Oxidation of Oil Sands Process-Affected Water by Potassium Ferrate

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Oxidation of Oil Sand Process-affected Water by Potassium Ferrate(VI) Chengjin Wang, Nikolaus Klamerth, Rongfu Huang, Haitham Elnakar, and Mohamed Gamal El-Din Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.5b04829 • Publication Date (Web): 23 Mar 2016 Downloaded from http://pubs.acs.org on March 28, 2016

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Environmental Science & Technology

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Oxidation of Oil Sand Process-affected Water by

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Potassium Ferrate(VI)

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Chengjin Wang, Nikolaus Klamerth, Rongfu Huang, Haitham Elnakar, and Mohamed Gamal El-Din*

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Department of Civil and Environmental Engineering, University of Alberta, Edmonton, Alberta, Canada,

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T6G 1H9

6 7

* Corresponding author:

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Mohamed Gamal El-Din

9

7-285 Donadeo Innovation Centre for Engineering

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Department of Civil and Environmental Engineering

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University of Alberta

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Edmonton, Alberta, T6G 1H9

13

Canada

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Tel.: +1 780 492 5124

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Fax: +1 780 492 8198

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E-mail: [email protected]

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KEYWORDS: oil sands process-affected water, ferrate(VI), naphthenic acids, toxicity,

18

coagulation, electron paramagnetic resonance

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ABSTRACT

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This paper investigates the oxidation of oil sand process-affected water (OSPW) by potassium

25

ferrate(VI). Due to the selectivity of ferrate(VI) oxidation, two-ring and three-ring fluorescing

26

aromatics were preferentially removed at doses 95%) was purchased from Tokyo Chemical Industry (TCI) (Tokyo, Japan). All

98

other chemicals were of analytical grade unless specified otherwise.

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Experimental Design.

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OSPW Study.

101

The investigated ferrate(VI) dose was between 1 and 400 mg/L calculated as iron. Dry potassium

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ferrate(VI) was added into 200 mL OSPW under rapid mixing with RT 15 power IKAMAG®

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magnetic stirrers (IKA Works, Inc., Wilmington, NC, USA). After 3 hours, rapid mixing was

104

stopped, and sodium thiosulfate was added to quench the residual ferrate(VI). The treated OSPW

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was then filtered through 0.45 µm nylon filter (Supelco Analytical, Bellefonte, PA, USA) and

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preserved in amber glass bottles. During the reaction, pH was monitored, and OSPW samples

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(0.5 mL) were taken at different times to measure residual ferrate(VI) concentration by

108

spectrophotometry at 510 nm, with a molar absorption coefficient ɛ510nm = 1150 M-1cm-1.29 As a

109

comparison, experiments with ferric chloride were conducted in a similar way with same iron

110

concentrations. All experiments were done in triplicates.

111

Water Quality Analysis for the OSPW Study.

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pH was monitored using an Accumet Research AR20 pH/conductivity meter (Fisher Scientific,

113

Ottawa, ON, Canada). UV254 absorbance was measured with a UV-Vis spectrophotometer

114

(Varian Inc., Santa Clara, CA, USA). Synchronous fluorescence spectra (SFS) of the filtered

115

water were recorded with Varian Cary Eclipse fluorescence spectrometer (Mississauga, ON,

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Canada). Excitation wavelengths ranged from 200 to 600 nm, and emission wavelengths were

117

recorded from 218 to 618 nm. Scanning speed was 600 nm/min and the photomultiplier (PMT)

118

voltage was 800 mV. Regarding the different response factors (i.e., different quantum yields) of

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different fluorescing aromatics under UV irradiance, the quantification of the removal of these

120

aromatics based on the peak area is not adopted in this study.

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Naphthenic acids (NAs) were quantified by ultra-performance liquid chromatography time-of-

122

flight mass spectrometry (UPLC TOF-MS). The injection solution was prepared with 500 µL

123

sample, 100 µL of 4.0 mg/L internal standard (myristic acid-1-13C) in methanol, and 400 µL

124

methanol to reach a final sample volume of 1 mL. Raw OSPW was used as the control sample

125

for quality control. Chromatographic separations were performed using a Waters UPLC Phenyl

126

BEH column (1.7 lm, 150 mm ×1 mm), with mobile phases of 10 mM ammonium acetate in

127

water (A) and 10 mM ammonium acetate in 50/50 methanol/acetonitrile (B). The elution


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gradient was 0–2 min, 1% B; 2–3 min, increased from 1% to 60% B; 3–7 min, increased to 70%

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B; 7–13 min, 95% B; 13–14 min, 1% B and hold 1% B until 20 min to equilibrate column with a

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flow rate of 100 µL/min. The column temperature was set at 50 ºC and the sample temperature at

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10 ºC. Samples were analyzed with a high-resolution time-of-flight mass spectrometer (Synapt

132

G2, Waters, Milford, MA, USA) with the electrospray ionization (ESI) source operating in

133

negative ion mode and TOF analyzer in high-resolution mode. The negative ESI has been

134

commonly used to determine naphthenic acids species based on previous comprehensive studies

135

of OSPW.

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extraction from spectra was performed using TargetLynx (Waters, Milford, MA, USA). The

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identification of NAs (O2) and oxy-NAs are based on the empirical formula CnH2n+zOx (x=2, 3, 4,

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5) with n ranging from 7 to 26 and z from 0 to -18. And the quantification of each NA species

139

was done by comparing its peak area with that of the internal standard. The UPLC TOF-MS

140

quantification method was developed previously and has been verified in a number of

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publications10, 30, 31, 33 as a semi-quantification method, given the full quantification method is not

142

currently available due to the lack of individual NA standard compounds. The method for

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acquisition of ion mobility (IM) spectra was the same as described by Wang et al.10

144

1

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Unity Inova Spectrometer (Agilent Technologies, Santa Clara, CA, USA) at 399.950 MHz. The

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extracted organic matter was reconstituted in 1 mL deuterated dimethyl sulfoxide (DMSO) and

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transferred into 5 mm OD (outside diameter) NMR tubes for analysis. Detailed description of

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sample preparation is provided in the SI. Key parameters of the instrument were set as follows:

149

1

30-32

The data acquisition process was controlled using MassLynx (Waters) and data

H nuclear magnetic resonance (NMR) spectra were acquired with 3-channel Agilent/Varian

H resonance frequency: 399.950 MHz; pw30 (i.e., 30 degree pulse): 3.3 ms; acquisition time: 5s;



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spectral width: 10000 Hz or 25 ppm; transmitter offset: 989.9; relaxation delay: 0.1 s. The data

151

were processed with software VNMRJ 4.2A.

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Fourier transform infrared (FT-IR) spectra were acquired with PerkinElmer Spectrum 100 FT-IR

153

Spectrometer (PerkinElmer Life and Analytical Sciences, Woodbridge, ON, CA). The samples

154

for FT-IR analysis were prepared as follows: the pH of 100 mL OSPW was adjusted to 2.0, and

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the organic fraction was extracted with 50 mL carbon tetrachloride (Sigma-Aldrich, St. Louis,

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MO, USA), and the extract was evaporated to dryness, reconstituted in 5 mL CCl4 and injected

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into the KBr cell for analysis.

158

The electron paramagnetic resonance (EPR) signals were obtained with Bruker Elexsys E500

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spectrometer (Billerica, MA, USA). α-(4-Pyridyl N-oxide)-N-tert-butylnitrone (POBN) (Sigma-

160

Aldrich, St. Louis, MO, USA) was used as the spin trap. All the spectra were obtained at 150 K

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using a liquid nitrogen dewar with the following settings: temperature = 150 K, center filed =

162

336 G, sweep width = 20 G, microwave frequency = 9.436 GHz, modulation frequency = 100

163

kHz, and power = 20 mW.

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Acute toxicity towards Vibrio ﬁscheri was measured with the Microtox® 81.9% screening test

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protocol. In vitro toxicity assessment was done using goldfish primary kidney macrophages

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(PKMs), and detailed procedure is provided in the SI and can also be found in Shu, et al. 34.

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Model Compound Study.

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Only classical and oxidized NAs were selected for model compound study, and sulfur-containing

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NAs and aromatics were not covered in the current study due to the lack of information on their

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molecular structure. Three NA model compounds were selected to investigate their structure-

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reactivity

towards

ferrate(VI)

oxidation.

These

three

model


compounds,

including

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cyclohexanecarboxylic acid (CHA, C7H12O2), cyclohexanepentanoic acid (CHPA, C11H10O2),

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and 1,3-adamantanedicarboxylic acid (ADA, C12H16O4), embodied some typical structure

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properties of NAs, including single-ring and multiple-ring, long and short side chains, and

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different number of tertiary carbons and α and β carbons relative to the carboxyl groups. Both the

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degradation kinetics and by-products will be investigated. Sodium bicarbonate buffer (1.18 g/L,

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pH adjusted to 8.5) with a single NA model compound concentration of 60 mg/L was used as the

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model mixture. The buffer solution was prepared to simulate the 600 mg/L (as CaCO3) alkalinity

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in OSPW. The dose of ferrate(VI) applied was 200 mg/L Fe(VI) and the experiment was

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conducted by adding solid ferrate(VI) under rapid mixing as described in the OSPW study.

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Samples were taken at 0 min, 2 min, 5 min, 10 min, 20 min, 30 min, 45 min, 60 min, 90 min, 120

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min, and 180 min, and immediately quenched with sodium thiosulfate. The concentrations of the

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model compounds after ferrate(VI) oxidation were measured with Agilent 1100 binary HPLC

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coupled with G1946D MSD mass spectrometer (Santa Clara, CA, USA). End products were

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identified with Agilent 1200 HPLC coupled with Agilent 6220 oaTOF mass spectrometer (Santa

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Clara, CA, USA). Mass spectrometry was performed in the negative scan/SIM (selected-ion

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monitoring) mode. Detailed information about the instruments and the setups are provided in the

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SI.

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RESULTS AND DISCUSSION

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Analyses of Raw OSPW.

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Organic matter in OSPW contains both aromatic and aliphatic organic compounds. Analyses of

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OSPW by various analytical methods (1H NMR, SFS, UV-Vis, FT-IR) show a variety of

193

different signals like aromatic and aliphatic hydrogen, and acid groups. 1H NMR spectrum



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provides the information about the structures and functional groups of the organic matter in

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OSPW by detecting the chemical shift of protons in different environments. The 1H NMR

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spectrum (Figure 1a and S2a) clearly shows three different resonances of hydrogen atoms

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bonded in organic molecules: at 0.3-2.2 ppm are the hydrogen atoms attached to aliphatic

198

carbons, at 6.5-8.0 ppm are the hydrogen atoms on aromatic rings, and at 11.4-12.5 ppm are the

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hydrogen atoms in carboxyl functional groups.35 No peak was detected from 4.5 to 5.5 ppm,

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which represented hydrogens attached to aliphatic carbon-carbon double bonds 35. The aromatic

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hydrogen peak is broad, indicating that aromatic organic matter in OSPW is a complex mixture;

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it might include benzene rings with aliphatic branches, oxygen or sulfur-containing heteroatomic

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aromatics, and even small amount of humic and fulvic acids.4, 35 So far there is no quantitative

204

information about the total concentration of the aromatics in OSPW. However, based on the peak

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area shown in the 1H NMR spectrum, the molar ratio of total aromatic compounds over total

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organic acids (including classical NAs, oxy-NAs, and other organic acids) in OSPW can be

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roughly estimated to be in the range of 0.08-0.32 depending on the number of carboxyl

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functional groups in each organic acid and the number of hydrogen atoms on each aromatic ring

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(detailed calculation is provided in SI). FT-IR spectrum is an effective way to study functional

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groups by detecting the stretching and bending of bonds. FT-IR spectrum (Figure S1b) confirms

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the existence of both organic acids and aromatics in OSPW: the sharp peak at 1706 cm-1 was

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assigned to the C=O stretch of carboxyl functional groups, while the peaks at 3100-3000 cm-1

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and at 1400-1500 cm-1 are the signals of the aromatic C-H stretch and aromatic in-ring C-C

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stretch, respectively.35 The UV-Vis absorption spectrum (Figure S1c) shows low absorption

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above 400 nm, consistent with our previous report of the low color of OSPW36. SFS of OSPW

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(Figure 1b and S2d) provides more specific information on fluorescing compounds: peak I at 267


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nm is assigned to one-ring aromatics (which refer to not only aromatics containing one aromatic

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ring but also aromatics containing two or more unfused aromatic rings), while peak II at 310 nm

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and peak III at 330 nm are assigned to aromatics with two and three fused rings respectively.37, 38

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Considering that the peaks for two-ring and three-ring aromatics are weaker than one-ring

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aromatics and the quantum yields for florescence of two-ring and three-ring aromatics are much

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higher than those of one-ring aromatics (e.g., the quantum yield of anthracene is ten times that of

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benzene)39, it can be concluded that one-ring aromatics predominate over two-ring and three-ring

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aromatics in OSPW.

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Consumption of Potassium Ferrate(VI) in OSPW.

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The reaction of ferrate(VI) with organic compounds in OSPW is a slow process and it took

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around 2-3 hours to consume more than 95% of ferrate(VI) depending on the initial doses. The

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consumption rate of ferrate(VI) was the highest at the first five minutes of the reaction, with 1

229

mg/(L·min) for a dose of 25 mg/L and 5 mg/(L·min) for a dose of 200 mg/L Fe(VI); the

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consumption rate decreased to values of 0.002-0.087 mg/(L·min) at the end of the reaction,

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depending on the initial doses (Figure S2a). This can be explained by two aspects: (1) Ferrate(VI)

232

oxidation is a second-order reaction as reported by Jiang et al21, which means that the reaction

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rate is related to the concentrations of both oxidants and reductants; with increasing reaction time,

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the concentration of ferrate(VI) and contaminants in OSPW decreased, leading to low

235

consumption rate. (2) The pH in OSPW increased from pH 8.6 to pH 8.7-10.9 (depending on the

236

doses) during the course of the experiments after ferrate(VI) was added (Figure S2b-c), which

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was mainly due to the generation of hydroxide ions during the reaction and the self-

238

decomposition of ferrate(VI).40 Due to the increase of pH, the proportion of protonated



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ferrate(VI) ions, which has a higher redox potential than unprotonated ferrate(VI), is lower, thus

240

decreasing the reactivity of ferrate(VI) in OSPW.

241

Organic matter removal.

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Removal of Aromatics.

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As mentioned in section 3.1, there are several ways to characterize the aromaticity of the organic

244

matters in OSPW such as UV-vis spectrum, SFS, FT-IR spectrum, and 1H NMR spectrum. To

245

investigate the removal of aromatics in OSPW, we employed only SFS and 1H NMR spectra,

246

because SFS is a more selective and sensitive method to measure the aromaticity of the water

247

than UV-vis spectrum41, and 1H NMR spectrum provides qualitative and quantitative information

248

of both aromatics and organic acids.35

249

Applying ferrate doses lower than 100 mg/L Fe(VI) removed two and three ring aromatic

250

compounds while no removal of one ring compounds could be observed (Figure 1b). By

251

increasing the ferrate(VI) dose from 100 to 400 mg/L Fe(VI), one ring aromatic compounds were

252

also significantly removed, while two and three ring aromatic compounds were almost

253

completely removed as can be seen in Figure 1b. This might be due to the fact that aromatic

254

compounds with two or more fused rings had a higher electron density and thus were more

255

reactive towards ferrate(VI) attack while the single benzene ring was usually unactivated and

256

thus more stable to oxidation.42 Additionally, the degradation of two and three ring compounds

257

led to the formation of one ring compounds, which might explain the apparent stability of one

258

ring aromatic compounds.

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The significant removal of aromatic organic matter was confirmed by 1H NMR spectra (Figure

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1a), which was acquired for samples treated at a representative dose of 200 mg/L Fe(VI). The 12 ACS Paragon Plus Environment

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peak assigned to hydrogen atoms on aromatic rings at the shift of 6.5-8.0 ppm decreased by 44.4%

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after ferrate(VI) oxidation, indicating that 44.4% of aromatic hydrogen atoms disappeared after

263

treatment. However, this does not mean that only 44.4% aromatics were oxidized, because some

264

aromatics can keep their aromaticity after oxidation. For example, fluorescing aromatics might

265

be oxidized to other non-fluorescing aromatics with partial or all the aromaticity preserved. This

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also explains why SFS showed an almost complete decrease of peak areas while 1H NMR

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spectrum only showed 44.4% removal of aromatic hydrogens at 200 mg/L Fe(VI). Similar

268

phenomenon is found in the oxidization of NAs, namely a high removal of NAs was achieved

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while only a moderate removal of the carboxyl functional groups was observed (29.7% removal

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of carboxyl functional groups at 200 mg/L Fe(VI)).

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Removal of NAs.

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Analysis of OSPW with ion mobility provides qualitative information about the present NAs.

273

Three different regions according to retention and drift time can be distinguished: classical NAs,

274

oxy-NAs, and heteroatomic NAs.10, 30 High intensities of all three clusters can be seen in raw

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OSPW (Figure 2a). After treatment with 200 and 400 mg/L of ferrate(VI), the classical NAs are

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significantly reduced, and heteroatomic NAs are completely removed (Figure 2b&c). The

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preferential removal of heteroatomic NAs can be attributed to the electron-rich moieties that

278

usually come with the heteroatoms.43, 44 No significant removal of oxy-NAs can be seen, this

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however can be attributed to the fact that oxy-NAs are removed but also produced during

280

treatment. This finding is further confirmed by UPLC TOF-MS.

281

Analysis by UPLC TOF-MS showed that the initial concentration of classical NAs was 39.8

282

mg/L which changed to 14.3 mg/L and 8.6 mg/L after treatment by 200 mg/L Fe(VI) and 400



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mg/L Fe(VI), corresponding to the removal of 64.0% and 78.4%, respectively (Figure 2d-2f).

284

NAs+O (oxy-NAs with one additional oxygen) had an initial concentration of 12.5 mg/L and it

285

changed to 14.8 and 15.6 mg/L after 200 mg/L Fe(VI) treatment and 400 mg/L Fe(VI) treatment,

286

corresponding to a negative removal (increase) of those species of -18.4% and -24.8%,

287

respectively; NAs+O2 (oxy-NAs with two additional oxygens) had an initial concentration of

288

13.3 mg/L, and the concentration changed to 15.7 and 14.4 mg/L after 200 mg/L Fe(VI)

289

treatment and 400 mg/L Fe(VI) treatment, corresponding to a negative removal of -18.0% and -

290

8.3%, respectively. Again this can be attributed to the fact that oxy-NAs are produced during

291

treatment.

292

In order to investigate the impact of carbon number and hydrogen-deficiency on the oxidation

293

process, classical NAs removal versus carbon number and hydrogen-deficiency was plotted in

294

Figure 2g and 2h. There is a clear trend that a high carbon number leads to a high removal

295

(Figure 2g), with complete classical NAs removal when n≥18. The reason is that larger carbon

296

skeleton structure means more potential sites for the attack of ferrate(VI), resulting in

297

preferential removal of NAs with high carbon number, which is in accordance with work done

298

by Pérez-Estrada, et al.45 Similarly, as shown in Figure 2h, high hydrogen-deficiency also leads

299

to high NAs removal, as higher z number NAs have a higher reactivity for the following two

300

reasons: first, high z number might lead to higher number of rings, branches, and thus increasing

301

number of tertiary carbons, which tend to be more reactive than primary and secondary

302

carbons;35 in addition, high hydrogen-deficiency might also imply more double bonds, which are

303

more reactive than the saturated hydrocarbons.45

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For NAs+O, their concentration increase at low n and z number outweighed their concentration

305

decrease at the region of high n and z number (Figure S3a-c). The reason behind this behavior is 14 ACS Paragon Plus Environment

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due to the fact that in general NAs+O with higher n and z numbers are more reactive towards

307

oxidation. This leads to degradation of oxy-NAs together with NAs resulting in lower n and z

308

number oxy-NAs (Figure S3) and is confirmed by other research.18 For NAs+O2 an additional

309

high n and z number cluster in the range of z=12-18 and n=15-24 could be observed (Figure S3d-

310

f). This leads to the conclusion that NAs+O2 were another main group of transformed products

311

from NAs oxidation, which was also confirmed by the fact that classical NAs in this range were

312

almost completely removed. In general, during the ferrate(VI) oxidation process, there were both

313

oxygen insertion which led to more “NAs+Ox” species and the breakdown of the large molecules

314

which led to more small molecules (i.e., molecules with small carbon number n). Both processes

315

actually increased the solubility of the total NAs fraction in oxidized OSPW, which was

316

confirmed by Figure S4 showing the OSPW before and after ferrate(VI) oxidation: the

317

ferrate(VI)-treated OSPW was transparent when pH was adjusted to 2.5, while colloids formed

318

in the untreated OSPW after pH adjustment.

319

Removal Mechanisms.

320

Oxidation vs Coagulation.

321

Although ferrate(VI) is a dual agent as both oxidant and coagulant, the removal of organic matter

322

can mainly be attributed to the oxidation process. In order to separate the contributions of these

323

two processes, a comparison experiment was conducted with ferric chloride and potassium

324

ferrate(VI). At FeCl3 concentrations between 1 and 400 mg/L calculated as iron,

325

coagulation/flocculation achieved negligible removal of fluorescing organic matter (no peak

326

decrease in the SFS), UV254-absorbing organic matter (0-0.7%), and organic acids (0-1.7%)

327

(Figure S5a-c), indicating that both aromatic and NA compounds were not removed through the



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coagulation/flocculation process. This negligible removal of organic matter by the coagulation in

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OSPW treatment was also reported by Alpatova, et al.46 The reasons behind this limited removal

330

of organic matter by the coagulation process are the follows: firstly, OSPW has a high pH of 8.6

331

(Figure S2b), making the floc surface highly negative. Organic acids in OSPW exist mostly in

332

the anion form which is also negatively charged, and therefore electrostatic repulsion exists

333

between the organic ions and the flocs; secondly, organic matter in OSPW has low molecular

334

weight in the range of 150-350 Da36, and the physical adsorption through van der Waals force is

335

not enough to overcome the electrostatic repulsion; thirdly, the surface of ferric hydroxide flocs

336

is too hydrophilic to be a good adsorbent for organic matter removal. Therefore, we can conclude

337

that the organic matter removal in the ferrate(VI) process predominantly comes from the

338

oxidation process, with the coagulation process making negligible contribution.

339

Oxidative Species.

340

At a pH above 8.6, the predominant ferrate(VI) species are FeO42- and HFeO4- (Figure S6).

341

Increasing the pH from 8.6 to 10.0, the FeO42- percentage increases from 95.2% to 99.8%, and

342

HFeO4- decreases from 4.8% to 0.2%. Although the concentration of FeO42- is 20 times higher

343

than HFeO4- initially, the main contributor to the oxidation in OSPW treatment is hypothesized

344

to be HFeO4- for the following reasons. (1) The redox potential of HFeO4- is around 1.4 V;

345

compared with unprotonated form FeO42-, protonated form HFeO4- has a higher spin density on

346

the oxo ligands which makes the metal center more electron deficient, and thus increases the

347

oxidation ability of HFeO4-.47, 48 FeO42-, on the other hand, has a redox potential of only 0.7 V,

348

and its contribution to the oxidation of organic matter especially classical NAs might be

349

insignificant49, 50, although having a high concentration in the solution. (2) Reported specific rate

350

constants for the reaction between organic compounds and HFeO4- are two to four orders of 16 ACS Paragon Plus Environment

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magnitude higher than those between organic compounds and FeO42-.47, 51 (3) At pH range of the

352

OSPW, the concentration of HFeO4- was high enough to make the oxidation process proceed

353

because HFeO4- and the abundant FeO42- exist in equilibrium; more HFeO4- will be generated

354

from FeO42- once consumed.

355

Oxidation Mechanisms of Aromatics and NAs.

356

As the mechanisms of ferrate(VI) oxidation varies with different organic compounds, the focus

357

in the following discussion will be on three groups of organics which were significantly removed

358

in the ferrate(VI) oxidation process: aromatics, heteroatomic naphthenic acids, and classical NAs.

359

(1) The reduction of aromaticity of the treated OSPW arose from the transformation of the

360

aromatic rings. The opening of the aromatic rings was reported by Sharma, et al.52 and

361

Anquandah, et al.53 It might occur through 1,3-dipolar cycloaddition and result in two aldehyde

362

groups53, which can be further oxidized by ferrate(VI) to acid groups. (2) The most likely

363

pathway for the removal of NAs with heteroatoms is through the formation of a complex

364

between the ferrate(VI) and target compounds, followed by inner-sphere electron transfer.51

365

Heteroatoms such as sulfur constitute electron-rich moieties, which are good ligands for

366

ferrate(VI) to form inner-sphere complexes, where electron transfer occurs between the electron-

367

rich moieties and the electron-deficient iron center of ferrate(VI). Rate constants of reactions

368

between HFeO4- and organosulfur compounds ranged from 103-106 M-1s-1

369

enough to lead to preferential removal of heteroatomic NAs. (3) For the removal of classical

370

NAs by ferrate(VI), the mechanisms have not been reported. Classical NAs lack electron-rich

371

moieties and are relatively more stable to oxidation than the other two organic fractions

372

mentioned above. The most probable attack sites for ferrate(VI) are reported to be α and β

373

carbons at the acid branches27, 55, 56 and tertiary carbons on the rings and the branches.45 1H NMR 17 ACS Paragon Plus Environment

54

, which is high


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spectra of raw OSPW and ferrate(VI)-treated OSPW showed that there was reduction of the α-C

375

hydrogen (Figure 1a), which indicates the attack on the α carbon as one of the possible oxidation

376

pathways for NAs by ferrate(VI). However, we can not quantify the reduction as the peak was

377

overlapped with other peaks. In addition, consistent with the order of reactivity of C-H (tertiary

378

hydrogen > secondary hydrogen > primary hydrogen), other studies also showed that the

379

bridgehead position of bicyclic aliphatic organics is vulnerable to the oxidation by metal

380

oxidants.57 To elucidate the oxidation mechanism of NAs by ferrate(VI), more studies with

381

model compounds are needed.

382

Electron Paramagnetic Resonance Analysis.

383

EPR analysis of the reaction mixture of POBN (the spin trap agent), OSPW and ferrate(VI)

384

detected triplet signals at g1=1.992, g2=2.001, and g3=2.015. On the other hand, no such signals

385

were detected with POBN/ferrate(VI) mixture or with POBN/OSPW mixture (Figure 3). It is

386

safe to assume that these signals derive from the generated radicals from the reaction between

387

OSPW and ferrate(VI). Due to the high g values, it can be stated that these radicals are of organic

388

nature.58 The existence of organic radicals as oxidation intermediates indicated that one of the

389

probable reaction mechanisms is one-electron transfer, which often occurs via hydrogen

390

abstraction by ferrate(VI) and results in organic radicals and ferrate(V). These organic radicals

391

can react with dissolved oxygen in the water, generating alcohols, aldehydes, and ketones which

392

can be further oxidized by ferrate(VI) or ferrate(V) until all oxidative iron species are reduced to

393

Fe(III).49 Similar “Fe(VI)Fe(V)Fe(III)” oxidation sequence was also reported in ferrate(VI)

394

oxidation of phenol.59

395

Model Compound Study.


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396

Three different model compounds which show a different structure to evaluate structure-

397

reactivity were selected for this study: CHA, with the carboxylic acid group directly attached to

398

the ring; CHPA, in the same homologous series with CHA but with four more CH2 increments

399

between the acid group and the ring; and ADA, a dicarboxylic acid with tricyclic diamond-like

400

structure and reported to exist in OSPW.14 As shown in Figure 4, CHA exhibited the highest

401

persistence towards ferrate(VI) oxidation, and its concentration only decreased by 26.3% during

402

the 3-hour experiment. In the first 30 minutes of ferrate (VI) oxidation, the concentration

403

decreased by 13.5% for CHA, 29.3% for CHPA and 32.4% for ADA. This is mainly because

404

CHA only has one possibly reactive tertiary carbon (which is also the α carbon to the carboxylic

405

acid group). On the other hand, CHPA is more reactive than CHA, with a total removal of 47.6%

406

at the end of the oxidation. This higher reactivity can be attributed to the long side chain which

407

provides more α and β carbons as reactive sites. Among the three model compounds, ADA is the

408

most reactive one with 57.1% removal at the end of the oxidation. Although ADA has fewer

409

hydrogens than CHPA, its activity is enhanced by the four tertiary carbons. These results are in

410

consistence with our previous explanation that tertiary carbons, α and β carbons to the carboxylic

411

acid groups, mainly contributed to the oxidation of classical NAs27, 55, 56, as well as the fact that

412

higher -z and carbon number NAs are more reactive towards oxidation.

413

Analysis of degradation by-products was also attempted with high resolution HPLC TOF-MS

414

(the mass spectra are provided in Figure S7-9). With oxidized CHPA (C11H20O2) solution, a

415

CHPA peak was detected at 4.82 min with m/z being 183.13849, and a composite product peak

416

was detected at 4.22 min. The product peak contained two ions with corresponding m/z being

417

199.13352 and 197.11835, which fit molecular ions of C11H20O3 and C11H18O3 respectively.

418

Most likely, after ferrate(VI) oxidation, a tertiary hydroxyl group was generated on the tertiary 19 ACS Paragon Plus Environment


Page 20 of 32

419

carbon, corresponding to the generation of C11H20O3. Tertiary alcohols are stable to ferrate(VI)

420

oxidation26 and can thus be detected. The α carbon reaction might also generate a hydroxyl

421

group in the intermediate product, but it will soon be further oxidized by ferrate(VI) to ketones,

422

which might be the second product we detected as C11H18O3. Additionally, a ketone group on the

423

β carbon might also contribute to the formation of C11H18O3 (this is supported by the fact that the

424

ketone peak is a composite peak combining at least two distinct peaks (Figure S8)). Ketone

425

groups are also resistant to ferrate(VI) oxidation, and can therefore survive and be detected at the

426

end of the reaction26. Similarly, two products, C7H12O2 and C7H10O3, were identified as oxidized

427

products of CHA (Figure S8); another two products, C12H16O5 and C12H14O5 were detected as

428

oxidized products of ADA (Figure S9). Most likely, these products keep the carboxylic acid

429

groups intact and have additional hydroxyl or ketone groups generated as discussed in CHPA

430

oxidation. However, we cannot rule out the possibilities of the generation of other products,

431

including products that are not carboxylic acids (which might not be detected under the negative

432

ESI mode). Nevertheless, these results are consistent with our previous conclusion from OSPW

433

study that more “NAs+Ox” were generated after classical NAs were oxidized.

434

Toxicity

435

As a prescreening tool for acute toxicity, Microtox® 81.9% screening test was adopted to

436

evaluate the toxicity of OSPW on Vibrio fischeri. The inhibition level of raw OSPW was

437

42.8±0.6%, which was close to the previously reported values of 40-45%.34, 60 Overall, ferrate(VI)

438

oxidation decreased the inhibition effect significantly to 19.6±1.6% at 200 mg/L Fe(VI) but a

439

further increase of ferrate(VI) dose to 400 mg/L Fe(VI) only reduced the inhibition to 17.6±0.1%

440

(Figure 5a). At the same time, NAs removal increased from 64.0% to 78.4%, indicating poor

441

correlation between NAs removal and acute toxicity reduction. Klamerth et al.’s study18 also 20 ACS Paragon Plus Environment

Page 21 of 32


442

showed that low concentration of residual NAs after ozonation did not guarantee low toxicity.

443

Similar poor correlation was reported by Zhang, et al.61 in biological treatment of OSPW: after

444

treatment, toxicity was reduced by 49.3% although there was no significant NAs removal. On the

445

other hand, there is a linear correlation (empirical) between toxicity reduction and SFS

446

(fluorescing compounds) peak area decrease (Figure S10). At low doses of 1-50 mg/L of Fe(VI)

447

the inhibition effect decreased from 42.8±1.6% at 1 mg/L to 30.7±1.8% at 50 mg/L with

448

decreasing peak area II and III (peak area decrease was calculated to be 2% at 1 mg/L Fe(VI) and

449

75% at 50 mg/L Fe(VI)), while no significant decrease in peak area I could be observed. At

450

higher ferrate(VI) doses from 200 mg/L to 400 mg/L Fe(VI), there was no further decrease of

451

peak II and III while only 10% more decrease of peak area I was observed, resulting in no

452

significant reduction of toxicity (Figure 5a). Therefore, it can be concluded that although the

453

concentration of the fluorescing compounds is relatively low compared to that of NAs, the

454

toxicity effect from them might be significant. This conclusion is consistent with the widely

455

reported toxicity of the aromatics especially PAHs.62-64 In OSPW studies, Rowland, et al.65 found

456

that the estrogenic effect of OSPW arose from the aromatics in the water, and Scarlett, et al.14

457

also suggested that aromatics in OSPW were at least as toxic as NAs. Further studies are needed

458

to compare the toxic effects of aromatics and NAs in OSPW.

459

The effects of the OSPW treated samples on goldfish primary kidney macrophages (PKMs) were

460

examined by measuring the production of reactive nitrogen intermediates (RNIs) as shown in

461

Figure 5b. Macrophages are large immune cells capable of destroying pathogens by producing

462

toxic RNIs, such as nitric oxide; the production of RNI by macrophages was quantified using the

463

Griess reaction.66 High RNI production indicates high functional activity of PKMs and

464

correspondingly, low toxicity of the test sample. The exposure of PKMs to the raw and



Page 22 of 32

465

ferrate(VI)-treated OSPW samples caused a reduction in RNI production compared to positive

466

control (heat-killed A. salmonicida) (38.9 ± 4.3 µM RNI), showing higher toxicity induced by

467

raw OSPW (8.8 ± 1.3 µM RNI) and significantly lower toxicity induced by treated OSPW (18.7

468

± 1.1 µM RNI at 100 mg/L Fe(VI)). This results, together with the Microtox® results, indicate

469

that ferrate(VI) treatment is effective in reducing toxicity towards prokaryotic organisms such as

470

Vibrio ﬁscheri and eukaryotic cells such as goldfish PKMs.

471

pH Adjustment. To study the pH effect on ferrate(VI) oxidation, the initial pH of OSPW was

472

adjusted to 6 and the oxidation experiments were repeated as at natural pH. The results were

473

summarized in Figure S11. This adjustment improved the performance of the ferrate(VI)

474

oxidation in terms of both aromatics removal (Figure S11a&b) and NAs removal (Figure S11c).

475

The trend of the SFS peak reduction was similar to that achieved at natural pH: peak II and III

476

were preferentially removed, and peak I was only significantly reduced at high doses. At the

477

same dose of ferrate(VI), oxidation at initial pH 6 achieved higher removal of fluorescing

478

aromatics than at natural pH, which was mainly caused by the higher removal of single-ring

479

aromatics (peak I) at initial pH 6 (Figure S11a&b). At the same time, inhibition effect on Vibrio

480

fischeri was also reduced to 16.3±0.8% at initial pH 6 at the dose of 400 mg/L Fe(VI), lower

481

than achieved at natural pH (17.6±0.1%) (Figure S11d). However, these improvements on

482

removal of aromatics and reduction of toxicity were not significant compared with the

483

enhancement of NAs removal. At initial pH of 6, ferrate(VI) achieved 78.4% removal of NAs at

484

200 mg/L Fe(VI), equivalent to the removal at 400 mg/L Fe(VI) at natural pH. As mentioned in

485

section 3.4, this also indicates the poor correlation between toxicity effect and NAs removal.

486

Environmental Implications. Although NAs and aromatics are two main fractions of organic

487

matter in OSPW, most of the current studies of OSPW treatment focus on the removal of NAs; 22 ACS Paragon Plus Environment

Page 23 of 32


488

the toxicity of OSPW has been presumably attributed to NAs. However, as shown in this study,

489

the toxicity effect of the aromatics might also be significant. More work is warranted to

490

investigate the concentrations, species, and the environmental impact of the aromatics in OSPW.

491

The results of this study demonstrated that ferrate(VI) oxidation is an effective process in

492

transforming NAs and the aromatics and reducing the acute toxicity of OSPW. However, when

493

high doses of ferrate(VI) (e.g., 200 mg/L and 400 mg/L Fe(VI)) are applied, large quantity of

494

sludge will be generated. The disposal of the sludge in an environmentally and economically

495

feasible way is a challenging topic to be addressed.

496

ASSOCIATED CONTENT

497

Supporting Information. Supplementary graphs, and detailed description of analytical methods

498

and calculation procedures are provided in the Supporting Information. This material is available

499

free of charge via the Internet at http://pubs.acs.org.

500

AUTHOR INFORMATION

501

Corresponding Author

502

*Phone: +1 780 492 5124; fax: +1 780 492 8190; email: [email protected].

503

ACKNOWLEDGMENT

504

This research was supported by a research grant from the Alberta Water Research Institute, the

505

Helmholtz-Alberta Initiative, and a Natural Sciences and Engineering Research Council of

506

Canada (NSERC) Senior Industrial Research Chair (IRC) in Oil Sands Tailings Water Treatment

507

through the support by Syncrude Canada Ltd., Suncor Energy Inc., Shell Canada, Canadian

508

Natural Resources Ltd., Total E&P Canada Ltd., EPCOR Water Services, IOWC Technologies

509

Inc., Alberta Innovates - Energy and Environment Solution, and Alberta Environment and 23 ACS Paragon Plus Environment


Page 24 of 32

510

Sustainable Resource Development. The authors would also like to thank Dr. Kerry N.

511

McPhedran for his help on experimental design Dr. Arvinder Singh for toxicity measurement, Dr.

512

Randy Whittal for model compound detection, and Dr. Selamawit Ashagre Messele for

513

manuscript revision.

514 515

516

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Sharma, V. K.; Smith, J. O.; Millero, F. J., Ferrate(VI) Oxidation of Hydrogen Sulfide. Environ. Sci. Technol. 1997, 31, (9), 2486-2491. Ho, C.-M.; Lau, T.-C., Lewis acid activated oxidation of alkanes by barium ferrate. New J. Chem. 2000, 24, (8), 587-590. Zimmermann, S. G.; Schmukat, A.; Schulz, M.; Benner, J.; Gunten, U. V.; Ternes, T. A., Kinetic and mechanistic investigations of the oxidation of tramadol by ferrate and ozone. Environ. Sci. Technol. 2012, 46, (2), 876-884. Sharma, V. K.; Mishra, S. K.; Nesnas, N., Oxidation of sulfonamide antimicrobials by ferrate(VI) [FeVIO42-]. Environ. Sci. Technol. 2006, 40, (23), 7222-7227. Sharma, V., K.; Noorhasan, N., N.; Mishra, S., K.; Nesnas, N., Ferrate(VI) oxidation of recalcitrant compounds: Removal of biological resistant organic molecules by ferrate(VI). In Ferrates, American Chemical Society: 2008; Vol. 985, pp 339-349. Anquandah, G. A. K.; Sharma, V. K.; Panditi, V. R.; Gardinali, P. R.; Kim, H.; Oturan, M. A., Ferrate(VI) oxidation of propranolol: Kinetics and products. Chemosphere 2013, 91, (1), 105-109. Sharma, V. K., Ferrate(VI) and ferrate(V) oxidation of organic compounds: Kinetics and mechanism. Coord. Chem. Rev. 2013, 257, (2), 495-510. Sharma, V. K.; Bielski, B. H. J., Reactivity of ferrate(VI) and ferrate(V) with amino acids. Inorg. Chem. 1991, 30, (23), 4306-4310. Drzewicz, P.; Afzal, A.; Gamal El-Din, M.; Martin, J. W., Degradation of a model naphthenic acid, cyclohexanoic acid, by vacuum UV (172 nm) and UV (254 nm)/H2O2. J. Phys. Chem. A 2010, 114, (45), 12067-12074. Wiberg, K. B.; Foster, G., The stereochemistry of the chromic acid oxidation of tertiary hydrogens1. J. Am. Chem. Soc. 1961, 83, (2), 423-429. Huang, H.; Sommerfeld, D.; Dunn, B. C.; Lloyd, C. R.; Eyring, E. M., Ferrate(VI) oxidation of aniline. J. Chem. Soc., Dalton Trans. 2001, (8), 1301-1305. Huang, H.; Sommerfeld, D.; Dunn, B. C.; Eyring, E. M.; Lloyd, C. R., Ferrate(VI) oxidation of aqueous phenol: Kinetics and mechanism. J. Phys. Chem. A 2001, 105, (14), 3536-3541. Afzal, A.; Drzewicz, P.; Perez-Estrada, L. A.; Chen, Y.; Martin, J. W.; Gamal El-Din, M., Effect of molecular structure on the relative reactivity of naphthenic acids in the UV/H2O2 advanced oxidation process. Environ. Sci. Technol. 2012, 46, (19), 10727-34. Zhang, Y.; McPhedran, K. N.; Gamal El-Din, M., Pseudomonads biodegradation of aromatic compounds in oil sands process-affected water. Sci. Total Environ. 2015, 521-522, 59-67. Kupryianchyk, D.; Rakowska, M. I.; Grotenhuis, J. T. C.; Koelmans, A. A., Modeling trade-off between PAH toxicity reduction and negative effects of sorbent amendments to contaminated sediments. Environ. Sci. Technol. 2012, 46, (9), 4975-4984. Lee, J.-H.; Landrum, P. F.; Koh, C.-H., Toxicokinetics and time-dependent PAH toxicity in the amphipod Hyalella azteca. Environ. Sci. Technol. 2002, 36, (14), 3124-3130. Sayles, G. D.; Acheson, C. M.; Kupferle, M. J.; Shan, Y.; Zhou, Q.; Meier, J. R.; Chang, L.; Brenner, R. C., Land treatment of PAH-contaminated soil: performance measured by chemical and toxicity assays. Environ. Sci. Technol.1999, 33, (23), 4310-4317. Rowland, S. J.; West, C. E.; Jones, D.; Scarlett, A. G.; Frank, R. A.; Hewitt, L. M., Steroidal aromatic 'naphthenic acids' in oil sands process-affected water: structural comparisons with environmental estrogens. Environ. Sci. Technol. 2011, 45, (22), 9806-9815. Stone, W.; Yang, H.; Qui, M., Assays for nitric oxide expression. In Mast Cells, Krishnaswamy, G.; Chi, D., Eds. Humana Press: 2005; Vol. 315, pp 245-256.

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Figures 27 ACS Paragon Plus Environment


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Figure 1. (a) 1H NMR spectral for 200 mg/L Fe (VI)-treated OSPW (the numbers besides the peaks indicate the relative peak area when the DMSO solvent peak was set to be 100 as internal standard) and (b) SFS of raw and treated OSPW.

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707 708 709 Figure 2. Ion mobility spectra for (a) raw OSPW, (b) ferrate(VI)-treated OSPW (200 mg/L), and (c)

ferrate(VI)-treated OSPW (400 mg/L); NAs distribution in (d) raw OSPW, (e) ferrate(VI)-treated 710 OSPW (200 mg/L) and (f) ferrate(VI)-treated OSPW (400 mg/L); and NAs removal with respect of (g) 711 carbon number and (h) –z number.



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Figure 3. EPR spectrum of the reaction mixture of OSPW, 100 mg/L Fe(VI) and 10g/L POBN.

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Figure 4. Model compound degradation by ferrate(VI) (dose: 200 mg/L Fe(VI)) at different reaction time (Note: initial concentration of each model compound is 60 mg/L).

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Figure 5. (a) Acute toxicity to Vibrio fischeri of raw OSPW and ferrate(VI)-treated OSPW (Note: error bars indicate mean±SD with n=3); (b) Nitrite production by PKMs exposed to medium control, raw OSPW, and OSPW treated with ferrate (VI) with the dose of 25 mg/L and 100 mg/L Fe(VI). Positive and negative controls were heat killed A. salmonicida and culture medium, respectively. Nitric oxide production was determined using the Griess reaction and RNI concentration was determined using RNI standard curve. Data are represented as RNI production by macrophages from separate cultures established from three individual fish. Statistical analysis was done using one-way ANOVA with a Dunnett’s post hoc test. Asterisks (*) indicate statistically significant increases (p

Oxidation of Oil Sands Process-Affected Water by Potassium Ferrate

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