Environ. Sci. Technol. 1082. 16. 852-857
hydrolysis rate constant of 2,4-DOE were observed. These observations are mutually consistent with a rapid, reversible partitioning of 2,4-DOE between water and “dissolved”humic substances. The humic-bound 2,4-DOE is not adsorbed to glass and is protected from alkaline hydrolysis. On the basis of this model, with aquatic humus concentrations expressed as g of C/g of H20, the distribution coefficient of 2,4-DOE between water and aquatic humus (&) was calculated from both the glass adsorption data and the alkaline hydrolysis data, yielding log Kh values of 4.4 and 4.8, respectively. From the magnitude of the distribution coefficient of 2,4-DOE between water and “dissolved”humic substances, a significant interaction is generally expected for hydrophobic compounds in natural waters. Rates of alkaline hydrolysis, volatilization, etc., would be expected to be reduced in proportion to the fraction of a hydrophobic solute that is associated with humic substances. If any general acid-base catalysis of the hydrolysis reaction is attributable to humic substances, that contribution is completely masked by the partitioning phenomenon, which strongly decreases the rate of alkaline hydrolysis of 2,4-DOE. Acknowledgments
We thank J. H. Reuter, School of Geophysical Sciences, Georgia Institute of Technology, Atlanta, GA, for the humic and fulvic acids from the Altamaha River sediments in south Georgia. Literature Cited (1) Khan, S. U. Pestic. Sci. 1978, 9, 39. (2) Struif, B.; Weil, L.; Quentin, K.-E. Vom Wasser 1976,45, 52.
Wolfe, N. L. In “Dynamics, Exposure and Hazard Assessment of Toxic Chemicals”; Haque, R., Ed.; Ann Arbor Science: Ann Arbor, MI, 1980; p 163. Sharom, M. S.; Miles, J. R.; Harris, C. R.; McEwen, F. L. Water Res. 1980, 14, 1089. Bailey, G. W.; Thruston, A. D., Jr.; Pope, J. D., Jr.; Cochrane, D. R. Weed Sci. 1970, 18, 413. Drevenkar, V.; Fink, K.; Stipvevic, N.; Stengl, B. Arh. Hig. Rada Toksikol. 1975, 26, 257. Perdue, E. M.; Reuter, J. H.; Ghosal, M., Geochim. Cosmochim. Acta 1980,44, 1841. “Standard Methods for the Examination of Water and Wastewater”, 12 ed.; American Public Health Association: New York, 1965, pp 406, 592. Karickhoff, S. W.; Brown, D. S. ”Determination of Octanol/Water Distribution Coefficients,Water Solubilities, and Sediment/ Water Partition Coefficients for Hydrophobic Organic Pollutants”; U.S. Environmental Protection Agency, Athens, GA, EPA-600/4-79-032, 1979. Hassett, J. J.; Means, J. C.; Banwart, W. L.; Wood, S. G. “Sorption Properties of Sediments and Energy-Related Pollutants”; U.S. Environmental Protection Agency, Athens, GA, EPA-600/3-80-041, 1980. Banerjee, S.; Yalkowsky, S. H.; Valvani, S. C. Environ. Sci. Technol. 1980,14, 1227. Means, J. C.; Wood, S. G.; Hassett, J. J.; Banwart, W. L. Enuiron. Sci. Technol. 1980, 14, 1524. Zepp, R. G.; Wolfe, N. L.; Gordon, J. A.; Baughman, G. L. Environ. Sci. Technol. 1975, 9, 1144. Perdue, E. M., unpublished results, Portland State University, 1981. O’Brien, R. D. “Environmental Dynamics of Pesticide”; Haque, R., Freed, V., Eds; Plenum Press: New York, 1975. Kenega, E. E.; Goring, C. A. I. Proceedings of the Third Aquatic Toxicology Symposium, American Chemical Society, Miami Beach, FL, 1978. Received for review September 28, 1981. Revised manuscript received July 13, 1982. Accepted August 10, 1982.
Kinetics of Decomposition of Tetrathionate, Trithionate, and Thiosulfate in Alkaline Media Ernest Rolla Metallurgical Chemistry Sectlon, Mineral Sclences Laboratories, CANMET, Department of Energy, Mines and Resources Canada, Ottawa, Ontario K I A OGI,Canada
Chunl L. Chakrabartl Department of Chemistry, Carleton University, Ottawa, Ontario K l S 586, Canada
rn At pH >10 tetrathionate decomposes to thiosulfate and
trithionate; the rate of reaction is first order with respect to both tetrathionate and hydroxide. In the temperature range 15-45 “C, the activation energy, E,, is 115.5 kJ mol-l. At pH 5.5-12 trithionate reacts with water to give thiosulfate and sulfate as reaction produds, the rate of reaction is not influenced by either hydroxide or dissolved oxygen. In the temperature range 70-85 “C, the activation energy is 91.7 kJ mol-l. Thiosulfate is oxidized to sulfate by dissolved oxygen in an alkaline solution; the rate of oxidation is first order with respect to thiosulfate, order 1.1 with respect to hydroxide ion, and order 1.66 with respect to oxygen pressure. In the temperature range 100-138 “C, the activation energy is 85.8 kJ mol-l. The oxidation reaction is characterized by an induction period whose duration increases with increasing pH and decreases with rising temperature and increasing pressure.
Introduction
During milling and flotation of sulfide ores, some sulfur 852
Environ. Sci. Technol., Vol. 16, No. 12, 1982
goes into solution in the form of partially oxidized sulfur compounds such as thiosulfate and polythionates, which are collectively known as thio salts. Sulfate ion can also be produced in significant quantities. Thio salts are eventually oxidized to sulfuric acid, but the oxidation is so slow in the flotation plant tailings ponds and in biological holding ponds that decomposition can be incomplete during the residence time. Thus the total potential sulfuric acid may not be neutralized prior to discharge of the effluent to receiving streams. On the other hand, oxidation is sufficiently rapid in the receiving body of water so that considerable acid is produced and the pH of the receiving rivers and lakes may decrease to about 3-4. As a consequence of these low pH values, environmental effects invariably include aquatic damage and fish kills (1). The above rapid oxidation in the receiving water bodies is attributed to the action of various thiobacilli species of bacteria (usually T. thiooxidans and T . ferrooxidans), which biologically convert thio salts to sulfuric acid. The formation of thio salts during ore processing and
0013-936X/82/0916-0852$01.25/0
0 1982 American Chemical Society
the need for their removal from mill tailings water to avoid subsequent environmental pollution were first recognized by Schmidt and Conn (2,3).Work at Canada Centre for Mineral and Energy Technology (CANMET) showed that the thio salt problem was caused chiefly by the presence of three thio salt anions: thiosulfate, trithionate, and tetrathionate. Tetrathionate concentrations gradually increased in the acidic (pH -3) biostabilization pond during the 20-day detention period, but no tetrathionate was found after lime treatment in the alkaline settling pond (4). Hence, tetrathionate was decomposed in alkaline media, leaving a solution containing thiosulfate and trithionate. Preliminary laboratory observations showed that the rate of trithionate decomposition was greatly increased by increases in temperature. It was also noted that during the decomposition of trithionate, a small amount of thiosulfate was oxidized to sulfate by dissolved oxygen. The above observations suggested that oxidation in alkaline media might be used to solve a serious environmental problem. Accordingly, the objectives of the present study were to study the reaction kinetics for the decomposition of tetrathionate, trithionate, and thiosulfate in alkaline solutions in the hope that information obtained from this study can be used to develop an alkaliie oxidation process. Experimental Section Apparatus. The experiments were done in a 500-mL glass reaction vessels enclosed in a water jacket. The solution was stirred magnetically and was kept at uniform temperature, f0.5 "C. A Metrohm Herisau Titrator E526 with Dosimat E535 was used to maintain a constant pH by incremental additions of 1N sodium hydroxide solution. Thiosulfate and polythionates were determined spectrophotometrically (5). An autoclave manufactured by Parr Instrument Co. was used to determine the effect of high temperatures and high oxygen pressures on the rate of thiosulfate oxidation. Reagents. Potassium tetrathionate was purchased from Eastern Chemical, Hauppauge, NY 11787. A weighed quantity of the potassium tetrathionate gave a clear solution in distilled water. The solution was analyzed spectrophotometrically (5). The potassium tetrathionate was found to be 100 f 1% pure. Potassium trithionate was prepared by the method of Stamm et al. (6). Analysis by heating it at 700 "C to convert it quantitatively to K2S04indicated it to be 99.3% pure. A weighed quantity of the potassium trithionate gave a clear solution in distilled water. The solution was analyzed spectrophotometridy and was found to be pure. Sodium hydroxide solution (1 N) was prepared from reagent grade pellets. The solution concentration was determined by titration with a standard hydrochloric acid solution. Sodium thiosulfate, Na2S2O3.5H20,was of analytical reagent purity. Cylinders of oxygen (99.99% pure), compressed air (medical grade), and a mixture of 47.9% oxygen and 52.1% nitrogen were used as a source of oxygen. Procedure. Distilled water (500 mL) in the reaction vessel was brought to constant temperature, and the pH was adjusted to the required value by the automatic titrator. Theoretical weights of reagents were added; aliquots (0.5-2.0 mL) of solution were then withdrawn periodically for determination of thiosulfate, tetrathionate, and trithionate. In some experiments, samples of solution (- 10 mL) were withdrawn periodically by pipette and were transferred to glass vials that were kept in cold water (5-15 "C) until the end of the test. The vials were then transferred to a water bath kept at room temperature before aliquots were taken for analysis. The volume of
26
I-
24
I-
22 20 1 18 \
-I l6
14
8
E 12
/s,o,-
10
YIELD
8 6 4
L
2
0.5
1.0
1.5
2.0
TIME / H
Figure 1. Changes in the concentrations of reactants and products as a function of time in the decomposition of tetrathionate at pH 12.0 and 25 "C.
sodium hydroxide solution required to keep the pH constant was recorded at the time of sampling. Reaction temperatures and the initial concentrations of the reactants were chosen so as to minimize analytical errors. Results and Discussion A. Decomposition of Tetrathionate. Fava and Bresadola (7)studied the catalytic effect of thiosulfate on the decomposition of tetrathionate in a neutral solution. In dilute alkaline solution the stoichiometry of tetrathionate decomposition is 4S40e2-
+ 60H-
-
5S20s2-+ 2s30s2-+ 3H2O
(1)
The reaction has been reported to be catalyzed by thiosulfate (8,9). Attack by hydroxide ions yields thiosulfate and trithionate in dilute alkaline solutions; however, thiosulfate and sulfite are formed in hot concentrated alkaline solutions. Stoichiometry. Figure 1 shows the changes in the concentrations of reactants and products with time during the decomposition of tetrathionate at pH 12.0 and 25 "C. After 2.0 h of reaction, the S40s2-reacted/OH- con~umed/S~O yield/S3Os2~~yield ratio is 1.0/1.5/1.25/0.5. This ratio corresponds to the accepted stoichiometry for the decomposition reaction (eq 1). Effect of Concentration. Semilogarithmic plots of concentration against time gave straight lines (Figure 2). Since the reaction is carried out a t constant pH, straight-line plots indicate pseudo-first-orderkinetics. The slope of the lines gives the value of the pseudo-first-order rate constant, k', in the expression -d[S40G2-]/dt = kqS40,j2"]
(2)
Effect of pH. Figure 3 shows a logarithmic plot of k' vs. [OH-]. The slope of the straight line (=1.0) gives the order of reaction with respect to hydroxide ion concentration. The rate expression can now be written for the decomposition of tetrathionate at constant temperature: -d[s40s2-] /dt = k[S40s2-][OH-]
(3)
The constant, k, for the reaction has the value
k = k'/[OH-] (4) The calculated values for k at different pH values are Environ. Sci. Technol., Vol. 18, No. 12, 1982 853
0
12t/
1 0.5
2
3
4
5
6
INITIAL (S, O:-)
I 0.5
I 1.0
1.5
7
8
9
1
0
rn M O L / L
Figure 4. Effect of the Initial thiosulfate concentration on the tetrathbnate decomposition in alkaline solutions ([S40t-] = 5 mmoi/L, pH 11.0, 35 "C).
2.0
TIME / H
Flgure 2. Semilogarithmic plots of the decrease of tetrathlonate as a function of time at pH 12.0 and 25 O C .
IO-^
According to the transition-state theory, the enthalpy of activation AH*is related to the experimental energy of activation,E,, in the Arrhenius equation. In liquid systems
AH* = E,- RT
(7)
and at 298 K, AH* = 113 kJ mol-l. The entropy of activation of the complex is given by the relation AS* = R(ln A - In (RT/Nn) - 1)
lo-'
10-2
[OH-] M O L / L
Figure 3. Logarithmic plot of k' (35 "C)vs. molar concentration of
and at 298 K, AS* = 33.5 J mol-l K-l. Effect of the Initial Thiosulfate Concentration. A plot of k'vs. [S202-] shows a first-order dependence of k' on the thiosulfate concentration (Figure 4). The effect of the initial thiosulfate concentration is probably diminished because of the rapid formation of thiosulfate from the decomposition of tetrathionate. Extrapolation of the data in Figure 4 shows that in the absence of thiosulfate the pseudo-first-order rate constant for the reaction would be s-l. The rate constant It'increases linearly with the amount of thiosulfate added and can be expressed as follows:
k' = ko + k,[S20,2-]
OH-. Table I. Effect of pH on the Kinetic Parameters for the Decomposition of Tetrathionate ([ S,0,2-] = 5 mmol/L, 35 "C)in Alkaline Solution
PH
t,,2,h
10.0 10.5 11.0 11.5
4.8 1.o 0.4
9.4
105k',s-I 2.0 4.8 18 43
k, L rno1-ls-l 0.20 0.1 5
0.18 0.14
shown in Table I. The average value is 0.17 f 0.03 L mol-l S-1.
Effect of Temperature. Arrhenius plots (log k'vs. 1/T (K)) were made for decomposition of tetrathionate in the temperature range 15-45 "C at pH 11.0. The activation energy was calculated by using regression line values for k< and It; from the relationship -E, = ((ln k{ - In k2/)/(1/T2 - l/Tl)) (5) At pH 11.0, E, has the value 115.5 kJ mol-l. The preexponential factor, A , of the Arrhenius equation is obtained from the relationship In k' = In A - E,/RTl (6) and it has the value 9 X 1014s-l. 854
Environ. Sci. Technol., Vol. 18, No. 12, 1982
(8)
(9)
s-l. The slope of the line in Figure 4 gives where ko = the value of kt, i.e.
k, = 126 X lom4mol-' s-l The rate constant, k (eq 41, at pH 11.0 is
k=
10-4
+ 126 x 10-4[s o 10-3.
2-
= 10-l
+ 12.6[S20,2-]
At pH 10-12 and 35 OC,the constant k has the value 0.17 f 0.03 L mo1-ls-l. The rate constant can be used to predict the rate of reaction under different conditions of concentration and temperature. At constant temperature and at a given pH, the pseudo-first-order rate constant k ' is related to the half-life by the relationship til2 = 0.693/k' (10) It is interesting to compare the half-life at pH 11.0 for three different temperatures, 5, 25, and 85 "C, approximating winter, summer, and process temperatures. The corresponding tlI2values are 137.5 h, 4.85 h, and 7.1 s, respectively. The results show that the reaction is extremely sensitive to temperature. In cold solutions, many days are required for the reaction to go to completion; in
Table 11. Effect of the Initial Concentration of Thiosulfate on the Kinetic Parameters for the Decomposition of Trithionate added 105kf, temp, 0 10 20 0 1
2
3
4
2.5 5.0 10.0
5
"C
tu*, h
70 70 70 85 85 85 85
3.07 3.25 2.76 0.85 0.826 0.742 0.686
S-1
6.27 5.92 6.97 22.6 22.3 25.9 28.1
TIME / H
do-fiist-orderdependence with respect to trithionate. The slope of the line gave the pseudo-first-orderrate equation
Flgure 1. Curves showing the decrease of trlthionate concentration and the yield of thiosulfate as a function of time in the decomposition of trithionate at pH 10.0 and 80 OC.
(13) -d[S30e2-]/dt = kTS302-1 Since the decomposition of trithionate is a hydrolysis reaction, the rate equation at constant temperature is (14) -d[S&"] /dt = k[S30e2-][HzO]@
20
10
\
1 0
Naito et al. (10) investigated the effect of water concentration by adding ethanol to the reaction solution. They found a first-order dependence on water; therefore, /3 = 1 in expression 14. In dilute aqueous solution, [H20] is nearly constant so that k = k'/[H20] = k'/55.5 M-l s-l (15)
5
z
E
ci z 0 0
0.5
1 t
\
0'
~
l
1
l 2
l
~
3
l
4
~ 5
1
c
6
1
~
7
TIME / H
Figure 6. Semilogarithmic plots of the decrease of trlthlonate concentration as a function of time at pH 11.0 and 70 "C.
hot solutions the reaction is complete in seconds. B. Decomposition of Trithionate. Naito et al. (10) studied the kinetics of the cleavage of trithionate ion in water over the temperature range 40-80 "C in ammoniacal media. They reported the rate law -d[S@e2-] /dt = k[HzO] [s@62-]
(11)
Stoichiometry. Figure 5 shows the changes in trithionate concentration and the yield of thiosulfate as a function of time during the hydrolysis of trithionate (5 mmol/L) a t pH 10.0 and temperature 80 "C. The glass electrode used for setting pH was standardized with a buffer solution held at 80 "C. After 4.$h, analysis showed the reaction products to be thiosulfate (4.8 mmol/L). Sodium hydroxide solution was added (8 mmol/L) to keep the pH constant. A sample taken 1 h later assayed 6.0 mmol/L sulfate. Air was not excluded from the reaction vessel. The results agree closely with the stoichiometry of the hydrolysis reaction (eq 12).
S30e2- + H,O
-
+
S2032- + Sod2- 2H+
(12)
When the reaction was left to proceed overnight, the final thiosulfate concentrations were less than predicted by eq 12, and the sulfate concentrations were higher. Most probably, oxidation of some of the thiosulfate occurred during the course of reaction. Effect of Concentration. Semilogarithmic plots of trithionate concentration vs. time are shown in Figure 6. Straight-line plots were obtained that indicated a pseu-
1
Effect of Temperature. An Arrhenius plot for the decomposition of trithionate in the temperature range 70-85 "C gave the following activation parameters: E, = 91.7 kJ mol-l; A = 5 X lo9 M-l s-l; AHs = 89.2 kJ mol-'; A' Sl = -67 J mol-' K-l. Effect of t h e Initial Thiosulfate Concentration. The kinetic parameters are summarized in Table 11. The results in Table I1 show that at 70 "C and pH 11.0, with the added thiosulfate, the values obtained for k 'lie within the range of experimental uncertainty of the values obtained for k'without the added thiosulfate. To minimize experimental error, the effect was studied at near-neutral pH, 5.5-8. The results show that at pH 5.5-8 and 85 "C the presence of initial thiosulfate accelerates trithionate decomposition. Increase in the rate of trithionate decomposition with increase in the initial thiosulfate concentration in neutral solutions has been observed by Naito et al. (10). These authors have concluded that a thiosulfate-trithionate reaction competes with the solvolysis of trithionate. It is interesting to determine to what extent the rate of reaction is increased as the temperature is increased from 85 to 100 "C. At 86 "C, k' = 22.6 X 10" s-' and t112= 0.56 s-l and tl = 0.37 h- From eq 5, at 95 "C, k'= 5.22 X h; at 100 "C, k' = 7.79 X s-l and tl/2 = 0.25 Thus at 100 "C, trithionate is almost completely hydrolyzed in -1h. C. Oxidation of Thiosulfate. Air oxidation of thiosulfate under normal pressures and temperatures is a very slow process. Solutions of thiosulfate and polythionates, at pH 7, were aerated for 4 months under sterile laboratory conditions at Noranda Research Centre with less than 10% change observed in the thio salt concentration (11). Oxidation of thiosulfate under elevated pressures of air or oxygen has been reported. Thiosulfate was oxidized to sulfate by Gluud at 100 "C under an air pressure of 980 kPa (12). Forward et al. (13) found that 88% of the thio salt was oxidized to sulfate at pH 7.3, 120 OC, and 3040 kPa of oxygen pressure. The authors concluded that kinetic considerations are more significant than thermody-
h.
Environ. Sci. Technoi., Voi. 16, No. 12, 1982 855
lo-,
3 4 5
2
7
Id'
INITIAL [OH-'] MOL/L
Flgure 8. Logarithmic plot of k' vs. initial hydroxide concentration at 125 "C and 345 kPa of oxygen pressure.
2l
/
Id'
TIME / H
Figure 7. Effect of the initial thiosulfate concentration on the rate of thiosulfate oxidation at 125 OC and 345 kPa of oxygen pressure.
namic factors in determining the ultimate course of reactions that led to, or through, intermediate sulfur species. Naito et al. (14) found that the oxidation of ammonium thiosulfate with oxygen in aqueous ammoniacal solutions was an almost zero-order reaction. Ammonium sulfamate and ammonium sulfate were the chief products. Initially in the present program, tests to oxidize thiosulfate by molecular oxygen in alkaline solutions were carried out in the temperature range 75-87 "C (15). Only a summary of this work will be given since, in this temperature range, there was a long induction period and it was difficult to interpret the complex oxidation curves. The results showed that the final product of the oxidation of thiosulfate by molecular oxygen in alkaline solutions was sulfate; there was no reaction in the absence of oxygen. Tetrathionate was not detected in any test. In the temperature range 75-85 "C, the induction time decreased and the rate of oxidation increased with increasing temperature. There was a linear relationship between log tl/2 and the reciprocal of absolute temperature. The stoichiometry of the reaction is S2032- 202 + H 2 0 --* 2S042- + 2H+
+
Autoclave Tests. Since the induction time was decreased by increasing the temperature, tests were carried out in an autoclave at temperatures >lo0 "Cand at greater than atmospheric pressures of oxygen. It was found that with the use of a baffle in the reaction vessel to improve the solution mixing, the induction time could be decreased to about 3 min. There was no oxidation of thiosulfate during heat-up in the autoclave experiments. Reaction occurred only when oxygen was introduced into the solution. Samples of solution ( 25 mL) were taken periodically from the autoclave and cooled to the room temperature for analysis. Effect of the Thio Salt Concentration. The effect of the initial thiosulfate concentrations in the range 6-18 mmol/L on the rate of thiosulfate oxidation at 125 "C and 345 kPa of oxygen pressure is shown in Figure 7. Because 2 mL of 1M NaOH was added initially for each mmol of S2032-to neutralize the acid that would be generated during the reaction, the initial hydroxide concentrations are higher for the more concentrated thiosulfate solutions, and the curves in Figure 7 exhibit different induction
-
858
Environ. Sci. Technol., Vol. 16, No. 12, 1982
I I
I00
1 1 , 1 1 1
300 500 1 x 0 0, GAS k Pa
1
mo
Figure 9. Logarithmic plot of k'vs. oxygen pressure at 125 "C and for 0.025 M initial hydroxide concentration.
times. To determine the rate of reaction, we subtracted the induction times and plotted semilogarithmically the points on the new time scale. The results indicated pseudo-first-order kinetics with respect to thiosulfate, tl/2 = 486 s, and k / = 1.48 X s-l. Effect of the Hydroxide Concentration. The effect of the hydroxide concentrations in the range 0.015-0.145 M NaOH on the rate of oxidation of 10 mM S203'-was studied at 125 "C and 345 kPa of oxygen pressure. The experimental data were plotted semilogarithmically, and straight lines drawn through the points gave induction times which were substracted to determine the half-life values. The logarithmic plot of k' vs. initial hydroxide concentration is shown in Figure 8. The computed slope of the regression line is 1.10. Effect of the Oxygen Pressure. The effect of the oxygen pressures in the range 345-1034 kPa on the oxidation rate of thiosulfate (10 mM) at 125 "C and for an initial hydroxide concentration of 0.025 M was studied. The data were plotted semilogarithmically,and the straight lines drawn through the points showed the same induction time, 0.06 h. After substraction of the induction time, the reaction half-life and the rate of reaction were determined. The logarithmic plot of k vs. oxygen pressure is shown in Figure 9. A linear relationship is observed with a slope of 1.66. Effect of Temperature. The effect of temperature in the range 100-138 "C on the rate of oxidation of thiosulfate (10 mM) was studied. The oxygen pressure was maintained at 690 kPa, which is -3 atm above the water vapor pressure at 138 OC (-345 kPa). The experimental data were plotted semilogarithmically, and straight lines drawn through the points indicated approximately zero induction time. The tlI2,k f ,and the rate of reaction were calculated; the results are given in Table 111. The computed slope of the regression line in the Arrhenius plot gave the following activation parameters: E, = 85.8 kJ mol-l; AH*=
Envlron. Scl. Technol. 1882, 16, 857-861
Schmidt, J. W.; Conn, K. Can. Min. J . 1969, 90, 54. Schmidt, J. W.; Conn, K. Can. Min. J . 1971, 92,49. Rolia, E., Division Report MRP/MSL 78-46 (TR); CANMET, Energy, Mines and Resources Canada; 1978. Rolia, E.; Barbeau, F. Talanta 1980, 27, 596. Stamm, H.; Goehring, M.; Feldman, U. 2.Anorg. Allg. Chem. 1942, 250, 266; Chem Abstr. 1943, 37, 3:5329. Fava, A.; Bresadola, S. J. Am. Chem. SOC.1966, 77,5792. Nickless, G., Ed. "Inorganic Sulfur Chemistry";Elsevier: New York, 1968; p 528. Byerley, J. J.; Fouda, S. A.; Rempel, G. L. J. Chem. Soc., Dalton Trans. 1973, 889. Naito, K.; Hayata, H.; Mochizuki, M. J . Inorg. Nucl. Chem. 1975,37, 1453. Cotton, M. L.; Spira, P.; Wheeland, K. G., Noranda Research Centre, Pointe Claire, Quebec, Canada, private communication. Gluud, W. Ber. Dtsch. Chem. Ges. B 1921,54B, 2425. Forward, F. A.; Peters, E.; Majima, H., paper presented at the AIME Annual Meeting, Feb 1964, New York. Naito, K.; Yoshida, M.; Shieh, M.; Okabe, T. Bull. Chem. Soc. Jpn. 1979,43, 1365. Rolia, E. M.Sc. Thesis, Carleton University, Ottawa, Canada, 1981.
Table 111. Effect of Temperature on the Reaction Half-Life for the Oxidation of Thiosulfate (10 mM) at 690 kPa of Oxygen Pressure and for an Initial Hydroxide Concentration of 0.025 M IO6 x rate of oxid of
temp, "C 100 124 130 138
t,,,, s 2070 432 241 162
10w, S-'
0.33 1.60
2.87 4.28
s,o,,-,
M s-' 3.33 12.5 17.8 21.6
83.3 kJ mol-'; AS*= -89.6 J mor1 K-l. The rate equation for the oxidation of thiosulfate by molecular oxygen in an autoclave can be written as follows: -d[S,032-]/dt = Iz[S~032-][OH-]1.'(Po,)'.66 (16) From the observed rates, after subtraction of induction periods, the average value of the rate constant k, based on six calculations, was 1.66 X 10-6M-1~1(PoJ-1~es 8'. The range was (1.3-2.0) X lo4, standard deviation was 0.32 X lo*, and the relative standard deviation was 199%. Literature Cited (1) Paine, P. J. M.Eng. Thesis, University of Ottawa, Canada; 1978.
Received for review September 16, 1981. Revised manuscript received June 21, 1982. Accepted July 19, 1982.
A Photoreactor for Investigations of the Degradation of Particle-Bound Polycyclic Aromatic Hydrocarbons under Simulated Atmospheric Conditionst Joan M. Dalsey,' Catherine 0. Lewandowskl, and Mllena i o r i Institute of Environmental Medicine, New York University Medical Center, New York, New York 10016
A fluidized-bed photochemical reactor has been developed for laboratory studies of reactions of particulateadsorbed polycyclic aromatic hydrocarbons (PAH) under simulated atmospheric conditions. The reactor consists of a glass column in which particles are suspended by the flow of air through a fritted disc at the base of the column. A quartz mercury vapor lamp, aligned with the column of suspended particles, is used for irradiation. Although the spectral wavelength distribution of the glass-filtered light is not identical with those of sunlight, it is a reasonably good approximation in the actinic region of the spectrum. The reactor is inexpensive and simple to construct and operate. It is three-dimensional and permits unlimited reaction time. Reaction conditions can be varied. Rate constants determined in the reactor have been shown to be reproducible to within f20% a t the 95% confidence level. Introduction Polycyclic aromatic hydrocarbons (PAH) in industrial and ambient atmospheres have long been of concern as a human health hazard (1). Many of the individual PAH compounds are potent carcinogens while others act as cocarcinogens or tumor promotors in mammalian bioassays. Evidence of tumorigenic potential in humans has come from epidemiological studies of workers exposed to high levels of these compounds (1-4). Presented at the Symposium on Atmospheric Chemistry, 182nd National Meeting of the America1 Chemical Society, August 23-28, 1981, New York; Abstr. PHYS 184. 0013-936X/82/0916-0857$01.25/0
The polycyclic aromatic hydrocarbons are environmentally ubiquitous compounds, which are produced by combustion of fossil fuels and certain industrial processes such as coke production and petroleum refining. Emissions from fossil fuel combustion can vary over several orders of magnitude depending upon the particular fuel and combustion conditions (1,5). Emissions of PAH per BTU from coal or wood burning for residential space heating are several orders of magnitude greater than for gas or oil burning (1,5). Thus, the trend toward increased use of coal and wood in the United States can have a substantial impact on the concentrations on airborne PAH. Although airborne PAHs have been studied for almost 30 years, our knowledge of the chemical lifetimes of these compounds in the atmosphere is very limited and based entirely on studies in model systems (6-13). Much of the early work on the degradation of PAH suggested lifetimes of the order of hours (1).More recent work by Korfmacher et al. (14) and by Butler and Crossley (15) indicates that the chemical lifetimes of particle-bound PAH are of the order of days. The degradation rates estimated from the available model studies vary widely (16) because many of the parameters of the model systems have not been well defined, e.g., light intensity, particle surface areas and surface composition, and concentration of PAH per gram of substrate. Many of the studies were done in two-dimensional model systems, i.e., PAH on a thin-layer chromatography plate. The flow-through chamber of Tebbens and coworkers (11-13) more closely approximates an environmental system. However, the flow-through system is relatively complex and has some of the same problems as
0 1982 American Chemical Soclety
Environ. Scl. Technol., Vol. 16, No. 12, 1982 857
