Kinetics of Catalytic Reduction of Sulfur Dioxide with Hydrogen David L. Murdockl and Glenn A. Atwood* The University of Akron, Akron, Ohio 44325
T h e kinetics of t h e reduction of sulfur dioxide with hydrogen over an activated bauxite catalyst was studied at 775 mm Hg and temperatures ranging from 345 to 39OoC. Concentrations were varied from 1.33 to 4.0 mol YO for SO2 and 2.67 to 8.0 mol % for H z . A mechanistic reaction model which agreed with t h e experimental data was developed using a randomized three-level factorial experimental design. T h e reaction rates in mol/hr/g of catalyst for SO2 and H 2 S were given by rsoz = -(0.007 f 0 . 0 0 0 5 ) p ~ 2a n d rH2S = (2.9 f 0.3) ”x 10-4 P H z 3 / ’ ps02-1/2. T h e st12 reaction’rate applies for all catalyst to feed ratios studied while t h e HzS rate is limited to catalyst to feed ratios below 11.8. It was concluded that t h e SO2 is probably reduced to elemental sulfur which in turn reacts with hydrogen to form HzS. T h e experimentally measured ratio of S/HZS was always greater than the equilibrium value.
Inroduction Over the past decade a large number of processes have been proposed for the removal or recovery of sulfur dioxide from flue gases. When this work was started, several processes under consideration involved adsorption of sulfur dioxide on a solid adsorbent and then reduction and regeneration of the adsorbent to form hydrogen sulfide, carbonyl sulfide, or carbon disulfide; finally sulfuric acid or elemental sulfur was to be recovered by further processing of the spent reducing gas. During the preparation of a study design of the U. S. Bureau of Mines Alkalized Alumina Process it became obvious that a reduction in both capital equipment costs and raw materials might be realized if sulfur dioxide were reduced directly to elemental sulfur. The reduction of sulfur dioxide with reducing agents such as hydrogen, carbon, carbon monoxide, or natural gas was practiced during the 1930’s and early 1940’s to recover sulfur in ore smelting operations but was abandoned during or shortly after World War 11. Despite the commercial use of direct reduction methods to recover sulfur during the pre-war period, there seems to be no published description of a kinetic model for the reaction. Bacon (1932) reported that hydrogen and sulfur dioxide did not react at temperatures below 500°C in the absence of a catalyst, but that reaction did occur at temperatures as low as 200°C in the presence of molten sulfur which served as a catalytic agent. Boswell (1930, 1931, 1932, 1934, 1936) patented methods of preparation and use of metallic oxide catalysts for the reduction of sulfur dioxide by either hydrogen or carbon monoxide. Doumani (1944) proposed a means of catalyst activation and Doumani, et al. (1944), studied the reduction of sulfur dioxide with hydrogen. They reported that hydrogen sulfide was always formed, in contrast to Bacon and Boswell, who did not comment on the hydrogen sulfide formation in their work. Shakhtakhtinskii and Guliev (1965, 1966) investigated the reaction of undiluted sulfur dioxide and hydrogen over bauxite and reduced alunite catalysts. They found conditions which resulted in greater than 98% reduction of the sulfur dioxide. However, no reaction mechanisms were proposed and the su1fur:hydrogen sulfide ratio was not reported. The recent studies of Haas, et al. (1971), have given insights into the reaction mechanism for sulfur dioxide with carbon monoxide while Querido and Short (1973) have investigated the technological feasibility of reduction of sulfur dioxide by carbon monoxide. Present address, Goodyear Tire and Rubber Co., Akron, Ohio 44316
254
Ind. Eng. Chern., Process Des. Develop., Vol. 13,No. 3,1974
Lepsoe (1938) presented a comprehensive treatment of the thermodynamics of the reaction of sulfur dixoide with hydrogen. In that work he assumed that the sulfur species present in the vapor phase were S2, Sg, and S S . More recent work (Meyer, 1965) indicates that the sulfur species present in the vapor consist of SS,Sq, Sg, and S7 as well. With an initial mixture of two moles of water and one-half mole of Sz, and assuming that the reactions SO,
+ 2H2 === 2H,O + %S, H, + %S, HZS
(1)
applied, Lepsoe calculated that the equilibrium ratio of S:HzS was approximately 4 at 375°C. The present work was undertaken to answer the following questions. (1) Which of several possible reaction mechanisms are consistent with experimentally observed reaction kinetics for the reduction of sulfur dioxide by hydrogen over an activated bauxite catalyst? (2) What are the kinetic rate constants in the rate expressions obtained from experimental data? (3) What factors govern the rate of formation of hydrogen sulfide relative to the rate of formation of elemental sulfur?
Thermodynamics The eight independent reactions shown in Table I formed the basis for computation of the thermodynamic equilibrium composition of the species involved in the reduction of sulfur dioxide by hydrogen. The equations which were used to calculate the equilibrium constants as a function of temperature for each reaction are also given in Table I. Figure 1 provides a comparison between the equilibrium compositions and sulfur yields calculated in this work and by Lepsoe (1938). The differences are caused primarily by the inclusion of the St, Sq, Sg, and S7 species in addition to the S2, S S , and SS species and by minor differences in the equilibrium constants. The thermodynamic calculations in this work were made for several different nonstoichiometric feed compositions in the presence of nitrogen diluent, while Lepsoe’s computations were made for the initial reactants undiluted and in stoichiometric proportions only. It was found that lower temperatures favor the formation of elemental sulfur over hydrogen sulfide as shown in Figure 1; however, even at the low temperatures significant quantities of hydrogen sulfide are predicted. Table I1 gives the calculated equilibrium conversion of sulfur dioxide and hydrogen, the concentration of hydrogen sulfide, and the ratio of sulfur to hydrogen sulfide for hydrogen to sulfur dioxide feed ratios of 0.67/1 to 8/1. It is seen that very little elemental sulfur should be produced
Table I. Equations Used to Calculate the Equilibrium Constants for the Reactions in the SOz/Sa/Hz/H2S/H20System at Different Temperatures Reaction 2s2 21/2s2
-T! s s
e S6
3s2
* s,
31/2s2 4S2
Hz0 Hz
+
K 1 = exp(13.3 - 0.0188T)/RT K 2 = exp(28.2 - 0.0367T)/RT K 3 = exp(47.8 - 0.0557T)/RT K 4 = exp(66.3 - 0.0751T)RT Kg = exp(78.9 - 0.0893T)/RT Kg = exp(96.8 - 0.1103T)/RT K7 = exp(-7944/T - 0.50661nT 1 . 7 5 1.525 X lO+T - 2.648 X 10-7T2) Ks = exp(19.4 - 0,00771 T In T 1.30 X 10bTZ 0.0125T)
2 s3
* s,
11/2s2
4
Reference
Equilibrium constant (T in OK).
s ss e NP + '/zSOz
'/&I
+ '/zSz
-T! NnS
R is 0.0019869 kcal/g mol
+ +
+
+
Detry, et al. Detry, et al. Detry, et al. Detry, et al. Detry, et al. Detry, et al. Doumani, et
(1967) (1967) (1967) (1967) (1967) (1967) al. (1944)
Kelley (1937)
OK.
Table 111. Temperature Dependence of the Thermodynamic Equilibrium Conversion of Hydrogen and Sulfur Dioxide and the Yield of Hydrogen Sulfide for the Hz/SOz/Sr'H&3/Hz0 Nz System a t 1Atmo
Table 11. Thermodynamic Equilibrium Conversions of Hydrogen and Sulfur Dioxide and the Yield of Hydrogen Sulfide for the Hz/SOZ/S,/H~S/HZO/NZ System at 375 "C and 1Atm
~~
Feed composition, mole % Hz
SO2
Product
Product % Conversion
Hz
% Conversion
Mole fraction HzS
SO2
Ratio
Temp, "C
S I :HnS
H2
so2
100.0
93.9 92.9 90.7 89.6
Mole fraction HzS
Ratio Si :H2S ~
2.67 1 . 3 3 5.33 1.33 1.33 8.0 2.67 2.67 5 . 3 3 2.67 2.67 8.0 2.67 4 . 0 5.33 4.0 4.0 8.0
100.0
74.9 49.9 100.0 100.0
99.9 100.0 100.0 100.0
9 0 . 1 2.73 x 100.0 1 . 3 3 X 100.0 1.33 X 48.2 9 . 5 1 x 90.7 4.87 x 100.0 2.65 X 32.5 7.07 x 6 3 . 8 2.22 x 9 1 . 4 6.90 x
3.38 5.18 x 10-7 5.79 x 10-8 10-4 1 2 . 5 10-3 3.98 lo-* 5 . 8 1 x 10-3 10-4 17.4 10-3 1 0 . 5 10-3 4 . 3 ., 10-3
9
1 +
t
"
+ X
+
+
x
x
SULFUR
VELD
-
345 360 375 390
100.0 100.0 100.0
3.13 X 3.71 X 4.87 X 5.47 X
~~~~
6.90 5.68 3.98 3.37
a Feed composition: 5.33 mole % H2, 2.67 mole % sulfur dioxide, 92.0 mole % nitrogen.
The equilibrium calculations showed that it is thermodynamically possible to obtain essentially complete reduction of sulfur dioxide to elemental sulfur and hydrogen sulfide, especially at temperatures below 400°C. It remained for the kinetic studies to determine a reaction mechanism and whether it would be possible to suppress the formation of hydrogen sulfide to increase the elemental sulfur yields.
X
symbds - Calcubbd. mls w a k Lines - Leproe
Theory The reaction rate for a differential reactor is given by
3
(2)
2 or in terms of conversion
s
C TEMPERATURE,
OC
Figure 1. Equilibrium gas composition for the reduction of sulfur dioxide with hydrogen at 760 mm Hg pressure: 0 ,Hz (Murdock, 1973); 0 , SO,; A, H2S; +, HzO; x, sulfur yield; -, Lepsoe (1938). when the initial hydrogen to sulfur dioxide ratio is greater than 3/1, yet for hydrogen to sulfur dioxide ratios of less than 2/1 the formation of elemental sulfur is favored. At constant hydrogen to sulfur dioxide ratios the sulfur to hydrogen sulfide ratio increases as the sulfur dioxide concentration increases. This indicates that the sulfur dioxide concentration should be maximized in a commercial process. Table 111 gives the equilibrium conversion of sulfur dioxide and hydrogen, the concentration of hydrogen sulfide, and the ratio of sulfur to hydrogen sulfide as a function of temperature for the case where the hydrogen to sulfur dioxide ratio in the feed is 2/1. This table shows that the ratio of sulfur to hydrogen sulfide varies inversely with temperature. Thus thermodynamic considerations suggest that the process should be operated a t relatively low temperatures.
(3) The reaction rates for sulfur dioxide and hydrogen were obtained by fitting the experimentally measured conversions with a hyperbolic tangent function in which the independent variable was ( w / N o ) (Corrigan, 1955). This function was differentiated and substituted into eq 3 to obtain the initial reaction rates. The initial hydrogen sulfide reaction rates were obtained using an integrated form of eq 2 and the experimentally measured hydrogen sulfide concentration at the lowest catalyst to feed ratio, (w/No). The initial rates were then used to develop the kinetic models which were tested using experimental data. Experimental Reactor Figure 2 is a schematic diagram of the small tubular reactor and associated equipment used in the experimental work, and Figure 3 shows reactor details. The gas flow rates were measured using calibrated capillary flow tubes mounted across water-filled manometers. The feed gases were mixed in a manifold where provision was made to divert a portion of the stream to a gas chroInd. Eng. Chem., Process Des. Develop., Vol. 13, No. 3, 1974 255
1 Figure 2. Experimental appartus for the catalytic reduction of SO*. 1/16. O.D. SAMPLE LINE
-.
TO G . C .
2.0 CW. I . D . RLACTOR TUBE
REACTION ZONE BAUXITB CATALYST
matograph for analysis. The gases were then warmed by passage through 316 stainless steel tubing which was wrapped with asbestos-insulated chrome1 heating wire. From the preheater the gases entered a 316 stainless steel reactor tube which was packed at the bottom with inert quartz chips to ensure uniformity of the temperature of the gas entering the catalyst bed located immediately downstream. The reactor was suspended in a Marshall tubular electric furnace equipped with shunt resistors which could be adjusted to provide a flat axial temperature profile. Thermocouples were placed in the center of the catalyst bed, at the bed entrance and at the bed exit. Temperatures were controlled with electronic controllers operating with proportional plus integral modes. Nearly all of the reaction product gases exited through a tee a t the top of the reactor. When desired, a small portion of the reactants could be sampled at any desired central axial position by means of a movable tube, Table IV gives the source and composition of the catalyst and gases used in this work.
Analytical Procedure R(WI!AT QUARTZ
ZOWL
The gaseous feed and products from the reactor were analyzed by a Varian Aerograph Model 202/1B gas chromatograph equipped with thermal conductivity detector cells; the gas chromatograph conditions are given in Table
CHIPS
v.
SEALBD END 8 A n m TUBE
m T CONTROL T.C. ----------
Figure 3. Reactor. 256
Ind. Eng. Chern., Process Des. Develop., Vol. 13, No. 3, 1974
Before chromatographic analysis of product gases, sulfur vapor in the product gas mixture was condensed in a replaceable 15 cm long x 8 mm i.d. glass tube packed with glass wool both to protect the chromatograph from deposition of solid sulfur and to enable a gravimetric determination of the amount of sulfur produced. The temperature of the line leading from the reactor outlet to the sulfur trap
Table IV. Source and Properties of Catalysts and Chemicals Materials
Source
Catalyst
Engelhard Minerals and Chemical Corp., Menlo Park, Edison, N. J.
Nitrogen Hydrogen sulfide Hydrogen Helium Sulfur dioxide
Linde Division Union Carbide Corp. Linde Division Union Carbide Corp. Linde Division Union Carbide Corp. Linde Division Union Carbide Corp. Matheson Gas Products
Table V. Chromatograph Operating Conditions Injector temperature Oven temperature Detector temperature Carrier gas Carrier gas flow rate Inlet pressure Filament current Analytical columns (two in series) First column Second column Reference column
100°C 95T 15OOC
Helium 100 cma/min 65 psig 200 mA
2.9 m long x 5 mm i.d. packed with Porapak Q-S 0.76 m long x 5 mm i.d. packed with Porapak R 0.3 m long X 1.5 mm 0.d. empty
was maintained at 220°C to prevent premature sulfur condensation, while the temperature of the line leading from the sulfur trap outlet to the chromatograph inlet was kept at 40°C to prevent water condensation. Chromatographic response factors relative to nitrogen were determined for each of the gaseous reaction components by calibration with known gas blends. It was found that the relative response factors for hydrogen, hydrogen sulfide, and sulfur dioxide varied linearly with composition, but the response factor for water was too erratic to be used. T o obtain a consistent sulfur dioxide analysis it was necessary to precondition the column with sulfur dioxide before each sample injection. Concentrations of the gases except for water were directly calculated from chromatogram peak heights by a n iterative procedure. The water concentration was calculated indirectly through elemental balances. Details of the procedures are given by Murdock (1973). Experimental Results A series of preliminary experimental reactions were carried out, first to determine if the reactor surfaces catalyzed the reaction, then to investigate the extent that diffusion limited the rate of the catalyzed reaction, and finally to obtain an estimate of feasible operating conditions for the reaction-model identification experiments. From results of the preliminary runs, a random-order three-level three-factor experimental design with variable feed composition and flow rates and constant temperature was planned as shown in Table VI to provide data for identification of the reaction model. Experiments at conditions of the central point of the experimental design
Properties Porocel activated bauxite Recovery Catalyst Mesh Particle Density g/cc Packed bed density g/cc Surface area, mz/g Approx. composition wt %
6/8 9/20 1.68 1.68 0.90 0.93 138 138 A1203 FezOs TiOz SOz
88 2.5 3.0 6.5
88 2.5
3.0 6.5
99.9% 99.9% 99.9% 99.9% 99.8%
were replicated at equal intervals both to provide a n estimate of experimental error and to monitor the activity of the catalyst. Also Table VI shows another set of experiments at the central replicate feed composition and flow rate, but a t several different temperatures. This temperature variation series was carried out to enable determination of the temperature dependence of kinetic rate constants in the reaction rate model. The preliminary studies showed that the reactor did not catalyze the reaction between 25 and 600°C. In all cases the sulfur dioxide conversion across the empty reactor was less than 3%. A set of experiments was performed to measure the effect of external diffusion and pore diffusion on the reaction rate. The reaction conditions and sulfur dioxide and hydrogen conversions are given in Table VII. For the first and third experiments in Table VI1 the conversions were essentially the same for linear velocities differing by a factor of 2 while gas residence time was the same. From this it was concluded that external diffusion was negligible. In the first two experiments of Table VII, all variables except particle size were held constant to test for the importance of pore diffusion limitations on the reaction rate. Higher conversion was observed for the experiment in which particle size was larger, which is contrary to the result expected if pore diffusion rate limitations were significant. The difference in conversions between these two experiments is attributed to unequal states of catalyst activity. Table VI presents the results of the reaction model identification experiments. By comparing results of the first three central replicate experiments, it was found that the catalyst activity increased and then stabilized. Experiments which had been performed during the initial period of catalyst instability were later rerun. Figure 4 shows the catalyst activity as a function of the run number. From the data in Table VI, the initial rates of reaction for sulfur dioxide and hydrogen sulfide were determined and then rate laws were developed for sulfur dioxide and hydrogen sulfide. The reaction rates of all other species are related by material balance to these two compounds. Table VI11 gives the initial reaction rates for sulfur dioxide and hydrogen sulfide along with the rate constant, k , which is explained later. For the range of conditions used in this work, the sulfur dioxide reaction rate was found to be independent of the Ind. Eng. Chem., Process Des. Develop., Vol. 13, No. 3, 1974
257
Table VI. Experimental Conditions and Conversions for the Experiments Used to Develop the Reaction Model
Feed conditions Flow rate, cma/min 1500
Concentration w/No
Temp, OC
Hz
so2
Nz
so2
11.8
375
0.0267
0.0133 0.0267 0,0400 0.0133 0.0267 0.0400 0.0133 0.0267 0.0400 0.0133 0.0267 0.0400 0.0133 0.02670 0.0400 0.0133 0.0267 0.0400 0.0133 0.0267 0.0400 0.0133 0.0267 0.0400 0.0133 0.0267 0,0400 0.0267 0.0267 0.0267
0.179 0.174 0.216 0.156 0.176 0.182 0.150 0.146 0.144 0,322 0.344 0.331 0.348 0.3245 0.342 0.237 0.313 0.293 0.621 0.563 0.521 0.504 0.491 0.504 0.428 0.484 0.486 0.180 0.246 0.405
0.178 0.074 0.056 0.296 0.151 0.091 0.534 0,225 0.168 0.311 0.196 0.107 0.644 0. 281a 0.229 0.731 0.433 0.263 0.559 0.269 0.166 0.841 0.449 0.311 0.980 0.653 0.453 0.169 0.243 0.368
0.0533 0.0800 750
23.6
375
0.0267 0.0533 0.0800
375
47.2
375
0.0267 0.0533 0.0800
750
a
Fractional conversion
23.6
345 360 390
0.0533 0.0533 0.0533
Hydrogen sulfide concn, mole fraction 0.015 0.010 0.009 0.034 0.025 0.018 0.071 0.045 0,033 0.035 0.027 0.019 0.120 0.061a 0.051 0.15 0.11
0.081 0.098 0.053 0.038 0.22 0.13 0.11
0.17 0.23 0.48 0,018 0.036 0.100
SdHnS ratio 10.9 16.4 23.1 7.7 13.4 18.3 6.6 10.8 18.2 8.7 18.2 20.1 5.1
lo.@ 16.8 4.3 8.4 11.5 6.1 12.3 15.4 3.8 7.5 10.8 1.5 5.8 9.7 22.4 15.9 7.7
Central replicate conditions.
Table VII. Experimental Conditions Used to Measure the Effect of Pore Diffusion and External Diffusion on the Reaction Rate
Feed composition, mole fraction
Catalyst
Feed flow rate, cma/min
Temp, "C
Size mesh
Wt, g
H2
so2
Nz
Hydrogen conversion, %
600 600 300
400 400 400
9/20 6/8 9/20
44.3 44.3 22.1
0.08 0.08 0.08
0.04 0.04 0.04
0.88 0.88 0.88
34 49 36.9
Table VIII. Initial Reaction Rates for Sulfur Dioxide and Hydrogen Sulfide at 375 "C
Inlet conditions Y80a
YE2
0.0133 0.0267 0.0400 0.0133 0,0267 0.0400 0.0133 0,0267 0.0400
0.0267 0.0533 0.0800
Initial sulfur dioxide rate, mol hr-1 (g of cat.) -1
Rate constant, k~% mol , hr -1 (g of cat.) -1 atm-l
Initial hydrogen sulfide rate, mol hr -l (g of cat.)
0.000190 0.000226 0.000200 0.000406 0.000348 0.000403 0.000604 0.000543 0.000518
0.0140 0.0166 0.0147 0.0149 0.0128 0.0148 0.0148 0.0133 0.0126
1.29 x 10-5 0.869 X 0.766 X 2.95 x 10-5 2.16 x 10-5 1.52 x 10-5 6.15 x 10-5 3.87 x 10-5 2.85 x 10-5
sulfur dioxide concentration and proportional to the hydrogen concentration. The sulfur dioxide rate law a t 375°F is
Rate constant, K H ~ S ,mol hr -l (g of cat.) -1 atm-l 3.35 x 3.19 x 3.45 x 2.71 x 2.81 x 2.42 x 3.08 x 2.74 x 2.46 x
10-4 10-4 10-4 10-4 10-4 10-4 10-4 10-4 10-4
ide concentration in the feed, it was postulated that the rate of hydrogen sulfide formation could be expressed as
(4) where kH2 = 0.014 f 0.001 mol hr-l (g of c a F ) - l atm-I. Using the initial rates for hydrogen sulfide and the observations that the hydrogen sulfide formation rate increased rapidly with increasing hydrogen concentration in the feed and decreased slowly with increasing sulfur diox258
Ind. Eng. Chem., Process
Des. Develop., Vol. 13, No. 3, 1974
From the data in Table VI11 it was found that a = 312 and b = yz and that the rate constant k was equal to 2.9 x mol hr-l (g of cat.)-l atm-l. The f 0.3 x error estimates are based on 95% confidence limits. From stoichiometric considerations for reactions 8 and 9 in
Table IX.Comparison of Measured and Calculated Reactions of Sulfur Dioxide and Hydrogen Sulfide at 375 ~
Feed Composition mole fractions
Flow rate cc/min
H2
so2
1500
0.0267
0.0133 0.0267 0.0400 0.0133 0.0267 0.0400 0.0133 0.0267 0.0400 0.0133 0.0267 0,0400 0.0133 0.0267 0.0400 0.0133 0.0267 0.0400 0.0133 0.0267 0.0400 0.0133 0.0267 0.0400 0.0133 0.0267 0.0400
0.0533 0.0800 0,0267
750
0.0533 0.0800 375
0.0267 0.0533 0.0800
~~
~~~
Hydrogen sulfide, mole %
Sulfur dioxide conversion Exptl
Calcd
17.8 7.4 5.6 29.6 15.1 9.06 53.4 22.5 16.8 31.1 19.6 10.7 64.4 28.1 22.9 73.1 43.3 26.3 55.9 26.9 16.6 84.1 44.9 31.1 98.0 65.3 45.3
~~
Exvtl 0.015 0.010 0.009 0.034 0.025 0.018 0.071 0.045 0.033 0.035 0.027 0.019 0.012 0.061 0.051 0.15 0.11 0.081 0.098 0.053 0.038 0.22 0.13 0.11 0.48 0.23 0.17
16.1 8.0 5.4 32.1 16.0 10.7 48.0 24.0 16.0 29.5 14.7 9.8 58.7 29.3 19.6 88.0 44.0 29.4 50.0 25.0 16.7 99.3 49.8 33.3 99.9 74.5 49.9
Calcd 0.012 0.0084 0.0068 0.036 0.024 0.020 0.070 0.046 0.036 0.022 0.015 0.012 0.070 0.044 0.035 0.15 0.085 0.066 0.038 0.025 0.020 0.16 0,075 0.058 0.22 0.15 0.11
Table X.Proposed Steps in the Reaction Mechanism and Corresponding Reaction Rates Reaction
0
;
J
I
9
16 23 RUN NUMBER
33
Figure 4. Catalyst activity time dependence. Table I, the rate of reaction for each of the other species, based on initial reaction kinetics, is
(7) ~ H . O=
Reaction
~ H ~ P H ~
(8)
Rate laws for the reacting species, eq 4, 5, 6, 7, and 8, were derived from experimentally determined initial reaction rates and stoichiometric relationships. Simultaneous numerical integration of these laws was carried out over the range 0 < w/No < 47.2 to test their ability to accurately predict reactant conversions comparable to those obtained experimentally for W I N O greater than zero. Conversions obtained by rate law integrations are compared with conversions obtained experimentally in Table IX. The rate law for the primary reaction sequence which produces sulfur and water is seen to apply over the entire range of sulfur dioxide conversions, 0-98%, but the predicted yields of hydrogen sulfide fall short of experimental yields when W I N O is greater than 11.8. Possibly an uni-
dentified secondary reaction is responsible for the higherthan-predicted hydrogen sulfide yields. The temperature dependence of the overall reaction rate constants, k~~ and k H I S were measured in experiments which repeated the central replicate conditions a t 345, 360, and 390°C, and were then correlated by the Arrhenius function
k =
A~-E'RT
(9) The resulting preexponential factor and activation energy for kH2 are
A = (4.8 f 0.2) X lo3 mol ( g of cat.)-' hr-' atm-' E = 16.3
* 0.3 kcal mol-I OK-'
and for k H z sare
A
=
1.5 f 0.1 X 108 mol (g of cat.)-' hr-' atm-I E = 34.3 0.6 kcal mo1-l "K-I
*
Proposed Reaction Mechanism Table X shows reaction rate expressions for each of the four steps of the proposed reaction mechanism. At steady state the rate of formation of oxysulfide sites must equal their rate of disappearance, i e . , 7-10 = -lhr13; otherwise the Ind. Eng.
Chem., Process Des. Develop., Vol. 13, No. 3,1974 259
catalyst surface would ultimately saturate with either the sulfur sites or the oxysulfide sites, and in either case the reaction would cease. With the definition which must be constant, since [OSaX] is constant at steady state, and noting that the rate of sulfur dioxide reaction is equal to -r10, it follows that
- r10 = - %kH*PH2
(11)
which is identical with the rate law for sulfur dioxide reaction obtained experimentally. Assuming that and [SbX] are equal, and combining the rate expressions for reactions 10, 11, and 13gives
[sax]
or (12) which is identical with the experimentally determined rate law for the formation of hydrogen sulfide. Summary The activated bauxite catalyst was found to be activated by exposure to sulfur dioxide and hydrogen, which were the reacting feed materials. During use at 375" the catalyst activity doubled, then stabilized. Results of experiments designed to directly measure the extent of pore diffusion limitation of the overall reaction rate were confounded by changes in the catalyst activity; however, subsequent reaction kinetics experimental results showed consistency with a proposed reaction mechanism in which the chemical reaction rate rather than mass transfer rate was controlling. External film diffusion was ruled out as a possible ratecontrolling step by experiments for which no significant change in reaction rate occurred when the velocity of reactant gases flowing past the catalyst particles was changed. Empirical initial-reaction-rate correlations indicated that the rate of the primary reaction route, which leads to elemental sulfur and water, is first order in hydrogen and zero order in sulfur dioxide. The integrated rate law for the primary reaction was found to be valid for predicting sulfur dioxide conversion over the range 0-9870. The initial rate of the reaction which leads to hydrogen sulfide was found to be three-halves order in hydrogen and negative-one-half order in sulfur dioxide. Integration of the second reaction rate law predicted hydrogen sulfide yields which were 50-80% of the experimental yields for w/No greater than 11.8. The ratio of elemental sulfur to hydrogen sulfide in the product was experimentally found to vary inversely with total conversion and directly with the ratio of sulfur dioxide to hydrogen in the feed.
260
Ind. Eng. Chem., Process Des. Develop., Vol. 13,
No. 3, 1974
The rate, law temperature dependences were satisfactorily correlated by means of the Arrhenius function from 345 to 390°C. A reaction mechanism involving oxidation, reduction, desorption, and regeneration of sulfur sites was postulated and was shown to be consistent with observed initial rate kinetics. By comparing the equilibrium ratio of sulfur to hydrogen sulfide with the experimentally measured ratio of sulfur to hydrogen sulfide, it apqears that the reduction of sulfur dioxide with hydrogen progresses in two steps, the first to elemental sulfur and then the reaction of sulfur with hydrogen to form hydrogen sulfide. Nomenclature A = Arrhenius frequency factor, mol/(g of cat.) atm hr E = energy of activation, kcal/mol "K ki = reaction rate constant for reaction i, mol/(g of cat.) atm hr Ki = thermodynamic equilibrium constant for reaction i NO= molar flow rate into reactor, mol/hr p j = partial pressure of speciesj, atm r, = reaction rate for speciesj, mol/hr (g of cat.) T j = average reaction rate for speciesj, mol/hr (g of cat.) R = universal gas constant, kcal/mol O K T = temperature, O K w = mass of catalyst bed, g w/No= space time, hr (g of cat.)/mol X = adsorption site on catalyst X S a = adsorption site occupied by elemental sulfur X S a O = adsorption site occupied by SQO [XI = concentration of an active site [XSQ] = concentration of sites occupied by sa [XSQO]= concentration of sites occupied by sao X, = fractional conversion of species j y , = mole fraction of speciesj Literature Cited Bacon, R. F . , U. S.Patent 1,830,076 (Jan 5,1932), Boswell, M. C.. Can. Patent 301,554 (July 1. 1930) Boswell, M. C., Can. Patent 308,238 (Feb 3, 1931) Boswell, M. C., U. S. Patent 1,880,741 ( O c t 4 , 1932). Boswell, M . C., Can. Patent 338,600 (Jan 16, 1934). Boswell, M. C., U . S.Patent 2,026,819 (Jan 7, 1936). Corrigan, T. E., Chem. Eng., 62, 203 (1955) Detry, D . , Drowart, J., Goldfinger, P., Keller, H.. Rickert, H.,Z. Phys. Chem. (Frankfurf am Main), 55, 314 (1967). Doumani. T. F., U. S. Patent 2,361,825 (Oct 31, 1944). Doumani, T. F., Deery, R. F., Bradley, W. E., Ind. Eng. Chem., 3 6 , 329 (1944), Haas, L. A,, McCormick, T. H., Khalafalls, S. E., Bur. Mines Rep. Invest., R I 7483 (March 1971). Kelley, K. K . , Bur. Mines Boll. No. 406 (1937). Lepsoe, R . , Ind. Eng. Chem., 30, 92 (1938) Meyer, B., Ed., "Elemental Sulfur," Interscience, New York, N. Y., 1965. Murdock, D. L . , Ph.D. Dissertation, The University of Akron, Akron, Ohio, 1973. Querido, R., Short, W. L., Ind. Eng. Chem., Process Des. Develop., 12, 10 (1973). Shakhtakhtinskii, G. B., Guliev, A. I., Azerb. Khim. Zh., 2 , 104 (1965). Shakhtakhtinskii, G. B.. Guliev, A. I . , Azerb. Neft. Khoz., 45, 39 (1966).
Received for reuiew August 30,1973 Accepted J a n u a r y 24,1974
