Recovery of Hydrogen and Sulfur by Indirect Electrolysis of Hydrogen

Jul 29, 2009 - Hydrogen sulfide (H2S) absorption and conversion to hydrogen and sulfur were carried out in an acidic aqueous vanadium dioxide (VO2)+ s...
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Energy Fuels 2009, 23, 4420–4425 Published on Web 07/29/2009

: DOI:10.1021/ef900424a

Recovery of Hydrogen and Sulfur by Indirect Electrolysis of Hydrogen Sulfide Haiyan Huang,*,† Ying Yu,† and Keng H. Chung‡ †

State Key Laboratory of Heavy Oil Processing, China University of Petroleum, Beijing 102249, People’s Republic of China, and ‡ Department of Chemical Petroleum Engineering, The University of Calgary, 2500 University Drive Northwest, Calgary, Alberta T2N 1N4, Canada Received May 7, 2009. Revised Manuscript Received July 8, 2009

Hydrogen sulfide (H2S) absorption and conversion to hydrogen and sulfur were carried out in an acidic aqueous vanadium dioxide (VO2)þ solution coupled with indirect electrolysis. In this paper, the mechanisms of absorption and electrochemical reactions of the process are discussed. Parametric studies were conducted to determine the effects of operating parameters on absorption and electrochemical reactions. The results showed that the H2S absorption increased with temperature; greater than 90% of H2S absorption occurred at 50 °C. The absorption reaction was mass-transfer-limiting. In the electrolysis reaction, the current efficiency reached 97% at 45 °C after an extended electrolysis time. The optimal reaction conditions were at a proton concentration of 7 mol/kg of H2O in the electrolyte, (VO2)þ concentration of lower than 0.65 mol/kg of H2O in the electrolyte, and (VO2)þ concentration of higher than 0.55 mol/kg of H2O in the absorbent. Sulfur particles that are produced can be easily recovered. Some aspects related to design optimization of the absorption process and electrochemical reactor are also discussed.

In recent years, various attempts have been made to explore the recovery of hydrogen and sulfur from H2S. These include thermal,5,6 thermochemical,7 electrochemical,8-16 photochemical,17,18 and plasmochemical methods.19 Electrolysis of

Introduction Hydrogen sulfide (H2S) is one of the main byproducts in natural gas plants, refineries, heavy oil upgraders, and metallurgical processes. It is a toxic gas and classified as a hazardous industrial waste. Since the 1970s, tremendous progress has been made to use H2S as a feedstock for the production of sulfur and hydrogen. Sulfur is an important base chemical feedstock, especially in fertilizer production. Hydrogen is a valuable commodity in heavy oil upgrading and refining operations.1 The standard commercial technology for treating H2S consists of a wet H2S absorption process followed by the Claus process, in which H2S is converted to elemental sulfur and water. In the Claus process, the reactions are conducted at high temperatures and require costly equipment.2,3 In addition, the Claus process recovers only element sulfur; hydrogen is oxidized to form water. In 2006, about 55 million tons of sulfur was produced from treating industrial waste gases that contain sulfur (H2S and SOx). This accounted for 80% of the total sulfur production in the world.4 If the hydrogen could be recovered, it would represent an enormous additional amount of hydrogen supply.

(7) Noring, J. E; Fletchers, E. A. High temperature solar thermochemical processing;Hydrogen and sulfur from hydrogen sulfide. Engery 1982, 7 (8), 651–666. (8) Anani, A. A.; Mao, Z.; White, R. E.; Srinivason, S.; Appleby, A. J. Electrochemical production of hydrogen and sulfur by low temperature decomposition of hydrogen sulfide in an aqueous alkaline solution. J. Electrochem. Soc. 1990, 137 (9), 2703–2709. (9) Mao, Z.; Anani, A.; White, R. E.; Srinivasan, S.; Appleby, A. J. A modified electrochemical process for the decomposition of hydrogen sulfide in an aqueous alkaline solution. J. Electrochem. Soc. 1991, 138 (5), 1299–1303. (10) Hyun, S. L.; Winnick, J. Electrochemical removal and concentration of hydrogen sulfide from coal gas. J. Electrochem. Soc. 1984, 131 (3), 562–568. (11) Weaver, D.; Winnick, J. Electrochemical removal of H2S from hot gas streams. J. Electrochem. Soc. 1987, 134 (10), 2451–2458. (12) Kalina, D. W.; Maas, E. T. Indirect hydrogen sulfide conversion;I. An acidic electrochemical process. Int. J. Hydrogen Energy 1985, 10 (3), 157–162. (13) Kalina, D. W.; Maas, E. T. Indirect hydrogen sulfide conversion;II. An basic electrochemical process. Int. J. Hydrogen Energy 1985, 10 (3), 163–167. (14) Olson, D. C. Method of removing hydrogen sulfide from gases utilizing a polyvalent metal chelate solution and electrolytically regenerating the solution. U.S. Patent 4,443,423, 1984. (15) Yu, Y.; Wang, C. Z.; Zhao, Y. F.; Zhu, Y. J. Hydrogen production from hydrogen sulfide by a combination technique of oxidation electrolysis process. Acta Energ. Sol. Sin. 1997, 18 (4), 400–407. (16) Mizuta, S.; Kondo, W.; Fujii, K. Hydrogen production from hydrogen sulfide by the Fe-Cl hybrid process. Ind. Eng. Chem. Res. 1991, 30 (7), 1601–1608. (17) Cervera-March, S.; Borrell, L.; Gimenez, J.; Simarro, R.; Andujar, J. M. Solar hydrogen photoproduction from sulphide/sulphite substrate. Int. J. Hydrogen Energy 1992, 17 (9), 683–688. (18) Sabate, J.; Cervera-March, S.; Simarro, R.; Gimenez, J. A comparative study of semiconductor photocatalysts for hydrogen production by visible light using different sacrificial substrates in aqueous media. Int. J. Hydrogen Energy 1990, 15 (2), 115–124. (19) Zaman, J.; Chakma, A. Production of hydrogen and sulfur from hydrogen sulfide. Fuel Process. Technol. 1995, 41, 159–198.

*To whom correspondence should be addressed. Telephone: þ86-1089733444. E-mail: [email protected]. (1) Fletcher, E. A.; Noring, J. E.; Murray, J. P. Hydrogen sulfide as a source of hydrogen. Int. J. Hydrogen Energy 1984, 9 (7), 587–593. (2) Pieplu, A.; Saur, O.; Lavxlley, J. C. Claus catalysis and H2S selective oxidation, catalysis reviews. Sci. Eng. 1998, 40 (4), 409–450. (3) Eow, J. S. Recovery of sulfur from sour acid gas: A review of the technology. Environ. Prog. 2002, 21 (3), 143–162. (4) U.S. Department of the Interior U.S. Geological Survey. 2006 Minerals Yearbook: Sulfur, http://minerals.usgs.gov/minerals/pubs/ commodity/sulfur/myb1-2006-sulfu.pdf. (5) Bandermann, F.; Harder, K. B. Production of H2 via thermal decomposition of H2S and separation of H2 and H2S by pressure swing adsorption. Int. J. Hydrogen Energy 1982, 7 (6), 471–475. (6) Faraji, F.; Safakik, I.; Strausz, O. P.; Yildikim, E.; Torress, M. E. The direct conversion of hydrogen sulfide to hydrogen and sulfur. Int. J. Hydrogen Energy 1998, 23 (6), 451–456. r 2009 American Chemical Society

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H2S is a promising process used to generate hydrogen and sulfur. It can be carried out by direct electrolysis, hightemperature electrolysis, and indirect electrolysis. The direct electrolysis process requires lower power input. However, the issues related to passivation of electrodes by sulfur, difficult separation and handling of sulfur, and undesirable secondary electrochemical reactions are yet to be resolved.8,9 The high-temperature electrochemical process has the advantage of being able to convert H2S from hot gas streams.10,11 The indirect electrochemical process has more operational flexibility.12-16 The selection of the electrochemical intermediate is a key criterion for the indirect electrolysis process, because each electrochemical medium has unique characteristics. Kalina and Maas12,13 compared two similar indirect electrochemical processes for converting H2S to elemental sulfur and hydrogen. One used a soluble tri-iodide (I3-) oxidant in acidic solution at pH 0-1, and one used soluble iodate (IO3-) oxidant in basic solution at pH 13-14. High current density and efficiency were achieved for the basic system, but the voltage requirement (4.9 V) was much higher than for the acidic system (2.1 V). Although the amount of sulfur produced in the basic system was high, because of the thermodynamic instability of sulfur in basic solution, the overall reaction efficiency in the process was lower. In contrast, in the acidic system, the sulfur product was more difficult to recover and purify but the overall reaction efficiency was high (>90%). Indirect electrochemical processes using iron complexes and iron chloride have also been investigated.14-16 Olson14 described a process involving a chemical step in which an iron chelate was reacted with H2S to produce sulfur, followed by an electrochemical step to recover the chelate and generate hydrogen. The chemical step involved complex reactions of H2S with iron chelate. A hybrid process consisted of H2S absorption in iron chloride (FeCl3) aqueous solution and subsequent electrolysis of aqueous FeCl2 solution.15,16 The process was carried out under strong acidic conditions, using an excess amount of hydrochloric acid (HCl). At 70 °C, almost 100% H2S absorption was achieved with a low electrolysis voltage (0.7 V) at a current density of 100 mA/cm2. However, the gummy sulfur product was difficult to recover. In this work, an indirect electrochemical process is used to convert H2S using acidic aqueous vanadium dioxide [(VO2)þ/(VO)2þ] as the electrochemical intermediate. Recovery of hydrogen and sulfur by indirect electrolysis of hydrogen sulfide in a new redox system has not been performed. The (VO2)þ/(VO)2þ couple is capable of oxidizing sulfide to elemental sulfur but not to higher oxidation states and does not undergo side reactions with H2S. In addition, the (VO2)þ/(VO)2þ couple is stable and has a high solubility for both the oxidized and reduced species under the operating window considered. The electrolysis reaction of the (VO2)þ/(VO)2þ couple is simple and easy to operate. The objectives are to determine the electrolysis reactions at high current density and efficiency and to examine H2S absorption in acid solution and sulfur production. The underlying absorption and electrochemical reaction mechanisms will be discussed.

Figure 1. Schematic diagram of indirect electrolysis of the hydrogen sulfide process.

of H2S.20-23 In acidic solutions, (VO2)þ oxidizes H2S to elemental sulfur in the absorption reactor and sulfur can be separated by filtration. The (VO2)þ/(VO)2þ couple meets the requirements of the electrochemistry process: (VO2)þ can be regenerated below the electric potential of the released oxygen, and the protons (Hþ) can be reduced to hydrogen gas simultaneously. A generalized process flow diagram is shown in Figure 1. The process consists of two main unit operations: chemical absorption and electrochemical conversion.15,16 It is applicable to H2S-containing gas streams, such as those produced from natural gas sweetening and tail gas cleanup. In the absorption reactor, a H2S-containing gas stream is contacted with aqueous (VO2)þ solution. The H2S reacts with (VO2)þ and is oxidized to sulfur and Hþ. The (VO2)þ is reduced to (VO)2þ in the presence of excess acid. In the electrochemical reactor, (VO)2þ is oxidized to (VO2)þ at the anode and Hþ permeating through the Nafion membrane is reduced to H2 at the cathode. The simultaneous anodic and cathodic reactions that occur at the respective electrodes in the electrochemical cell make up the overall electrochemical process. The overall reactions can be expressed as follows. Absorption reaction: 2Hþ þ 2ðVO2 Þþ þ H2 S f 2ðVOÞ2þ þ SV þ 2H2 O Electrolysis reaction: anode : 2ðVOÞ2þ þ 2H2 O f 2ðVO2 Þþ þ 4H þ þ 2ecathode : 2Hþ þ 2e- f H2 v

ð2Þ ð3Þ

Overall: H2 S f H2 v þ SV

ð4Þ þ

2þ

Because the energy potential, E°, of ((VO2) /(VO) ) is 0.991 V, E° of (S/S2-) is 0.142 V,24 E° of H2S absorption is 0.849 V, and the standard Gibbs free energy change (G°) of (20) Oriji, G.; Katayama, Y.; Miura, T. Investigation on V(IV)/V(V) species in a vanadium redox flow battery. Electrochim. Acta 2004, 49 (19), 3091–3095. (21) Skyllas-Kazacos, M.; Menictas, C.; Kazacos, M. Thermal stability of concentrated V(V) electrolytes in the vanadium redox cell. J. Electrochem. Soc. 1996, 143 (4), 86–88. (22) Kausar, N.; Howe, R.; Skyllas-Kazacos, M. Raman spectroscopy studies of concentrated vanadium redox battery positive electrolytes. J. Appl. Electrochem. 2001, 31 (12), 1327–1332. (23) Zhong, S.; Skyllas-Kazacos, M. Electrochemical behaviour of vanadium(V)/vanadium(IV) redox couple at graphite electrodes. J. Power Sources 1992, 39 (1), 1–9. (24) Lide, D. R. CRC Handbook of Chemistry and Physics; CRC Press: Boca Raton, FL, 2005; pp 8-27-8-28.

Theory þ

ð1Þ

2þ

The (VO2) /(VO) couple is an attractive candidate for the electrochemical intermediate in the indirect electrolysis 4421

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reaction 1 is negative, -163.86 kJ/mol, it suggests that H2S absorption can occur spontaneously to completion. In the electrochemical process, the theoretical voltage requirement is slightly more than 1.0 V. Hence, it is theoretically feasible to use (VO2)þ/(VO)2þ as the electrochemical intermediate, which is the basis for the experimental work.

Results and Discussion Hydrogen Sulfide Absorption Reaction. The absorption rate and reaction efficiency are dependent upon operating parameters, such as temperature, concentration of (VO2)þ and Hþ, and H2S. Temperature. The H2S absorption reaction is a gas-liquid reaction, consisting of mass-transfer and chemical reactions. The H2S diffuses from the gas phase to the gas-liquid interface, followed by H2S reacting with (VO2)þ in solution. Theoretically, the diffusion and chemical reaction rate increase with temperature. The experimental data in Figure 2 were in agreement with the theory, indicating the H2S absorption efficiency increased with temperature. Other factors that could contribute to the increased H2S absorption as the temperature increased is the reduced viscosity of the absorbent and increased reaction rate constant. Figure 3 showed the Arrhenius plot of ln[vH2S] as a function of 1/T for the H2S absorption reaction. From the slope of the Arrhenius plot, the apparent activation energy of the absorption reaction, Ea, of 8.862 kJ/mol was calculated. The apparent activation energy indicates that the reaction rate increases with the temperature. However, because the apparent activation energy was less than 30 kJ/mol, it suggests that the chemical reaction was rapid and the overall absorption reaction was controlled by mass transfer. Therefore, an increased temperature is beneficial to the absorption reaction because it increased the diffusion rate of the reactant. The mass transfer can also be improved from better reactor design, such as improved gas distribution and gasliquid flow regime.

Experimental Section The absorption reaction of (VO2)þ/(VO)2þ solution and H2S was carried out in a 40 cm tall and 5 cm diameter annular cylindrical glass reactor. The (VO2)þ/(VO)2þ solution was the continuous liquid phase. The H2S-containing gas was injected into the liquid phase through a sintered-glass gas dispersion tube. The gas and liquid inlets were located at the top of reactor. The gas bubbles flowed upward, countercurrent to the absorbent flow, and exited at the top of reactor. At the termination of each experimental run, the product slurry was placed in a thermal bath at an absorption temperature for 20 min. The sulfur particles were recovered by filtration and analyzed by a Coulter Laser Granularity analyzer. The off gas, which contained unreacted H2S from the chemical reactor, was routed to a caustic scrubber. The amount of H2S in the scrubber was determined by iodometry. The ability of the absorbent solutions to remove H2S from a gas sample was determined from the (VO2)þ concentration by titration using standardized Fe2þ solutions. The total concentration of (VO2)þ and (VO)2þ in the absorbing solution was 1 mol/kg of H2O. In this study, H2S absorption efficiency was defined as the amount of H2S reacted with respect to that in the feed. The H2S absorption rate, vH2S, is defined as the molar quantity of absorbent solutions required to remove H2S from a gas sample in an 8 min experimental interval. The H2S absorption reaction process is operated under strongly acidic conditions. For H2S absorption, excess Hþ is known to hinder H2S solubility in the absorbent. On the other hand, a low Hþ concentration will reduce the solubility of (VO2)þ and (VO)2þ in the absorbent, causing (VO2)þ and (VO)2þ to precipitate. To select the appropriate Hþ concentration used for this work, experiments were carried out by varying the Hþ concentration from 3.5 to 11 mol/kg of H2O in the absorbing solution. Subsequently, the Hþ concentration of the absorbing solution was set at 7 mol/kg of H2O. The electrochemical unit operation was based on a dual fluid cell. The cell used a parallel plate electrode configuration with an anode-cathode separation. The electrode surface area was 4 cm2. Platinum electrodes were used for this work not only for their electrical properties but also for their inert chemical and stable physical properties, which do not have adverse effects on the electrolyte and electrolysis processing. This cell allowed for delivery of separate anolyte and catholyte feed streams and maintained electrolyte isolation inside the cell via an ionexchange membrane between the electrodes. A CHI650A electrochemistry test station was used to measure the performance of the electrodes and maintained a constant cell voltage (E = 1.2 V) and current. Electrolyte temperatures were monitored in-line using thermocouples. The (VO2)þ concentration was determined by titration using standardized KMnO4 solutions. In situ cyclic voltammetry was conducted using a tripolar system. The potential, E, was cycled at a 30 mV/s scanning rate between 0.4 and 1.6 V. The E values are the relative platinum working electrode potential to that of the Ag/AgCl reference when a stable voltammogram was obtained. The total concentration of (VO2)þ and (VO)2þ of the electrolyte was set at 1 mol/kg of H2O, and the Hþ concentration was 7 mol/kg of H2O. The effective current, Qeff, was calculated for various concentrations of (VO)2þ by Faraday’s law. The sum of the current, Qsum, was calculated by the Ampere law. The current efficiency is defined as Qeff/Qsum.

Figure 2. Absorption of H2S as a function of the temperature for the system of m=[(VO2)þ/((VO2)þ þ (VO)2þ)]=1 and 100% H2S gas at 120 cm3/min.

Figure 3. Arrhenius plot of ln vH2S as a function of 1/T for the H2S absorption reaction.

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Concentration of Absorbent. Figure 4 shows the H2S absorption efficiency as a function of the (VO2)þ molar fraction, m = [(VO2)þ/((VO2)þ þ (VO)2þ)]. The results showed that H2S absorption efficiency increased with the (VO2)þ molar fraction up to 0.55 (VO2)þ molar fraction. After that, the H2S absorption efficiency leveled off at 90%. Because the H2S absorption reaction consists of mass transfer and chemical reactions, the absorbent with a low (VO2)þ molar fraction does not have sufficient (VO2)þ to react with H2S. Hence, the H2S absorption reaction is chemical-reaction-limiting for the absorbent with a low (VO2)þ molar fraction. As the (VO2)þ molar fraction of absorbent increased, more (VO2)þ is available to react with H2S, resulting in an increased H2S absorption efficiency. At a relatively high (VO2)þ molar fraction of absorbent (exceeding a critical concentration), the (VO2)þ is in excess with respect to the H2S. Hence, the absorption reaction is H2S-diffusion-limiting and independent of the (VO2)þ molar fraction of the absorbent. The results show that the critical (VO2)þ molar fraction was 0.55. H2S-Containing Gas. Figure 5 shows H2S absorption as a function of the acid gas rate. The results show that the H2S absorption rate increased with the acid gas rate. This is due to enhanced H2S solubility at high H2S concentrations and higher concentrations of HS-/S2- in the absorbing solution. However, the H2S absorption efficiency decreased with the acid gas rate. This is attributed to the chemical equilibrium shift. With a constant (VO2)þ concentration, an increased H2S concentration results in a chemical reaction equilibrium shift in the opposite direction to absorption. Sulfur Recovery. The operating conditions of the absorption process are known to have a direct influence on the particle size of sulfur produced and affect the ease of sulfur recovery. In the run with 50% H2S/50% N2 gas at 120 cm3/ min, (VO2)þ molar fraction of 0.55, and Hþ concentration of 7 mol/kg of H2O, the size of sulfur particles was 2-4 μm. The size of sulfur particles decreased as the temperature increased. Below 50 °C, the sulfur was gummy and adhered to the reactor wall, which is unfavorable from the viewpoint of sulfur recovery. At 50 °C, an increased acid gas rate triggered the agglomeration of sulfur particles, because of turbulent mixing. Subsequently, nitrogen gas was introduced to stir the product slurry, while it was placed in a thermal bath. Larger sulfur particles were easily recovered by filtration. Electrochemical Reaction. Simultaneous oxidation of (VO)2þ to form (VO2)þ and reduction of Hþ to yield hydrogen gas was carried out in a dual fluid cell separated by a membrane. The cell voltage, electrolysis temperature, and electrolyte concentration are the key parameters to be considered. Typical current-voltage characteristics of platinum anodes without adverse effects are shown in Figure 6. The cyclic voltammogram showed that the (VO2)þ/(VO)2þ couple underwent quasi-reversible one-electron reduction and oxidation. The peaks a and b represent oxidation and reduction, respectively. The data showed that, to oxidize (VO)2þ to form (VO2)þ, a cell voltage of higher than 1.2 V is needed. As a result, the cell voltage chosen for electrolysis in this study was set at 1.5 V. Temperature. Figure 7 shows the current density, i, as a function of the temperature. The current density increased with the electrolysis temperature. Figure 8 shows the Arrhenius plot of ln i as a function of 1/T for the electrolysis reaction. From the slope of the Arrhenius plot, the apparent

Figure 4. Absorption of H2S as a function of the (VO2)þ molar fraction for the system of ((VO2)þ þ (VO)2þ)=1 mol/kg of H2O, 50 °C, and 50% H2S/50% N2 gas at 120 cm3/min.

Figure 5. Absorption of H2S as a function of the acid gas rate for the system of m=[(VO2)þ/((VO2)þ þ (VO)2þ)]=0.55, 50 °C, and 50% H2S/50% N2 gas.

Figure 6. Cyclic voltammograms at the platinum electrode for a tripolar cell (electrolyte initial concentration m = [(VO2)þ/((VO2)þ þ (VO)2þ)] = 0, Hþ concentration of 12 mol/kg of H2O, Ag/AgCl as the reference electrode, 22.5 °C, and scanning rate of 0.03 V/s).

activation energy of the electrolysis reaction, Ea, of 16.18 kJ/ mol was obtained. Again, the apparent activation energy indicates that the reaction rate increases with the temperature. Because the apparent activation energy is lower than 30 kJ/mol, this suggests that the electrolysis reaction is also mass-transfer-limiting. An increased temperature not only improves the reaction on the electrode but also accelerates the diffusion of the reactant. Furthermore, the viscosity 4423

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Figure 7. Current density as a function of the temperature for a bipolar cell (electrolyte initial concentration m=[(VO2)þ]/([(VO2)þ] þ [ (VO)2þ])=0, Hþ concentration of 12 mol/kg of H2O, platinum electrode, and cell voltage of 1.5 V).

Figure 9. Current density as a function of the (VO2)þ molar fraction for bipolar cell of ((VO2)þ þ(VO)2þ)=1 mol/kg H2O, Hþ concentration of 13 mol/kg of H2O, platinum electrode, cell voltage of 1.5 V, and 45 °C.

Figure 8. Arrhenius plot of ln i as a function of 1/T for the electrolysis reaction.

Figure 10. Current density as a function of the Hþ concentration for bipolar cell of an initial electrolyte concentration, m = [(VO2)þ]/([(VO2)þ] þ [(VO)2þ]) = 0.6, platinum electrode, cell voltage of 1.5 V, and 45 °C.

and resistance of the electrolyte decreased with the temperature. Hence, a high temperature improves the mass transfer. On the other hand, a high temperature increases the power input and may trigger undesirable electrochemical reactions.20,21 Concentration of the Electrolyte. Figure 9 shows the current density as a function of the (VO2)þ molar fraction. The current density was relatively constant up to a (VO2)þ molar fraction of 0.65 and then dropped dramatically. The anodic reaction of (VO)2þ with water is shown in eq 2. At high (VO)2þ concentrations that exceed the critical concentration of 0.35, i.e., (VO2)þ concentration less than 0.65, the anodic reaction was mass-transfer-controlled and independent of (VO)2þ concentrations.21-23 Below the critical concentration of (VO)2þ, all of the (VO)2þ ions were consumed by the anodic reaction. The availability of (VO)2þ ions limits the extent of the anodic reaction. Hence, the current density was dependent upon the (VO)2þ concentration. Protons are the reactants to yield hydrogen on the cathode (cf. eq 3) and maintain the acidity of the electrolyte. Hence, the Hþ concentration is an important parameter influencing the current density.25 Figure 10 shows the current density as a function of the Hþ concentration. As the Hþ concentration increased, the current density increased, reached a maximum, and decreased. The peak current density was at a Hþ concentration of 7 mol/kg of H2O. This was consistent with the screening test results that were used to select the Hþ

concentration used in the absorption reaction experiments. At low Hþ concentrations, the current density increased with the Hþ concentration. This is because sulfuric acid improves conductivity and reduces the solution resistance when it used as a cathode electrolyte. However, excess sulfuric acid hinders Hþ hydration, because of strong interactions between sulfuric acid and water. It interferes with the first step of the hydrogen evolution reaction,26 as shown below. Hþ þ H2 O f H3 Oþ

ð5Þ

H3 Oþ þ e f MH þ H2 O

ð6Þ

MH þ MH f H2

ð7Þ

MH þ H3 Oþ þ e f H2 þ H2 O

ð8Þ

Current Efficiency. The current efficiency was 97% at various extended electrolysis times. There were a few factors that could cause the loss of current efficiency. First, water in the electrolyte vaporized, because of the extended electrolysis time. Second, gas evolving on the cathode could change the electrolyte conductivity. Third, the electrolyte became a gas-liquid mixture, because of gas dispersion in the bulk (26) Azizi, O.; Jafarian, M.; Gobal, F.; Helia, H.; Mahjania, M. G. The investigation of the kinetics and mechanism of hydrogen evolution reaction on tin. Int. J. Hydrogen Energy 2007, 32 (12), 1755–1761.

(25) Rahman, F.; Skyllas-Kazacos, M. Solubility of vanadyl sulfate in concentrated sulfuric acid solutions. J. Power Sources 1998, 72 (2), 105–110.

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liquid phase, leading to decreased conductivity and increased solution resistance. Lastly, gas bubbles generated, coalesced, grew in size, and adhered on the surface of the electrode, reducing its active surface.27-29 To enhance the current efficiency, the design of the electrochemical reactor should be improved to increase mass transfer for the reaction and with a more effective gas evacuation system.

vanadium dioxide (VO2) solution coupled with indirect electrolysis. The mechanisms of absorption and electrochemical reactions of the process were discussed. Parametric studies were conducted to determine the effects of operating parameters on absorption and electrochemical reactions. The results showed that the H2S absorption increased with temperature with a greater than 90% H2S absorption at 50 °C. The absorption reaction was mass-transfer-limiting. In the electrolysis reaction, the current efficiency reached 97% at 45 °C after an extended electrolysis time. The optimal reaction conditions were at an Hþ concentration of 7 mol/kg of H2O in the electrolyte, the (VO2)þ concentration of lower than 0.65 mol/kg of H2O in the electrolyte, and the (VO2)þ concentration of higher than 0.55 mol/kg of H2O in the absorbent. Sulfur particles produced can be easily recovered. Some aspects related to the design optimization of absorption and electrochemical reactor were also discussed.

Conclusions The feasibility of H2S absorption and conversion to hydrogen and sulfur was demonstrated in an acidic aqueous (27) Bisang, J. M. Effect of mass transfer on the current distribution in monopolar and bipolar electrochemical reactors with a gas-evolving electrode. J. Appl. Electrochem. 1993, 23 (10), 966–974. (28) Gabrielli, C.; Huet, F.; Keddam, M.; Macias, A.; Sahar, A. Potential drop due to an attached bubble on a gas-evolving electrode. J. Appl. Electrochem. 1989, 19 (5), 617–629. (29) Vogt, H. Mechanisms of mass transfer of dissolved gas from a gas-evolving electrode and their effect on mass transfer coefficient and concentration over potential. J. Appl. Electrochem. 1989, 19 (5), 713–719.

Acknowledgment. The financial support was provided by the China National Advanced Technology Research and Development Program.

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