Photosensitizer Method to Determine Rate Constants for the Reaction

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Environ. Sci. Technol. 2005, 39, 9182-9188

Photosensitizer Method to Determine Rate Constants for the Reaction of Carbonate Radical with Organic Compounds S I L V I O C A N O N I C A , * ,† T A M A R K O H N , †,§ M A R E K M A C , ‡,| F R A N C I S C O J . R E A L , †,⊥ JAKOB WIRZ,‡ AND URS VON GUNTEN† Swiss Federal Institute of Aquatic Science and Technology (EAWAG), Ueberlandstrasse 133, Postfach 611, CH-8600 Du ¨ bendorf, Switzerland, Departement Chemie, Universita¨t Basel, Klingelbergstrasse 80, CH-4056 Basel, Switzerland

Carbonate radical (CO3•-) is a powerful oxidant that is present in sunlit surface waters and in waters treated by advanced oxidation processes. The production of CO3•in aqueous solution through oxidation of carbonate anion by excited triplet states of aromatic ketones was investigated in this study to provide new methods for the determination of rate constants and to explore a possible photoinduced pathway of CO3•- formation in the aquatic environment. Rate constants for triplet quenching by carbonate anion of up to 3.0 × 107 M-1 s-1 and CO3•- yields approaching unity, determined using laser flash photolysis, allowed us to conclude that such a formation mechanism might be significant in sulit natural waters. Kinetic methods based on either flash photolysis or steady-state irradiation and on the use of aromatic ketones as photosensitizers gave bimolecular rate constants in the range of 4 × 106 to 1 × 108 M-1 s-1 for the reaction of CO3•with several s-triazine and phenylurea herbicides. For various anilines and phenoxide anions, rate constants determined by these methods agreed well with published values. Moreover, it could be shown for the first time by a direct method that dissolved natural organic matter (DOM) reduces the lifetime of CO3•- and a second-order rate constant of (280 ( 90) (mg of C/L)-1 s-1 was obtained for Suwannee River fulvic acid.

Introduction The carbonate radical (CO3•-) is a reactive intermediate that likely plays an important role in aquatic oxidation reactions. While CO3•- has been studied by fast kinetic methods for about four decades (1-11) and its environmental relevance has been recognized (12), only a few papers have been devoted to its aquatic chemistry (13-16). Recently, CO3•* Corresponding author phone: +41-1-823-5453; fax: +41-1-8235210; e-mail: [email protected]. † Swiss Federal Institute of Aquatic Science and Technology (EAWAG). ‡ Universita ¨ t Basel. § Present address Department of Civil and Environmental Engineering, University of California; Berkeley, Berkeley, CA 94720. | On leave from Jagiellonian University, Chemistry Department, Ingardena 3, PL-30060 Krakow, Poland. ⊥ On leave from Universidad de Extremadura, Departamento de Ingenieria Quimica y Energetica, E-06071 Badajoz, Spain. 9182

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has received much attention as a possible intracellular oxidant (17), because it is formed efficiently under physiological conditions by reaction of peroxynitrite with carbon dioxide (18-20). The reaction of hydroxyl radical with the carbonate and bicarbonate anions (3, 21) is considered to be a major source of carbonate radical in aquatic systems (12). Concentrations of CO3•- in sunlit surface waters and in the ozonation of drinking water have been estimated to be about 2 orders of magnitude higher than hydroxyl radical concentrations under the same conditions (22). Thus, oxidation by CO3•- may be more important than oxidation by hydroxyl radical for compounds bearing easily oxidizable moieties, such as anilines or phenols. In contrast to hydroxyl radical, which reacts very rapidly with almost any organic compound, CO3•- is very selective and the corresponding second-order rate constants cover a range of many orders of magnitude. About 180 rate constants for CO3•- reactions with individual chemical compounds are known from a classic compilation (23) and additional ones have been determined more recently (13, 14, 24, 25). Rate constants for many organic contaminants, however, are still missing. In this study, the capacity of several aromatic ketone photosensitizers to form CO3•- was investigated, both to provide a convenient CO3•- source to be used in laboratory experiments and to assess the relevance of such a process in aquatic systems. Kinetic methods using flash photolysis and steady-state irradiation of solutions containing a photosensitizer and carbonate/bicarbonate anions were developed for the determination of reaction rate constants of CO3•with target organic compounds. Finally, scavenging of CO3•by dissolved natural organic matter (DOM) was quantified by laser flash photolysis.

Background Various chemical reactions have been employed to generate CO3•- in aqueous solutions. These include one-electron oxidation of bicarbonate or carbonate by hydroxyl radical (3, 21), the sulfate radical anion (eqs 1a, b) (2, 7, 10), or the excited triplet state of a sensitizer, 3Sens*, (eqs 2a, b) (26, 27). Photolysis of the cobalt(III)-tetraaminocarbonate ion has also been employed (28).

SO4•- + HCO3- f H+ + SO42- + CO3•-

k1a ) 2.8 × 106 M-1 s-1 (1a)

SO4•- + CO32- f SO42- + CO3•-

k1b)4.1×106 M-1 s-1 (1b)

3

Sens* + HCO3- f H+ + Sens•- + CO3•-

(2a)

Sens* + CO32- f Sens•- + CO3•-

(2b)

3

By analogy to the carbonate and bicarbonate anions, CO3•was believed to undergo protonation in the neutral to basic pH range (28, 29). However, a later study showed that no protonation of CO3•- (eq 3) could be observed in the pH range of 0-10 (11), which leaves CO3•- as the only relevant carbonate radical species to be considered in this study and in most aquatic chemistry applications.

HCO3• f CO3•- + H+ 10.1021/es051236b CCC: $30.25

pKa < 0

(3)

 2005 American Chemical Society Published on Web 11/02/2005

FIGURE 1. Relevant reactions involved in the formation of carbonate radical by ketone photosensitization and subsequent reaction with a target compound, Pi. In this study: (A) -R1 ) phenyl, -R2 ) 4-carboxyphenyl (4-carboxybenzophenone); (B) -R1-R2- ) sC(CH3)dC(CH3)sC(dO)sC(CH3)dC(CH3)s (duroquinone). We have chosen to generate carbonate radical mainly by means of an excited triplet photosensitizer (eqs 2a, b), an efficient process which was known for decades (26, 27, 30) but, to our knowledge, was never employed for the determination of reaction rate constants. An important motivation behind this choice is, beside the advantage of having a lightactivated CO3•- production, to extend our knowledge on the oxidative behavior of excited triplet aromatic ketones that have been used as surrogates for excited triplet states of the DOM (31, 32) and shown to induce oxidation of phenols and phenylurea herbicides (33). The main reactions involved in such a system are represented in Figure 1. The aromatic ketone in its ground state is promoted to its excited singlet manifold by absorption of a photon. Rapid and efficient intersystem crossing (typically well within 1 ns) leads to the excited triplet state, which is sufficiently longlived (typically >1 µs) to interact with dissolved compounds. Both excited singlet and triplet states can be deactivated to the ground state. If the one-electron reduction potential of the triplet state of the aromatic ketone is high enough, electron transfer from dissolved carbonate takes place, and this reaction is predominant at high carbonate concentrations. The products of this reaction are the corresponding aromatic ketyl radical anion, which, in aerated solution, is recycled to the ground-state ketone by reaction with oxygen and production of superoxide radical anion, and CO3•-, which is used to study the reaction kinetics with a given target compound, Pi. The dashed arrow indicates the oxidation of the target compound by the excited triplet ketone, a possible side-reaction that has to be minimized. The principle just illustrated can be applied to the determination of rate constants either by using flash photolysis and observing quenching of CO3•- by the target compound, or by employing steady-state irradiation and monitoring depletion of the target compound in the presence of a reference compound (competition kinetics). All these methods are described in detail in the next section. As an alternative, to study the reaction of CO3•- with DOM we employed a method based on eqs 1a, b, using laser flash photolysis of the peroxydisulfate anion to produce CO3•- via the sulfate radical anion (9, 13).

Experimental Section Materials and Solutions. The aromatic ketone sensitizers 4-carboxybenzophenone (CBBP, Aldrich, 99%) and duroquinone (DQ, Fluka, purum) were used as received. Anilines, phenols, and pesticides were obtained from common commercial sources and used without purification (for details on purity, see Section S1 of the Supporting Information). Ethanol, used as a competitor in steady-state irradiation experiments,

was from Fluka (puriss. p.a., g99.8% vol/vol). Solvents used for high-performance liquid chromatography (HPLC) were purchased from Scharlau (Barcelona, Spain) and were used as received. Suwannee River fulvic and humic acid (SRFA and SRHA, respectively; type standard) were purchased from the International Humic Substances Society. Aqueous solutions were prepared using deionized water treated with a Milli-Q water purification device (Millipore). This water was also used as an HPLC eluent. All other chemicals used were analytical grade. Steady-State Irradiations. A DEMA 125 merry-go-round photoreactor (Hans Mangels, Bornheim-Roisdorf, Germany) equipped with a Hanau TQ 718 medium-pressure mercury lamp (maximum power 700 W) and appropriate glass filters and filter solutions (main irradiation wavelength at 366 nm, band-pass filter 308-410 nm) was used. The experimental setup is described in detail elsewhere (31, 34). Under these irradiation conditions, direct phototransformation of most target compounds was negligible within the exposure times required for the kinetic experiments. The aromatic ketones CBBP and DQ, used as photosensitizers, were relatively inert to the ambient laboratory illumination (fluorescent tubes). By limiting exposure to ambient light to the time needed for solution and sample handling (usually less than 1 h) and using brown glass vials for storage, no depletion of target compounds outside the photoreactor could be detected in sample solutions. Temperature was kept at 25.0 ( 0.2 °C. Flash Photolysis. Experiments were performed using a kinetic and spectrographic setup described elsewhere (35), at an ambient temperature of 23 ( 1 °C. Transient absorption measurements were performed by excitation with 248-, 308-, or 351-nm pulses (pulse width ∼25 ns, 100-200 mJ per pulse) from an excimer laser (Lambda-Physik EMG 101 or COMPEX 205). Alternatively, to follow CO3•- relaxation kinetics on the millisecond time scale, a home-built conventional flashphotolysis apparatus, equipped with two parallel flash lamps (quartz tubes of 16-cm length and 1-cm diameter, 6 kPa air pressure, 1 kJ electrical discharge energy, pulse width ∼50 µs) and a cylindrical quartz sample cell of 10.0-cm optical path length, was also used. Analytical Methods. The depletion of the target compound in steady-state irradiation experiments was followed by HPLC using a previously described system (36) and a Nucleosil C18-5 µm, 125 × 4 mm reverse-phase column (Macherey-Nagel, Oensingen, Switzerland). Details of the HPLC methods used for each compound are available in the Supporting Information. Prior to injection, the pH of the samples was corrected, when necessary, to ≈8 by addition of a well-defined small amount of 2 M aqueous hydrochloric acid. Electronic absorption spectra were recorded on a Kontron Instruments Uvikon 930 spectrophotometer, using quartz cuvettes of 1-10 cm optical path length. Measurement of the pH was performed at room temperature by means of a Metrohm pH meter (model 632 or 691) equipped with a combined glass electrode (either Metrohm 6.0204.100 or Orion 8115SC) that was calibrated before each series of measurements using standard buffers of pH 7.0 and 9.0. Determination of Reaction Rate Constants. Two methods of competition kinetics were implemented when using steady-state irradiation. In the first, the depletion of at least one target compound and of a reference compound, both present in the same solution at an initial concentration of 0.5-2.0 µM, was followed by HPLC analysis. Aqueous solutions additionally contained sodium carbonate/bicarbonate (0.1-0.5 M) and one of the photosensitizers (2-10 µM) and the pH was adjusted to the desired value by addition of aqueous sodium hydroxide or hydrochloric acid. The solutions were irradiated in the quartz tubes as described above and 6-12 samples (0.5 mL) were withdrawn at VOL. 39, NO. 23, 2005 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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exposure times between 10 min and 5 h. Assuming that the reaction of CO3•- with each target compound, Pi, and with the reference compound, R, leads to their irreversible transformation with unit yield, the following competition kinetic equation may be used:

( )

ln

[Pi]0

[Pi](t)

)

( )

kCO3•- ,Pi [R]0 ln kCO3•- ,R [R](t)

(4)

where the terms in square brackets denote molar concentrations, t is irradiation time, the subscript “0” denotes concentrations before irradiation, and kCO3•-,Pi and kCO3•-,R are the second-order rate constants for the reaction of CO3•with target compound and reference compound, respectively. Knowing the second-order reaction rate constant for the reference compound, the corresponding rate constant for the target compound was obtained from the slope of ln ([Pi]0/ [Pi](t)) vs ln([R]0/[R](t)), which was determined by linear regression. Depletion of target and reference compound by direct photolysis and reaction with the excited triplet state (the side reaction in Figure 1) was checked by performing irradiations under the same conditions but with carbonate replaced by phosphate buffer (50 mM). Where necessary, a correction of the depletion rates of Pi and R was performed. The second method of competition kinetics consisted of determining the initial depletion rate of the target compound in the presence of different concentrations of ethanol, which was used as a competitor. This method assumes that no scavengers of CO3•- other than the target compound and ethanol are present in significant concentration. Solutions initially contained 10 µM DQ, 1 µM target compound, and 0-40 mM ethanol. Initial depletion rates, rin, for the target compound were determined by linear extrapolation of the depletion time course and analyzed using eq 5.

kCO3•-,R rin([R] ) 0) )1+ [R] rin kCO •-,P [Pi]0 0 3

(5)

i

Using the second-order reaction rate constant for the reference compound, ethanol (≡R), the corresponding rate constant for the target compound was obtained from the slope defined by eq 5, which was calculated by linear regression. Using flash photolysis, absolute second-order rate constants were determined by monitoring the transient absorbance decay of CO3•- (λ ) 600 nm) on the microsecond to millisecond time scale. Aqueous solutions contained CBBP or DQ as the photosensitizer (50-200 µM), carbonate (0.10.5 M), and various concentrations of the target compounds (0-100 µM). Single-exponential functions were fitted to the absorbance decay traces and the first-order decay rate constants, keff, so obtained were fitted to eq 6

keff ) k0 + kCO3•-,Pi[Pi]0

(6)

to provide the second-order rate constants kCO3•-,Pi as the slopes, and the first-order rate constants k0 as the intercepts, which represent the reaction rate constant of CO3•- in the absence of target compound. The latter term may include reaction of CO3•- with various species, such as the photosensitizer, possible impurities, photoproducts, and superoxide radical anion.

Results and Discussion Formation of Carbonate Radical (CO3•-) by Reaction of the Excited Triplet State of Aromatic Ketones with Carbonate Anion. Excitation of a solution containing 50 µM CBBP and 0.5 M Na2CO3 (pH adjusted to 11.0 by addition of HCl) using 9184

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FIGURE 2. Laser flash photolysis traces of aqueous CBBP (50 µM) in the presence of Na2CO3 (0.5 M, pH adjusted at 11.0 by addition of HCl) measured at different observation wavelengths. 308-nm laser pulses resulted in the transient absorption signals displayed in Figure 2. The excited triplet state of CBBP, 3CBBP*, is formed within the duration of the laser pulse. The traces measured at the probe wavelengths of 415 and 540 nm show a fast initial decay of 3CBBP*, which occurs with a first-order rate constant of 9.7 × 105 s-1. In these traces, absorption due to reaction products is already evident, namely, there is an apparently constant residual absorption at delay times >10 µs after the laser pulse. The traces at 600 and 650 nm confirm the formation of products, one decaying on the time-scale of a few microseconds (first-order decay rate constant 4.0 × 105 s-1) and identified with the ketyl radical anion of CBBP, CBBP•-, and another one, which is much longer lived, identified with the carbonate radical. Our assignment of the traces in Figure 2 is in agreement with the known UV-visible spectra of each postulated transient species {absorption maxima for 3CBBP* at 350 nm (8800 M-1 cm-1) and 545 nm (5200 M-1 cm-1); for CBBP•- (pKa ) 8.2) at 360 nm (14 300 M-1 cm-1) and 660 nm (7100 M-1 cm-1) (37); for CO3•- at 600 nm (1860 M-1cm-1) (1)}. The residual absorption, attributed to CO3•-, decays on the millisecond time scale. Formation of CO3•- by laser flash photolysis of aqueous carbonate solutions containing SRFA or SRHA as the photosensitizer was also attempted but no CO3•- could be detected. This result is not surprising in view of the low efficiency of photoproduction of oxidative excited triplet state from DOM (31) and let us conclude that impractically high concentrations of DOM would be required for this purpose. Quenching of Excited Triplet States of Aromatic Ketones by Carbonate Anion. Although many rate constants are available for the reductive quenching of excited triplet states of organic chromophores by inorganic anions in aqueous solution (38), little is known about the quenching by carbonate and bicarbonate anions. Quenching of the excited triplet state and formation of CO3•- has been reported for anthraquinone sulfonates (26) and duroquinone (27, 30). Table 1 summarizes the second-order rate constants for the quenching of five triplet ketones by the carbonate anion, kq(CO32- ) (see Supporting Information for further details). Quenching rate constants by bicarbonate could not be determined with sufficient precision, but they were found to be at least an order of magnitude lower than for carbonate. These results are supported by measurements of the fluorescence quenching of 1-cyanonaphthalene by bicarbonate and carbonate, which show a 10-fold increase in the quenching constant when the pH is increased from 8 to 12 (the experimental method is described elsewhere (39)). As expected, quenching rate constants increase with increasing one-electron standard reduction potential of the excited

TABLE 1. Rate Constants for Quenching of Excited Triplet States of Aromatic Ketones by Carbonate and Hydroxide Anions

a

ketone (X)

E°′red(3X*)a V vs NHE

kq(CO32-)b M-1 s-1

kq(OH-)a M-1 s-1

duroquinone 4-carboxybenzophenone (carboxylate form) benzophenone acetophenone 3′-methoxyacetophenone

2.14 1.83 1.79 1.63 (>1.64) c

2.7 ( 0.8 × 107 1.3 ( 0.2 × 106 1.2 ( 0.2 × 106 2.9 ( 0.6 × 105 1.5 ( 0.6 × 104

1.5 × 109 1.0 × 107 4.9 × 106 1.3 × 105

Calculated from ref 38 except when noted.

b

This study, errors are 95% confidence intervals. c Estimated from ref 32.

triplet state E°′red(3X*), but never reach the diffusioncontrolled limit. This is the case even for DQ, which has a potential that largely exceeds the potential of the couple (CO3•-/CO32-), i.e., 1.59 V vs NHE (40). The highest rate constant, observed for the quenching of DQ, is about half of the value of 7 × 107 M-1s-1 previously obtained for the same compound dissolved in water/ethanol (2:1 vol/vol) (27). These results are analogous to those obtained with sulfite anion as a quencher (38) and underline the pronounced inner-sphere character of electron transfer from carbonate anion to triplet aromatic ketones. The efficiency of carbonate radical formation, η(CO3•-), which is the fraction of CO3•- formed per quenching event (see Supporting Information for its determination), was found to be high. For CBBP and benzophenone, values of 0.86 ( 0.17 and 0.69 ( 0.18 were found, respectively. For DQ a value of ≈1 was obtained. These results agree well with the high, but not quantified, carbonate radical yields previously observed (26, 27, 30, 41). It is remarkable that in most cases carbonate anion is a less efficient quencher than hydroxide anion (see rate constants from the literature in Table 1), while one would expect the opposite behavior in view of the lower one-electron standard reduction potential of the carbonate radical compared to the hydroxyl radical (1.90 V vs NHE (42)). However, formation of hydroxyl radical by this reaction should be negligible due to low production yield (41) and to the generally much lower concentration of hydroxide anion compared to carbonate anion. For the determination of carbonate radical rate constants, mainly CBBP was used owing to its good water solubility as well as its resistance to photodegradation. Despite its faster photodegradation, DQ was also employed with success. Reaction Rate Constants of Carbonate Radical (CO3•-) with Organic Compounds. The second-order rate constants determined in this study by flash photolysis and the other steady-state irradiation methods (vide infra) are summarized in Table 2. The flash photolysis method (see Section S3 of the Supporting Information) proved to be reliable to determine rate constants in the range ≈5 × 106 to 2 × 109 M-1 s-1. Within the concentration range used for target compounds, best results could be obtained with fast-reacting compounds. Only fresh solutions were employed for the determination of rate constants. In some cases it was observed that the decay of CO3•- in solutions containing compounds of low reactivity was accelerated in pre-irradiated solutions, probably as a consequence of products that reacted with CO3•faster than their parent compounds. Most of the second-order rate constants were determined by the first method of competition kinetics (eq 4). Figure 3 shows typical depletion curves of target and reference compound as well as the corresponding competition kinetics plot. The depletion of target and reference compound generally followed zero-order kinetics, at least up to about 50% depletion of both substrates. This concurs with the assumption that target and reference compound are the main scavengers of CO3•- for a substantial initial phase. The rate constants presented (Table 2 and Table S2 of the Supporting Information) are, with few exceptions, in very good agreement

TABLE 2. Second-Order Rate Constants for the Reaction of Carbonate Radical with Organic Compoundsa this studyb(107 M-1 s-1) flash photolysis

organic compound (Pi)

competition kineticsc

(A) Anilines N-methylaniline 255 ( 45 N-ethylaniline 220 ( 40 N,N-dimethylaniline 185 ( 35 4-methylaniline 115 ( 20 4-chloroaniline 62 ( 13 aniline 67 ( 18 61 ( 13 3,4-dichloroaniline 41 ( 9 4-cyanoaniline 18.0 ( 6.0 12.0 ( 3.5 4-aminobenzenesulfonate 24 ( 13 8.7 ( 1.9 4-nitroaniline 7.7 ( 3.4 6.3 ( 2.0

literature values (107 M-1 s-1)

180e 91g, 91i 43g 54e, 50g, 46i

7.3g

(B) Phenoxide ions

310 ( 60 120 ( 25 60 ( 20 35 ( 12 25.5 ( 5.6 11.8 ( 2.7 10.4 ( 3.5 4.0 ( 1.3

4-methoxyphenoxide 4-methylphenoxide 4-hydroxyphenylacetate 4-chlorophenoxide phenoxide vanillinate 4-carboxyphenoxide 4-cyanophenoxide

52f, 130h 48f 19f 5.5e, 24f, 27h 7.9f 6.5h

(C) Pesticides metoxuron isoproturon chlorotoluron monuron diuron fenuron fluometuron propanil ametryn irgarol prometryn terbutryn atraton atrazine

8.1 ( 0.6 ∼2.5

11 ( 5d 3.0 ( 0.4 2.6 ( 0.6d 1.7 ( 0.4 2.2 ( 1.8d 1.5 ( 0.4 0.83 ( 0.24 0.54 ( 0.06 0.60 ( 0.30 0.40 ( 0.30 0.42i 1.4 ( 0.7 0.74 ( 0.38 0.73 ( 0.12 0.61 ( 0.12 0.49 ( 0.13 0.43 ( 0.09 0.37 ( 0.18 0.40i

a k b Errors represent 95% confidence intervals. c First method •CO3 ,Pi. of competition kinetics, eq 4, except when noted. d Second competition kinetics method, eq 5. e Ref 4. f Ref 5. g Ref 6. h Ref 8. i Ref 14, competition kinetics method.

with those determined by flash photolysis. However, it should be noted that, when another aniline was used as a reference compound, reliable constants for anilines could only be obtained at high pH. We believe that the failure, observed for example at pH ) 8.2, is due to oxidation of the more oxidizable aniline by the radical cation of another aniline. Some of the deviations from literature values, observed for a few rate constants of phenoxide ions determined here, might also involve analogous radical mechanisms. A few experiments were also performed using the second method of competition kinetics. Ethanol was used as a competitor in very large excess because of its much lower second-order rate constant {kCO3•-,Ethanol ) 2.20 ( 0.08 × 104 VOL. 39, NO. 23, 2005 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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+ + where ∑σ + ) 0.66(σ+ p2 + σp6 ) + (σm3 + σm5) + σp4 was used to account for multiple substituents (44) (the number in the subscript indicates the position of the substituent on the phenyl ring). Taking into account the multiple substitution and the presence of charged substituents, the values of F+ found in this study agree reasonably well with those mentioned above for uncharged p-substituted compounds. Phenylureas exhibit a Hammett relationship with F+ lying in the same range as for anilines and phenoxides, but rate constants are about one and a half decades lower than for anilines:

log kCO3•-,phenylureas ) 7.20((0.18) - 1.21((0.62) Σσ+ n ) 12, r2 ) 0.65, s ) 0.28 (10) FIGURE 3. Depletion kinetics of 3,4-dichloroaniline (target compound, 1.0 µM) and aniline (reference compound, 0.5 µM) in an aqueous solution containing 2.5 µM CBBP and 0.5 M Na2CO3 (pH)11.8) under steady-state irradiation. The inset shows the competition kinetics plot and regression analysis according to eq 4.

FIGURE 4. Hammett plots of second-order rate constants for the reaction of carbonate radical with anilines, phenoxides, and phenylureas, respectively (eqs 8-10). M-1 s-1 (standard deviation), measured in aqueous 1 M Na2CO3 (9)} than the investigated target compounds (>106 M-1 s-1). Agreement with the results from other methods was good (Table 2 and Table S2). Quantitative structure-activity relationships (QSARs) for the reaction rate constants of CO3•- are known for only two classes of compounds, namely anilines and phenoxide anions, besides a QSAR involving seven monosubstituted benzene derivatives (4, 43). A recent analysis based on literature data yielded Hammett relationships with reaction constants (F+, the slope of the linear Hammett relationship) of -0.98((0.15) and -0.91((0.03) for p-substituted anilines and phenoxide ions, respectively (errors indicate 95% confidence intervals) (43). For homologous series of the above compounds including some other members than the ones used for the previous analysis, rate constant data from the present study gave the Hammett plots shown in Figure 4. The following QSARs for ring-substituted compounds could be derived using all the rate constants determined in this study with the exception of the competition kinetic values found for 4-cyanoaniline and 4-aminobenzenesulfonate (the values from flash photolysis were preferred):

log kCO3•-,anilines ) 8.82((0.16) - 1.05((0.33) Σσ+ n ) 9, r2 ) 0.898, s ) 0.16 (8) log kCO3•-,phenoxides ) 8.40((0.24) - 1.24((0.55) Σσ+ n ) 8, r2 ) 0.84, s ) 0.27 (9) 9186

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No QSAR was attempted for the triazine herbicides because the rate constants vary only within a factor of 2, which means that these compounds are relatively insensitive to substitution as regards their reactivity toward CO3•-. Reaction of Carbonate Radical (CO3•-) with Dissolved Natural Organic Matter (DOM). Carbonate radical in nonpolluted natural waters is supposed to react mainly with DOM (12). To our knowledge, the only published rate constant for reaction of CO3•- with DOM, kCO3•-,DOM, was 40 (mg of C/L)-1 s-1, determined by competition kinetics with aniline for a “strong, hydrophobic acid extract” of water from the Suwannee River, GA (12). In the present study, we initially attempted to employ a similar competition kinetics method using N-methylaniline, 3,4-dichloroaniline, and 4-cyanoaniline, respectively, as the competitor of DOM for carbonate radical. However, this procedure was abandoned due to the observation that oxidation of these anilines, induced in the absence of carbonate by direct reaction with excited triplet state of CBBP and performed by the methods described in our previous studies on phenols and phenylurea herbicides (31-33), was inhibited by DOM. Presumably, the aniline radical cation, (Ar-NH2)•+, resulting from one-electron oxidation of the corresponding aniline by CO3•-, is reduced by DOM to its parent compound (eq 11), which prevents quantification of kCO3•-,DOM by such a method.

(Ar-NH2)•+ + DOM fAr-NH2 + DOM•+

(11)

The direct method of flash photolysis was employed to determine kCO3•-,DOM. For this purpose, CO3•- was generated by reaction of 0.5 M carbonate/bicarbonate with sulfate radical, produced by 248-nm photolysis of 1 mM potassium peroxodisulfate (see eqs 1a and 1b). This method was preferred to the photosensitizer method developed here because the carbonate radical decay is slower, which allows avoidance of unpractically high concentrations of scavenger. A value of 280 ( 90 (mg of C/L)-1 s-1 was obtained for kCO3•-,DOM at pH ) 8.0, which is a factor of 7 higher than the available literature value cited above. This discrepancy underlines the need for more extensive investigations to reliably assess reaction rate constants of CO3•- with DOM. Because CO3•- is much more selective than hydroxyl radical, its reactivity toward different types of DOM is expected to vary in a wider range than for hydroxyl radical {1.4-3.3 × 104 (mg of C/L)-1 s-1 (45)}. Environmental Relevance. In the absence of a reliable experimental method to measure CO3•- formation rate and concentration, we perform here a tentative estimation of CO3•- production by excited triplet states of DOM chromophores, based on the results obtained for the DOMphotoinduced transformation of phenols (31, 32). Accordingly, we assume that excited triplet states of DOM with oxidizing character, 3DOM*, have reduction potentials near the ones of 3′-methoxyacetophenone and benzophenone,

and tentatively set the reaction rate constant of such triplet states with carbonate anion to roughly 1 × 105 M-1s-1 (the average of the values in Table 1 was taken on a logarithmic base). We neglect the reaction of 3DOM* with bicarbonate anion. In the following example, environmental formation rates and steady-state concentrations (denoted by the subscript “ss”) of reactive transients are discussed based on the conditions occurring in the top layer (1-m thickness) of Greifensee, a eutrophic Swiss lake, exposed to midday summer sunlight (46). Taking [3DOM*]ss ≈ 8 × 10-15 M (31) with a carbonate anion concentration of 1 × 10-5 M (approximate concentration at pH ) 8.0), one obtains a CO3•production rate of ∼1 × 10-14 M s-1. For comparison, the CO3•- formation rate induced by hydroxyl radical under the same conditions ([•OH]ss ≈ 4 × 10-17 M (46)) can be calculated to be ∼8 × 10-13 M s-1. At pH ) 9.0, which can be considered as an upper limit for the pH in such a lake during summer (47), formation rates by 3DOM* and hydroxyl radical are ≈1 × 10-13 and ≈2 × 10-12 M s-1, respectively. For open ocean surface waters exposed to sunlight around noon near the tropics, [•OH]ss ) 1.1 × 10-18 M was measured (48), which leads to a CO3•- formation rate of ∼1.4 × 10-13 M s-1. Under the same conditions, the triplet-induced formation rate can be estimated also to be ∼1.4 × 10-13 M s-1 assuming that triplet concentration is proportional to DOM concentration. Details of the calculations are shown in Table S3 of the Supporting Information. Considering that the uncertainty in the calculation of triplet-induced formation rates is high (probably (1 order of magnitude), it can be concluded from these two examples that excited triplet states, although to a likely lesser extent than hydroxyl radical, might be significant sources of CO3•- in sunlit natural waters at pH values higher than ∼8. In sunlit surface waters, CO3•- competes with many other oxidants for the transformation of organic contaminants. The relative importance of each oxidant can be estimated based on their steady-state concentrations and second-order rate constants for reaction with the selected contaminants. Taking the values for kCO3•-,DOM of 40 (12) and 280 (mg of C/L)-1 s-1 (this study), one can estimate [CO3•-]ss values of 5 × 10-15 and 7 × 10-16 M, respectively, for Greifensee water at the above conditions and pH ) 8.0, considering only the hydroxyl radical formation pathway. At pH ) 9.0, values of 1.4 × 10-14 and 2.0 × 10-15 M, respectively, are obtained. These values are comparable (at pH ) 8.0 somewhat lower) to the estimated concentration of excited triplet states, which are more reactive than CO3•-. Therefore, both excited triplet states and CO3•- are expected to have a comparable impact on compounds that react with carbonate radical at rate constants above 1 × 109 M-1 s-1, while excited triplet states, owing to their higher reaction rate constants (see for instance their reaction with a series of substituted phenols (31, 32)), will be the predominant oxidants for many other compounds, probably including various pesticides investigated in this study. Water Treatment. In natural waters and wastewaters subjected to advanced oxidation processes (AOPs), oxidation of contaminants initiated by CO3•- is always in competition with oxidation by hydroxyl radical and, depending on the type of AOP, with oxidation by ozone, oxidation by oxidizing iron species, or phototransformation by UV light. Formation rates of CO3•- cannot exceed those of hydroxyl radical, while the main known scavengers of both radicals are DOM and, if present, hydrogen peroxide. Except for highly polluted waters, where CO3•- scavenging might be dominated by specific pollutants, in waters with almost neutral pH (∼8) the ratio [CO3•-]ss/[•OH]ss will not exceed 2-3 orders of magnitude if DOM controls scavenging, while it will not exceed 63 if hydrogen peroxide controls scavenging (as it may occur in the H2O2/UV AOP). Considering that hydroxyl

radical reaction rate constants for most aromatic compounds are close to 5 × 109 M-1 s-1, CO3•- may override hydroxyl radical only in the oxidation of those compounds having kCO3•-,Pi higher than 5 × 106 M-1 s-1. This conclusion is also valid for sunlit natural waters. In ozone-based AOPs (such as ozonation at higher pH or the ozone/H2O2 process), ozone is very likely to dominate the oxidation of easily oxidizable compounds such as amines, activated aromatic compounds, and olefins (49), while hydroxyl radical will mainly degrade the more resistant compounds. Whether CO3•- plays a role in the presence of ozone should be checked for any particular water, type of AOP, and contaminant using the Rc parameter (which corresponds to the ratio [•OH]ss/[O3] and varies between 10-9 and 10-7 for a great variety of conditions (49)), the ratio [CO3•-]ss/[•OH]ss just discussed, and the rate constants for the reaction of the contaminant with CO3•- and ozone.

Acknowledgments We are grateful to Bruno Hellrung and Yavor Kamdzhilov for technical assistance in flash photolysis experiments, and to Elisabeth Salhi, Erich Meister, and Wendelin Stark for their contribution during preliminary experiments. This work was supported by the Swiss National Science Foundation, Project 20020-105219.

Supporting Information Available Additional details on experimental methods and conditions (chemical compounds, HPLC analysis methods, determination of rate constants), and on calculation of carbonate radical production rates and steady-state concentrations in sunlit natural waters. This material is available free of charge via the Internet at http://pubs.acs.org.

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Received for review June 28, 2005. Revised manuscript received October 3, 2005. Accepted October 4, 2005. ES051236B