8005
J . Phys. Chem. 1993,97, 8005-801 1
Dissociative Adsorption of Sulfur Dioxide on ?-Alumina Investigated by TPD and Mass Spectrometry I. G. Dalla Lana,? H. G. Karge,' and Z. M. George* Fritz-Haber-Institut der Max-Planck-Gesellschaft, 14195 Berlin, Germany Received: December 30, 1992; In Final Form: May 10, 1993
Prior studies on the adsorption of SO2 on y-alumina are not particularly revealing about the S o r a l u m i n a surface interactions. Chemisorbed SO2 on alumina was examined using a TPD-MS approach. The use of W 0 2 clearly demonstrated that both molecular and dissociated forms of SO2 chemisorb on the surface of y-alumina. The most likely mode of adsorption apparently involves interaction between oxygen atoms in the gaseous molecule and oxygen vacancies on the surface of alumina. The experiments also reveal the freedom of surface protons to form hydroxyl groups on the alumina surface.
Introduction Chemical reactions catalyzed by y-alumina have been studied over many years. To date, no single comprehensive view of the surface of this alumina clearly explains its somewhat diverse catalytic roles. The presence of base centers as well as both Br~ansted-and Lewis-acid centers have been used in various reaction mechanisms proposed. In the case of the modified Claus reaction, 2H,S
+ SO,
"/$,
S,,
3/2S, + 2H,O
(1)
n = 4,6,8
catalyzed by y-alumina, considerableeffort has been directed to elucidate the surface chemistry. While it is known that adsorbed H2S dissociates upon the surface of ?-alumina,' surprisingly, whether or not SO2 adsorption is clearly dissociative had not been demonstrated. This study not only was directed primarily at this uncertainty but also was used to examine, if possible, the role of the alumina surface during the chemisorption of SOz. The interactionbetween SO2 and alumina has been investigated using infrared spectroscopy'" and also by electron spin resonance1+5+'.10 and UV-visible.11-14 Several forms of adsorbed SOz, both weakly and strongly held to the alumina surface, have been proposedin thesestudies. 1R and UV-visiblespectroscopicstudies contributed significantly to the identification of the nature of those adsorbed species derived from H2S and SO2 and to the elucidationof their role in Claw cataly~is.~.'~ The interpretation was, to a large extent, based on the models of the surface of y-alumina as developed by Peril6and Kndzinger and Ratnasamy.17
Experimental Details 1. Materials. The y-alumina used in these experiments was ALON, a very finely dispersed industrial alumina (from Cabot Corporation). The 0.01-0.02-g samples (area of 100 m2-g1) were compressed into thin self-supportingwafers at about 1 MPa before being cut into 0.01-g samples for the experiments. The sulfur dioxide (99.995%) was purchased from MesserGriesheim, Diisseldorf. Isotopically labeled S*Oz was obtained from Isotron GmbH, Berlin (atom% content: 99.57 180, 0.02 17O,4.2234s). The adsorptionof SO2 upon alumina was reported in units of mmol SO2/g ALON. Experiments with water employed triple-distilled water subjected to three cycles of freezing to liquid N2 temperaturesfollowed by pumping to bS(10-8) mbar, after which the ice was isolated and thawed. Department of Chemical Engineering, University of Alberta, Edmonton, Canada. 8 Alberta Research Council, Edmonton, Canada.
0022-3654/93/2097-8005$04.00/0
MASS SPECTROMETER
TEMP€RANRE PROGRAMMER
f
A Figure 1. Description of TPD-MSequipment.
2. Equipment. Figure 1 describes schematicallythe apparatus used in the temperature-programmed desorption (TPD) experiments. A more detailed description is provided in ref 18. All connections used stainless steel fittings and vacuum flanges sealed with either copper or gold-plated copper gaskets (the latter were used in the permanent connections). The mass spectrometer (MS) was a Balzers Model QMG 3 11 quadrupole instrument with a normal scanning time of 1 s/amu, requiring about 1 min per cycle from 15 to 70 amu range. The MS responses, at assigned mass numbers, were plotted against temperature (or time) with the height of the deflection denoting the relative amount of that mass number. By plotting the consecutive responses within consecutivescanningcycles, the desorption patterns for each mass number of interest could be followed graphically during the course of the experiment. The best vacuum level attained was about 8 ( 1 t 9 ) mbar. 3. Experimental Procedure. Each experimental run varied from 24 to 60 h depending upon the number of steps involved and the rapidityof evacuationofthe apparatus at thestart. A weighed wafer of alumina was placed on a small horizontal heater, which was then insertedinto the vacuum chamber and its vacuum flange connected. A typical procedure follows: (i) The sealed apparatus was evacuated overnightwith external heating of surfaces (-425 K), followed by heating the alumina (at 5 K/min) to an activation temperatureof 675 K (in most experiments)until a stablevacuum pressure developed (about l . 2 ( l t 7 ) mbar). (ii) After cooling the apparatus to room temperature, meanwhile maintaining the alumina wafer at 675 K (to minimize readsorption of residual gases), the pressure stabilized to about 2-3(10-8) mbar. The quadrupole MS filament current was also always switched on for sometime prior to use to desorb gases and ensure their removal. 0 1993 American Chemical Society
Dalla Lana et al.
8006 The Journal of Physical Chemistry, Vol. 97, No. 30, 1993
64
300
400
500 600 TEMPERATURE [K]
700
Figure 2. Temperature-programmeddesorption of Sl6O*from alumina; activation temperature, 675 K adsorption temperature, 295 K.
(iii) The TPD heater was cooled from the activation temperature to the desired adsorption temperature. With a stable vacuum (1-1,2(10-9 mbar), the MS scan of residual gases usually indicated negligiblecontent; usually only a small water peak (mass 18) could be discerned. (iv) SO2 (1.100 mbar), measured in a calibrated volume manifold (546.2 cm3) at room temperature, was expanded into the sample chamber (2286 cm3 total) and allowed to stand until adsorption equilibrated. By recording the various pressures, the total mmol SOz/g ALON adsorbed could be calculated. (v) The liquid Nz-cooled trap then condensed all gaseous SO2 as well as the physically adsorbed SO2 which had desorbed. By the difference between the total and condensed amounts the residual SO2 chemisorbed on the ALON could be determined. The entire chamber and manifold were then reevacuated and another baseline MS scan recorded. (vi) The TPD then commenced using a 10 K/min heating rate from the initial temperatureof adsorption to the final temperature, usually 825 K. The repetitive 1-min MS scans were recorded throughout theTPD program. Toestablish the fragmentation levels (enabling the calculation of relative abundances) for gaseous S02, MS spectral scans were recorded at several low pressures of S02, straddling pressures encountered during the various TPD scans.
Results and Observations 1. Temperature-Programmed Desorption of 5'602. Figure 2 shows a typical plot of the experimental data obtained from a TPD run. The fragmentation of the SO2 in the quadrupole MS occurs via
- -
-
+
+ lag+ s+ l 6 0 + (2) 48 32 with the mass numbers, 64, 48, and 32, being monitored and plotted. The broad distribution for mass 64 as a function of S'602(g)
s'602+
64
s'60+
temperature is indicative of a heterogeneous surface. All such plots exhibit a maximum which increases with temperature as the adsorption temperature is increased. At the cut-off temperature of 775 or 825 K, some of the SO2 remains chemisorbed. Since the sample was previously activated by heating at 675 K, additional dehydroxylation (mass numbers 17 and 18) and C02 (mass number 44) desorption occurs only above 675 K. Figure 3 superimposesTPD curves for the 64 species from five runs ranging in adsorption temperature from room temperature to 525 K. The SO2 adsorbate could not be removed simply by
vacuum pumping nor entirely by heating to 775 K; thus, such strong bonding to the ALON surface is likely to be chemical bonding. Since all other experimental conditions were kept constant, including the 0.01 g of ALON, these traces clearly indicate that the adsorption temperature is the significant parameter and SO2 adsorption is an activated process in that, at higher adsorption temperatures, sites with higher adsorption energies (deeper potential wells) became accessiblefor the surface species. This in turn will be reflected in the desorption kinetics, i.e., the peak temperature in the TPD spectra shifts to higher temperatures as demonstrated in Figure 3. Figure 4 contrasts the influence of both activation and adsorptiontemperatures upon the 64 mass trace. With adsorption at 375 K, it seems that the additional dehydroxylation from activating at 975 K rather than 675 K has slightly reduced the total adsorption capacity, but the maximum around 540 K remained almost unchanged. With adsorption at 475 K, a change of adsorption capacity, if any, cannot be recognized from the respective traces since these are rather incomplete. However, the spectrum of SO2 desorption from alumina activated at the temperature increased to 975 K exhibits a maximum at a higher desorption temperature, viz., at 750 K compared to 650 K after activation at 675 K. This shift is due to the higher population of particularly strong adsorption sites created at 975 K and accessible for the surface species only at higher adsorption temperatures (vide supra). Since thevacuum pumping equipmentandoperating conditions were essentially constant for each run, one expects that the MS peak intensitiesfor a given mass are proportional to the population of this species being generated at that instant. Accordingly, the areas under these curves for the same mass number should be proportional to the amounts of that species desorbing. Therepeatability of the TPD-MS scans was excellent. Several runs were duplicated with nearly identical scans resulting. Becauseof this experimentalreliability,thefiveTPDcurvesshown in Figure 3 could be plotted on the same coordinate system for visual comparisons. Furthermore, the quantitative adsorption measurements from two runs on the same sample as plotted in Figure 5 also exhibit consistency. 2. Quantitative Measurements of 5'602 Adsorption. Figure 5 shows the adsorption isotherm for SO2 chemisorbed at 375 K after activation at 675 K. The isotherm was reproduced in two independent experiments (samples 1 and 2) and also after a TPD
The Journal of Physical Chemistry, Vol. 97, No. 30, 1993 8007
Dissociative Adsorption of SO2 on ?-Alumina
1
1
I
300
400
1
I
600 TEMPERATURE [K]
500
I
I
700
800
Figure 3. Temperature-programmeddesorption of S1602from alumina influence of adsorption temperature; activation; 12 h at 675 K, PSO,= 0.200 mbar; m / e = 64 (SOz+).
I
1
400
500
600
700
800
TEMPERATURE [K]
Figure 4. Temperature-programmeddesorption of Sl602from alumina influence of activation or adsorption temperature (a) Tad = 375 K and (b) Tad = 475 K.
to 825 K (sample 2). When sample 2 was heated only to 675 K, not all of the SO2 desorbed as shown by its subsequent lower uptake of S02. By tracking the progress of the TPD in 100 K increments and measuring the desorbing mass increments up to 675 K, the data of Table I were collected. The liquid nitrogen condenser was used to trap the mass increments and to maintain the equipment manifold under vacuum during each heating step. The volume of gas collected in the condenser was measured and then recondensed after each heating step. The two increments of desorbed gas collected above 675 K were not entirely due to SO2 because desorbing H2O and COz also contributed to the mass. As a result these two values for chemisorbed SO2 are too small. Karge and Dalla Lana4 show a correlation between SO2 adsorbed and absorbance of a strongly bonded species (indicated by a 1065-cm-1 IR absorption band) at room temperature. The values in Table I appear to be of the same order of magnitude as the published values at the end of their linear plot, uiz. 1.5 X 10-1 mmol S02/g Al203; the two sets of data were taken at similar
adsorption temperatures, the adsorption temperature obtained by the sample in the IR beam having been about 375 K. Table I1 contrasts the total adsorbed and chemisorbedamounts of SO2 at several combinations of calcination and adsorption temperatures. The total amount adsorbed does not appear to increase significantly with increase in activation temperature at the given equilibrium pressure of SOz. However, the chemisorbed fraction of the total increases with calcination temperature, suggesting that dehydroxylationof alumina increasesthe capacity of the surface to chemisorb S02. Dehydroxylation at elevated temperatures apparently increases the number of adsorption sites likely via
20H-
-
H 2 0 t + 02-+ [oxide vacancy]
(3)
with the products being favorable for SO2 chemisorption. At the calcination temperature of 1025 K, however, the amount chemisorbed now decreases with increasing adsorption temper-
8008 The Journal of Physical Chemistry, Vol. 97, No. 30, 1993
Dalla Lana et al.
I
1
1
0.20
0.10
0.40
0.30
PRESSURE OF SO2 [MBAR 1 Figure 5. Adsorption isotherm for SO2 on alumina at 375 K; activation temperature, 675 K.
TABLE I: Adsorption of SO2 at 375 K on Alumina Activated at 675 K, Followed by Heating and Desorption in 100 K Steps PSO,at temp (K) equilibrium adsorption of SO2 (mmol S02/g A1203) initial final 375 375 475 575 675 750
375 475 575 675 750 775
(mbar)
total
chemisorbed
physisorbed
0.083
0.464
0.138 0.124 0.089 0.062 0.048 0.017
0.326
1.0-
0.9-
48 I 6 4
OD
P N-
m
a-
0.8-
r
II
0.7-
3
qa $ 5 0.6II
TABLE 11: Effect of Calcination Temperature upon SO2 Adsorption adsorption of SO2 temp (K) Pso, (mmol S02/g A1203) calcination adsorption (mbar) total chemisorbed 675 675 1025 1025 675 1025 675
375 315 375 375 475 475 525
0.083 0.185 0.174 0.179 0.187 0.184 0.188
0.464 0.439 0.512 0.530 0.487 0.494 0.475
0.138 0.180 0.308 0.252 0.205 0.242 0.219
ature, e.g., from 0.252-0.308 at T(ads) = 375 to 0.242 mmol SO2/g A1203 at T(ads) = 475 K. At the very high T(ca1c) = 1025 K, the crystal structure of the y-alumina may have been, a t least partially, altered via phase transformations to the 6- or @-aluminaforms. Thus, the surface state and, correspondingly, the nature of the adsorption sites will differ to some extent from those of alumina surfaces activated at lower temperatures. This, in turn, should affect the SO2 adsorption. 3. Dissociation of Adsorbed S”Q. Figure 2 showed that the TPD-MS desorption patterns for parent and fragment ions closely tracked each other. By comparing the relative abundances of the fragment ions obtained from TPD of adsorbed SO2 with relative abundances from comparable TPD experiments with no alumina present, both a t the same pressure of S02, evidenceof the presence of SO+ ions in excess of that from normal fragmentation (no alumina) would support the case for dissociative adsorption. Figure 6 shows a semilogarithmic plot of the relative abundances, SO+/SO2+,S+/SO2+, and O+/SO2+, encountered as a function of SO2 pressure, no alumina being present. As anticipated, the three relative abundances parallel each other; however, at SO2 pressures above 2( le7) mbar, the SO+/SO2+
34
030
;$
e 0.5-
aF
:5
0.4-
0% w
p
a 0 z
3
m U
0.332/64 0.2-
16/64 0.1 -
0.04 10.7
10.8
1
P R E S S U R E O F SO, [ M B A R ]
Figure 6. Fragmentation pattern for Sl6O2; relative abundances as a function of SO2 pressure during MS fragmentation of gaseous SO2 no
alumina being present.
(48164) ratio appears to be excessive. The same SO*+signal was used in calculating each of the three mass ratios: hence, it would seem that the SO+ intensities appear to have become too large. Since the mass abundance levels for masses 48 and 64 at SO1 gas pressures above le7mbar tended to saturate the quadrupole MS counter, the large attenuations needed as well as the increasing inaccuracy of the quadrupole limited these background measurements. In this upper range, it seemed that the upturn in the 48/64 trace was not significant. Table I11 comparesthese relative abundance ratios fromvarious experiments with ratios from SO2 gas fragmentation in the same
The Journal of Physical Chemistry, Vol. 97, NO. 30, 1993 8009
Dissociative Adsorption of SO2 on y-Alumina
TABLE 111: Examination of MS Relative Abundances for Evidence of Presence of Desorbed SO experimental TPD-MS fragmentation of SO2 gas desorption calcination temp (K) SO+/S02+ s+/so2+ SO+/S02+ s+/so2+ adsorption Talcin 475 0.94 0.31 0.63 0.15 0.64 0.76 0.97 675 K 575 0.22 T*& 375 K Talcin
975 K T*& 375 K
0.33 0.32 0.23 0.34 0.31 0.21
0.99 0.91 0.93 0.97 0.96 0.59
675 775 475 575 675 775
0.73 0.73 0.70 0.75 0.72 0.71
50
I
0.104 0.104 0.104 0.104 0.104 0.104 0.104 0.104
I
li'
.J-.-..
*'I..,
','
I
s+/s02+
0.493 0.493 0.493 0.493 0.493 0.493 0.493 0.493
0.21 0.21 0.19 0.23 0.20 0.19
5K
.,.
fragmentation published
so+/so2+
TEMPERATURE [K] Figure 7. Temperature-programmeddesorption of S1802from alumina; activation temperature, 675 K adsorption temperature, 375 K.
apparatus. Both sets of ratios greatly exceed values calculated using publishedIg relative intensities for S02. Irrespective of possible common errors, the experimental values of SO+/SO2+ (TPD-MS) always exceed the background values at the same SO2 pressures, suggesting that SO+is being generated from both SO2 and SO species being desorbed from the alumina surface; qualitatively, this is indicative of SO2 dissociation. 4. Temperature-Programmed Desorption of S1*O*. TPD experiments identical to those using S I 6 0 2 were conducted using 180-labeled S1802. Figure 7 shows a rather detailed desorption pattern involving initially
with oxygen atoms in the surface of alumina, and the presence of mobile H atoms enables surface hydroxyls to be formed with the exchanged 0 atoms. The two sequential exchange reactions exhibit 68,66, and 64 plots like those of a classical two-step sequential reaction. Sequentially, mass 68 maximizes at about 550 K, mass 66 about 650 K, and mass 64 is still rising at 875 K.
po, po2+ p o ++ 1 8 0 + s++ 1 8 0 +
SO2 is irreversibly adsorbed on alumina unless the temperature
--+
--*
68
50
18
-
(4)
32
Surprisingly, the plot of mass 50 exceeded that for the parent species S1802+(68). The plot also shows the Occurrence of the exchange reactions with the alumina surface
-
-
Po, s180160P o 2 68
66
(5)
64
The parent species S 1 8 0 1 6 0 (66) also contributes S80+ (50) to that from S1802, the sum of which exceeds the intensity of the 68 species. The appearance of S 1 8 0 1 6 0 convincingly demonstrates that S I 8 0 2 bonds to the surface of alumina during chemisorption and exchanges 1 8 0 for 1 6 0 , for which dissociation of SO2 is required. This experiment alone confirms the tentative arguments for dissociation presented earlier herein. Of equal interest, and also proof of SO2 dissociation, is the appearance of H2180 as a desorbing species. Thus, SO2 upon adsorption easily exchanges oxygen
Discussion The above evidence enables certain inferences: (i) Chemisorbed is increased above the adsorption temperature. To some extent, SO2 remains chemisorbed even at 875 K. (ii) Chemisorption of SO2 on alumina is an activated process. The surface of alumina exhibits heterogeneous character in energy of bonding between SO2 and surface sites. (iii) Chemisorbed SO2 may exist in both undissociated and dissociated forms since both SOzand SO desorb during TPD. (iv) The TPD of chemisorbed S I 8 0 2 clearly demonstrates that chemisorbed SO2 can dissociate on alumina. (v) Exchanged 0 atoms on the alumina surface are easily accessible to mobile H atoms, enabling surface ISOH groups or I60Hgroups to form equally. (vi) The continued dehydroxylation of alumina to form water upon heating suggests that OH groups are either heterogeneous in bonding energies or water adsorbed within alumina migrates to the surface to be desorbed. The parallelism between the 18 and 20 mass TPD curves suggests the former is more likely. Tentative Mechanism. The inference (iv) and a preliminary description of a mechanism involving oxide sites was briefly expressed earlier.5 This mechanism is expanded herein. Because
8010 The Journal of Physical Chemistry, Vol. 97, No. 30, 1993
Dalla Lana et al.
of the form of heterogeneity of the alumina surface sites involved A in the chemisorption of S02, indicated by the single broad TPD peak, a large distribution of one type of site appears most likely. These sites could vary with minor energy variations between the different atomic arrangements associated with these sites. The easy access and exchange capability of oxygen atoms with the 0.64 nm alumina surface suggests that oxygen vacancies are available, ranging in their location (edges, planes, etc.) and/or in the type f of nearest neighbors (OH, 0,either 0 or AI vacancies, e t ~ . ~ ~ J ~ ) . adsorption on vacancy desorption of S'80'6 0 (exchange) The many possibilities create the heterogeneous character of oxygen vacancies. In such a model, the oxygen atoms adjacent to the A1 vacancy would have an excess of negative charge. Normally, such charges would be satisfied by surface protons which arevery mobile and easily form hydroxyl groups a t locations surface diffusion structure I a of SIEO structure I b adjacent to vacancies. Using a simple layered atomic model of the alumina surface, Figure 8A indicates the formation of structure I, involving a single bond between an oxygen vacancy and an oxygen atom 0 . @ O from the gaseous SO2 molecule. Figure 8B shows an additional 32s AI l80 '60 step, whereby a different adsorbed structure I1 is formed, in which two oxygen vacancies have participated. This may occur prior B to formation of structure Ia shown in Figure 8A. The evidence presented supports the view that dissociation of SO2 on a single vacancy and mobility of the resulting fragment, S80, on the alumina surface are likely. It would be anticipated that the mobile 4 SI80is able to occupy any oxygen vacancy and not just the nearby vacancy shown in Figure 8B. Moreover, one has to keep in mind that Figure 8 provides a very schematic representation of the alumina surface. A detailed quantum mechanical treatment would have to take into account the particular orbitals and their orientation when interacting with thoseof the adsorbate molecule. Earlier ESR experiments'O suggested that two types of S02surface diffusion of sulphur anion species may exist on the surface of alumina. It is conceivable I that local net charge transfer from the oxygen vacancy environment on the surface could generate two paramagnetic species with the general character of I and 11. Species I most likely I O .n~d desorbs as SO2 and species I1 could desorb either as SO2 or SO desorption of Sl80 structure II dissociation during the vacuum conditions within the TPD-MS apparatus. These processes may occur with or without oxygen exchange C with the alumina surface (see Figure 8A,B). With this dissociative adsorption mechanism, sulfur atoms in SO2 are less attractive to oxygen vacancies than the oxygen atoms. They could, however, form bonds with surface oxygen atoms (probably in arrangements with HS03-, S032-, or S042- electronic character; see, e.g., structures IIIa or IIIb in Figure 8C). To examine the feasibility of the surface vacancy model in describing the adsorption of SO2 on alumina, some simple formation of HSO; formation of HSO; calculations were performed. A gas-phase molecule of SO2 has a bond angle of 120° and a S-0 bond length of 0.142 nm20(center to center from S to 0 atoms). From this information, one calculates a 0.246 nm oxygen to oxygen distance in the SO2 molecule. For the surface of y-alumina, Peril6 suggests that each oxide ion occupies an area of about 0.08 (nm)2. Assuming this area to be circular, the radius of an oxide vacancy would structure llla structure lllb approximate 0.16 nm. From these simple geometrical considerations, it seems evident that either of the oxygen atoms in SO2 0 0 may occupy a single surface oxide vacancy. Alternatively, the a 0 AI 0 H S two oxygen atoms in SO2 cannot simultaneously occupy two single Figure 8. Simplified scheme for (A) oxygen -exchange, (B) SO2 oxide vacancies separated by one oxide ion without dissociation dikociation, and(C) formation of HSO3- structures on alumina surface: of S02. Also, the SO2 gas molecule would be unlikely to adsorb (A) adsorption of SI802 and 1 8 0 l60 exchangewith surface of A I 2 . 0 3 onto a linear double oxidevacancy without extension of the oxygen rhsulting in desorbed S8O160,(B) adsorption of SI802and dissociation to oxygen distance. under formation of S160 and/or S80,and (C)adsorption of SO2 and formation of HSO3- structures. The strength of such a proposed model of the alumina surface would be its ability to predict the adsorption behavior for other because of the bond angle between the two hydrogen atoms, two oxygen-containing adsorbate gases. For example, in an experadditional hydrogen bonding links may be formed. These three iment involvingthe introduction of H2O molecules into the vacuum linkages would generate a very strong H20-alumina adsorption. chamber, water molecules may occupy oxygen vacancies. But This is a well-known property of y-alumina.
b
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-
6
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1
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&-
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-
Dissociative Adsorption of SO2 on ?-Alumina
Acknowledgment. Financial support in the form of a NATO Travel Grant (to I.G.D.L. and H.G.K.) and from the Fritz-HaberInstitut der Max-Planck-Gesellschaft (to I.G.D.L. and Z.M.G.) is gratefully acknowledged. References and Notes (1) Karge,H.G.;Tower,R. W.;Dudzik,Z.;George,Z.M. Inproceedings
of the 7th International Congress on Catalysis, Tokyo, Japan, June 30-July 4, 1980 Seiyama, T., Tanabe, K., Eds.; Kodansha: Tokyo, 1980; p 643. (2) Deo, A. V.; Dalla Lana, I. G.; Habgood, H. W. J. Catal. 1971,21, 270. (3) Chang, C. C. J . Catal. 1978, 53, 374. (4) Karge, H. G.; Dalla Lana, I. G. J. Phys. Chem. 1984, 88, 1583. (5) Karge, H. G.; Dalla Lana, I. G.; Trevizan de Suarez, S.;Zhang, Y.
In Proceedings of the 8th International Congress on Catalysis, Berlin, Germany, July 2-9, 1984; Dechema, Ed.; Verlag Chemic: Vol. 111, p 453. (6) Datta, A.; Cavell, R. G.; Tower, R. W.; George, Z. M. J. Phys. Chem. 1985,89, 443; 1985,89,454. (7) Lavalley, J.-C.; Janin, A.; Preud-Homme, J. React. Kinet. Catal. Lett. 1981, 18, 85. (E) Babaeva, M. A,; Tsyganenko, A. A,; Filimonov, V. M. Kinet. Katal. 1984, 25,921. (9) Khulbe, K. C.; Mann, R. S . J. Catal. 1978, 51, 364. (10) Gutsze,A.;George, Z.; Dalla Lana, I. G.; Karge, H. G. In Proceedings ofthe 9th International Congress on Catalysis, Calgary, Canada, June 26-
The Journal of Physical Chemistry, Vol. 97, No. 30, 1993 8011 July 1 , 1988; Phillips, M.,Ternan, J., Eds.;Chemical Institute of Canada: Ottawa, Vol. 1, p 1791. (11) Karge, H. G.; Ziokk, M.; Laniecki, M. Zeolites 1987, 7, 197. (12) Karge, H. G.; Laniecki, M.;Zidek, M. Proceedings of the 7th International Zeolite Conference, Tokyo, Japan, August 17-22, 1986; Murakami, Y., Iijima, A., Ward, J. W., Eds.; Kodansha: Tokyo, and Elsevier: Amsterdam, 1986; p 251. (13) Karge, H. G.; Laniecki, M.; Ziokk, M. J. Catal. 1988, 109, 252. (14) Karge, H. G. In Recent Developments in Catalysis-Theory and Practice. Proceedings of the 10th NationalSymposium on Catalysis andlth Indo-Soviet Seminar on Catalysis, Madras, India, Dec. 18-21, 1990, Viswanathan, E., Pillai, C. N., Eds.;Narosa Publ. House: New Delhi, India, 1990; p 1. (15) Karge, H. G.; Zhang, Y.; Trevizan de Suarez, S.;Ziokk, M. In Proceedingsof the Conference'StructureandReactivityofModiiflcdZeolitesn, Prague, Czechoslovakia, July 9-13, 1984; Jacobs, P. A., Kazansky, V. E., Jaeger, N. I., Jiru, P., Schulz-Ekloff,G., Eds.; Elsevier Scientific Publ. Co.: Amsterdam, 1984; Studies Surf. Sci. Catal. 1984, 18, 49. (16) Peri, J. B. J. Phys. Chem. 1965, 69, 220. (17) Knbinger, H.; Ratnasamy, P. Catal. Rev.-Sci. Eng. 3978, 17, 31. (18) Karge, H. G.; Dondur, V. J. Phys. Chem. 1990, 94, 765. (19) Cornu, A.; Massot, R. Compilation of Mass Spectral Data; Heyden 6t Son: London, 1966; pp 5B, 151D. (20) Hargittai, I. In Sulphur Molecular Structures, Springer-Verlag: Berlin, Heidelberg, New York, 1978; p 74.
