7550
J. Phys. Chem. 1996, 100, 7550-7557
Adsorption and Reaction of Sulfur Dioxide on Alumina and Sodium-Impregnated Alumina Mark B. Mitchell* and Viktor N. Sheinker Department of Chemistry, Clark Atlanta UniVersity, Atlanta, Georgia 30314
Mark G. White School of Chemical Engineering, Georgia Institute of Technology, Atlanta, Georgia 30302 ReceiVed: July 11, 1995; In Final Form: NoVember 30, 1995X
The adsorption and oxidation of SO2 on alumina and sodium-impregnated alumina has been examined using thermogravimetric analysis and diffuse reflectance infrared Fourier transform spectroscopy. Sulfur dioxide chemisorbs initially at basic sites to form an adsorbed sulfite, which is quantitatively converted to sulfate on oxidation. It has been observed that at low coverages, ∼2.6 µmol/m2, sodium acts as a promoter for the formation of an adsorbed sulfite and sulfate which have structures similar to those of aluminum sulfite and sulfate, respectively. At higher sodium loadings, a second type of adsorbed SO2 is formed, similar to sodium sulfite and sulfate. The species with the aluminum sulfate structure appears to be more easily decomposed than does the sodium sulfate species and accounts for the regenerable adsorption capacity. Formation of the sodium sulfate species appears to account for the loss of adsorption capacity as the number of adsorption/ regeneration cycles increases. Oxidation of the sulfite form to the sulfate form can occur in the absence of added oxygen, but it is an activated process and begins to occur in measurable amounts at temperatures between 150 and 300 °C. Partitioning of adsorbed SO2 between aluminum and sodium forms is not a function of temperature and depends on only sodium loading.
Introduction The interaction of sulfur dioxide with alumina has considerable economic and environmental importance. Sulfur dioxide emission from coal-fired power plants, which leads to acid rain, is one of the main environmental concerns related to coalderived electric power. A number of different metal oxides have been screened for use as regenerable sulfur dioxide sorbents.1 Alkalized-alumina, alumina doped with sodium and/or lithium, has been investigated as a regenerable SO2 sorbent.2-4 The NOXSO process, which uses alumina doped with sodium carbonate, is being evaluated in pilot plant studies.5,6 Alumina is also the major catalyst used for the Claus process, for the oxidation of hydrogen sulfide to elemental sulfur and water using sulfur dioxide. Coal gasification and the processing of natural gas result in the production of large amounts of hydrogen sulfide, and the Claus process is one method for converting H2S into a marketable byproduct. For these reasons, a number of different groups have invesigated the interaction of alumina with sulfur dioxide. In the majority of these investigations, infrared spectroscopy was used to attempt to elucidate the interaction mechanism.7-13 Karge and Dalla Lana studied the interaction of SO2 with γ-alumina using infrared spectroscopy.8 Exposure of the surface to SO2 followed by desorption yielded infrared bands assigned to weakly adsorbed SO2 and strongly adsorbed SO2. H2O displaced the weakly held SO2 sites at room temperature and at 200 °C, but the strongly held SO2 was not affected. Similarly, water blocked the weakly held SO2 adsorption sites but not the strongly held sites. Preadsorbed ammonia and pyridine had similar effects. Adsorption of the Lewis acid BF3 effectively blocked the strong adsorption sites while having little effect on the weak adsorption sites. Thus, they concluded that the interaction of SO2 with basic sites on the surface leads to X
Abstract published in AdVance ACS Abstracts, February 15, 1996.
S0022-3654(95)01922-8 CCC: $12.00
formation of chemisorbed SO2 while adsorption at acid sites leads to the physisorbed form of SO2. From the infrared spectra generated during adsorption/ desorption experiments, Chang9 concluded that SO2 is initially adsorbed as a sulfite species. He observed that further exposure yields first another form of chemisorbed SO2 and finally physisorbed SO2. Oxidation of adsorbed sulfite/SO2 in oxygen at 400-500 °C gave rise to two strong, intense absorptions. Through the comparison of these spectra with those from γ-alumina doped with Al2(SO4)3, it was concluded that the oxidation of adsorbed SO2 resulted in the formation of aluminum sulfate-like species on the surface. Also on the basis of adsorption/desorption experiments using infrared spectroscopy, Datta et al.10 arrived at a detailed mechanism for the adsorption process involving initial adsorption either at Lewis acid sites (coordinately unsaturated aluminum atoms) or at Lewis base sites (exposed oxygen atoms). Adsorption at the Lewis acid sites gives rise to relatively weakly held SO2. Adsorption at the basic sites followed by cleavage of an Al-O bond was proposed as the source of surface-bound sulfite-like species (see below). Once the initial adsorption sites have been saturated, further adsorption results in the formation of physisorbed SO2 primarily at OH sites.
Saur et al.11 found results similar to those found by Chang9 in that when SO2, adsorbed on alumina, was oxidized at 450 °C in an excess of oxygen, the infrared spectrum was identical to that generated by doping alumina with Al2(SO4)3‚18H2O and dehydrating at 450 °C. They noted the disagreement between the observed spectrum and the infrared absorption spectrum of © 1996 American Chemical Society
Adsorption of SO2 on Na-Doped Alumina pure aluminum sulfate. This observation and the results of 18O2exchange experiments and H218O-exchange experiments led them to propose the structure
for the sulfate species present in an anhydrous environment which can hydrolyze to
J. Phys. Chem., Vol. 100, No. 18, 1996 7551 Once in solution these species ionize, with the sodium acetylacetonate ionizing to a greater extent than the lithium species, and the adsorption results are not significantly more uniform than those which might be obtained using other methods. For the studies which are reported here, sodium methoxide was finally chosen as the sodium source for impregnating alumina, since the impregnation could easily be carried out in methanol. In the work which is reported here, diffuse reflectance infrared Fourier transform spectroscopy (DRIFTS) and thermogravimetric analysis (TGA) have been used to examine the effects of added sodium on the adsorption and reaction of sulfur dioxide on alumina. Our results provide clear evidence that sodium acts as a promoter for the formation of one type of sulfite species on the sorbent at low sodium loadings. A second type of sulfite is indicated at higher sodium loadings. The sulfite which is formed is quantitatively converted to sulfate on oxidation. Experimental Section
which would correspond to a more normal aluminum sulfate. Nam and Gavalas13 investigated the behavior of the SO2/ alumina complex at temperatures in excess of 500 °C and observed infrared spectra very similar to those found by Chang9 and by Saur et al.11 A recent study by Dalla Lana et al.14 using temperatureprogrammed desorption-mass spectrometry led these investigators to propose a single type of adsorption site for SO2 on alumina, an oxygen vacancy, with a broad distribution of interaction energies depending on the location of the vacancy and the nearest neighbors. It has been shown that the impregnation of alumina with different alkali or alkaline earth metals affects the basic site distribution in different ways.15 Since strong SO2 adsorption on alumina occurs at basic sites, alkalized alumina is an obvious choice for an alumina-based sorbent for SO2. Methods for the preparation of alkali-doped aluminas include the following: (a) melting and sintering of alumina with the hydroxides16 or carbonates3,17 of alkali metals at 650-1100 °C; (b) mixing dried sodium acetate with alumina and calcining the mixture at 450 °C;4,18 (c) incipient wetness using aqueous sodium sulfate or acetate followed by calcining at 700-800 °C;4,18 (d) soaking alumina in aqueous solutions of sodium-containing salts, followed by filtration of the solid, drying, and calcining.3,19,20 Other methods of impregnating alumina with alkali and alkaline-earth metals are described elsewhere.21-23 Our group has been involved in the use of nonaqueous impregnation techniques for the preparation of well-dispersed metals on metal oxide substrates.24 Our earlier work using magnesium acetylacetonate for aluminum impregnation with magnesium,25 and copper acetylacetonate for copper impregnation on silica,26,27 showed that it was possible to generate welldispersed metals on these oxide surfaces, with predictable uptake from nonaqueous solution, methanol, or acetonitrile. One of the initial goals of this investigation was to examine whether or not the use of nonaqueous impregnation techniques for sodium and other alkali metals could yield a sorbent which was more efficient for sulfur dioxide uptake or more easily regenerable. Our experience with the alkali-metal acetylacetonate species, however, indicated that they were not predictable precursors. Divalent copper and magnesium acetylacetonate species yielded much more reproducible and predictable results than were observed in our early studies with monovalent lithium and sodium acetylacetonate. The alkali-metal complexes do not dissolve in acetonitrile and dissolve only slightly in methanol.
Sources of Materials. The γ-alumina used for these experiments is a high-purity (99.995%) product from Goodfellow with less than 0.1 µm average particle diameter and a surface area of ∼150 m2/g. It was pretreated before impregnation by heating in a vacuum oven to 150 °C for 24 h. Sodium methoxide was purchased from Aldrich as a 25% solution in methanol. The methanol used was HPLC grade from Aldrich. Method of Sample Preparation. All aspects of the impregnation were carried out in a drybox or on a Schlenk line as appropriate. Before impregnation, 5 g of the dehydrated γ-alumina were put in 50 mL of distilled and dried methanol. A specific amount of the 25% solution of sodium methoxide was then added to the solution, and the mixture was stirred for 5 h. The solid was filtered under dry nitrogen, washed with methanol, filtered again, dried under vacuum, and calcined for 6 h at 600 °C in air. Elemental analysis of the samples for sodium was carried out by Applied Testing Labs of Marietta, GA. The specific amounts of the sodium methoxide solution to be used for the impregnations are calculated as follows. The γ-alumina as prepared in these studies should contain approximately 10 OH groups per nm2.28 It is assumed that the mechanism for sodium impregnation under these conditions is
AlO-H + Na-OCH3 f AlO-Na + CH3OH Three different impregnated aluminas were prepared using different mole ratios of sodium to surface OH in the impregnation solution. Sample B was prepared using a 1:2.5 calculated mole ratio of sodium to surface OH, while samples C and D were prepared using 2:1 and 4:1 calculated mole ratios of sodium to surface OH, respectively. For the calculations, it was assumed that 25% of the volume of the methanol solution corresponded to NaOCH3, and from a measurement of the density of the solution, the appropriate volume of the 25% solution to give the desired number of moles of sodium methoxide per gram of alumina for each sample was calculated. TGA Method and Instruments. The TGA experiments were carried out using 5.5-8.5 mg samples of the alumina or alkali-doped alumina. The TGA instrument used was a PerkinElmer TGS-2 system, interfaced to a personal computer for data acquisition. The instrumental precision was (0.002 mg. A combined gas flow of 150 cm3/min was used for all cycles. The pretreatment cycle involved heating the sample in a gas stream of 90% O2 and 10% N2. The sample was heated from 25 to 700 °C at the rate of 30 °C/min, held at 700 °C for 5 min, and then cooled to 30 °C. The adsorption cycle involved flowing the SO2 (1100 ppm SO2, balance N2) over the sample
7552 J. Phys. Chem., Vol. 100, No. 18, 1996 for 20 min while the temperature was held at 30 °C. The sample was then purged with pure N2 for 20 min to remove loosely held/physisorbed SO2. Some weight was lost during this part of the experiment. The desorption cycle involved heating the sample from 25 to 700 °C at 15 °C/min, maintaining the sample at 700 °C for 1 min, and then cooling to room temperature. The adsorption/desorption cycle was repeated three times for each sample. The weights of the samples after adsorption of SO2 and after desorption of SO2 that are reported for each sample were determined at 30 °C in a flowing nitrogen stream, before and after heating, respectively. FT-IR Method and Instrument. Samples were prepared for infrared spectroscopy by mixing presifted and calcined (muffle furnace, 500 °C in air) γ-alumina (sample A) or its sodium-doped derivatives (samples B-D), 6% by weight in KBr. The powder mixture was loaded into the DRIFTS sample cup, and the surface was carefully leveled with minimal pressure. The sample was then heated to 400 °C in O2 to oxidize any remaining organic species, followed by cooling to a target adsorption temperature under vacuum. A spectrum taken at the target temperature was used as the pristine substrate spectrum and was subtracted from all subsequent spectra obtained at that particular temperature. Fresh samples were used for every temperature investigated. The temperatures which were studied were 30, 150, 300, and 400 °C ( 5%. Exposure of the samples to sulfur dioxide was accomplished by delivering SO2 (Matheson, Research Purity) to the cell via a polyethylene tube from a glass vacuum manifold. An MKS capacitance manometer attached directly to the cell measured the pressure in the cell. Adsorption of sulfur dioxide was accomplished by exposing the sample to 12.0 ( 0.1 Torr of SO2. The DRIFT spectra were obtained before and after exposure of the samples to sulfur dioxide or O2. Tests indicated that 5 min were enough for complete adsorption or desorption, although 15-20 min were used in these experiments. A spectrum of sulfur dioxide vapor was subtracted from the spectra where appropriate. Oxidation of the adsorbed sulfur dioxide was accomplished by exposing samples to O2 at 400 °C. The infrared spectra were measured using a Nicolet 5-DXB FT-IR spectrometer with an upgraded 680SX data station, an 18 bit A/D converter, a narrow-band MCT/A detector, and a high-intensity source. The spectra were measured using DRIFTS, using an optical accessory (HVC-DR2) and associated controlled environmental cell from Harrick Scientific. With this method, the maximum amount of contact between the adsorbing gas and the solid is achieved and the sample most closely mimics the actual mode in which a catalyst would be used. Results Adsorption of Sodium. The concentration of surface OH on the γ-alumina before impregnation is approximately 2.5 mmol OH/g Al2O3. If the substitutions take the place only at the OH sites, the sodium concentration should be approximately 2.5 mmol Na/g Al2O3 or 16.6 µmol Na/m2 Al2O3. The loading determined by elemental analysis for sample B was found to be 0.48 mmol Na/gr Al2O3 or 3.2 µmol Na/m2. Sample C (2:1 solution Na/surface OH mole ratio) was found to contain 9.9 µmol Na/m2, and sample D (4:1 solution Na/surface OH mole ratio) was found to contain 17.9 µmol Na/m2 (1.49 and 2.69 mmol Na/g Al2O3, respectively). Figure 1 is a plot of the adsorption isotherm for these three impregnation concentrations. As can be seen from the figure, the uptake of sodium from the solution is linear with concentration, and it does not appear that the surface concentration is approaching a limiting value. The plot does not extrapolate to zero, which may be the result of an initial, rapid uptake of
Mitchell et al.
Figure 1. Adsorption isotherm for adsorption of sodium onto alumina using sodium methoxide as the sodium source.
sodium, which is then followed by the more gradual adsorption seen in these results. If adsorption at only OH sites was the active mechanism for this adsorption process, the concentration of Na on the surface of the most highly doped sample would be very close to the predicted surface saturation. The fact that Figure 1 shows no evidence of saturation indicates that the proposed mechanism for sodium attaching to the surface is insufficient to account for all of the active sites for adsorption. It is possible for metal alkoxides to react with Lewis acid sites via the following reaction scheme.
This is the most likely alternate mechanism for the formation of surface sodium in these experiments. TGA Results. Figure 2 shows the weight percent uptake of sulfur dioxide as a function of adsorption cycle for samples A-D. This figure shows in a fairly dramatic way the effects of sodium loading. The initial uptake is much greater for the highly impregnated aluminas, but the loss of capacity with cycling is also much greater for these materials. Figure 3 is a graph of the data which shows that the uptake of SO2 is linear with sodium content for each cycle but that the slope of the line, the relative increase in SO2 uptake as a function of added sodium, decreases as the cycle number increases. The slope of the line calculated from the first adsorption cycle turns out to be 1 additional SO2 molecule adsorbed for every 2.7-2.8 sodium atoms on the surface. The slope of the line for the later adsorption/desorption cycles is obviously less than this, and the loss of capacity as a function of cycle is greatest for the highest loadings. We will return to this point later. FT-IR Results. Effect of Temperature. In all earlier studies of SO2 on alumina, the infrared measurements were carried out on pressed pellets of pure alumina. In order to evaluate the effect of the KBr diluent used in the DRIFTS experiments on the spectroscopic and absorption results, control experiments were performed. Sulfur dioxide adsorption/desorption experiments were carried out on pure KBr and on pure alumina. We were able to generate a small amount of what appeared to be K2SO4 by carrying out the experiments with pure KBr, but only insignificant amounts were generated. The spectrum for this material matched reasonably well the spectrum of potassium sulfate and was clearly different from that which we observed
Adsorption of SO2 on Na-Doped Alumina
J. Phys. Chem., Vol. 100, No. 18, 1996 7553
Figure 4. DRIFT spectra for alumina sorbent in the presence of 12 Torr of SO2 at different adsorption temperatures.
Figure 2. Percent weight changes observed for alumina and sodiumimpregnated aluminas for sequential sulfur dioxide adsorption/desorption cycles.
Figure 5. DRIFT spectra for alumina sorbent after evacuation of SO2 at different adsorption temperatures.
Figure 3. Plot of mmol of sulfur dioxide adsorbed as a function of mmol of sodium on the sorbent for different adsorption cycles.
when alumina was present. No contribution from this species is observed in these experiments. The spectra measured for SO2 on pure alumina were reasonably similar to those measured for the diluted samples, although the pure samples provided no useful information at frequencies less than 1100 cm-1 due to the strong absorption of alumina. Figure 4 shows the DRIFT spectra measured for the undoped alumina samples at the four adsorption temperatures studied under sulfur dioxide (p ) 12 Torr). They all show new absorption bands in the region 1400-1000 cm-1. SO2 is present in the cell during this measurement, and its spectrum was
subtracted from the DRIFT spectrum of the sample to yield the spectrum of the sorbent surface which is shown. The spectrum of the pristine sorbent surface is used as the background spectrum, and so the spectrum which is shown is due to adsorbed SO2. Figure 5 shows these same samples after evacuation of the SO2. A strong band at 1340 cm-1 and a weaker band at 1145 cm-1 disappear soon after evacuation is begun, whereas the other bands show virtually no change in intensity after evacuation of the gas. These two bands at 1340 and 1145 cm-1 are assigned to physisorbed SO2 and have been observed by several other groups. One of the most obvious features which remains after evacuation is a broad, intense band at approximately 1060 cm-1. This band is a chemisorbed SO2 and has also been observed by other workers. Unlike results found by Chang,9 we did not observe the disappearance of the 1060 cm-1 band from the spectra, even after evacuation of some of the samples for 24 h. As can be seen from the figure, adsorption at 150 °C yields the greatest quantity of chemisorbed SO2. As the adsorption temperature is increased, the amount of physisorbed SO2 decreases relative to the amount of chemisorbed SO2, indicated by the intensity of the band at 1060 cm-1. This is just the behavior expected for physisorbed species. Figure 6 shows the effects of increasing adsorption temperature at the highest sodium loading. As can be seen from this plot, the effect of temperature at the high sodium loading is very similar to that observed in the absence of sodium. The sorbents with the lower sodium loading exhibit very similar behaviors. Adsorption at 150 °C yields the largest amount of
7554 J. Phys. Chem., Vol. 100, No. 18, 1996
Mitchell et al.
Figure 6. DRIFT spectra for alumina sorbent with a surface coverage of 17.9 µmol/m2 sodium after adsorption of SO2 and after evacuation, at different adsorption temperatures.
Figure 8. DRIFT spectra of different sorbents for adsorption at 400 °C after evacuation of SO2. Sorbents for spectra A-D are as in Figure 7.
Figure 7. DRIFT spectra of different sorbents for adsorption at 150 °C, after evacuation of SO2. Spectrum A is the spectrum for the undoped alumina, spectrum B is for the 3.2 µmol/m2 sodium sorbent, spectrum C is for the 9.9 µmol/m2 sodium sorbent, and spectrum D is for the 17.9 µmol/m2 sodium sorbent.
Figure 9. DRIFT spectra of the different sorbents at 150 °C in the presence of SO2 to show the dependence of the physisorbed SO2 concentration on sodium loading.
adsorbed SO2, with the 30 and 300 °C adsorption temperatures yielding about the same amount of adsorbed SO2 and the 400 °C adsorption temperature yielding the smallest amount. Effect of Sodium. The effects of increasing sodium content are seen more easily if the data is plotted in a different way. Figure 7 contains four different spectra, one representing the infrared absorption spectrum for chemisorbed SO2 at each sodium loading, for adsorption at 150 °C. Figure 8 is similar, except that adsorption was carried out at 400 °C. As the sodium loading increases, the absorbance of the 1060 cm-1 band increases. In Figure 7, it can be seen that spectra from samples C and D of alkali-doped γ-alumina show a well-pronounced low-intensity band at 1215 cm-1. This band is also observed for adsorption at 30 °C. This band is not observed in the spectra of the pure γ-alumina (sample A), and it is not apparent on top of the broad absorption at 1060 cm-1 in a spectrum of sample B which contains a small amount of sodium. The sodium loading also affects the amount of physisorbed SO2. Figure 9 shows this dependence at 150 °C. This figure is similar to Figure 7 above, except the data was obtained before evacuation of the SO2. As can be seen, the amount of physisorbed SO2 increases with increasing sodium content, although not nearly as much as the chemisorbed SO2. Oxidation of Chemisorbed SO2. The spectra change significantly after oxidizing the sulfur dioxide-exposed samples at 400
Figure 10. DRIFT spectra of the sorbents for adsorption at 150 °C, after oxidation of the adsorbed SO2. Spectrum A is the spectrum for the undoped alumina, spectrum B is for the 3.2 µmol/m2 sodium sorbent, spectrum C is for the 9.9 µmol/m2 sodium sorbent, and spectrum D is for the 17.9 µmol/m2 sodium sorbent.
°C and cooling them to their target temperature (Figures 10 and 11). In this case, the broad, intense band at 1060 cm-1, observed in the spectra of all unoxidized samples, disappears and a new intense band at 1190 cm-1, with a shoulder at 1240-1250 cm-1, appears. The frequency of this band does not depend on the initial temperature of exposure of the samples to sulfur dioxide.
Adsorption of SO2 on Na-Doped Alumina
J. Phys. Chem., Vol. 100, No. 18, 1996 7555
Figure 11. DRIFT spectra of the sorbents for adsorption at 400 °C, after oxidation of the adsorbed SO2. Sorbents for spectra A-D are as in Figure 10.
TABLE 1: Summary of Frequencies (cm-1) Observed for Different Experimental Conditions with SO2 in cell after evacuation for T > 150 °C after oxidation a
1340
1215
X
Xa Xa Xa
1190
1145
1125
X X X
Xa Xa
1060 X X X
This absorption is only observed with sodium-impregnated samples.
The frequency and intensity of the band at 1190 cm-1 do not change even after evacuation for several hours or heating to 400 °C. Simultaneously with the appearance of the bands at 1190 cm-1, samples C and D show a second new band at 11201130 cm-1. This band does not appear in oxidized samples of pure γ-alumina (A) and appears only weakly, if at all, in spectra of sample B. The same weak shoulder was found for samples C and D after their exposure to sulfur dioxide at temperatures above 300 °C, even before oxidation. Table 1 summarizes all of these observations. Discussion Pure Alumina. With regard to adsorption on γ-alumina, at least three groups have concluded that the adsorption of the strongly held SO2 occurs on the basic sites, while weakly held or physisorbed SO2 is held at acidic sites. Two bands at approximately 1330 and 1150 cm-1 are observed for the physisorbed species, while an intense absorption at approximately 1060 cm-1 is due to the chemisorbed SO2.8-10 Our results are completely consistent with these earlier studies. The form of chemisorbed SO2 responsible for the 1060 cm-1 absorption has been assigned to a sulfite species attached to an aluminum through the sulfur atom.10 The results from the current study involving undoped γ-alumina which conflict with those observed by others involve the infrared spectra of the oxidized form of adsorbed SO2. There is little doubt that this form must correspond to an adsorbed sulfate. Our observed spectra for this form of adsorbed SO2 are very similar to the spectrum of bulk aluminum sulfate, with an intense absorption band at approximately 1190 cm-1. Three earlier infrared studies observed two intense bands for oxidized SO2 on alumina at 1380-1400 and 1045-1090 cm-1. Chang9 and Saur et al.11 explain the discrepancy between their observations and the spectrum of aluminum sulfate by suggesting that the two intense bands are due to a surface sulfate.
The three earlier infrared studies of oxidized SO2 on γ-alumina used pressed pellets of pure alumina as the infrared sample onto which the SO2 was adsorbed and oxidized. It may be that the action of pressing the pellets creates a particular phase of alumina which forms a sulfate distinct from that which is observed in the DRIFTS experiments which do not use pressed pellets. On the other hand, it could be that the alumina used in the current study was giving rise to a form of oxidized SO2 different from that observed in the earlier studies. In order to test these ideas, an experiment was conducted which used a pressed pellet of the Goodfellow alumina in the DRIFTS cell. The pellet was placed on top of a layer of KBr in the DRIFTS cell, heated as usual, exposed to SO2, and oxidized. The spectrum which was obtained was not one which was particularly useful; the pellet reflected very little or no radiation below 1200 cm-1 due to its thickness. However, a strong 1380 cm-1 band was observed, consistent with the other pressed pellet studies and completely different from the normal DRIFTS spectra. This provided convincing evidence that it is the mode of sample preparation, and not the nature of the sample itself, that accounts for the differences in the observed spectra of oxidized SO2 on γ-alumina. Sodium-Doped Alumina. While the interpretation of the infrared results of SO2 adsorption on undoped alumina can be compared with several studies, the spectra of SO2 on sodiumimpregnated alumina have been published by only one other group to our knowledge,7 and the results of that study do not correspond well with the results observed here. For chemisorbed SO2 on sodium-doped alumina, two bands were observed at the lower adsorption temperatures: the broad absorption at 1060 cm-1 and a weaker absorption at 1215 cm-1 (see Figure 7). The 1060 cm-1 band is assigned to an aluminum sulfite species, due to the similarity with the broad absorption observed for SO2 chemisorbed on undoped alumina. The new band at 1215 cm-1 observed in the spectra of chemisorbed SO2 on the sodium-doped samples is interpreted as indicating the presence of sodium sulfite. The vibrational spectrum of sodium sulfite has been investigated in detail.29 An intense band at 965-980 cm-1 and a less intense absorption at 1210-1230 cm-1 are commonly observed. The intense, lower frequency absorption of sodium sulfite is not observed because of the strong infrared absorption of γ-alumina at frequencies less than 1000 cm-1. Only two bands are observed in the spectra of the oxidized samples. The strong 1190 cm-1 band is assigned to an aluminum sulfate species, since its frequency and shape are identical to those observed in the spectra of oxidized, chemisorbed SO2 on undoped alumina. The shoulder at 1125 cm-1 observed for the sodium-doped samples can be correlated with the formation of sodium sulfate. Its frequency is in good agreement with a DRIFT spectrum of pure sodium sulfate generated by us and with spectra published by others,29,30 who assign this band either to the fundamental symmetric vibration of the sulfate anion or as a combination band. The 1125 cm-1 band could also be generated by oxidizing model mixtures of sodium sulfite in KBr or γ-alumina and sodium sulfite in KBr in the DRIFTS controlled environment cell. Evidence is seen for the formation of the sulfate species even before addition of oxygen. At the two highest adsorption temperatures, two new bands appear: at 1190 cm-1 for all samples and a broad, poorly-defined shoulder at 1125 cm-1 for samples C and D (Figures 10 and 11). The intensities of both bands grow with increasing adsorption temperature and sodium content. Given the above assignments, it is concluded that small amounts of aluminum and sodium sulfate are formed at adsorption temperatures in excess of 150 °C.
7556 J. Phys. Chem., Vol. 100, No. 18, 1996
Figure 12. Plot of the integrated intensity of the infrared band due to the aluminum sulfite species for each sample versus the integrated intensity of the corresponding aluminum sulfate band after oxidation.
It is believed that the sulfite or sulfate form of adsorbed SO2 may be formed depending on whether sulfur dioxide attaches to one or two oxygen atoms on the sorbent surface. A similar phenomenon was observed for adsorption on the surface of calcium carbonate.31 Disproportionation of surface sulfite species to form sulfate species and sulfide species in the absence of oxygen has been shown to occur on CaO surfaces32 and is another possible mechanism for the formation of the surface sulfate before the addition of oxygen. If any disproportionation was occurring, formation of the sulfide was not observed in the infrared spectra, due to the low frequency of the corresponding absorptions. Partitioning between Aluminum and Sodium Forms. Our results indicate that both sodium and aluminum oxide participate in the interaction with sulfur dioxide. The alumina is not simply acting as an inert support. The bands at 1215 and 1060 cm-1 indicate the formation of sulfite species of both sodium and aluminum, respectively. However, upon examination of the 150 °C adsorption spectra (Figure 7), it can be seen that there is a significant increase in the aluminum sulfite absorption intensity on going from sample A to sample B and little, if any, intensity at the frequency of the sodium sulfite band. This is clear indication that the presence of sodium at low loadings promotes the formation of an aluminum sulfite species, rather than sodium sulfite. Further increases in sodium content have less dramatic effects on the intensity of the aluminum sulfite absorption, and the formation of sodium sulfite is clearly indicated by the presence of the 1215 cm-1 absorption. Oxidation converts the sulfite species to sulfate species. The near-quantitative conversion of the aluminum sulfite species (1066 cm-1) to aluminum sulfate (1190 cm-1) by oxidation is demonstrated by the linear correlation observed between the intensities of the corresponding infrared absorption bands (Figure 12). This figure is a plot of the integrated intensity of the absorption band corresponding to the sulfate species versus the integrated intensity of the sulfite band, at each temperature and sodium loading. Points on the graph corresponding to the different sodium loadings tend to cluster together on the graph, indicating that the sodium loading has an overall greater effect on uptake than does adsorption temperature. However, for the two highest loadings, the sulfate and sulfite uptakes overlap almost completely as a function of temperature. For all loadings, the 150 °C adsorption temperature yields the greatest uptake. At the two highest sodium loadings, 30 °C adsorption tends to
Mitchell et al.
Figure 13. DRIFT spectra of the different sorbents after oxidation of SO2 adsorbed at 150 °C and after subtraction of the contribution from the aluminum sulfate species, using the scaled spectrum from sample A. The spectra correspond to the sodium sulfate species formed on oxidation.
Figure 14. Plot of the formation of sodium sulfate, using the integrated infrared intensities as a measure of the concentration, versus the surface sodium concentration.
give greater amounts of sulfate formation relative to sulfite than do the other temperatures. The remarkable similarity between Figures 10 and 11, and the virtually identical sets obtained for adsorption at 30 and 300 °C, indicates that the partitioning of the oxidization products of the adsorbed SO2 between an aluminum sulfate form and a sodium sulfate form does not depend strongly on temperature and only appears to be a function of the sodium loading. Assuming that the sulfite species do not form new cation associations during oxidation, and Figure 12 indicates that this is a valid assumption, this observation can be extended to the initial chemisorbed SO2. The spectrum of sample A after oxidation can be subtracted from the spectra of samples B-D after oxidation to yield the spectrum of the sodium sulfate form. The result of this subtraction is shown in Figure 13. Since integrated areas can be used as quantitative measures of the concentration of the species, the integrated areas of these peaks were plotted versus sodium loading to yield Figure 14. The line on Figure 14 is a first-order fit to the three points shown on the graph. The three points fit the line quite well, and the results indicate that for sodium coverages less than about 2.6 µmol/m2, it is predicted that no sodium sulfate will be formed on the surface.
Adsorption of SO2 on Na-Doped Alumina One interpretation of these results is that at low loadings, the sodium atoms fill a set of preferred sites, which by their nature do not facilitate the formation of sodium sulfate. At higher loadings, a second type of sodium site is created, which does allow the formation of sodium sulfate. One possibility is that in this second type of site, two or more sodium atoms are in close proximity on the surface, which could lead to the formation of Na2SO4. In the absence of this second type of site, the oxidized form of adsorbed SO2 presumably exists on the sorbent as Al2(SO4)3 in the absence of sodium, perhaps with the formation of some amount of NaAl(SO4)2 when sodium is present at low loadings. If this latter species is present to any significant extent on the surface, it must have an infrared spectrum very similar to that of Al2(SO4)3. Conclusion Impregnating γ-alumina with sodium increases the numbers of both the basic and the acidic sites, as demonstrated by an increase in both physisorption and chemisorption of SO2 as a function of sodium loading. This result has been observed in the past using other titration methods to examine basic and acidic sites on alkali-doped aluminas.15 From our results it appears that two different forms of adsorbed SO2 are formed on sodium-doped alumina. The predominant form at low loading corresponds to a species which has a structure similar to that of aluminum sulfate. This form of adsorbed SO2 can be decomposed, and the adsorption sites can be regenerated. The second form of adsorbed SO2 has a structure similar to that of bulk sodium sulfate and leads to adsorption sites which are not regenerated under the conditions used in our TGA experiments. At low sodium loadings, approximately 1wt/wt %, it appears that sodium acts as a strong promoter for the aluminum sulfite and sulfate-like species but that the formation of sodium sulfite and sulfate is minimal, if it occurs at all. Increasing the sodium loading beyond these levels yields initially higher levels of SO2 uptake, but the formation of sodium sulfate causes the increased capacity to be rapidly lost as the material is cycled between adsorption and regeneration. Oxidation of the sulfite species leads to the quantitative formation of sulfate, and the structural identities of the species appear to remain intact. That is, those species which start out with a structure similar to aluminum sulfite are converted to species which look like aluminum sulfate, while those species which look like sodium sulfite before oxidation look like sodium sulfate after oxidation. Partitioning of adsorbed SO2 between aluminum and sodium forms of sulfite, before oxidation, or sulfate, after oxidation, depends on the sodium loading and does not appear to depend strongly on the adsorption temperature. The use of pressed pellets to examine the adsorption characteristics of SO2 on γ-alumina using transmission infrared spectroscopy generates significantly different infrared spectra for the oxidized form of chemisorbed SO2 than do DRIFTS studies using powdered samples. Whether or not this is an important distinction depends on whether or not it forces a
J. Phys. Chem., Vol. 100, No. 18, 1996 7557 change in the proposed mechanisms or identifications. However, it is certainly a result that must be noted. Acknowledgment. The authors gratefully acknowledge support from DOE Grants DE-FG22-90PC90292 and DE-FG2292MT92017. The financial assistance of the EPA through Cooperative Assistance ID No. CR 818689 is acknowledged. References and Notes (1) Lowell, P. S.; Schwitzgebel, K.; Parsons, T. B.; Sladek, K. J. Ind. Eng. Chem. Process Des. DeV. 1971, 10, 384. (2) Town, J. W.; Paige, J. I.; Russell, J. H. Chem. Eng. Prog., Symp. Ser. 1970, 105, 260. (3) Schlesinger, M. D.; Illig, E. G. Chem. Eng. Prog., Symp. Ser. 1971, 67, 46. (4) Gavalas, G. R.; Edelstein, S.; Flytzani-Stephanopoulos, M.; Weston, T. A. AIChE J. 1987, 33, 258. (5) Ma, W. T.; Haslbeck, J. L. EnViron. Prog. 1993, 12, 163. (6) Ma, W. T.; Haslbeck, J. L.; Neal, L. G.; Yeh, J. T. Sep. Technol. 1991, 1, 195. (7) Deo, A. V.; Dalla Lana, I. G.; Habgood, H. W. J. Catal. 1971, 21, 270. (8) Karge, H. G.; Dalla Lana, I. G. J. Phys. Chem. 1984, 88, 1538. (9) Chang, C. C. J. Catal. 1978, 53, 374. (10) Datta, A.; Cavell, R. G.; Tower, R. W.; George, Z. M. J. Phys. Chem. 1985, 89, 443. (11) Saur, O.; Bensitel, M.; Mohammed Saad, A. B.; Lavalley, J. C.; Tripp, C. P.; Morrow, B. A. J. Catal. 1986, 99, 104. (12) Preud’homme, J.; Lamotte, J.; Janin, A.; Lavalley, J. C. Bull. Soc. Chim. Fr. I 1981, 433. (13) Nam, S. W.; Gavalas, G. R. Appl. Catal. 1989, 55, 193. (14) Dalla Lana, I. G.; Karge, H. G.; George, Z. M. J. Phys. Chem. 1993, 97, 8005. (15) Hombek, R.; Kijenski, J.; Malinowski, S. Preparation of Catalysts, II; Delmon, B., et al., Eds.; Elsevier: Amsterdam, 1979; p 595. (16) Dmitruk, G. F.; Yashchenko, N. N. Ukr. Khim. Zh. 1986, 52, 1015. (17) Neiman, A. Y.; Tigeeva, I. E. Russian J. Inorg. Chem. 1986, 31, 490. (18) Waqif, M.; Saad, A. M.; Bensitel, M. J .Chem. Soc., Faraday Trans. 1992, 88, 2931. (19) Kordulis, C.; Vordonis, L.; Lycourghiotis, A.; Pomonis, P. J. Chem. Soc., Faraday Trans. 1987, 83, 627. (20) Kordulis, C.; Voliotis, S.; Lycourghiotis, A.; Vattis, D. Appl. Catal. 1984, 11, 179. (21) Duisterwinkel, A. E.; Doesburg, E. B. M.; Hakwoort, G. Thermochim. Acta 1989, 141, 51. Hakwoort, G.; Van Den Bleek, C. M.; Schouten, J. C.; Valkenburg, P. J. M. Thermochim. Acta 1989, 114, 103. (22) Duisterwinkel, A. E.; Hakvort, G. Bull. Soc. Chim. Belg. 1989, 98, 439. (23) Meyer, A. European Pat. 0090994, 1987. (24) Beckler, R. K.; White, M. G. J. Catal. 1986, 102, 252. Beckler, R. K.; White, M. G. J. Catal. 1988, 112, 157. Babb, K. H.; White, M. G. J. Catal. 1986, 98, 343. Beckler, R. K.; White, M. G. J. Catal. 1988, 109, 25. (25) Rao, S. N. R. Master’s Thesis, Georgia Institute of Technology, Aug 1993. Rao, S. N. R.; E. Waddel, E.; Mitchell, M. B.; White, M. G. Submitted for publication in J. Catal. (26) Kenvin, J. C.; White, M. G.; Mitchell, M. B. Langmuir 1991, 7, 1198. Kenvin, J. C.; White, M. G. J. Catal. 1991, 130, 447. Kenvin, J. C. Ph.D. Thesis, Georgia Institute of Technology, Atlanta, GA, 1991. (27) Mitchell, M. B.; Chakravarthy, V. R.; White, M. G. Langmuir 1994, 10, 4523. (28) Kno¨zinger, H.; Ratnasamy, P. Catal. ReV. - Sci. Eng. 1978, 17, 31. (29) Hajek, B.; Muck, A.; Santava, E. Collect. Czech. Chem. Communs. 1985, 50, 2715. (30) Darue, R. A.; Milne, J. W. Spectrochim. Acta, Part A 1978, 34, 215. (31) Thompson, M. M.; Palmer, R. A. Appl. Spectrosc. 1988, 42, 945. (32) Marshi, D. W.; Ulrichson, D. L. Chem. Eng. Sci. 1985, 40, 423.
JP9519225
