PLENARY ACCOUNT Sulfide Oxidation: Sulfox System of Pollution Control Robert H. Rosenwald; Robert J. J. Hamblin, Peter Urban, and Robert P. Zimmerman Universal Oil Products Company, Corporate Research Center, Des Plaines, Illinnols 60016
Robert H. Rosenwald is a consultant to the Corporate Research Center of UOP Inc., Des Plaines, I l l . He received the Bachelor of Arts degree from North Central College i n Naperuille, Ill., in 1932. He was awarded the Master of Science degree i n 1934 and his Doctorate in 1936 both from the University of Minnesota at Minneapolis. Dr. Rosenwald joined UOP i n 1936 and has worked in the field of organic chemistry i n relation to the treating of motor fuels (inhibitors) and the isomerization and alkylation of isoparaffins He has done work on the chemistry and use of additives for the improvement and stabilization ofpetroleum products, rubber, polymers, and fats and oils. He has also conducted investigations of the chemistry of conversion o f sulfur derivatives, such as mercaptans and hydrogen sulfide, as required in refining technology. Dr. Rosenwald is a member of Phi Lambda Upsilon and Sigma X i i n addition to holding membership in the American Chemical Society.
Robert P. Zimmerman is assistant manager of deuelopmental research i n the Corporate Research Center of UOP Inc., Des Plaines, Ill. He is responsible for all pilot plant operations in catalvst evaluation resear.cn. nejoinea uvr in 1 ~ 4 as 1 a researcn cnemist after graduating from the Central YMCA, Chicago, Ill., with a Bachelor of Science degree. He earned the Master of Science'degree i n 1950 from the Illinois Institute of Technology. I n 1956 he was promoted to assistant research coordinator and he became research coordinator i n 1961. He assumed his present position in 1974.
218
Ind. Eng. Chem.. Prod. Res. DBY., Vol. 14. No. 4. 1975
R. J. J. Hamblin is currently a consultant i n the Chicago area. He received the €3ochelor o f Science (Hans.1 degree in chemi-
Master of Science degree i n chemical engineering i n 1951 from the University of Wisconsin at Madison. Mr. Hamblin was a project engineer with Trinidad Leasehold, London in 1949. He joined Procon (Great Britain) Limited i n 1953 working on project management. I n 1959 he transferred to the UOPProcess Division, engineering research and development department, and i n 1968 moved t o the Corporate Research Center of UOP Inc. working in the field of engineering aspects of chemical research. Mi-. Hamblin is a member of the American Institute o f Chemical Engineers.
Peter Urban is an assistant to the director of process research and deuelopment i n the Corporate Re.;.ii:; search Center of UOP Inc., Des Plaines, ILL He received his Bachelor of Science degree from Northwestern University, Evanstc,n, Ill., in .7948 and . . , uur rln-.. . joinea tnac. same year as a memoer or< the gasoline components research department. I n the twenty-seven years that he has been with the company, Urban has worked on antioxidants, metal deactivators, flue gas desulfurization, pollution control, and various problems in refinery treating processes. Mr. Urban is a member of the American Chemical Society.
Results on the conversion of ammoniacal solutions of hydrogen sulfide to elemental sulfur by oxidation with atmospheric oxygen are reported. Parameter studies in relation to the yields of sulfur and by-products were the effect of temperature, oxygen supply, catalyst, and a system of two-stage oxidation. Preliminary study was made with bench-top investigation followed by continuous flow experimentation. Catalyst studies considered the effect of carbon, metal sulfide, and metal complexes. Oxidation at temperatures to produce molten sulfur gave varying yields of sulfur and water-soluble products. A more efficient method was to oxidize a substantial portion of the sulfide under mild conditions to a polysulfide solution followed by liberation of sulfur by either (1) oxidizing the remaining sulfide under more severe conditions or (2) by steam stripping the solution.
Introduction The proper conversion of hydrogen sulfide is one of the more critical problems facing industry in its efforts against pollution. In the petroleum industry this is particularly true with recent trends toward hydrotreating and hydrocracking of crude oil stocks (Meredith, 1968; Yamato, 1968; Haensel, 1967; Steenberg, 1969). The problem is further accentuated by the ammonia liberated as formed from the nitrogen-containing constituents in stocks undergoing hydroprocessing. Likewise, this ammonia must be properly handled. This text describes investigations of conversion reactions of hydrogen sulfide. The entire discussion to be considered is that of oxidation with atmospheric oxygen. There are other capable oxidants for the conversion of sulfide to sulfur, but none can compete with air as a common, universal source of oxidant. However, any successful oxidation process must be selective in the products formed. The nature of these products depends on the particular conditions employed; products from elemental sulfur to sulfate salts can be obtained in varying degrees. This discussion and accompanying experimental results are presented on the basis that the most desirable product is elemental sulfur. This tenet is logical conclusion in view of industrial preference, the current relative monetary value of sulfur, ease of handling, and minimum pollution difficulties. The particular requirements for different refineries as to the nature of the product can vary. The favored position of elemental sulfur as product is of prime importance to the large-scale producer of hydrogen sulfide. There are other cases in which mere disposal of a minor sulfide stream is a problem. In these cases, oxidation to products such as sulfates can be the most convenient form for disposal. In order to discuss sulfur chemistry, there is a need to indicate the various forms of sulfur. The nomenclature as outlined in Table I will be used. All values of concentration and yield will be on the basis of sulfur in whatever form present. In a similar manner, ammonia concentrations will be calculated as NH3 in whatever form present. A further description of the sulfur compounds involved would concern valence and acidity of sulfur in the species under consideration. A chronicle of sulfur removal and conversion is an account of the chemistry of these two fundamental properties. The basic reaction of prime importance is the formation of sulfur as indicated by eq 1 2NH4SH
+0 2
2s'
+ 2NH4OH
(1)
in which the oxygen-to-sulfide-atom ratio ( O W - ) is 1.0. The preceding equation is written with the sulfur in the form of ammonium acid sulfide; it is this form which is to be considered throughout this paper. Oxidation to the polysulfide can likewise proceed with the same requirement, O/S2- ratio of 1.0, for the sulfide converted as shown in eq 2. 5NH4SH
+ 202
+
+
(NH4)2S4Sp 3NH4OH
+ H20
(2)
Further oxidations to thiosulfate, sulfite, and sulfate re-
Table I. Nomenclature System Symbol
ST So
Compound
= Total sulfur, includes all forms of sulfur = Elemental sulfur, uncombined.
No indication a s to extent of polymerization S2-
= Sulfur in polysulfide form, S,S,2-. Maximum value of x is 4. = Sulfide sulfur a s H,S o r salts,
S,-'
= Sulfide in polysulfide form, -S:,
SO ,,
Value of p is 1. = Sulfur a s oxides, and includes salts of sulfurous,
S,
e.g., (NH4lzSo r NH4SH
thiosulfuric, sulfuric, and polythionic acids. The ionic charge of the anion is not indicated. so4'- = Sulfur a s sulfate! S Z O ~= ~ ' Sulfur a s thiosulfate SO,,= Sulfur a s sulfite NH, = Ammonia found a s free NH3, ammonium hydroxide, and ammonium salts of HzS and above sulfur oxides quire O/S2- ratios of 2, 3, and 4,respectively, as indicated by the following equations. 2NH4SH
+ 202
+
+ 302 + 2NH4OH NH4SH + 202 + NH4OH
2NH4SH
+ H20 2(NH&S03 + 2H20 (NH4)2S04 + H20
(NH4)2S203 -+
(3) (4)
(5) The oxidation of sulfide solutions has been a subject of considerable study for many years (Sulfur, Vol. XIII, 1960). The oxidation process is subject to many variables including temperature, pH, concentration, catalysis, and means of contact. In view of this condition, it is not surprising that there is no mass of concordant data to describe rates and products of oxidation. The observations and results here presented include those from a preliminary bench-top experiment, but mainly deal with continuous flow experimentation. Of the possible means to obtain high selectivity to elemental sulfur by the oxidation procedure, regulation by thermodynamic considerations offers limited means for such control. If equilibrium conditions were realized in the presence of sufficient oxygen, the end-product would be essentially all sulfate. Calculations of the oxidation-reduction potentials of the sulfur-oxygen system point out that oxygen will be consumed until sulfur is converted to sulfate (Valensi, 1950). Attempts to regulate the extent of oxidation in the laboratory merely by limiting accessibility of oxygen were ineffective. Consequently, the problems associated with the sulfide oxidation process were considered mainly those of rates or kinetic control. Determination of the equilibrium constants for the following oxidation-reduction system H2S
-+
+ 4 H 2 0 F? S042-+ 2HC + 4H2
Ind. Eng. Chem., Prod. Res. Dev., Vol. 14, No. 4, 1975
219
also possibly acts as oxygen transfer agent, did not noticeably affect the rate. The oxidation of a polysulfide proceeds mol min-' l.-l), a t a considerably reduced rate (2.9 X indicating that the polysulfide sulfide (S,) reacts more slowly. The variations in catalytic effects and product selectivity indicate that the oxygen-sulfide reaction is a multistep process. The initial oxygen-sulfide interaction has been proposed by Abel (1957) as formation of a peroxy-sulfenic acid
0
HS0
40
SO
120
160
200
C o l o l y s t Concentration , p p m
Figure 1. Oxidation of H2S (initial concn. 0.5 M )in NHlOH solution (2.0 N ) ; air flow 0.425 1. min-'. verifies the strong tendency for sulfide to convert to sulfate, particularly at high temperatures in alkaline solutions (Ryzhenko, 1966). Preliminary, Bench-Top Experiments It was readily evident in the initial stages of this work on oxidation of solutions of sulfide that high selectivity was not obtained by oxidations in strong caustic solutions. Oxygen absorption studies revealed that oxidation proceeded to the thiosulfate and sulfite stages. The control of pH with buffer solutions in the range of about 8 to 10 gave encouraging results in regulating the extent of oxidation. I t was realized that ammoniacal solutions possessed suitable pH values and did give improved selectivity. In view of the fact that ammonia is a by-product from refinery processing, this discussion will concern only work using ammoniacal solutions as the preferred medium for oxidation. The preliminary experiments were made in order to characterize the oxidation process as to rates and selectivity. Included in Table I1 are results for runs with and without catalysts for ammoniacal sulfide solutions. Two wellknown catalysts for sulfide oxidation are included, namely carbon and a cobalt phthalocyanine. The promoting effect of carbon has been documented by Busch (1964), Kundo and Keier (1968, 1969, 1970), and Savel'ev (1969). Several laboratories have studied the catalytic properties of the metal phthalocyanines for this reaction. The oxidation of a well-aerated NH4SH solution was mol min-l l.-l), and was subject to slow (ca. 0.6-0.8 X catalysis. The catalyzed rates were found to be zero order in respect to sulfide concentration for over 50% of the run. Plotting duration against concentration gave a straight line, the slope of which is the rate value presented. The catalyst-promoted runs gave rate values about 10 times the uncatalyzed rates. The rates for various cobalt phthalocyanine tetrasulfonate concentrations, expressed as parts per million (ppm) by weight of solution, are shown graphically in Figure 1. The oxidation is somewhat dependent on the rate of air supply but, more important, the use of pure oxygen in place of air gave a large increase in rate (41.7 X mol min-' l.-l). I t is not known whether the rate is dependent on oxygen concentration for chemical reasons or whether oxygen transfer is the controlling factor. The products formed are to a great extent a function of the catalyst. This is illustrated by the results obtained with carbon and cobalt phthalocyanine tetrasulfonate, two materials which are similar kinetically. The phthalocyanine catalyst gave predominantly elemental sulfur (SO), the other water soluble products (&Om). The presence of xylene, which dissolves the sulfur and 220
Ind. Eng. Chem., Prod. Res. Dev., Vol. 14, No. 4, 1975
+
0 2
+
HSOO-
There is also the possibility of electron transfer to form a radical anion of the structure .OOS-, which enters into a chain sequence (Bowers et al, 1966). The catalytic effect of a cobalt phthalocyanine has been ascribed to the oxidation-reduction cycle of the various valence states of cobalt (Kundo and Keier, 1970). In this scheme there is no need for a direct interaction of sulfide and molecular oxygen. The catalytic effect of phthalocyanines has also been explained by Randin (1974) as the interaction of the metal chelate with oxygen, illustrated as follows. metal phthalocyanine
+0 2
-
[metal phthalocyanine]+
+02-
The donation of electrons from the chelate to oxygen produces two active species. These intermediates, which can be still considered as oxidants, are subject to reduction by sulfide and thereby give rise to a sulfide oxidation sequence. Catalyst Development On the basis of bench-top experiments, as previously described, a series of catalysts has been developed. These catalysts were designed for use in a solid bed type of reactor, the preferred mode of operation. These catalysts, identified as Sulfox-I, Sulfox-11, and Sulfox-I11 types, are based in most cases on the properties of three types of components, namely carbon, metal sulfides, and metal chelates. The following comments concern results obtained in a continuous flow reactor with the above three components. Carbon. The catalytic effect of carbon has been well documented by Krezil (1937), Benzuglyl (1937), Brusset (1950), Fischer (1964), Siedlewski (1966) and Sreeramamurthy (1975). Explanations have been proposed by Walker et al. (1968) and Coughlin (1969) to account for catalysis in general as observed with carbon, but these proposed theories have not been satisfactory in accounting for carbon differences as encountered experimentally. The catalytic properties of carbons varied considerably, depending on the source. This was particularly noticeable by the wide variations in product selectivity. Laboratory bench experiments, similar to those described in Table 11, showed a variation in elemental sulfur yield (SO) from a low of 22% of theoretical to a high of 89%. Those carbons in a granular form, upon evaluation in a continuous flow system, showed a similar but less pronounced variation in selectivity. Metal Sulfides. Metal salts are known catalysts for sulfide oxidation (Abegg and Elster, 1962). The oxidation of alkali metal sulfide solutions a t 20 and 105OC is accelerated by heavy metals. Commercial processes for removal of hydrogen sulfide from gas streams often rely on the presence of some heavy metal as catalyst (Goar, 1968). By the use of bench-top techniques, a number of metal sulfides, including sulfides of Co, Ni, Mn, V, Bi, Sb, Fe, As, and W, were screened as to activity and selectivity. Of
2.0N; Temperature, 40°C; Air Flow, 0.425 1. min-1
Table 11. Oxidation of Sulfide Solution (ca. 500 ml). "3,
Initial
Catalyst
Concn, PPm
None Carbon' Cobalt phthalocyanine tetrasulf onate tetrasulfonate tetrasulfonate tetrasulfonate tetrasulfonate tetrasulfonate tetrasulfonate tetrasulfonate * tetrasulfonate' tetrasulfonated tetrasulfonatee
0
.
Oxidation rate, x mol l.-' min-'
s2- concn, mol 1.
.
... 5 10 25 50 50 96 193 50 50
50 50 50
-'
% of s2charged as found SO ~"0,
...
0.6-0.8 ... 6.0 28 58 4.5 6.0 6.5 7.2 70 17 7.1 8.4 9.3 6.6 9.8 41.7 7.6 2.9 Oxygen in place of air. Presence of 100 ml of xylene. e Solution of poly-
ca 0.5 0.513 0.477 0.493 0.492 0.509 0.459 0.488 0.498 0.955 0.988 0.487 0.499 0.095
Commercial powdered carbon, 2.5 g. Air flow, 0.85 I. min-'. sulfide, 0.095 N in S P 2 - . a
... ...
e
.
.
...
... ... ... ... ... ... ... ... ...
... ... ... ... ... ...
... *..
...
Table 111. High Temperature Oxidation Charge Reaction conditions Temp, Press., "C wig LHSV
Air atom. ratio
Solution, wt % Catalyst
9-
3"
5.34 Sulfox -1 90 1.5 148 2.35 NiS' 50 1.0 125 2.35 NiSb 50 1.0 125 2.35 cosc 50 1.0 125 20% Ni. COSon A1203,20% co. a NiS on carbon, 63% Ni. NiS on AI&, these sulfides, the more promising ones, generally of the iron group, were evaluated in the continuous flow plant. These runs were conducted with a catalyst charge in the range of 50 to 100 ml. The liquid hourly space velocity (LHSV) values represent volume of liquid per hour per volume of catalyst. The duration for each run was a t least 32 hr with important runs extended to 4 weeks. In the use of a metal sulfide in a continuous flow plant, leaching of the metal by the water stream can occur. This is particularly severe if the sulfide (S2-) is completely oxidized, thereby allowing conversion of the metal sulfide to water-soluble salts. Metal Chelates. Some of the more efficient catalysts were metal derivatives of organic chelating reagents. The catalytic properties of a chelate and a sulfide of a given metal generally were different; the chelate structure brings about definite changes as to both activity and selectivity. By a proper selection of chelate and attachment to a substrate, a number of proprietary catalysts were developed and added to the Sulfox series. Continuous Flow Operation The proper design of a hydrogen sulfide unit is beset with the problems fundamental to the materials and reactions involved. The high selectivity desired is favored by low reaction temperatures, but this condition deposits elemental sulfur in the reaction and subsequent zones, and difficulties result. Actual continuous runs at temperatures below 12OOC were persistently plagued with either plugging problems or with catalyst deactivation due to sulfur deposi-.
3.17 0. a4 0.84 0.84
Product S2- charged converted, to, % S2-
o/s2-
%
2.1 1.0 1.0 1.0
100 85 68 56
so
+ s,
66 59 79 29
S"0,
37 8 11 24
tion. However, three means of operation were developed which allowed continuous plant operation and the collection of quantitative data as to conversions. These three modes of operation are: (1) oxidations a t temperatures above the melting point of sulfur (12OoC), (2) the use of an organic, water-insoluble solvent to remove sulfur, and (3) the two-stage process of initial low temperature followed by further processing in a second step. High-Temperature Operation. Oxidations in the temperature range of 12O-15O0C were investigated in the hope that complete sulfide conversion could be obtained without extensive overoxidation. This goal was not realized. With the Sulfox-I catalyst (Table 1111, an excess of oxygen (0/ S2- ratio of 2.1) was needed at a temperature of about 15OOC for complete conversion. The yield of water-soluble products (S,O,) was considered excessive (37%). The nickel and cobalt sulfides under less severe conditions (125O, 50 psig, O B 2 - ratio 1.0)gave incomplete conversions. Operations were erratic in this system; sulfur retention and release in the plant prevented quantitative yield calculations as is evident from the results presented in Table 111. These data, along with other unreported results, indicate that the simple direct oxidation at high temperatures is not a solution to the sulfide conversion problem. Furthermore the SO ,, products are not composed of a single compound, but are mixtures of various sulfur-oxyacid salts. S u l f u r Solvent. The use of a solvent such as toluene did indeed allow operation without the mechanical difficulties often encountered. As indicated in Table IV, complete sulfide conversion over Sulfox-I catalyst can be realized at a space velocity of 3 at a temperature of about 130°, but exInd. Eng. Chem., Prod. Res. Dev.. Vol. 14, No. 4. 1975
221
Table IV. Oxidation over Sulfox-I Catalyst in Presence of Toluene Solvent Charge Product
Air
Reaction conditions Temp, Press., "C psig LHSV
Solution, wt % S2-
atom. ratio
o/s2-
"3
S2- charged to, Yo
S2'
converted,
?h
Solventc
so + s,
sno,
124-1 32 60 3.0 1.92 1.02 2.2 3.9/1 100 88 67-77 60 6.0 1.95 2.0 1/3 62 46 0.82 80-94 15 15 5.87 6.46 1/3 50 37 0.97 63-89 60 3 10.1 10.8 1/3 67 57 1.19 a Inlet and outlet temperatures. Total LHSV, including aqueous solution and toluene. Volume ratio, aqueous to solvent.
22 25 12 13
Table V. First Stage Oxidation of Sulfide under Mild Conditions for Incomplete Conversion (Pressure, 5 psig) Charge Product
Air
Reaction conditions
Solution, wt ?b
Temp,"
a
converted,
S2- charged
o/s2-
B
s,
SnOm
2.7 2.4 3.5
S2-
to, %
"C
LHSV
Catalyst
S2'
35-75 35-80 42-64 3540 34-74 60-75
1.0 2.0 1.0 1.0 2.0 1.0
Sulfox -I Sulfox -I Sulfox -I Sulfox -n1 Sulfox -111 Carbon Glass Helices
11.1 11.3 6.03 6.15 6.15 6.55
6.2 6.2 3.38 3.59 3.59 3.66
1.04 0.98 0.90 0.88 0.92 0.86
66 58 77 80 74 67
63 59 73 74 57
1.0 12
6.43
3.65
0.88
25
15
5
60-65 1.0 Inlet and outlet temperatures.
cess oxygen (O/S2- ratio of over 2) was required. Sulfur (SO) yields of 80% and higher can be obtained. The use of a solvent permitted operation at reduced temperatures, that is, below 120' under conditions of incomplete reaction, low residence time, and low oxygen/sulfide ratios, but the yields of sulfur (So) and S,O, were not improved. It is evident that while this system does offer an advantage in operation at lower temperatures, the selectivity is still less than desired. Two-Stage Conversion Process. As the investigation progressed, the concept of a two-stage operation received more favorable consideration. In this mode of operation, the first stage, involving controlled oxidation, converts the ammonium sulfide solution to a solution of ammonium polysulfide [(NH4)2S,SP], followed by a second processing step to obtain sulfur. The successful use of this concept requires an efficient first stage oxidation to give 70-75% oxidation of the sulfide. This conversion value of less than 80% is limiting as the extent of oxidation for the maximum sulfur (S,) in soluble form corresponds to (NH&S4S,. The most favorable conditions for the first stage oxidation were indeed found to be mild, as indicated by data summarized in Table V. Results are presented with two different type catalysts, Sulfox-I and Sulfox-111. Oxidation of sulfide solutions at concentrations of 6 and 11% sulfide proceeded at inlet temperatures of 35', with outlet temperatures going to 60-80°. Reaction pressure of 5 psig was not only sufficient, but the more favorable. The O/S2- ratio need not be rigorously controlled at 0.8 to avoid overoxidation; ratios in the range of 0.8 to 1.0 were permissible. As previously indicated, overoxidation is avoided due to a greater stability of the polysulfide toward oxygen. Literature data of Schwarzenbach (1960), Giggenbach (1972), and Tedder (1969) reveal the distinct chemical properties of the sulfide and the polysulfide species. The oxidation as described proceeded to give a 95% yield of s,, based on the sulfide converted with about 5% going to S,O,. 222
atom. ratio
Ind. Eng. Chem., Prod. Res. Dev., Vol. 14, No. 4, 1975
"3
70
3.0
The importance of catalysis and of proper selectivity is demonstrated by runs with carbon and glass helices. Carbon did promote oxidation very similar to the conversion as obtained with a Sulfox catalyst, but the yield of S,O, products was high. With glass helices, the extent of oxidation was greatly reduced and the yield of S,O, was excessive. The conversion and yields in the first-stage oxidation can be accepted as reasonably sufficient. The next task is to take the effluent from stage one and devise suitable means to recover the sulfur (So) from the polysulfide ( S x ) . Two means to accomplish this will be considered. The first is decomposition by distillation, the second an oxidation procedure. Distillative Decomposition. The effluent from the first-stage oxidation can be subjected to a distillation process so the polysulfide sulfur (S,) will be thrown out of solution. This can be accomplished by stripping out ammonia and hydrogen sulfide, thereby decomposing the polysulfide (eq 6), a process for which the chemistry and technology have already been considered by Bretsznajder (1962). Preliminary data obtained in bench-top experiments indicated this approach was a reasonable one.
(NH1)2S,Sp + 2H20
-
2NH4OH
+ H2S + xSO
(6)
The most favorable conditions for operation of the column were under pressure with a pot temperature above the melting point of sulfur (120°C), and with provisions to adjust the rate of overhead take-off. With a pot temperature of 140' and a pressure of 38 psig, data were obtained as to the amount of free sulfur (So) liberated from the polysulfide solution in relation to the volume of overhead removed. These experiments were made with a plant effluent from the first-stage oxidation of the following weight-percent composition: S2-, 1.9; S,O,, 0.03; S,, 3.7; NH3,3.0. The relationship obtained is summarized in Figure 2. One observes an overhead of about 10 vol % releasing 90%
I'Or
1000 pslg
70 a c >
s
50
30
20
0'
40
60
10
80
V o l - % Overhead
Table VI. Second-Stage Oxidation of Polysulfide Solution to Liberate SulfuP
'
to S,O,
1.2 74 75 9.5 40 1.0 130 97 88 11.5 1.4 1.0 40 130 99 91 10.4 1.6 1.0 40 130 a Solution composition, Wt-% as S:S2-,1.3;S x , 3.73; SnOm, 0.247; SO4*-, 0.013; ST,5.35; wt '70"3, 3.0; catalyst, Sulfox-11. of the sulfur, with more thorough stripping releasing up to about 98% of the sulfur. As one approached the 20% overhead value, the sulfide (S2- and S P 2 - ) in the remaining aqueous solution in the pot was reduced to a low concentration; sulfide concentrations in the range of from less than 50 to 800 ppm were obtained. In conclusion, the distillative decomposition of the ammonium polysulfide solution proceeded to give a 95% or better recovery of sulfur by stripping out 20 vol % of the SOlution. Stage Two Oxidation t o Yield Sulfur. The second scheme was to oxidize the remaining sulfide from the Stage One in a second oxidation stage by increasing severity until sulfide was eliminated, at which point the polysulfide sulfur (S,) was released according to equation 7 .
+0 2
2(NH4)2SXS
-+
+
2 ( ~ 1)s'
+ 2H20 + 4NH3
(7)
A large number of runs was made to define those conditions which would lead to high sulfide conversion without ,, formation. Among the more favorable conexcessive SO ditions are those presented in Table VI. One observes that an excess of oxygen is required to realize high sulfide conversion. Under these conditions about 11%of the sulfide (S2-) and polysulfide (S,) is converted to SO ,, products. The yield of sulfur (SO) is thus about 90%. I t is essential to carry out the Stage Two oxidation without a high SO ,, yield at the expense of sulfur (So) produc,, can be considered tion. The conversion of Soor S, to SO as the known interaction of the sulfur with the ammonia liberated in the oxidation process. The following reactions (eq 8 and 9) have been observed by Schulek and Koros (1953) to proceed a t temperatures below 100OC. 6NH4OH 6NH40H
-
+ 4s'
+ 12s'
+ (NH4)2S203 + 3H20 2(NH4)2Ss + (NH4)2S203 + 3Hz0 2(NH4)2S
I
(8)
(9)
300 psig
1 0
Figure 2. Recovery of free sulfur (SO) from ammonium polysulfide solution by distillative decomposition.
Air Reaction conditions atom. to Temp. Press.. ratio verted, OC-' LHSV psig o/s~- YG . SO + s,
r
/
100
150
250
200
Reaction Temperature,
O C
Figure 3. Reduction of thiosulfate with hydrogen over metal sulfide catalyst. The solution process of converting sulfur to water-soluble products has been shown by Hartler et al. (1967) to be facilitated by the presence of sulfide ions (S2-) and polysulfide ions (SxSp2-) a t temperatures of 50 and 80'. Sulfur dissolution under comparatively high oxygen pressure has also been reported by Shiek et al. (1965). The SO ,, formation, as observed in our work, is to a large extent a product of the oxidation process and not merely the result of the aforementioned sulfur-ammonia interaction. Oxidation of a polysulfide solution under con,, (130°, 90 psig, O/S2- ratio of ditions leading to high SO 1.6, an alumina-type catalyst) was found to produce little S,O, if the oxygen was eliminated. This condition permits processing of polysulfide solutions without disproportionation to yield S,Om. Conversion of Overoxidized Products In the discussion so far, no detailed description has been ,., The given of those higher products designated as SO water-soluble, nonsteam distillable product is mainly ammonium thiosulfate. These were identified by thin-layer chromatographic inspection by Blasium and Kramer (1965), with further characterization by polarographic wave measurements and iodine titration. The presence of sulfites and polythionates has not been observed with normal operating conditions. The possible formation of sulfate was carefully followed. The formation of this product is a serious consideration because regenerative methods for sulfate removal are limited. I t has been our general experience that sulfate formation is not a problem. High-temperature oxidation has been observed to yield sulfate, but proper choice of reaction conditions and catalyst avoid the problem. It is evident that, regardless of the precautions taken, the formation of SO ,, (thiosulfate) in limited amounts will occur in any Sulfox arrangement. In order to prevent build,, in a recycle system, suitable means must be up of SO found to remove or convert these products. The approach taken to this problem was reduction to either sulfur or hydrogen sulfide. Two means to carry out this reduction will be described, namely, with hydrogen and with sulfide. Hydrogen Reduction. The reaction of thiosulfate with hydrogen produced sulfide as follows.
+
(NH4)2Sz03 + 4H2 2NH4SH 3H20 (10) Polythionates are reduced in a similar manner but sulfate resisted reduction under conditions to be described. The conversion of thiosulfate to sulfide at two pressures in relation to temperature is presented in Figure 3. The catalyst was a metal sulfide on a support, with an LHSV of -+
Ind. Eng. Chem., Prod. Res. Dev., Vol. 14, No. 4, 1975
223
g
The above reduction has been incorporated in the patent literature by Keller (1956) and Gerrard (1964), although it has not been studied in detail. The attractive features of using sulfide in this manner prompted a study of the interaction in a continuous flow unit. Passage of an ammonium thiosulfate solution with sulfide through a heated reactor with a 6-min residence time and a pressure of 650 psig gave reduction which depended on the thiosulfate/sulfide ratio. Results obtained with a solution of 4.50 wt % S2-, 2.04 wt % S2032-, and 4.36% NH3 (sulfur ratio of S2-/Sz032-, 2.2) are summarized in Figure 4. I t is observed that temperatures of 150’ and higher are preferred with near 100% reduction at 200’ with an excess of sulfide. The sulfur formed is obtained as polysulfide sulfur (S,) and soluble in the effluent.
20
a
z
ae
120
0
140
160
200
I60
T e m p e r a t u r e , C’
Figure 4. Thiosulfate reduction with sulfide (S2-). Air
f
S”
Figure 5. Sulfox-simple
sour water treater.
.
Mnt
“3
t
t
S’
H20
Figure 6. Sulfox-single-stage
t Aq Ammonia
plus thermal decomposition and
recycle. 2.0, and a &/ST mole ratio of 1O:l. A reduction of over 90% can be realized by operating a t a temperature of 200’ and a pressure of 1000 psig. Sulfide Reduction. The second method of thiosulfate reduction is the use of ammonia-hydrogen sulfide solutions according to eq 11 (NH4)2S203
+ 2NH4SH
4S0
+ 4NH3 + 3H20
(11)
With this stoichiometry, two atoms of thiosulfate sulfur consume two atoms of sulfide sulfur with the liberation of four atoms of sulfur (SO). In the presence of excess sulfide, elemental sulfur is not liberated but forms polysulfide sulfur ( S I ) , as shown in equation 12. (NH4)2S203 + 3NH4SH
224
*
(NH4)2S4Sp+ 3NH3
+ 3H20
Ind. Eng. Chem., Prod. Res. Dev., Vol. 14. No. 4, 1975
(12)
Process Configurations The chemical steps developed as described above can be regarded as building blocks which can be combined in various ways to fit almost any sulfide disposal situation. Contaminants in general present no problems; heavy hydrocarbons and solid matter must be removed, of course, and some materials, such as mercaptans and disulfide oils, will contaminate the elemental sulfur product unless special precautions are taken. The Sulfox process produces a regenerated ammoniacal absorbent stream which can be used for removal of H2S by vapor-liquid or liquid-liquid contacting. This stream has a greater capacity for H2S than conventional amines and the rich product is directly processable in the Sulfox unit without stripping. Sulfox can thus replace an existing amineClaus system, using existing equipment with minimal additions, offering higher capacity and elimination of Claus offgas problems. Some typical applications are described below. Direct Sour Water Treatment, Figure 5. This is a simple once-through operation from which the nonsmelling effluent is often suitable for disposal, having low BOD and COD, or for limited reuse. We should expect its prime application to be on low volume, low concentration streams (
