Environ. Sci. Technol. 1996, 30, 3408-3417
Kinetics of the Reduction of Hexachloroethane by Juglone in Solutions Containing Hydrogen Sulfide J U D I T H A . P E R L I N G E R , * ,† WERNER ANGST,‡ AND R E N EÄ P . S C H W A R Z E N B A C H ‡ Swiss Federal Institute for Environmental Science and Technology (EAWAG) and Swiss Federal Institute of Technology (ETH), CH-6047 Kastanienbaum and CH-8600 Du ¨ bendorf, Switzerland
Hexachloroethane was converted to tetrachloroethene and unknown products in solutions containing juglone and hydrogen sulfide. Measured rates of disappearance in solutions containing hydrogen sulfide alone were approximately a factor of 10 slower than rates containing micromolar concentrations of juglone and hydrogen sulfide. Electrochemicallyreduced juglone was unreactive with respect to hexachloroethane reduction. Reaction of hexachloroethane with polysulfides produced in the reaction of elemental sulfur with hydrogen sulfide in the experimental solutions also could not account for the rate observed in solutions containing juglone and hydrogen sulfide. Evidence is provided to indicate the reaction of hexachloroethane with the Michael addition product of hydrogen sulfide and juglone. No conclusions can be drawn from the present results as to whether the reaction mechanism is a one- or a two-electron transfer. This study points out the importance of the geochemistry of sulfur and organic matter in the transformation of halogenated alkane pollutants in reducing environments.
Introduction The transformation of halogenated alkanes in reducing sediments has been shown in batch experiments to proceed with half-lives of minutes to days (1, 2). A number of abiotic components of aquatic, reducing environments are capable of transforming halogenated alkanes. The most abundant components are water, redox-active and nucleophilic functional groups in natural organic matter, and various inorganic and organic forms of sulfur and iron (3). The relative reactivity of these various components depends highly on their speciation. This speciation is in turn * Corresponding author present address: Civil & Environmental Engineering Dept., Michigan Technological University, 1400 Townsend Dr., Houghton, MI 49931-1295; e-mail address:
[email protected]. † Kastanienbaum. ‡ Du ¨ bendorf.
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influenced by environmental factors such as pH, temperature, microbial activity within a given environment, and hydrogeologic characteristics. Halogenated alkanes undergo three types of transformation reactions with the above species in the absence of oxygen and light: electron transfer, nucleophilic substitution, and dehydrohalogenation. Nucleophilic substitution and HX elimination products have been found for the reaction of halogenated alkanes with water (4) and sulfur species (5-8). In the case of one-electron transfer reactions with halogenated alkanes, the identification of reaction mechanisms from reduction products is more difficult due to followup reactions of the radicals produced in the initial one-electron reduction reaction. For example, in the presence of HS-, carbon tetrachloride was converted to CO2 in the presence of vermiculite or biotite (9). In the presence of HS- and freshly ground pyrite, however, 50% of the total CCl4 added was converted to CHCl3 and traces of CS2, CO2, and formate to give a mass balance of 78% (10). Such reaction products are consistent with an initial singleelectron transfer reaction mechanism. The objective of this study was to mimic a sulfatereducing, organic matter-rich aqueous environment in order to evaluate the relative reactivities of the various species in solution with respect to transformation of polyhalogenated alkanes. Hexachloroethane was used as a probe molecule representative of polyhalogenated alkanes. It was chosen because it is commercially produced and thus could itself be an environmental pollutant; it is fully halogenated, and so it is reduced more rapidly than less-halogenated congeners; and finally because previous work in a lake water system suggested that it might react with polysulfides via a “nucleophilic elimination” reaction to produce tetrachloroethene (11). Such a mechanism has been employed to synthesize dichlorophosphoranes from phosphines and hexachloroethane (12):
RR′R′′P + C2Cl6 f RR′R′′PCl2 + C2Cl4
(1)
The experimental solutions contained juglone (5-hydroxy1,4-naphthoquinone) at various concentrations and hydrogen sulfide buffered at various pH values. Juglone in these solutions mimics quinone functional groups in natural organic matter. Anthraquinone disulfonic acid (AHQDS) has been similarly used to examine the reduction of hexachloroethane (13). Curtis and Reinhard reported a significant increase in the reaction rate of hexachloroethane in the presence of AHQDS and sulfide compared with the reaction with AHQDS or sulfide alone. Various species may be produced in the reaction of juglone with hydrogen sulfide (Table 1), and each has a given reactivity with respect to hexachloroethane transformation. Reduction of juglone by HS- produces the hydroquinone form, which may reduce the halogenated alkanes. As discussed in detail elsewhere (14), juglone may undergo a Michael addition reaction with hydrogen sulfide, as was the case in the solutions tested in this work. The oxidized form of the addition product(s), which will from hereon be referred to as mercaptojuglone(s), can be reduced to the semiquinone form (eq 3, Table 1) or to the hydroquinone form (eq 4, Table 1) by HS-. Note that in Table 1 the most likely formed isomer is shown. Polysulfides
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1996 American Chemical Society
TABLE 1
Reactive Species in Solutions Containing Juglone and Hydrogen Sulfidea OH
OH
O
OH
+ H2S
+ S0 OH
O OH
OH
O
+ H2S
+ HS• O•
O OH
(2)
OH
OH
O
OH
(3) + H2S
+ HS• SH
SH O•
O OH
OH
O
OH
(4) + H2S
+ S0 SH
SH O
OH
HS- + 3S0 h -S-S-S-S- + H+ a
(5)
The reduction potential of the juglone depends on pH and is given
by:
Eh ) 0.428 +
+ 3 + 2 + RT [Jug0] RT {H } + Kr1{H } + Kr1Kr2{H } ln ln + + nF [Jugr] nF {H } + K0
where R is the ideal gas constant (8.314 J mol-1 K-1), T is absolute temperature (K), F symbolizes Faraday’s constant (96 490 C mol-1), Jugo and Jugr are the total concentrations of oxidized and reduced juglone, respectively, {H+} is hydrogen ion activity, and K0, Kr1, and Kr2 are the acidity constants of the oxidized form of juglone and the reduced form of juglone, respectively, which have values of 1.4 × 10-9 (41), 2.5 × 10-7, and 2.5 × 10-11 (35). The computed reduction potential of a solution having a pH of 7.0 and Na2S ) 1 mM is -181 mV (3). Substituting this reduction potential into the above equation leads to the conclusion that under these conditions, juglone would be expected to be virtually completely reduced.
are also produced by the reaction of hydrogen sulfide with the elemental sulfur produced in the reduction of quinones (eqs 2 and 4). Thus, numerous species (including sulfide species, reduced juglone species, reduced mercaptojuglone species, and various polysulfides) could be reactive with respect to the tranformation of halogenated alkanes in solutions containing juglone and hydrogen sulfide.
Methods Chemicals. The purities and producers of the chemicals used in this work were as follows: hexachloroethane, 98%, Aldrich; tetrachloroethene, 99.5%, Fluka; 1,1,1,2-tetrachloroethane, 99%, Aldrich; 1,1-difluoro-1,2,2,2-tetrachloroethane, 97%, Aldrich; 3-chlorofluorobenzene, 99%, Aldrich; 4-(2-hydoxyethyl)-1-piperazineethanesulfonic acid (HEPES), microselect, Fluka; glycine, puriss, Fluka; sodium sulfide, puriss, Fluka; erythro-1,4-dimercapto-2,3-butanediol, 99%, Aldrich; 5-hydroxy-1,4-naphthoquinone (juglone), purum, Fluka; methanol, nanograde, Burdick & Jackson; n-hexane, n-pentane high purity, Burdick & Jackson. Preparation of Samples. Samples were prepared by autoclaving 50-mL (nominal volume) serum flasks, buffer, stoppers, and all glassware used in preparing the samples
at 121 °C for 30 min. Viton stoppers (Maagtechnik, Du ¨ bendorf, Switzerland) were used in this work. Viton is a copolymer of vinylidene fluoride and perfluoropropylene (15) and is resistant to diffusion of the halogenated alkanes. The ability of viton to prevent the loss of hexachloroethane and tetrachloroethene from the serum flasks during a time period of 30 days was demonstrated (16). The buffer was stirred and purged with argon gas (99.999%; used without further purification) while it cooled after autoclaving. Each 300 mL of buffer was purged for at least 3 h. The buffer was then transferred to a glovebox, and the samples were prepared beyond that point in the glovebox. Either 50 mM HEPES buffer or 61 mM glycine buffer was added to the serum flasks depending on the desired pH. At pH values e8.5, HEPES buffer was used, and at pH values g8.5, glycine buffer was used. Experiments comparing reduction rates of hexachloroethane in solutions containing HEPES and glycine buffers at pH 8.5 showed no difference in rate. An aliquot of 1 M HCl was added to compensate for the sulfide introduced later. An Na2S stock solution was prepared by rinsing Na2S crystals with argon-purged distilled water, weighing them, and finally diluting them with argon-purged distilled water to give the desired final concentration (typically 0.5 M). Aliquots of the Na2S stock solution followed by aliquots (0.285-1.43 mL) of a concentrated solution of juglone in methanol (0.002 or 0.02 M) were added with a digital pipet. The total volume of the solutions was 57 mL. There was approximately 1 mL of headspace in the flasks after the vials were stoppered in the glovebox. The flasks were immersed in a 25 °C water bath (temperature control (1 °C) that was covered to prevent light-induced reactions, and the samples were allowed to equilibrate for at least 12 h. Polysulfides were prepared by equilibrating elemental sulfur with a 0.5 M Na2S solution for 1 month. The final measured concentrations of the sulfur species were 0.160 M H2ST ) (H2S + HS- + S2-) and 0.342 M ∑n)2-6(H2Sn + HSn- + Sn2-). The measured hydrogen sulfide and polysulfide concentrations in the stock solution were within 6% of the predicted equilibrium values. Aliquots of this stock solution were then added to serum flasks along with an aliquot of 1 N HCl. Note that, although colloidal elemental sulfur may have been present in the aliquots of stock solution added to the serum flasks, it is possible that the solutions in the serum flasks were undersaturated with respect to elemental sulfur, and therefore polysulfides may be present at less than saturation concentrations. Kinetic Experiments. The kinetic experiments were performed by adding 570 µL of a solution containing the substrate and the internal standard (3-chlorofluorobenzene) in methanol to the reaction medium and sampling at appropriate time intervals. The initial concentration of substrate in all experiments discussed here was ∼1 µM. The final methanol concentration in the samples was e0.5 M. The methanol for the spike solution was oxygen saturated, but hydrogen sulfide reduces oxygen. For a saturation concentration of oxygen in methanol of 9.81 × 10-3 M [1 bar O2, 25 °C; (17)], the initial concentration of oxygen in the samples would be ∼20 µM. At the lowest HS- concentrations employed here at a pH of 7.0 (5 × 10-4 M HS-), the half-life of the oxygen is 21 h (18). At higher hydrogen sulfide concentrations and pH values, the oxygen would be consumed more rapidly. Samples were withdrawn into a syringe after injection of an equal volume of nitrogen gas. The purity of the
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nitrogen gas was 99.995%, and it was used without further purification. The sample was placed in a 3-mL glass vial with a plastic screwtop containing hexane or pentane. Usually 0.25 mL of the sample was added to 0.2-1.0 mL of organic solvent. The solution was then shaken for 30 s on a vortex mixer in order to extract the sample. After extraction, the hexane (pentane) phase was transferred to autosampler vials containing sodium sulfate added to dry the hexane (pentane). Eight to 12 samples were taken from one serum flask over 1-3 half-lives. The volume of the headspace increased in the serum flasks with each sample taken. Aqueous concentrations were corrected for the amount of solute that partitioned into the headspace according to:
[RX]corr ) [RX]measured +
[RX]measuredKHVn(gas) Vn(liquid)RT
(6)
where [RX]corr is the corrected aqueous concentration of a halogenated alkane RX (mol L-1), [RX]measured is the measured aqueous concentration of RX (mol L-1), KH is the Henry’s law constant of the halogenated alkane (atm L mol-1), Vn(gas) ) Vo(gas) + n2.5 × 10-4 L, Vn(liquid) ) Vo(liquid) - n2.5 × 10-4 L, R is the ideal gas constant (0.082058 L atm mol-1 K-1), n is the sample number, and T is absolute temperature. The values of Henry’s law constant at 25 °C used for hexachloroethane, tetrachloroethene, 1,1-difluorotetrachloroethane, and 1,1,1,2-tetrachloroethane, respectively, were 8.35 (19), 17.7 (20), 14.1 (21), 2.76 atm L mol-1 (22). Kinetics measurements of the disappearance of hexachloroethane were also performed by headspace analysis. In a sample containing 57 mL of solution and approximately 1 mL of headspace, 200 µL of gas from the headspace of the samples was withdrawn by syringe and injected directly into the gas chromatograph. Aqueous concentrations were determined by comparison with headspace analysis of external standards consisting of aqueous solutions containing the identical buffer to that used in the samples and added to vials with the same sample headspace and water volumes. Analyses. Gas chromatography (GC) was carried out on either a Carlo Erba HRGC 5160 or a Fisions Instruments GC 8165, both of which were equipped with autosamplers and Cryo 520 CO2 cooling systems. The injection volume was 1.5 µL. Two types of fused silica columns were used, either a 30 m DB-624 (i.d. 320 µm, 0.25 µm film thickness) or a 30 m DB-5.625 (i.d. 320 µm, 1.8 µm film thickness; both from J&W Scientific). The injector and detector temperatures were 200 and 300 °C, respectively. The carrier gas was H2 at a flow rate of 1.5 mL min-1. The detector used was an electron capture detector (ECD), Carlo Erba ECD 400, with a 63Ni source. The makeup gas for the ECD was a mixture of 90% Ar and 10% CH4. Injection was splitless. The response of a given substrate was computed relative to the internal standard in the sample and compared with the relative response of the substrate in standards in order to determine the aqueous concentration. Sulfide concentration was determined by iodometric titration (23). The concentration of polysulfides in the stock solution was measured as the difference in iodometrically determined hydrogen sulfide concentration and colorimetrically determined (methylene blue) hydrogen sulfide concentration (16).
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FIGURE 1. Sample plot of the disappearance of HCE from a solution containing 50 mM HEPES buffer, 1 mM Na2S, and 200 µM juglone, pH 7.05. The line is the exponential fit of the data.
The electrochemical reduction of juglone and the determination of the structure and concentration of radicals in the solutions by electron paramagnetic resonance spectroscopy (EPR) are discussed in detail elsewhere (14). Briefly, 100 µM juglone in approximately 250 mL of 50 mM HEPES buffer (ionic strength 0.1 M by KCl addition) at pH 6.95 was reduced at -200 mV at a platinum net working electrode (50 × 50 mm, 0.0762 mm o.d. fine-mesh blackened platinum net) while stirring. The reduction required approximately 3 h to reach completion. The complete reduction was verified by comparison of the UV-Vis spectrum with that of the reduced hydroquinone (λmax of 343 nm). EPR measurements of juglone and mercaptojuglone semiquinone species were carried out on a Varian E-9 Century Series spectrometer with a 9-in. magnet, an E-102 microwave bridge, an E-204 low frequency module, and a data aquisition system. Calibration was performed using 4-hydroxy-2,2′,6,6′-tetramethylpiperidinyloxy free radical at concentrations of 10 µM-1 mM. Equilibrium concentrations of hydrogen sulfide and polysulfides were determined by modifying the Turbopascal version of MINEQL (24) to consider reactions between elemental sulfur and sulfide species. Equilibrium constants for the formation of polysulfides from elemental sulfur and hydrogen sulfide are reported by Williamson and Rimstidt (25; Table 1), and acidity constants for hydrogen sulfide (pKa1 ) 6.98; pKa2 ) 18.5) and the polysulfides from refs 25-28 were used. For the reactions as written in eq 5 of Table 1, pK(n ) 3) ) 9.56, pK(n ) 4) ) 9.45, pK(n ) 5) ) 9.66, where n is the number of S0 atoms in the polysulfide.
Results The disappearance of hexachloroethane (HCE) in solutions containing Na2S, Na2S + S0, and Na2S + juglone was pseudofirst-order with respect to HCE concentration, and firstorder rate constants were determined according to
ln[HCE]t ) -kobst + ln[HCE]0
(7)
A sample plot of concentration vs time for a solution containing Na2S and juglone is shown in Figure 1 along with the exponential curve fit of the data.
kobs ) 9.7 ((8.4) × 10-6 [H2S] + 3.5 ((0.7) ×
TABLE 2
10-5 [HS-]
Rates of Hexachloroethane Transformation Measured in Solutions Containing Various Concentrations of Sulfide
Finally, when reaction with H2S is omitted, the equation is
[Na2S] (mM)
pH
kobs (10-7 s-1)
mass balancea
µb (mM)
time (h)
1 10 20 30 10 10 10 10 5d
6.85 7.01 6.81 6.80 7.98 8.43 8.90 9.90 7.01
2.3 ( 0.1c 3.3 ( 0.7 4.2 ( 0.3 5.8 ( 0.4 4.7 ( 0.3 3.1 ( 0.4 2.5 ( 0.1 3.3 ( 0.9 2.6 ( 0.2
50, 46 22, 23 62, 29 66, 16 33, 38 23, 50 23, 60 20, 65 3, 71
7.8 32.8 60.0 88.0 46.1 48.0 44.9 38.5 50.0
935 383 935 935 675 675 675 675 313
a [PCE] /[HCE] , [HCE] /[HCE] in %; t is the last time a sample was t o t o, taken, as reported in the last column. b µ ) 0.5 ∑cz2. c (1 SD. d Results from Roberts (11).
Reaction of Hexachloroethane with H2S and HS-. Firstorder rate constants measured in solutions containing 1-30 mM Na2S at pH 6.88 ( 0.13 increased with increasing hydrogen sulfide concentration (Table 2). Included in Table 2 is a rate constant measured by Roberts (11) for the disappearance of hexachloroethane at pH 7.01 and 5 mM Na2S. Also reported in Table 2 are the fractions of the initial hexachloroethane concentration, [HCE]0, that were converted to tetrachloroethene according to Cl Cl Cl
C
C
Cl Cl
Cl
Cl Cl + 2e–
C Cl
+ 2Cl–
C
(8)
Cl
The mass balances at time t, determined as 100% × ([PCE]t + [HCE]t)/[HCE]0 where [HCE]0 was determined from the y-intercept in eq 7, were less than 100%, indicating that some of the hexachloroethane was transformed to products not measured in our analytical scheme. Although pentachloroethane would have been detected with our analytical approach, it was not detected in any of the experimental solutions discussed in this work. Note, however, that β-elimination of pentachloroethane to form tetrachloroethene is over 1 order of magnitude faster than the reaction of hexachloroethane to form tetrachlorothene (29), thus causing the maximum accumulation of pentachloroethane to be too low to be detected. A preliminary analysis of the results in Table 2 indicated that the hypothesis that H2S reacts with HCE cannot be rejected. When the contributions of H2S and HS- are separately quantified:
kobs ) kH2S[H2S] + kHS-[HS-]
(9)
Estimates for the second-order rate constants kH2S and kHSwere determined by three multivariable regressions of the form in eq 9. When the two terms plus a constant for any additional reaction are included in the regression the equation is ((95% C.I.)
kobs ) 1.1 ((0.5) × 10-5 [H2S] + 1.2 ((0.7) × 10-5 [HS-] + 2.2 ((0.5) × 10-7
R2 ) -0.037 (11)
R2 ) 0.741 (10)
When reaction with the additional reagent is omitted the equation is
kobs ) 1.9 ((0.8) × 10-5 [HS-] + 2.1 ((0.7) × 10-7 R2 ) 0.473 (12) The best fit is obtained when reaction with H2S and an additional reagent are considered (eq 10). For the purpose of the following analysis, we will use values for kH2S and kHS- from eq 10 of 1.1 ((0.5) × 10-5 and 1.2 ((0.7) × 10-5 M-1 s-1, respectively. However, it is somewhat counterintuitive that kH2S and kHS- would be nearly equal. Because HS- has additional unpaired electrons, it would be expected to have a higher reactivity. Furthermore, the constant in eq 10 is of similar magnitude as compared to the kobs values in Table 2. This would indicate that there is a reaction of HCE with some reagent in the solution other than H2S or HS-. As demonstrated below, this reagent is not polysulfides. More measurements of the rate of reduction in the presence of Na2S (especially at pH < 7) are needed to determine if H2S is a reactant and sources, if any, of the additional reactant in these solutions. Reaction of Hexachloroethane with Polysulfides. Reactivity increased with increasing ∑Sn2- (Table 3). The fraction of the initial hexachloroethane that was converted to tetrachloroethene at each pH is reported in Table 3. In general, these fractions are lower than the values reported in Table 2. The influence of polysulfides on the reaction rate can be demonstrated according to the approach taken by Roberts et al. (8). The observed rate constant can be described by three rate constants plus a constant:
kobs ) kH2S[H2S] + kHS-[HS-] +
∑k
2-
Sn2-[Sn
]+
2.2 × 10-7 (13)
If the expressions for formation of polysulfides are written as in eq 5 in Table 1, kobs in eq 13 can be described in terms of hydrogen ion activity according to
kobs/[HS-] ) kH2S {H+}γ HS-/Ka1 + kHS- + γ HS- {H+}-1
∑k
-1 Sn2-KSn2-γSn2-
+ 2.2 × 10-7/[HS-] (14)
A plot of kobs/[HS-] vs γHS- {H+}-1 should yield an estimate of (kHS- + 2.2 × 10-7/[HS-]) from the y-intercept and ∑kSn2KSn2- from the slope when the first term in eq 14 is negligibly small (Figure 2). The results from Table 2 in which the pH was varied at constant total sulfide concentration are included in Figure 2. Also included in Table 3 and Figure 2 are two data points from Roberts (11). In the case in which no elemental sulfur was added, the first term in eq 14 was insignificant ( HS- ≈ SO32- g S2O32- > RSR′ . H2O [see review by Barbash and Reinhard (6) and references therein]. Phenyl-sulfur compounds are weaker nucleophiles than polysulfides by a factor of 1-3 orders of magnitude. Comparing reactivities on the basis of second-order rate constants, the second-order rate constant for the reaction of hexachloroethane with the polysulfides was 6.8 ((1.9) × 10-3 M-1 s-1. The total mercaptojuglone produced in the Michael addition can be deduced from the UV-Vis spectrum of the solutions to be equal to 2% of the total juglone under the experimental conditions shown in Figure 3a (14). Assuming that this amount is reduced to the hydroquinone form, the second-order rate constant for the transformation of hexachloroethane by mercaptojuglone hydroquinone is equal to (2.5 × 10-6 s-1-3.3 × 10-7 s-1)/ (0.02 × 2 × 10-4 M) ) 0.55 M-1 s-1. This rate constant indicates that the mercaptojuglone is more reactive than the polysulfides by nearly 2 orders of magnitude. According to these considerations, the computed reactivity of the
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mercaptojuglone would be too high to conclude that it functions as a nucleophile in SN reactions at carbon (eqs 16 and 19 in Figure 7) in these solutions. However, the three electron-donating hydroxy groups in mercaptojuglone hydroquinone would make it a somewhat better nucleophile than phenyl-S-, and thus the computations above cannot be regarded as conclusive. Furthermore, nucleophilic attack at chlorine (eq 18 in Figure 7) would be expected to be quite different than such attack at carbon. Correlation of the disappearance rate of a series of polyhalogenated alkanes with the change in free energy upon reaction with mercaptojuglone according to Marcus theory indicates that transfer of the first electron to these compounds is rate-limiting (16). This supports the reaction of hexachloroethane according to eq 16 in Figure 7 in the initial, rate-limiting step. However, the one-electron oxidation potential for mercaptojuglone used in the Marcus plot could be expected to be a surrogate measure of the nucleophilicity of the mercaptojuglone, as discussed above. Thus, no conclusions can be drawn from the present results as to whether the reaction mechanism is a one- or a twoelectron transfer. Environmental Implications. The results presented here indicate that completely reduced mercaptojuglone was the reactive species in the reduction of the halogenated alkanes in solutions containing juglone and hydrogen sulfide. Evidence for this hypothesis include the kinetics of these solutions compared to the kinetics in solutions containing species such as hydrogen sulfide, polysulfides, or juglone hydroquinone alone and compared to reported rates of reaction of these species with polyhalogenated alkanes. The possible reaction of nitro aromatic compounds and halogenated alkanes with mercaptoquinones formed in addition reactions in model systems (37) and in solutions containing natural organic matter and hydrogen sulfide (13, 38) needs to be reconsidered. The results presented here are especially applicable to sulfate-reducing environments. Organosulfur compounds and polysulfides will be formed by the mechanisms discussed here in such environments when humic substances come into contact with hydrogen sulfide. The addition of inorganic sulfur species onto organic matter appears to occur in marine and freshwater sediments and peat (14). Quinone functional groups in humic substances, which alone may not be efficient electron transfer agents, may become more efficient electron transfer agents through the addition of hydrogen sulfide. Polysulfides are formed in the environment through, among other pathways, the oxidation of hydrogen sulfide in the reduction of humic substances. Concentrations of 0.2 mM S0 have been reported in marine porewater and salt marsh water (39). Although the results presented here demonstrate that polysulfides are not more highly reactive with respect to hexachloroethane reduction than hydrogen sulfide at neutral pH, at higher pH their equilibrium concentrations increase and reaction rates are greatly increased. Halogenated alkanes entering sulfate-reducing environments could be expected to be transformed at faster rates due to the reactivity of hydrogen sulfide, organosulfur, and polysulfide species. The results of the model systems provide hints regarding possible reactants and mechanisms in the lake water system studied by Roberts (11). Roberts compared the reduction of hexachloroethane in lake water containing hydrogen sulfide (H2ST ) 7.1 mM) and organic matter (DOC ) 43
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ppm; L. Roberts, personal communication) with that in distilled water containing 5 mM hydrogen sulfide. Rates of reduction in the lake water were a factor of 12 higher than measured rates in the clean system. Traces of pentachloroethane were detected in the lake water samples, indicating that pentachloroethyl radicals were formed. The reaction mechanism was postulated to be a two-electron reductive elimination (eq 18, Figure 7). The elevated rates in the lake water were explained by invoking reaction with polysulfides. This work demonstrates that quinone functional groups in natural organic matter may serve as electron transfer mediators in solutions containing hydrogen sulfide and that, at the pH of the lakewater (6.8), equilibrium concentrations of polysulfides would not be high enough to account for a 12-fold increase in the rate. It is possible that electron transfer reactions produced pentachloroethyl radicals in that system. However, the second electron transfer and elimination reactions to form tetrachloroethene may have been faster than hydrogen-atom abstraction such that little pentachloroethane was observed, in analogy to the observations made in the present work. These results demonstrate the need to study the interplay between the biogeochemistry of natural systems and the transformation of xenobiotic compounds. This work shows that naturally occurring organosulfur compounds transform halogenated alkanes. The extent of formation of organosulfur compounds has been shown to vary systematically with environmental parameters such as lake trophic state and oxygen availability (40). Hence, systematic variations in the rates of reduction of xenobiotics in sediments also may occur.
Acknowledgments The authors wish to thank R. Stierli, who performed a number of the kinetic measurements. N. Urban, J. Buschmann, and L. Roberts kindly reviewed the manuscript.
Literature Cited (1) Jafvert, C. T.; Wolfe, N. L. Environ. Toxicol. Chem. 1987, 6, 827837. (2) Peijnenburg, W. J. G. M.; De Beer, K. G. M.; De Haan, M. W. A.; Den Hollander, H. A.; Stegeman, M. H. L.; Verboom, H.; Wolfe, N. L. Sci. Total Environ. 1991, 109/110, 283-300. (3) Schwarzenbach, R. P.; Gschwend, P. M.; Imboden, D. M. Environmental Organic Chemistry; Wiley Interscience: New York, 1993. (4) Roberts, A. L.; Jeffers, P. M.; Wolfe, N. L.; Gschwend, P. M. Crit. Rev. Environ. Sci. Technol. 1993, 23 (1), 1-39. (5) Schwarzenbach, R. P.; Giger, W.; Schaffner, C.; Wanner, O. Environ. Sci. Technol. 1985, 19, 322-327. (6) Barbash, J. E.; Reinhard, M. In Biogenic sulfur in the environment; Cooper, W., Ed.; American Chemical Society: Washington, DC, 1989; pp 101-138. (7) Barbash, J. E.; Reinhard, M. Environ. Sci. Technol. 1989, 23 (11), 1349-1357. (8) Roberts, A. L.; Sanborn, P. N.; Gschwend, P. M. Environ. Sci. Technol. 1992, 26 (11), 2263-2274. (9) Kriegman-King, M. R.; Reinhard, M. Environ. Sci. Technol. 1992, 26 (11), 2189-2206. (10) Kriegman-King, M. R.; Reinhard, M. Environ. Sci. Technol. 1994, 28 (4), 692-700. (11) Roberts, L. Ph.D. Dissertation, Massachusetts Institute of Technology, 1991. (12) Appel, R.; Schoeler, H. Chem. Ber. 1977, 110, 2382-2384. (13) Curtis, G. P.; Reinhard, M. Environ. Sci. Technol. 1994, 28, 23932401. (14) Perlinger, J. A.; Angst, W.; Schwarzenbach, R. P. Environ. Sci. Technol. Submitted for publication. (15) Hudlicky, M. Chemistry of organic fluorine compounds; Ellis Horwood/PTR Prentice Hall: New York, 1992. (16) Perlinger, J. A. Dissertation, ETH-Zu ¨ rich, 1994.
(17) Pospisil, J.; Luzny, Z. Collect. Czech. Chem. Commun. 1960, 25, 589-592. (18) Zhang, J.-Z.; Millero, F. J. In Environmental Geochemistry of Sulfide Oxidation; Alpers, C. N., Blowes, D. W. Eds.; American Chemical Society: Washington, DC, 1994. (19) Ashworth, R. A.; Howe, G. B.; Mullins, M. E.; Rogers, T. N. J. Hazard Mater. 1988, 18, 25-36. (20) Gossett, J. M. Environ. Sci. Technol. 1987, 21, 202-208. (21) Hine, J.; Mookerjee, P. K. J. Org. Chem. 1975, 40, 292-298. (22) Mackay, D.; Shiu, W. Y. J. Phys. Chem. Ref. Data 1981, 10, 11751199. (23) American Public Health Association. Standard methods for the examination of water and wastewater; APHA: Washington, DC, 1980. (24) Westall, J.; Zachary, J. L.; Morel, F. MINEQL, a computer program for the calculation of chemical equilibrium composition of aqueous systems; Technical Report 18; Ralph Parsons Lab, MIT: Cambridge, MA, 1976. (25) Williamson, M. A.; Rimstidt, J. D. Geochim. Cosmochim. Acta 1992, 56, 3867-3880. (26) Schwarzenbach, G.; Fisher, A. Helv. Chim. Acta 1960, 18, 13651390. (27) Millero, F. Mar. Chem. 1986, 18, 121-147. (28) Schoonen, M. A. A.; Barnes, H. L. Geochim. Cosmochim. Acta 1988, 52, 649-654. (29) Roberts, A. L.; Gschwend, P. M. Environ. Sci. Technol. 1991, 25, 76-86. (30) Zhang, J.-Z.; Millero, F. J. Geochim. Cosmochim. Acta 1993, 57, 1705-1718.
(31) March, J. Advanced Organic Chemistry; Wiley Interscience: New York, 1992. (32) Assaf-Anid, N.; Hayes, K.; Vogel, T. Environ. Sci. Technol. 1994, 28, 246- 252. (33) Pross, A. Acc. Chem. Res. 1985, 18, 212-219. (34) Michel, C. Acta Chem. Scan. 1992, 45, 695-706. (35) Clark, W. M. Oxidation-reduction potentials of organic systems; Waverly Press: Baltimore, 1960. (36) Eberson, L. Electron Transfer Reactions in Organic Chemistry; Springer-Verlag: Berlin, 1987. (37) Schwarzenbach, R. P.; Stierli, R.; Lanz, K.; Zeyer, J. Environ. Sci. Technol. 1990, 24, 1566-1574. (38) Dunnivant, F. M.; Schwarzenbach, R. P.; Macalady, D. L. Environ. Sci. Technol. 1992, 26 (11), 2133-2141. (39) Luther, G. W.; Giblin, A. E.; Varsolona, R. Limnol. Oceanogr. 1985, 30, 727-736. (40) Urban, N. R.; Ernst, K. H. Org. Geochem. Submitted for publication. (41) Mukherjee, T. Radiat. Phys. Chem. 1987, 29, 455-462.
Received for review October 12, 1995. Revised manuscript received July 22, 1996. Accepted July 26, 1996.X ES950759O X
Abstract published in Advance ACS Abstracts, October 15, 1996.
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