Activation of Manganese Oxidants with Bisulfite for Enhanced

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Activation of Manganese Oxidants with Bisulfite for Enhanced Oxidation of Organic Contaminants: The Involvement of Mn(III) Bo Sun, Xiaohong Guan, Jingyun Fang, and Paul G. Tratnyek Environ. Sci. Technol., Just Accepted Manuscript • Publication Date (Web): 30 Sep 2015 Downloaded from http://pubs.acs.org on October 5, 2015

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Environmental Science & Technology

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Activation of Manganese Oxidants with Bisulfite for

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Enhanced Oxidation of Organic Contaminants:

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The Involvement of Mn(III)

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Bo Sun1,2, Xiaohong Guan1*, Jingyun Fang3, and Paul G. Tratnyek4* 1

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State Key Laboratory of Pollution Control and Resources Reuse, Tongji University, Shanghai 20092, P. R. China

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2

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State Key Laboratory of Urban Water Resource and Environment, Harbin Institute of Technology, Harbin 150090, P. R. China

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3

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Sun Yat-Sen University, Guangzhou 510275, China

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School of Environmental Science and Engineering,

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Institute of Environmental Health, Oregon Health & Science University,

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3181 SW Sam Jackson Park Road, Portland, Oregon 97239-3098, USA

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*Contact/Corresponding author contact information:

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Email: [email protected] (X.H. Guan); Phone: +86-21-65980956.

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E-mail: [email protected] (P.G. Tratnyek); Phone: 503-346-3431.

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9/26/15 9:25 AM

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1 ACS Paragon Plus Environment


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ABSTRACT: MnO4- was activated by HSO3-, resulting in a process that oxidizes

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organic contaminants at extraordinarily high rates. The permanganate/bisulfite

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(PM/BS) process oxidized phenol, ciprofloxacin, and methyl blue at pHini 5.0 with

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rates (kobs ≈ 60-150 s−1) that were 5-6 orders of magnitude faster than those measured

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for permanganate alone, and ~5 to 7 orders of magnitude faster than conventional

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advanced oxidation processes for water treatment. Oxidation of phenol was fastest at

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pH 4.0, but still effective at pH 7.0, and only slightly slower when performed in tap

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water. A smaller, but still considerable (~3 orders of magnitude) increase in oxidation

27

rates of methyl blue was observed with MnO2 activated by HSO3- (MO/BS). The

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above results, time-resolved spectroscopy of manganese species under various

29

conditions, stoichiometric analysis of pH changes, and the effect of pyrophosphate on

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UV absorbance spectra suggest that the reactive intermediate(s) responsible for the

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extremely rapid oxidation of organic contaminants in the PM/BS process involve

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manganese(III) species with minimal stabilization by complexation. The PM/BS

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process may lead to a new category of advanced oxidation technologies based on

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contaminant oxidation by reactive manganese(III) species, rather than hydroxyl and

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sulfate radicals.

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INTRODUCTION

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Advanced oxidation processes (AOPs) are effective at removing many organic

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contaminants from water because they generate strong radical oxidants such as

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hydroxyl radical (HO•) and sulfate radical (SO4•−). Hydroxyl and sulfate radicals are

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known to rapidly oxidize organic compounds at nearly diffusion controlled rates (i.e.,

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second-order rate constants, k″ ≈ 109-1010 M−1 s−1).1, 2 However, even in optimized

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AOPs, the maximum concentrations of HO• and SO4•− are low (10-12-10-14 M),3 so the

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apparent degradation rates of some contaminants can be slow. Perhaps the most

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promising approach to overcoming this limitation of conventional AOPs would be a

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process that is mediated by strongly oxidizing species that can be generated at higher

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concentrations than HO• and SO4•−. Here we describe a novel, and very promising

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example of this approach, which involves the activation of manganese oxidants

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(permanganate and manganese dioxide) to form reactive Mn(III) species.

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The oxidation of contaminants by permanganate has been studied extensively

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because of its environmental applications for remediation of contaminated soil and

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groundwater,4-6 and water treatment to control various organic pollutants,7-10

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dissolved manganese/iron, taste/odor, etc.11 Compared to other chemical oxidants, the

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advantages of permanganate include modest cost, easy and safe storage and delivery,

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applicability over a wide range of conditions, and no tendency to form chlorinated or

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brominated byproducts.4 Another common manganese oxidant, manganese dioxide

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(MnO2), has also been used to oxidize various organic compounds including

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antibiotics,12 anilines,13 phenols,14 steroid estrogens,15 etc.

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Although permanganate is widely regarded as a strong oxidant, the rates of

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oxidation by permanganate are highly variable, and moderate to slow for some

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important contaminants.16 Many recent studies have investigated the kinetics of

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contaminant oxidation by permanganate, including the determination of new rate

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constants,7,

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relationships,19, 21 and enhancement of reaction rates by catalysis.22, 23 In most cases,

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the experimental work for these studies has been done under pseudo first-order

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conditions, with permanganate in excess, so any residual permanganate must be

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quenched to stabilize the test contaminant concentrations until they are analyzed. The

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reductants that have been used for this purpose include Na2S2O3,24, 25 Na2SO3,26 and

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NH2OH•HCl.9, 27 However, very little effort has been devoted to evaluation of these

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choices, even though protocols for quenching permanganate are likely to be affected

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by the same issues that are well documented in studies of oxidation by peroxides.28-30

17, 18

constructing more powerful kinetic models,19,

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structure-activity

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As part of our work on the kinetics of contaminant oxidation by permanganate,

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we tested several inorganic reductants, over a range of conditions, to evaluate their

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suitability as quenchers of residual permanganate. Surprisingly, we discovered that

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adding Na2SO3 resulted in greatly accelerated rates of organic contaminant oxidation.

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To our knowledge, this effect has never been recognized previously in the context of

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water treatment processes using permanganate. Since sulfite and bisulfite (SO32−, 26-Sep-15


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HSO3−) are species that could be applied in environmental engineering processes, the

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effect described here could become the basis for a novel advanced oxidation process

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(AOP) involving activation of permanganate. Therefore, the objectives of this study

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were to (1) investigate the kinetics of organic contaminant degradation by HSO3-

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activated MnO4- (PM/BS) over a range of relevant solution conditions; (2) explore the

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possibility of oxidizing permanganate-refractory contaminants by the PM/BS process;

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(3) provide a preliminary characterization of the contaminant oxidation products and

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pathway, and (4) identify the intermediate oxidant(s) responsible for the enhancement

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of contaminant oxidation in this process.

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EXPERIMENTAL SECTION

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Materials.

Potassium

permanganate

(GR

grade),

sodium

thiosulfate

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pentahydrate (GR grade), phenol (99% pure) and sodium oxalate (AR grade) were

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purchased from the Tianjin Chemical Reagent Co., Ltd. (Tianjin, China). Sodium

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bisulfite (AR grade), and sodium persulfate (AR grade) were obtained from Chinasun

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Specialty Products Co., Ltd. (Jiangsu, China). Methyl blue (AR grade), hydrogen

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peroxide (30 wt%), nitrobenzene (AR grade) and sodium pyrophosphate (PP, AR

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grade) were supplied by the Sinopharm Chemical Reagent Co., Ltd. (Shanghai,

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China). Ciprofloxacin (AR grade) and caffeine (AR grade) was purchased from the

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Aladdin Industrial Corporation. Methanol (99.9% pure) was supplied by Merck KgaA

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(Germany). All chemicals were used as received.

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The KMnO4 crystal was dissolved in Milli-Q water to make a 50 mM stock

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solution. The stock solution of NaHSO3 (250 mM) and MnCl2 (50 mM) were freshly

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prepared for each set of experiments to avoid oxidation by oxygen. The stock solution

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of phenol (1.0 mM), ciprofloxacin (1.0 mM), methyl blue (1.0 mM), Na2S2O3 (100.0

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mM) and Na4P2O7 (50.0 mM) were prepared in Milli-Q water every day. A stable

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colloidal MnO2 stock solution was prepared freshly before use following the

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procedure in literature31 by mixing the appropriate amounts of Mn(VII) and Na2S2O3

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stock solutions. The Mn(III)-PP stock solution was synthesized by mixing 500 µL

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KMnO4 (50 mM) and 500 µL NaHSO3 (250 mM) in the presence of 50 mL Na4P2O7

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(50 mM) and 449 mL Milli-Q water at pH 5.0.

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Experimental Procedures. A stopped-flow spectrophotometer (SFS, Model

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SX20, Applied Photophysics Ltd., Leatherhead, UK) was used to conduct the rapid

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kinetic experiments. A photodiode array for acquisition of multi-wavelength

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absorption, a UV-visible spectrophotometer and a fluorimeter were used as the

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detectors with a 150 W xenon lamp as the light source. An HP computer workstation

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was employed to control the stopped-flow and acquire the kinetic data.

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Before the stopped-flow kinetic experiments, two working solutions were

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prepared. One working solution contained 100 µM permanganate and the other

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contained 500 µM bisulfite and substrates of interest. The solutions were adjusted to

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the target pH levels by adding HCl or NaOH. Reactions were initiated by

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simultaneously injecting an equal volume of two working solutions into the optical 26-Sep-15


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cell of the SFS with two automatic syringes driven by compressed nitrogen. Three

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organic substrates—including methyl blue, phenol, and ciprofloxacin—were selected

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as the test contaminants in the SFS experiments because they could be detected in situ

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by UV and fluorescence methods without interference of MnO4-/MnO2, and they have

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frequently been detected in wastewater.

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The 3-D UV-visible spectrum of the reaction between HSO3- and MnO4- in the

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presence or absence of pyrophosphate or phenol was conducted using photodiode

126

array at 350-725 nm. To verify the generation of Mn(III) in the reaction between

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HSO3- and MnO4- or MnO2, the change of UV absorbance at 258 nm with time was

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determined with the stopped-flow spectrophotometer. In separate batch experiments,

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the reaction products generated in the reaction between bisulfite and KMnO4 or MnO2,

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in the presence or absence of pyrophosphate or phenol, after reacting for 10 minutes

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were characterized with the UV-vis spectrum collected with a Purkinje TU-1902

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automatic scanning UV-visible spectrophotometer at 200-800 nm with wavelength

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program controllers.

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The kinetic experiments with the reaction time > 5 minutes were conducted with

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glass bottles open to the air. Reactions were initiated by quickly spiking KMnO4,

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MnO2, or Mn(III)-PP into the solution containing the test contaminant (phenol,

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nitrobenzene, or caffeine) and the constituents of interest at initial pH (pHini) 5.0

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while they were being mixed with a magnetic stirrer. Sodium acetate (1 mM) was

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used as a buffer for the reactions at pH 5.0 in the oxidation of contaminants by 26-Sep-15



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KMnO4/MnO2 alone. The negligible effects of acetate on these reaction have been

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discussed previously.25 Periodically, 10.0 mL of sample was rapidly transferred into a

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25 mL beaker, immediately quenched with 100 µL of a sodium thiosulfate stock

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solution, and then subject to analysis with an

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chromatography (UPLC). To determine the mineralization of oxalate in the PM/BS

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process, pH of the solution containing 200 µM sodium oxalate and 1000 µM NaHSO3

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was adjusted to 5.0 before application of 500 µM KMnO4. Samples were collected

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periodically, quenched and analyzed for total organic carbon (TOC). All kinetic

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experiments were carried out in at least triplicate at 18±2 oC, and the data were

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averaged with the standard deviations < 5% unless otherwise noted. During all of the

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experiments, the initial pH was adjusted to the pre-determined value, but pH was not

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controlled constant in any other way during the reaction without buffer.

ultra-performance liquid

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Chemical Analysis. In the SFS experiments, the concentration of methyl blue

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was continuously monitored at 626 nm with a UV-Visible spectrophotometer. The

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change in concentration of phenol and ciprofloxacin were continuously detected by

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fluorimetry at Ex/Em = 272 nm/298 nm and 277 nm/450 nm, respectively. The

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concentration of phenol, nitrobenzene, and caffeine in the samples taken from batch

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experiments was quantified by UPLC (Waters Co.). Separation was accomplished

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with an UPLC BEH C18 column (2.1 × 100 mm, 1.7 µm; Waters Co.) at 35 ± 1 oC

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and a mobile phase of methanol-0.1% formic acid aqueous solution (from 40:60 to

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70:30). The flow rate was 0.3 mL min-1 and the largest volume injection was 10 µL. 26-Sep-15


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Concentrations of compounds were determined by comparing the peak area at

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254-273 nm with that of the corresponding compounds standards. The TOC analysis

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was performed with a TOC analyzer (TOC-LCPH, SHIMADZU). The variation of pH

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values of the reaction solutions were measured by a pH meter with a saturated KCl

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solution as an electrolyte. Daily calibration with proper buffer solution (pH 4.00 and

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6.86) was performed to ensure its accuracy.

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DMPO (5,5-dimethyl-1-pyrrolidine-N-oxide) was used as the spin-trapping agent

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in the electron spin resonance (ESR) experiments. KMnO4 or MnCl2 was added into

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the solution containing DMPO and NaHSO3 at pH 5.0, and the mixed solution was

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then inserted into the cavity of the ESR spectrometer (Bruker EMX-8/2.7).

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Measurements were carried out under the following conditions: a center field of 3517

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Gs, a sweep width of 100 Gs, a microwave frequency of 20 mW, a modulation

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amplitude of 1 Gs and a sweep time of 41.96 s.

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RESULTS AND DISCUSSION

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Contaminant Oxidation in the PM/BS Process. The disappearance kinetics of

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phenol, ciprofloxacin, and methyl blue in the PM/BS process were determined with

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stopped-flow spectroscopy and the results are shown in Figure 1(A-C). At pHini 5.0,

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complete disappearance of these three test contaminants was observed in 40-80 ms.

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MnO4- alone oxidized these test contaminants much more slowly, as shown by the

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inset graphs in Figure 1, and there was no detectable reaction with NaHSO3 alone (not 26-Sep-15



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shown). To model the kinetics for the PM/BS process and the MO/BS process (MnO2

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and NaHSO3, introduced below), we fit the majority of the data to pseudo first-order

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kinetics after excluding a few data points that suggest an initial lag phase (Figure

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1A-D). Kinetic data obtained with MnO4− and MnO2 in the absence of HSO3− (Figure

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1A-D insets) exhibit no lag, so these data were fit as pseudo first-order initial rates.

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The pseudo first-order rate constants (kobs) obtained by fitting the data in Figure

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1 are given in Table 1. For the three contaminants tested, kobs is on the order of 100 s−1

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with PM/BS, 0.1 s−1 for MO/BS, and 0.0001 s−1 for MnO4− and MnO2 only. Thus, the

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enhancement of contaminant oxidation rates by activating MnO4− (and MnO2) with

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bisulfite is roughly three to six orders-of-magnitude, under the conditions tested (e.g.,

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pHini = 5.0). For comparison with these very high reaction rates, we compiled the

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previously published kinetic data for degradation of the test contaminants by

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conventional and established advanced oxidation processes, and summarized those

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values in Tables S1-S2 (for phenol and ciprofloxacin only, insufficient data are

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available on methylene blue). The literature values of kobs typically fall in the range

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10−3 to 10−5 s−1, which are 5 to 7 orders of magnitude slower than those reported here

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for the PM/BS process. Additional exploratory experiments showed that the PM/BS

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process at pHini 5.0 can even oxidize contaminants that are relatively recalcitrant to

200

oxidation by permanganate: such as nitrobenzene and caffeine, which gave t1/2 on the

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order of ms (Figure S1).

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The influence of initial pH on the degradation of phenol in the PM/BS process

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was investigated and the results are shown in Figure 2. The lag phase in the pHini 5.0

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data (Figure 1A, Figure 2B), is even more prominent at pHini 6.0, but not evident in

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the higher or lower pH data, which is fully explained by the reaction stoichiometry as

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described in the next section. At pHini 4.0-7.0, 5 µM phenol was completely degraded,

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but only ~50% was removed at pHini 8.0-9.0. The inefficiency of the PM/BS process

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for phenol decomposition at pHini 8.0-9.0 may be due to a shift in speciation from

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bisulfite to sulfite (pKa = 7.2)32 and the higher disproportionation rate of Mn(III)

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under alkaline conditions.

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Neglecting the initial lag and final plateau data, the remaining data in Figure 2

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were fit to pseudo first-order kinetics and the resulting values of kobs are given in

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Table 1. The rate constants for the PM/BS process decreased about 100-fold with

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pHini increased from 4.0 to 8.0. The very fast decomposition of phenol over the pHini

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range of 4.0-7.0 in the PM/BS process indicates that the PM/BS process may be

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effective for removing organic contaminants under neutral and acidic conditions.

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The performance of AOPs can be affected by other solution conditions (besides

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pH) in a variety of ways, such as the scavenging of hydroxyl radicals by carbonate33

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and the chlorination of organics by chlorine radical formed from chloride.34 To

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provide a preliminary and practical assessment of the sensitivity of the PM/BS

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process to solution conditions, the kinetics of phenol oxidation were studied in

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solutions made from tap water of our lab (DOC = 2.4 mg C L−1, alkalinity = 0.34 mM 26-Sep-15



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as HCO3−, [Fe3+] + [Fe2+] = 2 µM).35 The results, shown in Figure S2, indicate there is

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negligible difference in the kinetics of phenol oxidation between tap and DI water at

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pHini 5.0, and this reaction is only slightly slower in tap water at pHini 7.0.

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Measurements of TOC at the end of the stopped-flow experiments showed no

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evidence of mineralization, but complete oxidation of the contaminants was not

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expected under the conditions of these experiments (based on stoichiometric

229

calculations). However, as with other AOPs, the PM/BS process is expected to

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produce intermediate organic oxidation products that include carboxylic acids,

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aldehydes, etc.36,

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conventional AOPs have shown that oxalic acid accumulates (because it is relatively

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unreactive with •OH), which makes oxalate formation a key barrier to

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mineralization.38 A preliminary experiment designed to test for oxalate degradation by

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the PM/BS process showed ~90.0% mineralized in 60 min at pHini 5.0 (Figure S3).

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This suggests that the PM/BS process may provide more complete oxidation of

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organics than conventional AOPs, which could be a major advantage of the former,

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and so will be investigated further in detail in a future study.

37

Detailed studies of products/pathways of phenol oxidation by

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Reactive Intermediates in the PM/BS Process. To characterize the overall

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reaction involved in the PM/BS process and identify the intermediate oxidant

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responsible for the reaction, stopped flow experiments were performed with

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photodiode array detection, and the resulting time-resolved absorbance spectra are

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presented in Figure 3. In all four cases, the spectrum of KMnO4 was dominant at the 26-Sep-15


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beginning (cf., reference spectra in Supporting Information, Figure S4), but its strong

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absorbance at 300-350 and 500-570 nm disappeared by about 100 ms, consistent with

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the rapid oxidation of phenol in the presence of NaHSO3. By 200 ms, a broad

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absorbance shoulder developed at < 500 nm (characteristic of colloidal MnO239) that

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was strong for the reaction of permanganate with bisulfite alone and with a low

249

concentration of phenol (Figure 3A, B), but less strong for the high phenol

250

concentration (Figure 3C). This trend was consistent with a rapid but multi-step

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reduction of KMnO4 to MnO2 via reactive intermediates that could be

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scavenged/diverted by oxidizing organics (such as phenol). In the presence of

253

pyrophosphate, the generation of MnO2 was suppressed (Figure 3D), suggesting a

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strong interaction between the reactive species and pyrophosphate.

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A variety of reactive intermediate species might be formed in the PM/BS process,

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one or more of which undoubtedly are responsible for the rapid oxidation of the test

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contaminants shown in Figure 1 and 2. Plausible reactive intermediates include

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various forms of the less stable oxidation states of Mn (VI, V, IV, and III),

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lower-valent species of sulfur (SO3•−, SO4•−, SO5•−), and reactive oxygen species (HO•

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etc.). Some of these species are known to oxidize organic contaminants at rates that

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are nearly diffusion controlled (i.e., second-order rate constants, k″ ≈ 109–1010 M−1 s−1)

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(e.g.,1, 2), although others are not as well studied. Assuming that k″ ≈ 109 M−1 s−1 is

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representative of the specie(s) responsible for contaminant oxidation in the PM/BS

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process, then the range of kobs measured in this study (Table 1) suggests that the total 26-Sep-15



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reactive species concentration during the measurement time period was ~0.05 to ~0.2

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µM. Such concentrations are unrealistic for highly reactive species such as HO• and

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SO4•−,3 suggesting that the intermediates responsible for the high rates of contaminant

268

oxidation are relatively direct products of the initial reaction between BS and PM,

269

such as sulfite radical (SO3•−) or intermediate-valence Mn species.

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To provide direct evidence regarding the free radical species that might be

271

involved in the PM/BS process, electron spin resonance (ESR) spectroscopy was

272

performed using 5,5-dimethyl-1-pyrolin-N-oxide (DMPO) as a spin trap. The PM/BS

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process after 2 min of reaction time and the MO/BS process after 10 min of reaction

274

time gave ESR spectra (Figure S5A, B) with hyperfine coupling constants αN = 14.53

275

and αβ-H = 16.12, which is consistent with the previously published ESR spectrum for

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the DMPO/SO3•− adduct.40 Distinctly different spectra were obtained for HO• from

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Fe0 activated H2O2 and for SO4•− from heat activated S2O82− (Figure S5E, and F,

278

respectively). Prior work on Mn2+/NaHSO3 process has shown that the reaction

279

proceeds via multiple chain propagation steps involving reactive sulfur species (SO3•−,

280

SO4•−, SO5•−, HSO5−, and S2O72−),41 and the ESR spectrum of this process is also

281

consistent with SO3•− (Figure S5C). Although the intensity of SO3•− peaks in the ESR

282

spectrum of the Mn2+/NaHSO3 system increased with increasing reaction time, phenol

283

added into this system disappeared negligibly in 10 min (data not shown). Thus, the

284

initial, very rapid oxidation of the test contaminants is not due to reactive sulfur

285

species obtained by Mn(II) catalyzed oxidation of NaHSO3. 26-Sep-15


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The remaining candidates for reactive intermediates that might account for the

287

very fast contaminant oxidation by the PM/BS process are species of Mn(VI), Mn(V),

288

Mn(IV), and/or Mn(III). However, negligible absorbance at 660 nm (corresponding to

289

Mn(V) reported by Simandi42) and 610 nm (corresponding to Mn(VI) reported by

290

Hassan43) was observed in the process of permanganate reduction by bisulfite, as

291

shown in Figure 3. Moreover, MnO2 generated from permanganate reduction by

292

bisulfite degraded organic contaminant negligibly even in 1 s (not shown). The

293

possibilities are further narrowed by the results shown in Figure 1D, where MnO2 was

294

used in place of KMnO4. In this case, disappearance of methyl blue was > 90% in ~10

295

s, with kobs ≈ 0.3 s−1 (Table 1), which is about 3000-fold faster than methyl blue

296

oxidation by MnO2 without activation by HSO3−. Compared with permanganate, the

297

reduction rate of MnO2 is much slower implying the gradual generation of active

298

intermediates. The analogous effect of NaHSO3 on the oxidation of organic

299

contaminants by KMnO4 and MnO2 suggests that similar reactive intermediates might

300

be responsible. Considering that MnO2 is Mn(IV), reduction of Mn(IV) cannot form

301

Mn(V) or Mn(VI), and the UV absorbance corresponding to Mn(V) or Mn(VI) was

302

not detected, so Mn(III) is the only candidate among the remaining reactive

303

intermediates that the PM/BS and MO/BS processes have in common.

304

Mn(III) stabilized by complexation with PP has an absorbance peak at 258 nm

305

(Figure S4), which is commonly used in studies of the role of Mn(III) in various

306

processes.44 We used SFS to measure absorbance at 258 nm over time in the PM/BS 26-Sep-15



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process under different conditions (at pHini 5.0), and the results are shown in Figure

308

S6. In the presence of PP but without organic contaminant, the absorbance increased

309

for about 0.2 s and then leveled out at a value consistent with [Mn(III)-PP] ≈ 46.7 µM

310

(assuming a molar absorption coefficient at 258 nm = 6750 M−1 cm-1), which is in

311

reasonable agreement with the initial concentration of KMnO4 (50 µM).44 As

312

illustrated in Figure S6(B), without stabilization by PP and the presence of organic

313

contaminant, the absorbance at 258 nm increased sharply and then decreased, forming

314

a peak at ~10 ms, followed by a gradual rebound to a plateau after 0.2 s. Although the

315

molar absorptivity of Mn(III) is not known for environmentally-relevant conditions, it

316

is likely to be significantly larger than PM at 258 nm, so the peak at 10 ms most likely

317

reflects the maximum transient concentration of Mn(III). After this peak, the

318

absorbance dropped moderately, which may be ascribed to (1) the decrease in

319

permanganate concentration; (2) the greater consumption rate of Mn(III) by bisulfite

320

and disproportionation than the generation rate of Mn(III). As the reaction proceeds,

321

the observed increase in absorbance at 258 nm could be due to the generation of

322

MnO2 because of Mn(III) disproportionation.

323

We tried testing the effect of adding phenol to the results in Figure S6(B), but

324

absorbance at 258 nm from phenol and its degradation products interfered detection of

325

Mn species at this wavelength. To avoid this interference, we used 20 mM methanol,

326

and found that this suppressed the peak concentration of Mn(III) at ~10 ms and

327

lowered the absorbance at 258 nm to ~0 (Figure S6(B)). This result helps to confirm 26-Sep-15


Page 17 of 32


328

that the transient at 10 ms is from the reactive Mn intermediate and suggests that the

329

disproportionation of Mn(III) is inhibited by the excessive methanol.

330

For additional insight into the chemistry species responsible for contaminant

331

oxidation in the PM/BS process, more complete UV-vis absorbance spectra were

332

recorded after 10 min for the PM/BS and MO/BS processes (Figure 4). Both

333

processes (PM/BS and MO/BS) show predominantly the spectrum for MnO2 without

334

PP, but in the presence of PP an absorbance peak at 258 nm is evident, which is the

335

characteristic peak of Mn(III)-PP (Figure S4). In the presence of both PP and phenol,

336

the absorbance arising from Mn(III)-PP is less (Figure 4A), confirming that Mn(III)

337

initially formed in the PM/BS process oxidizes phenol more rapidly than it complexes

338

with PP, and that both reactions are faster than the disproportionation of Mn(III).

339

The competition among the reactions that consume Mn(III) evidenced by the

340

spectra in Figure 4 is further supported by the kinetic data for phenol oxidation

341

(Figure S7). With PP present, the PM/BS process oxidized phenol with kobs = 5.5±0.5

342

s−1 (Figure S7A), which is 11-fold slower than that in the absence of PP. This could

343

be ascribed to the stable pH during the reaction (which changed less than 0.1) because

344

of the buffering capacity of PP, and the competition of PP with phenol for Mn(III).

345

Mn(III)-PP oxidized phenol very slowly with kobs ≈ 0.0016 min−1 (Figure S7B),

346

which is negligible compared to the rates obtained with the PM/BS process without

347

PP (Table 1).

26-Sep-15



Page 18 of 32

348

Pathways and Stoichiometry of the PM/BS Process. Based on all of the data

349

described above, a simplified reaction scheme for contaminant oxidation by HSO3−

350

activated MnO4− is proposed in Eqs. 1-5. The initial, activation of permanganate

351

involves reduction by (bi)sulfite to form Mn(III) (Eqs. 1 and 1’). The Mn(III) is

352

rapidly transformed to Mn2+ by disproportionation (Eq 2), and competing reactions

353

with (bi)sulfite (Eqs 3 and 3’) or organic contaminants (Eq 4). Some of the residual

354

(bi)sulfite can be oxidized by oxygen (Eqs 5 and 5’) since the experiments were

355

performed open to the air.

356

2 HSO3− + MnO4− → Mn(III) + 2 OH− + 2 SO42−

(1)

357

2 SO32− + MnO4− + 2H2O →Mn(III) + 4 OH− + 2 SO42− (at high pH)

(1’)

358

2 Mn(III) + 2 H2O → Mn2+ + MnO2 + 4 H+

(2)

359

HSO3− + 2 Mn(III) + H2O → 2 Mn2+ + SO42− + 3 H+

(3)

360

SO32− + 2 Mn(III) + H2O → 2 Mn2+ + SO42− + 2 H+ (at high pH)

(3’)

361

Mn(III) + contaminant → Mn2++ products

(4)

362

2 HSO3− + O2 → 2 H+ + 2 SO42−

(5)

363

2 SO32− + O2 →2 SO42− (at high pH)

(5’)

364

Because the initial concentrations of NaHSO3 and KMnO4 were 250 and 50 µM,

365

respectively, 100 µM NaHSO3 should be consumed by the reaction with KMnO4 (Eq

366

1) and the residual NaHSO3 (slightly less than 150 µM due to the oxidation of

367

bisulfite by oxygen during the pH adjustment) should be oxidized by oxygen (Eq 5)

368

and Mn(III) (Eq 3). Thus, about 100-150 µM H+ should be generated (without 26-Sep-15


Page 19 of 32


369

considering the consumption of bisulfite during pH adjustment) if no contaminant is

370

present, so the only sinks for Mn(III) are bisulfite oxidation and/or disproportionation

371

to Mn(II) and MnO2 (following Eq 3 and Eq 2, respectively). If these stoichiometric

372

assumptions hold, the theoretical pH of the mixture at the end of reaction (starting

373

with pHini = 5.0 and without contaminant present) should be 3.8-4.0, which matches

374

the experimentally determined value of ~4.0 shown in Figure S8.

375

When an organic contaminant (e.g., phenol) is present, some of the Mn(III) will

376

be used in oxidation of the organic contaminant, leaving less Mn(III) for

377

disproportionation (Eq 2) or oxidation of bisulfite (Eq 3), so the amount of H+

378

generated should be less. The decrease of H+ generated in this process would be

379

50-100 µM if all the Mn(III) was consumed by organic contaminants, which is

380

consistent with the progressive increase in pH from 4.0 to 4.9 (equivalent to the

381

decrease in H+ concentration by 87.4 µM) with increasing the initial phenol

382

concentration from 0 to 100 µM (Figure S8A). The same effect on pH variation was

383

found in the oxidation of other contaminants by the PM/BS process (not shown),

384

which is consistent with hypothesis that pH changes during oxidation of different

385

concentrations of contaminants in the PM/BS process is mainly determined by the

386

utilization of Mn(III).

387

The above analysis of pH effects on the PM/BS reactions (Eqs 1-5) also helps to

388

explain the cause and variability in the lag phase evident in some of the kinetic data

389

shown in Figures 1-2. The lag phase observed for all three contaminants, at C0 = 5 26-Sep-15



Page 20 of 32

390

µM, in the PM/BS process at pHini 5.0 (Figure 1, A-C) likely reflects the decrease in

391

pH (to pHend ≈ 4.1, Figure S8) during this time-period and the associated increase in

392

contaminant oxidation rate. No lag phase was observed in experiments performed

393

with phenol at higher initial concentrations (25 µM) (Figure S9), because there was

394

little change from the pHini of 5.0 (Figure S8A). Experiments initiated at higher pH

395

(pHini = 6.0) gave a larger drop in pH during the reaction (Figure S8B), which resulted

396

in the most prominent lag phase (Figure 2).

397

At pHini 8.0-9.0 bisulfite was mainly present as sulfite anion. Under these

398

conditions, the oxidation of sulfite anion by permanganate (Eq 1’) and Mn(III) (Eq 3’)

399

released OH- and H+, respectively, while no protons or hydroxide ions were generated

400

in the process of sulfite oxidation by oxygen (Eq 5’). The upward pH drift in

401

experiments that started at pHini 8.0-9.0 (Figure S8B) is most likely because the

402

amount of OH- generated in Eq 1’ was larger than that of H+ generated in Eq 3’.

403

It should be noted that the PM/BS system described in this work is highly novel

404

and many aspects of the chemistry involved will require further investigation. In

405

particular, it was not feasible within the scope of this initial study to fully characterize

406

the exact form of the Mn(III) species that is responsible for the high rates of organic

407

contaminant oxidation. There is prior work that is relevant to this issue, such as that

408

conducted by Moens et al.45 and Sisley et al.,46 which reported that the first hydrolysis

409

constant of Mn(OH2)63+ is 10-0.08 in ~4 M H+, Mn2+, ClO4- at 20 oC. The reaction

410

conditions to obtain this hydrolysis constant are very different from those of this study, 26-Sep-15


Page 21 of 32


411

but while the exact form of Mn(III) under the conditions is not known, it is likely to

412

be partially hydrolyzed. Further work will be required to fully characterize the

413

speciation of Mn(III), quantify Mn(III) generation and consumption rates, determine

414

second-order rate constants for oxidation of organic contaminants oxidation by

415

Mn(III), investigate the mechanism and products of contaminant oxidation, and

416

evaluate whether secondary oxidants (e.g., •OH) are significant under any conditions.

417

Environmental Implications. This is the first time that Mn(III) with minimal

418

stabilization has been implicated in chemistry performed under environmental

419

relevant conditions or has been shown to be reactive with environmentally-relevant

420

organic contaminants. The extraordinarily fast rates of contaminant oxidation in the

421

PM/BS process described in this study suggest that “activation” of permanganate

422

might lead to a new class of advanced oxidation processes (AOPs) in water treatment.

423

The contact times required to oxidize substantial concentrations of organic

424

contaminants with activated permanganate are very short, even when compared with

425

conventional AOPs that have been optimized and validated over many years of

426

research and development. The logistics for achieving sustained and complete

427

contaminant oxidation during water treatment will be addressed in future work. The

428

reagents required for the PM/BS process are practical for engineering scale

429

applications and the ultimate products formed from the reagents are SO42- and MnO2,

430

both of which can be accommodated within conventional water treatment processes.

431

26-Sep-15



Page 22 of 32

432

ASSOCIATED CONTENT

433

Supporting Information

434

The Supporting Information is available free of charge on the ACS Publications

435

website.

436

Degradation of refractory contaminants in PM/BS process, the influence of water

437

matrix on phenol decomposition, TOC removal, ESR spectra, UV-vis spectra of

438

various species, variation of absorbance at 258 nm in PM/BS process, influence of PP

439

on phenol removal kinetics, variation of pH under different conditions, degradation

440

kinetics of phenol with high concentration in PM/BS process and summary of the rate

441

constants of ciprofloxacin and phenol oxidation in various processes. This material is

442

available free of charge at http://pubs.acs.org.

443

AUTHOR INFORMATION

444

Corresponding Author

445

*Phenol: +86-21-65980956 (X.H. Guan), 503-346-3431 (P.G. Tratnyek)

446

Notes

447

The authors declare no competing financial interest.

448

ACKNOWLEDGMENTS

449

This work was supported by the Major Science and Technology Program for Water

450

Pollution Control and Treatment (Grant 2012ZX07403-001), the National Natural

451

Science Foundation of China (Grant 21522704) and the Fundamental Research Funds 26-Sep-15


Page 23 of 32


452

for the Central Universities. We thank Shanghai Institute of Organic Chemistry

453

(Chinese Academy of Sciences), Research Center for Eco-Environmental Sciences

454

(Chinese Academy of Sciences), Peking University and Lanzhou Institute of

455

Chemical Physics (Chinese Academy of Sciences) for providing access to SFS

456

instrumentation. The authors acknowledge helpful conversations regarding the

457

interpretation of these data with Profs. Alan Stone, George Luther, Bradley Tebo,

458

Matt Jones, Ninian Blackburn, Pierre Moënne-Loccoz, and Urs von Gunten.

459

REFERENCES

460

1. P. Neta, R. E. H., Free-radical chemistry of sulfite. Environ. Health Perspect. 1985, 64, 209-217. 2. G. V. Buxton, C. L. G., W. P. Helman, A. B. Ross, Critical review of rate constants for reactions of hydrated electrons, hydrogen atoms and hydroxyl radicals in aqueous solution. J. Phys. Chem. Ref. Data 1988, 17, 513-886. 3. Schwarzenbach, R. P.; Gschwend, P. M.; Imboden, D. M., Environmental Organic Chemistry, 2nd ed. John Wiley & Sons: 2005. 4. Siegrist, R. L.; Urynowicz, M. A.; West, O. R.; Crimi, M. L.; Lowe, K. S., Principles and practices of in situ chemical oxidation using permanganate. Battelle Press Columbus, OH: 2001; Vol. 505. 5. Petri, B. G.; Thomson, N. R.; Urynowicz, M. A., Fundamentals of ISCO using permanganate. In In Situ Chemical Oxidation for Groundwater Remediation, Springer: 2011; pp 89-146. 6. Lee, E. S.; Schwartz, F. W., Characteristics and applications of controlled–release KMnO4 for groundwater remediation. Chemosphere 2007, 66, (11), 2058-2066. 7. Guan, X.; He, D.; Ma, J.; Chen, G., Application of permanganate in the oxidation of micropollutants: a mini review. Front. of Environ. Sci. Eng. in China 2010, 4, (4), 405-413. 8. Rodríguez-Álvarez, T.; Rodil, R.; Quintana, J. B.; Triñanes, S.; Cela, R., Oxidation of non-steroidal anti-inflammatory drugs with aqueous permanganate. Water Res. 2013, 47, (9), 3220–3230. 9. Hu, L.; Stemig, A. M.; Wammer, K. H.; Strathmann, T. J., Oxidation of antibiotics during water treatment with potassium permanganate: reaction pathways and deactivation. Environ. Sci.Technol. 2011, 45, (8), 3635-3642.

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10. Zhang, J.; Sun, B.; Guan, X.; Wang, H.; Bao, H.; Huang, Y.; Qiao, J.; Zhou, G., Ruthenium nanoparticles supported on CeO2 for catalytic permanganate oxidation of butylparaben. Environ. Sci. Technol. 2013, 47, (22), 13011-13019. 11. Ma, X.; Hu, S.; Wang, H.; Li, J.; Huang, J.; Zhang, Y.; Lu, W.; Li, Q., Kinetics of oxidation of dimethyl trisulfide by potassium permanganate in drinking water. Front. of Environ. Sci. Eng. in China 2012, 6, (2), 171-176. 12. Pedersen, J. A.; Iv, K. F. R., Kinetics of oxytetracycline reaction with a hydrous manganese oxide. Environ. Sci. Technol. 2006, 40, (23), págs. 7216-7221. 13. Laha, S.; Luthy, R. G., Oxidation of aniline and other primary aromatic amines by manganese dioxide. Environ. Sci. Technol. 1990, 24:3, (3), 363-373. 14. Stone, A. T., Reductive Dissolution of manganese(III/IV) oxides by substituted phenols. Environ. Sci. Technol. 1987, 21, (10), 979-988. 15. Xu, L.; Xu, C.; Zhao, M.; Qiu, Y.; Sheng, G., Oxidative removal of aqueous steroid estrogens by manganese oxides. Water Res. 2008, 42, (20), 5038–5044. 16. Waldemer, R. H.; Tratnyek, P. G., Kinetics of contaminant degradation by permanganate. Environm. Sci. Technol. 2006, 40, (3), 1055-1061. 17. Tratnyek, P. G., Comment on "evaluation of the kinetic oxidation of aqueous volatile organic compounds by permanganate" by M. G. Mahmoodlu, S. M. Hassanizadeh, and N. Hartog, in Science of the Total Environment (2014) 485-486: 755-763. Sci. Total Environ. 2015, 502, 722–723. 18. Mahmoodlu, M. G.; Hassanizadeh, S. M.; Hartog, N., Evaluation of the kinetic oxidation of aqueous volatile organic compounds by permanganate. Sci. Total Environ. 2014, 485, 755-763. 19. Du, J.; Sun, B.; Zhang, J.; Guan, X., Parabola-like shaped pH-rate profile for phenols oxidation by aqueous permanganate. Environ. Sci. Technol. 2012, 46, (16), 8860-8867. 20. Hu, L.; Martin, H. M.; Strathmann, T. J., Oxidation kinetics of antibiotics during water treatment with potassium permanganate. Environ. Sci. Technol. 2010, 44, (16), 6416-6422. 21. Ogino, T.; Yaezawa, H.; Yoshida, O.; Ono, M., An experimental and theoretical study on the substituent effect of the permanganate oxidation of styrenes. Org. Biomol. chem. 2003, 1, (15), 2771-2776. 22. Zhang, J.; Sun, B.; Guan, X. H.; Wang, H.; Bao, H. L.; Huang, Y. Y.; Qiao, J. L.; Zhou, G. M., Ruthenium nanoparticles supported on CeO2 for catalytic permanganate oxidation of butylparaben. Environ. Sci. Technol. 2013, 47, (22), 13011-13019. 23. Zhang, J.; Sun, B.; Xiong, X. M.; Gao, N. Y.; Song, W. H.; Du, E. D.; Guan, X. H.; Zhou, G. M., Removal of emerging pollutants by Ru/TiO2-catalyzed permanganate oxidation. Water Res. 2014, 63, 262-270. 24. Li, L.; Shao, C.; Lin, T.-F.; Shen, J.; Yu, S.; Shang, R.; Yin, D.; Zhang, K.; Gao, N., Kinetics of cell inactivation, toxin release, and degradation during permanganation of Microcystis aeruginosa. Environ. Sci. Technol. 2014, 48, (5), 2885-2892.

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25. Sun, B.; Zhang, J.; Du, J.; Qiao, J.; Guan, X., Reinvestigation of the role of humic acid in the oxidation of phenols by permanganate. Environ. Sci. Technol. 2013, 47, (24), 14332-14340. 26. Xie, P.; Ma, J.; Fang, J.; Guan, Y.; Yue, S.; Li, X.; Chen, L., Comparison of permanganate preoxidation and preozonation on algae containing water: cell integrity, characteristics, and chlorinated disinfection byproduct formation. Environ. Sci. Technol. 2013, 47, (24), 14051-14061. 27. Jiang, J.; Pang, S.-Y.; Ma, J., Oxidation of triclosan by permanganate (Mn (VII)): Importance of ligands and in situ formed manganese oxides. Environ. Sci. Technol. 2009, 43, (21), 8326-8331. 28. Liu, J.; Zhang, X., Effect of quenching time and quenching agent dose on total organic halogen measurement. Int. J. Environ. Anal. Chem. 2013, 93, (11), 1146-1158. 29. Liu, W.; Andrews, S. A.; Stefan, M. I.; Bolton, J. R., Optimal methods for quenching H2O2 residuals prior to UFC testing. Water Res. 2003, 37, (15), 3697-3703. 30. Keen, O. S.; Dotson, A. D.; Linden, K. G., Evaluation of hydrogen peroxide chemical quenching agents following an advanced oxidation process. J. Environ. Eng. 2012, 139, (1), 137-140. 31. Perez-Benito, J. F.; Arias, C.; Amat, E., A kinetic study of the reduction of colloidal manganese dioxide by oxalic acid. J. Col. Int. Sci. 1996, 177, (2), 288–297. 32. Tartar, H. V.; Garretson, H. H., The thermodynamic ionization constants of sulfurous acid at 25° 1. J.am.chem.soc 1941. 33. Kochany, J.; Lipczynska-Kochany, E., Application of the EPR spin-trapping technique for the investigation of the reactions of carbonate, bicarbonate, and phosphate anions with hydroxyl radicals generated by the photolysis of H2O2. Chemosphere 1992, 25, (12), 1769–1782. 34. Anipsitakis, G. P.; Dionysiou, D. D.; Gonzalez, M. A., Cobalt-mediated activation of peroxymonosulfate and sulfate radical attack on phenolic compounds. implications of chloride ions. Environ. Sci. Technol. 2006, 40, (3), 1000-1007. 35. Zhang, J.; Sun, B.; Guan, X. H., Oxidative removal of bisphenol A by permanganate: Kinetics, pathways and influences of co-existing chemicals. Sep. Purif. Technol. 2013, 107, 48-53. 36. Zazo, J. A.; Mohedano, A. F.; Gilarranz, M. A.; Rodríguez, J. J.; Casas, J. A., Chemical pathway and kinetics of phenol oxidation by fenton's reagent. Environ. Sci. Technol. 2005, 39, (23), 9295-9302. 37. An, T.; Yang, H.; Li, G.; Song, W.; Cooper, W. J.; Nie, X., Kinetics and mechanism of advanced oxidation processes (AOPs) in degradation of ciprofloxacin in water. Appl. Catal. B-Environ. 2010, 94, (3), 288–294. 38. Garcia-Segura, S.; Brillas, E., Mineralization of the recalcitrant oxalic and oxamic acids by electrochemical advanced oxidation processes using a boron-doped diamond anode. Water Res. 2011, 45, (9), 2975-2984.

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39. Jiang, J.; Pang, S.; Ma, J., Role of ligands in permanganate oxidation of organics. Environ. Sci. Technol. 2010, 44, (11), 4270-4275. 40. Zamora, P. L.; Villamena, F. A., Theoretical and experimental studies of the spin trapping of inorganic radicals by 5, 5-dimethyl-1-pyrroline N-oxide (DMPO). 3. Sulfur dioxide, sulfite, and sulfate radical anions. J. Phys. Chem. A 2012, 116, (26), 7210-7218. 41. Connick, R. E.; Zhang, Y. X., Kinetics and mechanism of the oxidation of HSO3by O2. 2. The manganese(II)-catalyzed reaction. Inorg. Chem. 1996, 35, (16), 4613-4621. 42. Simandi, L. I.; Jaky, M.; Schelly, Z. A., Short-lived manganate(VI) and manganate(V) intermediates in the permanganate oxidation of sulfite ion. J.am. chem. Soc. 1984, 106, (22), 6866-6867. 43. Hassan, E. M.; Belal, F., Kinetic spectrophotometric determination of nizatidine and ranitidine in pharmaceutical preparations. J. Pharmaceut. Biomed. 2002, 27, (1-2), 31–38. 44. Webb, S. M.; Dick, G. J.; Bargar, J. R.; Tebo, B. M., Evidence for the presence of Mn(III) intermediates in the bacterial oxidation of Mn(II). Proc. Natl. Acad. Sci. U.S.A. 2005, 102, (15), 5558-5563. 45. Moens, J.; Seidel, R.; Geerlings, P.; Faubel, M.; Winter, B.; Blumberger, J., Energy levels and redox properties of aqueous Mn2+/3+ from photoemission spectroscopy and density functional molecular dynamics simulation. J. Phys. Chem. B 2010, 114, (28), 9173-9182. 46. Sisley, M. J.; Rb., J., First hydrolysis constants of hexaaquacobalt(III) and -manganese(III): Longstanding issues resolved. Inorg. Chem. 2006, 45, (26), 10758-10763.

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TOC Art

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For “Activation of Manganese Oxidants with Bisulfite for Enhanced Oxidation of

595

Organic Contaminants: The Involvement of Mn(III)” by Bo Sun, Xiaohong Guan,

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Jingyun Fang, and Paul G. Tratnyek

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Page 27 of 32


597 598

26-Sep-15



1.0

1.0

1.0

1.0

0.8

0.9

C/C0

0.8

0.6

0.6 0.5

0.4

0

15

30

45

60

Time (min)

0.2 0.0

A

0.00

0.6 0.4

0.7

0.6

C/C0

0.8

C/C0

0.8

0.2 0.0

0.4

0

0.06

0.09

0.12

0.015

1.0

C/C0

0.0

C/C0

C/C0

0.6

0.6 0.5 0

10

20

30

Time (min) C

0.060

Time (s)

0.8 0.7 0.6 0.5

0.4

0

20

40

60

80

Time (min)

0.2 D

0.0

0.00 0.02 0.04 0.06 0.08 0.10

599

0.045

0.9

0.7

0.2

0.030

1.0

0.8

0.8

0.4

60

B

0.000

0.9

0.6

45

Time (s)

1.0

0.8

30

Time (min)

Time (s) 1.0

15

0.2 0.0

0.03

Page 28 of 32

0

3

6

9

12

15

Time (s)

600

Figure 1. Disappearance kinetics for (A) phenol, (B) ciprofloxacin, and (C) methyl blue in

601

PM/BS process; as well as (D) methyl blue by NaHSO3 activated MnO2 at pHini 5.0. The insets

602

show the corresponding degradation kinetics of contaminants by KMnO4 or MnO2 without

603

NaHSO3. Reaction conditions: [KMnO4]0 or [MnO2]0 = 50 µM, [NaHSO3]0 = 250 µM. [Test

604

Contaminants]0 = 5 µM, T = 18±2 oC.

605 606 607 608 609 610 611 612 613

26

ACS Paragon Plus Environment

Page 29 of 32


C/C0

1.0 0.8

0.8

pH = 5.0 1.0 0.8

0.6

0.6

0.6

0.4

0.4

0.4

0.2

0.2

0.2

0.0

0.0

0.0

pH = 4.0

0.00

0.02

0.04

0.06

0.08

1.0

0.00 0.03 0.06 0.09 0.12 0.15

Time (s)

C/C0

1.0

1.0

pH = 8.0

0.8

0.6

0.6

0.6

0.4

0.4

0.4

0.2

0.2

0.2

0.0 1.0

1.5

Time (s)

2.0

0.14

0.21

1.0

0.8

0.5

0.07

0.28

Time (s)

0.8

0.0

614

0.00

Time (s) pH = 7.0

0.0

pH = 6.0

pH = 9.0

0.0 0.0

0.5

1.0

1.5

2.0

Time (s)

0.0

0.5

1.0

1.5

2.0

2.5

Time (s)

615

Figure 2. Disappearance kinetics for of phenol by NaHSO3 activated KMnO4 at different initial

616

pH levels. Reaction conditions: [KMnO4]0 = 50 µM, [NaHSO3]0 = 250 µM, [Phenol]0 = 5 µM.

617

27



0.10

0.10

A

0.06

0.06 0.04

0.04 0.02

0.02

0.00

0.00 400

400

m) λ (n

600 700 1.0

0.8

0.4

0.6

0.2

m) λ (n

500

500 600 700 1.0

) Time (s

0.10

0.06

0.06

Abs

0.08

Abs

0.8

0.6

0.04

0.4

0.2

) Time (s

0.10

C

0.08

D

0.04

0.02

0.02

0.00

0.00

400

400

m) λ (n

m) λ (n

500 600 700 1.0

0.8

0.6

) T ime (s

618

B

Abs

0.08

Abs

0.08

Page 30 of 32

0.4

0.2

500 600 700 1.0

0.8

0.6

0.4

0.2

Time (s)

619

Figure 3. The evolution of three dimensional UV-visible spectra in the process of NaHSO3

620

activated KMnO4 at pHini 5.0 (A) without phenol; (B) [Phenol]0 = 5.0 µM; (C) [Phenol]0 = 50

621

µM; (D) without phenol but [PP]0 = 5 mM. Reaction conditions: [KMnO4]0 = 50 µM;

622

[NaHSO3]0 = 250 µM.

623

28


Page 31 of 32


A

.8

PP: 0 mM, Phenol: 0 µM PP: 5 mM, Phenol: 0 µM PP: 5 mM, Phenol 50 µM

Abs

.6

.4

.2

0.0 200

300

400

500

600

700

.8

B (1) PP: 0 mM (2) PP: 5 mM (2) - x(1)

.6

Abs

800

.4

.2

0.0 200

300

400

500

600

700

800

Wavelength (nm)

624 625

Figure 4. (A) Influence of PP and phenol on the UV-visible spectrum of the products generated

626

in the PS process at pHini 5.0; (B) Influence of PP on the UV-visible spectrum of the products

627

generated in the MnO2/NaHSO3 process at pHini 5.0. The line (1) represents the absorption of

628

MnO2 and line (2) represents the absorption of MnO2 and Mn(III)-PP, x was a coefficient used to

629

subtract the absorption of MnO2 from line (2). Reaction conditions: [KMnO4]0 or [MnO2]0 = 50

630

µM; [NaHSO3]0 = 250 µM; T=18±2 0C.

631

29



635

634

oxidant

KMnO4

KMnO4

KMnO4

KMnO4

KMnO4

KMnO4

KMnO4

KMnO4

MnO2

Phenol

Phenol

Phenol

Phenol

Phenol

Phenol

Ciprofloxacin

Methyl blue

Methyl blue

5.0

5.0

5.0

9.0

8.0

7.0

6.0

5.0

4.0

pH

0.29±0.03

80.7±3.3

147±9

1.3±0.08

1.2±0.09

3.9±0.05

25.8±3.4

62.4±0.8

119.8±2.8

k1 (s )

-1

with NaHSO3

30

6.9 × 10 4 4.8 × 10 5 9.35 × 105 4.9 × 10 3 1.0× 10 3 6.0× 10 2 1.5 × 10 6 3.6 × 10 5 3.0 × 10 3

1.30±0.07 × 10 -4 2.76±0.85 × 10 -5 7.91±0.99 × 10 -4 1.20±0.03 × 10 -3 2.17±0.07 × 10 -3 9.50±0.08 × 10 -5 2.22±0.20 × 10 -4 9.50±1.33 × 10 -5

k1/k2

1.72±0.35 × 10 -3

k2 (s )

-1

without NaHSO3

= 50 µM; [NaHSO3]0 = 250 µM; [Test Contaminants]0 = 5.0 µM)

633

Contaminants

Table 1 Influence of NaHSO3 on the rate constants of contaminants oxidation by KMnO4 or MnO2 at pHini 5.0 ([KMnO4]0 or [MnO2]0

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Environmental Science & Technology Page 32 of 32

Activation of Manganese Oxidants with Bisulfite for Enhanced

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