Remediation of Chemically-Contaminated Waters Using Sulfate

Downloaded by STANFORD UNIV GREEN LIBR on August 6, 2012 | http://pubs.acs.org Publication Date (Web): September 2, 2011 | doi: 10.1021/bk-2011-1071.ch012

Chapter 12

Remediation of Chemically-Contaminated Waters Using Sulfate Radical Reactions: Kinetic Studies Stephen P. Mezyk,1,* Kimberly A. Rickman,1 Garrett McKay,1 Charlotte M. Hirsch,1 Xuexiang He,2 and Dionysios D. Dionysiou2 1Department

of Chemistry and Biochemistry, California State University Long Beach, 1250 Bellflower Blvd., Long Beach, CA 90840 2School of Energy, Environment, Biological and Medical Engineering, University of Cincinnati, Cincinnati, OH 45221-0012 *Phone: 562-985-4649, Fax: 562-985-8557, Email: [email protected]

The quantitative removal of chemical contaminants in water is one of the most pressing problems facing water utilities today. To augment traditional water treatments that are usually based on adsorptive and chemical-physical processes, radical-based advanced oxidation and reduction processes (AO/RPs) are now being considered. While most AO/RPs utilize the hydroxyl radical in treatment the use of oxidizing sulfate radicals is also gaining interest. To help assess the applicability of sulfate radical based AO/RPs in remediating contaminated waters, here we have determined absolute rate constants and reaction mechanisms for SO4-• reaction with four β-lactam antibiotics (amoxicillin, penicillin-G, piperacillin, tircarcillin), three estrogenic steroids (ethynylestradiol, estradiol, and progesterone) and one personal care product (isoborneol). For the four antibiotics of this study the relatively fast rate constant values suggests that the majority of the SO4-• oxidation occurs at the sulfur atom in the ring adjacent to the β-lactam moiety, as opposed to the hydroxyl radical reaction which occurs at peripheral aromatic rings. The measured sulfate radical rate constants for estradiol and progesterone are identical, with the slightly faster value for ethynylestradiol suggesting significant oxidation occurring at its ethynyl moiety. For isoborneol, the

© 2011 American Chemical Society In Aquatic Redox Chemistry; Tratnyek, P., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 2011.


sulfate radical reactivity was slightly lower, but still fast enough that AO/RP treatment utilizing this radical might be feasible at large-scale. Piperacillin was also chosen for a detailed investigation of its degradation by both SO4-• and •OH in a laboratory scale homogeneous UV photochemical system. It was found that although the absolute reaction rate constant for piperacillin reaction with SO4-• was lower than for •OH, the overall removal of this antibiotic was more effective when using UV/S2O82- than UV/H2O2. For an initial oxidant dose of 1 mM and an antibiotic concentration of 50 µM, percentage removals of 65.2% and 33.0%, respectively, at a UV fluence of 320 mJ/cm2 were obtained. This difference was attributed to the higher quantum yield of sulfate radical production from persulfate under UV 254 nm irradiation.

Introduction The adverse ecological impacts of endocrine-disrupting compounds, personal care products, antibiotics, and pesticides or herbicides in water supplies (1–6) and wastewater effluents are causing concern amongst regulatory groups and the public. Traditional water treatment relies primarily upon adsorptive and chemical-physical processes to remove or transform these unwanted organic contaminants. However, these treatment processes may sometimes not be sufficient (6), as quantitative removal of low (ng L-1) concentrations of dissolved chemicals may be complicated by the presence of much higher levels of other water constituents such as dissolved organic matter (DOM) and carbonate. In order to augment traditional water treatment processes, the use of in-situ generated radical species, such as the oxidizing hydroxyl radical (•OH) and/or reducing electron (eaq-) and hydrogen atoms (H•), to react with and destroy trace contaminants following standard water treatment processes could be a viable approach. These additional treatments are generally referred to as advanced oxidation/reduction processes (AO/RPs) (7–11). These radicals can be created using a variety of techniques (12); for example, the hydroxyl radical (•OH) can be generated through using a combination of O3/H2O2, O3/UV-C, or H2O2/UV-C, and mixtures of •OH, eaq-, and H• are produced from the UV irradiation of titanium dioxide, sonolysis, or the irradiation of water via electron beams or γ rays. The utilization of reducing radicals to destroy chemical contaminants in real-world waters is problematic due to the presence of dissolved oxygen, which preferentially scavenges these radicals to create the much less reactive superoxide radical, O2-• ([O2] ~ 2.5 × 10-4 M, k = 1.9 × 1010 M-1 s-1 (13)). Therefore, the most widespread AO/RPs are based on only the •OH radical production. However, another AO/RP that is gaining interest utilizes sulfate radical (SO4-•) reactions (14). The sulfate radical is also strongly oxidizing (Eo = 2.3 V, (15)) which means that it can react with most organic chemical contaminants. It typically reacts by electron abstraction from electron-rich centers in molecules, in contrast to 248 In Aquatic Redox Chemistry; Tratnyek, P., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 2011.


hydroxyl radical based oxidations that mainly occur by hydrogen atom abstraction and/or addition to aromatic moieties. Sulfate radicals can be readily formed through persulfate (S2O82-) decomposition induced by UV light (UV/ S2O82-), the presence of a catalyst, or higher temperatures. In addition, they can be formed in water by the addition of S2O82- to selected AO/RPs, where the reducing radicals will react with this species according to (13):

Persulfate addition itself has been shown to be effective in treating subsurface soils contaminated with chemicals such as diesel (16), PCBs (17), PAHs (18, 19), chlorinated hydrocarbons (20, 21) and VOCs (22). These remediation processes are considerably enhanced by persulfate activation, which increases its rate of decomposition to form sulfate radicals. However, much less investigation of sulfate radical use in chemically-contaminated water treatment has been reported. The sulfate radical reaction has previously been shown to have high efficiency in degrading model organic chemicals in water (14), but little is known about its chemistry with contaminants of higher molecular weight in real-world waters. The optimal, quantitative, removal of water contaminants through the use of AO/RPs requires a thorough understanding of the redox chemistry occurring between free radicals and the contaminant chemicals of concern under the conditions of use. This can be accomplished through kinetic computer models, that give the most information and provide the best test of the proposed treatment (23) as all the chemistry in the system is considered. A critical component for kinetic modeling of any free radical based process is the full understanding of the kinetics and mechanisms of the radical reactions occurring. These fundamental data allow for quantitative computer modeling of AO/RP systems to establish the feasibility and large-scale efficiency of using radicals for specific contaminant removal under real-world conditions. Therefore, in this work we describe absolute rate constant measurements for the sulfate radical reaction with four typical antibiotics, three representative estrogenic steroids, and one personal care compound isoborneol (see Figure 1) measured using an electron pulse radiolysis system. In addition, the radical-induced degradation of one specific antibiotic, pipericillin, was investigated in a laboratory scale homogeneous photochemical system using UV 254 nm /S2O82- and UV 254 nm/H2O2 to provide further insight into the radical chemistry occurring.

Experimental Kinetic Studies Chemicals used in this study were purchased from Sigma-Aldrich Chemical Company at the highest purity available (steroids, >98%, KSCN, 99%, K2S2O8, 99%, Na2S2O8, 98%, β-lactam antibiotics >98%, isoborneol >98%). m-Toluic 249 In Aquatic Redox Chemistry; Tratnyek, P., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 2011.


acid (3-methylbenzoic acid) was purchased from Fisher Scientific (99%). All were used as received. There are many possible methods of producing AO/RP radicals (12) but the use of an electron beam for kinetic studies is optimal, as its energy deposition quantitatively generates a mixture of •OH, eaq- and H• directlyfrom breaking bonds and ionizing water molecules (13) according to the stoichiometry:

The numbers preceding each species in Equation (3) are their absolute yields (G-value) in units of µmol J-1. These yields are constant for solutions containing relatively low concentrations (< 0.10 M) of solutes irradiated in the pH range 3-10. Electron pulse radiolysis allows generation of all these species in nanoseconds, while the secondary reaction between the produced radicals and any added solute molecule typically occurs on a microsecond timescale. Hydrogen peroxide reactions occur at much longer times (milliseconds or greater), and so do not interfere with radical kinetic measurements. The study of sulfate radical kinetic measurements by this technique requires the prior removal of hydroxyl radicals in order to isolate the reducing species. Therefore, these kinetic experiments were conducted using a constant high concentration of tert-butanol, (CH3)3COH, as a co-solvent (0.5-2.0 M), which immediately scavenges the radiolytically produced hydroxyl radicals and most hydrogen atoms (Equations 4 & 5) to produce the relatively inert •CH2(CH3)2COH alcohol radical (13):

The isolated hydrated electron quantitatively reacts with added persulfate (in our experiments 5.0 mM) to give the oxidizing sulfate radical. The sulfate radical will also slowly react with the added tert-butanol (13),

but by careful selection of added concentrations a significant fraction of sulfate radicals will react with the added chemical solute. All rate constant data were collected using the linear accelerator facilities at the Radiation Laboratory, University of Notre Dame. This irradiation and transient absorption detection system has been described in full detail previously (24). Absolute radical concentrations (dosimetry) were determined using the hydroxyl radical oxidation of N2O-saturated 1.0 x 10-2 M thiocyanate (KSCN) solutions at natural pH, whose efficiency has been previously established (25). These measurements were performed daily. The presence of the high alcohol concentration means that significant intraspur radical scavenging will occur, which will increase the initial hydrated electron 250 In Aquatic Redox Chemistry; Tratnyek, P., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 2011.


Figure 1. Structures of four antibiotics (amoxicillin, penicillin-G, piperacillin, tircarcillin), three estrogenic steroids (estradiol, ethynylestradiol, progesterone) and one personal care compound (isoborneol) of interest in this study. and sulfate radical yields (26). However, this mixed solvent solution also allowed higher concentrations of the steroids to be dissolved (isoborneol and the β-lactam antibiotics were sufficiently soluble in water) which meant that good pseudo-firstorder conditions ([Solute]:[SO4-•] > 20:1) were maintained. This considerably simplified the data analysis. 251 In Aquatic Redox Chemistry; Tratnyek, P., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 2011.


Solution flow rates were adjusted so that each pulse irradiation was performed on a fresh sample, and multiple traces (5-15) were averaged to produce a single kinetic trace. Typically, 3-6 ns pulses of 8 MeV electrons generating radical concentrations of 2-10 µM SO4-• per pulse were used in these experiments. All of these experiments were conducted at ambient room temperature (20 ± 2°C) with the temperature variation in any given experiment being less than ± 0.3oC. Rate constant error limits reported here are the combination of experimental precision and compound purities.

Homogeneous Photochemical Degradation of Piperacillin In addition to the directly measured kinetic parameters of this study, one antibiotic, pipericillin, was also chosen for further investigation into its radical-induced degradation using a homogeneous UV 254 nm photochemical system as an example of the application of AO/RPs to remediate chemically contaminated waters. Two different radicals, SO4-• and •OH, were generated by UV irradiation (15 W low-pressure UV lamps by Cole-Parmer, λmax = 254 nm) of added S2O82- and H2O2, respectively. The experiment was conducted in a collimated system made according to Bolton and Linden (27). The irradiance was determined by three different methods, iodide/iodate actinometry (28), ferrioxalate actinometry (29), and a radiometer. Before each experiment, the lamps were allowed to warm up for at least 30 minutes. A Pyrex® glass petri dish (60 mm × 15 mm) with a quartz cover was used as the reactor, to which a solution of 10 mL was added and mildly mixed with a magnetic stirrer bar. During the experiment, 0.1 mL samples were taken at specific fluence levels, and then mixed with 0.1 mL methanol to quench all of the radical reactions occurring. The solution concentration of piperacillin was determined by HPLC. An Agilent 1100 Series quaternary LC and a Nova-Pak C18 Waters 5-µm (3.9 mm × 150 mm) column was used with the photodiode array detector set at 238 nm. The mobile phase was the combination of 0.1% acetic acid in Milli-Q water (A) and acetonitrile (B) with a gradient mode of 95% A and 5% B as the initial, gradually changing to 65% A and 35% B in 12 minutes, 85% and 15% in the following 6 minutes, and return to the original combination at 20 minutes. The flow rate was set at 0.5 mL/min, the injection volume was 20 µL, and the temperature of the column was 25 °C. The rate constant of piperacillin with SO4-• was determined in this homogeneous UV 254 photochemical system using a competition approach with m-Toluic acid as the standard. The initial concentrations of pipericillin, Toluic acid, sodium persulfate, tert-butanol, and phosphate buffer (pH=7.4) were 50 µM, 50 µM, 10 mM, 500 mM, and 5 mM, respectively. The high concentration of persulfate was to ensure the sufficient production of SO4-• while the high concentration of tert-butanol was used again to quench any hydroxyl radical produced under UV irradiation of the reaction solution.

252 In Aquatic Redox Chemistry; Tratnyek, P., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 2011.

Results and Discussion


Kinetic Studies The sulfate radical has a broad absorption spectrum, with a maximum near 450 nm (30). The reactions of sulfate radicals with four β-lactam antibiotics, amoxicillin, penicillin-G, piperacillin and ticarcillin, were investigated in this work. These experiments were conducted in solutions containing 0.5 M tert-butanol which were adjusted to a pH of 7.4 using 5.00 mM phosphate buffer. Maintaining a constant near-neutral pH was necessary for reactivity comparison as a pH dependence has been previously reported for sulfate radical reaction with carboxylic acids and TCE in aqueous solution (31, 32). Typical antibiotic concentrations were 100-500 µM. From the rate of change of the first-order decay kinetics observed with varying antibiotic concentration (see Figure 2a) second order reaction rate constants could be readily determined (Figure 2b). Our measured rate constants are summarized in Table 1.

Figure 2. a) Decay of SO4-• radical at 450 nm for 92.7 (□), 300.0 (▿) and 500.0 µM (▵) added amoxicillin. Kinetic curves are offset in order to aid visibility. Solid lines through data points are fitted first-order kinetics, with pseudo-first-order rate constants of (7.06 ± 0.l2) × 105, (1.30 ± 0.03) × 106 and (2.09 ± 0.05) × 106 s-1, respectively. b) Second order transformation of first-order fitted values plotted against amoxicillin concentration. Solid line corresponds to second order rate constant, k = (3.48 ± 0.05) × 109 M-1 s-1. Inset: Oxidized amoxicillin product species absorbance in arbitrary units taken at end of sulfate radical reaction. 253 In Aquatic Redox Chemistry; Tratnyek, P., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 2011.

Table 1. Summary of sulfate radical reaction rate constants determined in this study Compound

Measured kSO4-• a M-1 s-1

k•OH M-1 s-1

amoxicillin

2.9 × 109

(6.94 ± 0.44) × 109 (39)

penicillin G

1.4 × 109 (30)

(8.70 ± 0.32) × 109 (40)


piperacillin

(1.74 ± 0.11) × 109 (30) 1.2 × 109 (30) 1.85 × 109 b 1.16 × 109

b

(7.84 ± 0.49) × 109 (40)

8.0 × 108 (30)

(8.18 ± 0.99) × 109 (40)

EE2

(3.01 ± 0.28) × 109 (38)

(1.52 ± 0.23) × 1010 (38)

estradiol

(1.21 ± 0.16) × 109 (38)

(1.15 ± 0.28) × 1010 (38)

progesterone

(1.19 ± 0.16) × 109 (38)

(8.5 ± 0.9) × 108 (38)

(5.28 ± 0.13) × 108

______

ticarcillin

isoborneol a

Rate constants for antibiotics are zero ionic strength corrected values. for antibiotics with SO4-• are not corrected for zero ionic strength.

b

Rate constants

It is important to note that at this pH these antibiotics are negatively charged, like the sulfate radical, and so the measured reaction rate constants will be dependent upon the solution total ionic strength. As our stock solution had a relatively high ionic strength of 0.025 M, we corrected our measured values to zero ionic strength using the standard equation (33):

where k is the measured second-order rate constant, ko is the corresponding zero ionic strength value, z1 and z2 are the charges of the sulfate radical (-1) and antibiotic, respectively, and I is the solution ionic strength. The antibiotic charges at pH 7.4 were calculated based on literature pKa values (34–36). For amoxicillin, with its reported pKa values of 2.4 and 7.49 (37), this reduces the SO4-• reaction rate constant from its measured value of 3.48 × 109 to 2.96 × 109 M-1 s-1. Previously (38) we had used an incorrect high pKa2 value of 9.6 for this correction which resulted in an erroneously high SO4-• rate constant for this antibiotic. The zero ionic strength rate constants for penicillin G, piperacillin, and tircarcillin were found to be 1.44 × 109, 1.17 × 109, and 0.80 × 109 M-1 s-1, respectively (see Table 1). While amoxicillin has the fastest sulfate radical reaction rate constant, the other three antibiotics show effectively the same value within experimental error. Moreover, the transient absorption spectra (see for example Figure 2b, Inset) obtained for the initial species produced in sulfate radical oxidation of all these antibiotics as well as their parent (+)-6-aminopenicillanic acid compound (38) 254 In Aquatic Redox Chemistry; Tratnyek, P., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 2011.


are similar, which implies a consistent reaction mechanism for these antibiotics. Based on these data we infer that the predominant oxidizing SO4-• reaction occurs at the S atom in the five-member ring immediately adjacent to the β-lactam ring in these antibiotics. In contrast, it has previously been reported (39, 40) that the corresponding hydroxyl radical reactions for these antibiotic species occur predominately at peripheral aromatic rings, producing hydroxylated species with an intact β-lactam core. While the hydroxyl radical oxidations are considerably faster than those of the sulfate radical (see Table 1) the closer site of reactivity of the latter (Figure 3) implies that it could be more efficient in destroying antibiotic activity than the corresponding hydroxyl radical reaction.

Figure 3. Suggested initial reaction mechanisms of sulfate radical oxidation of amoxicillin, EE2, and isoborneol. SO4-• Reaction with Estrogenic Steroids The sulfate radical reactivity with three typical contaminant estrogenic steroids (Figure 1) was also determined in this work. The considerably 255 In Aquatic Redox Chemistry; Tratnyek, P., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 2011.


decreased aqueous solubility of these chemicals meant that much lower steroid concentrations had to be used. However, following the same methodology as for the antibiotics, sulfate radical reaction rate constants were obtained (38) for ethynylestradiol (EE2), estradiol and progesterone as 3.01 × 109, 1.21 × 109 and 1.19 × 109 M-1 s-1, respectively (see Table 1). These second-order kinetic data are shown in Figures 4(a-c). For both progesterone and estradiol the SO4-• rate constant was slower than for EE2, implying that different reaction mechanisms were occurring. While these three rate constants do not allow the specific mechanism of oxidation for progesterone and estradiol to be quantitatively determined, the significantly faster rate constant for EE2 suggests that significant SO4-• oxidation occurs at the ethinyl bond in this molecule (Figure 3).

Figure 4. Second order kinetic plots for SO4-• radical reaction with a) ethynylestradiol (EE2), b) estradiol, c) progesterone, and d) isoborneol. Solid lines are weighted linear fits, corresponding to rate constants of (3.01 ± 0.28) × 109 M-1 s-1 and (1.21 ± 0.16) × 109 M-1 s-1, (1.19 ± 0.16) × 109 M-1 s-1, and (5.28 ± 0.13) × 108 M-1 s-1, respectively. All these sulfate radical oxidations are slower than measured for the hydroxyl radical, (see Table 1) indicating that the initial reaction mechanisms for these two oxidizing radicals differ. As observed for the antibiotic oxidations (38) the •OH reaction is anticipated to add to the constituent phenyl ring in these steroids. This would give dihydroxy stable product species, and for estradiol these dihydroxy species have been shown to have comparable steroidal activity and higher aqueous solubility (41). Therefore, the use of sulfate radicals which preferentially reacts at other electron-rich centers in these molecules may prove advantageous for the total removal of estrogenic activity in the treatment of large-scale real-world waters. 256 In Aquatic Redox Chemistry; Tratnyek, P., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 2011.


SO4-• Reaction with Isoborneol We have also investigated the sulfate radical reaction of a common personal care chemical, isoborneol (Figure 1). This aliphatic ring based chemical is a perfume agent, food additive, and moth repellant, and can be a skin, eye, and respiratory irritant (42). While isoborneol itself is not a major contaminant of concern, this compound was chosen for study because it is a saturated monoterpene that is more soluble in water than 2-methyl isoborneol, an analogous contaminant that has a major water quality impact due to its low odor threshold in the ppb range. The second-order reaction rate constant for the sulfate radical with isoborneol (Figure 3) was found to be (5.28 ± 0.13) × 108 M-1 s-1. While its reactivity is slower than for the estrogenic steroids and antibiotics of this study (Table 1), it was still surprisingly fast when compared to simple aliphatic alcohols such as tert-butanol (Equation 6, k6 = 8.4 × 105 M-1 s-1). A previous, systematic, investigation of other sulfate radical reaction rate constants with other (non-ring) aliphatic alcohols in water (43) has established a relationship with number and type of C-H bonds in these molecules. Based on the principle of additive reactivity, individual rate constants for primary, secondary, and tertiary C-H bond were calculated as 7.4 × 105, 1.8 × 107 and 9.8 × 107 M-1 s-1, respectively. Using these values for isoborneol, an overall sulfate radical rate constant of 3.1 × 108 M-1 s-1 is predicted, only slightly lower than its measured value. In addition, this analysis suggests that the predominant oxidation occurs by hydrogen atom abstraction from a tertiary carbon atom, as shown in Figure 3. Unfortunately, no hydroxyl radical rate constant for isoborneol could be found in the literature, precluding any further quantitative comparison between these two oxidizing radicals. However, it would be expected that •OH reaction would also occur by hydrogen atom abstraction, with a rate constant in the 108 - 109 M-1 s-1 range (13). Similar values would be expected for the 2-methyl isoborneol compound as based on their close structures. Overall, these sulfate radical data suggest that treatment of aqueous waste streams containing isoborneols by AO/RPs could be a viable treatment method, with both the oxidizing hydroxyl and sulfate radicals assisting in the total removal of these chemicals. Homogeneous Photochemical Degradation of Piperacillin Persulfate is a strong inorganic oxidant; however, it can be activated to form sulfate radicals (44–47) by use of a transition metal, UV irradiation or by increasing the temperature. In this study the application of sulfate radicals was shown by the degradation of a β-lactam antibiotic, piperacillin, in a laboratory scale homogeneous UV photochemical system (Figure 5). Initial concentrations of pipercillin and persulfate were 50 µM and 1 mM, respectively. At a UV fluence of 320 mJ/cm2 the degradation of piperacillin was 4.7%, 33.0% and 65.2% for UV, UV/S2O82- and UV/H2O2, respectively. Under dark conditions, there was no significant removal of pipericillin either by S2O82- or by H2O2 during the same time interval. Previously, the second order rate constants for SO4-• and •OH radicals with pipericillin were determined to be 1.2 × 109 M-1 s-1 (30) and (7.84 ± 0.49) ×109 M-1 s-1 (40), respectively, (the sulfate radical value was corrected to 257 In Aquatic Redox Chemistry; Tratnyek, P., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 2011.


Figure 5. Measured degradation of piperacillin in the homogeneous photochemical system, with the initial concentration of the antibiotic and the oxidant to be 50 µM and 1 mM, respectively, in Milli-Q water. zero ionic strength using Equation (7)). The slightly faster degradation of PIP by the UV/S2O82- AO/RP in comparison to the UV/H2O2 was therefore attributed to the rate of formation of the radicals, which was shown in the difference in the quantum yield of the oxidants under UV-254 nm irradiation (Equations 8 & 9, (48–50))

In this study, a chemical competition kinetics approach was also evaluated to determine the second order rate constant for sulfate radical reaction with pipericillin. The competitor, m-Toluic acid, was chosen because of its similar rate constant to pipericillin with SO4-•, 2 × 109 M-1 s-1 (51). As shown in Figure 6, neither pipericillin nor m-Toluic acid underwent directly photolysis; however, both of them followed UV fluence-based pseudo-first-order reaction kinetics, with kobs(PIP) = 1.74 × 10-3 cm2/mJ and kobs(TA) = 1.88 × 10-3 cm2/mJ (m-Toluic acid) when sulfate radicals were generated in the AO/RP. The absolute second order rate constant could thus be determined by equation (10)



to be kSO4-•(PIP) = 1.85 × 109 M-1 s-1. Correcting this value for ionic strength effects (the pKa of pipericillin was 4.14 and the total ionic strength at pH 7.4 was 0.041 M) according to equation (7) gave a limiting value of kSO4-•(PIP) = 1.16 × 109 M-1 s-1, in excellent agreement with our previous result (30).

Figure 6. Determination of second order rate constant of piperacillin with SO4-• through m-Toluic acid competition kinetics. For the UV/S2O82- experiment the initial concentrations of piperacillin, m-Toluic acid, sodium persulfate, tert-butanol, and phosphate buffer (pH=7.4) were 50 µM, 50 µM, 10 mM, 500 mM, and 5 mM, respectively. The chemical concentrations were the same for the UV-only experiment in the absence of persulfate.

Conclusions Rate constants for the reactions of oxidizing sulfate radicals have been determined for four common antibiotics, three estrogenic steroids, and one personal care product in water. The relatively fast rate constants extrapolated to zero ionic strength for SO4-• reaction with amoxicillin (2.9 × 109 M-1 s-1), penicillin-G (1.4 × 109 M-1 s-1), piperacillin (1.2 × 109 M-1 s-1), and tircarcillin (8.7 × 108 M-1 s-1), suggests a consistent mechanism, believed to be electron abstraction from the sulfur atom in the five-member ring adjacent to the β-lactam ring in these compounds. For the three estrogenic steroids, ethynylestradiol ((3.01 ± 0.28) × 1010 M-1 s-1), estradiol ((1.21 ± 0.16) × 1010 M-1 s-1, and progesterone ((1.19 ± 0.16) × 109 M-1 s-1) the lower values for the latter two suggests that significant sulfate radical reaction at the triple bond occurs for EE2. For the saturated aliphatic isoborneol ((5.28 ± 0.13) × 108 M-1 s-1), the sulfate radical reaction is still relatively fast suggesting that AO/RP treatment utilizing this 259 In Aquatic Redox Chemistry; Tratnyek, P., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 2011.

radical may be feasible at large-scale. Although the absolute rate constant of piperacillin with SO4-• was lower than that for •OH, the removal of the compound was faster by UV/S2O82- than by UV/H2O2, with a percentage loss of 65.2% and 33.0%, respectively, at UV fluence of 320 mJ/cm2. This finding is attributed to the higher quantum yield of persulfate under UV 254 nm irradiation.


Acknowledgments Rate constant measurements were performed at the Radiation Laboratory, University of Notre Dame, which is supported by the Office of Basic Energy Sciences, U.S. Department of Energy.

References 1.

Ternes, T. A. Occurrence of drugs in German sewage treatment plans and rivers. Water Res. 1998, 32, 3245–3260. 2. Synder, S. A.; Westerhoff, P.; Yoon, Y.; Sedlak, D. L. Pharmaceuticals, personal care products, and endocrine disruptors in water: implications for the water industry. Environ. Eng. Sci. 2003, 20, 449–469. 3. Westerhoff, P.; Yoon, Y.; Snyder, S. A.; Wert, E. Fate of endocrine-disruptor, pharmaceutical, and personal care product chemicals during simulated drinking water treatment processes. Environ. Sci. Technol. 2005, 39, 6649–6663. 4. Kummerer, K. Antibiotics in the aquatic environment: A review Part I. Chemosphere 2009, 75, 417–434. 5. Kummerer, K. Antibiotics in the aquatic environment: A review Part II. Chemosphere 2009, 75, 435–441. 6. Kolpin, D. W.; Furlong, E. T.; Meyer, M. T.; Thurman, E. M.; Zaugg, S. D.; Barber, L. B.; Buxton, H. T. Pharmaceuticals, Hormones, and Other Organic Wastewater Contaminants in U.S. Streams, 1999−2000: A National Reconnaissance. Environ. Sci. Technol. 2002, 36, 1202–1211. 7. Ikehata, K.; Nagashkar, N. J.; El-Din, M. C. Degradation of aqueous pharmaceuticals by ozonation and advanced oxidation processes: A review. Ozone: Sci. Eng. 2006, 28, 353–414. 8. Burbano, A. A.; Dionysiou, D. D.; Suidan, M. T. Effect of oxidant-tosubstrate ratios on the degradation of MTBE with Fenton reagent. Water Res. 2008, 42, 3225–3239. 9. Ning, B.; Graham, N.; Zhang, Y. P.; Nakonechny, M.; El-Din, M. G. Degradation of endocrine disrupting chemicals by ozone/AOPs. Ozone: Sci. Eng. 2007, 29, 153–176. 10. Lee, Y.; Escher, B. I.; von Gunten, U. Efficient removal of estrogenic activity during oxidative treatment of waters containing steroid estrogens. Environ. Sci. Technol. 2008, 42, 6333–6339. 11. Huber, M. M.; Ternes, T. A.; von Gunten, U. Removal of estrogenic activity and formation of oxidation products during ozonation of 17 alpha-ethinylestradiol. Environ. Sci. Technol. 2004, 38, 5177–5186. 260 In Aquatic Redox Chemistry; Tratnyek, P., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 2011.


12. Cooper, W. J.; Cramer, C. J.; Martin, N. H.; Mezyk, S. P.; O’Shea, K. E.; von Sonntag, C. Free radical mechanisms for the treatment of methyl tertbutyl ether (MTBE) via advanced oxidation/reductive processes in aqueous solution. Chem. Rev. 2009, 109, 1302–1345. 13. Buxton, G. V.; Greenstock, C. L.; Helman, W. P.; Ross, A. B. Critical review of rate constants for reactions of hydrated electrons, hydrogen atom and hydroxyl radicals (•OH/•O-) in aqueous solutions. J. Phys. Chem. Ref. Data 1988, 17, 513–886. 14. Anipsitakis, G. P.; Dionysiou, D. D. Transition metal/UV-base advanced oxidation technologies for water decontamination. Appl. Catal. B 2004, 54, 155–163. 15. Huie, R. E.; Clifton, C. L.; Neta, P. Electron transfer rates and equilibria of the carbonate and sulfate radical anions. Int. J. Radiat. Phys. Chem. 1991, 5, 477–481. 16. Do, S.-H.; Kwon, Y.-J.; Kong, S.-H. Effect of metal oxides on the reactivity of persulfate/Fe(II) in the remediation of diesel-contaminated soil and sand. J. Hazard. Mater. 2010, 182, 933–936. 17. Yukselen-Aksoy, Y.; Khodadoust, A. P.; Reddy, K. R. Destruction of PCB 44 in spiked subsurface soils using activated persulfate oxidation. Water, Air, Soil Pollut. 2010, 209, 419–427. 18. Andreottola, G.; Bonomo, L.; De Gioannis, G.; Ferrarese, E.; Muntoni, A.; Polettini, A.; Pomi, R.; Saponaro, S. Lab-scale feasibility tests for sediment treatment using different physico-chemical techniques. J. Soils Sediments 2010, 10, 142–150. 19. Tsai, T. T.; Kao, C. M.; Hong, A. Treatment of tetrachloroethylenecontaminated groundwater by surfactant-enhanced persulfate/BOF slag oxidation – a laboratory feasibility study. J. Hazard. Mater. 2009, 171, 571–576. 20. Teel, A. L.; Cutler, LM.; Watts, R. J. Effect of sorption on contaminant oxidation in activated persulfate systems. J. Environ. Sci. Health, Part A 2009, 44, 1098–1103. 21. Liang, C.; Lin, Y.-T.; Shih, W.-H. Treatment of trichloroethylene by adsorption and persulfate oxidation in batch studies. Ind. Eng. Chem. Res. 2009, 48, 8373–8380. 22. Huang, K.-C.; Zhao, Z.; Hoag, G. E.; Dahmani, A.; Block, P. A. Degradation of volatile organic compounds with thermally activated persulfate oxidation. Chemosphere 2005, 61, 551–560. 23. Crittenden, J. C.; Hu, S.; Hand, D. W.; Green, S. A. A kinetic model for H2O2/UV process in a completely mixed batch reactor. Water Res. 1999, 33, 2315–2328. 24. Whitman, K.; Lyons, S.; Miller, R.; Nett, D.; Treas, P.; Zante, A.; Fessenden, R. W.; Thomas, M. D.; Wang, Y. Linear accelerator for radiation chemistry research at Notre Dame 1995. Proceedings of the ’95 Particle Accelerator Conference & International Conference of High Energy Accelerators, Dallas, TX, 1996. 25. Buxton, G. V.; Stuart, C. R. Re-evaluation of the thiocyanate dosimeter for pulse radiolysis. J. Chem. Soc. Faraday Trans. 1995, 91, 279–282. 261 In Aquatic Redox Chemistry; Tratnyek, P., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 2011.


26. LaVerne, J. A.; Pimblott, S. M. Yields of hydroxyl radical and hydrated electron scavenging reactions in aqueous solutions of biological interest. Rad. Res. 1993, 135, 16–23. 27. Bolton, J. R.; Linden, K. G. Standardization of methods for fluence (UV dose) determination in bench-scale UV experiments. J. Environ. Eng. 2003, 129 (3), 209–215. 28. Rahn, R. O. Potassium iodide as a chemical actinometer of 254 nm radiation: use of iodate as an electron scavenger. Photochem. Photobiol. 1997, 66, 450–455. 29. Murov, S. L.; Carmichael, I.; Hug, G. L. Handbook of photochemistry, 2nd ed.; Marcel Dekker: New York, 1993; pp 330−336. 30. Rickman, K. A.; Mezyk, S. P. Kinetics and mechanisms of sulfate radical oxidation of β-lactam antibiotics in water. Chemosphere 2010, 81, 359–365. 31. Criquet, J.; Nebout, P.; Karpel Vel Leitner, N. Enhancement of carboxylic acid degradation with sulfate radical generated by persulfate activation. Water Sci. Technol. 2010, 61, 1221–1226. 32. Liang, C.; Zih-Sin, W.; Bruell, C. J. Influence of pH on persulfate oxidation of TCE at ambient temperatures. Chemosphere 2007, 66, 106–113. 33. Connors, K. A. Chemical kinetics: The study of reaction rates in solution; VCH Publishers Inc.: New York, NY, 1990. 34. Florey, E., Ed. Analytical profiles of drug substances; Academic Press: New York, 1978. 35. Hanch, C., Sammas, P. G., Taylor, J. B., Eds. Comprehensive medicinal chemistry; Pergamon Press: Oxford, 1990; Vol. 6. 36. Alkseev, V. G.; Volkova, I. A. Acid-base properties of some penicillins. Russ. J. Gen. Chem. 1991, 73, 1616–1618. 37. Andreozzi, R.; Canterino, M.; Marotta, R.; Paxeus, N. Antibiotic removal from wastewaters: The ozonation of amoxicillin. J. Hazard. Mater. 2005, 122, 243–250. 38. Mezyk, S. P.; Abud, E. M.; Swancutt, K. L.; McKay, G.; Dionysiou, D. D. Removing steroids from contaminated waters using radical reactions; Contaminants of Emerging Concern in the Environment: Ecological and Human Health Considerations; ACS Symposium Series; American Chemical Society: Washington, DC, 2010; Vol. 1048, Chapter 9, pp 213−225. 39. Song, W.; Weisang, C.; Cooper, W. J.; Greaves, J.; Miller, G. E. Free radical destruction of β-lactam antibiotics in aqueous solution. J. Phys. Chem. A. 2008, 112, 7411–7417. 40. Dail, M. K.; Mezyk, S. P. Hydroxyl-radical-induced degradative oxidation of β-lactam antibiotics in water: Absolute rate constant measurements. J. Phys. Chem. A 2010, 114, 8391–8395. 41. Zhao, Z.; Kosinska, W.; Khmelnitsky, M.; Cavalieri, E. L.; Rogan, E. G.; Chakravarti, D.; Sacks, P. G.; Guttenplan, J. B. Mutagenic activity of 4hydroxyestradiol, but not 2-hydroxyestradiol, in BB Rat2 embryonic cells, and the mutational spectrum of 4-hydroxyestradiol. Chem. Res. Toxicol. 2006, 19, 475–479. 42. MSDS sheet. http://msds.chem.ox.ac.uk/IS/isoborneol.html. 262 In Aquatic Redox Chemistry; Tratnyek, P., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 2011.


43. Huie, R. E.; Clifton, C. L. Rate constants for hydrogen atom abstraction reactions of the sulfate radical, SO4-. Alkanes and ethers. Int. J. Chem. Kinet. 1989, 21, 611–619. 44. Anipsitakis, G. P.; Dionysiou, D. D. Degradation of organic contaminants in water with sulfate radicals generated by the conjunction of peroxymonosulfate with cobalt. Environ. Sci. Technol. 2003, 37 (20), 4790–4797. 45. Bandala, E. R.; Peláez, M. A.; Dionysiou, D. D.; Gelover, S.; Garcia, J.; Macías, D. Degradation of 2,4-dichlorophenoxyacetic acid (2,4-D) using cobalt-peroxymonosulfate in Fenton-like process. J. Photochem. Photobiol., A 2007, 186 (2−3), 357–363. 46. Waldemer, R. H.; Tratnyek, P. G.; Johnson, R. L.; Nurmi, J. T. Oxidation of chlorinated ethenes by heat-activated persulfate: kinetics and products. Environ. Sci. Technol. 2007, 41, 1010–1015. 47. Kronholm, J.; Metsälä, H.; Hartonen, K.; Riekkola, M.-L. Oxidation of 4-chloro-3-methylphenol in pressurized hot water/supercritical water with potassium persulfate as oxidant. Environ. Sci. Technol. 2001, 35, 3247–3251. 48. Baxendale, J. H.; Wilson, J. A. The photolysis of hydrogen peroxide at high light intensities. Trans. Faraday Soc. 1957, 53, 344–356. 49. Yu, X.-Y.; Bao, Z.-C.; Barker, J. R. Free radical reactions involving Cl•, Cl2-•, and SO4-• in the 248 nm photolysis of aqueous solutions containing S2O82- and Cl-. J. Phys. Chem. A 2004, 108, 295–308. 50. Hori, H.; Yamamoto, A.; Koike, K.; Kutsuna, S.; Osaka, I.; Arakawa, R. Perfulfate-induced photochemical decomposition of a fluorotelomer unsaturated carboxylic acid in water. Water Res. 2007, 41, 2962–2968. 51. Neta, P; Madhavan, V.; Zemel, H.; Fessenden, R. W. Rate constants and mechanism of reaction of SO4-• with aromatic compounds. J. Am. Chem. Soc. 1977, 99 (1), 163–164.


Remediation of Chemically-Contaminated Waters Using Sulfate

Recommend Documents