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J. Phys. Chem. B 2001, 105, 3872-3877
Kinetics of SO2 Adsorption on Photoexcited r-Fe2O3† David S. Toledano and Victor E. Henrich* Department of Applied Physics, Yale UniVersity, P.O. Box 208284, New HaVen, Connecticut 06520-8284 ReceiVed: September 16, 2000; In Final Form: December 8, 2000
SO2 is an atmospheric pollutant whose oxidation leads to acid rain, while R-Fe2O3 (hematite, a naturally occurring component of atmospheric aerosol particles) has a charge-transfer band gap (about 2.2 eV) smaller than the cutoff of solar radiation in the troposphere (about 4.3 eV) and is thus able to participate in photochemical reactions. The interaction of SO2 with UV-irradiated, single-crystal R-Fe2O3 was examined by using UPS, XPS, and Auger spectroscopies. Between 261 and 331 K, SO2 adsorbs on the R-Fe2O3(0001) surface with a very low sticking coefficient in the absence of UV irradiation. The adsorbed species resembles SO42-; its heat of adsorption is estimated to be 2.4 eV. UV irradiation of the R-Fe2O3 substrate during SO2 exposure leads to a significantly increased rate of adsorption. All of the cations in stoichiometric R-Fe2O3 are Fe3+; UV irradiation produces Fe2+ cations through the creation of electron-hole pairs. The Fe2+ sites are found to be much more reactive to SO2 adsorption than are Fe3+ sites. A model is proposed in which SO2 adsorbs only at Fe2+ sites. According to this model, the small number of adsorption sites which are present in thermal equilibrium are greatly enhanced by UV irradiation, leading to an increased rate of SO2 adsorption.
1. Introduction 1.1. Motivation. SO2 is a prevalent smokestack emission whose oxidation leads to acid rain, and its photochemistry is thus a subject of great environmental interest.1,2 However, the interaction of gas-phase SO2 with radiation is very small for photon energies below the high-energy cutoff of solar radiation in the troposphere (about 4.3 eV).3 The ground state of SO2 is 1A , and the first singlet (1B ) and triplet (3B ) states can be 1 1 1 photoexcited above 3.6 eV; however, photolysis to SO3 and SO requires energies above 4.3 eV for pure gas-phase SO2.3 Despite this, photooxidation of SO2 to SO3 is observed in the troposphere3 (with quantum yields ∼ 10-3), suggesting that catalysts may be involved. Atmospheric aerosol particles (∼1 µm in size) can act as substrates to catalyze heterogeneous photochemical oxidation reactions in the troposphere. Most tropospheric aerosols are silicates, aluminosilicates, and salts whose band gaps are larger than 4.3 eV; they are thus unable to participate directly in photoexcited reactions. However, transition-metal oxides that have much smaller band gaps (predominantly iron and manganese compounds) also occur as aerosols, and these materials may thus undergo charge-transfer excitations in the presence of sunlight through electron-hole pair creation, which permit photochemical reactions for which the material would otherwise be inert. Some of the earliest work on such heterogeneous photocatalysis using transition-metal oxides was performed by Yates’ group using TiO2; ref 4 is an excellent review of the photochemistry of this material. We have studied the effect of photoexcitation on the chemisorption activity of hematite (the most thermodynamically stable phase of iron oxide at room temperature5 and the most prevalent iron oxide aerosol in the troposphere) in the lowpressure regime using single-crystal R-Fe2O3(0001) surfaces as model catalysts. In a previous paper,6 we reported that irradiation †
Part of the special issue “John T. Yates, Jr., Festschrift”. * Corresponding author. Phone: 203-432-4399 (voice); 203-432-4283 (fax). E-mail:
[email protected].
of R-Fe2O3 with UV photons induced reversible electronic effects that were similar to reduction. In that paper we also examined the interaction of photoexcited R-Fe2O3(0001) with sulfur dioxide (SO2) and found a significant increase in the rate of chemisorption while under UV irradiation. In the current work we have examined the adsorption, desorption, and kinetics of SO2/R-Fe2O3(0001), with and without UV irradiation, in greater detail and also propose a model for the interaction. 1.2. Background. R-Fe2O3 is the most thermodynamically stable iron oxide phase at ambient temperatures.5 It has the corundum crystal structure (space group R3hc) and is an antiferromagnetic insulator below 953 K.7,8 Several studies have identified R-Fe2O3 as having a charge-transfer band gap of about 2.2 eV,9-11 indicating that the lowest-energy optical excitations involve electron transfer from occupied O(2p) states to unoccupied Fe(4s) or -(3d) states rather than transitions between Fe(3d) states (termed a Mott-Hubbard band gap); chargetransfer band gaps are discussed in ref 12. The highest occupied states (located approximately 2 eV below EF) are of largely O(2p) character, with some hybridized Fe(3d) contribution.13 Irradiation of R-Fe2O3 with UV photons having energies > 2.2 eV can create electron-hole pairs. In a previous study,6 the authors reported that UV irradiation of R-Fe2O3 (which has nominal Fe3+ valency) produces effects similar to the creation of transient Fe2+ cations. The effects of UV irradiation were found to be repeatedly reversible in ultrahigh vacuum (UHV), indicating that defect creation (such as the photoinduced desorption of oxygen anions) is not responsible for the Fe2+ density. Studies of the photovoltage induced by UV irradiation of Fe2O3 in solution indicate a surface accumulation layer of electrons with decay time g milliseconds;14,15 such long recombination times are consistent with the creation of a measurable, steady-state density of transient Fe2+ cations. These findings are significant because Fe2+ cations display very different chemistry than do the more inert Fe3+. The interaction of SO2 with R-Fe2O3 has been studied by several groups.6,16-18 SO2 appears to adsorb on R-Fe2O3 as an
10.1021/jp003327v CCC: $20.00 © 2001 American Chemical Society Published on Web 02/14/2001
Kinetics of SO2 Adsorption
J. Phys. Chem. B, Vol. 105, No. 18, 2001 3873
SO32- or SO42- complex with a low (∼10-3) sticking coefficient at room temperature.6,16,18 The magnitude of SO2 adsorption on R-Fe2O3 decreases significantly with increasing temperature: Inoue et al.16 report a 50% decrease between 298 and 308 K, while we observed a similar decrease between 267 and 300 K.6 Previously we reported that UV irradiation enhances the adsorption of SO2 on R-Fe2O3 by roughly a factor of 3.6 Since SO2 adsorption decreases with temperature, the increased adsorption under UV irradiation cannot be attributed to an increase in the R-Fe2O3 surface temperature. As discussed below, we propose that the creation of highly reactive, transient Fe2+ cations by UV irradiation is responsible for the increased rate of adsorption. This conclusion is supported by a siteselective study of adsorption on oxidized Fe foil,19 which found that S adsorbs only at Fe2+ sites. 2. Experimental Details All experiments were conducted in a stainless steel UHV chamber with base pressure 1 × 10-10 Torr that is equipped for inert gas ion bombardment, UPS, XPS, and Auger spectroscopy, and low-energy electron diffraction (LEED). UPS spectra were obtained using the He II line (40.8 eV) of an unmonochromated helium discharge lamp, and XPS spectra were obtained with an unmonochromated Mg KR source; all spectra were corrected for the presence of satellite lines in the photon source. Electron energies were measured with a doublepass cylindrical-mirror analyzer (CMA) having an instrumental resolution of 0.24 eV for UPS and 0.8 eV for XPS, referenced to the Fermi level or 4f7/2 level (84.0 eV), respectively, of a clean gold foil in electrical contact with the sample holder. Auger spectra were excited with a 3 keV electron beam coaxial with the CMA, using a modulation of 1 V p-p. Ion bombardment was performed using 99.9995% pure argon. Oxygen for annealing was 99.997% pure, filtered with a zeolite moisture trap, and SO2 of 99.99% purity was used for adsorption experiments. Photoelectron spectra are presented as the number of photoelectrons emitted (on an arbitrary, linear scale), N(E). When differences are taken between photoelectron spectra, they are indicated by ∆N(E). Samples were fastened onto a copper mounting plate that could be heated to >1000 K and cooled to 110 K. Sample temperature was monitored by chromel/alumel thermocouples attached to the plate close to the sample. Solar radiation was simulated by a focused 200 W Hg(Xe) lamp (Oriel Corp.) which provided broad-spectrum irradiation up to about 5.5 eV; the output of the lamp was directed onto the sample through a UVquartz window on the UHV chamber. A 7.5 cm thick recirculating water filter was employed to remove most of the infrared and reduce sample heating. The integrated intensity reaching the sample for hν > Egap (2.2 eV) was roughly 70 times the solar flux in that wavelength range,20 and the total irradiation reaching the sample was 35-40 kW/m2. The copper mounting plate was maintained at constant temperature regardless of UV irradiation, and while the sample surface temperature during irradiation could not be directly measured, a calculation using the thermal conductivity of R-Fe2O3 (6.9 W/m K at 300 K)21 indicated a surface temperature rise of 875 K or T < 575 K were the same as at 875 or 575 K, respectively.
beam used for Auger did not, itself, desorb SO2 from the R-Fe2O3 surface at 300 K; therefore, the magnitude of the SLMM Auger peak was used as a quantitative measure of adsorbed SO2. Figure 5 displays a series of Auger spectra taken as R-Fe2O3(0001), exposed to SO2 at 300 K, is heated to over 875 K in 100 K increments (with approximately 4 min intervals between spectra). No changes are observed in the SLMM peak amplitude up to 675 K. SO2 thus remains bound to the surface at 675 K (for at least 4 min); it is only partly desorbed after 4 min at 775 K. No sulfur signal is present in the Auger spectrum taken at 875 K, and no additional changes were observed upon heating to higher temperatures (not shown). It is unlikely that adsorbed SO2 dissociates at high temperature, since the shape
(4)
Using the values for τ and νo given above leads to Edes ≈ 2.4 ( 1 eV (55 ( 23 kcal/mol). The greatest source of error in this calculation is the value of νo, which may differ significantly from 1013 s-1. Vibrational spectra of SO2 adsorbed on other materials26-28 give a wide range for νo, leading to the large uncertainty in Edes. Also, our data are not sufficient to prove that desorption does indeed obey first-order kinetics, although that would be expected. 4.2. Mechanism for SO2 Adsorption. In this section we consider the interaction of SO2 with R-Fe2O3(0001) on the atomic scale. So far we have determined the following: (i) The activation energy for desorption of SO2 from R-Fe2O3 is roughly 2.4 eV. (ii) Large doses (g102 L) are required for adsorption at all temperatures. (iii) The adsorption of SO2 decreases with temperature. (iv) Adsorption is enhanced both by UV irradiation and the presence of oxygen-vacancy defects. If there were a significant energy barrier to chemisorption of SO2 on R-Fe2O3(0001), the amount adsorbed should increase
3876 J. Phys. Chem. B, Vol. 105, No. 18, 2001 with increasing temperature. The observed decrease in adsorption with temperature thus suggests that there is at most a very small barrier in this case. In the absence of such a barrier, Edes, estimated above, is the same as the heat of adsorption, Eads. We thus conclude that Eads for SO2 on R-Fe2O3(0001) is also close to 2.4 eV. Once adsorbed, SO2 bonds strongly to R-Fe2O3. If the majority of surface sites were suitable for adsorption, such strong bonding would lead to a probability of adsorption approaching unity. However, the actual probability of interaction is surprisingly low, as indicated by initial sticking coefficients in the range of 10-2-10-5. A precursor-mediated adsorption model25 is consistent with such observations. Suppose that a large fraction of the SO2 molecules incident on the surface enter a weakly bound precursor state and that only a small fraction of surface sites (i.e., Fe2+ cations) will chemisorb SO2. [Some Fe2+ sites will be present even on our stoichiometric R-Fe2O3(0001) surfaces (without UV irradiation) due to thermodynamic considerations;29 by comparing UPS spectra in ref 13 for stoichiometric R-Fe2O3 and Fe3O4 (which contains 1/3 Fe2+ and 2/3 Fe3+ cations), we have determined that the UPS intensity between 0 and 2 eV binding energy on our stoichiometric R-Fe2O3(0001) surfaces (without UV irradiation) is consistent with a maximum Fe2+ density near 300 K of e1012 sites per cm2, or less than 1 per 1000 surface cations.] Precursor molecules are presumed to quickly thermalize to the R-Fe2O3 surface temperature. They can then diffuse along the surface, which consists of almost all Fe3+ cations, in search of suitable (Fe2+) chemisorption sites. Residence times in such precursor states are typically very low since the precursor-substrate bonding is very weak (comparable to physisorption25); thus most precursor molecules will desorb without finding a suitable surface chemisorption site. Also, the precursor state will be fully depleted almost immediately after SO2 exposure, so only those molecules that have managed to enter the chemisorbed state will be detected by UPS or XPS. As discussed above, precursor molecules that managed to locate an Fe2+ site before desorption presumably face a negligible barrier to chemisorption. Since this is a small fraction of the molecules incident on the surface, the apparent sticking coefficient for chemisorption would also be very small, as observed. A precursor state is also consistent with the observed temperature dependence of SO2 chemisorption. Increasing temperature will give both a smaller trapping probability into the precursor state and a shorter residence time in that state. Thus the number of SO2 molecules diffusing across the surface in search of active adsorption sites, and thus the number of surface sites sampled before desorption, will decrease with increasing temperature, yielding the temperature dependence observed in Figure 4. We thus propose that SO2 chemisorption on R-Fe2O3 takes place predominantly at Fe2+ cation sites, which may be created either by ion bombardment or UV irradiation of the surface. This is consistent with the increased activity of the ionbombarded surface (Figure 2) and evidence for the creation of transient Fe2+ cations through UV irradiation.6 Additional support for this hypothesis is provided by a study of the adsorption behavior of H2S, Cl2, and H2O on oxidized Fe foil (which contained both Fe2+ and Fe3+ cation sites);19 Cl and S were reported to compete for the same adsorption sites, which were determined to be strictly Fe2+ cation sites. To adsorb as SO32- or SO42-, the chemisorbed SO2 must interact strongly with surface O anions. We are not able to determine just what role surface Fe2+ cations play in facilitating
Toledano and Henrich the SO2-O interaction. However, it is clear from the very small sticking coefficients observed that O anions adjacent to Fe3+ cations are not capable of interacting strongly with SO2. An alternative to SO2 molecule diffusion across the surface could be charge transfer between cation sites, enabling the Fe2+ sites themselves to diffuse across the surface; this “hopping” mode of conduction is observed in bulk Fe3O4, which contains a mixture of Fe2+ and Fe3+ cations.30 However, the frequency of hopping should increase with increasing temperature, thus making chemisorption more likely at higher temperature, in disagreement with the observed behavior. We must also consider the possiblilty that direct UV excitation of SO2 molecules before they adsorb on the surface might create an excited SO2 species which would have a higher probability of chemisorption. This effect is not likely to be significant, since the UV absorption by gas-phase SO2 is very small over the energy range of the Hg(Xe) arc lamp used for photoexcitation (hν < 5.5 eV),3 while UV absorption by R-Fe2O3 approaches unity above 2.2 eV.31,32 To confirm that the effect of UV excitation on gas-phase SO2 can be disregarded, we have used data from ref 3 to determine the maximum fraction of SO2 molecules that could be UV-excited to the first singlet and triplet states. The calculated value (e10-6) is far too small to account for the observed 3-fold UV-induced increase in adsorption, since the sticking coefficient of SO2 is g10-5 even in the dark: even if photoexcited SO2 molecules adsorb with unity sticking coefficient, this would increase the total chemisorption of SO2 by no more than 10% (i.e., g10-5 of the SO2 molecules stick in the dark plus at most an additional 10-6 representing the number of photoexcited SO2 molecules). In addition, the ionization energy of SO2 is about 12.5 eV,33 which is far larger than the high-energy cutoff of the Hg(Xe) arc lamp. Note that while the UV-adsorption characteristics of adsorbed SO2 might differ somewhat from those of gas-phase SO2, we are concerned here only with the kinetics of SO2 molecules which have not yet chemisorbed to the surface rather than those which have become bound. 5. Summary The interaction of SO2 with single-crystal R-Fe2O3(0001) surfaces with and without UV irradiation was studied between 267 and 331 K. SO2 adsorbs on both stoichiometric and reduced surfaces as an SO32- or SO42- species which desorbs above 675 K. Stoichiometric surfaces exhibit very low sticking coefficients in the range of 10-2-10-5, which decreased with increasing T. The heat of adsorption was determined to be about 2.4 eV. The small sticking coefficient for chemisorption and the decrease in amount adsorbed with increasing temperature are consistent with a precursor-mediated adsorption model. A much greater quantity of adsorption was observed both on reduced surfaces and on stoichiometric surfaces that were irradiated with UV during adsorption, although the nature of the adsorbed species remained the same in all cases. SO2 adsorption is believed to take place only at reduced Fe2+ cation sites, which may be created either by ion bombardment (which reduces the surface through permanent oxygen-vacancy defects) or UV irradiation (which induces transient effects similar to Fe2+ cation creation). Acknowledgment. The authors thank Gary Haller and John Tully for valuable discussions. This research was partially supported by the Petroleum Research Fund (Grant 28797-AC5) and by NSF Grant CTS-9610140.
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