Ind. Eng. Chem. Res. 1993,32, 2490-2494
2490
Decomposition of Spent Alkylation Sulfuric Acid To Produce Sulfur Dioxide and Water Stephen Sung,? Gabor Szechy? and Lyle F. Albright' School of Chemical Engineering, Purdue University, West Lafayette, Indiana 47907 When sulfuric acid is used as the catalyst for the alkylation of isobutane with c3-C~olefins, highquality gasoline is produced in addition to small amounts of conjunct polymers, sulfur dioxide, and water. The latter two compounds are undesired byproducts formed by reactions between conjunct polymers and sulfuric acid. Both compounds are also formed during storage and transportation of the used sulfuric acid before it is regenerated to produce concentrated acid. T h e kinetics of the reactions between conjunct polymers and sulfuric acid were investigated from 10 to 60 "C for acids of different compositions; a kinetic equation was then developed. The formation of sulfone and hydroxyl groups on the conjunct polymers is also discussed. Up to several percent of a used acid can decompose during storage or transportation, if the acid is not handled correctly.
Introduction Alkylation of isobutane with C3-C5 olefins is generally considered to be the best method for producing high octane number gasolines for motor vehicles. This gasoline is cleanburning resulting in relatively low levels of partially oxidized compounds in the exit gases. In addition, light hydrocarbons including isobutane and olefins are sometimes added to gasolines causing undesired vapor emissions. Alkylation is a preferred method of converting such compounds into heavier and less volatile ones. Since the Clean Air Act legislates undesired gas emissions must be reduced, alkylation will almost certainly become increasingly important in the next several years (Chem. Systems, 1990). In the United States, sulfuric acid is used as a catalyst for production of about 53% of the alkylate or gasoline produced by alkylation (Albright, 1990). This percent likely will increase in the future since HF has, in the last several years, been found to be more dangerous than previously known. Under select conditions, HF can form aerosol clouds which contain lethal levels of HF and travel at ground level for several miles downward. Alkylation plants using sulfuric acid often consume 0.4-0.6 pounds acid per gallon of alkylate, and produce approximately 500000 barrels alkylate per day in the United States. Hence, about 10 million pounds of sulfuric acid are consumed each day. Most of this spent acid is transferred to sulfuric acid plants where it is regenerated to produce 98.5 to 99.5% fresh acid. Sulfur dioxide is formed to a small extent in analkylation unit, and in addition, it is produced and evolved from spent acids during storage or transportation. The following overall reaction is thought to be the predominant method of producing sulfur dioxide and water. H,SO,
+ CP
-
SO,
+ 2H,O + CP'
(1)
where CP is the initial conjunct polymer and CP' is a conjunct polymer that has two less hydrogen atoms and that probably has an additional double bond. Such a reaction is, of course, undesired. First, the evolution of sulfur dioxide and the production of water decrease the acidity of the catalyst and promote acid consumption. Second, if suitable precautions are not taken, part of the
* To whom correspondence should be addressed.
+ Present
address: University of Connecticut, Storrs, CT. Present address: Technical University, Budapest, Hungary. 0888-5885/93/2632-2490$04.00/0
sulfur dioxide can be lost to the atmosphere causing environmental problems and more sulfur must then be added when the acid is regenerated. Conjunct polymers are generally thought to be produced mainly from olefins during alkylation (Hofmann and Schriesheim, 1962;Den0 et al., 1964;Albright and Kranz, 1992). These polymers are also sometimes referred to as acid-soluble oils or red oil. They contain numerous double bonds, conjugated dienes, C5 rings, and branching (Miron and Lee, 1963). The conjunct polymers also react with sulfuric acid to form conjunct polymer sulfates (Albright, Spalding, et al., 1988). There is a need to clarify the chemistry of the decomposition of the spent acids and to investigate the kinetics of the reactions. This type of information was obtained in the present investigation.
Experimental Details A 250-mL three-neck glass flask was used as the reactor for the spent acids. A 250-300-g portion of acid was added and weighed in the flask a t the start of each run. The flask was immersed in a water bath that was controlled to within f0.3 "C. The temperature could be adjusted as needed in the range of 10-60 OC. Nitrogen gas was passed at a rate of about 300 mL/h through the gas space in the reactor in order to remove sulfur dioxide evolved from the acid. The resulting gas mixture was bubbled through a gas-washing bottle containing a dilute aqueous solution of hydrogen peroxide. Tests indicated that all of the sulfur dioxide was absorbed and converted to in this bottle. The reactor plus acid were weighed after a decomposition run. Two spent acids from a refinery and a synthetic spent acid were blended with 98% fresh HzS04 and water to produce acids containing various levels of water and conjunct polymers. The syntheticspent acid was prepared by bubbling l-butene through 98% fresh acid. The following analytical procedures were employed. The aqueous solution in the gas-washing bottle was generally titrated twice a day using a standardized caustic solution of about 0.25 M in order to determine the amount of sulfur dioxide absorbed. Both before and upon completion of a run, the acid was titrated for the acidity using the caustic solution and for the water content usinga Karl Fischer procedure. The hexane-soluble conjunct polymers were recovered from the spent acids by the procedure reported by Miron and Lee (1963),and water-solublepolymers were recovered 1993 American Chemical Society
Ind. Eng. Chem. Res., Vol. 32, No. 11, 1993 2491 -1.5L'
-3.5
0
40
80
120
160
200
Time [hours]
Figure 1. Sulfur dioxide evolution for duplicate runs with spent industrial acid at 40 O C . 3
-y2 Q
2.5
w
2
4-
if1.5 23 v)
1
0.5 0 0
4
8 rime)"
12
16
20
[(iiours)'
Figure 2. Sulfur dioxide evolution from spent industrial acid at 30-60 O C .
by a procedure developed by Hauser (1987). More details on the experimental procedures and equipment are reported by Sung (1992).
Experimental Results The amount of sulfur dioxide evolved from the sulfuric acids was investigated as a function of temperature (1060 "C), time (0-745 h), and acid composition. Several acids were investigated whose acidities varied from 80.7 to 96.4 w t % ;the water contents of the acids varied from 1.9 to 6.6 wt % ,and the conjunct polymers (CP) contents, as determined by difference, varied from 1.4 to 13 w t % . Figure 1shows the amount of sulfur dioxide evolved for duplicate runs a t 40 "C using an acid containing 2.5 w t % water, 7.4 w t % CP, and having an acidity of 90.1 w t % . The wt % sulfur dioxide (expressed as g of SO2 evolved/ 100 g of starting acid) was plotted versus time in hours. The results of these two runs agreed within about 5% on a relative basis, and 0.86 w t 7% sulfur dioxide was evolved after 170 h. The rate of sulfur dioxide evolution, as indicated by the slope of the curves, decreased significantly as the run progressed. Similar decreases in the rate of sulfur dioxide evolution occurred in all runs of this investigation. The following empirical equation modeled the kinetic results of all runs made to within approximately 10% on a relative basis. total w t % SO, evolved = k(time)0.5 (2) Figure 2 shows plots of the wt 5% sulfur dioxide evolved versus (timeP.5 for runs a t 30, 40, 50, and 60 "C. The results are represented fairly well by straight lines determined by a least-squares fit in which the lines are not forced to pass through the origin. Failure of the lines to pass through the origin (which indicates an induction period) can be explained as follows. First, when the starting acid was not saturated with dissolved sulfur dioxide as was often the case, 1-2 h of operation were
'
I
"
'
r -
k
00028
00030
0003:
(Temperalure)
00034
00036
[K ' 1
Figure 3. Arrhenius plots for spent industrial acid with different ages at 10-60 "C.
frequently requiredto saturate the acid; little sulfur dioxide evolved until the acid became saturated. When, however, the starting acid was already saturated at a low temperature with sulfur dioxide, part of the absorbed sulfur dioxide desorbed while the acid was heated to the run temperature; hence, there was rapid rate of sulfur dioxide evolution during start-up. A second explanation for the induction periods is the relatively slow rate of nitrogen flow; the nitrogen-sulfur dioxide mixture remained in the reactor flask on the average about 0.5 h before entering the gas-washing bottle. The slopes of the straight lines, shown in Figure 2, are the kinetic rate constants or k values. The rate of sulfur dioxide evolution can be determined by differentiating the above equation with respect to time: d(wt % SO, evolved) = 0.5k(time)-'*s (3) dt The experimental data plotted on Figure 2 are slightly better represented with S-shaped curves as compared to straight lines. With S-shaped curves, the slope (or values of k) decreases after the induction period slightly with time. A similar phenomenon was noted for most runs of the present investigation. Figure 3 indicates that the Arrhenius equation, k = Acm/RT,models the kinetic data for two sets of runs within about 5 % on a relative basis. The two sets were made using a spent acid having an acidity of 90.1 wt % and containing 2.5 wt % water and 7.4 w t % CP. The first set of runs conducted at temperatures from 10 to 60 "C was made when the spent acid was relatively new, approximately 1-4 months after being obtained from the refinery. Although all acids used in this investigation were stored at -15 "C, their acidities decreased slightly during storage while their water content increased slightly. When the acid was investigated at 30-60 "C approximately 1 year later, there were slightly smaller rates of sulfur dioxide evolution. The slopes of the two straight lines on Figure of 8.2 and 8.6 kcal/ 3 represent activation energies (a) (mol K). An average value of 8.4 is considered to a typical value for this and other used acids. Three decomposition runs were made at 50 "C for a spent alkylation acid during a period of about 1year. The first run was made about 30 days after the acid had been withdrawn from the alkylation unit. The other two runs were made about 130 and 360 days after the first run. As indicated by Figure 4, the k values at 50 "Cdecrease with time of storage at -15 "C. A k value of 0.135(time)-'.6 seems to be a reasonable extrapolation to zero time (or the time when the acid was in the alkylation unit). When the water contents of the acids varied from 1.9 to 6.6 w t %, the rates of sulfur dioxide evolution were essentially unaffected by the water content. Based on a series of three runs at 50 "C, the k values are, however,
2492 Ind. Eng. Chem. Res., Vol. 32, No. 11, 1993
O.OSOL,
"
'
I
'
"
100
0
'
I
'
"
200
'
'
'
"
'
400
300
Time [days]
Figure 4. Effect of aging spent industrial acid at -15 'C on kinetic constant at 50 "C. W'" " " ' ---I
0.1
0.08
1 0
1
2
3
4
5
6
WI% of Coiijuncl Polymer
Figure 5. Effect of conjunct polymer contents on kinetic constant
Material balances were performed in order to clarify the chemistry of the decomposition reactions. The results are summarized as follows: (a) The decrease in the weight of the acid was approximately equal to the calculated weight of sulfur dioxide absorbed in the gas-washing bottle. (b)The moles of sulfuric acid decomposed as determined by analytical measurements were essentially equal to the calculated moles of sulfur dioxide evolved. (c) The calculated molar ratio of water apparently produced to sulfur dioxide evolvedwas 2 or slightly greater. The actual ratio is, however, probably less than the calculated ratio. Although precautions were taken to minimize contact between the acid and humid air, small amounts of water were likely absorbed in all experiments from the air resulting in an apparent ratio that is somewhat too high. The reactor and connecting lines also had to be dried since water was found to be adsorbed on the surfaces in preliminary experiments. Conjunct polymers were extracted in several cases from the acids both before and after the decomposition runs. The following results were obtained: (a) The color of the extracted CP's was darker after the decomposition runs. (b) Appreciably less hexane-soluble CP was obtained following a decomposition run while relatively higher percents of water-soluble CP were obtained.
at 50 O C .
Discussion of Results
essentially directly proportional to the CP content of the acid, as shown in Figure 5 in which the CP content varied from 0 to 5 w t % CP. No significant amount of sulfur dioxide evolved for an acid containing no CP, and k values for these runs are modeled as follows: kmec = O.O193(wt % CP) (4) Seven additional runs were made at 50 "C using blends prepared from the synthetic spent acid containing 8.1 w t 5% CP, 89.0 w t % sulfuric acid, and 2.9 wt % water. The water content was again found to have little effect on the rates of sulfur dioxide evolution, but the rate was strongly dependent on the CP content. The runs were, however, made over aperiod of about 6 months. In order to "correct" the k values to the same age of acid, the results of Figure 4 and the following equation were used.
The decompositions at especially higher temperatures such as 40 and 50 "C can be considered as accelerated tests of the degradation that occurs when used alkylation acids are stored for extended periods of time. During storage or transportation, acid temperatures often approach ambient temperatures. Presumably in most industrial units or transportation vehicles, the temperatures will be 40 "C or lower. Under unusual conditions, significantly higher temperatures may however occur. The decrease in the rate of sulfur dioxide evolved during the course of a run is explained mainly by the partial depletion of the more reactive hydrogen atoms in the conjunct polymers during the initial stages of the decomposition. Various hydrogen atoms attached to the conjunct polymers have different reactivities. The more reactive hydrogen atoms are likely the tertiary and allylic hydrogens. Such hydrogens are initially present in relatively large amounts since the conjunct polymers contain numerous branches and C-C double bonds (Miron and Lee, 1963). The less reactive hydrogens probably are the secondary and primary ones. As sulfur dioxideis produced, the H/C ratio of the CP undoubtedly decrease. Eventually a carbon-rich solid deposit forms; industrial personnel have reported the presence of solid black particles in the acid after extended storage periods. Some black globules of a tarlike material were noted in one run in this investigation. The conjugated double bonds or single double bonds known to be present (Miron and Lee, 1963) are also reactive. For example, the following reactions certainly occur: polymerization of CP to form higher molecular weight polymers; formation of conjunct polymer sulfates when sulfuric acid reacts with a double bond of the conjunct polymer (Albright, Spalding, et al., 1988); and sulfone formation. The sulfones can be produced by reactions with sulfur dioxide or sulfuric acid with double bonds of the CP. Some sulfones are thought to have been produced during the decomposition runs. Such sulfones would explain in part the color changes of the recovered polymers and relatively increased recovery of the water-soluble
(") =(") Corrected k values for the seven runs were essentially
(5)
kydays
Fig4
'ydays
other acids
directly proportional to the CP content for values up to 8 w t 9%.
The synthetic spent acid mentioned above partially froze at -15 "C. The unfrozen portion contained almost 13 wt % CP. This portion was tested in two runs at 50 "C, and the k values found were about 50 % higher than the values predicted based on extrapolation of the straight line model developed at CP values up to 8 w t % . Partial freezing is, hence, a technique that partially separates CP polymers and also apparently results in more reactive polymers. The following empirical equation models the results to within about 10% on a relative basis with different acids at 10-60 "C. wt 5% SO, evolved = 0.193(wt % CP)(time)o.se x p [4.2 E - 4.2 (6)
71
where time is in hours, T is in K, and 4.2 = 8.4/R, since R = 1.987 kcal/mol.
Ind. Eng. Chem. Res., Vol. 32, No. 11, 1993 2493
4 T-40C
6-
$3 1 5 5 0
0
1
2
3
4
5
6
7
8
Storage Time [day$]
Figure 6. Estimated costa of decomposedsulfuric acid during storage or transportation. Baais: lo00barrels of alkylate produced per day.
polymers. Some hydroxyl groups likely also were produced by the following overall reaction: OH
(7)
Similar reactions are used to produce ethanol and 2-propanol from ethylene and propylene, respectively. Reactions such as indicated above would suggest that overall the molar ratio of the net water produced to sulfur dioxide produced may be slightly less than 2. The empirical kinetic equation developed here is likely sufficiently accurate in most cases. A more theoretical equation might however be developed using second-order kinetics. In such an equation, terms could be provided for reactions between sulfuric acid and both the more reactive and less reactive hydrogen atoms. As the decomposition progressed, the H/C molar ratio of the polymers likely decreases from a starting ratio of about 1.75/1 to much lower ratios. A preferred kinetic model should include terms for the formation of sulfones, sulfonic acid, higher molecular weight CP, and possibly hydroxyl groups. Sufficient information for such side reactions is, however, currently unavailable. The results obtained here can be used to estimate the amount and cost of sulfuric acid that decomposes during storage or transportation. The following assumptions are used to demonstrate the calculation procedures to be used: (a) acid consumption, 0.5 lb/gallon of alkylate; (b) feed acid composition, 99.5 wt 9% H2S0.4and 0.5 wt 9% water; spent acid composition, 90.1 wt 9% H2S04 and 7.4 w t 9% CP; (c) the k value of the spent acid at 50 "C is 0.135 h-0.5 (at zero time as shown on Figure 4), the k values at 10,20,30,and 40 "C are then 0.021,0.035,0.057,and 0.089 h-0.5 respectively, as calculated using the Arrhenius equation; (d) the cost of fresh sulfuric acid is $0.06/lb. Based on the above conditions and using 1000 bbl of alkylate/day as a basis, the amounts and costs of sulfuric acid decomposed during storage and transportation were calculated for temperatures in the range of 10-40 "C and for times up to 7 days. The costs per day are shown in Figure 6 as a function of temperature and time. As indicated, the costs become appreciable especiallyat higher temperatures and longer times. For a large alkylation plant that Droduces 20 000 barrels of alkvlate/dav, the annual cost may vary from about $15 000 to $176000. It was assumed that the total storage time varied from 1 to 7 days. In this regard, one refinery indicates that they do not transport their spent acid immediately since it is too "active". In other plants, however, the spent acid is piped rather quickly to the acid regeneration plant. In the latter case, the cost of acid decomposition is apparently greatly reduced. There is an obvious incentive to decrease the times and temperature in order to reduce the amount of
acid decomposition. Reducing acid consumption to less than 0.5 lb/gal would also decrease these costs. In this regard acid consumption values as low as 0.2 lb/gal have been reported in a commercial unit (Knepper et al., 1977). Provisions to minimize if not prevent any sulfur dioxide losses to the atmosphere likely will always be cost-effective. The results can also be used to estimate the amount of sulfuric acid that decomposes in the decanter of an alkylation reactor. The following conditions are used in calculating the costs of the acid lost to decomposition: (a) temperature in decanter, 10 "C; (b) total residence time of acid in decanter, 20 h; the acid is assumed to recycle 20 times and to remain in the decanter for 1 h each time; (c) the average acid composition is assumed to be 94 wt % H2S04 and 3.7 wt % CP, such a composition is approximately an average between the compositions of the feed and spent acids; (d) the k value at 10 "C for this acid is 0.0106 has5. Based on the above assumptions, 16.2 lb of acid decomposes/day for every 1OOO bbl of alkylate/day. For an alkylate unit producing 20 OOO bbl/day, the daily cost is $19.50. Reducing the residence time in the decanter, as is practiced in certain refineries, or using centrifuges would reduce the amount of acid that decomposes. Centrifuges would not be cost effective based solely on the savings due to reduced decomposition of the acid, but they do offer two additional benefits. First, there would be reduced decomposition of trimethylpentanes to form lower quality alkylates (Doshiand Albright, 1976). Second, acid consumption would be reduced. There would be less formation of conjunct polymers from dissolved sec-butyl acid sulfates and isopropyl acid sulfates (Albright, Doshi, et al., 1977). The results of this investigation help explain the socalled "acid runaway" phenomenon which occurs infrequently during a deviation from the preferred operating conditions in an alkylation plant. Plant personnel report that the following all occur in at least some runaways. First, and probably the most important, the acid strength decreases rapidly causing a sharp decrease in the catalytic activity of the acid phase. Second, both the quality and amount of alkylate decrease. Third, sulfur dioxide formation increases significantly sometimes resulting in a substantial increase in pressure. Fourth, tarlike materials sometimes form. Runaways occur only when the acid strength is low and substantial amounts of conjunct polymers are present in the acid phase. Both n-butenes and propylene, however, still continue to react readily with the acid to form sulfate esters which further increase the concentration of organic materials dissolved in the acid phase. During the runaway, the sulfuric acid obviously acts as an oxidizing agent for these organic materials as occurred in the present investigation. Gardner (1991) has, however, stated that carbon dioxide and hydrogen are also produced presumably in the reactor and/or the decanter during runaways. He has however provided no supporting evidenceor details. The results of the present investigation do not support his claim nor is any evidence known to support it.
Summary During storage and transportation of spent alkylation acids, time, temperature, and conjunct polymer content are important factors affecting the evolution of sulfur dioxide. By reducing the value of these three variables in an alkylation unit, the amount and cost of acid decomposition will be decreased.
2494 Ind. Eng. Chem. Res., Vol. 32, No. 11, 1993
Acknowledgment Rhone Poulenc Basic Chemicals Co. provided financial support for this project. Literature Cited Albright, L. F. Alkylation Will be Key Proceea in Reformulated Gasoline Era. Oil Gas J. 1990,88 (46), 79-82. Albright, L. F.; Kranz, K. Alkylation of Isobutane with Pentanes Using Sulfuric Acid as a Catalyst. Znd. Eng. Chem. Res. 1992,31, 475-481.
Albright, L. F.; Doshi, B.; Ferman, M. A.; Ewo, A. Two-step Alkylation with Cd Olefins: Reactions of Cd Olefins with Sulfuric Acid, Industrial and Laboratory Alkylations; Albright, L. F., Goldsby, A. R., E&.; ACS Symposium Series 55; American Chemical Society Washington, DC, 1977; pp 96-108. Albright,L. F.; Spalding,M. A.; Kopser, C. G.; Eckert, R. A. Alkylation of Isobutane with Cc Olefins: Production and Characterization of Conjunct Polymers. Znd. Eng. Chem. Res. 1988,27,386-391. Chem Systems, Huge Alkylation Expansion Forecast. Hydrocarbon Process. 1990, 70 (9), 19. Deno, N. C.; Boyd, D. B.;Hodge, J. D.; Pittman, C. U.; Turner, J. 0. Fate of t-Butyl Cations in 98% Sulfuric Acid. J. Am. Chem. SOC.1964,86, 1745-1748.
Doehi,B. M.; Albright, L. F. Degradation and IsomerizationReactions during Alkylation of Isobutane with Light Olefins. Znd. Eng. Chem., Process Des. Dev. 1976,15, 53-60. Gardner, A. S. Today’s Refinery 1991, Dec, 20-21. H a w r , R. Extraction of Water-Soluble Conjunct Polymers. Report, Purdue University, West Lafayette, IN, 1987. Hofmann, J. E.; Schriesheim, A. Ionic Reactions Occurring during Sulfuric Acid Catalyzed Alkylation. J. Am. Chem. SOC. 1962,84, 953-961.
Knepper, J. C.; Kaplan, R.D.; Tampa, G. E. Computer Modeling of Alkylation Units, Industrial and Laboratory Alkylation. Albright, L. F., Goldaby, A. R., Eds.; ACS Symposium Series 55; American Chemical Society Washington, DC, 1977; pp 260-270. Miron, S.; Lee, R. J. Molecular Structure of Conjunct Polymers. J. Chem. Eng. Data 1963,8,150-160. Sung,S. The Decompceition Kinetics of Spent SulfuricAcid Obtained from Alkylation Plants. MS Theais, Purdue University, West Lafayette, IN, Aug 1992. Received for review March 1, 1993 Revised manuscript received July 26, 1993 Accepted August 2, 19930 0
Abstract published in Advance ACS Abstracts, October 1,
1993.
