Heterogeneous Uptake and Oxidation of SO2 on Iron Oxides - The

C , 2007, 111 (16), pp 6077–6085 .... The scanning range for acquiring spectroscopy was from 0 to 1200 eV. ..... 4.3. Proposed Mechanism of SO2 Upta...
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J. Phys. Chem. C 2007, 111, 6077-6085

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Heterogeneous Uptake and Oxidation of SO2 on Iron Oxides Hongbo Fu, Xiao Wang, Hongbo Wu, Yong Yin, and Jianmin Chen* Center for Atmospheric Chemistry Study, Department of EnVironmental Science & Engineering, Fudan UniVersity, Shanghai 200433, China ReceiVed: January 5, 2007; In Final Form: February 14, 2007

Heterogeneous oxidation of gas-phase SO2 on different iron oxides was investigated in situ using a White cell coupled with Fourier transform infrared spectroscopy (FTIR) and diffuse reflectance infrared Fourier transform spectroscopy (DRIFTS). The results revealed that adsorbed SO2 could be oxidized on the surface of most iron oxides to form a surface sulfate species at ambient temperature. Additional support for this hypothesis was provided by X-ray photoelectron spectroscopy (XPS) measurements and ion chromatogram (IC) analysis. The spectroscopic results further revealed that the surface hydroxyl species on the iron oxides was the key reactant for this heterogeneous oxidation. Furthermore, the bidentate sulfate species were the predominant surface species with R-Fe2O3, R-FeOOH, and Fe3O4, while the monodentate surface Fe(III)sulfato complexes were available in the case of γ-Fe2O3. Using the BET area as the reactive surface area, the samples showed the varied reactivity in the order of R-Fe2O3 > γ-Fe2O3 > Fe3O4 > β-FeOOH > R-FeOOH. Some preliminary experiments indicated a significant acceleration during SO2 uptake on the surface of R-Fe2O3 in the presence of oxygen. In contrast, no significant formation of sulfate was seen on the surface of R-Fe2O3 prereduced by H2 at 523 K at the absence of O2, suggesting that the concentration of adsorbed oxygen over catalyst surfaces may be the key factor contributing to the oxidizing activities. On the basis of these results, the atmospheric implications of these studies on SO2 uptake on Fe-rich mineral aerosol were discussed.

1. Introduction Particulate matter present in the Earth’s atmosphere provides reactive surfaces for heterogeneous chemistry. A large contribution to the tropospheric aerosol budget is mineral aerosol.1 Iron is one of the most important components in mineral aerosol.2 Field measurements of the chemical composition of aerosol particles showed that the major source of iron in dry aerosol appears to be in the form of highly insoluble iron oxides and oxyhydroxides, mainly including R-FeOOH, R-Fe2O3, Fe3O4, γ-Fe2O3, and β-FeOOH.1-5 Hoffmann and co-workers5 have established that the total iron concentrations in urban fogs and clouds are unusually high. Iron was estimated to account for 7% w/w of an urban aerosol sample. The concentration of R-Fe2O3, R-FeOOH, and Fe3O4 in Fe-containing particles was 7.5, 60.8, and 9.8%, respectively. Given the relative abundance of iron oxides in aerosols, a thorough investigation of ironinvolved redox behavior in the atmosphere is warranted. SO2 is a prevalent smokestack emission whose oxidation leads to acid rain, and its transport and transformation is thus a subject of great environmental interest. It is well documented that gaseous SO2 could react directly with mineral aerosol to form particulate sulfate.6-8 Atmospheric sulfate particles are known to play a critical role in the global climate by scattering solar radiation and increasing the number of cloud condensation nuclei and thereby affecting climate indirectly.8,9 Nearly half of the global emissions of SO2 are converted to particulate sulfate, and this sulfate is often associated with mineral aerosol.6-11 Recently, the heterogeneous reactivity of SO2 has been observed on metal oxides,12,13 China Loess,14 and Saharan mineral.15,16 Laboratory studies of the heterogeneous reactions of SO2 were * Address correspondence to this author. Tel.: (+86)21-6564-2521. Fax: (+86)21-6564-2080. E-mail: [email protected].

undertaken with the recognition that gas-solid reactions could be important in the atmosphere. Ullerstam et al.14,15 reported sulfate formation on mineral dust as a surface product and the uptake coefficient in the range of 10-5-10-10, differing by a factor of 5. Usher et al.13 have determined that the initial uptake coefficient, γBET, was 7 ( 2 × 10-5 for R-Fe2O3.12 Faust et al.11,16 reported that irradiated iron oxides could oxidize sulfite and the photooxidation rate varied by 2 orders of magnitude among the different iron species. Although laboratory studies have quantified these processes to some extent, large uncertainties remain concerning the effect of the chemical composition and the surface properties of the iron mineral on this heterogeneous reaction kinetics. Because both the sulfate and Fecontaining aerosol play an important role in global environmental change, a better understanding of the origin and the overall mechanism of sulfate formation on iron oxides is highly desirable. The interaction between SO2 and the oxide surface is very complicated, depending on the nature of the adsorption sites. Determining the structure of the sulfate species more conclusively requires more knowledge about the sulfate species bonding to surface iron atoms. Such information is readily accessible through the profound analysis of in situ systems at the microscopic level accompanied by modern spectroscopic techniques. However, currently only a few spectroscopy studies are available to probe the interaction of SO2 and iron oxides.4,15,17 Our work so far has demonstrated that atmospheric particles are active in promoting the transformation of the S-containing compound.18-21 On the basis of this experience, we have now reinvestigated the formation and structure of the sulfate species on iron oxides in more detail, using a White cell coupled with in situ FTIR and DRIFTS. This made it possible to characterize, kinetically, both the formation of the surface product and the

10.1021/jp070087b CCC: $37.00 © 2007 American Chemical Society Published on Web 03/30/2007

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TABLE 1: Characterization of Iron Oxides

formula

mineral name

crystal structure

R-Fe2O3 γ-Fe2O3 R-FeOOH β-FeOOH Fe3O4

hematite maghemite goethite akaganeite magnetite

corundum inverse spinel rhombic rhombic spinel

crystal BET surface area size hydroxyl (nm) (m2/g) (×10-7 M) 39.7 8.9 38.8 14.9 34.6

10.2 116.0 39.7 16.6 94.3

7.5 7.0 3.6 11.8 11.7

gas-phase loss of the reactive species. XPS, IC, and HPLC were also used to characterize and identify the chemical properties of the surface-bound products. On the basis of the spectral data, the possible mechanism of SO2 uptake on the oxide surface was proposed. 2. Experimental Section 2.1. Materials. R-Fe2O3, R-FeOOH, Fe3O4, γ-Fe2O3, and β-FeOOH were chosen to be model oxides and were synthesized according to a procedure published previously.22 X-ray diffraction (XRD) of the iron oxides was identical to the standard JCPDS cards. The nitrogen adsorption-desorption isotherms of iron oxides were obtained at 77 K over the whole range of relative pressures using ASAP 2010 automatic equipment. Specific areas were estimated from these isotherms by applying the Brunauer-Emmett-Teller (BET) method. The mineral name, crystal structure, particle size, and BET surface area of the iron oxides are shown in Table 1. The specifications of the gases used in this experiment are as follows without further purification: SO2 (96 ppm, SO2/N2, Shanghai Yunguang Specialty Gases Inc.); N2 and O2 (99.999% purity, Wuxi Xinnan Chemical Gas Company Ltd.). All chemicals used were reagent grade. 2.2. Analytical Methods. In Situ DRIFTS Experiment. In situ DRIFT spectra were recorded on a Nicolet Avatar 360 FTIR equipped with a diffused refection chamber and a highsensitivity MCT detector cooled by liquid N2. The schematic diagram of the DRIFTS apparatus was shown in our previous report.18,23 Before the reaction gas was introduced, argon was flowed into (150 mL/min) the chamber to wash off water, physisorbed impurities, and the powder on the surface of the chamber. The sample for the in situ DRIFTS studies was finely ground and placed into a ceramic crucible in the chamber. A single-beam spectrum was used as a reference spectrum, which was collected prior to exposure of SO2, to obtain the absorbance spectra of the gas phase. After collecting the background spectrum, a mixture of SO2 (2 ppm) and O2 (21% v/v) with the N2 carrier (78.99% v/v) was introduced into the chamber at a flow rate of 125 mL/min. All spectra reported here were recorded at a resolution of 4 cm-1 for 100 scans unless otherwise noted. The losses of those species depleted in the surface reactions would show negative bands, whereas those of the products would show positive ones. In Situ FTIR Experiment. For FTIR measurements, approximately 20 mg of oxide sample was prepared by pressing it into a pallet with a diameter of 13 mm and then secured inside the infrared cell (the White cell reactor-the variable-path long path cell, Model 19-V, Infrared Analysis, Inc.) through the sample holder jaws made of Teflon. The details about the White cell have been given in our previous study.18,23 FTIR spectra were run on a Nicolet Avatar 360 FTIR equipped with a N2cooled MCT detector. The spectra were collected and analyzed using a data acquisition computer installed with OMNIC 6.0 software (Nicolet Corp.). Typically, each spectrum was recorded by 100 scans on average at a resolution of 4 cm-1 over the

range extending from 600 to 4000 cm-1. A single-beam spectrum was used as a reference spectrum, which was taken prior to exposure to SO2. The quantification method of SO2 was reliable with a linear correlation coefficient of 0.999 for the standard curve. All of the measurements were repeated at least twice. 2.3. Analysis of Surface-Bound Products. The solid products on the surfaces of the oxide samples were analyzed by XPS (Model PHI5000C, P.E. Inc.). The analytical system used Al KR as the irradiation source with a working power of 250 W at a voltage of 14 kV and was calibrated against C1s spectrum with Eb ) 284.50 eV used as the internal reference. The scanning range for acquiring spectroscopy was from 0 to 1200 eV. The minute spectroscopies of three elements (O1s, S2p, and Fe2p) were acquired. The special software PHI-MATLAB was used for data analysis. A high-performance liquid chromatograph (HPLC, Waters) was used to qualify and quantify the concentration of Fe(II). All of the analyses were performed on a Waters1525 Binary HPLC Pump and a 250 × 4.6 mm2, 5 µm capcell PAK C18 column, equipped with a Waters 2487 dual λ absorbance detector. The solid products on the surface of the samples were also analyzed using a Diones DX 500 ion chromatography equipped with a Dionex AS 14 analytical column. 3. Results 3.1. In Situ DRIFTS Study of SO2 Uptake on Iron Oxides. SO2 uptake on the R-Fe2O3 sample in the presence of O2 was observed first. The R-Fe2O3 sample was exposed to a mixture flow of SO2 (2 ppm), O2 (21% v/v), and N2 (78.99% v/v) at a rate of 125 mL/min. The in situ DRIFT spectra on the R-Fe2O3 sample were recorded as a function of time, as shown in Figure 1. It was evident that the intensities of the negative broad peaks from 3750 to 3150 cm-1 and from 1700 to 1500 cm-1 increased drastically with time. These bands were attributed to the vibrations of the surface hydroxyl species (OH). This result implies that the surface OH might be the reaction active site for SO2. Upon adsorption of SO2 on the surface of R-Fe2O3, three new bands were readily observed at 1225, 1149, and 1042 cm-1, and the intensities of the prominent peaks in the spectra increased with time until the surface of the oxide was saturated. The assignment of the vibrational bands of adsorbed SO2 has been discussed previously.24,25 These new bands are assigned to the stretching motion of adsorbed bisulfate, HSO4-, and/or sulfate, SO42-, on the surface. With the FTIR technique, there are two infrared sulfate vibrations that are accessible to spectroscopic investigation. They are the nondegenerate symmetric stretching V1 band and the triply degenerate asymmetric stretching V3 band.24 When SO2 uptake occurs on the surface of R-Fe2O3, the V1 band is weakly active and appears at 976 cm-1. The V3 band splits into three peaks, 1225, 1149, and 1042 cm-1, respectively. This indicated that the symmetry was lower when SO2 was transformed into the bound sulfate product on the surface of the oxide. SO2 exposed to R-Fe2O3 was bound through two oxygen atoms, resulting in the formation of the bidentate-surface complexes. These absorptions grew in intensity with decreasing O-H stretching and H-O-H bending, indicating that the amount of the complexes formed between the sulfate species and the surface hydroxyl by ligand exchange increased.23 These new bands remained when a blow-off process with argon was carried out, suggesting that the surface-adsorbed species with the broad absorbance bands between 800 and 1300 cm-1 was chemisorbed. The same set of experiments was also carried out using the R-FeOOH, γA-Fe2O3, β-FeOOH, and Fe3O4 samples, respec-

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Figure 1. DRIFT spectra of R-Fe2O3 following exposure to gaseous SO2 and O2 for 90 min in the range of 4000 to 3000 cm-1 (a) and 1800 to 900 cm-1 (b).

Figure 2. DRIFT spectra of SO2 uptake on R-FeOOH (a), γ-Fe2O3 (b), β-FeOOH (c), and Fe3O4 sample for 90 min (d).

tively. The in situ DRIFT spectra on these samples were recorded as a function of time, as shown in Figure 2. Compared to that of R-Fe2O3, SO2 uptake on R-FeOOH and Fe3O4 showed a similar spectral character. Three new absorption bands between 1000 and 1300 cm-1 appeared when exposed to SO2, suggesting that the bidentate sulfato-metal surface complexes formed when SO2 reacted with these iron oxides. However, the surface reaction of SO2 with γ-Fe2O3 appeared to follow a somewhat different mechanism. The DRIFT spectrum obtained following adsorption of SO2 on γ-Fe2O3 only showed two new bands between 1000 and 1300 cm-1. In the model proposed by Peak,24 these two peaks were split by the V3 band. This was a strong indication that coordination with surface Fe(III) sites has occurred. In this case, an inner-sphere complex formed only by one oxygen atom. The conversion of SO2 at the surface of

β-FeOOH was not well understood. Only two broad peaks appeared between 900 and 1300 cm-1 when β-FeOOH reacted with SO2. Sulfate weakly bound as outersphere complexes by electrostatic attraction would be expected. 3.2. Effect of O2 on the Heterogeneous Oxidation of SO2. To identify the role of oxygen in SO2 uptake on the iron oxides, the R-Fe2O3 sample was exposed to a flow of 1 ppm SO2 + the varied concentration of O2 from 0 to 39% (φ). Figure 3 shows the area of the characteristic peaks from 1062 to 1321 cm-1 as a function of the oxygen concentration ([O2]) in the system. It was evident that the peak area became greater with the increase of [O2]. When [O2] rose up to 21%, the peak area of the product was greater by 31, 11, and 3 times compared to the ones when [O2] was 0%, 5%, and 11%, respectively. This result implied that oxygen could enhance SO2 uptake greatly

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Fu et al. TABLE 2: Rate Parameters and Uptake Coefficients for the Heterogeneous Reaction sample R kf (×10-7 s-1) γobs (×10-5) γBET (×10-10)

Figure 3. Variation of the relative peak area of the product with O2 concentration.

Figure 4. DRIFT spectra of SO2 uptake on the R-Fe2O3 sample prereduced by H2 at different temperatures in the absence of O2.

on the surface of R-Fe2O3. However, to further enhance [O2] up to 39%, the area of the product peak decreased slightly. This passive effect of oxygen could be attributed to its competed absorption for surface active sites with the gaseous SO2, resulting in the lower formation of sulfate. The control experiments showed that oxygen played a critical role in sulfate formation. It was clear that iron oxides were much less reactive for SO2 uptake in the absence of oxygen. Considering that the H2 pretreatment of the sample was expected to weaken the function of O2, the experiment was further performed on the R-Fe2O3 sample prereduced by H2 at ifferent temperatures. The result is shown in Figure 4. In the case of 473 K, two absorption bands at 1186 and 1079 cm-1 appeared when the sample was exposed to SO2, suggesting that the surface products formed. In contrast, the product peaks were not visible in the presence of the samples treated at 523 and 653 K. At 473 K, both the lattice oxygen and the surface active oxygen of the sample remained. At 523 K, the lattice oxygen still remained while the surface active oxygen dropped. At 653 K, a fraction of the lattice oxygen dropped. This result implied that the lattice oxygen was not a main oxidant in this system. When surface oxygen was consumed, oxygen in the gas phase could be a supplement so that the oxidation reaction could continue until the oxide surface was covered by HSO4-/SO42species.23,25 3.3. Uptake Coefficient of SO2 on Iron Oxides Determined by in Situ FTIR. The R-Fe2O3 sample was exposed to a flow

R-Fe2O3 γ-Fe2O3 R-FeOOH

Fe3O4

β-FeOOH

0.9723 0.35 0.84 5.38

0.9655 1.27 3.08 2.13

0.9604 0.15 0.34 1.34

0.9583 2.37 5.8 3.28

0.9204 0.17 0.41 0.68

of 45 ppm SO2 + 21% O2 in N2 carrier gas at 298 K for 20 min, and then the inlet and outlet were closed. The in situ FTIR spectra on the R-Fe2O3 sample were recorded as a function of time and are shown in Figure 5. The strong bands at 1373, 1359, and 1346 cm-1 and the weak ones appeared at 1164 and 1135 cm-1 were assigned to the characteristic peaks of the gas-phase SO2.23 Besides, there was some weak absorption near 1550 cm-1, which was assigned to the feature of H2O produced from the reaction of SO2 with surface hydroxyls.23 With the increase of reaction time, the intensities of the SO2 peaks decreased in the spectra until the oxide surface was saturated. No other detectable peaks in the spectrum were visible, which implied that there were no gaseous products containing sulfur. The same set of experiments was also carried out using the R-FeOOH, γ-Fe2O3, β-FeOOH, and Fe3O4 samples, respectively. These samples showed a similar adsorption with the only difference being in the consumption rate of SO2. The SO2 concentrations shown in Figure 6 were determined by comparing the area of the absorbance peak of SO2 with the calibration curve of its concentration. To distinguish the effect of the system surface on the loss of SO2, a control experiment was performed under the same conditions by replacing the oxides with a gold mirror. In this case, only a very weak change of the gas-phase SO2 concentration was observed during the experiment, indicating that the system consumed a small quantity of SO2 through the adsorption or the heterogeneous reaction catalyzed by the surface of the system. It could be clearly seen that the reactivity of the sample varied with the various oxides. The concentration of the gaseous SO2 decreased drastically with γ-Fe2O3, whereas it decreased mildly with Fe3O4. The conversion of SO2 on the surface of γ-Fe2O3, Fe3O4, and R-Fe2O3 in 3 h was up to 17.8%, 11.9%, and 4.7%, respectively. R-FeOOH and β-FeOOH showed relatively low activity for SO2 uptake. The surfaces of R-FeOOH and β-FeOOH were almost saturated when they were exposed to gaseous SO2 for 3 h. This heterogeneous reaction was most likely irreversible because the reactants of SO2 and O2 were gaseous and one of the products was sulfate in the solid phase. The concentration of O2 could be regarded as constant because it was very abundant as compared to that of SO2. Ullerstam et al.15 reported that the heterogeneous oxidation of SO2 on the oxide surface was a pseudo-first-order reaction. The kinetics equation for the pseudo-first-order reaction should be as follows

-

dC ) kf C0 dt

(1)

C ) kf t C0

(2)

ln

where C0 is the initial concentration of SO2, C is the concentration of SO2, and kf is the apparent rate constant. Indeed, straight lines were obtained by plotting ln[C/C0] versus time (t). Table 2 shows that the correlation coefficients of ln[C/C0] versus time (t) for iron oxides were greater than 0.90. The result indicated that SO2 uptake on the iron oxides was fitted for pseudo-firstorder kinetics.

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Figure 5. In situ FTIR spectra for the heterogeneous reaction of R-Fe2O3 with SO2 for 480 min in the presence of O2.

Figure 6. Time evolution of the concentration of SO2 in the presence of different iron oxides.

The total uptake coefficient (γobs) was defined as the fraction of collisions with the surface that results in loss of the molecule from the gas phase divided by the total number of surface collisions per unit time.15 From the experiments, this can be calculated from the apparent kinetic constant, the collision frequency (Z), and the reactive surface area (As ) 0.785 cm2).

γobs ) Z)

kf Z As

1 x8RT Mπ 4V

(3)

(4)

where R is the gas constant (J mol-1 K-1), T is the temperature (298 K), V is the volume of the reactor (0.014 m3), and M is mol mass of SO2. The BET uptake coefficient (γBET) was deduced from eq 4.

γBET )

γobs A ABET s

(5)

where ABET is the BET surface area. The total uptake coefficient and BET uptake coefficient for each iron oxide are shown in Table 2. It is evident that the uptake coefficient of the iron oxide varied greatly, depending on its kind. Using the BET surface area as the reactive surface area, the activity of these oxides with SO2 was in the order of R-Fe2O3 > γ-Fe2O3 > Fe3O4 > β-FeOOH > R-FeOOH. The BET uptake coefficients of the iron oxides ranged from 0.68 × 10-10 to 5.38 × 10-10, which is consistent with Ullerstam’s report.15 The large variation of uptake rates for the different iron oxides was attributed to the differences in surface and structure properties because there was no correlation between the rates and the hydrodynamic diameter, crystal parameter, or surface area, as shown in Table 1. 3.4. Analysis of Surface-Bound Products. To identify the chemical composition on the reacted R-Fe2O3, the highresolution XPS spectrum of each component atom was obtained and the results are shown in Figure 6. O1s XPS spectra were peaked at 530.5 eV (Figure 6a) and ascribed as the bulk oxygen atom of R-Fe2O3. No S signals (Figure 6b) could be identified, indicating that there were no sulfur species before the reaction. With the inspection of the Fe2p spectrum as shown in Figure 6c, two characteristic peaks of Fe2p3/2 and Fe2p1/2 were observed at 710.6 and 724.1 eV, respectively, with a separation of 13.5 eV.26,27 After exposure to SO2, the XPS spectra of the reacted R-Fe2O3 apparently changed. As shown in Figure 6d, the O1s spectrum comprised two peaks. The intense peak at 531.4 eV was attributed to the bulk oxygen of R-Fe2O3. However, the small peak at 532.9 eV could be due to the oxygen origin from adsorbed HSO4-/SO42-.27 The ratio of two types of O atoms in R-Fe2O3 and surface complexes was 78:22. Signals from S2p photoelectrons were clearly observed when R-Fe2O3 was exposed to gas-phase SO2, as shown in Figure 6e. The XPS spectra for the S2p3/2 center were at 167.5 and 169.5 eV, respectively. The former was in good agreement with the binding energies of S2p3/2 in SO32- and the latter with that of S2p3/2 in SO42-, rather than sulfur-metal bonding, such as in FeS (161 eV).27 The peak fitting results for the XPS spectra from two individual samples indicated that the mixing ratio of SO32-/SO42- was 35:65, indicating that 65% of the total gas-

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Figure 7. XPS spectra of O, S, and Fe species on the R-Fe2O3 sample before (a-c) and after (d-f) uptake of SO2.

phase SO2 adsorbed on the surface of R-Fe2O3 was converted to sulfate, SO42-, and the rest of 35% of S(IV) was not converted to S(VI); instead, they existed in the form of the species SO32-. In Figure 6f, the spin-orbit doublet, Fe2p3/2 was composed of two peaks, one at 710.3 eV and the other at 712.2 eV, which was attributed to the characteristic peaks of Fe3+ and Fe2+, respectively. This observation was also supported by those literally reported.27 XPS analysis of R-Fe2O3 before and after the reaction provides the convincing evidence for the formation of SO32- and SO42- on the surface of R-Fe2O3, as well as the valence change of surface iron atoms from Fe(III) to Fe(II). The XPS spectra of the other iron oxides showed similar characters, the only difference being in the individual atom ratio, which could be due to the different reaction ability of the sample for SO2 uptake. Following XPS measurements, the amount of sulfate was further measured using IC. A content of 2.8 mg/g of soluble SO42- on the reacted R-Fe2O3 was determined. This confirmed the irreversible adsorption of SO2 on the surface and subsequent sulfate formation. Concurrently, 0.65 µg of dissolved Fe(II) for every milligram of reacted R-Fe2O3 was identified by HPLC, which coincided with the XPS analysis. 4. Discussion 4.1. Conformational Analysis of Surface Fe(III)-Sulfato Complexes. FTIR studies of gaseous or aqueous sulfurcontaining species absorbed to minerals have been con-

ducted.24,28 Sulfate adsorbs to a mineral surface through either one or two oxygen atoms, making the monodentate or bidentate sulfato-metal surface structures possible. The infrared spectrum of the sulfate depends on its symmetry, and the symmetry of the adsorbed sulfate changes depending in whether it is monodentate or bidentate. The free sulfate ion in solution belongs to the Td point group, in which only the triply degenerate V3 and V4 vibrational modes are infrared-active. The V3 vibrational mode occurs only above 1000 cm-1, especially at 1100 cm-1. The symmetry of the sulfate ion is reduced when it is adsorbed to a surface. A monodentate surface structure results in C3V symmetry, for which the V1 vibrational mode becomes active, and the V3 and V4 vibrational modes each split into two modes. Thus, for a monodentate surface sulfato-metal complex, two peaks are present above 1000 cm-1 due to the splitting of the V3 mode. The symmetry of the sulfate ion is further reduced to C2V when it is bound through two oxygen atoms, resulting in a bidentate surface ion. For this symmetry, the V3 vibrational mode of the sulfate ion splits into three bands.24 For a bidentate surface sulfato-metal complex, three vibrational modes are expected above 1000 cm-1. Alternatively, sulfate weakly bound as physisorption by electrostatic attraction and/or hydrogen bonding would not be expected to show spectral changes as large as those for the irreversible complexes.28 On the basis of peak splitting information supplied by the DRIFT spectra, we proposed that the bidentate complex formed when R-Fe2O3, R-FeOOH, and Fe3O4 were exposed to gaseous SO2. However,

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the monodentate surface sulfato-Fe species was available in the case of γ-Fe2O3. In contrast, SO2 did not specifically interact with the β-FeOOH surface. Its effect toward the β-FeOOH surface was a weaker interaction and a minimal distortion from Td symmetry, as indicated by the absence of distortion in the infrared band above 1000 cm-1. 4.2. Influence of Iron Phases on SO2 Uptake. Conformation of the surface Fe-sulfate complexes varied with the iron species. With R-Fe2O3, R-FeOOH, and Fe3O4, sulfate only formed bidentate complex species while monodentate surface complexes were available in the case of γ-Fe2O3. The structures of functional groups at the metal oxide surface could account for the different behavior observed in the two systems. The relative site densities for all functional groups present on various crystal faces for the iron oxides have been determined.11 With R-Fe2O3, the 110 faces predominate and the surface hydroxyl mainly presents oxygen-twin coordinated ones, which are capable of ligand substitution reactions with sulfate species and thus form bidentate complexes. As for γ-Fe2O3, the surface functional groups behave somewhat differently. The 311 faces are dominant, and the sites that readily protonate are the singly coordinated (Fe-O). On the central surface, the singly coordinated surface groups exist as Fe-OH. This explains the observed monodentate complex formed on the surface of γ-Fe2O3 when it exposed to SO2.22 The activity of iron oxides for SO2 uptake is correlated to the population of surface hydroxyl sites. This is quite understandable because these active sites are involved as the principal reactive sites in this heterogeneous reaction. The numbers of surface hydroxyl could be estimated from eq 6.

[>OH] )

(ΜFe/dFeV)SAdOH(1018 nm2/m2) (M) N

(6)

where MFe is the mass of the oxide (kg), dFe is the density of the oxide (kg/m3), V is the volume of the oxide (m3), >OH is the density of surface hydroxyl group, and N is Avogadro’s constant.22 The calculated results shown in Table 1 did not correspond to the activity orders of the iron oxides, suggesting that the factors in addition to the number of surface hydroxyl groups should be taken into account. The role of the surface hydroxyl groups as the principal reactive sites on metal oxide surfaces and as the source of the amphoteric properties of the metal oxides has been clearly established. Spectroscopic studies have shown that there are a variety of types of hydroxyl sites present on hydrolyzed metal oxide surfaces.11,22 Furthermore, these surface hydroxyl groups are not identical and are of varied relative energy. Compared to the population of surface functional groups, their properties could play a more important role.11 Jame et al.29 elaborated that adsorption of SO2 on Lewis acid sites (coordinately unsaturated metal atoms) resulted in weakly adsorbed SO2. However, the adsorption of SO2 on Lewis base sites (exposed oxygen atoms) were prone to a strong interaction and followed by rearrangement where the sulfite species attached to the metal atom through the sulfur atom resulted in chemisorbed sulfate. Karge and Dalla Lana30 further developed this theory. They studied the interaction of SO2 with γ-Al2O3 using siteblocking experiments with ammonia, pyridine, and boron trifluoride and determined that the interaction of SO2 with basic sites on the surface led to the formation of chemisorbed SO2 while adsorption at acid sites led to physisorbed SO2. Also, Pacchioni et al.31 observed systemically SO2 uptake on the surface of the MgO powder. They suggested that four-

coordinated oxide anions commonly found on steps and corners of MgO(100), which have more basic character, may provide favorable sites for sulfate formation. In the present study, we could grossly estimate the total number of surface hydroxyl groups on the basis of eq 6. However, further information about the surface sites of different energies and acid-base properties still remains uncertain. Thus, the conversion of SO2 on the different iron phases that was observed here was not fully understood. 4.3. Proposed Mechanism of SO2 Uptake on Iron Oxides. Several prior studies have investigated the surface sites responsible for adsorbed SO2 following the reaction of SO2 on Al2O3, MgO, and dust particles.29-31 The gas-phase SO2 could not react with O2 in the air directly under ambient temperature, whereas it could react with O2 on the surface of those oxides and was transformed to the products HSO3-/SO32-, followed by the formation of the surface HSO4-/SO42- species. The surfaces of iron oxides covered with different types of hydroxyl groups were most likely involved in the production of both HSO4- and SO42- species while they interacted with SO2. In this study, the experiments were performed under “dry” conditions. However, some adsorbed H2O or O2 will be present because the samples were not subjected to high temperature or freeze-drying.23 Gaseous H2O and O2 are prone to forming active oxygen and hydroxyl on the surface. As shown in the DRIFT spectra, SO2 reacted with Fe2O3 particles and produced a larger fraction of the strongly adsorbed and surface-coordinated bisulfate and/or sulfate on the surface. Thus, a mechanism could be postulated for the reaction of SO2 on Fe2O3 particles (βFeOOH was an exception). First, the gaseous SO2 could be adsorbed on the surface of the oxides, followed by rearrangement where the sulfite species attached to the Fe atom through the sulfur atom, resulting in chemisorbed sulfite. The reactions were as follows:

]-FeIII(OH) + ]-H2O + SO2 f ]-FeIII(OH)•HSO3- + H+ (7) ]-FeIII(OH)•HSO3- f ]-FeIIIOSO2- + H2O

(8)

Sulfite was further transformed into the sulfate species by a series of free-radical propagation, termination, and productformation reactions. On the basis of the radical mechanism reported by Backstro¨m,32 the heterogeneous reactions were as follows:

]-FeIIIOSO2- f ]-Fe(II) + SO3•-

(9)

]-Fe(III) + SO3•- + ]-H2O f ]-Fe(II) + HSO4- + H+ (10) SO3•- + O2 f SO5•-

(11)

]-Fe(II) + H+ + SO5•- f ]-Fe(III) + HSO5- (12) ]-Fe(II) + HSO5- f ]-Fe(III) + SO4•- + OH-

(13)

6084 J. Phys. Chem. C, Vol. 111, No. 16, 2007

Fu et al.

]-Fe(II) + SO4•- + H+ f ]-Fe(III) + HSO4- (14) The total reaction could be expressed as Fe3+

HSO3- + 1/2O2 98 SO42- + H+

(15)

As shown above, surface Fe(III) gains one electron from HSO3- and is deoxidized to Fe(II). Likewise, Fe(II) could be oxidized readily to give Fe(III) by SO5•-, HSO5-, and SO4•-, which is generated via the reaction of SO3•- and O2. Fe(III) could in turn catalyze the autoxidation of S(IV). An Fe(III)Fe(II) redox cycling was thus established. In the presence of oxygen, the redox cycle is controlled by the formation of the peroxomonosulfate radical, SO5•-. This species is a much more reactive oxidant (E°(SO5•-/HSO5-) ) 1.1 V, pH ) 7) than O2 (E°(O2/O2•-) ) -0.33 V, pH ) 7) and may oxidize either sulfite ion or Fe(II).33 It has been reported that Fe(III) could be reduced by low-valence sulfur compounds such as SO2.34-36 The direct evidence that the occurrence of Fe(II) was involved in this process has been supplied by HPLC analysis. It is noteworthy that the Fe(II) concentration that was detected in our experiment accurately correlated to the previous report.34 Zhuang et al.36 observed that the concentration of Fe(II) in the dust increased continuously when Fe(III) was reduced by low-valence sulfur compounds during the long-range transportation. Because there is a thin aqueous layer resist on the sample surface, the dissolved Fe ions may be involved in this process. When the sample was exposed to SO2, the ferric oxide was partially dissolved to give Fe(III) ions in the watery surface at the lower pH, which in turn oxidized HSO3- to SO42-.37-39 The reaction is outlined below:

SO2 + H2O f H2SO3

(16)

H2SO3 f H+ + HSO3-

(17)

6H+ + Fe2O3 f 2Fe3+ + 3H2O

(18)

H2O + Fe3+ + HSO3- f Fe2+ + SO42- + 3H+ (19) However, the absence of any significant reaction between Fe2O3 and gaseous SO2 in the absence of O2 ruled out this proposed mechanism being operative and made the alternative associative mechanism, which is based on the adsorption of O2 and its subsequent dissociation on the oxide surface either in a separate step or concerted with S(IV), more probable. Berresheim and Jaeschke,40 in their study on the kinetics of SO2 removal by aerosols of CuSO4, CuCl2, and Cu(NO3)2, ascribed the catalytic activity of aerosols to the activity of dissolved metal ions present in the aqueous phase, although the correlation was quite poor. Quite likely, surface catalysis might be important in their case too. 5. Conclusions From the compositional and structural analyses for the iron oxides exposed to SO2 using various techniques such as DRIFTS, FTIR, XPS, IC, and HPLC, it was found that SO2 could be oxidized on most of the iron oxides to produce surface sulfate species at ambient temperature. The conformation of surface Fe(III)-sulfate species corresponded to the type of iron oxides. The order of catalytic reactivity for the heterogeneous oxidation of SO2 was R-Fe2O3 > γ-Fe2O3 > Fe3O4 > β-FeOOH > R-FeOOH, when the BET area was used to be the reactive surface area. In the absence of O2, the heterogeneous oxidation

of SO2 slowed greatly, implying that oxygen could be the key reactant of the reaction. The kinetics results suggested that the particle surface was the major cause of the catalysis of SO2 autoxidation on iron oxides. This is in agreement with the previous indications that the mineral particles may play an important role in adsorption and subsequent oxidation of SO2.6-8 Because the real atmospheric particle in the given region, such as the Loess tableland in China, presents relatively high Fe content, its influence on the conversion of SO2 in the atmosphere should be not neglectable. Further experimental work should be directed to observe the field-collected Fe-rich aerosol and develop a better understanding of these reactions. This will permit realistic extrapolation that can be used to perfect atmospheric mineralogy, which will be a topic in our future communication. Acknowledgment. This work is financially supported by the National Natural Science Foundation of China (Key project Grant Nos. 40533017, 20377008, and 40605001). References and Notes (1) Laurent, D.; Maud, L.; Karine, D.; Gilles, M.; Christian, G.; Nadine, C. Chem. ReV. 2005, 105, 3388. (2) Jambor, J. L.; Dutrizac, J. E. Chem. ReV. 1998, 98, 2549. (3) Leland, J. K.; Bard, A. J. J. Phys. Chem. 1987, 91, 5076. (4) Weckler, B.; Lutz, H. D. Eur. J. Solid State Inorg. Chem. 1998, 35, 531. (5) Hoffmann, P.; Dedik, A. N.; Ensling, J.; Weinbruch, S.; Sinner, T.; Gutlich, P.; Ortner, H. M. J. Aerosol. Sci. 1996, 27, 325. (6) Jickells, T. D.; An, Z. S.; Andersen, K. K.; Baker, A. R.; Bergametti, G.; Brooks, N.; Cao, J. J.; Boyd, P. W.; Duce, R. A.; Hunter, K. A.; Kawahata, H.; Kubilay, N.; Laroche, J.; Liss, P. S.; Mahowald, N.; Prospero, J. M.; Ridgwell, A. J.; Tegen, I.; Torres, R. Science 1995, 308, 67. (7) Adams, J. W.; Rodriguez, D.; Cox, R. A. Atmos. Chem. Phys. 2005, 5, 2679. (8) Zhuang, G. S.; Yi, Z.; Duce, R. A.; Brown, P. R. Nature 1992b, 355, 537. (9) Laskin, A.; Gaspar, D. J.; Wang, W.; Hunt, S. W.; Cowin, J. P.; Colson, S. D.; Finlayson-Pitts, B. J. Science 2003, 301, 340. (10) Usher, C. R.; Michel, A. E.; Grassian, V. H. Chem. ReV. 2003, 103, 4883. (11) Faust, B. C.; Hoffmann, M. R.; Bahnemann, D. W. J. Phys. Chem. 1989, 93, 6371. (12) Goodman, A. L.; Li, P.; Usher, C. R.; Grassian, V. H. J. Phys. Chem. A 2001, 105, 6109. (13) Usher, C. R.; Al-Hosney, H.; Carlos-Cuellar, S.; Grassian, V. H. J. Geophys. Res. 2002, 107, 4713. (14) Ullerstam, M.; Vogt, R.; Langer, S.; Ljungstrom, E. Phys. Chem. Chem. Phys. 2002, 4, 4694. (15) Ullerstam, M.; Johnson, M. S.; Vogt, R.; Ljungstrom, E. Atmos. Chem. Phys. 2003, 3, 2043. (16) Faust, B. C.; Hoffmann, M. R. EnViron. Sci. Technol. 1986, 20, 943. (17) Toledano, D. S.; Henrich, V. E. J. Phys. Chem. B 2001, 105, 3872. (18) Wang, L.; Zhang, F.; Chen, J. M. EnViron. Sci. Technol. 2001, 35, 2543. (19) Wang, L.; Song, G. X.; Zhang, F.; Chen, J. M. Chem. J. Chin. UniV. (Chinese) 2002, 23, 1738. (20) Wang, L.; Zhang, F.; Chen, J. M. Chem. J. Chin. UniV. (Chinese) 2002, 23, 866. (21) Wu, H. B.; Wang, X.; Chen, J. M.; Yu, H. K.; Xue, H. X.; Pan, X. X.; Hou, H. Q. Chin. Sci. Bull. 2004, 49, 1231. (22) Schwertmann, U.; Cornell, R. M. Iron Oxides in the Laboratory: Preparation and Characterization; Wiley-VCH: New York, 2000. (23) Zhang, X.; Zhuang, G.; Chen, J.; Ying, W.; Wang, X.; An, Z.; Zhang, P. J. Phys. Chem. B 2006, 110, 12588. (24) Peak, D.; Ford, R. G.; Sparks, D. L. J. Colloid Interface Sci. 1999, 218, 289. (25) Liu, J.; Yu, Y.; Mu, Y.; He, H. J. Phys. Chem. B 2006, 110, 3225. (26) Ni, Y.; Ge, X.; Zhang, Z.; Ye, Q. Chem. Mater. 2002, 14, 1048. (27) Kim, Y. J.; Park, C. R. Inorg. Chem. 2002, 41, 6211. (28) Hug, S. J. J. Colloid interface Sci. 1997, 188, 415. (29) Jams, W.; William, F. B. Thermochim. Acta 1996, 288, 179. (30) Karge, H. G.; Dalla Lana, I. G. J. Phys. Chem. 1984, 88, 1538. (31) Pacchioni, G.; Clotet, A.; Ricart, J. M. Surf. Sci. 1994, 315, 337. (32) Backstrom, H. J. Z. Phys. Chem. 1934, 25B, 122. (33) Christian, B.; Istvan, F.; Rudi, V. Inorg. Chem. 1994, 33, 687.

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