2324
Ind. Eng. Chem. Res. 1995,34, 2324-2333
Study of Calcium-BasedSorbents for High-Temperature H2S Removal. 1. Kinetics of H2S Sorption by Uncalcined Limestone Lament A. Fenouilt and Scott Lynn* Department of Chemical Engineering and Lawrence Berkeley Laboratory, University of California, Berkeley, California 94720
In about 160 experiments, +6 mesh (1-2 mm diameter) limestone samples were exposed to a n atmosphere of COS (87-88%), H2 (1-2%), CO (4-5%), H2O (4-5%), and H2S (0.5-1.85%) for various temperatures (570-860 "C) and durations (5-240 min). Limestone sulfidation was followed quantitatively as well as qualitatively to elucidate the reaction mechanism. The morphology of the inside and of the outside of reacted limestone samples was observed with a scanning electron microscope, and X-ray maps of the sulfur content of the reacted samples were generated. The reaction between H2S and limestone is first-order with respect to the H2S partial pressure in the 560-660 "C range, where the kinetics is controlled by the chemical reaction with a n activation energy of about 39 kcaymol. Between 660 and 710 "C the reaction rate decreases as the temperature increases and the apparent order of the reaction changes from first- to half-order. Finally, in the 710-860 "C range, the reaction becomes controlled by solidstate diffusion through the C a s product layer with an activation energy of 30-40 kcaymol.
Introduction and Previous Work One of the most important needs in coal gasification is the development of reliable methods of removing solid and gaseous contaminants from coal gas at high temperatures and pressures. Gaseous contaminants include H2S, NH3, and volatilized alkali salts, all formed from trace components of the coal. Suspended solids are mainly fines formed during gasification. It is necessary that all of these be removed prior to combustion of the coal gas to prevent damage t o turbine equipment and infringement of emissions legislation. Squires et al. (1971)showed that cleaning the gas at high temperature and pressure offers numerous advantages, most notably higher thermal efficiency and reduced volume (hence cost) of processing vessels. Current industrial practice is to add limestone to the gasifier. This achieves moderately successful removal of H2S, but requires a significant excess over the stoichiometric quantity of limestone theoretically necessary to attain close-to-equilibrium removal of H2S and produces large amounts of mixed-limestonehlag-solid waste. Sulfur is present in two main forms in coal: pyritic and organic. Pyritic sulfur is contained as mineral inclusions inside the organic matrix of coal. Thus, fine grinding of the raw coal followed by separation of the organic from the mineral phase by flotation removes a large fraction of the pyritic sulfur prior to burning or gasifying (Lynch, 1981;Brown, 1962). However, half or more of the sulfur, along with some nitrogen and oxygen, is organically bound to the coal (Nowacki, 1981). Consequently, it cannot be extracted before the coal is burned. In a gasification process the residual sulfur in both forms ends up as H2S because of the reducing conditions present in the gasifier. A process under investigation in this laboratory involves removal of H2S from coal gas by sorption into a moving bed of limestone (or lime) particles. This paper describes the kinetics of the reaction of H2S with limestone and the evolution of the structure of the solid + Current address: Chemical Development Department, Shell Chemical, 3333 Highway 6 South, Houston, TX 77082.
during the course of the reaction. A differential quartztube reactor system was used to study these reactions in the temperature range of interest at atmospheric pressure. As previously reported, the reaction between large particles of limestone and H2S, under noncalcining conditions, is slow and limited (Fenouil et al., 1994). This paper identifies the causes of the poor reactivity. Many studies of the reaction between calcined limestone (CaO) and half- or fully-calcined dolomites (respectively CaCO3-MgO and CaO-MgO) with H2S have been published (Abbasian et al., 1990;Borgwardt et al., 1984; Freund, 1981; Yen, 1979; Ruth et al., 1972; Squires et al., 1971). These works were often concerned with the addition of limestone or dolomite to the coal in the gasifier where the conditions are such that the calcination of the MgC03, and frequently the CaC03, might occur rapidly (Freund, 1981). This paper treats the direct sulfidation of the noncalcined CaC03
+
CaC03(s) H2S(g)
CaS(s)
+ H,O(g) + CO,(g)
(1)
because it would permit coal gas cleanup below the calcination temperature of the limestone. Limestone is a sedimentary rock of quite variable composition, mainly consisting of calcium carbonate with some calcium sulfate and magnesium carbonate along with other impurities. Chemical compositions of various limestones from different origins can be found in Chan et al. (1970),Chang et al. (19841,Borgwardt and Roache (1984),Borgwardt et a2. (19871,and Fuller and Yoos (1987). The specific surface area of most of these stones is low (generally less than 1 m2/g)and the natural porosity ranges from 0 to 8% (Borgwardt and Roache, 1984;Hartman et al., 1978;Borgwardt et al., 1987). Scanning electron microscopy pictures of a few rock samples taken during this study reveal that they may contain small transparent crystal inclusions as big as a few microns in diameter. Borgwardt and Roache (1984)noted that the external dimensions of the limestone particles remained the same during calcination or sulfidation. Thus, in the course of reaction 1 the porosity of the particle should increase by 0.25because of the molar volume difference between CaC03 (2.71g/cm3)and CaS (2.61g/cm3). This
0 1995 American Chemical Society 0888-5885/95/2634-2324$09.QQ~Q
Ind. Eng. Chem. Res., Vol. 34, No. 7, 1995 2326 situation is fundamentally different from that of the sulfidation of dense calcium oxide, CaO (2.32 g/cm3), which occurs with reduction of the pore volume because of the larger volume of S2- relative to 02-.Thus, the reaction should proceed via solid-state diffision after a crust of nonporous CaS is formed (Borgwardt et al., 1984). However, contrary to the conclusions of this analysis, Borgwardt and Roache (1984) found that sulfidation of CaC03 almost stops after 11%conversion for large limestone pellets (D, '15 pm). They also studied the sulfidation of uncalcined limestone with particles ranging from 1.6 to 100 ,um at temperatures between 570 and 850 "C under an atmosphere of COZ (70%),Nz (29.5%), and HzS (0.5%). They assumed that the loss of porosity of the surface limestone at these high temperatures prevents the gaseous reactant from diffusing further toward the center of the solid particles. They found that the sulfidation kinetics of the uncalcined limestone with particle sizes ranging from 1.6 to 10 pm is well described by d[CaC031 dt
kr
=,
= -[CaC031[HzSl
where t is the time (in min), D, is the diameter of the particle (in cm), [CaC031 is the unreacted fraction of CaC03, [HzSI is the gas-phase concentration of HzS (in g-mol&) and kr is a constant (that varies with the temperature) being equal to 0.66 (cm*L)/(molof HzSnin) at 750 "C. The dependency on D, indicates that the reaction is chemically controlled. Ruth et al. (quoted in Borgwardt and Roache (1984)) found a similar expression, without the D, dependency, for the sulfidation kinetics of 60-pm diameter, half-calcined dolomites. Squires et al. (1971) also found that the reaction was first-order with respect to the HzS partial pressure. Borgwardt reported that COS and HzO enhanced the sulfidation reaction whereas HZ slowed the reaction rate. However, Ruth et al. (1972) found no effect of HZ on the rate of sulfidation. No completely satisfactory explanation for such poor reactivity has been suggested. Borgwardt and Roache (1984)proposed that limestone sintering caused the poor conversion. They observed that 1.6-pm-diameter limestone particles having an initial surface area of 4.5 m2/g sinter t o give 3.5-pm particles with an average specific surface area of 2.0 m2/g after 20 min a t 850 "C under 1 atm of COz. No CaO formed during their experiments, so this loss of surface area could only be attributed to a physical change of the stones. However, similar experiments showed no significant surface area loss with millimeter-size limestone pellets (Fenouil et al., 1994), eliminating CaC03 sintering as the cause of the poor reactivity of HzS with millimeter-size particles of limestone. Attar and Dupuis (1979)observed that the sulfidation of calcite (trigonal crystals of calcium carbonate) was controlled by the chemical reaction on the flat calcite crystal surface until about 80 CaS layers had formed. Then the newly-formed CaS layer limits the solid-state gas diffusion. These results are consistent with the Borgwardt and Roache (1984) observations of the steep decline of the sulfidation rate on large pellets after 11% conversion: 80 molecular layers corresponds to about 10%of the volume of a 1-pm-diameter CaC03 grain. The CaS layer obtained at relatively low temperature is not thermodynamically stable because the sulfide ions have just replaced the carbonate ions in their previous sites
without any structural rearrangement. If the temperature increases, the Ca2+ and S2- ions on the surface diffise fast enough to form a more stable CaS crystalline structure. Thus, the formerly flat CaS crust lets some cracks appear and exposes more fresh CaC03 to HzS. Differential scanning calorimetry data from Attar and Dupuis (1979) reveal that the rate of ionic diffision becomes relatively fast a t 635 "C: the time scale of crystalline rearrangement becomes of the order of the other experimental time scales, such as the characteristic time for gas diffusion or for the chemical reaction. Everything that can trigger this recrystallization^' process (temperature, presence of oxygen to form some s04'- ions to break the metastable CaS crust, impurities, etc.) will be favorable to a higher conversion of CaC03 into Cas. These predictions are consistent with the observations made by Ruth et al. (1972) of a sudden increase in the reaction rate after sending pulses of 0 2 past partially sulfided, half-calcined dolomite samples. Another physical phenomenon must account for the dramatic drop in rate of limestone sulfidation after about 11%conversion for large limestone particles a t temperatures as high as 750 "C. Sintering experiments on CaS powder showed that CaS suffers a significant, rapid loss of surface area under conditions typically encountered in coal gasification (Fenouil et al., 1994). The poor reactivity of the limestone could thus be attributed to sintering of the CaS layer around the limestone grains so that, instead of cracking to allow more HzS t o reach the core of the grain, the CaS coats the grain with a nonporous, quasi-impermeable layer. Potential gas-phase reactions and gas-phase masstransfer limitations can have a major influence on reaction 1 (Fenouil et al., 1994). For instance, Sun et al. (1978) and Borgwardt and Roache (1984) studied reaction 1in a C O D 2 atmosphere in which CaS04, and not Cas, is the stable sulfur-containing compound. Thermodynamic calculations show that, in the 600-900 "C range, it is necessary to have about 1%of CO in the gas phase to prevent CaS oxidation to Cas04 by COZ (Fenouil, 1995). Attar and Dupuis' (1979)results on the kinetics of the sulfidation of calcite are only reliable below 700 "C; above this temperature their kinetic measurements become limited by the rate of gas-phase mass transport.
Experimental Section Reactor. All experiments were carried out with a differential tube reactor originally designed by Towler (1992). The reactor, sketched in Figure 1, consists essentially of two quartz tubes having outer diameters of 30 and 7 mm (4.2 mm internal diameter). The gas mixture enters the outer tube, flows through the furnace, flows over the solid sample located a t the opening of the inner tube, and finally exits from the cold end of the inner tube. An inlet tube located a t the hot end of the reactor can be used to sample the process gas or to purge the reactor with pure COZbefore and after every run. A gas-tight seal located in the cold-end stopper allows the inner tube to be inserted and withdrawn from the furnace to quench the samples rapidly without the need of moving the outer tube. Moreover, the cold-end stopper and the inner tube can be separated from the outer tube to remove and replace the sample while the furnace heater is still running. The entire apparatus is shown in Figure 2. This reactor configuration ensured the following characteristics:
2326 Ind. Eng. Chem. Res., Vol. 34, No. 7, 1995 Retractable Sample
Quark 'hermowell
Sampling/
ep Gar Port
Figure 1. Differential tube reactor.
m
n
Knockout
Tank
NaOH AbmrpUon
M
Scrubbers
1
Yeasurw" t
and Control
m
I I I
R.1l.l
Valve
'
U
bactor
r
+=r+ U U L I U
Point
C02 N 2 H2S
A2
Figure 2. Experimental apparatus for preparing solid samples.
(i) The temperature profile was flat (k2 "C) over a sufficient length t o give a constant temperature at the solid sample. (ii)A relatively large gas flow rate (5-6 mUs (STP)) suppressed any significant external heat- and masstransfer effects. (iii) The solid samples were heated or quenched in less than 2 min by sliding the inner tube in or out of the furnace. (iv) The residence time of the process gas was about 1 min in the furnace before it entered the inner tube containing the solid sample. This provides enough time for the water-gas-shift reaction to equilibrate. Through the shift reaction it was also possible to have a controlled amount of water vapor in the gas phase. Analyses. The scanning electron microscopy (SEM) pictures were obtained on an ISI-DS 130 dual stage scanning electron microscope from International Scientific Instruments, Inc. (Santa Clara, CAI. All the samples were coated with a 200-250 nm conducting gold or graphite layer, because neither CaS nor CaC03 has sufficient electron-conducting properties to permit good-quality pictures. The energy dispersive spectroscopy ( E D 8 equipment came from EDAX International, a division of North American Philips Corporation (Mahwah, NJ). Most analyses were coupled with SEM pictures from an ISIDS 130C 144-10 dual stage scanning electron microscope from International Scientific Instruments, Inc. (Santa Clara, CA). EDS spectrometers can qualitatively identify all chemical elements with an atomic number larger than 6 on the surface of a solid. The scanned surface area can be as low as a fraction of 1pm2 and penetration is about 2.5 pm, making this technique
1 co
Temperature Detector
Sample
Table 1. Chemical Analysis of the Great Lakes Calcium Limestone (wt %) CaC03
M&03
97.80
1.63
Si02 0.28
Fez03
A1203
S
Ca
Mg
0.15
0.13
0.065
39.15
0.47
particularly adapted for micron-size grain analysis. Quantitative results are possible for relatively heavy elements like calcium and sulfur but are, at best, semiquantitative for lighter components such as carbon or oxygen. Sorbents. Samples of industrial-quality limestone were provided by the Great Lakes Calcium Corporation (Green Bay, WI).The average chemical composition of this limestone is given in Table 1. However, some variation can be expected in the composition, since the samples are from natural quarried rock. The gases used in the reactor experiments (C02, CO, H2, N2, and H2S) were industrial grade (99.9% pure) and provided by Matheson Gas Products (East Rutherford, NJ). Experimental Procedures. (a) Sample Preparation. It was found necessary t o half-calcine the limestone samples to convert MgCO3 to MgO before conducting sulfidation experiments to obtain reliable conversion data. The MgCO3 content of the limestone used throughout this research is variable and relatively low (1-6 wt %). By half-calcining the samples a t 800 "C under C02 before they were weighed initially, all of the weight change during sulfidation could be ascribed to converting CaC03 to Cas. The halfcalcination of the limestone samples was shown to have no noticeable impact on the sulfidation rate (compared to samples that were not previously half-calcined). (b)Sulfidation Experiments. The pretreated limestone samples were weighed, introduced a t the end of
Ind. Eng. Chem. Res., Vol. 34,No. 7,1995 2327 the inner tube, and maintained in place between two pieces of quartz wool. The inner tube was kept outside the furnace until the furnace reached thermal equilibrium at the desired temperature under a flow of pure CO2. Then the inner tube was inserted into the furnace to allow the solid sample to reach thermal steady state under an inert COz atmosphere. Meanwhile, the flows of the process gases, controlled by a set of calibrated valves and rotameters, were set while bypassing the reactor. When the process gas flow had stabilized, and the reactor had reached thermal steady state, the process gas was sent to the reactor and the flow of pure COZwas stopped. The process gas needed 60-90 s t o reach the limestone in the reactor, and we estimated the exposure time of the sample to the simulated coal gas accordingly. The pressure inside the reactor was maintained between 15 and 25 in. of water gauge by a dip-tube relief seal. The samples were reacted for the desired time and then rapidly removed from the furnace by withdrawing the inner tube and cooled under the reaction-gas mixture or CO2, depending on the type of experiment. After the solid had cooled, the reaction gas was switched back to the scrubber and the reactor was purged with COZ;the sample was then removed and replaced by a new one. A piping and instrumentation diagram and detailed instructions on the sequence of valve switching required to operate the apparatus are described elsewhere (Towler, 1992). Long-duration experiments (4h or more) with high HzS concentration must be carefully monitored. Sulfur deposits inside the reactor pipes or the valves can cause plugging severe enough t o create a large pressure drop in the reactor and to result in the diversion of the process gas to the pressure relief seal instead of the reactor. Pressure buildup within the reactor is also often associated with partial loss of samples from the inner reactor tube to the outer reactor tube. The effect of the H2S decompositionwas taken into account in the determination of the exact proportion of this gas in the gas phase using the kinetic model for thermal HzS dissociation developed by Towler (1992). ( c ) Measurements of the Conversion of CaCOs to Cas. The Great Lakes limestone contains few impurities other than MgC03, which was converted to MgO before sulfidation experiments. Thus, the conversion t o Cas (starting from limestone) can be estimated by a gravimetric measurement assuming that CaC03 and Cas are the only calcium compounds present in the samples:
9%
initial weight - final weight = loo( 0.72 x initial weight
)
(3)
Several titrations of the sulfide ions (using a standard iodometric titration) were also performed to confirm the conversion data based on gravimetric measurements. Agreement between titrimetric and gravimetric measurements was always found t o be within 2% (Fenouil, 1995). (d) Precision of the Results. In all o u r experiments on limestone sulfidation, the conversion was found to be below 12%, which corresponds to a relative weight change of 5% or less in the samples. In most experiments, this corresponds to an absolute weight change of only 1-5 mg out of a total of about 300 mg. The precision of the weighings before and after sulfidation was only 0.1 mg. Each experiment was repeated three to six times to increase the precision of the conversion measurements. This procedure also dimin-
1
1 A
0
30
60
90
120
150
180
210
240
Time (min)
Figure 3. Limestone conversion as a function of time (about 1.8% HzS in the gas phase).
ished the impact of the variation in composition or shape of the stones in each sample. For clarity, only the averages are presented in this paper. Finally, it was determined that less than 0.1 mg of solid was lost in the form of entrained "dust" during the sulfidation experiments, thus ensuring that stone attrition had a negligible influence on the determination of the conversion.
Experimental Results The kinetics of limestone sulfidation was investigated over the entire range of temperatures for which reaction 1is thermodynamically favorable. With about 2% HzS, 5% HzO, and 88% COZ in the gas phase, the lowest temperature a t which the sulfidation takes place a t atmospheric pressure is roughly 560 "C, and the upper temperature limit is set by the incipient calcium carbonate calcination a t about 870 "C. The values of the conversion of limestone to Cas as a function of time for several temperatures are reproduced in Figure 3. For clarity, data at only four different temperatures are shown in Figure 3. The results for several other temperatures are available elsewhere (Fenouil, 1995). As discussed below, it appears that the mechanism for the limestone sulfidation changes fundamentally as the temperature rises. High-Temperature Sulfidation (710-860 "C). About 70 experiments on limestone sulfidation for temperatures ranging from 710 to 860 "C were performed to determine the reaction rate as a function of temperature and time for a given concentration of H2S (nominally 1.8%). The rest of the gas phase was C02 (87.5%),CO (4.6%),HzO (4.6%>,and Hz (1.5%)). The COZfraction in the gas feed was kept very high so that the variations of the reactor temperature would have minimal influence on the gas-phase composition over the limestone samples aRer thermodynamic equilibrium was reached (Fenouil, 1992). As mentioned above, each point in Figure 3 is the average of four to six experiments. The error bars show the standard deviation of the results. When no bar is visible, the standard deviation is covered by the symbol. It was found that,
2328 Ind. Eng. Chem. Res., Vol. 34,No. 7, 1995 2 0 ,
I
I
t
I
1
-1
\T
I model -1
5 08
~
1
I
1
1
09
10
1 1
12
2 13
1000/T
(T in
OK)
Figure 4. Initial reaction rate as a function of temperature (about 1.8%H2S in the gas phase).
up to the calcination temperature (about 870 "C), the conversion of CaC03 to CaS stays below 10% when the stones are exposed t o about 2% H2S for up to 4 h. Low-TemperatureSulfidation(560-660 "C).The same experiments were also conducted a t temperatures ranging from 560 t o 660 "C. The results for 560 and 660 "C are presented in Figure 3. Once again it was not possible to obtain more than 10%conversion to CaS in less than 4 h, confirming the results of Borgwardt and Roache (1984). It was also confirmed that the reaction rate increases as the temperature increases within this range. However, the reaction rate was observed to be significantly higher at 660 "C than at 710 "C. Intermediate-TemperatureSulfidation (660-710 "C). The initial reaction rate was investigated more fully after we observed this unusually high reaction rate at 660 "C. The initial reaction rate first increases as the temperature increases (up t o about 670 "C), and then drops sharply (between 670 "C and 720 "C) before increasing again (from 720 "C up to the calcination temperature of CaC03). Figure 4 displays the initial reaction rate as a function of the inverse of the absolute temperature in an Arrhenius plot, using conversion data obtained after 15 and 30 min. It appears that a significant phenomenon occurs around 700 "C that causes a sharp decrease in the reaction rate. Apparent Order of the Sulfidation Reaction with Respect to HzS. Figure 5 displays the limestone conversion after 30 min at 665 "C under several different partial pressures of H2S, the rest of the gas being mainly composed of C02 (about 88%), H20 (about 5%), CO (about 5%), and H2 (about 1%).Since in this time range the conversion is still proportional to the exposure time, the limestone conversion is proportional t o the initial reaction rate and we can conclude that the initial reaction rate is proportional to the H2S partial pressure. Hence, the apparent order of the reaction with respect to H2S is 1. Figure 5 also contains data that were obtained under the same conditions as those obtained at 665 "C,except for the temperature, which is now 870 "C. The limestone conversion becomes proportional to the square root of the partial pressure of HzS, so the apparent order of the reaction with respect to H2S is
O Y 00
' 0
05 Z ' of HS ,
10
1.5
20
in t h e gas phase
Figure 5. Influence of H2S partial pressure on limestone conversion at 665 "C and at 870 "C (1bar total pressure).
0.5. Each data point in Figure 5 represents the average value of five experiments; the vertical error bars correspond to the standard deviation of these five measurements. Morphological Study. SEM pictures of the external surface of the particles reveal the same "smoothness" previously observed in the pictures of sintered CaS powder (Fenouil, 1992). The original limestone-particle surface has lost most of its sharp angles. Moreover, many small pores have disappeared or have significantly shrunk, and small cracks, which were expected to appear on the grain surface because of the difference in molecular volume between Cas and CaC03, are not present. However, a few large fractures (with a width of the order of 1pm) exist on the stone surface. These fractures are not the result of a purely mechanical effect, such as shocks during the sample handling. The internal surface of the faults appears very smooth. These kinds of fractures have been observed in several different sulfided samples, whereas they were never seen to this extent in fresh limestone samples that were simply heated to the same temperature under an inert atmosphere (Fenouil, 1992). Sulfur Distribution. A few samples were polished to allow X-ray mapping of the sulfur in the reacted samples. The sections obtained with a razor blade were too rough to allow good mapping (the electron scattering due to the coarse surface was too large to get a clean signal). It was then crucial to obtain a surface as smooth and flat as possible for a better resolution of the analysis. The h a l polish was obtained with a METADI1-pm-diamondpaste (from Buehler) and provided a very flat cross section of the center of the particles. A close correlation between the color, or rather the tone of gray, and the chemical composition of the samples was demonstrated earlier (Fenouil, 1992). A white spot coincides with an area containing at least 1%sulfur atoms: a lighter tone of gray corresponds to a zone very rich in Cas, whereas the darkest gray is associated with the original calcium carbonate. These tones are very useful for determining the distribution of CaS among the CaC03 crystallites on the SEM
Ind. Eng. Chem. Res., Vol. 34, No. 7,1995 2329
Figure 6. SEM picture of the inside of a limestone sample partially sulfided at 610 "C.
Figure 8. SEM picture of the inside of a limestone sample partially sulfided at 810 "C.
"bumps" that begin to appear. Finally, Figure 8 presents another limestone sample treated under the same conditions as the two preceding ones, the temperature being 810 "C this time. At this temperature, the surface of the grains is completely covered by "bumps". However, we could not correlate these morphological changes with any variation in the reaction kinetics because of limitations in our experimental equipment.
Figure 7. SEM picture of the inside of a limestone sample partially sulfided at 660 "C.
pictures. All the SEM pictures we took of the center of partially sulfided limestone reacted at 610-860 "C under an atmosphere composed of CO2 (87.6%),H2 (l%), CO (4.8%),H20 (4.8%),and H2S (1.8%)during 30 min show that the sulfur atoms are relatively homogeneously distributed around all the pores, even in the core of the limestone particle (Fenouil, 1995). The sulfide ions surround the pores but are not present in the core of the individual CaC03 grains forming the particle. This proves that H2S has not been prevented from diffusing deep into the particle. Figures 6-8 are SEM pictures of the inside of partially sulfided limestone particles obtained with the same magnification (5000x ; reproduced a t 70% of the original size). The reacted limestone particles were sectioned with a razor blade to induce a fracture that carried through the pellet to give two parts of roughly the same size. Thus, the stone separated following the weakest points of its grain-pore network, revealing the aspect of the inside grains of the stones. Figure 6 was taken after the limestone had been exposed to the same conditions as those described for Figure 4. The grain surface is very smooth and does not display any "bumps". Figure 7 is a picture of a particle exposed to similar conditions, except for the temperature, 660 "C, which marks the point at which the reaction rate begins to decline with further temperature increase. The aspect of the grain surface is quite similar to that observed in Figure 6 with the exception of a few little
Interpretation of the Experimental Data At low temperatures (up to about 670 "C) the chemical reaction between CaC03 and H2S is the limiting step in the sulfidation kinetics (at least up to 5% conversion). This is confirmed by the three following facts: First, the apparent reaction rate is first-order with respect to H2S, ruling out any diffusion limitation in the reaction rate (see part 2 in this series (Fenouil and Lynn, 1995)). Second, conversion values for 30 min are almost exactly twice those for 15 min, which is consistent with the hypothesis of a chemically-controlled initial reaction rate. Finally, the apparent activation energy of 39 kcaV mol is very similar to previous values published by Squire et al. (1971)and Borgwardt et al. (1984)on activation energy of the sulfidation reaction in the absence of diffusion limitations. At higher temperatures, the sharp decrease in reaction rate indicates the appearance of a new limiting step in the limestone sulfidation kinetics. Figures 6-8 show that this reduction in reaction rate is associated with the rearrangement of the Cas product layer around the calcium carbonate grains. When the reaction temperature increases above 660 "C, the originally metastable flat Cas layer (visible in Figure 6)breaks to nucleate new, more stable crystals (observe the rugged surface of the grain in Figures 7 and 8). This phenomenon was observed on calcite crystals at 635-650 "C by Attar and Dupuis (1979). However, the surface rearrangement is not associated with an increase in the reaction rate, which one would expect with the exposure of new surface of unreacted CaC03. Instead, there is a significant decrease in the reaction rate probably caused by the sintering of the Cas crystals, which creates a new, denser Cas coating layer around the calcium carbonate grains. This sintered layer appears to prevent CO2 or H20 from diffusing out of the limestone grains (or H2S from diffusing into the grains) and to be responsible for the poor conversion of limestone to Cas at high temperatures.
2330 Ind. Eng. Chem. Res., Vol. 34, No. 7, 1995 In the 820-880 "C temperature range the reaction is half-order with respect to HzS, which is characteristic of kinetics limited by diffusion through a growing product layer. Moreover, the apparent behavior of the reaction rate indicates that the diffusion follows an activation process that is characteristic of solid-state diffusion processes. These two observations, combined with the fact that Cas sinters quickly at these temperatures (Fenouil et al., 1994) and that significant surface changes can be observed on the grain surface, lead to the conclusion that the reaction is now limited by the solid-state diffusion of the reacting gas (HzS)through the sintered product layer of Cas. We are possibly in the presence of ionic diffusion of S2-and C032- through the product layer, as advanced by Borgwardt for the mechanism of sulfidation of CaO (Borgwardt et al., 1984). Kinetic Model. Numerous mathematical models have been developed in recent years to describe the kinetics of gas-solid reactions. One of them, the grain model (Szekely et al., 1976), an extension of the wellknown shrinking core model (Levenspiel, 1972), accounts for the following potential rate-limiting mechanisms in the reaction kinetics: (i)external mass transfer (from bulk gas to the sorbent's surface); (ii) pore diffusion; (iii) diffusion through the individual grains composing the matrix of the sorbent particles; (iv)chemical kinetics. The exact mathematical development of the model depends on the geometry assumed for the sorbent particle as well as for its grains (namely, flat, cylindrical, or spherical in shape). However, in the present experiments the values of the conversion of CaC03 to Cas stayed below 5%within the first 30 min of the reaction. For these low values of the conversion, all the kinetic expressions can be approximated by the same equation regardless of geometric considerations. Furthermore, the mathematical analysis is greatly simplified when the chemical kinetics is first-order with respect t o the concentration (or partial pressure) of the reacting gas. Under these conditions, the relationship between time and conversion becomes
grain. For the other terms in eqs 5-8 refer to the Nomenclature. We operated all our experiments a t a relatively large Reynolds number (between 3 and 10) so that the external mass-transfer resistapce had no impact on the overall reaction kinetics. Using a value of 26.5 kmoV m3 for e, (typical for our limestone), 0.27 moVm3 for C mz/s (about 1.8%HzS in the gas phase), and 2 x for the molecular diffisivity of HzS in the gas mixture, the time for complete conversion for l-mm-diameter limestone particles would have been a mere 20 s (even assuming a Sherwood number of 2, characteristic of a single solid sphere in an infinite, quiescent gaseous medium), if external mass transfer had been the limiting resistance in the overall reaction kinetics. Furthermore, pore diffusion throughout the entire sorbent particle could also not have been rate-limiting, since the sulfur distribution is actually homogeneous over the entire particle. This could have been predicted since ZDP is only about lo3 s for our limestone (based on 1.8%HzS in the gas phase and an effective diffusivity, De, of about 5 x m2/s). These considerations indicate that ZMT and ZDP are negligible so that eq 4 becomes
(4)
which reduces to ( ~ ~ c t if~ZR) = - ~ 0 and , ~ to (zR)-~if ZDG = 0 (to being equal to 30 min in our case). Study of the Asymptotic Cases. On one hand, if Z R is ~ much larger than 4tZDG, the reaction rate is given by the inverse of ZR, which is a constant. So, if e,, rg, and C are known, an Arrhenius plot of the data where ZR dominates will yield ko and Ea&. On the other hand, ~ much smaller than 4tZDG, the instantaneous when Z R is reaction rate is given by (4tq~3)-"~.The reaction rate is now a function of time. One should note that for any kinetics described by eq 9 the chemical reaction is always the initial limiting step until the time to = Z& (4ZDG) is reached. For times larger than to the diffusion becomes the limiting step in the reaction kinetics. However, if ZDG is much larger than ZR, the switch from chemical to diffusion control takes place very early and one can consider that the reaction is basically always diffusion-controlled. In the temperature range where the reaction is diffision-controlled, the conversion is not a linear function of time, as in the case of chemical-controlled reaction, but is proportional to the square root of the time. The experimental data for the 15-min samples above 800 "C used to generate Figure 4 were normalized accordingly. Figure 4 is a plot of percent CaC03 sulfidation for the first 30 min of the reaction (Le. to =
with
ZMT
= (XC
CeqJ(E)
e, ZDp
= (18(C- Ceq))E)
(6)
(7)
where ZDP is the characteristic time for diffusion through the entire sorbent particle, TMT the characteristic time for external mass transfer from the bulk gas to the surface of the pellet (film diffusion), ZDG the characteristic time for diffusion through the grain, and ZR the characteristic time for chemical reaction at the interface between the unreacted core and the reacted layer of the
(9)
Equation 9 can be easily differentiated t o give the rate of reaction: reaction rate = W d t = (t;
+ 4tq)G)-1'2
(10)
This is the equation that should describe our reaction kinetics for the first 30 min of exposure to HzS. However, we did not really measure the instantaneous value of Wdt in our kinetics experiments, but its average value during a given period of time (here, 15 and 30 min). We can rearrange eq 10 to obtain
Ind. Eng. Chem. Res., Vol. 34, No. 7, 1995 2331
TOP VIEW 'Cas I
CrC03
cas
O.67h
g " I)
1
cac03
hx I
caco3
CaC03
RECRYSTALLIZATION
h(l-X)
SINTERING
(ABOVE 006 C)
Figure 9. Schematic description of the CaS product layer affer recrystallization and sintering.
30 min.). For the 15-min samples above 800 "C we multiplied the values of the conversion by the square root of 2 (and not by 2 as we did for the 15-min samples a t lower temperatures) to normalize them with the 30min samples. Determination of the Parameters of the Sulfidation Kinetics. The parameters for eq 11 were obtained as follows: (i)The inverse of the molar volume of the solid was taken as 26.5 kmol/", based on a density of 2.65 g/cm3 for a typical limestone. (ii) The HzS concentration, C, averaged 0.23 mol/m3 (f15%) in our experiments. The value of Cq was significantly smaller than C throughout our experiments and thus was neglected throughout this work. (iii)The value of r, is obtained by assuming that the limestone particles are a collection of agglomerated spherical grains. From the experimentally determined specific surface area of 0.25 m2/g, the average grain radius is 4.5 pm. (iv) In the 560-670 "C range, where ~ D Gappears negligible, the logarithm of the initial rate is proportional to the inverse of the absolute temperature: ln(reaction rate) = -ln(t,) =
(since It, is equal to KO exp(-EaCJRgT). Fitting eq 12 to the low-temperature data of Figure 4, Eact % 39 kcall mol and KO w 0.52 x lo4 d s . This value of the activation energy is very close to that observed by Borgwardt and Roache (1984), 42 kcallmol, when they studied the sulfidation of limestone powder (diameter between 0.1 and 10 pm), but differs significantly from the values published by Attar and Dupuis (1979) for calcite and dolomite a t temperatures below 640 "C (between 17 and 19 kcaYmo1). (v) In the 820-870 "C range, where ZR is assumed negligible, the logarithm of the initial rate (based on a reaction time of 30 min) is also proportional to the inverse of the absolute temperature: ln(reaction rate) = -(1/2) ln(z,,t,)
(13)
Fitting eq 13 to the high-temperature data of Figure 4 yields a value of Eo somewhere between 64 and 81 kcaY mol. The precision is not high because the regression
was based on only six points (three based on 30-min conversion data, and three based on 15-min conversion data) over a temperature range of only 50 "C (set by the calcination temperature, around 900 "C). The lower value, 64 kcallmol, was used t o plot the curve presented in Figure 4; however, using the higher value does not have much impact on the curve's location. The diffusion of HzS through Cas was studied by Borgwardt et al. (1984) in their work on the sulfidation of grains of CaO. The value of the diffusion coefficient of HzS through the Cas product layer should not greatly depend on whether the starting material is CaO or CaC03. Borgwardt et al. (1984) found an activation energy for the diffusion process of the order of 62 kcallmol. (vi) The determination of the diffusion coefficient of HzS through the grains over the entire temperature range is modeled as follows: At low temperatures (below 665 "C), the diffusion of HzS through the nonrecrystallized Cas layer did not appear to be rate-limiting, at least in the very early stages of the reaction (up to about 5% conversion). However, above 665 "C, the Cas crust formed around the grains recrystallized and then started to sinter. The recrystallization of the formerly flat Cas crust produces a sharp-angled layer, thus providing a significant driving force for sintering. The two phenomena, Le., recrystallization and sintering, are actually taking place simultaneously. Furthermore, the average thickness of the Cas layer also varies with time, as the reaction proceeds. A comprehensive modeling of these three simultaneous phenomena would be a rather complex task and would require a good knowledge of the intrinsic kinetics of recrystallization and sintering of Cas layers formed on micron-size CaC03 grains, for which little is known. Nonetheless, we can offer a semiquantitative approach that will account for the physical rearrangements we observed on the grain surface, by "decoupling" the three physical and chemical phenomena. The overall process is then crudely approximated in Figure 9. It is assumed that all the "bumps" are cubic and all have the same height (hd. This lattice then proceeds to sinter, the loss of surface area resulting from the transfer of Cas from the cubic crystals to the void space located between them. The total volume of Cas is conserved during the sintering process (no densification of the Cas). This results in reducing the height of the original cubes, and in filling the spaces between them. Ultimately, as the sintering is completed, the Cas surface would become flat again, with a thickness of hd2.
2332 Ind. Eng. Chem. Res., Vol. 34,No. 7, 1995 The sharp-edged cubic shape of the partially sintered crystals does not provide a very realistic description of the actual shape of the CaS surface, which quickly loses its sharp features (Fenouil, 1992). However, it approximates the surface area evolution while sintering is occurring. Right aRer the recrystallization, and before any sintering has taken place, the surface area of the nascent CaS is equal to 6h02per unit cell (a cube plus a void space) and decreases to 6h02(l - 4d3) as the size of the original cube decreases to ho(1 - a) and the formerly void space is now filled by a CaS layer of hoa thickness. It is now possible to determine the flux of H2S through the CaS layer. Considering the diffusion to be onedimensional, perpendicular to the CaC03 surface, the flux per unit cell (area 2h02)is given by
[(A)(F+q)]x 1 -2a
Jcell=
+
a
-1
AC
simultaneously with the sulfidation reaction. Upon examination of eq 17 it seems reasonable to postulate the following functionality for a: a=Aexp
(--R:T)
-t
(18)
where A and B are two constants. Within the time frame of our experiments on the initial sulfidation rate (30min), the loss of surface area remains negligible at T, = 665 "C (yielding a = 01, but will reach its asymptotic value well before 30 min if the temperature is sufficiently high (Le., yielding a = 0.5 in our model for an "infinite" temperature). Based on these facts, the average value of a as a function of the absolute temperature is given by
(14)
for T L T,
which leads t o the following effective diffusivity Dg
(19)
where DK represents the Knudsen diffusivity of HzS through the void space between the CaS crystals, and D, the solid state diffusion coefficient of H2S through CaS (D, = Do exp(-EdRT)). The first term in equation 14, DJ(1- a),accounts for the diffusion of HzS through the bumps of height ho(1 - a), while the second term accounts for the diffusion of H2S through the pore of height ho(1 - 2a) followed by the diffusion through the CaS layer of thickness hoa (the diffusion resistances are in series). The value Of DK can be estimated with the following equation:
D,
M
9i'h0(T/34)1'2
(16)
As mentioned earlier, the value of Eo is about 64 kcaY mol. However, the value of DOcannot be easily measured. Furthermore, solid-state diffusion coefficients are extremely sensitive t o factors such as the gas-phase composition and the presence of impurities in the solid. Therefore, the value of Do will be kept as an adjustable parameter in this analysis. Finally, we now need to relate a to the sintering conditions. As mentioned earlier, a is obviously timedependent. However, for the purpose of modeling the initial sulfidation rate of limestone with the grain model described earlier, we will assume that a time-averaged value of a &e., the surface area), still a function of the experimental conditions, can adequately describe the overall reaction kinetics during the first 30 min. The kinetics of the sintering of Cas under simulated coal gas was studied by Fenouil et al. (1994). It was established that the sintering kinetics of CaS is well described by the following equation:
so- s SO =
[ko,
l/W
exp( - &)t]
= 4a
(17)
where SOis the initial surface area, S is the surface area, t is the time, T is the absolute temperature, R, is the ideal gas constant, w is equal to 4.6, K O s is equal to 33333 s-l, and E , is equal to 48 kcaymol. Equation 17 still contains time as variable, reflecting the fact that sintering is an ongoing phenomenon that takes place
In conclusion, above 665 "C, the observed average reaction rate is given by eq 11,where an average value for D, can then be obtained by inserting the average value of a given by eq 18 into eq 15. Below 665 "C, the CaS layer is much more permeable, and the diffusion of HzS through it is not rate limiting. Thus, eq 11 reduces to
Figure 4 shows the relatively good agreement between the experimental results and the crude model developed above. It should be noted that the model contains virtually no adjustable parameters (besides the preexponential factor in the expression of D, which was taken at 0.0035 mz/s in Figure 4) and accounts for all the physical phenomena observed during the early stages of the reaction between limestone and H2S in simulated coal gas. The model is also fairly insensitive to the value of ho ( i e . , DK): no change is observed in the overall reaction rate for values of ho ranging from 1 nm to 10 pm, the entire possible range for ho.
Conclusions The kinetics of the reaction between HzS and large particles of uncalcined limestone under a simulated coal gas was studied in the entire temperature range for which the reaction is thermodynamically favorable. We concluded that the conversion of CaC03 to CaS is limited to about 10%because of the sintering of the Cas product layer. We also identified two regimes for the initial reaction rate: at low temperatures (up to 660 "C), the chemical reaction is the limiting step of the kinetics, whereas, at higher temperatures, the diffusion through the CaS layer becomes limiting. A simple model describes the reaction kinetics over the entire temperature range, and is consistent with the structural changes observed in the sulfided samples. Nomenclature A = see eq 18 B = see eq 18
Ind. Eng. Chem. Res., Vol. 34, No. 7, 1995 2333 C = concentration of the gaseous reactant (mol/m3) Ceq = equilibrium concentration of the gaseous reactant
(moYm3) De = effective diffusivityof H2S within the sorbent particles (m2/s) D, = diffisivity of H2S in the grains of the sorbent particles (m2/s) DK = Knudsen diffusion coefficient (m2/s) DO = preexponential factor of the solid-state diffusion coefficient (m2/s) D, = diameter of the sorbent particle (m) D, = solid-state diffusioncoefficient = D Oexp(-EdR,T) (m2/ S)
= activation energy of the chemical reaction (J/mol) Eo = activation energy of the solid-state diffusion coefficient (J/mol) E, = activation energy for the sintering mechanism of Cas (J/mol) Jcell = flux of H2S to the grain surface (mol/(m%)) ho = initial height of the Cas crystal (m) k, = external mass transfer coefficient in the gas film around the sorbent particles (m/s) ko = preexponential factor of the reaction rate constant (m/ Eact
S)
ko, = preexponential factor for the Cas sintering kinetics (see eq 17) k, = reaction rate constant (see eq 2) ((cmL)/(mol of HzSnmin)) k, = reaction rate constant = KO exp(-EacJR,T) ( d s ) r = linear regression coefficient R = radius of the sorbent particles (m) R, = ideal gas constant (8.31439 J/(mol*K)) r, = average radius of the grains (m) S = specific surface area (m2/kg) So = initial specific surface area (m2/kg) t = time (s) to = duration of the exposure of limestone to simulated coal gas (SI T = absolute temperature (K) Tc = transition temperature (665 "C) X = conversion of the solid reactant (X varies between 0 and 1 ) a = fractional decrease in height of the Cas crystals during sintering E" = porosity of the sorbent e, = molar density of the solid sorbent (mol/") ZDG = characteristic time for diffusion through the grain (S)
characteristic time for diffusion through the entire sorbent particle (s) M= T characteristic time for external mass transfer from the bulk gas to the surface of the pellet (film diffusion)
t ~ = p
~
(9)
characteristic time for chemical reaction at the interface between the unreacted core and the reacted layer of the grain (s). o = see eq 17 t~ =
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Attar, A.; Dupuis, F. The Rate and the Fundamental Mechanisms of the Reaction of Hydrogen Sulfide with the Basic Minerals in Coal. Znd. Eng. Chem. Process Des. Dev. 1979,18 (4), 607-618. Borgwardt, R. H.; Roache, N. F. Reaction of H2S and Sulfur with Limestone Particles. Znd. Eng. Chem. Process Des. Dev. 1984, 23, 742-748. Borgwardt, R. H.; Roache, N. F.; Bruce, K. R. Surface Area of Calcium Oxide and Kinetics of Calcium Sulfide Formation. Environ. Progr. 1984,3 (2), 129-135. Borgwardt, R. H.; Bruce, K. R.; Blake, J. An Investigation of Product-layer Diffusivity for CaO Sulfation. Znd. Eng. Chem. Res. l987,26,1993-1998. Brown, D. J. Froth Flotation; Fuerstenau, D. W., Ed.; AIME and Petroleum Engineering, Inc.: New York, 1962. Chan, R. K.; Murthi, K. S.; Harrison, D. Thermogravimetric Analysis of Ontario Limestones and Dolomites. I. Calcination, Surface Area, Porosity. Can. J. Chem. 1970, 48, 2972-2978. Chang, E. Y.; Thodos, G. Complex Nature of the Sulfation Reaction of Limestones and Dolomites. AIChE J . 1984,30 (3),450-457. Fenouil, L. A. Structural Studies in Limestone Sulfidation. Master of Science Thesis, University of California at Berkeley, 1992. Fenouil, L. A. Kinetic and Structural Studies of the Sulfidation of Large Particles of Lime and Limestone in Coal Gas. Ph.D. Dissertation, University of California at Berkeley, 1995. Fenouil, L. A.; Lynn, S. Study of Calcium-Based Sorbents for HighTemperature H2S Removal. 2. Kinetics of H2S Sorption by Calcined Limestone. Znd. Eng. Chem. Res. 1995,34,2334-2342. Fenouil, L. A,; Towler, G. P.; Lynn, S. Removal of H2S from Coal Gas Using Limestone: Kinetic Considerations. Znd. Eng. Chem. Res. 1994, 33 (21, 265-272. Freund, H. Kinetics of Limestone/Dolomite with H2S under Rich Combustion Conditions. Combust. Sci. Technol. 1981,26, 8388. Fuller, E. L., Jr.; Yoos, T. R. Surface properties of Limestones and their Activation Products. Langmuir 1987, 3, 753-760. Hartman, M.; Pata, J.; Coughlin, R. W. Influence of Porosity of Calcium Carbonates on their Reactivity with Sulfur Dioxide. Znd. Eng. Chem. Process Des. Dev. 1978, 17 (41, 411-419. Levenspiel, 0. Chemical Reaction Engineering, 2nd ed.; John Wiley & Sons: New York, 1972; Chapter 12. Lynch, A. J.; Johnson, N. W.; Manlapig, E. V.; Thorne, C. G. Mineral and Coal Flotation Circuits. Developments in Mineral Processing, Vol. 3; Elsevier Sci. Publishing: Amsterdam-OxfordNew York, 1981. Nowacki, P. Coal Gasification Processes; Noyes Data Corporation: Park Ridge, NJ, 1981. Ruth, L. A.; Squires, A. M.; Graff, R. A. Desulfurization of Fuels with Half-Calcined Dolomite: First Kinetic Data. Environ. Sci. Technol. 1972, 11 (5), 488-491. Squires, A. M.; Graff, R. A,; Pell, M. Desulfurization of Fuel with Calcined Dolomite. I. Introduction and First Kinetic Results. Chem. Eng. Prog. Symp. Ser. 1971, 67 (115), 23-34. Sun, C. C.; O"eil1, E. P.; Keairns, D. L. The Sulfidation and Regeneration of Half-Calcined Dolomite. Thermochim. Acta 1978,26,283-296. Szekely, J.; Evans, J. W.; Sohn, H. Y. Gas Solid Reactions; Academic Press: New York, 1976. Towler, G. P. Synthesis and Development of Processes for the Recovery of Sulfur from Acid Gases. Ph.D. Dissertation Part I, University of California at Berkeley, 1992. Yen, J. H. Kinetic and Structural Studies in the Sulfidation of Dolomite. Ph.D. Dissertation, Carnegie-Mellon University, 1979.
Received for review July 7, 1994 Revised manuscript received April 18, 1995 Accepted April 28, 1995* IE940422F
* Abstract published in Advance ACS Abstracts, J u n e 15, 1995.
