Ind. Eng. Chem. Process Des. Dev. 1982, 21, 569-574
nane, isopropylcyclohexane, or trimethylbenzene. These mixtures present a wide range in boiling points, aromaticity, and acentric factor to test the correlation. Results of the generalized correlation for these wideboiling systems are shown in Table VII. The relative poor accuracy most likely is due to the limited number (18)of experimental points used for this comparison. Comparison with Other Predictive Methods A comparison of the modified RK method was made with the Soave and the Peng-Fbbinson methods. The data sets of Tables V and VI for paraffins, olefin, H2, N2, CO, C02, and H2S were used. Results of the evaluation are given in Table VIII. This shows that the errors for K values by this work were approximately 50% of those predicted by the Soave or the Peng-Robinson methods. Conclusions The modified RK procedure has been shown to be a versatile tool for computer calculation of phase equilibrium behavior for hydrocarbons, with and without the common noncondensable gases. The procedure is fully generalized and requires only a knowledge of the critical properties, acentric factor, and the molar volume of each component. It has been tested on binary and multicomponent systems of wide-boiling, sub-, and supercritical mixtures but should be used with caution on close boiling systems or systems known to form azeotropes until further study and testing is performed on these mixtures. Nomenclature A = aP/R2P a = energy constant, eq 2 B = bP/RT b = volume constant, eq 3 d = error, % f = fugacity, psia k = binary interaction coefficient K = equilibrium value = y / x
569
P = pressure, psia R = gas constant = 10.7335 t = temperature, O F T = temperature, O R u = pure-component volume V = specific volume x = mole fraction in liquid y = mole fraction in vapor 2 = compressibility factor 0 = parameters in the "a" and 'b" terms of the RK equation of state w = acentric factor u2 = variance Subscripts a, b = general reference to the two constants in the RK
equation of state c = critical property i j .... = component, i j , etc. m = molar Superscripts L = liquid phase V = vapor phase
* = effective binary temperature Literature Cited
Chueh, P. L.; Prausnltz. J. M. AIChE J . 1987, 73. 1099. Eakin, B. E.; DeVaney, W. E., Chem. Erg. frog. Symp. Ser. 1974, 740(70), 80. Graboski, M. S.; Daubert, T. E. Ind. Eng. Chem. Process Des. Dev. 1978, 77, 443. Graboski, M. S.; Daubert, T. E. Ind. €ng. Chem. Process. Des. Dev. 1979, 78, 300. Lenoir, J. M. Pet. Refiner 1980, 3 9 , 135. Peng, D.-Y.; Robinson, D. B. Ind. Eng. Chem. Fundam. 1978, 75, 59. Soave, 0. Chem. Eng. Sci. 1972, 27, 1197. Starling K. E. "FIuM Thermodynemlc Properties for Light Petroleum Systems". Gulf Publishing Co.: Houston, 1973. Wilson, G. M. A&. Cryog. Eng. 1984. 9 , 168. Zudkevitch, D.; Joffe, J. AIChEJ. 1970, 76, 112.
Received for review April 7, 1980 Revised manuscript received November 19, 1981 Accepted March 29, 1982
Thermal Degradation of Sodium Citrate Solutions Containing SO2 and Thiosulfate Gary T. Rochelle" and Mlchael Y. Glbson Department of Chemlcal Engineering, The Universlty of Texas at Austln, Austin, Texas 78712
The stability of citrate buffered sulfur dioxide solutions, with and without thiosulfate, was studied between 70 and 180 O C . Solutions were sealed in a stainless steel vessel, heated, and sampled periodically. The samples were analyzed for citrate, sulfur dioxide, sulfate, thiosulfate, and hlgher polythlonates. Two predominate reactions have been identified. A reaction involving both citrate and sulfur dioxide, in a ratio of 1:1, has an activation energy of 43 kcai/moi. Sulfotricarbailylic acid has been identified as a possible product. A second reaction involves disproportionation of sulfur dioxide to thiosulfate and sulfate. This reaction has an activation energy of about 45 kcal/moi and is catalyzed by thiosulfate. Rate equations and reaction mechanisms are proposed for both reactions. The rate of degradation of 100 O C and pH 4.0 should pose no significant problem for stack gas desulfurization by absorption/stripplng.
Introduction Aqueous scrubbing followed by steam stripping is a potentially attractive method of desulfurizing stack gas with the production of concentrated SO2 (Rochelle, 1977; 0196-4305/82/ 1121-0569$01.25/0
Johnstone et al., 1938). SO2 is absorbed from stack gas containing 500 to 5000 ppm of SO2by an aqueous solution at 30 to 60 OC. The solution is regenerated by stripping with steam a t 80 to 150 "C. Liquid water is easily con@ 1982 American Chemical Society
570
Ind. Eng. Chem. Process Des. Dev., Vol. 21, No. 4, 1982
densed from the stripper overhead vapor, leaving concentrated SO2. Sodium citrate was recognized as a potential aqueous absorbent for absorption/stripping as early as 1934 (Boswell, 1934; Applebey, 1937). It has recently reappeared in work by the US.Bureau of Mines (Nissen, 1976), in process development sponsored by Peabody, Inc., and in a process offered by Flakt, Inc. (Bengtason and Farrington, 1979). The Electric Power Research Institute (EPRI) is supporting development of the Flakt process. This paper is a report of work sponsored by EPRI on the stability of sodium citrate buffer solutions at stripping conditions. Irreversible side reactions may determine the maximum operating temperature of an absorption/stripping system using sodium citrate buffer. Disproportionation of bisulfite will produce other sulfur species such as sulfate and thiosulfate that accumulate in the solution and must be removed as soluble waste products. Direct reaction of bisulfite with citrate may degrade buffer capability and affect other solution properties such as solubilities. Significant disproportionation has been observed in the Wellman-Lord process (Bailey, 1974) and the NH,-steam stripping process (Slack, 1972) when regeneration is carried out at temperatures greater than 100 "C. In these cases disproportionation is minimized by operating at lower temperature. With the aluminum sulfate absorption/ stripping process, Applebey (1937) observed that disproportionation was catalyzed by thiosulfate, one of the disproportionation products. They eliminated thiosulfate by reaction with copper. Rochelle (1977) proposed a two-step reaction that involves intermediate formation of trithionate (S306,-)with net formation of thiosulfate (S20,2-)and sulfate (SO:-) 2H+ + 4HS03-
+
+ SZO3'-
2S3062- 2Hz0
-
-
2S3OGZ+ 3Hz0
2S042-+ 2S2032-+4H+
4HSOL -,2S042-+ SZO3'-
+ 2H+ + HzO
The rate of formation of trithionate has been measured at 70 OC and is given by (Battaglia and Miller, 1968) HS03-] rate = k[H+]3[Sz0,2-]3[ k70oC
= 2.2 X lo6 M4 s - ~
The hydrolysis of trithionate appears to be independent of pH and varies directly with [S302-] rate = k [ S 3 0 2 - ]
kloo0c= 1.43 h-' (Foerster and Hornig, 1922) On the basis of other data at 30 "C (Foerster and Hornig, 1922), 40 "C (Goehring et al., 1947), and 50 "C (Kurtenacker et al., 1935), the activation energy of hydrolysis should be 15-20 kcal/g-mol. In a batch system, this reaction sequence would be autocatalytic, since the rate increases with thiosulfate concentration which in turn is produced by the reaction. Nawisky and Sprenger (1943) patented the production of sulfotricarballyate acid by the two-step reaction of sodium bisulfite and citric acid. At temperatures above 160 "C, Usel'tseva et al. (1970) and Popov and Micev (1962) have shown that thermal decomposition of citric acid produces aconitic acid as well as itaconic, citraconic, and mesaconic acids HOOCCHZ-HOCCOOH-CHzCOOH (citric acid) HOOCCH=CCOOH-CH&OOH + H2O (aconitic acid) --+
THERYQIETER
w
H 1
SAURE OUTLET
REACTION
VESSEL
Figure 1. Experimental apparatus.
By analogy to the reaction of maleic acid with bisulfite (Hagglund and Ringbom, 1926), aconitic acid should react with bisulfite to give sulfotricarballylic acid.
-
HOOCCH=CCOOH-CH2COOH + HS03(aconitic acid) HOOCCHz--"03-CCOOH-CH2COOH (sulfotricarballylic acid) The objective of these experiments was to quantify the rates of these side reactions at 100 to 160 "C. Rates of SOz (total S4+species) disappearance from sodium citrate solutions were measured with and without thiosulfate. The measured rates are used to estimate the maximum stripper temperature as a function of solution composition and residence time in the stripper. Experimental Methods Rates of SO2disappearance have been measured using the stainless steel reaction apparatus shown in Figure 1. Three hundred milliliters of solution was prepared gravimetrically and sealed in the reactor. The reaction vessel was heated to the reaction temperature in 30 to 60 min. Temperature was then controlled at the specified level f 2 "C for the duration of the experiment. The main disadvantages of the apparatus were the possibility of corrosion and the lack of visual observation. Samples (10 mL) were withdrawn at appropriate times and analyzed for total sulfite, thiosulfate, sulfate, and citrate. Sulfite and thiosulfate were determined by iodimetric titration with and without formaldehyde (Kolthoff and Belcher, 1957). Sulfate was determined in runs 1-N through 27-N by titration with barium perchlorate using sulfonazo I11 as an endpoint indicator. In later runs, an ion chromatograph was used to determine citrate and sulfate quantitatively and to give a qualitative analysis for sulfite, thiosulfate, and polythionates. Results of typical experiments with and without thiosulfate are given in Figures 2 and 3. Except during the initial warmup period and after almost complete SOz disappearance,the total concentration of dissolved SO2was a linear function of time. For each run a regression analysis was performed to determine the slope of this line, taken to the rate of SO2 disappearance for each run. A summary of these data is presented in Table I. Results Without Thiosulfate Experiments were conducted in sodium citrate buffer without thiosulfate over the ranges: [citrate] = 0.2 to 1.0 M; [SO,] = 0.05 to 0.2 M; T = 413 to 453 K (140 to 180 "C); pH = 3.36 to 4.63. Selected data from these exper-
Ind. Eng. Chem. Process Des. Dev., Vol. 21, No. 4, 1982
571
-
0.20
010-
0
2
4
6
E 1 TIME (hrl
0
Figure 2. Run 11-N, 0.5 M citrate, 0.7 M Na+, 159 "C. 01
I
1
I
I
3
6
9
12
TIME
lhrl
Figure 4. Run 30-N, 0.2 M citrate, 0.9 M Na+, 180 "C.
06
1
@
Citrate
sop
TIME (hr)
Figure 3. Run 9-N, 0.5 M citrate, 1.4 M Na+, 0.2 M S 2 0 t - , 161 "C.
iments (indicated in Table I) were correlated by the expression 1 0
where In A = 48.1 f 11.48%;Hcit = 43 X lo3 (cal/g-mol) f 10.57%;a = 0.92 f 17.8%;b = 0.178 f 75.5%;Rcit is the rate of reaction of SO2 with citrate (M/h). The standard deviation of the estimate of In Rat is f0.29. This corresponds to an error of about 30%. The rate of SO2 disappearance appears to be independent of the SO2 concentration. Five runs (1, 2 - 4 8, 19, 29-N) a t constant citrate concentration and pH but with SO2 concentrations from 0.05 to 0.20 M show no variation of rate with SO2 concentration. The linearity within individual runs such as run 11-N (Figure 2) is also consistent with a zero order dependence on SO2. The approximate first-order dependence of rate on citrate concentration is substantiated by runs 31, 33, and 34-N. Participation and consumption of citrate is also confirmed in run 3-N (Figure 4) where the rate of SO2
I
I
I
I
3
6
9
12
TIME (hr)
Figure 5. Run 31-N, 1.0 M citrate, 2.5 M Na+, 163 "C.
disappearance decreases drastically as the amount of SO2 lost approaches the initial concentration of citrate. However, most runs with 150% excess citrate showed no effect of citrate loss on the linearity of SO2 disappearance with time. Three runs with acetate buffer (25,26,27-N)confun that little reaction occurs in the absence of both citrate and thiosulfate. The activation energy of the reaction was found to be 43 kcal/g-mol. The dependence of the reaction rate on pH is not well-defined, but appears to be quite small. Analysis for citrate during runs 31-N (Figure 5) and 33-N and the previously discussed effects of run 30-N show that citrate and SO2 disappear on a one to one stoichiometric
572
Ind. Eng. Chem. Process Des. Dev., Vol. 21, No. 4, 1982
Table I. Summary o f Experimental Data on SO, Disappearance from Sodium Citrate Buffers composition, M
rate X 10' M/h
citrate
Na'
SO,
S20,2-
acetate
pHh
T,"C
0.5
1.2
0.2
0
0
4.17 (4.63) 4.7 (4.63) (4.63) 4.7 3.36 (4.61)
140 146 150 156 158 168 159 135 150 150 150 163 168 150 161 140 147 158 180 130 137 150 167 166 180 132 163 163 164 152 150 152 150 168 163 70 168 157 170 160 152
0.5 0.5
0.7 1.3
0.2 0.2
0 0.05
0 0
(4.61) 4.61 0.5 0.5 0.5
0.5 1.4 1.6
0.2 0.2 0.2
0.05 0.1 0.2
0 0 0
0.5
1.35
0.2
0.2
0
0.5 0.5 0.5 0.2 0.214 0.75 1.0 0.1 0 0 0 0 0 0 0 0.5 0.5 0.5 0.5 0.5
1.75 1.8 0.84 0.9 0.496 1.796 2.5 0.95 1.2 1.03 0.838 0.8 1.1 1.15 3.1 1.3 1.2 1.2 1.2 1.3
0.35 0.4 0.4 0.5 0.023 0.2 0.5 0.2 0.2 0.05 0.14 0.2 0.2 0.2 0.81 0.2 0.2 0.2 0.2 0.2
0.2 0.2 0.02 0 0.0225 0 0 0 0
0 0 0 0 0
0 0 0.05 0.05 0.2 0.85 0.05 0.0 0.0 0.0 0.05
0 0 1.0 1.5 1.81 6.05 0.5 1.356 1.0 4.1 0.0 0.0 0.0 0.0 0.0
3.58 (4.61) 4.58 4.58 4.63 4.58 4.16 4.16 4.65 3.30 (4.63) (4.63) 4.88 5.16 4.74 3.57 5.7 4.7 4.87 4.1
4.50
measd 1.17 3.43 2.8 9.04 10.9 20 17.2 0.63 > 90 5.15 4.40 16.5 36.4 12.8 22.3 5.86 12.5 47.7 >171 4.42 8.26 26.8 145 179 19.2 0.519 26.5 26.5 4.22 0.93 1.125 10.5 0.522 3.77 24.4 1.57 30.5 11.5 38.7 15.0 6.93
calcd 1.20 2.12 3.33 7.07 8.94 27 .O 16.8 0.615 4.09 4.09 4.09 19.4 32.3 39.2 23.9 4.18 11.8 38.7 566 5.03 13.0 15.5 107 141 44.3 0.176 23.1 30.1 3.65 N .A. N.A. N.A. 0.012 4.45 22.5 0.00653 33.6 7.95 34.9 11.8 5.22
run no. notes 12-N 19-N 1-A 29-N 8-N 2 -A 11-N 12-A 3 -A 4-A 5 -A 1-N 6-A 18-N 9 -N 13-N 15-N 7 -N 14-N 24-N 23-N 28-N 38-N 37-N 30-N 3 -N 34-N 31-N 33-N 25-N 26-N 27-N 8-A 9 -A 32-N 13-A 7-A 4 -N 5 -N 10-N 2 -N
a a a a a
b b
b b b b b b b
b b a a
a
C
d
e
f g
Runs used in citrate rate regression analysis. Runs used in thiosulfate rate regression analysis. 0.1% hydroquinone. 1 mM Cr,(SO,), . e 0.05 M Fe. f 1 mN FeSO, ; 1 mN Cr,(SO,), ; 1 mN Mn(S0,); 1 mN Ni(NH,),(SO,). 1 mM FeSO, . Parentheses indicate an estimated value.
ratio. No other species, including sulfate, thiosulfate, or polythionates, were detected by the ion chromatograph. However, the ion chromatograph would probably not detect a new species with a negative charge of four or greater. Analysis of the initial and final solutions of run 30-N for total carbon accounted for 89% of the initial citrate. Therefore, little of the citrate was lost by volatilization or precipitation. Complete degradation of the citrate/S02 solutions has little effect on their buffer capacity. In run 36-N, solution containing 0.2 M Na2HCitand 0.21 M NaHS03 was heated at 170 OC until less than 0.07 M SO2was left. The initial and final solutions were acidified and sparged with air to remove SO2 and COB. Titration of these solutions by NaOH showed that the final decomposed solution had about 20% less buffer capacity from pH 2.0 to 7.0. The shape of the titration curves is very similar (Figure 6). These results are consistent with rate-limiting dehydration of citric acid followed by fast sulfonation of aconitic acid to give sulfotricarballylic acid. The overall reaction consumes citrate and SO2 in a ratio of one to one. The product retains all of the initial carbon and has a charge of -4, which would make it indetectible by the ion chromatograph. Sulfotricarballylic acid should have about the same buffer characteristics as citric acid. The reaction kinetics would be consistent with slow dehydration of the
citrate followed by fast sulfonation with bisulfite. Corrosion of the reaction vessel was occasionally significant. Runs 2,4, 5, and 10-N tested the effect of dissolved Fe, Cr, Ni, and Mn on the reaction rate. Changes in thiosulfate and sulfate concentrations were found to be insignificant. The rate of sulfur dioxide disappearance was slightly above the rate predicted by eq 1. Results with Thiosulfate The presence of thiosulfate in sodium citrate buffers results in a second independent reaction for SO2 disappearance. Experiments with solutions containing thiosulfate were performed over the following range of conditions: [citrate] = 0.214 to 0.5 M; [SO,] = 0.023 to 0.4 M; [Na2S203]= 0.025 to 0.2 M; pH = 3.3 to 4.63; T = 405 to 453 K (132 to 180 "C). The rate of SO2 disappearance attributed to the thiosulfate reaction was estimated by subtracting the calculated rate for reaction of SO2with citrate (eq 1) from the total measured rate with thiosulfate. For selected runs (indicated in Table I) the rate of reaction of SOz with thiosulfate was correlated by
where In B = 69.4 f 14.0%,Hni0 = 45 x lo3 (cal/g-mol)
Ind. Eng. Chem. Process Des. Dev., Vol. 21, No. 4, 1982 573 Table 11. Stoichiometry of SO, Disappearance by Disproportionation a run no.
s,o, '-1 A [so,1Thio
1-N 2-N 28-N
1.03 26.7 1.12
4A [
4A [so,"1 A [so2 1Thio
1.92
1.88 1.76
a A I S o , ] ~ o= T A [ S O , ] T ( R , T @ ~ / R ~ ~ )L. O W thio sulfate concentration resulted in inaccurate analysis.
I
1
I
1
0.2
0
0.6
0.4
EQUIV. OF BUFFER CAP / LITER
Figure 6. Buffer capacity of initial and final solutions, run 36-N, 0.2 M citrate, 0.61 M Na+,170 O C , initial SOz,0.21 M final SO2,0.07
M.
t
I
I 0
d
CIT:SOr CIT',
SOz,Sp;
Ad.SO,,
s,o;
Metal I o n s
0
4HSOc
B
I 0"
is linear with time, even with 98% loss of SOz. However, thiosulfate concentration should increase as SO2decreases. Furthermore, reliable rate data over a large range of initial SO2 concentration were not available, so there is uncertainty in the dependence of rate on SOz concentration. The reaction of SO2with thiosulfate is approximately second order in thiosulfate and 1.6 order in H+. It has an activation energy of 45 X lo3 cal/g-mol, which is very near that of the reaction of SO2with citrate. The independence of the citrate and thiosulfate reaction was established by two experiments with solutions containing thiosulfate and acetate buffer, but no citrate (runs 9-A and 32-N). As shown in Table I, the calculated total rates for these experiments agree well with the measured rates. A third experiment, run 17-N, showed that no reaction occurs between citrate and thiosulfate in the absence of s02. The stoichiometry of the thiosulfate reaction was quantified by careful solution analysis in runs 1-N, 2-N, and 28-N (Table 11). Sulfur dioxide disappearance attributed to the thiosulfate reaction agrees with the stoichiometry of the net disproportionation of bisulfite to thiosulfate and sulfate
c
J I
IO-'
R,,
(colc.)
Idg (mlhr)
Fuure 7. Measured and calculated total rates of SO2disappearance.
i 15.990,c = 2.0 f 17.790,and d = 1.63 f 18.290. The standard deviation of the estimate of In RThio is *0.79. This corresponds to an error of a factor of 2.2. The total rate of SO2disappearance is correlated by the sum of rates calculated from eq 1 and 2
(3) Figure 7 compares the measured and calculated rates of SO2disappearance for all runs giving rate data. Six of the runs deviated more than 50% from the empirical correlation. As in solutions without thiosulfate, the rate of SO2 disappearance does not appear to depend on SO2 concentration. Run 9-N (Figure 3) shows that SO2disappearance
-
S2032-+ 2SOd2-+ 2H+ + H20
Both S2032-and SOZ-increased as SO2disappeared. The stoichiometry is further verified by the decrease in pH during runs 13-N and 15-N, from 4.58 to 4.49 and 4.41, respectively. The thiosulfate reaction kinetics are similar to those measured by Battaglia and Miller (1968) at 70 "C for the reaction of bisulfite and thiosulfate to give trithionate. They found a third-order dependence on H+and S2032and a first-order dependence on dissolved SOz in contrast to our orders of 2, 1.6, and 0.0, respectively. If Battaglia's rate equation is extrapolated to 150 "C with our measured activation energy, the calculated rate of the thiosulfate reaction is 0.106 M/h with 0.2 M S2032-and 0.2 M SO2at pH 4.5. Our calculated rate at the same conditions is 0.0184 M/h. If a lower activation energy of 39 kcal/g-mol is used with Battaglia's data, his correlation predicts the same rate as ours.
Design Implications Loss of bisulfite by disproportionation is potentially a more significant problem than reaction with citrate. Disproportionation produces thiosulfate and sulfate which accumulate in the system as acids, requiring NaOH or Na2C03makeup to maintain pH. Sulfate can be removed by refrigerated crystallization as Glauber's salt. Thiosulfate is more difficult to remove by crystallization. It may be removed by oxidation/reduction reactions or by acidification to give sulfur and sulfite. Thiosulfate must be removed to minimize its catalytic effect on the rate of disproportionation. In systems with H2S regeneration, thiosulfate is a necessary intermediate of system chemistry and reaches a steady-state concentration balancing reactions of H2S (Rochelle and King, 1979).
574
Ind. Eng. Chem. Process Des. Dev., Vol. 21, No. 4, 1982
acid. This reaction is first order in total citrate, zero-order in bisulfite, and has an activation energy of 43 kcal/g-mol. 2. The disproportionation of bisulfite to thiosulfate and sulfate is catalyzed by thiosulfate and is accelerated at lower pH. It has an activation energy of 45 kcal/g-mol and appears to be zero order in bisulfite. 3. The estimated rate of SOz disappearance in 1.5 M citrate is 0.02 M/h at 150 "C/pH 4.3/0.1 M NazS2O3,at 130 OC/pH 4.3/0.5 M NaZS2O3,or at 100 "C/pH 3.1/0.5 M Na2S203.The rate of SO, disappearance at 100 "C and pH 4.0 should pose no significant problem for simple absorption/stripping unless the thiosulfate concentration is greater than 0.5 M.
I40
T
("C) 120
Nomenclature
'00
I 5
3
PH
Figure 8. Conditions giving a calculated total rate of SO2 disappearance of 0.02 M/h in 1.5 M citrate.
Reaction of bisulfite with citrate is not necessarily undesirable. The resulting sulfotricarballyic acid is a buffer with properties similar to those of citrate, except perhaps for solubilities. However, the reaction must be taken into account in setting up analytical and operating procedures for systems using citric acid makeup. In a commercial SO2 absorption/stripping process the maximum tolerable rate of SO2 disappearance should correspond to about 1% of the SO2 absorbed. With a scrubber make-per-pass of 0.20 M and a liquid residence time at elevated temperature of 0.1 h this constraint is equivalent to a rate of about 0.02 M/h. Figure 8 presents sets of operating conditions which give a calculated rate of 0.02 M/h with 1.5 M total citrate. The reaction of citrate with SO2 is significant only at temperatures above 150 "C. Below 150 "C the disproportionation reaction is the dominant mechanism of SOz disappearance. It appears that operation at temperature as high as 150 "C would not be feasible. The maximum rate of SO2 disappearance (0.02 M/h) would be exceeded a t pH 4.3 with a thiosulfate concentration greater than 0.1 M. At 130 "C, as much as 0.5 M thiosulfate could be tolerated at pH 4.3. At 100 "C, the rate of SO2 disappearance would be unacceptable only if the pH was less than 3.1 and the thiosulfate concentrate was greater than 0.5 M. These constraints would be more stringent if less decomposition could be tolerated or if the liquid residence time at stripper temperature was greater than 0.1 h. Simple designs for absorption/stripping with sodium citrate buffers requires operation at 100 "C and pH 4 with times less than 1 h. Under these conditions disproportionation rates would be unacceptable only if the thiosulfate concentration exceeded 0.5 M. Conclusions 1. At 130-170 "C, citrate reacts directly with bisulfite to give a buffer product which is probably sulfotricarballyic
a, b, c, d = exponents in eq 1 and 2, dimensionless A, B = constants in eq 1 and 2, M/h [citrate] = total concentration of citric acid and its anions, M Hcit,HThio = activation energy in rate eq 1 and 2, cal/g-mol k = rate constant, units vary R = gas constant, 1.987 cal/(g-mol K) Rcit = rate of SO2 disappearance by reaction with citrate, M/h Rmo = rate of SO2 disappearance by reaction with thiosulfate (disproportionation), M/h Rbd = total rate of SO, disap earance, M/h [SO,] = total concentration of S species (primarily bisulfite), M T = temperature, "C or K [I = concentration, M
E
Literature Cited Applebey, M. P. J. SOC.Chem. Ind. Trans. 1937, 56, 139. Bailey, E. E. "Proceedings: Symposium on Flue Gas Desulfurizatlon-Atlanta, November 1974", US. EPA Report NO. 650/2-74-136b, 1974 p 745. Battagk. C. J.; Miller, W.J. Phot. Sci. Eng. 1968, 72, 46. Bengtsson, S.;Farrington, J. F. "Flakt-Boliin Absorptlon/Steam Stripping FGO Process", presented at 35th ACS Southwest Regional Meeting, Austin, TX, Dec 5-7, 1979. Boswell, M. C. U S . Patent 1972074, Sept 4, 1934. Foerster, F.; Hoenig, A. Z . Anorg. AI@. Chem. 1922, 125, 86. Gcehrlng, M.; Helbing, W.;Appel, I.Z . Anorg. A/@. Chem. 1947, 254, 185. Hagglund, E.: Ringbom, A. Z . Anorg. A/@. Chem. 1928, 750, 231. Johnstone, H. F.: Read, H. F.: Blankmeyer, H. C. Ind. Eng. Chem. 1938, 30, 101. KoRhoff, I.M.; Belcher, R. "Volumetric Analysis"; Interscience Publishers, Inc.: New York, 1957; Vol. 111, p 300. Kurtenacker, A.; Mutschin, A.: Stastny, F. Z . Anorg. A/@. Chem. 1935, 224, 399. Nawisky, P.; Sprenger, G. E. US. Patent 2315375, Mar 30, 1943. Nissen, W. I.; Elklns, D. A,; McKinney, W. A. "Proceedings: Symposium on Flue Gas Desulfurization-New Orleans"; US. EPA Report No. 600/2-76136b, 1976; p 843. Popov, A.; Micev, I . C. R . Acad. Bu@. Sci. 1982, 15, 37-40. Chem. Abstr. 1962, 5 7 , 14414. Rochelle, G. T. "Proceedings: Symposium on Flue Gas DesuifurizationHollywood, FL", U.S. EPA Report No. 600/7-78-058b, 1977; p 902. Rochelle, G. T.; King, C. J. AIChE Symp. Ser. 1979, No. 788, 75, 48-61, Slack, A. V. "Sulfur Oxide Removal from Stack Gases: Visits in Europe, June 12-July 8, 1972"; TVA. Division of Chemical Development, Muscle Shoals. AL, July 20, 1972. Usel'Tseva, Y. A.; Pobedinskaya, A. I.; Kobenma, N. M. Izu. Vyssh. Ucheb. Zaved., Khim. Tekhnol. 1970, 73, 507-11. Chem. Abstr. 1970, 73, 116514. R e c e i v e d for r e v i e w J u n e 17, 1980 A c c e p t e d M a y 3, 1982 T h i s p a p e r was p r e s e n t e d a t t h e ACS S o u t h w e s t R e g i o n a l M e e t i n g , A u s t i n , TX, D e c 5-7, 1979.
