Oxidation of sulfur dioxide to sulfur trioxide over honeycomb DeNoxing


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I n d . Eng. Chern. R e s . 1993,32, 826-834

826

Oxidation of SO2 to SO3 over Honeycomb DeNoxing Catalysts Jiri Svachula; Louis J. Alemany,* Natale Ferlazzo, Pi0 Forzatti,' and Enrico Tronconi Dipartimento di Chimica Zndustriale e Ingegneria Chimica "G. Natta" del Politecnico, Piazza Leonard0 da Vinci 32, 20133 Milano, Italy

Fiorenzo Bregani Centro di Ricerca Termica e Nucleare, ENEL-DSR, Via Monfalcone 18, 20132 Milano, Italy

A systematic study addressing the effects of the operating conditions (contact time, temperature), of feed composition ( 0 2 , SOZ, H20, "3, NO,, NH3 NO, concentrations), and of catalyst design parameters (wall thickness, V content) in the oxidation of SO2 to SO3 over honeycomb commercialtype DeNoxing catalysts is described. Data are presented which refer to transient behavior of the catalysts, indicating that long conditioning times are required associated with the buildup of surface sulfate species. The steady-state reaction rate is of variable order in SOz, the order increasing with SOz concentration as long as this is below 200 ppm and then decreasing; it is asymptotically independent of oxygen, depressed by water, strongly inhibited by ammonia, and slightly enhanced by NO,. The apparent activation energy changes from -50 t o -20 kcal/mol on increasing the reaction temperature. A redox steady-state kinetic model is presented which accommodates qualitatively all of the observed effects. Once properly modified the same model has the potential to explain also transient effects during conditioning of the catalyst.

+

Introduction The selective catalytic reduction (SCR) of nitrogen oxides with ammonia is widely used for the control of NO, emission in flue gases from thermal power plants (Bosch and Janssen, 1988; Nakatsuji and Miyamoto, 1991). Commercial SCR catalysts consist of homogeneous mixtures of anatase TiO2, tungsta, and vanadia, along with minor amounts of silicoaluminates as mechanical promoters, and are employed in the form of monoliths or plates. Commercial catalysts are required to have several characteristics, including high DeNO, activity in a wide temperature window, high stability, and low SO2oxidation activity. In particular sulfur trioxide produced through the oxidation of SO2 along with SO3 already present in flue gases is known to react with ammonia and water to form ammonium bisulfate (NH4HSO4) or ammonium sulfate ((NH4)2S04) which may deposit in the cold equipment downstream of the SCR reactor causing corrosion problems and reduced performances, the deposition being controlled by the equilibrium between ammonium sulfates and NH3 + SO3 + H2O. This problem is so important that industrial specifications for SCR processestypically include upper bounds both on the NH3 slip and on the concentration of SO3 exiting the SCR reactor. The admissible limits of outlet SO3concentration (of the order of 10 ppm) correspond to practical SO2 conversions as low as 1-2 % . While several scientific and technical publications deal with DeNO, activity and stability of tungsta-vanadiatitania catalysts (EPAJEPRI, 1991; Bosch et al., 1986; Tuenter et al., 1986; Binder-Begsteiger et al., 1990; Chen and Yang, 1992; Tronconi et al., 1992; Svachula et al., 19931,papers addressing the oxidation of SO2 to SO3 over the same catalytic systems are scarce in the scientific literature.

* To whom correspondence should be addressed.

On leave from Department of Physical Chemistry, University of Chemical Technology, 53210 Pardubice, Czechoslovakia. f On leave from Department of Chemical Engineering, Campus Teatinos, 29071 University of Malaga, Spain.

In this paper we present a systematic study of the effects of the operating conditions (contact time, temperature), of feed composition (02, S02, H20, "3, NO,, NH3 + NO, concentrations), and of catalyst design parameters (wall thickness, V content) on the oxidation of SO2 to SO3 over honeycomb commercial-type SCR catalysts. Data are presented which refer to the conditioning of the catalysts and to the steady-state behavior of the catalysts. A kinetic model is derived to describe in a comprehensive way the results obtained by changing the operating conditions, the feed composition, and the catalyst characteristics. The capability of the model to explain transient effects during the conditioning of the catalyst is also addressed.

Experimental Methods Apparatus. Figure 1shows the schematic diagram of the apparatus used for the measurements of the oxidation of SO2 to S03. The reactant gas typically consisted of 1000 ppm SOz, 2% Oz,lO% H20, and balance Nz. The SO2 concentration is typical for flue gasesfrom combustion of fuels with 1.5-2 % w/w S content. In some experiments the reactant gas also contained NO, and/or NH3 as specified in the text; when employed, NH3 was directly injected at the top of the reactor and mixed to prevent side reactions such as the direct oxidation of ammonia. The flow rates of the individual gaseous streams were controlled by Brooks mass flowmeters; water was supplied by a metering micropump (GILSON Model 302). The reaction mixture was preheated and premixed. The reactor was an electrically heated stainless steel tube provided with an internal glass tube (i.d. 3 cm) to prevent the reactor wall from catalyzing the oxidation of SO2 to S03. The reactor was loaded with catalyst samples with square cross section consisting of 9 or 16 channels and with 15-cm length and was operated isothermally, as confirmed by a thermocouple sliding inside a capillary tube. The catalyst sample was wrapped with quartz wool and a bandage of ceramic material that prevented the outer surface of the catalyst from catalyzing the reaction and then was forced into the reactor to secure that no gas would

0888-5885/93/2632-0826$04.00/00 1993 American Chemical Society

Ind. Eng. Chem. Res., Vol. 32, No. 5, 1993 827

I

Table I. Catalyst Characteristics Vcontent

pitch (mm)

wall thickness (mm)

no. of channels

A B C D E

medium medium medium low high

=7 =7 -7 -7 =6

=1.4 =1.4 -1.1 -1.4 -1.1

9 9 9 9 16

Table 11. Results of SCR Activity catalyst

A B C D

E

SO

w

catalyst

3

ANALYSIS

Figure 1. Schematic diagram of the apparatus used for the measurements of the oxidation of SO2 to SOs: F, mass flowmetere; M, mixer; P, pump; R, reactor; C, catalyst bed; S, ammonia scrubber.

bypass the catalyst. The area velocity was 10 m/h(NTP) in most experiments. The area velocity AV is defined as AV = Q/(V&),where Q is the reactant gas flow rate in m3/h (NTP), V , is the total catalyst volume in m3, and a is the catalyst geometric surface area per unit volume in m2/m3. SO3 concentration in the product gas was determined through condensation of sulfuric acid at 90 "C in a glass spiral, followed by analysis with a Ionic Chromatograph Dionex Model Quic. It was found impossible to directly measure the very low conversions of Son, since they were typically limited to a very few percent. Blank experiments confirmed that the contribution of the apparatus to the formation of SO3could be safelyneglected. Special care was taken to ensure that the measurements refer to true equilibrium conditions rather than to transient conditions: transient effects were generally observed after changes in the experimental conditions and particularly in the reaction temperature and in the SO2 concentration, and typically lasted for several hours. The same apparatus was used to study the activity of the catalysts in the selective catalytic reduction of NO, with NH3 (Svachula et al., 1992). In this case the reaction mixture consisted of 500 ppm N,, 550 ppm "3,500 ppm S02,2% 02,10% HzO, and Nz balance, and the system was operated with AV = 33 m/h(NTP). In this case the gas exiting the reactor was scrubbed with a 6% aqueous solution of phosphoric acid to trap unconverted ammonia, cooled to 4 - 1 0 "C to condense water vapor, split into two streams, and eventually analyzed for NO/NO,, SOZ,and 02 contents using a chemiluminescence NO/NO, analyzer (Beckman, Model 955), a ND IR analyzer (Beckman, Model 865), and a paramagnetic oxygen analyzer (Beckman, Model 755). Catalysts. Different honeycomb commercial-type SCR catalysts were used in this study. The V loading of commercial SCR catalysts, expressed as Vz05 content, ranges between 0.3 and 2% w/w, and is homogeneously distributed across the thickness of the monolith walls by virtue of the preparation procedure. Such a uniform distribution of V is generally required in SCR applications in order to secure a constant level of activity in spite of possible abrasion by the particulate in the gas stream. The relevant characteristics of the catalysts, including V content, pitch, and wall thickness, are listed in Table I.

V content medium medium medium low high

T ("C) 320 320 320 320 320

NO, (m/h(NTP)) 33.4 35.5 35.2 23.5 49.1

Results DeNO, Activity. The DeNO, activity of the investigated catalysts is quoted in Table I1in terms of first-order overall catalyst activity constant NO, according to the expression k,ol = -AV ln(1- xNO,) (1) where

XNO, represents the fractional conversion of NO, = [([NOxlin - [NOxlout)/[NOxlin)l. With reference to the experimental conditions of this study the reactor operates under combined intraparticle and external diffusion control so that (Svachula et al., 1993; Beeckman and Hegedus, 1991; Lefers et al., 1991) (XNO,

l/kNOz = l/kc + l/kg

(2)

where k, is the pseudo-first-order rate constant of the surface chemical reaction and k, represents the interphase gas-solid mass-transfer coefficient. Notice that k, is an effective rate constant incorporating the influence of mass transfer in the catalyst pores, the DeNO, reaction being strongly limited by intraparticle diffusion (Tronconi et al., 1992;Beeckman and Hegedus, 1991;Lefers et al., 19911, and that it is referred to the geometric catalyst surface. The results indicate that the SCR activity depends primarily on the vanadium content and that it increases on increasing the V loading. A more detailed discussion of the effects of the operating parameters and of the catalyst characteristics on NO, removal over honeycomb SCR catalysts is presented by Svachula et al. (1993). Oxidation of SO2 to SO3 1. Conditioning of the Catalysts. To obtain significant and reproducible results in the oxidation of SO2 to SO3, the catalyst must be conditioned. The conditioning procedure implies operation at the reaction temperature with gaseous mixtures containing sulfur dioxide, oxygen, water, and nitrogen. Figure 2 shows the results of a typical conditioning experiment: the concentration of SO3 at the reactor exit increaseswith time until a steady-state level is approached. When the flow of sulfur dioxide is stopped, the concentration of SO3 decreases slowly, indicating that SO3 is released from the catalyst (Figure 3). If the SO2 supply is resumed, the outlet concentration of SO3 gradually recovers its original value. These data prove that the conditioningof the catalyst is associatedwith a slow process of buildup of metal oxide sulfates onto the catalyst surface (it may take several hours up to -70 h depending on the experimental conditions and the catalyst) and proceeds until a steady-state concentration of sulfates is reached. The steady-state concentration of sulfates appears to be limited by the formation-decomposition equilibrium of metal oxide sulfates via so3 adsorption-desorption processes.

828 Ind. Eng. Chem. Res., Vol. 32, No. 5, 1993

I 0

I

I

I

I

10

20

30

40

1

TIME (h)

1

3

2

TI=

4

5

(h)

Figure 3. Concentration of SO3 at the reactor exit after stopping the flow of SO2 (standard conditions): catalyst E, T = 330 "C. 30 -I4

I/ 0

500

lo00 SO

1

2

3

4

5

6

7

O2 CONCBNTRATION (w1.S)

Figure 2. Results of a typical conditioning experiment during the oxidation of SO2 (standard conditions): catalyst E, T = 330 O C .

0

0

1500

COWCBNTFWTION (ppm)

Figure 4. Effect of SO2 concentration on the oxidation of SO2 (standard conditions): catalyst B, T = 380 "C.

2. Effect of SO2. Figure 4 shows the results of experiments performed with different SO2 concentrations

Figure 5. Effect of 0 2 concentration on the oxidation of SO2 (standard conditions): catalyst B, T = 380 "C.

in a wide range over catalyst B. Diagnostic calculations performed by invoking the analogy with the GraetzNusselt heat-transfer problem to compute local masstransfer coefficients for SO2 (Hawthorn, 1974; Tronconi and Forzatti, 1992) showed that negligible SO2 concentration gradients are present at the gas-olid interface. Likewise, also intraporous limitations can be neglected on the basis of the Weisz-Prater criterion (Froment and Bishoff, 19791, since the SO2 oxidation reaction is considerably slower than SO2 diffusion within the catalyst pores. Accordingly the process is controlled by the chemical reaction, which is in line with very low measured SO2 conversions. The plot of SO2 conversion in Figure 4 exhibits a maximum, suggesting an apparent kinetic order in SO2 higher than 1for low SO2 concentrations (0-200 ppm) and a decreasing fractional order at higher inlet SO2 concentrations. Such a peculiar behavior is typical of SCR catalysts, since it has been observed in our laboratory in the case of several honeycomb catalysts when the effect of SO2 concentration was carefully investigated. It is worth noticing that, for practical purposes, a firstorder dependence on SO2 may provide a reasonable approximation in the range 0-1O00 ppm SOZ,in line with several technical reports (EPA/EPRI, 1991; Bosch et al., 1986). 3. Effect of 0 2 . The effect of oxygen concentration on the oxidation of SO2 is plotted in Figure 5. It appears that the rate of reaction is almost independent of oxygen concentration above 0.5-1 % v/v oxygen levels, due to the excess of oxygen. A dependence becomes manifest only for 02 concentrations comparable to those of SO2 (=lo00 ppm = 0.1% v/v), as expected. However, it is worth stressing that the specific oxygen concentrations where a dependence of the SO2 conversion becomes apparent depend both on the experimental conditions and on the catalyst redox properties. In any case the conversion of SO2 is almost independent of oxygen partial pressure over SCR commercial catalysts for representative operating conditions (02> 2% v/v). 4. Effect of Water Vapor. The addition of water to the reaction mixture results in a significant decrease of , 9 0 2 conversion (Figure 6). The inhibiting effect of water is apparent at low water concentrations and diminishes above 5% vlv water content: accordingly the rate of reaction is practically independent of the concentration of water vapor in the range of practical interest ( 6 1 5 % v/v). The original activity measured in the absence of water is restored within -1 h and then keeps constant if

Ind. Eng. Chem. Res., Vol. 32,No. 5, 1993 829 6

4

U

I

5

3

5IN

Y

gN

C

4

't I

I

1

5

10

4

3 1.4

H20 CONCENTRATION (%r)

1.6

1.5 lOOO/T

Figure 6. Effect of H20 concentration on the oxidation of SO2 (standard conditions): catalyst A, T = 380 O C . the flow of water is stopped: this time is much shorter than the characteristic time of catalyst conditioning as deduced from Figure 1 (1h << 2&70 h). The inhibition of water is likely explained in terms of reversible blockage of the surface vanadyl sulfates, which are the most acidic vanadium sites present on the surface and which are associated with the active sites for the oxidation of SO2. The inhibiting effect of water can hardly be explained in terms of destruction of the active sites since its restoration by reconditioning of the catalyst would require a much longer time than observed. 5. Effect of Reaction Temperature. The temperature dependence of the oxidation of SO2 to SO3 was investigated using a reactant mixture consisting of 1000 ppm S02,2% 02,10% H20, and balance N2. The data were analyzed by assuming a simplified plug-flow reactor model and a rate equation first order in SO2 and zero order in both 0 2 and H2O in view of the results previously discussed. Accordingly the pseudo-first-order rate constant of SO2 oxidation, kSOn, can be expressed as follows:

(3) where Q is the total gas flow rate (m3/h(NTP)),V , is the total catalyst volume (m3), and xsol is the conversion of SO2 (0 I%soyI1). Notice that V,includes the volumes of both the ceramic catalytic material and the empty channels of the catalyst. Equation 3 implies that the inhibiting effect of water vapor is incorporated into the kinetic rate constant kSOz and that the volume of the whole ceramic catalytic matrix participates effectively in the reaction due to the absence of intraparticle, diffusional limitations. This point is in line with the effect of catalyst wall thickness that will be discussed below. The effect of temperature for one of the investigated catalysts is shown in Figure 7. The apparent activation energy changes from -50 to =20 kcal/mol in the temperature range 36&420 "C. The change in the energy of activation with temperature was also observed for several other SCR catalysts and similar values of the activation energies were calculated; the transition occurs at lower temperature on increasing the V content. The values of the activation energies are consistent with an overall process controlled by the chemical reaction, in contrast with DeNO, reaction where apparent activation energies as low as -5 kcal/mol were typically measured (Svachula

Figure 7. Effect of temperature on the oxidation of SO2 (standard conditions): catalyst D.

0

500

1OOO

NOxCOWCENTFUTION (ppm)

Figure 8. Effect of NO, concentration on the oxidation of SOz (standard conditions): catalyst A, T = 380 O C .

et al., 1993). As a matter of fact, the rate of the oxidation of SO2 is much slower than that of NO reduction by "3, which eventually explains why the former reaction is chemically controlled and the latter one is limited by both interphase and intraparticle diffusions. It is worth mentioning that the break in the Arrhenius plots with temperature was also reported for commercial sulfuric acid catalysts with similar values of the activation energies (Boghosian et al., 1989;Balzhinimaevet al., 1986;Livbjerg and Villadsen, 1972;Xie and Nobile, 1985;Doering and Berkel, 1987). 6. Effect of NO,. The effect of NO, concentration on SO2 conversion is shown in Figure 8. It appears that the conversion is slightly enhanced in the presence of NO,: this effect was confirmed over other catalysts and could possibly be related either to the contribution of the gasphase oxidation of SO2 by NO,, which is known to take place in the chamber process for the production of sulfuric acid (West and Duecker, 19741,or to the role of NO, in the catalyst redox process. 7. Effect of "3. The addition of small amounts of ammonia to the reaction mixture strongly inhibits the oxidation of SO2 as shown in Figure 9,the activity being almost suppressed in the presence of 100ppm "3. When the flow of ammonia is stopped, the activity is restored

830 Ind. Eng. Chem. Res., Vol. 32, No. 5, 1993

4

w u

3

s E

8

2

si" 1

0 50

0

100

"ti CONCENTRATION (ppm) 3

Figure 9. Effect of NH3 concentration on the oxidation of SO2 (standard conditions): (+) catalyst E, T = 330 O C ; ( 0 )catalyst E, T = 340 'C.

I 0

10

I

I

20

30

0.2

0.4

0.6

0.8

NH3/N0

150

I

TIWE ( h )

Figure 10. Transient effects of NH3 inhibition on the oxidation of SO2 for catalyst E, T = 330 'C: lo00 ppm S02,2% 02,10% H20, 100 ppm "3, NZbalance; lo00 ppm SOZ, 2% 02,10% HzO, N2 balance; lo00 ppm SOZ,2% Oz,lO% HzO, 100 ppm NO,, Nz balance.

within a couple of hours but the final activity is slightly lower than the original one as shown in Figure 10. Noticeably the characteristic time of this process is much shorter than the characteristic time of catalyat conditioning (2-3 h << 20-70 h). This eventually indicates that the active sites responsible for the oxidation of SO2 are deactivated but not destroyed upon adsorption of ammonia so that the recovery of the original activity is reasonably fast and does not imply reconditioning of the catalyst. The catalytic activity can be restored completely upon admitting NO, into the reactant mixture, and it remains constant if the flow of NO, is stopped afterward (Figure 10). These data are explained assuming that some residual ammonia is retained adsorbed onto the catalyst surface, which results in the deactivation of a small number of active sites. However, this type of adsorbed ammonia is destroyed upon admission of NO, because it reacts with NO, according to the SCR reaction. This could eventually

Figure 11. Effect of NH3 + NO, in the oxidation of SO2: catalyst A, T = 380 O C , lo00 ppm SOZ,2% Oz,lO% HzO, 500 ppm NO,, NH3 = variable, Nz balance.

suggest that different active sites are involved in SO2 oxidation and that those reversibly poisoned by ammonia are either most active or most abundant and somehow less acidic. The fact that the catalytic activity remains essentially constant after the NO, flow is stopped is apparently in contrast with the previously reported promoting effect of NO, upon SO2 oxidation. However, a comparison with the data in Figure 8 suggest that such an effect is scarecely significant in this case due to the small concentration of NO, employed in the experiment of Figure 10. 8. Effect of NH,+ NO,. The combined effect of NH3 + NO, displayed in Figure 11shows that (i) a t very low NHs/NO, values the SO2 conversion is greater than in the absence of the two species, indicating that the behavior is dominated by the promoting effect of NO, (see Figure 8);(ii) for growing values of NHdNO, the inhibiting effect of ammonia becomes predominant, progressivelyreducing the SO2 conversion; and (iii) the NH3 inhibition is lower than in the absence of NO, (seeFigure 91, because ammonia is consumed according to the DeNO, SCR reaction. 9. Effect of Area Velocity. The relation between SO2 conversion and contact time (VdQ= 1IAV) obeys eq 3, as shown in Figure 12. Considering that the volume of the catalyst V is proportional to the volume of the ceramic catalytic material, it turns out that the whole volume of the ceramic catalytic material effectively oxidizes SO2 to SO3. 10. Effect of Catalyst Wall Thickness. Table I11 shows an approximate linear dependence of SO2 conversion on wall thickness for catalysts with comparable medium vanadium content. This is consistent with an overall process controlled by chemical reaction so that the entire volume of the catalytic material effectively participates in the oxidation process, and compares well with the effect of the area velocity as already discussed. An important conclusion from these results is that the oxidation of SO2 to SO3 can be lowered by reducing the thickness of the ceramic monolith. It is alsoworth stressing that such a reduction does not significantly affect the efficiencyof the DeNO, reaction since this reaction occurs essentially a t the catalytic geometric surface,being heavily controlled by intraparticle diffusional resistances. This is confirmed in Table 11, where the comparable NO, are reported for the catalysts with similar vanadium contents

Ind. Eng. Chem. Res., Vol. 32, No. 5, 1993 831 silica carrier, and due to the presence of alkali metals they operate as supported liquid-phase catalysts under reaction conditions (Kenney, 1980; Urbanek and Trela, 1980). Indeed between 440 and 600 O C the alkali metal sulfatevanadium pentoxide mixtures in equilibrium with SO2/ Sodair mixtures are all liquid (Tandy, 1956).

N

2

0

50

100

150

l W / A V (h/Nm)

Figure 12. Effect of AV in the oxidation of SO2 (standard conditions): catalyst C, T = 380 O C . Table 111. Effect of Catalyst Wall Thickness on SO2 Oxidation (Reaction Temperature = 360 "C) catalyst wall thickness (mm) conversion ( % ) 21.4 1.4 A ~1.4 1.5-1.6 B =l*l 1.1 C Table IV. Effect of V Loading on SO2 Oxidation (Reaction Temperature = 360 "C) catalyst V loading conversion ( % ) D low 0.2 A medium 1.4 6.3a E high a Calculated by extrapolating the result of 4.0% conversion of SO2 measured at 340 "C and using the value EaCt = 24 kcal/mol secured by regression of experimental data in the range 300-340 "C.

but with different wall thicknesses. Along these lines improved SCR commercial catalysts have recently been developed by the catalyst suppliers (EPA/EPRI, 1991). 11. Effect of Vanadia Content. In Table IV we report the SO2 conversion for catalysts with different vanadium loadings and similar wall thickness (1.2-1.35 mm). Although the catalysts might differ in other relevant properties, it is apparent that the conversion of SO2 to SO3increases markedly with V in line with previous reports in the technical literature (Drews et al., 1989). This indicates that the oxidation of SO2 depends primarily on the V content and that it can be properly controlled by using catalysts with low V loadings. It is worth mentioning that in the case of commercial sulfuric acid catalysts it has been proposed that the active sites are constituted by dimericvanadyl species (Balzhinimaevet al., 1986;Glueck and Kenney, 19681,which would imply a second-order dependence of the SO2 oxidation rate constant kSOn on V content. In conclusion,it seems reasonableto assume that the active sites for the oxidation of SO2 are constituted by vanadyls with sulfate ligands, as already proposed for commercial sulfuric acid catalysts (Balzhinimaev et al., 1984;Ivanov and Balzhinimaev, 1987). Although commercial sulfuric acid catalysts represent an important reference for SCR commercial systems for what concerns the activity in the SO2 oxidation reaction, it must still be stressed that they are rather different. As a matter of fact, the commercial sulfuric acid catalysts consist of VzO5 and alkali metal oxides deposited onto an inactive porous

Discussion Analysis of Transient Effects during SO2 Oxidation. The data in Figures 2 and 3 indicate that the characteristic time for the activation of SO2 oxidation (Figure 2) and the characteristic time for the decay of SO3 evolution from the catalyst (Figure 3) are both much greater than the characteristic time of reaction. Such results can be interpreted by identifying the active sites for SO2 oxidation with dimeric vanadyls with sulfate ligands. This identification compares well with the consolidated picture of active sites in commercial sulfuric acid catalysts (Ivanovand Balzhinimaev, 1987). They are quoted in the followingas V2-SO3. These sites are formed through adsorption with reaction of SO3 formed by SO2 oxidation onto oxidized binuclear vanadyl centers. Thus, the process responsible for the formation of active sites can be described by the reaction (4)

with Vp,oxrepresenting oxidized dimeric V sites. Accordingly,the concentration of active sites is expected to increase with time during catalyst conditioning due to the forward reaction (4), until the formation of these sites becomes limited by their formation-decomposition equilibrium. Correspondingly,the catalyst activity and hence the SO3evolution follow a similar development,increasing with time. Likewise, in the absence of ,902 in the feed mixture it is expected that the concentration of active sitesdecreases due to the reverse of reaction 4, sincegaseous SO3 is continuously removed from the system. Due to the fact that the characteristic time of this process is larger than that of reaction 4, SO3 monitored at the reactor exit decreases slowly with time. It must be considered however that so3could have been adsorbed onto surface sites other than the active ones or the most active ones for SO2 oxidation. These sites could be possibly associated with tungstyl groups or other basic groups (e.g., titanyl groups) and might account for some of the SO3 released in Figure 3. As an alternative explanation of Figure 2, one may assume that the rate of SO2 oxidation is virtually constant throughout the conditioning time, being independent of sulfate concentration, whereas the evolution of 803 increases due to progressive saturation of the catalyst surface. According to this picture, a plot of SO2 conversion versus conditioning time should be invariant. However, if ,902 oxidation promotes the formation of active sites, as supposed previously, a plot of SO2 conversion should parallel the growing trend exhibited by gas-phase SO3 concentration in Figure 2. Unfortunately, discrimination between the two hypotheses was prevented by the impossibility of obtaining reliable measurements of the low SO2 conversions. It is worth noting, however, that the identification of the active sites with sulfate species provides a rationale also for the observed variation of the kinetic order of SO2, as discussed in the following section. A further point worthy of note is that the conditioning of the catalyst is required in the case of SO2 oxidation but apparently not in the case of the DeNO, reaction. Considering that the conditioning procedure involves accumulation of sulfates, that sulfates are known to promote the removal of NO, up to ~ 2 w/w, % and that

832 Ind. Eng. Chem. Res., Vol. 32, No. 5, 1993 -1% w/w of stable sulfates is already present in fresh commercial SCRcatalysts (Drewset al., 1989),a significant promoting effect would be expected during conditioning under a purely chemical regime for the removal of NO,. However, under SCR conditions this process is limited by both interphase and intraparticle diffusion so that the expected change in the overall catalyst activity is small and can hardly be appreciated. As discussed in Svachula et al. (1993), this is primarily due to the presence of stable sulfates in fresh commercial catalysts which ensures high catalytic activity already before the catalyst has been conditioned. As for the sulfur distribution within the catalyst, we notice that the characteristic times of conditioning for SO2oxidation (tdo,), of SO2oxidationreaction ~, (tso,), of SO2 or NO intraparticle diffusion ( ~ D s otDNO,) and of NO, reduction reaction (tNO,) obey the following relative order of magnitude:

tcso, >

SO, > ~ D S O ,

> NO*

~ D N O ~

(5)

Therefore, uniform intraparticle sulfate concentrations are always expected in the course of the conditioning procedure of SCR catalysts. It follows that the fact that conditioning is not required in the case of NO, removal cannot be explained in terms of a preferential sulfation of the outer layer of the catalyst along with the low catalyst effectiveness factors for the DeNO, reaction (Tronconi et al., 1992). Steady-State Kinetics of SO2 Oxidation. The reported effects of the different operating and design variables can be reconciled on a qualitative basis according to a comprehensive kinetic equation along the lines described below. On the basis of the previous discussion, the active sites for the oxidation of SO2 to SO3 can be identified with dimeric vanadyl sulfates. At steady state their surface concentration can be modeled assuming adsorption equilibrium: [v2-s031,x

= ~so,cso,[v2lox

(6)

where [V2-S0310x, [V2lOx,and KSO,represent the steadystate concentrations of sulfated and not sulfated oxidized dimeric V sites and the equilibrium constant for SO3 adsorption, respectively. The inhibition of water and ammonia is accounted for by assuming adsorption equilibria on the vanadyl sulfates which eventually result in the reversible deactivation of the active sites that are no longer available for the SO2 oxidation reaction. Noticeably, vanadyl sulfates correspond to the most acidic sites containing vanadium that are present at the catalyst surface, so it seems reasonable to neglect the adsorption of NH3 and of H20 at, e.g., oxidized dimeric V sites. The two adsorption processes above can be described by

stand for adsorption equilibrium constants for reactions 7 and 8. On the basis of the different concentrations of the poisoning agents (7% v/v of H2O vs ppm of "3) and on the inhibition effects shown in Figures 6 and 9, it can be deduced that K", >> KH?Oin line with the basicity strengths of NH3 and H2O. Under the experimental conditions typical of the SCR process the rate of SO2oxidation appearsto be independent of oxygen partial pressure, exhibits a variable kinetic order in SO2 concentration (see Figures 4 and 51, and is far from equilibrium. These results can be accounted for by assuming that the reaction proceeds through a redox mechanism, where SO2 is first adsorbed at a sulfated dimeric vanadium site, (V2-S03),, + SO2 (V2-S03-S02)o, (11) followed by a slow step involvingreduction of the vanadium site accompanied by evolution of SO$ Q

(V2-S03-S02),, -,(V2-SO3Ird + SO3 (12) Finally, reoxidation of the catalytic site results in the regeneration of the active sites:

-

(V2-s03),ed + '/202 (V2-SO3)ox (13) Assuming adsorption equilibrium of SO2 at the active vanadium sites, we obtain from reaction 11 By requiring that the rate of catalyst reduction equals the rate of reoxidation, we obtain the concentration of the reduced vanadium sites as

where kred and k,, are the rate constants for reactions 12 and 13, respectively. Then, the following site balance equation can be written:

+ [V~-SO~]O, + [V2-S031red + [

[v210x

~

~

-

+~

[V~-SO~-H2010x + [V2-S03-NH3],, = C, (16) which upon substitution from eqs 6,9,10,14, and 15 yields

[ v , ] =~ Ctot/{l ~ + KSO,cSO,[l

+ KS02cS02(1 + kred/

(k,xCoz1'2))+ KH,OCH,O + K",C",lJ

(17)

Cbtrepresents the concentration of total active sites. Since they involve vanadyl dimeric species, it is therefore proportinal to the square of the vanadium percent (w/w) in the catalyst. Upon imposing that the rate of reaction equals the rate of catalyst reduction 14, (18) rso, = rred = kre,[V2-S02-S0310x and substituting for [V2-SOrS0310, in eq 18, after in thenumerator introducingan empiricalterm (1+ ~CNO,) accounting for the promoting effect of NO,, we obtain eventually rSO,

= [(kredKS02KS0,Cb~)CS02CS03(1 + bcNO)]/{l + Ks03cS03[1+ KsO,cSO,(1 + ~red/(koxcOz1'2)) +

+ KNH3CNH3]~ (19) Effects of Operating Variables and of Catalyst Characteristics. Equation 19 is able to explain the [VrSO~H2010x and [ V d 0 d & l a r represent the steadyvariable kinetic order of the SO2 oxidation reaction state concentrations of the sites poisoned through adsorption of water and of ammonia, and KH,Oand K N H ~ suggested by the conversion curve in Figure 4. Indeed the KH,OcH,O

~

~

-

Ind. Eng. Chem. Res., Vol. 32, No. 5, 1993 833 term Kso,Cso,in the numerator accounts for an apparent order in SO2 partial pressure higher than 1 at low SO2 (and SO3) concentrations, since under these conditions the terms KSO~CSO~ and Kso3Cso3appearing in the denominator can be neglected (Kso2Cso2<< l ; Kso3Cso3<< 1). On increasing the SO2 and SO3 concentrations, however, such terms become comparable with unity and cannot be neglected any longer; accordingly the apparent order in SO2 is reduced below 1. Notice that the rate of reaction is nil in the absence of SO3;accordingly a finite althoughvery small concentration of SO3must be assumed at the beginning of the conditioningprocedure, which could be likely associated with the presence of stable sulfates in fresh commercial SCR catalysts (Drews et al., 1989). It is worth stressing that, because of the term [V2-SO&,x = Kso,Cso,[V210xin eq 6, the concentration of active sites is assumed to increase with an increasing concentration of SO3, i.e., of a reaction product. The ~ ~ ~ I I I K H ~ OinCthe H ~denominator O of eq 19accounts for the inhibiting effect of water reported in Figure 6: the inhibition is more important at low water partial pressures and weakens above 5 1 0 % water. Considering that flue gases typically contain 10-20% v/v water vapor, in practice eq 19could be simplified by incorporating the dependence on the partial pressure of water into an overall rate constant kso,. The term KNH~CNH, accounts for the strong inhibiting effect of ammonia displayed in Figure 9. The inhibition is almost suppressed when the flow of ammonia is stopped since the sites poisoned due to the absorption of ammonia can be easily regenerated via the reverse of reaction 8. The term [l kre,J(koxC~zl/2)] in the denominator of eq 19 is suitable to represent the asymptotic effect of the oxygen concentration displayed by the data in Figure 5. The promoting effect of NO, over SO2 oxidation is included in eq 19 by the empirical term (1+ ~CNO,); as already mentioned, this term could likely be associated with the contribution of the gas-phase oxidation of SO2 by NO, or to the role of NO, in the catalyst redox process. It is worthy of note that in the presence of ammonia the promoting effect of NO, is overcompensated by the inhibiting effect of ammonia, since a small overall inhibition is measured. Indeed, under typical SCR conditions the extents of NO, promotion and NH3 inhibition appear to be controlled by the interaction between the kinetics of SO2 oxidation and the kinetics of the DeNO, reaction, as suggested by the differences observed between the individual effects of the two species (Figures8 and 9) and their combined effect NH3 are cofed to the reactor (Figure 11). when NO, Such an interaction is of a complex nature, since the experimental conversion of SO2 in this case is an integral result determined by both axial and intraporous concentration gradients of NO, and NH3 in the monolithic catalyst. This point will be further investigated in a forthcoming paper. Concerning the effect of temperature, apparent activation energies of =20 and =50 kcal/mol were calculated in the high- and low-temperature regions, respectively. The transition between the two regions was observed to shift to lower temperatures on increasing the V content. The values calculated in the two regions are consistent with those reported by several authors for the oxidation of SO2 over sulfuric acid catalysts. It is worth noticing that a decrease in the overall apparent activation energy with temperature can be anticipated from eq 19 due to the temperature dependence of the adsorption equilibrium constants Kso,. Besides,the shift of the Arrhenius "break"

+

+

at lower temperature on increasing the V content is explained by the larger values of Cbt on increasing the V loading. As for the catalyst characteristics, the effect of the vanadia content (Table IV) is reflected in eq 19 through the term Cbt which represents the concentration of total active sites, namely of oxidized and sulfated dimeric V sites. The results of Figure 10 indicate that at least two types of sites are likely involved in the reaction, and that those reversibly poisoned by ammonia appear to be either the most active or the most abundant. These results can be explained by assuming that the most active sites are constituted by dimeric vanadyls, whereas the other sites are constituted by tungstyls that are known to be less active and more acidic in nature (Ramis et al., 1992) and therefore more strongly poisoned by ammonia. Inspection of Table I1 reveals that also the activity in the removal of NO, depends primarily on the V content and increases with increasing V loadings. According to the analysis reported in a companion paper (Svachula et al., 1993),the DeNO, reactivity would depend linearly on the V loading, whereas a quadratic dependence is foreseen by eq 19 for the SO2 oxidation reactivity. Such a difference may provide a rationale for the lower levels of V adopted in manufacturing recent commercial SCR catalysts, which would in fact result in a more marked reduction of SO3 formation than of NO, removal. Finally the data in Table I11 indicate that the SO2 conversion is approximately proportional to the catalyst wall thickness, which is in line with the observation that the oxidationprocess is controlled by the chemical kinetics so that the whole volume of the ceramic catalytic material does participate effectively in the reaction. Preliminary results of a systematic kinetic analysis of SOz oxidation for a given SCR honeycomb catalyst, which will be presented in a forthcoming paper, confirm the capability of eq 19to account quantitatively for the effects of the operating variables under steady-state conditions. For the purposes of the present paper however it is worth stressing that the proposed kinetic equation was found to describe qualitatively the catalytic behavior of all the investigated SCR catalysts and therefore lends itself as the adequate kinetics for SO2 oxidation over commercial catalysts under representative SCR conditions. Conclusions The following effects were observed for the oxidation of SO2 to SO3over honeycomb tugsta-vanadia-titania SCR catalysts: 1. The overall oxidation process is controlled by the chemical reaction as opposite to the DeNO, process due to the very small activity in SO2 conversion. 2. Conditioning of the catalyst is required to obtain significant and reproducible data, with the conditioning process lasting 10-70 h and being associated with the buildup of sulfates on the catalyst surface until adsorptiondesorption equilibrium is established. Transient effects of comparable length are observed also after changes of the experimental conditions and are similarly related to the reequilibration of surface sulfates. 3. The steady-state reaction rate is of variable order in SO2, with the order increasing with SO2 concentration as long as this is below 200 ppm and then decreasing. The rate is asymptotically independent of oxygen, depressed by water, strongly inhibited by ammonia, and slightly enhanced by NO,. The apparent activationenergy changes from =50 to -20 kcallmol on increasing the reaction temperature.

834 Ind. Eng. Chem. Res., Vol. 32, No. 5, 1993

4. The transient effects observed in SO2 oxidation both during catalyst conditioning and following changes of the

experimental conditions, as well as the rate dependence on SO2 concentration, can be explained by assuming that the oxidation of SO2 is catalyzed by the surface sulfate species, whose steady-state concentration is determined b y adsorption-desorption equilibrium of SO3, the reaction product. 5. The rate of SO2 oxidation depends linearly on the catalyst wall thickness, indicating that the whole volume of the catalytic monolith is active in the reaction, and increases markedly on increasing the V content, in line with the assumption of catalytic active sites involving sulfated dimeric vanadyl species. 6. Based on the above assumption, a redox steady-state kinetic model is derived which accommodatesqualitatively the experimental effects both of the operating variables and of the catalyst characteristics. Once properly modified, the same model has potential to also explain the transient effects observed during catalyst conditioning and upon modifying of the experimental settings.

Acknowledgment

Binder-Begsteiger, I.; Herzog, G. W.; Megla, E.; Tomann-Rosos, M. Kinetics of the DeNO, reaction over Ti02-WO3 honeycomb catalysts. Chem. Zng. Tech. 1990,62,60. Boghosian, S.;Fehrmann, R.; Bjerrum, N. J.; Papatheodorou, G. N. Formation of Crystalline Compounds and Catalyst Deactivation during SO2 Oxidation in VzOb-MzSz07 (M=Na,K,Cs) Melts. J . Catal. 1989,119,121. Bosch, H.; Janssen, F. Catalytic Reduction of Nitrogen Oxides: A Review on the FundamentalaandTechnology. Catal. Today 1988, 2,369. Bosch, H.; Janasen, F.;vander Kerkhof, F.; Oldenziel,J.; van Ommen, J.; Ross, J. The Activity of Supported Vanadium Oxide Catalysts for the Selective Reduction of NO with Ammonia. Appl. Catal. 1986,25,239. Chen, J. P.; Yang, R. T. Role of WO3 in Mixed VzOb-W03/Ti02 Catalysts for Selective Catalytic Reduction of Nitric Oxide with Ammonia. Appl. Catal. A: Gen. 1992,80,135. Doering, F. J.; Berkel, D. A. Comparison of Kinetic Data for K/Vand Cs/V Sulfuric Acid Catalysts. J . Catal. 1987,103,126. Drews, R.; Hess, K.; Hdlderich, W.; Ruppel, W.; Scheidsteger, 0. Effectively prevent deactivation. Chem. Ind. 1989,47. EPAIEPRI. 1989Joint Symposium on Stationary CombusionNO, Control, San Francisco, March 6-9;E P A Research Triangle Park, NC, 1989. EPAIEPRI. 1991 Joint Symposium on Stationary Combusion NO, Control, Washington, DC, March 25-28; EPA: Research Triangle Park, NC, 1991.Froment. G. F.: Bishoff. K. B. ChemicalReactor Anulvsis and Design: Wiley: New York, 1979. Glueck, A. R.; Kenney, C. N. The Kinetics of the Oxidation of Sulfur Dioxide over Molten Salts. Chem. Eng. Sci. 1968,23,1257. Hawthorn, R. D. Afterburner Catalysts. Effects of Heat and Mass Transfer between Gas and Catalyst Surface. AZChE Symp. Ser. 1974,137,428. Ivanov, A. A.; Balzhinimaev,B. S. New Data on Kinetics and Reaction Mechanism for SO2 Oxidation over Vanadium Catalysts. React. Kinet. Catal. Lett. 1987,35,413. Kenney, C. N. In Catalysis; Chemical Society. Specialist Periodical Repork, Wright: Bristol, 1980;p 123. Lefers,J. B.; Lodder, P.; Enoch, G. D. Modelling of SelectiveCatalytic Denox Reactors-Strategy for Replacing Deactivated Catalyst Elements. Chem. Eng. Technol. 1991,14,192. Livbjerg, H.; Villadsen, J. Kinetics and Effectiveness Factor for SO2 Oxidation on an Industrial Vanadium Catalyst. Chem. Eng. Sci. 1972,27,21. Nakatsuji, T.; Miyamoto,A. Removal Technologyfor Nitrogen Oxides and Sulfur Oxides from Exhaust Gases. Catal. Today 1991,10, 21. Ramis, G.; Busca, G.; Cristiani, C.; Lietti, L.; Forzatti, P.; Bregani, F. Characterization of Tungsta-Titania Catalysts. Langmuir 1992, 8, 1744. Svachula, J.; Ferlazzo, N.; Tronconi, E.; Forzatti, P.; Bregani, F. Selective Reduction of NO, by NH3 over Honeycomb DeNoxing Catalysts. Znd. Eng. Chem. Res. 1993,in press. Tandy, G. H. The Role of Alkali Sulfates in Vanadium Catalysts for Sulfur Dioxide Oxidation. J . Appl. Chem. 1956,6, 68. Tronconi, E.; Forzatti, P. Adequacy of Lumped Parameter Models for SCR Reactors with Monolith Structure. AZChE J . 1992,38, 201. Tronconi, E.; Forzatti, P.; Gomez Martin, J. P.; Malloggi,S. Selective Catalytic Removal of NO,: a Mathematical Model for Design of Catalyst and Reactor. Chem. Eng. Sci. 1992,47,2401. Tuenter, G.; van Leeuwen, W. F.; Snepvangers, L. J. M. Kinetics and Mechanism of the NO, Reduction with NH3 on V205-WO3-Ti02 Catalyst. Ind. Eng. Chem. Prod. Res. Dev. 1986,25,633. Urbanek, A.: Trela, M. Catalvtic Oxidation of Sulfur Dioxide. Catal. Rev. Sci. Eng. 1980,21,j3. West. J. R.: Duecker. W. W. In Riezel's Handbook of Industrial Chemistry, 7th ed.;Kent, J. A., Ed.yVan Nostrand Reinhold: New York, 1974;p 62. Xie, K. C.; Nobile, A,,Jr. Discontinuities in the Rate of Sulfur Dioxide Oxidation on Vanadium Catalysts. J . Catal. 1986,94,323. I

This work was performed under contract with ENEL/ DSR/CRTN Milano.

Nomenclature a = geometric surface area per unit volume of the catalyst, mYm3 b = adaptive parameter in eq 14 AV = Q/V area velocity, m/h(NTP) [V-i-j] = concentration of unoccupied active sites and of active sites occupied by species i and j C,,, = concentration of total active sites NO, = overall catalyst activity constant, m/h(NTP) k, = gas-solid mass-transfer coefficient, m/h(NTP) k, = effective rate constant of the chemical reaction, m/h(NTP)

kso2 = rate Constant of SO2 oxidation (eq 3) Ki = adsorption equilibrium constant for component i [NO,]i, = inlet concentration of NO,, ppm [NOxlout= outlet concentration of NO,, p p m Q = reactant gas flow rate, mYh(NTP) &SOp = time of conditioning of catalyst for SO2 oxidation, h tso2 = time of SO2 oxidation reaction, h ~ D S O=~time of SO2 intraparticle diffusion, h t D N 0 = time of NO intraparticle diffusion, h t N O = time of NO reduction reaction, h T = temperature, "C V, = catalyst volume, m3 XNO, = fractional conversion of NO, xso2 = fractional conversion of SO2

Literature Cited Balzhinimaev, B. S.; Ponomarev, V. E.; Boreskov, G. K.; Ivanov, A. A. Studies of fast Relaxations in SO2 Oxidation on Active Components of Vanadium Catalysts. React. Kinet. Catal. Lett. 1984,25, 219. Balzhinimaev, B. S.;Ponomarev, V. E.; Belyaeva, N. P.; Ivanov, A. A.; Boreskov, G. K. Steady State Kinetic Equation for SO3 Oxidation onVanadiumCatalysts. React. Kinet. Catal.Lett. 1986, 30,23. Beeckman, J. W.; Hegedus, L. L. Design on Monolith Catalysts for Power Plant NO, Emission Control. Ind. Eng. Chem. Res. 1991, 30,969.

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Received for review July 30, 1992 Revised manuscript received December 7, 1992 Accepted January 4, 1993