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To develop a highly active calcium oxide high-temperature desulfurization sorbent, a method of preparation of macroporous calcium oxides from lime was...
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Ind. Eng. Chem. Res. 1998, 37, 3943-3949

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Preparation of Macroporous Lime from Natural Lime by Swelling Method with Acetic Acid for High-Temperature Desulfurization Eiji Sasaoka* and Norimasa Sada Faculty of Environmental Science and Technology, Okayama University, Tsushima-naka 2-1-1, Okayama 700, Japan

Md. Azhar Uddin Faculty of Engineering, Okayama University, Tsushima-naka 3-1-1, Okayama 700, Japan

To develop a highly active calcium oxide high-temperature desulfurization sorbent, a method of preparation of macroporous calcium oxides from lime was studied. This method is composed of two steps: swelling of the lime and calcination of the swelled sample. Swelling occurred when lime was exposed to the vapor of acetic acid. The swelling resulted from calcium acetate formation in the sample. The swelling rate was at a maximum in the presence of acetic acid and depressed by the presence of water vapor. The swelled sample was converted to macroporous calcium oxide by heating to 850 °C. The reactivity of the macroporous calcium oxide for the removal of SO2 or H2S in the presence of H2O vapor was higher than that of the calcined raw limestone. In particular, its SO2 removal capacity and the oxidative character of CaS to CaSO4 and CaO were greatly improved by this swelling method. These characteristics were also compared with those of a sample prepared from limestone by this swelling method. Introduction Limestone is a very important material as a hightemperature desulfurization sorbent: limestone is used for SO2 capture in fluidized bed combustors of coal and can also be used in coal gasifiers for in-bed removal of H2S. In coal combustors, limestone usually decomposes into CaO and CO2 and then reacts with SO2: CaO formed from limestone is converted to CaSO3 and then CaSO4.1 In gasifiers of coal, decomposition of limestone into CaO and CO2 depends on the pressure of CO2 and the temperature.2-4 Therefore, sulfidation reactions of CaO and/or that of CaCO3 have to be considered. Furthermore, the product CaS has to be converted to CaSO4 before disposal because H2S is released from the reaction between CaS and water.5 The conversion of CaS to CaSO4 is a complex reaction.6 If lime (or limestone) absorbs SO2 or H2S, the internal structure of the lime is changed when these solid products are produced. Because the molecular volumes of CaCO3, CaO, CaS, and CaSO4 are 36.9, 16.9, 28.9, and 46.0 cm3/g, respectively, the intraparticle pores become smaller in size and may be finally plugged by the product if the molecular volume of the product is larger than that of the reactant.3,7 These problems have been indicated by some previous researchers. Hartman et al. and Zevenhoven et al. reported pore plugging due to the molar volume change on conversion of CaO to CaSO4 (the molar volume ratio CaSO4/CaO: ca. 2.7).8,9 Anderson et al. reported that the CaSO4 from CaO constituted an adhering barrier that limited the extent of the reaction.1 Efthimiadis and Sotirchos reported that the sulfidation rate of lime was affected by the * To whom correspondence should be addressed. Telephone: 086-251-8442. Fax: 086-251-8442. E-mail: sasaokae@ cc.okayama-u.ac.jp.

connectivity of the intraparticle pores (molar volume ratio CaS/CaO: ca. 1.7).10 Fenouil and Lynn reported that the reaction kinetics is controlled by the diffusion of H2S through the pores of the CaS product layer around the lime particle.2 Yrjas et al. supposed that the high conversion of lime to CaS resulting from the calcination of limestone caused increased particle porosity.3 In the case of the conversion of CaS to CaSO4 (the molar volume ratio CaSO4/CaS ca. 1.6), Torres-Ordonez et al. reported that if the layer of CaSO4 covered the inside of the CaS pore, the CaS was not converted to CaSO4 until the CaSO4 decomposed into CaO, SO2, and O2 at a high temperature.11 We have found a method for the preparation of a highly active macroporous CaO from coarse limestone particles.12 In the present work, a preparation method of macroporous CaO from a coarse natural lime particle was examined. The lime was prepared from a coarse natural limestone particle by heating to 850 °C. It was found that the lime swelled when exposed to a vapor of acetic acid-water. The swelled sample was dried and calcined to a macroporous CaO. In the present paper, the swelling of lime and the pore structure of the macroporous lime thus obtained are described and compared to lime prepared from limestone. The reactivity of the macroporous lime prepared from natural lime is also reported in comparison with that of the macroporous lime from limestone that was previously reported.12 Experimental Section Swelling and Calcination Procedure. Table 1 shows the composition of the raw limestones used in this study. The two limestones originated from different areas of Japan (Okayama and Chubu). The lime was mainly obtained by heating the limestones at 850 °C

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3944 Ind. Eng. Chem. Res., Vol. 37, No. 10, 1998 Table 1. Chemical Composition (wt %) of Limestones limestone

CaO

MgO

SiO2

Al2O3

Fe2O3

Ig loss

Okayama Chubu

54.7 55.1

0.23 0.75

0.05 0.02

0.02 0.02

0.03 0.03

44.3 43.6

for 1 h (ca. 36 min of heating needed to reach 850 °C) in an air atmosphere. The heating temperature was set at 700, 750, or 800 °C in some cases. Approximately 2 cm3 of the sample lime (0.7 mm particle diameter) was placed in a cylindrical glass sample tube (i.d. ) 27 mm), which was placed unplugged in a cylindrical polystyrene sample tube (i.d. ) 60 mm; volume ) 200 cm3). An acetic acid aqueous solution, ca. 30 cm3, was put into the polystyrene sample tube, and this tube was plugged. This polystyrene sample tube was placed in an incubator at a controlled temperature. The sample lime particles, exposed to acetic acid-water vapor, were stirred occasionally. The sample volume was measured by a tapping method using a measuring cylinder. The fractional swelling of the sample was calculated according to following equation:

Figure 1. Effect of concentration of CH3COOH(aq) on the swelling of lime (Okayama) at 40 °C.

fractional swelling of sample, f ) volume of swelled sample (1) volume of nonswelled sample The swelled samples were dried (60 °C for 4 h + 110 °C for 21 h) and then calcined in a muffle oven in air. The calcination temperature of 850 °C, measured above the samples in the muffle oven, was reached after about 36 min of heating and retained for 1 h. After the calcination, the sample particles had become partially sintered in the ceramic dish and were crushed and sieved to average diameter 0.7 mm. The bulk density of the sample was measured and compared with that of the nonswelled sample. For the comparison of the macroporous lime from natural lime (in the present work) and the macroporous lime from limestone (in the previous work), macroporous lime from limestone was used. The details of the preparation method have been previously reported.12 Characterization of the Sample. The thermal decomposition characteristics of the swelled samples were examined using a TPD apparatus equipped with a quadrupole mass spectrometer. The pore volume and pore size distribution of the calcined samples were measured using a mercury penetration porosimeter. Direct observation of the pores of the calcined samples was carried out using a field emission scanning electron microscope (FE-SEM). The reactivities of the macroporous CaO with SO2 were measured using a flow-type packed-bed tubular reactor system under atmospheric pressure at a constant temperature of 800 °C. The microreactor consisted of a quartz tube, 1.5 cm i.d., in which 0.2 g of the sorbent was packed. In these sulfation experiments, a mixture of SO2 (1500 ppm), O2 (3%), CO2 (10%), and H2O (9.2%), with the remainder being He, was fed into the reactor at 500 cm3/min at STP. SO2 concentrations of the inlet and outlet gases were measured using a wet absorption method {Arusenazo III method (JIS K103)}.13 The reactivity of a sample with H2S and oxidation of the CaS formed was examined using a flow type thermogravimetric apparatus equipped with a qualz tubular reactor (1.5 cm inside diameter).14 About 25 mg of the reactant particles was placed in a platinum wire net sample holder (1.3 cm diameter; net opening, ca. 0.3 mm). The sulfidation experiments were done at 800

Figure 2. Comparison of the swelling of lime and limestone at 40 °C.

or 900 °C under atmospheric pressure and a gas flow of 500 cm3/min at STP (linear velocity of the gas: 4.7 cm/s at STP). The inlet gas composition during sulfidation was 1500 ppm H2S, 9.2% H2O, and the remainder N2. In the oxidation experiments of the CaS produced from the CaO sample, the inlet gas was 1% O2, 9.2% H2O and the rest N2. The amount of SO2 evolved from the sample during the oxidation was measured using the same wet absorption method as in the sulfation of CaO.

Results and Discussion Swelling Rate of Limestone. The swelling rates of the two samples of lime were examined after exposure to the vapors of acetic acid aqueous solution at different concentrations at 40 °C. As shown in Figure 1, the swelling rate of the lime (Okayama) was increased at higher concentrations of acetic acid in the solutions: the maximum value of the rate was obtained at ca. 100% CH3COOH for both samples (the results of lime (Chubu) are not shown here). In another case of the swelling of limestone, the limestone did not expand under a vapor of ca. 100% CH3COOH.12 Furthermore, the dependency on the acid concentration of the swelling of the limestone differed from that of the lime, as shown in Figure 2. Therefore, it was concluded that the swelling mechanism of the lime was not the same as that of limestone. The dependency of the swelling rate of the limestone (Okayama) on the temperature was examined using 80 vol % CH3COOH aqueous solution. The swelling rate increased with rising temperature, as shown in Figure 3.

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Figure 3. Effect of temperature on the swelling of the lime using 80% CH3COOH(aq).

Figure 4. Effect of the precalcination temperature of the original limestone on the swelling of the lime formed using 80 vol % CH3COOH(aq).

The effect of the precalcining temperature on the swelling rate was examined using 80 vol % CH3COOH aqueous solution, but no effect was observed, as shown in Figure 4. The data reported hereafter in this study were obtained using lime prepared by calcination at 850 °C. Characterization of the Swelled Limestone by TPD. It is well-known that calcium acetate Ca(CH3COO)2 is produced from the reaction of Ca(OH)2 and CH3COOH, and it can be further supposed that Ca(CH3COO)2 is formed on exposure of CaO to a vapor of CH3COOH and H2O. Figure 5 shows the TPD spectra of reagent Ca(CH3COO)2‚H2O, the limestone (Okayama) CaCO3, a partially swelled sample of limestone, and a partially swelled sample of lime. From the compound Ca(CH3COO)2, CH3COCH3 and CO2 were evolved as a result of the following reactions:

Figure 5. Temperature-programmed decomposition of the limestone (Okayama), the partially swelled sample, and reagent Ca(CH3COO)2‚H2O. The partially swelled samples were prepared using 80 vol % CH3COOH(aq) at 40 °C and their fractional swelling (f) was 2.2.

ature as that for Ca(CH3COO)2. From the swelled sample, CH3COOH, which seemed to be adsorbed on the sample, was desorbed, as shown in Figure 6. Very similar TPD spectra (not shown) were obtained from the swelled sample (Chube). From the results in Figures 1, 3, 5 and 6, the swelling of lime by CH3COOH can be explained by the following reaction:

CaO(s) + 2CH3COOH(g) f Ca(CH3COO)2 + H2O

Ca(CH3COO)2 f CH3COCH3 + CaCO3

(2)

(4)

CaCO3 f CaO + CO2

(3)

Drying and Calcination of the Swelled Limestone. The bulk density of the 0.7 mm calcined samples was measured. From the bulk density of the swelled sample, the ratio of bulk density (the bulk density of the nonswelled sample/the bulk density of swelled sample) was calculated in order to ascertain the degree of expansion of the sample. The ratio of bulk density of the samples prepared using the 80 vol % and ca. 100 vol % CH3COOH(aq) are shown in Figure 7 in comparison with their fractional swelling. In Figure 7, a dotted line shows the theoretical relationship between the bulk density and fractional swelling assuming that the

The high-temperature position of CO2 evolved from Ca(CH3COO)2 was lower than that of the limestone. The cause of this shift is unknown, but it may be supposed that the CaCO3 formed from the decomposition of Ca(CH3COO)2 was more unstable than that of the limestone because of the difference of the two sample histories. From the partially swelled limestone, CH3COOH was evolved at low temperature in addition to CH3COCH3 and CO2. From the partially swelled lime, CH3COCH3 and CO2 were evolved at the same temper-

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Figure 8. Pore size distribution of the swelled-calcined sample (Okayama). The samples were swelled using 80 vol % CH3COOH(aq) at 40 °C and calcined at 850 °C except for one data point (*: swelled using 50 vol % CH3COOH). ( ), fractional swelling of sample; [ ], ratio of bulk density ()nonswelled-calcined sample/ swelled-calcined sample).

Figure 6. Temperature-programmed decomposition of differently swelled lime (Okayama). The swelled samples were prepared using 80 vol % CH3COOH(aq) at 40 °C. Fractional swelling, f.

Figure 7. Relationship between the fractional swelling of the sample and the ratio of bulk density of the sample after the calcination. The dotted line is the theoretical value. The sample lime was obtained from the limestone (Okayama) by precalcination at 700-850 °C. The swelled-calcined samples were prepared using 80 or 100 vol % CH3COOH(aq) at 40 °C, except for one data point (750 °C*: swelled at 30 °C), and calcined at 850 °C.

swelled sample was unaltered by the drying-calcination. The positions of the data of the samples lie on the right side of the lines except for one instance: this means that the samples expanded during the dryingcalcination. This result was inconsistent with the results of the swelling of limestone: the samples

exposed to 80 vol % CH3COOH(aq) were not greatly changed by the drying-calcination.12 The causes of the difference in these two results were unknown. The clarification of the causes needs much more study. Pore Size Distribution of Lime Produced from the Swelled Sample. The pore size distributions were examined using calcined samples prepared using 80 vol % and ca. 100 vol % CH3COOH(aq). Figure 8 shows the differential pore volume of the samples (Okayama): the pore size distributions of the nonswelledcalcined sample observed almost agree with those previously reported;10,15 the size of pore formed by the swelling method is distributed in three classes, from 10 to 100 nm, from ca. 0.2 µm to ca. 4 µm, and the largest class ranging over ca. 6 µm. The pore size distributions of the samples prepared from limestone (Okayama) using the 80 vol % and 50% CH3COOH(aq) are also shown in Figure 8:12 the pore size distributions of the samples prepared from the limestone were similar to those of the samples prepared from the lime. The peak position of the smallest pore (ranging from 10 to 100 nm) of the samples prepared from the lime was in the same position as the original lime, but those of the partially swelled samples prepared from limestone were in the same position as those of the swelled samples. In the case of the samples prepared from lime, calcium acetate formation is believed to occur on the surface of the intraparticle pore; therefore, the pore of the original lime remains after the partial swelling and calcination; as the size of the smallest pore of the partially swelled samples prepared from limestone was the same as that of the swelled samples, it appears that the pore forms from Ca(CH3COOH)2 formed from lime and/or limestone. The pores formed by the swelling method were directly measured by SEM using the sample (Okayama). The size of the macropores in the SEM photographs (not shown) were consistent with the pore sizes measured by the mercury penetration porosimeter. Reactivity of the Macroporous Lime Prepared. The reactivity test reactions of the macroporous lime

Ind. Eng. Chem. Res., Vol. 37, No. 10, 1998 3947

Figure 9. Reactivity of the samples prepared from lime (Okayama and Chubu) with SO2. The samples were the same those in Figure 6. The values of the fractional conversion of CaO to CaSO4 were measured 2 h after the start of the reaction.

Figure 10. Reactivity of the samples prepared from limestone (Okayama and Chubu) with SO2. The swelled samples were prepared using 50-80% CH3COOH(aq) at 18-40 °C. The calcination temperature of the samples was 850 °C. The values of the fractional conversion of CaO were measured 2 h after the start of the reaction.

prepared are as follows:

CaO + SO2 + 1/2O2 f CaSO4

(5)

CaO + H2S f CaS + H2O

(6)

The samples shown in Figure 7 were used for the reactivity test for SO2 removal. Furthermore, samples prepared from the limestones were used for a comparison with the samples prepared from lime. The reactivities of the swelled samples prepared from lime and limestone for the removal of SO2 are shown in Figures 9 and 10, respectively, in comparison with those of the nonswelled sample: the fractional conversion of the sample CaO to CaSO4 was increased with the increase in the swelling ratio; especially, the effect of the swelling on the reactivity of the sample seemed to be large at low ratios of bulk density in the case of both of the samples prepared from the lime and the limestone. The difference of the dependency of the fractional conversion on the ratio of bulk density between that of the two samples was small: in the low ratio bulk density range, the dependency of the sample prepared from lime seemed to be slightly larger than that of the sample prepared from limestone. In previous reports, the improvement of the reactivity of lime for SO2 capture by the swelling method was thought to be induced by the formation of macropores shown in Figure 8. Naruse et al. reported that pores

Figure 11. Weight gain of the swelled-calcined sample and the nonswelled-calcined sample during the sulfidation followed by oxidation. The swelled-calcined sample (Okayama) was prepared using 80 vol % CH3COOH(aq) at 40 °C and calcined at 850 °C. The ratio of bulk density of the swelled-calcined sample was 7.3.

with an average diameter between 1 and 1.6 µm are responsible for the high SO2 capture capacity of calcined shells.15 Harman and Coughlin also reported that pores with diameters larger than 0.796 µm were responsible for the high capacity of calcined limestone.8 This pore size range curiously coincided with that of the second class macropore produced by the swelling in this study. The reactivity test for H2S removal and the oxidative character test of formed CaS were performed serially at 800 or 900 °C using three samples. Figure 11 shows the results of the sample [prepared from the lime (Okayama) using 80% CH3COOH(aq); ratio of bulk density 7.6] at 800 °C, in comparison with that of the nonswelled sample. The reaction rate of the swelled sample was considerably improved by the swelling; and the final conversion, which was measured during the 2 h reaction, also increased from 0.81 to 0.92. The oxidation character was drastically changed: the weight gain of the swelled sample by the oxidation was ca. 2 times larger than that of the nonswelled sample. In the oxidation of CaS, the following reactions occur.16

CaS + 2O2 f CaSO4

(7)

CaS + 3/2O2 f CaO + SO2

(8)

A considerable amount of SO2 evolution was measured in the present experiment. The sigmoid curve of the second reaction in Figure 11 can be explained if the weight loss by the reaction of eq 8 affects the weight gain of the reaction of eq 7. The comparison of the reactivity of the macroporous limes from the lime and that from the limestone (Okayama) were examined using two samples: the ratio of bulk density of the sample from lime and limestones was 7.6 and 6.5, respectively. As shown in Figure 12, the characteristics of the sulfidation followed by oxidation of the two samples were almost identical. From this result and the results in Figures 9 and 10, it may be concluded that the effect of the initial material (lime or limestone) on the samples is relatively minor. The CaO sorbent for removal of H2S is sometimes used at a temperature higher than 800 °C. Figure 13 shows the character of the sample at 900 °C in comparison with that at 800 °C. The reaction time of the experiment at 900 °C was set at 3 h to obtain almost the same conversion as was measured in the case of the experiment at 800 °C. The characteristics of the oxida-

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Figure 12. Comparison of the weight gain of the swelled-calcined sample from lime and that from limestone during the sulfidation followed by oxidation. The swelled-calcined sample (Okayama) was prepared using 80 vol % CH3COOH(aq) at 40 °C and calcined at 850 °C. The ratio of bulk density of the sample from lime was 7.3, and that of the sample from limestone was 6.3.

Figure 14. Composition of the sample after the oxidation. The samples were sulfided and then oxidized as shown in Figures 1113. The composition was calculated on the basis of the CaS formed from the calcined samples (Okayama).

diffusibility of the SO2 formed to the outside of the sample particle. In the practical oxidation in a fluidized bed reactor, this SO2 may react with the CaO formed, and finally convert it to CaSO4. If a packed bed reactor is used, the following reaction eq 9 may occur:6 Figure 13. Effect of temperature on the weight gain of the swelled-calcined sample from lime during the sulfidation followed by oxidation. The swelled-calcined sample (Okayama) was prepared using 80 vol % CH3COOH(aq) at 40 °C and calcined at 850 °C.

tion of CaS formed also drastically changed with the reaction temperature: at 800 °C, the sample weight sigmoidally increased by the oxidation, but at 900 °C, the weight of the sample decreased just after the injection of the oxidative mixed gas and then increased. From this weight loss just after the oxidation, it was supposed that the reaction of eq 8 predominantly occurred compared with the reaction of eq 7 at that time. From the weight gain and the amount of evolved SO2, the composition of the oxidized sample could be calculated. Figure 14 shows the composition of the samples after the oxidation. In this figure, the composition was calculated on the basis of the CaS formed from the calcined sample during the sulfidation: the unreacted part (CaO) of the sample was not counted in the calculation of the composition. The oxidation of CaS was accelerated by swelling: the ratio of the remaining CaS was drastically decreased by the swelling. It could be supposed that this difference in the samples was caused by the different character of the pore structure: the CaS formed from the macroporous CaO was easily converted to CaSO4 and CaO, because the CaSO4 formed could not plug the large pores and O2 easily diffused into the inside of the sample particle through the pores. The formation of CaO was greatly affected by the temperature: the conversion of CaS to CaO was about 35% at 800 °C and about 60% at 900 °C. The CaO formation seems to be controlled by the ratio of the rate of reaction eq 8 to the rate of the reaction eq 7 and the

CaS + 2SO2 f CaSO4 + S2

(9)

This reaction seemed to compete with the reaction of eq 8. The effect of water vapor is unknown. If the reaction of eq 9 could be depressed, the CaO formed via eq 8 could be reused for desulfurization: it may be possible to develop a regenerable CaO sorbent for hightemperature desulfurization. Conclusion In the present study, swelling was found when lime was exposed to acetic acid vapor. The swelling of the sample resulted from an increase of calcium acetate formation in the sample. The swelling rate of the sample reached a maximum when the sample exposed to vapor of pure acetic acid and is depressed in the presence of water vapor. The swelled sample was converted to macroporous calcium oxides by heating to 850 °C. The reactivity of the macroporous calcium oxide for the removal of SO2 or H2S in the presence of H2O vapor was greatly improved in comparison with that of the calcined raw limestone. In particular, the SO2 removal capacity and oxidative characteristics of CaS to CaSO4 and CaO were drastically improved by this swelling method. Characteristics of the macroporous lime produced from lime were also compared with those of a sample prepared from limestone by the swelling method, which has been presented by us. In-bed decomposition of the calcium acetate is a major concern for the energy efficient preparation of CaO; therefore a study of the decomposition in flue gas and/ or coal-derived fuel gas is necessary for the practical

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application of calcium acetate. Furthermore, the existence of a possibility of preparing regenerable CaO sorbent was suggested by this study. Therefore, the study of the regeneration of CaS to CaO in a packed bed is a very important future problem. Acknowledgment The authors gratefully acknowledge financial support from Steel Industry Foundation for Advancement of Environmental Protection Technology. Literature Cited (1) Anderson, D. C.; Anderson, P.; Galwey, A. K. Surface Textural Change during Reaction of CaCO3 Crystals with SO2 and O2(Air). Fuel 1995, 74, 1024. (2) Fenouil, L. A.; Lynn, S. Study of Calcium-Based Sorbent for High-Temperature H2S Removal. 2. Kinetics of H2S Sorption by Calcined Limestone. Ind. Eng. Chem. Res. 1995, 34, 2334. (3) Yrjas, K. P.; Cornelis, A. P.; Hupa, M. M. Hydrogen Sulfide Capture by Limestone and Dolomite at Elevated Pressure. 1. Sorbent Performance. Ind. Eng. Chem. Res. 1996, 35, 176. (4) Khinast, J.; Krammer, G. F.; Brunner. Ch.; Staudinger, G. Decomposition of Limestone: The Influence of CO2 and Particle Size on The Reaction Rate. Chem. Eng. Sci. 1996, 51, 623. (5) Ninomiya, Y.; Sato, A.; Watkinson, A. P. Oxidation of Calcium Sulfide in Fluidized Bed Combustion/Regeneration Conditions. Int. Conf. Fluidized Bed Combustion, 13th 1995, 1027. (6) Anderson, D. C.; Galwey, A. K. A Kinetic and Mechanistic Investigation of the Formation of Calcium Sulphate in Reactions that may be of use in Flue-Gas Desulphurization. Proc. R. Soc. London A 1996, 452, 585. (7) Zevenhoven, C. A. P.; Yrjas, K. P.; Hupa, M. M. Hydrogen Sulfide Capture by Limestone and Dolomite at Elevated Pressure. 2. Sorbent Particle Conversion Modeling. Ind. Eng. Chem. Res. 1996, 35, 943-949.

(8) Hartman, M.; Coughlin, R. W. Reaction of sulfur Dioxide with Limestone and Influence of Pore Structure. Ind. Eng. Chem., Process Des. Dev. 1974, 13, 248-253. (9) Zevenhoven, R.; Yrjas, P.; Hupa, M. How Does Sorbent Particle Structure Influence Sulfur Capture under PFBC Conditions? Int. Conf. Fluidized Bed Combustion, 13th 1995, 13811392. (10) Efthimiadis, E. A.; Sotrichos, S. V. Sulfidation of LimestoneDerived Calcines. Ind. Eng. Chem. Res. 1992, 31, 2311. (11) Torres-Ordonez, R. J.; Wall, F. T.; Longwell, P. J.; Sarofim, A. F. Sulfur Retention as CaS(s) during Coal Combustion: a Modelling Study to Define Mechanisms and Possible Technologies. Fuel 1993, 72, 633. (12) Sasaoka, E.; Uddin. M. A.; Nojima, S. Novel Preparation Method of Mcroporous Lime from Limestone for High-Temperature Desulfurization. Ind. Eng. Chem. Res. 1997, 36, 3639-3646. (13) Sasaoka, E.; Tanaka, K.; Inami, Y.; Sakata, Y.; Kasaoka, S. Development of Catalyst for Simultaneous Oxidative Adsorption of SO2 and NO. Kagaku Kogaku Ronbunshu 1994, 20, 880-888. (14) Sasaoka, E.; Iwamoto, Y.; Hirano, S.; Uddin, M. A.; Sakata, Y. Soot Formation over Zinc Ferrite High-Temperature Desulfurization Sorbent. Energy Fuels 1995, 9, 344-353. (15) Naruse, I.; Nishimura, K.; Otake, K. Characteristics of Desulfurization Reaction by Shells. Kagaku Kogaku Ronbunshu 1995, 21, 904-908. (16) Miura, K.; Mae, K.; Inatomi, J. Study of Macroporous CaO Particles for High-Temperature H2S Removal. Preprint of the 61th Annual Meeting of the Society of Chemical Engineering Japan, Nagoya, 1996, M204, p167.

Received for review March 2, 1998 Revised manuscript received June 10, 1998 Accepted July 9, 1998 IE980137M