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Ind. Eng. Chem. Res. 1997, 36, 4236-4242
Recovery of Calcium Carbonate and Hydrogen Sulfide from Waste Calcium Sulfide Mark W. Brooks† and Scott Lynn* Department of Chemical Engineering and Lawrence Berkeley National Laboratory, University of California, Berkeley, California 94720-1462
The removal of H2S from hot coal gas using a limestone-based sorbent generates a large quantity of calcium sulfide waste (e.g., 67 tons/day from the gasification of 1000 tons/day of 3% S coal) that cannot be landfilled because of potential H2S evolution and sulfide leaching. In the process proposed here, the CaS is dissolved by reaction with H2S complexed with aqueous methyldiethanolamine (MDEA) or other alkanolamines. In the second step, the highly soluble Ca(HS)2 reacts readily with CO2, also complexed with aqueous MDEA, to precipitate pure CaCO3 of uniform crystal size and to form MDEA-complexed H2S in solution. The solution from step 2 is splitshalf is recycled to step 1 and half is sent to a stripper to recover H2S and from there to a column where the needed CO2 is absorbed from a combustion gas. Introduction and Previous Work Currently, most of the coal consumed in power generation is burned to produce high-pressure steam, which is expanded through a steam turbine system to drive electric generators. The last few decades have seen intense research on cleaner and more efficient coalfired power generation. The technique of gasifying coal prior to combustion has received a great deal of this research attention. A benefit of coal gasification at high temperature and pressure is the greater ease of sulfur removal. The sulfur is in the form of H2S and can be removed by a solid sorbent with relatively little cooling of the gas. By contrast, following combustion the sulfur is in the form of SO2, is greatly diluted, and is at near-atmospheric pressure. This combination increases cleaning costs. Of the many possible solid sorbents, the least expensive is limestone. Fenouil and Lynn (1995a,b) found that limestone, which is typically about 97 wt % CaCO3, can reduce H2S to a level of 100-200 ppm, well below that mandated by the Clean Air Act, with conversion of 90+% of the calcium to CaS:
CaO + H2S f CaS + H2O
These characteristics make limestone an attractive sorbent for H2S removal from coal gas. However, CaS cannot be placed in a landfill directly because of the danger of H2S evolution or leaching of the sulfide ions into the ground water. Practical possibilities for CaS treatment include converting the sulfide back to CaCO3 or CaO, thus regenerating the H2S sorbent, and direct oxidation of the sulfide to sulfate, which would allow safe disposal of the waste. Direct oxidation of the CaS with air produces CaO by the following reaction:
2CaS + 3O2 f 2CaO + 2SO2
Unfortunately, if sulfate formation is to be prevented, the reaction must be conducted above 1400 °C (Schwerdtfeger and Barin, 1993). However, after five cycles their CaO retained only 15% of its initial H2Ssorption capacity, probably as a result of sintering. Keairns et al. (1974) proposed reversing the sorption reaction at a lower temperature:
CaS + H2O + CO2 f CaCO3 + H2S
(1)
The sorption is done in a separate step, just ahead of a combustion turbine and within 25 °C of the calcination temperature, which occurs at 800-900 °C and depends on pressure and CO2 content. (Limestone is frequently mixed with the coal in the gasifier, but there is a significant loss in both calcium conversion and H2S reduction and a corresponding increase in ash volume.) Limestone is much cheaper than zinc-based sorbents (even with allowance for recycling the latter). Moreover, its optimal temperature being 200-300 °C higher allows an improved thermal efficiency of power generation. Even when the lower H2S level obtainable with zincbased sorbents is deemed essential (10-20 ppm), bulk desulfurization with limestone followed by polishing with a zinc-based sorbent may prove to be economically attractive. * Author to whom correspondence is addressed. Telephone: (510) 642-1634. FAX: (510) 642-4778. E-mail: Lynn@ cchem.berkeley.edu. † Present address: Exxon Research and Engineering, P.O. Box 101, Florham Park, NJ 07932. S0888-5885(97)00151-6 CCC: $14.00
(2)
(3)
However, they reported that, after 21 cycles, sintering reduced the conversion of CaCO3 to CaS from 69% to 13%. The argument could be made that the cost of limestone is so small there is no need to regenerate the sorbent. This may be the case, but CaS is considered hazardous and must be rendered inert before disposal. Schwerdtfeger and Barin (1993) studied oxidation to CaSO4 but concluded that it was not possible to convert CaS quantitatively to CaSO4 in a single-step process. They found that at elevated temperatures (>1100 °C) the sulfation reaction was fast but that some of the CaS was converted to CaO with the release of SO2. Evolution of SO2 was avoided at lower temperatures, but the oxidation was too slow for an industrial process. Van der Ham et al. (1996) proposed a complex reaction sequence that featured reaction (1) followed by
CaS + 2SO2 h CaSO4 + S2 and © 1997 American Chemical Society
(4)
Ind. Eng. Chem. Res., Vol. 36, No. 10, 1997 4237
CaS + 3CaSO4 h 4CaO + 4SO2
(5)
in which sulfidation (reaction (1)) is best carried out near 850 °C, reaction (4) near 700 °C, and reaction (5) at 1100-1200 °C. They postulated that reaction (5) occurs in the eutectic melt formed by CaS and CaSO4 above 1100 °C. Although such a mechanism would normally lead to sintering, van der Ham et al. reported a drop in capacity of only about 35% after eight cycles. However, their equipment limited the temperature for reaction (5) to 980 °C, so the loss in capacity may thereby have been minimized. All of the proposed regeneration and disposal schemes described above produce dilute streams of SO2 requiring further processing and/or encounter sintering problems that hinder further cycles of regeneration. The production of fresh CaCO3 crystals from an aqueous slurry of CaS would eliminate both problems. As early as 1869, Bechamp studied reaction (3) between CO2 and solid CaS particles suspended in water (Gmelins Handbuch der Anorganishen Chemie, 1957). However, the reaction was incomplete due to the formation of a CaCO3 layer around the CaS. Neither CO2 nor CaS is very soluble in water. Calcium sulfide has a solubility of less than 1 g/L at 25 °C, and CO2 at 1 atm of partial pressure has a solubility of only 1.7 g/L at 20 °C (Perry and Green, 1984). Reaction (3) occurs between the dissolved CO2 and solid CaS at the liquid-solid interface and is limited by how fast the CO2 can dissolve and diffuse. The production of CaCO3 at this interface eventually encapsulates the remaining CaS, preventing further reaction. Biswas et al. (1976) achieved 98% carbonation of 150µm CaS particles in an aqueous slurry by controlling the CO2 flow, bubble size, and slurry depth. By using roughly 3 times the stoichiometric amount of CO2, they produced a gas containing about 33% H2S, just high enough for sulfur recovery in a Claus plant. They proposed that a series of reactions occurs: reaction (3) followed by
CaS + H2S f Ca(HS)2
(6)
Ca(HS)2 + CO2 + H2O f CaCO3 + 2H2S
(7)
and
Thus, CaS reacts initially with CO2, with the formation of aqueous H2S and the precipitation of CaCO3. The aqueous H2S that does not escape into the gas phase reacts with more CaS to form the highly soluble Ca(HS)2, which can then react in solution with CO2 to precipitate CaCO3 and produce more H2S, perpetuating the reaction. The rate at which Ca(HS)2 and CaCO3 are formed is dictated by how fast the CO2 can dissolve and diffuse to a reaction site and also by the concentration of aqueous H2S that can be maintained. A process in which the aqueous solubilities of both CO2 and H2S were significantly increased would greatly improve the rates of reactions (3), (6), and (7), the utilization of the CO2, and the concentration of the product H2S gas. In the research presented below, aqueous methyldiethanolamine (MDEA) is used to increase the solubilities of CO2 and H2S by forming the corresponding substituted ammonium complexes. Since alkanolamines are weaker bases than CaS, solubilizing CO2 and H2S in this way does not inhibit their reactions with CaS and Ca(HS)2. The reactions are as follows:
MDEA‚H2S + CaS f MDEA + Ca(HS)2
(8)
MDEA‚H2CO3 + MDEA + Ca(HS)2 f 2MDEA‚H2S + CaCO3 (9) In reaction (8), MDEA has complexed H2S, keeping it in solution and readily available for reaction with solid CaS. The products of this liquid-solid reaction are free MDEA and a solution of Ca(HS)2. (Calcium hydrosulfide is very soluble in water, 248.7 g/L at 20 °C (Seidell, 1940).) Carbon dioxide, bound in solution by MDEA as the bicarbonate, then reacts in a separate step with the aqueous Ca(HS)2 to precipitate CaCO3 and produce H2S, which is complexed by the free MDEA. A portion of the MDEA‚H2S stream is used to produce more Ca(HS)2 via reaction (8), and the remaining portion is stripped to produce a concentrated H2S gas, which may be converted to elemental sulfur by known technology. The stripped MDEA solution is reloaded with CO2 in an absorber, with a combustion gas as the CO2 source, and recycled. This system has several advantages compared to other CaS-regeneration techniques: (1) “Fresh” CaCO3 is obtained after each regeneration cycle. This eliminates crystal-structure problems such as sintering and attrition that are experienced in many high-temperature sorption/regeneration cycles. Pure CaCO3 has other uses as a product but is also a suitable material for landfill. (2) A concentrated stream of H2S is generated from the stripping operation, which facilitates its conversion to elemental sulfur. Utilization of CO2 is correspondingly high. (3) The higher H2S and CO2 solubilities increase reaction rates and decrease reactor volumes. (4) The formation of the highly soluble Ca(HS)2 eliminates CaS encapsulation during carbonation, resulting in nearly stoichiometric conversion of CaS to CaCO3. The goal of this research was to study the critical steps in this method for converting CaS into CaCO3 and H2S. The effect of CaS particle size and prior CaS oxidation on the kinetics of Ca(HS)2 generation, reaction (8), was determined. Reaction (9) was studied by determining the effects of MDEA‚H2CO3 and Ca(HS)2 concentrations on the crystal-size distribution of the precipitated crystals and the amount of excess CO2 required for complete Ca(HS)2 conversion. The data were then used to design a complete process for the production of an enriched H2S stream and a relatively pure (>95%) calcium carbonate product from waste CaS. Experimental Equipment and Procedures The limestone used in the experiments described below was provided by Great Lakes Calcium Corp., Green Bay, WI; Table 1 shows the average chemical composition. The MDEA was supplied by Union Carbide and had a purity of 99 wt %, with the balance being water. The CaS was obtained from three sources. One batch was made from limestone by the procedure described below. The second was a sample of technicalgrade CaS supplied by Fisher Scientific Co. that had been exposed to air periodically over a period of several years. The third was a fresh sample of reagent-grade (“Alpha”) CaS supplied by Johnson Mathey Co. Table 2 gives the chemical compositions of these calcium sulfide samples.
4238 Ind. Eng. Chem. Res., Vol. 36, No. 10, 1997 Table 1. Limestone Composition (wt %) (Great Lakes Calcium Corp., Green Bay, WI) CaCO3
MgCO3
SiO2
Fe2O3
Al2O3
S
97.8
1.63
0.28
0.15
0.13
0.01
Table 2. Calcium Sulfide Compositions (wt %) and Particle Sizes source Alfa Towler reactor Fisher
CaS CaSO4 CaCO3 MgCO3 SiO2 Fe2O3 Al2O3 size (µm) 98 95 95 74
2 3 3 12
