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Reactions between Hydrogen Sulfide and Sulfuric Acid: A Novel Process for Sulfur Removal and Recovery Qinglin Zhang, Ivo G. Dalla Lana, Karl T. Chuang,† and Hui Wang Department of Chemical & Materials Engineering, University of Alberta, Edmonton, Alberta T6G 2G6, Canada
Reactions between hydrogen sulfide and sulfuric acid were investigated in a laboratory batch reactor. At 120 °C, sulfur is a visible product of reactions between H2S and H2SO4. The amount of SO2 formed in the reaction depends on the sulfuric acid concentration and the reaction temperature. The overall reaction was observed to be exothermic. Based on the experimental results and thermodynamic analysis, the following two independent reactions were identified: (i) H2S + H2SO4 ) S + SO2 + 2H2O and (ii) 2H2S + SO2 ) 3S + 2H2O. The experimental results can be well explained according to the above-proposed reaction scheme. A novel process is suggested for hydrogen sulfide removal and sulfur recovery. Introduction Although the reaction between H2S and H2SO4 has been reported as early as 1858,1-3 very little information is available concerning its detailed reaction chemistry. It was observed that elemental sulfur, sulfur dioxide, and water are the products of the interaction between hydrogen sulfide and concentrated sulfuric acid.4 Various possible reactions have been postulated to explain how these products are formed; however, the independent reactions and the reaction chemistry involved in the H2S-H2SO4 system were not established. In particular, no data are available in the literature regarding the influences of temperature and acid concentration on the product distribution. Actually, this reaction system is very important in practice, for example, in some hydrometallurgical metal extraction processes.5,6 In these processes, the leaching of sulfide minerals in aqueous sulfuric acid involves the formation of H2S in the sulfuric acid medium. The knowledge of reactions between H2S and H2SO4 or SO2 in a sulfuric acid medium is very important to minimize pollution from the metal extraction processes. The reaction of H2S and H2SO4 can also be used to regenerate the active carbon used in a sulfur dioxide adsorption and recovery process.7 We have also observed the formation of sulfur, sulfur dioxide, and water while carrying out acid gas dehydration by concentrated sulfuric acid. Because of the possibility of removing hydrogen sulfide and recovering sulfur through the interaction between hydrogen sulfide and concentrated sulfuric acid, this study was carried out to clarify the reaction chemistry of this system. To eliminate the difficulties from sulfur being solidified in the reaction vessel, our studies focused on the temperature range above 120 °C, so that the sulfur was maintained in the liquid state (the melting point of sulfur is 119.3 °C). The sulfur dioxide being produced was measured at various reaction conditions in a batch reactor. The stoichiometry encountered could not be reconciled with use of the single reaction equation * To whom all correspondence should be addressed. Tel: (780) 492 4676. Fax: (780) 492 2881. E-mail: karlt.chuang@ ualberta.ca.
relating hydrogen sulfide consumption solely to oxidation by sulfuric acid.
H2S + H2SO4 f S + SO2 + H2O To reconcile the experimental results, additional reactions were proposed and tested to establish their relevance. Based on these test results, the independent reactions in the H2S-H2SO4 system were established. Conditions for complete conversion of H2S without SO2 emission were identified. A novel process for H2S removal is proposed. Thermodynamic Analysis To estimate the conditions under which various components present in the reaction system may coexist thermodynamically, a thermodynamic analysis was carried out based on the thermodynamic properties and free-energy minimization using a commercially available software called Outokumpu HSC Chemistrty for Windows.8 The thermodynamic calculations examined equilibria at atomospheric pressure and temperature ranging from 0 to 200 °C and sulfuric acid concentration up to 98 wt %. The possible reactions between H2S and H2SO4 proposed in the literature5,7,9 are summarized as follows:
S + H2SO4(a) ) 3SO2(g) + H2O(l) ∆H120 °C ) 417.26 kJ/mol (1) H2S(g) + 3H2SO4(a) ) 4SO2(g) + 4H2O(l) ∆H120 °C ) 517.54 kJ/mol (2) H2S(g) + H2SO4(a) ) S + SO2(g) + 2H2O(l) ∆H120 °C ) 100.27 kJ/mol (3) 2H2S(g) + SO2(g) ) 3S + 2H2O(l) ∆H120 °C ) -216.71 kJ/mol (4) 3H2S(g) + H2SO4(a) ) 4S + 4H2O(l) ∆H120 °C ) -116.44 kJ/mol (5) It is known that sulfur may present as S2 or S8 and often in the mixture of S2 and S8. In this study we
10.1021/ie990717z CCC: $19.00 © 2000 American Chemical Society Published on Web 05/12/2000
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Figure 1. Influence of temperature on ∆G of the reactions between H2S and H2SO4. 1. S + 2H2SO4(a) ) 3SO2(g) + H2O(l), ∆H120 °C ) 417.26 kJ. 2. H2S(g) + 3H2SO4(a) ) 4SO2(g) + 4H2O(l), ∆H120 °C ) 517.54 kJ. 3. H2S(g) + H2SO4(a) ) S + SO2(g) + 2H2O(l), ∆H120 °C ) 100.27 kJ. 4. 2H2S(g) + SO2(g) ) 3S + 2H2O(l), ∆H120 °C ) -216.71 kJ. 5. 3H2S(g) + H2SO4(a) ) 4S + 4H2O(l), ∆H120 °C ) -116.44 kJ. (a): aqueous. (g): gas. (l): liquid.
Figure 3. Equilibrium compositions at 120 °C after interaction of 1 mol of H2S with 1 mol of H2SO4 of different concentrations.
150 °C. Below this temperature, sulfur is the dominant product. These results are consistent with the ∆G calculations and the experimental results reported in the literature.4 The effect of the H2SO4 concentration on the equilibrium composition at 120 °C was also calculated using HSC for a system initially containing 1 mol of H2S gas and 1 mol of a sulfuric acid solution of various concentrations. The results illustrate that H2S may be converted to sulfur completely at an acid concentration above 30 wt % (Figure 3). To determine the independent reactions involved in this H2S-H2SO4 reaction process, further experimental studies were carried out in a laboratory batch reactor. Experimental Section
Figure 2. Equilibrium compositions at different temperatures after 1 mol of H2S interacts with 1 mol of 98 wt % H2SO4.
followed the HSC representation of sulfur as S in our calculations. The free energies, ∆G, of reactions (1)(5) calculated using HSC for the temperature range of 0-200 °C are shown in Figure 1. This figure shows that sulfur is unlikely to be oxidized at temperatures below 150 °C (reaction (1)). Sulfur dioxide as the sole product is also unlikely below 150 °C (reaction (2)). Reaction (3) is increasingly likely to occur as the temperature increases above about 50 °C. Reactions (4) and (5) are both indicated to be likely reactions, reaction (5) more so as the temperature increases. These results suggest that reaction (3) is likely to
H2S(g) + H2SO4(a) ) S + SO2(g) + 2H2O
(3)
occur even at temperatures as low as 50 °C. Additional consecutive reactions of the sulfur dioxide produced with hydrogen sulfide in the solution (reaction (4)) may also be anticipated at all temperatures examined. Reaction (5) also appears possible over a wide temperature range. The influence of the reaction temperature on the equilibrium distribution of sulfur-containing components was estimated using HSC for a system initially containing 1 mol of hydrogen sulfide and 1 mol of sulfuric acid (98 wt % concentration). The results show that dithiosulfate and bisulfate ions, sulfur, and sulfur dioxide may be present (Figure 2). Further oxidation of sulfur by the concentrated sulfuric acid is possible above
Materials and Analysis. The acid of specific concentration was prepared batchwise by diluting ∼96% concentrated sulfuric acid (Anachemia Canada Inc.). The concentrations of both fresh and spent acids were measured by titration with a standard 0.1 N NaOH solution (Fisher). Hydrogen sulfide (C.P. grade) was provided by Matheson Gas Products Inc., Canada. The final gas composition was analyzed using a GOW MAC GC with a thermal conductivity detector (TCD). The GC column was 2 m of a 1/8 in. o.d. stainless steel tube packed with Poropak Q (100-120 mesh). The GC column temperature was kept at 100 °C and the detector temperature at 125 °C. Helium (Praxair Canada) was used as a carrier gas with a flow rate of 30 mL/min (STP). Batch Experimental Procedures. The schematic diagram of the experimental setup is shown in Figure 4. The heated Pyrex glass reactor is 122 cm3 in volume, containing a magnetic stirrer. A sensitive Heiss pressure gauge was used to measure the pressures either in the gas reservoir (334 cm3) or in the reactor. For most of the batch experimental runs, the temperature was controlled at 120 ( 0.2 °C using an OMERON temperature controller. The heat of the overall reaction was estimated qualitatively by the temperature change observed during the reaction course. During the measurements, the heater output was kept at a constant level at which the temperature in the reactor could be maintained at 120 ( 0.2 °C if no H2S was introduced into the reactor. Under this condition, the change in the acid temperature as a result of the reaction between H2S and H2SO4 was recorded using a jacketted thermocouple.
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Figure 4. Schematic diagram of batch experimental setup: (1) gas supply cylinder; (2) gas reservoir; (3) pressure gauge; (4) reactor; (5) heater and magnetic stirrer; (6) needle valve; (7) threeway valve; (8) on-off valve.
The experimental procedure involved accurately weighing the amount of acid solution (around 70 g) charged to the reactor. The solution was bubbled using nitrogen (Praxair Canada) and then evacuated under vacuum for 20 min to remove the dissolved oxygen. Pure H2S gas was introduced into the reservoir to a certain pressure. After the initial pressure and temperature were recorded, H2S gas was introduced from the reservoir into the reactor. The final pressure of the reservoir was recorded, and through the difference of the initial and final pressures of the reservoir and its volume, the amount of H2S fed into the reactor was measured. Immediately following the introduction of H2S into the reactor, the stirrer was turned on. The pressure change in the reactor against time was recorded, and the final pressure was taken when the pressure was observed unchanging in the period of half an hour. The final gasphase composition in the reactor was sampled and analyzed by using the gas chromatograph. The amount of SO2 dissolved in the acid was estimated using the SO2 solubility data or correlation.10 These measurements showed the amount of H2S consumed and the amount of SO2 produced. The mixture of sulfur and the acid was purged with nitrogen for 3 h to remove any dissolved SO2 gas before it was weighed. The mixture was then heated at 140 °C for 5 h and cooled, making sulfur form big solid flakes and easily separable by filtration. The separated sulfur was washed, dried, and weighed. The concentration of the reminded acid solution was determined by titration using a 0.1 N standard sodium hydroxide solution (Fisher). The concentrated acid was diluted with distilled water before titration. As a result, the acid consumption and sulfur production data were obtained. Batch Experimental Results. 1. Reaction between Sulfur and Aqueous Sulfuric Acid. It has been shown in the previous thermodynamic analysis that sulfur can be oxidized with concentrated acid at temperatures above 150 °C. Several batch runs were carried out to evaluate the reactivity of sulfur in aqueous sulfuric acid at 120 °C. After interaction of H2S with 96 wt % sulfuric acid, the mixture of sulfur and acid was purged for 3 h with nitrogen to remove the dissolved SO2. The mixture was then heated at 120 °C overnight. No pressure increase was observed for the gas phase, and GC analysis could not detect any amount of sulfur dioxide produced within the detection limit of the GC which was 0.075 mol % for SO2. The same result was observed even at 140 °C. Because the oxidizability of the aqueous sulfuric acid solution declines with the decrease of its concentration, this observation indicates that the sulfur oxidation by aqueous acid with concen-
Figure 5. Influence of the acid concentration and multiple feed on H2S conversion [)(H2S fed - H2S remaining)/H2S fed × 100%] and SO2 selectivity ()net SO2 produced/H2S consumed × 100%). H2SO4 concentration: 88 wt %; 120 °C. b,O: Single feed. Feeding H2S in sequence four times.
Figure 6. Temperature of the acid solution and initial pressure drop versus reaction time. H2SO4 concentration: 96 wt %.
trations of 96 wt % and less at temperatures e 140 °C is insignificant. This conclusion is in good agreement with the thermodynamic analysis (Figures 1 and 2) as well as the observations of Snurnikov et al.4 Consequently, sulfur was then separated by filtration after the spent acid was heated at temperature ∼140 °C for several hours. The colloided sulfur formed large particles upon heating which resulted in easy filtration. 2. Influence of the Acid Concentration and the Heat of Reaction. Complete conversion of H2S was obtained over a wide range of acid concentrations from 65 to 96 wt % (Figure 5), though the time it took to reach it was different for various acid concentrations. However, SO2 production depended strongly on the acid concentration. At an acid concentration of 80 wt % or lower, SO2 was not detected. Even in the solutions of acid concentration higher than 80 wt %, where SO2 was observed in the remaining gas phase, the mass balance showed that the amount of SO2 measured did not agree with that calculated stoichiometrically based on any single reaction proposed in the literature. SO2 produced by the oxidation reaction had to be consumed somehow in additional reactions. This conclusion was supported by the experimental results of multiple introduction of hydrogen sulfide into the reactor system containing 88 wt % acid solution. If there were not any additional reactions to consume the SO2 produced by H2S oxidation, SO2 would have built up when more H2S was introduced into the reactor system. In fact, multiple introduction of more H2S into a 88 wt % sulfuric acid
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solution significantly reduces the SO2 selectivity, defined as the mole ratio of SO2 formed over the total H2S reacted (Figure 5). Compared to a single H2S feed, multiple injection of more H2S reduced the SO2 selectivity from 20.1 to 6%. This result also indicates that further reaction of H2S with SO2 takes place in the sulfuric acid solution. To obtain further insight into the reactions between H2S and H2SO4, the influence of the reaction on the observed acid temperature was measured. An immediate temperature rise (about 8 °C) was observed once H2S was introduced. It should be pointed out that the overall heat is related to reaction conditions. Part of the temperature rise may be attributed to the heat released from dilution of acid because H2O is one of the reaction products. Both consumption of H2SO4 and production of H2O during the reaction will lead to a decrease in the acid concentration. Our estimation of the temperature rise related to dilution of 88 wt % sulfuric acid is less than 4 °C based on the amount of water produced during the reaction. It is evident that the overall reaction is exothermic. Because the reactions leading to formation of sulfur dioxide are all endothermic but the reaction between hydrogen sulfide and sulfur dioxide is exothermic, the exothermic characteristics of the overall reaction also support our view of the involvement of an H2S-SO2 reaction in the H2S-H2SO4 system. The stoichiometry observed for reactant consumption (H2S) and SO2 generation from various steady-state continuous operations of a packed column (not described herein) was compared for consistency with various combinations of reactions (1)-(5). Mathematically, one can show that two independent reactions are sufficient to describe the H2S-H2SO4 reaction system. Of the various combinations tested, only that of reactions (3) and (4) was shown to be independent and to be satisfactory in describing the influence of the acid concentration and the mode of H2S feeding on the product distribution. The results from the above analysis and the exothermic character of the reaction system provide the basis for our view that SO2 reacts consecutively and H2S reacts in parallel in reactions (3) and (4). Discussion The formation of stable sulfur and SO2 at 120 °C was predicted by thermodynamic analysis and was also confirmed by the experimental results. The amount of SO2 produced during the interaction between H2S and H2SO4 cannot be reconciled with use of the single reaction equation which relates hydrogen sulfide consumption solely to oxidation by sulfuric acid. The fact that both sulfur and sulfur dioxide are formed and that the overall reaction is exothermic cannot be explained by use of a single reaction equation from those postulated in the literature. Sulfur dioxide must be an intermediate reaction product, which has been indicated by the dependence of the amount of SO2 production on the sulfuric acid concentration and results of SO2 production in different H2S feeding modes. It has been well understood that hydrogen sulfide is a powerful reducing agent and concentrated sulfuric acid is a strong oxidant. The higher the acid concentration, the stronger the oxidizing power. When H2S contacts with H2SO4, H2S can be oxidized into S and H2SO4 reduced into SO2 according to the following reaction:
H2S + H2SO4 ) S + SO2 + H2O
(i)
The SO2 produced further reacts with H2S according to the following reaction:
2H2S + SO2 ) 3S + 2H2O
(ii)
Reaction (ii) is called the Wackenroder reaction when it takes place in liquid media6 and the modified Claus reaction when it occurs on the surface of a catalyst. Dry H2S and SO2 do not react when they are mixed, but they react slowly in nonaqueous media.11 However, the reaction between them proceeds more rapidly over any wet surface even at room temperature.12 This reaction has been used commercially for natural gas desulfurization in the Townsend process.13 The reaction in aqueous media differs from the modified Claus reaction in that it uses water as the medium instead of a catalyst. In our laboratory, we have observed the formation of sulfur at room temperature by simply introducing the mixture of H2S and SO2 into water. Sulfur produced in this reaction is in colloidal form and appears as a milky white liquid product stream. The increase in H2S conversion and the decrease in SO2 production resulted from multiple injection of more H2S into the reactor confirmed that the Wackenroder reaction was taking place in the reactor (Figure 5). The above reaction scheme, parallel for hydrogen sulfide and consecutive for sulfur dioxide, explains well the reaction stoichiometry encountered in the experiments as well as the overall exothermic characteristics of the reaction. Thermodynamic analysis indicated the possibility of formation of dithiosulfate ions at a temperature < 40 °C (Figure 2). We have attempted to identify these species by analyzing the solution after interaction of H2S with sulfuric acid of various concentrations at different temperatures using UV-vis. Thiosulfate ions were detected only with a diluted acid solution (
