Oxidation of Chlorinated Ethenes by Heat-Activated Persulfate

In some cases, data were collected for more than 2 half-lives, and the later time-points exhibited the ... Figure 1 Pseudo-first-order disappearance o...
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Environ. Sci. Technol. 2007, 41, 1010-1015

Oxidation of Chlorinated Ethenes by Heat-Activated Persulfate: Kinetics and Products RACHEL H. WALDEMER, PAUL G. TRATNYEK,* RICHARD L. JOHNSON, AND JAMES T. NURMI Department of Environmental and Biomolecular Systems, Oregon Health & Science University, 20000 NW Walker Road, Portland, Oregon 97006

In situ chemical oxidation (ISCO) and in situ thermal remediation (ISTR) are applicable to treatment of groundwater contaminated with chlorinated ethenes. ISCO with persulfate (S2O82-) requires activation, and this can be achieved with the heat from ISTR, so there may be advantages to combining these technologies. To explore this possibility, we determined the kinetics and products of chlorinated ethene oxidation with heat-activated persulfate and compared them to the temperature dependence of other degradation pathways. The kinetics of chlorinated ethene disappearance were pseudo-first-order for 1-2 half-lives, and the resulting rate constantssmeasured from 30 to 70 °Csfit the Arrhenius equation, yielding apparent activation energies of 101 ( 4 kJ mol-1 for tetrachloroethene (PCE), 108 ( 3 kJ mol-1 for trichloroethene (TCE), 144 ( 5 kJ mol-1 for cis1,2-dichloroethene (cis-DCE), and 141 ( 2 kJ mol-1 for trans1,2-dichloroethene (trans-DCE). Chlorinated byproducts were observed, but most of the parent material was completely dechlorinated. Arrhenius parameters for hydrolysis and oxidation by persulfate or permanganate were used to calculate rates of chlorinated ethene degradation by these processes over the range of temperatures relevant to ISTR and the range of oxidant concentrations and pH relevant to ISCO.

Introduction In situ chemical oxidation (ISCO) of subsurface contamination is performed by injection of a chemical oxidantsusually hydrogen peroxide, permanganate, ozone, or persulfates sometimes in combination and sometimes with other adjuvants that favor the formation of reactive intermediates such as hydroxyl radical or sulfate radical (1). The variety of ISCO methods makes this approach applicable over a range of settings. For example, hydrogen peroxide (H2O2) premixed with chelated iron can be effective when injected directly into a contaminant source zone, which allows the hydroxyl radicals (•OH) produced by iron-catalyzed degradation of H2O2 to react with the contaminants before they react with other aquifer components (e.g., bicarbonate or natural organic matter (1)). In contrast, contaminants in hard to reach areas may be treated more effectively with permanganate, which is less reactive and more selective than •OH, thereby allowing more time for the oxidant to be delivered to areas with contamination. * Corresponding author phone: (503)748-1023; fax: (503)748-1273; e-mail: [email protected]. 1010

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Persulfate (S2O82-) offers some of the advantages of activated H2O2 and permanganate (2). S2O82- is relatively stable and therefore can be delivered considerable distances in the subsurface (like permanganate), but it can be activated to produce sulfate radicals (SO4•-), which are very reactive with a wide range of contaminants (much like •OH). There are two general ways of activating S2O82-: homolysis of the peroxide bond using heat or light (eq 1) (3) and an oxidationreduction process (analogous to the Fenton reaction) with electron donors, including e- (aq) from radiolysis of water (4) or low-valent metals (Mn+) such as Fe2+ and Ag+ (eq 2) (3, 5). ∆ or hν

S2O82- 98 2SO•4

(1)

2S2O82- + Mn+ f SO•+ Mn+1 4 + SO4

(2)

For in situ applications, recent attention has mostly focused on activation by eq 2 with (chelated) Fe2+ (6-8), although heat activation by eq 1 continues to be used, especially in bench-scale studies (8). Recently, it has been suggested that S2O82- can be activated by H2O2 (8, 9), but this may be due mainly to homolysis of the peroxide bond (as in eq 1) by heat evolved from decomposition of the H2O2 (9). Once S2O82- is activated (eq 1 or 2), the resulting SO4•initiates a chain of reactions involving other radicals and oxidants (eqs 3-8), some of which are potentially reactive intermediates, such as •OH and peroxymonosulfate (HSO5-).

SO4•- + H2O a HO• + H + + SO42-

(3)

2SO4•- f S2O82-

(4)

SO4•- + HO• f HSO5-

(5)

2HO• f H2O2

(6)

H2O2 f H2O + 0.5O2

(7)

H2O2 + S2O82- f 2H + + 2SO42- + O2

(8)

The variety of intermediate oxidants generated in activated S2O82- systems complicates the kinetics of contaminant oxidation. In principle, the contaminant disappearance should be describable with a pseudo-first-order rate constant that is the sum of second-order terms for each oxidant

kobs ) k′′SO4•-[SO4•-] + k′′HO•[OH•] + k′′S2O82-[S2O82-] + k′′other[other] (9) where k′′ represents the second-order rate constants for the reaction of the contaminant with each reactive intermediate. Under most conditions, the dominant term in eq 9 is presumed to be the one involving SO4•- (10-12), but it is not known how the relative significance of these terms varies with system parameters, such as temperature. As mentioned earlier, a potential advantage of S2O82- is its stability before activation, which may allow delivery of the oxidant to contamination in hard to reach places. To take full advantage of this property, the best method of activating S2O82- may be the remote, localized, and directed heating provided by some of the technologies used in ISTR. (For a review of ISTR heating technologies, see ref 13.) This approach to activating S2O82- contrasts favorably with 10.1021/es062237m CCC: $37.00

 2007 American Chemical Society Published on Web 12/15/2006

activation methods that involve mixing reagents before injection (e.g., S2O82- with chelated iron) because the latter inevitably result in reaction of some of the oxidant before it reaches the contaminated zone. In the case of S2O82activation with chelated iron, we expect that some of the SO4•- produced will react with excess Fe2+ (k′′ ) 3-9.9 × 108 M-1 s-1 (14, 15)) thereby lowering the concentration of SO4•that is available to degrade contaminants. This effect may account for the observation that iron activation of S2O82- is not always effective at degrading some contaminants that are degraded with S2O82- activated by other methods (8). Combining ISTR with ISCO using S2O82- may enhance remediation performance in more ways than just thermolysis of S2O82- to SO4•- (eq 1). For example, k′′ for all the terms in eq 9 will increase with temperature, resulting in faster oxidation of contaminants if the oxidant concentrations are not decreased. The latter, however, cannot be assumed because the concentrations of SO4•- and other oxidants shown in eq 9 are the net result of many reactions (eqs 1-8, the reactions of each oxidant with the contaminants and their byproducts, etc.), each of which will have its own dependence of rate on temperature. An additional uncertainty is the effect of temperature on the relative rates of contaminant degradation by chemical oxidants other than S2O82- and background degradation processes such as hydrolysis. This study addresses these issues for the chlorinated ethenes by determining the temperature dependence of the kinetics and products of their oxidation by S2O82- and comparing these results to calculations made from previously published data on the degradation of the chlorinated ethenes by permanganate and hydrolysis.

Materials and Methods Materials. Sodium persulfate (Na2S2O8, 98+% purity), tetrachloroethene (PCE, 99.9% purity), cis-1,2-dichloroethene (cis-DCE, 97% purity), and trans-1,2-dichloroethene (transDCE, 98% purity) were obtained from Sigma-Aldrich (St. Louis, MO). Trichloroethene (TCE, 99.9% purity) was obtained from Fluka Chemical (Switzerland), and sodium bicarbonate (NaHCO3, 99+% purity) was obtained from EM Science (Gibbstown, NJ). All chemicals were used without further purification. Experimental. Saturated stock solutions of the chlorinated ethenes were made by allowing the pure nonaqueous phase liquid to equilibrate with deionized, air-saturated water overnight with gentle stirring. Organic cosolvents were never used. Sodium persulfate solutions were made not more than 3 h before use in experiments. Reactions were carried out in 40-mL volatile organic analysis (VOA) vials with Mininert valves (VICI Precision Sampling, Inc., Baton Rouge, LA) incubated without shaking in water baths at 30, 40, 50, 60, and 70 °C ( 1 °C. Most experiments were performed without buffers or ionic strength compensation to avoid potential complications due to reaction between these additives and SO4•-. Although some studies of reactions with activated S2O82- have been done with pH and ionic strength control (e.g., ref 16), the risk of side reactions is great because SO4•- reacts readily with many common inorganic anions (e.g., hydrogen phosphate (k′′) 1.2 × 106 M-1 s-1 (17)), dihydrogen phosphate (k′′ e 7 × 104 M-1 s-1 (17)), bicarbonate (k′′ ) 2.8-9.1 × 106 M-1 s-1 (18, 19)), carbonate (k′′ ) 4.1 × 106 M-1 s-1 (20)), acetate (k′′ ) 3.9-5.0 × 106 M-1 s-1 (18, 21)), and chloride (k′′ ) 2.5 × 108 M-1 s-1 (18))). Furthermore, the risk of significant pH or ionic strength effects on the results reported here is small because the chlorinated ethenes are uncharged, and S2O82- and SO4•are unprotonated even in highly acidic solutions (22, 23). Even product yields and ratios may not be greatly affected by pH, as has been reported for oxidation of alcohols by SO4•- over the pH range of 1.5-11 (24).

For most kinetic experiments with PCE, 1.3 mL of the saturated PCE stock was added to 40 mL of unbuffered deionized water in a VOA vial with a Mininert valve. After the solution was heated to the desired temperature ((1 °C), 1 mL of 0.019 M Na2S2O8 was added, and the syringe plunger was flushed ∼3 times to mix the contents of the vial. At this point, there was little to no headspace in the vials ( TCE > trans-DCE > cis-DCE and correlates with ln A and EA (shown in Figure 3, for 30 °C, but the same holds at the other temperatures). Additionally, ln A, EA, and kobs all correlate fairly well with the energy of the highest occupied molecular orbital (EHOMO) for the chlorinated ethenes. This is to be expected because oxidationsregardless of the exact mechanismsgenerally involves a shift of electron density from the HOMO of the reductant (in this case, the chlorinated ethenes) to an electron acceptor (in this case, SO4•-) (26). These correlations suggest that the rate-determining step for chlorinated ethene oxidation by heat activated persulfate involves reaction with SO4•- and is not, for example, the formation of SO4•- (i.e., eq 1). Our value of EA for TCE is within 10% of the value reported previously (97.7 kJ mol-1) by Liang et al. (16). The only other data on EA for oxidation of chlorinated ethenes by heatactivated S2O82- are from Huang et al. (27). These values (PCE: 46.4, TCE: 60.7, cis-DCE: 41.4, trans-DCE: 49.8, 1,1dichloroethene: 64.9, and vinyl chloride: 56.9 kJ mol-1) are significantly lower (2-3 fold) than the values reported here (and by Liang et al. (16)), and the two sets of data do not correlate well (not shown). A likely explanation for this disagreement lies in the fact that Huang et al. measured degradation kinetics in a mixture of 59 contaminants (27). In such a situation, all of the contaminantssand possibly their degradation intermediates and productsscan compete for SO4•- (and any other reactive species). If the various

FIGURE 4. Effect of bicarbonate on the oxidation of PCE by persulfate. Experimental conditions: 30 °C, 4.5 × 10-5 M PCE, 4.5 × 10-4 M Na2S2O8. Replicate analysis are shown for each sample. The control did not contain persulfate or bicarbonate. reactions have different activation energies, some contaminants will have proportionally higher reaction rates at the same temperature than others, and the contaminants with higher activation energies will become more effective competitors at higher temperatures. This competition effect could explain why the apparent activation energies for the chlorinated ethenes are lower in pseudo-first-order reactions with many other species present than when they are the only reductant in solution. Effect of Bicarbonate. SO4•- oxidizes bicarbonate (HCO3-) to carbonate radical (CO3•-) according to eq 11 with reported rate constants for k′′) 2.8 × 106 M-1 s-1 (18) and 9.1 × 106 M-1 s-1 (19).

SO4•- + HCO3- f SO42- + CO3•- + H +

(11)

The equivalent reaction for carbonate (CO32-) also proceeds rapidly (k′′ ) 4.1 × 106 M-1 s-1 (20)), and the pKa of CO3•< 0 (28), so CO3•- will be produced regardless whether the reaction is initiated with HCO3- or CO32-. These reactions are of concern because (bi)carbonate concentrations that are typical of groundwater (50-400 mg/L or 0.82-6.6 M (29)) could result in suppressed concentrations of SO4•-, and therefore slower degradation of chlorinated ethenes, during ISCO with activated S2O82-. This problem is analogous to the effect of (bi)carbonate on •OH when Fenton-like reactions are used for ISCO (1). To determine the degree that environmental concentrations of carbonate species affect the kinetics of oxidation by heat-activated S2O82-, PCE was used as the model compound, along with 100, 300, and 500 mg L-1 of bicarbonate at 30, 50, and 70 °C. The data were fit to pseudo-first-order kineticss as shown in Figure 4 for 30 °Csand the resulting values of kobs are given in the Supporting Information (Table S2). At all temperatures, bicarbonate inhibited the oxidation of PCE by activated S2O82- (Figure S2). The inhibitory effect of bicarbonate appears to be steeper than a simple exponential relationship and equivalent at all three temperatures, but we did not attempt to model this behavior. The inhibitory effect of bicarbonate probably reflects competition between the PCE and bicarbonate for SO4•-, but changes in pH could also be contributing. For all the experiments shown in Figure S2, the final pH decreaseds possibly due to the release of protons when SO4•- reacts with water (eq 3)sand the degree of pH decrease was less

pronounced with increased bicarbonate concentration (Figure 4). However, even with bicarbonate present, the oxidation of PCE still follows the Arrhenius model (Supporting Information, Figure S3), and the marginal changes in the Arrhenius parameters are consistent with simple competition between PCE and bicarbonate for SO4•- (Table S2). Products of Oxidation. Little has been reported on the products of oxidation of the chlorinated ethenes by S2O82-, and there is no information on how the reaction pathways and products are affected by temperature. Anticipating some degree of mineralization, we measured Cl- present at the end of experiments performed with the four chlorinated ethenes at 70 °C. For this purpose, we chose to run each experiment for a period anticipated to encompass 4 halflives of contaminant degradation (∼94%), which was estimated from the average of the pseudo-first-order rate constants listed in Table S1. Overall, Cl- recoveries were consistent with complete dechlorination of 80-90% of each chlorinated ethene (data not shown). The only comparable literature data are for TCE, and they include Cl- recoveries of 75% at 40 °C, 80% at 50 °C, and 85% at 60 °C (16). For systems containing SO4•-, however, the interpretation of Cl- data is complicated by the possibility that they react according to eq 12 (k′′ ) 2.47 × 108 M-1 s-1 (18)). Then, the resulting chlorine radical, Cl•, can react with additional Cl-, establishing an equilibrium with Cl2•- according to eq 13 (k′′ ) 8 × 109 M-1 s-1 (18), K ) 1.4 × 105 M-1 (30)).

SO4•- + Cl - f Cl• + SO42-

(12)

Cl• + Cl- H Cl2•-

(13)

Equation 13 suggests that the predominant chlorine radical species at the initiation of the Cl- recovery experiments was Cl•, when the initial Cl- concentration was very low (below detection), and that Cl2•- should have become predominant by the end of each experiment (because degradation of the chlorinated ethenes produced final Cl- concentrations of 2-5 mM, which should shift eq 13 to favor Cl2•-). One consequence of eqs 12 and 13 could be depletion of Cl-, but of greater interest is the possibility that Cl• and Cl2•- react with the chlorinated ethenes or their intermediate degradation products to produce more highly chlorinated products, as has been reported previously for chlorinated phenols (31). The kinetic data necessary to evaluate how favorable chlorination of ethenes might be do not appear to be available, but we did detect trace quantities of hexachloroethane as intermediates in the oxidation of TCE and PCE. The only chlorinated compounds that were identified as major intermediates were cis-DCE during trans-DCE oxidation and trans-DCE during cis-DCE oxidation. This evidence for isomerization is shown in Figure S4 (Supporting Information) for 70 °C, but the effect was seen at all temperatures. Isomerization presumably occurs after formation of a singlebonded intermediate, which might be the radical cation formed by electron-transfer to SO4•- (32), or the SO4•--radical adduct, formed by asymmetric addition of SO4•- to the double bond. Evidence for the latter is strong, with many SO4•-radical adducts having been identified during reactions of SO4•- with alkenes by electron spin resonance (21, 33, 34). It is less clear, however, what factors control the elimination of SO4•- to form the isomerized DCE’s. Comparison of Remediation Technologies at Different Temperatures. To evaluate the potential benefits of coupling ISCO and ISTR for chlorinated ethenes, it is necessary to consider how temperature affects a variety of contaminant degradation processes. For hydrolysis and oxidation by permanganate, it is straightforward to calculate the respective values of kobs vs temperature (up to 100 °C) using eq 10, previously published Arrhenius parameters (35, 36), and VOL. 41, NO. 3, 2007 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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FIGURE 5. Comparison of three degradation processes for PCE as a function of temperature. The hydrolysis band was calculated using Arrhenius parameters obtained from ref 36. The bottom of the permanganate (pink) band corresponds to 100 mg/L (0.63 mM) permanganate and the top of this band corresponds to 40 000 mg/L (253 mM), calculated using Arrhenius parameters obtained from ref 35. The persulfate band (light blue) corresponds to Arrhenius parameters obtained from experiments with 0.45 mM Na2S2O8/0.045 mM PCE. The data points correspond to experimental data for varying concentrations of S2O82- while keeping the PCE concentration at 0.045 mM. The three concentrations represented by the data points are 4.5 mM, 45 mM, and 450 mM Na2S2O8. The dashed line (blue) shows the rate of reaction with PCE at the upper limit of this S2O82concentration range data extrapolated to higher temperatures assuming Arrhenius behavior. appropriate assumptions regarding the dose of permanganate (100-40 000 mg L-1, 0.0007 to 0.28 M (37)) and groundwater pH. For PCE, we show the results of these calculations in Figure 5, with the thickness of the bands reflecting the relevant ranges of permanganate concentration and pH. For the other chlorinated ethenes, we used only a representative permanganate concentration and pH, and the results are shown in Figures S5-S7. Over the whole range of conditions considered, oxidation of the chlorinated ethenes is much faster than hydrolysis, although the opposite would have been the case for some of the more highly chlorinated ethanes. Additional factors must be considered for temperatures over 100 °C (38). For contaminant oxidation by activated S2O82-, it is more difficult to develop a general description of kobs vs T because there are many reactions that can contribute to the concentration of the reactive intermediate SO4•- (eqs 1-8, 11, 12, etc.), each with its own dependence on temperature. (Note that calculating the temperature dependence of contaminant oxidation by activated H2O2 would pose a similar challenge.) To overcome this, we used the EA for PCE that was determined in this study (Table 1) to fix the slope of the temperature dependence and estimated the intercept (ln A) from replicate experiments done at initial concentrations of S2O82- ranging from 107 to 107 000 mg/L (0.00045-0.45 M). This amounts to assuming that the apparent EA is unaffected by S2O82concentration. The result, shown in Figure 5, reveals that the temperature at which the rate of PCE oxidation by heatactivated S2O82- becomes greater than oxidation by permanganate is approximately 40 °C; and, by 100 °C, heatactivated S2O82- oxidizes PCE about 400 times faster than permanganate. In contrast, the other chlorinated ethenes (Figures S5-S7) are so much more reactive with permanganate that oxidation with activated S2O82- is slower even at 100 °C. The relative reactivity of PCE with these two oxidants is consistent with the relative reactivity of the chlorinated ethenes with each oxidant: PCE > TCE > cis-DCE > trans1014

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DCE with activated S2O82- (Figure 2) and the reverse with permanganate (39). While the comparisons facilitated by Figures 5 and S5S7 should be qualitatively accurate, the absolute rates we calculated may differ from those observed in the field for a variety of reasons. Prominent among the factors that suggest slower degradation rates under in situ conditions is the inefficiency of mixing of aqueous-phase oxidants with aqueous- and nonaqueous-phase contaminants because of aquifer heterogeniety (40), limited dispersion (especially vertical mixing (41)), and displacement of nonsorbing contaminants by the injected fluids. Mixing inefficiencies might be ameliorated by coupling ISCO with ISTR, because ISTR should encourage mixing of the oxidant plume with the contamination (by increasing flow within the heated zone due both to reduced water viscosity and increased buoyancy of the heated water). Another reason to expect slower oxidation rates in the field is that “natural oxidant demand” will suppress SO4•- concentrations (42) leaving less SO4•- to oxidize contaminants. Here, again, coupling ISCO with ISTR may be advantageous, because S2O82- will be relatively stable in unheated zones, and heat-activation can be focused on contaminated zones. Additional factors remain to be investigated, such as thermal effects on geochemical processes that might, in turn, influence contaminant oxidation rates.

Acknowledgments This work was supported by the Strategic Environmental Research and Development Program (SERDP) under Projects CU-1289 and ER-1458. This report has not been subject to review by SERDP and therefore does not necessarily reflect their views and no official endorsement should be inferred.

Supporting Information Available Two tables (details on experimental conditions and results) and seven figures (further validation of our results). This material is available free of charge via the Internet at http:// pubs.acs.org.

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Received for review September 19, 2006. Revised manuscript received November 3, 2006. Accepted November 7, 2006. ES062237M

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