Ind. Eng. Chem. Res. 1988,27, 268-213
268
Direct Sulfation of Calcium Carbonate Michael J. H. Snow,+John P. Longwell,* and Adel F. Sarofim Department of Chemical Engineering, Massachusetts Institute of Technology, 77 Massachusetts Avenue, Cambridge, Massachusetts 02139
This work reports studies of sulfation of limestone under conditions such that the CaC03 does not decompose to CaO prior to sulfation. Direct sulfation of CaC03 is shown to proceed to high conversion, in contrast to the sulfation of a precalcined limestone (CaO). The sulfation rates are consistent with, and modeled by, a shrinking-core reaction with reaction rate control. The diffusion resistance through the product layer is significantly lower than that encountered in the usual sulfation of CaO. This is attributed to formation of a porous product layer caused by the COz generated by the reaction between CaC03 and SO2 and 02.This porous layer, for the CaC03 crystals studied, can physically expand to accommodate the larger molar volume of CaS04. Limited studies of highly porous CaC03prepared by calcium oxalate decomposition also showed rapid and complete conversion to CaS04. Limited studies of highly porous CaC03 prepared by calcium oxalate decomposition also showed rapid and complete conversion to CaS04 a t temperatures below the CaC03 decomposition temperature. The rate of direct sulfation of this process is consistent with its use in fluidized bed combustion to achieve high stone conversion. Though the concept of using limestone to capture acidic gases such as sulfur dioxide is old, interest in this process has been heightened by projected increases in coal usage and the associated increases in sulfur oxide emissions. An efficient method of minimizing these emissions into our environment is sought, especially from coal-fired power plants. When limestone is used as a sulfur absorbent in either fluidized bed or pulverized coal combustors,full conversion to calcium sulfate is not reached. In fluidized-bed combustion systems, less than 50% utilization of each particle is achieved and Ca/S ratios of more than 2 M must be used. In pulverized coal systems, utilizations are usually below 20-25%. In both of these systems, temperatures and concentrations are generally such that the reaction proceeds sequentially: CaCO,(s) CaO(s) + CO,(g) (1)
-
-
CaO(s) + SO&) + ‘/@Z(g) CaS04(g) (2) Several factors account for the low chemical conversions. The solid volume must expand significantly with sulfate formation since the molar volume of calcium sulfate is 34.5% larger than that of the carbonate and is 172% larger than that of the oxide. In addition, if particle shrinkage occurs during the decomposition, less volume is available for subsequent reaction. In large particles, internal concentration gradients cause calcium sulfate to form preferentially in pores near the surface, which then close or “plug”. Further reaction is very slow and involves solidstate diffusion through this outer shell (McClellan et al., 1970; Pigford and Sliger, 1973; Borgwardt and Bruce, 1984; Chrostowski and Georgakis, 1978; Bhatia and Perlmutter, 1981; Hartman et al., 1978). In smaller particles, the sulfation occurs more uniformly throughout the particle. In this case, the rate of sulfation is controlled at the internal oxide grain level, usually by solid-state diffusion through the growing product layer of calcium sulfate. When the latter mechanism prevails, the rate is strongly affected by the diameter of the oxide grains and therefore by the internal specific surface area. Without particle expansion, the final conversion is limited by the porosity of the oxide particles (Borgwardt and Harvey, 1972; DeLucia, 1985). Current address: Poloroid Corporation, 1265 Main Street, Waltham, MA 02254. 0888-5885/88/2627-0268$01.50/0
At temperatures below the decomposition temperature of CaCO,, the following “direct” reaction is believed to occur: CaCOAs) + SO&) + 1/202(g) CaS04(s)+ COz(g) (3) This reaction differs from (1) and (2) by avoidance of formation of CaO crystals and by generation of COz at the reaction interface. Because of these differences, one would expect differences in the relationships between sulfation rate and conversion. The literature on this reaction regime is, however, quite limited. Van Houte et al. (1981) explored the direct sulfation reaction at low temperatures at which the sulfite (CaSO,) is also a stable product. They observed that below 650 “C, reagent-grade CaCO, (4-15 Fm) must be impregnated with CaC1, to proceed to complete sulfation and concluded that the oxidation of the sulfite was rate-limiting. Unimpregnated CaC03sulfated at slower rates, initially first-order in [SOz + O,] (sic) concentration. They did not report findings above 650 “C, but their data show considerable acceleration of rate with temperature. Another work (Van Houte et al., 1978) at higher temperatures (600-900 “C) shows higher sulfations (>60% in 60 min at 900 “ C ) , though, again, not as high as their impregnated samples. No particle size variation was reported. Glasson and O’Neill (1980) reported conversions of CaC0, and Ca(OH), to CaS04by heating the particles in a TGA in atmospheres containing various amounts of SOz and 02.They concluded that the simultaneous formation of “quicklime” (CaO) causes the “greater reactivity” observed. Their hydroxide samples converted 100% to sulfate, though most of t h e reaction occurs after t h e decomposition. Their carbonate seemed to react during the decomposition, but to less than 50%. No mechanistic conclusions were drawn. Heating rates, particle diameters, and limestone source were not reported. Fee et al. (1982) carried out a similar sulfation in their study of a hydration process to enhance calcium utilization. For comparison purposes, they sulfated a raw limestone ( 0
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decompmition a t higher temperatures. The weight traces in Figure 2 have been interpreted to obtain extents of sulfation, assuming that reaction 3 proceeds prior to calcination (plus symbols in Figure 4) and that reactions 1 and 2 occur after its onset (triangles). During calcination (eq I), the weight data alone are insufficient to calculate the extent of sulfate formation. The weight gain Seen in Figure 2 below the calcination temperature is seen to correspond to nearly 20% sulfation. The sulfation continues after the decomposition of the remaining 80% of the CaC03, though the rate is seen to slow at this point. This is seen more clearly in Figure 5. Figure 5 shows sulfation conversion as a function of temperature for the 10-12-pm particles, for CO, concentrations ranging from 2% to 95%. CO, partial pressure has no effect on the sulfation rate until carbonate decomposition is attained. The rate of direct sulfation rises with temperature, and continues along the model curve until the carbonate decomposition temperature is reached. The rate of sulfation then drops noticeably. Thus, the major effect of increasing the CO, partial pressure is to delay the decomposition reaction until a higher temperature is attained so that the direct sulfation given by reaction 3 can proceed to a greater extent. The advantage of avoiding decomposition is clearly illustrated. Figure 6 presents scanning electron microscope photographs of sulfated particles. The sample which was sulfated after carbonate decomposition appears to consist of tightly packed sulfate crystals on the particle surface. The sample prepared by direct sulfation, in contrast, appears quite porous. This porosity is probably caused by the flow of CO, produced by reaction 3 through the product sulfate layer and is believed to account for the improved access
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1 Figure 6. (a) SEM photomicrograph of Iceland Spar limestone, decomposed at 970 "C in nitrogen and then s u l f a t d at the same temperature for 80 min with 3w0 ppm SO2. in 5% O2 and 95% nitrogen (1.5 cm = 1.0 p m ) . (b) SEM photomicrograph of Iceland Spar limestone, sulfated directly by heating at 20 "C/min with 3000 ppm SO2 in 5 % O2in nitrogen (1.5 cm = 1.0 pm).
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of SO, and 0,to the reacting CaCO, surface. In the above experiments, the carbonate decomposition temperature was modified by changing the CO, partial pressure. An alternative means to change decomposition temperature is use of different salts of calcium. Figure 7 shows decomposition and sulfation conversion results obtained when reagent grade (1-5-i" Ca(OH), crystals were heated in helium or in the presence of SO, and O2 Very little direct sulfation of the hydroxide occurred before it
Ind. Eng. Chem. Res., Vol. 27, No. 2, 1988 271 1.0
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decomposed because of the low temperature at which Ca(OH), decomposes. The CaO formed at this low temperature was apparently very reactive since it sulfated to about 90%. In contrast, CaC03of the same particle size sulfated substantially before decomposition and sulfated to completion by the time the temperature had reached 950 "C. Figure 8 presents data on decomposition and sulfation of calcium oxalate. Decomposition occurs as follows (Tanaka et al., 1981) under an inert atmosphere: CaC204.H20(s) CaC204(s)+ H20(g) at 125 "C (5) CaC204(s) CaCO,(s) + CO(g) at 475 "C (6) CaC03(s) CaO(s) + C02(g) at 650 "C (7)
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On the loss of the water of hydration, no sulfation of the oxalate occurs. However, when the oxalate decomposes to porous carbonate and CO, this carbonate sulfates essentially to full conversion as it is formed. The molar volume of calcium oxalate is 57.2 cm3/mol,larger than that of calcium sulfate, 46.0 cm3/mol, so particle expansion is not a problem, and the resulting sulfate could even be porous. The measured weight at temperatures greater than 650 OC correspond to complete formation of CaS04. Figure 9 shows that sulfation occurs predominately in the 400-500 "C temperature range and before decomposition of CaCO3 This sulfation is more rapid than for 1-5-pm CaC03 crystals and illustrates the importance of the high porosity resulting from calcium oxalate decomposition in addition to the sulfation prior to decomposition of the resulting carbonate. Discussion The sulfation results on Iceland Spar crystals show large effects of particle size and C02 concentration. These processes are modeled by considering each of the particles to react as a spherical shrinking core with reaction control at the core interface. The measured direct sulfation conversion-time relationships for the different conditions reflect the combination of a temperature-activated rate constant and the slowing of the reaction with conversion due to the shrinking of the available surface area of the core. For reaction control, the initial rate is inversely proportional to particle diameter, whereas, for productlayer control, the rate is inversely proportional to the square of particle diameter. The effect of particle size for the 2-3- and 10-12-pm particles was found to be well represented as the inverse of particle diameter. The larger particles showed even higher rates than predicted by this diameter effect. This is believed to be caused by development of cracks in the larger particles which result in an effectively smaller size. While some diffusion control is probable, these data were
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272 Ind. Eng. Chem. Res., Vol. 27, No. 2, 1988 K = 71 B x EXPI-15, 300/1 987XT1
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the integral can be solved using the second-exponential integral function, E2, which is defined as
decomposition temperature, is not surprising. The following expression for the direct sulfation reaction is found to hold for the range of particle diameters tested: k = 71.8 exp(-15300/1.987T] (cm/s) (22)
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By use of a more accurate approximation for the exponential integral involved, a better representation could be achieved analytically. However, the simpler equation given in (21) works very well, describing the key characteristics of the nonisothermal data of Figure 5 . To determine the experimental rate constant, the nonisothermal sulfation conversion data were differentiated numerically. k was determined as a function of the instantaneous temperature by using (9)-(11). It is assumed that p is that of calcite (2.71 X mol/cm3), R is the nominal particle radius (cm), and C is the experimental SO2 concentration (3.41 X mol/cm3). An Arrhenius plot of the rate constant is presented in Figure 10 for the six particle diameters tested in 95% COz. The three largest particle diameter samples had a higher rate factor (KO) than the others, though the activation energy was the same. For these, effective diameters smaller than the nominal diameters, where chosen in order to bring these data in line with the results from smaller particles. The reduction in diameter is presumed to be caused by internal cracking. The diameters used for the K O calculation are in parentheses. The mean activation energy of 15.3 kcal/mol has a standard deviation of approximately 1.0 for the six different particle sizes in 95% C02. Since this activation energy is much smaller than that found for calcite decomposition (49 kcal/mol), the relatively rapid rate of decomposition observed, once the carbonate was above its
The nonisothermal model (eq 21), with k from expression 22, is compared to two sets of data in Figure 11. The deviation of the data from the theoretical curves below 700 "C is largely due to the difference between the experimental heating rate in this temperature range of 55 "C/ min and the model assumption of a rate of 20 "C/min over the entire temperature range. While agreement above 700 "C is reasonably good, the model curve rises at a slightly higher rate than the experimental results. This could be due to an increase in diffusional resistance as conversion increases. The rates of the sulfation found indicate that 80-9070 sulfation can be achieved for 10-12-pm crystals in about 4 h, for an isothermal reaction in 1-atm partial pressure of C02 and 850 "C. This time, while too long for use of this process in a standard pulverized coal-fired furnace, is in the range of sorbent residence time for fluidized bed combustion. Direct sulfation, therefore, offers an approach to attainment of higher conversion to calcium sulfate than is currently attained by the direct limestone injection in fluidized-bed combustors where local conditions can cause calcination to precede sulfation.
Conclusions 1. Direct sulfation of 2-50-pm calcite crystals can be described as a shrinking-core process with relatively small diffusional resistance and an activation energy of 15.3 kcal/mol. 2. Formation of calcium oxide by raising the temperature above the calcium carbonate decomposition temperature results in a lower sulfation rate, with diffusion resistance increasing with conversion. 3. Photomicrographsshow that the calcium sulfate layer formed during direct sulfation is porous, while the sulfate layer formed from calcined stone shows no porosity. This is probably due to generation of COz at the surface during direct sulfation. 4. The time required for 80-90% conversion of 11-pm calcium carbonate crystals is predicted to be 4 h for 1atm of C02 pressure and 850 "C. This time is consistent with the use of direct sulfation to achieve high levels of stone conversion in fluidized bed coal combustion. 5 . The rate of CaC03 decomposition t o CaO is much faster than the rate of direct sulfation so that even brief exposure to calcination conditions will cause decomposition and deactivation with respect to sulfation.
I n d . Eng. Chem. Res. 1988,27, 213-219
Acknowledgment Exxon Research and Engineering is gratefully acknowledged for their support of this study. Registry No. SOz, 7446-09-5.
Literature Cited Bhatia, S. K.; Perlmutter, D. D. AIChE J. 1981,27,226-234. Borgwardt, R. H.; Bruce, K. R. USEPA (IERL), 1984. Borgwardt, R. H.; Harvey, R. D. Environ. Sci. Technol. 6, 1972, 350-360. Chrostowski, J. W.; Georgakis, C. ACS Symp. Ser. 1978,65,1. DeLucia, D. E. “The Cyclic Use of Limestone to Capture COz”. Master’s Thesis, Massachusetts Institute of Technology, Cambridge, 1985. Fee, D. C. Wilson, W. I.; Myles, K. M. “The Applicability of the ANL Hydration Process to Enhance to Calcium Utilization of Three Lowellville Limestone Sorbent Product Streams upon Being Recycled Back through the BABCOCK & WILCOX AFBC”. Argonne National Laboratory, May 1982.
273
Floess, J. K. “The Effect of Calcium on the Gasification Reactions of Carbon”. Doctoral Diss., Massachusetts Institute of Technology, Cambridge, 1985. Glasson, D. R.; O’Neill, P. Proc. 6th Int. Coni. Thermal Anal. 1980b, 1, 517-522. Hartman, M.: Pata, J.: Coughlin, R. W. Ind. Enp. Chem. Process Des. Dev. i978,i7,411-419. McClellan. G. H. Hunter. S. R.: Scheib. R. M. “X-Rav and Electron Microscope Studies of Calcined and Sulfated Limestones”. In The Reaction Parameters of Lime; Special Technical Publication No. 472; ASTM: New York, 1970; pp 32-66. Pigford, R. L.; Sliger, G. Ind. Eng. Chem. Process Des. Dev. 1973, 12, 85-91. Tanaka, H.; Ohshima, S.; Ichiba, S.; Negita, H. Thermochim. Acta 1981,48,137-146. Van Houte, G.; Maon, J. CL.; Dumont, P. H.; Delmon, B. J . Air Pollut. Control Assoc. 1978,28, 1030-1033. Van Houte, G.; Rodrigue, L.; Genet, M.; Delmon, B. Enuiron. Sci. Technol. 1981,15, 327-332. Received for review January 21, 1987 Revised manuscript received September 9, 1981 Accepted September 22, 1987
An in Situ, Multibeam, Spectrophotometric, Transient Analysis Method for Multiple Species in Chemical Reactors Yung C. Lin, Robert J. Adler,* and Robert V. Edwards Chemical Engineering Department, Case Institute of Technology, Case Western Reserve University, Cleveland, Ohio 44106
A spectrophotometric method is reported for measuring simultaneously the transient concentrations of several species at several locations in a chemical reactor. Locally averaged composition is measured noninvasively along each of several light beams. The method employs multiple light beams, optical multiplexing, fiber optics, several filter frequencies, time-shared detection, and computerized fast data acquisition. In a version constructed for testing, dual filter frequencies are employed to follow the concentrations of two absorbing species simultaneously along each of eight light beams. Concentration transients as short as 0.2 s are handled. An IBM personal computer stores 600 readings/s (2 wavelengths, 8 light beams, and 30 samples/s for each wavelength along each light beam, plus 120 auxilliary readings). Tests confirm that the method is robust, noninvasive, and inexpensive. Minor modifications can extend the number of species monitored, and trade-offs are possible between the number of positions monitored and speed: We recently desired to monitor the transient composition of a reacting mixture with several stoichiometric degrees of freedom, where the composition transients due to mixing and reaction were on the order of a few tenths of a second. The reaction was to take place in a volume of a few liters, and it was also desired to observe simultaneously spatial composition variations due to incomplete mixing. Spectrophotometry is attractive because it is noninvasive and in principle can satisfy all of our needs, but a survey of commercially available spectrophotometers revealed no instrument capable of handling several species simultaneously with the desired speed of response. In addition, non could probe-or could be readily modified to probe-several locations simultaneously. Conventional spectrophotometers have the capability of analyzing for multiple components, but since they mechanically scan a range of wavelengths, their speed of response is typically 1-5 min. Spectrophotometers of this type are not suitable for reactive mixtures. Conventional stopped-flowspectrophotometers have fast transient capability down to the millisecond range, but they probe only one wavelength along one beam path and 0888-588518812627-0213$01.50/0
are therefore unsuitable for multicomponent mixtures and multiposition analysis. The type of spectrophotometer which came closest to meeting our needs is based on array diodes. An example is the Hewlett-Packard Model 8450 spectrophotometer. I t uses 400 diodes in parallel to measure the spectrum virtually instantaneously (0.1 s/spectrum). This philosophy can in principle handle multicomponent mixtures with rapidly changing composition. However, data processing and storage requirements limit the sampling rate to one analysis per a few seconds. Further, modification for multiposition measurements is impractical. We concluded that we had to construct a special type of spectrophotometer to do composition analysis of a few-component mixture with transient times on the order of a few tenths of a second where multiple positions are to be observed simultaneously. We describe here a spectrophotometric method developed in our laboratories to meet the above needs. It fills the gap between the conventional stopped-flow type (very fast speed of response-but limited to a single component and position) and the diode array type (many components-but limited to about 1-s response speed and
0 1988 American Chemical Society
