J . Phys. Chem. 1985,89, 2411-2481
2411
width) one gets the peculiar situation that for a fixed amount of energy dose the degree of polarization may increase with increasing power. We finally stress that the state of polarization is important with respect to several applications, most of all in the technique of frequency-modulated polarization spectroscopf but in other fields, e.g., Stark spectro~copy,~ as well.
is low in the center and high in the wings. We further found that the degree of polarization is little affected by the photoreversibility of the hole burning reaction, contrary to the width of the hole which is strongly affected. The bleaching phenomena in the width and degree of polarization of a photochemical hole show a good agreement between theory and experiment in case photoreversibility is taken into account. Under the restrictive condition of a quasi-two-level system we investigated the influence of power on the degree of polarization in the presence of photochemical bleaching. Since strong power changes simultaneously the bleaching time (via an increase of the
Acknowledgment. We thank D. Haarer for helpful discussions. A grant from the VW-Stiftung is greatfully acknowledged. Registry No. DAQ, 81-64-1; MeOH, 67-56-1; EtOH, 64-17-5.
Spectroscopic Characterization of Molybdenum Oxalate in Solution and on Alumina K. Y. S . Ng, X. Zhou, and E. Gulari* Department of Chemical Engineering, The University of Michigan, Ann Arbor, Michigan 481 09 (Received: October 30, 1984)
The change in structure of molybdenum oxalate as a function of pH in solution and on alumina support has been investigated by Raman and FTIR spectroscopy. Both Raman and FTIR spectroscopyshow that the molybdenum oxalate complex changes from a dimeric bidentate structure to a monomeric unidentate structure in solution as pH increases. Molybdenum oxalate complexes are stable when supported on alumina, as evidenced by the fact that they retain their characteristic Mo=O vibration frequency in solution. This is in sharp contrast to those samples prepared from an ammonium heptamolybdate solution. This study demonstrates the advantage of using both Raman and FTIR spectroscopy to study supported catalyst systems.
Introduction A number of studies on the spectroscopic characterization of molybdena on alumina systems have been made in the past few years. Contradictory results were reported by different researchers probably due to different preparative procedures.' Such discrepancies have lead to different surface models proposed for the molybdena-alumina system.2 Wang and Hall2 pointed out that the usual way to prepare the molybdena catalysts by incipient wetness has caused most of these complications. This is because a change of p H is unavoidable in the small volume of solution impregnated into the pores of support, and the equilibrium of the molybdate solution is pH dependents3 Consequently, Wang and Hall2 proposed a static equilibrium adsorption method to prepare a better defined catalyst. By using the equilibrium adsorption method, they showed that for pH 1 8 , octahedrally coordinated molybdate is the dominant species that binds to the alumina support. A small fraction of tetrahedrally coordinated molybdate species is also found in these catalysts. Cheng and Schrader4 investigated the effects of pH of the impregnating solution, but they did not observe a strong pH effect. All of their dried spectra indicate deposition of the Mo70246ion, even for the impregnation at pH 9 (which has mainly simple M a d 2 -ions in solution). They explained the deposition of M070246-ion as being due to the equilibrium shift when water was driven out during the drying process. Houalla et al.? on the other hand, observed a dependence of ESCA peak intensity and hydrodesulfurization activity on the pH of the impregnating solution. The complications in the molybdena-alumina system were not only caused by the method of incipient wetness. Another underlying problem is the stability and solubility of the ammonium heptamolybdate solution (AHM), the most commonly used precursor. The relatively low solubility of the AHM and its sensitivity (1) Massoth, F. E. Adv. Caul. 1978, 27, 265. (2) Wang, L.; Hall, W. K. J . Cutal. 1980, 66, 252. (3) Ng, K. Y. S.; Gulari, E. Polyhedron 1984, 8, 1001. (4) Cheng, C. P.;Schrader, G. L. J. Catal. 1979, 60,276.
(5) Houalla, M.; Kibby, C. L.; Petrakis, L.; Hercules, D. M. J . Catal.
0022-3654185 I2089-2417SOl . SOJO ., ,
~
~
Experimental Methods Material and Catalyst Preparation. Molybdenum oxalate solution was prepared from molybdenum oxalate salts (Climax Molybdenum Co.) by using doubly distilled water. Solutions of higher pHs were obtained by adding ammonium hydroxide. Catalysts were prepared by the method of incipient wetness. Support material was y-alumina (BDH Chemicals, Ltd.) with a surface area of 100 m2/g. Samples were air-dried at 110 OC for 12 h. (6) Tsigdinos, G. A.; Chen, H. Y.; Streusand, B. J. Ind. Eng. Chem. Prod. Res. Deu. 1981, 20, 619. (7) Streusand, B. J.; Husby-Coupland, K. J. Report L-287-81/83; Climax Molvbdenum Co.: Ann Arbor. MI. ~,19R4 _ _ _ .. (g) Beltran, A.; Caturla,Fl;Cervilla, A,; Beltran, J. J. Inorg. Nucl. Chem. 1981, 43, 3277.
1983, 83, 50. I
to change in hydrogen and hydroxyl ion concentration undoubtedly account for the inhomogeneity of most of the impregnated catalysts. Tsigdinos et aL6 studied the stability and adsorption properties of molybdate solutions as a function of pH, concentration, time, temperature, and the presence of oxalate ligands. They found that the molybdenum oxalate solution has a unique adsorption behavior in that the amount adsorbed on alumina is independent of concentration. More recently, Streusand et aL7 studied the adsorption behavior of molybdenum oxalate in a bimetallic solution and found similar results. They further suggested that the adsorption mechanism of molybdenum oxalate onto alumina is different from that based on A H M solutions. The purpose of this work was to characterize the molybdenum oxalate in solution and on alumina by Raman and FTIR spectroscopy. The equilibrium adsorption technique has the merit of producing more disperse and homogeneous catalysts, but it may have drawbacks in dealing with higher metal loadings and bimetallic solutions. Molybdenum oxalate is well-known to form stable complex species in aqueous solution.* In this study we want to investigate whether a more stable molybdenum oxalate solution of higher buffer capacity will lead to a more homogeneous catalyst even under incipient wetness conditions.
-8
Q 1985 American Chemical Society
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Ng et al.
The Journal of Physical Chemistry, Vol. 89, No. 12, 1985
TABLE I: IR Peak Frwuencies of the Molybdenum Oxalate Solution as a Function of pH“
O s
wavenumber, cm-’
solution PH
800-1 100
1.1 1.3 2.5 4.0 6.0
916 (m), 951 (m) 916 (m), 951 (m) 910 (m), 947 (m) 858 (m), 907 (m). 941 (sh) 856 (m), 905 (m)
IIOI-1400 1256 (w), 1400 (m) 1258 (w), 1400 (m) 1267 (w) 1285 (w) 1287 (w), 1306 (sh)
1401-1 700 1694 (sh) 1437 (sh), 1694 (sh) 1406 (m), 1437 (sh), 1678 (sh), 1690 (s) 1438 (s), 1686 (s) 1443 (s), 1570 (m), 1684 (s)
1701-2000 1715 (s) 1715 (s) 1712 (s)
= strong, m = medium, w = weak, sh = shoulder.
Raman Spectroscopy. Raman spectra were obtained with a Spex Triplemate spectrograph coupled to a Tracor Northern 1024 large area intensified diode array detector. The 488-nm line of a Lexel Model No. 95 Ar+ laser was used as the excitation source. A grating monochromator was used to reject any spurious lines and background from the laser before the radiation entered the spectrometer. Typical power at the pellet surface was on the order of 50 mW. All the spectra were taken with 2-cm-’ resolution. Since the spectrum is totally digital and no gratings were moved, spectral subtraction and signal averaging were used to improve the signal-to-noise ratio. The catalyst samples were pressed into the form of pellets and rotated at 2000 rpm in a Spex Model 1445A sample rotator to avoid local heating. ZR Spectroscopy. All the IR spectra were taken with a Digilab FTS 20/C FTIR spectrometer with 4-cm-I resolution. In order to make the pellets transparent, the catalysts were thoroughly mixed and diluted with KBr. Using the digital subtraction capability of the spectrometer, we subtracted the support spectrum from the sample spectra to obtain only the spectrum of molybdenum oxalate on the surface. Solution spectra were obtained by using an ATR cell and substracting off the water spectrum. Results and Discussion Molybdenum Oxalate Solution. Figure 1 shows a representative number of Raman spectra of 6 wt % molybdenum oxalate solution at pHs 2.0, 3.5, 4.0, and 6.0. A general trend for the M e 0 stretch region (900-950 cm-’) can be observed as a funqtion of pH. The main peak frequency shifts from 938 cm-I at pH 2.0 to 901 cm-’ at pH 6.0. The decrease in Mo=O stretch frequency can be attributed to a decrease in the degree of polymerization4 or a difference in Mo coordination for different molybdenum oxalate complexes. Figure 2 shows the main peak frequency of M e 0 stretches as a function of pH. At least two different molybdenum oxalate complexes can be deduced from the curve if one uses the Mo=O characteristic frequency as a criterion. Beltran et ale8studied the formation equilibrium for the molybdenum oxalate complexes by spectrophotometry and saline cryoscopy. They concluded that there are one monomeric molybdodioxalate and two dimeric molybdooxalates in aqueous solution, and their relative concentrations depend only on the solution’s pH. The formulas they deduced are as follows: ( M o O ~ ( O H ) ~ ( C ~ O ~ )6.50 ~ ) ~>- pH
> 4.25 ( M o ~ O ~ ( O H ) ~ ( C ~ O ~4.25 ) ~ )>~ pH - > 2.75 ( M o ~ O ~ ( O H ) ( C , O , ) ~ ~ ~2.75 - > pH > 1.25 Our results agree with their assignment that in the more acidic medium there are indications of polymeric species. Further evidence can be seen in the lower wavenumber region of the Raman spectra (Figure 1). Upon addition of ammonium hydroxide, a peak at 354 cm-I begins to develop (pH 3.5). This peak can be assigned to the terminal Mo=O bending vibration which is an indication of breaking up of the dimeric molybdooxalates into monomeric molybdodioxalate and possibly “simple molybdate”. The carbon-oxygen stretches in the oxalate group can also be observed in the 1200-1 800-cm-’ region by Raman spectroscopy, even though the intensity of these peaks is much less than the corresponding Mo=O stretches. Raman bands are observed in the 1700-, 1400-, and 1250-cm-] regions which correspond to symmetric and asymmetric C = O vibration, v ( C - 0 ) and v(CC) vibration, and v(C-0) and a(OC0) vibration, respectively.
-
m
m
m
o
m
p H 6.0
P H 4.0
p H 3.5
pH
2.0
WAVENUMBER c m - 1 Figure 1. Raman spectra of molybdenum oxalate solution as a function of pH. As pH increases, the Mo=O vibration frequency shifts downward, indicating a decrease in the degree of polymerization.
M d Frequency
8 0 0 0 0
910
0
8 0
0
4
5
890
1
2
3
6
PH Figure 2. Raman peak frequencies of molybdenum oxalate in solution (0)and on alumina (0) as a function of pH. Note that they are in excellent agreement with each other. This is in sharp contrast to impregnation catalysts made from A H M solution.
As pH increases, it is observed that the C = O vibrations shift down and the C-0 vibrations shift upward. Since IR spectroscopy should be a more sensitive probe of these carbop-oxygen stretches, we would like to report our IR result first and then continue the discussion. Solution spectra of molybdenum oxalate obtained by IR spectroscopy as a function of pH are shown in Figure 3 and Table I. Once again, absorption peaks are observed in the 1700-, 1400-, 1250-cm-’ regions for all pHs. The 1700-cm-’ region can be assigned to the C = O vibration. For pH 1.1 and pH 1.3 spectra, a splitting gives two absorption bands at 1715 and 1694 cm-I. As
Molybdenum Oxalate in Solution and on Alumina
The Journal of Physical Chemistry, Vol. 89, No. 12, 1985 2479
O,.
Q
0
0
?,'
4
2-
t2OH
Moo4
+
H2°
0
Figure 4. Proposed scheme of transition from a bidentate structure to a unidentate structure for molybdenum oxalate. "Simple molybdate" species are formed as a result of this transformation. m 0 .?
2000
1700
1400
1100
- 0 0 9
m m
800
WAVENUMBERS (cm -1)
Figure 3. FTIR absorbance spectra of molybdenum oxalate solution as a function of pH. The downward shift of the C=O vibrations (- 1700 cm-I) and the upward shift of the C-0 vibrations (-1400 and 1200 cm-I) indicate that the influence of the metal on the oxalate group is decreasing.
p H 2.0
v
f \
pH increases to 2.5, the structure of this band becomes complicated. It consists of a mixture of C=O stretches of different environments. Further addition of ammonium hydroxide suppresses the 1715- and 1694-cm-' peaks and leaves a rather uniform 1686-cm-' peak. The 1400-cm-' region provides another piece of evidence about the C-O stretch. At pH 1.1, there is only one 1400-cm-' peak; as pH increases, a shoulder at 1437 cm-' begins to develop. The transformation is completed at pH 4.0, leaving 1438 cm-' as the only absorption band. The 1280-cm-' region can be assigned to v(C-0) and b ( O C 0 ) ; the same trend of shifting to higher wavenumbers is also observed with pH. I R results on the oxalate groups are essentially the same as the Raman results discussed above. The fact that the C=O stretching frequency shifts downward and the (2-0 stretching frequencies shift upward indicates that the influence of the metal on the oxalate group is decreasing? Tsigdinos et aL6 suggested that as more ammonium hydroxide is added, the hydrogen ion is simply substituted by the ammonium ion, and monoammonium and diammonium salts of molybdenum oxalate result. According to this scheme, the molybdenum and oxalate groups do not undergo any structural transformations. However, the changes in molybdenum-oxygen stretch and oxalate group frequencies observed in this study as a function of pH suggest a more radical structural change than just a simple ammonium ion substitution. These results agree with the proposed change from a bidentate structure to a unidentate structure for the molybdenum oxalate as pH increases.8 As shown in Figure 4, the oxalate group in the unidentate structure is less influenced by the Mo than in the bidentate structure. Further evidence of this transformation can be seen with additional bands at 1570 and 1306 cm-l at pH 6.0 which are due to the free -COz group.8 The 900-cm-' region bands of the IR spectra are due to Mo=O stretches. The decrease of a higher wavenumber band (951 cm-I) and the emergence of a new band a t a lower wavenumber (858 cm-I) as pH increases also parallel our Raman results. This again
is in agreement with the above proposed transformation. Molybdenum Oxalate on Alumina. Figure 5 shows the Raman spectra of MoOx/alumina in the region of 500-1000 cm-' as a function of pH. The region below 500 cm-' has a high background noise, probably due to fluorescence in the alumina support. The main peak frequency in the 900-1000-cm-' region is also plotted in Figure 2 in order to compare with peak frequencies of the molybdenum oxalate in solution. We see that the M-0 peak positions in solution are in excellent agreement with those on the alumina surface. This is very different from the catalysts made from A H M ~ o l u t i o n .A~ peak at 610 cm-I is also observed for pHs 2.0 and 3.5 which can be assigned to symmetric Mc-0-Mo stretches present in the dimeric structure. For the pH 3.5 spectra, a band at 709 cm-' begins to develop. This band increases in intensity so much that it becomes the most intense peak for pH 4.0 and up. However, this band is not observed if we use sodium hydroxide to adjust the pHs. When compared with other typical amine complexes,1° it is likely that the 709-cm-' band is due to the p(NH,) vibration of molybdenum oxalate ammine complexes. It is interesting to note that this band is observed neither in the solution spectra (Figure 1) nor in the solid spectra of those catalysts prepared from AHM solutions (Figure 6). The Raman spectra of MoOx/alumina catalysts for the 1000-2OOO-cm-' region have a poor signal-to-noise ratio, probably due to the poor scattering properties of the oxalate group. On the other hand, the IR spectra of the MoOx/alumina catalysts (Table 11) all show strong, typical bands in the 1700-, 1400-, and 1300-cm-' regions for the oxalate group as observed
(9) Nakamoto, K.; McCarthy, P. J. 'Spectroscopy and Structure of Metal Chelate Compounds"; Wiley: New York, 1968.
(10) Nakamoto, K. "Infrared Spectra of Inorganic and Coordination Compounds"; Wiley-Interscience: New York, 1970.
WAVENUMBER cm- 1 Figure 5. Raman spectra of molybdenum oxalate on alumina as a function of pH. The strong band at 709 cm-I is due to the p(NH,) vibration of molybdenum oxalate ammine complexes.
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The Journal of Physical Chemistry, Vol. 89, No. 12, 1985
I
-*
I
: I
I
pH
5.2
WAVENUMBER c m - 1 Figure 6. Raman spectra of A H M solution and Mo/alumina catalysts from AHM as a function of pH by the method of impregnation. The resulting Mo/alumina catalysts have the same spectra regardless of impregnating solution’s pH.
TABLE 11: IR Peak Frequencies of the Molybdenum Oxalate Supported on Alumina as a Function of pH’ wavenumber, cm-’
preparation PH
1200-1 400
1401-1600
1.3 2.0 2.5 3.1 3.5 4.0 4.5 6.0
1296 (m), 1385 (m) 1298 (m), 1385 (m) 1298 (m),1385 (m) 1298 (m), 1385 (m) 1296 (m), 1385 (m) 1292 (m), 1387 (m) 1294 (m), 1387 (m) 1296 (m),1385 (m)
1406 (s) 1404 (s) 1408 (s) 1406 (s) 1408 (s) 1408 (s) 1414 (s) 1410 (s)
1601-1800 1697 (s), 1697 (s), 1697 (s), 1697 (s), 1697 (s), 1695 (s), 1695 (s), 1697 (s),
1719 (s) 1719 (s) 1720 (s) 1719 (s) 1719 (s) 1718 (s) 1718 (s) 1718 (s)
“The 1385-cm-I peak is due to potassium bromide. s = strong, m = medium.
in the solution spectra (Figure 3). However, there are no significant shifts in the carbon-oxygen vibrations as a function of pH, as shown in Table 11. These results suggest that once the molybdenum oxalate is supported, the oxalate group is in more or less the same chemical environment. The region below 1000 cm-l cannot be subtracted off reliably because of the strong absorption of IR radiation by the alumina support. Adsorption Mechanisms. The adsorption behavior of molybdenum oxalate solution onto alumina differs significantly from the A H M solution.’ The fact that the supported molybdenum oxalate can retain its M e 0 characteristic frequency observed in solution indicates that there is a strong interaction between the MoOx ion and the alumina surface. This also shows that molybdenum oxalate is a stable ion and that it is not degraded by the local pH of the support surface as suggested by Jeziorowski and Knozinger” for AHM. Even when water is driven out during the drying process, the equilibrium shift will not cause deposition of only polymeric species as suggested by Cheng and S ~ h r a d e r . ~ In order to show that the results obtained in this study were not due to a different alumina support, we prepared two Mo/A1203 catalysts of the same Mo loading from A H M solution using the same support and same drying treatment as our MoOx/AI20, catalysts. The pH of one of the impregnating solutions was adjusted to 9.8 by addition of ammonium hydroxide, and the other was at its natural pH of 5.2. The species in pH 9.8 sample are mainly monomeric molybdates MOO^^-) with the main peak frequency at 894 cm-l, and the species in pH 5.2 sample are mainly polymeric molybdates ( M O , O ~ ~ with ~ - ) a main peak frequency (1 1) Jeziorowski, H.; Knozinger, H. J. Phys. Chem. 1979, 83, 1166.
Ng et al. of 941 cm-I in solution. Figure 6 shows the Raman spectra of these catalysts after drying at 110 OC for 12 h and their corresponding impregnation solutions. The main peak position for both catalysts is at 941 cm-l despite their difference in the impregnating solution pH. This result agrees with Cheng and Schrader’s4 and also proves that our findings for MoOx/alumina are not due to differences in support, Mo loading, and drying conditions. The stability of the MoOx ion on the support surface can be explained in terms of the size and charge of the ion. From a study of anion adsorption on alumina, D’AnielloI2suggested that increasing size and charge of the adsorbed anion can enhance the stability of the surfaceanion pairing. Compared to simple molybdates, molybdenum oxalate has a larger size and higher charge which can account for its stability on the alumina support. The structural difference of unidentate and bidentate oxalate complexes of molybdenum may also suggest a difference in the adsorption path onto the alumina support. For the bidentate structure, the oxalate ligands are neutral and the possible adsorption will be via the terminal molybdenum-oxygen groups which carry negative charges. For the unidentate structure, both oxalate ligands and terminal M d groups carry negative charges which can be attracted toward the protonated sites on the alumina surface. However, the fact that we do not observe much change for the C=O and C-0 stretches in the IR spectra of MoOx/ alumina (Table 11) strongly favors the mechanism of adsorption via the oxalate groups. When the oxalate groups of the unidentate structure are adsorbed onto the alumina surface, they will be under the influence of the aluminum cation and therefore have a bidentate-like environment as the dimeric bidentate complex. It has been found that, for A H M solution^,'^ the amount of Mo loading adsorbed increases with decreasing pH (up to pH 1.O) on the alumina support because of the increase in degree of protonation. However, for molybdenum oxalate solution, the equilibrium loading at pH 6.0 is higher than the loading at pH 1.0.6 Our proposed adsorption mechanisms of molybdenum oxalate complex as a function of pH may well explain these results: for high pH, molybdenum oxalate exists as a monomeric bidentate structure where adsorption is via the oxalate groups. The oxalate group is more electronegative when compared with the terminal Mo=O group; therefore, it is more stable on the alumina surface and thus the amount adsorbed is more. However, another possible explanation is that the increase in the amount adsorbed a t high pH is due to the fact that some “simple molybdate” species are formed when the pH increases and they are very reactive toward the support. The actual amount of Mo adsorbed from molybdenum oxalate solution is considerably less than the amount of Mo adsorbed from A H M solution,6 especially for the very acidic region (pH 1). This decrease may be explained in terms of the steric effect of the molybdenum oxalate. For AHM solution, the dominant species are polymeric octamolybdates at pH 1. When this species is adsorbed onto an alumina surface, it is likely to form clusters of molybdates having a three-dimensional structure.2 For molybdenum oxalate, the degree of polymerization is limited to a dimer, and the charges on the complex prevent cluster formation, thus limiting the amount of Mo adsorbed to a one-to-one basis (one Mo per adsorption site on alumina).
Conclusions This study shows the changes of molybdenum oxalate structure in solution and on alumina as a function of pH by Raman and FTIR spectroscopy. The advantage of using both Raman and FTIR spectroscopy to study a metal complex system is demonstrated: Raman spectroscopy is more sensitive to metal-oxygen vibration while FTIR spectroscopy is more sensitive to carbonoxygen vibration. The following conclusions can be made: (1) Molybdenum oxalate exists as a dimer in the pH range 1.0-3.0 and as a monomer in the pH range 4.0-6.0. (2) The dimeric molybenum complex has a bidentate structure, while the monomer (12) DAniello, Jr., M. J. J . Cafal. 1981, 69, 9. (13) W a n g , L.; Hall, W. K. J . Cafal. 1982, 77, 232.
J. Phys. Chem. 1985,89, 2481-2486 complex has a unidentate structure. (3) Both monomer and dimer complexes are very stable on alumina surfaces as evidenced by the fact that they retain their characteristic Mo=O frequency in solution. (4) The adsorption mechanism for the dimer is via the terminal molybdenum oxygens, while for the monomer it is via the oxalate group. In summary, it is concluded that molybdenum oxalate, because of its stability on the alumina surface even under incipient wetness conditions, is a good candidate to be further explored as a possible
248 1
substitution for ammonium heptamolybdate as Mo-based catalyst precursor.
Acknowledgment. Financial support of this research by the Dreyfus Foundation, the National Science Foundation (CPE8107724 and DMR-8100130), and the donors of the Petroleum Research Fund, administered by the American Chemical Society, is gratefully acknowledged. Registry No. Alumina, 1344-28-1.
XPS Study on Surface and Bulk Palladium Oxide, Its Thermal Stability, and a Comparison with Other Noble Metal Oxides Marcell Peuckert Institut fur Grenzfliichenforschung und Vakuumphysik, Kernforschungsanlage Jiilich, Postfach 9 D-5170 Jiilich, Federal Republic of Germany (Received: November 6, 1984)
The surface of a Pd crystal was oxidized in 0.1-MPa flowing oxygen gas at 900 K and according to X-ray photoemission spectra a thick uniform layer of PdO had formed. It showed the same narrow Pd 3dSp and 0 1s peaks at 337.0- and 530.3-eV electron binding energy, respectively, as a bulk PdO sample. PdO thermally decomposed in ultrahigh vacuum in two distinct steps at 420 K and at about 750-800 K. In the transitional range over more than 300 K XPS signals of both the metal and the oxide were recorded. Similar behavior has also been observed with other noble metals and was interpreted as the formation of a cluster oxide phase.
Introduction In oxidation catalysis on noble metal surfaces the reactivity of the metal itself toward oxygen plays an important role, as a reduced or an oxidized surface will exhibit quite different activity and selectivity in the catalyzed reaction.I4 Especially in the case of palladium, at least three different states of oxygen (beside a probably existent but so far not reported molecularly adsorbed dioxygen species at low temperature) have been reported to exist in the surface region.+l0 Dissociative atomic oxygen adsorption at 300 K on Pd(ll1) and Pd(100) surfaces shows up in (2x2) LEED patterns corresponding to 0.5 monolayer saturation c o ~ e r a g e ;an ~ ?adsorption ~ energy of about 230 kJ mol-’ in accordance with a flash desorption peak at 900 K was meas ~ r e d . ~Exposure .~ to oxygen at low pressures and at temperatures between 500 and 1000 K leads to penetration of oxygen into the bulk and formation of some kind of subsurface compound, which is thought to be a transition state to Pd0.54 At about 800 to 900 K in 0.1 MPa of oxygen or air, three-dimensional PdO grows to a thick adlayer according to ref 9 and 10. While the chemical nature of adsorbed oxygen on one side and bulk crystalline PdO” on the other side are rather well understood, the intermediate so-called “subsurface oxygen” is still subject to debate. It is not clear whether this state should be described as a solid solution of oxygen in Pd as proposed by Campbell et al.,’ as some kind of subsurface ~ x i d e , or ~ .as ~ associated with oxygen bound to Si impurities that are contained in the Pd substrate.6 Whatever the exact nature of this species is, from the previous studies it appears certainly to be different from bulk PdO and adsorbed oxygen, since its decomposition temperature is much higher than the thermodynamic data of PdO would predict,I2-l4 and it proved to be inert toward H2 and C0.5*8 The experimental approach in the present investigation was to prepare a deep oxide adlayer on a Pd metal sample by oxidation in 0.1 MPa of O2gas and to follow the various stages of its thermal decomposition in ultrahigh vacuum (UHV) by X-ray photoemission spectroscopy (XPS). The results are compared with an ‘Present address: Hoechst AG, Keramilcforschung G 864, Postfach 800320, D-6230 Frankfurt 80, Federal Republic of Germany.
identical series of experiments on bulk PdO, and they are discussed in relation to the similar behavior of superficial oxide layers on Rh,15 1r,l6 Pt,” and Aula that have also recently been studied.
Experimental Section A high-purity (checked by XPS and Auger spectroscopy) Pd single crystal with (1 10) surface orientation was oxidized in a quartz tube furnace in 0.1 MPa of flowing oxygen gas at 900 K for 1 h. The sample was cooled in flowing oxygen, then clamped on a Mo holder, and transferred in an UHV system for XPS analysis (Leybold-Heraeus LHS 10). Commercial anhydrous palladium( 11) oxide (Ventron GmbH, 99.9%, 100 mesh) was used as a reference sample. The powder (1) V. I. Savchenko, G . K. Boreskov, A. V. Kalinin, and A. N. Salanov, Kine?. Katal., 24, 983 (1983). (2) C. F. Cullis and B. M. Willatt, J . Carol., 83, 267 (1983). (3) J. A. Gates and L. L. Kesmodel, J . Catal., 83, 437 (1983). (4) Ya. V. Salyn, M. K. Starchevskii, I. P. Stolyarov, M. N. Vargaftik, V. I. Nefedov, and I. I. Moiseev, Kine?. Katal., 24, 631 (1983). (5) H. Conrad, G . Ertl, J. Kiippers, and E. E. Latta, Surf. Sci., 65, 245 (1977). (6) T. W. Orent and S. D. Bader, Surf. Sci., 115, 323 (1982). (7) C. T. Campbell, D. C. Foyt, and J. M. White, J. Phys. Chem., 81,491 (1977). (8) D. L. Weissman, M. L. Shek, and W. E. Spicer, Surf. Sci., 92, L59 (19801. (9fP. Ugarl, L. Hilaire, G . Maire, G . Krill, and A. Amamou, Surf. Sci., 107, 533 (1981). (10) K. S.Kim, A. F. Gossmann, and N. Winograd, Anal. Chem., 46, 197 (1974). (1 1) A. F. Wells, ‘Structural Inorganic Chemistry”, Clarendon Press, Oxford, 1975, p 446. (12) J. S. Warner, J . Electrochem. SOC.,114, 68 (1967). (13) H. Kleykamp, Z . Phys. Chem. (Frankfurt am Main), 71, 142 (1970). (14) E. Fromm and E. Gebhardt, Eds., ‘Gase und Kohlenstoff in Metallen”, Springer, Berlin, 1976, and literature cited therein. (15) M.Peuckert, Surf. Sci., 141, 500 (1984). (16) M. Peuckert, Surf.Sci., 144, 451 (1984). (17) M. Peuckert and H. P. Bonzel, Surf. Sci., 145, 239 (1984), and literature cited therein. (18) M.Peuckert, F. P. Coenen, and H. P. Bonzel, Surf. Sci., 141, 515 (1984).
.
0022-3654/85/2089-248 l$Ol .50/0 @’ 1985 American Chemical Society
