Oxidative transformations of environmental pharmaceuticals by Cl2, ClO2, O3, and Fe(VI): Kinetics assessment

Chemosphere 73 (2008) 1379–1386

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Review

Oxidative transformations of environmental pharmaceuticals by Cl2, ClO2, O3, and Fe(VI): Kinetics assessment Virender K. Sharma * Chemistry Department, Florida Institute of Technology, 150 West University Boulevard, Melbourne, FL 32901, USA

a r t i c l e

i n f o

Article history: Received 2 June 2008 Received in revised form 23 August 2008 Accepted 25 August 2008 Available online 11 October 2008 Keywords: Chlorination Ozonation Chlorine dioxide Ferrate Drugs Antibiotics

a b s t r a c t Several pharmaceuticals have been detected globally in surface water and drinking water, which indicate their insufficient removal from water and wastewater using conventional treatment methods. This paper reviews the kinetics of oxidative transformations of pharmaceuticals (antibiotics, lipid regulators, antipyretics, anticonvulsants, and beta-blockers) by Cl2, ClO2, O3, and ferrate(VI) ðFeVI O2 4 ; FeðVIÞÞ under treatment conditions. In the chlorination of sulfonamide antibiotics, HOCl is the major reactive Cl2 species whereas in the oxidation by Fe(VI), HFeO 4 is the dominant reactive species. Both oxidation processes can oxidize sulfonamides in seconds at a neutral pH (t1/2 6 220 s; 1 mg L1 HOCl or K2FeO4). The reactivity of O3 with pharmaceuticals is generally higher than that of HOCl (kapp,pH 7 (O3) = 1–107 M1 s1; kapp,pH 7 (HOCl) = 102–105 M1 s1). Ozone selectively oxidizes pharmaceuticals and reacts mainly with activated aromatic systems and non-protonated amines. Oxidative transformation of most pharmaceuticals by O3 occurs in seconds (t1/2 6 100 s; 1 mg L1 O3) while half-lives for oxidations by HOCl differ by at least two orders of magnitude. Ozone appears to be efficient in oxidizing pharmaceuticals in aquatic environments. The limited work on Fe(VI) shows that it can also potentially transform pharmaceuticals in treatment processes. Ó 2008 Elsevier Ltd. All rights reserved.

Contents 1. 2.

3.

4.

Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Aqueous chemistry of oxidants . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.1. Chlorine . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.2. Chlorine dioxide . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.3. Ozone . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.4. Ferrate(VI) . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Oxidation kinetics . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.1. Antibiotics . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.1.1. Sulfonamides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.1.2. Macrolides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.1.3. b-Lactums . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.1.4. Fluoroquinolones . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.1.5. Others . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.2. Fibrate lipid regulators and metabolites . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.3. Antipyretics and non-steroidal anti-inflammatory drugs . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.4. Anticonvulsants and anti-anxiety agents . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.5. Beta blockers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Conclusions. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Acknowledgments . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Appendix A. Supplementary material . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

* Tel.: +1 321 674 7310; fax: +1 321 674 8951. E-mail address: vsharma@fit.edu 0045-6535/$ - see front matter Ó 2008 Elsevier Ltd. All rights reserved. doi:10.1016/j.chemosphere.2008.08.033

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1. Introduction

Table 1 Redox potentials for the oxidants/disinfectants used in water treatment

Micropollutants in the aquatic environment are of great concern because of their influence on freshwater systems (Schwarzenbach et al., 2006). Increased levels of micropollutants in surface and groundwaters may cause long term effects on aquatic ecology and on human health. Among emerging micropollutants, pharmaceuticals in the aquatic environment have drawn attention in the last decade (Halling-Sorensen et al., 1998; Cunningham et al., 2006; Khetan and Collins, 2007). Pharmaceuticals are designed to treat and prevent diseases and their sale has increased by 62% over five years (2000–2004) in the United States (Khetan and Collins, 2007). Because of growth and an inverting age structure of the population, increased use of pharmaceuticals is expected in the future. The intake of pharmaceuticals leads to the adsorption, distribution, metabolism, and ultimately excretion of the original or modified drugs from the body. Significant proportions of molecules are excreted in unmetabolized form or as active metabolites through urine, which may or may not be treated (Halling-Sorensen et al., 1998; Ternes et al., 2004). Additionally, unused medicines and drug-containing waste from manufacturing facilities may also contribute to environmental contamination. The fates of human and veterinary pharmaceuticals after excretion are different (HallingSorensen et al., 1998). Excreted human drugs pass through sewage treatment plants (STP) before entering streams or rivers. Comparatively, veterinary pharmaceuticals do not undergo STP treatment and may directly enter soil and groundwater. Animal manure containing these pharmaceuticals is usually applied to soil directly and may cause contamination of surface water and groundwater by runoff after rainfall. STP are not designed to treat pharmaceuticals, hence plant treatment processes do not adequately remove these compounds (Miao et al., 2004; Ternes et al., 2004). A wide range of biologically active compounds including antibiotics, painkillers, anticonvulsants, lipid regulators, beta-blockers, cytostatic drugs, and antihistamines have been detected in the range of ng L1 to lg L1 in STP effluents and in surface waters (Kolpin et al., 2002; Andreozzi et al., 2003; Snyder et al., 2003; Anderson et al., 2004; Miao et al., 2004; Ternes et al., 2004; Jiang et al., 2005; Westerhoff et al., 2006; Batt et al., 2007; Watkinson et al., 2007). Unfortunately, mixtures of pharmaceuticals even at ng L1 can inhibit cell proliferation (Pomati et al., 2006). In general, drinking water utilities abstract water from various sources such as ground water, rivers, streams, springs, or lakes in a watershed; small communities generally receive water from aquifers, while large metropolitan areas receive water from surface sources. In most cases, source waters require treatment before use in order to meet national quality standards for regulated compounds. Pharmaceutical compounds are not currently regulated but their treatment is desired. Pharmaceuticals can be treated by membrane filtration (nanofiltration or reverse osmosis) or filtration over activated carbon (Benner et al., 2008). The adsorption or retention capacity of these methods decreases with operation time due to buildup of organic matter, which causes clogging of filters (Schwarzenbach et al., 2006). Moreover, membrane processes and activated carbon adsorption are energy and/or material intensive for application to wastewater treatment. Inherent tests of antibiotics proved an occurrence of ultimate biodegradation of Penicillin G (Gartiser et al., 2007). In this study, certain ultimate biodegradation of amoxicillin, imipenem, and nystatin was observed (Gartiser et al., 2007). In testing of biological degradation of pharmaceuticals in municipal wastewater treatment, more than 90% transformation occurred for ibuprofen, paracetamol, 17bestradiol, and estrone (Joss et al., 2006). UV light irradiation techniques for the disinfection of drinking water and purification of

Oxidant

Reaction

Ferrate(VI)

FeO2 4 FeO2 4

Ozone Hypochlorite Chlorine Chlorine dioxide

Eo (V) +



3+

+ 8H + 3e , Fe + 4H2O + 4H2O + 3e , Fe(OH)3 + 5OH O3 + 2H+ + 2e , O2 + H2O O3 + H2O + 2e , O2 + 2OH HClO + H+ + 2e , Cl + H2O ClO + H2O + 2e , Cl + 2OH Cl2(g) + 2e , 2Cl  ClO2 ðaqÞ þ e () ClO2

2.20 0.70 2.08 1.24 1.48 0.84 1.36 0.95

wastewater (Hijnen et al., 2006) can induce transformation of pharmaceuticals (Andreozzi et al., 2003; Boreen et al., 2004). Advanced oxidation processes such as O3/H2O2, UV/H2O2, Fenton/ photo-Fenton, and UV/TiO2 have also been studied to degrade pharmaceuticals and results are recently reviewed (Ikehata et al., 2006; Esplugas et al., 2007). The use of chemical oxidants before or after biological treatment may be a feasible approach for treating water. Chemicals selectively oxidize pharmaceuticals to readily biodegradable and less toxic compounds. In treatment systems, Cl2, HOCl, ClO2, and O3 are frequently applied for oxidative treatments because of their high reduction potentials (Dodd et al., 2006; Ikehata et al., 2006; Esplugas et al., 2007) (Table 1). In recent years, the use of ferrate(VI) ðFeðVIÞ; FeO2 4 Þ has also been proposed (Jiang, 2007; Sharma, 2007). In acid solution, the redox potential of Fe(VI) is the highest of oxidants commonly used in water treatment (Table 1) (Jiang and Lloyd, 2002; Sharma, 2002). However, O3 is a more powerful oxidant in basic solution compared to other oxidants (Table 1). This study gives an overview of the aqueous chemistry of different oxidants and their reaction kinetics with pharmaceuticals. The structures of pharmaceuticals are presented in the Supplementary material (Figs. S1–S9). The rates of the reactions are discussed based on reactive functional groups. The summary of rate constants will provide information on the dose consumption of an oxidant if matrix component of water compete for applied oxidant in treatment processes. 2. Aqueous chemistry of oxidants 2.1. Chlorine The reactivity of compounds with Cl2 depends on the speciation of Cl2 as a function of pH (Fig. 1). HOCl and OCl are both present in the pH range of 6–9. HOCl is the major reactive Cl2 species in oxidation processes. The kinetics of the Cl2 reaction with compounds is first-order in the [HOCl]Total and first-order in the total concentration of compound, i.e. overall second-order. The reactivities of HOCl and OCl for a particular compound vary significantly; therefore, the second-order rate constants (k) vary with pH in chlorination reactions (Deborde and Gunten, 2008). The reactivity of Cl2 with inorganic molecules is generally derived from an initial electrophilic attack of HOCl. The k for organic compounds varies from <0.1 to 109 M1 s1 and possible pathways of reactions include oxidation, addition, and electrophilic substitutions (Deborde and Gunten, 2008). 2.2. Chlorine dioxide Chlorine dioxide is a stable free radical, a powerful oxidant, and does not produce trihalomethanes. Chlorine dioxide decomposes slowly in neutral aqueous solution (Odeh et al., 2002). However, its decay is accelerated in basic solution and kinetic studies have shown three concurrent pathways: pathway 1 is first-order in

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2.4. Ferrate(VI)

1.0

Fraction of Species

0.8

HFeO4-

Fe(VI) is easily produced by oxidizing a basic solution of Fe(III) salt by hypochlorite and O3 (Thompson et al., 1951; Perfliev et al., 2006). Fe(VI) ion has a characteristic violet color much like that of permanganate in aqueous solutions. The spectra of aqueous Fe(VI) solutions show one maximum at 510 nm (e = 1150 ± 25 M1 cm1) and a shoulder between 275 and 320 nm. The decomposition rate is strongly dependent on the initial Fe(VI) concentration, pH, temperature, and to some extent on the surface characteristics of the hydrous iron oxide formed upon reaction. Dilute solutions of Fe(VI) are more stable than concentrated solutions. The spontaneous decomposition of Fe(VI) in water forms molecular oxygen (Eq. (1)) (Sharma, 2002).

OCl-

0.6

0.4

2-

HOCl

FeO4

0.2

3þ 2FeO2 þ 3=2O2 þ 10 OH 4 þ 5H2 O ! 2Fe

At neutral and alkaline pH, hydrogen ferrate ion to give ferrate ion ðFeO2 4 Þ (Eq. (2))

0.0 5.0

6.0

7.0

8.0

9.0

10.0



Fig. 1. The species of aqueous Cl2 and Fe(VI) as a function of pH at 25 °C.



[ClO2], pathway 2 is also first-order in [ClO2] and forms ClO2 ; and pathway 3 is second-order in [ClO2] and produces equal amounts   of ClO2 and ClO3 . The use of ClO2 is restricted to high quality water such as treated surface water (Gates, 1998). Dosing of ClO2 must be kept low; for example, in the United States, dosages ranging from 1.0 to 1.4 mg L1 are used mainly for the preoxidation of surface water (Gates, 1998). Importantly, reduction of ClO2 produces the  ClO2 ion, which is considered a blood poison. Higher dosages of ClO2 (>1.4 mg L1) are likely to produce chlorite levels that exceed the USEPA standard of 1 mg L1. The reactivity of ClO2 with inorganic and organic compounds obeys first-order kinetics with respect to ClO2 and is also first-order with respect to the compound (Hoigne and Bader, 1994). The k varies over a wide range (105–105 M1 s1). The reactivity of ClO2 with Fe(II), O3, and H2O2 is high (Hoigne and Bader, 1994; Wang et al., 2004). Under water treatment conditions, aromatic, hydrocarbons, carbohydrates, and molecules containing primary and secondary amines, aldehydes, and acetone are un-reactive. However, phenolic and tertiary amino group containing compounds are reactive with ClO2. The reactivity with these compounds is governed by the pH; phenoxide ion and neutral species of the amine are much more reactive than either the neutral phenol or the protonated amine (Hoigne and Bader, 1994; Tratnyek and Hoigne, 1994). 2.3. Ozone Ozone is unstable and has a half-life in the range of seconds to hours depending on the water quality (Gunten, 2003). The decomposition of O3 is initially fast, followed by first-order kinetics, producing a strong oxidant, OH. Therefore, O3 oxidation processes consider both O3 and OH species. Ozone is a very selective oxidant while OH reacts indiscriminately with organic molecules (Gunten, 2003). The k for the reactivity of O3 with compounds varies between 1 and 107 M1 s1 at pH 7. The environmentallyrelevant inorganic compounds such as Fe(II), Mn(II), H2S, cyanide, and nitrite react fast with O3 through an oxygen transfer mechanism. In reaction with organic compounds, O3 attacks double bonds, activated aromatic systems, and neutral amines (Ikehata et al., 2006).

dissociates

HFeO4 () ðHþ þ FeO2 4 ; pK a;HFeO4 ¼ 7:23 at 25 C ðSharma et al:; 2001Þ:

pH

ð1Þ ðHFeO 4Þ

ð2Þ

The speciation of the Fe(VI) ion is shown in Fig. 1. The by-product of Fe(VI) is non-toxic ferric ion, Fe(III) (Eq. (2)). This fact makes Fe(VI) an ‘‘environmentally friendly” oxidant (Sharma, 2002; Jiang and Lloyd, 2002). Additionally, the ferric oxide produced from ferrate(VI) acts as an effective coagulant that is suitable for removal of metals, non-metals, radionuclides, and humic acids (Jiang and Wang, 2003; Lee et al., 2002; Sharma, 2002; Yngard et al., 2008). Fe(VI) can also achieve disinfection at relatively low dosages, over wide ranges of pH (Sharma, 2007). The reactions of Fe(VI) with pollutants are first-order for each reactant. The oxidations are strongly pH dependent and the k vary from 3.0  102 to 1.7  107 M1 s1 (Rush et al., 1996; Sharma, 2002). The reactions of ferrate(VI) with sulfur- and nitrogen-containing inorganic compounds are completed in seconds to minutes with formation of less toxic products than parent compounds (Sharma et al., 2005). The reactivity of Fe(VI) with organic compounds shows that Fe(VI) is a strong but selective oxidant (Sharma, 2002). Fe(VI) selectively oxidizes primary and secondary alcohols to aldehydes (not acids) and ketones, respectively, and primary amines to aldehydes. 3. Oxidation kinetics 3.1. Antibiotics 3.1.1. Sulfonamides The reactivity of sulfonamides with HOCl and Fe(VI) has been determined as a function of pH. The rates were found to be first-order with each reactant and the k varies with pH (Fig. 2). The k increases for reactions of sulfonamides with HOCl with increasing pH. Comparatively, rate constants decrease non-linearly with increasing pH for Fe(VI) as an oxidant. The pH dependence of k can be attributed to the combined speciation effects of oxidant and sulfonamides. Sulfonamides have two dissociation constants, one corresponding to protonation of the aniline N ðSHþ 2 () Hþ þ SH; pK a;SH2 ¼ 1:5—2:9 and the other involves the protonation of sulfonamide ðSH () Hþ þ S ; pK a;SH ¼ 5:0—7:4Þ (Table 2). Four 2  reactions, HOCl and OCl (or HFeO 4 and FeO4 ) with SH, and S are possible in the reactions of sulfonamides with oxidants in the studied pH range (6.1–9.7) (Chamberlain and Adams, 2006; Sharma et al., 2006). Since HOCl has been shown to be much more reactive than OCl with sulfonamides (Gallard and Gunten, 2002; Dodd and Huang, 2007), only two specific rate constants, kHOCl,SH and kHOCl;S were considered to explain the variation of rates for the oxidation

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V.K. Sharma / Chemosphere 73 (2008) 1379–1386

The values of kapp vary from 2.0  102 to 5.63  103 M1 s1 for the oxidation of sulfonamides by HOCl at pH 7.0 (Table 2). The range of kapp is from 0.85 to 1.50  103 M1 s1 for the oxidation by Fe(VI) under similar conditions (Table 2). The reaction half-lives for the oxidation of sulfonamides by both oxidants are estimated to be up to 200 s at a dose of 1 mg L1 (Table 2). The oxidation of sulfonamide by other oxidants, ClO2 and O3, has been determined only for sulfamethoxazole (SMX) (Huber et al., 2003, 2005a). The kapp values were found to be 2.2  103 and 2.5  106 M1 s1 for oxidation of SMX by ClO2 and O3, respectively, at pH 7.0 and 20 °C. An order of reactivity is thus HOCl < Fe(VI) < ClO2 < O3. At a dose of 1 mg L1, SMX can be transformed by O3 in less than a second while other oxidants would require seconds. (See Table 3).

k, M-1s-1

104

103

HOCl 102

3.1.2. Macrolides The reactivity of macrolide antibiotics has been determined with ClO2 and O3 (Huber et al., 2003, 2005b; Qiang et al., 2004; Dodd et al., 2006). The k values were determined as a function of pH, which increased with an increase in pH. The pH dependence was explained by protonation of a tertiary group of macrolides. Oxidants selectively attack the tertiary amino group of the macrolides with a preference for deprotonated amines. The apparent rate constants for both oxidants are of the order of 104–105 M1 s1 at pH 7.0 and 20 °C. For water treatment conditions of pH 7–8, halflives for the transformation of macrolides by ClO2 and O3 are <1 s at a dose of 1 mg L1 (Fig. 3).

k, M-1s-1

103

102

101 Fe(VI)

100 6.0

7.0

8.0

9.0

3.1.3. b-Lactums The reactivity of b-lactums (amoxicillin (AM), cefalexin (CP), and Penicillin G (PG)) with O3 has been studied (Andreozzi et al., 2003; Dodd et al., 2006). The k for the oxidation of AM varies from 4.0  103 at pH 2.5 to 6.0  106 M1 s1 at pH 7.0. PG contains only a thioether moiety and the reactivity is thus expected to be independent of pH. CP showed no pH dependence up to pH 7 and a slight increase was observed at pH > 7 due to possible attack on the primary amine (pKa = 7.1). AM had the highest reactivity with O3 among the studied b-lactums (Table 2). Half-lives for the reactivity of AM are thus in milliseconds. CP and PG react with O3 relatively fast and half-lives are <1 s and <10 s, respectively (Fig. 3).

10.0

pH Fig. 2. The pH dependence of the k for the reactions of HOCl and Fe(VI) with sulfonamides at 25 °C. s – SDM; h – STZ; D – SMN; e – SMR; . – SMX; j – SML; d-SIL. See Table 2 for abbreviated sulfonamide names. Data for the reactivity with HOCl and Fe(VI) were taken from Chamberlain and Adams (2006) and Sharma et al. (2006), respectively.

of sulfonamide by free Cl2 (Chamberlain and Adams, 2006) (Table 2). The anionic form (S) of sulfonamides was determined to react faster with HOCl than the neutral form (SH). However, the protonated form of Fe(VI) ðHFeO 4 Þ was estimated to react faster with SH species of sulfonamides than the S species (Table 2). The estimated rate constants in Table 2 suggest that the HFeO 4 species re act faster than the deprotonated FeO2 4 species. HFeO4 has a larger spin density on the oxo ligands than does FeO2 4 , which increases the oxidation ability of the protonated Fe(VI) species relative to the deprotonated species. The fraction of HFeO 4 species increases with decrease in pH (Fig. 1) and thus contributes to an increase in the rate with a decrease in pH (Fig. 2).

3.1.4. Fluoroquinolones The Cl2 species reactivity towards ciprofloxacin (CF) and enrofloxacin (EF) has been determined as a function of pH (Dodd et al., 2006). Chlorine is expected to attack mainly the basic amine group of fluoroquinolones. HOCl reacts rapidly with the secondary N(4) amine site of CF. EF contains a tertiary amine group and has a smaller rate constant than CF with HOCl (Table 2). This gives a large variation in half-lives in the reactivity of CF and EF with HOCl at neutral pH (Fig. 3).

Table 2 Thermodynamic parameters of sulfonamides and second-order rate constants for reactions of sulfonamides with HOCl and Fe(VI) at 25 °C Sulfonamide

pK1

pK2

25 °C Sulfamethoxazole (SMX) Sulfamethazine (SMN) Sulfamethizole (SML) Sulfadimethoxine (SDM) Sulfisoxazole (SIL) Sulfamerazine (SMR) Sulfathiazole (STZ)

1.7 2.3 2.1 2.9 1.5 2.1 2.0

DHSH (kJ mol

5.6 7.4 5.3 6.1 5.0 6.9 7.1

1

)

36.9 ± 3.0 – 39.4 ± 2.4 – 21.3 ± 1.1 – –

kHOCla (103 M1 s1) SH S

kapp (103 M1 s1) pH 7.0

1.10 1.45 0.36 10.3 – 0.77 3.47

0.65 0.79 0.20 5.63 – 0.42 1.89

2.40 2.89 0.73 20.7 – 1.55 6.94

Half-lives for oxidant dose of 1 mg L1 HOCl and K2FeO4 at pH 7.0 and 25 °C. a From Chamberlain and Adams (2006). b From Sharma et al. (2006).

(s)

kHFeO4b (103 M1 s1) SH S

kFeO4b (101 M1 s1) SH

kappb (103 M1 s1) pH 7.0

(s)

56 46 181 6 – 86 20

30.0 1.90 22.3 18.8 11.0 – –

0.12 22.5 – – – – –

1.50 0.87 0.64 – 0.85 – –

91 157 214 – 161 – –

t1/2

0.17 0.55 0.22 0.38 2.42 – –

t1/2

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V.K. Sharma / Chemosphere 73 (2008) 1379–1386 Table 3 Second-order rate constants for the reactions of pharmaceuticals with HOCl, ClO2, and ozone at pH 7 Pharmaceutical

kHOCl (M1 s1)

Temp. (°C)

Reference

kClO2 (M1 s1)

Temp. (°C)

Reference

Azithromycin Clarithromycin Lincomycin Roxithromycin

– – – –

– – – –

– – – –

– – – 1.4  104

– – – 20

1.1  105 4.0  104 3.3  105 6.3  104

20 20

Tylosin tartrate Amoxicillin

– –

– –

– –

– –

– –

– – – Huber et al. (2005b) – –

5.1  105 6.0  106

20 25

Cefalexin Penicillin G Ciprofloxacin Enfrofloxacin Trimethoprim Tetracycline Spectinomycin Triclosan Bezafibrate

– – 7.6  105 5.1  102 5.8  101 – – 4.7  102 –

– – 22 25 25 – – – –

– – Dodd et al. (2005) Dodd et al. (2005) Dodd and Huang (2007) – – Rule et al. (2005) –

– – – – – – – – <1.0  102

– – – – – – – – 20

8.7  104 4.8  103 1.9  104 1.5  104 2.7  105 1.9  106 1.3  106 3.8  107 5.9  102

20 20 20 20 20 20

Clofibric

–

–

–

<2.0  101

20

Gemfibrozil

9.0  101a

23

Pinkston and Sedlak (2004) –

–

–

– – – – – – – – Huber et al. (2005b) Huber et al. (2005b) –

Diclofenac

–

–

4

1.0  10

Ibuprofen

–

–

–

<1.0  10

20

Naproxen

2.5  100a

23

–

–

Huber et al. (2005b) Huber et al. (2005b) –

Paracetamol

1.5  102a

23

–

–

–

2

<1.5  10

20

<2.5  102

20

and Sedlak

– –

– –

Huber et al. (2005b) Huber et al. (2005b) – –

and Sedlak

–

–

–

and Sedlak

–

–

–

and Sedlak

–

–

–

2

Carbamazepine

–

–

Pinkston and Sedlak (2004) Pinkston and Sedlak (2004) –

Diazepam

–

–

–

Acebutanol Atenolol

– 1.3  102a

– 23

Metoprolol

1.4  102a

23

Nadolol

2.3x101a

23

Propranolol

6.1  101a

23

– Pinkston (2004) Pinkston (2004) Pinkston (2004) Pinkston (2004)

a

20

kO3 (M1 s1)

Temp. (°C)

– – 1.0  10

20

23 20

Dodd et al. (2006) Lange et al. (2006) Qiang et al. (2004) Dodd et al. (2006) Dodd et al. (2006) Andreozzi et al. (2005) Dodd et al. (2006) Dodd et al. (2006) Dodd et al. (2006) Dodd et al. (2006) Dodd et al. (2006) Dodd et al. (2006) Adams et al. (2002) Suarez et al. (2007) Huber et al. (2003)

–

–

–

–

6

20

Huber et al. (2003)

0

20

Huber et al. (2003)

9.6  10 –

–

–

1.41  103 3.0  10

Reference

5

7.5  101

20

Andreozzi et al. (2005) Huber et al. (2003)

20

Huber et al. (2003)

1.9  103 1.7  103

20–22 20–22

Benner et al. (2008) Benner et al. (2008)

2.0  103

20–22

Benner et al. (2008)

–

–

–

1.0  105

20–22

Benner et al. (2008)

The values were calculated from the half-lives given for a total Cl2 concentration of 10 mg L1.

The reactivity of CF and EF with O3 exhibited strong pH dependence (Dodd et al., 2006). The k values of the reactions were determined to be in a range from 4  102 to >105 M1 s1 between pH 3 and 8 for CF. Comparatively, EF reacts much faster than CF in a similar pH range. This is possibly related to the presence of a higher fraction of reactive anionic species in the case of EF compared to CF (pK2(EF) = 7.7; pK2(CF) = 8.8). For example, the molar fraction of anionic EF is 0.15 compared to 0.01 for CF. Both fluoroquinolones react quickly with O3 and have half-lives of less than 5 s at a dose of 1 mg L1 O3 (Fig. 3). 3.1.5. Others The reaction of trimethoprim (TMP) with Cl2 is governed by its 2,4-diamino-5-methylpyrimidinyl moiety at pH 7 and by its 3,4,5trimethoxybenzyl moiety at pH < 7 (Dodd and Huang, 2007). Triclosan reacts very rapidly with O3 with an initial attack on the phenolic moiety. Both reactions are about one-order of magnitude apart in k (Table 2). The half-lives are <100 s and <1000 s for TMP and triclosan, respectively, at a dose 1 mg L1 HOCl. Ozone reacts rapidly with these molecules and the other antibiotics, tetracycline and spectinomycin. The values of k were determined to be >105 M1 s1 (Adams et al., 2002; Dodd et al., 2006; Suarez

et al., 2007). The half-lives for transformation of these antibiotics were estimated to be less than 0.2 s at 1 mg L1 O3 dose. 3.2. Fibrate lipid regulators and metabolites Bezafibrate did not show any reactivity with ClO2 (Table 2). Ozone has a relatively high reactivity with bezafibrate to give t1/ 2 < 100 s (Fig. 3). No reactivity of ClO2 with clofibric was found (Table 2). The reactivity of HOCl with the aromatic ether, gemfibrozil, is very low, hence the half-life for the oxidative transformation is more than 4  104 s (Fig. 3). 3.3. Antipyretics and non-steroidal anti-inflammatory drugs Diclofenac is an aniline derivative and has relatively high rate constants with both ClO2 and O3 (Table 2). However, no pH dependence in rate constants was observed with either oxidant (Huber et al., 2003, 2005b). The high reactivity yields half-lives of <5 s for the oxidation of diclofenac by both ClO2 and O3 at a dose of 1 mg L1 (Fig. 3). No significant reaction of ClO2 with ibuprofen was observed (Table 2). Ozone, however, has some reactivity with ibuprofen, but the rate constant is low (Table 2). Hence, the

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V.K. Sharma / Chemosphere 73 (2008) 1379–1386

10-3

10-2

10-1

100

101

102

103

104

105

106

107

Propranolol Nadolol Metoprolol Atenolol Paracetamol Naproxen Gemfibrozil Triclosan Trimethoprim Enfrofloxacin Ciprofloxacin

HOCl

Propranolol Metoprolol Atenolol Acebutanol Diazepam Carbamazepine Paracetamol Ibuprofen Diclofenac Bezafibrate Triclosan Spectinomycin Tetracycline Trimethoprim Enfrofloxacin Ciprofloxacin Pencillin G Cefalexin Amoxicillin Tylosine Tartrate Roxithromycin Lincomycin Clarithromycin Azithromycin

O3

10-3

10-2

10-1

100

101

102 t1/2 (s)

103

104

105

106

107

Fig. 3. Half-lives (t1/2) of pharmaceuticals oxidation by HOCl and O3 at dose = 1 mg L1. The values of t1/2 were calculated using rate constants given in Table 3.

half-life for transformation of ibuprofen by O3 is more than 3.4  103 s (Fig. 3). The oxidation of ibuprofen with Fe(VI) is about one-order magnitude slower than with O3 (Sharma and Mishra, 2006). The reactivity of naproxen with HOCl is slow with a half-life greater than 1  104 s (Table 2). Because of the phenolic group in paracetamol, appreciable reactivity with HOCl and O3 was observed (Table 2). This resulted in half-lives of 242 and 24 s for the oxidative transformation of paracetamol by HOCl and O3, respectively. 3.4. Anticonvulsants and anti-anxiety agents Chlorine dioxide did not give any appreciable reactivity with carbamazepine and diazepam (Table 2). However, O3 reacts with both of these molecules. The rate constant for the reaction of diazepam with O3 is very low (Huber et al., 2003). Carbamazepine showed high reactivity with O3 at a neutral pH (Table 2). This is re-

lated to the attack of O3 at the double bond connecting the two phenyl moieties. Carbamazepine can thus be decomposed efficiently by O3 (t1/2 = 0.1 s; 1 mg L1). 3.5. Beta blockers Rate constants for the reaction of Cl2 and O3 with beta blockers have been investigated (Pinkston and Sedlak, 2004; Benner et al., 2008). Beta lockers contain two reactive sites: the aromatic ring and the secondary amine moiety. Since the reactions of aromatic structures are independent of solution pH, oxidation of beta blockers is explained by the pKa of the amine group. The pKas of amines are in the range of 9.2 to 9.7, hence beta blockers show increased reaction rates above pH 8. The reactivity of HOCl with beta blockers besides propranolol is slow at pH 7 (Table 2). Comparatively, the k for reactions of O3 with acetbutolol, atenolol, and propranolol were 2  103 M1 s1 (Table 2). Similar to HOCl, O3 also has high

V.K. Sharma / Chemosphere 73 (2008) 1379–1386

reactivity with propranolol. The difference in structure explains the difference in reactivity. Propranolol has a naphthalene moiety whereas other beta blockers have phenolic groups. HOCl does not seem to be a practical reagent for the oxidation of beta blockers (t1/2 = 6  102–2.8  106 s). Oxidation by O3 is highly effective with half-lives less than 20 s (Fig. 3). 4. Conclusions The reactivity of pharmaceuticals with Cl2, ClO2, O3, and Fe(VI) follows second-order kinetics. Most studies were carried out with O3 and the values of kapp,pH 7 vary from 1 to 107 M1 s1. The values of kapp,pH 7 for the oxidation by HOCl were in the range of 102 to 105 M1 s1. The half-lives of O3 with most pharmaceuticals were less than 100 s at a dose of 1 mg L1 under treatment conditions. Comparatively, half-lives for oxidations carried out by Cl2 were > 100 s. Chlorine dioxide showed slower reactivity than O3 and Cl2 (kapp,pH 7 < 102–104 M1 s1). The reactivity of HOCl and Fe(VI) with sulfonamides showed opposite behavior with pH and the values of kapp,pH 7 were higher for Fe(VI) than HOCl. The halflives for oxidation of sulfonamides using 1 mg L1 dose of these oxidants vary in a range from 20 to 200 s at a neutral pH. Overall, significant progress has been made in understanding the trends in the reactivity of pharmaceuticals with O3, however, similar studies using Cl2, ClO2, and Fe(VI) are needed to characterize and understand the behavior of pharmaceuticals with these oxidants. Acknowledgments The author wishes to thank Dr. Mary Sohn, Dr. Ria Yngard, George A. K. Anquandah, and Antoine Zufferey for their useful comments. The author also wishes to thank two anonymous reviewers for useful comments. Appendix A. Supplementary material Supplementary data associated with this article can be found, in the online version, at doi:10.1016/j.chemosphere.2008.08.033. References Adams, C., Wang, Y., Loftin, K., Meyer, M., 2002. Removal of antibiotics from surface and distilled water in conventional water treatment processes. J. Environ. Eng. 128, 253–260. Anderson, P.D., D’Aco, V.J., Shanahan, P., Chapra, S.C., Buzby, M.E., Cunningham, V.L., Duplessie, B.M., Hayes, E.P., Mastrocco, F.J., Parke, N.J., Rader, J.C., Samuelian, J.H., Schwab, B.W., 2004. Screening analysis of human pharmaceutical compounds in US surface waters. Environ. Sci. Technol. 38, 838–849. Andreozzi, R., Canterino, M., Marotta, R., Paxeus, N., 2005. Antibiotic removal from wastewaters: the ozonation of amoxicillin. J. Hazardous Mater. 112, 243–250. Andreozzi, R., Raffaele, M., Nicklas, P., 2003. Pharmaceuticals in STP effluents and their solar photodegradation in aquatic environment. Chemosphere 50, 1319– 1330. Batt, A.L., Kim, S., Aga, D.S., 2007. Comparison of the occurrence of antibiotics in four full-scale wastewater treatment plants with varying designs and operations. Chemosphere 68, 428–435. Benner, J., Salhr, E., Ternes, T., Gunten, U.V., 2008. Ozonation of reverse osmosis concentrate: kinetics and efficiency of beta blocker oxidation. Water Res. 42, 3003–3012. Boreen, A.L., Arnold, W.A., Mcnell, K., 2004. Photochemical fate of sulfa drugs in the aquatic environment: sulfa drugs containing five-membered heterocyclic groups. Environ. Sci. Technol. 38, 3933–3940. Chamberlain, E., Adams, C., 2006. Oxidation of sulfonamides, macrolides, and carbadox with free chlorine and monochloramine. Water Res. 40, 2517–2526. Cunningham, V.L., Buzby, M., Hutchinson, T., Mastrocco, F., Parke, N., Roden, N., 2006. Effects of human pharmaceuticals on aquatic life: next steps. Environ. Sci. Technol. 40, 3457–3461. Deborde, M., Gunten, U.V., 2008. Reactions of chlorine with inorganic and organic compounds during water treatment – kinetics and mechanisms: a critical review. Water Res. 42, 13–51. Dodd, M.C., Buffle, M.-O., Gunten, U.V., 2006. Oxidation of antibacterial molecules by aqueous ozone: moiety-specific reaction kinetics and applications in ozonebased wastewater treatment. Environ. Sci. Technol. 40, 1969–1977.

1385

Dodd, M.C., Huang, C-H., 2007. Aqueous chlorination of the antibacterial agent trimethoprim: reaction kinetics and pathways. Water Res. 41, 647–655. Dodd, M.C., Shah, A.D., Gunten, U.V., Huang, C.-H., 2005. Interaction of fluoroquinolone antibacterial agents with aqueous chlorine: reaction kinetics, mechanisms, and transformation pathways. Environ. Sci. Technol. 39, 7065– 7076. Esplugas, S., Bila, D.M., Krause, L.G.T., Dezotti, M., 2007. Ozonation and advanced oxidation technologies to remove endocrine disrupting chemicals (EDCs) and pharmaceuticals and personal care products (PPCPs) in water effluents. J. Hazard. Mater. 149, 631–642. Gallard, H., Gunten, U.V., 2002. Chlorination of natural organic matter: kinetics of chlorination and of THM formation. Water Res. 36, 65–74. Gartiser, S., Urich, E., Alexy, R., Kummerer, K., 2007. Ultimate biodegradation and elimination of antibiotics in inherent tests. Chemosphere 67, 604–613. Gates, D., 1998. The Chlorine Dioxide Handbook. American Water Works Association, Denver, Colorado. Gunten, U.V., 2003. Ozonation of drinking water: Part I. Oxidation kinetics and product formation. Water Res. 37, 1443–1467. Halling-Sorensen, B., Nielsen, S.N., Lanzky, P.F., Ingerslev, F., Luzhoft, H.C., Jorgensen, S.E., 1998. Occurrence, fate, and effects of human pharmaceutical substances in the environment – a review. Chemosphere 36, 357–393. Hijnen, W.A.M., Beerendonk, E.F., Medema, G.J., 2006. Inactivation credit of UV radiation for viruses, bacteria and protozoan (oo) cysts in water: a review. Water Res. 40, 3–32. Hoigne, J., Bader, H., 1994. Kinetics of reactions of chlorine dioxide (ClO2) in water – I. Rate constants for inorganic and organic compounds. Water Res. 28, 45–55. Huber, M.C., Canonica, S., Park, G.-Y., Gunten, U.V., 2003. Oxidation of pharmaceuticals during ozonation and advanced oxidation process. Environ. Sci. Technol. 37, 1016–1024. Huber, M.M., Goebel, A., Joss, A., Hermann, N., Loffler, D., McArdell, C.S., Ried, A., Siegrist, H., Ternes, T.A., Gunten, U.V., 2005a. Oxidation of pharmaceuticals during ozonation of municipal wastewater effluents: a pilot study. Environ. Sci. Technol. 39, 4290–4299. Huber, M.M., Korhonen, S., Ternes, T.A., Gunten, U.V., 2005b. Oxidation of pharmaceuticals during water treatment with chlorine dioxide. Water Res. 39, 3607–3617. Ikehata, K., Naghashkar, N.J., El-Din, M.G., 2006. Degradation of aqueous pharmaceuticals by ozonation and advanced oxidation processes: a review. Ozone Sci. Eng. 28, 353–414. Jiang, J.Q., 2007. Research progress in the use of ferrate(VI) for the environmental remediation. J. Hazard. Mater. 146, 617–623. Jiang, J.Q., Lloyd, B., 2002. Progress in the development and use of ferrate salt as an oxidant and coagulant for water and wastewater treatment. Water Res. 36, 1397–1408. Jiang, J.Q., Wang, S., 2003. Enhanced coagulation with potassium ferrate(VI) for removing humic substances. Environ. Eng. Sci. 20, 627–635. Jiang, J.Q., Yin, Q., Zhou, J.L., Pearce, P., 2005. Occurrence and treatment trials of endocrine disrupting chemicals (EDCs) in wastewater. Chemosphere 61, 544– 550. Joss, A., Zabczynski, S., Gobel, A., Hoffmann, B., Loffler, D., McArdell, C.S., Ternes, T.A., Thomsen, A., Siegrist, H., 2006. Biological degradation of pharmaceuticals in municipal wastewater treatment: proposing a classification scheme. Water Res. 40, 1686–1696. Khetan, S.K., Collins, T.J., 2007. Human pharmaceuticals in the aquatic environment: a challenge to green chemistry. Chem. Rev. 107, 2319–2364. Kolpin, D.W., Furlong, E.T., Meyer, M.T., Thurman, E.M., Zaugg, S.D., Barber, L.B., Buxton, H.T., 2002. Pharmaceuticals, hormones, and other organic wastewater contaminants in US streams, 1999–2000: a national reconnaissance. Environ. Sci. Technol. 36, 1202–1211. Lange, F., Cornelissen, S., Kubac, D., Sein, M.M., Sonntag, J.V., Hannich, C.B., Golloch, A., Heipieper, H.J., Moder, M., Sonntag, C.V., 2006. Degradation of macrolide antibiotic by ozone: a mechanistic case study with clarithromycin. Chemosphere 65, 17–23. Lee, Yunho, Um, I-H., Yoon, J., 2002. Arsenic(III) oxidation by iron(VI) (ferrate) and subsequent removal of arsenic(V) by iron(III) coagulation. Environ. Sci. Technol. 37, 5750–5756. Miao, X., Bishay, F., Chen, M., Metcalfe, C.D., 2004. Occurrence of antimicrobials in the final effluents of wastewater treatment plants in Canada. Environ. Sci. Technol. 38, 3533–3541. Odeh, I.N., Francisco, J.S., Margerum, D.W., 2002. New pathways for chlorine dioxide decomposition in basic solution. Inorg. Chem. 41, 6500–6506. Perfliev, Y.D., Benko, E.M., Pankratov, D.A., Sharma, V.K., Dedushenko, S.K., 2006. Formation of iron(VI) in ozonolysis of iron(III) in alkaline solution. Inorg. Chim. Acta 360, 2789–2791. Pinkston, K.E., Sedlak, D.L., 2004. Transformation of aromatic ether- and aminecontaining pharmaceuticals during chlorine disinfection. Environ. Sci. Technol. 38, 4019–4025. Pomati, E., Castiglioni, S., Zuccato, E., Fanelli, R., Vigetti, D., Rossetti, C., Calamari, D., 2006. Effects of a complex mixture of therapeutic drugs at environmental levels on human embryonic cells. Environ. Sci. Technol. 40, 2442–2447. Qiang, Z., Adams, C., Rao, S., 2004. Determination of ozonation rate constants for lincomycin and spectinomycin. Ozone Sci. Eng. 26, 525–537. Rule, K.L., Ebbett, V.R., Vikesland, P.J., 2005. Formation of chloroform and chlorinated organics by free-chlorine-mediated oxidation of triclosan. Environ. Sci. Technol. 39, 3176–3185.

1386

V.K. Sharma / Chemosphere 73 (2008) 1379–1386

Rush, J.D., Zhao, Z., Bielski, B.H.J., 1996. Reaction of ferrate(VI)/ferrate(V) with hydrogen peroxide and superoxide anion: a stopped-flow and premix pulse radiolysis study. Free Rad. Res. 24, 187–198. Schwarzenbach, R.P., Escher, B.I., Fenner, K., Hofstetter, T.B., Johnson, C.M., Gunten, U.V., Wehrli, B., 2006. The challenge of micropollutants in aquatic systems. Science 313, 1072–1076. Sharma, V.K., 2002. Potassium ferrate(VI): an environmentally friendly oxidant. Adv. Environ. Res. 6, 143–156. Sharma, V.K., 2007. A review of disinfection performance of Fe(VI) in water and wastewater. Water Sci. Technol. 55 (1–2), 225–230. Sharma, V.K., Burnett, C.R., Millero, F.J., 2001. Dissociation constants of monoprotic ferrate(VI) ion in NaCl media. Phys. Chem. Chem. Phys. 3, 2059–2062. Sharma, V.K., Kazama, F., Jiangyong, H., Ray, A.K., 2005. Ferrates (iron(VI) and iron(V)) as environmentally friendly oxidants and disinfectants. J. Water Health 3, 42–58. Sharma, V.K., Mishra, S.K., 2006. Ferrate(VI) oxidation of ibuprofen: a kinetic study. Environ. Chem. Lett. 3, 182–185. Sharma, V.K., Mishra, S.K., Nesnas, N., 2006. Oxidation of sulfonamide antimicrobials by ferrate(VI) ½FeVI O2 4 . Environ. Sci. Technol. 40, 7222–7227. Snyder, S.A., Westerhoff, P., Yoon, Y., Sedlak, D.L., 2003. Pharmaceuticals, personal care products, and endocrine disruptors in water: implications for water industry. Environ. Eng. Sci. 20, 449–469.

Suarez, S., Dodd, M.C., Omil, F., Gunten, U.V., 2007. Kinetics of triclosan oxidation by aqueous ozone and consequent loss of antibacterial activity: relevance to municipal wastewater ozonation. Water Res. 41, 2481–2490. Ternes, T.A., Joss, A., Siegrist, H., 2004. Scrutinizing pharmaceuticals and personal care products in wastewater treatment. Environ. Sci. Technol. 20, 393A– 399A. Thompson, G.W., Ockerman, L.T., Schreyer, J.M., 1951. Preparation and purification of potassium ferrate(VI). J. Amer. Chem. Soc. 73, 1379–1381. Tratnyek, P.G., Hoigne, J., 1994. Kinetics of reactions of chlorine dioxide (ClO2) in water – II. Quantitative structure–activity relationships for phenolic compounds. Water Res. 28, 57–66. Wang, L., Odeh, I.N., Margerum, D.W., 2004. Chlorine dioxide reduction by aqueous iron(II) through outer-sphere and inner-sphere electron-transfer pathways. Inorg. Chem. 43, 7545–7551. Watkinson, A.J., Murby, E.J., Costanzo, S.D., 2007. Removal of antibiotics in conventional and advanced wastewater treatment: implications for environmental discharge and wastewater recycling. Water Res. 41, 4164–4176. Westerhoff, P., Yoon, Y., Snyder, S., Wert, E., 2006. Fate of endocrine-disruptor, pharmaceutical, and personal care product chemicals during simulated drinking water treatment processes. Environ. Sci. Technol. 40, 6649–6663. Yngard, R.A., Sharma, V.K., Philips, J., Zboril, R., 2008. Ferrate(VI) oxidation of weak acid dissociable cyanides. Environ. Sci. Technol. 42, 3005–3010.