on the Consumption of Desulfurizing Agent Ca(OH) - American

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Ind. Eng. Chem. Res. 2010, 49, 1444–1449

Study on the Effect of CO2 on the Consumption of Desulfurizing Agent Ca(OH)2 in Flue Gas Desulfurization Wenlong Wang,* Mingqiang Hu, Yong Dong, and Chunyuan Ma National Engineering Laboratory for Coal Combustion Pollutant Control, School of Energy and Power Engineering, Shandong UniVersity, Jinan, China, 250061

In view of the ambiguity concerning the effect of CO2 on the consumption of desulfurizing agent and based on our previous findings when studying flue gas desulfurization (FGD) byproduct, a specific study on the influence of CO2 has been carried out here with Ca(OH)2 being employed as a desulfurizing agent. This study is based on compositional analyses of the reaction products, and combined measurements have provided relatively precise compositions of the reaction products derived from Ca(OH)2 and flue gases. It has been found that the presence of CO2 does have an effect on the desulfurization reaction and on the consumption of desulfurizing agent when Ca(OH)2 is employed as the sorbent for SO2. Also, the formation of CaCO3 is inevitable at about 70 °C and is enhanced under conditions of high humidity. Although over 90% SO2 removal efficiency could be achieved with a relatively high Ca/S mole ratio, the effective utilization of desulfurizing agent would be low because CaCO3 is not the targeted product. Therefore, in order to improve semidry FGD technology, measures must be taken to prevent the reaction between CO2 and Ca(OH)2 so as to increase the utilization ratio of the desulfurizing agent. 1. Introduction Flue gas desulfurization (FGD) technologies are the chief means of SO2 removal in coal-burning power plants. According to differences in their SO2 sorbents and desulfurization byproducts, FGD technologies can be classified into wet, dry, and semidry types. In all types, calcium-based sorbents are the most widely used. In particular, in dry or semidry FGD technologies, Ca(OH)2 often serves as the desulfurizing agent. Usually, the effect of the CO2 in flue gas is seldom considered when Ca(OH)2 reacts with SO2. This neglect is based on the knowledge of better solubility of SO2 than that of CO2, stronger acidity of sulfurous acid compared to that of carbonic acid, and a higher equilibrium constant for SO2 than that for CO2 in their reactions with Ca(OH)2 (1032.8 and 1010.1, respectively).1 In fact, in many previous experiments, the simulated flue gas employed did not contain CO2 at all.2-4 Nevertheless, in most actual flue gases, the concentration of CO2 is tens or even hundreds of times higher than that of SO2. Can its influence really be neglected? When our research group focused on study of the utilization of FGD byproduct, a very interesting problem was encountered.5 That is, quite a large amount of CaCO3 was found in the byproduct of semidry FGD systems. Five samples of FGD byproduct, which were taken from five different power plants, were analyzed. Table 1 gives their mineral compositions. It can be seen that the contents of CaSO3 and CaCO3 were almost equal, which meant that a big portion of desulfurizing agent had not been utilized effectively. Through analysis, we excluded the possibility of low quality of the agent materials. Therefore, this finding can only be attributed to the desulfurization reaction process. With this in mind, we have focused on the competition between the reactions of Ca(OH)2 with SO2 and CO2. In fact, some researchers have already studied the effect of CO2 in desulfurization reactions to some extent. For instance, Rochelle found that, in the range 10-40%, CO2 did not seem * To whom correspondence should be addressed. Tel.: 86-53188399372-603. Fax: 86-531-88395877. E-mail: [email protected].

to have any influence on the sulfur removal reaction when 2000 ppm SO2 was present in the gas.6 Hartman found that at temperatures between 170 and 580 °C only a small proportion of CaO reacted with CO2 to form CaCO3 and that this had little influence on the SO2 removal reaction.7 Klingspor et al. investigated the influences of CO2 and O2 on the reaction rates for limestone and lime in a sand-bed reactor. They reported that CO2 and O2 had negligible influence on the rates of the reactions between SO2 and limestone or lime.8,9 Similarly, the experiments of Irabien et al. showed that CO2 (12%) and O2 (2%) had no influence on the removal of SO2 from flue gas.10 Stevens found that an appreciable fraction of the unsulfated sorbent from the initial Ca(OH)2 was present in the waste solids as CaCO3 when CO2 was present in the flue gas, and considered that CaCO3 was unreactive under the conditions prevailing in FGD processes at low temperatures.11 Ortiz found some CaCO3 at the beginning of the reaction, but not at higher conversion when 10 000 ppm SO2 and 12% CO2 were present in the reaction gas. When only 100 ppm SO2 was present, they detected some CaCO3 at all conversion levels.12 To sum up, these conclusions would seem to disfavor a strong influence of CO2. On the other hand, some researchers (refs 13 and 14 and Garea et al., 2000) have pointed out that a deeper understanding is needed when CO2 is included in the composition of the flue gas.13-15 It was reported by Gmelin that at a temperature below 100 °C the reaction of CaO with CO2 was enhanced when water vapor was present.16 Ho et al. investigated the effects of CO2 (12.6%) and O2 (5.4%) in flue gas on the reaction of Ca(OH)2 with SO2 using a fixed-bed differential reactor. Their experimental results showed that the apparent sulfation rate, the total reaction rate, and the final total conversion of Ca(OH)2 were Table 1. Mineral Compositions of the Samples of Desulfurization Residues (mass %) sample

CaSO3

CaCO3

CaO

balance

1 2 3 4 5

35.75 36.90 35.51 11.78 9.86

27.77 29.89 31.45 15.57 8.36

3.27 5.34 11.99 1.17 2.37

33.22 27.87 21.05 71.49 79.41

10.1021/ie901401j  2010 American Chemical Society Published on Web 12/15/2009

Ind. Eng. Chem. Res., Vol. 49, No. 3, 2010

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Table 2. Experimental Conditions (T ) 70 °C)

Figure 1. Experimental setup for the competitive reactions between SO2 and CO2: 1, high purity nitrogen; 2, synthetic flue gas; 3, rotameter; 4, humidifier; 5, U-shaped reactor; 6, water bath; 7, dropper; 8, air-washing bottle.

greater than those in the absence of CO2.17 Krammer et al. investigated the reaction of Ca(OH)2 with SO2 at low temperature (353 K) by means of a thermogravimetric analyzer. Their experimental results showed that the conversion of Ca(OH)2 was enhanced when CO2 (8.9%) was added to a gas mixture of SO2, O2 (7.1%), H2O, and N2 due to the reactions of Ca(OH)2 with SO2 and CO2.2 They surmised that the formation of CaCO3 was certainly influenced by SO2, but it remained unclear as to whether the presence of CO2 also had some impact on the reaction of Ca(OH)2 with SO2. In summary, it can be stated that, although the influence of CO2 on SO2 removal has been taken into consideration, the experimental results have been inconsistent and do not entirely account for the effects of CO2 when Ca(OH)2 is used as a sorbent for SO2 in semidry FGD technologies. Especially the effect of CO2 on the consumption of desulfurizing agent is unclear. In view of the difficulty of explaining the existence of so much CaCO3 in our findings and the ambiguity about the effect of CO2 in previous literature, a specific study on the competition mechanism between SO2 and CO2 has been carried out using Ca(OH)2 as the desulfurizing agent, the results of which are reported here. 2. Experimental Section 2.1. Experimental Setup. Figure 1 shows a sketch of the experimental setup. The experiments were carried out in a glass U-shaped reactor of 2.5 cm inlet diameter. A packed sand bed was prepared to contain a dispersal of Ca(OH)2, and the bed was supported by a glass wool pipe section in the reactor. The U-shaped reactor was placed in a water bath, which could provide a constant water temperature in the range 20-100 °C. A reaction temperature of 70 °C was chosen in order to mimic the commonly used semidry desulfurization conditions. In order to ensure the exact reaction temperature, the sand bed was placed at the outlet end so that the flue gas could be preheated sufficiently. Different flue gases were prepared ahead of the reactor, and the tail gas was absorbed by NaOH solution in an air-washing bottle. 2.2. Atmosphere. Three kinds of atmosphere were designed for contrast in our experiments. One synthetic flue gas was composed of 2000 ppm SO2, 14% CO2, and a balance of N2. Its volume proportion of CO2 to SO2 was 70:1, which is close to that of the actual flue gas from coal-burning boilers. The other two atmospheres were pure CO2 and SO2/N2, which corresponded to the cases when CO2 or SO2, respectively, took effect independently. The synthetic flue gases did not contain O2 because it does not show a remarkable effect on the consumption rate of Ca(OH)2 in the reaction with CO2 or SO2, although the presence of O2 can promote the oxidation of CaSO3 to CaSO4.18 The gases were passed through the sand bed at a rate of 1 L/min.

no.

SO2 (ppm) (dry)

CO2 (%) (dry)

N2

initial sample mass (mg)

humidity

1 2 3 4 5 6

2000 2000 0 0 2000 2000

0 0 14 14 14 14

balance balance balance balance balance balance

500 500 500 500 500 500

60% RH wetted 60% RH wetted 60% RH wetted

Two humidity levels were applied, 60% relative humidity (RH) and wetted, respectively. In the case of 60% RH, the gases were passed through the upper space of a water evaporating bottle, which was placed in another water bath. By selecting the appropriate water temperature, the required relative humidity could be obtained. In the case of wetted conditions, a water dropper was employed to keep the sand bed wet during the reaction process. Although only two cases were designed, they could, basically, reflect the actual state in the desulfurization process: the former stood for the typical reaction conditions in the desulfurization process, while the latter meant some sorbent particles might meet more water drops and become entirely wet. The series of experimental conditions are listed in Table 2. 2.3. Experimental Methods. Analytically pure Ca(OH)2 was used as the experimental agent. Table 3 describes its properties and chemical composition. The Ca(OH)2 agent was dispersed in a bed of inert silica sand (particle size range 450-900 µm) in order to avoid channeling and to improve gas/solid contact. In each of the experiments, Ca(OH)2 (100 mg) was mixed with silica sand (5 g) and the ratio was kept at 1:50. The flue gas from the cylinder was passed through the sand bed after preheating. The aeration was carried out for over 2 h to ensure sufficient reaction between the Ca(OH)2 and the flue gas. Then, before the end of the experiment, the atmosphere was switched to N2 until the sand bed was dried. In this way, the possible reaction between wet Ca(OH)2 left in the sample and CO2 in air could be avoided. The sand bed was subsequently removed and the product was separated from the silica. Finally, the collected product samples were immediately preserved in vacuum bags. According to the different reaction conditions, six product samples were obtained in our experiments, and our study was based on the compositions of these samples. Three analysis methods were jointly applied to determine the chemical compositions. X-ray diffraction (XRD) analyses were employed to ascertain the minerals that were formed in the reactions, thermogravimetric analyses (TGA) were employed to determine the content of unreacted Ca(OH)2, and sulfur/carbon analyses by far-infrared heating were carried out to distinguish the contents of CaSO3 · 1/2H2O and CaCO3. 3. Results and Discussion 3.1. XRD Analyses. In the experiments, the main possible reactions are as follows: SO2 + Ca(OH)2(liquid) f CaSO3 · 1/2H2O + 1/2H2O

(1) CO2 + Ca(OH)2(liquid) f CaCO3 + H2O

(2)

CaCO3 + SO2 + 1/2H2O f CaSO3 · 1/2H2O + CO2

(3) Therefore, the samples of the final residues should contain CaSO3 · 1/2H2O, CaCO3, and unreacted Ca(OH)2. Of course, if

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Table 3. Properties and Chemical Composition of Ca(OH)2 Used properties

chemical composition (wt %)

BET surf. area (m2/g)

mean pore diam (nm)

true density (kg/m3)

average diam (µm)

Ca(OH)2

CaCO3

SiO2

Fe2O3

Al2O3

MgO

15.3

1950

2215

31

95.83

0.27

0.51

0.09

0.08

1.96

O2 is contained in the reaction gases, CaSO4, which usually amounts to a small portion in the semidry FGD byproducts, might also be detected. However, this kind of confusion was excluded in our experimental design. To confirm the mineral compositions, XRD analyses were carried out and Figures 2-7 show the diffractograms. It is noticeable that the peaks of Ca(OH)2 predominate in all of the diffractograms, indicating that the conversions of Ca(OH)2 were low in all of the experiments. Nevertheless, the experimental method adopted corresponds to the typical and optimal way when employing a fixed bed, and other literature has also reflected a low conversion. In Figure 2, for which theoretically only reaction 1 applies, it can been seen that the peaks of CaSO3 are very inconspicuous. In Figure 3, CaSO3 becomes noticeable, indicating its greater formation in the wetted case. In Figures 4 and 5, the diffraction peaks of both Ca(OH)2 and CaCO3 can be seen, and CaCO3 is much more noticeable than CaSO3 in Figures 2 and 3. Moreover, in Figure 5, when the sorbent was fully saturated, the CaCO3 peaks become stronger, showing that the carbonation reaction was also enhanced in this case. In

Figures 6 and 7, for which the atmosphere was SO2/CO2/N2, the peaks of Ca(OH)2 and CaCO3 are conspicuous, while those of CaSO3 can only just be detected. Nevertheless, it can be concluded that the intensities of the CaSO3 and CaCO3 peaks tend to be stronger in the saturated case. 3.2. TGA and Sulfur/Carbon Analyses. Of particular interest to us was the content of each mineral in the product samples, especially the contents of CaSO3 · 1/2H2O and CaCO3, which reflect the competitive reactions of CO2 and SO2 with Ca(OH)2. TGA could determine the content of unreacted Ca(OH)2 but could not distinguish CaSO3 and CaCO3 because their decompositions occur in almost the same temperature regions. This problem was solved by the joint employment of sulfur/carbon analysis by far-infrared heating, since this method can precisely determine the sulfur and carbon contents through heating the samples up to 1600 °C. The TGA curves are shown in Figures 8-10, all of which were obtained in an N2 atmosphere. Clearly, each weight loss process occurs in two dominant stages. The first stage, at about 400 °C, can be mainly attributed to the decomposition of

Figure 5. X-ray diffractogram of product sample 4. Figure 2. X-ray diffractogram of product sample 1.

Figure 3. X-ray diffractogram of product sample 2.





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Table 4. Results of Sulfur/Carbon Analyses sample

S (mass %)

C (mass %)

S/C molar ratio

1 2 3 4 5 6

4.43 9.47 0.06 0.07 3.82 8.89

0.61 0.66 2.70 6.77 2.20 3.43

2.71:1 5.38:1 0.01:1 0.01:1 0.65:1 0.97:1

Table 5. Mineral Contents of Samples and Conversion sample

CaSO3 · 1/2H2O (mass %)

CaCO3 (mass %)

Ca(OH)2 (mass %)

conversion (%)

1 2 3 4 5 6

17.86 38.18 0.24 0.28 15.40 35.84

5.08 5.50 22.50 56.42 18.33 28.58

67.3 50.14 62.65 39.89 59.53 34.11

17.23 34.12 21.13 51.24 27.34 55.01

Figure 8. Results of thermogravimetric analysis (nos. 1 and 2).


result may be attributed to the mass-transfer resistance between the flue gas and unreacted Ca(OH)2 agent caused by the product layers on the sorbent surface. However, in the wetted case, there were clear increases in Ca(OH)2 conversion in all three atmospheres tested. Thus, both the carbonation and sulfurization reactions are enhanced when the Ca(OH)2 agent is wetted. These results are in accordance with those of the XRD analyses. Table 4 gives the results of the sulfur/carbon analysis by farinfrared heating. Therein, the direct data are the mass contents of the elements S and C, on the basis of which the mass percentages of CaSO3 · 1/2H2O and CaCO3, respectively, could be evaluated. Table 5 gives the mineral contents of each sample, and the conversion was calculated according to the following formula: conversion ) mCaSO3·1/2H2O /MCaSO3·1/2H2O + mCaCO3 /MCaCO3 mCaSO3·1/2H2O /MCaSO3·1/2H2O + mCaCO3 /MCaCO3 +


Figure 11. Mass fractions of Ca(OH)2 in samples.

Ca(OH)2, while the second stage, in the range 500-800 °C, can be attributed to the decomposition of CaSO3 and/or CaCO3. According to the weight loss ratios, we can roughly estimate the mass percentages of Ca(OH)2 in each sample, which are shown in Figure 11. Based on Figure 11, it can be seen that the reaction was quite insufficient, with a high proportion (about 35-70%) of Ca(OH)2 not participating. Especially in the case of 60% RH, samples 1, 3, and 5 each contain nearly 60% of unreacted Ca(OH)2. This

× 100%

mCa(OH)2 /MCa(OH)2 where mx is the mass percentage of the corresponding mineral and Mx is its molar mass. In fact, the numerator is the number of moles of Ca(OH)2 that participate in sulfurization and carbonation reactions, and the denominator is the total number of moles of Ca(OH)2. In Table 5, when SO2 was solely employed as reaction gas, samples 1 and 2 should theoretically contain a maximum of 0.27% CaCO3. The measured amount of ∼5% must be due to the reaction with CO2 when Ca(OH)2 before the experiment has been exposed to air. The presence of sulfite in samples 3 and 4 when only CO2 was employed as reaction gas can only be attributed to impurities in the Ca(OH)2 agent. Nevertheless, these amounts of unwanted carbonate or sulfite are not too high, and should have little influence on our conclusion. 3.3. Discussion. Combining the results of the three analyses, we can first conclude that both SO2 and CO2 react with Ca(OH)2 in the designed conditions, and, in view of its high concentration, CO2 probably has a fairly high conversion. Consequently, the consumption of this desulfurizing agent might be expanded because of the unwanted carbonation reaction. Second, the level of humidity has a pronounced effect on the reactions between Ca(OH)2 and SO2 or CO2. Figures 12 and 13 reflect this effect more clearly. In the wetted case, there are great increases in the CaSO3 · 1/2H2O mass percentage in SO2 or SO2/CO2 atmosphere, and in the CaCO3 mass percentage in CO2 or SO2/CO2 atmosphere. This suggests that the carbon-

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CaSO3 · 1/2H2O and CaCO3 in samples 5 and 6 are also in accordance with the data in Table 1, which resulted from our study concerning the competition. The nonnegligible consumption of Ca(OH)2 caused by CO2 would surely lower the utilization ratio of the desulfurizing agent. However, a new reaction mechanism is needed to explain the observations and further research must be carried out. 4. Conclusions

Figure 12. Mass percentages of CaSO3 · 1/2H2O in samples 1, 2, 5, and 6.

In contrast to other research work, this study has been based on compositional analyses of reaction products. Combined measurements have provided relatively precise compositions of the reaction products derived from Ca(OH)2 and flue gases. It can be concluded that the presence of CO2 does have an effect on the desulfurization reaction when Ca(OH)2 is employed as a sorbent for SO2. The formation of CaCO3 is inevitable at about 70 °C and is enhanced under conditions of higher humidity. This conclusion is in accordance with our findings from a previous study, that quite a large amount of CaCO3 is formed in the desulfurization residues of semidry FGD technologies. Although over 90% SO2 removal efficiency could be achieved with a relatively high Ca/S mole ratio, the effective utilization of desulfurizing agent would be low because CaCO3 is not the targeted product. Therefore, in order to improve semidry FGD technology, measures must be taken to prevent the reaction between CO2 and Ca(OH)2 so as to increase the utilization ratio of the desulfurizing agent. Acknowledgment

Figure 13. Mass percentages of CaCO3 in samples 3-6.

ation and sulfurization reactions are greatly enhanced when the Ca(OH)2 sorbent is wetted, which would be in accordance with the intuitive assumption that the rate of ionic reactions is usually faster than that of gas-solid reactions. In semidry FGD technologies, Ca(OH)2 slurry is sometimes sprayed into the desulfurization reactor, or, more generally, Ca(OH)2 powder and water are sprayed in individually. In both cases, there is contact between wet Ca(OH)2 and SO2/CO2, but in the latter case this process only occurs locally and its duration may be shorter. Finally, our experiments have proved that the carbonation reaction does compete with the sulfurization reaction when Ca(OH)2 is used as a sorbent for flue gas containing both CO2 and SO2. In Table 5, the conversion of Ca(OH)2 is seen to increase from 21.13% (sample 3) to 51.24% (sample 4) when solely CO2 is applied, while it only changes from 17.23% (sample 1) to 34.12% (sample 2) when solely SO2 is applied. This indicates that wetting has a greater accelerating effect on the carbonation reaction. Actually, the S/C molar ratio in Table 4 reflects the competition directly. In samples 5 and 6, the relevant ratios are 0.65:1 and 0.97:1, respectively, which means that the number of moles of Ca(OH)2 that react with SO2 is lower than the number of moles that react with CO2. Nevertheless, under conditions of increased humidity, the sulfurization reaction is clearly enhanced, since the S/C molar ratio increases from 0.65 to 0.97. Therefore, in view of the above finding that wetting can enhance the carbonation reaction more markedly, coupled with the much lower concentration of SO2 compared to CO2, we can affirm that the reaction between Ca(OH)2 and SO2 is more competitive than that between Ca(OH)2 and CO2. Even so, most probably because of the great disparity in concentration between CO2 and SO2, 70:1, similar amounts of Ca(OH)2 react with both gases. The similar percentages of

The authors thank the support of National High-Tech Research and Development Program of China (863 Program) (Grant 2009AA05Z303), the National Natural Science Foundation of China (Grant 50906046), and the National Natural Science Foundation of Shandong Province (Grant Q2008F10). Literature Cited (1) Wu, Z. B. Studies on wet and spraying dry flue gas desulfurization. Dissertation, Zhejiang University, Hangzhou, 1993. (2) Krammer, G.; Brunner, Ch.; Khinast, J.; Staudinger, G. Reaction of Ca(OH) 2 with SO2 at low temperature. Ind. Eng. Chem. Res. 1997, 36, 1410–1418. (3) Cortabitarte, F.; Ortiz, M. I.; Irabien, A. Flue gas desulphurization at low temperatures: thermogravimetric characterization of the reaction product. Thermochim. Acta 1992, 207, 255–264. (4) Rice, R. W.; Bond, G. A. Flue gas desulfurization by in-duct dry scrubbing using calcium hydroxide. AIChE J. 1990, 36, 473. (5) Wang, W. L.; Xu, X. R.; Ren, L.; Dong, Y.; Ma C. Y. Study on utilization ratio of SO2 sorbents in semidry flue gas desulfurization technology. Presented at the Asia-Pacific Power and Energy Engineering Conference (APPEEC), Wuhan, China, 2009. (6) Rochelle, G.; White, W.; Jozewicz, W.; Chang, J. Reaction of hydrated lime with SO2 in humidified flue gas. In Proceedings 1990 SO2 Control Symposium, New Orleans; Sedman, C., Radcliffe, P., Eds.; U.S. Environmental Protection Agency, Air and Energy Engineering Research Laboratory: Research Triangle Park, NC, 1991; Vol. 4, Paper 7A-119. (7) Hartman, M.; Trnka, O. Reaction between calcium and flue gas containing sulfur dioxide at lower temperatures. AIChE J. 1993, 39, 615– 624. (8) Klingspor, J.; Karlsson, H. T.; Bjerle, I. A kinetic study of the dry SO2-limestone reaction at low temperature. Chem. Eng. Commun. 1983, 22, 81–103. (9) Klingspor, J.; Stro¨mberg, A.; Karlsson, H. T.; Bjerle, I. Similarities between lime and limestone in wet-dry scrubbing. Chem. Eng. Process. 1984, 18, 239–247. (10) Irabien, A.; Cortabitarte, F.; Ortiz, M. I. Kinetics of flue gas desulfurization at low temperature: nonideal surface adsorption model. Chem. Eng. Sci. 1992, 47, 1533–1543.

Ind. Eng. Chem. Res., Vol. 49, No. 3, 2010 (11) Stevens, N. J. Dry SO2 scrubbing pilot test results. Proceedings of the Symposium on Flue Gas Desulfurization; American Chemical Society: Washington, DC, 1981; PB81-243164. (12) Ortiz, M. I.; Garea, A.; Irabien, A.; Cortabitarte, F. Flue gas desulphurization at low temperatures. Characterization of the Structural Changes in the Solid Sorbent. Powder Technol. 1993, 75, 167–172. (13) Diffenbach, R. A.; Hilterman, M. J.; Frommell, E. A.; Boohev, H. B.; Hedges, S. W. Characterization of calcium oxide-fly-ash sorbents for SO2 removal. Thermochim. Acta 1991, 189, 1–24. (14) Jorgensen, C.; Chang, J. C. S.; Brna, T. G. Evaluation of sorbents and additives for dry SO2 removal. EnViron. Prog. 1987, 6, 26–32. (15) Garea, A.; Herrera, J. L.; Renedo, M. J.; Fernandez, J.; Irabien, A. Thermogravimetric determination of the influence of water vapour in the FGD in-duct injection at low temperatures. J. Chem. Technol. Biotechnol. 2000, 75, 484–490.

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(16) Gmelin, L. Handbuch der anorganischen Chemie; Verlag Chemie: Heidelberg, Germany, 1961; Band Ca, Teil B. (17) Ho, C. S.; Shih, S. M.; Lee, C. D. Influence of CO2 and O2 on the reaction of Ca(OH)2 under spray-drying flue gas desulfurization conditions. Ind. Eng. Chem. Res. 1996, 35, 3915–3919. (18) Liu, C.-F.; Shin, S.-M. Effects of Flue Gas Components on the Reaction of Ca(OH) 2 with SO2. Ind. Eng. Chem. Res. 2006, 45, 8765– 8769.

ReceiVed for reView September 8, 2009 ReVised manuscript receiVed November 24, 2009 Accepted November 29, 2009 IE901401J

on the Consumption of Desulfurizing Agent Ca(OH) - American

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