Ind. Eng. Chem. Res. 1995,34, 2334-2342
2334
Study of Calcium-BasedSorbents for High-Temperature H2S Removal. 2. Kinetics of H2S Sorption by Calcined Limestone Laurent A. Fenouilt and Scott Lynn* Department of Chemical Engineering and Lawrence Berkeley Laboratory, University of California, Berkeley, California 94720
Samples of calcined limestone particles having a diameter of about 1 mm were exposed to simulated coal gases containing between 500 and 18 000 ppm HzS for temperatures ranging from 560 to 1100 "C in a differential tube reactor. The formation of C a s was followed quantitatively as well as qualitatively to elucidate the reaction mechanism. Contrary to the limited conversion of CaC03 to Cas, it was found that the limestone particles could be completely converted to C a s by 1%H2S in about 1h if the particles are precalcined or ifthe rate of calcination is higher than the rate of sulfidation. The reaction then takes place between CaO and H2S and follows a shrinking-core mechanism. The reaction kinetics is controlled by the diffusion of HzS through the pores of the C a s product layer formed around the lime particle (effective diffusivity and 5.1 x m2/s>. The kinetics of the sorption of H2S by CaO is relatively between 2.8 x insensitive to the reaction temperature, and the reaction rate does not decrease significantly when the CaO is severely sintered for several hours a t 1050 "C prior to sulfidation.
Introduction and Previous Work It was established in part 1(Fenouil and Lynn,1995) that "large" particles of limestone (1-2 mm diameter) cannot be completely converted to Cas. The sintering of the Cas product layer formed at the surface of the pores originally present in the limestone forms an impermeable barrier when the temperature of the gas stays below the calcination temperature of CaC03 (about 900 "C under 1 bar of C02) (Fenouil, 1992; Fenouil et aZ.,1994). However, it was found that the CaC03 is completely converted to CaS if the sulfidation is carried out above the calcination temperature. This paper describes the kinetics of the reactions of HzS with both lime and limestone a t atmospheric pressure, with emphasis on the effects of the evolution of the structure of the solid during the course of the reaction. Below the calcination temperature of limestone the main reaction of importance is CaCO,
+ H2S
Cas
+ H20 + CO,
(1)
which becomes increasingly favorable as the temperatures increases. At a given fugacity of COS,limestone also undergoes calcination if the temperature is high enough: CaCO,
t CaO
+ CO,
(2)
The external dimensions of the solid are observed to remain constant during calcination (Boynton, 1980; Fuertes et aZ.,1991,1993), suggesting that the porosity of the solid increases with conversion, which would lead to a fast reaction and high conversion. "he lime'formed can then react with HzS: CaO
+ H2Sz= C a s + H 2 0
(3)
However, at the pressure and composition of the effluent from most coal gasifiers, lime may sinter. Once sintered (or "dead burned"), it has been feared that lime would sorb H2S relatively slowly, as is observed with Current address: Chemical Development Department, Shell Chemical, 3333 Highway 6 South, Houston, TX 77082. +
0888-5885/95/2634-2334$09.OOJQ
SO2 sorption by sintered CaO. It was thus of interest to investigate the effects of the CaO sintering on the sulfidation rate. Reaction 1is endothermic (about +165 kJ/mol at 900 "C) whereas reaction 3 is exothermic (about -65 kJ/mol at 900 "C). Thus the lowest level of H2S thermodynamically possible in a coal gas in contact with lime(stone) occurs at the calcination temperature of the CaC03, which is only a function of the partial pressure (i.e., fugacity) of COz. Thermodynamic calculations show that, for most gasifiers, the level of removal of HzS could be well above the 90% mandated by the Clean Air Act, and often as high as 98% (Figure 1). Below the calcination temperature of CaC03, COz, and H2S will compete to react with CaO. The feasibility of a process that utilized CaO to remove H2S from coal gas below the calcination temperature of CaC03, would thus depend on the relative rates of reaction 3 and the reverse of reactions 1 and 2. Finally, it should be remembered that the composition of a coal gas is dictated by the initial C:H:O ratio and the equilibrium of the water-gas-shift reaction (WGS) as a function of temperature:
CO
+ H20t CO, + H,
(4)
Failure to account for calcination and WGS equilibration has cast doubt on the validity of many previous studies of this system; therefore a fresh evaluation of the kinetics is being made. The influence of the WGS reaction, the effect of the thermal decomposition of H2S at high temperature into SZ and H2, and the oxidation of CO to COS will be carefully considered in this work. More particularly, the kinetic model for thermal H2S dissociation developed by Towler (1992) was used to estimate the exact composition of the gas reacting with the sample in the reactor.
Experimental Section The differential tube reactor described in part 1 was used to produce sulfided samples and to gather kinetic data for reactions 1 and 3. Preliminary experiments were designed primarily to demonstrate the feasibility of sulfiding large lime particles (around 1 mm radius) 0 1995 American Chemical Society
Ind. Eng. Chem. Res., Vol. 34,No. 7, 1995 2335 600
500
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CO, c 5% , H,O 5% c CO, 3% , H,O 4%
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0
800
900
Temperature,
1000 O C
Figure 1. Equilibrium concentration of H2S in coal gas as a function of the temperature and the C02 and steam content of the gas.
to a very high conversion with simulated coal gas. They also allowed rapid assessment of the influence of experimental conditions on the conversion to Cas. The reactor was operated under packed-bed, rather than differential, conditions. Operating under nondifferential conditions allowed us to react a larger quantity of limestone sample in each experiment, increasing the precision of our conversion measurements, which are based on the weight change during the reaction. Most experiments were performed on minibeds of CaO (0.2-0.3 g of lime) with a gas containing about 1.7% HzS. The mass-average diameter of the limestone particles was measured a t 1.75 mm with a standard deviation of about 10%. In most cases, the initial reaction rate was about 5 % conversion to CaS per minute or less. The gas flow past the bed in all the experiments described in this section was between 5 and 6 mL/s (STP), which corresponds t o about 4 x to 5 x mows HzS (for 1.7% HzS in the gas phase). A n overall rate of 5%/min CaS formation amounts t o 1.6 x to 2.5 x mol CaS formed per second for a bed weight of 0.2-0.3 g (typical quantity of solid used during this series of experiments). The H2S concentration thus dropped by 30-65% over the length of the bed during the initial stage of the reaction. At later times the bed became closer to a true differential reactor. The Great Lakes limestone used in part 1 was also used here. It contains only about 2.2% of a few inert impurities plus magnesium carbonate, MgC03. If the only products of the limestone reaction are CaS and CaO, the conversion to CaS (starting from CaC03) can be estimated by a gravimetric measurement: %
when CaC03 was also present in the final products of the reaction. These titrations, combined with a gravimetric measurement, were sufficient to determine unequivocally the relative proportion of the three calcium species a t the end of the reaction. Titrations were also performed to check the validity of eq 5 in cases where the only calcium compounds were known to be CaO and Cas. Agreements between the results of the titration and the values obtained from eq 5 were always within 5%. All experiments were operated under sufficiently reducing conditions that Cas03 and Cas04 are not thermodynamically favored (Fenouil, 1995). Industrial quality (99.9%pure) gases were supplied by Mattheson Gas Products (East Rutherford, NJ). The kinetics of the limestone calcination was followed by gravimetric measurements considering limestone as pure calcium carbonate:
final weight - 0.56 x initial weight = loo( (0.72- 0.56) x initial weight (5)
Allowing for the impurities and the presence of MgO/ MgC03 introduces an uncertainty of only about 1%in the conversion value. This is below our total experimental accuracy of about 3%, mainly caused by inaccuracies in our weight measurements, except for conversions below 5% or above 95%. Titration of the sulfide (using a standard iodometric titration) was required
)
(6)
As in part 1, analysis of reacted or sintered samples was made using energy dispersive spectroscopy ( E D 9 and scanning electron microscopy (SEM). To determine the distribution of CaS in the reacted limestone particles, some samples were polished to allow X-ray mapping of a cross section of the solid samples. CaS is fluorescent under X-rays, which allows clear identification of regions that have undergone sulfidation. Comparison of X-ray and SEM pictures of the same sample showed that lighter gray regions on the SEM pictures correspond t o zones rich in Cas, whereas darker gray areas correspond to unreacted limestone or lime.
Minibed Kinetics Experiments Limestone Calcination. Several limestone samples were calcined under an HzS-free simulated coal-gas atmosphere to establish the basis for comparing the kinetics of lime sulfidation with that of limestone calcination. We used similar operating conditionssample size, temperature, gas flow rate, and composition (except for HB-in the calcination and the sulfidation experiments described later in this paper. It was found that freshly calcined limestone samples can reabsorb COz if they are cooled from above 900 "C to ambient temperature under an atmosphere rich in COz. In our experiments the lime samples needed about 2-4 min to reach ambient temperature and during this period 20-30% of the CaO reacts back t o CaC03. This explains the large differences in CaO conversion observed in Figures 2-4 between samples that were calcined under the same conditions but cooled under different ones (Nz in one case and COz in the other). Figure 2 shows that limestone calcination is practically complete after 15 min if the reaction takes place 45-50 "C above the calcination temperature (890-900 "C under 1bar of COz). However, the time for complete conversion increases significantly as we approach the calcination temperature. At 10-15 "C above the calcination temperature, the calcination rate becomes very slow and complete conversion requires more than 2 h (Figure 4). It is also worth noting that a small change in the CO2 partial pressure in the gas above the limestone particles has a very significant effect on the calcination rate: a decrease of the C 0 2 partial pressure from 1to 0.88 bar (the gas composition being CO2 (88%),Hz (0.5%),H2O (5.75%),and CO (5.75%))more than doubles the initial calcination rate at 930 and 950 "C (Figures 2 and 3)).
2336 Ind. Eng. Chem. Res., Vol. 34, No. 7, 1995 1 GO
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';4ol 5
1
v V calcined under 88% COz. cooled under N2 calcined under 100%CO,, cooled under N 2 0 calcined under 88%C O P , cooled under COz
0
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20
0 calcined under 88%C02, cooled under C02
0
0 0
5
0
10
15 20 time (min)
25
1 I 1 I 150 183 210 240 time ( m i n ) Figure 4. Limestone calcination under various conditions a t 915 "C (see text for more details).
0
30
Figure 2. Limestone calcination under various conditions at 950 "C (see text for more details).
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5
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20
V calcined under
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V 0
40
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0 1
15 20 time (min)
1
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25
30
Figure 3. Limestone calcination under various conditions at 930 "C (see text for more details).
We observed a sharp difference in the morphology of the center of limestone particles calcined under Werent conditions. A much smaller number of small pores was observed on SEM pictures when the calcination was performed under the C O D 2 mixture for 240 min. This is the result of the much stronger sintering undergone by these samples. Limestone Sulfidation under Calcining Conditions. In contrast to the reaction between CaC03 and HzS (see part l),the reaction rate between CaO and HzS is rapid enough to allow complete conversion of CaO t o Cas in about 1-2 h for lime particles of 1 mm in diameter exposed to about 1.7%H2S in the gas phase (Figure 5). The data a t 945 "C were obtained using limestone and the following gas phase composition: CO2,88%;CO and Hz0,4.75%;H2,0.8%;H& 1.7%.At 945 "C the rate of limestone calcination is larger than the observed sulfidation rate (see Figure 2), and the actual chemical reaction is between CaO and H2S, not between CaC03 and HzS. For the experiments at 560
20
40
60 80 time ( m i n )
1
100
120
Figure 5. Effect of temperature on lime conversion, sulfidation taking place under the following conditions: 87.5% C02,5.25% CO and H20, 0.7% Hz, 1.3% HzS (1015 "C) (no precalcination); 88%C02, 4.75% CO and H20,0.8% Ha, 1.7% HzS (945 "C) (no precalcination); 93.3% Nz,5.0% Ha,1.7% H2S (560"C)(precalcined for 30 min under N2 at 865 "C).
"C it was necessary first to calcine the limestone (30 min a t 865 "C under pure nitrogen) and then react the
freshly formed lime with the following gas mixture (free of COz to prevent any back reaction between CaO and CO2): 93.3%N2; 5% H2; 1.7%H2S. Figure 5 not only demonstrates that the reaction can go to completion but also shows that the effect of temperature on the reaction rate is quite small under the conditions of our experiments. When the data presented in Figure 5 are plotted in a log(x)us logW graph, the experimental data points appear to form a straight line with a slope between 0.5 and 0.6, which is characteristic of a rate controlled by diffusion through a product layer (here Cas) that is growing as the reaction proceeds. When the limestone is exposed t o H2S above the calcination temperature, the sulfidation reaction occurs in two stages: first, the limestone calcines and then the
loo
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-7 V
01
1
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precalcined limestone
I
0
20
40
60 80 time (min)
100
120
Figure 7. Effect of limestone precalcination on the sulfidation rate a t 945 "C (gas composition: COz (88%), CO and H2O (4.75%), Hz (0.8%),and HzS (1.7%)).
freshly formed CaO reacts with HzS to form Cas. Figure 6 compares the sulfidation kinetics a t 915 "C between precalcined limestone particles (30 min at 865 "C under Nz) and particles in which the calcination and the sulfidation are simultaneous. In both cases the gas composition was similar t o that used in the 945 "C experiments of Figure 5. At 915 "C the intrinsic calcination rate, from CaC03 to CaO, is lower than that of sulfidation, from CaO to Cas. At 860 "C, in contrast, CaC03 cannot calcine and the conversion t o CaS stays below 10% even after 240 min (Figure 6). One concludes that the conversion of CaC03 to CaS is limited by the rate of lime formation and not by that of lime sulfidation. In limestone samples reacted at 915 "C without precalcination, precise gravimetric measurement as well as precise sulfur titrations reveal that less than 4% of the calcium atoms are present as CaO at any given time, those remaining being divided between CaC03 and Cas. This confirms that the reaction is limited by CaO formation and that CaO reacts with HzS almost as soon as it is formed. At 945 "C (which corresponds to about 50 "C above the limestone calcination temperature for a gas containing 0.88 atm of COz) the calcination rate is rapid enough not to interfere with the sulfidation reaction and precalcining the limestone has very little effect on the overall conversion of the stones to CaS (Figure 7, same simulated coal gas as in Figure 5). Figure 5 also contains data that were obtained under the same conditions as those shown in Figure 7 except for the temperature (1015 "C instead of 945 "C). At this temperature the HzS decomposition into HZ and SZ is such that the H2S concentration in the gas phase is between 1.2 and 1.4% instead of about 1.7% as in the other experiments. Nonetheless, complete conversion is still attained in 2 h and the calcium carbonate content of the samples is always under 6%at any point during the reaction. This confirms the weak effect of temperature on the sulfidation rate and also shows that CaO sintering, much more significant at 1015 "C than at 945 "C (Borgwardt, 1985), has little effect on the reaction kinetics. This also proves that sulfidation still goes t o completion a t 1015 "C, even if the sintering of CaO is very severe at this temperature. Thus, lime sulfidation
can be carried out a t temperatures well above the calcination temperature of CaC03. Limestone particles were calcined for 15 min under Nz a t 915 "C and then respectively exposed to 9000 ppm HzS in Hz for 15 min or to 900 ppm HzS in 90% NZand 9.1% HZ for 360 min. SEM pictures of the inside of the partially sulfided limestone particles reveal a sharp interface between the CaS outer shell and the CaO core, the interface being even sharper for the stones exposed to 900 ppm HzS (Fenouil, 1995). This sharp CaO/CaS interface can also be observed with limestone particles directly exposed at 1015 "C to the simulated coal gas described for Figure 5: this is not surprising, since the limestone calcination is completed in less that 10 min under these conditions. Parts a, b, and c of Figure 8 show the evolution of the CaO/CaS interface with time when the reaction takes place at 1015 "C (respectively after 5, 30, and 240 min). Close-up SEM photographs of the center of Figure 8c also reveal that the CaS product contains a large number of micron-size pores (Fenouil, 1995). The situation is slightly different when the limestone samples are exposed t o the same simulated coal gas at 915 "C. As mentioned above, the overall conversion of the limestone to CaS is now controlled by the formation of CaO from the limestone. CaO-rich regions located near the original pores of the limestone grow inside the whole particle. It should, however, be noted that the distribution of the CaO-rich regions is not completely homogeneous within the limestone particles: their concentration appears to decrease as we approach the center of the stones. Consequently, early in the reaction, the CaS is located only near the surface and cannot be found deep inside the stone. However, the C a s outer layer contains a substantial fraction of unreacted CaC03 as observed in Figure 9 (electron dispersive microscopy (EDS) picture of the sulfur atoms in the outer layer of a stone exposed for 15 min t o the same simulated coal gas as in Figure 5). Therefore, under these conditions, the reaction mechanism is controlled by the calcination mechanism and follows neither a sharp-interface,shrinking-core mechanism (as observed at higher temperatures) nor a homogeneous reaction throughout the entire
2338 Ind. Eng. Chem. Res., Vol. 34, No. 7, 1995
Figure 9. Cross-sectional sulfur map of a partially sulfided limestone sample at 915 "C (EDS) under 88%C02,4.75% CO and H20,0.8% H2, and 1.7% H2S. The region with the white spots (sulfur atoms) corresponds to the surface of the stone, its center being toward the lower left corner of the picture.
Figure 8. Cross section of a lime particle exposed to 9000 ppm H S in H2 at 1015 "C (SEM) (lime obtained from limestone calcined 15 min under N2 a t 915 "C) (a, top) for 5 min, (b, middle) for 30 min, and (c, bottom) for 240 min.
stone, but rather a combination of both. SEM pictures obtained from samples exposed for 2 h a t 915 "C to simulated coal gas reveal that Cas patches can be found throughout the stone (Fenouil, 1995). EDS sulfur maps on reacted limestone samples confirm the previous observations. Sulfur maps of the outer shell of particles obtained at 945 "C and the same gas composition as in Figure 5 reveal a much sharper CaO/CaS interface than when the samples are sulfided a t 915 "C, and an even sharper profile is observed when the limestone is calcined prior to sulfidation proving
that the reaction between CaO and H2S is totally controlled by the diffision of H2S through the Cas crust (Fenouil, 1995). Before being exposed to 9000ppm H2S in Ha for 5 min a t 915 "C, some samples were calcined a t that temperature under pure N2 for 15 min and some others for 240 min under an equimolar mixture of CO2 and N2. It was shown by EDS analysis that very little sulfur had reached the center of the stones in either case, further confirming the shrinking-core mechanism of the reaction. CaO Sdfidation Kinetics under Very Low Concentrations of H2S. Calcined limestone samples were exposed to very low levels of H2S to determine the rate of reaction 3 near thermodynamic equilibrium conditions. The purpose of these experiments was to determine whether there was a fundamental change in the apparent reaction kinetics as the level of H2S was reduced to near that desired in a cleanup process. We expect that H2S diffision through the product layer will still be the limiting step in the reaction kinetics. Consequently, we also expect a sharp decrease in the reaction rate (due to a decrease in the diffusion driving force) as the level of H2S in the gas phase approaches thermodynamic equilibrium. The results, presented in Figures 10 and 11, should be considered semiquantitative at best, especially a t very low concentrations of HBS,for the following reasons: (i) The reaction was not carried out under truly differential conditions: it was established that the H2S concentration drop over the bed length during the first 100 min of the reaction a t 915 "C averaged from 25 to 30% for a nominal inlet concentration of 1500 ppm to 40-45% for a nominal inlet concentration of H2S of 500 PPm. (ii)Moreover, as the concentration of H2S approaches the equilibrium concentration under the experimental reaction conditions, a slight fluctuation in these conditions results in a significant variation of the driving force for the reaction, namely the diffision gradient CHzS bulk - CHzS eq. (iii)The determination of CHZsbulk - CHzS eQ is further complicated by the relatively poor knowledge of the exact extent of the reactions taking place in the gas phase to form Sa, SO2, and COS. Our simulated 500 ppm H2S coal gas would in fact result in a mixture of
Ind. Eng. Chem. Res., Vol. 34, No. 7, 1995 2339
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nominal H,S concentration
0 500 ppm V 1000 ppm
F(x)
t=8480
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time (min)
Figure 10. Influence of H2S concentration on limestone conversion at 915 "C. (F(X)is defined in eq 7). T
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400
500
600
700
800
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time (min)
Figure 11. Influence of H2S concentration on limestone conversion at 1015 "C. (F(X)is defined in eq 7).
about 200 ppm H2S, 175 ppm SOz, 125 ppm COS, and 100 ppm SOif enough time were provided for the gas mixture to reach thermodynamic equilibrium a t 915 "C. At 1015 "C, the thermal decomposition of H2S into H2 and SZ is even more pronounced and the major sulfurcontaining species becomes S02. However, the residence time of the process gas was not sufficient to ensure that the thermal decomposition of HzS was complete, even a t 1015 "C. For a thorough discussion of the kinetics and equilibria of these reactions see Towler and Lynn (1993). The extent of the formation of COS and S2 could be estimated with the help of a kinetics computer code developed by Towler (1992). However, little is known about the formation of SO2 resulting from the interaction of H2S and Sz with the main constituents of the simulated coal gas (mainly C02). Consequently, the distribution of the sulfur-
containing species cannot be known with precision for a nominal (Le., inlet) HzS concentration of 1000 ppm or less. Above this level, the concentrations of SOz, COS, and S2 become quite negligible compared to that of H2S, and the actual H2S concentration can be approximated with the inlet H2S concentration with an uncertainty less than 10%. Finally, to complicate things further, COS and S2 can also react with lime t o produce Cas. The kinetics of these reactions have been studied by Borgwardt et al. (19841, Yang and Chen (19791, and Borgwardt and Roache (1984). Nonetheless, Figures 10 and 11 demonstrate that the sulfidation of large particles of CaO still takes place at HzS levels comparable to those expected when thermodynamic equilibrium is reached in o u r simulated coal gas (100-200 ppm). There is no sign of a plateau being reached in the plot of the conversion versus time at 915 "C, even when the lime is exposed to coal gas for more than 10 h. As the temperature increases, the extent of the conversion of CaO to CaS decreases (Figure 11).This can be explained by the combination of the following causes: (i) The extent of the H2S thermal decomposition is much larger at 1015 "C than a t 915 "C. Thus, a given nominal H2S concentration corresponds to a lower equilibrium level of H2S in the bulk a t 1015 "C than at 915 "C, thus reducing the driving force for the reaction. (ii) An increase in the temperature results in an increase in the equilibrium partial pressure of HzS in the coal gas in the presence of CaO since reaction 3 is exothermic. This also decreases the diffusion driving S - CH~S eq, which lowers the conversion force C H ~bulk rate. A nominal concentration of H2S of 500 ppm at 1015 "C in our simulated coal gas results in an actual H2S level slightly below 200 ppm, which is the thermodynamic equilibrium value for CaO sulfidation at this temperature and gas composition. Hence, we observed no conversion of CaO into Cas, even after 15 h (Figure 11). (iii) An increase of temperature also results in an increase of the CaO/CaS sintering rate which can ultimately influence the diffision coefficient of H2S (or HzO) through the CaS product layer. A decrease in this diffision coefficient is associated with a decrease of the reaction rate. Finally, it is clear that at 915 "C the reaction is still diffusion-controlled a t an H2S level of about 500 ppm: the shrinking-core mechanism of the reaction was confirmed by SEM and EDAX analysis. It also seems that the reaction rate is still proportional to the difference between the H2S concentration in the gas phase and its equilibrium concentration (in the presence of CaO) as predicted by the kinetic model based on diffusion limitations that is developed in the section Single-Pellet Kinetics Data. Comparison of the Relative Rates of Reaction 3 and the Reverse of Reaction 2. Figures 12 and 13 explore the possibility of running the CaO sulfidation at temperatures well below the calcination temperature of CaC03, i.e., under conditions for which the backreaction between CaO and COZis thermodynamically favorable. Since the reaction between CaO and H2S is exothermic,it is advantageous to perform the sullidation at the lowest possible temperature to get the lowest possible H2S level in the cleaned coal gas. Figures 12 and 13 show that about 20-30% of the initial CaO reacted to form Cas rather than CaC03 in a simulated coal gas containing CO2 (88%),H2O and CO (4.5%),H2
L,
2340 Ind. Eng. Chem. Res., Vol. 34, No. 7, 1995
100
40 30
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time (min)
Figure 12. CaO to C a s conversion as a function of temperature (gas composition: 88%COz, 4.75%CO and HzO, 0.8% Hz,and 1.7% H2S).
e
0
I
coal gas (Fenouil, 1995). Then, a second bed of zinc titanate sorbents could be operated in the 500-650 "C temperature range to polish the gas to the parts per million level. The main advantage of this hybrid cleanup system would be the extension of the lifetime of the very expensive zinc titanate sorbent by an order of magnitude because of the reduced sorptioddesorption cycle frequency, a t a relatively small cost.
0
e
Figure 14. Conversion of CaO to C a s as a function of time for a single-pellet reactor. (F(X)is defined in eq 7.)
0
Single-Pellet Kinetics Data
v
i?
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v
v
v
500
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700
v 800
900
0
Temperature ( C)
Figure 13. CaCO$CaO/CaS distribution as a function of temperature under a simulated coal gas containing 88%COz, 4.75% CO and H20, 0.8% H2, and 1.7%H2S.
(1.2%), and H2S (1.8%). Therefore, it may be possible to operate a sorption bed of lime in the vicinity of 500600 "C and remove H2S from a coal gas to a level comparable to that obtained by a zinc titanate bed (ie., a few parts per million). However, the sorbent utilization would be quite low (since most of the final product will be CaC03) and a workable scheme might involve at least two moving beds operating a t two different temperatures. (The first would operate at 30-50 "C above the calcination temperature to remove the major fraction of the H2S and the second between 500 and 600 "C to further decrease the H2S concentration in the coal gas. The spent sorbent from the second bed could then be fed to the first one.) A third bed to calcine the limestone needed for the second bed would also be required. A simpler scheme might involve a moving bed of limestone particles operating slightly above the calcium carbonate calcination temperature. This bed could remove up to 90% of the H2S initially present in the
Having established the feasibility of the removal of H2S by CaO, we then changed our experimental procedure from a "mini-bed" reactor to a "single-pellet" differential reador to gather detailed and precise kinetic data for design purposes. The mass-average diameter of the limestone particles used for these experiments was found to be 1.75 mm with a standard deviation of about 10%. Each experiment was actually carried out with batches of three to six stones, corresponding t o a total weight of 20-30 mg, to guarantee the differential character of the reactor (except in the very early stages of the reaction). Moreover the gas flow was increased to its largest possible value to prevent any significant gas-phase mass-transfer resistance (about 5 mLJs (STP)). Calculations show that even if external mass-transfer resistance cannot be made completely negligible in the early stage of the reaction, it does not play a large role in the conversion data obtained using a gas containing either 9000 or 900 ppm of H2S in H2 and N2. The samples were first calcined for 15 min under Nz (5 mL/s (STP))at 915 "C before being exposed t o one of the two following process gases: N2 (go%), H2 (9.1%), and H2S (900 ppm) or Hz (99.1%)and HzS (9000 ppm). Each experiment was repeated three t o six times t o obtain a better estimate of the conversion. Figure 14 displays the results of these experiments; the error bars correspond to one standard deviation. It was also determined that the calcination procedure had little effect on the subsequent sulfidation rate. Figure 15 shows that we lose some initial reactivity if we increase the duration of the exposure of CaO t o high temperature and increase the fraction of C02 during the calcining stage. CaO is then more strongly sintered, which decreases the surface area initially available to
Ind. Eng. Chem. Res., Vol. 34, No. 7, 1995 2341 0
9000 ppm H,S
balance
80 -
H2&
CT) (d
u bp
v
I
I
h
60
t=94 2 F(x)
El
; 5
I
T
calcination procedure N,
8
17 m m at Q15'C under
v
32 mln a t 815'C under 5 0 Z C0,/50%
Figures 14 and 15 also display the best fit on the experimental kinetic data using eq 7. The values of z1 obtained from these graphs are probably slightly overestimated since the reaction was partially limited by external mass transfer of HzS during the first 5-10 min of the reaction. The values of De we extract from z1 will thus be slightly underestimated, leading to conservative values of the kinetic parameters of the reaction. From m2/s for the 9000 Figure 14 we obtain De = 5.1 x ppm data and De = 2.8 x m2/s for the 900 ppm data (es = 27 k m ~ l * m and - ~ R = 0.875 mm). The effective diffusivity of HzS through the CaS layer can be determined using the random pore model developed by Wakao and Smith (Froment and Bischoff, 1990):
+
De = E : ( D ~ - ~DM-l)-'
N,
mln a t 915'C under 50% C0,/50Z N, 0 1 IO min aL 1O6O0C under 50% CO,/SOX N,
(9)
V 252
0
10
20
30
40
50
60
time (min) Figure 15. Conversion of CaO to CaS as a function of time for a single-pellet reactor. (F(X)is defined in eq 7.)
H2S. However, we still obtain nearly complete conversion t o CaS in 1 h. We conclude from this that the structure of the lime is not very important and the reaction is controlled by the diffusion of HzS through the CaS product layer. This is true except for the earliest stages of the reaction, when an increase in reactive surface area is associated with an increase in reaction rate. It was also shown that the presence of a simulated coal gas instead of the N2/H2/HzS mixture had no noticeable influence on the reaction kinetics for a 900ppm level of H2S in the gas phase. The following gas composition, C02 (24.7%), CO (8.5%), HzO (8.5%),HZ (2.1%), N2 (56.1%), and HzS (900 ppm), instead of 900 ppm of HzS in 90% Nz and 9.1% H2, produced no measurable changes in the reaction kinetics (provided the same calcination procedure was followed). As mentioned in the section Minibed Kinetics Experiments, SEM pictures and sulfur maps (EDS) on the reacted limestone particles confirm that the reaction takes place following a shrinking-core mechanism, provided we are a t least 20-30 "C above the calcination temperature of CaC03. Thus, the experimental data were compared to the shrinking-core model (which is identical to the gain model developed in part 1 when no kinetic limitation is present a t the grain level) for spherical particles in the case where the diffusion through the product (i.e.,Cas) layer is the limiting step. Under these assumptions the shrinking-core model yields the following kinetic equation (Levenspiel, 1972):
t = tl[l - 3(1 -
mW3+ 2(1 - X)]= t,F(X) (7)
The other terms are defined in the Nomenclature. Traditionally, z1 is defined neglecting C,, (Levenspiel, 1972). This is correct when Ceq is significantly smaller than C . However, to describe the sulfidation kinetics of the lime for the full range of HzS concentrations, including the values near thermodynamic equilibrium, it is necessary to include C,, in the expression for z1.
where EO is the porosity of the CaS shell, D K is the Knudsen diffusivity of H2S given by
DK = 9 7 R p O r e m
(10)
and DM is the molecular diffusivity of H2S in the gas mixture which can be estimated, using the Fuller et al. correlation (Reid et al., 1987) to be 7.5 x m2/s in Hz and 2.0 x m2/sin a mixture of 90% NZand 9.1% Hz at 915 "C. Using the two values of De obtained for 900 and 9000 ppm of H2S, one obtains EO = 0.15 and D K = 3.2 x m2/s &e., an average pore radius of 0.55 pm for the CaS layer). Based on the values of the densities of CaC03 (2.70 g/cm3),CaO (3.31 g/cm3),and CaS (2.61 g/cm3), EO should be equal t o 0.25 plus the initial porosity of the limestone (generally very low, between 0 and 0.08) if no shrinkage occurs during the calcination or the sulfidation stage. However, sintering has been observed with both CaO and CaS in coal gas conditions (Borgwardt, 1989; Fenouil et al., 1994) and shrinkage of the order of 10% has also been.observed during limestone calcination (Fuertes et al., 19931, so an average value of 0.15 for the CaS porosity is quite plausible. For cylindrical pores of equal diameter, the specific surface area of the CaS layer can be estimated from the specific void volume and the average pore radius:
(11) which leads to a value of 11 m2/g. There is however another, slightly different way to interpret the data. The shrinkage of the stones is more and more significant as the severity of sintering increases. Generally, the surface area is first lost rapidly and then reaches an asymptotic value (ranging from 1 to 10 m2/g) before the loss of porosity begins to be noticeable (Coble, 1961). One can thus assume that the average radius of the pores in the Cas layer quickly reaches its final value and remains roughly constant while the porosity is decreasing. In this case the value of the Knudsen diffisivity remains constant throughout the sulfidation of the lime and only EO varies as the reaction proceeds. The lower value for De in the experiments using 900 ppm HzS could then be explained by the longer duration of the experiments (compared to those with 9000 ppm H2S) and, hence, of the larger shrinkage experienced by the stones. The difference in the two values of De could be explained by a loss of just 0.07 (from 0.25 to 0.18) in the porosity of the Cas layer. Using EO = 0.25 for the experiments carried out under 9000 ppm H2S, one gets DK = 1.4 x m2/s from eq
2342 Ind. Eng. Chem. Res., Vol. 34,No. 7, 1995
9, which corresponds to an average pore radius of 0.24 pm in the CaS layer. Both analyses result in quite similar values for €0 as well as for the effective diffusivity of H2S in the C a s layer, and for design purposes one might want to choose the lowest estimate of the Knudsen diffisivity (1.4 x m2/s)for a conservative estimate of the sulfidation kinetics.
Conclusion In contrast to the direct conversion of CaC03 to Cas, which was limited to less than 10%in 1h, it was found that limestone particles could be completely converted to Cas by 1% H2S within this time if the particles are either precalcined or if the rate of calcination is higher than the rate of sulfidation. The reaction then takes place between CaO and H2S and follows a shrinkingcore-mechanism. The reaction kinetics is controlled by the diffusion of H2S through the pores of the C a s product layer formed around the lime particle (effective diffisivity between 2.8 x and 5.1 x m2/s), and the kinetics of the sorption of H2S by CaO is relatively insensitive to temperature.
Nomenclature C = concentration of H2S (mol of H2S/m3 of gas) C,, = equilibrium concentration of H2S (mol of HzS/m3 of gas) De = effective diffisivity of H2S in the Cas layer (m2/s) DK = Knudsen diffusivity of H2S (m2/s) D M = diffisivity of H2S in the gas mixture (m2/s) R = effective average radius of 1imeAimestonepellet (m) Rpore= average pore radius in the CaS layer (m) S, = specific surface area of the CaS layer (m2/kg) T = absolute temperature (K) t = time (s) X = fraction of lime converted to Cas (0 < X 1 ) V, = specific void volume in the Cas layer (m3/kg) Greek Letters €0 = porosity of the CaS layer es = molar density of the solid (mol of solidm3 of solid) tl = characteristic particle diffusion time: eBR2/(6D,(CCeq))
(s)
Literature Cited Borgwardt, R. H. Calcination Kinetics and Surface Area of Dispersed Limestone Particles. N C h E J. 1985, 31 (11, 103-
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. Prog. l 9 8 4 , 3 (21, 129-135. Boynton, R. S. Chemistry and Technology of Lime and Limestone, 2nd ed.; John Wiley & Sons, Inc.: New York, 1980; Chapters 6 and 7. Coble, R. L. Sintering Crystalline Solids. I. Intermediate and Final State Diffusion Models. J.Appl. Phys. 1961,32 (51, 787-792. 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. 1. Kinetics of H2S Sorption by Uncalcined Limestone. Znd. Eng. Chem. Res. 1996,34, 23242333. Fenouil, L. A.; Towler, G. P.; Lynn, S. Removal of H2S from Coal Gas Using Limestone: Kinetic Considerations. Znd. Eng. Chem. Res. 1994,33 (2), 265-272. Froment, G. F.; Bischoff, K. B. Transport Processes with Reactions Catalyzed by Solids. In Chemical Reactor Analysis and Design, 2nd ed.; John Wiley & Sons: New York, 1990; pp 148-149. Fuertes, A. B.; Alvarez, D.; Rubiera, F.; Pis, J. J.; Marban, G. Surface Area and Pore Size Changes During Sintering of Calcium Oxide Particles. Chem. Eng. Commun., 1991,109, 7388. Fuertes, A. B.; Alvarez, D.; Rubiera, F.; Pis, J. J.; Marban, G. Simultaneous Calcination and Sintering Model for Limestone Particles Decomposition. Chem. Eng. Res. Des. 1993,71(A),6976. Levenspiel, 0. Chemical Reaction Engineering, 2nd ed.; John Wiley & Sons: New York, 1972; Chapter 12. Reid, R. C.; Prausnitz, J. M.; Poling, B. E. Diffusion Coefficients. In The Properties of Gases and Liquids, 4th ed.; McGraw-Hill Company: New York, 1987; pp 587-597. Towler, G. P. Synthesis and Development of Processes for the Recovery of Sulfur from Acid Gases. Ph.D. Dissertation, University of California at Berkeley, 1992. Towler, G. P.; Lynn, S. Development of a Zero-Emissions Sulfur Recovery Process. 1. Thermochemistry and Reaction Kinetics of Mixtures of H2S and C02 at High Temperature. Ind. Eng. Chem. Res. 1993,32, 2800-2811. Yang, R. T.; Chen, J . M. Kinetics of Desulfurization of Hot Fuel Gases with Calcium Oxide. Reaction Between Carbonyl Sulfide and Calcium Oxide. Environ. Sci. Technol. 1979,13, 549-553.
Received for review July 7 , 1994 Revised manuscript received April 18, 1995 Accepted April 28, 1995@ I39404238
111.
Borgwardt, R. H. Calcium Oxide Sintering in Atmospheres Containing Water and Carbon Dioxide. Znd. Eng. Chem. Res. 1989,28,493-500.
* Abstract published in Advance ACS Abstracts, J u n e 1, 1995.
