrick J. Roach, W. Hunter Woodward, and A. W. Castleman, Jr.* Departments of Chemistry and Physics, PennsylVania State UniVersity, UniVersity Park, PennsylVania 16802 ReceiVed: NoVember 23, 2009; ReVised Manuscript ReceiVed: March 29, 2010

The reactivity of aluminum anion clusters with water was found to exhibit variations with size, with some clusters exhibiting negligible reactivity, others absorbing one or more water, while even others releasing H2 with addition of multiple waters. (Roach, P.J., Woodward, W.H. et al. Science, 2009, 323, 492). Herein, we provide further details on the role of complementary active sites in the breaking of the O-H bond on the cluster. We examine the reactions of Aln- + H2O where n ) 7-18, and show how the complementary active sites may be best identified. The clusters with active sites are found to be reactive, and clusters with barriers to reactivity have an absence of paired active sites. The role of charge in the reactivity is considered, which could account for the observed increase in reactivity at large sizes. The H2 release in the reactivity of Al17with multiple water molecules is also studied by comparing multiple reaction pathways, and the selective H2 production is explained by the first water inducing a new active site. A mechanism for transferring hydroxyl groups on the surface of the cluster is also discussed. 1. Introduction A basic understanding of factors controlling the stability, physical properties, and chemical reactivity of small atomic clusters is not only important for understanding fundamental science at reduced sizes, but also exploiting unique properties of clusters in applications.1–14 Early experiments on mass spectra of sodium clusters revealed that cluster stability fluctuates with size and that certain “magic clusters” were more stable than others.15 These evolutions were explained using a simple model where valence electrons responded to a uniform positive background formed by smearing the charge of the nuclei into a spherical region determined by the cluster size.16,17 For such a background charge, quantum confinement results in an electronic spectrum marked by discrete bunches of levels (1S, 1P, 1D, 2S,...) in analogy to the atomic levels (1s, 2s, 2p,...) but with different combinations of quantum numbers.18 The magic clusters corresponded to clusters with filled electronic shells showing the role of the electronic structure on cluster stability. Later experiments also showed that at larger sizes, the more stable clusters correspond to sizes with filled geometric shells.19 Together, these findings indicated that the stability of clusters have both an electronic and geometric component, with the electronic component dominating at small sizes and the geometry becoming important at larger sizes. These developments raised the question if the electronic and geometric features control other properties, in particular the chemical character5 including reactivity and catalysis. The appearance of electronic shells raised the intriguing question if the partial or complete filling of the electronic shells could lead to chemical behaviors as in atoms. In particular, could one assign a chemical valence to clusters that would predict * To whom correspondence should be addressed. E-mail: snkhanna@ vcu.edu (S.N.K.), [email protected] (A.W.C. Jr.).

the nature of compounds they form and ultimately allow their classification as in the periodic table of elements? This idea started several investigations into the properties of clusters and, in particular, understanding of the observed reactivity of aluminum clusters with oxygen. The etching experiments involving reaction of aluminum cluster anions with oxygen had shown that clusters such as Al13-, Al23-, and Al37- did not react,5,7–10 thus exhibiting an “inert” tendency even though bulk aluminum reacts readily with oxygen.20–27 These clusters have 40, 70, and 112 valence electrons, which correspond to filled electronic sub shells and a large HOMO-LUMO gap within the model of a nearly confined electron gas resulting in effective zero valence as in inert gas atoms.9 Furthermore, the removal of an electron led to an Al13 with an electron affinity of 3.34 eV and an effective chemical valence of one much like the halogen atoms. Further work on such chemical analogies led to the proposition that selected clusters could be regarded as superatoms (originally termed unified atoms5) forming a third dimension to the periodic table, and over the past few years superatoms mimicking inert,28–32 halogen,33–35 alkaline earth,36,37 and magnetic atoms38 have been proposed, and recent findings have led to a quantification of their characteristics.39 While superatoms and elemental atoms share a commonality of electronic shells controlling chemical characteristics, the superatoms do possess additional attributes, namely, a nonspherical geometrical structure. Can this additional degree of freedom lead to defining chemical features? In this paper, we focus on our findings that the reaction of water with aluminum cluster anions, which highlight the role of geometric factors on reactivity. In particular, we show (1) that some clusters with open electronic shells such as Al20- do not reveal any significant reactivity with water whereas (2) selected clusters, even including some with closed electronic subshells, can exhibit strong reactivity toward water. Theoretical

10.1021/jp911136s  2010 American Chemical Society Published on Web 04/13/2010

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analysis shows that these behaviors are rooted in the nonuniform distribution of charge densities on cluster surfaces and hence have a geometric rather than an electronic origin. In an ideal spherical metal particle, the charge density will be evenly distributed over the surface of the sphere; however, in the cluster size regime, clusters are rarely spherical. Jahn-Teller effects40,41 caused by an incomplete electronic shells cause distorted geometries, and the stability of certain geometric motifs result in a preference for structures with one or more adatoms or defects on the surface of a stable core.42,43 The distortions and defects break the degeneracy in the expected shell structure and result in an uneven charge distribution where different surface sites on the cluster accept or donate charge unequally. This allows even the electronically closed-shell clusters such as Al23and Al37- to be reactive. Our initial findings, reported recently in Science,1 drew from the fact that the reactivity of aluminum clusters with nucleophiles, such as water, depends strongly on the charge transfer tendencies of the surface sites and varies depending on the cluster’s Lewis acidity at different sites on the cluster.44–47 As we show, the reactivity entails a pair of complementary active sites, one of which plays the role of Lewis base while the other acts as a Lewis acid,48 and that their close proximity to each other is the reason for the splitting of the O-H bond.49,50 We believe that our results illustrate a model case to identify active sites51–57 that have a propensity to accept or donate charge on the surface of the cluster and reveal how these sites may lead to novel reactive patterns. While some of our initial findings were reported in a recent paper,1 we provide here a more detailed description of the reactivity of aluminum clusters with waters. We explore the correlation between some general features of aluminum anion clusters and reactivity such as the absolute position of the HOMO and LUMO, the gap, the dipole moment, and the polarizability. Furthermore, we examine the binding energies, transition states, and frontier orbitals of Aln- n ) 7-18, and show the presence of prominent complementary active sites in Al12- and Al16-18-, which also display the highest reactivity. Al11- and Al13- have especially high energy transition states, and consequently decreased reactivity. The use of charge densities and integrated charges are used in order to understand what tools best indicate the presence of active sites. We also look at the effect of the charged state on the reactivity by carrying out theoretical studies on neutral and cationic species. Numerous multiple-water reaction pathways are examined to understand the selective production of H2 after binding two water molecules in Al16-18-. The lowest energy pathway for H2 production is given for Al17-, (The pathways for Al16- and Al18are given in the supplement) and many alternate reaction pathways are considered. The selective H2 production58–66 is found to be caused by an induced active site on the opposite side of the cluster. We also examine Eley-Rideal and direct H transfer reactions, which are found to be poor because the H atom serves as a poor Lewis base. Finally, details are presented of a mechanism we propose that accounts for transferring H atoms and moving the hydroxyl groups bound to a cluster. 2. Methods 2.1. Experimental Methods. Our experimental setup has been described in detail previously,1 and will only be briefly discussed here. Clusters were created in a laser vaporization source consisting of a rotating and translating aluminum rod (99.999%, Puratronic) which was ablated using the second harmonic (532 nm) of a 30 Hz Nd:YAG laser. Clusters were formed in the presence of a helium backing gas (High Purity,

Reber et al. Praxiar, Inc.) which flowed continuously at approximately 8000 standard cubic centimeters per minute. Clusters were carried out of the source through an expansion nozzle into a laminar flow tube that was maintained at a pressure of approximately 0.7 Torr by a high volume pump. Deionized water was introduced into the flow tube from a glass bubbler via a reactant gas inlet. Water concentration was controlled using a needle valve and by heating the bubbler. Reactants and products were sampled through a 2 mm orifice and guided through a series of differentially pumped chambers via RF octopoles and electrostatic lenses before being analyzed using quadrupole mass spectrometry (Extrel, CMS). The ion optics and detector were controlled and monitored using a personal computer. 2.2. Theoretical Methods. First principles electronic structure investigations were carried out to understand the nature of reaction and the release of hydrogen in selected cases. The studies used a first-principles molecular orbital approach within a gradient-corrected density functional framework. The molecular orbitals are expressed as a linear combination of atomic orbitals that were, in turn, formed via a linear combination of Gaussian functions located at the atomic sites. The exchange correlation contributions are included within the GGA-PBE gradient corrected density functional formalism.67 The calculations were carried out, at an all-electron level, using the Naval Research Laboratory Molecular Orbital Library (NRLMOL) set of codes developed by Pederson and co-workers.68,69 The basis sets are built from a variable number of primitive Gaussians and are based on a total-energy minimization for free atoms, further optimized for all-electron density functional calculations.70 The basis set for Al had 6s, 5p, and 3d Gaussians; that for O had 6s, 5p, and 3d Gaussians; and the basis set for H had 4s, 3p, and 1d Gaussians. The basis sets were supplemented with a diffuse d-Gaussian. Some supplemental calculations were done using the deMon2K density functional code,71 including some QMMD (quantum mechanical molecular dynamics) simulations. The presented results are either local minima, in which the geometry was fully optimized from a number of predicted geometries, or are transition states found through constrained optimizations; as in most cases either the O-H bond distance or H-H bond distance were used as the reaction coordinate. Numerous calculations involving a constrained optimization with a fixed O-H bond distance at different sites and orientations were compared to identify the sites with the lowest transition state energies. Binding energies were calculated using

Eb ) E(reactants) - E(products)

(1)

Reactants are the pure Aln- cluster and H2O, unless otherwise specified. Zero point energy corrections and basis set correction errors were found to be negligible for comparing reactivity between different clusters and are not included, except for the Al17- mechanisms given in Figures 7-9. Extensive structural studies of the aluminum clusters have been previously performed,42,43,72–74 and, using these results along with additional structures generated by a genetic algorithm, we have found the same geometries as previous workers identifying the ground state structures of these rather large clusters. 3. Results and Discussion 3.1. Experimental Results. In a recent paper, we presented our experimental findings, in which Aln- clusters from n ) 7-73 were reacted with water at low pressures. In both the

Reactivity of Aluminum Cluster Anions with Water present and earlier study, D2O and H2O were used interchangeably with no discernible difference in reactivity. The main findings were as follows. (1) Al12- was found to be reactive, forming an Al12(H2O)- species in certain situations and at other times Al12(H2O)2-. (2) Al16-, Al17-, and Al18- also were observed to form complexes with one or more water molecules. (3) At larger sizes, Al24-, Al27-, Al29-, and Al37- were found to be most reactive. (4) In Al16-, Al17-, and Al18- the addition of multiple waters resulted in products that were deficient by exactly two or four hydrogen atoms. We have now performed experiments at increased water concentrations, and we present these results in Figure 1. Figure 1A shows the Aln- spectra where no water is introduced into the fast flow reactor; Figure 1B is with a low amount of heavy water, and Figure 1C is the same spectrum zoomed in on the series AlnD2O-, where n ) 17-19, shown as a gray box in Figure 1B. It is easily observed in Figure 1B that Al12- is the first smaller cluster (n < 20) to adsorb a water molecule, and this peak is labeled. Al24- is also observed to absorb one or two water molecules toward the higher end of the spectrum. The purpose of Figure 1C is to demonstrate the observed hydrogen molecule loss from specific aluminum clusters, and so the peaks labeled represent all instances where D2 is missing from a cluster that has absorbed two or more water molecules. Exactly four m/z units to the right of each labeled peak is another peak belonging to the complex before any D2 is lost, and all other unlabeled peaks can be attributed to odd-numbered water additions or bare aluminum clusters. It was observed that as more water is added, the larger clusters (n > 20) would quickly etch away, followed by the size selective reactivity of the smaller clusters. Therefore, before introducing large amounts of water, the instrument was reoptimized around the cluster distribution n ) 5-26, and this spectrum is shown in Figure 1D. Figure 1E was taken at a “maximum flow,” with the water (nonheavy) flow rate high enough that further increase did not appear to modify the spectrum. At these high concentrations Al8-, Al9-, and Al10- reveal modest reactivity; Al11-, Al13-, and Al20- exhibit enhanced resistance to water etching; and Al23is entirely etched. This is interesting, as Al23- has a closed electronic jellium subshell, yet it reacts with water, whereas clusters without shell closings such as Al20- do not. It is also apparent that no Aln(H2O)m- peaks remain at large water concentrations. In this work, we primarily focus on Aln- where n ) 7-18, because this region has the most variable reactivity. 3.2. The Reactivity of Aln- Clusters (7 e n e 18) with Water. To understand the observed size selectivity, theoretical investigation of the reactivity of Aln- with water, where 7 e n e 18 were undertaken. We evaluated binding energies, transition state energies for the breaking of the O-H bond, and binding energies following dissociative chemisorptions. We also examined a variety of factors that might be expected to control the reactivity, including the HOMO-LUMO gap, the dipole moment, polarizability, and absolute position of the LUMO. We have found that the explanation for the observed spectra is dissociative chemisorption, in which water adducts are observed when the water has sufficient energy from binding and its environment to break the O-H bond to form a covalently bound HAln(OH)- species. These cases are marked by the presence of neighboring active sites, one functioning as a Lewis acid and the other like a Lewis base, which facilitate the splitting of the O-H bond. In cases where the aluminum cluster appears resistant to water addition, it does not necessarily imply that the cluster is immune to nucleophilic attack; rather this observation supports our mechanism in that without the

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Figure 1. Experimental mass spectra of aluminum cluster reactivity with water demonstrating the difference between a partial and full etching reaction. (A) Aluminum cluster distribution before introducing water; (B) Aluminum cluster distribution after introducing a small amount of water; (C) Zoomed-in section of grayed box in B focusing on D2 loss; (D) Retuned aluminum cluster distribution before introducing a large amount of water; (E) Aluminum cluster distribution after introducing a large amount of water. D2O and H2O were used interchangeably with no discernible difference in reactivity. The specific mass spectra shown were acquired using D2O for A, B, and C, and H2O for D and E. The colored lines are to aid the reader in identifying aluminum cluster peaks, with Al12-, Al17-, Al20-, and Al23- represented by green, red, blue, and purple lines, respectively.

complementary active sites and resulting dissociative chemisorptions, the water-cluster complex has enough energy to eject the water before reaching the detector. We now examine how we came to this conclusion by considering both hypotheses. Water has a permanent dipole moment of 1.85 D and therefore the interaction with aluminum clusters is marked by ion-dipole contributions. In addition, the permanent dipole moment can

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Reber et al. identifying higher densities of states near the Fermi energy as a higher density of available states leads to higher polarization.76 Next, we examined if the relative binding energy due to chemical bonding of H2O to the aluminum cluster anions could be responsible for the observed trends and if the rules proposed by Chreˆtian et al. for the binding sites and relative energies for binding of propene to gold clusters77 could be extended to the present case. For each cluster, a H2O molecule was placed in various locations, and the energy was minimized to determine the most stable configurations for an intact water molecule. In Figure 2B we show the progression of the calculated binding energy (black line) with size.77–80 Al12- has a particularly high binding energy of 0.52 eV, which is consistent with the observed Al12(H2O)- peak. The binding energy is defined as

BE ) E(Aln-) + E(H2O) - E[Aln(H2O)-]

Figure 2. (A) The calculated dipole moment of the Aln- clusters from n ) 5-18. B) The binding energy (BE) of an intact water molecule to an Aln- cluster, the transition state energy (TS) for breaking an O-H bond on the cluster, and the binding energy (Rlx) after the O-H bond has been broken. (C) The absolute energy of the HOMO, LUMO, and LUMO+1 of the Aln- clusters where 7 e n e 18.

polarize the aluminum motifs. Consequently, we proceeded to examine if the variation in the dipole moment or the polarizability of anionic clusters74 could be related to the observed trends. First principles theoretical studies were carried out to calculate the dipole moment (Figure 2A) and the average polarizability of the anionic clusters containing up to 17 atoms. The dipole moment was calculated with respect to the center of mass, and while the dipole moments of charged species vary with the choice of coordinate, it gives some information about size-dependent variations. The theoretical results showed that Aln- clusters (n ) 8-11) have the largest dipole moments; however, water does not appear to bind to these species. Alternatively, water was observed to bind with Al12-, which we calculated as having a much smaller dipole moment. The calculated polarizability increased roughly linearly with size, indicating that the observed irregular variations in reactivity are different from the trends in dipole moment and/or polarizability.75 We do note that local polarizability, in which the site specific polarizability is estimated, may prove useful for

(2)

However, the second largest binding energy is Al11-, which has no observed reactivity at thermal energies. Al7- and Al8also have larger binding energies than Al16-18-, whereas the former form no observable complexes but the latter do. The absolute level of the LUMO or LUMO+1, as shown in Figure 2C, is a way of identifying which clusters bind undissociated water most strongly, because binding results in a transfer of negative charge to the cluster. Al13- has a strikingly high LUMO and binds water quite weakly, whereas Al11- and Al12- have low LUMOs versus those around them and bind water most strongly.81 In clusters with an odd number of electrons, the LUMO+1 becomes more important for predicting the binding site. The clusters usually form chemisorbed complexes in which the lone pair of the oxygen atom interacts with the aluminum cluster. If the cluster has an even number of electrons, then the lone pair may mix with the site with the largest LUMO isosurface, thereby forming a weak bond. We refer to sites with the largest LUMO isosurface as the LUMO site, and the site with the largest HOMO charge density as the HOMO site. In most metal clusters with an odd number of electrons, the HOMO and LUMO are half-filled versions of the same level, so that the two electrons in the lone pair cannot mix with the half-filled HOMO/LUMO state. This may be thought of as identifying sites on the cluster which act as Lewis acids; sites which accept a lone pair most readily are best for binding water, whereas the LUMO of odd electron clusters may only accept a single electron. In Figure 3, the LUMO/LUMO+1 charge density is shown under the first column and binding sites for H2O absorption in the second column. In the case of Al9- and Al17-, two even electron clusters, the lone pair on water binds most strongly to the LUMO. In the odd electron clusters Al8-, Al10-, and Al12-, the lone pair binds most strongly with the LUMO+1. Although water binds through the lone pair on oxygen in larger clusters, at smaller sizes or in selected cases the water molecule more favorably forms a hydrogen bond with the cluster. For example, Al7-, and Al13- have the lowest energy states in which the water molecule is intact and whereupon the water is hydrogen bonded to the cluster. These structures are not shown in Figure 3. The strength of the hydrogen bond is enhanced by increased electronic density on the surface of the cluster and exhibits smaller variations in binding energy than lone pair-cluster interactions as all metallic clusters have charge density at the surface. For example, smaller Aln- cluster anions n e 7 will all be more stable in the hydrogen-bonded orientation, as the anionic charge is more concentrated on the smaller

Reactivity of Aluminum Cluster Anions with Water


Figure 3. The geometry of Aln- clusters 7 e n e 14 and the LUMO wave function if n is odd, or LUMO+1 wave function if n is even or n ) 11, are given in the first column. The Initial column is the geometry where the intact water molecule has access to the transition state with the HOMO charge density of the complex. The T.S. column is the geometry of the transition state for splitting the O-H bond, and the Final column is the local minimum geometry after the O-H bond has been broken. An isosurface of 0.03 was used.

clusters, which enhances the strength of the hydrogen bond, while weakening the strength of the lone pair-cluster interaction. Next we consider the possibility of O-H bond cleavage in the interaction of the aluminum clusters with water. As a water molecule binds to the cluster, it gains an initial binding energy. After this initial absorption, the cluster could undergo O-H bond dissociation. As the O-H bond is stretched, the energy first increases upon reaching the highest point (transition state) until the O-H bond is broken and the OH and H bond to different sites. Hence, the system must have enough energy to overcome the barrier for the process to proceed. We calculated the transition state energies for splitting a water molecule on the surface of the cluster as well as the resulting binding energy after the O-H bond is broken. We give the transition state energy as the total electronic energy of the transition state with respect to the cluster and H2O.

T.S. ) Etot(reactants) - Etot(transition state)

(3)

The T.S. of the lowest identified reaction paths are shown in Figure 2B, and the geometry of the transition states and relaxed geometries are shown in Figures 3 and 4. Al12- has the lowest transition state energy of any of the clusters, followed by Al18-, Al17-, and Al16-. All of these clusters have T.S. energies that are at least 0.08 eV below the energy provided by the binding of the water molecule, they are observed to form Aln(H2O)complexes, showing that the ordering of the transition state energies are consistent with the mass spectra. Al14- and Al9have the next lowest transition states, followed by Al15-, Al7-, Al10-, and Al8-, respectively. Finally Al11- and Al13- have activation energies that are more than 0.1 eV greater than what is provided by binding a water molecule. We note that density functional theory generally underestimates barriers, although we also do not include zero point energy effects, and these two errors have opposite signs, so there will be some compensation. This analysis strongly suggests that the observed complexes are due to one or more dissociative chemisorbed water molecules on the surface of the aluminum cluster.

Figure 4. The geometry of Aln- clusters 15 e n e 18 and the LUMO charge density if n is odd, or LUMO+1 charge density if n is even, or Al15-. Initial is the geometry once the intact water molecule is bound at the active site with the HOMO charge density. T.S. is the geometry of the transition state, and Final is the local minimum geometry after the O-H bond has been broken. An isosurface of 0.03 was used.

To understand the size selective reactivity we next need to examine the factors that might facilitate the splitting of the O-H bond. The transition state of Al12- with water, as shown in Figure 3, shows that an Al-O bond is formed on the top site of the cluster, where the LUMO+1 charge density lies, and a Al-H bond begins to form on an adjacent aluminum atom. The HOMO charge density of the cluster-water complex is plotted in the Initial column, and there is significant charge density on the adjacent atom. The LUMO+1 also indicates where a nascent Al-O bond is most likely to form. In all of the 7 e n e 18 clusters, our studies indicate that the transition state is located

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such that the O atom binds to an Al atom that has significant LUMO or LUMO+1 charge density. This shows that the Lewis acid active sites on a cluster are an important part of the reactivity of a cluster with water. A second Al site may also facilitate O-H bond cleavage by having HOMO charge density and preferably accepting the H atom. In Al12-, the HOMO charge density on the Al atom indicates that charge density on the adjacent site facilitates reactivity. A filled frontier orbital’s charge density may then form a nascent Al-H bond, which lowers the energy of the transition state, and further enhances reactivity. This charge density may be thought of as a Lewis Base site, so that the reactivity of aluminum clusters is enhanced by complementary active sites comprised of two adjacent atoms, one which serves as a Lewis Acid, and the other serves as a Lewis Base. The reactivity of Al16-, Al17-, and Al18- with water demonstrates the importance of complementary active sites. As shown in Figure 4, all three clusters have an icosahedral core, with a two atom extrusion on one side, and a one, two, and three atom extrusion on the opposite side of the cluster. The LUMO and LUMO+1 are located on the icosahedra at sites that are adjacent to the two atom defects. These serve as the Lewis acid sites in all three clusters and indicate where the lone pair on water binds to the cluster and the nascent Al-O bond in the transition state. All three clusters also have HOMO charge density on the two atom defect sites, as shown in the Initial column of Figure 4. These function as Lewis base sites and further enhance the reactivity. We note that all four clusters, n ) 12, 16, 17, and 18, which are found to react in the experiment, have a welldefined set of complementary active sites. The two clusters with the largest barrier to reactivity are Al11and Al13-. Note that these clusters survive under high concentration water etching. Al13 is known as a superhalogen33 with a large electron affinity of 3.34 eV, because Al13- has a closed jellium electronic shell that results in a HOMO-LUMO gap of 1.87 eV and resistance to etching with O2.5,7,33 In Al13-, all the surface sites are equivalent, and its outer electronic subshell is filled, while complementary active sites require two inequivalent sites, one which serves as a Lewis acid and one which serves as a Lewis base. We have only plotted (see Figure 3) one LUMO level, while in fact there are 5 nearly degenerate states that combine to form a spherically symmetric shell with uniform charge density around each site. Once water binds, the symmetry is broken, and the HOMO becomes a 2P jellium level that is perturbed in energy by the presence of the lone pair, much like it is observed when the iodides and K3O are bound to the cluster.35 The adjacent site does not show any additional charge density, such that the transition state energy required for water to react with Al13- is the largest of any cluster studied here. Al11- does not have the spherical symmetry of Al13-; however, it is also lacking the required complementary active sites in adjacent locations. The LUMO and LUMO+1 of Al11- are on one of the capping atoms, indicating the best location for Lewis acidity, as shown in Figure 3. However, once water is bound to the LUMO, the adjacent sites do not have significant HOMO charge density. Al11- and Al13- show the largest barriers to reactivity, do not have complementary active sites, and do show the highest resistance to etching by water. The remaining Aln- clusters (n ) 7-10, 14, 15) have reaction barriers that are intermediate in size and exhibit minimal reactivity in low pressure experiments, but do react under high water concentrations. For Al7- and smaller clusters, the anionic charge is distributed over a relatively small number of atoms, reducing the Lewis acidity. Al7- has an octahedral core (see

Reber et al. Figure 3) with an atom capping one face. The LUMO Lewis acid site sits on the opposite side of the core as the defect; and a small amount of HOMO charge density sits on the adjacent atom, allowing for a weak set of complementary active sites and a transition state that is only 0.01 eV below the energy available for reaction. Al8- has an octahedral core with atoms on each side of the core, and has a Lewis acid site adjacent to a defect. It also has HOMO charge density on an adjacent site; however, the energy of the transition state is also only 0.01 eV below the energy available. Al9- and Al10- both have similar structures with active sites on an octahedral feature with threeand four-atom caps. Their transition states also look similar, although Al9- has HOMO charge density on an adjacent site, whereas Al10- does not. As a result, the Al9- transition state is 0.04 eV lower than that of Al10-. Al14- has an icosahedral core with a single-atom defect on the surface. The LUMO+1 site is on the opposite side of the icosahedral core as the defect. There is a HOMO site on an adjacent atom, which indicates the position of the transition state; however, the energy of the transition state is higher than all those that exhibit reactivity. Al15- does not have adjacent HOMO and LUMO sites; however the LUMO+1 site is adjacent to the HOMO, resulting in a weak binding of the initial H2O molecule indicating that the cluster is also a relatively poor Lewis acid. The above analysis shows that the reactivity of aluminum cluster anions with water is consistent with the presence of complementary active sites, or an uneven charge distribution on the surface of the cluster. This allows us to make some generalizations about the experimental results observed at larger sizes. Adatoms either on the surface of a cluster with an complete geometric shell, as extruded atoms above an icosahedral core as seen on Al16-18-, or as defects that can be seen in Al12-, where the cluster is nearly an icosahedron with a vacancy, are the most effective means of generating uneven charge density. So the reactivity of water may be used as a probe whether a cluster has a closed geometric shell. The role of geometry may also be instructive in studying the reactivity of clusters at high temperature.82,83 A second way in which charge can be made nonuniform is on the edges between two surfacelike facets of larger clusters. 3.3. Using Local Charges and Frontier Orbitals to Determine Lewis Acidity. We have used the position of the frontier orbitals, HOMO and LUMO, as proxies to predict the position and the absolute position of the LUMO to predict the strength of the Lewis acidity. Next we investigate the effectiveness of local charges in predicting the position of complementary active sites. To examine this, in Figure 5, A and B, we plot the Hirshfeld charges84,85 and LUMO, LUMO+1, and LUMO+2 charge densities of Al17-. The charge densities show that the four adatoms on the icosahedral core have relatively high charge densities of -0.12 e-, which is consistent with our contention that these sites may serve as Lewis bases. There are two inequivalent Lewis acid sites adjacent to the defects: one which lies next to a single defect atom in the dimer defect with a charge of -0.03 e-, indicated in red, and a second with a charge of -0.02 that lays between the dimer defect sites, in green. The charge densities show that the LUMO in blue is on the first Lewis acid site, the LUMO+1, in red, is on an equatorial site, and LUMO+2, in green, and is on the second Lewis acid site. The first Lewis acid site is favored by the charge densities, and the second Lewis acid site is slightly favored by the Hirshfeld charge. Next we consider the relative transition state energies of the two competing sites. The reaction pathway for the first site is


Figure 5. (A) The geometry of Al17- with the Hirshfeld charges on the cluster labeled. (B) The Al17- cluster with the LUMO charge density plotted in blue, LUMO+1 charge density plotted in orange and red, and the LUMO+2 charge density plotted in green. (C) The binding site, (D) transition state, and (E) dissociated state for splitting water on the complementary active site. (F) The binding site, (G) transition state, and (H) dissociated state for splitting water on a secondary site.

shown in Figure 5C-E, and for the second site, it is shown in Figure 5F-H. On the first site, the binding energy of an intact water molecule is 0.36 eV, and the transition state energy is 0.09 eV below the break-even energy, indicating the feasibility of breaking the O-H bond. On the second active site, the binding energy of the intact water is only 0.22 eV, as shown in Figure 5F. This indicates that the binding is markedly stronger on the site indicated by the frontier orbital analysis, rather than the site indicated by the Hirshfeld charges. Further evidence is seen in Figure 5G where the transition state for splitting water on the second site is 0.42 eV higher than on the first site, as shown in Figure 5D. Some of the increase in energy of the transition state at the second site is likely due to the first site having better access to the charge density of the HOMO. As the charge density is a more physical representation of the local Lewis acidity of the cluster, we found it to be a more useful tool for identifying active sites than local atomic charges. 3.4. Effect of Charge on Reactivity. This work has focused on the reactivity of aluminum anion clusters, which raises the question of what role the extra electron plays in the reactivity of the clusters. To this end, we examined the binding and reactivity of Al13-/0/+ with water in Figure 6. As mentioned earlier, Al13- binds weakly to water with a hydrogen-bond based binding, and has a significant barrier to reacting with water. The icosahedral structure results in all the surface sites being identical, and the closed electronic structure results in a highlying LUMO. In neutral Al13, the excess electron is removed and, although the icosahedral geometry is still evident, the intact water binds by 0.51 eV. The reason for the large increase in binding energy is that the interaction of the lone pair of the oxygen atom to the cluster requires the transfer of charge to the cluster. Al13 is known as a superhalogen33 in part for its affinity for an extra electron and the fact that it readily accepts extra charge whereby the binding energy to water is close to that of Al12-. A charge analysis of the transition state indicates that there is little charge transfer between the cluster and H2O


Figure 6. The pathway for breaking the O-H bond on Al13-, Al13, and Al13+, with binding energies given by eq 1.

such that the energy difference between the anion and neutral at the transition state is less pronounced compared to the initial binding energy. This shows that the neutral Aln clusters accept charge and bind water more strongly than anionic clusters, which is consistent with larger Aln- clusters being generally more reactive than the smaller clusters, and the observed high reactivity on aluminum nanoparticles with water.86–88 Al13+ has a different structure than the anion and neutral clusters, which makes it a probe of both the effect of charge and geometry on its reactivity with water. Water binds very strongly with a 1.10 eV binding energy. This large binding energy has previously been attributed to the shell closure at 40 electrons, and the extra large affinity for charge of Al13+.16 Al13+ also has a respectable HOMO-LUMO gap of 1.19 eV, and charge is easily transferred along the LUMO site sitting on the capping Al atom. The transition state for splitting water on the cation is lower than for the neutral and anion, although the difference is much smaller than the dissociative binding case. This is due to the cation readily accepting charge, while the transition state depends on the paired active sites, which are enhanced by the breaking of the icosahedral symmetry of the cluster. Also, in the cation the dipole moment of the water is aligned with the electric field, while it is opposed in the anion, although the electric field is not large enough to create a barrier that prevents the water molecule from rotating. Hence, the charge has a significant influence on reactivity. 3.5. Mechanisms for H2 Release in Al17-. 3.5.1. SelectiWity in the Langmuir-Hinshelwood Mechanism. The experimental mass spectra show that Al16-, Al17-, and Al18- show products with the addition of multiple water molecules that are deficient in mass by exactly two or four hydrogen atoms from the stoichiometry of water. From this, we conclude that molecular H2 is being formed on the surface of selected hydrogen clusters and released:

Aln- + 2H2O f [Aln(OH)2]- + H2

(4)

We examine several reaction pathways that either compete with or could lead to the formation of H2, with Al17-. The pathways that we find most likely are shown in Figures 7-9.

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Reaction coordinate for the formation of H2 from Al17- and 2H2O (A) Al17-. (B) An intact water molecule bound to the cluster. (C)

Figure 7. molecules. The transition state for splitting the first water. (D) The cluster after water is split, with the LUMO+2 charge density to show the induced Lewis acid site. (E) After the second water is bound. (F) The transition state for splitting the second water. (G) After the breaking of the second O-H bond, (H) A rearrangement of the cluster to facilitate the release of H2. (I) The transition state for releasing H2. (J) The final cluster after H2 is released. The total energies include zero point energies.

Reber et al.

Figure 9. (A) Reaction pathway for splitting the first water molecule with two intact water molecules on Al17-. (B) A reaction pathway for direct H2 release from two water molecules on adjacent sites (includes zero point energies).

Figure 10. A low-barrier proton transfer reaction between an intact water molecule on the surface of the Al17- cluster and hydroxyl bound to the cluster.

Figure 8. (A) Reaction pathway for splitting a second water on the opposite site of [HAl17(OH)]-. The initial column indicates the intact water molecule bound to the cluster, TS is the transition state for breaking the cluster, and Final is the local minimum after the O-H bond is broken. Energies are calculated from eq 3. (B) Reaction path for the breaking of an O-H bond on the adjacent active site. (C) Reaction pathway for breaking a second water on the diagonal active site. (D) The pathway for breaking a second water where the H atom is used as a Lewis base in a Eley-Rideal type mechanism (includes zero point energies).

We have included zero point energies in this pathway, although the primary effect is only to reduce the barriers. It begins with the first water reacting as shown previously in Figure 7A-D.

Once the first water is bound to the Al17- cluster, it induces an active site on the Lewis acid site on the opposite side of the cluster, in accordance with the charge density shown in Figure 7D. A second water then preferentially splits on this new Lewis acid site as shown in Figure 7E-G, resulting in H atoms bound to adjacent Al sites with a significant amount of available energy, from the splitting of the water. We found that a rearrangement of Figure 7G from an icosahedral core to a 14-atom core (Figure 7H) caused a drop in the transition state energy (Figure 7I) for releasing H2. The final product is shown in Figure 7J, and this cluster also contains a second two-atom defect on the surface of the 14-atom core, which can split water and lead to the observed Al17(OH)4- species. Al16- and Al18- have similar potential pathways for H2 release, as shown in Figure S1 and Figure S2 of the Supporting Information. This is a LangmuirHinshelwood mechanism in which both the final product and the main intermediate are observed, lending strong support to this being the most likely mechanism. The mass spectrum, Figure 1C, shows that the product which is concluded to release H2 is much more abundant than the peak for the binding of two waters. To understand the origin of this selectivity in the products, we examine several mechanisms for the splitting of the second water molecule in Figure 10. This step is the key to understanding what controls whether or not H2 is released, because if the H atoms are near each other, then the H2 formation and release is likely to be rapid, whereas if they are separated, then the cluster is likely to cool through collisions with helium atoms in the flow reactor before the H atoms may find each other and be released as H2. This is complicated by the fact that Al17- has four equivalent complementary active sites. Figure 8, row A, shows the Langmuir-



Hinshelwood type second step, where the water splits on the opposite site; the initial binding energies, transition state energies, and final energies are given. Row B shows what occurs if the second water attempts to split on the site adjacent to where the initial water split, and row C shows the geometries and energies for splitting water on the diagonal site on the Al17cluster. The energies are calculated using eq 5, so that pathways with slightly different initial states may be easily compared.

BE ) E(Al17-) + 2E(H2O) - E[Al17(H2O)2]-

(5)

First, we note that the binding energy at the adjacent site (row B) is relatively high compared to the other binding energies. This is due to the hydrogen bond between the water molecule and the lone pair on the hydroxyl group, an arrangement that has some interesting consequences that we discuss later. The binding energy on the diagonal site is 0.31 eV less than on the opposite site, which is evidence of the importance of the induced Lewis acid active site on the opposite position, as shown in Figure 7D. The transition state energies show that the opposite site has a much lower transition state energy for splitting the second water than the other two complementary active sites. The transition state energy for the opposite site is 1.71 eV, whereas it is 0.18 eV higher for the adjacent site, and 0.25 eV higher for the diagonal active sites. This suggests that the induced active site is why the abundance of Al17(OH)2- is much larger than Al17(H2O)2-. Figure 8D shows an Eley-Rideal mechanism89 for the release of H2 with the addition of the second water. In this case, the complementary active sites are the opposite Lewis acid site and the H atom as a Lewis base site. This mechanism results in the nearly direct release of H2; however, the transition state is much higher than pathway A. If a detached H atom could serve as an effective Lewis base it would be expected that nearly all clusters which display appreciable reactivity to water would also show H2 release. However, only Al16-18- and a few much larger clusters show experimental evidence of H2 release. Because the Langmuir-Hinshelwood type pathway shown in Figure 8A has a lower transition state for the splitting of the second water, this pathway is most likely to explain the H2 release. 3.5.2. Direct ReactiWity. A second potential pathway for the production of H2 is a direct process in which two intact water molecules are bound to adjacent Al atoms on the surface of the cluster. The concerted splitting of one or both O-H bonds may then result in the formation of two hydroxyls on the surface of the cluster and a released H2 molecule. In this case, the Al sites act as the Lewis acid sites, and the hydrogen in the water molecule acts as the Lewis base. For comparison, in Figure 9A we plot the splitting of the O-H bond where a second intact water molecule is bound to the opposite site. Figure 9B shows the direct pathway. Although the binding of the intact water molecule is respectable, only 0.13 eV below the energy at the opposite site, the transition state energy for splitting from the adjacent site is 0.8 eV higher in energy. This is due to the oxygen atom withdrawing charge from the hydrogen atom, making it a poor Lewis base. Although this barrier is quite high, the process seems logical at first glance as the adjacent Lewis acid sites suggest such a concerted process to be possible. The direct process and the Eley-Rideal type mechanism shown in Figure 8D also raise the possibility that the H atom’s ability to tunnel may enhance these direct H-H bond formation mechanisms. We do not believe that such processes are particularly likely, as generally H atom tunneling does not give

rise to the formation but rather involves a jump from a stable configuration to a neighboring lone pair. In any case, we point out that the role of tunneling in the release of H2 has not been investigated here and may be an interesting subtlety to these processes. We note that we do not need it to explain the primary results of the experiments. 3.5.3. Hydrogen Transfer Processes. A secondary reaction in which a hydrogen atom is transferred from an intact water molecule to a hydroxyl bound to the cluster is found to have the lowest barrier for any reaction that we investigated. The stoichiometry of the reactant and product are the same, so it is impossible to confirm that this reaction is occurring in the present experiment. As shown in Figure 10, the reaction begins with a water molecule hydrogen bonded to a hydroxyl group on an adjacent Al atom. The barrier for the H atom to jump to the hydroxyl is only 0.02 eV, resulting in the hydroxyl being formed on a different site. This process is similar to the Grotthuss mechanism, which has been proposed to play a role in accelerating the hydrogen production in a recent ab initio molecular dynamics study.90 This is likely to contribute to the formation of Al17(H2O)2-, as it moves the hydroxyl and decreases the transition state energy at the diagonal site, so the dissociated H atoms are on opposite sides of the cluster. It also suggests that the sequential reaction of water and then glycol may potentially lead to the tethering of clusters. 4. Conclusions The present work shows how complementary active sites facilitate the cleavage of the O-H bond and how paired sets of active sites result in the preferential release of H2 after binding multiple waters. We show that the results of a flow tube experiment in which the Aln- clusters are exposed to water, results in size selected reactivity. The reactivity is controlled by complementary active sites that are made of a Lewis acid site that forms a nascent Al-O bond with water, and a Lewis base site that accepts the hydrogen atom. The best Lewis acid sites on a cluster are identified using the LUMO charge density for even electron systems and LUMO+1 for odd electron systems, allowing a determination of the position where a lone pair may bind. We also show that the excess charge of an anion reduces the reactivity of the cluster. Further, we note that uneven charge density which leads to complementary active sites is generally a feature of clusters that have defect-like structures, and that relatively unreactive species indicate the presence of closed geometric shells, and high reactivity indicates defects. We have also studied several possible mechanisms for the release of H2 after the binding of multiple water molecules. We found that a Langmuir-Hinshelwood type mechanism in which the successive binding of water followed by the formation of H2 from two H atoms bound directly to the cluster at adjacent sites is most likely. The binding of the first water molecule induces a Lewis acid site on the opposite site, which makes this reactive selective. We also examined Eley-Rideal and direct mechanisms but found that the H atom is a relatively poor Lewis base. We also found a mechanism for hydrogen exchange between a surface-bound hydroxyl and a chemisorbed water molecule that may open occupied active sites and may be used to bind other molecules to the cluster. This work demonstrates the role of geometry on the reactivity observed at the nanoscale, in which paired complementary active sites result in the preferential production of H2 from water. The role of complementary active sites in reactions with other small molecules remains to be explored.

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Reactivity of Aluminum Cluster Anions with Water: Origins of Reactivity

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