PLENARY ACCOUNT Sulfur Dioxide Absorption and Conversion to Sulfur Robert H. Rosenwald. and Peter Urban Corporate Research Center, UOP Inc, Des Plamms, llllmis 60016
Daniel E. Smith A#rCorrectionO!vismn, UOP Inc., Der Plamms, lllmls 60016
Robert € Rosenwald I. is a consultant t o the Corporate
Research Center of UOP Inc., Des Plaines, Ill. He received the Bachelor of Arts degree from North Central College in Naperuille, Ill., in 1932. He was awarded the Master of Science decree in 15I34 and his Doctorate in 1936 both from the Uniuersity of Minnesota at Minneapolis. Dr. Rosenwald joined UOP in 1936 and has worked i n the field of organic chemistry in relation to the treating of motor fuels (inhibitors)and the isomerization and alkylation of isoparaffins. He has done work on the chemistry and use of additives for the improuement and stabilization of petroleum products, rubber, polymers, and fats and oils. He has also conducted investigations of the chemistry of conuersion of sulfur derivatives, such as mercaptans and hydrogen sulfide, as required in refining technology. Dr. Rosenwald is a member of Phi Lambda Upsilon and Sigma X i in addition to holding membership in the American Chemical Society.
Peter Urban is an assistant to the director of process re- ! search and deuelopment i n the Corporate Research Center of UOP Znc., Des Plaines, Ill. He receiued his Bachelor of Science degree from Northwestern liniuerk t y , Evanston, Ill., in 1948 and joined UOPthat same year as a member of thegasoline components research department. I n the twenty-seven years that he has been with the company, Urban has worked on antioxidants, metal deactivators, flue gas desulfurization, pollution control, and uarious problems i n refinery treating processes. Mr. Urban is 'a member of the American Chemical Society. ~
Ind. Eng. Chem.,Prod. Res. Dev., Vof. 16, No. 3, 1977
Daniel B. Smith is manager of aduanced technology at the Air Correction Diuision of UOPZnc., DesPlaines, Ill. He is responsible for investigating and recommending for license all forms of aduanced technology, and directs the internal research and development program for promising technology. He is a graduate of Southern Illinois Uniuersity with a Bachelor of Science degree in chemical engineering. He worked for the State of Illinois, Department of Highways, prior to joining UOP in 1969. At UOP Mr. Smith worked as project engineer until 1974, when he was promoted to manager of system engineering. He was promoted to chief engineer, systems engineering, in 1975, and assumed his present position in 1976. Mr. Smith is a member of the American Institute of Chemical Engineers.
This paper explains the chemistry involved in sulfur dioxide extraction from a flue gas using a sodium carbonatebicarbonate absorption with the subsequent reduction to elemental sulfur. The absorption is carried out at pH 6.2 and in the presence of sodium bisulfide/polysulfide to convert the sulfite to sodium thiosulfate. The formation of thiosulfate controls the oxidation of sulfite to sulfate, an undesirable by-product. The sodium thiosulfate is reduced with hydrogen sulfide, with or without the presence of carbon dioxide, to form elemental sulfur and a polysulfide solution. The unreacted thiosulfate and the polysulfide sulfur are catalytically reduced with hydrogen-carbon monoxide to yield hydrogen sulfide, carbon dioxide, and a carbonate-bicarbonate solution for recycle as absorbent. Chemical factors of concern in the sequence of steps are reviewed.
Introduction The quest for abundant energy sources involves not only consideration of the fuel itself but also the effect on the environment in the utilization of that fuel. Such fuel search is an intricate one, for the entire situation is subject to social acceptability, economic justification, and political approval. In response to this energy search, investigators have pointed out that coal is one of the more promising fuels. In general, coal meets the requirements of availability, ease of usage with efficiency, and cost. However, coal is contaminated with some form of sulfur compounds that yield sulfur oxides on combustion. Consequently, the large-scale use of coal, particularly with the more accessible high sulfur-bearing coals (ca. 4% sulfur), requires a means to avoid sulfur oxide pollution of the atmosphere. Sulfur oxide emissions have also been a problem in the petroleum industry, as crude oil is also contaminated with sulfur compounds. The need to provide emission control technology in an oil refinery has prompted UOP Inc. to develop, over a period of years, a capability and expertise in sulfur chemistry. With petroleum fractions, the removal of sulfur from a fuel prior to combustion can be accomplished to varying extents by normal processing technology, but extensive removal becomes progressively difficult as the fuel (e.g., petroleum residues) becomes heavier. Consequently, an increased emphasis has been placed on sulfur oxide removal from stack gases. About two-thirds of the electrical energy produced in the United States is generated by coal-fired plants. Therefore, development of such processes is of concern to the electric utility industry ( 1 5 ) . The development of a successful flue gas desulfurization process requires the ability to remove sulfur oxides with the production of an environmentally acceptable product. The more desirable scheme is to recover elemental sulfur, an easily handled, saleable product. Furthermore, such a scheme allows regeneration of the absorbent for reuse in the process. This text describes the chemistry of a process that achieves absorption of the sulfur oxides followed by conversion to elemental sulfur. In brief, the process extracts sulfur by use of an aqueous solution of sodium salts to produce oxy-sulfur derivatives, followed by direct reduction to produce elemental sulfur and a solution for reuse in the absorption step. The overall chemical reaction can be summarized as 2S02 + 2(H2 + CO) --* 2S0 + 2(H20 + CO2) For purposes of discussion, the entire process is considered as comprising three individual steps, each of which is briefly described by the following equations and statements. Subsequent discussion will elaborate on each step with emphasis on the chemical reactions involved. Absorption. The sulfur dioxide is absorbed in a sodium carbonate-bicarbonate solution in the presence of a reductant to produce a thiosulfate solution.
SO2 + HzO
+ Na2C03 + NaHC03 reductant Na2S203
The formation of sodium thiosulfate controls the oxidation of the absorbed sulfur dioxide to sulfate.
Table I. Nomenclature System Symbol
Total sulfur, includes all forms of sulfur Elemental sulfur, uncombined. No indication as t o extent of polymerization. Sulfur in polysulfide form, S,S, 2-. Maximum value s x of X is 4. Sulfide sulfur as H2S or salts, e.g., (NH&S or S2NH4SH. Sulfide in polysulfide form, SxS,z-. Value of p is 1. S,2 Sulfur as oxides, and includes salts of sulfurous, S,O, thiosulfuric, sulfuric, and polythionic acids. The ionic charge of the anion is not indicated. SO4*- Sulfur as sulfate. S2O3*- Sulfur as thiosulfate. SO3’Sulfur as sulfite.
Reduction. First Step. The thiosulfate solution is subjected to incomplete reduction with sulfide to produce elemental sulfur and a polysulfide solution.
Second Step. The remaining thiosulfate is completely reduced with hydrogenxarbon monoxide, the polysulfide sulfur (S,) incompletely reduced to form H2S and a sulfide-polysulfide solution with x’ < x . Na2S203 NaS, S H
+ NaSH + Cos2-
HzS+ NaS,, SH
Stripping. The resulting polysulfide solution is stripped to remove H2S for reduction in Step 2, and carbon dioxide to participate in the deposition of sulfur. The remaining solution is recycled to the absorption step. In the following discussion of the individual steps, the nomenclature and concentration values are summarized in Table I, in accordance with a previous publication (21). Sulfur concentrations and charge rates are expressed as moles or weight of sulfur regardless of species. All temperatures are in degrees centigrade. Absorption Sulfur dioxide can be effectively absorbed in solutions with pH -6 or higher to form sulfite-bisulfite solutions. This is because of the acidity (pK = 1.7 X 10-2) of the sulfurous acid which is formed by a hydration step of the sulfur dioxide ( 5 ) . A major obstacle with this procedure is the formation of sulfate by the oxidation of sulfite. Sodium sulfate in solution is very resistant to reduction. The rate of sulfite oxidation is subject to transfer processes in a two-phase system ( I , 7,26). The proposed scheme converts the sulfite to thiosulfate in the absorption step and thereby inhibits sulfite oxidation. This conversion is a reduction that can be carried out using sulfide as a reductant. At pH 6-7, the absorbed sulfur dioxide is present primarily Ind. Eng. Chem., Prod. Res. Dev., Vol. 16, No. 3, 1977
Table 11. Scrubber Operating Conditions Temperature in scrubber
Flue gas Rate SO2 content Charge Exit NO content Oxygen content CO2 content LIQUID PRODUCT FRESH WATER
BOOSTER F A N
3 LIQUID LEVEL C O N T R O L &
FRESH FEED A D D I T I O N O N 1pH C O N T R O L
M A K E UP W A T E R O N D l N S l T Y C O N T R O L
1 IN OF 1 1 IN P A C K I N G
PLEXIGLAS PIPE 10 I N O D
ClRCULLTlNG P U M P
as the bisulfite and is converted by the following equation:
Absorption liquid pH (charged) Rates Fresh feed Product withdrawal Stoichiometric reductant ( S x To form thiosulfate, %
Figure 1. Experimental sulfur dioxide scrubber.
+ 2S0 + 2NaHC03
+ 2NaS2SH + 4NaHC03 7NazS203
+ 5H20 + 2CO2 + 7H20 + 4 c o 2
The conversion of a sulfite solution to thiosulfate with all three reductants was initially studied in the laboratory a t a temperature of 55 OC. Elemental sulfur worked best in the presence of a surfactant (16). The most reactive reagent was the S, of a polysulfide; sulfur in this form was consumed within 30 s. The effectiveness of the thiosulfate concept was established by absorption studies with a flue gas scrubber. In this evaluation, several scrubbers were investigated. The data presented here were obtained with equipment and conditions as follows. The absorption studies were run in a 10-in. 0.d. Plexiglas pipe with a supported 4-in. bed of berl saddles and spray nozzles to ensure even distribution of the solution. A diagram of the unit is shown in Figure 1. The unit was automated for withdrawal of absorption solution on level control, the addition of fresh solution by pH, and the addition of fresh water (to compensate for evaporation) by density control. The flue gas was supplied by a furnace fired with oil containing sulfur. Provisions were made for the metered additions of sulfur dioxide and nitric oxide (NO), if desired. Variations in sulfur dioxide content in the flue gas were obtained either by addi196
0.428 0.214 0.00 0.903 1.805 1.14 1.15 0.86 100.00
+ 2NaSSH + 2NaHC03
1900 mL/h 2500 mL/h S2-) Varied
Ind. Eng. Chem., Prod. Res. Dev., Vol. 16, No. 3, 1977
The use of polysulfide solution accomplishes reduction with both sulfide and sulfur (Sx). The equation defining amounts of reactants depends on the value of X in the polysulfide (NaS,SH). For example, with an X value of 1and 2, the following equations describe the complete conversion (a reductant stoichiometry of 100%). 6NaHS03
Charge rates, mol/h
Sulfur dioxide Solution constituents Na2S04 NazSz03 NanSOa NaHC03 Na2C03 NazS
Elemental sulfur likewise reduces bisulfite 2NaHS03
3000 ppm 10-50 ppm 800 ppm 10 f 2% 10 f 2%
Table 111. Scrubber Operation
@ I N L E T SPRAY 5500 M L M I N U T E @ INLET FOR 5 0 2 A N D N O I N J E C T I O N A S DESIRED
0.65 5.98 0.97 0.00 0.00
tions of sulfur compounds to the fuel or by injections of sulfur dioxide; both procedures gave identical results with the latter generally more convenient. The solution charged (fresh feed) was prepared from commercial chemicals, giving a recycled absorbent with pH -6.2. The composition is similar to that expected from an absorbent as would be regenerated in an integrated plant. Typical flow rates of sulfur dioxide, composition of the fresh absorbent, and products as withdrawn are summarized in Tables I1 and 111. The overall equation for absorption and conversion may be represented as follows with 100% reductant stoichiometry and an X value of 1.0: 6S02
+ 8NaHC03 + 2NaSSH
+ 5H20 + 8C02
Operation of the scrubber with the flow rates as presented in Tables I1 and I11 gave a reductant stoichiometry of 70%. Other values were obtained by proper control of charge rates of the reactants involved. Satisfactory stable operation and weight balances of solutions and constituents were accomplished by automatic regulation of the multiple rates involved. The absorption capacity for sulfur dioxide, over 1lb/gal, was a typical value for the solutions studied. The operation described was an efficient process for SO2 removal; the SOz concentration in the exit gas was in the range of 10-50 ppm as indicated in Table 11. A series of runs was made with varying amounts of reductant. As shown in Figure 2, the formation of sulfate as percent of sulfur dioxide charged depended on the reductant stoichiometry. With no provisions for thiosulfate as intermediate (zero stoichiometry),sulfate formation was 25% to 40%. With a stoichiometry of loo%, sulfate formation decreased to 2%. The concept of thiosulfate as an intermediate to control sulfite oxidation is considered established by these data.
REDUCTANT STOICHIOMETRY, 40
N O CONCENTRATION IPPMI
~i~~~~2. Effect of reductant stoichiometry or, sulfate formation, Nitric oxide concentration, 800 ppm.
Figure 3. Effect of nitric oxide concentration on sulfate production; 70% stoichiometric sulfur.
In the citrate process, thiosulfate had been proposed as the primary factor for the efficiency of sulfur dioxide extraction and the inhibition of oxidation (20). It is believed that the results reported herein do not involve this particular thiosulfate effect. Thiosulfate not only serves as a protection against sulfite oxidation, but also lends itself well to reduction reactions. The reductions were carried out with ease and efficiency because of the ease of handling sodium thiosulfate. In our investigations on the absorption variables, a number of observations were made. Oxygen Concentration. The formation of sulfate during absorption is a result of the reaction of sulfites with the oxygen in the flue gas. In the uninhibited absorption (no thiosulfate as intermediate), the extent of sulfate formation was related to the oxygen concentration; low oxygen concentrations gave low sulfate formation. With thiosulfate as the intermediate, the oxygen concentration was not a critical factor; similar results were obtained with either 4%or 10%oxygen concentrations. In fact, a certain amount of sulfate was formed regardless of oxygen concentration. Sulfate formation up to 2% was obtained in runs made with simulated flue gas streams containing no oxygen. Sulfate formation in the absence of oxygen can be attributed to some form of disproportionation. Possible reactions, which may proceed at a low but appreciable rate, are those of bisulfite
a reductant stoichiometry of 70%. Sulfate formation was reduced but not eliminated by low nitric oxide concentrations. There was no evidence that NO or its oxidation products were absorbed in the aqueous phase; the nitrogen contents in the absorbent were low (less than 25 ppm). The Reductant. According to data in Figure 2, sulfate formation can be repressed to about 2% by use of the reductant at 10096 stoichiometry. This is not practical because with that amount of polysulfide, hydrogen sulfide was released in small amounts. Experience established that 70% stoichiometry did not lead to sulfide release and allowed a safety factor. The reductant of preference was found to be the polysulfide (NaS,SH) rather than the sulfide (NaSH). More effective sulfate control was realized with the polysulfide. Metals. The oxidation of sulfite solutions is known to be catalyzed by the presence of salts of such metals as iron (9) copper ( 3 ) manganese (14), and cobalt ( 1 7 ) . No effect was observed with an absorbing solution containing the usual promoting metals (Fe, Mn, Cr, Ni, V, etc.). pH. Variations in pH of about one unit did not affect the formation of sulfate. The inherent instability of thiosulfate solutions to acids did not allow application of the thiosulfate concept a t low pH values of below about 5. Polythionates. The reaction of sulfur dioxide with hydrogen sulfide is known to produce polythionic acids and their salts according to the following equation where X may have values from 2 to 6 and even higher ( 2 ) .
of thiosulfate (12,22)
S ~ 0 3 ~ -H20
The reaction is commonly called the Wackenroder reaction +
and of polythionates ( 1 3 ) . These data indicate that some sulfate, at least 2%, will be formed regardless of absorption conditions. Since some sodium sulfate will be formed, provisions must be made for removal to avoid a build-up. This can be accomplished by cooling a portion of the absorption liquid and crystallizing the sodium salt. T o avoid interference with sodium sulfite, provisions must be made to complete the reaction to sodium thiosulfate. This can be carried out by bringing in a sulfide stream, either NaSH or H2S, after the absorbent leaves the scrubber. Nitric Oxide. The data in Figure 2 were obtained with metered amounts of nitric oxide (NO) to give a concentration of 800 ppm. It was observed that the extent of sulfate formation was dependent on the nitric oxide concentration. Figure 3 presents data on this relationship in a series of 6 runs with
(18).Polythionates have been reported to be formed in solutions containing sulfite, thiosulfate, and sulfide ( 4 , 8, 27).
Upon examination of the absorbent to determine if these compounds were formed, the presence of di- and tripolythionates was detected with the trithionate being the predominant species. Concentrations were indefinite because of their decomposition with time to yield various products, including sulfate. However, concentrations with normal operation were low, often traces, and subsequent reducing conditions were found to decompose those polythionates. Reduction The next step in the processing scheme is the reduction of the thiosulfate, as formed in the scrubber and subsequent treatment, to form elemental sulfur and eventually a solution suitable for recycle absorbent. These experiments on reduction were run under continuous flow experiments. The gases and liquids as indicated were charged to heated reactor tubes Ind. Eng. Chem., Prod. Res. Dev., Vol. 16, No. 3, 1977
Table IV. Reduction of NazS203 and Polysulfide Solutions (1st Stage, Reductant HzS; 2nd Stage, Reductant Hz/CO) Reduction stage Reaction conditions Temperature, "C Pressure, psig LHSV Catalyst
175 750 1.34 Carbon
215 750 1.4 Metal Sulfide
200 750 1.1 Metal Sulfide
Flow rates mol/h Charge Gases HzS
coz H2 co
0 0 0
0.273 0.273 0.032 0
0.347 0 0 0
0.141 0.074 0.120 0.289
with provisions to control temperature, pressure, and rates. The solutions in all cases were aqueous and generally prepared from commercial chemicals. However, in a sequential series of reaction, all have been run with solutions and gases as produced from a preceding step(s). The reduction step will be considered as occurring in two successive reactions; the first reduction involves mainly the thiosulfate conversion to elemental sulfur followed by a more complete reduction. Reduction, Step I. Since sulfide solutions were needed as a component of the scrubbing solution, hydrogen sulfide was selected as a logical reductant. A prolonged investigation was made of the reduction as to conditions, yields, catalysts, etc. It was essential from engineering considerations to run this reaction at a temperature of 125 "C or higher in order to yield sulfur in a liquid form. From this study two reaction schemes were developed to carry out this step. Sulfur could not be satisfactorily produced by the following direct interaction as follows: NazSz03
+ 2H2S s 4S0 + 2NaOH + HzO
It appears that the sodium hydroxide as formed caused the well known sulfur-caustic inteiaction to occur, or the reverse of the desired reaction. This is in contrast to the ammonia system in which this was not of consequence (21). It was observed that the reduction with an excess of hydrogen sulfide (4 moles to one of sodium thiosulfate) gave a polysulfide solution which by proper temperature control liberated sulfur. Na&03
+ 4H2S * 2NaSzSH + 3H20
2NaSZSH F? 2NaSSH + 2So The more favorable conditions for the reduction were a temperature of 175 "C and a pressure of 750 psig. Under these conditions, a soluble polysulfide was formed to which can be assigned the formula NaS2SH. A decrease in temperature to 125 "C a t 750 psig prompted a deposition of sulfur and formation of a polysulfide NaSSH. Thus, 2 atoms of sulfur in the thiosulfate were converted to sulfur. These conditions gave, as summarized in Table IV, a high thiosulfate consumption with about a 50 wt % yield of sulfur based on thiosulfate charged. Complete conversions were not realized since the system is in equilibrium. 198
Ind. Eng. Chem., Prod. Res. Dev.. Vol. 16, No. 3, 1977
0.106 0.115 0.161 0
0.001 0.262 0.070 0
Thiosulfate converted, % Soyield based on NazSz03 charged, %
0.176 0.037 0.279 0.186
Liquids 0.072 0.361 0.572 0.248
The above NazS203-HzS reaction was conducted by passage of a Na&03 solution (9 wt % sulfur as NaZS203) with H2S. The reactor effluent was cooled to 125 "C under pressure and the two liquid phases (aqueous and sulfur) were separated. In the first of two runs presented, the reaction was run in the presence of a granular carbon while in the second run a group 8 metal containing solid bed catalyst was used. The Na&03-H2S reaction was subject to catalysis, with granular carbon showing a moderate promoting effect that declined within one week on stream. The exact species involved and the course of the reactions concern the chemistry of polysulfides (19,23). The temperature effect on sulfur release has been suggested in the literature (6, 10,25). This work clearly demonstrated that there is a temperature effect on polysulfide composition and its ability to retain or release sulfur. The effect of the excess sulfide can be attributed to the difference between attack on sulfur of the OH- anion in contrast to the bisulfide (SH-) or polysulfide (S,SH-) anion. The concentrations of the SH- anion, as well as the OHanion, are both major factors in defining the composition of the polysulfide (11). The other reaction scheme involves the use of carbon dioxide with a deficiency of hydrogen sulfide. The passage of one mole of sodium thiosulfate with about 1.5 mol of hydrogen sulfide and 1.5 mol of carbon dioxide at a temperature of 215 "C and a pressure of 750 psig gave a 60% conversion of the thiosulfate and over 80 wt % yield of sulfur based on the thiosulfate charged (Table IV). The favorable effect of carbon dioxide can be explained by its buffering capability, thereby decreasing the basicity of the system. As with the temperature effect, the use of carbon dioxide can be considered an extension of the basicity factor. Reduction, Step 2. The Na&03-HzS reduction produced a solution containing unreacted thiosulfate and a polysulfide with an X value in the range of 1 to 2. The purpose of the second stage reduction is the conversion to sulfide of the remaining thiosulfate and a large portion of the s,. This was accomplished by the use of an H&O reductant in 60/40 mole ratio. Carbon monoxide was used not only as an excellent reductant but as a source of anions to carry basicity (sodium carbonates) back to the scrubber. The reductions involved are:
Table V. Stripping a Sodium Sulfide Solution with and without Carbon Dioxide (Solution Charge Rate ca. 200 -1 /I.\ lllrr, a,,
Stripping Conditions 126 125 Reboiler temperature, "C Pressure, psig 16 17 Charge Rates, mol/h NaHC03 0.244 0.243 S20.309 0.316 0 0.011 NaOH Hz0 10.979 10.934 0 0.25 COz for stripping Mole ratio COp/S20.78 Charge Overhead, % Water 35 33 H2S 34 64 coz 23 47
+ + + + + 4Hz
0.240 0.243 0.310 0.286 0.008 0.032 10.959 11.076 0.375 0.50 1.2 1.75 31
32 86 58
Table VI. High Pressure Stripping of Polysulfide Solution Striming conditions ~ b i u i temperature, h "C Reboiler temperature, "C Pressure, psig Solution charge rate, mL/h Charge Rates, mol/h S2-
COZ for Stripping Mole ratio C02/S2Product, S, Charged, % Sz0s2- (net) SO
S, remaining S2- Charged in overhead, %
130 125 750 200
200 125 750 200
0.18 0.38 0.02 0.67 3.7 3.5 64 22 61
0.18 0.38 0.02 0.67 3.7 5.9 62 24 64
"C and 90 psig) upon evaporation of 35% of the water, gave in the overhead about one-third of the sulfide and one-quarter Naps203 4CO Hp0 2NaSH 4co2 of the carbonate. This was considered inadequate. The use of carbon dioxide as a stripping agent, countercurrent to the NaS, SH XHp NaSH XHpS solution, increased the efficiency so that up to 86% of the sulfide and 58% of the carbonate were liberated. NaS,SH X C O X H z O NaSH XHzS XCO2 One concludes that by the use of carbon dioxide, a sulfide It was found that the catalyzed reductions did occur as solution can be stripped to liberate sulfide. shown in Table IV. From the numerous variables studied, the The stripping of a polysulfide solution resulted in the following conditions were selected as suitable: temperature conversion or back-hydrolysis of the S, to thiosulfate. The 200 " C , pressure 750 psig, and a residence time of about 1 h action of carbon dioxide on a polysulfide solution at temperwith a metal sulfide catalyst. The general equation can be atures below 50 "C liberated SO in good yields. However, upon expressed as follows: attempting to liberate SO at higher temperature (120 "C) under conditions described for the sulfide solution, thiosulfate l.O(SI.ISH-) l.l(S2032-) + 4.4(CO H2) 2.5(So&H-) Y H ~ S formation was appreciable due to S, disproportionation. Hydrogen sulfide likewise liberated So at 25 "C, but at The polysulfide was collected in the aqueous solution while higher temperatures, no sulfur separated. The problem was the remaining hydrogen sulfide (yH2S) was released in the gas not thiosulfate formation but high So solubilities in the sulfide phase. solution. The above conversions involve a number of factors, a few It was discovered that polysulfide solutions can be stripped of which merit comment. The thiosulfate is more readily reby raising the pressure to the range of 600-800 psig as sumduced than the s,, particularly with carbon monoxide. It is marized in Table VI. The use of pressure with carbon dioxide not known whether the reductants directly attack the thiodid not eliminate thiosulfate formation but did reduce it to sulfate or if the hydrogen sulfide initiates the attack followed a tolerable amount. The stripping of a polysulfide, ca. 1.0 M, by involvement of the hydrogen-carbon monoxide. The latter with an X value of about 2 at 750 psig and at 130 or 200 "C route is suggested by the fact that the rate is markedly incolumn temperature did strip out over 60% of the sulfide. creased by the incorporation of some sulfide in the feed. The Thiosulfate formation was calculated as less than 6%. composition of the product was greatly dependent on the The data in Table VI1 were obtained by the high-pressure amount of reductant. Limited amounts of reductant(s) constripping of an effluent as prepared by a second stage reducverted the thiosulfate first with little conversion of the s,. tion. Stripping was observed to proceed in the presence of Excess reductant attacked the S, and decreased the X value sodium thiosulfate with improved recoveries of So. Conseas severity of processing and the amount of reductant were quently, the operation was run to give incomplete thiosulfate increased. An excess of reductant, 120-15W0 of the theoretical, reduction with an X value between 1 and 2. The stripping of was required for large conversion of S,. this solution with considerable carbon dioxide (ca. 11mol/mol I t was by proper adjustment of these and related factors of sulfide) gave high sulfide removal from the solution (over that the selected and controlled reduction to yield H2S and 80%)and Sorecoveries evidently depending on the thiosulfate a polysulfide solution with a low S, value was achieved. This content of the solution. solution was then in condition for further processing. An Integrated P l a n t Stripping The individual steps here described can be fitted together The third and final processing step is that of stripping. The into an integrated process. Such a general and quantitative second stage reduction yields a solution from which hydrogen scheme is illustrated in Figure 4, disregarding side reactions sulfide and carbon dioxide are to be obtained for recycle to the such as sulfate formation. Such features as rates, conversions, first stage reduction step. The remaining solution is to be used process conditions, etc., are not indicated. The function of as carbonate-sulfide absorbent for the scrubber. The following each of the 5 units is self-evident. stripping experiments were run in a packed stainless steel The process here described has been run in an integrated pipe. plant to obtain engineering information. The operation conThe first consideration was the direct stripping by boiling firmed the chemistry presented for the individual steps as of a sulfide-bicarbonate solution. As summarized in Table V, developed from laboratory experimentation. a carbonate-sulfide solution (prepared in an autoclave at 120 From an engineering point of view, it is desirable to simplify Na2Sp03
Ind. Eng. Chem., Prod. Res. Dev., Vol. 16, No. 3, 1977
Hp. CO H V D R O G E N SULFIDE
Table VII. Stripping of Effluent from 2nd Stage Reduction (Solution Charge Rate ca. 100 mL/h) 2nd Stage Reduction
Mol/mol S as S ~ 0 3 (CO ~ - + H2) Reducer Effluent, mol/h
SCRUBBER SULFITE CONVERTER ISTSTAGEREDUCER
S20& Converted, %
Stripper Conditions Column temperature, “C Pressure, psig CO2 for stripping, mol/h Mole ratio, C02/S2Liquid effluent, mol/h S2-
CO. Hz. Hz S. CO2
200 200 680 680 1.48 1.47 11.6 11.7 0.024 0.036 0.098 0.127
0.45 0.382 1.16
0.128 0.122 0.218 0.177 0.073 0.128 0.163 0.157 81 66
Figure 4. Integrated plant schematic flow.
0.45 0.387 1.31
1 2 3 -
Temperature, “C Pressure, psig Charge rates, mol/h S2-
S2- Charged in overhead, % S, Recovered as SO, %
0.014 0.019 0.126 0.130
Figure 5. Two unit system.
Table VIII. Two Unit System. Reduction of Sodium Thiosulfate in Unit 1; Stripping in Unit 2
the 5-component plant with the visionary hope to attain an “ultimate process” (24). Toward this end, a two unit system-one reactor and one stripper-was devised as a simplification. Such a system eliminated the necessity of concentrated hydrogen sulfide, a hazardous chemical, and reduced the number of flow schemes to be regulated. This elimination is suggested by the fact that the following equations, which summarize the two fundamental reactions involved, do not show the necessity of hydrogen sulfide as such.
Reaction Conditions Temperature, OC Reducer Stripper Settler Pressure, psig Reducer Stripper Reducer catalyst Stripper packing Thiosulfate solution, LHSV
6S02 + 2NaSSH + 8NaHC03 5Na2S203
+ 5H20 + 12CO
125 125 750 680
Metal sulfide Metal helices ca. 1.1 Product
+ 5H20 + 8 C 0 2
Rates, mol/h s2032-
+ 2NaSSH + 8NaHC03 + 4 c o 2
This simplification was accomplished by the reaction scheme illustrated in Figure 5 with experimental results shown in Table VIII. This is one of the possible arrangements that, by the control of two main streams (reductant and carbon dioxide), produced sulfur and a product solution for recycle without the concentration or isolation of hydrogen sulfide. One observes that the reductant gases assist as a stripping agent to remove hydrogen sulfide and carbon dioxide from the liquid stream. The carbon dioxide injection was a t a high temperature (200 “C) a t the reducer exit. Most of this carbon dioxide was recovered in the exit gas and couid be recycled. The yield of sulfur represented about 50% of the thiosulfate converted. Higher conversions were obtained by the use of carbon monoxide in place of the hydrogen-carbon monoxide mixture, because of its greater reactivity as a reductant. The Ammonia System The work previously described was exclusively with sodium salts. The same capability and chemistry have been shown to apply to the system with ammonium salts. The volatility of ammonia and certain derivatives necessitated some processing changes, but the overall operations were similar. The main chemical difference observed was that the ammonium salts 200
Ind. Eng. Chem., Prod. Res. Dev., Vol. 16, No. 3, 1977
Charge 0.390 0.384
0.394 Converted, %
0.118 0.083 0.042 0.139
were more reactive; lower temperatures could be used to carry out the reactions. Conclusion This account represents the chemistry in aqueous solutions of sulfur and its compounds as encountered in a project on sulfur dioxide control. In the more simple terms, the chemistry is merely a chronicle of the valence and acidity of the sulfur species involved. However, the many valence states and their numerous combinations create a complex system subject to chemical equilibria. Accordingly, a process designed to yield sulfur, which is in an intermediate valence state, will likely involve under- and over-reacted products. In view of this situation, the experimental approach has been used to evaluate the chemical aspects of a possible reaction scheme. The process and chemistry here described resulted from the accu-
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Received for revieu; March 18, 1977 Accepted June 6,1977
Presented a t the 173rd National Meeting of the American Chemical Society, San Francisco, Calif., Aug 29-Sept 3, 1976.
Ind. Eng. Chem., Prod. Res. Dev., Vol. 16, No. 3, 1977