This research has elucidated a partial understanding of the behavior of aerially absorbed bromine in pine tissues. Pine needles continued to accumulate bromine during their second year at about the same rate as during their first year. But only a small proportion of bromine was translocated from first-year needles into first-year twigs. Second-year twigs did not accumulate bromine during their second year, or if they did, it was translocated out a t an equal rate. Elemental bromine is a strong oxidant and one might expect more injury than was observed. The strong neg,ative exponential character of the accumulation curves would tend to limit possible injury to vegetation growing close to the sources. Although the degree of visible injury resulting from a particular bromine emission episode was not determined, the 800-1000 ppm concentrations in needles exhibiting no visible injury suggest that needle tissues can detoxify and accumulate relatively large concentrations of bromine with no visible ill effects. Bromine did not accumulate in the organic litter layer to a concentration greater than that found in living pine needles in the canopy above the plots. Loblolly and shortleaf pine needles are normally cast a t the end of 2 years of age, so that bromine content of organic litter should reflect the content of a mixture of primarily 2-year-old needles and some twigs and other organic debris. In general, this content ranged from less than to less than 1/3 of that of 2-year-old needles sam-
pled in the same plots. This finding and the low levels of bromine in soil suggest that, as the litter was decomposed, the resulting bromides were quickly leached away. The negative exponential power curves of bromine accumulations are compatible with Berge’s (2) prediction. He reasoned that, since bromine gas was five times as dense as air, it should be sharply confined to the immediate vicinity of the sources. This is strikingly shown by our regression equations. Most of the bromine accumulated was within 1 or 2 km from the source. For example, even though plots around GLE and VCC had accumulations of 550 ppm a t 0.48 km, compared to 150-250 ppm for ACI, GLM, and ETH, at 2 km accumulations by l-year-old needles at all five sources were reasonably tightly grouped at 40-100 ppm. L i t e r a t u r e Cited ( 1 ) Yaron, F. In “Bromine and Its Compounds”; Jolles, Z. E., Ed.; Academic Press: New York, 1966; pp 3-42. ( 2 ) Berge, H. “Phytotoxische lmmissionen (Gas-Rauch-und Staubschaden)”; Paul Parey: Berlin and Hamberg, 1963; Chapter 10, pp 53-54. (3) A m . SOC. Test. Mater., Book A S T M Tentatiue Stand. 1975,26, $24-737.
Receiced for review August 14, 1978. Resubmitted November 8, 1979. Accepted March 10, 1980. This work was supported i n part by Arkansas Chemicals Inc., Ethyl Corporation, Great Lakes Chemical Corporation, and Velsicol Chemical Corporation.
Removal of Thiosulfate/Sulfate from Spent Stretford Solution Tsoung-Yuan Yan” and Wilton F. Espenscheid Mobil Research and Development Corporation, P.O. Box 1025, Princeton, N.J. 08540
now containing tetravalent vanadium and reduced ADA is regenerated by oxidation with air and recycled to the gas absorber. The ADA also acts as a catalyst for the regeneration step. The net reaction is the indirect oxidation of hydrogen sulfide to form water and sulfur, which is recovered. The reactions can be represented as follows:
Thiosulfate accumulation presents a serious disposal problem in the Stretford process for removing hydrogen sulfide from contaminated gas streams. A method for treating spent Stretford solution to dispose of thiosulfate and recover the chemicals that it contains is presented. In this technique the solution is acidified with sulfuric acid to decompose the thiosulfate to sulfur and sulfur dioxide, and limed to remove added sulfate and restore the pH of the solution. The extent of reaction and recovery of chemicals has been investigated for each step of the method, and the feasibility of the process scheme as a whole has been investigated.
.
oxidation regeneration
Low concentrations of sulfur contaminants occur in gas streams such as coke oven gas, natural gas, and the tail gas from the Claus process ( I ) . The latter process, which is widely used in petroleum refineries to convert hydrogen sulfide byproduct to sulfur, generates large volumes of waste tail gas. Direct discharge of this gas, which contains residual HzS, to the atmosphere results in pollution levels unacceptable to an increasing number of communities. A highly effective process for removing low concentrations of hydrogen sulfide from contaminated gas streams is the Stretford process (2). In this method, the gas stream is contacted with aqueous sodium carbonatelbicarbonate solution that contains pentavalent vanadium and anthraquinonedisulfonic acids (ADA). The hydrogen sulfide is oxidized to sulfur with accompanying reduction of the vanadium and ADA. After the sulfur is separated, the spent aqueous solution 732
Environmental Science & Technology
+
16V5+ 8H2S 4V4+
-
+
16V4+ 16H+
+ Sa
ADA + 4H+. + 02 --+ 4V5+ + 2H20
In practice, the process performs its intended function of removing hydrogen sulfide from waste gas streams or other gas streams such as natural gas extremely well. Nonetheless, the process suffers one serious drawback. During each cycle of the process, a small percentage of sulfide is converted irreversibly to thiosulfate and, to a lesser extent, sulfate. This accumulation of thiosulfate reduces the solubility of vanadium and ADA in the solution, and decreases the rate of oxidative regeneration. To maintain these salts a t an acceptable level of 20-30 wt %, a continuous purge becomes necessary. This results in a loss of the valuable chemicals: ADA, sodium vanadate, and sodium carbonate. More important, however, high thiosulfate concentrations present a serious problem for disposing this purge stream because of their high chemical oxygen demand (COD). Thiosulfate is stable and has high water solubility. The methods for its destruction andlor removal from the process stream are costly. Proposed processes for disposal of this effluent have included evaporation, incineration, and even
0013-936X/80/0914-0732$01 .OO/O
@ 1980 American Chemical Society
-
Clear
Regenerated S t r e t f o r d S o l u t i o n
I
#
*
Gas Sulfur
Sulfur
T
I
-
Oxidatior
Holding Tank
1 H2S Con t a in ing Gas
-
1
SO2 t o Claus P l a n t /
Tank
:Strea'
Reaction
'Ni zeauttiroanl -
1 1
'~.
Steam
Regeneration
.
Lime S l u r r y
Purq
Caus t i-
' cizer
:
Figure 1. Schematic flow of the p r o p o s e d scheme
biodegradation (3). But, by and large, all are unattractive economically. Process S c h e m e
The objective of this work was to devise an inexpensive scheme for disposal of thiosulfate/sulfate with recovery of as much of the chemicals as possible. The proposed scheme ( 4 ) is shown in Figure 1, and the steps involved are outlined below. Step 1: Decomposition of Thiosulfate. The purge stream from the Stretford process is neutralized and further acidified with sulfuric acid to decompose the thiosulfate according to Equation 1.
+
8 S ~ 0 3 ~ -16H'
+
8S02 t
+ Ss 1 + 8H20
(1)
This overall reaction takes place in steps according to the following stoichiometric equations ( 5 ) : Protonation of thiosulfate ions:
S20:3Z-+ H-I s HSz03-
(pK2(H2S203)= 1.7) (2)
+ Hf 2 H2S203
(pKl(H2S203) = 0.6) (3)
HS203-
Internal redox reaction: HzS203
+
H2S03
+S1
(4)
Evolution ofSO2: H2S03 + H20
+ SO2 t
(5)
The redox reaction (Equation 4) also produces polythionate as the byproduct. The rate of decomposition of thiosulfuric acid has been studied by Schmidt (6). Step 2: Removal of Products. The acidified solution is stripped of SO2 using steam. The SO2 is recycled to the Claus plant, and the sulfur, which floats on the top, is recovered by skimming or filtration. Step 3: Recovery of Stretford Solution. The clear solution is reacted with lime to raise the p H and precipitate sulfate/sulfite ions as CaSO&aS03. In this step, the sulfate resulting from both byproduct formation in the Stretford process and sulfuric acid addition in the acidification step is removed. Incomplete stripping of SO2 will result in excess lime
consumption through formation of sulfite. The slurry is settled in a thickener, and the solution is recycled to the Stretford process. The calcium salts are disposed of after washing. Materials and Methods
Except for the used Stretford solution, the chemicals were reagent grade. The former was sampled from a tail-gas unit of a petroleum refinery. One of the samples contained 1.79 g/L mixed salts of ADA, 0.84 g/L of vanadium, and 15.8 w t % Na2S20.3. The pH of the used solution was 8.7. Thiosulfate contents of used Stretford solutions and stock thiosulfate preparations were determined by titrating the iodine liberated from a weighed KI03/KI solution to a starch end point in the presence of formaldehyde to complex
SO.?. The anthraquinonedisulfonic acid content of the Stretford solution was determined by filtering the sample to remove sulfur and other solids. The solution was diluted 2000-fold and made alkaline with 6 N NaOH. The ADA is oxidized by sodium dithionite to anthraquinone, and the quinone was measured photometrically a t 440-nm wavelength. ADA obeys Beer's law and has a molar absorptivity of 1.56 X lo4 L/mol.cm as determined from pure samples of the 2,6 and 2,7 salts. Experimental Results
a. Acid Decomposition of Thiosulfate. The decomposition products from dilute thiosulfate solutions are generally complex ( 7 ) and strongly dependent on acid strength. I t is believed that low acidities favor the desirable reaction (i.e., decomposition of thiosulfate to sulfur and sulfur dioxide by Equation l ) ,and medium strength acids favor the formation of a mixture of polythionic acids (8).Experiments with stock thiosulfate solutions and several acids of varying strengths showed that this was indeed true for pHs slightly below 7, or for weak acids such as acetic. A t pHs 5-6, some sulfur is formed and some sulfur dioxide evolves, but the decomposition of thiosulfate is far from complete. If the pH is lowered to the 4-5 range, a mixture of polythionic acids begins to form (7). However, strong acids capable of lowering the pH to 2-3, such as sulfuric or phosphoric, drive Reaction 1 t o completion. Since the pK1 of H P S Ois~ about 1.9, removal of SO2 might be V o l u m e 14, N u m b e r 6, June 1980
733
100
r
I
Table 1. Solubilities of Calcium Salts in Water
salt
CaS04 Ca(W2 CaS03 CaC03 (calcite) Ca vanadate Ca anthraquinone1,5-disulfonate a
solubllity product PKsp at 25 " C
solubility, w l %
4.6
0.22 (0 oC)a
5.43
0.19 (0 "c) 0.0043 (18 'C) 0.001 (25 "C)
6.6
6.2
0.26 (50 "C)" 0.08 ( l o o O C ) 0.027 (90 "C) 0.002 (100 "C)
unknown 3.0
As CaS04.2H20
Table II. Material Balances on Treatment of Spent Stretford Solution (Basis: 500 g) material
0
0.2
0.4
0.6
0.8
1 .o
H2S04/S203= Molar R a t i o
Figure 2. Conversion of thiosulfate vs. H2S04/S20z2-molar ratio l!
S2OS2-
Sod2-(ti2So4)
Ca2+(CaO) Cas04
input, g
output, g
% conversion
55.96 58.0
3.78
93.5
0.55
97.6
23.3
98.9
76.97
15.9 31.8
S
so2
% recovery
ADA V
0.79 0.33
0.72 (0.2)a
91 (60)a
Estimated.
0 Experimental
bo4=]
0 Experimental
2
1(
0,
c
.-Y0 L Y c
"
c u c
I
Solution pH
Figure 3. Solubility of calcium and sulfate ions vs. solution pH
more effective a t pHs of
