The Oxidation of Methanol to Formaldehyde on TiO2 (110)-Supported

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J. Phys. Chem. B 2001, 105, 1366-1373

The Oxidation of Methanol to Formaldehyde on TiO2(110)-Supported Vanadia Films G. S. Wong, D. D. Kragten, and J. M. Vohs* Department of Chemical Engineering, UniVersity of PennsylVania, Philadelphia, PennsylVania 19104-6393 ReceiVed: October 6, 2000; In Final Form: December 11, 2000

The reaction of methanol on TiO2(110)-supported vanadium oxide was studied using temperature-programmed desorption (TPD) and high-resolution electron energy loss spectroscopy (HREELS). TPD results show that methanol is oxidized to formaldehyde on monolayer and submonolayer vanadia films on TiO2(110), whereas both clean TiO2(110) and multilayer vanadia films supported on TiO2(110) are relatively inactive for this reaction. HREELS results demonstrate that methoxides are the primary surface intermediates in the oxidation of methanol to formaldehyde on the supported vanadium layers. The reactivity trends obtained for the model catalysts used in this study are similar to those observed for high surface area analogues. This suggests that vanadia films supported on single-crystal metal oxide substrates are excellent model systems for studying the relationships between the structure and reactivity of supported oxide catalysts.

Introduction Vanadium oxide supported on a second metal oxide such as TiO2, CeO2, Al2O3, or ZrO2 is an active catalyst for a variety of reactions, including selective oxidation (e.g., o-xylene to phthalic anhydride and methanol to formaldehyde), ammoxidation of aromatic hydrocarbons, and the selective catalytic reduction of NOx with ammonia (SCR).1-5 It is well established that the active form of vanadium in these catalysts consists of monolayer or submonolayer coverages of an oxidized vanadium species. The monolayer vanadia complex has a structure distinct from that of bulk V2O5 and exhibits higher activity and selectivity than the unsupported oxide.1-3 It has also been shown that the identity of the support significantly influences catalytic activity.5,6 For example, Deo and Wachs have reported that the turnover rate for methanol oxidation over vanadia supported on various metal oxides varies by as much as 3 orders of magnitude.6 Although the synergism between vanadia and an underlying oxide support is well documented, many questions still remain concerning the specific interactions at the interface and how they influence reactivity. In an effort to better understand the relationships between the structure and activity of supported vanadia catalysts, we have been studying well-defined models of supported vanadia catalysts composed of mono- and multilayer coverages of oxidized vanadium on single-crystal TiO2(110) substrates. The use of single-crystal substrates allows for the application of a wide range of surface-sensitive spectroscopic probes, including high-resolution electron energy loss spectroscopy (HREELS) and temperature-programmed desorption (TPD). In a previous letter, we presented preliminary TPD results that showed that the reactivity of the model catalysts for the selective oxidation of methanol exhibited trends similar to that observed for high surface area analogues.7 In this paper, we present results of a more thorough investigation of the structure and reactivity of models of supported vanadia/titania catalysts. Experimental Section Experiments were performed in two separate ultrahigh vacuum (UHV) chambers. The UHV chamber used for the

temperature-programmed desorption studies was equipped with a quadrupole mass spectrometer (UTI), an ion sputter gun (Physical Electronics), a quartz crystal film thickness monitor (Maxtek, Inc.), and a retarding field electron energy analyzer (Omicron) which was used for both low-energy electron diffraction (LEED) and Auger electron spectroscopy (AES). The other UHV chamber was used for high-resolution electron energy loss spectroscopy (HREELS) studies and was equipped with an LK Technologies model 3000 HREEL spectrometer, a quadrupole mass spectrometer (UTI), an ion sputter gun (Physical Electronics), a quartz crystal film thickness monitor (Maxtek, Inc.), and AES/LEED optics (OCI). The TiO2(110) single crystal was obtained from Commercial Crystal Laboratories and was cleaned by repeated cycles of sputtering and annealing. An evaporative vanadium metal (Alfa Aesar, 99.8%) source was used to grow the vanadia thin films. The source consisted of a tungsten wire (0.2 mm in diameter), wrapped with a short length of vanadium wire. The tungsten wire was attached to an electrical feed-through on the UHV chamber, allowing it to be heated resistively. Monolayer or multilayer coverages of vanadia (as determined by the film thickness monitor) were grown by first depositing a layer of vanadium atoms on a clean TiO2(110) substrate, followed by annealing in 1 × 10-6 Torr of O2 (Matheson, >99.6%) at 500 K for 30 min. Methanol (Fisher, HPLC grade) was purified using repeated freeze-pump-thaw cycles prior to use and was admitted into the vacuum system using a variable leak valve. Formaldehyde was produced by heating p-formaldehyde (Aldrich, 95%) and was also admitted into the vacuum system using a variable leak valve. For TPD experiments, the sample was exposed to 2 L of CH3OH or H2CO at 300 K and then heated at 2 K/s. The mass spectrometer was multiplexed allowing multiple masses to be monitored simultaneously. All HREEL spectra were collected in the specular direction using a 6 eV electron beam directed 60° from the surface normal with the sample held at 200 K. A typical HREEL spectrum of the clean TiO2(110) surface is displayed in Figure 1 and contains three intense peaks centered at 369, 450, and 755 cm-1 that are due to excitation of the Fuchs-Kliewer optical phonon modes

10.1021/jp003691u CCC: $20.00 © 2001 American Chemical Society Published on Web 01/25/2001

TiO2(110)-Supported Vanadia Films

J. Phys. Chem. B, Vol. 105, No. 7, 2001 1367

Figure 1. (a) Raw and (b) Fourier deconvoluted HREEL spectra of TiO2(110). The fwhm for the elastic peak was 50 cm-1. The electron beam energy was 6 eV, and all spectra were collected at 200 K.

of the TiO2(110) surface.8,9 The series of smaller peaks at higher energies are due to multiple scattering events and occur at integer multiples of the phonon energies. Unfortunately, for adsorbate-covered surfaces, these combination peaks can obscure adsorbate vibrational losses that occur at energies greater than those of the phonon fundamentals. To alleviate this problem, we used a Fourier deconvolution procedure that removed the peaks due to multiple scattering events from the spectrum. The details of this procedure and its application to HREEL spectra of metal oxides has previously been described in detail.10-14 A deconvoluted spectrum of TiO2(110) which was obtained using this method is displayed in Figure 1. Note that after deconvolution the spectrum is essentially flat at energies greater than 1000 cm-1.

Figure 2. (a) AES spectra for various coverages of vanadia on TiO2(110). (b) The ratio of the Ti(LMM) to V(LMM) AES signals vs vanadia coverage.

Results and Discussion Growth and Characterization of Vanadia Layers. The growth of vanadia thin films on the TiO2(110) substrate was characterized using both AES and HREELS. AES spectra and the ratio of the Ti(LMM) peak-to-peak height to that of V(LMM)-versus-vanadia coverage are displayed in panels a and b, respectively, of Figure 2. Note that the peak height ratio decreases from 7.0 to 0.36 as the vanadia coverage is increased from 0.25 to 2 ML. In the spectrum of the 8 ML vanadia film, the Ti(LMM) peak was not observed. The rapid decrease in the Ti:V ratio suggests a Frank-van der Merve (i.e., layer-bylayer) growth mode for the vanadia film, although StranskiKrastanov growth (i.e., monolayer formation followed by growth of three-dimensional particles) cannot be completely ruled out. The data are inconsistent, however, with three-dimensional Volmer-Weber growth. This result is similar to that reported in a previous study of the growth of vapor-deposited vanadia films on TiO2(110) by Madix et al.15 In that study based on XPS and STM data, it was concluded that growth occurs via either simultaneous multilayer formation or layer-by-layer. Although photoelectron spectroscopy was not available to characterize the vanadia films used in this study, Madix and co-workers have used both XPS and near-edge X-ray absorption

Figure 3. HREEL spectra of vanadia/TiO2(110) as function of vanadia coverage. (a) Clean TiO2 and (b) 0.25, (c) 0.5, (d) 1, and (e) 8 ML of vanadia.

fine structure (NEXAFS) to characterize vanadia films on TiO2(110) that were grown using a procedure nearly identical to that used here.16 The results of that study show that the vanadium cations in the vapor-deposited vanadia films are in the +3 oxidation state. HREEL spectra as a function of vanadia coverage are displayed in Figure 3. Spectrum a in this figure corresponds to the clean TiO2(110) substrate. As noted above, the intense peaks centered at 369, 450, and 755 cm-1 are due to excitation of the

1368 J. Phys. Chem. B, Vol. 105, No. 7, 2001

Figure 4. Fourier deconvoluted HREEL spectra of (a) 0.5 and (b) 1 ML vanadia films on TiO2(110).

three optically active surface phonon modes of TiO2(110). The smaller peaks that appear at higher energies are phonon-phonon combination peaks that result from multiple scattering events. Spectra b and c were obtained immediately after depositing 0.25 and 0.5 ML of vanadia, respectively. Except for a noticeable decrease in the intensity of the phonon modes, these spectra are similar to those obtained from the clean TiO2(110) surface. Further increases in the vanadia coverage produced more dramatic changes in the HREEL spectra. In addition to a decrease in intensity, a pronounced broadening of the phonon peaks is evident in spectra d and e in Figure 3, which were obtained after depositing 1 and 8 ML of vanadia, respectively. A slight downward shift in the energy of the large phonon peak at 755 cm-1 in the spectrum of the clean TiO2(110) surface is also evident. This peak is centered at 738 cm-1 in the spectrum for the 1 ML film, and it shifts to 710 cm-1 in the spectrum for the 8 ML film. This shift is most likely due to an overlap of the substrate phonon peaks with peaks due to V-O stretching modes. On the basis of Raman spectra of supported vanadia catalysts, the V-O stretching modes would be expected to occur between 700 and 1030 cm-1.5,6 Note, however, that individual V-O stretching modes cannot be resolved in this region of the HREEL spectrum. Previous HREELS studies of epitaxial metal overlayers on oxides have shown that a 1 or 2 ML metal film is sufficient to completely shield the substrate phonon modes from the incident electron beam.17 As noted above, XPS studies have shown that the vanadium cations in the vanadia films are most likely in the +3 oxidation state, suggesting the formation of a V2O3 layer. Since V2O3 is metallic at temperatures above 160 K,18 it is therefore somewhat surprising that TiO2 phonon peaks are still present in the spectrum of the sample with the 8 ML film. This indicates that even though the vanadium cations are in the +3 oxidation state, the vanadia film does not have electronic properties consistent with those of bulk V2O3. Figure 4 displays deconvoluted versions of the HREEL spectra of the 0.5 and 1.0 ML films. Note that several peaks are evident in these spectra: a broad peak between 3100 and 3700 cm-1 and two narrower peaks centered at 1170 and 1610 cm-1. The two higher-energy peaks can be assigned to O-H stretching and O-H bending modes of H2O that adsorbed during

Wong et al.

Figure 5. CH3OH TPD from TiO2(110); shown here are the signals for m/e 29, 30, and 31.

collection of the spectra.19 These spectra were collected immediately after film growth at a sample temperature of 200 K. The low sample temperature was used in order to obtain higher resolution by freezing out the low-energy plasmon modes of the conduction electrons in the TiO2 substrate. Unfortunately, the high O2 pressures used while oxidizing the film produced an increase in the background level of H2O in the chamber. These factors, coupled with the fact that it took 1-2 h to collect each spectrum, made adsorption of background water difficult to avoid. The peak centered at 1170 cm-1 in the deconvoluted spectra of the 0.5 and 1 ML vanadia films is not at an energy consistent with that of the vibrational modes of adsorbed water. It is possible that this peak is due to a V-O stretching mode. Note that in Raman spectra of high surface area supported vanadia catalysts, a terminal VdO stretching mode appears near 1030 cm-1.6 Additional evidence, however, is needed before a definitive assignment can be made. Interaction of CH3OH with TiO2(110). Before discussing the reaction of CH3OH on the supported vanadia films, it is useful to consider the interaction of CH3OH with clean TiO2(110). Figure 5 displays TPD spectra for several m/e values obtained from a TiO2(110) surface dosed with 2 L of CH3OH at 300 K. Peaks were observed only for m/e values in the cracking pattern of CH3OH (i.e., 31, 30, 29, and 28). The relative intensities of these peaks were also consistent with the cracking pattern of methanol. Thus, the only desorbing species detected during CH3OH TPD on TiO2(110) was the parent molecule. As shown in the figure, CH3OH desorbs in a series of three overlapping peaks centered at 360, 390, and 420 K. In addition to these peaks, there is a broad tail on the CH3OH desorption signal which extends up to ∼600 K. This result is qualitatively similar to that reported previously by Henderson et al.9,20 In that study, the TPD spectra obtained following exposure of a vacuum-annealed TiO2(110) surface to CH3OH at 135 K contained peaks centered at 165, 295, 350, and 480 K, which were assigned to desorption of CH3OH multilayers, molecularly adsorbed CH3OH, recombinative desorption of methoxide species on nonvacancy sites, and recombinative desorption of methoxide species on oxygen vacancy sites, respectively. The

TiO2(110)-Supported Vanadia Films assignment of the two higher-temperature peaks was based on both HREELS results and comparisons to H2O TPD studies.19 Henderson et al.9 have also reported that during CH3OH TPD on oxygen-predosed TiO2(110) surfaces, some formaldehyde is formed at 625 K. The fraction of adsorbed CH3OH that reacted to form H2CO on these surfaces was very small and was estimated to be 0.04 ML. Although the mechanism for H2CO production on the oxygen exposed surfaces is not completely clear, the following possibility was suggested. Molecular oxygen adsorbs on surface oxygen vacancy sites, resulting in vacancy filling and the formation of an oxygen adatom. The oxygen adatoms react with methanol at room temperature, forming a terminal OH group and an adsorbed methoxide intermediate. A fraction of the terminal OH groups formed in this manner react at room-temperature, producing water that desorbs at 300 K. Since this removes hydrogen from the surface, a small fraction of the methoxide species becomes stranded without an adsorbed hydrogen to recombine with upon heating. It is these stranded methoxides that decompose to form formaldehyde at 625 K. We also performed CH3OH TPD experiments on TiO2(110) samples that were annealed in 10-8 Torr of O2 at 600 K for 10 min. The results of these experiments were essentially identical to those obtained on the vacuum-annealed samples, and formaldehyde production at 625 K was not observed. It is not clear why the formaldehyde production reported by Henderson et al. on O2-dosed samples was not observed in the present study. Assuming that surface oxygen vacancies are required in order to form the oxygen adatoms upon exposure to O2, it is possible that differences in the sample preparation procedures may have produced different concentrations of surface oxygen vacancies in the two studies. Differences in the CH3OH dosing temperatures between the two works (300 K in this study and 135 K in the Henderson study) may also play a role. HREELS was also used to characterize the interaction of CH3OH with the clean TiO2(110) surface. Deconvoluted HREEL spectra obtained after dosing a clean TiO2(110) surface at 200 K with 20 L of CH3OH and after flashing the dosed surface to 300 K are displayed in Figure 6. These experiments were done at least 1 day after the films were deposited in order to allow the background pressure in the chamber to recover. Thus, unlike the HREELS results reported above, the adsorption of background water was not observed. The spectrum obtained after dosing at 200 K contains peaks due to the adsorbed species centered at 1014, 1064, 1163, 1236, 1366, 1470, 2848, 2965, and 3082 cm-1. The spectrum obtained after heating to 300 K is nearly identical, except that a few of the peaks are difficult to resolve and appear as shoulders. A spectrum was also collected after heating to 500 K, and peaks due to adsorbates could not be resolved from the background. Table 1 presents a comparison of the HREEL spectra of adsorbed CH3OH on TiO2(110) with the IR spectrum of both gaseous and liquid CH3OH.21 With the exception of the small peak near 1230 cm-1 and the lack of an O-H stretching mode in the HREEL spectrum, there is good agreement between the spectrum for gaseous CH3OH and that for the adsorbed species. Although the lack of a discernible O-H stretching mode may suggest that the methanol has dissociated to form methoxide, it is more likely due to hydrogen bonding between adsorbed molecular species, resulting in a broad O-H stretching peak that cannot be resolved. This explanation is also consistent with the fact that a prominent O-H bending mode is present in the HREEL spectra of the CH3OH-dosed surfaces at 1366 cm-1. This peak appears in the HREEL spectra of molecular methanol

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Figure 6. HREEL spectra of TiO2(110) dosed with 20 L of CH3OH and flashed to various temperatures: (a) clean TiO2 at 200 K, (b) TiO2 dosed with CH3OH at 200 K, then flashed to (c) 300, (d) 400, and (e) 500 K.

adsorbed on metal surfaces and, as would be expected, disappears upon dissociation to from methoxides.22 The HREELS results, therefore, indicate that the majority of the methanol is adsorbed molecularly on the TiO2(110) surface at temperatures below 300 K; however, the possibility that a small fraction of the adsorbed methanol dissociates to form methoxide cannot be ruled out. As noted above, Henderson et al. have also used HREELS to characterize the interaction of CH3OH with TiO2(110).9 The results of that study were similar to those reported here, except that they were able to resolve a small peak at 3660 cm-1 in the spectra of CH3OH-dosed surfaces that had been heated to 400 K. They attributed this peak to the O-H stretching mode of surface hydroxyl groups. In the Henderson study, the TiO2(110) sample was dosed at 135 K, which was low enough to produce multilayers of condensed CH3OH. The methanol coverages in that study may therefore have been higher than those used here, making methoxides easier to detect. Interaction of CH3OH with Vanadia/TiO2(110). As noted above, we have previously reported that during the TPD of CH3OH on a 1 ML vanadia film on TiO2(110), H2CO is produced at 660 K.7 For comparison purposes, a representative set of TPD spectra obtained following exposure of a 1 ML vanadia film on TiO2(110) to 2 L of methanol at 300 K is displayed in Figure 7. In addition to CH3OH, peaks are observed for H2CO, H2O, and CO. The CH3OH desorption spectrum is similar to that obtained from the clean TiO2(110) surface and contains three overlapping peaks at low temperature, centered at 368, 392, and 420 K. The CH3OH desorption curve also contains a broad tail that extends up to 700 K. Two additional broad peaks centered at ca. 550 and 660 K appear superimposed on this tail. Formaldehyde was produced in a single peak centered at 660 K. The curve for H2CO in the figure corresponds to the total m/e 29 signal minus the contribution from the m/e 29 methanol cracking fragment. A small m/e 28 peak due to CO desorption was detected near 660 K. This spectrum has also been corrected for the contribution due to the m/e 28 methanol cracking fragment. The H2O desorption signal rises throughout the TPD

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TABLE 1: Vibrational Frequencies of Gaseous and Liquid CH3OH and of CH3OH-Dosed TiO2(110)a CH3OH gas, IRb

CH3OH liquid, IRb

CO str CH3 rock

1033 1060, 1165c

1030 1115, 1165c

OH bend CH3 s-deform CH3 d-deform CH3 s-str CH3 d-str OH str

1345 1455 1477 2844 2960, 3000 3681

1418 1450 1480 2834 2946, 2980 3328

vibrational mode

a

CH3OH/TiO2(110) 200 K, HREELS

CH3OH/TiO2(110) 300 K, HREELS

1015 1064, 1163 1236 1366

1015 (1064), 1162 (1224) 1366 1451 (1490) (2850) 2985, (3050)

1470 2848 2965, 3082

All values in cm-1, and all parentheses denote shoulders. b From ref 21. c From Raman.

Figure 7. CH3OH TPD from 1 ML vanadia film on TiO2(110); shown here are CO, H2CO, H2O, and CH3OH. Unlike the CH3OH TPD on TiO2(110), there is formaldehyde production at 660 K.

run, indicating that water is also produced via the reaction of CH3OH on the vanadia film. We have previously attributed the formaldehyde peak at 660 K to the reaction of methanol on the supported vanadia layer. In light of the recent work of Henderson et al., which shows that methanol can react to produce formaldehyde on oxygen-dosed TiO2(110),9 this assignment warrants further discussion. Since the growth mode for the vanadia film on TiO2(110) may not completely be layer-by-layer, a small portion of the TiO2(110) surface may have been exposed for the 1 ML vanadia sample. Since this sample was annealed in O2 during vanadia deposition, it is possible that the formaldehyde peak at 660 K results from the reaction of methanol with oxygen adatoms on exposed regions of the TiO2(110) surface. Several pieces of data provide evidence, however, against this scenario. First, the formaldehyde peak temperature from the 1 ML vanadia sample is 35 K higher than that reported for oxygen-dosed TiO2(110). Second and more importantly, the fraction of adsorbed methanol that reacts to form H2CO on the vanadia film is 14%, which is substantially higher than that observed for reaction on oxygendosed TiO2(110). Since the total amount of methanol that adsorbed on the clean and vanadia-covered samples was roughly equivalent, this result is opposite to what would be expected if reaction on exposed portions of the substrate were responsible for formaldehyde production on the VOx/TiO2(110) sample.

Figure 8. H2CO desorption spectra obtained from a CH3OH-dosed 1 ML vanadia film on TiO2(110). Spectrum a is from the first run with a freshly oxidized sample, while spectra b and c are from the second and third runs, respectively. Spectrum d was obtained after reoxidizing the sample used in spectrum c in 1 × 10-6 Torr of O2 for 30 min at 500 K.

These results therefore support the assignment of H2CO production to the reaction of adsorbed methanol on the supported vanadia layer. Additional evidence will be presented below that provides further verification for this assignment. To determine whether the formaldehyde peak at 660 K was due to a reaction- or desorption-limited process, we performed formaldehyde TPD experiments. TPD spectra obtained for a 1 ML vanadia film dosed with 2 L of H2CO at 300 K contained a single formaldehyde desorption peak centered at 380 K. This result confirms that in the case of the CH3OH-dosed sample, the H2CO peak at 660 K is due to a reaction-limited process most likely involving the breaking of a C-H bond in an adsorbed methoxide intermediate. Figure 8 displays the H2CO peak for a series of CH3OH TPD experiments with the 1 ML vanadia sample. Spectrum a in the figure was obtained after dosing the freshly prepared film with CH3OH and contains a H2CO peak centered at 660 K. Spectra b and c correspond to the second and third CH3OH TPD runs in the series. Note that H2CO production was not observed in these runs. Since H2O was also produced during the first TPD experiment, this change in reactivity can be attributed to reduction of the vanadia layer during the first TPD run. Spectrum d in the figure was obtained after reoxidizing the vanadia film in 1 × 10-6 Torr of O2 at 500 K. The H2CO desorption peak reappeared in this run.

TiO2(110)-Supported Vanadia Films

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Figure 10. Deconvoluted HREEL spectra obtained from a 1 ML vanadia film on TiO2(110) (a) dosed with 20 L of CH3OH at 200 K and then flashed to (b) 300, (c) 500, (d) 600, and (e) 700 K. Figure 9. The upper panel contains H2CO desorption spectra from CH3OH-dosed VOx/TiO2(110) samples for various vanadia coverages. The relative yields of H2CO and CH3OH in the TPD spectra as a function of vanadia coverage are plotted in the lower panel.

In addition to the 1 ML vanadia film, CH3OH TPD experiments were also performed with TiO2(110) samples being covered with 0.5 and 2.5 ML of vanadia. The methanol desorption curves obtained for these vanadia coverages were similar to those for the 1 ML film. Differences were observed, however, in the H2CO desorption curves for the various vanadia coverages. The H2CO peak (m/e 29) as a function of vanadia coverage is displayed in Figure 9. The relative yield of H2CO as a function of vanadia coverage is also plotted in this figure. Several trends are apparent in the data. First, the area of the H2CO peak varies with vanadia coverage and goes through a maximum near 1 ML. Only a very small H2CO peak was detected in the spectrum from the 2.5 ML vanadia film, indicating that multilayer vanadia films are nearly inactive for the oxidation of methanol to formaldehyde. The fact that the amount of formaldehyde produced increased with vanadia coverage up to 1 ML provides further evidence that the reaction which produces formaldehyde occurs on the vanadia. If oxygen adatoms on the TiO2(110) surface were the active site for this reaction, the opposite trend would be expected, since the amount of exposed TiO2(110) surface decreases with increasing vanadia coverage. The second trend observed in Figure 9 is that the H2CO peak temperature decreases with increasing vanadia coverage. For the 0.5 ML vanadia film, the H2CO peak was centered at 700 K, and for the 1 ML vanadia film, this temperature decreased to 660 K. For the 2.5 ML film, the H2CO peak was centered at 640 K. Thus, the reaction that produces formaldehyde is structure-sensitive, and the activation energy is a function of the vanadia coverage. If first-order kinetics and a pre-exponential factor of 1013 s-1 are assumed, the TPD peak temperatures correspond to activation energies of 187, 176, and 170 kJ/mol for the 0.5, 1.0, and 2.5 ML vanadia films, respectively. Previous studies of the structure of vanadia overlayers on TiO2(110) may provide some clues as to the origin of this

variation in activation energy with vanadia coverage. Madix et al. have proposed on the basis of both the STM results and the analysis of attenuation effects in XPS data that the vapordeposition growth of vanadia on TiO2(110) proceeds via the formation of isolated vanadia nuclei at low coverage which coalesce into a textured film as the coverage is increased.15 The results of the present study indicate that the activation energy for the reaction of methoxide to formaldehyde on the isolated vanadia nuclei is greater than that on the more two-dimensional films that are formed at higher vanadia coverages. Deconvoluted HREEL spectra obtained from a CH3OH-dosed 1 ML vanadia sample as a function of temperature are displayed in Figure 10. These spectra were collected after dosing 20 L of CH3OH at 200 K and then after briefly annealing to various temperatures, as indicated in the figure caption. The spectrum obtained after dosing at 200 K is similar to that for CH3OHdosed TiO2(110) at 200 K (see Figure 6) and can therefore be assigned to molecularly adsorbed CH3OH. Unlike with TiO2(110), however, heating the CH3OH-dosed 1 ML vanadia film to 300 K resulted in several changes in the spectrum. These include a decrease in the relative intensities of the peaks between 1300 and 1600 cm-1 and the appearance of a new peak at 1329 cm-1. Further heating to 500 K produced relatively few changes in the spectrum, while heating to 600 K resulted in a marked decrease in the intensity of the adsorbate related features. HREEL spectra of CH3OH-dosed samples as a function of vanadia coverage are displayed in Figure 11. These samples were dosed with 20 L of CH3OH at 200 K and then briefly heated to 500 K before the spectra were collected. For comparison purposes, the HREEL spectrum of CH3OH-dosed TiO2(110) flashed to 300 K is also included in the figure. With the exception of slight variations in peak intensities, the spectra of the CH3OH-dosed vanadia films did not change significantly with vanadia coverage, indicating that the same surface species were formed on all of the samples. As noted above, comparison of the spectra from the vanadia films with that for CH3OH adsorbed on TiO2(110) reveals several differences. One of the most noticeable is the intensity of the peak at 1366 cm-1. This peak is much less intense in the spectra obtained from the

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TABLE 2: Vibrational Frequencies of Methoxides on Metals and of CH3OH-Dosed VOx/TiO2a CH3O/Pd(110)b

vibrational mode

CH3O/Ni(10)c

CO str CH3 rock

1010 1130

1030 1150

CH3 def

1460

1440

CH str a

2800, 2975 -1 b

2790, 2910

CH3OH/0.25 ML VOx/TiO2, 500 K

CH3OH/0.5 ML VOx/TiO2, 500 K

CH3OH/1 ML VOx/TiO2, 500 K

1060 1160 1360 1460 1620 2830, 2940, 3010

1065 1155

1089 1165 1345 1440 1590 2953

1420 1612 2960

c

All values in cm . From ref 22. From ref 23.

asymmetrical O-C-O stretching modes which appear near 1350 and 1650 cm-1, respectively. Since these modes have relatively large dipole moments, they are easy to observe in HREELS. This conclusion is also consistent with IR and Raman studies of the interaction of methanol with high surface area vanadia/titania catalysts where surface formate intermediates have been observed.27 The small CO desorption peak in the TPD spectra from the CH3OH-dosed, 1 ML vanadia film may have been due to the decomposition of these formate intermediates. The data in Figure 11 also suggest that methoxide intermediates are formed on the 4.0 ML vanadia film and that these species are stable up to at least 500 K. This is an interesting result, since multilayer vanadia films were nearly inactive for the oxidation of methanol to formaldehyde. This demonstrates that the lack of methanol oxidation activity for the multilayer films is not due to an inability of these films to activate the O-H bond in adsorbed methanol. Rather, the primary reaction pathway for adsorbed methoxides on the multilayer films is recombination with surface hydrogen and desorption. Summary Figure 11. Deconvoluted HREEL spectra of samples dosed with 20 L of CH3OH at 200 K. Spectrum a is for a CH3OH-dosed TiO2(110) sample that was flashed to 300 K. The remaining spectra correspond to CH3OH-dosed vanadia films on TiO2(110) flashed to 500 K. The vanadia coverages are (b) 0.25, (c) 0.5, (d) 1, and (e) 4 ML.

vanadia thin films than in that from TiO2(110). A discrete O-H stretching mode centered near 3370 cm-1 is also apparent in the spectra for the vanadia films. Both of these observations are consistent with dissociation of adsorbed methanol on the supported vanadia film to form methoxide intermediates. The peak positions in the spectra of the CH3OH-dosed samples with 0.25, 0.5, and 1.0 ML vanadia coverages are listed in Table 2. Additionally, the peak positions in the spectra of methoxide intermediates adsorbed on Pd(110)22 and Ni(110)23 are also presented in the table. Note that with the exception of the peaks centered near 1360 and 1600 cm-1, which are only present for the VOx/TiO2(110) samples, there is good agreement between the spectra of methoxides on metal surfaces with those from the vanadia films. This result is again consistent with dissociation of CH3OH on the vanadia films to form methoxide intermediates. There are several possibilities for the origin of the peaks at 1360 and 1600 cm-1, including the presence of molecularly adsorbed CH3OH and/or H2O. Since the samples were heated to 500 K, which should be sufficient to desorb molecular species, and the background pressure in the chamber was ∼3 × 10-10 Torr during data collection, both of these possibilities seem unlikely. A more plausible explanation is the formation of a small number of formate intermediates. The two most intense peaks in the HREEL spectra of formates adsorbed on both metal24 and metal oxide25,26 surfaces are the symmetrical and

The results of this study demonstrate that monolayer and submonolayer vanadia films supported on TiO2(110) in which the vanadium cations are in the +3 oxidation state are active for the oxidation of methanol to formaldehyde. Methanol adsorbs dissociatively on the vanadia films to form surface methoxide intermediates, which are stable up to 600 K. At higher temperatures, abstraction of a hydride from the methoxide results in the formation of gaseous H2CO. This reaction is structure-sensitive, and the activation energy decreases with increasing vanadia coverage. Although CH3OH also adsorbed dissociatively on multilayer vanadia films on TiO2(110), these films were found to be nearly inactive for the oxidation of methanol to formaldehyde. The primary reaction pathway for methoxides on the multilayer films was recombination and desorption as molecular methanol. Under the reaction conditions, high surface area vanadia/ titania catalysts contain vanadium cations in both +3 and +5 oxidation states.28,29 It is generally accepted, however, that the active sites involve V+5. In light of this, one must be careful in using the results of the present study to provide insight into structure-activity relationships for high surface area catalysts. It is interesting, however, that the reactivity trends for the VOx/ TiO2(110) model catalysts are similar to those observed for high surface area analogues. For both systems, monolayer and submonolayer coverages of vanadia are active for the oxidation of methanol to formaldehyde, while multilayer coverages are nearly inactive. Differences in the reactivity of isolated vanadyl species and polyvanadates have also been reported for high surface area vanadia/titania catalysts.5 These results suggest that vapor-deposited vanadia monolayers on single-crystal metal oxide supports may be good model systems for studying the

TiO2(110)-Supported Vanadia Films relationships between the structure and activity of supported monolayer oxide catalysts. Acknowledgment. We gratefully acknowledge the financial support of the National Science Foundation (Grant CTS9712774) and the Laboratory for Research on the Structure of Matter at the University of Pennsylvania for the use of its facilities. References and Notes (1) Bond, G. C. Appl. Catal., A 1997, 157, 91. (2) Centi, G. Appl. Catal., A 1996, 147, 267. (3) Deo, G.; Wachs, I. E.; Haber, J. Crit. ReV. Surf. Chem. 1994, 4, 141. (4) Went, G. T.; Leu, L.; Rosin, R. R.; Bell, A. T. J. Catal. 1992, 134, 492. (5) Khodakov, A.; Olthof, B.; Bell, A. T.; Iglesia, E. J. Catal. 1999, 181, 205. (6) Deo, G.; Wachs, I. E. J. Catal. 1991, 129, 307. (7) Wong, G. S.; Kragten, D. D.; Vohs, J. M. Surf. Sci. Lett. 2000, 452, 293. (8) Fuchs, R.; Kliewer, K. L. Phys. ReV. 1963, 140, 2076. (9) Henderson, M. A.; Otero-Tapia, S.; Castro, M. E. Faraday Discuss. 1999, 114, 313. (10) Cox, P. A.; Flavell, W. R.; Williams, A. A.; Egdell, R. G. Surf. Sci. 1985, 152, 784-90.

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