Environ. Sci. Technol. 2000, 34, 804-811
Electrolytic Reduction of Low Molecular Weight Chlorinated Aliphatic Compounds: Structural and Thermodynamic Effects on Process Kinetics ZHIJIE LIU,† ERIC A. BETTERTON,‡ AND R O B E R T G . A R N O L D * ,† Department of Chemical and Environmental Engineering and Department of Atmospheric Sciences, The University of Arizona, Tucson, Arizona 85721
A series of chlorinated low molecular weight alkanes and alkenes was transformed electrolytically using a porous nickel cathode at potentials from -0.3 to -1.4 V (versus standard hydrogen electrode). Kinetics were first-order with respect to the concentration of the halogenated targets. The dependence of the first-order rate constants on cathode potential followed the Butler-Volmer equation, modified to account for mass transfer resistance to reaction. The masstransfer-limited rate constant for reaction of all species was about 1.55 L m-2 min-1. Log-transformed reaction rate constants for reduction of chlorinated alkanes, derived via experiments at the same cathode potential (Ec ) -1.0 or -1.2 V vs SHE), were linearly related to carbonhalogen bond enthalpies, as expected based on a physical model that was developed from transition state theory. The chlorinated ethenes reacted much faster than predicted from bond enthalpy calculations and the alkane-based correlation, suggesting that alkenes are not transformed via the same mechanism as the chlorinated alkanes. Dihaloelimination was the predominant pathway for reduction of vicinal polychlorinated alkanes. For chlorinated alkenes and geminal chlorinated alkanes, sequential hydrogenolysis was the major reaction pathway.
Introduction The reductive dehalogenation of chlorinated solvents by zerovalent metals has been successfully incorporated into funneland-gate systems for groundwater protection and treatment (1, 2). Although questions remain regarding the mechanisms of dehalogenation and kinetics are difficult to anticipate due to the variable conditions of metal/metal oxide surfaces, the technology is fit for field application. Incentives for its application include the low cost of suitable metals (primarily iron) and the passive nature of the treatment system once the barrier/gate has been installed. A number of investigators have examined the electrolytic reduction of chlorinated solvents in order to overcome certain perceived shortcomings of zero-valent metal, nonelectrolytic reactions (3-7). Electrolysis of halogenated solvents such as trichloroethene (TCE) and carbon tetrachloride (CT), for * Corresponding author telephone: (520)621-2410; fax: (520)6216048; e-mail:
[email protected]. † Department of Chemical and Environmental Engineering. ‡ Department of Atmospheric Sciences. 804
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example, can result in much higher area-specific rates of transformation than those observed in reactions with elemental metals (5-7). Process kinetics are a function of cathode potential, at least until mass transfer limitations are encountered, and therefore subject to engineering control. Deterioration of process kinetics due to metal corrosion and surface passivation can be avoided, and potentially toxic metal ions are not released, thus expanding the range of metals that may serve as cathode materials. The electrochemical reduction of chlorinated aliphatic compounds has been studied for decades, although seldom in aqueous systems. Most investigations involved organic solvents such as acetonitrile and N,N-dimethylformamide (DMF) and/or mercury-drop electrodes (8-11). Interest in applying electrolysis for the destruction of chlorinated solvents in water is relatively new. Previous studies concentrated on process kinetics, electrode selectivity, and product distribution as functions of cathode material and potential, target identity, bulk solution pH, and ionic strength (3, 5, 12). Understandably, the a priori prediction of contaminant reaction rates is of paramount importance to the design of electrolytic treatment systems. The dependence of heterogeneous reaction rate constants for CT reduction on electrode potential follows Butler-Volmer kinetics, modified to account for mass transfer limitation at the electrode surface (5). While the dependence of rate on electrode potential is of some utility, the relationship between reaction rate and characteristics of target compounds (structure, degree of halogenation, reduction potential) is imperfectly understood despite at least two successful attempts to find correlations between reduction kinetics and compound-specific parameters in nonelectrolytic situations. Peijnenburg et al. (13, 14) developed quantitative structure-activity relationships on the basis of pseudo-first-order rates of disappearance for 45 halogenated aromatic hydrocarbons in anoxic sediments. Independent variables included such molecular descriptors as carbon-halogen bond strength, the summed Hammett σ constants for additional substituents, the summed inductive constants for these substituents, and certain steric factors. Scherer et al. (15) developed satisfactory linear free energy relationships (LFERs) using estimated energies for lowest unoccupied molecular orbitals, calculated for a series of chlorinated alkanes and alkenes, to predict rate constants for reaction with a single metal, elemental iron. Correlation between log-transformed rate constants and one-electron reduction potentials was also excellent for the same series of compounds, implying that rate limitation arises from an initial electron transfer (i.e., RX + e- f R• + X-), as suggested previously (15). Following these examples, we developed a LFER between potential-dependent rate constants for the electrolytic dehalogenation of randomly selected, short-chain chlorinated aliphatic compounds and their minimum carbon-chlorine bond dissociation energies. Selection of bond dissociation enthalpy as a primary predictor of dehalogenation kinetics was motivated by a physical model developed from transition state theory. Model predictions were tested empirically using data for the electrolytic reduction of alkanes and alkenes at a porous nickel cathode. The chlorinated ethenes reacted much faster than predicted on the basis of a relationship that was derived for the chlorinated alkanes. Such differences have not been reported previously. Findings offer general guidance for electrolytic system design and remediation of groundwater contaminated with halogenated solvents. 10.1021/es991049b CCC: $19.00
2000 American Chemical Society Published on Web 01/29/2000
Previous work (5) indicated that copper and nickel are superior to a variety of other metals as cathode materials for dehalogenation of CCl4. Since the performance of copper cathodes is comparatively well characterized, we concentrated on nickel cathodes here. The source of metaldependent differences in observed dehalogenation rates is reserved for future studies.
Materials and Methods Electrolytic Cell. The glass reaction vessel (Electrosynthesis Inc.) consisted of two major compartments. The 330-mL cathodic compartment contained ports for gas- and liquidphase sampling, addition of acid and purge gases, a pH probe, and the working electrode. A side compartment was fitted with the reference electrode (single-junction Ag/AgCl with a saturated KCl/AgCl solution, Orion 90-02 with outer sleeve removed). In the anodic compartment, a platinized platinum foil electrode (25 × 21 × 0.1 mm) or a graphite rod (d )10 mm; L ) 50 mm) served as anode. Cathode and anode compartments were separated by a cation-permeable membrane (Nafion 117, Du Pont). Cathode and reference compartments were separated by a sintered glass frit. Cathodes were 100 × 30 × 10 mm porous nickel slabs (20 pores/in., Electrosynthesis) that were soldered to a silver wire. Surface area was 290 ( 20 cm2, per manufacturer’s data. Cathode potential was fixed throughout dehalogenation experiments using a potentiostat (model 410, Electrosynthesis) and a 24-V power supply (model 6201A, Harrison Laboratories). Potentials are reported relative to the standard hydrogen electrode (SHE). The potential difference between the cathode and the anode was normally in the range of 2-15 V. A glass combination electrode (Brinkmann 39843) was used for the continuous measurement of pH in the cathodic compartment, and pH was maintained at 7.0 using a pH-stat (691 pH meter, 614 Impulsomat, 665 Dosimat, Brinkmann) for the automatic addition of 1.0 M HCl. Gas leaks were negligble, as indicated by the development and maintenance of internal pressure during the course of an experiment and by direct measurement. To maintain identical hydrodynamic conditions, electrodes were oriented in the same direction for each experiment, and reactor contents were stirred at the same constant rate. Kinetic Experiments. Porous electrodes were rinsed with dilute HCl and acetone and dried with lens paper immediately before use. For kinetic studies, the cathode compartment was filled to near capacity with 0.1 M Na2SO4 to minimize headspace masses of volatile compounds and the influence of mass transport limitation at the gas-liquid interface on kinetic observations. Initial concentrations of the chlorinated targets were in the range of 0.1-1.5 mM, depending on analytical limits and compound solubilities. To eliminate dissolved oxygen, cathode electrolyte was purged with He (99.995%, U.S. Airweld) for approximately 1 h while the precleaned cathode was protected by maintaining its potential at -0.2 to -0.3 V. After being purged, 5-30 µL of the target compound (solubility dependent) was injected directly into the electrolyte. Contents were mixed continuously using a 50-mm, Teflon-coated magnetic stir bar and a stirring rate of about 150 rpm. During a 30-min period that was provided for dissolution of the target, the cathode potential was not controlled. No detectable reaction product was present at the conclusion of this period, and concentrations of reactants were stable. After the target compound was dissolved and the liquid-phase concentration was steady, the desired cathode potential was set to initiate the experiments. Liquid-phase samples (20 µL) were withdrawn at 1020-min intervals for measurement of chlorinated organics. Samples were extracted in pentane and analyzed by gas chromatography (HP-5790 with ECD; capillary column DB-
624, 30m × 0.53 mm, J & W Scientific; isothermal, 40 °C; He carrier gas, 8 psi; injector, 150 °C; detector; 250 °C). Standards were prepared by injecting known volumes of pure liquid reagents (Aldrich) into water-filled 165-mL vials that were crimp sealed over Teflon-coated rubber septa. Calibration was repeated for each experiment. Analytical limits using the extraction/ECD procedure were in the range of 0.1-10 µM, depending primarily on degree of chlorination. Chloride ion concentrations were measured using a Dionex 100 ion chromatograph with an ion-exchange column (HPIC AS4A) and 20 mM borate elution buffer at 2 mL min-1. Product Identification. In experiments designed to identify volatile organic products, 185 mL of electrolyte was added to the 330-mL cathode compartment. After being purged with He for 1 h, 5-30 µL of pure target compound was injected into the liquid phase and stirred to establish equilibrium between liquid and headspace. When the gasphase concentration of the halogenated target was stable, the desired cathode potential was set. Gas-phase samples (20 µL) containing both the reactant and products were withdrawn at 10-20-min intervals for GC analysis (HP-5890 with FID; capillary column DB-624; isothermal, 30 °C; He carrier gas, 8 psi; injector, 150 °C; detector, 275 °C). Standards were prepared by injecting known volumes of liquid-phase chemicals into 165-mL serum vials that contained electrolyte and nitrogen at the same gas/liquid ratio as the cathode compartment. Methane, ethene, and ethane standards were purchased (Aldrich). Standards for C2H2, C3H8, C4H10, and C2H3Cl were developed by injecting known volumes of the pure gases (99.5+%) into sealed glass vessels. Analytical limits were in the range of 0.1-5.0 µM, again depending on degree of chlorination. Additional gas and liquid samples were periodically withdrawn from the reactor for product identification via GC/MS. The GC (Varian 3400, Varian Assoc., Inc.) was equipped with a capillary column (DB5, 30 m × 0.25 mm i.d., J & W Scientific). Injection port temperature was 150 °C. Ultrapure He was used as the carrier gas (30 psi at injection port). The oven temperature was held at 40 °C for 5 min, ramped (10 °C/min) to 200 °C, and held for 10 min. Compounds were identified using an ion-trap detector (Finnigan MAT 700, Finnigan Co.) and comparison of fullscan spectra with library spectra (NIST Library, Finnigan Co.). Computational. Standard enthalpies for the chlorinated methanes and ethenes and free radicals derived from these compounds were computed at the G2MP2 level of theory. The G2MP2 method uses the modified version of the Gaussian-2 (G2) computational method and Moller-Plesset perturbation theory carried to a second order (MP2) as the basis for set extension corrections. The method is nearly as accurate as the full G2 method at substantially lower computational cost. All calculations were performed with the Gaussian 94 programs (16) using an IBM RS6000 computer at the Center of Computer and Information Technology, The University of Arizona.
Results and Discussion Stoichiometry. Representative data derived for the electrolytic decomposition of perchloroethene (PCE) and TCE on a nickel electrode (Ec ) -1.0 V; Figure 1a,b) indicate that disappearance of the target compound is accompanied by the stoichiometric appearance of chloride ion. Mass balances on total chlorine were generally (5%, suggesting that all major chlorine-containing intermediates were accounted for. In PCE experiments at potentials less negative than -0.8 V, TCE accumulated as a transient intermediate. Dichloroethene (DCE) isomers, although detectable, remained at trace levels. Vinyl chloride never reached the analytical detection limit, estimated at 0.1 µM. Final products of PCE/TCE reduction following the 3-h experiments included primarily ethene, VOL. 34, NO. 5, 2000 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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FIGURE 3. Dependence of overall first-order rate constants (kS) for CT (2), 1,1,2,2-TeCA ([), PCE (b), and TCE (9) reductions on cathode potential using a porous nickel electrode. Replicate experiments were carried out for CT, PCE, and TCE reductions. TeCA experiments were performed one time only. Error bars represent 2 SD in kS values, the measured first-order rate constants for chemical transformations (n ) 3-5). Error bars lie within data symbols for the PCE and TCE experiments.
FIGURE 1. Representative kinetic data for the electrolytic decomposition of (a) PCE and (b) TCE at a porous nickel electrode. Ec ) -1.0V (vs SHE) (b, PCE; 9, TCE; 2, Cl-; (, total chlorine; ], ethane; O, ethene; 4, all DCE isomers). Total chlorine is 4PCE + 3TCE + 2DCE + Cl-.
selected compounds is illustrated in Figure 3. Results were reproducible in repeat experiments conducted when CT, PCE, and TCE were the target compounds. Error bars corresponding to two standard deviations in measured rate constants are shown on the figure. Experiments with other test compounds were not repeated. Reduction kinetics conformed to the Levich equation over the full experimental range of cathode potential. That is
1 1 1 ) + kS kL kc
(1)
where kS is the overall heterogeneous, first-order rate constant (L m-2 min-1; experimentally obtained); kL is the mass transfer coefficient; kc is the rate constant for electrochemical reduction at cathode surface. Since electrochemical transformations of chlorinated targets are irreversible throughout the range of potentials used here, kc should obey ButlerVolmer kinetics, and the backward reaction may be neglected. Equation 1 may be rewritten as
1 1 1 ) + kS kL koc exp(-nRFEc/RT)
FIGURE 2. Log-transformed PCE (b) and TCE (9) concentrations versus time. Lines of best fit to the transformed data (linear regression analysis) were used to estimate compound-dependent kS values. See Table 2. The original (untransformed) PCE and TCE data are from Figure 1, panels a and b, respectively. Porous nickel cathode; Ec ) -1.0V (vs SHE). ethane, butene, butane, hexane, and unidentified hydrocarbons that were probably of greater molecular weight. The sum of the identified carbon-containing products accounted for more than 95% of TCE transformed. Dependence of Kinetics on Potential. Disappearance of the halogenated target compounds was typically first-order with respect to the aqueous-phase concentration, as indicated by log-linear plots of concentration versus time (Figure 2). First-order rate constants were estimated via linear regression analysis on the log-transformed data. Typically, r 2 > 0.98 for 12-15 data points obtained more or less uniformly over the 3-h experiments. Heterogeneous rate constants were normalized using the nominal electrode surface area and liquid volume in the cathode compartment. Dependence of the heterogeneous rate constants on cathode potentials for 806
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(2)
where koc is the area-specific rate constant for electroreduction at Ec ) 0.0 V; n is the number of electrons transferred in the rate-determining reaction step; R is a transfer coefficient (the fraction of electrochemical energy or applied potential that is used to overcome the reaction activation energy); F is the Faraday constant (96 500 °C mol-1); and RT is the Boltzmann temperature (2.48 × 103 J mol-1). When Ec is sufficiently negative, 1/kL . (1/koc) exp(nRFEc/RT) and mass transport limits the observed reaction rate. At less negative potentials, 1/kL may be negligible and Butler-Volmer kinetics is predicted, i.e., reduction rate increases exponentially with decreasing potential. Equation 2 can be rearranged to yield:
(
ln
)
1 1 nRF E ) ln koc + kS kL RT c
(3)
The linearized form offers a method for estimating the parameters nR and koc once kS is measured as a function of Ec. Over the range of potential in which data were obtained, resistance to target transformation due to transport of PCE or TCE to the cathode was modest in relation to the cathode polarization resistance (Figure 3). Therefore, 1/kL was not empirically available for PCE or TCE. For the analysis of PCE
TABLE 2. Rate Constants (kS) for the First-Order Disappearance of Halogenated Aliphatic Target Compounds at a Porous Nickel Electrodea first-order rate constant (L m-2 min-1)b
FIGURE 4. Estimation of koc and nr using the transformed Levich equation and potential-dependent data for the conversion of CT (2), TeCA ([), PCE (b), and TCE (9). Goodness of fit indicates that transformation kinetics are well represented by the Levich equation (eq 1). Parameter estimates are summarized in Table 1.
TABLE 1. Estimates of Intrinsic Kinetic Parameters, nr and koc, Using Potential-Dependent Transformation Data and the Substituted Levich Equation (eq 2)a compd
nr
log koc
CT TCE PCE 1,1,2,2-TeCA
0.27 ( 0.02 0.14 ( 0.01 0.13 ( 0.005 0.20 ( 0.01
-2.96 ( 0.24 -3.66 ( 0.18 -2.99 ( 0.08 -3.34 ( 0.16
a Intervals represent 2 SD in the estimated mean values for nR and log koc. Parameter estimates were obtained via regression analyses using transformed, voltage-dependent transformation rate data. See Figures 4 and 5.
and TCE data, a kL value corresponding to the transportlimited reduction of CCl4 (1.55 L m-2 min-1) was adopted instead. While the validity of the substitution was not verified experimentally, hydrodynamic conditions were the same for all experiments, and the molecular diffusion coefficients for the compounds are similar. The linear relationship between ln(1/kS - 1/kL) and Ec (Figure 4) confirmed that transformation kinetics were in fact governed by the Levich equation and yielded estimates of kinetic parameters (Table 1). A similar set of voltage-dependent experiments provided estimates of nR and koc for the reduction of 1,1,2,2tetrachloroethane (1,1,2,2-TeCA). Assuming that n is invariant among the halogenated target compounds, it is evident that the transfer coefficients for PCE and TCE reduction are significantly lower than those for the reduction of CCl4 and 1,1,2,2-TeCA. These results suggest that the mechanisms for reductions of chlorinated alkanes and alkenes at nickel electrodes differ in some fundamental way. Values of koc for PCE and TCE reduction were 2 and 10 times smaller than koc for CCl4 reduction, accounting for the significantly more negative potentials that were required to achieve measurable rates of PCE/TCE reduction (Figure 3). Reactivity. A total of 15 chlorinated methanes, ethanes, and ethenes were screened to establish kS values for reductive transformation on a nickel cathode at -1.0 and -1.2 V (Table 2). In each case, reduction kinetics were first order with respect to the target compound throughout the experiment. At the highest rates observed, reaction kinetics were apparently limited by mass transfer to the cathode surface. Values of kL so estimated were similar to the mass transfer coefficient for CT transformation (see above). Consequently a single kL value (1.55 L m-2 min-1) was adopted for all test compounds. Compare, for example, rate constants for reductions of perchloroethane (PCA) and TeCA isomers at Ec ) -1.2 V with transport-limited kS values for CCl4 reduction (Figure 3). At a fixed cathode potential, reaction rate constants were
compd
Ec ) -1.0 V
Ec ) -1.2 V
PCE TCE trans-DCE cis-DCE 1,1-DCE PCA 1,1,1,2-TeCA 1,1,2,2-TeCA 1,1,1-TCA 1,1,2-TCA 1,1-DCA 1,2-DCA CT CF DCM
(1.7 ( 0.3) × 10-1 (5.0 ( 1.3) × 10-2 (4.0 ( 0.9) × 10-2 (5.5 ( 1.3) × 10-2 (5.8 ( 1.4) × 10-2 1.35 ( 0.14 1.45 ( 0.10 (7.6 ( 0.9) × 10-1 (4.1 ( 0.4) × 10-1 (2.0 ( 0.2) × 10-1 (1.5 ( 1.0) × 10-2 (3.8 ( 1.1) × 10-3 1.40 ( 0.16 (2.1 ( 0.8) × 10-1 (4.6 ( 2.3) × 10-3
(3.5 ( 0.2) × 10-1 (1.1 ( 0.1) × 10-1 (8.0 ( 1.0) × 10-2 (1.0 ( 0.2) × 10-1 (1.0 ( 0.1) × 10-1 1.53 ( 0.13 1.52 ( 0.12 1.21 ( 0.12 (9.0 ( 0.8) × 10-1 (5.4 ( 1.2) × 10-1 (1.6 ( 0.5) × 10-1 (2.8 ( 0.9) × 10-2 1.52 ( 0.10 0.91 ( 0.10 (3.7 ( 1.1) × 10-2
a Experiments were conducted at E ) -1.0 and -1.2 V (vs SHE). c Estimated mean values and standard deviations were obtained by analyzing the results of multiple identical experiments (n ) 3-5). b Mean ( 1 SD.
directly related to the degree of chlorine substitution of the target compound. For Ec ) -1.0 V, rate constants increased by a factor of 3 from the twice-substituted DCE isomers to PCE and by 2 orders of magnitude or more from the DCA isomers to the TeCA isomers and PCA. Trends in reactivity generally conformed to widely observed relationships between degree of chlorine substitution and facility as an electron acceptor (17). It has been suggested that, for a homologous series such as the chlorinated ethanes, dehalogenation rate is functionally related to the standard reduction potential for halogen removal via hydrogenolysis (17). Furthermore, a roughly linear relationship between reduction potential and dehalogenation rate constant was observed previously for the reduction of chlorinated aliphatic compounds by elemental iron (18). A similar relationship is apparent here among the chlorinated alkanes tested (Figure 5a). Dehalogenation rates for chlorinated ethenes were poorly predicted using the resultant LFER. Because analysis of reaction intermediates (see below) suggested that some of the starting compounds tested were transformed primarily by dihalo-elimination rather than hydrogenolysis, free energies of formation were used to calculate pathway-specific standard reduction potentials, and correlations were built on free energy changes for the primary transformations. Results indicate that rate expectations are improved when separate, pathway-dependent linear free energy relationships (LFERs) are developed for compounds reduced by hydrogenolysis and dihaloelimination pathways (Figure 5b). An analysis based on bond dissociation energies, however, should make this distinction unnecessary inasmuch as rate limitation arises from a oneelectron transfer/C-X bond cleavage in each mechanism. Vogel et al. (17) speculated, with some empirical support, that rate constants for hydrogenolysis reactions are inversely related to carbon-halogen bond energy. The assertion is physically grounded for reaction mechanisms in which disruption of the C-X bond controls the overall rate of the reductive process, as indicated subsequently. LFER Development. Transition state theory rests on a relationship between reaction rate constant and the energy required to reach a postulated transition state. That is
kc ) A exp(-∆Gq/RT)
(4)
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FIGURE 6. Effect of changes in free energies of reactants and products in a homologous series on the free energy of the transition state. G0R, G0P, and Gq are the standard free energies of reactants, products, and transition state for a reference reaction in a homologous series of reactions (e.g., reductive dehalogenations of chloroalkanes). G0R and G0P are changes in free energy (from the reference reaction) for another compound in the homologous series. See eq 8 in the text. FIGURE 5. Relationships between ln(1/kc) and standard reduction potentials for (a) hydrogenolysis of halogenated alkanes and alkenes (RCl + H+ + 2e- f RH + Cl-) and (b) reductive dechlorinations (hydrogenolysis or dihalo-elimination) based on the (empirically) dominant reaction pathway. Experiments were conducted using a porous Ni cathode at -1.0 V. Values were calculated from kS values (Table 2) using kL for CT and eq 1. Primary reaction pathways were selected on the basis of observed reaction intermediates (Table 4). Symbols: O, hydrogenolysis of alkenes; ], hydrogenolysis of alkanes; [, dihalo-elimination of alkanes. Regression lines were derived for the alkane reduction data only.
coefficient, θ, is a dimensionless parameter ranging from 0 to 1 that reflects similarity between the transition state and reaction products. Since ∆Gq ) Gq - G0R and ∆G0 ) G0P - G0R, for our system it can be shown that
δ∆Gq ) θδ∆G0
(9)
Assuming that θ is a constant (independent of reaction within the reaction series), then integration of eq 9 yields
limiting step on the overall reaction pathway and A is the reaction-dependent preexponential factor. When rate limitation arises from the electron transfer step (polarization resistance), the activation energy is functionally related to the cathode potential (19) by
where the constant C1 contains information regarding ∆Gq and ∆G0 for an arbitrary reference reaction. Substituting ∆G0 ) ∆H0 - T∆S0 in eq 10
∆Gq ) ∆Gq0 + RnFEc
∆Gq ) θ∆H0 - θT∆S0 + C1
(5)
∆Gq0
where is the activation energy at an electrode potential of 0.0 V on any convenient electrochemical scale, and Ec is the cathode potential on the same scale. Thus
kc ) A exp(-∆Gq0/RT) exp(-RnFEc/RT)
(6)
For the purpose of subsequent discussion, it is assumed that rate limitation in electrolytic dechlorination arises from cleavage of the carbon-halogen bond. That is, from -
•
RCl + e f R + Cl
-
(7)
We will examine the validity of this assumption later. Leffler (20-22) postulated that a change in the energy of the transition state can be represented as a linear combination of changes in the standard free energies of reactants and reaction products (Figure 6). That is
δGq ) θδG0P + (1 - θ)δG0R
(8)
The equation applies to changes in stability of a homologous series of compounds due to small alterations in chemical structure or medium composition. The subscripts P and R correspond to products and reactants, and the sensitivity 808
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∆Gq ) θ∆G0 + C1
(10)
(11)
and if entropy changes are about the same for all reactions of the form selected (eq 7):
∆Gq ) θ∆H0 + C2
(12)
Since, in this context, ∆H0 is the standard enthalpy change for the reaction in eq 7
∆H0 ) ∆H0f (R•) + ∆H0f (Cl-) - ∆H0f (RCl)
(13)
where ∆H0f (i) is the standard enthalpy of formation of species i. If the carbon-chlorine bond dissociation enthalpy (23) is defined by
D(C-Cl) ) ∆H0f (R•) + ∆H0f (Cl•) - ∆H0f (RCl)
(14)
then ∆H0 differs from D(C-Cl) by the difference in enthalpies of free chloride ion and atomic chlorine radical:
∆H0 ) D(C-Cl) + ∆H0f (Cl-) - ∆H0f (Cl•)
(15)
Combining eqs 15, 12, and 4 yields the following relationship between reaction rate constant and the enthalpy of
TABLE 3. Summary of Standard Heats of Formation for Chlorinated Alkanes and Alkenes and for Free Radicals Derived from Those Compounds via Carbon-Chlorine Bond Dissociationa
species
produced radical
∆H0f (kcal/mol)
abbrev
Methanes -24.6 ( 0.6 (24, 31, 32) CCl3• -24.5 ( 0.2 (24, 31, 32, 35) CHCl2• -22.8 ( 0.2 (24, 31, 32, 35) CH2Cl• -20.1 ( 0.5 (24, 31, 32, 35) CH3• Cl•
CCl4 CHCl3 CH2Cl2 CH3Cl
CT CF DCM CM
CCl2dCCl2 CCl2dCHCl
PCE TCE
-2.9 ( 2.0 (31, 32, 35) -1.4 ( 1.6 (31, 32, 35)
CCl2dCH2 trans-CHCldCHCl cis-CHCldCHCl CHCldCH2
1,1-DCE trans-DCE cis-DCE VC
0.52 ( 0.4 (31, 32, 35) 1.1 ( 2.1 (31, 32, 35) 0.85 ( 2.1 (31, 32, 35) 8.6 ( 0.3 (31, 32, 35)
C2Cl6 C2HCl5
HCA PCA
CHCl2CHCl2 CCl3CH2Cl
1,1,2,2-TeCA -36.4 ( 1.5 (30-32, 35) 1,1,1,2-TeCA -36.4 ( 0.3 (30)
CH2ClCHCl2
1,1,2-TCA
-34.8 ( 1.3 (30-32, 35)
CH3CCl3 CH2ClCH2Cl CH3CHCl2 C2H5Cl
1,1,1-TCA 1,2-DCA 1,1-DCA CA
-34.5 ( 0.2 (30, 35) -31.0 ( 0.8 (30-32, 35) -30.9 ( 0.7 (30-32, 35) -26.5 ( 0.4 (31, 32, 35)
-34.1 ( 1.4 (30-32, 35) -34.1 ( 1.7 (30-32, 35)
∆H0f (kcal/mol) 18.4 ( 1.5 (24, 33, 34, 40, 41) 24.3 ( 1.2 (27, 34) 30.3 ( 1.1 (24, 27, 34) 34.4 ( 0.4 (24, 29, 33) 28.992 ( 0.002 (42)
exp (C-Cl) (kcal/mol)
theor calcd D(C-Cl) (kcal/mol)
72.0 ( 2.1 77.8 ( 1.4 82.1 ( 1.3 83.5 ( 0.9
72.65b 77.54b 81.85b NA
91.0 (39) 93.3 (39)
Ethenes CCl2dCCl• Cl•CdCHCl CCl2dCH• Cl•CdHCl CH2dCCl• CHCldCH• CHCldCH• CH2dCH•
NA NA NA NA >60 ( 0.3 (28) NA NA 70.0 ( 2.0 (36-38)
93.8 (39) 88.7 (39) 88.2 (39) 107.6 ( 2.3
94.52b 93.56b 98.52b 95.40b 93.67b 97.33b 98.98b NA
Ethanes C2Cl5• CHCl2CCl2• CCl3CHCl• CHCl2CHCl• CH2ClCCl2• CCl3CH2• CH2ClCHCl• CHCl2CH2• CH3CCl2• CH2ClCH2• CH3CHCl• C2H5•
8.1 ( 1.9 (26, 44) 5.3 ( 1.9 (26, 45) NA NA NA NA NA NA NA 22.8 ( 2.0 (41) 19.3 ( 2.0 (41) 28.9 ( 0.4 (43)
71.2 ( 3.3 68.4 ( 3.6 NA NA NA NA NA NA NA 82.8 ( 2.8 79.2 ( 2.7 84.4 ( 0.8
68.83c 68.95c 74.08c 74.65c 70.19c 80.84c 76.04c 81.44c 73.60c 82.23c 79.12c 84.13c
a Data are derived from the sources indicated. Experimental enthalpies for carbon-chlorine bond dissociation were derived from enthalpies of formation using eq 13. Standard errors shown were obtained by adding the standard errors of component enthalpies. Theoretically calculated D(C-Cl) values were determined via the procedures indicated. b Calculation was performed at G2MP2 level using Gaussian 94. c Data from ref 25.
carbon-halogen bond cleavage:
kc ) A′ exp[-θD(C-Cl)/RT]
(16)
where A′ ) A exp{-θ[∆H0f (Cl-) - ∆H0f (Cl•)]/RT - C2/RT}. Clearly, everything that contributes to A′ is independent of the reaction selected from the homologous series, and relative reaction rates in the series are functionally related to the C-Cl bond dissociation enthalpy alone. The derivation suggests that the relative magnitudes of rate constants for the reduction of chlorinated targets in a homologous series can be predicted from the enthalpies of C-Cl bond dissociation, a finding that we tested using experimental data for the chlorinated alkanes. Correlation between Rate Constants and CarbonChlorine Bond Dissociation Enthalpy: Empirical Verification. While measurement and/or theoretical prediction of heats corresponding to radical formation from the chlorinated methanes has met with considerable success (24), there is a scarcity of thermochemical data corresponding to the formation of polychloroethyl and polychlorovinyl radicals (Table 3). Fortunately, there are at least four methods for theoretical prediction of C-Cl bond dissociation energies, based primarily on differences in computational detail and computational demands. Here, experimentally determined heats of formation of chlorinated species are considered to be the most reliable source. Where multiple empirically determined values were available, arithmetic mean values were used for the development of correlations. When no such value existed, values were calculated to supplement the database and extend the same correlations (Table 3). The level of uncertainty among computed values is on the order of 1.0 kcal
mol-1 (4.2 kJ mol-1), as indicated by comparison of values computed by alternative means and by differences between theoretical and empirical values when both exist. There are relatively few experimental data for enthalpies of formation of the chlorinated ethenes. Consequently, a series of calculations of D(C-Cl) for chlorinated ethenes was conducted at the G2MP2 level using Gaussian 94 (16). The G2MP2 performance is consistent with the results of some experiments, e.g., TCE and 1,1-DCE dehalogenation, but the errors in other cases, e.g., trans- and cis-DCE, were as high as 11 kcal mol-1 (46 kJ mol-1) (Table 3). Calculated dissociation enthalpies for the chlorinated ethenes should, therefore, be used cautiously. Calculated D(C-Cl) enthalpies for the chlorinated ethanes are from Cioslowski et al. (25). What seems apparent from Table 3 values is that energies necessary for removal of chlorine from an unsaturated carbon atom are much greater than bond energies for chlorine and alkyl carbons. Various explanations have been offered, perhaps the most plausible of which involves π orbital effects (46). In the chlorinated ethenes, halogen substituents attract electrons not only from the carbon-halogen σ bond but also from the π orbital between carbons. The deformed π bond contributes to the relatively high energy of the carbonhalogen bond. The partial double-bond character of carbonhalogen bonds in corresponding vinyl radicals accounts for their characteristically high heats of formation (>250 kJ mol-1). As further evidence, reorganizational energies are typically high (>50 kJ mol-1) for vinyl radicals, whereas reorganizational energies for alkyl groups are low or negative (47). At cathode potentials of -1.0 and -1.2 V, a linear correlation was obtained between log-transformed rate VOL. 34, NO. 5, 2000 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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TABLE 4. Intermediates and Stable Products from the Electroreduction of Chlorinated Ethenes and Ethanesa starting compd PCE TCE trans-DCE cis-DCE 1,1-DCE PCA 1,1,1,2-TeCA 1,1,2,2-TeCA 1,1,1-TCA 1,1,2-TCA 1,1-DCA 1,2-DCA
major intermediates
trace intermediates
TCE
trans-DCE, C2H2 trans-DCE, cis-DCE, 1,1-DCE C2H2
TCE, 1,2-DCA 1,1-DCE, TCE cis-DCE, trans-DCE 1,1-DCA VC
PCE, 1,1-DCA, 1,1-DCE PCE, trans-DCE, cis-DCE PCE, TCE 1,1-DCE, 1,3-dichloro-2-butene 1,1-DCA, 1,1-DCE
major products
trace products
C2H6, C2H4 C2H6, C2H4 C2H6, C2H4 C2H6, C2H4 C2H6, C2H4 C2H6 C2H6 C2H6, C2H4 C2H6 C2H6, C2H4 C2H6 C2H6
propane, C4, C6 propane, C4, C6 C4, C6 C4, C6 C4, C6 C2H4, C4, C6 C4, C6 C4, C6 C2H4, 2-butyne, C4, C6 C4, C6
a Experiments were conducted using a porous nickel electrode at E ) -1.0 or -1.2 V vs SHE. C : 1-butene, trans-2-butene, cis-2-butene, c 4 isobutane. C6: 2-hexene, 3-hexene, trans-4-methyl-2-pentene, cis-methyl-2-pentene.
FIGURE 7. Relationship between ln(1/kc) and enthalpies of C-Cl bond dissociation at (a) Ec ) -1.0 V and (b) Ec ) -1.2 V for the chlorinated alkanes ([) and chlorinated ethenes (O) listed in Table 3. When thermodynamic data were available, D(C-Cl) values were calculated using eq 13. Otherwise values were computed as indicated (Table 3). Regression lines are for the chlorinated alkanes only. constants ln(1/kc) and C-Cl bond dissociation enthalpies for the chlorinated methanes and ethanes tested (Figure 7). The relationship was anticipated on the basis of theoretical considerations (eq 16) and buttresses assumptions related to commonality of mechanism, the nature of rate limitations, etc., for reactions of this class. Sensitivity coefficients (θ) derived from the line of best fit to the Figure 7 data were 0.35 and 0.31 for data corresponding to Ec ) -1.0 and -1.2 V, respectively, suggesting that compound-dependent changes in transition state energies are more closely related to changes in enthalpies of the reactants than those of the radical products. The relationship between rate constants and C-Cl bond strength for chlorinated ethenes is inconclusive, presumably due to uncertainty in experimental energies for C-Cl bond cleavage or the narrow window (∼25 kJ mol-1) of the exhibited 810
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bond strengths. Figure 7 reveals two additional differences between chlorinated ethenes and ethanes: (i) unlike the alkanes, reaction rates for the ethenes were insensitive to the number of chlorine substituents and (ii) kc values for the chlorinated ethenes were orders of magnitude larger than would be predicted based on the relationship between C-Cl bond dissociation enthalpies and kinetic data derived for the chlorinated alkanes. Results indicate that dehalogenations of the chlorinated alkanes and alkenes occur via different reaction mechanisms, as suggested by others (15). Pathways. TCE was a major intermediate in the electroreduction of PCE. The only other chlorinated product detected was trans-DCE, which accounted for less than 1% of the PCE transformed. Both were observed transiently under these experimental conditions. 1,1-DCE, cis-DCE, and vinyl chloride were not detected. Cathode potential was an important determinant of intermediate distribution. For example, TCE concentration peaked at 20% and 10% of the initial PCE concentration at Ec ) -1.0 and -1.2 V, respectively. More negative potentials apparently minimize the accumulation of potentially toxic, chlorine-containing reaction intermediates. Among the final products identified were ethene, ethane, trans- and cis-2-butene, 1-butene, propane, isobutane, and hexenes. Trace amounts of other compounds were detected but not identified. Mass spectrometry results suggest that unidentified products were chlorine-free hydrocarbons containing more than four carbons per molecule (Table 4). The formation of four- or six-carbon molecules from chlorinated ethenes indicates that free radicals are produced during the reaction and that radicals are sufficiently long-lived for polymerization to occur. The distribution of TCE reduction products was more complex. All three dichloroethene isomers were detected, although trans-DCE was the dominant form. Again, vinyl chloride was never observed. The sum of all measurable chlorinated intermediates accounted for less than 0.5% of the TCE transformed. Avoidance of vinyl chloride formation was confirmed via experiments in which DCE isomers were the original target compounds. The only products of DCE reduction on nickel electrodes were chlorine-free hydrocarbons that were similar in identity and distribution to those produced in PCE/TCE experiments. The ratio of ethane to ethene increased over the course of most experiments, suggesting that ethane is produced via ethene reduction. To test this hypothesis, we introduced pure ethene into the reactor headspace in the absence of chlorinated targets. Conversion to ethane was apparent (data not shown). Hydrogenolysis, which is generally held to be a major pathway for the dechlorination of polychloroethenes, was observed here in the conversion of PCE to TCE and TCE to DCE isomers. Alternatively, a concerted two-electron transfer
can displace two chloride ions from polychloroethenes forming an alkyne (48-50). The latter species, however, typically has a more positive redox potential than its halogenated parent and may be rapidly reduced at the same applied potential. That is
which is identical in appearance to the results of successive hydrogenolysis reactions. When pentachloroethane (PCA) was the target compound, TCE was the major intermediate, suggesting that dihaloelimination was the dominant pathway. 1,1-Dichloroethane (1,1-DCA) was generated but to a much lesser extent. PCE and 1,1-DCE were also among the minor products, suggesting that PCA is transformed via dehydrohalogenation as well as hydrogenolysis and dihalo-elimination. There is ample evidence that all intermediates formed were also transformed to chlorine-free products such as ethene, ethane, butene, etc. For tetrachloroethane isomers, dihalo-elimination was again a principal transformation pathway. 1,1-DCE, for example, accounted for 35% of 1,1,1,2-TeCA transformed, and 1,2-DCE was a major intermediate during 1,1,2,2-TeCA reduction. TCE, which was present at low levels during these experiments, indicated that dehydrohalogenation contributed to the transformation of both isomers. Absence of less chlorinated ethanes implied that hydrogenolysis is not a principal pathway for transformation of TeCA isomers. When only geminal chlorine atoms were present in the halogenated target (e.g., 1,1,1-TCA), sequential hydrogenolysis was the expected transformation pathway. Although 1,1DCA was among the 1,1,1-TCA transformation products, chloroethane was not, and dehydrohalogenation was evident from the formation of an alkene product (1,1-DCE). The mechanism of formation of 1,3-dichloro-2-butene and -2butyne from 1,1,1-TCA is not clear. Vinyl chloride was detected only during the reduction of 1,1,2-TCA, presumably via dihalo-elimination. 1,1-DCA and 1,1-DCE were also observed as minor intermediates, indicating that 1,1,2-TCA transformation occurred via several pathways simultaneously. The production of trace levels of TCE as a 1,1,2-TCA transformation intermediate suggests that two hydrogens can be eliminated from vicinal carbons with the formation of a double bond.
Acknowledgments This publication was made possible by Grant P42 ESO 4940 from the National Institute of Environmental Health Sciences, NIH, with funding provided by the Environmental Protection Agency. The manuscript was prepared by Ms. Cecilia Flowers.
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Received for review September 13, 1999. Revised manuscript received November 30, 1999. Accepted December 10, 1999. ES991049B VOL. 34, NO. 5, 2000 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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