Reaction Kinetics for the Catalytic Oxidation of Sulfur Dioxide with

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Ind. Eng. Chem. Res. 2007, 46, 80-86

Reaction Kinetics for the Catalytic Oxidation of Sulfur Dioxide with Microscale and Nanoscale Iron Oxides Yonghui Shi and Maohong Fan* Center for Sustainable EnVironmental Technologies and Department of CiVil, Construction, and EnVironmental Engineering, Iowa State UniVersity, Ames, Iowa 50011

This paper contains a study of the oxidation of sulfur dioxide using microscale and nanoscale iron oxide as catalysts. A comparison of the catalytic performance of microscale and nanoscale iron oxides showed that nanoscale iron oxide generally is more effective than microscale iron oxide in regard to catalyzing the oxidation of sulfur dioxide. The reaction orders, with respect to the reactants sulfur dioxide and oxygen, were 1 and 0.24, and 1 and 0.30 when microscale and nanoscale iron oxides, respectively, were used as catalysts. Furthermore, the activation energy of catalytic oxidation of sulfur dioxide with microscale iron oxide is 88.9% higher than that with nanoscale iron oxide. Introduction Sulfur dioxide (SO2) is the main source of acid rain, and its emission is strictly restricted by the Clean Air Act (CAA) amendments of 1990. SO2 in the atmosphere originates from many sources, such as coal-fired power plants, petroleum refineries, and diesel engines that use high-sulfur fuel. Among all of the SO2 sources, the power plants contribute more than 60% of the pollution and are considered the major concern in the United States Environmental Protection Agency (USEPA)’s SO2 reduction effort. Far more than a contributor of acid rain, SO2 is also associated with human respiratory disease. Those individuals, especially the elderly and children, and patients with heart or lung disease, are most vulnerable to the effects of SO2 pollution.1,2 Despite its negative effect on the environment and human health, SO2 is also an important raw material to produce a variety of chemicals in industries. Basically, SO2 can be converted to other chemicals through three routes: reduction, non-redox, and oxidation reactions. SO2 is often reduced to sulfur by a modified Claus process, which has been widely used for sulfur recovery in the petroleum refining or chemical industries.3 The non-redox SO2 utilization approaches are relatively simple and direct. For example, SO2 can be absorbed by sodium hydroxide or sodium carbonate to produce sodium sulfite. This approach can also be used to extract organic acids from anaerobic fermentation broth.4,5 The oxidation of SO2 is the most frequently used method for its utilization. Many chemicals can be generated with the oxidation-based approach. For example, SO2 can be oxidized to SO3 to produce fuming sulfuric acid through a contact process. Other oxidation-based SO2 utilization includes the production of sulfur-containing fertilizers and sodium sulfate, which is an important chemical that is used in the manufacture of soap, paper, and glass.6,7 Another oxidation-based application was developed in our research group where SO2 was used to synthesize polymeric ferric sulfate,8 which is a new coagulant for water and wastewater treatment.9-11 The key step in oxidation-based SO2 utilization is expressed as

SO2 + 0.5O2 98 SO3

(R1)

* To whom correspondence should be addressed. Tel.: 515-2943951. E-mail address: [email protected].

Under the standard-state conditions (i.e., 25 °C and 1 atm), the change in Gibbs free energy of the reaction (∆G0) is -71 kJ/ mol.12 Therefore, thermodynamically, the oxidation is feasible. However, the chemical kinetics of reaction R1 limits its applicability. The oxidation is too slow at room temperature to be valuable for industry. To accelerate the rate of SO2 oxidation, catalysts and high temperature are used. The currently used catalyst is vanadium pentoxide. There are two problems that are associated with the use of vanadium pentoxide. First, a high temperature is needed to achieve a high SO2 oxidation efficiency.13 Second, the application of vanadium pentoxide in industry raises a concern because research has shown that it is a pulmonary carcinogen.14,15 Precious-metal catalysts can be used to avoid these problems, but they are very expensive and easily vitiated by certain impurities in SO2. A catalyst that is active at a low temperature, inexpensive, and environmentally friendly would be very desirable. Iron oxide (Fe2O3) is an inexpensive alternative to the catalysts currently used. It occurs naturally as a mineral called hematite and is fairly active and selective for several heterogeneous catalytic reactions. Microscale Fe2O3 has been successfully used as a catalyst for the oxidation of many air pollutants such as polychlorinated dibenzodioxin/dibenzofuran16 and SO2.17,18 Compared to microscale Fe2O3, nanoscale Fe2O3 features a smaller particle size, higher specific surface area, and greater concentration of catalytic active sites. These characteristics make nanoscale Fe2O3 a more promising catalyst with significantly improved catalytic performance over its microscale counterpart. Unlike the expensive noble-metal-based catalysts, nanoscale Fe2O3 catalyst is also expected to provide a more economical solution for modern industries. Recent studies have been made on the catalytic performance of nanoscale Fe2O3 in many oxidation processes. For example, using nanoscale Fe2O3 as a catalyst, Li et al. reported that CO can be removed efficiently through catalytic oxidation.19 However, nanoscale Fe2O3 has never been tested as a catalyst for the oxidation of SO2. This paper focuses on deriving the chemical kinetics, reaction orders, and Arrhenius expressions of SO2 oxidations using microscale and nanoscale Fe2O3 as catalysts. The main objective of the research is to provide a basis for comparing microscale and nanoscale Fe2O3 in catalyzing the oxidation of SO2, in terms of economic and operational benefits.

10.1021/ie060889d CCC: $37.00 © 2007 American Chemical Society Published on Web 12/05/2006

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Figure 1. Experimental setup of the oxidation of SO2. (Legend: 1, certified SO2; 2, oxygen gas; 3, nitrogen gas; 4, flowmeter; 5, Lindberg/Blue M TF55030A-1 tube furnace; 6, quartz tube reactor; 7, catalyst powder; 8, quartz wool; 9, Permapure MD-110-48F-2 Nafion concentric tube dryer; 10, particulate filter; 11, California Analytical ZRF NDIR SO2 analyzer; and 12, data acquisition system.)

Experimental Section Materials. Microscale Fe2O3 was procured from Bailey-PVS Oxides, LLC (Canonsburg, PA) and used as-received. BrunauerEmmett-Teller (BET) surface analysis indicated that the microscale Fe2O3 has a specific surface area of 4.0 m2/g. The nanoscale Fe2O3 was purchased from Mach I, Inc. (King of Prussia, PA). It is a brown superfine powder with a large specific surface area of 240.0 m2/g, which was specified by the manufacturer and confirmed by BET surface analysis that was conducted at Iowa State University. Pure O2 and N2, as well as certified SO2 and O2, were purchased from Linweld, Inc. (Des Moines, IA) and used in the experiments. The certified SO2 and O2 had concentrations of 5000 and 2500 ppm, respectively, and were balanced with N2. Structure Characterization with Transmission Electron Microscopy (TEM). The structures of both microscale and nanoscale Fe2O3 were characterized with transmission electron microscopy (TEM) bright-field imaging and selected-area electron diffraction (SAED) methods. The characterizations were performed with a Philips CM30 electron microscope equipped with a LaB6 electron source that was operated at an acceleration voltage of 300 kV. Apparatus and Operation Procedures. The laboratory setup used for the research is shown in Figure 1. The SO2 oxidation experiments were performed using a quartz tube reactor with an inner diameter of 6 mm. In the middle of the quartz tube, a 2-cm-long column of catalyst powder was located with quartz wool on both sides to avoid catalyst loss that was caused by gas flow. For microscale Fe2O3, the amount used in each test was 400 mg, and that of nanoscale Fe2O3 was 40 mg. The quartz tube with catalyst was then placed inside a TF55030A-1 tube furnace that was made by Lindberg/Blue M (Asheville, NC). The temperature of the tube furnace was controlled with a UT150 temperature controller (Yokogawa M&C Corporation, Newnan, GA). The flows of the SO2, O2, and N2 were controlled with a C-03217-52 150-mm Teflon correlated flowmeter (Cole Parmer, Vernon Hills, IL), a C-03229-11 150-mm correlated flowmeter (Cole Parmer), and a Gilmont Instrument 150 mm direct reading flowmeter (Cole Parmer), respectively. The flowmeters were calibrated with a bubble meter before the experiments. The total flow rate of the mixture gas was 50 mL/ min, which was used for all the tests. Before entering the quartz tube, SO2, O2, and N2 were mixed to produce a gas mixture with predetermined concentrations of SO2 and O2. The gas mixture was balanced with nitrogen in

each test. To remove water vapor and particles that may be present in the system, the effluent gas from the quartz tube first passed a Permapure Model MD-110-48F-2 Nafion concentric tube dryer (Permapure, Toms River, NJ) and then through a Cole Parmer 0.2-µm in-line particulate filter. Finally, the SO2 concentration after the reaction in the effluent gas stream was measured with a California Analytical Model ZRF NDIR gas analyzer. The analyzer was connected to a computer system that recorded the SO2 concentrations, which were used to calculate the oxidation efficiencies of SO2. A continuous packed-bed catalytic reactor was used in all the SO2 oxidation tests in this study. The degree to which the volumetric flow rate of gas stream before and after oxidation is changed should be evaluated, to determine whether the changes have significant effects on the accuracies of kinetic model derivations. Assuming that (i) the flow rate of the gas mixture (expressed in units of mol/s) fed into the reactor is F, (ii) the molar ratios of SO2, O2, and N2 in the feeding gas mixture are ySO2,0, yO2,0, and yN2,0, respectively, and (iii) the SO2 oxidation ratio to SO3 is X, then the effluent stream consists of SO2, O2, SO3, and N2, with flow rates (expressed in units of mol/s) of ySO2,0F(1 - X), (yO2 - 1/2XySO2,0)F, XySO2,0F, and yN2,0F, respectively. Therefore, the flow rate of gas mixture leaving the reactor (expressed in units of mol/s) can be expressed as

1 nTotal ) 1 - XySO2,0 F 2

(

)

(1)

The maximum SO2 concentration used in the entire research was 2000 ppm, which resulted in only 0.1% change in the volume of gas mixture. Therefore, the change in volume of gas mixture before and after reaction can be considered to be negligible. Baseline tests were conducted under different temperatures to check whether oxidation of SO2 occurred without the presence of catalysts. The tested temperature range is 500-1000 °C, with an interval of 50 °C. The flow rates of 30, 40, 50, 60, and 70 mL/min were used to study the effects of the total flow rate on the SO2 oxidation efficiency under the temperatures of 470 and 330 °C for microscale and nanoscale Fe2O3, respectively. The concentration of SO2 used in this particular study was 1200 ppm. To derive the reaction order for SO2, experiments were arranged with SO2 concentrations of 400, 800, 1200, 1600, and 2000 ppm, while the initial concentration of O2 was 50 vol %. In all the experiments for calculating the SO2 reaction order, the concentrations of O2 in the gas mixture were at least 250

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Figure 2. (a) Bright-field transmission electron microscopy (TEM) image of the microscale Fe2O3 particles; (b) selected area electron diffraction pattern of one of the microscale Fe2O3 particle along the [001] zone axis.

Figure 3. (a) Bright-field TEM image of the nanoscale Fe2O3 particles; (b) selected area electron diffraction pattern of the nanoscale Fe2O3 particles.

times greater than the stoichiometric amount of O2 needed for complete oxidation of SO2. To establish the reaction order, with respect to O2, the initial concentrations of O2 used ranged from 200 ppm to 1000 ppm with an interval of 200 ppm, and the initial concentrations of SO2 were always twice that of O2. The microscale and nanoscale Fe2O3-based catalytic oxidation tests used to obtain the reaction order for SO2 were undertaken at 500 °C and 320 °C, respectively. The temperatures selected for conducting tests to calculate the reaction order for O2 were 650 and 400 °C, with respect to microscale and nanoscale Fe2O3, respectively. Results and Discussion Structure of Microscale and Nanoscale Fe2O3. According to Figures 2a and 3a, the bright-field TEM images showed that the size of microscale Fe2O3 is in the range of 100-200 nm, whereas that of the nanoscale Fe2O3 is only ∼3 nm. Figure 2b, which is the SAED pattern of one of the microscale Fe2O3 particles along the [001] zone axis, showed that Fe2O3 has a hexagonal structure with lattice parameters of a ) 5.0 Å and c ) 13.6 Å. The SAED pattern of nanoscale Fe2O3 in Figure 3b indicated that the Fe2O3 particle is in the polycrystalline form, which also has a hexagonal structure. Baseline. The results of tests on the noncatalytic oxidation of SO2 are presented in Figure 4. This figure shows that SO2 cannot be oxidized at all by O2 without the help of a catalyst, even at a temperature of 1000 °C. Therefore, catalysts are

Figure 4. Effect of temperature on SO2 oxidation efficiencies, with respect to microscale and nanoscale Fe2O3. (Conditions: total flow rate, 50 mL/ min; SO2 concentration, 2000 ppm; mass of microscale Fe2O3, 400 mg; mass of nanoscale Fe2O3, 40 mg; and running time, 10 h.)

necessary for the oxidation of SO2 to occur in an economically feasible temperature range. Effects of Flow Rate. The effects of flow rate on the efficiencies of SO2 oxidation catalyzed by nanoscale and

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of both nanoscale and microscale Fe2O3, which consequently decreased the catalytic capacities of both materials. Generally, nanoscale Fe2O3 was more effective than microscale Fe2O3 in reagrd to catalyzing the oxidation of SO2. Reaction Orders for SO2 and Oxygen. Assuming thatCSO2,0, CO2,0, and CSO3,0 are the initial concentrations (given in terms of molarity, M) of reactants SO2, O2, and SO3, and CSO2, CO2, and CSO3 represent the concentrations (given in terms of molarity, M) of SO2, O2, and SO3 in the effluent stream when the reaction is in steady state, then the concentration of consumed SO2 in the steady state can be expressed as CSO2,0 CSO2. Based on the stoichiometric relationship between SO2 and O2 that has been shown as reaction R1, the consumed O2 and produced SO3 at the steady state can be expressed as 0.5(CSO2,0 - CSO2) and CSO2,0 - CSO2. Assuming that the volumetric flow rate of gas mixture is V0 (expressed in units of L/s), then the rates of SO2 consumption, O2 consumption, and SO3 generations ∆FSO2, ∆FO2, and ∆FSO3, respectively (expressed in units of mol/s)scan be expressed as follows:

Figure 5. Effect of the total flow rate on SO2 oxidation efficiencies, with respect to microscale and nanoscale Fe2O3. (Conditions: SO2 concentration, 1200 ppm; temperature for the reaction catalyzed with microscale Fe2O3, 470 °C; temperature for the reaction catalyzed with nanoscale Fe2O3, 330 °C; mass of microscale Fe2O3, 400 mg; and mass of nanoscale Fe2O3, 40 mg.)

microscale Fe2O3 are shown in Figure 5. This figure reveals that the SO2 oxidation efficiencies obtained under the catalysis of nanoscale and microscale Fe2O3 were unaffected by the flow rate in the tested flow rate range at the given test conditions, even though there was a large difference in specific surface areas associated with nanoscale and microscale Fe2O3. The fact indicates that the rate-limiting step of SO2 oxidation might be the gas-solid mass-transfer process (affected by surface areas of nanoscale and microscale Fe2O3) at the beginning of the reaction; however, external mass-transfer resistance can be negligible if the gas flow rate is kept sufficiently high for the studied flow-through (rather than flow-over) gas-solid reaction systems. Effect of Temperature. The changes in the SO2 oxidation efficiency in the presence of microscale and nanoscale Fe2O3, in the temperature range of 100-1000 °C, are depicted in Figure 4. These results indicated that microscale and nanoscale Fe2O3 catalysts can substantially enhance the oxidation of SO2. Furthermore, four more facts are shown in Figure 4. First, the figure indicates that the efficiency of SO2 oxidation under catalysis of the microscale or nanoscale Fe2O3 was a function of temperature. Second, the onset temperature of SO2 oxidation for nanoscale Fe2O3 was ∼150 °C, whereas that for microscale Fe2O3 was 200 °C. The difference in the onset temperature may be explained by the fact that the nanoscale Fe2O3 is typically better than microscale Fe2O3 in reducing the activation energy for the oxidation of chemicals.19 Third, SO2 can be almost completely oxidized when nanoscale Fe2O3 was used as a catalyst at 450 °C, whereas the maximum oxidation efficiency of SO2 was only 84% under the catalysis of microscale Fe2O3, even at 650 °C. Fourth, the efficiencies of SO2 oxidation catalyzed by both Fe2O3 decreased as the temperature increased above 650 °C and 500 °C for microscale and nanoscale Fe2O3, respectively. This phenomenon might result from the sintering of nanoscale and microscale Fe2O3 after 500 and 650 °C, respectively. The sintering reduced the specific surface areas

∆FSO2 ) V0(CSO2,0 - CSO2)

(2)

∆FO2 ) 0.5V0(CSO2,0 - CSO2)

(3)

∆FSO3 ) V0(CSO2,0 - CSO2)

(4)

The reaction rates of reaction R1 (given in units of mol/(g s))s rSO2, rO2, and rSO3, in terms of SO2, O2, and SO3, respectivelys can be expressed as follows:

rSO2 ) rO2 ) -

∆FSO2 ∆W

∆FO2 ∆W

(5) (6)

and

rSO3 )

∆FSO3 ∆W

(7)

where ∆W is the mass of catalyst loaded (given in grams). It is obvious that the following relationships among rSO2, rO2, and rSO3 exist:

-rSO2 ) -2rO2 ) rSO3

(8)

In all the tests, the inlet and outlet concentrations of SO2 in the gas streams were measured. The reaction rate, in terms of SO2 (rSO2), was used to calculate the rate constant and reaction orders of reaction R1, using the following relationship:

-rSO2 )

V0(CSO2,0 - CSO2) ∆W

) kCSO2RSO2CO2RO2

(9)

where k represents the reaction rate constant and the exponents RSO2 and RO2 are the reaction orders, with respect to reactants SO2 and O2, respectively. Reaction orders RSO2 and RO2 are not necessarily integers, because reaction R1 is not an elementary reaction. Special initial reactant concentrations were chosen to derive the reaction orders RSO2 and RO2. When RSO2 was to be established, CO2,0 was designed to be much larger than CSO2,0, so that CO2, at any time of the reaction, can be considered to be

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a constant. The reaction rate of reaction R1, with respect to SO2, then can be expressed as

V0(CSO2,0 - CSO2) ∆W

) k′CSO2RSO2

(10)

where k′ ) kCO2RO2. Taking the logarithm of eq 10 yields eq 11:

ln

(

)

V0(CSO2,0 - CSO2) ∆W

) ln k′ + RSO2 ln CSO2

(11)

Based on eq 11, a series of tests were designed at a specified temperature T, where the initial concentrations of SO2 were varied and the initial concentrations of O2 were set at a level much higher than those of SO2. For each test under a given initial SO2 concentration (CSO2,01), two data (ln(V0(CSO2,0 - CSO2)/ ∆W)1 and CSO2,1) were collected. Based on the data obtained from different initial SO2 concentrations (CSO2,0), a graph of ln(V0(CSO2,0 - CSO2)/∆W) versus lnCSO2 can be plotted to calculate RSO2. Because of the limit of the concentration measurement range (0-10 vol %) of the SO2 analyzer used in the project, the reaction order RO2 cannot be established with the same method as that previously mentioned. To derive the reaction order RO2, reactions were designed in such a way where the initial concentrations of SO2 were always 2 times greater than those of O2 at any temperature T. Consequently, at any time of the reaction, the concentrations of SO2 and O2 in the effluent have the following relationship:

CSO2 ) 2CO2

(12)

The reaction rate, with respect to SO2, then can be expressed as

-rSO2 )

V0(CSO2,0 - CSO2) ∆W

( )

) kCSO2RSO2

CSO2

RO2

2 k

)

CSO2RSO2 +RO2 (13) 2RO2 Taking the logarithm of eq 13 yields eq 14:

ln(-rSO2) ) ln k - RSO2 ln 2 + (RSO2 + RO2) ln CSO2

(14)

Based on eq 14 and RSO2 obtained from eq 11, the reaction order RO2, along with the reaction rate constant (k), at a given temperature T, can be derived. According to eq 11, the reaction orders for SO2 for the microscale and nanoscale Fe2O3 catalysts (RSO2,Micro and RSO2,Nano) are shown as the slopes of the lines in Figure 6. This figure shows that the reactions are both first-order for SO2, regardless of which type of Fe2O3 catalyst was used. The reaction orders, with respect to O2, for the microscale and nanoscale Fe2O3 catalysts (RO2,Micro and RO2,Nano) are also shown in Figure 6. According to Figure 6, eq 14 is written as eq 15:

ln(-rSO2) ) -4.84 + 1.24 ln CSO2

(15)

The equation set (eqs 16 and 17) can be obtained as follows from eq 15:

RSO2,Micro + RO2,Micro ) 1.24

(16)

ln k - RO2,Micro ln 2 ) -4.84

(17)

Figure 6. Determination of the reaction orders for SO2 and O2 when microscale and nanoscale Fe2O3, respectively, were used as catalysts. Reaction orders for SO2: total flow rate, 50 mL/min; oxygen concentration, 50 vol %; SO2 concentrations, 400, 800, 1200, 1600, and 2000 ppm; temperature, 500 °C for microscale Fe2O3 and 320 °C for nanoscale Fe2O3. Reaction orders for O2: total flow rate, 50 mL/min; oxygen concentrations, 200, 400, 600, 800 and 1000 ppm; SO2 concentrations, 400, 800, 1200, 1600 and 2000 ppm; temperature, 650 °C for microscale Fe2O3 and 400 °C for nanoscale Fe2O3.

Therefore, the reaction order, with respect to O2 for using microscale Fe2O3 as a catalyst, RO2,Micro, is 0.24, and the natural logarithm of the reaction rate constant at a temperature of 650 °C is -4.67. Using the same method, the reaction order for using nanoscale Fe2O3 as a catalyst, RO2,Nano, is 0.30, and the natural logarithm of the reaction rate constant at a temperature of 400 °C is -2.19. Consequently, the rate of SO2 consumption for the microscale and nanoscale Fe2O3 as catalysts (-rSO2,Micro and -rSO2,Nano, respectively) can be expressed as

-rSO2,Micro ) kMicroCSO2CSO20.24

(18)

-rSO2,Nano ) kNanoCSO2CO20.30

(19)

where kMicro and kNano represent the reaction rate constants when microscale and nanoscale Fe2O3, respectively, are used as catalysts. Reaction Constant (k) and Apparent Activation Energy (Ea). The relationship between reaction rate constant (k) and reaction temperature (T) can be expressed in the empirical Arrhenius equation

( )

Ea k ) A exp RT

(20)

where A is the pre-exponential factor, Ea the apparent activation energy of the reaction, and R ideal gas constant. When the relationships of ln k versus 1/T are plotted, Ea and A in eq 20 can be derived using the slope (-Ea/R) and intercept (log A) of the plot, respectively. The Arrhenius relationship was determined by changing the reaction temperature while the reactant concentrations remained unchanged. For the reaction that used microscale Fe2O3 as a

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Figure 7. Determination of the apparent activation energy (Ea) and preexponential factor (A) for the reactions catalyzed by microscale and nanoscale Fe2O3, respectively. (Conditions: total flow rate, 50 mL/min; oxygen concentration, 800 ppm; SO2 concentration, 1600 ppm; temperature range, 450-650 °C for microscale Fe2O3, and 300-500 °C for nanoscale Fe2O3; and temperature interval, 50 °C.)

catalyst, the temperature range was 450-650 °C, with an interval of 50 °C. For the experiments that used a nanoscale Fe2O3 catalyst, the temperature range was 300-500 °C, with an interval of 50 °C. Fixing the input SO2 and O2 concentrations as 1600 and 800 ppm, respectively, a series of reaction rates (-rSO2) can be obtained according to eq 9 at different temperatures. Consequently, the reaction rate constants k, at different temperatures, can be calculated using eq 14. When the microscale Fe2O3 is used as a catalyst, the relationship between ln k and 1/T is plotted in Figure 7. From this figure and eq 20, the apparent activation energy (Ea,Micro) and pre-exponential factor (AMicro) for microscale Fe2O3 were calculated to be 32.8 kJ/mol and 0.7, respectively. In the case of a nanoscale Fe2O3 catalyst, the relationship between ln k and 1/T also is plotted in Figure 7. The apparent activation energy (Ea,Nano) and pre-exponential factor (ANano) for nanoscale Fe2O3 are 17.4 kJ/mol and 2.6, respectively. The Arrhenius equations for the microscale and nanoscale Fe2O3 catalysts can be written as

( 3945 T ) 2093 ) ) 2.6 exp(T )

kMicro (mol-0.24 dm3.72 s-1 g-1) ) 0.7 exp -

(21)

kNano (mol-0.30 dm3.90 s-1 g-1

(22)

The pre-exponential constant in eq 26 associated with nanoscale Fe2O3 is much larger than that with microscale Fe2O3, as shown in eq 27. This might be caused by the large difference in specific surface area between nanoscale and microscale Fe2O3. The larger specific area of nanoscale Fe2O3 could provide more chances for the collisions of the two reactants, SO2 and O2, on the surface of the catalyst, possessing the correct orientation, to lead to the formation of SO3. Equations 26 and 27 indicate that the apparent activation energy of the reaction decreased by ∼50% when using nanoscale Fe2O3 as a catalyst, in comparison with the case using microscale Fe2O3, which was mainly attributed to the difference

in structure associated with nanoscale and microscale Fe2O3. The polycrystalline structure of nanoscale Fe2O3 was more favorable to the reduction of activation energy of reaction between SO2 and O2 than that of microscale Fe2O3. The activation energy is a direct indication of whether a reaction is liable to happen. This explains the experimental results observed in this study, that the reaction is likely to occur at lower temperatures (