Transformation of Tetracycline Antibiotics and Fe(II) and Fe(III

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Transformation of Tetracycline Antibiotics and Fe(II)/(III) Species Induced by Their Complexation Hui Wang, Hong Yao, Peizhe Sun, Desheng Li, and Ching-Hua Huang Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.5b03696 • Publication Date (Web): 30 Nov 2015 Downloaded from http://pubs.acs.org on December 7, 2015

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Environmental Science & Technology

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Transformation of Tetracycline Antibiotics and Fe(II)/(III)

2

Species Induced by Their Complexation

3 4

Hui Wang†,‡, Hong Yao†,§,*, Peizhe Sun‡, Desheng Li†,§ and Ching-Hua Huang‡,*

5 6

†

7

Beijing 100044, Peoples Republic of China

8

‡

9

Georgia 30332, United States

Department of Municipal and Environmental Engineering, Beijing Jiaotong University,

School of Civil and Environmental Engineering, Georgia Institute of Technology, Atlanta,

10

§

11

Beijing Jiaotong University, Beijing 100044, Peoples Republic of China

Beijing Key Laboratory of Aqueous Typical Pollutants Control and Water Quality Safeguard,

12 13 14 15

*

16

*

Phone: 86-10-5168 2157; e-mail: [email protected] Phone: 404-894-7694; fax: 404-385-7087; e-mail: [email protected]

17 18 19 20

Revised manuscript submitted to Environmental Science & Technology

21 22

November 27, 2015 1 ACS Paragon Plus Environment


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ABSTRACT

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Tetracycline antibiotics (TCs) are frequently detected micropollutants and are known to have

25

strong tendency to complex with metal ions such as Fe(II) and Fe(III) in the aquatic

26

environment. Experiments with Fe(II) and TCs showed that complexation of Fe(II) with

27

tetracycline (TTC), oxytetracycline (OTC) or chlorotetracycline (CTC) could lead to

28

accelerated oxidation of Fe(II) and promoted degradation of TCs simultaneously. The

29

reaction started with complexation of Fe(II) with TC, followed by oxidation of the Fe(II)-TC

30

complex by dissolved oxygen to generate a Fe(III)-TC complex and reactive oxygen species

31

(ROS). The ROS (primarily ·OH) then degraded TC. The oxidation rate constants of Fe(II) in

32

the FeII-H2L and FeII-HL complexes were 0.269 min-1 and 1.511 min-1, respectively, at

33

ambient conditions (pH 7, 22°C and PO2 0.21 atm), which were about 60 and 350 times of the

34

oxidation rate of uncomplexed Fe(II). Humic acids (HA) compete with TCs for Fe(II) but the

35

effect was negligible at moderate HA concentrations (≤ 10 mg⋅L-1). Experiments with Fe(III)

36

and TCs showed that complexation of Fe(III) with TC could generate oxidized TC and Fe(II)

37

without the need of oxygen, at a relatively slower rate compared to the reaction involving

38

Fe(II), O2 and TCs. These findings indicate the mutually influenced environmental

39

transformation of TCs and Fe(II)/Fe(III) induced by their complexation. These newly

40

identified reactions could play an important role in affecting the environmental fate of TCs

41

and cycling of Fe(II)/Fe(III) in TCs-contaminated water and soil systems.

42 43

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INTRODUCTION

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Tetracycline antibiotics (TCs) are widely used in human medicine, and in livestock and

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aquaculture industry for controlling diseases and/or promoting growth due to their prominent

47

therapeutic values.1, 2 In 2011, according to the U.S. Food and Drug Administration (FDA),

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the domestic annual total sales of TCs were at 5,642,573 kg, making them the most selling

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pharmaceutical group.3 TCs were also reported to be the most widely used antibacterial

50

compounds in the U.K. in 2000 at the usage of 16,268 kg.4 Due to their high usage, TCs may

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enter the environment through discharge of wastewater effluent and animal manure, leading

52

to potential negative impact on aquatic and terrestrial organisms.4 Recent research has shown

53

the omnipresence of TCs as micropollutants in the environment at the concentrations around

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4.58 mg⋅kg-1 in animal dung samples, 86-199 µg⋅kg-1 in soils, and 0.13-0.51 µg⋅L-1 in surface

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waters.5, 6

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TC compounds, such as tetracycline (TTC), oxytetracycline (OTC) and chlorotetracycline

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(CTC), consist of a four-ring system with multiple ionizable functional groups (Figure 1).7

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Owing to the multiple O- and N-containing moieties, TCs have a strong tendency to complex

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with metals,8 which may greatly impact TCs’ fate in the aquatic environment, including

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interactions with mineral surfaces,9, 10 redox reactions and photodegradation.11, 12 For example,

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interactions of TTC with metal cations (e.g., Cd(II), Cu(II) and Pb(II)) promoted its

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adsorption in soils due to ion bridging effects.13 Complexation of TCs with dissolved Mn(II)

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and Cu(II) significantly enhanced TCs’ oxidative transformation in the presence of oxygen at

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pH 8.0-9.5 and pH 4.0-6.0, respectively.14 Ca(II) and Mg(II) ions in natural water samples

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influenced the direct photolysis11 and self-sensitized oxidation of TCs12 under light irradiation 3 ACS Paragon Plus Environment


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due to formation of Ca(II)- or Mg(II)-TC complexes. Fe(III) ions could bind strongly to TTC

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and act as a photoreactive species in the TTC photodegradation process, affecting the

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photodegradation rate at varying pH.15 Recently, our research team found that complexation

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of TCs with Fe(III) could promote oxidative degradation of TTCs at pH 5-9 without any light

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exposure. The promoted degradation occurred via complexation of Fe(III) with TCs’

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C4-dimethylamino group and subsequent oxidation by Fe(III) to yield Fe(II) and TC

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oxidation products.16 However, the impact of Fe(II) on TCs’ transformation had not been

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investigated.

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Ferrous ion, as a common form of iron, can be easily oxidized by O2 at neutral pH to form

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stable Fe(III). Compared to Fe(II), inorganic Fe(III) is highly insoluble, more

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thermodynamically stable, and can be subject to reduction to dissolved Fe(II).17 Fe(II)/Fe(III)

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is an important redox couple in the natural environment. The transformation between Fe(II)

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and Fe(III) can occur in the oxic-anoxic boundary of various global reservoirs including

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atmosphere,18, 19 oceans,20-22 soils,23, 24 and sediments.25-27 In natural waters, most dissolved

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Fe(II) and Fe(III) exist in organic species through organic ligand complexation.28 Some

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natural or synthetic organic compounds have been observed to accelerate or retard Fe(II)

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oxidation rate.20, 28-32

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The oxidation of dissolved Fe(II) by ambient O2 (i.e., autoxidation of Fe(II)) has been

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proposed to consist of four electron transfer reactions shown in eqs 1-4 (the Haber-Weiss

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mechanism).33 The autoxidation of an organic complex of Fe(II) (i.e., FeII-X, X represents the

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organic ligand) is believed to be similar as shown in eqs 5-9.20, 34

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FeII + O2 → FeIII + O2 ·

-

(1) 4 ACS Paragon Plus Environment

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FeII + O2 · + 2H+ → FeIII + H2 O2

(2)

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FeII + H2 O2 → FeIII + ·OH + OH-

(3)

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FeII + ·OH → FeIII + OH-

(4)

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FeII + X ↔ FeII -X

(5)

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FeII -X + O2 → FeIII -X + O2 ·

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FeII -X + O2 · + 2H+ → FeIII -X + H2 O2

(7)

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FeII -X + H2 O2 → FeIII -X + ·OH + OH-

(8)

95

FeII -X + ·OH → FeIII -X + OH-

(9)

-

(6)

-

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The favorable complexation of TCs to Fe(II) and Fe(III)10, 35, 36 will likely affect the

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environmental transformation of both TCs and Fe species, but such effects have not been

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examined thoroughly. Research on the effect of Fe(II) on TCs, in particular, is lacking.

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Therefore, the objective of this study was to investigate the complexation of Fe(II) with TCs

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and how such complexation influences the transformation of both TCs and Fe(II) under

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different reaction conditions to elucidate the mechanisms. The impact of Fe(II) was also

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compared to that of Fe(III) to develop a better understanding of how transformation of TCs

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and redox cycle of Fe(II)/Fe(III) may occur and influence each other simultaneously.

104 105

MATERIALS AND METHODS

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Chemicals and Reagents. TTC, OTC and CTC were obtained from Sigma at 95-97% purity

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and used directly. Humic acid (HA) was obtained from MP Biomedicals (Santa Ana, CA, US).

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Additional chemicals and reagents used in this study are detailed in the Supporting

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Information (Text S1). 5 ACS Paragon Plus Environment


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Oxidation Experiments. Unless specified otherwise, all the oxidation experiments were

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conducted by adding 40 µM Fe(II) into 50-mL buffered solution (10 mM buffer of

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2-(N-morpholino)ethanesulfonic acid (MES) for pH 5.5-6.5 or 4-morpholinepropanesulfonic

113

acid (MOPS) for pH 6.5-7.5) in plastic bottles, with magnetic stirring and gentle ultra-zero air

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purging throughout the reaction. The dissolved oxygen (DO) concentration was found to

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maintain at around 8.9 mg⋅L-1 measured by a DO meter (Orion Star A123, Thermo Scientific).

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The reaction bottles were kept from light. Prior to the reaction, air was bubbled into the

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solution for 5 min to ensure saturated dissolved oxygen before the addition of Fe(II). For the

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oxidation experiments involving both Fe(II) and TCs, different concentrations

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(4.4×103-8.9×104 µg/L or 10-200 µM) of TCs were added into the solution before addition of

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Fe(II). Most experiments were conducted with 40 µM (1.8×104 µg/L) of initial TC

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concentration. The reactor temperature was maintained at 22 °C. Aliquots were taken

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periodically from the reaction solution for the determination of Fe(II) and TC concentrations.

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Selected batch reaction experiments were conducted with Fe(III) and TCs, in which the

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reactors were continuously stirred by a stir bar in the ambient air and protected from light. All

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the sample aliquots for TCs’ measurements were immediately quenched by adding 0.1 M HCl

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to reach pH ~2.0. The quenched samples were stored in 2-mL amber glass vials at 5°C and

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analyzed within 2-3 days. All the experiments were conducted in at least two or more

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replicates.

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Analytic Methods. Fe(II) concentrations were determined by adding 0.5-mL sample into a

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cuvette (1.0 cm path length) that contained 0.4-mL 1 mM ferrozine (FZ).37 FZ is known to

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react extremely rapidly and selectively with Fe(II) to form a stable purple complex, Fe(II)FZ3, 6 ACS Paragon Plus Environment

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with maximum absorbance at 562 nm at neutral pH,38-40 and does not bind Fe(III)

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significantly.37, 41 Furthermore, preliminary experiments confirmed that the FZ concentration

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used in this study was sufficient to fully outcompete TCs, resulting in negligible free Fe(II)

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and Fe(II)-TC complexes (SI Text S2). The absorbance was measured on a Thermo

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Spectronic Genesys 20 visible spectrophotometer.

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TCs were analyzed by an Agilent 1100 high performance liquid chromatography (HPLC)

138

system with an Agilent Eclipse C18 reverse-phase column (250 mm × 4.6 mm, 5 µm) at 20°C.

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The mobile phase was a mixture of 80% 0.01 M oxalic acid and 20% acetonitrile with a flow

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rate of 0.8 mL-min-1. TCs were detected at 365 nm by a diode-array UV/vis detector. TCs’

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transformation products were analyzed by an Agilent 1100 HPLC/UV/MSD system with the

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same column at a flow rate of 0.5 mL-min-1. Mass spectrometric analysis was conducted by

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positive electrospray ionization with a mass scan range of m/z 50-1000.

144 145

RESULTS AND DISCUSSION

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Accelerated Fe(II) Oxidation by TCs. The autoxidation of Fe(II) was significantly

147

enhanced in the presence of TTC (Figure 2), OTC and CTC (SI Figure S1) at pH 6.5.The rate

148

enhancement increased with increasing TC concentration. About 8.9%, 15%, 23% and 36%

149

of Fe(II) were oxidized by O2 in 20 min when the TTC concentrations were 10 µM, 20 µM,

150

30 µM and 40 µM, respectively, whereas only 6.1% of Fe(II) were oxidized in the absence of

151

TCs (Figure 2). The enhancement by the three TCs was comparable, with CTC slightly

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higher than OTC, followed by TTC. Due to the compounds’ structural similarity and no

153

observed difference in their impact on Fe(II) autoxidation, TTC was chosen as the primary 7 ACS Paragon Plus Environment


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target compound for further experiments. In separate experiments, the solution was purged

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with nitrogen gas and little transformation of Fe(II) occurred in the presence or absence of

156

TCs (data not shown).

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pH could be an important factor influencing the reaction of Fe(II) autoxidation.42

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Experiments in the present study showed that the autoxidation rate of Fe(II) increased at

159

higher pH in both the presence and absence of TTC (Figure 3). Very similar trends were also

160

observed for OTC (SI Figure S2). Stumm and Lee demonstrated that the autoxidation rate of

161

inorganic Fe(II) is determined by the concentrations of dissolved oxygen and OH- (eq 10).43

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The results in Figure 3 and Figure S2 suggest that the oxidation rate of Fe(II) in the presence

163

of TCs follows similar pH dependence as uncomplexed Fe(II).

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d[FeII] dt

= -k[FeII][O2 ][OH-]

2

(10)

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Humic acids (HA) are important components of organic matter in natural waters and

166

possess a high density of carboxylate functional groups that may complex with Fe(II). The

167

Fe(II)-carboxylate complexes have faster oxidation rates than inorganic Fe(II) species.20, 31

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Experiments in the present study showed that the autoxidation rate of Fe(II) increased with

169

increasing HA concentration (SI Figure S3). To assess Fe(II)’s autoxidation in the presence of

170

both TTC and HA, 40 µM TTC was added into solutions containing different concentrations

171

of HA. Results revealed that at lower HA concentrations (5 and 10 mg⋅L-1), the autoxidation

172

rate of Fe(II) was close to the mathematic sum of each alone (Figure 4, SI Figure S4a,b),

173

meaning that at [HA] ≤ 10 mg/L, both TTC and HA promoted the autoxidation of Fe(II)

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without interference from each other. However, at higher HA concentrations (20 and 40

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mg⋅L-1), the oxidation rate of Fe(II) approached the oxidation rate in the presence of HA only 8 ACS Paragon Plus Environment

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(Figure 4, SI Figure S4c,d). This was likely caused by the dominance of Fe(II)-HA

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complexation at a high HA concentration, causing the effect of HA to overshadow the effect

178

of TTC. Hence, HA may only inhibit the effect of TTC on Fe(II) oxidation at high

179

concentrations, suggesting that strong complexation of TCs with Fe(II) plays an important

180

role in systems with low HA concentrations.

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Evidently, TCs may promote Fe(II)’s autoxidation under most of the reaction conditions

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tested. In the presence of Fe(II) and TC, the solution contained both Fe(II)-TC complex(es)

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and free Fe(II), and their relative abundance depended on the concentration ratio of Fe(II) and

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TC. Several carboxylate organic ligands, including EDTA and citric acid, were reported to

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accelerate autoxidation of Fe(II) by complexing Fe(II).28-31 The accelerated oxidation rate of

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Fe(II) by TCs may be due to several possible reasons: (1) the Fe(II)-TC complex(es) are

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better electron donors and thus more easily oxidized by O2 than uncomplexed Fe(II); (2) TCs

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may form strong complexes with Fe(III) which subsequently reduce the rate of Fe(II)

189

reformation from Fe(III) reduction; and (3) the intermediate products formed may be strong

190

oxidants of Fe(II).

191 192

Fe(II)-TC Complexation. Complexation of Fe(II) by TCs played a crucial role in the

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promoted autoxidation of Fe(II). However, the complexation constants of Fe(II) with TCs

194

under environmental conditions were not available in the literature and required further study.

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Experiments were conducted under anoxic conditions to study the Fe(II)-TC complexation

196

(SI Text S3). The Job’s method has been commonly employed in previous research to

197

determine the stoichiometry and stability constant of certain metal-ligand complex. 44, 45 It 9 ACS Paragon Plus Environment


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was employed in this study to determine the predominant complex species between Fe(II)

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and TTC at pH 6.5. The Fe(II) and TTC complexation was detected by difference in the UV

200

absorbance of the Fe(II)-TTC complex in comparison to the UV absorbance of the free Fe(II)

201

plus the UV absorbance of free TTC, as shown by eq 11 and Figure S5a:

202

∆Absorbance = Abs (FeII only) + Abs (TTC only) - Abs (FeII + TTC)

203

The ∆(Absorbance) was measured in a series of solutions in which the sum of total Fe(II) and

204

TTC concentration remained the same but their concentration ratio varied. The ∆(Absorbance)

205

was measured immediately after Fe(II) and TTC were mixed together. It was found that the

206

maximum ∆(Absorbance) occurred when the ratio of [Fe(II)]:[TTC] was close to 1 (SI Figure

207

S5b), indicating that the major complex was 1:1 Fe(II)-TTC. Thus, the complexation reaction

208

may be expressed by eq 12:

209

FeII + TTC ⇌ FeII-TTC

210

where [Fe(II)] and [TTC] are the concentrations of uncomplexed Fe(II) and TTC, respectively,

211

and K1:1 is the 1:1 complexation constant.

K1:1 =

[FeII-TTC] TTC[Fe(II)]

(11)

(12)

212

Next, to determine the complexation constant K1:1, the ∆(Absorbance) was measured in a

213

complexometric titration with a fixed total Fe(II) concentration (i.e., Fe(II)T) but varying TTC

214

concentration, and the results were applied to the Benesi-Hildebrand equation (eq 13)45, 46 to

215

obtain the value of K1:1 (details in SI Text S3 and Figure S6) by plotting 1/[TTC] versus

216

1/∆(absorbance), and the complex formation constant could be determined from the slope and

217

the y-axis intercept

218 219

1 ∆(Absorbance)

=

1 Fe(II)T K1:1 ∆ϵ1

×

1 [TTC]

+

1

(13)

Fe(II)T ∆ϵ1

where ∆ϵ1 = ϵ1 - ϵFe(II) - ϵTTC (the molar absorptivity of Fe(II)-TTC complex minus the 10 ACS Paragon Plus Environment

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220 221


absorptivities of uncomplexed Fe(II) and uncomplexed TTC). The value of K1:1 was found to vary with pH (SI Text S3), indicating that the speciation of

222

TTC needed to be taken into account in the complexation reaction. As shown in Figure S7,

223

the species of TTC can be expressed as H3L, H2L, HL and L (corresponding to three acid

224

dissociation constants),7 with H2L and HL species dominant in the moderate pH range of 5.5

225

to 7.5 employed in this study (eq 14). Thus, the Fe(II)-TTC complex formation could exist as

226

FeII-H2L and FeII-HL accordingly (eqs 15 and 16):

227

H2 L ⇌ HL + H+

Ka 2 =

228

FeII + H2 L ⇌ FeII -H2 L

KFeII -H2 L =

229

FeII + HL ⇌ FeII -HL

KFeII -HL =

230

Based on the values of K1:1 obtained at different pHs, the complexation constants of KFeII -H2 L

231

and KFeII -HL were determined to be 2.39×103 and 2.05×104, respectively (SI Text S3). These

232

results indicated that, at pH 6.5 for example, about 11% of Fe(II) was complexed with TTC

233

when Fe(II)T:TTCT = 1:1 (TTCT represents total TTC concentration), and the percentage

234

increased to 38% when Fe(II)T:TTCT = 1:5 (SI Text S3 and Figure S8).

HLH+ H2 L

= 10-7.78

(14)

FeII -H2 L H2 L[Fe(II)]

(15)

[FeII -HL] HL[Fe(II)]

(16)

235

Moreover, it was desirable to understand the binding sites of TTC with Fe(II).

236

Complexation of TTC with Fe(II) at varying Fe(II):TTC ratios from 1:4 to 16:4 was

237

measured by UV absorbance spectra immediately after Fe(II) and TTC were mixed together.

238

Thus, any absorbance change was due primarily to complexation and could be calculated

239

using eq 11. The results revealed that the ∆(Absorbance) increased significantly around

240

280-320 nm with an increase in the Fe(II):TC ratio, but did not change as much around

241

340-380 nm (SI Figure S9). Literature has reported that TCs’ A ring (Figure 1) chromophore 11 ACS Paragon Plus Environment


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contributes to the absorption band at 250-300 nm, whereas the BCD ring chromophore

243

contributes to the absorption bands at both 250-300 nm and 340-380 nm.14 Thus, Fe(II) likely

244

complexes with TTC at the A ring and exerts weaker effects on the BCD ring. This result is in

245

agreement with the study by Mikulski et al., which proposed TTC’s amide group and C3

246

oxygen at A ring as the most probable binding site to Fe(II).47

247 248

Oxidation Rate of Fe(II)-TTC Complex. At constant pH and oxygen partial pressure,

249

autoxidation of inorganic Fe(II) species and Fe(II)-organic complexes is expected to follow

250

pseudo-first-order kinetics according to eq 9. Hence, the eq 17 may be used to describe the

251

oxidation of Fe(II) by O2 involving both inorganic Fe(II) species and Fe(II)-organic

252

complexes29, 31, 34:

253

kobs = αi ki + αc kc

254

where kobs (in min-1) is the overall pseudo-first-order oxidation rate constant of Fe(II); αi and

255

αc are the molar fractions of inorganic Fe(II) species and Fe(II)-organic complex, respectively;

256

and ki and kc (in min-1) are the pseudo-first-order oxidation rate constants of the inorganic

257

Fe(II) species and Fe(II)-organic complex, respectively. Note that because the

258

ln{[Fe(II)]t/[Fe(II)]}versus time line was not always straight, the kobs was obtained from a

259

shorter reaction time period where the linear regression yielded R2 ≥ 0.99. The time range

260

was 3 min and 1 min for pH 6.5 and 7.0, respectively.

261

(17)

Since the complexation of Fe(II) with TTC involves FeII-H2L and FeII-HL species, eq 17

262

may be re-written as eq 18:

263

kobs (Fe(II)T ) = ki FeII + kFeII-H2L FeII-H2L + kFeII-HL FeII-HL 12 ACS Paragon Plus Environment

(18)

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By combining with eqs 14-16 and the mass balance of total Fe(II), eq 18 is converted to eq

265

19: k obs =

266

ki [Fe(II)] + k FeII -H L [Fe II -H 2 L] + k FeII -HL [Fe II -HL] 2

[Fe(II)] + [Fe II -H 2 L] + [Fe II -HL] ki + k FeII -H L K FeII -H L 2

= 1+

2

K FeII -H L 2

K a2

[H + ] [HL] + k FeII -HL K FeII -HL [HL] K a2

(19)

[H + ][HL] + K FeII -HL [HL]

267

The kobs was measured at a fixed initial Fe(II) concentration but with varying initial TTC

268

concentrations at pH 6.5 and 7.0, respectively (Figure 5, SI Figure S10). Applying the

269

KFeII -H2 L and KFeII -HL values obtained earlier, eq 19 was solved along with the mass balance

270

of total TTC (Text S4) against the kobs data for the unknown kFeII -H2 L and kFeII -HL using

271

Matlab. The kFeII -H2 L and kFeII -HL were calculated to be 0.269 min-1 and 1.511 min-1,

272

respectively, by the model fitting (Figure 5, SI Figure S10, Text S4). The FeII -HL complex’s

273

oxidation rate constant was approximately 5.6 times higher than that of the FeII-H2L complex,

274

indicating that FeII-HL was the major species in promoting Fe(II)’s oxidation in the presence

275

of TTC at near-neutral pH. In comparison, the oxidation rate constant of uncomplexed Fe(II)

276

(ki ) was only 2.60×10-3 min-1 and 4.40×10-3 min-1 at pH 6.5 and 7.0, about 1/100 and 1/600

277

(pH 6.5) or 1/60 and 1/350 (pH 7.0) of that of FeII-H2L and FeII-HL complexes, respectively.

278

Enhanced TTC Degradation. While there was enhancement of Fe(II) autoxidation by TTC,

279

TTC’s degradation was also significantly promoted. For example, 25.6% of 40 µM TTC was

280

degraded in the presence of 40 µM Fe(II) in 90 min at pH 6.5 versus no obvious degradation

281

in the absence of Fe(II) (Figure 6). Furthermore, there was much less degradation of TTC

282

when nitrogen was purged though the solution (Figure 6), indicating that Fe(II)’s autoxidation

283

played an important role in the promoted TTC degradation. Oxidation of Fe(II) could 13 ACS Paragon Plus Environment


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284

generate Fe(III) and previous research showed that dissolved Fe(III) can complex with TCs

285

and oxidize TCs in the Fe(III)-TC complex.16 However, the Fe(III)-induced degradation of

286

TTC had a much slower rate (less than 6.0% degradation of TTC in 90 min) compared to the

287

Fe(II)-promoted degradation (Figure 6). Thus, the enhanced degradation of TTC by Fe(II)

288

must include other reactions.

289

The Fe(II)’s autoxidation (eqs 1-9) produces three kinds of ROS (O2-·, H2O2, ·OH), which

290

may degrade TTC. To evaluate the role of ROS in both Fe(II) oxidation and TTC degradation,

291

three scavengers (superoxide dismutase (SOD), catalase (CAT), tert-butanol (TBA)) were

292

employed in this study (SI Figure S11). Results showed that addition of the three ROS

293

scavengers had negligible effects on Fe(II)’s oxidation rate in the presence of TTC (Figure

294

7a), indicating that the Fe(II) in Fe(II)-TTC complex was primarily oxidized by dissolved

295

oxygen (eq 6). Conversely, the addition of ROS scavengers affected TTC’s degradation

296

significantly (Figure 7b). TTC degradation was reduced by about 2.68%, 10.0% and 6.60%

297

(at 90 min) when 400 units-mL-1 SOD, 400 units-mL-1 catalase and 1.0 M TBA were added

298

into the solution, respectively. Both catalase (H2O2 scavenger) and TBA (·OH scavenger)

299

retarded TTC’s degradation, indicating that the H2O2 and ·OH formed during Fe(II)-TTC

300

oxidation contributed to TTC’s degradation. This finding is consistent with literature report

301

that TTC can be readily degraded by ·OH radicals at a second-order rate constant of (6.3 ±

302

0.1)×109 M-1s-1.48 In contrast, SOD barely affected TTC’s degradation; this result alone,

303

however, could not exclude the contribution of superoxide radical because SOD may react

304

with O2-· to form H2O2 (eq 20),49, 50 which could react with Fe(II) and TTC subsequently:

305

O2 · + SODFeII + 2H+ → H2O2 + SOD(FeIII)


(20)

Page 15 of 35


306

To verify if O2-· had an effect on TTC’s degradation, higher catalase concentrations and

307

simultaneous addition of both SOD and catalase were tested, respectively. When the catalase

308

concentration was varied from 400 to 2000 units-mL-1, TTC’s degradation decreased from

309

15.6% to 8.2% of that without catalase addition at 90 min. However, further increase of

310

catalase concentration to 4000 units-mL-1 did not result in additional inhibition effect on

311

TTC’s degradation (SI Figure S12a). Thus, the remainder 8.2% enhancement of TTC’s

312

degradation might be caused by O2-· or other reactive species. However, the degradation

313

profile of TTC in the presence of both catalase and SOD overlapped with the profile resulting

314

from the presence of only catalase (SI Figure S12b), which ruled out the impact of O2-· on

315

TTC’s degradation.

316

To further confirm the contribution of H2O2 to TTC’s degradation, three scenarios were

317

tested with addition of H2O2 to solutions containing (1) only TTC; (2) TTC, Fe(III) and TBA;

318

and (3) TTC and Fe(II). In scenario (1), results showed that H2O2 alone had little effect on

319

TTC’s degradation, suggesting that H2O2 cannot directly oxidize TTC. In scenario (2), less

320

than 5% of total TTC was degraded in the H2O2-Fe(III)-TTC system within 90 min (detailed

321

in SI Text S5), which indicated that H2O2 cannot directly oxidize Fe(III)-TTC complex. In

322

scenario (3), the addition of H2O2 with Fe(II), however, significantly enhanced TTC’s

323

degradation due to production of ·OH from the Fenton process (SI Figure S13a), and the

324

oxidation of Fe(II) exhibited the same trend as TTC (SI Figure S13b). Nearly identical

325

concentration consumptions of Fe(II) and TTC were observed in the Fe(II)-TTC-H2O2 system

326

(SI Figure S13c), which also supports that TTC’s degradation is caused by ·OH.

327

Overall, TTC’s degradation during the oxidation of Fe(II) is primarily caused by ·OH. The 15 ACS Paragon Plus Environment


328

contribution of H2O2 is because it can react with Fe(II) to produce ·OH. Other reactive

329

species (not O2-·) may degrade TTC but at a slight contribution.

330

TTC Degradation Products. Preliminary analyses were conducted to determine the

331

transformation intermediates and products of TTC’s degradation during Fe(II)’s oxidation

332

process. Only one major degradation product was observed on LC/MS; the peak area of the

333

observed product [M+34] (M represents the molecular weight of the protonated parent

334

compound) increased with time (Figure 8a). It was reported that for the ·OH radical reactions

335

with methyl methacrylate, the major reaction pathway involves the addition of ·OH to the

336

terminal carbon of C=C with the breakage of the double bond.51 Thus, the dominant product

337

of [M+34] indicated that there were two oxygen atoms added to the parent compound with

338

the breakage of C=C double bonds, consistent with hydroxylation of TTC by ·OH radicals.

339

OTC and CTC showed similar [M+34] product generation as TTC (data not shown).

340

Moreover, the product’s peak area increased sharply in the first 3 h and then dropped

341

gradually in the following 48 h (Figure 8b). The concentration of Fe(II) dropped to below the

342

method detection limit (MDL) of 0.5 µM in 3.26 h, which could explain why the area of the

343

M+34 product stopped increasing and began to drop. After Fe(II) was depleted, the M+34

344

product was unstable and might be oxidized by Fe(III) formed in the system.

345

Proposed Fe(II)/Fe(III) Redox Cycle in the Presence of TCs and Environmental

346

Implications. On the basis of all of the experimental results, the redox cycle of Fe(II)/Fe(III)

347

in the presence of TCs in aqueous environments is proposed in Figure 9 and the associated

348

reactions and constants listed in Table 1. In the oxic aquatic environment, Fe(III), as the

349

dominated iron species, can complex with TC and oxidizes TC to form Fe(II) and TC radical 16 ACS Paragon Plus Environment

Page 16 of 35

Page 17 of 35


350

as previously illustrated.16 The yielded Fe(II) can be oxidized back to Fe(III) by oxygen and

351

TC radicals can further degrade. This pathway of Fe(III) and TC transformation is a fairly

352

slow process. On the other hand, in the anoxic aquatic environment, ferrous ions, present as

353

free Fe(II) and Fe(II)-TC complexes, can be easily oxidized to Fe(III) and Fe(III)-TC species

354

when oxygen enters the system. The oxidation rate of total Fe(II) is enhanced by the

355

significantly higher oxidation rate of the Fe(II)-TC complexes. Meanwhile, the ROS

356

(primarily ·OH) formed during Fe(II)’s oxidation induces TC’s degradation. The resultant

357

Fe(III) can be reduced back to Fe(II) by other environmental reductants. More simply, the

358

Fe(II)/Fe(III) redox cycle in the presence of TCs contains two different electron transfer

359

pathways. For the Fe(III)-induced TC degradation process, Fe(III) readily complexes with TC,

360

and the TC molecule acts as the electron donor to Fe(III), forming Fe(II). For the

361

Fe(II)-induced TC degradation process, Fe(II) complexes with TC and the Fe(II)-TC

362

complexes act as the electron donor and transfer electrons to oxygen, forming superoxide

363

anions.

364

The transformation of Fe(II)/Fe(III) may occur in many geochemical environments, such

365

as the oxic/anoxic boundary in marine and freshwater basins, the oxycline in sediments, and

366

sediment-water interfaces.17 TCs, the most widely used antibiotics, enter the aquatic

367

environments via various wastes. When TC contaminants enter the environment, there is a

368

high possibility of TCs and Fe to mutually affect each other’s environmental fate. This study

369

sheds new mechanistic insights on such reactions and improves the knowledge basis to better

370

predict the environmental fate of TC contaminants. Furthermore, this study may provide

371

guidance for additional research on the transformation of transition metal ions and 17 ACS Paragon Plus Environment


372

environmental contaminants, as well as facilitate development of remediation strategies for

373

TC contaminants.

374 375

ASSOCIATED CONTENT

376

Supporting Information. Text S1-S5 and Figure S1-S13 are available free of charge via the

377

Internet at http://pubs.acs.org.

378 379

ACKNOWLEDGEMENTS

380

Funding sources of this work included the National Natural Science Foundation of China

381

(51578043), and the U.S. Department of Agriculture Grant 2009-65102-05843. The authors

382

thank the three anonymous reviewers for their helpful comments to improve this manuscript.

383 384

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the

reactions

of

OH

radicals

and

Cl

atoms

517 518


with

methyl

methacrylate

Page 25 of 35


Table 1. Proposed reactions involving Fe(II) or Fe(III) with TCs and associated constants. Stability constant/

Reaction

Ref

Rate constant FeII + H2L ⇌ FeII-H2L

2.39×103a

KFeII-H2L

Complexation

FeII + HL ⇌ FeII-HL

KFeII-HL

FeIII + TC ⇌ FeIII-TC

KFe(III)-TC

FeII + O2 → FeIII + ROS II

e--transfer reaction

105.3, 33

2.05×104a

k0

-

109.9, 33

9.35 M-1⋅min-1a

8.29 M-1⋅min-1,39c

(pH 6.5) 2

-1

(pH 6.5) -1

Fe -H2L + O2 → FeIII-H2L + ROS

k1

9.68×10 M ⋅min

a

-

FeII-HL + O2 → FeIII-HL + ROS

k2

5.44×103 M-1⋅min-1a

-

TTC TTC radical

k3

2.90×10-3 min-1a

-

FeIII-TTC → FeII + TTC radical

k4

1.11×10-4 min-1b

-

FeII, O2

a. Constants determined in this study where pH = 6.5-7.0, T = 22 °C, initial Fe(II)T = TTCT = 40 µM, [O2] = 2.78×10-4 M. b. Constants determined in this study where pH = 7.0, T = 22 °C, initial Fe(III)T = TTCT = 40 µM, solution protected from light. -3

-1

c. The first-order oxidation rate constant of Fe(II) was reported to be 2.33×10 min at pH 6.5, [O2] ~ 2.81×10-4 M and 20.5 °C.



Page 26 of 35

R1

R2

pKa1

pKa2

pKa3

Tetracycline (TTC)

H

H

3.32

7.78

9.58

Oxytetracycline (OTC)

H

OH

3.22

7.46

8.94

Chlorotetracycline (CTC)

Cl

H

3.33

7.55

9.33

Figure 1. Structures and properties of TCs.


Page 27 of 35


1.00

[Fe(II)]t/[Fe(II)]0

0.90

0.80

0.70 w/o TTC w/ 10 µM TTC w/ 20 µM TTC w/ 30 µM TTC w/ 40 µM TTC

0.60

0.50 0

4

8 12 Time (min)

16

20

Figure 2. Effect of initial TTC concentration on Fe(II)’s oxidation in 20 min. Condition: [Fe(II)]0 = 40 µM, pH = 6.5 (buffered by 10 mM MOPS), T = 22 °C, with air purging.



[Fe(II)]t/[Fe(II)]0

1.00

0.80

0.60 w/o TCs pH 5.5 w/o TCs pH 6.5 w/o TCs pH 7.5 w/ 40 µM TTC pH 5.5 w/ 40 µM TTC pH 6.5 w/ 40 µM TTC pH 7.5

0.40

0.20 0

5

10 15 Time (min)

20

Figure 3. Fe(II)’s oxidation in the absence and presence of TTC at different pH. Condition: [Fe(II)]0 = 40 µM, buffered by 10 mM MES for pH 5.5 and 10 mM MOPS for pH 6.5 and 7.5, T = 22 °C, with air purging.


Page 28 of 35

Page 29 of 35


0.0

ln{[Fe(II)]t/[Fe(II)]0}

-0.4

-0.8

-1.2

w/ 5 mg/L HA w/ 40 mg/L HA w/ 40 µM TTC w/ 5 mg/L HA + 40 µM TTC w/ 40 mg/L HA + 40 µM TTC sum of 5 mg/L HA and 40 µM TTC sum of 40 mg/L HA and 40 µM TTC

-1.6

-2.0 0

5

10 15 Time (min)

20

Figure 4. Oxidation of Fe(II) in the presence of TTC and HA. Condition: [Fe(II)]0 = [TTC]0 = 40 µM, [HA]0 = 5 or 40 mg/L, pH = 6.5 (buffered by 10 mM MOPS), T = 22 °C, with air purging. Dashed line: the mathematic sum of the effect of HA only and the effect of 40 µM TTC only.



Page 30 of 35

0.30 0.25

kobs (min-1)

0.20 0.15 0.10 Model fitting

0.05

Experimental data 0.00 0

2

4 6 [TC-] (µM)

8

10

Figure 5. Experimental data and model fitting of observed Fe(II) oxidation rate constants with varying initial TTC concentrations at pH 6.5. Condition: [Fe(II)]0 = 40 µM, [TTC]0 = 0, 40, 60, 90, 120, 150 and 200 µM, buffered by 10 mM MOPS, T = 22 °C, with air purging.


Page 31 of 35


[TTC]t/[TTC]0

1.00

0.90

0.80 TTC only TTC w/ 40 µM Fe(II), Oxygen TTC w/ 40 µM Fe(II), Nitrogen TTC w/ 40 µM Fe(III)

0.70

0.60 0

30

60 Time (min)

90

Figure 6. Degradation of TTC in the presence of Fe(II) or Fe(III) . Condition: [Fe(II)]0 = [Fe(III)]0 = [TTC]0 = 40 µM, pH = 6.5 (buffered by 10 mM MOPS) and T = 22 °C. Reactions with Fe(II) were purged with air or nitrogen, and no gas purging in the reaction with Fe(III).



ln{[Fe(II)]t/[Fe(II)]0}

0

-0.2

-0.4 w/o TTC w/ 40 µM TTC, w/o air w/ 40 µM TTC (a) w/ 40 µM TTC, TBA w/ 40 µM TTC, catalase w/ 40 µM TTC, SOD

-0.6

-0.8 0

4

8 12 Time (min)

16

20

[TTC]t/[TTC]0

1.0

0.9

0.8 w/o Fe(II) w/ 40 µM Fe(II) w/ 40 µM Fe(II), TBA w/ 40 µM Fe(II), catalase w/ 40 µM Fe(II), SOD

0.7

(b)

0.6 0

30

60 Time (min)

90

Figure 7. Scavengers’ effects on (a) Fe(II)’s oxidation and (b) TTC’s degradation in the Fe(II)-TTC system. Condition: [Fe(II)]0 = [TTC]0 = 40 µM, pH = 6.5 (buffered by 10 mM MOPS), T = 22 °C, with air purging; [TBA] = 1.0 M, [catalase] = 400 units-mL-1, [SOD] = 400 units-mL-1.


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parent compound product M + 34

12

160

8

120

4

80 (a) 40

Peak Area for M+34 (×105)

Peak Area for parent compound (×105)

16 200

0 0

15

30

45 60 Time (min)

75

90

1

4

[TTC]t/[TTC]0

0.8

product M + 34

3

0.6 2 0.4 [Fe(II)] below MDL

0.2

1 (b)

0

Peak Area for M+34 (×104)

parent compound

0 0

10

20

30

40

50

Time (h)

Figure 8. (a) Product evolution of TTC during Fe(II)’s oxidation process; (b) The evolution of parent TTC and product during Fe(II)’s oxidation process in a 48-hours period. Condition: [Fe(II)]0 = [TTC]0 = 40 µM, pH = 6.5 (buffered by 10 mM MOPS), T = 22 °C, with air purging.



Figure 9. Proposed Fe(II)/Fe(III) redox cycle in the presence of TC in aquatic environments


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TOC Art:


Transformation of Tetracycline Antibiotics and Fe(II) and Fe(III

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