Kinetics of the Chlorination of Biphenyl under Conditions of Waste Treatment Processes Eric H. Snider" and F. C. Alley Department of Chemical Engineering, Clemson University, Clemson, S.C. 29631
The chemical kinetics of the chlorination of low concentrations of biphenyl dissolved in water are presented. The pH range investigated was 6.2 to 9.0, in which range the chlorinating agent is molecular hypochlorous acid and its dehydration product, chlorine monoxide. The reactions in the neutral and alkaline ranges produced two monochlorobiphenyls, 0- and p-chlorobiphenyl, with a 1/2 o:p ratio of 0.8. Second-order dependence upon undissociated hypochlorous acid concentration and first-order dependence upon biphenyl concentration were observed. Reaction rate constants and activation energies were determined. Moderate ionic strength variations and potential metallic catalysts had no measurable influence on reaction rate. It was found that acid catalysis occurs below a pH of 6.2, where highly reactive molecular chlorine is present. Biphenyl (C12H10)has been used extensively for many years as a heat transfer medium and, more recently, as a dye carrier in the synthetic fiber industry ( I ) . Polychlorinated biphenyl (PCB) mixtures, produced by reacting molten biphenyl with gaseous chlorine, were used industrially until recently, when research indicated that the PCBs possess mutagenic, teratogenic, and potential carcinogenic characteristics. As a result of these studies, the industrial use of PCBs has been terminated ( 2 ) . The production of chlorinated organics by reaction of organic compounds dissolved in water with aqueous chlorine has been extensively studied (3-10). Most studies of such reactions have involved the use of fairly reactive organic compounds (ones which undergo electrophilic substitution easily) or the use of acidic reaction conditions in which molecular chlorine is present, or both (11-16). These studies have shown that easily measurable concentrations of chlorinated organics can be produced from a number of organic reactants under certain conditions. Chlorination of refractory compounds has not been widely studied, since many of these compounds are only sparingly soluble in water and the concentrations of chlorinated products are very low. However, improved analytical techniques and a growing awareness of the effects of some chlorinated species on health and the environment have shown the need for more work in this important area. The chlorination of biphenyl in water under acidic conditions has been studied by Beaven et al. ( 1 7 ) . Until recently, it was generally accepted that biphenyl and other refractory organics were unreactive in the presence of chlorine at neutral and alkaline pH values. Evidence that chlorination occurs under these conditions came in a report of Gaffney (18) on water samples taken from a treatment facility whose wastewater contained biphenyl. The treatment plant had been using 150 to 190 kg per day of gaseous chlorine for influent odor control and disinfection of the effluent. Analysis of material collected from the high-rate trickling filters showed a PCB concentration of 18 mg L-l. In further work, Gaffney spiked deionized water and wastewater samples with biphenyl and chlorinated at pH 8.0 (19).Analysis of the chlorinated samples showed a number of electron-capture gas chromatographic Present address, Department of Chemical Engineering, The University of Tulsa, 600 South College, Tulsa, Okla. 74104. 1244
Environmental Science & Technology
peaks corresponding to mono- and dichlorobiphenyls. Johnsen (20)and Carlson et al. (8)performed similar studies indicating that biphenyl could be chlorinated under conditions of water treatment processes. This study was performed to provide quantitative kinetic data on these reactions. Biphenyl was chosen as the organic reactant because of its wide use and the potential danger resulting from its chlorination products. Biphenyl is highly biorefractory and passes undegraded to the chlorination step in biological treatment plants. I t was the purpose of this project to study these chlorination reactions in distilled water with levels of biphenyl and chlorine not far removed from those encountered in treatment practice and at neutral and alkaline pH ranges. Experimental
Sample Preparation and Chlorination. Water was prepared by distillation over alkaline potassium permanganate and was stored in glass vessels. About 20 g of powdered biphenyl was added to a 5-gal vessel of distilled water, and the mixture was stirred for 24 h or longer. The mixture was filtered twice, first through glass wool and second through Whatman No. 1 filter paper. The filtrate was diluted as desired and buffered in all cases except the low ionic strength studies. Buffers used were potassium dihydrogen phosphate-sodium hydroxide and, in several runs, sodium dihydrogen phosphate-disodium hydrogen phosphate, with concentrations as suggested in Perrin and Dempsey ( 2 1 ) . The buffered sample was brought to the desired reaction temperature in a water bath controlled to f O . l "C and was divided among a number of 2.5-L glass vessels with screw tops. Sodium hypochlorite solution (5%) obtained from a laboratory supplier was pH adjusted with the sulfuric acid to within f0.5 pH unit of the buffered biphenyl-water solution. The desired volume of this chlorine solution was added to each 2.5-L vessel, and the vessel was filled to the top with buffered biphenyl-water, sealed, and placed in the water bath. I t was found to be important to have the reaction vessels filled to prevent vaporization of a portion of the low-concentration biphenyl into an air space. At some time during reaction, a 1.0-mL sample was withdrawn from each vessel for a chlorine analysis by the otolidine colorimetric method ( 2 2 ) .A reaction time of 12 h was used in all runs. This time was required to produce accurately analyzable product concentrations. For all reactions except the Arrhenius studies, a temperature of 40.0 "C was used. To terminate the reactions, 2 g of sodium thiosulfate was added to each vessel. A blank sample containing no chlorine accompanied each run. The pH was measured using a specific ion meter and pH electrodes and employing the instrument's built-in temperature correction adjustment. Sample Analysis. Immediately after reaction termination, a 1.5-L aliquot of each sample was extracted with 0.015 L of high purity hexanes. Two subsequent extractions were made, each with 0.015 L of fresh hexanes, and the three extracts were combined and dried by passing over 20 g of anhydrous sodium sulfate. The dried extracts were brought to 0.050 L volume with hexanes in volumetric ware. Sample analyses were performed on a Varian 204 gas chromatograph equipped with dual columns and dual detectors. A 10 f t X 1/8 in. (3.05 m X 0.0032 m) stainless steel column packed with 5% OV-101 on SO/lOO mesh Chromosorb 00 13-936X/79/09 13- 1244$0 1.OO/O
@ 1979 American Chemical Society
W-HP was used in conjunction with a flame ionization detector, and an identically sized stainless steel column containing 4% SE-30 on 100/120 mesh Chromosorb W-AW was used with a Sc 3H electron-capture detector. Operating conditions were selected so that peak heights were linear in concentration, and so peak heights were used in the quantitative analyses. Biphenyl analyses were performed by flame ionization using the OV-101 column. Chlorobiphenyl analyses were performed primarily by the same system, with routine verification by the electron-capture system. Normally a volume reduction was necessary for adequate sensitivity in chlorobiphenyl analysis. This concentrating step was performed by evaporating 95% of the hexane extract in a Kuderna-Danish evaporative concentrator, leaving a 0.005-L concentrate for analysis. For each analysis the chromatograph was calibrated using standards prepared in hexanes from reagent biphenyl and pure chlorobiphenyl isomers obtained from chemical suppliers. I t was found that recovery of organic material throughout the entire process exceeded 90% and was usually 95% or better.
Discussion and Results Chemistry of Chlorine in Water. Sodium hypochlorite, the reagent used in this investigation, ionizes in water as follows: NaOCl
+ HzO
-
HOCl
+ NaOH
The hypochlorous acid produced is a weak acid, and partially dissociates in water to form the hypochlorite ion as follows: HOCl e H +
+ OC1-
(2)
In the pH range from 5 to 9 this reaction is incomplete, and has an ionization constant:
K . = [H+][OCl-] (3) ' [HOC11 which has the numerical value 3.66 X mol L-l a t 40 "C, the temperature a t which this study was conducted. This value was obtained by extrapolating the acid ionization constant data of Morris (23),which covered the range of 5 to 35 "C, to 40 "C. It is apparent that the relative proportions of HOCl and OC1- will be strongly dependent upon pH, an observation borne out in this study. Several additional reactions of chlorine in water are important. First, a t pH values below about 6, molecular chlorine exists in solution in the following equilibrium: C12
+ H2O e HOCl + HC1
(4)
which has an equilibrium constant of 5.6 X lo-* a t 40 "C. Above a p H of 6, molecular chlorine concentrations are vanishingly small ( 2 4 ) . Second, there is a dehydration reaction to produce chlorine monoxide, which occurs to a limited extent in solution as represented by the following equation (25): 2HOC1+ C120
+ H20
(5)
with equilibrium constant:
a t 25 "C. Relative reactivities of the various chlorine-containing species influence the rates of chlorination reactions. Among the species known to exist in aqueous solutions of chlorine, Clz and ClzO are the most reactive (24,26), although molecular chlorine exists only in acidic solutions. Intermediate in reactivity is HOC1, and of very low chlorinating ability is
oc1-.
Products of Reaction. Three monochlorobiphenyls are possible from the reaction of chlorine with biphenyl: o-, m-,
Table 1. Representative Data for Chlorination of Biphenyla [CIZHIOI! [HOClt I, lo-ClzHgCl], [P-CizHgCI], Pg L-1
3770 3770 3770 3770 3770 3400 2720 2040 1360 680 3590 3590 3590 3590 3590 a
P g L-I
17.9 38.5 54.7 74.7 90.5 247 247 247 247 247 304 304 304 304 304
pH
7.00 7.00 7.00 7.00 7.00 6.78 6.78 6.78 6.78 6.78 7.06 7.47 8.04 8.31 9.17
PCg L-l
1.2 5.7 11.0 23.0 34.3 222 163 122 83.4 42.5 187 92.0 18.4 7.4 3.4
P9 L-1
0.3 3.8 8.3 15.8 21.7 134 99.1 62.5 47.7 23.8 125 56.5 12.2 4.8 2.9
Reactiontime 12.0 hat 40.0 O C . Experimentallymeasured solubility of in water was about 5000 pg L-'.
C12H10
and p-chlorobiphenyl. If reacting conditions are vigorous enough or the reaction is allowed to proceed for a long enough time, a number of dichlorobiphenyl isomers may also be produced. Even more reactive conditions produce tri-, tetra-, and higher polychlorinated biphenyls. In the reactions performed during this study, chromatographic peaks corresponding to the monochlorobiphenyls were observed. One peak apart from the others was obviously ochlorobiphenyl. A second peak was detected that, upon first inspection, could have been due to either m- or p-chlorobiphenyl or both, since the retention index values for these two isomers are almost identical (27). However, inspection of relative molar responses of the two chromatographic detectors used in this work to m- and to p-chlorobiphenyl provided a method of determining which isomer formed this peak ( 2 8 ) . By comparing the relative heights of the two monochlorobiphenyl peaks on the flame ionization detector and on the electron-capture detector, it was found that over 97% of the height of the second peak was due to p-chlorobiphenyl. This result was supported by Beaven et al., who found no measurable rn-chlorobiphenyl under acidic reaction conditions ( I 7). Therefore, this peak was taken to be entirely p- chlorobiphenyl in this study. The 1/2 o:p ratio was found to be 0.8, and did not vary significantly with reaction conditions. Except in the several acid-catalyzed runs, no measurable concentrations of dichlorobiphenyls or higher isomers were found. Representative data are presented in Table I. The Kinetic Method. Since the reactions studied are series parallel in nature, it was necessary to use a method of data collection that would minimize the effects of the competing reactions. The method chosen was that of initial rate as a function of initial concentration, described by Frost and Pearson (29). In this method experiments are performed a t varying initial concentrations of each reactant, and the reactions are terminated before a limiting reactant uptake of 10% is achieved. By performing studies in this manner, it introduces only a small error to assume that the reactant concentrations are constant a t their initial levels. Also, by producing only low product concentrations the tendency for the series reactions to proceed to a second step is minimized. A mass action form of rate expression was postulated, in which the reacting species are presumed to be HOC1, biphenyl, and hydrogen ions: Volume 13, Number 10, October 1979
1245
-8
I
I
-7
-6
I
-5
1
-4
In [HOClt] ,mol L - '
Figure 1. A log-log plot of chlorobiphenyl formation as a function of total chlorine concentration at pH 7.0 for 12-h reaction time at 40 "C: (0) O-CIZHQCI; ( 0 )p-CizHgCl
=e 8.5
8.0
7.5
70
PH
Figure 3. Dependence of reaction rate on pH for 12-h reaction time at 40 "C: (0)OCi2HgCI; ( 0 )P - C ~ ~ H Q C I
and x p were 2 (actual values obtained were x , = 1.97 and x p = 1.99) and the values of y o and y p were 1 (actual values obtained were yo = 1.02 and y p = 1.09). The calculated reaction orders were very close to integral values, and it was believed that the integral orders suggested above would represent the data as accurately as the orders predicted by the least-squares program. To verify this, error analyses were performed on the data for both integral and least-squares orders, and there was no significant difference in accuracy. Therefore, the integral values of 2 and 1 were accepted as correct for the values of PH reaction orders in chlorine and biphenyl, respectively. Plots Figure 2. Observed dependence of reaction rate on pH with total hyof the data for these first two sets of runs are shown in Figure pochlorous acid rate law for 12-h reaction time at 40 "C: (0)OC12HgCI; 1. In these runs, all reactions were performed a t 40 "C for 12 ( 0 )P - C ~ ~ H ~ C I h; the ordinate of Figure 1presents the molar concentration of products resulting after 12 h of reaction. With the data d[0-C12H&l] presented in this manner, a straight line should result in each = k , [HOC1]X~[C12H1~]Y~[H+]Zo (7) dt case with a slope of 2.0. That this is the case is confirmed by Figure 1, which indicates a good fit of the data to the postudb-C12HgC11 = kp[HOCl]X~[C~~Hlo]Y~[H+]z~ (8) lated expression. dt The next series of runs was performed to study the effects T h e values of the reaction orders x , y, and z were unknown of p H on reaction rate. Runs were made with pH values and were to be determined by the data. The first series of exranging from 6.8 to 8.5, with the pH for each run selected and periments was performed a t constant pH by adding buffers. maintained by buffers. It was found that reaction rate deIn these runs the effeqts of the hydrogen ion were incorporated creased with increasing pH. Two mechanisms seemed plauinto the rate constant. Also, the measured total chlorine sible and were investigated. In the first, it was postulated that concentration data were used, without taking into account the the increase in reaction rate with [H+]occurred due to direct proportion of HOCl to HOCl OC1-, since a t constant pH reaction of hydrogen ions, and not due to changes in the this proportion does not change. Therefore, the modified rate HOCI-OCl- equilibrium. Therefore the postulated rate exexpressions for the first sets of runs were: pressions for this case were written in terms of total chlorine concentrations and including a mass action term in [H+],as in Equations 7 and 8. By taking 12-h data a t 40 "C as described above and plotting logarithmically, straight lines of slope t oand z p should db-C12HgC11= k,,,,[HOCl~] X~[C12H1~]Y~(10) dt result if indeed there is a direct dependency of reaction rate on pH. As seen in Figure 2, the resulting lines are curved, inwhere the subscript t on HOCl refers to total measured chlodicating that the above assumption does not seem aderine concentration, and: quate. In the second mechanism, it was postulated that the change in reaction rate resulted entirely from the change in HOCIOC1- equilibrium from Equations 2 and 3 and not from a diWith these constraints one set of runs was made in which the rect reaction of hydrogen ions. In this case the rate expressions chlorine concentration was varied over a wide range, while were formulated with undissociated hypochlorous acid as the biphenyl concentration was held constant or nearly so, and active chlorinating species, and an exponential dependence a second set of runs varied the biphenyl concentration, while upon hydrogen ion concentration was retained in the event chlorine concentration was held nearly constant. The pH for that z,, and z p were nonzero, again shown by Equations 7 and both sets of experiments was 7.0, maintained by buffers. The 8 with [HOClIt replaced by [HOCl],, where the subscript u data obtained from these runs were fitted to a linear leastdenotes hypochlorous acid in the undissociated form. As besquares computer program to predict the orders x,, yo, x p , and fore, the data were plotted logarithmically in Figure 3. The y p . The results of this program showed that the values of x , ordinate again represents the molar concentrations of chlo-
+
1246
Environmental Science & Technology
5 r
200.-
A
4-
3-
N
J
u
1
-
0 05
E
OO 2-
1
'310
0 10
0 15
0 20
P Figure 5. Effect of ionic strength on reaction rate for 12-h reaction time at 40 "C: (0) o-C&I~CI; ( 0 )p-Cj2HgCI 325
340
X IO:
355
370
K-'
lo/
Figure 4. Arrhenius plots for chlorobiphenyl formation for 12-h reaction time: (0) o-Ci2HgCI; ( 0 )p-Ci2HgCI
robiphenyls produced after 12-h reaction time at 40 "C. The slopes of the resulting lines for 0-and p-chlorobiphenyl represent the values of zo and z,, respectively. As seen in Figure 3, both of these values are approximately zero. Least-squares fits gave slopes of +0.10 for the ortho line and +0.06 for the para. Hence, the value of zero was taken for both z, and 2,. This second explanation satisfied the data and the known relative reactivities of HOCl and OC1- much better than the first, and so this interpretation was accepted. The final rate expressions, based on concentrations of measurable species HOC1, C12H10, and the monochlorobiphenyls, were as follows:
By analyzing the data for all the above runs in light of these rate expressions, values of the rate constants k , and k , were obtained, and were found to be 130 and 82 L2 mol-2 h-l, respectively. Variation of rate constants with temperature from 0 to 40 "C was determined by the use of the Arrhenius equation:
which includes two constants, the preexponential or frequency factor A and the energy of activation E,. The energy of activation for each reaction was obtained by plotting In k VS. inverse temperature, the slope of the resulting line being negative energy of activation divided by the gas constant. The Arrhenius plots are presented in Figure 4. The resulting values of E , based on these data are: o-C12H9C1,E , = 14.7 kcal mol-I; p-C12H9C1,E , = 14.2 kcal molp1. Ionic strength variations from 0.02 to 0.20 were investigated by varying the composition and concentration of the buffers used to maintain pH. A plot of the data, shown in Figure 5 , shows that within this limited range ionic strength variations had no measurable effect on reaction rate. Iron metal and iron salts were added to several runs to investigate the effects on reaction rate of a common catalyst in aromatic halogenations. No changes in reaction rate were seen, a result that verified the absence of molecular chlorine, since the chief effect of iron catalysts seems to be the polarization of the chlorine molecule to produce increased electrophilic substitution (29). Acid catalysis was observed in the several runs that were performed at pH values below 6.2. In this pH range, molecular
J +/ 1
1
1
75
70
1
65
1
6 0
I
5 5
PH
Figure 6. Reaction rate as a function of pH for region in which acid catalysis begins for 12-h reaction time at 40 " C : (0)o-C12H9CI; ( 0 ) P-CizHgCI
chlorine is present, and so an increased reaction rate due to this highly reactive species is expected. A plot of observed reaction rate vs. pH is presented in Figure 6, indicating the rapidly increasing rate of reaction under increasingly acidic conditions. Probable Chlorinating Species. Several species are known to exist in aqueous solutions of chlorine. As described above, HOCl,OCl-, Cla, and C120 may be present, depending upon conditions. Other species, such as ClOHz+ and C120H+, have been postulated ( 3 1 ) . The species ClOH2+ is thermodynamically unfavorable (32), but the others are present. Since a second-order expression in chlorine concentration resulted from the data, the species Clp and ClzO (and its protonated form C120H+) seem most logical as the active chlorinating agents. Molecular chlorine is known to be absent in the pH range over which most of this study was performed. Therefore, ClzO and C120H+ may account best for the limited chlorination which occurred. The equilibrium constant from Equation 6 shows that a small but finite concentration of reactive ClzO is present in aqueous chlorine solutions. Several previous studies have concluded that aqueous C120 was the active agent in chlorinating allyl alcohol (33) and anisole (34).
Conclusions From the results above it is concluded that aqueous chlorination of biphenyl at pH values above 6.2 proceeds very slowly, with production of 0- and p-chlorobiphenyl. It appears from the rate expression and relative reactivities that chlorine monoxide rather than HOCl may be the active chlorinating species in reactions with such unreactive organic compounds as biphenyl. From this study it may be argued that the rate of formation of chlorobiphenyls is so slow that little concern for health or the environment is warranted. This has indeed been shown to be true for low chlorine concentration disinfection of water for short times, such as in tertiary treatment. However, since Volume 13, Number 10, October 1979
1247
reaction rate is proportional to the square of undissociated hypochlorous acid concentration, concern should be shown for cases in which higher chlorine dosages may be used. One instance is the chlorine oxidation of wastewater sludges as a stabilization procedure prior to dewatering (35). In such processes chlorine doses of 2000 mg L-l are routine. If a sludge contains biorefractory organics such as biphenyl either dissolved in the aqueous suspension or physically or chemically adsorbed to the sludge particles (36),extensive chlorination of the organics may occur. If, as is the usual case, the liquid effluent from the subsequent dewatering process undergoes biological treatment, further chlorination, and discharge, release of significant concentrations of such chlorinated organics as chlorobiphenyls may result. Further work is warranted to verify these suppositions. Also, work is needed in the area of reactions of chlorine with biphenyl in the presence of ammonia and chloramines, as will normally be the case in actual water treatment practice. In addition, studies of the chlorinating abilities and other properties of C120 in aqueous chlorine solutions are needed. Extension of this to real and complex wastewaters would be desirable. Several time-consuming and intricate analytical separation steps would likely be required in this case, and a more definitive identification tool such as a gas chromatograph-mass spectrometer combination would be required. Literature Cited
(1) Poffenberger, N.; Hubbard, H. L. In “Encyclopedia of Chemical Technology”, 2nd ed.; Standen, A,, Ed.; Wiley: New York, 1968;Val. 7. (2) Chem. Eng. (N.Y.)1976,83(18),66-7. (3) Bellar, T. A,; Lichtenberg, .J. J.; Kroner?R. C. “The Occurrence of Organohalides in Chlorinated Drinking Waters”, 1974, 1J.S. Environmental Protection Agency; EPA-fiSOi4-74-008. (4) Harrison, R. M.; Perry, R.; Wellings, R. A. Enciron. Sei. Technol. 1976,10, 1156-60. ( 5 ) Lee, G. F.; Morris, J. C. Int. J . Air Water Pollut. 1962, 6 , 41931. (6) Glaze, W. H.; Henderson, .J. E. IV J Water Pollut. Control Fed 1975,47, 2511-5. (7) Glaze. W. H.: Henderson. ,J. E. IV: Bell. J . E.: Wheeler. V. A. J . Chromatogr, Sci. 1973, 21, 580-4. (8) Carlson. R. M.: Carlson. R. E.: Komerman. H. L.: Caole, R. E n uvon. Sci Technol. 1975,9, 674-5.‘ (9) Rook, J. J. Water Treat Exam 1974,23, 234-43. (10) Morris, J. C.; McKay, G. M. “Formation of Halogenated Organics ~
by Chlorination of Water Supplies”, 1975, U.S.N.T.I.S., P.B. Report 1975, No. 241511. (11) Weil, I.; Morris, J. C. J . Am. Chem. SOC.1949, 71, 1664-71. (12) Robertson, P. W. J . Chem. Soc. 1954,1267-70. (13) de la Mare, P. B. D.; Harvey, J. T.; Hassan, M.; Varma, S. J . Chem. Suc 1958,2756-9. (14) Brown, H. C.; Stock, L. M. J . Am. Chem. Soc. 1957, 79, 51759. (15) Edmond, C. R.; Soper, F. G. J . Chem. Soc. 1949,2942-5. (16) de la Mare, P. B. D.; Ketleg, A. D.; Vernon, C. A. J . Chem. Soc. 1954,1290-7. (17) Beaven, G. H.; de la Mare, P . B. D.; Hassan, M.; Johnson, E. A.; Klassen. N. V. J . Chem. Soc. 1961.2749-55. (18) Gaffney, P . E. Science 1974,183, 367-8. (19) Gaffney, P . E. J . Water Pollut. Control Fed. 1977,49, 401-4. (, 2 0,) Johnsen. R. “Chlorination of Waters for Disinfection-A Studv of the Production of Undesirable Chlorinated Products”, Proceedings of the National Conference on PCB’s, Chicago, Ill., 1975; EPA-56016-75-004. (21) Perrin, D. D.; Dempsey, B. “Buffers for pH and Metal Ion Control”; Chapman and Hall, Ltd.; London, 1974. (22) “Standard Methods for the Examination of Water and Wastewater”, 13th ed.; American Public Health Association: New York, 1971; pp 117-23. (23) Morris, J . C. J . Phys. Chem. 1966,70, 3798-805. (24) White, G. C. “Handbook of Chlorination”; van Nostrand Reinhold: New York, 1972; p 186. ( 2 5 ) Skrabal, A.; Berger, A. Monatsh. Chem. 1937, 70, 168-92. (26) Palin, A. T. In “Disinfection-Water and Wastewater”; Johnson, J . D., Ed.; Ann Arbor Science Publishers: Ann Arbor, 1975. (27) Albro, P . W.; Fishbein, L. J . Chronatogr, 1972,69, 273-83. (28) Safe, S. “Overview of Analytical Identification and Spectroscopic Techniques”, Proceedings of the National Conference on PCB’s, Chicago, Ill., 1975; EPA-56016-75-004. (29) Frost. A. A,: Pearson. R. G. “Kinetics and Mechanism”. 2nd ed.: Wiley: New York, 1961; pp 43-6. (30) Morrison, R. T.; Boyd, R. N. “Organic Chemistrv”. :3rd ed.; Allvn and Bacon: Boston, 1973; p 349. (31) de IaMare, P. B. D.; Hilton, I. C.; Varma, S. J . Chem. Soc. 1960, 4044-54. (32) de la Mare, P. B. D.; Ridd, ,J. H. “Aromatic Substitution”; Academic Press: New York, 1959; p 117. (33) Israel, G. C.; Martin, J. K.; Soper, F. A. J . Chem. SOC.1950, 1282-6. 134) Swain, C. G.; Crist, D. R. J . Am. Chrm. SOC.1972, 94, 3195-
‘Loo.
(35) “Process Design Manual for Sludge Treatment and Disposal”, 1974, U.S. Environmental Protection Agency, Technology Transfer Series; EPA-62511-74-06;pp 5-29. (36) Choi, P. S. K.; Nack, H.; Flinn, J. E. Bull. Enuiron. Contam. TOX~C 1974,Il-, O ~ . 12-7.
Rereiced /or reuieic. April 9, 1979. Accepted June 15, 1979.
Dissolution of Tetragonal Ferrous Sulfide (Mackinawite) in Anoxic Aqueous Systems. 1. Dissolution Rate as a Function of pH, Temperature, and Ionic Strength James F. Pankow*l and James J. Morgan Environmental Engineering Science, California Institute of Technology, Pasadena, Calif. 91 125
A rather insoluble compound, mackinawite often forms in reducing sediments and is believed to form in anaerobically digested sewage sludge ( I ) . The trace metal constituents of these systems, which also form insoluble metal sulfides (e.g., Ag+, Cd2+,Cu2+,Pb2+,Zn2+),very likely associate with the much more predominant iron sulfides. This material may be expected to a t least partially dissolve: (a) when reducing sediments are disturbed by animal life, dredging, or by more subtle events, and (b) when anaerobic sewage sludge is disPresent address, Department of Environmental Science, Oregon Graduate Center. 19600 N.W. Walker Road. Beaverton. Orep. 97005. 1248
Environmental Science & Technology
posed of in the marine environment. The interest in this problem from the perspective of sewage sludge may be more easily understood when it is realized that, currently, approximately 5,7,9,95,and 120 metric tons per year of silver, cadmium, lead, copper, and zinc, respectively, are being discharged from an ocean outfall of the city of Los Angeles anaerobic sewage sludge facility ( 2 ) . Mackinawite has been characterized fairly recently ( 3 )and was first identified in recent sediments by Berner ( 4 ) . It is believed to be one of the initial iron sulfides to form in iron‘learing sediments the Onset Of production$with greigite (spinel Fe&d and Pyrite (cubic FeSd being other possible early minerals ( 5 ) .
0013-936X/79/0913-1248$01.00/0 @ 1979 American Chemical Society
