Hydrogen Reduction of Alkali Sulfate J . R. Birk, C. M. Larsen, W. G. Vaux, and R. D. Oldenkamp Atomics International, North American Rockuell Corporation, P.O. Box 309, Canoga P a r k , Calif. 91304 The reaction between sulfate, dissolved in an alkali carbonate, and hydrogen was found to be of zero order in sulfate and two-thirds order in hydrogen (500 to 840°C.).
The reaction i s autocatalyzed by sulfide with a reaction order of two-fifths and i s catalyzed by iron, iron compounds, and stannous sulfate. The average uncatalyzed reaction rate of 0.1 1 mole per liter hour (600°C.) quadrupled with each 100°C. increase in temperature while the iron-catalyzed rate of 0.65 mole per liter hour (600" C.) doubled for each 100" C. increase. The mechanism appears to involve reaction between sulfate and sulfide to form poylsulfide. The catalytic activity of certain transition metals appears t o result from oxidation of the metal b y the sulfate or reactive intermediate. The metal i s reformed with hydrogen.
T h e hydrogen reduction of sulfate dissolved in a molten alkali carbonate (Reaction 1) was studied in order to optimize the reduction step of the molten carbonate process for sulfur oxide removal from flue gases (Heredy et al., 1969).
M?SO?+ 4H?
600 to 800aC. +
M?S + 4HgO
(1)
The reaction yields the alkali metal sulfide (where M represents the eutectic carbonate mixture of lithium, sodium, and potassium cations). The reduction of sulfate to sulfide has been carried out routinely for over a century, first in the LeBlanc process, 1823, and later in the recovery furnaces of the Kraft pulping process. The reduction has been accomplished with several different reducing agents (Meyer, 1965); however, hydrogen is generally considered the most active (Nyman and O'Brien, 1947; White and White, 1936). Although considerable discrepancy concerning reduction temperatures and rates exists in the literature, temperatures between 750" and 900" C. are frequently cited as the most ideal. Temperatures in the 600°C. range have been utilized a t the expense of longer reaction times, while temperatures above 900" C. have caused side reactions. The phase of the sulfate-containing medium which proves to be the most reactive is that phase giving the best contact between the reducing agent and the sulfate. Thus, the solid phase is better for hydrogen reductions ( S y m a n and O'Brien. 1947; White and White, 1936). Catalysts for the reduction of sulfate with hydrogen fall into three general categories: Transition-metal compounds (Bagbanly and Mirbabaeva, 1959; Nyman and O'Brien, 1947), sulfide (autocatalytic) (Polyvyannyi and Demchenko, 1960); and carbon (Kunin and Kirillov, 1968). Of the transition-metal compounds, iron has been demonstrated to be the most useful (Bagbanly and Mirbabaeva, 1959; Nyman and O'Brien. 1947) catalyst for all of the reducing agents. Since the catalytic activity is universal, it seems that the iron-containing species must interact with the sulfur-containing species. Little information has been presented concerning the many parameters that influence the sulfate reduction reaction. As a result, mechanistic descriptions could only be speculative. Thus, the most discussed mechanism was the most logical one; that
is, the bimolecular reduction of sulfate with hydrogen to form sulfite and the subsequent disproportionation of sulfite to sulfide and sulfate (Fotiyev, 1960; Nikitin and Kunin, 1960). This mechanism, assuming the ratedetermining step is the sulfate-hydrogen interaction, results in a kinetic expression that would be first order with respect t o both sulfate and hydrogen. With use of the above background information as a guide and with the restriction of having a molten carbonate eutectic as a solvent, studies were carried out to optimize the reduction of sulfate with hydrogen. In particular, the reduction was investigated under a wide variety of conditions. Operating variables included temperature, catalyst, iron concentration, sulfide concentration, nature of iron compound, hydrogen pressure, etc. Experimental
The procedure used in all tests is described below. A graphite crucible (1J4 in. i.d. x 10 in.) was washed with 1-to-1 hydrochloric acid, distilled water. and acetone, and dried under a stream of air. The reactant salt mixture was weighed, mechanically mixed, and transferred to the graphite crucible. The equivalent amounts of salts used were 100.8 grams lithium carbonate, 100.8 grams potassium carbonate, 84.9 grams sodium sulfate, and 37.0 grams sodium carbonate (all chemicals were reagent grade). For catalytic experiments, about 0.14 mole of transition metal catalyst was added to the salt mixture. The graphite crucible was placed in a stainless steel container which was subsequently welded shut. A diagram of the resulting reaction vessel is shown in Figure 1. The container was then placed in a furnace and brought t o temperature under a continuous stream of argon or helium. After the melt was a t temperature, a sample was taken by drawing a portion (4 to 7 grams) of melt into a 4or 5-mm quartz tube. The vessel was sealed (except for the inlet and exit gas lines) and the reactant gases were passed through the melt a t a rate of 1.4 1. per min (unless otherwise indicated). Melt samples were taken periodically a t the same time the water trap (magnesium perchlorate) was weighed. During sampling an inert gas was passed through the melt. Samples were analyzed for sulfate by standard gravimetric barium sulfate precipitation. Total sulfur was done Ind. Eng. Chem. Process Des. Develop., Vol. 10, No. 1, 1971
7
r
by bromine oxidation in basic solution and subsequent barium sulfate precipitation. A recently developed gas chromatographic technique was used for the analyses of sulfite, sulfide, and carbonate (Birk et al., 1970). Accuracies of the techniques are +6% for sulfite and sulfide, and 1 2 % for carbonate, sulfate, and total sulfur. Iron was determined by the standard colorimetric procedure using 1,lo-phenanthroline. Results
Since sulfate analyses were the most reliable, sulfate data were used to calculate the extent of reduction. However, in every case, the per cent reduction based on sulfide results compared closely with that based on sulfate results. Water evolution data, however, tended to give results that were about 25% low. Since the initial amount of sulfate remained unchanged between tests, the results for completion of reaction and reaction rate are given as per cent reduction and per cent completion per hour, respectively. The per cent completion per hour-i.e., rate of reaction-was arbitrarily taken as the average rate between the time when 2 5 5 of the sulfate was reduced and the time when 7 5 5 of the sulfate was reduced. This was the most direct way of data handling since the results were used only in comparison with other tests. However, the average rate of reduction in per cent per hour can readily be converted to moles per liter per hour by multiplying by 0.038. Catalysis. The catalytic activity of several transitionmetal compounds and one alkali metal compound was investigated. Ferrous sulfate, cupric sulfate, chromic oxide, manganous sulfate, nickelous sulfate, stannous sulfate, molybdenum metal, vanadium oxysulfate, titanium sulfate, cobaltous sulfate, germanium oxide, and rubidium carbonate were tested a t 600" C. in concentrations of about 5 mole %. The results of some of these tests are shown in Figure 2. These results show that only tin and iron significantly catalyze the reduction. Although molybdenum appears catalytic from the figure, it was added as the metal and, as a result, undoubtedly took some
-GRAPHITE PURGE TUBE (3/16 in. ID)
-STAINLESS STEEL CONTAINER (2-1/2 in. PIPE 12 i n . LON
-GRAPHITE CRUCIBLE (1-3/4in. ID x 10 in.)
-MELT (APPROXIMATELY 4 in. DEEP)
Figure 1. Schematic of reaction vessel
100
z
0 I-
o 3
n w E I-
50
z W o cz W a
0
0
20
10
30
TIME (hours)
Figure 2. Comparison of catalysts for the reduction of sulfate at 600" C. 8
Ind.
Eng. Chem.
Process Des. Develop., Vol. 10,
No. 1, 1971
40
These results show a fairly steady, but not sharp, increase in reduction rate with temperature. I n these tests the rate of increase of the iron-catalyzed reduction was only about a factor of two for each 100°C. temperature increase, while for the uncatalyzed tests a factor of about four was observed. This result was not surprising since the iron undoubtedly lowered the activation energy of the reaction, thus making it less sensitive t o temperature change. Iron Catalysis. The comparatively high catalytic activity of iron as a reduction catalyst prompted two separate investigations in addition t o the temperature study discussed above; these were the effect of using various chemical compounds containing iron on the reduction rate and the effect of various concentrations of iron on the reduction rate. Three different iron compounds (FeS, FeSO,, Fer03) and elemental iron were employed as catalysts. The effect these species have on the reduction is shown in Figure 5 . The average reduction rate was found to be 19.45 per hour with FeS, 17.2% per hour with FeSO?, 15.65 per hour with Fe203, and 12.85 per hour with Fe as compared with 2.9% per hour for the uncatalyzed reduc-
Table I. Effect of Temperature on Reduction Rate Temperature
C.
600 600" 650 650" 700 700" 730 840 a
Rate, %/hr.
2.9 17.5 8.2 25.0 14.7 38.1 24.0 88
Iron catalyzed.
part in reducing the sulfate. This could account for part of its activity and also explain the atypical reduction curve noted in the molybdenum test. Metal compounds not shown in Figure 2-Le., compounds containing Cu, Cr, Mn, Ti, and Ni-were found not to enhance the reaction rate. Effect of Temperature. The effect of temperature on the rate of reduction is given in Table I and Figures 3 and 4.
100
z
0
Figure 3. Effect of temperature on the iron-catalyzed reduction rate
I-
o 3
n W
E
50
F
z W o E
W
a
0
2
4
6
8
TIME (hours)
Figure 4. Effect of temperature on the uncatalyzed reduction rate Ind. Eng. Chem. Process Des. Develop., Vol. 10, No. 1, 1971
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100
c 0 = l
n W
c z
50
W
0
K W
a
0 TIME (hours) Figure
0
5. Reduction at 600" C. using various iron compounds as catalysts
5
10 TIME (hours)
15
20
Figure 6. Effect of iron concentration on the reduction a t 650°C.
tion a t 600°C. Although the length of the initial slow period (the induction time) changes with the use of different forms of iron, the actual rates vary only 40%. The difference between the control and the average rate with iron catalysts is 5607. The differences among the rates using various iron compounds are probably attributable to the presence of sulfide (in the case of FeS and FeSO?, the latter of which decomposes forming sulfide), and may also result from the rate a t which the iron compound changes into the most stable form (and apparently the most active form) for the existing conditions. T o test the effect of iron concentration on the rate of reduction, iron concentrations of 0.01, 0.09, 0.23, 1.3, 2.4, 4.8, and 9.6% were used a t 650'C. The effect of the different amounts of iron is presented in Figure 6. The results show that increasing the iron concentration several orders of magnitude increased the reduction rate 10
Ind. Eng. Chem. Process Des. Develop., Vol. 10, No. 1, 1971
appreciably. The anomalous points, seen in the case where 9.6% iron was employed, can probably be attributed t o the very high apparent viscosity observed in this test. This high viscosity probably resulted from insolubility of the iron compounds and caused decreased mixing between the gas and liquid phases. Either of these phenomena could result in a lower reaction rate. Sulfide Catalysis. The shape-Le., increasing slope-of the reduction us. time curves indicated that the reaction was autocatalytic, in agreement with the literature (Polyvyannyi and Demchenko, 1960). T o substantiate autocatalysis a test was carried out a t 65OoC. with a melt initially containing 22.5% alkali sulfide (Le., the product of the reduction reaction) and compared with a test using no initial sulfide. The results of these tests are shown in Figure 7 , where the average reduction rate with 22.5% initial sulfide was 37.0% per hour and with no initial sulfide was 25.0% per hour. Figure 7 demon-
100
z
2
IW -I
a
2
0 0
50
I-
z W 0 IY
W
a
0
2
1
0
TIME (hours)
Figure 7. Effect of sulfide on the rate of reduction at 650°C. strates that the slow initial rate was completely eliminated, thus confirming the hypothesis that the reaction is autocatalytic in sulfide. Effect of Hydrogen Pressure. The effect of increasing hydrogen pressure was investigated. An exploratory experiment was done with five atmospheres of hydrogen a t a flow rate of 1.4 1. per min (ambient conditions). The rate of reduction was considerably slower than in the one atmosphere test done under otherwise identical conditions. These results suggest that decreased agitation and surface contact between the liquid and gas due to lower gas volumes a t the higher pressure was more important than the increase in pressure. Experiments were then designed to eliminate the variation in contact and agitation by keeping a constant volume throughout at all pressures employed. At pressures below one atmosphere, helium was mixed with hydrogen in the desired proportions; a t pressures above one atmosphere, the ambient flow rate was proportionately increased. The experiments were carried out at 600°C. with use of both sulfide (10% Na2S)
and iron (2.4% Fe added as FeS04) catalysts. The results are given in Figure 8. The rate of reduction was found to be 6.0% per hr a t '/1 atm, 10.0% per hr a t % atm, 17.0% per hr a t 1 atm, 29.9% per hr a t 2 atm, 42.4% per hr at 5 atm, and 68.5% per hr at 10 atm. These results show that the rate increases by about 1.6 for each twofold increase in hydrogen concentration or pressure. Effect of Other Gases. In conjunction with the pressure tests, two tests were done using water instead of helium to dilute the hydrogen and to determine if water had any effect-e.g., cause a reversal of the reduction reaction. The tests were done at 65OoC, with 0.4 and 0.7 atm water vapor pressure. The reduction rate with 0.6 atm hydrogen plus 0.4 atm H 2 0 was 3470 slower than with pure hydrogen. Based upon the pressure studies one would expect a rate decrease of 30% if water vapor acted only as a diluent. The close agreement points to the conclusion that water vapor simply acts as an inert diluent in the reduction of sulfate with hydrogen. The test using 0.7 atm water vapor incorporated a 10% sodium sulfide catalyst and was not followed with an identical test using no water; however, semiquantitative results based upon interpolated data indicated the same conclusion mentioned above. In a separate iron-catalyzed reduction test the effect of carbon dioxide on the reduction rate was investigated a t 7 0 P C . I n the experiment, a gas mixture containing 10% carbon dioxide and 90% hydrogen was used. I t was found that this gas mixture reduced the reaction rate by 34% (from 34.5 to 22.7% per hour), whereas less than a 10% reduction would be expected if carbon dioxide were inert. All the water from the oxidation of the hydrogen was recovered in this test. This was not noted in tests with no carbon dioxide present, where 75% recovery was typical. Discussion
Analysis of the Reduction Data. The first conclusion to be drawn from inspection of the data is that the hydrogen reduction of sulfate is zero order or very nearly so with respect to sulfate concentration. If the sulfate concentration influenced the reaction rate, its disappearance would
100
z 0, I-
u
3 0
50
I-
2 W 0 &
W
a
0 0
2
4
6 TIME (hours)
8
10
Figure 8. Effect of hydrogen pressure on the reduction at 600°C. Ind. Eng. Chem. Process Des. Develop., Vol. 10, No. 1, 1971
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cause the reaction rate to decrease. However, the reaction rate does not decrease; instead, it increases. Hence, the concentration of sulfate must have no influence upon the reaction rate. This conclusion is in agreement with data obtained in the reduction of solid sodium sulfate with hydrogen (Puttagunta, 1967). The second conclusion is that the sulfate reduction is autocatalyzed. Results from tests with high initial sulfide concentration and the observation that the reaction rate increased with time supported this conclusion. T o account for both of these conclusions, the reaction rate expression is written in the form
where h is the overall reaction rate constant. Integrating this expression and assuming that PH?is independent of time and that a t t = 0, (M2S) = 0, the following expression is obtained:
(M?S)'- n = h'(PH2)? (3) where h' = (1 - n ) k . Thus, a logarithmic plot of (M2S) or per cent reduction us. time will give a slope from which the reaction order, n, can be calculated. The reaction order with respect to sulfide concentration in the hydrogen reduction of sulfate was 0.371 i 0.071 (at the 95% confidence level). This number was independent of temperature and catalyst. The order of the reaction, m , can be derived, conventionally, as the slope of a plot of log (rate) us. log (hydrogen pressure). The least squares slope was found to be 0.646 0.058 (at the 90% confidence level). To account for iron catalysis, the basic rate expression (Equation 2 ) is modified in the following manner: d (M2S0,)= ~ ( M ~ S ) " ( P H , )+" [(Fe)'] C dt
--
(4)
where C is a constant. The last term in brackets accounts for the fact that the reaction proceeds without any iron present and that when iron is present it proceeds a t the uncatalyzed rate plus an additional amount which is a function of (Fe)P. Subtracting out the rate when no iron is present yields the expression
d d -_ (M?SO,)T + - (MeS04)o= dt
dt
Here, the first term represents the total actual observed rate and the second term represents the rate when no iron is present. The slope of a logarithmic plot of the left side of Equation 5 us. iron concentration equals the reaction order. The least squares slope-Le., the reaction order-was 0.41 at 650" C. Mechanistic Interpretation. The fact that the reaction is zero order or very nearly zero order with respect to sulfate demonstrates that sulfate itself is not involved in the rate-determining step and suggests that some other sulfur-containing compound is actually being reduced. Since the reaction is autocatalytic with respect to sulfide, and since sodium sulfate does not undergo appreciable simple self-decomposition or self-reaction at these temperatures, it is likely that sulfide and sulfate react to 12
Ind. Eng. Chem. Process Des. Develop., Vol. 10, No. 1, 1971
form the intermediate. This is strongly supported by the fact that isotopic exchange of sulfur has been observed between the oxidized and reduced species during sulfate reduction (Vinogradov and Belyi, 1967). Some possible reactions and intermediates are as follows:
The intermediate that appears most likely is polysulfide because of its stability up to temperatures of 13OO0C. (Ahlgren and Teder, 1967), and because formation of polysulfide via Reaction 8 involves only one sulfate for seven sulfides, explaining how the reaction order with respect to sulfate could be nearly zero. Further support for Reaction 8 is realized by noting similar well-known reactions; for example, the reduction of concentrated sulfuric acid by hydrogen sulfide
HzS04
+ H2S
-+
S + 2Hz0
+ SO2
(9)
or the reduction of sulfur dioxide by hydrogen sulfide
SO2 + 2H2S -+ 3s + 2H20
(10)
Additional evidence for polysulfides is the fact that sulfur was observed upon acidification of samples taken during the reaction. Thus, the reaction which satisfies all experimental results to date involves the reduction of sulfate with sulfide to form polysulfide. Although Reactions 6 to 8 suggest that alkali sulfide reacts with the sulfate, this may not be the case. In every test, except the one where hydrogen was diluted with carbon dioxide, a considerable amount of water seemed to remain in the melt. In fact, the water remaining in the melt constituted a quantity sufficient for a l-to1 molar ratio between the sulfide and water. This suggests that the two may be reacting to form hydroxide and bisulfide
MzS
+ H20
-+
MHS + MOH
(11)
Thus, bisulfide and not sulfide may be involved in the mechanism. There are at least two major ways in which iron could catalyze the reduction: act as a hetergeneous (surface) catalyst and (or) act directly in reducing the sulfate or the reactive intermediate. The former is unlikely since sulfide, usually a very severe poison for surface catalysis, acts to further increase instead of decrease the reaction rate. In addition, all iron compounds, including metallic iron, are nearly equally effective. Since heterogeneous catalysis is critically dependent on surface, it is likely that this type of catalysis is not occurring. The second role for iron involves a mechanism whereby iron itself reduces the sulfate or sulfur-containing intermediate, forming iron oxide, and reduction of the iron oxide with hydrogen back to metallic iron or a lower oxide. This mechanism has been previously suggested (Puttagunta, 1967) for solid sodium sulfate. For such a reaction to be feasible, the free energy for each of the two reactions
should be negative or close to zero. Thermodynamic calculations show this to be the case with four other metals [Mo, Ge, V(VO2 + VO), and Sn] as well as iron. I t is interesting to note that, of all the metals tried as catalysts, four (Fe, V, Mo, and Sn) of the five mentioned above were the most effective of all the metals that were investigated. In addition, Reaction 12 was found to take place in molten carbonate. I t was found that 4.8% iron oxide dissolved in the carbonate eutectic was, reduced by 31% in 6.75 hours. Reaction 13 has also been shown to take place (Vinogradov and Belyi, 1967). This indirect information suggests that iron is directly involved in the reduction. literature Cited
Ahlgren, P., Teder, A., Acta Chem. Scand. 21, 1119 (1967). Bagbanly, I. L., Mirbabaeva, F. Yu, Azerbaidzahn. Khim. Zhur. (Russ.) 111 (1959); CA 55, 10817c (1961). Birk, J. R., Larsen, C. M., Wilbourn, R . G., Anal. Chem. 42, 273 (1970). Fotivev. A. A., Izuest. Sibir. Otdel. Akad. Nauk S.S.S.R. (Russ.) No. 9, 107 (1960). Heredy, L. A., McKenzie, D. E., Yosim, S. J. (to North American Rockwell), U. S. Patent 3,438,722 (April 15, 1969).
Kunin, V. T., Kirillov, I. P., Izu. Vyssh. Ucheb. Zaued., Khim. Tekhnol (Russ.) 11, 569 (1968); C A 69, 64258x (1968). Meyer, R. J., “Gmelins Handbook of Inorganic Chemistry,” 8th ed. (Ger.), Vol. 21, pp. 184-205, Verlag Chemie, Weinheim, Germany, 1965. Nikitin, V. A., Kunin, T. I., Zhur. Vsesyuz. Khim. Obshchestua im D. I . Mendoleeua (Russ.) 5 , 350 (1960). Nyman, C. J., O’Brien, T. D., Ind. Eng. Chem. 39, 1021 (1947). Polyvyannyi, T. R., Demchenko, R. S., Izuest. Akad. Nauh, Kazakh., SSR, Ser. Met., Obogashchen. i Ogneuporou (Russ.) 34, (1960); CA 55, 198e (1961). Puttagunta, V. R., Ph.D. thesis, Univ. of Saskatchewan, Saskatoon, Saskatchewan, 1967. Vinogradov, V. I., Belyi, V. M., Izotopy Sery Vop Rudoobrazou. (Russ.) 118 (1967); CA 70, 5899e (1969). White, F. M., White, A. H., Ind. Eng. Chem. 28, 244 (1936). RECEIVED for review November 3, 1969 ACCEPTED August 3, 1970 The basic process is a proprietary development of North American Rockwell Corporation. Subsequent work was performed pursuant to contract P H 86-67-128 with the U. S. Public Health Service, Department of Health, Education and Welfare.
Adsorption of light Hydrocarbons from Nitrogen with Activated Carbon William C. McCarthy Phillips Research Center, Phillips Petroleum Co., Bartlesuille, Okla.
74003
Dynamic adsorption data were obtained on activated carbon for single, binary, and multicomponent hydrocarbons in nitrogen a t 100 and 300 psig. Component concentration, gas velocity, presaturation, and pressure were investigated and correlated with the length of the mass transfer zone. Presaturation of the carbon with the next lighter hydrocarbon component had the greatest effect on zone length. Bed length, component concentration, gas velocity, and pressure also affected the zone length substantially. Since the zone length varied considerably, depending upon conditions, it was not considered a good design tool for multicomponent systems. An exposure-recovery plot facilitates the design of an adsorption unit but a separate, experimentally-determined plot i s required for each feed composition, pressure, and velocity encountered.
T h e recovery of propane from natural gas using a shortcycle carbon adsorption unit has been demonstrated by Bray et al. (1965) to be an attractive processing method. To accurately design short-cycle adsorption units dynamic loading data for the adsorbed hydrocarbons are required. Dale et al. (1961) and Campbell et al. (1963) have experimentally determined the dynamic loading of several hydrocarbons on silica gel and have correlated the data through the use of the mass transfer zone concept. Dynamic loading data for carbon, however, has been seriously lacking in the literature. The purpose of these studies
was to obtain dynamic loading data for single, binary, and multicomponent hydrocarbon systems on carbon and to determine the effects of concentration, gas velocity, and pressure on the length of the mass transfer zone for the various hydrocarbons. Experimental
The adsorbent used was Columbia Grade NXC 4-6 mesh activated carbon. Nitrogen was used for the carrier gas. All hydrocarbons were Phillips Petroleum Co. pure grade. Ind. Eng. Chem. Process Des. Develop., Vol. 10, No. 1, 1971
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