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reactivity of HOCl can be related to the increased leaving group ability of OH- (from HOCl) relative to O2- (from OCl-) during reactions in which Cl+ is transferred to an organic compound (Scheme 1). In solutions containing chloride, molecular chlorine (Cl2) may also serve as a chlorinating agent, especially at low pH
Chlorine Monoxide (Cl2O) and Molecular Chlorine (Cl2) as Active Chlorinating Agents in Reaction of Dimethenamid with Aqueous Free Chlorine
HOCl + Cl-+ H+ T Cl2+ H2O
JOHN D. SIVEY, COREY E. MCCULLOUGH, AND A. LYNN ROBERTS* Department of Geography and Environmental Engineering, Johns Hopkins University, 313 Ames Hall, 3400 North Charles Street, Baltimore, Maryland 21218
Received December 22, 2009. Revised manuscript received February 22, 2010. Accepted February 24, 2010.
HOCl is often assumed to represent the active oxidant in solutions of free available chlorine (FAC). We present evidence that Cl2O and Cl2 can play a greater role than HOCl during chlorination of the herbicide dimethenamid. Reaction orders in [FAC] were determined at various solution conditions and ranged from 1.10 ( 0.13 to 1.78 ( 0.22, consistent with the concurrent existence of reactions that appear first-order and second-order in [FAC]. Solution pH, [Cl-], [FAC], and temperature were systematically varied so that the reactivity and activation parameters of each FAC species could be delineated. Modeling of kinetic data afforded calculation of second-order rate constants (units: M-1 s-1) at 25 °C: kCl2O ) (1.37 ( 0.17) × 106, kCl2 ) (1.21 ( 0.06) × 106, and kHOCl ) 0.18 ( 0.10. Under conditions typical of drinking water chlorination, Cl2O is the predominant chlorinating agent of dimethenamid. To the extent that Cl2O represents the active species in reactions with other disinfection byproduct (DBP) precursors, the influence of [FAC] and pH on DBP precursor reaction rates is different than if HOCl were the principal oxidant. Moreover, these findings call into question the validity of apparent rate constants (kapp) commonly reported in the chlorination literature.
Introduction A well-documented consequence of using free available chlorine (FAC) to disinfect drinking water and wastewater is the formation of halogenated disinfection byproducts (DBPs). In generating DBPs, the overall reactivity of FAC depends on the concentrations and chemical characteristics of the individual chlorine species present in solution. The most abundant constituents of FAC are hypochlorous acid (HOCl) and hypochlorite (OCl-) HOCl T OCl- + H +
(2)
K2 ) 2.3 × 103 M-2 (25 °C, ref 3; corrected to 0.0 M ionic strength via the Davies equation). As the concentration of Cl2 is typically low relative to HOCl at circum-neutral pH, the possible importance of this species is often discounted. Some researchers (e.g., refs 4 and 5), however, have observed that Cl2 is far more intrinsically reactive than HOCl. With very few notable exceptions (4, 6-9), the environmental literature is quiet regarding chlorine monoxide (Cl2O) as a chlorinating agent in solutions of FAC. The omission of Cl2O is curious: with a pKB of 6.5, OCl- should be a much better leaving group than OH- (pKB ) 0). We can, therefore, predict Cl2O to be more reactive than HOCl. Indeed, the oxidative potency of Cl2O has been recognized for over 170 years (10). Cl2O can be generated via dehydration of 2 equiv of HOCl 2HOCl T Cl2O + H2O
(3)
Accordingly, [Cl2O] is proportional to [HOCl]2. Roth (11) calculated an equilibrium constant (K3) of 8.70 × 10-3 M-1 at 19 °C for eq 3. Using a reaction enthalpy (∆H3°) of 0.60 kJ/mol (9), we calculate K3 as 8.74 × 10-3 M-1 at 25 °C. As shown in Figure 1, [Cl2O] is low relative to [HOCl] at conditions representative of drinking water (DW) treatment. A recent review of chlorine chemistry surmised that the low concentration of Cl2O renders the contribution of this oxidant negligibly small relative to HOCl (2). Indeed, the vast majority of reports in the environmental literature describing the reactivity of FAC do not mention Cl2O. This suggests that Cl2O is either being discounted due to its low concentration, or that its existence has been forgotten. Nonetheless, several literature reports indicate that Cl2O is a facile chlorinating agent of organic compounds, including anisole (14), p-xylene (4), biphenyl (7), and trans-2-butenoic acid (15) in aqueous solutions of FAC. In each report (4, 7, 14, 15), a second-order dependence of reaction rates on [HOCl] in nominally chloride-free solutions (or extrapolations to [Cl-] ) 0) provided evidence that Cl2O was the primary oxidant. Similar findings also implicated Cl2O as a powerful oxidant of allyl alcohol, with a reactivity far exceeding that of HOCl (16). As with HOCl, Cl2O appears to
SCHEME 1. Postulated Mechanism of Dimethenamid Reaction with Chlorinating Agents of the form XCl, Where X- = HO-, O2-, Cl-, or ClO- for HOCl, OCl-, Cl2, and Cl2O, Respectively
(1)
pK1 ) 7.5 (25 °C, 0.0 M ionic strength, ref 1). HOCl is often assumed to be the primary chlorinating agent of natural organic matter (NOM) and xenobiotic micropollutants during disinfection with FAC (2). The greater * Corresponding author phone: (410)516-4387; fax: (410)516-8996; e-mail:
[email protected]. 10.1021/es9038903
2010 American Chemical Society
Published on Web 03/22/2010
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FIGURE 1. Aqueous free chlorine equilibrium speciation for [FAC] ) 2.8 × 10-5 M (2 mg/L as Cl2, a typical value for DW treatment, ref 12) at 25 °C (solid lines only). Inclusion of NaCl at 3.0 × 10-4 M (the median chloride level in DW sources, ref 13) generates Cl2 (broken line), with no appreciable changes to the other species. All concentrations calculated from eqs 1-3. chlorinate aromatic compounds via electrophilic aromatic substitution (10, 17). Much attention has recently been devoted to FACmediated transformations of xenobiotic micropollutants, including pharmaceuticals, endocrine disruptors, pesticides, and pesticide degradates (e.g., refs 2, 18, and 19). Herein we examine reactions of the herbicide dimethenamid (DM) with FAC. Key independent variables of pH, [Cl-], FAC dose, and temperature were systematically varied to explore their effects on reaction rates. In interpreting our results, we take pains to evaluate several reactive species as potential chlorinating agents of DM, including HOCl, OCl-, Cl2, and Cl2O. DM was selected for study because it was previously shown to react to a single product (chlorodimethenamid, CDM, Scheme 1) upon exposure to FAC in laboratory studies (18). Moreover, at several water utilities in the Midwestern United States, DM levels decreased with concomitant generation of CDM in water disinfected with FAC (20). In both laboratory and field investigations, DM underwent chlorination at the sole unsubstituted carbon on the thiophene ring (18, 20). The presence of a single reactive electrophilic site leading to one chlorinated product as well as the absence of ionizable groups makes DM a useful probe for investigating the reactivity of FAC with aromatic moieties. Moreover, elucidating the reactivity trends of DM with various chlorinating species may be useful in predicting the broader chemical behavior of micropollutants and NOM during water treatment with FAC.
Experimental Section A list of reagents can be found in the Supporting Information (SI). All solutions were prepared in reverse osmosis water further purified by distillation and a Milli-Q (Millipore) system. NaOCl solutions were standardized iodometrically (21). Working solutions of FAC (70 mM) were prepared fresh daily from aqueous NaOCl (6% w/w). FAC concentrations in reactors were determined colorimetrically using the DPD method (21). FAC solutions were nominally chloride-free. Analysis of FAC working solutions via ion chromatography (see the SI for details) did not reveal chloride (detection limit ) 0.01 mM). In solutions of FAC (0.6 mM) prepared in phosphate or acetate buffer (10 mM) with 0.10 M NaNO3, chloride was detected at levels e0.9 mM. This suggests trace chloride contamination in buffers and/or supporting electrolytes 3358
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could represent a source of chloride. Back calculation of [Cl-] from data sets in which DM reaction rates were measured at variable NaCl fortification levels indicate [Cl-] ) 0.3 ( 0.1 mM in reactors to which no NaCl was added (see the SI for more details). Hereafter, we assume this [Cl-] for reactors prepared with no added NaCl. Kinetic Experiments. Reactors were prepared in 40-mL amber glass vials with Teflon-lined caps; vials were rinsed with FAC solution (0.6 mM) and distilled/deionized water before use. Reactors contained pH buffer (10 mM of either acetate (pH < 6), phosphate (6 e pH < 8), or borate (pH > 8)), supporting electrolyte (NaNO3), and DM (typically 10 µM) dissolved in methanol (final methanol concentration 7.5. (We justify our selection of Cl2 rather than H2OCl+ as a reactive FAC species at low pH in the SI, Figure S7). Log kobs profiles as a function of pH were also obtained at five chloride fortification levels (3, 10, 20, 30, and 50 mM). For clarity, only two chloride fortification levels (3 and 30 mM) are shown in Figure 3 (solid symbols). Data at all chloride levels are provided in Figure S8. At pH < 7.5, kobs values increase with increasing [Cl-], consistent with the formation of Cl2. At pH > 7.5, kobs values are not appreciably influenced by [Cl-]. The slope of the log kobs vs pH plots in the pH range of 7.5-8.8 is -1.6 ( 0.3, not the -1.0 predicted if HOCl were the sole reactive FAC species. A slope of -1.6 ( 0.3 implies that another species (namely, Cl2O) is influencing the reactivity of DM in this pH range. Recall from Figure 1 that at pH > 7.5, the log [Cl2O] versus pH plot decreases with a slope of -2, whereas that of log [HOCl] has a slope of -1. At pH > 9, the slope of the log kobs vs pH data approaches -1.0, consistent with the increasing importance of HOCl in this pH range. In light of these insights, attempts to model the entire pH data set (at all chloride levels) were performed using Scientist v3.0 (MicroMath), assuming Cl2O, Cl2 and/or HOCl as the reactive species 3360
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(9)
Best estimates for the second-order rate constants kCl2, kCl2O, and kHOCl were obtained by fitting measured values of kobs and concentrations of Cl2, Cl2O, and HOCl computed for each experiment to eq 9 (see SI for more details). The best-fit second-order rate constants (units: M-1 s-1) were computed as kCl2O ) (1.37 ( 0.17) × 106, kCl2 ) (1.21 ( 0.06) × 106, and kHOCl ) 0.18 ( 0.10. We note that these second-order rate constants are only as robust as the corresponding equilibrium constants used to calculate concentrations of individual FAC species. Specifically, difficulties in obtaining reliable thermodynamic parameters for Cl2O (including K3) have been discussed (9); this topic merits additional research and is addressed further in the SI. The nearly equal reactivity of Cl2O and Cl2 with DM is consistent with the findings of Voudrias and Reinhard (4), who reported statistically equivalent second-order rate constants for reactions with p-xylene (kCl2O ) 54 ( 4 M-1 s-1 and kCl2 ) 58 ( 9 M-1 s-1). Major structural differences between the nucleophilic moiety of DM (a substituted thiophene) and p-xylene (a substituted benzene) likely account for the greater reactivity of DM toward Cl2O and Cl2. No detectable reactivity of HOCl with p-xylene was observed (4). Thus, the reactivity trend reported for p-xylene is consistent with our findings for DM: kHOCl , kCl2 ≈ kCl2O. This trend is consistent with the greater nucleofugality (leaving group ability) of Cl- (from Cl2) and OCl- (from Cl2O) relative to OH- (from HOCl); see Scheme 1. Determination of Activation Parameters. To elucidate the activation parameters of Cl2, Cl2O, and HOCl, variable temperature experiments were performed at three sets of solution conditions. The first set was performed at pH 5.6 with [Cl-] ) 30 mM; under these conditions, Cl2 is responsible for >93% of the total FAC reactivity with DM. The second set was performed at pH 6.2 with no added chloride; under these conditions, Cl2O is responsible for >93% of the total FAC reactivity. The third set was performed at pH 9.7 with no added chloride, such that HOCl accounts for >95% of the total FAC reactivity. The subsequent Eyring plots are shown in Figure 4. The activation enthalpy (∆H+) and entropy (∆S+) can be calculated from the slope and intercept, respectively, of a linear regression of the form ‡
ln
( Tk ) ) - ∆H RT
+
( )
kb ∆S‡ + ln R h
(10)
where k, R, T, kB, and h are a second-order rate constant, universal gas constant, absolute temperature, Boltzmann constant, and Planck’s constant, respectively. For the system in which Cl2 is the predominant chlorinating agent, ∆H+ ) 33.1 ( 4.7 kJ/mol and ∆S+ ) -16.4 ( 1.4 J/(mol K). For the system in which Cl2O is the predominant FAC species, ∆H+ ) 19.7 ( 2.4 kJ/mol and ∆S+ ) -62.2 ( 3.8 J/(mol K). Finally, for HOCl, ∆H+ ) 53.3 ( 6.1 kJ/mol and ∆S+ ) -80 ( 13 J/(mol K). In going from Cl2 to Cl2O, a large decrease in ∆H+ is observed with a concomitant decrease in ∆S+ (i.e., a more negative ∆S+). The latter implies a more ordered transition state structure for Cl2O relative to Cl2. The entropic barrier of HOCl exceeds that of both Cl2 and Cl2O. This may result from H-bonding between solvent water molecules and HOCl in the activated complex. The large ∆H+ of HOCl likely reflects the weak leaving group ability of OH-. To our knowledge, this is the first report of all three sets of Eyring parameters for reactions of active constituents of FAC with an aqueous organic contaminant.
Environmental Significance In 2002 (the most recent year for which data are available), nearly 3.6 million kg of DM were applied in the United States,
FIGURE 4. Eyring plot of DM reactions with FAC under experimental conditions in which Cl2 (pH 5.6, [Cl-] ) 30 mM), Cl2O (pH 6.2, no added Cl-) or HOCl (pH 9.7, no added Cl-) is the predominant oxidant (responsible for >93% of kobs). Error bars denote 95% confidence intervals (smaller than symbols if not shown). Uniform conditions: [DM]o ) 1.0 × 10-5 M; acetate (pH 5.6), phosphate (pH 6.2), or borate (pH 9.7) [buffer] ) 0.01 M; [FAC]o ) 6 × 10-4 M (2 × 10-4 M for experiments at pH 9.7); [NaCl] + [NaNO3] ) 0.1 M. representing a 30% increase in use from 1997 (23). DM has been detected in surface waters (18), including DW sources (24). As DM is poorly removed via coagulation (18), reactions with disinfectants including FAC may influence the fate of this herbicide during DW treatment. Indeed, Hladik et al. detected CDM on exposure of DM to FAC in both laboratory (18) and field studies (20). The potential significance of Cl2O during these processes will be a function of solution chemistry conditions, most notably pH, [Cl-], and [FAC]. Shown in Figure 5 is a profile of the relative contributions of Cl2O, Cl2, and HOCl under conditions representative of treatment with FAC, assuming the kCl2O, kCl2, and kHOCl values reported above. Cl2O is especially important at circum-neutral pH (Figure 5A) and at low [Cl-] (Figure 5B). At [Cl-] levels typical of both DW and wastewater (WW), Cl2O is an important reactive FAC species with DM (Figure 5B).
The relative contributions of FAC species to reactions with DM are also a strong function of [FAC]. The range of [FAC] employed in our systems was 1.5 × 10-5 to 1.5 × 10-3 M (1 - 100 mg/L as Cl2), which spans the FAC ranges commonly used during treatment of DW (ca. 2 mg/L as Cl2, ref 12) and WW (up to 40 mg/L as Cl2, ref 25). As shown in Figure 5C, Cl2O is an important contributor to the reaction of aqueous chlorine at [FAC] typical of DW and WW treatment. Note that uncertainties in K3 do not influence the relative importance of Cl2O (see the SI). Results in Figure 5 were calculated at 25 °C, a temperature somewhat higher than that typically encountered in water treatment systems. As ∆H+ for Cl2O is less than that of HOCl and Cl2, the relative importance of Cl2O will increase at temperatures below 25 °C. One outcome of our studies is to call into question the utility of widely reported apparent rate constants (typically computed as kapp ) kobs/[FAC]o) for reactions of organic compounds with FAC. To the extent that Cl2O (as opposed to HOCl or Cl2) reacts with organic compounds in DW and WW, a nonlinear relationship between kapp and [FAC] is expected (eq 6). In these instances, it is inappropriate to use [FAC] as a normalizing factor when comparing the reactivities of organic compounds in solutions of FAC. To the extent that Cl2O and Cl2 may be important contributors to the reactivity of other organic compounds, our findings also suggest that previously published values of kHOCl should be interpreted with caution. Calculating kHOCl as kobs/[HOCl] may give erroneously high kHOCl values in such cases. For reaction of DM with FAC (0.6 mM) at pH 7.0 and no added chloride, kobs/[HOCl] ) 4.0 M-1 s-1. This value overestimates our more robust calculation of kHOCl by a factor of 22. Even greater discrepancies between the true secondorder rate coefficient kHOCl and an estimate obtained from kobs/[HOCl] would be anticipated at higher [FAC], lower pH, and higher [Cl-]. Particularly noteworthy are the potential ramifications of Cl2O on reaction rates with DBP precursors and hence, on rates of DBP formation (assuming the initial chlorination step is rate-limiting). When Cl2O is a significant reactive FAC species, DBP formation rates will not be linearly related to [FAC]. Moreover, rates of reaction may slow more rapidly with increasing pH than is anticipated from reactions with HOCl (the implicit assumption of kapp models). Indeed, kapp
FIGURE 5. Contributions of Cl2O, Cl2, and HOCl to DM reactivity (as a fraction of kcalc at 25 °C, where kcalc ) kCl2O[Cl2O] + kCl2[Cl2] + kHOCl[HOCl]): (A) as a function of pH under typical DW chlorination conditions ([FAC] ) 2.8 × 10-5 M ) 2 mg/L as Cl2) and [Cl-] ) 3 × 10-4 M (the median chloride level in DW sources, ref 13) and (B) as a function of [Cl-] at pH 7.0. Typical chloride concentrations in DW (interquartile range, ref 13) and WW (range reported by ref 25) are shown. In (C), the effects on kcalc of varying [FAC] (assuming pH 7.0 and [Cl-] ) 3 × 10-4 M) are depicted; typical [FAC] ranges employed during DW (1-3 mg/L as Cl2, ref 12) and WW (2 - 40 mg/L as Cl2, ref 25) treatment are shown. VOL. 44, NO. 9, 2010 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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may over- or underpredict formation rates of DBPs, depending on solution conditions. Our results indicate that the frequent (and often untested) assumption of HOCl as the major chlorinating agent during water treatment processes is not universally valid. Our findings demonstrate that the low concentrations of Cl2O and Cl2 (relative to HOCl) do not preclude their contribution to the overall reactivity of FAC. Indeed, contributions to FAC reactivity with other organic micropollutants have been reported for Cl2 (e.g., refs 5 and 26-28). Moreover, this work joins a very limited number of previous papers in the environmental literature (4, 7, 8) which identify Cl2O as a major oxidant of an organic compound under water treatment conditions. It is conceivable that other electron-rich moieties (including activated aromatic structures in other xenobiotics and NOM) are also susceptible to facile reaction with Cl2O. Whether other organic compounds are substantially more reactive with Cl2O than with HOCl certainly merits further investigation.
Acknowledgments We are indebted to Andrew M. Graham and three anonymous reviewers for their insightful comments on this work. We gratefully acknowledge funding support to C.E.M. from a JHU Provost’s Undergraduate Research Award and to J.D.S. from the American Water Works Association (Abel Wolman and LARS Doctoral Fellowships), ARCS Foundation (Endowment Fellowship), Maryland Water Resources Research Center (Summer Fellowship), and United States Environmental Protection Agency (EPA Science to Achieve Results Graduate Fellowship). The EPA has not officially endorsed this publication, and the views expressed herein may not reflect the views of the EPA.
Supporting Information Available Reagent lists, synthesis procedures, analytical methods, data modeling, and additional experimental results. This material is available free of charge via the Internet at http:// pubs.acs.org.
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(8) Voudrias, E. A.; Reinhard, M. A kinetic model for the halogenation of p-xylene in aqueous HOCl solutions containing Cl- and Br-. Environ. Sci. Technol. 1988, 22, 1056–1062. (9) Reinhard, M.; Redden, G. D.; Voudrias, E. A. The hydrolysis constants of chlorine monoxide and bromine chloride in water. Jolley, R. L., et al. Eds.; In Water Chlorination: Environmental Impact and Health Effects; Ann Arbor Science: Ann Arbor, MI: 1990; pp 859-870. (10) Renard, J. J.; Bolker, H. I. The chemistry of chlorine monoxide (dichlorine monoxide). Chem. Rev. 1976, 76, 487–508. (11) Roth, W. A. Zur Thermochemie des Chlors und der unterchlorigen Sa¨ure. Z. Phys. Chem. Abt. A 1929, 145, 289–297. (12) Montgomery Watson Harza. Water Treatment Principles and Design; John Wiley: Hoboken, NJ, 2005; p 1948. (13) Davis, S. N.; DeWiest, R. J. M. Hydrogeology; Wiley: New York: 1966; p 463. (14) Swain, C. G.; Crist, D. R. Mechanisms of chlorination by hypochlorous acid. The last of chlorinium ion, Cl+. J. Am. Chem. Soc. 1972, 94, 3195–3200. (15) Craw, D. A.; Israel, G. C. The kinetics of chlorohydrin formation. Part III. The reaction between hypochlorous acid and crotonic acid at constant pH. J. Chem. Soc. 1952, 550–553. (16) Israel, G. C.; Martin, J. K.; Soper, F. G. The kinetics of chlorohydrin formation. Part I. The reaction between hypochlorous acid and allyl alcohol in aqueous solution. J. Chem. Soc. 1950, 1282– 1285. (17) Marsh, F. D.; Farnham, W. B.; Sam, D. J.; Smart, B. E. Dichlorine monoxide: A powerful and selective chlorinating reagent. J. Am. Chem. Soc. 1982, 104, 4680–4682. (18) Hladik, M. L.; Roberts, A. L.; Bouwer, E. J. Removal of neutral chloroacetamide herbicide degradates during simulated unit processes for drinking water treatment. Water Res. 2005, 39, 5033–5044. (19) Duirk, S. E.; Collette, T. W. Degradation of chlorpyrifos in aqueous chlorine solutions: Pathways, kinetics, and modeling. Environ. Sci. Technol. 2006, 40, 546–551. (20) Hladik, M. L.; Bouwer, E. J.; Roberts, A. L. Neutral degradates of chloroacetamide herbicides: Occurrence in drinking water and removal during conventional water treatment. Water Res. 2008, 42, 4905–4914. (21) Standard Methods for the Examination of Water and Wastewater, 18th ed.; Greenberg, A. E., Clesceri, L. S., Eaton, A. D., Eds.; American Public Health Association, American Water Works Association, Water Environment Federation: Washington, DC, 1992. (22) Inaba, K.; Doi, T.; Isobe, N.; Yamamoto, T. Formation of bromosubstituted triclosan during chlorination by chlorine in the presence of trace levels of bromide. Water Res. 2006, 40, 2931– 2937. (23) Gianessi, L.; Reigner, N. 2006. Pesticide use in U.S. crop production: 2002. http://www.croplifefoundation.org/Documents/ PUD/NPUD%202002/Fung%20&%20Herb%202002%20Data% 20Report.pdf (accessed Sept 4, 2009). (24) Hladik, M. L.; Bouwer, E. J.; Roberts, A. L. Neutral chloroacetamide herbicide degradates and related compounds in Midwestern United States drinking water sources. Sci. Total Environ. 2008, 390, 155–165. (25) Tchobanoglous, G.; Burton, F. L.; Stensel, H. D. Wastewater Engineering: Treatment and Reuse; McGraw-Hill: Boston: 2003; p 1819. (26) Oyler, A. R.; Liukkonen, R. J.; Lukasewycz, M. T.; Heikkila, K. E.; Cox, D. A.; Carlson, R. M. Chlorine “disinfection” chemistry of aromatic compounds. Polynuclear aromatic hydrocarbons: Rates, products, and mechanisms. Environ. Sci. Technol. 1983, 17, 334–342. (27) Acero, J. L.; Rodriguez, E.; Meriluoto, J. Kinetics of reactions between chlorine and the cyanobacterial toxins microcystins. Water Res. 2005, 39, 1628–1638. (28) Rule, K. L.; Ebbett, V. R.; Vikesland, P. J. Formation of chloroform and chlorinated organics by free-chlorine-mediated oxidation of triclosan. Environ. Sci. Technol. 2005, 39, 3176–3185.
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