Upper Bound for the Efficiency of a Novel Chemical Cycle of H2S

2184

Energy & Fuels 2009, 23, 2184–2191

Upper Bound for the Efficiency of a Novel Chemical Cycle of H2S Splitting for H2 Production Guifen Yu,† Hui Wang,*,† and Karl T. Chuang‡ Department of Chemical Engineering, UniVersity of Saskatchewan, Saskatoon, Saskatchewan S7N 5A9, Canada, and Department of Chemical and Materials Engineering, UniVersity of Alberta, Edmonton, Alberta T6G 2G6, Canada ReceiVed October 15, 2008. ReVised Manuscript ReceiVed January 18, 2009

To develop sustainable hydrogen production to meet its needs in oil sands upgrading and refining, a chemical cycle of hydrogen production from splitting hydrogen sulfide, the waste product from the same industrial sector, has been proposed. There are two routes of chemical processes to conduct the chemistry. Process 1 takes 1 mol of H2S and converts it into H2 and S, and process 2 takes hydrogen sulfide, oxygen, and water (1:1:2 H2S/O2/H2O) as a feedstock and produces 2 mol of hydrogen and 1 mol of sulfuric acid. Similar to the sulfur-iodine cycle of water splitting for hydrogen, this cycle consists of an iodine-iodide loop and a sulfur dioxide-sulfuric acid loop. This study uses the thermodynamic data of the reactions involved to calculate an upper bound of the thermal efficiency of the hydrogen sulfide splitting cycle. With the enthalpy of the H2S oxidation reactions as part of the energy input to the cycle, the values of the maximum thermal efficiency for processes 1 and 2 are 0.41 and 0.36, respectively. However, if only external energy that meets the requirement of heat and work for reactions and pumping is taken into account, the thermal efficiency can be higher, at 0.66 and 0.70. It is found that the separations of sulfuric acid and hydrogen iodide respectively from their aqueous solutions generated within the cycle are still the most energy-consuming processes.

1. Introduction

S + O2 f SO2

(4)

The endeavor in converting industrial wastes into value-added products is always economically and environmentally beneficial and so are the efforts of converting hydrogen sulfide from gas and oil industries into hydrogen, particularly, in oil, including oil sands bitumen, upgrading and refining, where H2 is largely needed and H2S is produced. A chemistry cycle of producing H2 from H2S splitting was recently proposed.1 The chemistry was based on the sulfur-iodine (S-I) cycle of H2O splitting for H2, which has been studied for decades.2 The basic chemistry consists of three reactions and converts 1 mol of H2S into 1 mol of H2 or H2S f H2 + S

4H2O + 2I2 + 2SO2 f 2H2SO4 + 4HI

(5)

4HI f 2H2 + 2I2

(6)

H2S + H2SO4 f S + SO2 + 2H2O

(1)

2H2O + I2 + SO2 f H2SO4 + 2HI

(2)

2HI f H2 + I2

(3)

The modified chemistry additionally turns elemental S produced from reaction 1 into SO2; as a result, it converts 1 mol of H2S, 1 mol of O2, and 2 mol of H2O into 2 mol of H2 and 1 mol of H2SO4 or H2S + 2H2O + O2 f 2H2 + H2SO4 H2S + H2SO4 f S + SO2 + 2H2O

(1)

* To whom correspondence should be addressed. Telephone: +1-3069662685. Fax: +1-306-9664777. E-mail: [email protected]. † University of Saskatchewan. ‡ University of Alberta. (1) Wang, H. Int. J. Hydrogen Energy 2007, 32, 1907–3914. (2) Norman, J. H.; Besenbruch, G. E.; Brown, L. C. GA-A16713 Report, 1982.

In petroleum upgrading and refining, hydrotreating processes usually desire more H2 and produce less H2S on a molar basis because of simultaneous removal of sulfur, nitrogen, and aromatics, where they all consume hydrogen but only desulfurization leads to H2S production. In other words, if hydrogen produced from H2S splitting is recycled back for hydrotreating, its amount is insufficient to meet the need in hydrotreating. Therefore, it is anticipated that more moles of H2 are produced from 1 mol of H2S. From this point of view, the modified H2S splitting cycle (reactions 1, 4, 5, and 6) appears to be more advantageous. Currently, hydrogen used for hydrotreating processes is mainly from natural gas reforming. Because of relative shortage of natural gas reserves, the oil industry, particularly oil sands upgrading industry, will face the challenge to maintain itself as sustainable. It is obvious that if H2 can be recovered from H2S removal and circulated back to oil sands upgrading, the reliance of oil sands processing on natural gas will be reduced. In a more general perspective, the chemistry of H2S splitting for H2 production is CO2-free. Provided that the energy to facilitate the process is from nuclear power plants, a route in which H2 is produced without CO2 emission will be possible. In addition, H2, in turn, is a clean-energy medium. On the basis of this discussion, to convert H2S into hydrogen would not only make oil sand processing sustainable but also the process itself would be environmentally friendly in general. Therefore, technologies that are large-scale, cost-effective, and environmentally attractive H2 production processes are to be developed according to this chemistry.

10.1021/ef800895v CCC: $40.75  2009 American Chemical Society Published on Web 03/23/2009

Efficiency of H2S Splitting for H2 Production

Energy & Fuels, Vol. 23, 2009 2185

It was mentioned that the chemistry of H2S splitting was developed on the basis of the S-I cycle of H2O splitting for H2 production. The S-I cycle consists of three reactions3 H2SO4 f SO2 + H2O + 0.5O2

(7)

2H2O + I2 + SO2 f H2SO4 + 2HI

(2)

2HI f H2 + I2

(3)

To develop the engineering process, the S-I cycle is detailed as the following equations, in which the process temperatures are identified such that the Gibbs free energy change of the reaction is as close as possible to zero. The physical states of the involved substances and reaction temperatures are also specified. 16H2O (l) + 9I2 (l) + SO2 (g) f (H2SO4 + 4H2O)L1 + (2HI + 10H2O + 8I2)L2 (393 K) (8) L2 ) (2HI + 10H2O + 8I2)L2 f 2HI (g) + (10H2O + 8I2)L (500 K) (9) 2HI (g) f H2 (g) + I2 (l) (600 K)

(10)

L1 ) (H2SO4 + 4H2O)L1 f H2SO4 (l) + 4H2O (l) (570 K) (11) H2SO4 (l) f H2SO4 (g) (630 K)

(12)

H2SO4 (g) f SO3 (g) + H2O (g) (670 K)

(13)

SO3 (g) f SO2 (g) + 0.5O2 (g) (1140 K)

(14)

where Ta represents the ambient temperature and |∆HH0 2O(Ta)| ) 286 kJ at ambient T and P (298.15 K and 101.3 kPa), Q and W are the heat and work requirements to conduct the reactions, and ηr is the efficiency of the conversion system. Although Goldstein et al. did not use the absolute value for the enthalpy of hydrogen combustion, they omitted the negative sign before it. When calculating the upper bound of the thermal efficiency, they assumed that the energy requirements for every reaction are the minimum prescribed by the thermodynamics; i.e., all of the reactions are reversible, and they set the efficiency of the conversion system as 0.5, considering that the high-temperature nuclear reactors are used. They supposed each reaction to be performed at a temperature at which the Gibbs free energy change for the reaction is or is possibly close to zero, but they ignored the difference in heat capacities between reactants and products and assumed the heat to bring the reactants to reaction conditions and the heat to bring the products back to ambient condition are balanced. They used a reaction to represent a phase change. They only considered the heat that is needed for the endothermic reactions and the work required for increasing the pressures of the material streams in the cycle and neglected the heat released from the exothermic reactions and the work relieved from the depressing streams. The result of their calculation shows that the upper bound of the efficiency of the S-I cycle is 0.51. Following the procedure that Goldstein et al. used, this paper estimates an upper bound of the efficiency of H2S splitting cycles, which consist of reactions 1-3 and reactions 1 and 4-6, respectively. 2. Process for the H2S Splitting Cycle

(15)

A comparison between reactions 1-3, the reactions for the H2S splitting cycle, and reactions 7, 2, and 3, those for the S-I H2O splitting cycle, indicates that the two cycles share two chemical reactions. The difference is that the H2SO4 decomposition in the latter is replaced with the reaction between H2S and H2SO4 in the former. To produce more SO2, the cycle of reactions 1 and 4-6, in comparison to that of 1-3, is to oxidize the elemental S produced from reaction 1 into SO2 and, thus, reactions 2 and 3 take place in doubled stoichiometric coefficients, reactions 5 and 6. In practice, reaction 1 is conducted in excess of concentrated sulfuric acid (e.g., 96 wt %) such that the concentration of sulfuric acid in a single pass will not drop significantly and a rapid reaction rate is maintained.5 However, for the convenience sake of calculation, we assume that the initial H2SO4/H2S ratio is 1.5 and, thus, the resultant sulfuric acid concentration is the same as that produced from reaction 8, i.e., 1:4 H2SO4/H2O or H2SO4 in 57.7 wt %. The temperature of the reaction is at 393 K or higher to keep the produced elemental sulfur above its melting point for good flowability. Reaction 4 is specified at a temperature of 698 K, at which the self-sustainable ignition of elemental sulfur takes place.6 The conditions of the reactions or steps that are shared with the S-I cycle are kept the same as specified in refs 3 and 4. It is noticed that Goldstein et al.4 used 400 K for the Bunsen reaction (reaction 10) in their calculation. To maintain the consistency, we use 400 K for this reaction as well. Therefore, the detailed process of the H2S splitting cycle based on reactions 1-3 is described as follows:

(3) Vitart, X.; Le Duigo, A.; Carles, P. Energy ConVers. Manage. 2006, 47 (17), 2740–2747. (4) Goldstein, S.; Borgard, J. M.; Vitart, X. Int. J. Hydrogen Energy 2005, 30, 619–626.

(5) Wang, H.; Dalla Lana, I. G.; Chuang, K. T. Ind. Eng. Chem. Res. 2004, 43, 5846–5853. (6) Muller, T. L. In Kirk-Othmer Encyclopedia of Chemical Technologies; Kroschwitz, J. I., Howe-Grant, M., Eds.; John Wiley and Sons, Inc.: New York, 1997; Vol. 23, pp 363-408.

where g and l in parentheses represent gas and liquid states, respectively, and subscripts L, L1, and L2 represent liquids in different phases and compositions, respectively. To conduct reaction 2 or the Bunsen reaction, excess H2O and I2 are used such that two immiscible aqueous phases that respectively contain products HI and H2SO4 are formed. The compositions of the two phases are indicated in reaction 8. The L1 phase is aqueous sulfuric acid, and the L2 is the mixture of hydrogen iodide (HI), iodine (I2), and H2O, namely, HIx. Steps 9 and 10 are for conducting reaction 3, which includes HI separation from the HIx solution and the subsequent decomposition. Similarly, steps 11-14 are for SO2 production from H2SO4 decomposition, which consists of the concentration of the aqueous sulfuric acid solution, the vaporization of liquid sulfuric acid, and the twostep decomposition.3,4 Goldstein et al.4 defined the thermal efficiency, ηth, of the S-I cycle as “the ratio of the enthalpy of the hydrogen and oxygen recombination reaction at ambient temperature and pressure to the total heat requirement of the cycle”. In fact, the enthalpy should take the absolute value in this definition because it is thermodynamically negative. The definition should mathematically be described as ηth )

|∆HH0 2O(Ta)| Q + W/ηr

2186 Energy & Fuels, Vol. 23, 2009

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process 1: H2S f H2 + S 16H2O (l) + 9I2 (l) + SO2 (g) f (H2SO4 + 4H2O)L1 + (2HI + 10H2O + 8I2)L2 (400 K) (8) L2 ) (2HI + 10H2O + 8I2)L2 f 2HI (g) + (10H2O + 8I2)L (500 K) (9) 2HI (g) f H2 (g) + I2 (l) (600 K)

(10)

H2S (g) + 1.5H2SO4 (l) f S + SO2 (g) + (0.5H2SO4 + 2H2O)L1 (393 K) (16) L1 ) 1.5(H2SO4 + 4H2O)L1 f 1.5H2SO4 (l) + 6H2O (l) (570 K) (17) And the detailed process of the H2S splitting cycle based on reactions 1 and 4-6 is described as follows: process 2: H2S + 2H2O + O2 f 2H2 + H2SO4 32H2O (l) + 18I2 (l) + 2SO2 (g) f (2H2SO4 + 8H2O)L1 + (4HI + 20H2O + 16I2)L2 (400 K) (18) L2 ) (4HI + 20H2O + 16I2)L2 f 4HI (g) + (20H2O + 16I2)L (500 K) (19) 4HI (g) f 2H2 (g) + 2I2 (l) (600 K)

H2S (g) + 0.5O2 (g) f H2O (l) + S

(21)

L1 ) 2.5(H2SO4 + 4H2O)L1 f 2.5H2SO4 (l) + 10H2O (l) (570 K) (22) Following the nomenclature of general atomics (GA),7 we also use three sections to divide the H2S splitting processes, as shown in Figure 1. Section 1 involved the Bunsen reaction, leading to two separate liquid phases: one is diluted sulfuric acid, and the other is the iodine/iodide solution. Section 2, the only section that differs from the one in the S-I water splitting cycle, includes the sulfuric acid concentration and the reaction between concentrated sulfuric acid and H2S, and in process 2, it also has elemental sulfur oxidation. In addition, section 3 is associated with the HI/I2 separation and decomposition. The major difference between processes 1 and 2 is the addition of the elemental sulfur oxidation in the latter, and as a result, the according reactions occur in different stoichiometric coefficients, which result in the production of 2 mol of hydrogen from 1 mol of H2S. 3. Upper Bound of the Thermal Efficiency of H2S Splitting The definition of thermal efficiency in thermodynamics is described as the fraction of the heat input that is covered to net work output.8 A more general definition of thermal efficiency is given by the ratio of energy sought to energy that costs.9 For a thermochemical cycle, such as the S-I cycle, the net work output should be replaced by the recoverable energy from the cycle, with the latter equal to the absolute value of enthalpy of the hydrogen and oxygen recombination reaction or the formation of water, because hydrogen alone is the energy carrier

(23)

For process 2, the final product is sulfuric acid and the reaction is described as H2S (g) + 2O2 (g) f H2SO4 (l)

(24)

Similar to the S-I cycle, the H2S splitting cycle recovers only hydrogen as the energy carrier. Thus, the recoverable energy is also equal to the absolute value of the enthalpy of H2 combustion. However, different moles of H2 are produced from processes 1 and 2. The recoverable energy should be the product of the number of H2, n, and |∆HH0 2O(Ta)|. Therefore, eq 15 is modified as follows:

(20)

H2S (g) + 1.5H2SO4 (l) f S + SO2 (g) + (0.5H2SO4 + 2H2O)L1 (393 K) (16) S + O2 (g) ) SO2 (g) (698 K)

recovered from the cycle. For the S-I cycle, the sole reactant, H2O, does not bring any energy input into the cycle. However, for other chemical cycles, such as the one that we are dealing with, the possible energy that the reactants bring into the cycle should be taken into account as part of the input energy. The amount of this energy in the H2S splitting processes is equal to the absolute values of the enthalpy of burning or oxidizing 1 mol of H2S into the corresponding final product with higher oxidization numbers at ambient T and P, i.e., ∆HH0 2S oxid(Ta). For process 1, the final product is elemental S. Thus, the oxidizing reaction is

ηth )

n|∆HH0 2O(Ta)| |∆HH0 2S oxid | + Q + W/ηr

(25)

where Q is the heat to facilitate the endothermic reactions and W is the work required by the reactions in the cycle and transporting the reactant streams to a certain pressure. The enthalpy of the hydrogen and oxygen recombination reaction, ∆HH0 2O(Ta), and the enthalpy of the H2S oxidation reactions 23 and 24, ∆HH0 2S oxid(Ta), were calculated using HSC Chemistry 5.11 version. Their values at 298.15 K are -286, -265, and -767 kJ, respectively. It has been noted that, when HSC is used, the state of any material, except that of solid elements, must be identified, as shown in reactions 23 and 24. In addition to gas (g), liquid (l), and solid (s), some liquid species must be specified as aqueous solution (a), positive or negative aqueous species (+a or -a), etc., that the calculation can be performed. We use all of the assumptions that Goldstein et al. have used in the calculation of the heat Q and work W required by the cycle: (1) all of the reactions are reversible; (2) the conversion efficiency of the system, ηr, is 0.5; (3) the heat used for heating the reactants and cooling the products can be balanced; (4) heat from exothermic reactions and work from depressing processes are ignored; and (5) phase changes are represented by reactions. With the assumption that all reactions are reversible, the following formula could be obtained according to basic thermodynamic laws:4 W ) ∆H - T∆S

(26)

(7) Norman, J. H.; Besenbruch, G. E.; Brown, L. C.; O’Keefe, D. R.; Allen, C. L. DOE/ET/26225, 1981. (8) C¸engel, Y. A.; Turner, R. H. Fundamentals of Thermal-Fluid Science; McGraw-Hill: Boston, MA, 2001; pp 221-222. (9) Van Wylen, G. J.; Sonntag, R. E. Fundamentals of Classical Thermodynamics, 2nd ed.; John Wiley and Sons, Inc.: New York, 1976; p 170.



∆H ) Q + W

(27)

Q ) T∆S

(28)

Thus, the heat Q and work W requirements of each reaction are Q ) T∆S ) ∆H - ∆G

(29)

W ) ∆H - T∆S ) ∆G

(30)

For an exothermic reaction, there is no heat required. In addition, for the case where the system produces work, there is no work needed. Therefore, the requirement of the total energy for heat (Q) and work (W) for each reaction is ∆G + (∆H - ∆G), ηr

if ∆H - ∆G > 0

or

Figure 1. Sketch of the H2S splitting processes.

(31)

∆G , ηr

if ∆H - ∆G < 0

(32)

3.1. Thermodynamic Data for Section 2. Only reactions 16 and 21 are new and have not been discussed in ref 4. Reaction 16 can eventually be deemed as the addition of a chemical reaction H2S (g) + H2SO4 (l) f S + SO2 (g) + 2H2O (l)

(33)

and the dilution of sulfuric acid in water 0.5H2SO4 (l) + 2H2O f (0.5H2SO4 + 2H2O)L1

(34)

Using HSC Chemistry version 5.11, the thermodynamic data of reactions 33 and 21 at a broad temperature range were calculated and are reported in Figures 2 and 3, respectively. These figures indicate that, in the considered range of temper-


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Figure 2. Thermodynamic date for reaction 33: H2S (g) + H2SO4 (l) f S + SO2 (g) + 2H2O (l).

Figure 3. Thermodynamic data for reaction 21: S + O2 ) SO2 (g). Table 1. Summary of the Thermodynamic Data for the Involved Reactions in Process 1 section

reaction number

1

8

2

17 33 34 9 10

3

temperature (K) ∆H (kJ) ∆G (kJ) Q ) ∆H - ∆G (kJ)

reaction 16H2O (l) + 9I2 (l) + SO2 (g) f (H2SO4 + 4H2O)L1 + (2HI + 10H2O + 8I2)L2 L1 ) 1.5(H2SO4 + 4H2O)L1 f 1.5H2SO4 (l) + 6H2O (l) H2S (g) + H2SO4 (l) f S + SO2 (g) + 2H2O (l) 0.5H2SO4 (l) + 2H2O (l) f (0.5H2SO4 + 2H2O)L1 L2 ) (2HI + 10H2O + 8I2)L2 f 2HI (g) + (10H2O + 8I2)L 2HI (g) f H2 (g) + I2 (l)

400

-93

-72

-21

570 393 400 500 600

87 -60 -29 122 -24

60 -43 -20 77 34

27 -17 -9 45 -58

Table 2. Summary of the Thermodynamic Data for the Involved Reactions in Process 2 section

reaction number

1

18

2

22 33 34 21 19 20

3

reaction

temperature (K)

∆H (kJ)

∆G (kJ)

Q ) ∆H - ∆G (kJ)

32H2O (l) + 18I2 (l) + 2SO2 (g) f (2H2SO4 + 8H2O)L1 + (4HI + 20H2O + 16I2)L2 L1 ) 2.5(H2SO4 + 4H2O)L1 f 2.5H2SO4 (l) + 10H2O (l) H2S (g) + H2SO4 (l) f S + SO2 (g) + 2H2O (l) 0.5H2SO4 (l) + 2H2O (l) f (0.5H2SO4+2H2O)L1 S + O2 (g) ) SO2 (g) L2 ) (4HI + 20H2O + 16I2)L2 f 4HI (g) + (20H2O + 16I2)L 4HI (g) f 2H2 (g) + 2I2 (l)

400

-186

-144

-42

570 393 400 698 500 600

145 -60 -29 -306 244 -48

100 -43 -20 -299 154 68

45 -17 -9 -7 90 -116

ature, both reactions are spontaneous and exothermic. In particular, at the specified reaction temperatures, ∆H and ∆G are -60 and -43 kJ, respectively, for the reaction between H2S and sulfuric acid at 393 K and ∆H and ∆G are -306 and -299 kJ, respectively, for S oxidation at 698 K. Both of them meet

the condition ∆H - ∆G < 0 (eq 32). Figures 2 and 3 indicate that this condition will be satisfied in a broad range of temperatures. Therefore, the total heat requirement for these two reactions is zero (Q + W/ηr) ) ∆G/ηr ) 0) for all T < 1000 K. Goldstein et al.4 estimated the values of ∆H and ∆G for H2SO4



Table 3. Upper Bound of Thermal Efficiency for Hydrogen Production from H2S Splitting process 1 reaction number

Q + W/ηr (kJ)

section 1 section 2

8 17 33 34

0 147 0 0

section 3

9 10

199 68

pumping W/ηr ∑(Q + W/ηr) |∆HH0 2S oxid(Ta)| n|∆HH0 2O(Ta)| ηth

16.7 431 265 286 0.41

used. Therefore, the values of the upper bound of thermal efficiency of processes 1 and 2 are 0.41 and 0.33, respectively.

process 2 reaction number

Q + W/ηr (kJ)

18 22 33 34 21 19 20

0 245 0 0 0 398 136 33.4 812 767 572 0.36

dissolution in water at T ) 400 K and P ) 2 bar. Computed by Engels strong acid model, ∆H and ∆G are -58 and -40 kJ, respectively, for the process of 1 mol of H2SO4 dissolving in 4 mol of water. They supposed that temperature does not have a significant effect on the heat requirement of the dissolving process. Therefore, these data can be used for reaction 34 at 393 K, the sulfuric acid dilution, or dissolving to form a H2SO4/ H2O ) 1:4 solution, and reaction 22, the opposite process, reactions 17 and 22 at 570 K. The data were based on 1 mol of H2SO4. We recalculated them according to the corresponding stoichiometric coefficients of reactions 17, 22, and 34. 3.2. Thermodynamic Data for Sections 1 and 3. All of the reactions in sections 1 and 3 are involved in the S-I cycle, and their thermodynamic data were determined by Goldstein et al.4 In this paper, we use their data for these reactions. The data of ∆H, ∆G, and Q, which is equal to ∆H - ∆G, for the reactions and separation or dissolution steps in the three sections are shown in Tables 1 and 2. 3.3. Pumping Work. Goldstein et al.4 just considered the pumping work for the processes where pressure increase is needed. For instance, SO2 was produced at 1 bar, but the Bunsen reaction was supposed to take place at 2 bar. The HI separation and decomposition were carried out in a reactive distillation column at 50 bar. Therefore, SO2 would be compressed from 1 to 2 bar at a constant temperature of 400 K, and the heavier liquid phase of the Bunsen reaction product, the solution of HI, I2, and water in the ratio of 1:4:5, would be pumped from 2 to 50 bar. With the assumptions of ideal gas and uncompressible liquid and the stoichiometry specified for the S-I cycle in ref 4, the work required for the two compressing processes were 2.25 and 4 kJ, respectively. Considering mechanical efficiency of 0.75, the total work required was 8.33 kJ. The stoichiometry of process 1 indicates that the handling rates of SO2 and L2 are the same as in the S-I cycle and the total work is also the same, 8.33 kJ. However, process 2 will handle 2 times SO2 and L2, and the required work would be doubled as well. Thus, the pumping or compressing work is 16.7 kJ. Table 3 summarizes the results of calculations and shows the upper bound of thermal efficiency for processes 1 and 2 of the H2S splitting for hydrogen production. The Q + W/ηr for reactions is calculated by eqs 31 and 32 using ηr ) 0.5. For reactions 8, 18, 21, 33, and 34, where both ∆H and ∆G are negative, Q + W/ηr ) 0; for reactions 7, 17, 19, and 22, where both ∆H and ∆G are positive, eq 31 is used; and for reactions 10 and 20, where ∆H is negative but ∆G is positive, eq 32 is

4. Discussion 4.1. Assumptions and Simplifications. It has shown that the calculation of the upper bound of thermal efficiency of hydrogen production from H2S splitting is based on many assumptions and simplifications. First of all, the reactions and separations are assumed as reversible processes. In practice, however, this is not achievable. Therefore, the heat and work requirements are Q < T∆S or Q < ∆H - ∆G and W > ∆H - T∆S or W > ∆G. Second, the calculation ignored the difference of heat capacities between reactants and products, assuming that the energy can be balanced when heating the reactants from ambient temperature to the reaction temperature and cooling the products the other way. Third, no phase change is assumed to take place during heating the reactants or cooling the products because, in the calculation phase, changes or separations are treated as independent reactions, such as eqs 19 and 22. Lastly, the energy needed for transporting materials in the cycle, except the energy for increasing pressure, has not been considered. With these assumptions, the thermal efficiency calculated is the maximum value or upper bound and the cycle can theoretically be achieved. Goldstein et al. used the Engels strong acid model to calculate ∆H and ∆G for the dilution of 1 mol of H2SO4 with 4 mol of water (reaction 34).4 The values of ∆H and ∆G at 400 K are -58 and -40 kJ, respectively. The value of ∆H is very close to the enthalpy change between two sulfuric acid concentrations estimated using the enthalpy chart for sulfuric acid concentration, -53.0 kJ at 394 K.10 However, HSC Chemistry gives different values if the dilution is described as H2SO4 (l) f H2SO4 (ia)

(35)

where ia means ionic aqueous solution. The values of ∆H and ∆G for reaction 35 at 398 K are -156.3 and -16.45 kJ, respectively. The possible reason for the difference could be that HSC Chemistry does not exactly define the concentration of H2SO4 (ia), which might be different from that of the solution of 1 mol of H2SO4 and 4 mol of H2O. In this paper, we use the data from Goldstein et al.4 In determining the thermodynamic data for the Bunsen reaction, Goldstein et al. assumed reaction 8 is the addition of the following four reactions: SO2 (g) + 2H2O (l) f H2SO4 (l) + H2

(36)

I2 (l) + H2 f 2HI (g)

(37)

H2SO4 (l) + 4H2O (l) f (H2SO4 + 4H2O)l

(38)

2HI (g) + 8I2 + 10H2O (l) f 2(HI + 5H2O)l + 8I2 (39) By adding the data for these four reactions together, they found that ∆H and ∆G for reaction 8 are respectively -93 and -122 kJ. They even doubted the correctness of this addition because it did not consider the ionization of iodine molecules in the reaction. When using HSC Chemistry, the Bunsen reaction, with the consideration of iodine molecule ionization, can be written as (10) Buecker, W. W.; West, J. R. The Manufacture of Sulfuric Acid; Reinhold Publishing Corp.: New York, 1956; p 9.


Yu et al.

3I2 + SO2 (g) + 2H2O (l) f 2H (+a) + 3I3 (-a) + H2SO4 (ia) (40) where +a and -a represent cationic and anionic species in aqueous solutions. The ∆H and ∆G of reaction 40 are -240.9 and -39.9 kJ, respectively. Despite different stoichiometry, reactions 8 and 40 represent the Bunsen reaction. It appears that the difference in the values of ∆H and ∆G of the two reactions is significant, which might have resulted from the different stoichiometry and/or depended upon whether the ionization of iodine molecules is taken into account. This difference does not seem to affect the heat and work result for section 1 because ∆H and ∆G are negative in both cases, and thus, the value of Q + W/ηr is zero anyway. However, how to specify the product solution from the Bunsen reaction must have an effect on the result of the thermal efficiency of the whole cycle. This is because solutions of HI, I2, and water in different compositions require different amounts of energy in the separation of HI. Because of the complexity of interactions among HI, I, and water, how to quantify the effect remains in a debate. In our calculation, we still use stoichiometry that Goldstein et al.4 specified, such as shown in reactions 9 and 19. 4.2. Improvement of Energy Efficiency for Hydrogen Production. When the enthalpy of H2S oxidations are included as part of the input energy, the upper bound of the thermal efficiency of the H2S splitting is 0.41 and 0.36, respectively, for processes 1 and 2. These numbers are lower than that of the S-I cycle of H2O splitting for H2, which is 0.51 according to Goldstein et al.4 Nonetheless, if only Q + W/ηr, which must come from external energy sources, is taken into consideration for the energy input to the cycle, the thermal efficiency or also called the energy efficiency of the H2S splitting cycle for H2 production is in effect higher, at 0.66 and 0.70 respectively, for processes 1 and 2. This indicates that to produce 1 mol of hydrogen, splitting H2S requires less external energy than splitting H2O. This result is obvious because hydrogen is less tightly bound in H2S than in H2O. To further improve the energy efficiency of the H2S splitting cycle for hydrogen production, one should consider the recovery of the reaction heat generated from the exothermic reactions involved, for example, the elemental sulfur oxidation (reaction 4 or 21). In sulfuric acid manufacture, liquid sulfur is transported from the source at 408-428 K, at which its viscosity is a minimum, and then is injected into the burner, which is preheated at 600-628 K, with which temperatures the ignition device can be avoided. In the burner, the S is oxidized into SO2 with air. The temperature of the exiting gas mixture (10-12 vol % SO2 plus air) from the burner can reach as high as 1250-1400 K.6 The mass balance for sulfur combustion with air shows that this gas mixture of 10 vol % SO2 will contain 1 mol of SO2, 2.68 mol of O2, and 6.32 mol of N2. Its enthalpy at 1250 K relative to 298 K is 318 kJ. The enthalpy data for gases are found from Matheson Gas Data Book.11 This hightemperature energy source can be used to support the reactions and separations in section 3, where heat of high quality is needed. Doing this makes the entire cycle more energy-effective. 4.3. Further Analysis on Section 2. The actual energy efficiency of a process depends upon the process design. The process design of sections 1 and 3 has been discussed in the S-I cycle elsewhere.4,12 This paper focuses on the flowsheet (11) Yaws, C. L. Matheson Gas Data Book; Matheson Tri-Gas: Parsippany, NJ, 2001; pp 599, 648, 759. (12) Brown, L. C.; Lentsch, R. D.; Besenbruch, G. E.; Schultz, R. E.; Funk, J. E. AIChE Spring National Meeting, Conference Proceedings, 2003.

Figure 4. Flowsheet of section 2.

of section 2. One of the designs according to process 2 is shown in Figure 4. The reaction between H2S and sulfuric acid is conducted in a packed column reactor in concurrent flow.5 The excess sulfuric acid is used. Because of the production of water in the reaction between H2S and sulfuric acid, the concentration of sulfuric acid is downgraded. To simplify the calculation, the stoichiometry is shown as in reaction 16. However, to maintain a sufficient reaction rate, even more sulfuric acid is used for a single pass, such that the resultant sulfuric acid concentration is no less than 92 wt %.5,13 If the feedstock of H2S is pure (separated from sour gas, a mixture of H2S and CO2), pure SO2 is generated. Otherwise, SO2 generated from the reaction will coexist with CO2. After separation from a three-phase settler, gas, downgraded sulfuric acid, and sulfur will be produced. This sulfuric acid, together with the even diluted one resulted from the Bunsen reaction (section 1), is sent for upgrading (concentrating in a distillation column). In addition, the elemental sulfur is sent to generate more SO2. If air is used as the oxidant, N2 and unreacted O2 will coexist with this SO2. This design has the flexibility of using external sulfur, which can be pumped into the burner as well as to generate more SO2 and, of course, more H2 at the end. The other alternative is to conduct the reaction between H2S and sulfuric acid and the S oxidation in the same reactor with the help of a catalyst. This design saves a unit operation of sulfur burner but, meanwhile, loses the high-temperature heat from it as well as the flexibility of using external sulfur feedstock. The third design is to simply open the sulfur loop of the cycle, take H2S from all of the sour gas as the feedstock, and burn it into SO2. However, by doing this, only 1 mol of hydrogen will be produced from 1 mol H2S. Hydrotreating consists of not only sulfur removal but also nitrogen and aromatics removal, which all consume hydrogen. To supply sufficient H2 for oil sands upgrading, splitting H2S from the hydrotreating alone is not enough. This alternative has to add H2S from other sources, such as natural gas plants. However, the advantages of this change are avoiding concentrated sulfuric acid circulation in the system to support H2S oxidation and the separation between sulfur and sulfuric (13) Wang, H.; Dalla Lana, I. G.; Chuang, K. T. Ind. Eng. Chem. Res. 2002, 41, 6656–6662.


acid. The other benefit is that the diluted sulfuric acid from the Bunsen reaction does not necessarily need upgrading if such a concentration can find applications elsewhere. 5. Conclusions The H2S splitting cycle for H2 production was recently proposed to meet the needs of a sustainable oil sands development. As one of the initiations of technology development based on the chemistry, an upper bound of the thermal efficiency of the cycle was determined according to the two chemical routes, processes 1 and 2. When the enthalpy of the H2S oxidation reactions is included as part of the energy input to the cycle, the values for processes 1 and 2 are 0.41 and 0.36, respectively.


However, if only external energy that meets the requirement of heat and work for reactions and pumping is taken into account, the values of thermal efficiency can be higher, at 0.66 and 0.70. The analysis indicates that separations of H2SO4 and HI from their respective aqueous solutions generated from the Bunsen reaction are the most energy-intensive. Acknowledgment. The authors acknowledge the financial support from the Natural Science and Engineering Research Council of Canada (NSERC) with a Strategic Grant and the beneficial discussion with Dr. Jean-Marc Borgard of Commissariat a` l’Energie Atomique (CEA) in France. EF800895V

Upper Bound for the Efficiency of a Novel Chemical Cycle of H2S

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