Microcalorimetric Characterization of Structural and Chemical

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Langmuir 1997, 13, 888-894

Microcalorimetric Characterization of Structural and Chemical Heterogeneity of Superacid SO4/ZrO2 Systems† V. Bolis,* G. Magnacca, G. Cerrato, and C. Morterra Dipartimento di Chimica Inorganica, Chimica Fisica e Chimica dei Materiali, Universita` di Torino, Via P. Giuria 9, 10125 Torino, Italy Received October 26, 1995X The population of coordinatively unsaturated (cus) Zr4+ ions acting as strong Lewis acid sites on superacid sulfated zirconia catalysts, as well as the relevant heats of adsorption, were measured by means of CO adsorption at 303 K. The presence of charge-withdrawing SO4 groups slightly enhances the electronaccepting ability of the residual free surface cations, not involved in sulfate/metal complexes. H2O adsorbed at 303 K by a stepwise procedure progressively eliminates the Lewis acidic sites, through a mechanism that suggests a competition in strongly binding H2O between cus Zr4+ cations and SO4 groups. Quantitative and energetic data are reported and correlated to IR spectroscopic and catalytic data previously reported on the same catalysts. A model for the catalytic active site is proposed, based on the assumption that a strict cooperation between Lewis and Brønsted acidic sites is likely to exist.

Introduction Sulfate-promoted zirconia (SO4/ZrO2) is a potential solid superacid, in that it may exhibit an acidic strength stronger than 100% H2SO41 being thus active in the isomerization of n-butane at low temperature.2 The Hammet H0 acidic function has been reported to be as low as -16.3 Solid superacids are of greatest interest in catalysis. In fact, as superacids they are able to lower dramatically the reaction temperature, with consequent energy saving and increased selectivity to branched alkanes. As solid catalysts they are less environmentally hazardous and far more easily separated from the reaction products stream than the liquid ones. Not all SO4/ZrO2 systems are superacidic. A number of preparative variables must be carefully controlled; a great deal of work has been made in order to assess the best preparative conditions giving rise to superacidic materials.4-7 The effect of sulfate promotion on physicochemical and textural properties of the actual catalyst has been deeply examined.8-14 In particular, Morterra et * To whom correspondence should be addressed: e-mail, [email protected]; telephone, 39.11.670.7565; fax, 39.11.670.7855. † Paper presented at the Second International Symposium Effects of Surface Heterogeneity in Adsorption and Catalysis on Solids, held in Poland/Slovakia, September 4-10, 1995. X Abstract published in Advance ACS Abstracts, September 15, 1996. (1) Gillespie, R. J. Acc. Chem. Res. 1968, 1, 202. (2) (a) Hino, M.; Kobayashi, S.; Arata, K. J. Am. Chem. Soc. 1979, 101, 6439. (b) Hino, M.; Arata, K. J. Chem. Soc., Chem. Commun. 1980, 101, 851. (3) Misono, M.; Okuhara, Y. CHEMTECH 1993, November, 23. (4) Nascimento, P.; Akrotopoulou, C.; Oszagyan, M.; Coudurier, G.; Travers, C.; Joly, J.-F.; Ve´drine, J. C. Proceedings of the 10th International Congress on Catalysis; Guzci, L., Solymosi, F., Tetenyi, P., Eds.; Akade´miai Kiado`: Budapest, 1993; Vol. B, p 1185. (5) Morterra, C.; Bolis, V.; Cerrato, G.; Pinna, F.; Signoretto, M.; Strukul, G. Europacat-I (Montpellier, Sept 12-17, 1993), Commm, I-77, p 150-I. (6) Morterra, C.; Cerrato, G.; Bolis, V. Catal. Today 1993, 17, 505. (7) Morterra, C.; Cerrato, G.; Emanuel, C.; Bolis, V. J. Catal. 1993, 142, 349. (8) Parera, J. M. Catal. Today 1992, 15, 481. (9) Chen, F. R.; Coudurier, G.; Joly, J.-F.; Vedrine, J. C. J. Catal. 1993, 143, 616. (10) Morterra, C.; Bolis, V.; Cerrato, G.; Magnacca, G. Surf. Sci. 1994, 307, 1206. (11) Lunsford, J. H.; Sang, H.; Campbell, S. M.; Liang, C.-H.; Anthony, R. G. Catal. Lett. 1994, 27, 305. (12) Clearfield, A.; Serrette, G. P. D.; Khazi-Syed, A. H. Catal. Today 1994, 20, 295.

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al.15,16 have put in relation the structural and morphological features (studied by both XRD and HRTEM) and the surface properties (monitored by IR spectroscopy) of variously prepared materials with their catalytic activity toward n-butane at T ) 423 K. From their results it appears that the following features are required for the onset of superacidity in sulfated zirconia: (i) the tetragonal structure is preferred (but no experimental evidence exists that the monoclinic phase is not active); (ii) the optimum sulfate groups loading is comprised in the one-half to a (nominally) complete monolayer interval; (iii) the best calcination temperature of the sulfated amorphous precursor oxide is comprised in the very narrow interval of 823-923 K (the temperature must be sufficiently high to allow the formation of regular microcrystals, but must not exceed the decomposition temperature of the sulfates, i.e., T ) 973 K); (iv) the activation temperature of the catalyst must be at least 573 K, in order to reach the maximum dehydroxylation compatible with the stability of the sulfates. Several of the preparative variables leading SO4/ZrO2 systems to superacidity are now well established, but the actual nature of the catalytically active site is not yet fully understood. The main question under debate is whether the superacidic behavior is ascribable to the presence at the surface of the promoted oxide of Lewis or of Brønsted acidic sites or, perhaps, to the reciprocally assisted presence of both. The simultaneous presence of both types of sites has been proposed as necessary for best catalytic activity.4,14,16 Strong Lewis acidity at the surface of SO4/ZrO2 activated at T g473 K is not surprising, because of the presence of coordinatively unsaturated (cus) Zr4+ cations, bearing a high electron accepting ability (as generally do transition metal cations, as well as some non d metal cations such as Al3+ 17,18). The intrinsic electron accepting ability of the cus Zr4+ cations is expected to be enhanced by the nearness of the charge-withdrawing SO4 groups, that cause the actual charge of the cations and, consequently, (13) Ward, D. A.; Ko, E. I. J. Catal. 1994, 150, 18. (14) Babou, F.; Coudurier, G.; Ve´drine, J. C. J. Catal. 1995, 152, 341. (15) Morterra, C.; Cerrato, G.; Pinna, F.; Signoretto, M. J. Phys. Chem. 1994, 98, 12373. (16) Morterra, C.; Cerrato, G.; Pinna, F.; Signoretto, M.; Strukul, G. J. Catal. 1994, 149, 181. (17) (a) Bolis, V.; Morterra, C.; Fubini, B.; Ugliengo, P.; Garrone, E. J. Chem. Soc., Faraday Trans. 1992, 88, 391. (b) Bolis, V.; Morterra, C.; Fubini, B.; Ugliengo, P.; Garrone, E. Langmuir 1993, 9, 1521. (18) Morterra, C.; Bolis, V.; Magnacca, G. Langmuir 1994, 10, 1812.

© 1997 American Chemical Society

Heterogeneity of Superacid SO4/ZrO2 Systems

their electron-accepting capacity to increase, as shown by the following:

This effect is well-known and it was studied in detail by some of us in the case of some d0 metal cations by coupling IR spectroscopic and microcalorimetric techniques.17,19,20 As for the Brønsted acidity, some controversy does exist about its assignment to hydrated zirconium ions13,21

where the protonated zirconium hydroxyls are rendered more acidic than those in pure zirconia owing to the chargewithdrawing effect of the adjacent sulfate group. The Brønsted acidity may, at the opposite, be ascribed to the protonated sulfates as suggested by Lunsford et al.:11

a hypothesis also supported by other authors.14,22-24 The spectral features of superacidic tetragonal sulfated zirconia specimens have been deeply examined and discussed by Morterra et al.15 From the inspection of the spectral region in the interval 4000-900 cm-1, they revealed the formation of hydronium ion at the surface of the hydrated sample. The hydronium species could be revealed by IR spectra because they were virtually free from H-bonding interactions, in that they were sufficiently separated from one another. They were found to resist thermal destruction for outgassing at T e 473 K. The authors observed also a parallel evolution with outgassing temperature of the spectral features of the sulfate groups, that were observed to move from an ionic (SO42-) to a covalent (OdSdO) configuration. The presence of some residual Brønsted acidic sites at the surface after activation at rather high temperature (T ) 723-773 K, typical activation temperature for the working catalyst) was put in evidence by the formation, upon contact with pyridine at room temperature, of small but always detectable amounts of pyridinium ion. This behavior is peculiar of sulfated zirconia, in that the pure oxide outgassed at such temperatures does not exhibit Brønsted acidity and is (19) Bolis, V.; Fubini, B.; Garrone, E.; Morterra, C. J. Chem. Soc., Faraday Trans. 1 1989, 85, 1383. (20) Garrone, E.; Bolis, V.; Fubini, B.; Morterra, C. Langmuir 1989, 5, 892. (21) Arata, K. Adv. Catal. 1990, 37, 165. (22) Babou, F.; Bigot, B.; Sautet, P. J. Phys. Chem. 1993, 97, 11501. (23) Babou, F.; Coudurier, G.; Vedrine, J. C. J. Chim. Phys. 1995, 92, 1457. (24) Adeeva, V.; de Haan, J. W.; Ja¨nchen, J.; Lei, G. D.; Scu¨nemann, V.; van de Ven, L. J. M.; Sachtler, W. M. H.; van Santen, R. A. J. Catal. 1995, 151, 364.

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virtually completely dehydroxylated.25 In the OH stretching frequency region, a couple of bands at 3760 and 3640 cm-1, similar to those found on the pure oxide (of both monoclinic and tetragonal structures25), were extremely resistant to the thermal treatments in the case of the sulfated specimen, being still present after outgassing at T ) 723 K. For this reason it is virtually impossible to study the surface Lewis acidic properties and the catalytic behavior of a specimen deprived of any species suspected to exhibit Brønsted acidic activity. In fact, to have a sulfated zirconia surface completely dehydroxylated, one should activate it at a quite high temperature, certainly higher than the temperature at which the sulfate groups start to decompose (T ) 973 K). On the other hand, the same authors reported in a further paper16 that the activity of the catalyst dropped immediately to zero if small amounts of H2O were introduced in the catalytic bed: water appeared to poison irreversibly the catalyst. It was supposed to eliminate the strong Lewis acidity because of the rehydroxylation of the cus Zr4+ cations. In fact the initial catalytic activity was restored only after further outgassing the sample at T g 573 K and, thus, by destroying hydronium species and appreciable amounts of surface hydroxyls. The authors reported also that if a CO stream instead of an inert gas was used as a vector in the catalytic reactor, the reaction immediately stopped and started again only when the inert gas was re-employed as the proper vector. Thus, the presence of CO too does poison the catalyst, but in this case reversibly: CO is a soft Lewis base that is able to block the strong Lewis acidic sites not permanently, but making them inaccessible to the alkane molecules. This fact conclusively demonstrated that the presence of Lewis acidity is necessary for catalytic activity. However, no experimental evidence has ever been reported to either confirm or definitely discard the hypothesis of the catalytic utility of a Brønsted acidity. Ve´drine et al. suggest that HSO4- groups at the surface are actually responsible for the catalytic activity,14 but they seem to assign a role also to the Lewis cationic sites.23 Summarizing, the state of the knowledge on this subject seems to be that small quantities of residual water are needed in order to allow the catalytic reaction to occur by ensuring the presence of some Brønsted acidity, but it is not demonstrated (if not in the case of unsulfated zirconia) that without Brønsted acidic sites the catalytic reaction does not occur. A slight excess of water causes the activity to drop immediately to zero, destroying Lewis acidic sites, owing to the rehydroxylation of the cus Zr4+ cations. The suggestion arising from all data available in literature is that a cooperation between the two kinds of acidic sites is the most likely hypothesis. In order to give a contribution to the debate, the affinity of superacid sulfated zirconia toward water was studied by adsorbing at 303 K successive very small doses of H2O vapor. The Lewis acidity of the catalyst was monitored by adsorbing CO at the same temperature, first on the highly dehydroxylated surface (i.e., on the working catalyst) and then on the same surface, but progressively gradually rehydrated, in order to clarify the interplay of aprotic and protonic acidic sites. Quantitative and energetic data relative to the adsorption of CO and H2O were obtained by microcalorimetry, and correlated to the IR spectroscopic15 and catalytic16 data previously reported and summarized above. The evolution of the energy of interaction upon increasing coverage of CO and of H2O allowed putting in evidence (25) Morterra, C.; Cerrato, G.; Ferroni, L.; Montanaro, L. Mater. Chem,. Phys. 1994, 37, 243.

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the heterogeneity of the surface, that is, of both structural and chemical nature. Experimental Section Materials. Two sulfated ZrO2 samples, designated in the text as ZS8(823) and ZS8(923) were studied. The figures in brackets indicate the temperature T1 at which the catalyst (prepared by a sol-gel technique described in ref 14) was first fired in a dry air stream. The amount of surface sulfates was 2.6 SO4/nm2 on ZS8(823) (BET surface area 188 m2/g) and 1.4 on ZS8(923) (130 m2/g) and was evaluated as described in ref 26. The phase of the materials was checked by XRD and was pure tetragonal for both samples. Both catalysts were checked to be active toward the isomerization of n-butane at low temperature (T ) 423 K) as required for being classified as superacid.3,21 The details on the catalytic reaction conditions and results are described in ref 16. CO Specpure Matheson was employed; H2O was distilled in vacuo and rendered gas-free by several “freeze-pump-thaw” cycles. Methods. Microcalorimetry was employed to evaluate the enthalpy changes (∆adsH) related to the adsorption of CO and H2O as probe molecules on the catalyst surface: the heats of adsorption (qads ) -∆adsH) were measured at 303 K by means of a heat-flow microcalorimeter (standard Tian-Calvet type, Setaram-France) following a well-established stepwise procedure, previously described.19,27 The calorimeter was connected to a high vacuum gas-volumetric glass apparatus, which enabled simultaneous determination of the adsorbed amounts (na) and the heats evolved (Qint) for small increments of the adsorptive. The adsorbed amounts will be reported per unit surface area as a function of the equilibrium pressure of the adsorptive (na vs p) and will be taken as a measure of the population of the sites active toward the probe molecule considered. The differential heats of adsorption will be reported as a function of coverage (qads vs na). The initial heat value (q0) could be estimated from the slope at zero coverage of the curve of the integral heats evolved, plotted as a function of the adsorbed amounts (Qint vs na)]; the (Qint vs na) plots are not reported for the sake of brevity. The differential heats of adsorption will be taken as a measure of the energy of interaction between the probe molecule and surface sites. A first run (I) of adsorption was performed on the samples previously outgassed at a given temperature (T2, K), chosen in order to reach the desired degree of dehydration of the surface. The reversible component of the adsorptive was then evacuated by pumping off at the calorimeter temperature. A second run (II) was then performed in order to evaluate the contribution to the total adsorption of the reversible component. In the case of CO, the total adsorption was found to be entirely reversible and thus we will refer to the first and second runs as equivalent. The samples were thermally pretreated in vacuo for 2 h at the chosen temperature and then put in contact with some 100 Torr (1 Torr ) 133.3 N m-2) of O2, in order to ensure stoichiometry and to avoid the presence of reduced Zr3+ centers. After that, the samples were evacuated, cooled to room temperature and transferred into the calorimeter (kept isothermally at 303 K) without further exposure to the atmosphere. In one case [ZS8(823)] the adsorption of CO was first performed on the bare surface of the sample evacuated at 773 K; then, a very small amount (≈0.01 µmol/m2) of water vapor was admitted on the sample, previously evacuated of the fully reversible adsorbed CO. After the irreversible adsorption of the first dose of water was accomplished, another dose of CO (chosen in order to reach an equilibrium pressure of ≈30 Torr) was admitted on the partially rehydrated surface. The fully reversible adsorbed CO was then evacuated, and a second dose of water was allowed to be irreversibly adsorbed, causing a further small region of the surface to be rehydrated. This procedure was followed until the adsorption of the last dose of water was, at least in part, reversibly adsorbed. The irreversibility of the adsorption of water was checked by pumping off at the calorimeter temperature at each (26) Sarzanini, C.; Sacchero, G.; Pinna, F.; Signoretto, M.; Cerrato, G.; Morterra, C. J. Mater. Chem. 1995, 5 (2), 353. (27) Bolis, V.; Morterra, C.; Volante, M.; Orio, L.; Fubini, B. Langmuir 1990, 6, 695.

Figure 1. Adsorption of CO on ZS8(823)723 ([, ]), ZS8(823)423 (2), and ZS8(923)723 (9, 0): (a) volumetric isotherms (na vs pCO); (b) heat of adsorption as a function of coverage (qads vs na); full symbols, first run; open symbols, second run. step of the process: if neither calorimetric nor volumetric response was detected, the adsorption was considered (virtually) fully irreversible. When an endothermic calorimetric signal was detected upon evacuation, indicating that a no longer negligible desorption was occurring, the procedure was stopped. A 0-100 Torr transducer gauge (Baratron MKS) was employed to measure the pressure of the dose of gas (or vapor) admitted on the sample in the calorimetric cell of known volume, as well as the pressure of the gas phase in equilibrium with the solid.

Results and Discussion Adsorption of CO at 303 K: Strong Lewis Acidity. In Figure 1 data relative to the adsorption of CO on ZS8(T1) samples activated in vacuo at T ) 423 K and 723 K are shown. In Figure 1a the volumetric isotherms (na vs pCO) are reported, and in Figure 1b the heats of adsorption as a function of coverage (qads vs na) are reported. CO is known to σ-coordinate onto coordinatively unsaturated d0 metal cations (such as Zr4+ cations), as witnessed by a stretching frequency of adsorbed CO higher than that of CO gas (2143 cm-1).17-19 In the case of sulfated ZrO2, as well as in the case of unsulfated ZrO2,10,27 the interaction has been reported to be entirely reversible, as shown by adsorption-desorption cycles monitored by IR spectroscopy.16 CO is a soft Lewis base that interacts at room temperature only with the strongest fraction of surface acidic centers: the adsorbed amounts reported as adsorption isotherms (Figure 1a) may be thus taken as the direct measure of the population of strong Lewis sites present at the surface, i.e., of the sites most likely to be involved in the catalytic reaction.16

Heterogeneity of Superacid SO4/ZrO2 Systems

From inspection of Figure 1 the following can be observed: (i) The adsorption is confirmed to be fully reversible in all cases, as indicated by the close overlap of the first and second run curves. (ii) For a given activation temperature (i.e., dehydration degree) the higher the loading of sulfate groups, the lower the population of sites (per unit surface area) sufficiently acidic to bind CO at room temperature. A ≈10% of difference in the amounts of CO adsorbed does exist (for pCO g 10 Torr) between the two samples activated at 723 K. This difference corresponds to ≈0.05 µmol/m2, i.e., to ≈0.03 molecule/nm2; this means that this amount of sulfate is bound to Zr4+ cations that, when free, are strong enough to bind CO at room temperature, otherwise no differences would be observed in the two curves. (iii) Opposite to what observed at higher pressure, at the very beginning of the process (pCO e 1 Torr), the two isotherms overlap very closely, indicating the substantial identity of the strongest acidic sites in the two samples. (iv) The curves of the adsorption isotherms do not reach, in the pressure interval examined, the plateau corresponding to the accomplishment of the monolayer, as indicated by the curves still growing at the highest equilibrium pressure of CO attained; however, the monolayer values can be roughly evaluated by extrapolating the curves to a reasonable plateau. The estimated figures and ≈0.55 µmol/m2 for ZS8(923)723, ≈0.50 µmol/m2 for ZS8(823)723, and ≈0.06 µmol/m2 for ZS8(823)423, equivalent to ≈0.33, ≈0.30, and ≈0.036 CO molecules per nm2, respectively. Thus, even on the most populated surface [ZS8(923)723] the population of strongly acidic cus cations, measured from the amount of CO adsorbed at room temperature under the reasonable hypothesis that each adsorbed molecule corresponds to an acidic site, is actually very low. It is some 10-15% of the total cus cations available. (v) If the surface is highly hydrated, as in the case of ZS8(823)423, the population of Lewis sites is dramatically lower than in the case of a highly dehydroxylated surface, as expected. The working catalyst contacted by water, even at the temperature of work (T ) 423 K), was reported to behave as if it was poisoned:16 in fact the strong Lewis acidity is almost completely absent in the presence of water irreversibly held at T ) 423 K. Moving to Figure 1b, where the energetics of the process are illustrated, it can be noticed that the heats of adsorption for ZS8(823)723 and ZS8(923)723 decrease with coverage in a similar way, starting from an initial value (q0) of ≈70 kJ/mol and approaching a value of ≈50 kJ/mol. The latter value is very close to the value obtained for the (partially) hydrated ZS8(823)423 surface, that remains nearly constant in the whole interval of coverage examined. The heats of adsorption are related to the energies of the acid-base interaction and, in this case, may be taken as the measure of the Lewis acidity of the coordinatively unsaturated cations. The adsorption enthalpy was found to be correlated (in the case of d0, d10, and non-d metal cations, i.e., in all cases in which π back donation is prevented) to the blue shift of the stretching frequency of CO adsorbed with respect to the frequency of the gas.17-19 The acidity of cus Zr4+ cations on sulfated zirconia is enhanced by the presence in the nearness of the sulfates groups, as indicated by the initial heat of adsorption (slightly) higher than that measured in the case of pure zirconia: q0 ≈ 70 kJ/mol on sulfated zirconia instead of 62-65 kJ/mol on the unsulfated one.10,17 The Lewis acidity of cus Zr4+ cations is high, but not extremely high. In particular, with respect to pure zirconia samples, that do not exhibit activity toward isomerization of n-butane nor other superacidic properties; the energy of interaction is

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Figure 2. Population of CO sites (per unit surface area) on ZS8(823)723 as a function of the equilibrium pressure ([) and at pCO ≈ 30 Torr on the surface progressively stepwise rehydrated (]).

(in the whole interval of coverage examined) only less than 10 kJ/mol higher. The blue shift of the stretching frequency of CO adsorbed on sulfated ZrO2 is found to be no more than ≈10 cm-1 higher than in the case of pure ZrO2.10,16,28 Any comparison among different samples of sulfated and unsulfated zirconia, must be made with caution, in that substantial differences may exist in both structure and morphology of the samples. Both factors play in fact a very important role in affecting the surface acidic or superacidic properties of zirconia materials. The results discussed above indicate, however, that the Lewis acidity of sulfated zirconia does not appear to be as dramatically strong as one could expect, considering that the presence of sulfates makes such a great difference in the catalytic behavior of the oxide. Progressive Rehydration of the Surface: Gradual Elimination of Strong Lewis Acidity. The surface of ZS8(823)723 was progressively rehydrated by contacting the surface with very small amounts of water (0.01-1 molecule/nm2) following a stepwise procedure described in the Experimental Section, and already reported in ref 27 as employed for a specimen of pure monoclinic zirconia. Each dose of water put in contact with the initially dehydroxylated surface was irreversibly held (as witnessed by the absence of any equilibrium pressure and by the absence of desorption upon outgassing): the OH groups (thermally eliminated by outgassing at T ) 723 K) were supposed to be restored on cus Zr4+ sites, so causing the cus surface cations to lose their acidic properties. The residual Lewis acidity at each step of the gradual rehydration was monitored by the (reversible) adsorption of CO. The quantitative results obtained are reported in Figure 2 and referred to the adsorption isotherm of CO on the bare surface. As expected, the population of strong Lewis sites was progressively depressed upon rehydration, and eventually it was suppressed. The amount of water required to eliminate the strong Lewis acidity was ≈7 µmol/m2; in ref 16 it is reported that the catalyst was already irreversibly poisoned, with consequent suppression of the catalytic activity, by the contact with ≈3 µmol/ m2 of water. In Figure 3 the number of residual strong CO sites is reported as a function of the total number of water molecules irreversibly adsorbed up to that point. The relationship is not linear, opposite to that found for pure monoclinic zirconia.27 In the earliest stages of the (28) Morterra, C.; Cerrato, G.; Bolis, V.; Lamberti, C. J. Chem. Soc., Faraday Trans. 1995, 91 (1), 113.

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Figure 3. Population of CO sites (per unit surface area) on ZS8(823)723 as a function of the amount of H2O molecules irreversibly preadsorbed (per unit surface area) on the surface. The arrow points out the residual Lewis acidity on the surface after the adsorption of an amount of water equivalent to the one sufficient to suppress the catalytic activity when introduced into the catalytic bed (at 423 K), as reported in ref 16.

interaction of water with the sulfated surface, the elimination of the Lewis acidic sites was quite rapid. But, as the rehydration of the surface proceeded, more and more water was needed to eliminate the residual strong Lewis sites. At the beginning of the stepwise process, ≈5 H2O molecules were needed to suppress 1 Lewis site, whereas at the end of the process (when water admitted on the sample was no longer only irreversibly held at the surface), ≈20 H2O molecules were irreversibly adsorbed per Lewis site eliminated. Adsorption of H2O Vapor: Mechanism of Rehydration of the Surface. The adsorption of water was studied at 303 K on variously dehydrated surfaces, i.e., on sulfated surfaces outgassed at two different temperatures, 423 and 723 K, in order to investigate (on the basis of their energetics) the nature of the sites able to bind water present on surfaces in different stages of hydration. The 723 K outgassing temperature was obviously chosen (as in the case of CO adsorption) because that is the temperature at which the working catalyst is usually activated. The 423 K temperature was chosen because at this temperature the surface hydronium species are (almost) completely destroyed, although a virtually unmodified amount of acidic protons is still present at the surface, as clearly evidenced by the IR spectra of adsorbed pyridine reported in refs 15 and 16. The interaction with a completely hydrated surface was also studied by performing a second run of adsorption after the first run of both the -723 and -473 samples, i.e., on the rehydrated samples simply outgassed at the calorimeter temperature, 303 K. The adsorption process was studied in a very narrow interval of water vapor pressures, because we were interested in the earliest contact of water with the sulfated surface. In the case of the -723 sample, however, before outgassing and performing the second run, a virtually complete rehydration of the surface was achieved by keeping overnight the sample under ≈5 Torr of water vapor. The affinity toward water of this kind of surfaces was observed by IR spectroscopy to be extremely high.15 In Figure 4 data relative to the adsorption of water on ZS8(823)723 are reported. The volumetric isotherms refer to the first and second runs of adsorption (Figure 4a). The difference between the two curves is quite large and indicates that the first contact of water with the bare surface gave rise to a large amount of adsorbed species that were not removed from the surface by simple pumping

Figure 4. Adsorption of H2O on ZS8(823)723 for ([) first run and (]) second run: (a) volumetric isotherms (na vs pH2O). (b) heat of adsorption as a function of coverage (qads vs na).

off at 303 K. Also the shape of the two curves differs significantly: the curve relative to the first run sharply increases at the beginning, indicating that the adsorption was proceeding with a very small increase of the equilibrium pressure. This means that water molecules either adsorbed dissociatively or remained tightly bound to the surface, without participating in the gas phase equilibrium. After a short plateau the curve keeps again increasing almost vertically, but at pH2O ≈ 1 Torr the adsorption isotherm starts increasing more smoothly. This kind of shape in the adsorption isotherm is indicative of the presence of different steps and mechanisms within the adsorption process. At the opposite, the second run curve increases quite smoothly from the beginning: its shape is typical of a process largely dependent upon the equilibrium pressure and in which the adsorbateadsorbate interaction competes with the surface-adsorbate interactions. At pH2O ≈ 2 Torr, ≈7 µmol/m2 H2O was adsorbed on the bare dehydroxylated surface, and this figure corresponds to ≈4 molecules/nm2; at this stage the surface is almost completely rehydroxylated and no longer exhibits any strong Lewis acidity (see Figure 3). Under the same water vapor pressure, only ≈2 µmol/m2 (≈1 molecule/nm2) was adsorbed on the hydrated surface, as shown by the second run curve, through a completely different kind of process. In fact, the energetics of the two runs are completely different, as shown in Figure 4b, where the heat of adsorption of water on ZS8(823)723 is reported as a function of H2O coverage. The initial heat value is quite high (q0 ≈ 250 kJ/mol), typical of a dissociative adsorption,29 but much higher than that

Heterogeneity of Superacid SO4/ZrO2 Systems

Figure 5. Adsorption of H2O on ZS8(823)723 ([, ]) and on ZS8(823)423 (2, 4): (a) volumetric isotherms (na vs pH2O); (b) heat of adsorption as a function of coverage (qads vs na); full symbols, first run; open symbols: second run.

previously reported for pure (monoclinic) zirconia (q0 ≈ 160 kJ/mol30). Then the heat of adsorption decreases quite rapidly, owing to the structural and chemical heterogeneity of the surface. The curve exhibits a region of nearly constant heat around 130-140 kJ/mol, quite unusual for heats vs coverage plots relative to the adsorption of water on oxides,29 but not unusual for some pure zirconia specimens.30 Eventually, the heat of adsorption reaches the value of ≈120 kJ/mol, typical of molecular water adsorbed undissociatively on coordinatively unsaturated cations,29 and then progressively decreases below 100 kJ/ mol, reaching the region of reversible H-bonding formation. The lower curve shown in Figure 4b is relative to the second run. The heat values measured in this case were lower than those obtained in the first run (with the only exception of the first dose of water bound to a very small number of residual highly energetic sites) and are typical of the adsorption of water on an hydrated surface layer, via reversible H-bondings.29 In Figure 5 the first and second runs of adsorption of water on ZS8(823)423 are reported and compared to those of ZS8(823)723. On the ZS8(823)423 sample one can expect (on the basis of the IR spectroscopic evidence reported in ref 15) water molecules to find a surface still hydrated. In fact, outgassing at 423 K eliminates both (29) Fubini, B.; Bolis, V.; Bailes, M.; Stone, F. S. Solid State Ionics 1989, 32-33, 258. (30) Orio, L.; Bolis, V.; Fubini, B.; Morterra, C. Ceramics Todays Tomorrow’s Ceramics; Materials Science Monographs, 66C; Vincenzini, P., Ed.; Elsevier Science Publishers B. V.: Amsterdam, 1991; p 1789.

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weakly and strongly adsorbed molecular water (the former H-bonded on the hydrated surface layer, the latter coordinated on cus Zr4+ cations) as well as the hydronium species, but only a very small amount of OH groups (if any). In Figure 5a the volumetric isotherms are reported: the shape of the first run adsorption is somewhat similar to the first run curve of the ZS8(823)723 sample, but lower amounts of water were obviously adsorbed in the -423 case. The second run curve lies below the first one (as expected for the reversible component of a process involving, at the first contact of the surface with the adsorptive, the formation of irreversibly held adspecies), but it is different from the ZS8(823)723 second run in both shape and amount. In Figure 5b, the relevant heats of adsorption vs coverage plots are reported and compared. The first run curve starts from a still quite high heat value (q0 ) 150 kJ/mol) and rapidly decreases reaching a short plateau at qads ≈ 120 kJ/mol. Then it gradually decreases down to the second run values, below 100 kJ/mol. The second run curves relative to ZS8(823)723 and ZS8(823)423 samples are very close overlapping, indicating that the similar reversible interactions occur at the hydrated surfaces, irrespective of the dehydration degree initially reached by the sample. However, it is worth noting that the stability of the adspecies formed is not the same in the two cases examined: the relative amount of H2O adsorbed in the second run with respect to the first one is in fact much higher in the -423 case than in the -723 one. In other words, the irreversible component is far lower in the -423 case than in the -723 one. The non-coincidence of the two curves relative to the second run adsorption isotherms for the -723 and the -423 samples indicates that the (stepwise) rehydration of the two surfaces leads to the formation of surface species of higher lability in the case of the sample mildly dehydrated at 423 K. It is worth reminding that on the -723 sample a virtually complete rehydration was achieved, in that the sample was kept overnight under a pressure of water vapor before the outgassing between the two runs. The thorough rehydration of the surface reached by the -723 sample apparently created surface species of high stability, that could not be removed by simply pumping off at 303 K. At the opposite, the species formed at the surface upon rehydration of the mildly dehydrated -423 sample were eliminated in quite high proportion by the outgassing: a thorough rehydration process is more difficult in these conditions, and labile species are mostly formed. A different lability of surface adspecies following the hydration extent of the surface, was previously observed also for other adspecies such as bicarbonate groups in the case of various zirconia systems.31 In the case of ZS8(923)723, results similar to those for ZS8(823)723 were obtained, for both first and second runs of adsorption, as shown in Figure 6, where the heats of adsorption of the two dehydroxylated samples are compared. It is worth noting that, whereas the initial heat of adsorption is nearly the same (q0 ≈ 250 kJ/mol), the flat region on the heat vs coverage plot (first run) at 130-140 kJ/mol is less extended in the case of the ZS8(923) sample, i.e., on the sample bearing the lower loading of sulfates. Thus, being the plateau in the heat vs coverage plots is dependent upon the abundance of sites carrying hydrated sulfate groups, the heat value of 130-140 kJ/mol is reasonably assigned to the adsorption of molecular water on the Brønsted acidic sites, HSO4-. Also in the case of the low-loading sulfated zirconia, the second run of the adsorption curve was found to closely (31) Morterra, C.; Cerrato, G.; Ferroni, L. J. Chem,. Soc., Faraday Trans. 1995, 91 (1), 125.

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zirconia, play nevertheless a substantial role, that may be explained in analogy to what happens with the Olah “magic acid”, i.e., the 1:1 mixture FSO3H + SbF5 having a H0 Hammet acidic function lower than -20.3,32 The presence of a strong Lewis acid as SbF5 cooperates with the strong Brønsted acid (FSO3H) in developing superacidity through an effect of extrastabilization on the conjugate base SO3F-:

2HSO3F + SbF5 f HFSO3H+‚‚‚SbF5(SO3F)It is noteworthy to remind that the “magic acid” is sufficiently acidic to protonate methane.3 Thus, in the sulfated zirconia case, one could suggest that the release of the proton by the Brønsted acidic site is made easier by the extra stability of the conjugate base, i.e., of the surface species consequently formed: Figure 6. Heat of adsorption of H2O as a function of coverage on ZS8(823)723 ([, ]) and on ZS8(923)723, (9, 0): full symbols, first run; open symbols, second run.

overlap the curve obtained in the same conditions for the high-loading sulfated zirconia. Proposed Model for the Catalytic Site. To explain the quantitative and energetic results discussed in the previous sections, we can hypothesize that at the very beginning of the contact with the highly dehydrated sulfated surface (where different kinds of cus Zr4+ cations are present, among which a ≈10% are very strongly acidic), H2O dissociates on sites involving very strongly acidic cus Zr4+ cations, next to a dehydrated sulfate group. The interaction gives rise to a bridged OH group15,25 and to a protonated sulfate:

One Lewis acidic site is eliminated (see the earliest stage of the gradual rehydration and consequent elimination of strong Lewis acidity, shown in Figure 3) but one Brønsted acidic site is created. The enthalpy change associated to the reaction is ≈250 kJ/mol, consistent with both rehydroxylation of the surface and hydration of the sulfuryl group. In a second step, when the surface concentration of the Brønsted acidic sites (HSO4- groups bound to zirconium cations) has become sufficiently high, these latter start to compete with Lewis cationic sites in strongly binding water molecules, allowing the formation of oxonium ions (see the second part of the curve relevant to the suppression of Lewis acidity upon rehydration, reported in Figure 3):

This model is in agreement with the fact that even a slight excess of water causes the catalytic activity to drop to zero. The Brønsted acidic sites that bind water with consequent formation of hydronium will be inhibited in manifesting their catalytic activity as a result of the interaction, in that H2O is a stronger base than alkanes and will deny the access of n-butane to the catalytic site. As for the superacidity, the strong Lewis sites, unable alone to yield the extremely high acidic strength of sulfated

HRTEM studies combined with catalytic tests (Morterra et al., to be published) suggest that superacidity preferentially develops on sulfated zirconia preparations showing a regular morphology. This fact, if confirmed, would be in agreement with the hypothesis that a peculiar geometry of the surface terminations is required to allow the Lewis sites to properly interact with the Brønsted ones, as in the model proposed. Conclusions Both Brønsted and Lewis sites are confirmed to be present at the surface of superacidic sulfated zirconia. Oxonium ions, whose presence on hydrated sulfated surface was previously evidenced by IR spectroscopy,15 are formed at the Brønsted acidic sites upon contact with water vapor, in competition with the dissociation of water onto coordinatively unsaturated Zr4+ cations. Lewis acid sites present at the sulfated surface are strong, but not strong enough to explain alone the superacidic behavior of the catalyst. In the early stages of rehydration, Lewis acidic sites are eliminated in parallel with the simultaneous formation of Brønsted acidic ones. Then, H2O strongly binds on the protonated sulfates acting as Brønsted acids, inactivating them toward the n-butane isomerization. The catalyst is efficient if only very slightly hydrated, and this is compatible with the Brønsted acidity hypothesis. Only a very little excess of H2O is sufficient to suppress the Lewis acidity and to inhibit the Brønsted one. The role played by the Lewis acidic sites is substantial for the onset of superacidity. The coordinatively unsaturated cations, owing to their strong electron-accepting ability, facilitate the proton transfer and give an extra stabilization to the conjugate base of the Brønsted superacid. Acknowledgment. This work was supported with funds of the CNR (Rome), “Progetto Finalizzato Materiali Speciali per Tecnologie Avanzate”, and “Comitato Nazionale per le Ricerche Tecnologiche (94.00544.CT11)”. The authors are indebted to Professor F. Pinna’s group (University of Venice, Italy) for providing the ZS catalysts as well as for fruitful discussions.. LA950938J (32) Olah, G. A.; Surya Prakash, G. K.; Sommer, J. Superacids; John Wiley & Sons: New York, 1985.