Environ. Sci. Technol. 2006, 40, 3501-3507
Experimental Evidence for a Lack of Thermodynamic Control on Hydrogen Concentrations during Anaerobic Degradation of Chlorinated Ethenes AXEL C. HEIMANN* AND RASMUS JAKOBSEN Institute of Environment & Resources, Bygningstorvet, Building 115, Technical University of Denmark, DK-2800 Lyngby, Denmark
Hydrogen (H2) concentrations during reductive dechlorination of cis-dichloroethene (cDCE) and vinyl chloride (VC) were investigated with respect to the influence of parameters entering the Gibbs free energy expression of the reactions. A series of laboratory experiments was conducted employing a mixed, Dehalococcoides-containing enrichment culture capable of complete dechlorination of chlorinated ethenes. The objective was to investigate whether a constant energy gain controls H2 levels in dechlorinating systems, thereby evaluating the applicability of the partial equilibrium approach to microbial dechlorination at contaminated sites. Variations in the temperature between 10 and 30 °C did not affect the H2 concentration in a fashion that suggested thermodynamic control through a constant energy gain. In another set of experiments, H2 levels at constant ionic strength were independent of the chloride concentration between 10 and 110 mmol chloride per liter. These findings demonstrate that the partial equilibrium approach is not directly applicable to the interpretation of reductive degradation of chlorinated ethenes. We also present recalculated thermodynamic properties of aqueous chlorinated ethene species that allow for calculation of in-situ Gibbs free energy of dechlorination reactions at different temperatures.
Introduction In anaerobic environments the energy available from catabolic terminal electron-accepting processes (TEAPs) is often limited to a threshold gain scarcely allowing for microbial growth (1). This minimum metabolically convertible energy yield corresponds to the energy necessary for synthesis of 1/5 to 1/3 of an ATP unit during oxidative phosphorylation and amounts to approximately -10 to -20 kJ/(mol reaction) (2, 3). In such environments, hydrogen (H2) and short-chain acids (either as H2 releasing or direct substrates) act as the currency of energy flow within the microbial community. Anaerobic sediments commonly exhibit aqueous H2 concentrations in the range of nanomoles per liter (4-6), which is fairly low compared to concentrations of other reactants (e.g., electron acceptors, short-chain acids) and implies a very short residence time for H2. Consequently, the metabolic threshold energy gain should be reflected in a thermodynamic constraint on the H2 concentration resulting from this dynamic state of simultaneous production and consumption * Corresponding author phone: +45-4525-2172; fax: +45-45932850; e-mail:
[email protected]. 10.1021/es052320u CCC: $33.50 Published on Web 04/20/2006
2006 American Chemical Society
of H2. The occurrence of this constraint was demonstrated by Hoehler et al. (7) in a series of laboratory experiments investigating H2 concentrations during the TEAPs acetogenesis, methanogenesis, and sulfate reduction. H2 levels responded to changes in factors entering the Gibbs free energy expression (e.g., temperature, electron acceptor concentration) in a way that suggested thermodynamic control. Similarly, Jakobsen et al. (8) showed that free energy calculations could explain observations of different TEAPs, such as sulfate reduction and methanogenesis, in a landfill leachate plume. The partial equilibrium approach (implying TEAPs occurring close to thermodynamic equilibrium while fermentative H2 production is rate-limiting) has been successfully applied to the explanation of redox zonation and competition between different TEAPs in sediments and aquifers (9, 10). Extending this concept to anaerobic biodegradation of groundwater contaminants could provide a powerful tool for investigating aquifers contaminated with chlorinated ethenes, one of the most frequent and hazardous groundwater contaminants in industrialized countries (11, 12). However, presently there is a dearth of data on the potential influence of parameters entering the Gibbs free energy expression (temperature, activities of reactants) on the H2 level during reductive dechlorination of chlorinated ethenes. This might be due to the transient nature of dechlorination in which the concentrations of ethene species are often so low (in the micromolar range) that in most experiments product/educt ratios change too rapidly to allow for establishment of stable H2 levels corresponding to a specific reaction quotient. Also, in field situations competing H2-consuming TEAPs may occur concomitantly. An example of the Gibbs free energy of a dechlorination reaction (∆Gr) is the conversion of cis-dichloroethene (cDCE) to vinyl chloride (VC):
∆Gr ) ∆G0,T + RTln
(
[VC][Cl-][H+] [cis - DCE][H2]
)
(1)
where ∆G0,T is the standard state Gibbs free energy at in-situ temperatures, R is the gas constant, and activities of the reactants are indicated by square brackets. Values of ∆G0,25°C are generally more negative than -100 kJ/(mol reaction) for the dechlorination of chlorinated ethenes (13) leading to a peculiar result when applying the partial equilibrium concept with a threshold energy of -20 kJ/(mol reaction). Assuming a 1:1 ratio of cDCE and VC, a chloride activity of 0.01, and a pH of 7, the resulting aqueous H2 concentration should be around 10-18 nmol/L at 25 °C (Figure 1). This is in the order of one molecule of H2 per cubic meter water and obviously unrealistic. This reasoning prompted Dolfing (13) to postulate that H2 levels in dechlorinating systems appear to be under kinetic control rather than thermodynamic influence. However, the partial equilibrium concept could still be applicable to reductive dechlorination, but in that case with a much higher threshold energy (i.e., more negative ∆G). This higher energy threshold may be rationalized by metabolic dehalogenating efficiencies not having developed into an evolutionary climax, or nonoptimal configuration of existing pathways (14). Apparently, halorespiration involves rather simple respiratory chains which use the theoretically available Gibbs free energy quite inefficiently (15). The actual occurrence of variation in the minimum energy gains for different TEAP was demonstrated by Seitz et al. (16) for acetogenesis, methanogenesis, and sulfate reduction. Recently, it was shown that energy requirements can vary greatly even within VOL. 40, NO. 11, 2006 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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FIGURE 1. Relationship between H2 concentration and Gibbs free energy (∆Gr) for various steps in hydrogenotrophic degradation of chlorinated ethenes (abbreviations are explained in the text). Calculations were performed for a 1:1 ratio of each ethene species pair, a chloride activity of 0.01, and a pH of 7 (T ) 25 °C). The hatched area represents a range of H2 concentrations commonly found in dechlorinating environments (0.1-10 nmol/L). ∆G0 values for the calculations were taken from Table 2. the same microbial species (17), depending on the electron acceptor used for respiration (i.e., chlorophenol vs amorphous Fe(III)oxyhydroxides). If this high energy threshold for dechlorination existed, a compilation of in-situ geochemical data (i.e., activities of reactants, pH) would allow for a straightforward comparison of in-situ free energies with this threshold. However, experimental verification of either theory is yet lacking. In an effort to overcome this lack we designed experiments in a fashion that only one factor within the Gibbs free energy expression is varied at a time, while the other parameters remain either constant or change in a way that is common to all replicates. We studied the H2 levels during reductive dechlorination at (i) different temperatures and (ii) different concentrations of chloride. To reduce the effects of competing TEAPs through simultaneous dechlorination of 3-4 chlorinated ethene species we focused on the reductive dechlorination of cDCE and VC. This seems justifiable considering the similar thermodynamic properties of both compounds (note the close vicinity of lines for cDCEfVC and VCfethene (ETH) conversion in Figure 1). Another advantage of studying cDCE and VC dechlorination is that all known organisms degrading these compounds require H2 as the electron donor (15, 18). Moreover, the frequently observed accumulation of these two intermediates of perchloroethene (PCE) and trichloroethene (TCE) dechlorination (19-23), along with their high toxicity (especially VC), warrants investigations into the degradation of these compounds. To this end, we conducted batch experiments with the mixed dechlorinating culture KB-1 which contains dechlorinators of the genus Dehalococcoides (24, 25). These bacteria are capable of complete, hydrogenotrophic dechlorination of chlorinated ethenes resulting in accumulation of the end product ethene.
Materials and Methods Chemicals. The following chemicals were obtained in liquid form: trichloroethylene (GC grade 99.5+%, Merck) and cisdichloroethene (97%, Acros). Vinyl chloride gas was purchased from Gerling, Holz & Co. (99.97%), and ethene was obtained as pure gas from Mikrolab, Aarhus. DL-Lactic acid sodium salt solution (Fluka; purum, 50% in water), and sodium acetate (Merck, p.a., min. 99%) were used as electron donors. 3502
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Setup of Experimental Batches. Serum bottles (120 mL) were filled with 72-104 mL of sulfate-free mineral salts medium (26) and sealed with 1-cm-thick butyl rubber stoppers and aluminum crimp caps; headspace gas was N2/ CO2 (80/20%). The anaerobic microbial consortium KB-1 dechlorinator (kindly provided by SiREM, Guelph, ON) was used as inoculum. This enrichment culture was originally derived from TCE contaminated soil and groundwater and contains bacteria of the genus Dehalococcoides (24, 25), together with several other organisms (e.g., Geobacter sp., Methanomethylovorans sp.; http://www.siremlab.com). However, H2-consuming processes other than dechlorination were excluded by the growth medium which was free of alternative electron acceptors (e.g., nitrate, iron oxides, sulfate), apart from inorganic carbon. Hydrogenotrophic CO2 reducing TEAPs, such as methanogenesis or acetogenesis, typically require at least an order of magnitude higher H2 levels than those observed in chloroethene-dechlorinating systems (27), and therefore should not be of concern here either. In experiments at different temperatures, lactate was used for culture pre-growth as high-rate H2 releasing substrate at an initial concentration of 4.1 mmol/L. Easily fermentable substrates such as lactate commonly produce a fermentation burst creating high concentrations of H2 (28) which we took advantage of to stimulate dechlorinating activity in the culture. Four replicates were set up at three different temperatures: 10, 20, and 30 °C (total of 12 bottles). Each bottle initially received 0.8 µmol TCE resulting in an aqueous concentration of 7.2 µmol/L. TCE was used as the initial electron acceptor despite our exclusive interest in cDCE and VC dechlorination. This is because the here-employed culture was enriched from TCE-containing environments (see above). However, all H2 concentrations discussed hereafter in terms of thermodynamic control are exclusively taken from periods in which TCE was already completely consumed. The pregrowth phase was initiated by inoculation with 300 µL of KB-1 suspension (1011 Dehalococcoides cells per mL). Bottles were incubated upside down in the dark without agitation. After the initial amount of TCE had been completely dechlorinated to ethene (78 days of pre-growth) the main experiment was started by re-spiking with 6.6 µmol TCE producing an aqueous TCE concentration of 70 µmol/L. pH values were between 6.3 and 6.7. In the second experimental approach, various chloride levels were produced by supplementing the mineral salts medium with different solutions containing a mixture of magnesium chloride (MgCl2‚6H2O, Fluka, purum p.a.) and magnesium bromide (Br2Mg‚6H2O, Fluka, purum p.a.). Mgsalts were chosen in preference to Na-salts to avoid potential Na-dependent membrane effects. Bromide was included to mask potential effects of varying ionic strength resulting from addition of high amounts of chloride. Accordingly, bottles receiving low chloride spikes were supplemented with high amounts of bromide, and vice versa. Using this approach we set up four different media which featured the following concentrations of chloride and bromide (as determined by ion chromatography): chloride 10.2 ((0.0), 43.9 ((0.1), 77.7 ((0.9), and 109.9 ((2.0) mmol/L; bromide 95.9 ((15.7), 60.9 ((0.3), 31.1 ((0.3), and 0 mmol/L. Three replicates were produced for each concentration level and were inoculated with 7% v/v from the 20 °C culture that had previously received lactate and had completely dechlorinated TCE to ethene. For each chloride level, two controls were produced: one without addition of VC and ETH and one without addition of culture (abiotic control). Acetate was added as H2 releasing substrate at an initial concentration of 5.3 mmol/L. For simplicity, in this experiment only VC was used as electron acceptor, putting focus onto the final step of the dechlorination sequence, the formation of ethene. This allowed for
reducing the variation in the Gibbs free energy expression for this reaction by addition of ethene to produce an initial aqueous VC/ethene ratio of approximately 1 (see eq 1). Average initially measured aqueous vinyl chloride and ethene concentrations were 109.7 ((5.6) and 109.1 ((4.4) µmol/L, respectively (12.6 (0.6 and 33.9 (1.4 µmol per bottle, respectively). To check whether variations in ionic strength had any effect on the observed H2 levels we conducted another experiment, this time varying only the bromide concentration. Three different bromide levels (0, 50, and 100 mmol/L) were established by supplementing the mineral salts medium with different amounts of Br2Mg‚6H2O. For each level, triplicate bottles were prepared containing 5 mmol/L acetate and 107 mmol/L VC and were inoculated with 3.5% v/v of the culture. All bottles were incubated at 20 °C, upside down in the dark, without agitation. Analytical Methods. Chlorinated ethenes and ethene were analyzed with a gas chromatograph (Agilent 6890N) equipped with a mass selective detector (MS, Agilent 5973). Acidified aqueous samples (1 mL) in 21 mL GC vials were preheated to 80 °C and introduced to the GC by headspace sampling. Separation was achieved on a 25.0 m × 320 µm × 1.00 µm (nominal) capillary column (J&W GSQ) with helium (class 2) as carrier gas. Chloroform served as internal standard (0.5 mL of a 10 ppmv solution). Aqueous concentrations were converted to amount of substance per bottle using Henry’s law constants at different temperatures for chlorinated ethenes (29) and ethene (30). In the experiments featuring VC as sole electron acceptor, concentrations of VC and ethene were determined by headspace sampling directly from the culture bottles. Headspace gas (100 µL) was introduced into a Shimadzu 14A gas chromatograph equipped with a packed column (3% SP/500 Carbopack B) and a flame ionization detector (FID). Samples for acetate and lactate were filtered through 0.45 µm nylon filters, acidified with 50 µL of 17% H3PO4 per mL of sample, and kept frozen until analysis by suppressed ion chromatography on a Dionex ICE-AS1 9 × 250 mm ion exclusion column (eluent, 4 mM heptafluorobutyric acid; chemical suppression, 10 mM tetrabutylammonium). Chloride and bromide were analyzed by suppressed ion chromatography on a Dionex Ion Pac AS14 4 × 250 mm column with 3.5 mM Na2CO3/1 mM NaHCO3 as eluent. Headspace hydrogen (H2) was analyzed by a reduction gas detector (RGD2, Trace Analytical) as described elsewhere (31). As gaseous samples (sample size 400 µL) were allowed to attain atmospheric pressure prior to injection (via a 250 µL sample loop), measured concentrations of H2 were corrected for overpressure as determined with a portable manometer (Manofix X30D) to account for this loss in mass (e.g., a sample going from 1.1 to 1.0 atm experiences a loss of around 10% in a given volume). Headspace concentrations were converted to aqueous-phase concentrations using tabulated Henry’s law constants at different temperatures (30). Error bars on all graphs represent ( one standard deviation from the average value (triplicate cultures if not indicated otherwise). Thermodynamic Calculations. Calculations of the insitu Gibbs free energy for a given reaction at different 0 temperatures require the standard Gibbs free energy (∆Gf,aq ) 0 and the enthalpy (∆Hf,aq ) of formation of the aqueous chloroethene species. Together with tabulated data of the Gibbs free energy and enthalpy of formation of H2(aq) and H+(aq), the standard Gibbs free energy (∆G0r ) and enthalpy (∆H0r ) of reaction may be calculated for a given dechlorination step
∆G0r )
∑ ∆G
0 f,aq(products)
-
∑ ∆G
0 f,aq(educts)
(2)
∆H0r ) ∆G0r
∑ ∆H
0 f,aq(products)
-
∑ ∆H
0 f,aq(educts)
(3)
∆H0r
and can then be used to calculate the Gibbs free energy at a given temperature (∆G0,T r ) from the van’t Hoff equation (32):
∆G0,T r )
∆H 0r (298.15 - T) + ∆G0r T 298.15
(4)
0 where T is the temperature in Kelvin. We used ∆Gf,aq and 0 ∆Hf,aq values from Stumm and Morgan (33) for inorganic 0 species. ∆Gf,aq values of ethene species can be found in 0 Dolfing and Janssen (34) and Haas and Shock (35). ∆Hf,aq values for chloroethenes appear in Haas and Shock (35). However, due to an inaccurate value of the Henry’s law 0 constant for VC used by Haas and Shock (35) the ∆Hf,aq value is dubious (see Supporting Information). Also, this reference lacks thermodynamic data for ethene. Therefore, we recalculated thermodynamic properties of aqueous ethene species by using the standard Gibbs free energies and enthalpies of formation for the gaseous state together with Henry’s law constants for each compound. Henry’s law constants need to be expressed as
Kgfaq )
ai fi
(5)
where Kgfaq represents the equilibrium constant for the gas solubility reaction (e.g., VCgTVCaq), while ai and fi are the activity of the aqueous species and the fugacity of the gaseous species, respectively. The standard Gibbs free energy for this 0 gas solubility reaction (∆Ggfaq ) is then found through 0 ∆Ggfaq ) -RTlnKgfaq
(6)
and can be used to calculate the standard Gibbs free energy 0 of formation of the aqueous species (∆Gf,aq ) 0 0 0 ∆Gf,aq ) ∆Ggfaq + ∆Gf,g
(7)
0 where ∆Gf,g is the standard Gibbs free energy of formation of the gaseous species. Standard enthalpies of formation of 0 the aqueous species ∆Hf,aq can be calculated analogously from the standard enthalpy of the gas solubility reaction 0 (∆Hgfaq ) and the standard enthalpy of formation of the 0 0 gaseous species (∆Hf,g ), where ∆Hgfaq is the intercept from 0 the linear regression of ∆Ggfaq (derived from Kgfaq at different temperatures using eq 6) vs T (K) within a narrow interval around T ) 298.15 K (35). This is based on the relationship between the standard Gibbs free energy (∆G0), enthalpy (∆H0), and entropy (∆S0) of a system at constant temperature (∆H0 and ∆S0 are assumed to be constant in the narrow temperature interval):
∆G0 ) ∆H0 - T∆S0
(8)
In the calculations described above we used data from 0 0 the compilation of Yaws (36) (for ∆Gf,g and ∆Hf,g , both at T ) 298 K), Staudinger and Roberts (29) (Kgfaq of chlorinated ethenes at different temperatures), and Wilhelm et al. (30) 0 (Kgfaq of ethene at different temperatures). ∆Ggfaq values were calculated for 4, 10, 15, 20, 25, 30, and 40 °C for chlorinated ethenes, and for the same temperatures except 0 4 °C for ethene. Linear regressions of ∆Ggfaq vs T (K) had R 2 values of at least 0.999 for all species. The complete calculation is available as Supporting Information to this article.
Results and Discussion Thermodynamic Properties of Aqueous Ethene Species. Table 1 shows the recalculated standard Gibbs free energy VOL. 40, NO. 11, 2006 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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TABLE 1. Thermodynamic Properties of Ethene and Chlorinated Ethene Species ∆Gf0,g ∆Gg0 faq ∆Gf0,aq ∆Hf0,g ∆Hg0 faq ∆Hf0,aq speciesa (kJ/mol)b (kJ/mol)c (kJ/mol)d (kJ/mol)b (kJ/mol)e (kJ/mol)f PCE TCE cDCE tDCE 1,1-DCE VC ETH
22.6 16.1 19.7 22.0 25.4 42.9 68.1
7.0 5.6 3.6 5.9 8.4 8.0 13.3
29.6 21.7 23.2 27.9 33.8 51.0 81.4
-12.1 -9.6 -2.8 -0.4 2.4 28.5 52.3
-39.9 -38.3 -32.3 -34.4 -32.8 -25.9 -15.6
-52.0 -47.9 -35.1 -34.8 -30.4 2.6 36.7
a Abbreviations: perchloroethylene (PCE), trichloroethylene (TCE), cis-1,2-dichloroethylene (cDCE), trans-1,2-dichloroethylene (tDCE), 1,1dichloroethylene (1,1-DCE), vinyl chloride (VC), and ethene (ETH). b Data from Yaws (36), at T ) 298 K. c Calculated from Henry’s law constants for chlorinated ethenes and ethene given in Staudinger and Roberts (29) and Wilhelm et al. (30), respectively. d Sum of ∆Gg0 faq and ∆Gf0,g. e Intercept from linear regression of ∆G0 gfaq vs T (K) for T ) 277.15313.15 K (chlorinated ethenes), and T ) 283.15-313.15 K for ethene (based on data in Staudinger and Roberts (29) and Wilhelm et al. (30)). 0 f Sum of ∆H0 gfaq and ∆Hf,g.
TABLE 2. Thermodynamic Properties of Different Reactions in the Dechlorination Succession of PCE to ETH, and ∆Gr for In Situ Conditions in a Hypothetical, Contaminated Groundwater reactiona
∆G0 (kJ/mol)b
∆H0 (kJ/mol)b
∆Gr (kJ/mol)c
PCE + H2 f TCE + H+ + ClTCE + H2 f cDCE + H+ + ClcDCE + H2 f VC + H+ + ClVC + H2 f ETH + H+ + Cl-
-156.8 -147.4 -121.1 -118.4
-158.9 -150.2 -125.3 -128.9
-158.5 -149.1 -122.9 -120.2
a All reactants, including H , VC, and ETH, as aqueous species. 2 Calculated from Gibbs free energies of formation and enthalpies of formation from Table 1 (ethene species) and from Stumm and Morgan (33) (inorganic species). All reactants as aqueous species, including H2, VC, and ETH, and at unity activities. c Calculated according to the relationship ∆Gr ) RTln∏i[i]υi and assuming the following aqueous activities: PCE, TCE, cDCE, VC, and ETH 10-4; H2 2 × 10-9; H+ 10-7; Cl10-2; and T ) 25 °C. b
and enthalpy values of the aqueous ethene species. For the sake of completeness we also included the higher chlorinated ethenes PCE and TCE, and all three DCE isomers (the cisisomer is commonly the predominant metabolic isomer in microbial reductive chloroethene dechlorination (12)). From these values we derived Gibbs free energies and enthalpies of reaction for various dechlorination steps (Table 2). Please note that, for the sake of uniformity, all compounds appearing in this table are regarded as aqueous species, including H2, VC, and ETH (performing the same calculations with gaseous H2 would change, e.g., the ∆G0 value by approximately 17.6 kJ/mol). As already discussed above, the in-situ free energy (∆Gr) at H2 levels indicative of reductive dechlorination of chloroethenes (e.g., 2 nmol/L; (37)) is quite negative. This is not new, however, the recalculation of these data led to a significant difference when we calculated temperaturedependent H2 concentrations controlled by a constant free energy yield (discussed in a later section; Figure 3). While the conversion of VC to ETH exhibits an increasing H2 level with increasing temperature according to our data, the slope is negative using data in Haas and Shock (35) (data not shown). Dechlorination at Different Temperatures. As expected, pre-growth of the dechlorinating culture on lactate and TCE yielded rapid dechlorination and transiently high aqueous H2 concentrations well above 1 µmol/L within the first 15 days of incubation at all temperatures (data not shown). This level is about 3-4 orders of magnitude above what is presently considered to be a threshold hydrogen level for halorespi3504
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FIGURE 2. Sequential dechlorination and H2 concentrations (hatched area) at 30 °C after re-addition of TCE (data points represent averages from four replicate cultures). Curves were smoothed using the cubic B-spline connection in Microcal Origin 6.0 (Microcal Software, Inc., Northampton, MA). ration of chlorinated ethenes (37-39). However, lactate was depleted after 26 days, and around 5 mmol/L of acetate was produced in all bottles. The fermentation burst leveled off as H2 was consumed and aqueous concentrations decreased to below 30 nmol/L H2 after 15, 20, and 44 days, at 30, 20, and 10 °C, respectively (data not shown). Re-addition of TCE after pre-growth again resulted in complete, sequential dechlorination of TCE via cDCE and VC to ethene, as exemplified for 30 °C in Figure 2. At all temperatures, H2 concentrations dropped to below 1 nmol/L upon re-addition of TCE (Day 0), until all chlorinated species were consumed (data not shown). After complete dechlorination of all chlorinated compounds we observed a characteristic rise of H2 concentrations back to approximately the initial levels (as exemplified in Figure 2). This observation is consistent with findings of others (37, 40) and clearly demonstrates the presence of H2 donating substrates throughout the experimental period. This could be either (i) acetate, that is oxidized to CO2 and H2 at low hydrogen levels (31, 40), or (ii) biomass that is fermented to H2 (41, 42). Difficulties may arise from comparing different temperature levels in terms of H2 concentrations during dechlorination. This is mainly due to the fact that dechlorination rates are temperature-dependent (43) which may result in (i) different dechlorination phases occurring at different times, and (ii) different product/educt ratios occurring at the same time. Both may result in invalid comparisons with respect to dynamic or steady-state H2 levels. Therefore, two different approaches were used to compare values that best reflect a steady-state system. In the first approach, we opted for comparing the minimum H2 concentration at the different temperature levels (during cDCE and VC dechlorination) with the theoretical temperature-dependent H2 range for cDCE and VC dechlorination, assuming thermodynamic control and a constant energy yield. If there is a H2-temperature relationship based on a constant free energy yield this should be reflected in a temperature-dependent change in the minimum H2 levels. To this end, we determined the minimum H2 concentration for each replicate once TCE was completely depleted. Of the four values at each temperature level (four replicates) the highest number was eliminated to minimize the influence of extreme values at the upper end. The remaining three values were averaged. The results are shown in Figure 3 along with the theoretical temperature-H2 relationship for both cDCE and VC dechlorination assuming a constant ∆Gr. In the second approach, a different subset of H2 concentrations was considered to account for varying product/educt ratios at each minimum value. For the same
FIGURE 3. Average minimum H2 concentrations measured during dechlorination of cDCE and VC (closed symbols, dashed line), H2 concentrations at an approximate VC/ETH ratio of 2:1 (open symbols, dotted line), and the theoretical temperature-H2 relationship assuming a constant ∆Gr at all temperatures (solid lines). Note that experimental H2 concentrations for both approaches are identical at 20 °C. Experimental conditions entering the calculations are pH 6.3, chloride activity 0.012, and a 2:1 ratio of the educt/product ethene species. Constant Gibbs free energies were chosen to match the experimental data at 10 °C (lines converging at 10 °C) and were ∆GcDCEfVC ) -115.0 kJ/mol reaction, ∆GVCfETH ) -112.6 kJ/mol. Curves were smoothed as in Figure 2. replicates, H2 concentrations occurring at an approximate VC/ETH ratio of 2:1 (average VC/ETH concentration quotient of 2.2 ( 0.8, n ) 9) were determined. These values were averaged for each temperature and are also shown in Figure 3. The experimental data evidently do not follow the theoretical relationship very closely for either of the two dechlorination reactions. The temperature-dependent H2 variation predicted by a constant free energy yield is in itself fairly small, e.g., for cDCE degradation the variation is only between 0.2 and 0.28 nmol/L (between 10 and 30 °C). Increasing temperatures might stimulate hydrogen production from the organic substrate to a larger degree relative to hydrogen consumption by dechlorinators, potentially resulting in compensation concentrations that increase with temperature (44). Westermann (45) and Hoehler et al. (7) found that steady-state hydrogen concentrations decreased with decreasing temperature, both reflecting energetic requirements of the syntrophic H2 producing step and the terminal H2 consuming reaction: while, e.g., syntrophic acetate oxidation to CO2 and H2 yields less energy with decreasing temperature (more positive Gibbs free energy), the energy yield for H2 dependent methanogenesis increases (more negative Gibbs free energy) as the temperature decreases (2). Similarly, Chin and Conrad (46) showed that H2 concentrations in methanogenic paddy soil decreased upon a temperature shift from 30 to 15 °C. Consequently, the temperature-dependent increase in H2 concentrations in our experiment (Figure 3) may reflect energetic requirements of the H2 producing population. This potential effect, together with the small theoretical temperature-dependence of H2 levels during dechlorination, renders the temperature approach insufficient to conclusively rule out the possibility of thermodynamic control on H2 levels. Therefore, there is a need for further evidence, which was sought through a variation of the chloride activity. Dechlorination at Different Chloride Levels. Hydrogen levels during VC dechlorination are shown in Figure 4 along with the theoretical chloride-H2 dependence in case of thermodynamic control. Average H2 concentrations were around 1.3-1.8 nmol/L at all chloride levels while the H2
FIGURE 4. Aqueous concentrations of H2 in mineral salts medium containing different chloride levels (closed circles; data points represent averages from triplicates) and calculated chloridedependent H2 concentrations (solid line) assuming a constant free energy gain of -115.8 kJ/(mol reaction). This ∆Gr was chosen to match the experimental data at 10.2 mmol/L chloride; activities entering the equation: VC, ETH 10-4; H+ 10-6.3.
FIGURE 5. Concentrations of VC, ETH, and H2 in mineral salts medium containing 109.9 mmol/L of chloride (data points represent averages from triplicates). Curves were smoothed as in Figure 2. level predicted by a constant free energy yield should change by a factor of approximately 10 (the chloride activity changes the numerator in the logarithmic expression of the Gibbs free energy formulation approximately 10-fold in the range applied here; see also eqn 1). This observation unequivocally demonstrates the lack of thermodynamic control on the H2 level for this reaction. Dechlorination of VC to ethene proceeded at similar rates at all four applied levels of chloride. As an example the data at the highest chloride level (around 110 mmol/L) are shown in Figure 5. Time series of ethene species in this figure are plotted in µmol per liter (rather than amount of substance per bottle) since the aqueous concentration is of interest for thermodynamic considerations. As Henry’s law constants for both compounds differ, aqueous ethene concentrations do not increase to the same extent as vinyl chloride concentrations decrease. A rather long lag-phase of 40-50 days was followed by rapid VC dechlorination lasting approximately 35-40 days, accompanied by fairly stable and uniform H2 concentrations of around 1.3 nmol/L. For the period in which H2 concentrations stayed fairly constant at all chloride levels (active vinyl chloride dechlorination between day 51 and 79; Figure 5) the H2 levels were averaged for each chloride concentration (data shown in Figure 4). Control cultures that did not receive any VC and ETH displayed H2 levels of at least 1 order of magnitude above VOL. 40, NO. 11, 2006 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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free energy yield. Both dechlorination at different temperatures and at different chloride levels did not display the characteristic relationship between H2 and factors that enter the Gibbs free energy expression, as was shown for other TEAPs. This renders dechlorination reactions of these compounds inapplicable to partial equilibrium calculations that could have been a powerful tool in evaluating the state of contaminated aquifers with these industrial pollutants. Although ionic strength in itself was shown to have an influence, the control on H2 in these environments remains only poorly understood and needs to be investigated further.
Acknowledgments We thank SiREM (Guelph, ON) for providing the KB-1 culture used in this study. We also thank the Danish Research Training Council (FUR) for funding this research. FIGURE 6. Aqueous concentrations of H2 during dechlorination of VC in mineral salts medium containing different bromide concentrations: 0 mmol/L (open squares), 50 mmol/L (closed circles), and 100 mmol/L (open triangles). Curves were smoothed as in Figure 2. this level; abiotic controls showed no dechlorination (not shown). The aqueous concentration of both species was intentionally similar (107 vs 105 µmol/L for VC and ETH, respectively), thus establishing an approximate product/ educt ratio of 1. Clearly, this ratio changed considerably once rapid dechlorination proceeded, however, because of the similar dechlorination rates at all chloride levels the product/ educt ratios also changed in a similar fashion. The 10-fold higher H2 concentrations in nondechlorinating cultures together with the reproducible observation of rising H2 levels upon complete consumption of VC (observed at all chloride levels), demonstrate that there was no shortage of H2 donor in these experiments. In dechlorination, acetate oxidation apparently functioned as a viable H2 source, a process becoming thermodynamically feasible at low H2 levels (31, 40). Decay of biomass may be an alternative explanation for H2 production in these experiments (41). To check whether the ionic strength variation in itself produced any noticeable changes in the H2 concentration we investigated VC dechlorination at different bromide levels (without correction for bromide variation by addition of a different electrolyte, as was done in the experiment previously shown). Dechlorination rates were quite similar at 0 and 50 mmol/L bromide (on the same order as in the experiment with different chloride levels), but decreased to around 1/3 of this value at 100 mmol/L bromide (data not shown). The corresponding H2 levels are shown in Figure 6. An increase of 0 to 100 mmol/L bromide approximately doubled the H2 concentration (0.33 vs 0.67 nmol/L) while bottles at 50 mmol/L bromide showed intermediate values. This demonstrates that ionic strength can affect H2 levels and that this effect should be taken into account when conducting experiments such as the ones presented here. Stable H2 concentrations here are somewhat lower than the values measured during VC dechlorination at different chloride levels (0.3-0.7 vs 1.3-1.8). This could be due to the fact that the bromide experiment was conducted at a later point (after two more transfers of the acetate/VC utilizing culture into fresh medium). Thus, the dechlorinating population may have developed into a more competitive group, or the H2 producing population may have been less efficient, both leading to decreased H2 levels. However, as all bottles in the bromide experiment were inoculated with the same stock culture this should not alter the conclusions drawn here. H2 concentrations during reductive dechlorination of cDCE and VC are apparently not controlled by a constant 3506
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Supporting Information Available Compound and reaction information calculations and graphics. This material is available free of charge via the Internet at http://pubs.acs.org.
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Received for review November 17, 2005. Revised manuscript received March 6, 2006. Accepted March 17, 2006. ES052320U
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