Lewis Acid Site and Hydrogen-Bond-Mediated Polarization Synergy in

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Lewis Acid Site and Hydrogen-Bond-Mediated Polarization Synergy in the Catalysis of Diels-Alder Cycloaddition by Band-Gap Transition Metal Oxides Taha Salavati-fard, Efterpi S. Vasiliadou, Glen Richard Jenness, Raul F. Lobo, Stavros Caratzoulas, and Douglas J Doren ACS Catal., Just Accepted Manuscript • Publication Date (Web): 10 Dec 2018 Downloaded from http://pubs.acs.org on December 10, 2018

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Lewis Acid Site and Hydrogen-Bond-Mediated Polarization Synergy in the Catalysis of Diels-Alder Cycloaddition by Band-Gap Transition Metal Oxides Taha Salavati-fard1,2, Efterpi S. Vasiliadou2, Glen R. Jenness2, Raul F. Lobo2,3, Stavros Caratzoulas2* and Douglas J. Doren2,4 1

Department of Physics and Astronomy, University of Delaware, Newark, Delaware 19716, United States

2

Catalysis Center for Energy Innovation (CCEI), University of Delaware, Newark, Delaware 19716, United

States 3

Department of Chemical and Biomolecular Engineering, University of Delaware, Newark, Delaware 19716,

United States 4

Department of Chemistry and Biochemistry, University of Delaware, Newark, Delaware 19716, United

States

ABSTRACT We present evidence from kinetic studies and electronic structure calculations that monoclinic ZrO2 and HfO2 can catalyze Diels-Alder [4+2] cycloaddition between furan and methyl acrylate. The two oxides present the same apparent activation energies, in the range of 11.5 kcal/mol, but HfO2 seems intrinsically more active. We use DFT calculations to investigate catalytic pathways and the influence of surface hydration on activity. Partially hydroxylated surfaces are more reactive than the dehydrated or completely hydroxylated surfaces on account of a synergy between Lewis acidic metal centers and surface hydroxyl groups that lowers the LUMO of surface bound methyl acrylate. We argue that surface hydroxyl groups present a polar environment which, via hydrogen bonding, are solely responsible for the rate acceleration on completely hydroxylated surfaces without necessarily proton transfer. The Lewis metal centers of completely dehydrated surfaces are marginally effective active sites. Upon comparison with γ-Al2O3, which has been reported to catalyze the Diels-Alder of cyclopentadiene with methyl acrylate, we argue that without compromising catalytic efficiency, ZrO2 and HfO2 present the advantage of selective catalysis without need for surface activation or a controlled environment. Keywords: Diels-Alder cycloaddition; furan; methyl acrylate; zirconium oxide; hafnium oxide; density-functional theory; Lewis-Brønsted synergy.

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INTRODUCTION We report the first experimental and computational evidence that partially or fully hydroxylated transition metal oxides ZrO2 and HfO2 can catalyze Diels-Alder [4+2] cycloaddition between furan and an activated dienophile, methyl acrylate. Diels-Alder cycloaddition is a powerful tool in organic synthesis and in the case of furans it yields oxanorbornene derivatives that can be converted to a variety of molecules, ranging from molecules of biological interest to naturally occurring organic compounds. Lately, it has also been propounded as a promising alternative for the synthesis of aromatics by dehydration of appropriate oxanorbornene intermediates obtained from renewables.1-14 The latter strategy has been showcased with the synthesis of p-xylene by dehydration of the Diels-Alder product of dimethyl furan and ethylene over solid Brønsted acids (HY; H-BEA; phosphorous-modified, all-silica BEA) and solid Lewis acids (homomorphously substituted Sn-BEA and Zr-BEA).7-9, 11-12, 15-16 None of these catalysts, however, accelerated the Diels-Alder step of the tandem scheme (vide infra).11, 16-17 Furans are poor nucleophiles for Diels-Alder reactions due to their aromatic character. For example, Mahmoud et al.18 have reported no yield when furan was reacted with methyl acrylate, an activated dienophile, for 24 h at 298 K in the absence of catalyst. Thus, the development of efficient catalytic heterogeneous systems with improved yields and selectivity is highly desirable for the establishment of environmentally friendly technologies. Although Brønsted and Lewis acids are generally able to catalyze various Diels-Alder reactions19 via mechanisms readily explained in terms of frontier molecular orbital theory,20-21 the polar character of the transition state and the concepts of electronic hardness and electrophilicity,2227

recent electronic structure calculations have shown that they are quite ineffective with furans as

dienes7, 11, 15-16—unless the dienophile is already activated by polar, electron withdrawing groups that bind more strongly to the active site, in which case the catalyst promotes normal electron demand [4+2] cycloaddition.28-29 (In passing, when the furan is more polar than the dienophile and thus binds more strongly to the active site, one needs very strong Lewis acids to promote inverse electron demand [4+2] cycloaddition, even when the furan is activated with electron withdrawing groups.16 To the best of our knowledge, such a catalytic system has not been reported in the DielsAlder literature of furans.) Mahmoud et al.18 recently reported Lewis acid catalysis of Diels-Alder cycloaddition of methyl acrylate and furan (Scheme 1) over the solid Lewis acids Sn-, Zr-, and Hf-BEA at 298 K. Quantum


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chemical calculations29 showed that all three catalysts preferentially bind the dienophile, instead of the furan, enhancing its electrophilic character. Despite their different atomic electronic properties, Sn, Zr and Hf acquire very similar electrophilic characters when substituted into the BEA framework and, with similar electrophilicity indices in the range of 1.7 eV, Sn-, Zr-, and HfBEA behave as moderate Lewis acids, promoting modest charge transfer from the diene to the dienophile at the transition state (normal electron demand). Notwithstanding their mildly Lewisacidic power, they still achieve a reduction of about 12.5 kcal/mol in the cycloaddition apparent activation energy relative to the uncatalyzed reaction. This is a significant step toward efficient catalytic systems that overcome a number of undesirable or inconvenient features of traditional homogeneous Lewis acid catalysis, such as sensitivity to water or strong binding between the Lewis acid and the electron withdrawing group of the product, which can slow down the exchange. Because, however, zeolitic Lewis acids with homomorphously substituted metal centers come with certain limitations (complex hydrothermal synthesis, low amounts of metal incorporation and use of highly corrosive and poisonous hydrofluoric acid30), it is of significant import that naturallyoccurring, transition metal oxides such as ZrO2 and HfO2, which are easily synthesized and are thermally and chemically stable, can catalyze the Diels-Alder cycloaddition of furans with activated olefins. Upon comparison with γ-Al2O3, which has been reported to catalyze the DielsAlder of cyclopentadiene with methyl acrylate, dimethyl maleate and dimethyl fumarate31, we argue that without compromising catalytic efficiency, ZrO2 and HfO2 present the advantage of selective catalysis without need for surface activation or a controlled environment. METHODS Models and computational methodology. ZrO2 and HfO2 have, for the most part, similar bulk and surface properties.32 For both of them, we have modelled the thermodynamically most stable and catalytically most active (111) facet of the monoclinic phase.32-40 We have modelled both dehydrated and hydroxylated surfaces as both oxides adsorb water readily;35, 38, 41-43 water either physisorbs or adsorbs dissociatively, depending on coverage.35, 38, 44 We have run mechanistic studies for both partially and fully hydroxylated surfaces. Periodic DFT calculations were performed with 22 supercells with an 18Å vacuum in the direction normal to the surface. The partially hydroxylated surfaces were modelled by including four water models and the fully hydroxylated ones by including twelve water molecules, which amount to coverages


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of 2.2 and 6.7 water molecules/nm2. Figure 1 illustrates the dehydrated (1a), partially hydroxylated (1b; lowest structural energy among different configurations tested) and fully hydroxylated (1c) surfaces. Because the two oxides have wide band gaps, we employed the PBE+U+D3 theory level, as implemented in the Vienna ab initio simulation package (VASP, version 5.4.1)45-47 The core electrons were represented by the projector augmented wave function (PAW) method,48-49 while plane waves with a 400 eV energy cutoff were used for the valance electrons. A Hubbard U=3.0 eV was considered for Hf and Zr atoms50 to provide a localizing, on-site Coulomb interaction for electrons in the d-orbitals.51-52 The dispersion correction D3, introduced by Grimme, was also employed.53-55 The reciprocal space was sampled using a 3 × 3 × 1 grid, which provides 5 irreducible k-points in the Brillouin zone, based on the Monkhorst-Pack method.56 A Gaussiansmearing scheme with 𝜎 = 0.10 𝑒𝑉 was used to treat discontinuities at the Fermi level. Optimizations were performed with the conjugate-gradient algorithm, with a force threshold of 0.05 𝑒𝑉/Å. Transition states were located using the nudged elastic band (NEB) method57-58 with 12 images, followed by dimer calculations.59-60 The force thresholds for the NEB and dimer calculations were set to 0.2 and 0.05 𝑒𝑉/Å, respectively. In the case of alumina, we modelled the (110) facet of the non-spinel bulk structure. The fully hydrated 22 surface was built with 16 dissociated water molecules. The PBE+D3 method, using the same parameters as above, was employed to calculate reaction pathways on completely dehydrated and completely hydroxylated alumina. Thermal corrections to the electronic energy at 𝑇 = 293𝐾 were estimated within the quasirigid rotor harmonic oscillator approximation (qRRHO) of Grimme61 and Head-Gordon62 which addresses the low-frequency problem in the entropy and enthalpy of the harmonic oscillator. The vibrational frequencies of the reactants, intermediates and transition states were calculated with a tight energy threshold, 10 -8 𝑒𝑉. Bader analysis63 was performed using a VASP add-on.64 Density of states (dos) analysis was performed using GPAW65-66 after recalculation of the wave functions on a 6 × 6 × 1 k-grid, which provides 18 irreducible points in the Brillouin zone. A Gaussian broadening of 0.10 𝑒𝑉 was used for better illustration. The energy reference for the calculation of the dos and Fermi levels was taken to be the vacuum energy of each system. We should note that one needs dipole corrected electrostatic energy without exchange contribution, at different heights from the surface, to find


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the energy plateau that corresponds to the vacuum energy (work function) from the zero (Fermi) level. For instance, Figures S1 and S2 in the Supplementary Information show work functions for clean, dehydrated HfO2 and dehydrated HfO2 with surface bound methyl acrylate, respectively. The molecular orbitals of the isolated furan and methyl acrylate molecules were calculated at the PBE/aug-CC-pVTZ theory level in Gaussian G09.67 Materials and experimental methods. Commercially available HfO2 (Sigma-Aldrich, powder 98%), ZrO2 (Sigma-Aldrich, nanopowder, < 100𝑛𝑚 particle size TEM) and Al2O3 (AlfaAesar, powder) were used in this study. Nitrogen physisorption was performed in a Micromeritics ASAP 2020 system at 77𝐾 to determine the surface area of the oxides. Because the surface area was expected to be low, larger than typical amounts of sample was used for the analysis (1 ― 2 𝑔𝑟). Samples were degassed overnight at 523𝐾 and backfilled with dry nitrogen prior to analysis. The surface area of γ-Al2O3 is provided by the vendor and ranges between 80-120 m2/gr. X-ray diffraction (XRD) patterns were collected using a Bruker D8 diffractometer with Cu Kα radiation. The diffraction pattern was collected for 2 seconds at each increment of 0.02 degrees between 5 and up to 90 degrees 2θ (see Figure 2). Scanning electron microscopy images were recorded on a JEOL JSM 7400F at 10 μA. X‐ray fluorescence (XRF) using a Rigaku WDXRF was used for elemental analysis. X‐ray photoelectron spectroscopy (XPS) experiments were conducted by using a Thermo‐Fisher K‐alpha+X‐ray photoelectron spectrometer equipped with a monochromatic AlKα X‐ray source (400 μm). Furan (Sigma-Aldrich, > 99%) and methyl acrylate (Sigma-Aldrich, 99%, contains ≤ 100 ppm monomethyl ether hydroquinone as an inhibitor) were stored under an inert atmosphere and used as received. Diels−Alder reactions were performed in 12 mL Q-tube pressure glass tube reactors with maximum pressure limit of 180 psi (Q LabTech). The reactors were heated with a Q-block (Q LabTech) and stirred magnetically. Reactions were conducted under solvent-free conditions at 293, 303 and 313𝐾 for 24 ℎ; 14.1𝑚𝑚𝑜𝑙 of furan, 2.8𝑚𝑚𝑜𝑙 of methyl acrylate, and 0.50 𝑔 of catalyst were loaded into the reactor. For the experiments with γ-Al2O3, the amount of the reactants was three times higher, keeping reactant concentrations constant, to ensure efficient stirring. For all reactions, 0.163𝑚𝑚𝑜𝑙 of p-xylene (Sigma, ≥ 99%) was used as an internal standard, which was added before addition of the catalyst. Quantification was performed with 1H NMR using the internal standard and the peaks corresponding to the endo and exo isomers of the cycloadduct turned out to be at 3.76 and 3.67𝑝𝑝𝑚, respectively. NMR spectra were recorded at


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298 𝐾 with a Bruker AVIII400 NMR spectrometer. Details about the NMR analysis can be found elsewhere.68 RESULTS Catalyst characterization. X-ray diffraction patterns of HfO2, ZrO2 and Al2O3 are shown in Figure 2. The hafnium and zirconium oxides are monoclinic and highly crystalline. X-ray diffraction of Al2O3 verified the existence of the gamma phase;69 no other phases were observed. The surface area and the bulk chemical composition of the catalysts are presented in Table 1. There are significant differences in the specific surface area of the three oxides with HfO2 possessing the lowest and γ-Al2O3 the highest surface area. Because of the different surface areas and in order to compare the catalytic data on equal footing, we are expressing catalyst activity as area rate and turnover frequency (see discussion below). Variations in catalyst morphology are clearly distinguished with SEM (Figure 3 (a), (b) and (c)). SEM of HfO2 reveals well-formed polyhedronlike particles with sizes in the 50-200 nm range. On the other hand, ZrO2 consists of spherical agglomerates of particles with sizes between 20 and 50 nm. Finally, SEM images of γ-Al2O3 do not indicate well-defined particle shapes; particle agglomerates are clearly present. In addition, the texture of γ-Al2O3 is significantly more porous than HfO2 and ZrO2. These observations align with the specific surface area properties of the oxides. The chemical composition of HfO2 and ZrO2 indicates the presence of minor impurities accounting for about 3%wt in both oxides; Ti is the most abundant one in HfO2 and Hf in ZrO2. In contrast, γ-Al2O3 is highly pure, containing P at very low concentration. Surface compositions were studied by XPS. The XPS spectra of the a) Hf 4f and O 1s; b) Zr 3d and O 1s; and c) Al 2p and O 1s emissions are shown in Figure 4. For HfO2, the set at 17.1 eV and 18.8 eV corresponds to Hf 4f7/2 and Hf 4f5/2, respectively. This is in agreement with values reported in literature for Hf(IV) oxide.70 The O 1s peak can be deconvoluted into two contributing bands. The one located at binding energy 530.4 eV is assigned to lattice oxygen in HfO2 while the second, located at 532.2 eV, could be attributed to OH groups that arise from exposure of the sample to atmospheric moisture.71 The Zr 3d5/2 line is at 182.2 eV and the Zr 3d3/2 at 184.6 eV. This is in good agreement with that of bulk ZrO2.72 The O 1s spectrum can be separated into two features corresponding to two chemical states, one located at 530.2 eV and the other at 531.9 eV. Similarly to the finings for HfO2, these two peaks correspond to Zr-O bonds from bulk ZrO2 and surface Zr-OH bonds, respectively.72 Al 2p at 74.9 eV corresponds to Al in γ-


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Al2O3. From the O 1s spectra, we observe a broad peak is observed that involves two separate peaks at 531.3eV and 5324eV; these are attributed to Al-O and Al-OH bonds.71 The XPS results, although qualitative, clearly indicate that the surface of the catalyst prior to reaction is hydroxylated. Catalysis and kinetics of Diels-Alder cycloaddition of furan and methyl acrylate. As mentioned earlier, without catalyst, no reaction is observed over 24 h at 298K. Thermodynamic calculations in earlier work18 showed that the reaction is kinetically limited at temperatures below 300 K and equilibrium-limited at temperatures above 330 K. Operation at low temperatures (≤ 330 K) requires a highly active catalyst. We ran the reaction of furan and methyl acrylate in the presence of HfO2 and ZrO2 without solvent (to avoid unnecessary separation steps and environmental impact). We investigated HfO2 and ZrO2 as catalysts at 293, 303 and 313 K for 24 h. We found that both oxides catalyze the reaction and are highly selective, with the oxanorbornene adduct being the only product. Oxanorbornene carboxylic methyl ester yields and area specific rates are shown in Table 2. The area specific rates show that HfO2 is more active than ZrO2, because the surface area of ZrO2 is six times as large as that of HfO2. For a more accurate comparison, we also calculated turnover frequencies (TOF, h-1), by taking into account the different number of accessible active sites (8.89 Hf or Zr atoms per nm2). The HfO2 TOF is at least 7.5 times as high as that of ZrO2 at 313K (Figure 5). Although the Hf-BEA and Zr-BEA zeolites18 are intrinsically more active than HfO2 and ZrO2, the HfO2 TOF at 298K is twice as high as the TOF over silica-supported titanium chloride (TiCl4/SiO2 TOF=0.025 h-1).73 The latter catalytic system, apart from Hf-BEA and Zr-BEA, is the only published example of catalysis of Diels−Alder cycloaddition of furan and methyl acrylate. To further assess the catalytic efficiency of ZrO2 and HfO2 relative to an oxide that has strong and weak Lewis acid sites as well as Brønsted acid sites, we also tested γ-Al2O3, which has been reported to catalyze the Diels-Alder of cyclopentadiene with methyl acrylate.31 In our hands, nopretreated alumina exhibited area rates and turnover frequencies (calculated using 5.90 Al atoms per nm2) similar to ZrO2, but it was less active than HfO2 (Table 2). Using excess furan and assuming pseudo-first order kinetics in MA, we calculated apparent activation energies of Ea=11.3 kcal/mol on ZrO2 and Ea=11.7 kcal/mol on HfO2; Arrhenius plots are shown in Figure 6. The apparent rate constants were deduced from the initial methyl acrylate concentration and the initial rate of its disappearance. In contrast to the similar apparent activation


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energies, the pre-exponential frequency factors are quite different for the two oxides: A=151.3 s-1 .

gcat-1 for HfO2 and A=65.5 s-1 . gcat-1 for ZrO2. The difference in intrinsic activity between HfO2

and ZrO2 must therefore be attributed to entropic effects and to the number of sites on the catalytic surface. Similar observations have recently been reported by Gonell et al.74 for the MeerweinPonndorf-Verley (MPV) reduction of cyclohexanone with 2-propanol over different ZrO2 polymorphs and Zr-BEA. For γ-Al2O3, we found Ea=10.6 kcal/mol and A=73.9 s-1 . gcat-1, suggesting that ZrO2 and HfO2 are on par with Al2O3. However, ZrO2 and HfO2 seem to be more selective. Over alumina, we observe—even at room temperature— a change in color from clear to light brown upon addition of furan into the reaction mixture, strongly suggesting side reactions which most likely involve furan oligomerization. We suspect that activated alumina will be even more reactive toward furan oligomerization. Methyl acrylate binding geometry and surface interactions. Methyl acrylate binds more favorably than furan by at least 8 kcal/mol; On dehydrated ZrO2, the binding energy is 26.4 kcal/mol. In the optimized geometry (Figures 7(a) and 7(b)), the carbonyl oxygen coordinates to a surface metal atom 2.35 Å apart, and the plane of the vinyl group (plane of the C-C π bond) is parallel (perpendicular) to the plane of the surface and interacting with two metal atoms. We have found almost identical binding geometries but slightly stronger binding (28.5 kcal/mol) on HfO2. Binding is accompanied by very modest charge shift from MA to the surface. According to Bader analysis, MA acquires partial charge of +0.04e on both oxide surfaces (Table 3). Binding of MA to the partially hydroxylated surfaces is weaker, 22.8 and 22.1 kcal/mol for ZrO2 and HfO2, respectively. On both oxides, however, the carbonyl group coordinates to both a metal atom and a neighboring hydroxyl group via a hydrogen bond (Figures 7(c) and 7(d)). On ZrO2, the metal–carbonyl oxygen and surface OH–carbonyl oxygen bond lengths are 2.41 and 2.08 Å, respectively (2.35 and 2.15 Å for HfO2). We will see later that this bridging is critical to the reactivity of the partially hydroxylated surfaces. The plane of the vinyl π-bond is now tilted and more parallel to the surface, which should be partially responsible for the weaker binding, on account of weaker electronic interactions with the surface metal atoms. Despite the weaker binding, we find the same modest charge shift from MA to the surface, 0.03 and 0.04𝑒 for partially hydroxylated ZrO2 and HfO2, respectively. This should not be all that surprising because charge transfer to both the dehydrated and partially hydroxylated surfaces mainly takes place through the


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coordinated carbonyl group; the HOMO of MA is a non-bonding p-orbital of the carbonyl oxygen atom, acting as an electron donor in the H-bond. Unsurprisingly, on the fully hydroxylated surfaces (Figures 7(e) and 7(f)), there is minimal interaction between MA and the metal atoms and even less charge shift to the surface; the Bader partial charge on MA is only +0.01e (Table 3). Instead, the MA carbonyl coordinates to three surface OH groups, with typical H-bond lengths in the range of 2.1 Å for both oxides. On ZrO2, the MA binding energy is 25.9 kcal/mol, which is somewhat weaker than the dehydrated surface but stronger than the partially hydroxylated one. Binding to HfO2 is weaker by about 4 kcal/mol. Catalytic pathways. Surface-bound methyl acrylate can react either with surface-bound furan or with furan coming in from the gas phase. We will be using the shorthand “Langmuir-type” for the first mode and “Eley-Rideal-type” for the second mode, but with the understanding that only dynamical studies could distinguish between the two, as strictly speaking the former entails equilibration on the surface. For both ZrO2 and HfO2 and irrespective of the degree of hydroxylation, the activation energies for the Langmuir-type regime are higher. The energies of intermediates and transition states in both regimes are presented in Tables 4 and 5. On ZrO2, the Eley-Rideal-type regime is favored by 2.8, 8.7 and 7.9 kcal/mol for dehydrated, partially and fully hydroxylated surfaces; on HfO2, the corresponding figures are 3.3, 8.0 and 6.4 kcal/mol. In Figure 8, we show the geometries of transition state complexes and intermediates on the dehydrated surfaces. Although it is energetically more favorable to have both addends surface bound (Tables 4 and 5), visual inspection of the Diels-Alder transition state reveals that the adsorbed furan molecule must come off the surface before it reacts with adsorbed MA. Because of this extra energy cost, the Langmuir-type regime is less favorable. Similar transition state geometries are observed for the partially and completely hydroxylated surfaces, shown in Figures S3 and S4 of the SI. For the more favorable regime (Eley-Rideal-type), we have performed vibrational frequency analysis and computed thermal corrections to the electronic energy within the quasi-rigid rotor harmonic oscillator approximation.61-62 The enthalpy profiles are presented in Figures 9 and 10 for ZrO2 and HfO2, respectively; the enthalpy values used for the plots are also provided in Table S1 of the SI. The dehydrated surfaces seem to be weakly catalytic. On both ZrO2 and HfO2, the calculated activation enthalpies, in the range of 17 kcal/mol, are marginally lower than that of the uncatalyzed reaction in the gas phase (18.3 kcal/mol). The partially hydroxylated surfaces are


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substantially more active, as the barriers drop to 12.1 and 12.8 kcal/mol on ZrO2 and HfO2, respectively. Complete hydroxylation reduces their activity somewhat, as the respective enthalpy barriers are slightly higher, 13.0 and 14.0 kcal/mol. Within the error of the DFT calculations, one should not make much of these small differences and practically conclude that the partially and completely hydroxylated surfaces are equally reactive. Overall, the calculations are consistent with the experimental observation of very similar apparent activation energies on ZrO2 and HfO2. In addition, the computed enthalpy barriers for the hydrous oxides are very close to the experimental values. The (110) surface of activated (i.e., completely dehydrated) γ-alumina has strong and weak Lewis acid sites. With MA bound to the tri-coordinated Al atom (the strong Lewis acid site), the computed activation energy is 10.2 kcal/mol, that is, ca. 7 kcal/mol lower than on unhydrous ZrO2 and HfO2. This figure climbs up to 15.2 kcal/mol when MA is at a weaker Lewis acid site, a tetrahedrally coordinated Al atom. On the completely hydroxylated surface, the activation energy is 14.8 kcal/mol, which is higher than on ZrO2 and HfO2 (Table 6). Despite the small differences in activation energy, notable is the observation that the hydroxylated (110) surface of alumina seems slightly more active than the weak Lewis acid sites of the activated surface. In this respect, we discern a similarity between alumina and the other two oxides. It seems that hydrous ZrO2 and HfO2 are, overall, no less active than alumina while they present the additional advantage of selective catalysis without need for surface activation or a controlled environment. DISCUSSION How can we understand the catalytic activity of these wide-band-gap oxides and of the hydroxylated surfaces in particular? On account of the activated nature of the dienophile, we expect normal electron demand cycloaddition and this is indeed confirmed by Bader analysis and the charge shift from furan to methyl acrylate in the transition state (Table 3). In addition, our analysis reveals a correlation between the activation energy and the net electron density transfer to the dienophile—the greater the excess charge, the lower the barrier. On the dehydrated surfaces, 0.19e of electronic charge is transferred from furan to MA, which acquires a partial charge of 0.12e. This means that in the transition state the overall charge transfer to MA is 0.16e, since the surface bound MA has partial charge of +0.04e (Table 1). On the partially hydroxylated surfaces,


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these figures are higher; partial charges of +0.24e and -0.18e on furan and MA, respectively, imply that 0.22e of charge transfer from furan to MA. On the fully hydroxylated oxides, the respective partial charges are somewhat smaller, +0.19e and -0.16e, and in this case 0.17e of net charge transfer to MA (recall that the partial charge of surface bound MA is +0.01e). These figures are not dramatically different from the 0.15e of electron charge shift from furan to MA in the gas-phase transition state. Nevertheless, they suggest lowering of the LUMO of the dienophile on all three surface types. Indeed, the LUMO energy of an isolated, fully relaxed methyl acrylate molecule is -2.3 eV. Remarkably, it coincides with the center of the narrow d-band that dominates the conduction band of the oxides, shown in Figure 11(a) for clean (no surface MA) dehydrated ZrO2. Upon binding, mixing of the narrow LUMO band of MA with the narrow d-band of the oxide lowers the former and raises the latter, clearly evident in Figure 11(b), where we show the density of states (dos) of dehydrated ZrO2 with surface bound MA. The high peak in the conduction band represents the d-band of the metal atoms while the resonance at -3 eV in the band gap represents the LUMO band of adsorbed MA. These changes in the dos are also more clearly delineated in Figures 11(c)-(e). In Figure 11(c), we show contributions to the dos of clean dehydrated ZrO2 from the d-band of surface Zr atoms and from the p-band of surface O atoms that are in the immediate vicinity of MA when MA is bound to the surface (these atoms are marked in Figure S6 in the SI). In Figure 11(d), we show the same projected dos (pdos) as in Figure 11(c) but for dehydrated ZrO2 with surface-bound MA. In Figure 11(e), we show contributions from the surface-bound MA molecular orbitals, with the methyl acrylate LUMO band at -3 eV, that is, 0.7 eV lower than the gas-phase orbital. In the pdos shown in Figure 11(e), we see also significant broadening of the HOMO-1, HOMO-2 and HOMO-5 bands of MA due to interactions with the oxide metal and oxygen atoms. We shall return to these bands as, on account of their local πsymmetry with respect to the C=C plane, they participate in secondary C-C bond-forming orbital interactions. The corresponding dos and pdos for hydroxylated ZrO2 are provided in Figures 12(a)-(e); the surface metal and oxygen atoms used in the calculation of the pdos shown in Figures 12(c)-(d) are marked in Figure S7 of the SI. The LUMO band of MA is located even lower, at -3.4 eV, consistent with the higher excess charge on MA in the transition state and the significantly lower activation energy. The additional 0.4 eV stabilization of the MA LUMO on the partially hydroxylated surface needs further consideration. On both dehydrated and partially hydroxylated ZrO2, the MA carbonyl


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group is coordinated to a metal atom, with almost identical carbonyl oxygen-metal atom bond lengths that suggest similar electronic coupling strengths. On the hydroxylated surface, however, the carbonyl group is also coordinated to a surface OH. This hydrogen bond provides extra stabilization to the LUMO of the absorbate. This type of activation via hydrogen bonding becomes more evident on the fully hydroxylated surfaces, where the carbonyl group is coordinated to three OH groups. In Figures 13(a)-(e) we present the corresponding total dos and pdos; the surface metal and oxygen atoms used in the calculation of the pdos shown in Figures 13(c)-(d) are marked in Figure S8 of the SI. The lack of strong interactions between MA and the metal atoms is quite evident in Figures 13(c)-(d), where we see that the presence of MA on the surface does not influence the location of the conduction dband, which remains centered at around -2.3 eV. In contrast, the MA LUMO shifts lower and appears at -2.75 eV (Figure 13(e)). Thus, similarly to MA bound to the partially hydroxylated oxides, we see lowering of its LUMO by as much as 0.4 eV on account of hydrogen bonding. The lack of interaction with the surface metal atoms is also evident if one compares the structures of the HOMO-1, HOMO-2 and HOMO-5 bands (all of π-symmetry) across the three surface types (cf. Figures 11(e), 12(e) and 13(e)). All three bands are shifted in energy and dispersed due to coupling with the oxide bands below the Fermi level, but more so on the dehydrated and partially hydroxylated surfaces than on the fully hydroxylated surface. These arguments alone cannot explain why the completely hydroxylated surface is more active (lower computed barrier) than the dehydrated surface given that the MA LUMO is more stabilized on the latter. To understand that, we need to consider secondary bonding interactions (in addition to the leading (LUMOdienophile – HOMOdiene)-1 bonding term) and specifically the MA bands below the Fermi level that can mix with the π1*+π2* orbital (LUMO) of furan, where π1 and π2 denote its two π-bonds. These MA orbitals, shown in Figure S5 of the SI, are the HOMO-1 and HOMO-2 and to a lesser extent the low-lying HOMO-5 (the HOMO of MA is the carbonyl oxygen lone pair and thus irrelevant to Diels-Alder cycloaddition); all three orbitals are some linear combination of the C-C π-bond with oxygen lone pairs or with the carbonyl π-bond. The higher these bands lie, the more the secondary orbital interactions contribute to the lowering of the activation energy. In Table 7, we have compiled the centers of the HOMO-1, HOMO-2 and HOMO-5 bands for all three surface types of ZrO2 and HfO2. It is quite evident that, for both oxides, these bands lie higher in the fully hydroxylated surface than in the dehydrated one. In ZrO2, the HOMO-1, HOMO-2 and


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ACS Catalysis

HOMO-5 band centers of the fully hydroxylated surface lie 0.8, 0.36 and 0.48 eV higher than in the dehydrated surface. In HfO2, the respective figures are 0.92, 0.46 and 0.57 eV. For clarity of presentation, in the above we have chosen to analyze ZrO2 in detail, but the same arguments apply to HfO2 as well. The related densities of states are shown in Figures S9-S11 of the SI. Briefly, the LUMO band of MA is centered at ca. -3.0, -3.3 and -2.7 eV on the dehydrated, partially hydroxylated and fully hydroxylated HfO2 surfaces, respectively. Like MA bound to fully hydroxylated ZrO2, the valence bands of MA interact little with the surface atoms and lie higher than on the other surfaces, as evinced by the very modest broadening of the corresponding pdos. Thereby, we expect them to mix with the LUMO of furan and contribute to the stabilization of the Diels-Alder transition state on the fully hydroxylated surface. The comparison with alumina enhances our insight into these catalytic systems and quite pleasingly confirms the picture that emerges from the analysis presented above. The dos and pdos from binding at the undercoordinated Al atom of unhydrous Al2O3 (110) are provided in Figures S12(a)-(e) in the SI. The LUMO band of MA is located at -3.4 eV, which is lower by ca. 1.1 eV relative to the gas-phase molecule and by 0.4 eV relative to the MA LUMO band on unhydrous ZrO2 (or HfO2), consistent with the fact that the undercoordinated Al atom presents a stronger Lewis acid site compared to a fully coordinated Zr or Hf atoms. In Figures S13(a)-(e) of the SI, we present the dos and pdos analysis for the completely hydrated alumina surface. We find again clear evidence how important are the MA π-bands below the Fermi level in what we earlier referred to as secondary interactions between them and the LUMO of furan. The HOMO-1, HOMO-2 and HOMO-5 band centers of the fully hydroxylated surface lie 1.6, 1.0 and 0.95 eV higher than in the completely dehydrated surface (Table 8). It would thus seem that we have Lewis acid catalysis with marginal benefits on the dehydrated surfaces of ZrO2 and HfO2, and that the surface hydroxyl groups present a polar environment which, via hydrogen bonding, is solely responsible for catalysis on the completely hydrated surfaces. There is a clear synergy of the two—Lewis acidity and hydrogen-bond-mediated polarization—on the partially hydroxylated surfaces. CONCLUSION Diels-Alder cycloaddition with furans as dienes are slow reactions, even when the dienophile is activated by electron withdrawing groups, because of the aromatic nature of furans. Lewis acids


ACS Catalysis 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

are typical Diels-Alder catalysts as they provide an electrophilic environment that lowers the orbital energies of the dienophile. We have presented the first experimental and computational evidence that the band-gap metal oxides ZrO2 and HfO2 can catalyze Diels-Alder [4+2] cycloaddition between furan and methyl acrylate. Their effectiveness as well as the catalytic mechanisms seem to depend on the degree of hydroxylation of their surfaces. Analysis of the electronic states of three catalytic systems—dehydrated and hydrous (partially and fully hydroxylated) surfaces—reveals that the partially hydroxylated surfaces are more active than the rest. We have argued that this must be attributed to synergistic stabilization of the LUMO of methyl acrylate by both Lewis (metal atoms) and OH sites that coordinatively bind the acrylate’s carbonyl group. Via hydrogen bonding, the surface hydroxyl groups seem to polarize the carbonyl group of the dienophile and to be solely responsible for the rate acceleration on completely hydroxylated surfaces. The Lewis metal centers of the completely dehydrated ZrO2 and HfO2 surfaces turn out to be marginally effective. We have observed similar behavior by the completely hydrated (110) surface of γ-Al2O3, which is somewhat more active than the weak Lewis acid sites (tetracoordinated Al sites) of the activated surface. However, the under-coordinated Al sites of activated γ-Al2O3 (110) are more active than all the catalytic sites investigated here. According to reaction kinetics, ZrO2 and HfO2 present the same apparent activation energies, in the range of 11.5 kcal/mol, but HfO2 is intrinsically more active because of its smaller surface area and entropic phenomena that result in a twofold, higher pseudo-first order rate constant. The computed activation barriers are in good agreement with the experimental apparent values, especially in the case of partially hydroxylated surfaces. The kinetic studies with no-pretreated γAl2O3 revealed that it is as active as ZrO2 but intrinsically less active than HfO2. In addition, our experiments showed that alumina promotes side reactions and likely furan oligomerization. We conclude that, without compromising catalytic efficiency, ZrO2 and HfO2 present the advantage of selective catalysis without need for surface activation or a controlled environment. AUTHOR INFORMATION Corresponding Author * [email protected]


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ACS Catalysis

Present Addresses T.S.: Department of Chemical and Biomolecular Engineering, University of Houston, Houston, Texas G.R.J.: Environmental Lab, U. S. Army Engineer Research and Development Center, Vicksburg, Mississippi ASSOCIATED CONTENT Supporting Information Enthalpy values for states along reaction pathways Calculated work functions Images of optimized structures (intermediates and transition states) Frontier orbitals of methyl acrylate Density of states analysis for reaction on HfO2 and Al2O3 Cartesian coordinates of all structures This material is available free of charge via the Internet at http://pubs.acs.org ACKNOWLEDGEMENTS This material is based upon work supported as part of the Catalysis Center for Energy Innovation, an Energy Frontier Research Center funded by the U.S. Department of Energy, Office of Science, Office of Basic Energy Sciences under Award Number DE-SC0001004. This research used resources of the National Energy Research Scientific Computing Center, a DOE Office of Science User Facility supported by the Office of Science of the U.S. Department of Energy under Contract No. DE-AC02-05CH11231. TSF wishes to thank Dr. Jeffrey Frey for help with calculations.


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Scheme 1. Diels-Alder reaction of furan and methyl acrylate.

Table 1. Surface area and bulk chemical composition of the catalysts Catalyst

Surface area (m2 g-1)

Hf

Zr

Al

P

Cl

Ca

Ti

Fe

K

HfO2

3.9

97.3

0.372

-

0.771

0.101

0.321

1.16

-

-

ZrO2

23.9

2.487

97.043

-

-

0.098

0.277

-

-

0.092

γ-Al2O3

80-120a

-

-

99.071

0.577

-

0.199

-

0.081

0.071

a

Surface area provided by the vendor—an average number was used for the rate calculations

Table 2. Yields and area specific rates for Diels-Alder of furan and methyl acrylate at different temperatures. (𝑡r= 24h; furan : methyl acrylate molar ratio=5; in the experiments with γ-Al2O3, the amount of reactants was three times higher—at constant reactant concentration—to ensure efficient stirring.) Catalyst Oxanorbornene adduct yield ( Area specific rate (𝜇𝑚𝑜𝑙 . 𝑚 ―2 Temperature (K) %) ℎ ―1) 293 1.38 0.82 HfO2

ZrO2

γ-Al2O3

303

2.41

1.44

313 293 303 313 293 303 313

4.94 1.15 2.52 3.90 1.72 2.26 5.48

2.94 0.11 0.25 0.38 0.12 0.16 0.38


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ACS Catalysis

Table 3. Bader partial charges of surface-bound methyl acrylate (MA) and of the MA and furan moieties in the transition state complex, on dehydrated and hydrous ZrO2 and HfO2 (see Methods for details). (Charges in atomic units.) Species

ZrO2 Partially hydrous

Fully hydrous

Dehydrated

0.04

0.03

0.01

0.19

0.24

-0.12

-0.18

Dehydrated

Surface MA Furan in TS MA in TS

HfO2 Partially hydrous

Fully hydrous

0.04

0.04

0.01

0.19

0.19

0.24

0.19

-0.16

-0.12

-0.17

-0.16

Table 4. Energies of intermediates and transition states for “Langmuir”-type kinetics on dehydrated and hydrous ZrO2 and HfO2 (see Methods for details). (Energies in kcal/mol.) ZrO2 State Reactants Surface bound interacting complex Transition state Surface bound product Product intrinsic barrier

HfO2

0.0

Partially hydrous 0.0

Fully hydrous 0.0

0.0

Partially hydrous 0.0

Fully hydrous 0.0

-44.4

-37.4

-35.1

-47.3

-37.4

-33.3

-24.6

-16.5

-14.2

-27.0

-16.4

-12.9

-38.2

-32.7

-28.4

-40.5

-31.7

-28.1

-8.7 19.8

-8.7 20.9

-8.7 20.9

-8.7 20.3

-8.7 21.0

-8.7 20.4

Dehydrated

Dehydrated

Table 5. Energies of intermediates and transition states for “Eley-Rideal”-type kinetics on dehydrated and hydrous ZrO2 and HfO2 (see Methods for details). (Energies in kcal/mol.) State Reactants Surface bound interacting complex Transition state Surface bound product Product intrinsic barrier

0.0

ZrO2 Partially hydrous 0.0

0.0

HfO2 Partially hydrous 0.0

Fully hydrous 0.0

Fully hydrous 0.0

-31.7

-29.1

-32.2

-34.0

-28.6

-29.6

-14.8

-16.9

-19.2

-17.0

-15.6

-15.6

-38.4

-32.8

-37.4

-40.6

-32.5

-34.1

-8.7 16.9

-8.7 12.2

-8.7 13.0

-8.7 17.0

-8.7 13.0

-8.7 14.0

Dehydrated


Dehydrated

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Table 6. Energies of intermediates and transition states for “Eley-Rideal”-type kinetics on dehydrated and hydrous Al2O3 (see Methods for details). (Energies in kcal/mol.)

State Reactants Surface bound interacting complex Transition state Surface bound product Product intrinsic barrier

Dehydrated (3-coordinated) 0.0

Dehydrated (4-coordinated) 0.0

Fully hydrous 0.0

-52.1 -42.0 -54.2 -8.7 10.2

-46.5 -31.4 -48.3 -8.7 15.2

-21.1 -6.0 -28.7 -8.7 14.8

Table 7. Methyl acrylate HOMO-1, HOMO-2 and HOMO-5 band centers for dehydrated and hydrous ZrO2 and HfO2. (Energies in eV.)

HOMO-1

-8.2

ZrO2 Partially hydrous -8.3

HOMO-2

-7.94

-8.33

-7.58

-7.98

-8.31

-7.52

HOMO-5

-10.55

-10.8

-10.07

-10.58

-10.82

-10.01

Band

Dehydrated

Fully hydrous -7.4

-8.26

HfO2 Partially hydrous -8.35

Fully hydrous -7.34

Dehydrated

Table 8. Methyl acrylate HOMO-1, HOMO-2 and HOMO-5 band centers for dehydrated and hydrous Al2O3. (Energies in eV.) Band LUMO HOMO HOMO-1 HOMO-2 HOMO-3 HOMO-4 HOMO-5

Dehydrated Fully hydrous -3.48 -2.23 -8.10 -6.52 -8.36 -6.82 -8.05 -7.06 -9.66 -8.42 -9.67 -8.87 -10.45 -9.51


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ACS Catalysis

Figure 1. Dehydrated, partially (four water molecules) and completely hydroxylated (twelve water molecules) surfaces.


ACS Catalysis

ZrO2

Intensity (a.u.)

HfO2

Intensity (a.u.)

10

20

30

40

50

10

60

20

30

25

Al2O3

20

15

10

5

0 20

40

Angle (2theta)

Angle (2theta)

Intensity (a.u.)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

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40

60

80

Angle (2)

Figure 2. X-Ray diffraction patterns of HfO2, ZrO2 and γ-Al2O3 .


50

60

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ACS Catalysis

100nm

100nm

100nm

Figure 3. SEM images of HfO2, ZrO2 and Al2O3


ACS Catalysis

a) Hf 4f

a) O 1s

17.1eV

530.4eV

Intensity (a.u.)

Intensity (a.u.)

18.8eV

22

20

18

16

14

532.2eV

536

535

534

Binding energy (eV)

b) Zr 3d

532

531

530

b) O 1s

182.2eV

529

528

530.2eV

Intensity (a.u.)

184.6eV

Intensity (a.u.)

533

Binding Energy (eV)

531.9eV

188

186

184

182

180

536

535

534

Binding Energy (eV) c) Al 2p

533

532

531

530

529

528

529

528

Binding Energy (eV) c) O 1s

74.9eV

531.3eV

Intensity (a.u.)

532.04eV

Intensity (a.u.)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

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80

78

76

74

72

70

536

Binding Energy (eV)

535

534

533

532

531

530

Binding Energy (eV)

Figure 4. XPS spectra of a) Hf 4f and O 1s, b) Zr 3d and O 1s and c) Al 2p and O 1s.


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0.25

HfO2

Turnover frequency (h-1)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

ACS Catalysis

0.20

0.15

0.10

γ-Al2O3

0.05

0.00 290

ZrO2 295

300

305

310

Temperature (K) Figure 5. Temperature dependence of experimental turnover frequencies.


315

ACS Catalysis

-13.5

-13.5

HfO2 Ea=11.7kcal/mol

-13.8

-14.1

-14.1

ln(k)

-13.8

-14.4

-14.7

-15.0

-15.0

-15.3 3.15

3.20

3.25

3.30

3.35

ZrO2 Ea=11.3kcal/mol

-14.4

-14.7

-15.3 3.15

3.40

3.20

3.25

3.30

3.35

1000/T

1000/T

-13.5

Al2O3 Ea=10.6kcal/mol

-13.8

-14.1

ln(k)

ln(k)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

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-14.4

-14.7

-15.0

-15.3 3.15

3.20

3.25

3.30

3.35

3.40

1000/T

Figure 6. Arrhenius plots of the kinetics of Diels-Alder cycloaddition of furan and methyl acrylate.


3.40

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ACS Catalysis

Figure 7. Methyl acrylate equilibrium geometries on dehydrated (top), partially hydroxylated (middle) and fully hydroxylated (bottom) transition metal oxide surfaces. Top (on the left) and side (on the right) views are shown. For contrast, all oxide atoms are in purple.


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Figure 8. Optimized structures on dehydrated MO2 (M=Hf, Zr) surface for Eley Rideal (top) and Langmuir (bottom) kinetics. (a) Furan-methyl acrylate interacting complex, (b) transition state and (c) cycloadduct.


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ACS Catalysis

Figure 9. Reaction enthalpy profile Eley-Rideal kinetics on dehydrated and hydrous ZrO2 at 𝑇 = 293𝐾. The black dashed line represents the uncatalyzed reaction.


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Figure 10. Reaction enthalpy profile Eley-Rideal kinetics on dehydrated and hydrous HfO2 at 𝑇 = 293𝐾. The black dashed line represents the uncatalyzed reaction.


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Figure 11. (a) Total density of states (dos) of the clean dehydrated ZrO2. (b) Total density of states of dehydrated ZrO2 with surface bound methyl acrylate. (c) Contributions to the dos of clean dehydrated ZrO2 from the d-band of surface Zr atoms and from the p-band of surface O atoms that are in the immediate vicinity of methyl acrylate when it is bound to the surface. (d) Contributions to the dos of dehydrated ZrO2 with surface bound methyl acrylate from the d-band of surface Zr atoms and from the p-band of surface O atoms that are in the immediate vicinity of methyl acrylate. (e) Contributions to the dos of dehydrated ZrO2 with surface bound methyl acrylate from the molecular orbitals of methyl acrylate.


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Figure 12. (a) Total density of states (dos) of the clean partially hydroxylated ZrO2. (b) Total density of states of partially hydroxylated ZrO2 with surface bound methyl acrylate. (c) Contributions to the dos of clean partially hydroxylated ZrO2 from the d-band of surface Zr atoms and from the p-band of surface O atoms that are in the immediate vicinity of methyl acrylate when it is bound to the surface; O(h) denotes hydroxylated surface oxygen atoms (d) Contributions to the dos of partially hydroxylated ZrO2 with surface bound methyl acrylate from the d-band of surface Zr atoms and from the p-band of surface O atoms that are in the immediate vicinity of methyl acrylate. (e) Contributions to the dos of partially hydroxylated ZrO2 with surface bound methyl acrylate from the molecular orbitals of methyl acrylate.


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Figure 13. (a) Total density of states (dos) of the clean fully hydroxylated ZrO2. (b) Total density of states of fully hydroxylated ZrO2 with surface bound methyl acrylate. (c) Contributions to the dos of clean fully hydroxylated ZrO2 from the d-band of surface Zr atoms and from the p-band of surface O atoms that are in the immediate vicinity of methyl acrylate when it is bound to the surface; O(h) denotes hydroxylated surface oxygen atoms (d) Contributions to the dos of fully hydroxylated ZrO2 with surface bound methyl acrylate from the d-band of surface Zr atoms and from the p-band of surface O atoms that are in the immediate vicinity of methyl acrylate. (e) Contributions to the dos of fully hydroxylated ZrO2 with surface bound methyl acrylate from the molecular orbitals of methyl acrylate.


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Lewis Acid Site and Hydrogen-Bond-Mediated Polarization Synergy in

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