J, Phys. Chem. 1980, 84, 3075-3079
(12) (13) (14) (15) (16)
ment to the study of the picosecond dynamics of this phenomenon.2s>26 Acknowledgment. We are grateful to the Deutsche Forschungsgemeinijchaft for financial support within Sonderforschungsbereich 93 “Photochemistry with Lasers”.
(17) (18) (19)
References and Notes (1) (2) (3) (4) (5)
(6) (7) (8) (9) (10) (1 1)
. ,
P. Debye, Trans. Electrochem. Soc., 82, 265 (1942). H. Rosmari and R. M. Noyes, J. Am. Chem. Soc., 80,2410 (1958). S.Aditya and J. E. Willard, J . Am. Chem. Soc., 79, 2680 (1957). R. Marshall and N. Davidson, J. Chem. Phys., 21, 2086 (1953). S. R. Logan, R. Bonineau, J. Joussot-Dubien, and P. Fomler de Violet, J . Chem. SOC.Fnraday Trans. 7, 71, 2148 (1975). A. M. Halpern and K. Weiss, J . Phys. Chem., 72, 3863 (1968). R. L. Strong and J. 1E. Willard, J. Am. Chem. Soc.,79, 2098 (1957). R. L. Strong, J . Phys. Chem., 86, 2423 (1962). R. L. Strong, J. Am. Chem. Soc., 87, 3563 (1965). J. Frank and E. Rabinowitch, Trans. Faraday Soc., 30, 120 (1934). E. Rabinovvitch and W. C. Wood, Trans. Faraday Soc., 32, 1381 (1936).
(20) (21) (22) (23) (24) (25) (26)
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J. Zlmmerman and R. M. Noyes, J . Chem. Phys., 18,658 (1949). F. W. Lamp and R. M. Noyes, J. Am. Chem. Soc., 78,2140 (1954). D. Booth and R. M. Noyes, J . Am. Chem. Soc., 82, 1668 (1960). L. Meadows and R. M. Noyes, J. Am. Chem. Soc., 82, 1872 (1960). K. Luther and J. Troe, Chem. Phys. Left., 24,85 (1974); C.Dupuy and H. van den Bergh, IbM., 57, 348 (1978): H. Hippler, K. Luther, M. Maler, J. Schroeder, and J. Troe In “Laser Induced Processes In Molecules”, Vd. 6, Springer-Vedag, West Berlin, 1979, p 286. H. Hlppler, V. Schubert, and J. Troe, to be submitted. H. HiDDler. K. Luther, and J. Troe, Ber. Bunsenges. Phys. Chem., 77, i i 0 4 (1973). D. W. Brazier and G. R. Freeman, Can. J . Chem., 47, 893 (1969). J. Jonas, D. Hasha, and S. 0 . Huang, J . Chem. Phys., 71, 3996 (1979). A. D. OsborneandG. Porter, Proc. R. SOC.London, Ser. A , 284, 32 (1965). R. M. Noyes, Z . Elektrochem., 84, 153 (1960). J. Schroeder, J. Troe, and U. Unterberg, Nachr. Akad. WIss. Bttlnaen. Math.-Phvs. KI., 2, 1 (1980). T. J. ehuang, G. W. Hoffmann, and K. E. Eisenthal, Chem. Phys. Left., 25, 201 (1974). C.A. Langhoff, K. GnBdig, and K. B. Elsenthal, Chem. Phys., 46, 117 (1980).
Infrared Spectroscopy of Some Chemisorbed Molecules on Tungsten Oxide-Silica A. J. van Roosmalen,” D. Koster, and J. C. Mol Unlverslty of Amsterdam, Institute of Chemlcd Technology, Plantage Muklergracht 30, 1018 TV Amsterdam, The Netherhnds (Received March 4, 1980)
Transparent plates of a tungsten oxidesilica metathesis catalyst were prepared by hydrolyzing a WC&-Si(OC2H& mixture. From the infrared spectra of adsorbed pyridine and ammonia, it is concluded that the chemisorption sites on the catalyst surface are of the Lewis type. Strong Bransted acid sites could not be observed the surface hydroxyls appear to resemble the weakly Bransted-acidichydroxyls on dry silica gel. Temperature-programmed reduction showed that -95% of the tungsten on the calcined catalyst is present as surface compounds and less than 5% as metathesis-inactive“free” oxide. The Lewis acidity is ascribed to these surface compounds, which are thought to be coordinativelyunsaturated species resulting from the dehydration of W06 octahedra shariing edges or planes with the silica lattice.
Introduction Tungsten oxide-silica is a well-known catalyst for the metathesis of alkene!J.1v2 In recent years, much work has been done to elucidate the nature of the active sites on this and other oxidic metathesis catalysts by using a large variety of techniques.”“ From the results of laser-Raman spectroscopy and temperature-programmed reduction of tungsten oxide-silica systems, Thomas et ah6 concluded that the precursor for the active site in metathesis is a surface compound and not the “free” oxide as such. However, no details of this surface compound could be given. More knowledge about the surface structure of tungsten oxide-silica might be obtained from infrared spectroscopy, being a straightforward, nondestructive method that can be applied easily in situ. Unfortunately, tungsten oxidesilica is opaque below 1350 cm-l (transmission less than 1%)owing to the strong lattice vibrations of the support. As a result, the interesting surface modes v(W-0) and v(W-0-Si) are obscured. Infrared studies will, therefore, yield only indirect injformation. In this paper, we present the results of a study on the interaction of some probe molecules, viz., pyridine, ammonia, and hexamethyldisilazane, with the surface of tungsten oxide--silica. One reason for this study was the
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0022-3654/80/2084-3075$01 .OO/O
suggestion made by Laverty et aL7that strong Brernsted acid sites play an important role in the heterogeneously catalyzed metathesis. Because the force constants in ammonia and pyridine are strongly dependent on protonation and complex formation: the infrared spectra of these molecules adsorbed on the catalyst surface can prove or disprove the presence of such Brernsted acid sites. A second reason was the observation that ammonia and amines can greatly enhance the catalytic activity of tungsten oxide-silica in the metathesis of propeneagJO Hexamethyldisilazane was used as it is known to react almost quantitatively with the hydroxyl groups on dry silica gel.l1 Conventional catalysts are, in general, not easy to study by infrared spectroscopy: scattering usually reduces the transparency of pressed catalyst wafers considerably. Moreover, most tungsten oxide-silica catalysts hold substantial amounts of “free” tungsten oxide. Here, we will show that it is possible to prepare a catalyst with an excellent transparency and a low “free” oxide content. Temperature-programmed reduction12 was used to compare this catalyst with tungsten oxide-silica catalysts from other preparations in order to quantify the bonding strength between the tungsten and the support and to calculate the amount of “free” tungsten oxide. All catalysta 0 1980 American Chemical Society
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The Journal of Physical Chemistry, Vol. 84, No. 23, 1980
van Roosmalen et al.
were active for the metathesis of propene.13 Experimental Section Materials. Tungsten oxide-silica catalysts were prepared as follows: WSI (Impregnation, 4.2 wt % WO,). Silica gel (24 g) (Davison Grace 62, 180-220 pm, 364 m2 g-l, calcined for 5 h at 825 K in a stream of dry air) was soaked in a filtrated solution of 2.0 g of ammonium metatungstate (Koch-Light Laboratories, 99.9%) in 65 mL of doubly distilled water. The water was removed in a rotating film evaporator kept at lo4 P a and 325 K for 2 h and at 365 K for 0.5 h. WSR (Reaction,5.7 wt % WO,).A solution of 3.4 g of tungsten hexachloride (Merck, freshly sublimed under vacuum) in 250 mL of dry, deoxygenated benzene was added to a suspension of 24 g of silica (same as above, activated for 24 h at 825 K in dry nitrogen) in 250 mL of benzene.14 After 18-h reaction at room temperature, the supernatant was decanted. The remaining dark-brown solid was washed repeatedly with fresh benzene and the residual benzene was removed by evacuation. The tungsten chloride-silica catalyst was hydrolyzed in saturated water vapor. WSC (Cogelation, 3.7 wt % W 0 3 ) . Tungsten hexachloride (1.2 g, same as above) was dissolved in 40 mL of tetraethoxysilane (Merck, 97%). The dark-red solution was stirred in air until it became yellow. This solution, probably containing partly hydrolyzed (triethoxy)siloxytungsten, was filtrated and diluted with 60 mL of absolute methanol. Then, 60 mL of concentrated hydrochloric acid was added under vigorous stirring. Within 1h the clear, colorless mixture was completely gelated. The gel was washed and autoclaved according to Peril6 The autoclaved catdyst was milled in an agate mortar. In the experiments, the 150-220-pm sieve fraction was used. WSCT (Cogelation, Thin plates, 4.3 wt % WOJ. This catalyst was prepared as described for WSC, with the exception that after the addition of hydrochloric acid the solution was poured out on mercury to obtain samples suitable for infrared spectroscopy.15 The surface area of the calcined catalysts was -600 m2 g-’ for the cogelation catalysts and -360 m2 g-l for the Davison-silica based systems (BET nitrogen). WSI, WSR, and WSC were calcined for 4 h at 825 K in a stream of dry air, These calcined samples were used in the temperature-programmed reduction measurements. Pyridine (BDH, 99.5%) and hexamethyldisilazane (Aldrich-Europe, 98%) were degassed by repeated freezing and evacuation. Anhydrous ammonia (Baker, 99.99%) was used without purification. Equipment and Procedures. Details about the apparatus for infrared spectroscopy are given elsewhere.le Ammonia and pyridine desorbing from the walls of the infrared cell disturbed the measurements if the system was closed after evacuation. Therefore, in the present study no helium could be admitted after evacuation to cool the sample. Consequently, the temperature of the irradiated sample rose under vacuum to -360 K. The spectrometer was flushed with nitrogen and calibrated by using the 1601.4-, 2850.7-, and 3027.1-cm-l bands of polystyrene. The scanning rate was 12.5 cm-l m i d , the slit width -2 cm-l, and the estimated accuracy f l cm-l. The WSCT samples were calcined for 2 h at 875 K in 2 X lo4 Pa of oxygen, followed by 1-h evacuation at the same temperature. The equipment for temperature-programmed reduction was the same as that used by Thomas et al! Thus, 60-70 mg of catalyst was activated in air at 775 for 1h, cooled under vacuum to 475 K, and reduced in a 2:l hydrogen-
I
I
I
I
I
I
I
500
700
900
1100
I
T. K
Figure 1. Temperature-programmed reduction patterns of tungsten oxide-silica catalysts.
nitrogen mixture (Hoek-Loos) with a flow rate of -3 X m3s-l and a heating rate of 4.8 K rnin-l. The hydrogen consumption was measured with a thermal conductivity cell. The reproducibility of the reduction pattern was within 20 K. Results Temperature-Programmed Reduction. Figure 1shows the reduction patterns of the WSI, WSR, and WSC catalysts. Under similar conditions, pure silica gel had a negligible hydrogen consumption. Assuming complete reduction, the tungsten content of the catalysts can be calculated from the areas below the reduction peaks. This amounts to 4.6,6.3, and 3.8 wt % W03for WSI, WSR, and WSC, respectively, which agrees well with the values from elemental analysis (see the Experimental Section). The reduction patterns of WSI and WSR resemble the patterns of impregnation-type tungsten oxide-silica catalysts reported in the literature.6 The bands at 680-780 K have been ascribed to metathesis-inactive oxides related to tungsten trioxide and 12-silicotungstic acid; a surface compound, probably the precursor of the active site for metathesis, was thought to cause the broad band at 865 Ka6 Because WSR appears to contain large amounts of inactive “free” oxides, no attempts were made to prepare samples for infrared spectroscopy by reaction, Le., from tungsten chloride and silica gel thin ~ 1 a t e s . l ~ Figure 1 shows that the reduction pattern of WSC is distinctly different from that of the other two catalysts. The fraction inactive oxide (band around 730 K) is only 5%, and the “surface-compound” band lies at much higher temperatures, viz., 1050 K. Nevertheless, this catalyst is highly active for the metathesis of propene.13 From this, we conclude that the results obtained from the infrared spectroscopy of WSCT can be extrapolated safely to the other catalyst systems. Infrared Spectroscopy. The spectrum of calcined WSCT is much the same as that of calcined silica aerogel?6 No new bands are visible in the 1350-8000-~m-~ region. The only difference appears to be a somewhat broader u(OH), which resembles Peri’s findings with silica-alumina aeroge1.l’ Adsorption of Hexamethyldisilazane. Within the detection limits, hexamethyldisilazane reacts with all hydroxyl groups on the catalyst surface in the same way as with silica This means that silylation cannot be used to discriminate between Si-OH and W-OH groups. Adsorption of Pyridine. A sample of calcined WSCT was exposed for 1 h to 2 X lo3 P a of pyridine at 425 K, followed by 1-h evacuation at the same temperature.
The Journal of Physlcal Chemistry, Vol. 84, No. 23, 1980 3077
Chemisorbed Molecules on Tungsten Oxide-Silica 1.5
1.0 0
C
em
al 0
VI
P
m
1.0
0,
n m
0.5
1400
1600 frequency. cm-'
F W e 3. Spectra h the N-H deformatbn region of tungsten Oxkle-siliCa catalyst before (A) and after (B) reaction with ammonia, and after readmitting gaseous ammonia to B (C); (dotted line) difference between B and C, dlsplaced 1.0 absorbance units upward. Absorbance scale
0.5 frequency cm-' I
Flgure 2. Spectra of tungsten oxide-silica catalyst before (A) and after (B) reaction with pyridine, and after readmitting gaseous pyridlne to B (C). Absorbance scale refers to A.
refers to A. 1
.
1
I
I
TABLE I : Ring Vibrations of Pyridine Adsorbed on Tungsten Oxide-Silicaa b
mode Visb v~9a V8a 1490 1595 physisorbed ( C - .B) 1446 1490 1615 1453 chemisorbed (B) a For B and C, see Figure 2. Weak bands in Figure 2 near 1600 and 1650 c1y1-l are caused by atmospheric "he 8b mode was water vapor. Frequencies in cm-' too weak to be observed.
.
Figure 2B shows the gpectrum in the 1350-1750-~m-~ region after the treatment, together with the spectrum before 0 ~ exposition (A), and after readmitting ~ ~ Pa 1of pyridine to the treated sample (C). The difference between C and B of Figure 2 is obviously due to physisorbed pyridine. The new bands are closely similar to those reported for weakly adsorbed pyridine on silica and silica-based systems.lSz1 Table I, first line, gives the peak positions of the physisorbed pyridine in Figure 2C with the usual mode assignments for pyridine ring vibrations.8 Compared with the spectrum of pure pyridine,22*23 all bands appear to be shifted -9 cm-l toward higher frequencies, apparently the result of hydrogen bonding t o the surface hydroxyls. The spectrum of the more strongly adsorbed pyridine (Figure 2B) resembles Basila's spectra of pyridine chemisorbed on the Lewis acid sites on silica-alumina.lg Table I, second line, gives the frequencies of the bands in Figure 2B. It has been noted that the 19b and 8a modes of the pyridine molecule are sensitive to complex formation with Lewis acids and that thiese bands are shifted toward higher frequencies when the strength of the Lewis acid is increased.8Js Compared with data from other systems, it appears that the Lewis acid sites on the tungsten oxidesilica surface are stronger than those on chromium-silicaN and titania,24weaker than those on and of the same strength as aluminum trichloride22and the sites on sili~a-alurnina~~J~ and highly dehydroxylated Cabosil.21 A difference between pyridine chemisorbed on silicaalumina and on tungsten oxide-silica is the absence in our spectra of a distinct barid near 1540 cm-l (PyH+:vlg&8 It seems, therefore, that the calcined tungsten oxide-silica catalyst does not hold strong Brernsted-acidicsurface sites. Adsorption of Ammonia. The above-mentioned result of the pyridine adsorption, viz., the apparent absence of strong Brernsted acid sites on WSCT, is not conclusive. The pyridinium-ion band near 1540 cm-' is usually not
3100
3300 frequency, cm-'
3500
Figure 4. Spectra in the N-H stretching region. For assignment of curves,
see Figure 3. Transmission scale refers to A.
TABLE I1 : N-H Vibrations of Ammonia Adsorbed on Tungsten Oxide- Silica' mode 6 vs Vas amine groups (H) 1553 3446 3533 chemisorbed (C- B ) b 1615' 3298 3398 For B and C, see Figures 3 and 4. Frequencies in For assignment of 3186-cm-' band, see text. em-'. Asymmetric deformation (v4 ).
very intense and can be obscured by atmospheric absorptions and pyridine overtones. More definite proof can be derived from ammonia adsorption experiments: the asymmetric deformation of the ammonium ion gives rise to a very intense infrared band that lies in a rather "clean" spectral r e g i ~ n . ~ ~ J ' To this purpose, tungsten oxide-silica was exposed for 1 h to 5 X lo4 Pa of anhydrous ammonia a t 525 K and evacuated for 1 h at the same temperature. Figure 3, A and B, shows the infrared absorption in the N-H deformation region before and after the reaction with ammonia. Figure 4, A and B, shows the corresponding spectra in the N-H stretching region. Because the v(N-H) bands are rather weak, the latter spectra were recorded in the transmission mode. Place and shape of the observed bands correspond with the absorptions of the Si-NH2 groups on ammonia-treated s i l i ~ a ,so~ that ~ * ~an~ assignment could be made (Table 11, first line). Surprisingly, no chemisorbed ammonia is observed in Figures 3B and 4B. Because this might be caused by the thorough degassing of the ammonia-treated catalyst, -50 Pa of ammonia was readmitted to the sample. Infrared
3078 The Journal of Physical Chemistty, Vol. 84, No. 23, 1980
TABLE 111: Surface Hydroxylation of Tungsten Oxide-Silicaa IW031,wt ?h 0 4 11 50 100 [OH], pmol m-' 2.62 2.89 3.47 3.61 3.63 a After heating at 825 K under vacuum. Calculated from the weight increase on silylation with hexamethyldisilazane. 3 3
spectra after this treatment are given in Figures 3C and 4C. The dotted line in Figure 3 was obtained by graphically subtracting curve B from curve c. From Figure 3C two conclusions can be derived. In the first place, strong Brernsted acid sites are absent indeed, as is indicated by the absence of a band between 1400 and 1500 cm-l (NH4':v4).26 In accordance with this, no strong band can be seen around 3280 cm-l in Figure 4C (NH4+:~3).26 Secondly, because the maximum in the dotted line in Figure 3 lies at 1615 cm-l, and ammonia physically adsorbed on silica-based systems absorbs at 1630 cm-l, the bands appearing in Figures 3C and 4C are to be assigned to ammonia chemisorbed on Lewis acid sites (Table 11, second line). The v, and v, frequencies of the chemisorbed ammonia are higher than the corresponding bands of known Lewis-acid complexes, such as boron trifluoride-ammonia,29 This, however, does not appear to be exceptional: similar band positions have been observed for ammonia chemisorbed on nickel-~ilica,~~ zirc~nia-silica,~~ and silica-alumina aeroge1,17 and for some ammine complexes of rut h e n i ~ m The . ~ ~ asymmetry of the uaq and v, bands probably reflects a heterogeneity of the tungsten oxide-silica surface. Presumably, there exist adsorption sites on the catalyst surface with an acidity higher than is indicated by the band maxima from Table 11, second line.
Discussion Chemisorption of pyridine and ammonia on tungsten oxide-silica degassed at 875 K revealed the presence of Lewis acid sites that resemble the Lewis acid sites on silica-alumina. We can describe silica-alumina as a silica gel where part of the silicon atoms is replaced by aluminium, and assign the Lewis acidity of silica-alumina to tricoordinated aluminum ions on tetrahedron positions.32 Presumably, tungsten oxide-silica has a similar surface structure, which means that the Lewis acid sites can be envisaged as tungsten ions in a more or less tetrahedral silica gel matrix. However, tungsten is hexavalent, and in order to maintain electrical neutrality the tungsten ions should hold negatively charged species in their coordination spheres. On a freshly prepared, fully hydrated catalyst these species are obviously hydroxyl groups. But also on the calcined catalyst hydroxyl groups are retained; it has been observed that tungsten oxide-silica holds more hydroxyls than silica gel after calcination (see Table 111). The low tungsten content of WSCT, viz., -2 atom % if all tungsten is accumulated in the surface layer, probably prevents direct observation of the W-OH groups, although they might be the cause of the broader hydroxyl stretching vibration on WSCT relative to silica gel. Thomas et al.6 concluded from their laser-Raman spectra that the surface compound on nonactivated tungsten oxide-silica consists of distorted WO, octahedra. With increasing tungsten loadings, the maximum in the temperature-programmed reduction pattern assigned to this surface compound shifted to lower temperatures, whereas the position of the corresponding Raman band remained almost unchanged., Hence, it appears that the geometry of the tungsten species is not much affected by a change in the strength of the interaction with the support, so that
van Roosmalen et ai.
in all cases the structure of the surface complexes can be derived from fully hydrated octahedral trioxide, W (OH),. The W(OH)6unit can be attached to silica gel in three ways, viz., sharing a corner (I),an edge (ID, or a plane (111).
I
I1
I11
For the sake of clarity, the =Si-0bridges have been omitted. A more than threefold bonding is not probable, owing to the rigid tetrahedral structure of silica gel and silicates.* 35 Neighboring I units might react with each other under the formation of oxo bridges. This would result in a silicon-bound oligomeric tungsten oxide with a structure closely related to that of 12-silicotungsticacid.% Therefore, we assign the 680-780 K bands in the temperature-programmed reduction pattern of tungsten oxide-silica catalysts to compounds originating from monomers and oligomers of I. WSR was prepared by hydrolyzing a tungsten chloride-silica catalyst. It has been demonstrated that on the latter catalyst most of the tungsten atoms are present as (=Si-O-)2WC14,14 so that the hydrolyzed compound in WSR will resemble 11. The 865 K reduction peak, then, can be assigned to complexes having this structure. The mechanism for the formation of tungsten oxide-silica catalyst from ammonium metatungstate and silica gel (WSI) has not yet been resolved. From the similarity between the high-temperature reduction peaks of WSR and WSI, however, it follows that also on WSI the surface compound will have a II-like structure. WSC was prepared by reacting tungsten hexachloride with tetraethoxysilane and hydrolyzing the reaction product. The surface compound on this catalyst reduces at 1050 K (Figure 1). Hence, the tungsten appears to be even more encased than in WSR, so that its structure probably resembles 111. Heating tungsten oxide-silica will remove hydroxyl groups, with the formation of water: -W HO'
OH
IV
/ \OH
-t H20
V
Reaction with neighboring silanol groups is less likely, because this would give rise to the formation of highly strained oxo bridges. Perhaps for steric reasons, complete dehydration of I1 is not achieved at the usual calcination temperatures (825-875 K). We deduce this from the high hydroxyl concentration on tungsten oxide-silica (Table 111). As a result of the dehydration, the coordination number of the tungsten ion decreases from six to five. Probably, the coordinately unsaturated tungsten compound V is the source of the Lewis acidity of the WSCT tungsten oxide-silica catalyst. There is evidence that IV and V are associated with the active sites in the metathesis of alkenes. This will be discussed in a following paper.13 In Figure 4C, a weak but distinct band is visible at 3186 cm-l. This band might be due to an overtone of chemisorbed ammonia, viz., 2v4,97 but the large difference between 3186 and 2 X 1615 cm-l makes this explanation unlikely, The proposed structure of the Lewis acid site (V) makes it more reasonable to ascribe the 3186-cm-l band
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J. Phys. Chem. 1980, 84, 3079-3084
to a second-order interaction of the chemisorbed ammonia with the adjacent oxygen ion. A similar band has been observed during adsorption of ammonia on alumina.3s From the ammonia-adsorption experiments we concluded that the surface of tungsten oxide-silica is heterogeneous. This can account for the broadness of the high-temperature reduction peaks and be another reason for the absence of distinct W-OH bands in our infrared spectra. ObviousXy, a broad spectrum of interaction strengths is to be found on all catalysts. The structures IV and V represent, therefore, only an average stoichiometry. Acknowledgment. We thank M. C. Mittelmeijer-Hazeleger for the temiperature-programmed reduction measurements, and the Chemical Analysis Department of the Twente Teclhnical University for analyzing our catalyst samples.
References and Motes (1) (2) (3) (4) (5) (6)
(7) (8) (9) (10) (11) (12) (13)
Banks, R. L. Fortschr. Chem. Forsch. 1972, 25,39. Mol, J. C.; Moulljn, J. A. Adv. Catal. 1975 24, 131. Olsthoorn, A. A.; Woelhouwer, C. J . Catal. 1976, 4 4 , 197. Kerkhof, F . P. J. M.; Thomas, R.; Moulljn, J. A. Recl. Trav. Chim. Pays-Bas 1977, 86, M 121. Lombardo, E. A.; Lo Jacono, M.; Hall, W. K. J. Cafal. 1976, 51,243. Thomas, FL; Moulijn, J. A.; Medema, J., De Beer, V. H. J. J . Mol. Catal. 1980, 8 , 161. Laverty, D. T.; Roaney, J. J.; Stewart, A. J. Catal. 1976, 45,110. Knozlnger, H. Adv. Catal. 1976, 25, 184. Gangwal, S. K.; Wills, G. B. J. Catal. 1976, 52,539. Takahashi, T. Nippon Kagaku Kalshl1878, 3 , 418. Hertl, W.; Hair, M. L. J. Phys. Chem. 1971, 75,2181. Jenklns, J. W.; Mcfrlicol, B. D.; Robertson, S. D. CHEMTECH 1977, 316. Van Roosmalen, A. J.; Mol, J. C., to be submitted for publlcation.
(14) Van Roosmalen, A. J.; Polder, K.; Mol, J. C. J. Mol. Catal. 1980, 8, 187. (15) Perl, J. B. J . Phys. Chem. 1966, 70,2937. (16) Van Roosmalen, A. J.; Mol, J. C. J . Phys. Chem. 1078, 62.2748. 1171 Perl. J. B. J. Phvs. Chem. 1966. 70.3168. i i 8 j Pariy:~. B: J. 1963, 2, $71. ’ (19) Basila, M. R.; Kantner, T. R. J . Phys. Chem. 1966, 70, 1681. (20) . . Zecchina. A.: Garrone. E.: Ghiottl, G.; Coluccia. S. J . Phys. Chem. 1975, 79,972. (21) Morrow, B. A,; Cody, I.A. J . Phys. Chem. 1978, 80, 1995. (22) Terenln, A.; Fllimonov, V.; Bystrov, D. Z . Elektrochem. 1958, 62, 1080. (23) Reitsma, H. J. W.D. Thesis, University of Amsterdam, The Nethetlands, 1975. (24) Primet, M.; Pichat, P.; Mathieu, M. V. J. Phys. Chem. 1971, 75,1221. (25) Klvlat, F. E.; Petrakls, L. J . Phys. Chem. 1973, 77,1232. (26) UyttemoeVen, 4J. B.; christner, L. G.; Hall, W. K. J. Phys. Chem. 1985, 69, 2117. (27) Cant, N. W.; Little, L. H. Can. J. Chem. 1968, 46, 1373. (28) Morrow, B. A.; Cody, I. A.; Lee, L. S. M. J. Phys. Chem. 1976, 80, 2761. (29) Goubeau, J.; Mitschelen, H. Z . Phys. Chem. (Frankfurt am Main) 1958, 14, 61. (30) Peri, J. B. Discuss. Faraday Soc. 1986, 41, 121. (31) Allen, A. D.; Senoff, C. V. Can. J. Chem. 1967, 45, 1337. (32) Hair, M. L. “Infrared Spectroscopy In Surface Chemistry”; Marcel Dekker: New York, 1967; p 169 ff. (33) Moulijn, J. A,; \/an Cranenburgh, H.; Van Roosmalen, A. J.; Kerkhof, F. P. J. M. Lecture presented on the 1978 Catalysis Symposium of the Royal Dutch Chemlcal Society. (34) Doremus, R. ti. “Glass Science”; Wlley: New York, 1973; p 24. (35) Tungsten oxide might share more than three oxygen Ions with the support In the case of tungsten oxkle-alumlna catalyst, owing to the fact that y-alumina has a partly octahedral structure. This could account for the very hlgh reduction temperature of the surface compound on this catalyst, 1250 K (Kerkhof, F. P. J. M. Ph.D. Thesis, University of Amsterdam, The Netherlands, 1979). (36) Kepert, D. t. “The Early Transkim Metals”; Academic Press: London, 1972; Chapter 4. (37) Cannon, C. G. SDectrochim. Acta 1958. 10. 425. (38) Pichat, P.; Mathleu, M. V.; Imellk, B. J. Chim. Phys. Phys.-Chim. Biol. 1969, 66,845.
C&/.
Surface Reactions of Oxygen Ions. 5. Oxidation of Alkanes and Alkenes by 0,- on MgO Masakaru Iwamoto and Jack H. Lunsford* Depaifment of Chemkfty, Texas A & M University, College Statlon, Texas 77843 (Received: Aprll 10, 1980; In Final Form: July 28, 198Q)
Stoichiometric reactions between superoxide ions on MgO and simple hydrocarbons were observed at 175 “C, althclugh the 0, ions were much less reactive than either 0-or 0,-ions. Several types of oxygen-containing products were formed, as well as other hydrocarbons and COD The reaction of propylene with 0, gave no gaseous product at 175 “C; however, at elevated temperatures acetaldehyde and methanol were obtained. acetone was detected tat 175 “Cin addition to acetaldehyde and Following the reaction of propane with 02-, methanol which were observed at higher temperatures. With 1-butene as the reactant, 2-butanol was formed together with methanol, acetaldehyde, and acrolein above 300 “C. Infrared spectra of surface intermediates indicate that the reaction of 02-with propylene at 175 “C resulted in the simultaneous formation of formate and acetate ions which are consecutively converted to carbonate ions at elevated temperatures. It is proposed that hydrogen atom abstraction is the initial step in the reaction of 0, with simple hydrocarbons. The resulting radicals react with lattice oxygen ions forming carboxylate ions or with H02- ions forming alkoxy or epoxide intermediates.
Introduction The superoxide ion, 02-, on a variety of metal oxides has been studied by EPR techniques;l however, its importance in catalytic reactions has not been resolved. There exists evidence that 02-is the source of oxygen in epoxidation of ethylene over supported silver,2 but this is a difficult system to explore by using EPR. On ZnO the ion does not react at room temperature with hydrogen, carbon mon0022-3654/80/2084-3079$01 .OO/O
oxide, or ethylene and only reacts slowly with propylene.s Kugler and Gryder4proposed that n-allyl reacted with 0,to form 2,3-epoxypropanal (glycidaldehyde), which was an intermediate in the formation of acrolein on the surface. Reactions also have been reported between 02and butene on Sn02,Ti02, and Zn0,6as well as between 02-and butadiene on supported Moo3, for which the products were maleic anhydride and COD6 In addition, it has been 0 1980 Amerlcan Chemlcal Saclety
