Ind. Eng. Chem. Res. 1994,33, 1090-1097
1090
Hydrogen Recovery from Hydrogen Sulfide by Oxidation and by Decomposition Barry L. Yang and Harold H. Kung' Ipatieff Laboratory and Chemical Engineering Department, Northwestern University, Evanston, Illinois 60208
Selective oxidation of hydrogen sulfide t o hydrogen and sulfur oxides in a two-step process and the catalytic decomposition of hydrogen sulfide were studied for the recovery of hydrogen from hydrogen sulfide. Platinum, when adequately dispersed on a silica support, was found to be effective in reacting with hydrogen sulfide to produce hydrogen and platinum sulfide at 500 "C. The platinum sulfide could then be treated with oxygen at 400 O C to release sulfur oxides and regenerate the platinum. However, oxidation of sulfur dioxide t o trioxide, retention of oxygen by platinum, and adsorption of hydrogen sulfide by the silica support also occurred, which resulted in a minor loss in hydrogen yield. In the decomposition of hydrogen sulfide, platinum sulfide was found to be catalytically active. T h e equilibrium hydrogen yields were measured experimentally over the range 350-650 "C and compared with the values calculated on the basis of a model that included the S, allotropes, the H2S, sulfanes, and HS. The values agreed well at low temperatures but deviated from each other up to 20% a t high temperatures.
Introduction Hydrogen is consumed in large quantities in the oil refining process for sulfur removal. The hydrogen sulfide generated in the hydrodesulfurization (HDS) process is then reacted with air to elemental sulfur and water in a Claus plant. Currently, hydrogen is supplied primarily by steam reforming of light hydrocarbons, which is an energy intensive process. With increasing demand for cleaner fuels, a higher degree of sulfur removal from the fuel is expected in the future. This, together with the push to utilize the high molecular weight portion of oil, and the need to reduce aromatic content in gasoline, would further increase the demand for hydrogen in the future. One method to reduce this demand is to recover hydrogen from hydrogen sulfide. In principle, hydrogen can be recovered by decomposing hydrogen sulfide. A typical reaction would be H,S e H,
+ '/,S,
(1)
although other sulfur allotropes, sulfanes, and HS can all be formed. However, the decomposition of H2S is thermodynamically highly unfavorable. Thermodynamic calculation given in Appendix A suggests that the equilibrium hydrogen yields would be only 0.17,1.0, and 4.8% at 600, 800, and 1000 K, respectively. Various methods have been investigated to increase the yield of hydrogen. One approach is to remove selected products from the reaction mixture. Selective sulfur removal has been demonstrated by adding a cold trap to a circulating reaction system to continuously remove sulfur from the gas phase (Kotera et al., 1976; Fukuda et al., 1978). Using this method, a conversion higher than 95 % was achieved at a reactor temperature of 800 OC. However, the temperature difference between the reactor and the cold trap created heavy loads for heating and cooling,which easily upset the economics of the process. Selective hydrogen removal by employing a membrane reactor has also been proposed (Raymont, 1975)and tested (Edlund and Pledger, 1993). A conversion of 99.4% from a batch of 1.5%H2S in nitrogen was reported by applying apressure of8 atmacrossaPt/SiOz/V/SiOz/Pdmembrane
* To whom correspondence should be addressed.
at 700 "C. However, the membrane technology for hydrogen sulfide decomposition still needs further development. In addition to the limitations on membrane permeability and selectivity, the corrosive character of hydrogen sulfide to the membrane material is a potential problem. Other methods like pressure swing adsorption (Bandermann and Harder, 1982) and thermal diffusion (Chivers and Lau, 1987) have also been studied, but they are not yet ready for commercialization. Another approach is to couple reaction 1with another reaction such that the overall reaction becomes thermodynamically favorable. Earlier proposals involve reaction of hydrogen sulfide with carbon monoxide (Kotera, 1976) or with methane (Raymont, 1975):
+ CO + '/,O,e H, + '/,S, + CO, 2H,S + CH, + 0, s 4H, + S, + CO,
H,S
(2)
(3)
There are no reports indicating success of either method. In the former case, the undesirable side reaction
2cos s CO, + cs,
(4)
has been identified as the major problem (Kotera, 1976). Yet another approach is to selectively oxidize hydrogen sulfide to sulfur dioxide and hydrogen: H,S
+ 0,
H, + SO,
(5)
This reaction is thermodynamically very favorable, with AH~K and A G ~ Kbeing -272.9 and -252.6 kJ/mol, respectively (Barin, 1989). For this process to be technically viable, a selective catalyst is needed that oxidizes sulfur preferentially to hydrogen. Also, since sulfur dioxide is not marketable, means of on-site conversion of sulfur dioxide are needed to convert it into a marketable product, such as sulfuric acid. The stringent requirement of the selective catalyst for reaction 5 can be relaxed somewhat if the reaction is carried out in two separate stages to avoid direct contact of H2 with 02.In the first stage, HzS is reacted with a reagent to produce hydrogen. The reagent is then regenerated by reaction with oxygen to release SO,: 0 1994 American Chemical Society
Ind. Eng. Chem. Res., Vol. 33, No. 5, 1994 1091
+ M-S + 0 2
H2S M
+
H2 + M-S
(6)
M + SO,
(7)
+
Such a two-step approach has previously been studied by Kiuchi et al. (1982) using silver and molten lead as the reagent. Both metals reacted with hydrogen sulfide effectively to form hydrogen and the corresponding sulfides. However, difficulty was encountered in the regeneration step. Because the formation of sulfate is thermodynamically more favorable than reaction 7 for both silver and lead, regeneration by oxidation could not be accomplished effectively. This paper reports our investigation of this two-stage operation using platinum as the reagent. Platinum was chosen because its reaction with H2S to form PtS is thermodynamically favorable up to 900 "C, its oxide decomposes at relatively low temperatures, and there is no stable sulfate of platinum known at the temperatures of interest. The oxidation of PtS to SO2 and Pt is also thermodynamically very favorable with a AG~OOK of -235.4 kJ/mol (Barin, 1989). The results of the reaction of unsupported and silica-supported Pt with hydrogen sulfide and the regeneration by reaction with oxygen are described. Also studied was the long-puzzled-over equilibrium formation of hydrogen from the decomposition of hydrogen sulfide (Raymont, 1975; Chivers et al., 1980; Weil, 1983). Both experimental measurement and thermodynamic calculation of the equilibrium were performed in this study. The equilibrium hydrogen yields were measured experimentally under flow reaction conditions at temperatures between 350 and 650 "C. Possible experimental artifacts such as incomplete stoichiometric reaction of H2S with the catalyst and hydrogen formation from hydrocarbon impurities were considered. A reaction model was used for numerical evaluation of the equilibrium hydrogen yields that considered all three groups of decomposition products: sulfur allotropes, sulfanes, and HS. It was found that the measured and the calculated equilibrium values agree with each other at low temperatures. The deviation between the two increased with increasing temperature but did not exceed 20%.
Experimental Section The platinum black was from a commercial source (Alfa, 99.9 '% ). The 5 % Pt/SiOz was prepared by impregnating silica (Davison952,140-200 mesh, washed with 4 N "03) with an aqueous solution of tetraammineplatinum(I1) nitrate, followed by overnight drying at 100 "C and calcination at 250 "C for 1 h. The Pt dispersion as estimated from the sulfur coverage as surface sulfide formation was 3%. The 25% Pt/SiOz was similarly prepared. Iron oxide was made by calcining an iron oxide hydroxide powder at 300 "C for 4 h. The hydroxide was precipitated from an aqueous ferric nitrate solution with aqueous ammonia, followed by thorough washing and drying. The 5.8% Fe203/Si02 was prepared by incipientwetness impregnation of a Davison 62 silica gel (150/200 mesh) with an aqueous ferric nitrate solution, followed by drying and calcination at 400 "C for 12 h in flowing 20% Oz/He. Hydrogen sulfide (Aldrich, 99.5+ % ) was used as received. Its hydrocarbon contaminants were found to be C2 to Cq alkanes and alkenes at a total concentration of 0.04 vol %, with propane at 0.018% being the major component. A reaction system capable of performing flow reaction, pulse reaction, temperature-programmed reduction (TPR), and temperature-programmed desorption (TPD)was used.
A catalyst, after reaction, could be characterized by TPR without being exposed to the atmosphere. Hydrogen sulfide and oxygen (UHP) pulses could be introduced independently. For flow and pulse reactions, the catalyst was pretreated in a flowing mixture of 10% Hz in Ar at 150 "C for 10 min, and then in flowing He (UHP) or Ar (UHP) at 500 O C for 0.5 h. The He and Ar carrier gases were purified online by being passed through a reduced MnOJSiOz bed to remove oxygen and a molecular sieve bed to remove water. For the pulse reaction of hydrogen sulfide, H2S pulses (typically 4.2 X lo3 mol each) were carried by Ar (30 mL/ min) over the platinum solid reactant packed in a quartz U-tube reactor. The reactor effluent was directed to flow through a Porapak Q column (1/8 in. X 10 ft) at 100 "C for product separation and then to a thermal conductivity detector (TCD) where Hz, H2S, and H2O were quantified. For 0 2 pulse reaction (typically 4.2 X 10-5mol per pulse), He (30 mL/min) was used as the carrier gas, and the same gas chromatographic system was used. For flow reactions, undiluted H2S was used as the reactant. Hydrogen and H2S in the reactor effluent were analyzed by online gas chromatography using the same columns and the TCD, and Ar (30 mL/min) as the carrier gas. In some experiments, the TCD effluent was further directed to a flameionization detector (FID)for quantification of hydrocarbon impurities. TPR was conducted in a flow of 20 mL/min of 10% Hz/Ar, and TPD in a flow of 30 mL/min of He. For both TPR and TPD, the platinum sample was first purged with the carrier gas at room temperature for 2 h. The reactor was then heated at 20 "C/min from room temperature to 620 "C for TPD, or at 10 "C/min from room temperature to 1000 "C for TPR. The reactor effluent was directed to flow through a molecular-sieve trap to remove water and hydrogen sulfide before being detected by the TCD. Reduction of CuO (Aldrich, 99.999%) under the same TPR condition was used for calibration. Results Reaction of 0 2 with Pt. As prepared, the Pt/SiOz contained oxygen,although metallic platinum was the only phase identified by X-ray diffraction. The oxygen could be removed by treatment with H2 at room temperature, and no consumption of Hz could be detected by TPR afterward. Desorption of oxygen could be detected in TPD experiments from the 5% Pt/SiOz after different oxygen treatments, as shown in Figure 1. Oxygen desorbed as a broad peak beginning at 250 "C from samples treated in air or oxygen at 250 "C (curves 1A and 1B). After the TPD experiment, exposing the catalyst to a pulse of H2S at 500 OC did not result in detectable formation of water, which suggested that all of the oxygen had been desorbed. For the sample treated in 0 2 at 450 "C, the desorption peak was a t a much higher temperature (curve lC),and desorption was not complete at 620 "C, which was the end of the TPD experiment. Subsequent reaction of this sample with hydrogen sulfide resulted in formation of water. Although the oxygen content of the 450 "C 02treated sample was high, metallic platinum was the only phase detected by X-ray. Thus the platinum oxide phase was amorphous. From the amounts of oxygen desorbed, it was estimated that the stoichiometries of the 250 "C air- and the 250 "C 02-treated 5%Pt/SiOz samples were Pt00.36 and Pt00.47, respectively. The oxygen content of unsupported platinum black was below the detection limit. Treatment of any of these samples in 10% HdAr at 150 "C for 10 min followed by helium or argon purging at 500
1092 Ind. Eng. Chem. Res., Vol. 33, No. 5,1994 ,,
I
. .
v~
I
0
- --
A
200
400
600
DEGREE OF SULFURIZATION
i
I
1nr
Y
2
4
TEMP (C)
6
8 1 0 1 2
PULSE NUMBER
Figure 1. Oxygen desorption profiles during TPD of 5% Pt/SiOz treated in (A) air at 250 OC for 1 h, (B) additional 02 at 250 "C for 0.5 h, and (C) additional 02 at 450 OC for 0.5 h. The samples were heated at 20 OC/min to 620 OC. O C was found to be effective to reduce the sample to metallic platinum. Reaction of HzS with Pt. Production of hydrogen from the reaction of hydrogen sulfide with platinum (reaction 6) was studied in the pulse reaction mode on both the platinum black and the 5% Pt/SiOz. Figure 2 shows the results of three different experiments, plotted as the hydrogen production as a function of pulse number. The three sets of data showed the common feature that the yield of hydrogen was high in the first pulse, decreased with increasing pulse number, but maintained a low but constant value a t large pulse numbers. The high yields of hydrogen in the first few pulses were much higher than the calculated equilibrium yields for H2S decomposition, which would be 0.14,0.36, and 0.53 X 10-6mol at 400,500, and 550 O C , respectively (Table 1). These high yields of H2 indicated that platinum was highly reactive toward hydrogen sulfide. In fact, a substantial yield of H2 could be detected in the first pulse at as low as 100 "C,as a result of sulfurization of the platinum surface. Under these conditions, the silica support was inactive and there was no detectable hydrogen formation by pulsing H2S over Si02 up to 550 O C . The low but constant hydrogen yield in later pulses could be attributed to the catalytic decomposition of hydrogen sulfide (reaction 1). The yields were 0.2, 0.9, and 0.5 X 10-8 mol for the 5% Pt/SiOz at 400 "C,the 5% PtlSiO2 at 500 O C , and the platinum black at 550 O C , respectively. It is interesting that the yields on the 5 % Pt/SiOz were higher than the calculated equilibrium values (Table 1). Similar observation of hydrogen production in excess of equilibrium has long been reported (Raymont, 1975; Chivers et al., 1980;Weil, 1983). A chromatographic effect was suspected, as will be discussed later. With the assumption that the H2 yield due to catalytic decomposition was constant in every pulse, the degree of platinum sulfurization could be calculated from the cumulative hydrogen yield after correction for the portion due to catalytic decomposition. These values are shown in the figure as the percent of Pt being converted to PtS in helium, whose presence was confirmed by the X-ray diffraction (XRD) shown in Figure 3. The PtS formed from the 5% PtlSiO2 appeared to be thermally stable at the pulse reaction temperatures. Heating it at 500 OC in flowing helium for 13.5 h did not result in any indication of decomposition. It was found
DEGREE OF SULFURIZATION 100 %
"
2
4
6
8
1
0
PULSE NUMBER ~
MRXN (6) -
"1 3
4
3
3
62
'
- ..RRXXNN ( t )
9
Fz
5 w 0
2 2
DEGREE OF SULFURIZATION
1
1
I n "
2
4
6
8
10 12 14 16
PULSE NUMBER Figure 2. Production of hydrogen as a function of HgS pulse number: (A, top) 0.0108 g Pt black at 550 OC, (B,middle) 0.063 g 5% Pt/SiOz at 500 O C , and (C, bottom) 0.063 g 5% Pt/Si02 at 400 O C . The degrees of platinum sulfidation at the end of experiment were 30, 100, and 60% for A, B, and C, respectively. Each pulse contained 4.2 X 106 mol HzS.
that a temperature higher than 700 "C was needed to effectively decompose this PtS in helium, and the decomposition was completed in 30 min at 900 O C (Figure 3). The data in Figure 2A showed that only 30% of the platinum black could be converted to PtS rapidly a t 550
Ind. Eng. Chem. Res., Vol. 33, No. 5, 1994 1093 Table 1. Equilibrium Yield of
Hz from a HzS Pulse of 4.2 X
10* mol
~~~
temD (OC) 400 500 550
exDta 0.14 0.39 0.64
~~
equilibrium yield (1W mol) method Ab method Bb 0.14 0.06 0.36 0.24 0.53 0.42
a Calculated from equilibrium yield determined by steady-state reaction data. See text. b Method A used a comprehensive model; method B only considered SIto Ss. See Appendix A.
25
50
TWO THETA
Figure 3. XRD patterns of the 5% Pt/SiOz after different treatments (A) fresh; (B) after 500 OC H2S reaction; (C) after B and 400 O C 02 treatment; (D) after B and 500 "Chelium 13.5 h; (E) after D and 700 O C helium 0.5 h; (F)after D and 750 OC helium 0.5 h; (G) after D and 825 O C helium 0.5 h; (H) after D and 900 O C helium 0.5 h.
450
C
400 C
I 1
2
3
4
5
TIME (h)
Figure 4. Conversion of H2S to H2 in a flow reactor at various temperatures over 0.10 g of 25%Pt/SiOp. HzS flow rate 20 mL/min.
"C. Presumably, slow lattice diffusion prevented further sulfidation of the bulk platinum. The resistance to bulk sulfidation due to slow lattice diffusion was greatly reduced by dispersing the platinum on a silica support. As can be seen in Figure 2B, the 5% Pt/SiOz could be nearly completely sulfurized in three pulses at 500 "C. At 400 "C, complete sulfidation became difficult even for the dispersed platinum. The degree of sulfidation was only 60% after about 15 pulses (Figure 2 0 . Catalytic Decomposition of HzS over PtS and Fe&. The catalytic decomposition of HzS was further studied in a flow reactor. The results of a steady-state test are shown in Figure 4. In this experiment, HzS was passed over a 25% Pt/SiOz catalyst at 450 "C. The conversions were kept below the equilibrium conversions. After a steady state was attained, the temperature was changed. This was repeated a number of times after a new steady state was attained at each temperature. The activity of the catalyst at 450 "C was checked regularly. The data in Figure 4 show that a steady state could be reached rapidly at every temperature except at the
I
200
I
400
600
so0
1,000
TEMP (C)
Figure 5. Temperature-programmedreduction profiles of used catalysts (10% HdAr, 10 OC/min): (A) 5%Pt/SiO2 after 500 OC reaction; (B) unsupported Pt black after 650 OC reaction.
beginning of the experiment, when a higher conversion than steady state conversion was observed. This could be attributed to the formation of platinum sulfide at the beginning of the experiment. The data showed that sulfurization of the catalyst was completed in 1 h at 450 OC, and there was no indication of hydrogen production over the steady-state rate due to further stoichiometric reaction with Pt when the reaction temperature was raised to 500 and 550 OC. Also, there was no deactivation of the catalyst. The same conversion at 450 "C was obtained after the reactor temperature was changed to 500 and 550 "C and then returned to 450 "C. Similar results were obtained on the 5% Pt/SiOz, and on supported and unsupported iron oxide catalysts. However, water was formed at the beginning of the reaction over iron oxide catalysts. The bulk phase of the iron catalysts after the steady state was attained was found to be FeDSlo. Its XRD pattern was in agreement with the published standard (JCPDS File, 1990). The activity of Si02 was low, being less than 5% of that of Pt/SiOz a t 400 and 450 "C. For platinum black, sulfurization was not completed in 1h a t 450 "C. Raising the temperature to 500,550, and 600 "C all resulted in extra hydrogen production before the establishment of a new steady state. It was found that a reaction temperature of 650 "C was needed to complete the sulfurization in 1h. However, sintering was observed at that temperature. For both supported and unsupported Pt catalysts, platinum monosulfide (PtS)was the only phase observed in XRD after reaction. There was no metallic platinum or platinum disulfide (PtSz) phase detected. The sulfur contents of the sulfurized Pt catalysts were further determined by TPR. Figure 5 gives two typical TPR spectra, indicating a peak maximum at 570 "C for the sulfurized 5% Pt/SiOz and a broad and asymmetric peak for the unsupported catalyst. For the unsupported catalyst, the lack of a well-defined peak maximum and the abrupt termination of reduction suggested complex reaction kinetics that was most likely influenced by mass transport. After TPR, the catalysts showed only the metallic platinum phase in XRD. From the area under the TPR peaks, the sulfur to platinum molar ratio was found to be 1.0 f 0.05 for both silica-supported and unsupported catalysts. This again suggests the stoichiometry of PtS, in agreement with XRD results. An attempt to determine the S/Pt stoichiometry by oxidation was unsuccessful because of the formation of both SO2 and ,903, and the difficulties in quantifying the latter species. The equilibrium hydrogen yield under the flow reaction conditions was also measured. The constant yields at high W / F (weight of catalyst to reactant flow rate), as shown in Figure 6, were assumed to be the equilibrium yields. In this figure, the contribution to the Hz yield due to the decomposition of hydrocarbon contaminants has been corrected, since some cracking of CS to C:! occurred as
1094 Ind. Eng. Chem. Res., Vol. 33, No. 5, 1994
i
!-
-
A J
pulse no. temp (OC) SO2 (1o-Smol) (a) Pt black, 0.0108 g, 30% sulfurized (16.7 by H2S at 550 "C
650
-tf
yc,
Table 2. Oxidation of Sulfurized Platinum by (4.2 x 1od mol 0 s DB? D U h )
1 2 3 4 5 6 7 8 9 10 11
550
-=
500
350
375 400 450
X
0 2
Pulses
H20 (1o-S mol) 1W mol PtS)
0.30 0.54 0.73 0.81 0.85 1.63 1.50 1.39 3.08 1.12 0.16 12.1
total (b) 5% Pt/SiOz, 0.030 g, 100% Sulfurized (7.7 X 1W mol PtS) by HIS at 500 O c a "
0
40
20
60
W/F (g min/l) Figure 6. Hydrogen yield as a function of WIF for equilibrium determination: (A) 5% PtlSiOz; (B)FezOs; (C) 5.8% FezOs/SiOz.
1 2 3 4 5 total
400 450 500 550
3.81 0.03 0 0 0 3.84
0.45 0 0 0 0 0.45
(c) 5% Pt/SiOz, 0.063 g, 100% Sulfurized (16.2 X 1o-Smol PtS) by H2S at 500 O C *
4 7 I
1 2 3 4 5 6 7
100
200 300 400 500
0.10 0.03 0.12 0.04 0.33 6.34 0.06 7.02
0 0 0 0 0 2.92 1.75 4.67
3.16 3.95 1.05 0 8.16
0.62 1.08 0.10 0 1.80
total (d) 5% Pt/SiOz, 0.063 g, 55% Sulfurized (8.9 X 1o-S mol PtS) by HzS at 400 O C C 1 2 3 4
total
01 -----,
500
400
*
a Sample purged at 400 O C before 0 2 puleing. Sample exposed to HzS at 400 "C and purged at 100 "C before 02 pulsing. Sample purged at 350 O C before 02 pulsing.
k;' 300
350
700
TEMP (C) Figure 7. Measured and calculated equilibrium hydrogen yields for the decomposition of H2S: (A) measured; (B)calculated on the basis of a complete model of reaction; (C) calculated on the basis of Sj,j = 1-8 only.
shown by gas chromatography. They represented less than 0.03 and 0.15% hydrogen yield at 400 and 600 OC, respectively. The thermocouple and temperature controller/display used in equilibrium measurements were calibrated on a Setaram DSC 111 calorimeter against the melting points of In, Zn, and A1 at 156.6,419.6,and 660.4 "C, respectively. The error in temperature determination was less than 2 O C within the temperature range of reaction study. As can be seen in Figure 6, the sulfurized iron catalyst had a higher activity than the platinum catalysts at temperatures below 550 "C. However, it deactivated rapidly at 600 OC and above. Dispersing either platinum or iron oxide on silica resulted in satisfactory catalysts for high-temperature studies. The measured equilibrium hydrogen yields are plotted in Figure 7 (curve A), along with those (curves B,C) calculated from thermodynamic data (Appendix A). Oxidation of PtS. Oxidation of the sulfurized platinum was also studied in the pulse mode using 0 2 as the oxidant.
Sulfur dioxide was the only detected product from sulfurized platinum black. In particular, no water was detected. Results in Table 2a show that the oxidation reaction proceeded a t a noticeable rate a t 350 O C on sulfurized platinum black. The rate of reaction increased with increasing temperature. No detectable reaction was observed below 300 "C. In contrast, oxidation of the sulfurized Pt/SiOz produced sulfur dioxide and water (Table 2b-d). Since water was not detected from the unsupported samples, hydrogen sulfide adsorbed by silica, either molecularly or dissociatively, was suspectedto be the source of hydrogen for water formation. This argument was supported by the observation that the higher the purging temperature before pulse reaction with 0 2 the lower the water formation(Table 2c,d). High-temperaturepurging would reduce the amount of adsorbed hydrogen sulfide on the sample. The presence of adsorbed hydrogen sulfide on silica complicated the quantification of sulfur dioxide formed from platinum sulfide. The adsorbed hydrogen sulfide could be oxidized to either sulfur dioxide or elemental sulfur, and the latter could not be detected in our system. One should note that the presence of adsorbed H2S on silica did not invalidate the conclusion on the degree of sulfidation of Pt based on the data in Figure 2, for which the hydrogen production was measured.
Ind. Eng. Chem. Res., Vol. 33, No. 5, 1994 1095 Despite the uncertainty in the quantification of sulfur dioxide formed by the oxidation of PtS/SiOz, the data in Table 2b-d clearly show that there was less sulfur dioxide produced than that expected from complete conversion of PtS. For the unsupported sample, the sulfur dioxide production was about 75% of the theoretical value. For the supported ones, the value could be lower than 50% (Table 2c), even if one assumed that all sulfur dioxide detected was from the oxidation of platinum sulfide. However, the low value of SO2 production could not be due to incomplete oxidation of PtS/SiOz because when an 02-regenerated sample, such as those in Table 2, was used without further treatment for hydrogen production from H2S pulses, the total amount of hydrogen products (H2 and H20) detected was identical to that for a fresh sample. The amount of water observed was small, but ita presence indicated that some oxygen was held by the sample after oxygen regeneration. This oxygen might be in the form of Pt oxide or adsorbed oxygen. Water formation from the 400 "C 02-regenerated 5% Pt/SiOz represented about 6 % of the hydrogen-containing producta. Also, as revealed by X-ray diffraction, the PtS phase disappeared totally after 0 2 regeneration (Figure 3). Metallic platinum was the only phase observed. Thus the low value of SO2 production was attributed to the production of SO3 that was not detected.
Discussion The strong affinity of platinum toward hydrogen sulfide observed in the pulse reaction study (Figure 2) agrees well with literature reports (Bartholomew et al., 1982; Bonzel and Ku, 1973;Fischer and Kelemen, 1977;Koestner et al., 1986). Hydrogen sulfide, even at a very low concentration in the ppm range, exhibits a strong poisoning effect to many metal catalysts (Bartholomew et al., 1982). The sticking coefficientof hydrogen sulfide on a Pt(ll0) surface has been found to be close to unity throughout the range of surface sulfur coverage from 0 to 0.6 (Bonzel and Ku, 1973). Upon adsorption on a Pt(ll1) surface at 110 K, hydrogen sulfide was found to decompose into HS, S, and H species (Koestner et al., 1986). Furthermore, complete hydrogen evolution occurred below 100"C, suggesting that hydrogen sulfide dissociation on platinum is rapid and complete at 100 "C. Although reaction of H2S with the Pt surface is rapid, complete sulfidation of a Pt particle is controlled by lattice diffusion. Dispersing the platinum greatly facilitates the sulfidation process. Thus complete conversion of Pt to PtS could be achieved at 500 "C with the 5% Pt/SiOZ sample, but not with the Pt black sample even at 550 "C (Figure 2). Similar to the sulfurization of other transition metals (Bartholomew et al., 1982), diffusion of metal cations (Pt2+in this case) through the sulfide layer is considered the transfer mechanism. The S2-anion (0.184nm radius) is much larger than the Pt2+cation (0.080-nm radius), such that its contribution to the diffusion process would be small. Under the conditions studied here for the catalytic decomposition of hydrogen sulfide, platinum monosulfide was identified by XRD and TPR as the bulk phase of the catalyst. The fact that platinum converts to platinum monosulfide completelyupon exposure to hydrogen sulfide at high temperatures (450 "C for the supported and 650 "C for the unsupported catalyst) suggests that it is impractical to use platinum as either an electrode material, as was suggested in the study of the effect of non-Faradaic electrochemical modification of catalytic activity (NEMCA) (Alqahtany et al., 19921, or a membrane material for
the decomposition of hydrogen sulfide (Edlund and Pledger, 1993). In the electrocatalytic applications, the electric conductivity of platinum would be lost. In the membrane applications, the integrity of the platinum membrane would not be maintained. The equilibrium yields of hydrogen determined from the steady-state reaction measurements were higher than those calculated from published parameters (Figure 7). A similar observation has been reported over a wide range of temperatures on various catalysts (Raymont, 1975; Chivers et al., 1980; Weil, 1983) but was not understood. In one extreme case, a hydrogen yield of 3.6 % was reported over cobalt molybdate at 750 K (Weil, 1983) in contrast to the predicted 0.66%. Possible causes include experimental artifacts and inappropriate thermodynamic calculation. One common experimental artifact is the hydrogen formation from the stoichiometric reaction of H2S with catalyst. This possibility can be excluded if a true steady state reaction measurement is made (Figure 4). The fact that consistent hydrogen yields were obtained over different catalysts, as shown in Figure 6, further suggested that true equilibrium was reached. Another experimental artifact is the release of hydrogen from hydrogen-containing impurities in H2S, such as hydrocarbons. Hydrocarbons are common impurities found in commercial H2S. Because the equilibrium conversion of H2S decomposition is low, the contribution from impurities may not be negligible. Such a contribution from hydrocarbon impurities was monitored and corrected in this study. The calculation of the thermodynamic equilibrium composition in H2S decomposition involves the consideration of a number of reactions. Although the decomposition to S2 and H2 (reaction 1)has often been used as the representative reaction, other sulfur allotropes, sulfanes, and HS are formed also. The importance of including sulfur allotropes and sulfanes in the equilibrium calculation was first recognized by Raymont (1974). Kaloidas and Papayannakos (1987) performed a calculation considering the formation of S,, n = 1-8, in a temperature range of 750-1450 K. Berk et al. (1991) extended the calculation to include HzS,, n = 2-5, from 800 to lo00 K. Acalculation that considers all three groups of decomposition products, S, allotropes, H2S, sulfanes, and HS, is given in Appendix A. As is shown there and in Figure 7, there is a significant difference, especially at low temperatures, between whether or not sulfanes (H2S2 in particular) are included. The contribution of HS formation, however, is negligible at temperatures lower than 1000 K. The difference between calculated equilibrium hydrogen yields, depending on whether the decompositionto sulfanes was included in the calculation, could explain in part the literature reports on the production of hydrogen in excess of equilibrium amounts. However, when a more comprehensive reaction model was used, the calculated equilibrium hydrogen yields agreed well with the measured ones, especially at low temperatures. Thus, higher hydrogen yields reported in the literature than these values were likely due to experimental artifacts. The hydrogen yields observed in the pulse reactions at large pulse numbers (Figure 2b) were higher than the equilibrium values, Possible causes for this were the partial decomposition of PtS between pulses, partial oxidation of PtS to Pt and SO2 by 0 2 impurity in the carrier gas, and the local separation of reaction products due to chromatographic effects. The decomposition of PtS between pulses was not likely to be the cause, because a temperature
1096 Ind. Eng. Chem. Res., Vol. 33, No. 5, 1994 Table 3. Calculated Equilibrium Composition for the Decomposition of HsS at 1 atm species
600 K 0.9967 0.1678 X le2 0.2947 X 0.1645 X 0.3501 X 1od 0.2512 X 0.7795 X 10-18 0.1196 x lo-' 0.2054 X 10-8 0.2183 X le7 0.3842 X 0.7416 X 0.7078 X lo" 0.2791 X lo"
700 K 0.9918 0.4264 X 1k2 0.1956 X 0.3655 X le2 0.7973 X 106 0.5493 X lk7 0.1768 X lO-" 0.2852 X 10-3 0.5605 X 106 0.5633 X 10-8 0.2739 X 10-8 0.3573 X 10-8 0.3838 X lk7 0.8872 X lo"
higher than 700 "C was needed for the effective decomposition of PtS in He (Figure 3). The possible effect of 0 2 contaminant in the carrier gas was also ruled out because its maximum contribution would be less than 1%of the observed hydrogen yield. Therefore, the overequilibrium formation of HZ from H2S decomposition in the pulse reaction mode is most likely due to a chromatographic effect that caused local separation of the products and therefore the shift of equilibrium. Indeed, overequilibrium production of Hz from H2S decomposition has been demonstrated by applying pressure swing adsorption (Bandermann and Harder, 1982). When the sulfurized platinum was oxidized, less sulfur dioxide was produced than expected on the basis of complete conversion of PtS to Pt and SO,. There are several possible reasons for this low sulfur dioxide production. First, the oxidation of PtS might be incomplete because of slow bulk diffusion. Second, the oxidation might be complete, but other products might be produced in addition to sulfur dioxide, such as sulfur trioxide that was not detected. The possibility of thermal decomposition of some of the platinum sulfide could be excluded because of its thermal stability, as discussed earlier. It is unlikely that slow lattice diffusion was the main reason. The higher conversion to sulfur dioxide observed for the unsupported samples than for the supported samples argues against lattice diffusion limitation (Table 2). The fact that little additional sulfur dioxide was produced by increasing the oxidation temperature beyond 400 "C further supports the contention that the oxidation of platinum sulfide was complete, as does evidence by XRD (Figure 3) and the reaction tests, as reported in the Results. Therefore, the nonstoichiometry in sulfur dioxide production is explained by the formation of other sulfurcontaining products, most possibly sulfur trioxide. After all, platinum is known to be catalytically active for the oxidation of SO2 to SO3 (Twigg, 1989). It is interesting that the oxidation of PtS proceeds readily at 400 "C. Oxidation has been studied as a regeneration method for sulfur-poisoned metal catalysts. However, the removal of sulfur could be hindered by the formation of an oxide layer on nickel catalysts (Bartholomew et ai., 1982). The low tendency of platinum in oxide formation is likely the reason that the oxidation of platinum sulfide is effective. However, it was also observed that some oxygen was retained by the platinum after the oxidation regeneration. This is likely due to the high degree of coordination unsaturation of platinum at the surface (McCabe et al., 1988). It is demonstrated in this study that platinum, when dispersed on a support, can function as a reagent for hydrogen recovery from hydrogen sulfide in a two-step operation. The reaction of platinum with hydrogen sulfide
800 K 0.9818 0.1025 X 10-I 0.4204 X 10-8 0.5665 X 1W2 0.1212 x lo-' 0.8469 X le7 0.4968 X 10-l2 0.2210 x 10-2 0.4088 X lP 0.3349 X 106 0.5384 X 10-8 0.4501 X 10-8 0.4482 X le7 0.5883 X lo"
lo00 K 0.9249 0.4770 X 10-' 0.2557 X lo-' 0.7246 X 10-2 0.1262 X lo-' 0.6618 X 0.9673 X 10-0 0.1982 X 10-1 0.2436 X 10-3 0.1113 X lo-' 0.2741 X 10-8 0.8976 X 10-7 0.6006 x lo" 0.2631 X 10-9
900K 0.9615 0.2343 X 10-1 0.4217 X 106 0.6875 X 10-2 0.1340 X lo-' 0.8137 X lk7 0.3460 X 0.8061 X 10-2 0.1233 X 10-3 0.7517 X 106 0.4421 x 10-8 0.2253 X 10-8 0.1878 X le7 0.1371 X lo"
proceeded effectively at 500 "C, and the regeneration by oxygen was achieved at 400 "C. The two moderate operation temperatures are particularly attractive from a process point of view. Previous approaches to hydrogen sulfide decomposition required operating temperatures from 700 to lo00 "C (Raymont, 1975;Fukuda et al., 1978; Edlund and Pledger, 1993),and that to selective oxidation using molten lead required a regeneration temperature higher than 750 "C (Kiuchi et al., 1982). From the point of view of a process, one needs to have a satisfactory method to convert the sulfur dioxide into more acceptable products. There are three known routes for sulfur dioxide conversion, all of which are thermodynamically very favorable. First, sulfur dioxide can be further oxidized to sulfur trioxide and then converted to sulfuric acid: SOz + l/zOz + HzO HzSO4. This is a proven technology. The overall process starting with H S can be expressed as
-
H,S
+ 3/20, + H,O
-
H,
+ H,SO,
(8)
The hydrogen recovery would be 100% on the HzS basis, but would be only 50 5% if all hydrogen-containing reactanta (HzS and HzO) are considered. An alternative is to reduce sulfur dioxide with methane to elemental sulfur and water, SO2 + l/& + l/zCOz + HzO, and the overall process becomes
-
This process would have a hydrogen recovery of 100% based on H2S and 50%based on HzSand CHI. Reduction of SO2 by methane is a proven technology and was once operated commercially for a sulfide-ore roasting facility in Canada at about 800 "C with bauxite as the catalyst (Sarlis and Berk, 1988). The process was, however, less cost-effective than the route to sulfuric acid and was terminated in the mid 1970s. Recent progress in this process was made by using Co-Mo sulfide as a more active and selective catalyst which operates with a lower reaction temperature of about 700 "C and a higher selectivity to sulfur (Mulligan and Berk, 1992). The formation of HzS and COS as byproducts was a problem in the original hightemperature operation. The third method is to react sulfur dioxide with H2S as in the second stage of the Claus process: SO2 + 2HzS 3/2S2 + 2H20. The overall process is
-
3H,S
+ 0,
-
H, + 3/2S,+ 2H,O
(10)
The Claus process is a current technology. However, the maximum hydrogen recovery would be only 335%.
Ind. Eng. Chem. Res., Vol. 33, No. 5, 1994 1097 Several side effects associated with the use of Pt/SiOz were also identified in this study. They are oxidation of sulfur dioxide to sulfur trioxide, retention of oxygen by platinum, and retention of hydrogen sulfide by silica. Oxidation of SO2 to SO3 would have no negative impact if sulfuric acid were the desired end product. However, if the desirable product were elemental sulfur to be produced from SO2 by methane reduction or the Claus reaction, the formation of SO3 would lower the overall efficiency. Retention of oxygen by platinum and retention of hydrogen sulfide by silica would both lower the hydrogen recovery yield. Retention of oxygen by Pt would lower the yield of hydrogen because some of the hydrogen would be consumed to react with oxygen. Retention of H2S by silica would also lower the yield of hydrogen because this H2S would be oxidized to HzO. Losses due to these two side reactions were about 6 76 each under the experimental conditions of this study. It is possible that a proper choice of operating variables could minimize these losses.
Acknowledgment This work was supported by the Division of Chemical Sciences, U.S. Department of Energy, Basic Energy Science. Appendix A In the decomposition of hydrogen sulfide, hydrogen monosulfide (HS),sulfanes (H&, i = 2,3,4, ...), and sulfur allotropes (S,,j = 1,2,3,...) can be formed. These reactions and their corresponding equilibrium expressions are H2S
* HS + 1/2H2
H2S e 1/iH2Si+ ((i
(AI)
- l)/i)H2 i = 2,3, ...
H2S e l/jSj + H,
j = 1,2,3,
...
(A3)
(A51
where K is the equilibrium constant and P, is the partial pressure of component n in atmospheres. For reaction at 1 atm,
Also, the reaction stoichiometry requires
The equilibrium constants for HS, H2S2, and S j , j = 1-8, were calculated by using tabulated thermodynamic data (Barin, 19891, and those of H2S3 and H2S4 were calculated by using data given by Berk et al. (1991). The equation set was solved numerically by applying an iterative algorithm. The iteration was terminated when a given calculation tolerance (typically 1.0 X 10-9 was satisfied. The calculated equilibrium composition is given in Table 3, and the equilibrium hydrogen yield is plotted in Figure 7. As can be seen in Table 3, the major decomposition products were H2, H2S2, and S2 at low temperatures and Ha, SZ, and S3 a t high temperatures.
Throughout the temperature range studied, the presence of HS, H2S3, H2S4, and Sj, j = 1,4-8, was negligibly weak. It is therefore safe to assume that the presence of even higher sulfanes and sulfur allotropes can be neglected.
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Receiued for review October 13, 1993 Revised manuscript received January 21, 1994 Accepted January 31, 1994. e Abstract published in Aduance ACS Abstracts, March 15, 1994.
